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CHEMISTRY

 

CHEMISTRY TOPICAL QUESTIONS
FORM I WORK
TOPIC 1
SIMPLE CLASSIFICATION OF SUBSTANCES.
PAST KCSE QUESTIONS ON THE TOPIC.
1. The diagram below represents a paper chromatogram of pure w, X, and Y. A mixture K contains W and Y only.  Indicate on the diagram the chromatogram of K.               (1mk)
2. 
Study the information below and answer the question that follows. A mixture contains the solids; Alum camphor and sugar. The solubility of different liquids is shown in the table below.
 
solid Liquid
 Water Ethanol Ether
Alum Soluble Insoluble Insoluble
Camphor Insoluble Soluble Very soluble
Sugar Soluble Soluble Insoluble
Explain how you would obtain a sample of solid sugar from the mixture.
3. 
The equation below represents two processes that takes place without any change in temperature.
i) H2O(s) → H2O (l)
ii) CdCl2(s)   →     CD2+ (l)     + 2CL (l)     
a)    Explain why although heat is required for each of the processes to take place, the temperature remained constant in both processes.  (1mk)
b) Which of the two processes has a higher enthalpy change ∆H; Give a reason?        (2mks)
4. 
The table below gives some properties of gas D and E.   (2mks) 
Gas Density Effect on H2SO4 Effect on NaOH.
D Lighter than air React to form salt  Dissolve without reacting
E Heavier than air Not affected Not affected
a) Describe how you would obtain a sample of gas E from the mixture of gas D and E
b) Suggest a possible identity of gas D. Give reasons for your answer. (2mks)
5. 
Explain how you would separate a mixture of Nitrogen and Oxygen gases given that their boiling points are – 1960C and -1830C respectively.  (2mks)

6. 
In an experiment to separate a mixture of organic liquid “m” (B.P. 560C) and liquid “n” (B.P. 1180C) a student set up the apparatus shown below.

a) Identify two mistakes in the set up.     (2mks)
b) What method would the student use to test the purity of the distillates?
(1mk)
7. 
Some sodium Chloride was found to be contaminated with Copper (II) Oxide.  Describe how a sample of sodium chloride can be separated from the mixture.  
(3mks)

8. 
The set up below represents apparatus that may be used to separate a mixture of two miscible liquids “C” and “D” whose boiling points are 800C respectively.

a) Name B.
 b) What is the purpose of the thermometer?    (1mk)
 c) Which liquid is collected in the test tube?    (1mk)
 
9. Air was passed through several reagents as shown in the flow chart below.

a) Write an equation for the reaction which takes place in the chamber with magnesium powder.       (1mks)
b) Name one gas which escapes from the chamber containing magnesium.  Give a reason for your answer.     (2mks)
10. Dry Carbon (II) Oxide gas reacts with heated Lead (II) as shown in the equation below.
PbO(s) + CO (g)      →    CO2 (g)    + Pb(s)
a) Name the process undergone by the Lead (II) Oxide.   (1mk)
b) Give a reason for your answer (a) above.    (1mk)
c) Name another gas that can be used to perform the same function as Carbon (II) Oxide gas in the above reaction.    (1mk)
11. The diagram below shows a Bunsen burner when in use.

a) Name the region labelled C and D.       (2mks)
12. Samples of urine from three participants F, G and H at an international sports meeting were spotted onto a chromatography paper alongside two from illegal drugs A1 and A2.  A chromatogram was run using methanol.  The figure below shows the chromatogram.

 a) Identify the athelete who had used an illegal drug.   (1mk)
 b) Which drug is more soluble in methanol?    (1mk)
13. The graph below is a cooling curve of a substance from gaseous state to solid state.

Give the name of the:
 a) Process taking place between t0 and t1;     (1mk)
 b) Energy change that occurs between t3 and t4    (1mk)
14. For each of the following experiments give the observation, the type of change that occurs (physical or chemical) and the formula (e) of any substance(s) formed.  If no new compound (substance) is formed write no new compound formed.
Experiment Observation  Type of change  Formulae
Add few drops of concentrated sulphuric acid to small amount of sugar (C12H22O11)   
A few crystals of Iodine I2 are heated gently in a test tube.   
Few crystals of Copper (II) Nitrate are heated strongly in a test tube.   
Sodium hydroxide platettes in an evaporating dish are left in humid air for one day.   
15. 
 a) What method can be used to separate a mixture of ethanol and propanol?            (1mk) 
b) i) Explain how a solid mixture of sulphure and sodium chloride can
be separated into solid sulphur and solid sodium chloride. (4mks)
  ii) How can one determine that solid sulphure is pure?  (2mks)
c) The table below gives the solubilities of potassium bromide and potassium bromide and potassium sulphate at 00C and 400C.
Substances Solubilities in g/100g of water
Potassium bromide 00 400
 55 75
Potassium sulphate 10 12
When aqueous mixture containing 60g of potassium bromide and 7g of potassium sulphate in 100g of water at 800C, some crystals were formed.
i) Identity the crystals.       (1mk)
ii) Determine the mass of crystals formed.    (1mk)
iii) Name the method used to obtain the crystals.    (1mk)
iv) Suggest one industrial application of the method named in (c) (iii) above
           (1mk)
16. 
 Describe the process by which Nitrogen is obtained from air on a large scale.            (4mks)
17.   Name the methods by which the following substances could be separated.
 a) Kerosene from crude oil      (1mk)
 b) Coloured extract from grass dissolved in ethanol.   (1mk)
 c) Aluminium chloride from sodium chloride.    (1mk)
 d) Iron fillings from sulphur powder.     (1mk)
18. The diagram below represents three methods for collecting gases in the laboratory
  (i)   (ii)   (iii)
 a) Name the methods shown in the diagram    (3mks)
b) State with reasons the most suitable methods for collecting each of the following gases.
 i) Oxygen       (1mk)
 ii) Hydrogen       (1mk)
 iii) Carbon (IV) Oxide      (1mk)
19. A laboratory technician accidentally mixed liquids suspected to be benzene (B.P. 78 0C).  He has a problem of separating the mixture and seeks your help.  Describe to him.        (4mks)
 a) The method he should use
 b) The apparatus he should use
 c) The precautions he should take when carrying out the separation.
20. Study the following chart for laboratory preparation of dry nitrogen.
 a) State what happens in step I and II
 b) Name the compounds which can be used in step I and II respectively.
(2mks)
21. Explain how naphthalene could be separated from a mixture of naphthalene and common salt.         (2mks)
22. A student added some pure potassium nitrate crystals to cold water and stirred the mixture.  A few of the crystals did not dissolve at room temperature.
 a) i) Give a reason why some crystals did not dissolve.  (1mk)
ii) What would happen if the contents of the mixture in a beaker were warmed? Explain.      (2mks)
b)        i) Name two substances which can be reacted to give Copper (II) Sulphate.       (1mk)
           ii) Write the equation for the reaction between the substances named in b (i) above.       (1mk)
c) Some Copper (II) sulphate crystals were gently heated in a test tube until no more water was given off.
i) Draw a diagram of the apparatus that could be used to heat the crystals and collect the water given off.   (3mk)
ii) State what would be observed if the residue in the test tube is cooled and few drops of water is added to it.   (1mk)
23. The set up below was used to determine the melting point of naphthalene.

a) State precautions which should be taken into consideration when carrying out this experiment.     (3mks)
b) State the use of the following in this experiment.
 i) Thermometer.      (1mk)
 ii) Stirrer       (1mk)
 iii) Boilling water      (1mk)
c) The experimental value of the melting point of naphthalene is 780C and theoretical value is 800C. Suggest one reason for this difference.         (1mk)
24. The following diagram is used to show that air contains Carbon (IV) Oxide.
 a) Name liquid “p”       (1mk)
b) State the observation made on liquid “p” which will indicate the presence of carbon (IV) Oxide.       (1mk)
c) Write an equation for the reaction between “p” and Carbon (IV) Oxide.
(1mk)
25. Explain why potassium is kept under paraffin while phosphorous under water.           (2mks)

26. Study the information below and answer the questions that follow.
 
Solids Cold water Hot water
R Soluble Soluble
S Insoluble Insoluble
V Insoluble  Soluble
Briefly explain how you can separate a mixture of solid R, S AND V (3mks)
 
TOPIC 2
ACIDS, BASES AND INDICATORS
1. What would be observed when aqueous sodium hydroxide is added to aqueous Lead (II) Nitrate?        (1mk)
2. Explain why concentrated sulphuric acid is a weaker acid than dilute sulphuric acid?          (1mk)
3. When solid calcium carbonate is reacted with excess dilute hydrochloric acid, Carbon (IV) Oxide gas is evolved. The graph below shows a plot of the volume of carbon (IV) Oxide evolved against time.
     
  0              4                  8                  12                16                20                24              28
Explain how the evolution of carbon (IV) oxide varies with time.  (2mks)
4. Study the flow chart below and answer the question that follows.
Write the chemical formula for the complex ions in M and N.
5. Explain the following observations.  A molar solution of nitrous acid (Nitric (III) acid has a PH of 2 whereas a one molar solution of hypochlorous acid (Chloric (I) acid has a PH of 4.        (2mks)
6. Solutions may be classified as strong basic, weakly acidic, strong acidic.  The information below gives solutions and their PH values.  Study it and answer the questions that follow.
 
Solutions PH values
B 1.5
C 6
D 14
Classify the solutions in the table above using the stated classification (3mks)
7. 
Explain how you would distinguish between a carbonate and a sulphite using dilute acid and blue litmus paper.      (1mk)
8. 
In the equation below, identify the reactant that act as an acid and explain how you would arrive at your choice.
 NH +4(aq) + H2O (I) ↔ NH3 (g)      +   H3O + (aq)             (2mks)
9. 
Describe how the following reagents can be used to prepare Lead sulphate, solid potassium sulphate, solid lead carbonate, dilute nitric acid and distilled water           (2mks)
10. Distinguish between strong and weak acid.  Give an example of each. (2mks)
11. 
Describe how a solid sample of Lead (II) chloride can be prepared using the following reagents.  Dilute nitric acid (Nitric (V) acid), dilute Hydrochloric acid and lead (II) carbonate.       (2mks)
12. 
A bee keeper found that when stung by a bee, application of a little solution of sodium hydrogen Carbonate help to relieve the irritation from the affected area. Explain.         (2mks)
13. 
State and explain the observations that would be made when a few drops of concentrated sulphuric acid are added to a small sample of hydrated copper (II) sulphate.         (2mks)
14. 
Dg of potassium hydroxide were dissolved in distilled water to make 100cm3 of the solution required 50cm3 of solution.  50cm3 of 2m Nitric (V) acid for complete neutralization.  Calculate the mass of d of potassium hydroxide.
Relative molecular mass of KOH = 56
KOH(aq)    +    HN3(AQ)  → KNO 3(Aq)   + H2O(l)
15. 
a) A few drops of freshly prepared iron (II) sulphate solution was added to potassium Nitrate solution in a test tube.  Concentrated sulhuric acid was then carefully added to the mixture.  State the observation that was made.
           (1mk)
b) Write an equation for the reaction that occurs when solid potassium nitrate is strong heated.       (1mk)
16. The PH of a sample of soil was found to be 5.0.  An agricultural officer recommended the addition of calcium oxide in the soil.  State two functions of calcium oxide in the soil.       (2mks)
17. 
18.       In an experiment 30cm3 of 0.1M sulphuric acid were reacted with 30cm3 of
0.1M sodium Hydroxide.
 a) Write an equation for the reaction that took place   (1mk)
b) State the observations that were made when both blue and red litmus papers were dropped into the mixture.    (1mk)
c) Give a reason for you answer in (b) above.    (1mk)
19. 
 The following tests were carried out on separate portions of a colorless solution S.
 Tests Observation
i) Addition of dilute Hydrochloric acid to the  first  portion of S. No observable changes.
ii) Addition of aqueous ammonia to the third portion of s White precipitate was formed which dissolved in excess of aqueous ammonia.
iii) Addition of aqueous ammonia to the third portion of S. White precipitate was formed which dissolved in excess of aqueous ammonia.
(a)      From the information in test (i) name action which is not present in
solution S.         (1 mk)
(b)      Identify a cation which is likely to be present in solution S.   (1 mk)
(c)      Write an ionic equation for the reaction which takes place in test (II).             (1 mk)
20. 
In an experiment, equal amounts of magnesium powder were added into test tubes 1 and 2 as shown below.
Explain why the amount of Hydrogen gas liberated in test tube 2 is greater than in test tube 1 after 5 minutes. 3 mks
21.     
Ammonia gas was passed into water as shown below.
(a)     When a red litmus paper was dropped into the resulting solution, it turned
blue. Give a reason for this observation.     (1 mk)
(b)      What is the function of the funnel?      (1 mk)
22.     
Zinc (II) Oxide reacts with acid and alkalis.
(a)      Write the equation for the reaction between Zinc (II) Oxide and
(i)       Dilute sulphuric acid       (1 mk)
(ii).   Sodium hydroxide solution.                              (1 mk)
(b)      What property of Zinc oxide is shown above by the reaction (a) above?  
          (1 mk)
23.     
Equal volumes pf 1M monobasic acid L and M were each reacted with excess magnesium turnings. The table below shows the volumes of the gas produced after one minute.
Acids Volume of gas in cm3
L 40
M 100
Explain the difference in the volumes of the gas produced    (2mks)
24. 
When a few drops of aqueous ammonia were added to Copper (II) Nitrate   solution a light blue precjpitate was formed. On addition of more aqueous ammonia a deep blue solution was formed. Identify the substance responsible for the
(a)      Light blue precipitate       (1 mk)
(b)      Deep blue precipitate       (1 mk)
25.    
When a student was stung by a nettle plant, a teacher applied an aqueous solution of ammonia to the affected area of the skin and the student was relieved of pain. Explain.         (2mks)
26.    
In an experiment, a few drops of concentrated nitric acid were added to aqueous iron (II) sulphate in a test tube. Excess sodium hydroxide solution was then added to the mixture.
(a)     State the observations that were made when
(i)       Concentrated nitric acid was added to aqueous iron (II)
sulphate        (1 mk)
(ii)      Excess sodium hydroxide was added to the mixture.  (1 mk)
(b)     Write an ionic equation for the reaction which occurred in (a) (ii) above.             (1 mk)
27. The table below shows the tests that were carried out on solid N and the observation«made.
I Test Observations
II Dilute hydrochloric acid was added to
solid N. A colourless solution was formed.
III To the colourless solution obtained in
test II, excess sodium hydroxide solution was added. A white precipitate was formed
which dissolved to form a
colourless solution.
Write the formula of the anion in:
a) Solid N        (1mk)
b) The colourless solution formed in test II.    (1mk)
28. 
Zinc reacts with both concentrated and dilute sulphuric (VI) acid.  Write equations for the two reactions.      (2mks)

29. 
A compound whose general formula is M (OH)-4(aq)   reacts as shown by the equation below.
M (OH)3(s)   + OH-(aq)     → M (OH)-4(aq).
M (OH)3(s)   +  3H+(aq)     → M3+(aq)  + 3H2O(l)
a) What name is given to compounds which behave like M (OH)3(s)  in the two reactions below.       (1mk)
b)       Name two elements whose hydroxides behave like that of M. (2mks)
30. Study the flow chart below and answer questions that follow.

a) Give the name of the process that takes place in step 1.  (1mk)
 b) Give:
  i) The name of substance G1     (1mk)
  ii) One use of substance F1     (1mk)

31. a) Give the name of each of the processes described below which takes place
when the salt are exposed to air for some time.
 i) Anhydrous Copper (II) Sulphate becomes blue.  (1mk)
 ii) Magnesium chloride forms an aqueous solution.  (1mk)
 iii) Fresh crystals of sodium carbonate (Na2CO3:  10H2O become
covered with a white powder of formula Na2CO3: H2O. (1mk)
b) Write the formula of the complex ion formed in each of the reactions described below.
 i) Zinc Oxide dissolves in excess ammonia solution.  (1mk)
 ii) Copper hydroxide dissolves in excess ammonia solution. (1mk)
c) A hydrated salt has the following composition by mass;
 Iron 20.2%, Oxygen 23.0%, Sulphur 11.5%, water 45.3%.  Its relative formula mass is 278.
i) Determine the formula of hydrated salt were dissolved in distilled water and the total volume made to 250 cm3 of solution. Calculate the concentration of the salt solution in moles per litre. (2mks)
32. 
The reaction between bromine and mehanoic acid at 3000C proceeds according to the information given below.
Br2(l)  + HCOH(aq)            2Br-(aq)  + CO2(g)  +  2H+  (aq)

The table below shows the change in concentration of Bromine liquid against time.
Concentration of Br2(l)  mole/dm3 Time in minutes
10.0 x 103 0
8.1 x 103 1
6.6 x103 2
4.4 x 103 4
3.0 x 103 6
2.0 x 103 8
1.3 x 103 10
a) Plot a graph of concentration of bromine (vertical axis) against time.
(3mks)
 b) From the graph determine
  i) The concentration of bromine at the end of 3 minutes. (1mk)
  ii) The rate of reaction at t= 1 ½ minute.     (2mk)
 c) Explain how the concentration of bromine affects the rate of the reaction. 
           (2mks)
d) On the same axis, sketch the curve that would be obtained if the reaction was carried out at 200C and label the curve as curve II.  Give a reason for your answer.

33. 
The table below gives the volumes of gas produced when different volumes of
2M Hydrochloric acid were reacted with 0.6g of magnesium powder at room
temperature.

Volume of 2m HCL in cm3 Volume of gas (cm3)
0 0
10 240
20 360
30 600
40 600
50 600
 a) Write an equation for the reaction between magnesium and Hydrochloric
acid.         (1mk)
b) On the grid provided plot a graph of the volume of gas produced (vertical axis) against the volume of acid added (note that before the reaction comes to a completion the volume of gas produced is directly proportional to the volume of acid added.
c) From the graph, determine
i) The volume of the gas produced if 12.5cm3   of 2M Hydrochloric acid had been used.      (1mk)
ii) The volume of 2M Hydrochloric acid which react completely with 0.6g of magnesium powder.     (1mk)
d) State and explain the effect on the rate of production of the gas if
 i) 0.6g of magnesium ribbon was used instead of magnesium powder.
ii) 3m Hydrochloric acid was used instead of 2M Hydrochloric acid.         (2mks)
e) Given that one mole of the gas occupies 2400cm3 at room temperature.  Calculate the relative atomic mass of magnesium.   (3mks)
34. a) Name one ore from which copper of extracted.
b) The flow chart below shows a sequence of reactions starting with copper.  Study it and answer the questions that follow.

i) Identify gas P       (1mk)
  Reagent Q       (1mk)
  Solid R       (1mk)
 ii) Write an equation for the reaction that takes place in step 5. (1mk)
 iii) State the observations made in steps 4 and 7.   (2mks)
  Step 4________________________
  Step 7________________________
c) Bronze is an alloy of copper and another metal.
 i) Name the other metal
 ii) Give one use of bronze.
35. The graph below shows how the PH value of soil in a farm changed over a period of time.

 i) Describe how the PH of the soil is determined.  (2mks)
ii) State one factor that may have been responsible for the change in the soil PH in the time interval AB    (1mk)
36. The following data gives the PH value of solution P, Q and R.  
Solution PH value
P 13.6
Q 6.9
R 1.3
i) Which solution would produce Carbon (IV) Oxide when reacted with Copper (II) Carbonate?      (1mk)
ii) What would be the colour of solution “P” after adding a few drops of phenolphthalein indicator?      (1mk)
37. a) What is basicity of an acid?      (1mk)
 b) With reason write down the basicity of ethanoic acid.   (CH3COOH).
           (2mks)
38. An indicator established the following quilibrium when dissolved in water.
 OX(aq)  +  H2O(l)                           HOX(aq)  +   OH (aq)
            Blue            Yellow
 State and explain the colour of this indicator in
 i) Acidic medium       (1mk)
 ii) Alkaline medium       (1 mk)
39. Study the flow chart below and answer the questions that follow:

Write the formula of the ions in solid X.     (1mk)
40. The table below shows the PH values of certain solutions
 
Solution A B C D
PH values 8 5 7 11
 Which of the solutions is most likely to be solutions of
 i) Common salt        (1mk)
 ii) Lime water        (1mk)
 iii) Orange juice        (1mk)
 iv) Household soap       (1mk)
41. The table below shows PH values for some solutions. 
Solution A B C D
PH values 13.5 7 1 6.5
a) What solution reacts vigorously with magnesium metal?  (1mk)
b) Which solution forms complex ions with zinc (II) Oxide?  (1mk)
c) Which solution is likely to be that of lemon juice?   (1mk)
42. a) Freshly prepared iron (II) sulphate solution was reacted with a few drops
of Sodium Hydroxide solution.  State the observation made.  (1mk)
b) State and explain the observations made when the products formed in the above reaction stand for some time.     (2mks)
  i) Observation
  ii) Explain
43. Explain the differences between strong and weak acids.
44. The following table shows the PH values of solutions A, B, C and D. (2mks)
Solution PH values
A 9.8
B 2.0
C 5.2
D 12.0
Which one of the solutions, NaOH (aq), CH3COOH (aq),   HCL (aq) and NH3 (aq) correspond to solutions A, B, C and D.       (2mks)
45. When ion fillings were dissolved in dilute sulphuric acid a pale green solution formed and a colourless gas was given off.  The solution filtered and divided into two portions.
 a) Write an equation for the reaction.     (1mk)
b) To the first portion of the filtrate, aqueous ammonia was added drop wise until in excess.
 i) What was observed?      (1mk)
 ii) Write an ionic equation for the reaction   (1mk)
c) To the second portion of the filtrate, dilute sulphuric acid was added and warmed. A few drops of concentrated Nitric acid were added and a mixture heated.  Brown fumes were given off and a brown solution removed.
 i) Write an equation for this reaction.    (1mk)
 ii) What was the purpose of concentrated Nitric acid in this reaction? 
          (1mk)
d) To the brown solution formed in (c) above zinc metal was added.  The mixture was the left to stand for 30 minutes.
 i) What observations would be during and after 30 minutes?  (2mks)
 ii) What is the role of zinc metal?    (1mk)
 iii) Write an ionic equation for this reaction.   (1mk)

 
TOPIC 3
AIR AND COMBUSTION
1. 
Study the experiment set up represented by the diagram below and answer the question that follows.
a) Explain what would be observed if red and blue litmus papers were dipped into the water at the end of experiment.    (2mks)
b) Write an expansion in terms of X and Y to show the (%) percentage of gas used by the burning candle.      (1mk)
2. 
The diagram below represents two iron nails with some parts wrapped tightly with zinc and copper strips respectively.
What observations would be made at the exposed points A and B if the wrapped nails are left in the open for several months?  Explain.   (3mks)
3. 
In an experiment, rods of metals P, Q and R were cleaned with a sand paper and placed in a beaker containing water.  Another set of rods was also cleaned and placed in a beaker containing dilute acid.  After placing the rods in the two liquids bubbles of gas were seen around some of the rods as shown in the diagram below.

a) Why was it necessary to clean the rods with sand paper before dipping them into the liquids?       (1mk)
b) Arrange the three metals in order of their reactivity starting with the most reactive.        (1mk)
4. 
When magnesium is burnt in air it reacts with oxygen and nitrogen gas giving a white ash.  Write two equations for the two reactions that take place. (2mks)

5. 
Oygen reacts with the elements phosphorous, sulphur and chlorine to form oxides in which the elements is in its highest oxidation number.  The table below gives the oxide of sulphur and its highest oxidation number.  Complete the table for phosphorous and chlorine.  (Atomic number p=15, s=16, Cl= 17)  (2mks)
Elements Oxides Highest oxidation number
P  
S SO3 +6
Cl  
6. 
Write an equation for the reaction that takes place when carbon (II) Oxide gas is passed over heated Lead (II) Oxide.      (1mk)
7. 1997: pp 1A q. 1
 The set up below was used to study some properties of air

 State and explain two observation that would be made t the end of the experiment.            (3mks)
8. 
Give the formula of an oxide which reacts both dilute Hydrochloric acid and hot concentrated sodium hydroxide.
9. 
In an experiment a certain volume of air was passed repeatedly from syringe over heated excess zinc powder as shown in the diagram below.

The experiment was repeated using excess magnesium powder.  In which of the experiments was the change in volume of air greatest? Give reasons. (3mks)
10. 
State and explain the change in mass that occurs when the following substances are separately heated in open crucibles.  
i) Copper metal 
ii) Copper (II) Nitrate       (3mks)
11. 
The diagram below shows an iron bar, which supports a bridge. The iron is connected to a piece of magnesium metal.

Explain why it is necessary to connect the piece of magnesium metal to the iron bar.          (3mks)
12. 
Explain why magnesium continue to burn in a gas jar full of Sdulphur (IV) Oxude while burning splint would be extinguished.

13. 
 The diagram below is a set up for the laboratory preparation of oxygen gas.

a) Name solid R.        (1mk)
 b) Write an equation for the reaction that takes place in the flask. (1mk)
 c) Give one commercial use of oxygen.
14. Nitrogen (II) Oxide and nitrogen (IV) Oxide are some of the gases released from car exhaust pipes.  State these gases affect the environment.   (2mks)
15. The set up below was used to abtain a sample of iron.

Write two equations which occur in the combustion tube.   (2mks)
16. 
The low chart below outlines some of the process involved during extraction of copper from pyrites.  Study it and answer the questions that follow.
 a) i) Name gas K.       (1mk)
ii) Write an equation for the reaction that takes place in 1st roasting furnance.       (1mk)
iii) Write the formula of the cations present in the slag M (1mk)
iv) Identify gas P.       (1mk)
v) What name is given to the reaction that takes placein chamber N? Give a reason for your answer.    (2mks)
b) Copper obtained from chamber N is not pure.  Draw a labelled diagram to show the set up you would use to refine the copper by electrolysis. (2mks)
c) Given that the mass of copper obtained from the above extraction was 210 kg, determine the percentage purity of the ore (copper pyrite) if 810 kg of it was fed to 1st roasting furnance. Cu= 63.5, Fe= 56.0, S=32.0 (3mks)
d) Give two effects that this process could have on the environment. (2mks)
17. 
The table below gives the information about the major constituents of crude oil.  Study it and answer the questions that follow. 
Constituents Boiling point in 0C
Gases Below 40
Petrol 49-175
Kerosene 175-250
Diesel oil 259-350
Lubricating oil 350-400
Bitumen Above 400
i) Which one of the constituent of crude oil has molecules with the highest number of carbon atoms?      (2mks)
ii) Name the process you would use to separate a mixture of petrol and diesel and explain how the separation takes place.    (2mks)
iii) Explain why constituents of crude oil do not have sharp boiling points.
          (2mks)
iv) a) Name one gas that is likely to be a constituent of crude oil and
write its formula.      (2mks) 
b) What conditions could cause a poisonous gas to be formed when kerosene is burnt.  Explain.     (2mks)
c) Give one use of bitumen.     (1mk)
18. 
The diagram below shows a set up used by a student in an attempt to prepare collect oxygen gas
a) i) Complete the diagram by collecting the mistakes in it. (2mks)
  ii) Identify solid w.      (1mk)
b) A piece of phosphorous was burnt in excess air.  The amount of hot water to make a solution.
 i) Write an equation for the burning of phosphorous in excess air.          (1mk)
ii) The solution obtained in (b) above was found to have a PH of 2.0.  Give reasons for this observation.    (2mks)
c) Explain why cooking pots made of aluminium do not corrode easily when exposed to air.        (1mk)
d) The reaction between sulphure (IV) Oxide and oxygen to form Sulphur (VI) Oxide per day (condition for the reaction a catalyst, 2 atmospheric pressure and temperature between 4000 5000C)
2SO(aq)  +  O2(g)     2SO3(g)
Factory manufacturing sulphuric acid by contact process produces 350kg of sulphur trioxide per day (conditions) for the reaction catalyst. 2 atmospheres pressure and temperatures between 400 – 500 oC.
i) What is meant by an exothermic reaction?   (1mk)
ii) How would the yield per day of sulphur trioxide be affected  if temperatures lower than 400oC are used? Explain.  (1mk)
iii) All the sulphur (VI) Oxide produced was absorbed in concentrated sulphuric acid to form oleum.
 SO3(g)    + H4SO4(l)     →  H2S2O7(l)
Calculate the mass of oleum that was produced per day.
 (S+ 32.0, O= 16: H 1.0)     (3mks)
19. 
a) Fractional distillation of liquid air usually produces nitrogen and oxygen as the major by-product.
i) Name one substance that is used to remove carbon (IV) Oxide from air before it is changed into liquid.   (1mk)
ii) Describe how liquid Nitrogen gas is obtained from liquid air.
 Boiling points; Nitrogen = -1960C; Oxygen = -1830C. (1mk)
b) Study the flow chart below and answer the questions that follows

     i) Name element M.      (1mk)
ii) State and explain the change in mass that is likely to occur in tube N by the end of the experiment.    (2mks)
iii) Name two gases that come out through tube M.  (1mk)
iv) Write an equation for the reaction in stem 7.   (1mk)
v) Give one use of Ammonium –Nitrate.   (1mk)
         c) State and explain the observations that would be made if a sample of
sulphur is heated with concentrated Nitric  acid.  (Nitric (V) acid.
20. 
a) Candle wax is mainly a compound consisting of two elements.  Name the two elements
b) The up below was used to investigate the burning of candle.  Study it and answer the questions that follow.
i) What would happen to the burning candle if the pump were turned off? Give reasons.      (3mks)
ii) State and explain the change in mass that is likely to occur in tube N by the end of the experiment.    (2mks)
iii) Name another substance that would be used in place of calcium oxide.        (1mk)
21. Why is iron not used to make steam boilers?     (1mk)
22. Study the arrangement below and answer the questions that follows.

 Explain what happens to the lime water after some time.   (1mk)
23. When air is bubble through pure water (Ph 7.0).  The PH drops to 6.0.  Explain why.          (1mk)
24. A white compound was moistened with a little concentrated Hydrochloric acid and placed over a flame.  A yellow flame was observed.  Identify the metallic ions in the compound.        (1mk)
25. Magnesium ribbon was burned in a gas jar of Nitrogen.  A few drops of water were then added to the jar.  Write equation for the reactions in the jar.        (2mks)
 
26. The diagram below shows an experiment to compare the heating effect of luminous and non luminous flame.
 
a) What was observed at the bottom of each beaker at the end of the experiment?        (1mk)
b) Which sample of water boils first?  Give a reason for your answer. (2mks) 
c) Besides the amount of heat produced by the two flames, state other differences.        (2mks)
27. a) Study the equation below and answer the questions that follow.
  CO3-2(aq)    + H2O (l)    →     HCO (aq)   + OH (aq)
  Which substance is an oxidizing agent?  Give reasons.  (2mks)
 b) Identify the reducing agent in the equation below
  Fe2+(aq)   +  Cl2(g)    →  Fe3+(aq)    +  2CL –(aq)

28. A candle was burnt using the apparatus shown below.  The initial volume of measuring cylinder was 90cm3.  The apparatus was allowed to cool and the volume of air in the measuring cylinder had dropped to 70cm3.
 a) Why was the volume recorded when the air was cooled?  (1mk)
 b) What was the pupose of sodium Hydroxide?    (1mk)
 c) Use the results given to calculate the percentage of oxygen in air. (2mks)
29. The graph below shows the changes that occur when a pure and an impure substance are heated.
 a) Which curve represents pure substance?  Explain.    (2mks)
 b) Name one factor which affects the melting point of a solid and state
effects.         (2mks)
 
TOPIC 4
WATER AND HYDROGEN
1. 
Use the information shown in the diagram below to answer the questions that follows.
i) Explain why it is important to pass the hydrogen gas for some time before lighting it at point Z.       (1mk)
ii) Write an equation for the reaction that takes place when hydrogen burns at point Z.        (1mk)
2. 
The order of reactivity of metal p, R and T starting with the most reactive is R.T.P.  By using a tick () to indicate no reaction, complete the table below to show what happens when the metals of each are added to solutions containing ions of metal P, R and T. 

      (3mks)
 Aqueous solution containing ions of metal
Metal P R T
P   
R   
T   
3. 
In an experiment, soap solution was added to three separate samples of water.  The table below shows the volumes of soap solution required to form lather with 100cm3 of each sample of water before and after boiling.
 Sample 1 Sample 2 Sample 3
Volume of soap before water is boiled in cm3 27.0 3.0 10.6
Volume of soap after water is boiled in cm3 27.0 3.0 3.0
a) Which water sample is likely to be soft?  Explain.   (2mks)
b) Explain the change in volume of soap solution used in sample III (1mk)
4. 
Study the diagram below and answer questions that follow.

Write an equation for each of the two reactions that take place in the experiment represented by the diagram above      (2mks)
5. Zinc metal and hydrochloric acid react according to the following:
Zn(s) + 2HC (aq)     →        ZnCl2 (aq)   + H2 (g)
 1.9 g of zinc metal was reacted with 100cm3 of 0.2m Hydrochloric acid.
 a) Determine the reagent that was in excess.    (2mks)
b) Calculate the total volume of hydrogen gas that was liberated at S.T.P. (Zn= 65).  Molar gas volume= 22.4dm3   at STP.   (1mk)
6. The diagram below represents set-up that was used to react lithium with water vapour.  Study it and answer the questions that follow.
a) Write an equation for the reaction that takes place given that the atomic number of Lithium is 3.        (2mks)
b) Why would it not be advisable to use potassium in place of Lithium in the above set up?        (2mks)
7. A student set up the experiment below to collect gas K.  The glass wool was heated before heating the zinc powder.
a) Why was it necessary to heat the moist glass wool before heating zinc powder?        (1mk)
b) What would happen if the zinc powder was heated before heating the glass wool?         (1mk)
c) What property of gas K made it possible for it to be collected as shown in the diagram?        (1mk)
8. 
A sample of water drawn from a river passing through an agricultural district was divided into two portions.  The first portion gave a white precipitate when acidified barium chloride was added.  The second portion when warmed with aqueous sodium hydroxide gave a colourless gas which turned a moist red litmus paper to blue.
 a) Identify the ions present in the river water.    (2mks)
 b) Suggest the possible sources of the ions identified in (a) above     (2mks)
9. 
The column below was used to soften hard water.

a) Explain how the hard water was softened as it passed through column.          (1mk)
b) After sometime the material in the column is not able to soften hard water.  How can the material be reactivated?     (1mk)
c) Give one advantage of using hard water for domestic purposes.   (1mk)
10. 
The table below shows the test carried out on separate samples of water drawn from a well and results obtained.
Tests Results
I) Addition of excess ammonia solution White precipitate
II) Addition of two drops of dilute sulphuric acid No precipitate
III) Addition of dilute hydrochloric acid followed by few drops of Barium chloride White precipitate
 a) Identify the cation and anion present in the water.   (2mks)
 b) Write an ionic equation for the reaction which takes place in test (III)
           (1mk)
11. 
 Study the set up below and answer the question that follows.
a) Write an equation for the reaction which takes place in the combustion tube.         (1mk)
 b) What property of gas z allows it to be collected as shown in the diagram?
           (2mks)
12. 
10g of sodium hydrogen carbonate were dissolved in 20cm3 of water in a boiling tube.  Lemon juice was then added dropwise with shaking until there was no further observable change.
a) Explain the observation which was made in the boiling tube when the reaction was in progress.      (2mks)
b) What observation would have been made if the lemon juice had been added to copper turnings in a boiling tube?  Give a reason.  (1mk)
13. a) State one cause of temporary hardness in water.   (1mk)
 b) How does distillation remove hardness in water?   (3mks)
14. 
Explain why hydrogen forms compounds in which its oxidation state is either + 1 or -1 (atomic number of H =1)      (3mks)
15. 
An atom of hydrogen can form two ions.  Write two equations to show how a neutral atom of hydrogen can form the two ions.  In each case show the sigh of the energy changes.        (2mks)
16. 
When steam was passed over heated charcoal as shown in the diagram below hydrogen and carbon (II) oxide were formed.

 a) Write the equation for the reaction which takes place.  (1mk)
b) Name two uses of carbon (III) Oxide gas which are also the uses of Hydrogen gas.        (2mks)
17. 
The table below shows the test carried out on a sample of water and the results obtained.
Tests Results
i) Addition of sodium Hydroxide White precipitate which dissolves in excess
ii) Addition of excess ammonia solution Colourless solution obtained.
iii) Addition of dilute Hydrochloric acid and barium chloride White precipitate
 
a) Identify the anions present in water.     (1mk)
 b) Write an ionic equation for the reaction in (iii)   (1mk)
 c) Write the formula of the complex ion formed in (ii)    (1mk)
18. 
 The set up below used to demonstrate the effect of heat on hard water.

a) Name substance, A       (1mk)
 b) Explain why the heating of hard water produces substance A. (2mks)
19. 
The diagram below shows a student’s set up for the preparation and collection of hydrogen gas.

a) How would the final volume of hydrogen gas produced be affected if 80cm3 of o.7M hydrochloric acid was used?    (1mk)
b) Give a reason why helium is increasingly being preferred to hydrogen in weather balloons.       (1mk)

20. 
In a laboratory experiment hydrogen gas was passed over heated copper (II) oxide as shown in the diagram below.

 Describe a chemical test that can be used to identify the product E.  (2mks)
21. a) A student was supplied with a colourless liquid suspected to be water.
i) Describe one chemical test that could have been used to show that the liquid was pure water.     (1mk)
ii) How it could have been shown that the liquid was pure water.              (1mk)
b) The flow chart below shows the various stages of water treatment.  Study it and answer the question that follows.

i) Which substances are likely to be removed in filtration unit I?          (1mk)
ii) What is the name of process Y?    (1mk)
iii) What is the purpose of;
 a) Process Y
 b) Addition of sodium hypochlorite?   (1mk)
iv) It was confirmed that magnesium sulphate was present in the tap water.
a) What type of hardness was present in the tap water? (1mk)
   b) Explain how this hardness can be removed.  (2mks)
22. 
 The set up below was used to prepare hydrogen gas.
a) Complete the diagram to show how a dry sample of hydrogen gas can be collected.        (3mks)
b) Write an equation which takes place when hydrogen gas burns in air.
(1mk)
c) 1.2 litres of hydrogen gas was produced at room temperature and pressure when 3.27g of zinc were used.  Determine the relative atomic mass of zinc (molar gas volume is 24 litres).     (4mks)
d) State two industrial used of hydrogen gas.    (2mks)
23. 
The set up was used to collect gas F, produced by the reaction between water and calcium metal.

i) Name gas F.        (1mk)
ii) At the end of the experiment, the solution in the solution is a weak box.
          (2mks)
iii) Give one laboratory use of the solution formed in the beaker.    (1mk)
24. A piece of sodium was put into a beaker containing water.
 a) Write the equation for this reaction.     (1mk)
b) State the observations made in the above reaction.   (2mks)
25. When Na2CO3: XH2O is strongly heated, it loses 63.2% of its mass.  Find the value of X.         (2mks)
26. Study the diagram below and answer the questions that follows:
 a) State two observations that may be made in the combustion tube. (1mk)
 b) Write an equation for the reaction of hydrogen with Lead (II) Oxide.
(1mk)
27. The table below gives information on reactions of metals B, C, D and E. 
Metal Reaction with acid Action of heat on its nitrate  Reaction with water
B Hydrogen evolved Oxide formed No reaction
C No reaction Metal formed NO reaction
D Hydrogen evolved Oxide formed Hydrogen evolved
E NO reaction  Oxide formed  NO reaction
Arrange the metals in the order of decreasing reactivity starting with the least reactive.
28. The diagram below shows how lithium reacts with steam.

 i) Write an equation for the reaction.     (1mk)
 ii) Why is it not advisable to use potassium in place of lithium? (1mk)
29. Steam reacts with iron fillings to form tri-iron tetra oxide.
 3Fe(s)   + 4H2O (g)    →    3H2 (g)    +   4H2 (g)
 a) State one experimental condition that will make the reaction reversible.
(1mk)
 b) Give two commercial uses of Hydrogen gas.    (2mks)
30. When a metal oxide of element “W” reacts with hydrogen, the equation for the reaction is:
   WO3(s) + 3H2 (g)    →   W(s)   +     3H2O (2)
 Comment on the reactivity of element “W” with hydrogen gas.  (1mk)
31. The following observations were made during the investigation of the reaction of metal with water.
- When a piece of sodium metal was dropped in a bowl; of water, it reacted vigorously, darting over the surface of water.  Hydrogen gas was liberated.
- Iron metal did not react with cold water but red hot iron reacted with steam liberating hydrogen an tri- iron tetra oxide.
- Copper did not react with cold water but red hot iron reacted with steam liberating hydrogen a tri-iron tetra oxide.
- Copper did not react with water or steam.
Answer the following questions
 a) Which metal is;
  i) The most reactive?      (1mk)
  ii) The least reactive?      (1mk)
b)   i) What other product apart from hydrogen is formed in the reaction
between sodium and water?     (1mk) 
           ii) Write a chemical equation for the reaction in (b) above  (1mk)
c) Comment on the PH of the resulting solution in (b) above. (1mk)
d) Name any other two elements which react in similar way to sodium        (2mks)
e) Give the test for hydrogen gas.    (1mk)
32. What is the differences between a deliquescent and hygroscopic substance? 
(2mks)
33. When trying to put off an oil fire, water is not used.  Explain.  (2mks)

 
FORM 2 WORK
TOPIC 1
STRUCTURE OF THE ATOM AND THE PERIODIC TABLE
1. Complete the table below.      (1 ½ mks)
Isotope Number of
 Protons Neutrons Electons
59
    Co
27   
2. 
The electron arrangement of ions X3 + and Y-2 are 2:8 and 2:8:8 respectively.
 a) Write the electron arrangement of elements “X” and “Y”      (2mks)
b) Write the formula of the compound that would be formed between X and Y.         (1mk) 
3. 
With reference to its atomic number of one explain why hydrogen can be placed in either group I or VII on the periodic table.     (2mks)
4. 
 An element Y has the electronic configuration of 2:8:5
 a) Which period of the periodic table does the element belong.  (1mk)
b) Write the formula of the most stable anion formed when element Y ionizes.        (1mk)
c) Explain the difference between the atomic radius of element Y and ionic radius.         (1mk)
5.                 34
 An ion of phosphorous can be presented as    P -3
                15
Draw a diagram to show the distribution of the electrons and the composition in the nucleus of the ion of phosphorous.     (2mks)
6. 
The grid below shows part of the periodic table.  The letters do not represent the actual symbols of the element.
         
     G    
     H   I 
F         
a) Select
i) Element which has the largest atomic radius   (1mk)
ii) Most reactive non- metal
b) Show on the grid the position of element “J” which forms J-2 ions with electronic configuration 2:8:8:8     (1mk)
 
7. Study the information in the table below and answer questions that follows;
Ions Electron arrangement Ionic radius
Na+ 2,8 0.95
K2+ 2,8,8 0.133
Mg2+ 2,8 0.065
Explain why the ionic radius of
a) K+ is greater than that of Na+     (1mk)
b) Mg2+ is smaller than that of Na+     (2mks)
8. 
An atom of hydrogen can form two ions.  Write down two equations to show how the neutral atom of each case show the sign of the energy change involved. (2mks)
9. 
 Brass is an alloy of zinc and copper.  Give one used brass   (1mk)
10. 
Use the information in the table below to answer questions that  follows.  That follows.  The letters do not represent the actual symbols of the elements.
Elements B C D E F
Atomic numbers 18 5 3 5 20
Mass Numbers 40 10 7 11 40
a) Which two letters represent the same element? Give a reason   (2mks)
 b) Give the number of neutrons in an atom of element D  (1mk)

11. 
The table below gives some information about elements I, II, III and IV which are in the same group of the periodic table.
Use the information to answer the questions that follows.
Element First ionization energy K 5 mol -1 Atomic radius (nm)
I 520 0.15
II 500 0.19
III 420 0.23
IV 400 0.25
State and explain the relationship between the variation in the first ionization energies and the atomic radii.       (3mks)
 12. 
The table below shows the relative atomic masses and the percentage abundance of the isotopes L1, L2 of element L
 Relative atomic masses % abundance
L1 62.93 69.09
L2 64.93 30.91
 Calculate the relative atomic mass of element L.      (3mks)
13. 
Explain why there is general increase in the first ionization energies of the elements in period 3 of the periodic table from left to right.   (2mks)

14. 
The table below shows the number of valance electrons of the elements P, Q and R. 
Element P Q R
Number of valence electrons 3 5 2
a) Explain why P and R would not be expected to form a compound. (1mk)
b) Write an equation to show the effect of heat on the carbonate of R  (1mk)
c) Write the formula for the most stable ion or Q.   (1mk)
15. a) What are isotopes?       (1mk)
       18
 b) Determine the number of neutrons in    O
        8    (1mk)
16. 
The grid below is part of the periodic table.  Use it to answer the questions that follow.  (The letters are not the actual symbols of the elements)
         
       R S 
N Q       T U
P        
        
a) Indicate on the grid the position of an element represented by letter V whose atomic number is 14.      (1mk)
b) Select a letter which represents a monoatomic gas.   (1mk)
c) Write an equation for the reaction between Q and T.   (1mk)
17. 
The table below gives elements represented by letters T, U, V, w, x, Y their atomic numbers.
Elements T U V W X Y
Atomic numbers 12 13 14 15 16 17
Electronic arrangement      
Use the information in the table to answer the questions below
a) Complete the above table giving the electron arrangement of each of the element        (2mks) 
b) In which period of the periodic table do these elements belong? Give a reason.         (2mks)
c) How does the atomic radius of V compare with that of X. Explain? (2mks)
d) Give the formula of the compound that could be termed between “U” and “W”          (1mk)
e) What type of bonding will be present in a compound formed between T and Y?   Explain        (2mks)
f) Arrange the species T2+ T+    and T in increasing order of size
g) Which are the ions X+2 and X-2 is most suitable?  Explain   (2mks)
h) Give the fomula of
i) An acidic oxide formed when one of the elements in the table is heated in air        (1mk)
ii) A basic oxide formed when one of the elements in the table is heated in the air.      (1mk)
18. Study the table below and answer the questions that follows:-
Elements Atomic numbers Relative atomic mass Melting point 0C
Aluminium 13 27.0 1020
Calcium 20 40.0 850
Carbon - 12.0 3730
Hydrogen - 1.0 -249
Magnesium 12 24.3 650
Neon 10 - -249
Phosphorus 15 31.0 442 white
590 red
Sodium - 23 97.8
a) Complete the table by filling in the missing atomic numbers and atomic masses         (2mks)
b) Write the electron arrangement for the following ions
 i) Ca2+        (1mk)
 ii) P-3        (1mk)
c) What is the melting point of hydrogen in degrees Kelvin  (1mk)
d) Which of the two allotropes of phosphorous has a higher density? Explain
          (2mks)
e) The mass numbers of the three isotopes of magnesium are 24, 25 and 26.  What is the mass number of the most abundant isotope of Magnesium?  Explain         (2mks)
f) Give the formula to the compound formed between aluminium and carbon.         (1mk)
19 
The grid given below represents part of the periodic table.  Study it answer the questions that follows. The letters do not represent the actual symbols of the elements.
         
         
 C    B    
 F   D   E  
         
i) What name is given to the group of elements to which “C” and “F” belong?         (1mk)
 ii) Which letter represents the element that is least reactive? Explain. (2mks)
 iii) What type of bond is formed when B and E reacts?  Explain.  (2mks)
iv) On the grid indicate with a tick the position of an element G which is in the third period of the periodic table and terms G-3 ion.  (1mk)
 
20. 
Study the information in the table below and answer the questions that follow.  The letters do not represent the actual symbols of the elements 
Elements Electronic configuration Ionization energy kj mol -1
P 2,1 519
C 2,8,1 494
R 2,8,8,1 418
i) What is the general name given to the group which elements P, C and R belongs?        (1mk)
ii) What is meant by ionization energy      (2mks)
iii) Explain why element p has the highest ionization energy.  (2mks)
iv) a) When a piece of element “C” is placed on water.   It melts and
hissing sound is produced as it moves on the surface of the water.  Explain these observations     (2mks)
b) Distinguish between a strong and a weak base.  Give an example of each.       (2mks)
c) Neutralization is one of the methods of preparing salt
 i) What is meant by neutralization   (1mk)
ii) Describe how you would prepare crystals of sodium nitrate starting with 200 cm3 of 2m sodium hydroxide. (3mks)
iii) Write an equation for the reaction that takes place when a solid sample of sodium nitrate is heated.  (1mk)

21. a) The chart below is an outline of part of the periodic table
i) With the help of vertical and horizontal lines, indicate the direction of increasing metallic nature of elements.     (2mks)
ii) Which type of elements are represented in the shaded area? (1mk)
b) i) Element “A” is in the same group of the periodic table as chlorine. 
Write   the formula of the compound formed when “A” react with potassium metal       (1mk)
ii) What type of bonding exists in the compound formed in b (i) above? Give a reason for your answer    (3mks)
c) Starting with aqueous magnesium sulphate, describe how you would obtain a sample of magnesium oxide.     (3mks)
d) Write two ionic equations to show that aluminium hydroxide is amphoteric         (2mks)
22. Brine usually contain calcium and magnesium salts.  Explain how sodium carbonate is used to purify brine.      (2mks)
23. The table below gives information about elements A1, A2, A3 and A4
Element Atomic number Atomic radius (nm) Ionic radius (nm)
A1
A2
A3
A4 3
5
13
17 0.134
0.090
0.143
0.099 0.074
0.12
0.050
0.181
 i) In which period of the periodic table is element A2
  Give reason.        (2mks)
 ii) Explain why the atomic radius of:
  I. A1 is greater than that of A2     (2mks)
  II. A4 is smaller than its ionic radius.    (2mks)
 iii) Select the element which is in the same group as A3   (1mk)
iv) Using dots (.) and crosses (x) to represent outermost electrons, draw a diagram to show the bonding in the compound formed when A1 reacts with A4          (1mk)
24. Using the table below explain the following
Ions Na+ Mg2+ Al3+ K+
Ionic radius 0.086 0.073 0.064 0.097
 a) Ionic radius of Na+ is less than that of K+.  Explain    (1mk)
b) Sodium, magnesium and aluminium belong to the same period in the periodic table.  Explain the trend in their ironic radii.  (3mks)
25. Study the information in the table below and Answer questions that follows.
W X Y Z
Glows red hot when heated.
Does not react with water but turns red brown on surface when left outside over night. Forms a ball on the surface of water and react.
Produce a hissing sound.
Burning in air with a yellow orange flame. Burns with dazzling fame and does not react with cold water. Burns with a red flame and produce hydrogen with cold water.
 a) Identify the above metals     (1 ½mks)
b) Arrange the metals according to their reactivity starting with the most reactive.        (1mk)
26. Element Z in the second period of the periodic table forms Z 3+ ions using (x) to represent electrons; draw a complete structure of an isotope of “Z” having mass number 8.         (3mks)
The table below gives information on the some elements.  The letters are not actual symbols of the elements.  Study it and it to answer the questions that follow.

Elements Ionization energy
(k j) Atomic radius (NM) Ionic radius (NM)
L 410 0.154 0.091
G 380 0.192 0.097
Q 490 0.108 0.086
 a) Select the most reactive element and give reasons for your answer. (2mks)
 b) Do this element represent metallic or nonmetallic group. Explain.    (2mks)
27. The table below shows part of periodic table for some elements represented by Q, R, T, V, W, X, Y and Z.  The letters do not represent the actual symbols of the elements.  Study it and answer the questions that follows.
T1        T1 2
Q3 4   5 W6 7 8 V9 10
R11 12   13 14 15 16 X17 Y18
19 20        
a) i) Explain why element T has been placed in two positions in the
periodic table.       (1mk)
  ii) What is the name of the chemical family to which q and R belong?
  iii) Elements Y is generally unreactive.  Explain   (1mk)
   b) i) Explain the difference in atomic radius of atoms of elements X and
Y.        (1mk)
  ii) V is more reactive than W Explain     (1mk)
c) i) Draw cross (x) and dots (.) diagram to show bonding between
“W” and “T” to form compound WT4   (2mks)
  ii) Explain why WT4 have low melting point and does not dissolve in
water         (2mks)
d) Element X consist of two isotopes whose mass numbers are 35 and 37 exist in the ratio of 3:1 respectively.    
i) Draw the atomic structure of the isotope whose mass number is 35 and atomic structure of the isotope whose mass number is 35 and atomic number 17.      (2mks)
  ii) Determine the relative atomic mass of element X   (2mks)
28. a) What is an isotope?
 b) Determine the relative atomic mass of argon whose isotope mixture is 
  36. Ar (0.34%) 38Ar (0.06%) 40 Ar (99.6%)
  18     18    18
29. An element “z” has a mass number of 33 and has 18 neutrons
 a) What is the atomic number of element Z?     (1mk)
 b) Write an equation to show how atom of “z” forms an ion.  (1mk)

30. Study the flow chart below and answer the questions that follows-:
a) Name
  i) Gas P        (1mk)
  ii) Compound T       (1mk)
  iii) Gas U         (1mk)
 b) Give the chemical test that you would use to identify
  i) Gas P        (1mk)
  ii) Gas U        (1mk)
31. Element E has atomic numbers 15
 a) Write the electronic arrangement for an atom of “E”   (1mk)
 b) Explain why “E” forms a chloride which is a liquid of low boiling point.
           (2mks)
32. An element “H” consist of isotopes of mass “10” and “11” with a percentage composition of 18.7% and 81.3% respectively.  Determine the RAM of H. (2mks)

 
TOPIC 2
CHEMICAL FAMILIES.
1. 
The table below gives the atomic numbers of elements W X Y and Z.  The letters
do not represent the actual symbols of the elements.
Element W X Y Z
Atomic numbers 9 10 11 12
 a) Which one of the elements is less reactive? Explain.   (2mks)
 b) i) Which two elements would react most vigorously with each other
ii) Give the formula of the compound formed when elements in b (i) above react       (1mk)
2. 
The table below gives the energy required to remove the outer most electrons from same group
Elements I II III IV
Energy kj /Mole 494 418 519 376
Arrange the electrons in the order of their reactivity starting with the most reactive.            (2mks)
3. 
The information below relates to elements s, T, U, and x.  The letters do not represent the actual symbols of the elements.
i) “T” displaces “X” from aqueous solution containing ions of “X”
ii) Hydrogen gases reduces heated oxide of “s” but does not reduce the heated oxide of “X”
iii) “U” liberates hydrogen gas from cold water but “T” does not
a) Write an equation for the reaction between “T” and ions of “X” both T and X are in the group II of the periodic table  (1mk)
b) Arrange the elements in order of their increasing reactivity (1mk)
4. 
 The electronic structures for elements represented by letters A, B, C, and D are:-
 A= 2, 8, 6 B= 2, 8, 2 C= 2, 8, 1 D= 2, 8, 8
 a) Select the element which forms
  i) Double charged cation     (1mk)
  ii) A soluble carbonate      (1mk)
 b) Which element has the smallest atomic radius    (1mk)
5. 
The information in the table below relates to elements in the same group of the periodic table.  Study it and answer the questions that follows:-
Elements Atomic size (mm)
G1 0.19
G2 0.23
G3 0.15
Which element has highest ionization energy?  Give a reason.  (3mks)
6. The oxides of elements “A” and “B” have the properties shown in the table below.  The letters do not represent actual symbols of the elements.
A B
A gas at room temperature Solid normal temperature
Dissolves in water to form acidic solution Dissolves in water to form alkaline solution
 Give one example of element “A” and “B”     (2mks)
7. 
 An oxide of F has the formula F2O5
 a) Determine the oxidation state of “F”     (1mk)
 b) In which group of the periodic table is element “F”   (1mk)
8. 
Yellow phosphorus reacts with chlorine gas to form a yellow liquid. The liquid fumes when exposed to air.  Explain these observations.   (2mks)
9. 2003
 Explain why the reactivity of group (VII) elements decreases down the group.
           (3mks)
10. 
The atomic numbers of element “C” and “D” are 19 and 9 respectively.  State and explain the electro conductivity of compound CD in:-  
a) Solid state       (1 ½ mark)
b) Aqueous state       (1 ½ mark)
11. 
a) Explain why the metals magnesium and aluminium are good conductors of electricity.        (1mk) 
b) Other than cost, give two reasons why aluminium is used for making electric cables while magnesium is not.    (2mks)
12. 
The table below gives information on four elements represented by letters K, L, M and N.  Study it and answer the questions that follow.  The letters do not represent the actual symbols of the elements. 

Elements Electron arrangement Atomic radius (nm) Ionic radius
K 2,8,2 0.136 0.065
L 2,8,7 0.099 0.181
M 2,8,8,1 0.099 0.181
N 2,8,8,2 0.174 0.099
 a) Which two elements have similar chemical properties? Explain (2mks)
b) What is the most likely formula of the oxide of “L”     (1mk)
c) Which element is a non-metal?  Explain    (2mks)
d) Which one of the elements is the strongest reducing agent? Explain (2mks)
 e) Explain why the ionic radius of “N” is less than that of “M”  (2mks)
 f) Explain why the ionic radius of “L” is larger than its atomic radius. (2mks)

13. 
Study the information given in the table below and answer the questions that follow.  The letters do not represent the actual symbols of elements.
Elements  Atomic numbers  Boiling point
S 3 1603
T 13 2743
U 16 718
V 18 87
W 19 1047
 a) Select the element which belong to the same
  i) Group        (1mk)
  ii) Period        (1mk)
 b) Which element
  i) is in gaseous state at room temperature? Explain   (2mks)
   Take room temperature to be 298K
  ii) Does not form oxides      (1mk)
 c) Write the:-
  i) Formula of the nitrate of element T    (1mk)
  ii) Equation for the reaction between element “S” and “U”  (1mk)
d) What type of bond would exist in the compound formed when element “U” and “T” react? Give a reason for your answer   (2mks)
e) The aqueous sulphate of element “w” was electrolyzed using inert electrodes.  Name the products formed at the
 i) Cathode       (1mk)
 ii) Anode        (1mk)
14. The table below shows some properties of chlorine, bromine and iodine.
Elements Formulae Colour and state at room temperature Solubility in water
Chlorine Cl2 (i)………. Soluble
Bromine Br2 Brown liquid (ii)………
Iodine I2 (iii) ………… Slightly soluble
 a) Complete the table below by giving the missing information in (i) (ii)
           (3mks)
b) Chloride is prepared by reacting concentrated hydrochloric acid with Manganese (IV) oxide.
i) Write the equation for the reaction between concentrated hydrochloric acid and manganese (IV) oxide.
ii) What is the role of manganes (IV) oxide in this reaction (1mk)
c)         i) Iron (ii) chloride reacts with chlorine gas to form substance “E”.  Identify substance “E”     (1mk)
           ii) During the reaction in c (i) above, 6.30g of iron (II) chloride were converted to 8.06g of substance “E”.  Calculate the volume of chlorine gas used. (Cl=35.5) molar gas at room temperature = 24000 cm3 (Fe= 56)      (3mks)
d) Draw and name the structure of the compound formed when excess chlorine gas is reacted with ethane gas.   (2mks)
 Structure…………………………………..
 Name …………………………………….
15. The grid below represents part of the periodic table.  Study it and answer the questions that follows:- The letter given do not represent the actual symbols of the elements.
         
       A  
 B   C  D  E 
F G       
       H 
i) Select the element that can form an ion with a change of-2. Explain your answer.        (2mks)
ii) What type of structure would the oxide of C have? Explain your answer.
          (2mks)
iii) How does reaction of H compare with that of E?   (2mks)
iv) 1.3g of “B” react completely when heated with 1.21 litres of Cl2 (g) at STP.  (1 mole of gas of STP occupies 22.4 litres)
 I) Write a balanced equation for the reaction between B and Cl2(1mk)
  Ii)      Determine the relative atomic mass of B.   (2mks)
 v) Explain how you would expect the following to compare.
  a) Atomic radii of “F” and “G”     (1mk)
  b) The pH values of aqueous solution of oxides of B and D. (2mks)
vi) The table below shows some physical properties of some substances.  Use the information in the table to answer the questions that follow:-
    Electrical conductivity
Substances Melting Boiling point 0C Solid Solid
U 1083 2595 Good Good
V 801 1413 Poor Good
W 5.5 80.1 Poor Poor
X -114.8 -84.9 Poor Poor
Y 3550 4827 Poor poor
 i) Which substance is likely to be     (1mk)
         (I) A metal      (1mk) 
         (II) Liquid at room temperature    (1mk)
          ii)       Which substance is likely to have the following structures?
       (I) Simple molecular     (1mk)
       (II)    Giant atomic      (1mk)
16. Lithium, sodium and potassium belong to the same group of the periodic table
 i) Arrange the elements in the order of increasing ionization energy. (1mk)
 ii) Explain the trend in 2(i) above     (2mks)
17. When heated in a current of Nitrogen gas, magnesium reacts to form a compound magnesium nitride, Mg3N2
a) Calculate a volume of Nitrogen at s.t.p required to react with 8g of magnesium (Mg= 24) molar gas volume at s.t.p= 22.4 dm3)  (3mks)
b) Magnisium Nitrite reacts with water to form magnesium hydroxide and ammonia. Calculate the volume of ammonia produced at S.T.P, if all magnesium nitride formed reacts completely with water.  (3mks)
18. A student at Loreto Secondary school used 2g of calcium to prepare hydrogen gas in the laboratory.  He used the set up below.

 a) Write a chemical equation for the reaction that produced hydrogen (1mk)
b) Calculate the volume of hydrogen produced at room temperature (molar gas volume= 24,000cm3)       (2mks)
c) Explain why the same method cannot be used to prepare hydrogen using sodium in the laboratory       (2mks)
d) Explain why the same method cannot be used to prepare hydrogen using sodium in laboratory        (2mks)
e) Calculate the mass of the products formed if all the hydrogen produced in this experiment was burnt in excess air.    (3mks)
f) Explain how calcium is able to conduct electricity    (2mks)

19. The table below gives atomic and mass numbers of some elements represented by letters “T” to “Y”. The letters are not actual symbols of elements. Use it to answer questions that follows:-
Elements T U V W X Y
Atomic numbers 1 18 1 19 20 17
Mass numbers 2 39 1 39 40 35
 a) Which element has the lowest ionization energy?   (2mks)
b) Element “V” is uniquely positioned in the periodic table.  It has a tendency of forming compounds by either gaining or sharing electrons. Give the formula of a compound of “V” that is formed when V gain an electron.          (1mk)
20. When magnesium metal burn metal burn in air.  It reacts with both oxygen and Nitrogen gases giving a white ash- like substances.  Write two equations for the two reactions that take place.       (2mks)
21. Chlorine and iodine are elements in the same group in the periodic table.  Chlorine gas is yellow while iodine solution is brown.
a) What observations would be made if chlorine gas is bubbled through aqueous sodium iodide?  Explain using an ionic equation.  (1mk)
b) Under certain conditions chlorine and iodine react to give iodine trichloride (LCl3 (s)) . What type of bonding would you expect to exist in iodine trichloride? Explain.      (1mk)
22. It is not appropriate to refer to group VIII elements as “inert gases” Explain giving an example.        (2mks)
23. What observations will you make when chlorine gas is bubbled through
 i) Potassium bromide       (1mk)
 ii) Potassium chloride       (1mk)
 iii) Explain these observations      (3mks)
24. Explain why the reactivity of group (VIII) elements decreases down the group.
           (3mks)

 
TOPIC 3
STRUCTURES AND BONDING 
1. When electric current is passed through two molten substances “M” and “N” in different containers.  The observation in the table below were made. 
Molten “M” Conduct electricity current and is not decomposed.
Molten “N” Conduct electric current and gas is formed at one of the electrodes.
Suggest the type of bonding present in substances “M” and “N”  (2mks)
2. 
a) Using dot (.) and crosses (x) to represent electrons draw diagrams to represent the bonding in NH3 and NH4     (1mk)
b) State why Ammonia molecule NH3 can combine with H to form NH4
 (Atomic numbers: N= 7 and H= 1)
3. 
Explain why aluminium chloride is fairly soluble in organic solvents while anhydrous magnesium chloride is insoluble     (2mks)
4. 
Using (•) crosses (x) to represents electrons. Draw diagrams to show bonding in CO2 and H3O+ (atomic numbers) (H=1, C=6, O=8)    (2mks)

5. 
The table below shows some properties of substances C, D and E. Study it and answer the questions that follows:
 Elements M.P 0C Solubility in water Electrical conductivity
   Solid Molten
C -39 Insoluble Good Good
D 1610 Insoluble Poor Poor
E 801 Soluble Poor good
Select a substance
(a)  With a giant molecular structure     (1mk)
(b)  That is not likely to be an element     (1mk)
6. 
Diamond and graphite are allotropes of carbon in terms of structures and bonding. Explain the following
(a)  Diamond is used to drill through hard rock.     (1mk)
(b)  Graphite is used as a lubricant     (1mk)
7. 
A hydrocarbon slowly decolourises bromine gas in presence of sunlight but does not decolourise acidified potassium manganate (VII). Name and draw the structural formula of the fourth member of the series to which the hydrocarbon belongs         (2mks)

8. 
 What type of bond is formed when lithium and fluorine react?
 Atomic number (Li= 3 F = 9) Explain     (2mks)
9. 
When solid magnesium carbonate was added to solution of hydrogen chloride in methyl benzene, there was no apparent reaction on addition of water to the resulting solution/ mixture, there was vigorous effervescence. Explain these observations         (2mks)
10. 
Compound “Q” is a solid with a giant ionic structure. What forms would the compound conduct an electric current? Explain    (2mks)
11. 
The melting point of phosphorous trichloride is 900  C while that of magnesium chloride is 7150C in terms of structures and bonding.  Explain the differences in their melting points.        (3mks)
12. 
 Name one property of neon that makes it possible to be used in electric lamps.
           (1mk)
13. 
With reference to iodine distinguish between covalent bonds and van der waals forces.          (2mks)

14. 
The table below gives some information about electrical conductivity and likely bonding in substances N, P and Q.  Complete the table by inserting the missing information in spaces numbered I, II, and III     (3mks)
Substances Likely type of bonding Electric conductivity
Molten                     Solid
N Metallic  I                               Conduct
P II Does not conduct    Conduct
Q III Do not conduct        Does not conduct
15.
 a) What is meant by heat of vaporization?    (1mk)
 b) The boiling points of ethanol, propanal and butanol are 780C, 97.2oC and
1170C. Explain this trend.      (1mk)
16. 
Use dot (.) and crosses (x) to represent electrons, show bonding in the compounds formed when the following elements reacts (Si= 4, Na = 11, Cl = 17)
 a) Sodium and chlorine       (1mk)
 b) Silicon and chlorine       (1mk)
17. In terms of structures and bonding explain why graphite is used as a lubricant  
(2mks)

18. a) Distinguish between a covalent bond and a co-ordinate bond. (2mks)
 b) Draw a diagram to show bonding in ammonium ion.
  (N=7) (H=1)  (NH+4)       (1mk)
19. Explain why the boiling point of ethanol is higher than that of hexane.  Relative molecular mass of ethanol is 46 while that of hexane is 86.
20. Both chlorine and iodine are halogens
 a) What are halogens?       (1mk)
b) In terms of structure and bonding.  Explain why the boiling point of chlorine is lower than of iodine.     (2mks)
21. The diagram below is a section of a model of the structure of element t.
 a) State the type of bonding that exists in T.    (1mk)
 b) In which group of the periodic table does element T belong? Give reason.
           (2mks)

22. 
The table below gives atomic numbers of elements represented by the letters A,
B, C and D
Element A B C D
Atomic number 15 16 17 20
Use the information to answer the questions that follow.
a) Name the type of bonding that exists in the compound formed when A and D react.                                                                                              (1mk)
b) Select the letter which represents the best oxidizing agent.  Give a reason for your answer.       (2mks)
23. 
Study the information to answer the questions that follow. The letters do not represent the actual symbols of the elements.
Elements  Atomic number  Melting point (0C)
L 11 97.8
M 13 660
N 14 1410
C 17 -101
R 19 63.7
a) Write the electron arrangement for the ions formed by elements “ M” and  “C”         (2mks)
 b) Select an element which is
  i) The most reactive non-metal     (1mk)
  ii) A poor conductor of electricity    (1mk)
 c) In which period of the periodic table does element “R” belongs? (1mk)
 d) Element R loses its outermost electrons more readily than “L”. Explain
           (2mks)
e) Using dots and crosses to represent electrons, show bonding in the compound formed between N and Ca.
24. 
The following diagrams show the structures of two allotropes of carbon. Study them and answer the questions that follow:-

i) Name the allotrope
 M         (1mk)
 N         (1mk)
ii) Give one use of N       (1mk)
iii) Which allotrope conducts electricity? Explain   (2mks)
25. a) The diagram below represents part of the structure of a sodium chloride
crystal.  The position of one of the sodium ions in the crystal is shown as

  i) On the diagram, mark the positions of the other three sodium ions
           (2mks)
ii) The melting and boiling points of sodium chloride are 8010C and 14230C respectively.   Explain why sodium chloride does not conduct electricity at 250C and 14130C.     (2mks)
b) Give a reason why ammonia gas is highly soluble in water.    (2mks)
 c) The structure of an ammonium ion is shown below.
  Name the type of bond represented in the diagram by N H  (1mk)
26. Hydrogen reacts with iodine according to the equation give below.
 H2(g)  + I2                        2HL (g) ∆H =+ve
 In terms of bond energy explain why H is positive.   (2mks)
27. The molecular mass of hydrogen sulphide is 34 while that of water is “18”. 
Explain why the boiling of water is higher than that of hydrogen sulphide.  (2mks)
28. Using dots (.) and crosses(x) to represent electrons.  Draw a diagram to show bonding in carbon (II) oxide.  (C= 6, O = 8)     (2mks)
29. Explain what happens when atoms are bonded together by
 i) Ionic bond        (1mk)
 ii) Covalent bond        (1mk)
30. Explain the following statements
 i) Solid sodium conducts electricity but is not electrolyte  (1mk)
 ii) Solid iodine does not conduct electricity.    (1mk)
 iii) Solid sodium iodide has a giant ionic structure but does not conduct
electricity whereas liquid sodium iodide and aqueous solution of sodium iodide are both electrolytes.      (2mks)
31. A certain substance has a boiling point of 16800C. It does not conduct electricity when in solid form but conducts when molten. What is the most likely structure of the substance?  Explain.       (2mks)
32. Study the table below and answer the questions that follows:-
Substance  Formula Molar heat of vaporization kj/mole Melting points
Carbon disulphide CS2 27.2 -111
Calcium chloride Cacl2 149 782
Ethanol C2H5OH 43.5 -117
 a) Which of the substance above have crystalline structure?  Explain. (2mks)
 b) What is the best term to describe the structure of ethanol   (1mk)
c) Why is molar heat of vaporization of ethanol greater than that of carbon disulphide?        (2mks)
33. Study the table below and answer the questions that follows.
Substances  Mp0c BP0C Electrical conductivity
   Solid Liquid
U 1083 2595 Good Good
X 801 1413 Poor Good
W 5.0 80 Poor Good
V -115 -84 Poor Good
Y 355 4827 Poor Good
a) Which substances is likely to be
  i) A metal. Explain      (2mks)
  ii) A liquid at room temperature     (1mk)
b) Which substance is likely to have following structure?
  i) Simple molecular      (1mk)
  ii) Giant atomic structure      (1mk)
34. Explain why at room temperature hexane is a liquid while methane is a gas. 
(2mks)
35. Study the table below and answer the questions that follows. 
Substance A change heat in air  Melting point 0C Thermal conductivity
E Unreactive High  Poor
F Reactive High Poor
G Unreactive High Good
H Unreactive Low Good
Select the substance that would be most suitable.
 a) For making a cooking pot      (1mk)
 b) A thermal insulator       (1mk)
 
TOPIC 4
SALTS
1. Study the solubility curves below and answer the questions that follows-
a) At what temperature would equal amounts of potassium nitrate and calcium ethanoate dissolve in 100g of water?   (1mk)
b) Explain how you would prepare a saturated solution containing 80g of potassium nitrate in distilled water      (1mk)
c) A student added 30g of calcium ethanoate to 100g of boiling water and noticed that not all of it dissolved.  Explain what would happen if the student cools the mixture with stirring up a temperature of 100C. (1mk)
2. 
The table below shows how solubility of some substances in water varies with temperature.
Substances Change in solubility with temp in 100g
Temperature 00C 200C 400C 600
W 0.334 0.16 0.97 0.0058
X 27.60 34.0 40.0 45.5
Y 35.70 36.0 36.0 37.3
Which of the above substances is likely to be a gas?
3. 
Describe how the following reagents can be used to prepare lead sulphate, solid potassium sulphate, solid lead carbonate, and dilute nitric acid distilled water.           (2mks)
4.  
Study the information in the table and answer the question that follows:
Substances  Solubility g/100g water
A 1.26x102
B 1.0 x 10-2
Describe how a solid sample of substance A could be obtained from a solid mixture of A and B.       (2mks)
 
5. Study the chart below and answer the questions that follows:
 

 a) Name reagent used in
  i) Step 1        (1mk)
  ii) Name compound a      (1mk)
 b) Write an ionic equation for the reaction in the step (IV)   (1mk)
6. The table below shows the solubility of a salt at various temperatures.
Temperature 0C Solubility
0 36
40 30
80 25
110 25
What would happen if a sample of saturated solution of salt at 400C is heated to 800C? Explain.        (2mks)
7. Study the solubility curves below and answer the questions that follows:
What happens when a solution containing 40g of potassium chlorate in 100g of water at 900C is cooled to 400C? Explain    (3mks)
8. Sample solutions of salts were labeled as I, II, III and IV.  The actual solutions not in that order are lead nitrate, zinc sulphate, potassium chloride and calcium chloride.
a) When aqueous sodium carbonate was added to each sample, separately, a white precipitate was formed in I, III, IVonly. Identify solution II.  (1mk)
b) When aqueous sodium hydroxide was added to each sample, separately, a white precipitate was formed in III only. Identify solution III.
c) When excess aqueous sodium hydroxide was added to each sample, separately, white precipitate was formed in III only. Identify solution III.
          (3mks)
9. 
 State one use of sodium hydrogen carbonate.     (1mk)
10. 
a) Starting with magnesium oxide solid, describe how a solid sample of magnesium hydroxide can be prepared.    (2mks)
b) Give one use of magnesium hydroxide.    (1mk)
11. Study the flow chart below and answer the questions that follows:

a) Name reagent Z       (1mk)
 b) Describe the process which takes place in step 2.   (1mk)
 c) Identify the white solid       (1mk)
12. 
a) Name the process that take place when crystals of zinc nitrate change into solution when exposed to air.      (1mk)
13. 
Starting with sodium metal, describe how a sample of crystals of sodium hydrogen carbonate may be prepared.     (3mks)
14. 
Starting with copper metal, describe how a sample of crystals of copper (II) chloride may be prepared in the laboratory.     (3mks)
15. 
The flow chart below shows analysis of mixture “R” that contains two salts. 
a) Study the analysis and answer the questions that follows:-
i) What conditions are necessary for the process in step I to take place?         (1mk)
ii) Draw a labelled diagram to the set up that could be used to separate the mixture formed in step II.   (2mks)
iii) Write an ionic equation for the reaction between the cation in fitrate X and aqueous ammonia     (1mk)
iv) What observations would indicate the presence of NO2 (g) in step I.          (1mk)
v) State how the water vapour in step I could be identified. (1mk)
b)

i) What conclusion can be drawn from step (IV) only?  Explain           (2mks)
ii) Write the formula of an anion present in residue U. Explain.(2mks)
iii) Suggest the identity of the cations present in solution Z. (1mk)
 c) Name the two salts present in mixture R.    (2mks)
16. 
a) Give the name of each of the following processes described below which takes place when the salts are exposed to air for some time.
 i) Anhydrous copper sulphate becomes blue and wet.  (1mk)
 ii) Magnesium chloride forms an aqueous solution.  (1mk)
iii) Freshy crystals of sodium carbonate, Na2 CO310H2O, become covered with a white powder of formula, Na2CO3: H2O (1mk)
b) Write the formula of the complex ion formed in each of the reactions described below:-
c) A hydrated salt has the following compostion by mass iron 20.2%, Oxygen 23.0%, sulphur 11.5%, Water 45.3%. Its relative formula mass is 278. Determine the formula of the hydrated salts.    (3mks)
i) 6.95g of the hydrated salts were dissolved in distilled water and the total volume made to 250 cm3 of the solution.  Calculate the concentration of the salt solution in moles per litre.  (2mks)
17 
During the electrolysis of aqueous copper (II) sulphate using copper electrodes, a current of 0.2 amperes was passed through the cell for 5 hours.
i) Write an ionic equation that took place at the anode   (2mks)
ii) Determine the change in the mass of the anode which occurred as a result of the electrolysis process.  (Cu = 63.5, IF =96500 coulombs). (3mks)
18. 
The table below gives the solubilities of hydrated copper (II) sulphate in mol/ dm3 at different temperature.
Temperature (0C) Solubilities mol/dm3
20 8x10-2
40 12 x 10-2
60 16x10-2
80 22x10-2
100 30x10-2
i) On the the graph paper (provided) plot a graph of solubility of copper (II) sulphate (Vertical Axis) against temperatures   (3mks)
ii) From the graph, determine the mass of copper (II) sulphate deposited when the solution is cooled from 700C to 400C. (Molar mass of hydrated copper (II) Sulphate is 250g.)      (2mks)
b) In an experiment to determine the solubility of sodium chloride, 5.0 cm3 of a saturated solution of the sodium chloride, 5.0 cm3 of a saturated solution of the sodium chloride solution weighing 5.35g were placed in a volumetric flask and diluted to a total volume of 250cm.  25 cm3 of the dilute solution of sodium chloride completely reacted with 24.1 cm3 of 0.1m silver nitrate solution
  AgNO3 (aq) + NaCl(s) → AgCl (s) + NaNO3 (aq)
Calculate
 i) Moles of silver nitrate in 24.1cm3 of the solution.  (1mk)
 ii) Moles of sodium chloride in 25.0cm3 of solution.  (1mk)
 iii) Moles of sodium chloride in 250 cm3 of saturated sodium chloride.  
(1mk)
 iv) Mass of water in 5.0 cm3 of saturated sodium chloride.  (1mk)
 v) Mass of water in 5.0cm3 of saturated solution of sodium chloride.          (1mk)
 vi) Solubility of sodium chloride in g/100g water  (2mks)
19. 
a) At 250C, 50g of potassium nitrate were added to 100g of water to make a saturated solution. What is meant by saturated solution?  (2mks)
b) The table below gives the solubilities of potassium nitrate of different temperatures.
Temperature (0C) 12 20 28 36 44 52
Solubility in /100 water 22 31 42 55 70 90
i) Plot a graph of the solubility of potassium nitrate (Vertical axis) against temperature.
ii) Use the graph
a) Determine the solubility of potassium nitrate at 150C.         (1mk)
b) Determine the mass of nitrate that remained undissolved given that 80g of potassium nitrate were added to 100 cm3 of water and warmed to 40C.    (2mks)
20. 
a) The table below shows the solubility of ammonium phosphate in water at different temperatures.
Temperature (oC) Solubility of ammonium phosphate in g/100g water
10 63.0
20 69.0
30 75.0
40 82.0
50 89.0
60 97.0
i) On the grid provided, draw the solubility curve of ammonium phosphate.  (Temperature on x-axis).    (3mks)
ii) Using the graph, determine the solubility of ammonium phosphate at 250C.       (1mk)
iii) 100g of a saturated solution of ammonium phosphate was prepared at 250C  
 I) What is meant by a saturated solution?  (1mk) 
II) Calculate the mass of ammonium phosphate which was used to prepare the saturated solution.  (2mks)
 
21. Study the chart below and answer the questions that follows:
 a) Identify:-
  i) The cation in the solution K     (1mk)
  ii) The white precipitate “L”     (1mk)
 b) What property of white precipitate L is illustrated in steps I and II.  (1mk)
22. Study the flow chart below and answer the questions that follows:

a) Name  (i)   Compound T      (1mk)
   (ii)   Gas “U”      (1mk)
 b) Give a chemical test that you would use to identify gas U.
23. Potassium sulphite gave white precipitate with Barium Nitrate solution.  An addition of dilute Hydrochloric Acid, the white precipitate disappeare. 
 a) Write the formula of the compound which formed the white precipitate. 
           (1mk)
b) Write the equation for the reaction between dilute hydrochloric acid and the compound whose formula is written in(a) above.   (1mk)
25. 0.63g of lead powder dissolved in excess nitric acid to form Lead Nitrate solution.  All the Lead Nitrate solution was reacted with sodium sulphate solution was reacted with sodium sulphate solution.     (1mk)
a) Write an ionic equation for the reaction between lead nitrate and sodium sulphate solution.       (1mk)
b) Determine the mass of Lead salt formed in (a) above.
 (Pb= 207) (S= 32) (O=16)      (2mks)
26. Study the scheme below and answer questions that follow
 a) Write the formula of the cations present in F.   (1mk)
 b) What property of chlorine is shown in step I?   (1mk)
 c) Write an equation for the reaction which occurs in step (III). (1mk)
27. The scheme below shows some reactions sequence starting with solid N.
 a) Name solid N.        (1mk)
 b) Write the formula of complex ions present in solution Q.  (1mk)
28. When pellets of sodium hydroxide are exposed to air, a solution is formed which gradually disappears leaving a white powder.  Explain.   (2mks)
29. What observation is made when hydrated copper (II) sulphate is heated gently?
30. Study the scheme below for the laboratory preparation of precipitate “V” and answer the questions that follow

a) What observations is made when aqueous ammonia is added to 2 cm3 of solution “A”.  Explain.      (2mks)
b) State and explain the observation made when aqueous ammonia is added to spatula- end full of white precipitate “V”.    (2mks)
 
 
TOPIC 5
CARBON AND SOME OF ITS COMPOUNDS
1. Give two properties of carbon (IV) oxide which make it suitable for use in extinguishers.         (2mks)
2. 
Give a reason why calcium hydroxide solution us used to defect the presence of Carbon (IV) oxide gas while sodium hydroxide sodium is NOT   (1mk)
3. 
A sample of air contaminated with carbon (II) oxide and sulphur (IV) oxide was passed through the apparatus shown below.
Which contaminant was removed by passing the contaminated air through the apparatus. Explain         (2mks)
4. The decomposition of calcium carbonate can be represented by the equation.
 CaCO3(s)                                 CaO (s)   +CO2 (g)
 Explain how an increase in pressure would affect the equilibrium position (2mks)
5. 
Explain how you would obtain solid sodium carbonate from a mixture of lead carbonate powder.        (2mks)
6. 
When extinguishing a fire caused by burning kerosene, carbon (IV) oxide is used in preference to water.  Explain       (2mks)
7. 
When dilute nitric acid was added to a sample of solid “C” a colourless gas that formed a white precipitate with lime water was produced.  When another sample of solid “C” was heated strongly in a test tube, there was no observations changes.  Write the formula of the ions in solid “C”     (2mks)
 
8.  
The diagram below represents a charcoal burner.  Study it and answer the questions that follows:

 Write equations for the reactions taking place at I and II   (2mks)
9. When excess carbon (IV) oxide passed over heated lead (II) oxide in combustion tube, lead (II) oxide was reduced.
 a) Write an equation for the reaction which took place.   (1mk)
b) What observations was made in the combustion tube when the reaction was complete?        (1mk)
 c) Name another gas which would be used to reduce lead (II) oxide. (1mk)

10. 
The simplified flow chart shows some of the steps in the manufacturing of the sodium carbonate by the solvey process.
a) Identify substance L
 b) Name the process taking place in step II
 c) Write an equation for the reaction which take place in step III. (1mk)
11. Use the scheme below to answer the questions

 a) Identify the solids H and J      (2mks)
 b) State one commercial use of solid J.     (1mk)
12. State any two difference between luminous flame and non luminous flame.(2mks)
13. 
The apparatus shown below was used to investigate the effect of carbon (II) oxide on copper (II) oxide.
a) State the observation that was made in the combustion tube at the end of the experiment.           
b) Write an equation for the reaction that took place in the combustion tube
c) Why is it necessary to burn the gas coming out of tube K?
14. 
When carbon (IV) oxide gas was passed through aqueous calcium hydroxide a white suspension/ precipitate was formed.
 a) Write an equation for the reaction that took place in the combustion tube.
           (1mk)
b) State and explain the change that would occur when excess carbon (IV) oxide gas is bubbled through the white suspension.   (1mk)
15. 
A certain GCO3 reacts with dilute Hydrochloric acid according to the equation given below.
 G CO3(s)    + 2HCl (ag)     →       GCl2 (aq)      +    CO2 (g)   + H2O (1)
If 1 g of the carbonate reacts completely with 20cm3 of 1m Hydrochloric acid.  Calculate the relative atomic mass of G. (C= 12.0, O = 16.0)  (3mks)
16. In the industrial extraction of lead metal, the ore is first roasted in a furnace.  The solid mixture obtained is then fed into another furnace together with coke, limestone and scrap iron.  State the functions of the following in the process.
a) Coke         (1mk)
b) Lime stone           (1mk)
c) Scrap iron        (1mk)
17. .
 When calcium carbonate is heated, the equilibrium shown below is established.
 CaCO3(s)   CaO(s)   +     CO2 (g)
How would the position of the equilibrium be affected if a small amount of dilute potassium hydroxide is added to the equilibrium mixture.  Explain   (2mks)
18. 
The set up below was used to obtain a sample of iron

Write two equations for the reactions which occur in the combustion tube. (2mks)
19. 
Dry carbon (II) oxide gas reacts with heated lead (II) oxide as shown in the equation below.
PbO(s)   + CO (g)    →  Pb s)   + CO2 (g)
a) Name the process undergone by the lead (II) oxide.   (1mk)
b) Give a reason for your answer in (a) above.    (1mk)
c) Name another gas that can be used to perform the same function as carbon (II) oxide gas in the above reaction.     (1mk)
20. 
In an experiment to study the properties of concentrated sulphuric acid, a mixture of the acid and the wood charcoal was heated in a boiling tube.
a) Write the equation of the reaction that took place in the boiling tube.  
(1mk)
b) Using oxidation numbers, show that reduction and oxidation reactions took place in the boiling tube.      (2mks)
21. Name the process:-
 Solid carbon (IV) oxide (dry ice) changes directly into gas.   (1mk)
22. The diagram below represent part of the set up used to prepared and collect gas T.

a) Name two reagents that reacted to produce both carbon (IV) oxide and carbon (II) oxide.       (1mk)
b) Write the equation for the reaction which takes place in the wash bottle.            (1mk)
c) Give a reason why carbon (II) oxide is not easily detected.  (1mk)
23. The diagram below shows a jiko when in use.  Study is and answer the questions that follow.
a) Identify the gas formed at region A.
b) State and explain the observation made at region B.   (2mks)
24. The set- up below was used to collect a dry sample of a gas
Give two reasons why the set-up cannot be used to collect carbon (IV) oxide gas. 
           (2mks)
25. a) Explain why permanent hardness in water cannot be removed by boiling.
           (2mks)
b) Name two methods that can be used to remove permanent hardness from water.         (1mk) 
26. 
When solid B1 was heated, a gas which formed a white precipitate when passed through lime water was produce.  The residue was dissolved in dilute nitric (V) acid to form a precipitate which dissolved on warming was formed.
a) Write the formula of the:
 I. Cation in solid B1      (1mk)
 II. Anion in solid B1      (1mk)
b) Write an ionic equation for the reaction between the residue and dilute nitric (V) acid
27. 
The flow chart below llustrates the industrial extraction of lead metal.  Study it and answer the questions that follow
a) i) Name the ore that is commonly used in this process.  (1mk)
  ii) Explain what takes place in the roasting furnace.  (1mk)
  iii) Identify gas P       (1mk)
iv) Write the equation for the main reaction that takes place in the smelting furnace.      (1mk)
v) Give two environmental hazards likely to be associated with extraction of lead.      (2mks)
b) Explain why hard water flowing in lead pipes may be safer for drinking than soft water flowing in the same pipes.    (3mks)
c) State one use of lead other than the making of lead pipes.
28. In an experiment, carbon (IV) Oxide gas was passed over heated charcoal and the gas produced collected as shown in diagram below.
 i) Write an equation for the reaction that took place in the combustion tube.
           (1mk)
 ii) Name another substance that can be used instead of sodium hydroxide. 
           (1mk)
iii) Describe a simple chemical test that can be used to distinguish between carbon (II) oxide and carbon (IV) oxide.    (2mks)
iv) Give one use of carbon (II) oxide     (1mk)
29. 
The scheme below shows some reactions starting with calcium oxide.  Study it and answer the questions that follows.
30. 
Carbon exists in different crystalline forms.  Some of these forms were recently discovered in soot and are called fullerenes.
i) What name is given to different crystalline forms of the same element? 
          (1mk)
ii) Fullernes dissolve in methylbenzene while the other forms of carbon; describe how crystals of fullerenes can be obtained from soot. (3mks)
iii) The relative molecular mass of one of the fullerenes is 720.  What is the molecular of this fullerene? (C=12.0)     (1mk)
31. When extinguishing fire caused by petrol, carbon (IV) oxide is used in preference to water.  Explain.       (2mks)
32. Write an equation for the reaction that occurs when carbon (II) oxide is passed over heated Copper (II) oxide.
33. Use the flow chart below to answer the questions that follows.
 i) Name the process that take place in S and R     (2mks)
 ii) state one use of calcium chloride Cacl2    (1mk)
 iii) Write the equation for the reactions that take place in
Q.         (1mk)
  Slaker I        (1mk)
 b) Explain how sodium carbonate can be used to soften hard water. (1mk)
 c) Give one commercial use of sodium carbonate    (1mk)
d) X frams of sodium carbonate (Na2CO2(s) react completely with 30cm3 of dilute hydrochloric acid to produce 672cm3  of carbon (IV)oxide at STP (Ma=23)
 (i) Write the equation for the reaction.    (1mk)
 (ii) Calculate the concentration of the acid in moles per litre. (2mks)
 (iii) Calculate the value of “X”     (2mks)
34. a) Explain the following
  i) Temporary hardness in water     (1mk)
  ii) Permanent hardness in water.     (1mk)
 b) i) Draw a diagram and explain how ionic exchanger works. (3mks)
ii) Explain why hard water is recommended for healthy development of teeth.       (2mks) 
35. When a solid “T” is heated, a black solid is left and a colourless gas which form white precipitate with lime water is evaluated.  Identify “T” and write an equation for the decomposition of “T”.       (2mks)
36. State the confirmation test for the following gases:-
 i) Carbon (II) oxide       (1mk)
 ii) Carbon (IV) oxide       (1mk)
37. Explain why dilute sulphuric acid does not react fully with calcium carbonate while dilute Hydrochloric acid react fully with the same liberating carbon (IV) oxide.           (2mks)
38. Name the process in each case by which carbon (IV)is constantly being
 i) Added to the atmosphere      (1mk)
 ii) Removed from the atmosphere     (1mk)
39. A compound contains 40% carbon, 6.67% hydrogen and the rest is oxygen.  Find the simplest formula for this compound. (C=12) (H=1) (O=16)  (2mks)
40. Below is a set up used by a student to prepare gas n
 i) Identify gas “N”       (1mk)
 ii) Explain why it was possible to isolate gas N.    (1mk)
 iii) Comment of the PH of the water after the experiment  (1mk)
 
FORM 3 WORK
TOPIC 1
GAS LAW
1. 
Explain why the volume of a gas increases when its temperature is increased at a constant pressure.         (1 mk)
Cotton wool pads were socked with concentrated solutions of gas “p” and gas “Q” the
pads were then placed of the opposite ends of a long horizontal glass tube at the same
time. The tube was then immediately corked at both ends as shown the diagram
below.

After sometimes the gases were observed to meet at point “R” which of the two gases is dense? Explain your answer      (2 mks)
2. 
A mixture containing equal volumes of hydrogen and carbon (IV) oxide was introduced as shown below
Which gas would be detected at point “C” first? Explain  (2 mks)
3. 
In an experiment to study diffusion of gases a student set up the apparatus shown in the diagram I. After sometime the student noticed a change in the water level as shown in diagram II.

Give an explanation for the change in water level   (2 mks)
4. A fixed mass of gas has a volume of 250 cm3 at a temperature of 2700 and 750 mm Hg pressure. Calculate the volume the gas would occupy at 420c and 750 mm pressure.
5. A gas occupies a volume of 400 cm3 at 500k and atmospheric pressure. What will be the temperature of the gas when the volume and pressure of the gas is 100 cm3 and 0.5 atmospheric pressure respectively?    (2 mks)
6. A sealed glass tube containing air at S.T.P was immersed in water at 1000C. Assuming there was no increase in volume of the glass tube due to expansion of the glass. Calculate the pressure of the air inside the tube.
Standard pressure = 760mmHg: Standard temperature = 273 K. (2 mks)
7. 
A given volume of Ozone (03) diffused from a certain apparatus in 96 seconds. Calculate the time taken by equal volume of carbon (IV) oxide (Co2) to diffuse under the same condition  (O= 16) (C=12)  (2 mks) 
8. 
A few crystals of potassium manganate VII were carefully placed in a beaker at one spot. The beaker was left undisturbed for two hours. State and explain the observation that was made.      (2 mks)
9. 
The graph below shows the behaviour of a fix mass of a gas at constant temperature.

(a) What is the relationship  between the volume and the pressure of the gas
(1 mk)
(b) 3 litres of oxygen gas at one atmospheric pressure were compressed to two atmospheres at constant temperature. Calculate the volume occupied by oxygen gas.       (2 mks)
10. 
When a hydrocarbon was burnt completely in oxygen, 4.2 g of carbon (IV) oxide and 1.71 g of water were formed. Determine the empirical formula of the hydrogen (H= 1.0) (C=12)     (3 mks)
11. 
60cm3 of oxygen gas diffused through a porous partition in 50 seconds. How long would it take 60cm3 of sulphur (IV) oxide gas to diffuse through the same conditions?  (S= 32.0) (O=16.0).     (3 mks)

12. 
(a) State Charles law
(b) The volume of a sample of hydrogen gas at temperature 291K and 1.0 x 105 pascals was 3.5 x 10-2m3. Calculate the temperature at which the volume of the gas would be 2.8 x 10-2 m3 at 1.0 x 105 pascals. (2 mks)
13. A small crystal of potassium (VII) was placed in a beaker containing water. The beaker was left standing for two days without shaking. State and explain the observations that were made.
14.  (a)  State the Graham’s law of diffusion     (1 mark)
(b)  The molar masses of gases w and x are 16.0 and 44.0 respectively. If the rate of diffusion of w through a porous material is 12cm3S-1, calculate the rate of diffusion of x through the same material.  (2 marks)
15. Calculate the R.F.M of gas “A” given that the time taken for equal volumes of oxygen and gas “A” to diffuse through a hole is 20 seconds and 24 seconds respectively (O= 16.0)       ( 2 mks)
16. A certain volume of Co2 gas takes 200 seconds to diffuse through porous plug. How long would it take the same volume of HCL to diffuse under the same condition?
(3 mks)
17. What volume of a butane (C4H10) must be burnt in oxygen to give 11g of Co2 at r.t.p?
 The equation for the combustion of butane is given below
2 C4H10 (g) + 13 O2 (g) →8C02 (g) + 10H20 (l)
18. The set up shown below was used to investigate some properties of two gases “ P” and Q

When beaker A was filled with gas P the level of water in the glass tubing rose to level ll. When the experiment was repeated using gas Q, the level of water dropped to level III. Explain these observations.
19. Study the set up below and answer the questions that follows
(i)  What observations would be made in the tube  (1 mk)
(ii)  Indicate with a cross (x) on the diagram the likely position where the observation stated in (i) above would be made.  (1 mk)
(iii)  Write an equation for the reaction that takes places in the set up above
(1 mk)
20. 88 cm3 of gas K diffuse through a small hole in 40 seconds while 50cm3 of hydrogen gas diffuse through the same hole under the same conditions in 5 seconds. Calculate the RMM of the gas K       (3 mks)
21. 200 cm3 ammonia gas are burnt in 300 cm3 of oxygen gas (excess). 200 cm3 of nitrogen (II) oxides and 300 cm3 steam were formed. 50 cm3 of oxygen was left unused. Deduce the equation for this reaction.    (3 mks)
22. Sketch a demonstration graph showing variation of pressure of a gas against volume at a constant temperature.       (2 mks)
23. Nitrogen gas occupies a volume of 200 cm3 at 250C. What will be the temperature of nitrogen if it occupied a volume of 300 cm3?    (2 mks)
24. What will be the volume of a certain mass of nitrogen gas at 200 C if it occupies 200 cm3 at 250C pressure remain constant.     (2 mks)
25. 200 cm3 of gas “p” at s.t.p was cooled and the volume contracted to 160 cm3. Calculate the new temperature of the gas in 0C if the pressure is kept constant.
26. Form three students found that a mass of nitrogen gas occupies 330 cm3 at 2800C and 760 mm Hg pressure. At what temperature will the volume of the gas be 190cm3 and the pressure 800 mm Hg?

 
TOPIC 2
THE MOLE
1. When 34. 8g of hydrated sodium carbonate Na2 Co3   XH2O were heated to a constant mass. 15.9g of anhydrous sodium carbonate were obtained.  Find the value of “X” in hydrated carbonate (Na= 23), (O = 16), (C= 12), (H = 1.0)
(3 mks)
2. Hydrogen reacts with oxygen as shown in the equation
2H2 (g) + 02 (g) → 2H2O (g)
In an experiment 100cm3 of hydrogen gas was mixed with 100cm3 oxygen gas and the mixture heated to form H2O. Which of the gas was in excess and how much.         (2 mks)
3. Calculate the amount of calcium carbonate that would remain if 15.0g of calcium carbonate were reacted with 0.2 moles of hydrochloric acid. The equation for the reaction is.
CaCo3 + Hcl (aq) → CaCl2(s) + H2O (l) + C02 (g)
(C = 12), (O = 16) (Ca = 40)      (2 mks)
4. 1995: PP1 A Questions 14
A compound has an empirical formula C3H6O and a relative formula ¬mass of 116
(a)  Determine its molecular formula
 (H= 1.0) (C = 12.0), (0= 16.0)    (2 mks)
(b)  Calculate the percentage composition of carbon by mass in the compound
         (1 mk)
5. In an experiment 2.4 g of sulphur was obtained by reacting hydrogen sulphide and chlorine as shown by the equation below
H2S (g) + Cl(2) → S(s) + 2 Hcl(g)
(a)  Which of the reagent acts as a reducing agent? Explain  (1 mk)
(b)  Given that the yield of sulphur in the above reaction is 75%. Calculate the number of moles of H2S (g) used in the reaction (S= 32.0) (2 mks)
6. 
(a) The empirical formula of the hydrocarbon is C2H3. The hydrocarbon has a relative molecular mass of 54  (H= 1)
(i)  Determine the molecular formula of the hydrocarbon
(ii)  Draw the structural formula of the hydrocarbon (1 mk)
(iii)  To which homologous series does the hydrocarbon drawn in (ii) above belong      (1 mk)
(b) 90cm3 of 0.01M calcium hydroxide were added to a sample of water containing 0.001 moles of calcium hydrogen carbonate.
(i)  Write an equation for the reaction which took place  (1 mk)
(ii)  Calculate the number of moles of calcium ions in 90cm3 of 0.01M calcium hydroxide
(c) What would be observed if soap solution was added dropwise to a sample of the water after the addition of calcium hydroxide? Give a reason (1 mk)
7. Calculate the mass of nitrogen (IV) oxide gas that would occupy the same volume as 10g of hydrogen gas at the same temperature and pressure (H= 1.0), (N = 14.0) (O = 16)        (2 mks)
8. On complete combustion of a sample of hydrocarbon, 3.52 g of carbon (IV) oxide and 1.44g of water were formed. Determine the molecular formula of the hydrocarbon. (Relative molecular mass of hydrocarbon is 56) (Carbon (10) oxide = 44) water = 18) (H=1.0) (C= 12.0)     (4 mks)
9. 20.0cm3 of solution containing 4g per litre of sodium hydroxide was neutralized by 8.0cm3 of dilute sulphuric acid. Calculate the concentration of sulphuric acid in moles per litre (Na = 23.0) (O = 16.0) (H= 1.0)    (3 mks)
10. A weighed sample of crystalline sodium carbonate Na2, Co3: N: H20 was heated in a crucible until there was no further change in mass. The mass of the sample reduced by 14.5. Find the number of moles (N) of the water of crystallization (Na = 23.0), (O = 16.0), (C = 12.0), H = 1.0)   (3 mks)
11. In an experiment 30 cm3 of 0.1 sulphuric acid were reacted with 30 cm3 of 0.1M sodium hydroxide.
(a)  Write an equation for the reaction that took place (1 mk)
(b)  State the observations that were made when both blue and red litmus were dropped into the mixture    (1 mk)
(c)  Give a reason for your answer   (1 mk)
12. When excess dilute hydrochloric acid was added to sodium sulphite, 960cm3 of sulphur (IV) oxide gas was produced. Calculate the mass of sodium sulphite that was used. (Molar mass of sodium sulphite = 126g: and molar gas volume = 24000 cm3)         (3 mks)
13. When “X” cm3 of a solution of 0.5m magnesium nitrate were reacted with excess ammonium carbonate solution, the mass of magnesium carbonate formed was 8.4g.
(a) Write the ionic equation for the reaction that took place (1 mk)
(b) Calculate the value of = “X” (C= 12) (Mg= 24) (O = 16) (2 mks)
14. A certain carbonate of GCo3 react with dilute hydrochloric acid according to the equation given below.
GCO3(s) + 2HCl (aq) → Co2 (g) + H2O(l) + GCl2 (aq)
If 1 g of the carbonate reacts completely with  20 cm3 of 1 m  hydrochloric acid, calculate the atomic mass  of  G     (3 mks)
15. When 94.5g of hydrated – barium hydroxide Ba(OH) 2: nH2O were heated to a constant mass. 51.3g of anhydrous- barium hydroxide were obtained. Determine the empirical formula of the hydrated barium hydroxide. (Ba = 137.0) (O= 16), (H= 1.0)
16. 
15.0 cm3 ethanoic acid (CH3COOH) was dissolved in water to make 500 cm3 of solution. Calculate the concentration of the solution in moles per litre. (C= 12.0; H= 1.0; O = 16.0’ density of ethanoic is 1.05g/cm3   (3 mks)
17. 
An alkanol has the following composition by mass: Hydrogen 13.5%, oxygen 21.6% and carbon 64. 9%
(a) Determine the empirical formula of the alkanol (C= 12.0; H = 1.0; 0 = 16.0)       (2 mks)
(b) Given that empirical formula  and the molecular formula of the alkanol are the same, draw the structure of the alkanol   (1 mk)
18. 
6.84 of aluminium sulphate were dissolved in 150cm 3 of water. Calculate the molar concentration of the sulphate ions in the solution. (Relative formula mass of aluminum sulphate is 342)      (3 mks)
19. When a hydrated sample of calcium sulphate CaSO4.   XH2O was lost, the       following  data was recorded:
Mass of crucible   =  30.296g
Mass of crucible + hydrated salt  =  33.111g
Mass of crucible + anhydrous salt =  32.781g
Determine the empirical formula of the hydrated salt (relative formula mass of CaSO4= 136, H2O = 18)     (3 mks)
20. 
Phosphoric acid is manufactured from calcium phosphate according to the following equation.
Ca3 (PO4)2(s) + 3H2SO4 (l) →2H3PO4 (aq) + 3CaSO4(s)
Calculate the mass in (kg) of phosphoric acid that would be obtained if 155 kg of calcium phosphate reacted completely with the acid (Ca = 40, P= 31, S = 32, O = 16, H = 1)       (2 marks)
21. 
In an experiment to determine the percentage of magnesium hydroxide in an anti- acid, a solution containing 0.50g of the anti- acid was neutralized by 23.0cm3 of 0.10M hydrochloric acid. (Relative formula mass of magnesium hydroxide =58). Calculate the:
(a) Mass of magnesium hydroxide in the anti- acid  (2 mks)
(b) Percentage of magnesium hydroxide in the anti- acid (1 mark)
22. 
(a)  Name one raw material from which sodium hydroxide is manufactured           (1 mk)
(b)  Sodium hydroxide pellets were accidentally mixed with sodium chloride. 17.6g of the mixture were dissolved in water to make one litre of solution. 100 cm3 of the solution was neutralized by 40 cm3 of 0.5m sulphuric acid.
(i)  Write an equation for the reaction that took place    (1 mk)
(ii)  Calculate the
I. Number of moles of the substance that reacted with sulphuric acid.
II. Number of moles  of the substance that would react with acid in the one litre of solution     ( 1 mk)
(iii)  Mass of the un-reacted substances in the one litre of solution
23.  (i) A hydrated salt has the following composition By mass: iron 20.2% Oxygen 23.0% sulphur 11.5% and 45.3% water of = crystallization. If RMM= 278
(ii)  Determine the formula of the hydrate salt.   ( 2 mks)
(iii)  6.95g of the hydrated salt were dissolved in water and the total volume made up to 250c3 of solution. Calculate the concentration of the salt solution in moles per litres     ( 2 mks)
24. 1.9 g of magnesium chloride was dissolved in water. Silver nitrate solution was added till in excess. Calculate the mass of silver nitrate that was added for the complete reaction. (Rmm of magnesium chloride= 95 (m= 14 (O= 16) (Ag = 108). 
25. During welding of fractured railway lines by thermite reaction 12g of oxide of iron is reduced by aluminum to 8.4g of iron. Determine the empirical formula  of the oxide (Fe= 56) (O= 16)     ( 3 mks)
26. In a titration reaction a student was provided with 0.1M sulphuric acid solution labeled. (M.A) and a carbonate solution containing 13.8g/cm3 labeled (X.A). The student was required to calculate the formula mass of X2Co3 and atomic mass of x in the carbonate. She pipette 25 cm3 of XA and titrated against (MA) using methyl orange indicator, her results of the titration are shown in the table below.
Experiment 1 2 3
Final burette readings in cm3 25.0 25.0 25.0
Initial burette readings in cm3 0.00 0.0 0.0
Volume of solution MA cm3   
(i) Complete the table     ( 2 mks)
(ii) What is the average volume  of MA  used  ( 1 mk)
(iii) Calculate
(a) Moles of acid used
(b) Moles of  carbonate used
(iv) Calculate the molarity of the carbonate  ( 2 mks)
(v) Calculate  the formula  mass  of X2Co3  ( 2 mks)
(vi) Calculate the Ram of x    ( 2 mks)
27. Excess Co gas was passed over heated sample of oxide of iron as shown in the diagram. Study the information and answer the questions that follows:

Mass of empty dish = 10.98g
Mass of empty dish + oxide of iron = 13.30g
Mass of empty dish + residue = 12.66g
(i)  Determine the formula of the oxide of iron. (RMM of oxide of iron = 232) Fe = 56) (O=16)      ( 3 mks)
(ii)  Write an equation for the reaction taking place  ( 1 mk)
28. 12.5 cm3 of solution  containing 13.8g/cm3 of carbonate M2 Co3 required 12.3 cm3 of H2 SO4 containing 9.8g/ dm3 for complete neutralization.
(a)  Write the equation for the above reaction   ( 1 mk)
(b)  Calculate the molarity of the acid    ( 2 mks)
(c)  Calculate the molarity of the carbonate   ( 2 mks)
(d)  Calculate the molar mass of the carbonate   ( 2 mks)
(e)  Find the relative atomic mass of M    ( 2 mks)
29. Calculate the mass  of lead (ii) nitrate that must be heated to give 22.3g of lead (ii) oxide (Pb = 207) (M = 14) (O = 16)     ( 3 mks)
30. Solution “A” is NaOH containing 48g/dm3. Solution “B” is (CooH)z: nH2O containing 63g/dm3. 20cm3 of solution “A” was pipetted into a conical flask and titrated with solution “B”. The titration was done three times. The results are shown in table below. The equation for the reaction is:
(CooH2:nH2O(aq) + 2NaoH(aq) →CooNa2(aq) + (n+ H2(l)
Experiment 1 1 3
Final readings 24.1 24.1 49.0
Initial readings 0.0 0.0 25.0
Volume used 24.1 24.1 24.0
Find
(i)  The average volume of solution “B” used   (1 mk)
(ii)  The moles of solution “A” in 20 cm3 of solution  (1 mk)
(iii)  The number of moles of “B” in dm3 of the solution  (1 mk)
(iv)  The formula mass of (CooH)2 nH2O    (2mks)
(v) Value of n       (1 mk)
(C= 12) (O= 16 (H= 1)     
31. Calculate the volume of carbon (iv) oxide measured at S.T.P that is evolved when 1 mole of copper (II) carbonate is heated to a constant mass.
32. How many molecules  are there  in 360 cm3 of nitrogen as  r.t.p
33. Define the following terms
(a) Monatomic gas
(b) Diatomic gas
(c) Atomicity of an element

 
TOPIC 3
ORGANIC CHEMISTRY 1
1. 
Propane and chlorine react as sown below
CH3CH2CH3 + Cl2 →CH3 CH2 CH2Cl + Hcl
(a) Name the type of reaction that takes place   (1 mk)
(b) State the conditions under which this reaction takes place  (1 mk)
2. 
(a)  Name one substance used for vulcanization of rubber (1 mk)
(b)  Why is it necessary to vulcanize natural rubber before use (1 mk)
3. 
R COO – Na+ and R – C6 H5 So3 – Na+ represents two cleaning agents where “R” is a long hydrocarbon chain.
(a)  Write the formula of the salts that would be formed when each of those cleaning agents is added to water containing calcium ions (2 mks)
(b)  Explain how the solubility of the two calcium salts (a) above effect the cleaning properties of each of the cleaning agents.  (2 mks)
4. 
The general formula for a homologous series of organic compound is CnH2n+1 OH
(a) Give the name and structural formula of the forth member of this series
(1 mk)
(b) Write an equation for the complete combustion of the fourth member of this series       (1 mk)
5. 
(a)  Name one natural fibre     ( 1 mk)
(b)  Give one advantage of synthetic fibres over natural fibre ( 1 mk)
    6.
Study the table below and answer the questions that follow
Alkanes Formula Heat of combustion (DHC) kj mol¬-1
Methane CH4 -890
Ethane C2H6 -1560
Propane C3H3 -2220
Butane CuH10 -
(a)  Predict the heat of combustion of butane and write it in the space provided in the table above     ( 1 mk)
(b)  What does the sign ∆ Hc vatue- indicated about combustion of alkanes
         ( 1mk)
A compound C4H10O is oxidized by excess acidified potassium permanganate to form another compound C4H8O2. The same compound C4H10O react with potassium to produce hydrogen gas
(a) Draw the structural formula and name compound C4H10 O ( 2 mks)
(b) Write equation for the reaction between potassium and compound C4H10O
          ( 1 mk)
6. 
Explain how sample of CH3CH2OH could be distinguished from CH3COOH by means of chemical reaction.      ( 2 mks)
7. 
Methane react with oxygen as shown by equation I and II below
(I) CH4(g) + 202 → Co2 + 2 H2O(l)
(II) 2CH4 (g) + 302 (g) → 2CO(g) + 4 H2O(l)
Which one of the two reactions represents the complete combustion of methane? Explain        ( 2 mks
8. A polymer has the following structure
― CH2 ― CH ― CH2 ―CH ― CH2 ―CH2 ― CH
       CN          CN           CN           n
A sample of this polymer is found to have a molecular mass of 5194. Determine the number of monomers in the polymer. (H= 1.0), (C= 12.0), (N= 14, 0)
         (3 mks)
9. 
A mixture of pentane and pentanoic acid was shaken with 0.1m sodium hydroxide solution.  And let to separate as shown in the diagram below.

Name the main component in layer W. Give a reason for your answer
        ( 2 mks)
10. 
Name and draw the structure of the compound formed when methane react with excess chlorine in presence of U.V light   ( 1 mk)

11. 
State the observations that would make when a piece of sodium metal is placed in samples of pentane and pentanol.    ( 2 mks)
12. 
Compound “L” react with hydrogen bromide gas to give another compound whose structure is
 H H H Br H
H ― C   ― C   ― C   ― C   ― C   ― H
 H H H H H H
(a)  Give the structural formula and name of compound “L” ( 2 mks)
(b)  Write an equation for the reaction which takes place between enthyne excess chlorine gas      ( 2 mks)
(c)  Write an equation for the reaction which takes place between ethyne excess chlorine gas.      ( 1 mk)

13. 
One of the fuels associated with crude oil is natural gas. Name the main constituent of natural gas and write an equation for its complete combustion.
         ( 2 mks)
14. 
Bromine react with ethane as shown below
C2H6 + Br2 → C2 H5Br + HBr
(a)  What condition is necessary for this reaction to occur ( 1 mk)
(b)  Identify the bonds, which are broken and those which are formed 
( 2 mks)
15. 
A hydrocarbon “p” was formed to decolorize bromine water.  On complete combustion of 2 moles of “P” 6 moles of carbon (IV) oxide and 6 moles of water were formed
(a)  Write the structural formula of “p”    ( 1 mk)
(b)  Give the name of p      ( 1 mk)
(c)  Name one industrial source of “p”    ( 1 mk)
16. Pentane and ethanol are miscible. Describe how water could be used to separate a mixture of pentane and ethanol.     ( 2 mks)
17. 
In the presence of U.V light ethane gas undergoes substitution reaction with chlorine.
(a) What is meant by the term substitution reaction with chlorine?
(b) Give the structural formula and the name of the organic compound formed when equal volumes of ethane and chlorine react together.
18. 
But – 2- ene undergoes additional hydrogenation according to the equation given below
CH3CH = CH-CH3 (g) + H2 (g) → CH3CH2CH2CH3
(a) Name the product formed when  but -2-ene reacts with hydrogen gas
(b) State one industrial  use of hydrogenation   ( 1 mk)
19. 
Name the organic compound formed when  CH3CH2CH2OH is reacted with concentrated  sulphuric acid at 1700C.    (1 mk)
20. 
(a)  What is meant by isomerism?    ( 1 mk)
(b)  Draw and name two isomers of butane  ( 2 mks)
(b)  Propane can be changed into methane and ethane as shown in the equation below.
CH3CH2CH3 (g) high temperature     CH4(g)  + C2H4(g)
Name the process undergone by propane  ( 1 mk)
21. 
The relative formula mass of hydrogen is 58. Draw and name two possible structure of the hydrocarbon (C= 12.0; H- 1.0)  ( 3 mks)
22. 
The structure of a detergent is
(a)  Write the molecular formular of the detergent  ( 1 mk)
(b)  What type of detergent is represented by the formula? ( 1 mk)
(c)  When this type of detergent is used to wash linen in hard water, spots (marks) are left on the linen. Write the formula of the substance responsible for the spots     ( 1 mk)
23. 
The structure below represents a sweet smelling compound
CH3 ―CH2 ―CH2 ―C ― 0 ―CH2 ―CH2 ―CH3
Give the names of the two organic compounds that can be used to prepare this compound in the laboratory.      ( 2 mks)
24. 
Study the table below and answer the questions that follow:
Compounds Melting point 0C Boiling points 0C
C2H4O2 16.6 118
C3H6 -185.0 -47.7
C3H8O -127 97.2
C5H12 -130 36.3
C6H14 -95.3 68.7
(a)  (i)  Which of the compounds is a solid at 10.00C. Explain ( 1 mk)
(ii)  Choose two compounds which are members of the same homologous series and explain the difference in their melting points       ( 3 mks)
(iii)  The compound C3H8O is an alcohol. How does its solubility in water differ from the solubility of C5H12 in water. Explain         ( 2 mks)
(b)  Complete combustion of one mole of a hydrocarbon produces four moles of carbon (IV) oxide and four moles of water.
(i)  Write the formula of the hydrocarbon   ( 1 mk)
(ii)  Write the equation for the complete combustion  ( 1 mk)
(c)  (i)  in a reaction, an alcohol “J” was converted to hex -1-ene. Give the
structural formula of alcohol “J”    ( 1 mk)
(ii)  Name the reagent and conditions necessary for the reaction in C (ii) above       ( 1 mk)
(d)  Compound K reacts with sodium hydroxide as shown below
 CH2 – OOc –c17 H35
 CH2 – OOC-C17H35 + H35 + 3NaoH →CH2OH
         CH-OH + 3C17H35COO-Na+
 CH2- OOCC17H35
      CH2 - OH
(i)  What type of reaction is represented by the equation above ( 1 mk)
(ii)  To what class of compound does “K” belong?  ( 1 mk)
25. 
(a) Give the names of the following compounds
(i) CH3CH2CH2CH2OH      (1 mk)
(ii) CH3CH2COOH       (1 mk)
(iii) CH3-COO-CH2-CH3      ( 1 mk)
(b)  Study the information in the table below and answer the questions that follow
Number of carbon atoms per molecule Relative molecular mass of hydrocarbons
2 28
3 42
4 56
(i)  Write the general formula of the hydrocarbons in the table ( 1 mk)
(ii)  Predict relative molecular formula mass of hydrocarbon with 5 carbon atoms
(iii)  Determine the molecular formula of the hydrocarbon in (ii) above and draw its structural formula. (H=1.0), (C= 12.0)  ( 1 mk)
26. 
The following equations represent two different types of reactions
(a) (i) NC4H8(g) →(C4H8 n(g)
(ii) C2H6(g) + CL2 (g) →C2H5CL(g) + HCL(g)
State the type of reaction represented by (i) and (ii)   ( 2 mks)
(b) The fermentation of glucose produces ethanol as shown in the equation below.
C6H12O6(aq)  yeast   2 CH3CH2OH(aq) + 2CO2(g)
(i)  State how the concentration of ethanol produced could be increased      ( 1 mk)
(ii)  State and explain the observations that would be made when a piece of sodium metal is added to a sample of ethanol contained in a beaker.      ( 2 mks)
(iii)  Give two commercial uses of ethanol other than manufacturing of alcohol drinks      ( 2 mks)
(c) The molecular formula of a hydrocarbon is C6H14. The hydrocarbon can be converted into two other hydrocarbons as shown by the equation below.
C6H14 →C2H6+ x
(i)  Name and draw the possible structural formula of x  ( 1 mk)
(ii)  State and explain the observations that would be made if a few drops of bromine water were added to a sample of x. (2mks)
(iii)  Write an equation for the complete combustion of C3 H8 ( 1 mk) 
27. 
(a) Give the names of the following compounds
(i) CH3CH = CH CH2CH3      ( 1 mk)
(ii) CH3 CH2 COOH       (1 mk)
(b)  Ethane and Ethene react with bromine according to the following equations given below
(i) C2H6(g) + Br2 (g) →C2H5Br (l)  + HBr (g)
(ii) C2¬H4(g) + Br2 (g) → C2H4Br2 (l)
Name the type of bromination reaction taking place in (i) and (ii) above
(c)  Study the diagram below and answer the questions that follow
(i)  Write the equation for the complete combustion of butane ( 1 mk)
(ii)  The PH of substance K was formed to be less than 7 explain this observations.      ( 2 mks)
(d)  The polymerization of tetrafloureoethane (C2F4) is similar to that of   ethane (C2H4)
(i)  What is meant by the term polymerization?  ( 1 mk)
(ii)  Draw the structural formula of a portion of the polymer obtained from the monomers (C2F4)    ( 1 mk)
(e)  State any two advantages that synthetic polymers have over natural polymers       ( 2 mks)
28. 
(a)  In which homologous series do the following compounds belong?
(i) CH3CCH       (1 mk)
(ii) CH3CH2COOH      (1 mk)
(b) Raw rubber is heated with sulphur in manufacture of natural rubber.
 (i) What name is given to the process?  (1 mk)
 (ii) Why is the process necessary?  (1 mk
(c) Study the scheme given and answer the questions that follow

(i) Write an equation for the reaction between propan-1-ol  and  potassium metal       ( 1mk)
(ii) Name process I and II     ( 2 mks)
(iii) Identify the products  “A” and “B”   ( 2 mks)
(iv) Name ONE catalyst used  in process II  (1 mk)
(v) Draw the structural formula of the repeating unit in the polymer “C”
(d) State two uses industrial uses of methane   (2 mks)
29. 
(a) State how burning can be used to distinguish between ethane and ethyne. Explain your answer.
(b) Draw the structural formula of the third member of the homologous series of the ethyne.      ( 1 mk)
(c) The flow chart below shows a series of reaction starting with ethanol. Study it and answer the questions that follow.

(i) Name
I Process “A”
II  Substance “B” and C
(ii) Write the equation for the combustion of ethanol. ( 1 mk)
(iii) Explain why it is necessary to sue high pressure to change gas “B” into polymer      ( 1 mk)
(iv) State one use of methane    ( 1 mk)
30. 
(a)  Crude oil is a source of many compounds that contain carbon and hydrogen only       ( 1 mk)
(i)  Name the process used to separate the components of crude oil
(1 mk)
(ii)  On what two physical properties of the above components does the separation depend     ( 2 mks)
(b)  Under certain conditions hexane can be converted to two products. The formula of one of the products is C3H8
(i)  Write the formula of the other product  ( 1 mk)
(ii)  Describe a simple chemical reaction to show the differences between two products in b(i) above.   ( 2 mks)
(c)  Ethyne  (C2H2) is another compound found in crude oil. One mole of ethyne was reacted with one mole of hydrogen chloride gas and a product “P1” was formed. P1 was then reacted with excess hydrogen gas to form P2. Draw the structure of P1 and P2    ( 2 mks)
(d)  The set up below was used to prepare and collect ethane gas. Study it and answer the questions that follows:

(i)  Name substance “T”     ( 1 mk)
(ii)  Give the property of ethane that follows it to be collected as shown in the set up      ( 1 mk)
(e)  One of the reactions undergone by ethane is addition polymerization. Give the name of the polymer and one disadvantage of the polymer it forms 
( 2 mks)
31. 
(a) What name  is given to a  compound that contain  carbon hydrogen only
(b) Hexane is a compound that contain carbon and hydrogen only
(i)  What method is used to obtain hexane from crude oil?
(ii)  State one use of hexane    ( 1 mks)
      (c) Study the flow chart below and answer the questions that follows:

(i)  Identify reagent L     ( 1 mk)
(ii)  Name the catalyst used in step 5   ( 5 mks)
(iii)  Draw the structural formula of  “J”   ( 1 mk)
(iv)  What name is given to the process that takes place in step 5 
( ½  mk)
(v)  State
I.  One use of product “R”
II.  A commercial application of the process which take place in step 6
(d)  (i)  Write the equation for the reaction between aqueous
sodium hydroxide and aqueous ethanoic acid ( 1 mk)
(ii)  Explain why the reaction between 1g sodium carbonate and 2 m hydrochloric acid is faster than the reaction between 1 g of sodium carbonate and 2 M ethanoic acid (2mks)
32. 
(a)  Give the systematic names of the following compounds
(i)  CH2 = C ― CH3
  
CH3      (1 mk)
(ii) CH3CH2CH2C ≡CH     (1 mk)
33. 
(a) Biogas is a mixture of mainly carbon (IV) oxide and methane
(i)  Give a reason why biogas can be used as fuel
(ii)  Other than fractional distillation, describe a method that can be to determine the percentage of methane in biogas ( 3 mks)
(b)  A sample of biogas contains 35.2% by mass of methane. A biogas cylinder contains 5.0 kg of the gas.
(i)  Number of moles of methane in the cylinder. (Molar mass of methane = 16)      ( 2 mks)
(ii)  Total volume of carbon (IV) oxide produced by the combustion of methane in the cylinder (molar gas volume = 24.0 dm-3 at room temperature and pressure.    ( 2 mks)
(c) Carbon (IV) oxide, methane, nitrogen(I)oxide and trichlorofluoromethane are green- house gases
(i)  State one effect of an increased level of these gases to the environment      ( 1 mk)
(ii)  Give one source from which each of the following gases is released to the environment
I Nitrogen (I) oxide    ( 1 mk)
II Trichlorofluoromethane   ( 1 mk)
34. State what you understand by the following terms as used in organic chemistry
(i)  A hydrocarbon      ( 1 mk)
(ii)  A homologous series      ( 1 mk)
(iii)  Saturated hydrocarbons     ( 1 mk)
(iv)  Isomerism       ( 1 mk)  
35. Name the following compounds using the I.U.P.C rules
(i) CH3 – CH = CH – CH3
(ii) CH3 – CH2 – CH- CH2
   
    CH3
36. Below is a scheme of some reaction of ethyne
(i)  State the condition and reagents required to effect steps I and II (2mks)
(ii)  Give the formula of products A, B, C and D    (4mks)
37. Write down the structural formula of the following compounds
(i)  2, 2 – Dimethypropane      ( 1 mk)
(ii)  2 – Chloropropane       ( 1 mk)
38. Study the crude oil fractionating column in the diagram and answer the questions that follows
(a)  How would you except the temperature to vary from A to E  (2mks)
(b)  For each fraction below state at which position it will be collected compound with       (5mks)
- C15- C25 atoms
- C4- C12 atoms
- C20 – upwards
- C9- C16
- C1 – C4
39. The boiling points at 760 mg pressure of three alkanes are Butane, 273k pentane 309K and Hexane 342K. Account for the fact that the pentane has a higher boiling point than butane.       ( 2 mks)
40. Petrol is a mixture of hydrocarbon used as fuel and is obtained from crude oil by fractional distillation.
(i)  State the range of carbon atoms in the molecules of hydrocarbon in petrol
(ii)  Name two gases that pollute the atmosphere as a result of burning petrol in combustion engines      ( 2 mks)
41. What is the role of sunlight in substitution halogenations reaction ( 1 mk)
42. A,B,C are three homologous series of organic compounds
Series  General formula
A CnH2n-2
B CnH2n
C CnH2n + 2
(i)  What is the name given to series C    ( 1 mk)
(ii)  Write down the name and structural formula of the second member of series “B”
(iii)  Write down an equation and name the products of reaction between HBr with second member of series “B”    ( 2 mks)
43. The scheme below shows preparation of methane
CH3COOH  NaOH  CH3COONa   reagent  CH4 Cl2 T
        R   U.V
(i)  Name reagent “R”      ( 1 mk)
(ii)  Name substance “T”      ( 1 mk)
(iii)  Write an equation for the reaction between CH3COONa and reagent “R”
         ( 2 mk)
44. CH2 = CH2 Polymerize [-CH2 – CH2]n compound U
(i)  Name compound U     ( 1 mk)
(ii)  If the RMM of U is 42000 determine the value of n ( 1 mk)
45. The empirical formula a hydrocarbon is C2H3 it RMM is 54.
(a)  Determine the molecular of the hydrocarbon  ( 1 mk)
(b)  Draw the structural formula of this hydrocarbon ( 1 mk)
(c)  To which homologous series does the hydrocarbon draw above belong?
        ( 1 mk)
 
TOPIC 4
NITROGEN AND ITS COMPOUNDS
1. 
A student set- up apparatus to prepare and collect a sample of ammonia gas as shown in the diagram below. Study the set up and answer the question that follows
Identify the two mistakes in the set- up represented by the diagram ( 2 mks)
2. 
State two observations that would be made when solid lead (II) Nitrate is heated strongly.        ( 2 mks)

3. 
Dilute nitric acid reacts with  copper  according to the equation
3Cu(s) + 8H+ (aq) + 2NO-(3)(aq) →3cu21 (aq) + 2NO(g) + 4H2O
(a)  What is the oxidation number of nitrogen in NO-3 and No.? (2 mks)
(b)  With respect to nitrogen, explain whether the above reaction is an oxidation of reducing process.     ( 1 mk)
4. 
On strong heating, sodium nitrate liberates oxygen gas, draw a labeled diagram of set up that could be used for heating sodium nitrate and collecting the oxygen gas liberated.        ( 3 mks)

Complete the diagram below to show how sample of solution of ammonia can be prepared in the laboratory

6. 
Urea (NH¬3 + CO2 → (NH2)2 CO (aq) + H2O (l)
In one process 680 kg of ammonia were reacted with excess carbon (IV) oxide. Calculate the mass of urea that was formed. (H= 1.0) (C = 12.0) (N= 14.0) (0= 16.0) and relative molecular mass of ammonia = 17   ( 3 mks)
7. 
The scheme below show some reactions sequence starting with solid “N”
(a)  Identify solid “N”      (1 mk)
(b)  Write the formula of the complex ion present in solution C. (1 mk)
8. 
A study set up apparatus shown below to prepare ammonia gas and react it with copper (II) sulphate solution

(a) Identify solution “V”     ( 1 mk)
(b) State the observation which were made in the beaker ( 2 mks)
9. 
In an experiment, ammonium chloride was heated in a test tube. A moist red litmus was placed in a mouth of the test tube first change to blue then read. Explain these observations     ( 3 mks)
10. 
When potassium nitrate is heated it produce potassium nitrate and gas C1
(i) Identify gas C1
(ii) Name the type of reaction undergone by potassium nitrate
11. Ammonium nitrate was gently heated and the products collected as shown in the diagram below

Describe one chemical test and one physical property that can be used to identify gas G.        ( 3 mks)

12. 
(a)  When a red litmus paper was dropped into the resulting solution. It turns blue, give a reason for this observations   ( 1 mk)
(b)  What is the function of the funnel?


When ammonium Nitrate is heated in the set up below a colourless gas “A” is produced

(i) Identify gas “A”
(ii) State and explain the precautions that must be taken before heating is stopped       ( 2 mks)
14. 
The diagram below shows a set up that was used to prepare and collect a sample of nitric acid in the laboratory
(a)  Give a reason why it is possible to separate nitric acid from the sulphuric acid in the set up      ( 1 mk)
(b)  Name another substance that can be used instead of potassium nitrate
         ( 1 mk)
(c)  Give one use of nitric acid     ( 1 mk)
15. 
The first step in the industrial manufacture of nitric acid is the catalytic oxidation of ammonia gas
(a)  What is the name of the catalyst used     ( 1 mk)
(b)  Write the equation for the catalytic oxidation of ammonia gas ( 1 mk)
(c) Nitric acid is used to make ammonium nitrate. State uses of ammonium nitrate         ( 1 mk)
16. 
State and explain the observation made when excess ammonia gas reacts with chlorine gas        ( 3 mks)
17. When magnesium was burnt in air, a solid mixture was formed. On addition of water to the mixture a gas which turned moist rd litmus paper blue was evolved. Explain these observations.     ( 2 mks)
18. 
In an experiment, ammonia gas was prepared by heating ammonium salt with an alkali. After drying 120 cm3 of ammonia gas was collected at room temperature and pressure. All the ammonia gas was then reacted completely with 250 cm3 solution of phosphoric acid.
(a)  What is meant by the term alkali?     ( 1 mk)
(b)  Explain using the physical properties of the gas, why ammonia is not collected
(i)  Over water       ( 1 mk)
(ii)  By downward delivery     ( 1 mk)
(c)  Ammonia turns wet red litmus paper blue. Which ions are responsible for this reaction?        ( 1 mk)
(d)  Calculate the number  of moles  of ammonia gas that were collected in the above experiment given that one mole of gas occupied a volume  of 24000cm3 at room temperature  and pressure   ( 3 mks)
(e)  The equation below shows the reaction between ammonia and phosphoric acid.
3NH3(g) + H3PO4(aq) →(NH4)PO4(aq)
(i)  Explain how  crystals of ammonium  phosphate could be obtained in this  experiment       ( 2 mks)
(ii) Calculate  the maximum mass of ammonium  phosphate that could  be  obtained  in this experiment     ( 2 mks)
(N= 14.0) (0= 16.0) (P = 31.0) (H= 1.0)
19. 
(a) The diagram below shows a set up that can be used to obtain nitrogen gas in an experiment
(i)  Name liquid “L”      ( 1 mk)
(ii)  What observations would be made in tube “K” after heating for some time          (1 mk)
(iii)  Write an equation for the reaction that took place in tube “k”  (1 mk)
(iv)  If 320 cm3 of ammonia gas reacted completely with copper (II) oxide calculate
(i)  The volume of nitrogen gas produced  ( 1 mk)
(ii)  The mass of copper oxide that reacted  ( 3 mks)
(iii)  At the end of the experiment, the pH of the water in the beaker was found to be 10: Explain     ( 2 mks)
(b)  In another experiment a gas jar, containing ammonia was inverted over a burning splint. What observations would be made?  ( 1 mk)  
20. a) The diagram below represents a set up used to obtain nitrogen from air. Study and answer the questions that follow
(i)  Name solid Q      ( 1 mk)
(ii)  What is the purpose of sodium hydroxide  ( 1 mk)
(iii)  Write an equation for the reaction which took place in tube “P” 
( 1 mk)
(iv)  Give the name of one impurity in the nitrogen gas obtained ( 1 mk)
(v)  Give a reason why liquid nitrogen is upside for storage of semen for artificial insemination     ( 1 mk)
(b)  The set up below was used to prepare nitric acid

(i)  Give the name of liquid “R”    ( 1 mk)
(ii)  Write an equation for the reaction which took place in the glass retort       ( 1 mk)
(iii)  Explain the following
(i)  Nitric acid is stored in dark bottles  (1 mk)
(ii)  The reaction between copper metal with 50% nitric acid in an open tube gives brown fumes  ( 2 mks)
(c)  A factory uses nitric acid and ammonia gas as the only reactant for the preparation of the fertilizer. If the daily production of the fertilizer is 4800 kg, calculate the mass of ammonia gas used daily (N= 14.0), (0 = 16.0), (H= 1.0)       ( 3 mks)
21. The flow chart below shows the industrialization of ammonia and the process used in the manufacture of some ammonium compounds. Study it and answer the questions that follow

(a) Give the name of the
(i)  Process in step 1      ( 1 mk)
(ii)  Reaction that takes place in step 5    ( 1 mk)
(b) State one other source of hydrogen gas apart from natural gas  ( 1 mk)
(c) Explain why it is necessary to compress nitrogen and hydrogen in this process        ( 2 mks)
(d) Write an equation for the reaction which takes place in step 6  ( 1 mk)
(e) Name the catalyst and reagents used in step 3  ( 2 mks)
(f) Name compound Z1      (1mk)
(g) Give one commercial used of compound Z2   (1 mk)
22. 
(a) The flow chart below shows some reactions starting with lead (II) nitrate. Study it and answer the questions that follows.
(i)  State the conditions necessary in step 1  ( 1 mk)
 I.  Identify I: reagent K     ( 1 mk)
 II.  Gas Q       ( 1 mk)
 III.  Acid products “S” and “R”    ( 1 mk)
(ii)  Write
 I. The formula of the complex ion formed in step 3  ( 1mk)
 II. The equation of the reaction in step 4   ( 1 mk)
(b)  The use of materials made of lead in roofing and in water pipes is being   discouraged. State
(i) Two reasons why these materials have been used in the past
(ii) One reason why their use is being discouraged  ( 1 mk)
(c)  (i) The reaction between lead (II) nitrate and concentrate sulphuric
acid starts but steps immediately explain  ( 2 mks)
(ii)  Name one suitable reagent that can be reacted with concentrated sulphuric acid to produce nitric acid
23. Write an equation to show the effect of heat on the nitrate of:
(i)  Potassium        ( 1 mk)
(ii)  Silver        ( 1 mk)
24. 
(a)  Study the flow chart below and answer the questions that follow

(i)  Identify gas J       ( 1 mk)
(ii)  Using oxidation numbers, show that ammonia is the reducing agent in step (VI)      ( 2 mks)
(iii)  Write the equation for the reaction that occurs in step (V) ( 1 mk)
(iv)  Give one use of ammonia nitrate    ( 1 mk)
(b) The table below shows the observations made when aqueous ammonia was added to cations of element E, F, and G until in excess
Cations Addition of a few drops of aqueous ammonia Addition of excess aqueous ammonia
E White precipitate Insoluble
F No precipitate No precipitate
G White  precipitate Dissolve
(i)  Select the cation that is likely to be Zn2+   ( 1 mk)
(ii)  Given that the formula of the cations of element E is E2+, write the ionic equation for the reaction between E2+(aq) and aqueous ammonia       ( 1 mk)
25. Nitric (Nitric (V) acid is prepared in the laboratory by the action of concentrated acid on a suitable nitrate and distilled off nitric acid. The reaction is carried out in all glass apparatus.
(i)  Why is an all glass apparatus desirable in this preparation? (1 mk)
(ii) Pure nitric (v) acid is colourless liquid but the product in this preparation is yellowish in colour explain.                                                 (imk)
(iii) How can this yellow colour be removed from the acid.   (1 mk)
26. A dry gas X was passed over heart copper (ii) oxide. A brown residue, a colourless liquid “y” and a colourless gas “z” were formed. Gas “z” has no effect on litmus papers and does not support combustion
(a) Suggest the identities of x, y, z and a colourless liquid         (4mks)
(b)  Write an equation for the reaction above.
27. Study the chart below for the large scale production of nitrogen.
 (a) Explain briefly each of the process P and Q.     (2mks)
 (b)  How is nitrogen eventually obtained from step “C”.   (2mks)
28. The following is flow chart representing the manufacture of a fertilizer.

 (i )  Write an equation for the reaction in chamber A    (1mk)
 (ii)  Name the catalyst in chamber “B”                          (1mk)
29. Study the flow chart below an answer the question that follows.

Identify
(i)  Liquid Q
(ii)  Gas x 
(b)  Write the equation between the brown gas above and water.                                                                                                    
30. Study the apparatus and answer the Questions follow?
(a)  Why doses nitric (v) acid appears yellow?                                      (1mk)
(b) When strongly heated brown fumes are evolved. What are these fumes (1 mk)
(c)  Give the identity of gas Q and give its test.                         (1mk)
(d)  State the use of glass wool and the role of sand in the experiment.  (2mk)
(e)  Write an equation to show the decomposition of nitric acid when strongly 
heated                                                                                                (1mk)
31. The diagram below shows an investigation on a property of ammonia gas

(a)  The platinum wire is observed to glow. Explain the cause of that observation         (2 mks)
(b)  State the observations made when the rubber bang is removed.  (1mk)
32. The reaction below represents a major reaction in the industrial process.
 3H2 + N2 (g)    2NH3 (g)
(a)  Name the industrial process      (1 mk)
(b)  Name the catalyst used in above process    (1 mk)
(c)  Explain the following observations. When ammonia gas mixed with oxygen is sparked over platinum gauze wire, brown fumes are evolved (2 mks)
33.  The scheme below shows some reactions starting with salt “P” study it and answer the questions that follows
(a)  Which ions are contained in solution “P”   (1 mk)
(b)  Write the formula of solid Q and the brown precipitate (2 mks)
(c)  Write an equation for the formation of
 (i)  Brown precipitate     (1 mk)
 (ii)  Solid Q
34.  The flow chart below illustrates the major steps in the manufacture of nitric (v) acid. Study it and answer the question that follows.

(a)  Give reasons for purifying raw material “A” and “B”  ( 1 mk)
(b)  Name the substance D, E and F     (1 mk)
(c)  Name the parts labeled D, E and F    ( 3 mks)
(d)  Write chemical equations for the reactions taking place in
 (i)  Chamber D       ( 1 mk)
 (ii)  Chamber F       ( 1 mk)
(f)  A mixture that comes out is 65% Nitric (V) acid and 35% water. How could concentration of nitric acid be increased?   ( 1 mk)
(g)  Give one use of Nitric (V) acid     
(h)  When copper metal is reacted with concentrated Nitric (V) acid a brown gas is evolved, explain      ( 1 mk)
 
TOPIC 5
SULPHUR AND ITS COMPOUNDS
1. 
Study the flow chart below and answer the questions
2. 
The diagram below represents the extraction of sulphur by frash process
(a)  Name the substance that passes through tube I and II (2 mks)
(b)  What is the purpose of the hot compressed air in this process (10 mks)

3. 
State what would be observed when dilute hydrochloric acid is added to product formed when a mixture of iron fillings and sulphur are heated ( 1 mk
4. 
Study the flow chart below and answer the questions that follow
(a)  Name compound “T” and gas “U”    ( 2 mks)
(b)  Give a chemical test that you could use to identify gas “U” ( 1 mk)
5. 
Sulphur (IV) oxide and nitrogen (IV) oxide react as shown in the equation below
SO2(g) + NO2 →SO3(g) + NO(g)
(i)  Using oxidation numbers of either sulphur or nitrogen show that this is a redox reaction       ( 2 mks)
(ii)  Identify the reducing agent     ( 1 mk)
6. 
In an attempt to prepare – sulphur (IV) oxide gas, dilute sulphuric acid was reacted with Barium sulphite. The yield of sulphur (IV) oxide was found to be negligible. Explain       ( 2 mks)
7. 
When a solid sample of sulphur is heated in a test tube. It changes into a liquid, which flow easily. On further heating the liquid darkness and does not flow easily. Explain these observation     ( 3 mks)
8. 
A certain matchstick head contains potassium Chlorate and sulphur. On striking two substances react to produce sulphur (IV) oxide and potassium chloride. Explain the environmental effect of using such matches in large numbers
         ( 2 mks)
9. 
Describe a simple laboratory experiment that can be used to distinguish between sulphide and sodium carbonate.     ( 2 mks)
10. 
What observation would be made if hydrogen sulphide gas was bubbled though a solution of Zinc- nitrate?      ( 1 mk)
11. 
The apparatus shown below was set up to prepare and collect hydrogen sulphide gas
(a)  Name solid C       ( 1 mk)
(b)  Give a reason why warm water is used   ( 1 mk)
(c)  What observations would be made if hydrogen sulphide gas was bubbled into a solution of lead (II0 Nitrate    ( 1 mk)
12. 
Concentrated Nitric acid was added to iron (II) sulphate acidified with dilute sulphuric acid and the mixture was heated. The solution turned from pale green to yellow with evolution of brown gas. Explain this observation. ( 3 mks)
13. 
In an experiment 30cm3 of 1.0m. sulphuric acid were reacted with 30cm3 of 0.1m sodium hydroxide.
(a)  Write an equation for the reaction that took place  ( 1 mk)
(b)  State the observation made when both blue and red litmus papers were dropped with the mixture     ( 1 mk)
(c)  Give a reason for your answer in (b) above   ( 1 mk)
14. 
Sulphur exist in two crystalline forms
(a)  Name one crystalline form of sulphur   ( 1 mk)
(b) State two uses of sulphur     ( 1 mk)

15. 
Oleum (H2S2O7) is an intermediate product in the industrial manufacture of sulphuric acid
(a)  How is Oleum converted to sulphuric acid?   ( 1 mk)
(b)  Give one use of sulphuric acid    ( 1 mk)
16. 
Dilute hydrochloric acid and sodium sulphite were reacted as shown in the set up below

(a)  Name the gas produced in the flask    ( 1 mk)
(b)  Give two reasons why no gas was collected in the gas jar ( 2 mks)
17. 
Determine the oxidation state of sulphur in the following compounds
(a)  H2S        ( 1 mk)
(b)  Na2S2O2       ( 1 mk)
18. 
When hydrogen sulphide gas was bubbled into aqueous solution of iron (lll) chloride a yellow precipitate was formed
(a)  State another observation that was made   ( 1 mk)
(b)  Write an equation for the reaction that took place  ( 1 mk)
(c)  What type of reaction was undergone by hydrogen sulphide gas?
19. 
Study the flow chart below and answer the question that follows

(a)  Name reagent “Z”      ( 1 mk)
(b)  Describe the process which takes place in step 2   ( 1 mk)
(c)  Identify the white solid     ( 1 mk)

20. 
Below is a sketch of a graph showing the change in viscosity ((Ease to flow) with temperature when solid sulphur is heated.

Describe what happens to the sulphur molecules when sulphur is heated from 1500C to about 2000C.      ( 2 mks)
21. 
(a)  State the observation made at the end of the experiment when a mixture of iron powder and sulphur is heated in a test tube.  ( 1 mk)
(b)  Write an equation for the reaction between the product in (a) above and dilute hydrochloric acid.     ( 1 mk)
(c)  When a mixture of iron powder and sulphur is heated, it glows more brightly than that of iron fillings and sulphur. Explain this observation         ( 1 mk)

22. 
(a)  The graph below shows the solubility of sulphur (IV) Oxide gas at different temperatures. Use the information in it to answer the questions that follows
(i)  From the graph determine
I.  The lowest temperature at which 1,000 cm3 of solution would contain 116g of sulphur (IV) oxide  ( 1 mk)
II.  The maximum mass of sulphur (IV) oxide that would dissolves in 15 litres of solution at 100C
(ii)  Sodium hydroxide reacts with sulphur (IV) oxide according to the following equation
2NaOH (aq) + SO2 (g) → Na2SO3(aq) + H2O(l)
Using the information in the graph, determine the volume of 2m sodium Hydroxide  required to completely neutralize one litre of saturated  sulphur  (IV) oxide at 230C (S= 32.0 ) (0= 16.0) 
(3 mks)
(b)  Study the flow chart below and answer the questions that follow

(i)  Write equations for the reaction taking place at
I.  The roasting furnace    ( 1 mk)
II.  The absorption tower    ( 1 mk)
III.  The diluter     ( 1 mk)
(ii)  The reaction that takes place in chamber “K” is
SO2 (g) + ½ O2 (g)   SO3 (g)
I. Explain why it is necessary to use excess air in chamber “K’ ( 1 mk)
II. Name another substance used in chamber “k”    ( 1 mk)
23. 
The reaction between sulphur (IV) oxide and oxygen to form sulphur (VI) oxide in the contact process is exothermic
2SO2(g) + O2 (g)   2SO3 (g)
A factory manufacturing sulphuric acid by contract process produces 350 kg of sulphur (VI) Oxide per day. (Condition for the reaction: a catalysts 2 atmospheres pressure and temperature between 400 – 5000C
(i)  What is meant by an exothermic reaction?
(ii)  How would the yield per day of sulphur (VI) oxide be affected if temperatures lower than 4000 C is used explain  ( 3 mks)
(iii)  All the sulphur  (VI) oxide produced was absorbed  in concentrated sulphuric acid to form oleum
H2S04 (l) + SO3 (g) → H2S2O7 (l)
Calculate the mass of oleum that was produced per day ( 3 mks)
(S= 32.0) (O= 16.0) (H= 1.0)
24. 
(a)  Name one ore from liquid which copper metal is extracted ( 1mk)
(b)  The flow chart below shows a sequence of reaction starting with copper. Study it and answer the questions that follows
(i)  Identify gas “p” and reagent Q and “R”  ( 2 mks)
(ii)  Write an equation for the reaction that place in step 5 ( 1 mk)
(iii)  State the observation made in steps 4 and 7   ( 1 nk)
(c)  Bronze is an alloy of copper and another metal
(i)  Name the other metal      ( 1 mk)
(ii)  Give one use of bronze     ( 1 mk)

25. 
The diagram below illustrates how sulphur is extracted by frasch process
(a)  Label the pipe through which superheated water is pumped in ( 1 mk)
(b)  The equation below shows the oxidation of sulphur (IV) oxide to sulphur (VI) oxide in the contact process
2SO2 (g) + O2 (g) →2 SO3(g) ∆H = - 196KJ
(i)  Name the catalyst used in this process  ( 1 mk)
(ii)  State and explain the effect on the yield of sulphur (VI) oxide when
I.  The temperature is increased   ( 2 mks)
II.  The amount of oxygen is increased  ( 2 mks)
(iii)  Describe how sulphur (VI) oxide is converted to sulphuric acid in the contact process.     ( 2 mks)
(c)  Ammonium sulphate is a fertilizer produced by passing ammonia gas into concentrated sulphuric acid
(i)  Write the equation for the reaction   ( 1 mk)
(ii)  Calculate the mass in kg of sulphuric acid required to produce 25kg of fertilizer (S= 32.0) (0= 16.0) (N = 14.0) (H. 1.0) (3 mks)
26. 
(a)  The diagram below shows some processes that takes place during the industrial manufacture of sulphuric acid.

(i)  Write the equation for the reaction in which sulphur (IV) Oxide is produced
(ii)  Why is it necessary to keep the gas pure and dry?  ( 1 mk)
(iii)  Describe the process that takes place in chamber G  ( 1 mk)
(iv)  Name the gases that escape into the environment  ( 1 mk)
(v)  State and explain the harmful effect on the environment of one of the gases
(vi)  Give one reason why it is necessary to use 2- 3 atmospheric pressures and not more     ( 1 mk)
(b)  (i)  Complete the table below to show the observations made when 
concentrated sulphuric acid add to the substances shown 
Substances Observations
Iron fillings 
Crystals  of white  sugar 
(ii)  Give a reason for the observation made using
I.  Iron fillings      ( 1 mk)
II.  Crystal of white sugar     ( 1 mk)
(c)  Name one fertilizer made from sulphuric acid   ( 1 mk)
(d)  Suggest a reason why BaSO4 ( a pigment made from sulphuric acid) would be suitable  in making paint  for cars    ( 1 mk)
27. When sulphur is heated in a test tube, the yellow crystal melt to form a golden yellow liquid, which changes at 1800C. Into dark brown, very viscous liquid more heating to 400C a brown less viscous liquid.
(i)  What is the molecular mass sulphur in the yellow crystals ( 1 mk)
(ii)  If the brown liquid at 4000C is cooled rapidly at room temperature, which form of sulphur is produced?     ( 1 mk)
(iii) Explain why the molten sulphur becomes viscous  ( 2 mks)
28.  (a) State two observations  made  when  acidified  potassium permanganate is reacted  with hydrogen  sulphide     ( 2 mks)
(b)  Explain the observation made in (a) above   (1 mk)
(c)  Write an ionic equation for the above reaction  ( 1 mk)
29. Below  is a flow  chart showing  some of the  major  steps  involved  in the manufacture of  sulphuric (VI) acid by contact process

(a)  Identify
 (i)  Substance A      ( 1 mk)
 (ii)  Catalyst used in chamber “P”    ( 1 mk)
(b)  The conversation of S) 2 to SO3 I the contact process is shown by the equation
 2SO2 (g) + O2 (g) → 2SO3 (g) →∆H = -197KJ
 What would be the effect of?
 (i)  Increasing the concentration of oxygen  ( 1 mk)
 (ii)  Increasing the temperature    ( 1 mk)
(c)  Write an equation for
 (i)  The formation of Oleum    ( 1 mk)
 (ii)  Formation of sulphuric (IV) acid from Oleum  (1mk)
30. State and explain  the  observation  made when hydrogen – sulphide  gas  is  bubbled  in a  solution of iron (III) ions    ( 1 mk)
31. State all the changes that will be seen when concentrated sulphuric acid is added to cane sugar in a boiling tube.   ( 2 mks)
32. The set up below shows preparation of sulphur (VII) oxide study it and answer the questions that follows.
(b) Write an equation for the reaction taking place in the combustion tube. ( 1 mk)
33. When sulphur (IV) oxide is passed into aqueous solution of chlorine the greenish yellow colour of chlorine disappears. Write equation for the reaction taking place
           ( 1 mk)
34. Study the flow chart below and answer the question that follows

(a)  Name solid P        ( 1 mk)
(b)  Give the formula of sodium salt     ( 1 mk)
(c)  Name gas R        ( 1 mk)
(d)  Write an equation for the reaction between Nitric acid and solid “P”           ( 1 mk)
35. Sulphur is one of the  elements that exhibits allotropy
(i)  What is allotropy       ( 1 mk)
(ii)  Give another element other than sulphur that shows allotropy ( 1 mk)
(iii)  Name two allotropes of sulphur    ( 2 mks)
(iv)  State two major uses of sulphur    ( 2 mks)
36.  9.0g of zinc sulphide reacted with 100cm3 of 0.2m sulphuric acid. Determine the reagent that was in excess. (Zn = 65, S= 32)    ( 2 mks)
 
TOPIC 6
CHLORINE AND ITS COMPOUNDS
1. 
When excess chlorine gas is bubbled through dilute sodium hydroxide solution the resulting solution act as a bleaching agent.
(a)  Write an equation for the reaction between chlorine gas and sodium hydroxide solution       (1 mk)
(b)  Explain how the resulting solution acts as a bleaching agent (2mks) 
2. 
A solution of chlorine in Tetracloromethane turns colourless when propene gas is bubbled through it
(a)  What type of reaction takes place     ( 1 mk)
(b)  Write an equation for the above reaction    ( 1 mk)
3. 
The reaction of propane with chlorine gas gave a compound with formula C3H7CL
(a)  What condition is necessary for the above reaction to take place ( 1 mk)
(b)  Draw a structured formula of compound C3H7CL
4. 
In an experiment chlorine gas was passed into moist hydrogen sulphide  in a boiling as shown in the diagram
(a)  What observations was made in the boiling tube  ( 1 mk)
(b)  Write an equation for the reaction which took place in the boiling tube
(c)  What precautions should be taken in carrying out this experiment? Give a reason.        ( 1 mk)

The diagram below shows a set up for the laboratory preparation and collection of dry chlorine gas

(a)  Name
(i)  Substance G      ( 1 mk)
(ii)  Suitable drying agent     ( 1 mk)
(b)  What property of chlorine make it possible for it to be collected as shown in the diagram       ( 1 mk)
6. 
The following two sets were carried out on chlorine water contained in two test tubes
(a)  A piece of blue flower was dropped into the first test tube. Explain why the flower was bleached.     ( 2 mks)
(b)  The second test tube was corked and exposed to sunlight. After a few days it was found to contain gas that rekindled a glowing splint. Write an equation for the reaction which produced the gas.  (1 mk)
7. 
The set up below was used to prepare hydrogen chloride gas and react it with iron powder. Study it and answer the questions that follows

At the end of the reaction, the iron powder turned to light green solid
(a)  Identify the light green solid    ( 1 mk)
(b)  At the beginning of the experiment the pH of the solution in container “L” was about 14. At the end the pH was found to be 2. Explain.
8. 
Calcium Oxide can be used to dry ammonia gas
(a)  Explain why calcium oxide cannot be used to dry hydrogen chloride gas
(b)  Name one drying agent for hydrogen chloride gas  ( 1 mk)
9. 
The reaction between hot concentrated Sodium hydroxide and chlorine gas produces sodium chlorate (V), sodium chloride and water
(a)  Write the equation for the reaction    ( 1 mk)
(b)  Give one use of sodium chlorate (V)    ( 1 mk)
10. 
Water from a town in Kenya is suspected to contain chloride ions but not sulphate ions. Describe how the presence of chloride ions in the water can be shown          ( 2 mks)
11. 
The diagram below represents the set up that was used to prepare and collect dry hydrogen chloride gas in the laboratory.
(i)  State the purpose of concentrated sulphuric acid in the wash bottle ( 1 mk)
(ii)  Write an equation for the reaction between dry hydrogen chloride gas and heated iron        ( 1 mk)
(iii)  Hydrogen chloride gas is dissolved in water to make hydrochloric acid. State one use of hydrochloric acid     ( 1 mk)
12. 
Complete the following table by filling in the missing test and observations
No. Gas Test Observation
I Chlorine Put  a moist  red  litmus paper into  the gas 
II Sulphur (IV) Oxide  Paper turns green
III Butene Add drop of bromine water 
13. 
In an experiment, a test tube full of chlorine water was  inverted  in chlorine water as shown in the diagram below and the set up left in sunlight for one day.

After one day, a gas was found to have collected in the test- tube
(a)  Identify the gas
(b)  What will happen to the pH of the solution in the beaker after one day? Give an explanation.      ( 2 mks)
14. 
The diagram below is part of a set up used in the laboratory preparation of a gas
Complete the diagram to show how a dry sample of the gas can be collected 
         ( 3 mks)
15. 
The diagram below shows an incomplete set up of the laboratory preparation and collection chlorine gas. Study it and answer the questions that follows
(i)  Complete the set up to show how dry chlorine gas may be collected
         ( 3 mks)
(ii)  What is the function of the water in flask L
(iii)  The equation for the redox reaction that takes place is
 Mno2(s) + 4Hcl (aq) → Mncl2 (aq) + 2H2O (l) + Cl2 (g)
Explain using oxidation numbers which species is reduced ( 2 mks)
16. 
The set up below was used to prepare anhydrous chloride of a number of elements in laboratory where no fume cupboard was available. The chloride were to be collected in flask 1
The following table shows the melting and boiling points of the chloride that were prepared
Chloride Nacl Alcl3 Sicl4 Pcl3
Melting point in 0C 801 Sublime (178) -70 -91
Boiling point 0C 1413  58 76
(a)  Explain why it is necessary to pass dry chlorine gas through the apparatus before heating each element     ( 2 mks)
(b)  Give two reasons why tube II is filled with soda lime (a mixture of sodium hydroxide and calcium hydroxide    ( 2 mks)
(c)  Explain why it would not be possible to collect any sodium chloride in flask I        ( 1 mk)
(d)  Name one other substance that can be used in tube II ( 1 mk)
(e)  Write an equation for the reaction that forms phosphorous (III) chloride
(f)  Describe how you would separate a mixture of sodium chloride and aluminium chloride      ( 2 mks)
17. 
(a)
(i)  In the spacer provided sketch a diagram to show how hydrogen chloride gas can be prepared and collected in the laboratory using sodium chloride  and concentrated sulphuric acid (the gas need not be dry) ( 4 mks)
(ii)  Write an equation for the reaction that takes place
 (iii)  Name one drying for hydrogen chloride gas
(iv)  State and explain the observation that would be made when hydrogen chloride gas is bubbled through a solution of lead (II) nitrate ( 3 mks)
(v)  Concentrated hydrochloric acid is used for removing oxide from metals surfaces (pickling). Explain why concentrated nitric acid cannot be used for the same purpose
(b)  A sample of hydrogen chloride gas dissolved in water to make 250 cm3 of solution. 25 cm3 of the solution required 46 cm3 of 11.0m sodium hydroxide for complete neutralization.
(i)  Calculate the number of moles of hydrochloric acid in 25 cm3 solution
         ( 3 mks)
(ii)  Determine the mass of hydrogen chloride that was dissolved to make 250 cm3 of solution. (Cl = 35.5) (H= 1.0)    ( 2 mks)
18. 
(a)  Give the name of one reagent which when reacted with concentrated hydrochloric acid produces chlorine gas    ( 1 mk)
(b)  A student set out to prepare iron (lll) chloride using apparatus shown in the diagram below
(i)  Explain why it is necessary to pass chlorine gas through the apparatus  before heating begins?      ( 1 mk)
 (ii)  Calcium oxide would be preferred to calcium chloride in the guard tube
          ( 1 mk)
(iii)  What property of iron (III) chloride makes it possible to be collected as shown in the diagram       ( 1 mk)
(iv)  The total mass of iron (III) chloride formed was found to be 0.5g. Calculate the volume of chlorine gas that reacted with iron. (Fe = 560 (Cl = 35.5) and molar gas volume of 298k is 24,000 cm3 ( 3 mks)
(c)  When hydrogen sulphide gas passed through a solution of iron (III) chloride the following observation was made
The colour of the solution changed from reddish brown to green and yellow solid was deposited. Explain these observations ( 2 mks
(d)  State and explain the observations that would be made if a moist blue-litmus paper was placed in a gas jar full of chlorine gas ( 2 mks) 
. Study the flow chart below and answer the questions that follows

(a)  Identify substance A and B     ( 2 mks)
(b)  Name process “C”      ( 1 mk)
(c)  Give one use of P.V.C      ( 1 mk)
(d)  Write an ionic equation for the reaction in which chlorine gas is produced
         ( 1 mk)
(e)  State and explain the observation that would be made if chlorine gas was bubbled into an aqueous solution of sodium iodide   ( 1 mk)
(f)  In the preparation of a bleaching agent (Sodium hypochlorite) excess chlorine gas was bubbled into 15 litres of cold 2M sodium hydroxide
(i)  Write an equation for the reaction between chlorine gas and dilute sodium hydroxide       ( 1 mk)
(ii)  (a)  Calculate the number of moles  of sodium hydroxide  used ( 1 mk)
 (b)  Calculate the mass in kg of sodium hypochlorite produced 
  (Ma = 23) (cl = 35.5) (O=16)    ( 3 mks)
 
19. (a)
The table below shows some properties of chlorine, bromine and  iodine
Elements Formula Colour  and state at room temperature Solubility  in water
Chlorine Cl2 i……………. Soluble
Bromine Br2 Brown liquid ii…………..
Iodine L2 iii…………….. Slightly soluble
Complete the table by giving the missing information in (i) (ii) and (iii)
         ( 3 mks)
(b)  Chlorine gas is prepared by reacting concentrated hydrochloric acid with manganese (IV) oxide
(i)  Write the equation for the reaction between concentrated hydrochloric acid and manganese (IV) oxide    ( 1 mk)
(ii) What is the role of manganese (IV) oxide in this reaction ( 1 mk)
(c)  (i)  Iron (III) chloride react with chlorine  gas to form substance “E”
identify substance “E”    ( 1 mk)
(ii)  During the reaction in C (i) above 6.30g of iron (II) chloride were converted to 8.06g   of substance “E” Calculate the volume of chlorine gas used. (CL= 35.5) Molar gas volume a room temperature = 24000 cm3 (Fe = 56)    ( 3 mks)
(d)  Draw and name the structure of the compound formed when excess chlorine gas is reached with ethane gas   ( 3 mks)
(e)  Give one industrial use of chlorine    ( 1 mk)
20. 
The diagram below shows the set up used in an experiment to prepare chlorine gas and react it with aluminium foil. Study it and answer the questions that follow

(a)  In the experiment, concentrated hydrochloric acid and potassium manganate (VII) were used to prepare chlorine gas. State two precautions that should be taken in carrying out this experiment.  ( 2  mks)
(b)  Write the formula of another compound that could be used instead of potassium manganate (VII)     ( 1 mk)
 (c)  Explain why is necessary to allow the acid to drip slowly onto potassium manganate (VII) before the aluminium foil is heated. (2 mks)
(d)  State the property of the product formed in the combustion tube that makes it possible for it to be collected in the receiver. ( 1 mk)
(e)  When 1.08g  of aluminium foil were heated in a stream of chlorine gas, the mass of the product formed was 3.47g. Calculate the:
(i)  Maximum mass of the product formed if chlorine was in excess
 (AL = 27; Cl = 35.5)     ( 3 mks)
(ii)  Percentage yield of the product formed  ( 1 mk)
(f)  Phosphorous trichloride is a liquid at room temperature what modification should be made to the set up if it is to be used to prepare phosphorous trichloride       ( 1 mk)
21.  (i)  What is the action of chlorine on cold dilute  sodium hydroxide ( 1 mk)
(ii)  Write down the equation for the above reaction   ( 1 mk)
22. If chlorine gas is passed over  heated iron fillings and the products dissolved in water, a yellow solution is formed
(i)  Identify the yellow solution      ( 1 mk)
(ii)  What would be observed if aqueous sodium hydroxide solution was added to the yellow solution       ( 1 mk)
(iii)  Write an equation for the reaction between the yellow solution and sodium hydroxide        ( 1 mk)
23. A solution of hydrogen chloride in methylbenzene (toluene) has no effect on limestone. A solution of hydrogen chloride in water reacts with limestone to produce a gas explain       ( 1 mk)
24. The diagram below represents the industrial manufacturer of hydrochloric acid. Study it and answer the questions that follow.

(a)  Name the reagents “W” and “Y”     ( 1 mk)
(b)  Explain the role of the glass beads in the absorption chamber ( 1 mk)
(c)  Write an equation for the reaction in chamber “X”    ( 1 mk)
(d) Explain why hydrochloric acid formed appears yellow in colour ( 1 mk)
25. The diagram  below shows preparation of hydrochloric acid

(i)  State one mistake in the diagram
(ii)  Hydrogen chloride does not have any effect on litmus paper unlike hydrochloric acid. Explain      ( 1 mk)
26. The flowchart below summarizes the results of series of chemical reactions; study it and answer the questions that flows

(a)  Identify gas “A” gas “D” substance E and F, Gas J solution K and metal Q
         ( 4 mks)
(b)  What is the effect of solution “B” and a solution “C” on dry blue litmus paper?  Explain      ( 2 mks)
(c)  What would you observe if excess ammonia solution is added to the solutions of substance “E” and “F” separately, explain your observations         ( 2 mks)
(d)  What reagent would you use to convert?
(i)  Substance “E” to substance “F”   ( 1 mk)
(ii)  B to gas D      ( 1 mk)
(e)  State the condition required in the formation of substance E or F which is not given in the diagram      ( 1 mk)
27. Below is a set up of the apparatus used to prepare a dry sample of chlorine gas in the laboratory?

(a)  State two observation that were made in the reaction  ( 2 mks)
(b)  Suggest two collection that should be made on the above set up so that experiment is successful     ( 2 mks)
(c)  What is the role of water in this set up?    ( 1 mk)
(d)  (i)  Write an equation for the reaction which produces chlorine ( 1 mk)
(ii)  What is the role of water of MNO2 in this reaction  ( 1 mk)
(e)  Determine the mass of chlorine gas formed if 40 cm3 of 11.0 m hydrochloric acid was used in this reaction (Cl= 35.5) ( 3 mks)
(f)  0.53g of chlorine gas was reacted with iron to form 0.81 of product. Determine the molecular formula of the products given that its relative molecular mass is 162.5 (Fe = 56) (Cl = 35.5)  ( 4 mks)
(g)  Name two raw materials that are used with chlorine to produce hydrochloric acid on the large scale    ( 1 mk)
28. The experiment below was set up to prepare iron (iii) chloride tram chlorine

(a)  Name two reagents that could be used to prepare chlorine gas in the laboratory       ( 1 mk)
(b)  Why is it necessary to dry chlorine gas before using it here? ( 1 mk)
(c)  What property of iron (III) chloride makes it possible to collect it as shown?       ( 1 mk)
(d)  Give the names of solid J and state its functions  ( 1 mk)
(e)  Where should this experiment be carried out and why ( 1 mk)
(f)  Give the equation for the reaction that takes place in the combustion tube
(g)  What would be observed if some chlorine water is shaken in gas jar of hydrogen sulphate gas      ( 1 mk)
29. A student set up the apparatus below in the school laboratory to prepare and study the properties of a certain gas A.

(a)  Name gas A       ( 1 mk)
(b)  Write down a chemical equation for the reaction taking place to produces gas “A”
(c)  What major property of “gas” enables the student to collect the gas above as shown in the diagram     ( 1 mk)
(d)  Suggest a possible drying agent if the student want to collect dry sample of the gas       ( 1 mk)
(e)  Large qualities of the gas were bubbled into the same amount of water by passing the gas through an inverted filter funned placed on the surface of the water to prepare a solution “Q”
(i)  Give a reason why a filter funned is necessary in (e) above (1 mk)
(ii)  Some of the resulting solution Ce was mixed with silver nitrate- solution a white precipitate was observed. Name the white precipitate      ( 1 mk)
(iii)  Write down an ionic equation for the formation of the white precipitate in e (ii) above
(iv)  Suggest the identity of solution Q.   ( 1 mk)
 
FORM 4 WORK
TOPIC 1
ENERGY CHANGES
1. 
Below is the energy level diagram for the reaction
½ H2(g) + ½ F2(g) → HF(g)
(a) Calculate the heat of formation of HF (g)   ( 2 msk)
(b) Is this reaction exothermic  or endothermic?  ( 2 mks)
2. 
When excess magnesium powder was added to 100cm3 of 0.5m iron (III) sulphate solution, the pale green colour of solution faded and the temperature rose by 6.00C
(a)  Write an ionic equation for the reaction that takes place ( 1 mk)
(b)  Calculate the molar heat of reaction given that heat change = mass x temperature change x 4.2J/g/0C and the density of solution is 1g/cm3
( 2 mks)
3. 
Explain why the enthalpy of neutralization of ethanoic acid with sodium hydroxide is different from that of hydrochloric acid with sodium hydroxide
         ( 2 mks)
4. 
Use the information below to answer the questions that follows
Equation
Enthalpy = Formation
H2 + ½ O2 (g)     → H2O(l)
        ∆H = -286kj/mole
C (s) + O2 →  CO2
   ∆H = - 394kj/mol
2C(s) + 3H2 (g) + ½ O2(g) → C2H5OH(aq)
     ∆H3 = 277 KJ KJ/Mole
(a)  Define the term “enthalpy of formation of a compound”
(b)  Calculate the molar enthalpy of combustion ∆H4 of ethanol ( 2 mks)
 C2H5OH(aq) + 3O2 (g) →2 CO2 (g) + 3H2O(l)
5. 
When 0.6 g of element “J” were completely burnt in oxygen and all the heat evolved was used to heat 500 cm3 of water, the temperature of water rose from 230C to 330 C. Calculate the relative atomic mass of element “J” given that the specific heat capacity of water = 4.2J/g/k density of water = 1.0g/cm3 and molar heat of combustion of “J” is 380kj/mole    ( 3 mks)
6. 
Sulphur burns in air to form sulphur (IV) oxide. A simple energy level diagram for the reaction is given below. Study the diagram and answer the questions the follows
(a)  What do the following represents? ∆H1 and ∆H3  ( 2 mks)
(b)  Write an expression for ∆H3 and in terms of ∆ H1 and ∆H2 (1 mk)
7. 
Study the information given in the table below and answer the questions below the table
Bond Bond energy lJ/mole
C-H 414
CL- CL 244
C-CL 326
H- CL 431
Calculate the enthalpy change for the reaction
CH4 (g) + Cl2 (g) →CH3OCl (g) + HCl (g)
8. 
Ca(s) + ½ O (g) → CaO (s) ∆ H= -635 Kj/mole
C(s) + O2 →Co2(g) ∆H = -394 Kj/mole
Ca(s) + Co2(g) + 3/2O2 →CaCO3(s) ∆H = -1207Kj/mole
Calculate the enthalpy change for the reaction.  ( 2 mks)
CaO(s) + Co2 (g) →CaCo3(s)
9. 
Hydrogen and Flourine react according to the equation below
H2 (g) + F2 (g) →2HF (g)  ∆H = 538kj
(a)  Sketch an energy level diagram for the forward reaction ( 1 mk)
(b)  Calculate the molar enthalpy of formation of HF (g)  ( 1 mk)
10. 
State and explain the function of tartaric acid in baking powder ( 1 mk)
11. 
Use the equation below to answer the question that follows
K+ (g) + CL (g) → KLC(s): H1 = -701 Kj/mole
KCL(s)   H2O  K+(aq) + CL-(aq)
  H2 = + 14Kj/mole
(a)  What name is given to Hl?    ( 1 mk)
(b)  Calculate the heat change of the process  ( 2 mks)
K+(g) + CL(g) → K+(aq) + CL-(aq)
12. 
Use the following equations to determine the heat evolved when aluminium metal is reached with iron (III) oxide     ( 3 mks)
2Al(s) + 3/2 O2 → Al2O3 ∆H1 = -1673 Kj/mole
2Fe + 3/2 O2 →Fe2O3 ∆H2 = -836.8 Kj/mole
13.  (a)  What is meant by heat of vaporization   ( 1 mk)
(b)  The boiling point of ethanol, propanol and butanol are 780C, 97.20Cand 1170C. Explain this trend     ( 1 mk)

14. 
Copper (II) sulphate reacts with barium chloride according to the equation below
CaSo4 (aq) + BaCl2(aq) →CuCl2(aq) + BaSO4(s)
∆H = -17.7 Kj/mole
Calculate the temperature change when 900 cm3 of 1M copper (II) sulphate were added to 600 cm3 of 1 m barium chloride
Assume heat capacity of solution is 4.2j/g/k and density = 1g/cm3 ( 3 mks)
15. 
Below is a sketch of a reaction profile

(a)  On the diagram shown the heat of reaction ∆ H   ( 1 mk)
(b)  State and explain the type of reaction represented by the profile  (2mks)
16. 
The table below shows some information about element I, II, III and IV which are in the same group of periodic table. Use the information to answer the questions that follows
Element First ionization energy Kj/mole Atomic radius (nm)
I 520 0.15
II 500 0.79
III 420 0.23
IV 400 0.25
State and explain the relationship between the variation in the first ionization energies and the atomic radii.      ( 3 mks)
17. 
The scheme below shows the energy changes that are involved between water and steam. Study it and answer the questions that follows
H2O(s)       ∆H1 H2O(l)    ∆ H2       H2O(g)
       ∆H4     ∆H3
(a)  What name is glue to the energy change ∆ H4  ( 4 mks)
(b)  What is the sign of ∆ H3? Give a reason   ( 2 mks)
18. 
At 200C, No2 and N2O4 Gases exist in equilibrium as shown in the equation below
2NO2     N2 O4  ∆H = -ve
Brown  pale yellow
State and explain the observations that would be made when:
(a)  A syringe containing the mixture at 200C is immersed in ice cold water
(b)  Volume of gas in syringe reduced    (1 ½ )

19. 
The graph below shows a curve obtained when water at 200C was heated for 15 minutes

(a)  What happens to the water molecules between points “W” and “X”
(b)  In which part of the curve does a change of state occurs (1 mk)
(c)  Explain why the temperature does not rise between points X and Y
         (1 mk)

20. 
Study the diagram below and answer the questions that follows
(a)  What do ∆ H1 and ∆ H2 represents?    ( 2 mks)
(b)  Write an expression to show the relationship between ∆H1, ∆H2 and ∆H3
21. 
The thermo chemical equations for the formulation of hydrogen peroxide under standard conditions are:
H2(g) + O2 (g) →H2O (g); ∆Hθf = -133 kJ mol-1
H2 (g) + O2 (g) → H2O2 (l); ∆Hθf = -188kJmol-1
Write the thermo chemical equations for the molar heat of vaporization of hydrogen peroxide       ( 2 mks)
22. 
The diagram below is a sketch of the graph of the non- catalyzed decomposition of hydrogen peroxide.

On the same axis, sketch the graph for the decomposition of hydrogen peroxide when manganese (IV) oxide is added     ( 2 mks)
23. 
The table below gives the solubilities of substances J, K and L at different temperatures
Substance Solubility in grammes per 100g water at
 00C 200C 400C 600C
J
K
L 0.334
27.60
35.70 0.16
34.0
36.0 0.097
40.0
36.6 0.0058
45.5
37.3
Select the substance which, when dissolved in water, heat is given out. Give a reason         ( 2 mks)
24. 
In an experiment to determine the heat of combustion of methanol (CH3OH) a student used a set up like the one shown in the diagram below. Study the set- up and the data below it and answer the questions that follows

Volume of water = 500cm3
Final temperature of water = 27.00C
Initial temperature of water = 20.00C
Final mass of lamp + methanol = 22.11g
Initial mass of lamp + methanol = 22.98g
Density of water = 1.0/ cm3
Heat change = mass x temperature x 4.2j/g/C
(a)  Write an equation for the combustion of methanol  ( 1 mk)
(b)  Calculate
(i)  The number of moles of methanol used in the experiment (C=12), (O= 16) (H=1)      ( 2 mks)
(ii)  Heat change in this experiment
(iii)  The heat of combustion per mole of methanol ( 2 mks)
(c)  Explain why the value of molar heat of combustion for methanol obtained the theoretical value      ( 2 mks)
(d)  On the axis below sketch an energy diagram for the combustion of methanol

 Energy
   Reaction path
25. 
In order to determine the molar heat of neutralization of sodium hydroxide. 100 cm3 of 1m sodium hydroxide and 100 cm3 of 1 m hydrochloric acid both at the same initial temperature were mixed and stirred continuously with a thermometer. The temperature of the resulting solution was recorded after every 30 seconds until the highest temperature of the resulting solution was attained. Thereafter, the temperature of the solution was recorded for a further two minutes
(a) (i)  Why was it  necessary to stir the mixture of two solutions
 (ii)  Write an ionic equation for the reaction which took place ( 1 mk)
The sketch below was obtained when the temperature of the mixture were plotted against time.
Study it and answer the questions that follows

(I)  What is the significance of pointY2   ( 1 mk)
(II)  Explain why there is a temperature change between points
 Y1 and Y2      ( 1 mk)
 Y2 and Y3      ( 1 mk)
(III)  If the initial temperature for both solutions was 24.50C and the highest temperature attained by the mixture was 30.90C. Calculate
(I)  Heat change for the reaction (Specific heat capacity of solution= 4.2j/g/k and the density of the solution = 1.0g/cm3
(II)  Molar heat of neutralization of sodium hydroxide (2mks)
(III)  Explain how the value of the molar heat of neutralization obtained in this experiment would compare with the one that would be obtained if the experiment was repeated using 100 cm3 of 1M ethanoic acid instead of hydrochloric acid.      ( 2 mks)
(b)  On the grid provided below, draw an energy level diagram for the reaction between hydrochloric acid sodium hydroxide  ( 2 mks)

 Energy
   Reaction coordinate
26. 
(a)  Distinguish between exothermic and endothermic reaction ( 2 mks)
(b)  Change of state is either exothermic or endothermic. Name a change of state that is
(i) Endothermic      ( 1 mk)
(ii) Exothermic      ( 1 mk)
(c)  When pure  water is heated  at 1 atmospheric pressure at sea level, the temperature of the water does not rise beyond 1000C even when continued heating. Explain these observations.    ( 1 mk)

 (d) Study the energy cycle diagram below and answer the questions that
follows
     Fe(s)
      Cl2(g) ∆H1
       Cl2(g)    ∆H3  FeCl2(s)
       
       ∆H2
     FeCl3(s)
  (i)  What does ∆ H1 represents    ( 1 mk)
(ii)  Show the relationship between ∆ H1, ∆H2, and ∆H3 ( 3 mks)
(e)  Butane and propane are constituent of cooking gas. Which one produces more energy per mole on combustion? Explain   ( 2 mks)
27. 
(a)  In an experiment to determine the molar heat of reaction when magnesium displaces copper. 0.15g of magnesium powder was added to 25 cm3 of 0.2m copper (ll) chloride solution was 250C while that of the mixture was 430C.
(i)  Other than increase in temperature, state and explain the observation which were made during the reaction    ( 3 mks)
(ii)  Calculate the heat change during the reaction (Specific heat capacity of the solution = 4.J/g/k and the density of the solution = 1g/cm3 ( 2 mks)
(iii)  Determine the molar heat of displacement of copper by magnesium (mg = 24.0)
(iv)  Write the ionic equation for the reaction   ( 1 mk)
(v)  Sketch an energy level diagram for the reaction  ( 2 mks)
28. 
(a)  State two factors that should be considered when choosing fuel for cooking       ( 2 mks)
(b)  The diagram below represents a set- up that was used to determine the molar heat of combustion of ethanol
During the experiment, the data given below was recorded
Volume of water    450cm3
Initial temperature of water   250C
Final temperature of water   46.50C
Mass of ethanol + lamp before burning 125.5g
Mass of ethanol + lamp after burning  124.0g
Calculate the
(a)  Heat evolved during the experiment (Density of water = 1g/cm3), specific heart capacity of water = 4.2 Jg-1 k-1)   ( 2 mks)
(b)  Molar heat of combustion of ethanol (C= 12.0, O = 16.0, H = 1.0) ( 2 mks)
(c)  Write the equation for the complete combustion of ethanol  ( 1 mk)
(d)  The vale of the molar heat of combustion of ethanol obtained in (b) (ii) above is lower than the theoretical value. State two sources of error in the experiment       ( 2 mks)
29. 
(a)  Define the standard enthalpy of formation of a substance ( 1 mk)
(b)  Use the thermochemical equations below to answer the questions that follow
1. C2H6 + 7/2 O2 → 2CO2(g) + 3H2O(l) ; ∆ H1 – 1560 kJmol-1
2. C(graphite) + O2 (g) → CO2 (g) ; ∆H2 – 394 kJMol-1
3. C2 (g) + ½ O2 (g) → H2O (g);  ∆H3 – 286 kJmol-1
(i)  Name two types of heat changes represented by ∆H3
(ii)  Draw an energy diagram for the reaction represented by equation 1.
(3 mks)
(iii)  Calculate the standard enthalpy of formation of ethane (2 marks)
(iv)  When a sample of ethane was burnt, the heat produced raised the temperature of 500g of water by 21.5K. (Specific heat capacity of water = 4.2jg-1 K).
Calculate the:
I. Heat change for the reaction    ( 2 mks)
II. Mass of ethane that was burnt (relative formula mass of ethane = 30)
30. The heat of combustion of charcoal is 360kj/mole. Find the amount of charcoal that will produce 30 kj of energy. (C=12)    ( 1 mk)
31. When 5 grams of propanol (C3H7OH) is burnt in air, 167 kj of heat is produced. Calculate the molar heat of combustion of propanol (H=1), (C = 12) (O=16)
          (2mks)
32. In a class experiment 5.0 of ethanol (CH3CH2OH) was completely burnt and all the heat evolved was used to heat 500cm3 of water from 200C to 800C. Given that the specific heat capacity of water is 4.2j/g/k and the density of water is 1g/cm3.
(i)  Write the equation to show the reaction that takes place when ethanol is burnt        ( 1 mk)
(ii)  Calculate the heat energy observed by water   ( 2 mks)
(iii)  Find the molar heat of combustion of ethanol  ( 1 mk)
 C= 12 (H= 1) (O= 16)
33. When excess iron fillings were placed in 100cm3 of 0.1M copper (II) sulphate solution, this was a temperature rise of 40C. Find the molar heat of reaction. Take specific heat capacity of 4.2j/g/k and density of solution 1.0g/cm3. ( 3 mks)
34. Study the information in the table below and answer the  questions that follows
Bonds C.H CL-Cl C-CL H-CL
Bond energy 444 244 326 431
Calculate the enthalpy change for the reaction   ( 2 mks)
CH4(g) + Cl2(g) → CH3CL(g) + HCl (g)
35. The graph below shows part of temperature – time curve obtained when solid naphthalene was heated.

Explain what happen to the naphthalene molecules along the curve
(a)JK      (b) KI
36. Below is an energy level diagram for the combustion of ethanol. Use it to answer the questions that follows

(i)  State whatever the reaction is endothermic or exothermic. Give your reasons        ( 1 mk)
(ii)  What is the sign of ∆H? Give a reason   ( 1 mk)
37. Study the following  redox reactions
(a)  Mg(s) + Cu2+(aq) → mg2+ + Cu(s) ∆H = 526Kj/mole
(b)  Pb(s) + Cu2+ →Cu(s) + Pb(g)2+ ∆H= -63Kj/mole
Calculate the amount of heat liberated when
(i)  0.25 moles of copper is formed in reaction (a) ( 1 mk)
(ii)  0.5 moles of copper is formed in reaction (b)  ( 1 mk)
38. Given the following values of heat  of combustion, calculate the heat  of formation of ethane (C2H4)
 ∆HC ethane = -1432 Kj/mole
 ∆HC hydrogen = -272kj/mole
 ∆HC carbon = 406 kj/mole
39. The heat of neutralization of a strong acid is usually 57.4 kj/mole, whereas that of a weak acid usually less than 57.4 kj/mole. Explain

 
TOPIC 2
RATE OF REACTION
1. 
The table below gives factors which affect the value of reaction
Zn(s) + 2 HCl(aq) → ZnCl2(aq) + H2 (g)
Complete the table to show how the factors given affect the rate of reaction and give an explanation       ( 2 mks)
Factors Effect on rate Explanation
Using Zinc powder instead of granules  
Heat the reactants  
2. 
The equation below represents two processes that take place without any change in temperature
I.  H2O(s) → H2O (l)
II.  CdCL 2(s) → Cd2+ (l) + 2Cl¬ -(l)
(a)  Explain why although heat is required for each of the process to take place the temperature remains constant in both processes  (1 mk)
(b)  Which of the two has a higher enthalpy change (H)? Give a reason
( 2 mks)
3. The curves below represents the changes in the concentrations of substances “E” and “F” with tie in the reaction

(i)  Which curve represents the change in the concentration of substance F? Give a reason       ( 2 mks)
(ii)  Give a reason for the shapes of the curves after t minutes

The curves shown below were obtained when two equal volumes of hydrogen peroxide of same concentration were allowed to decompose separately in one case, manganese (IV) oxide was added to hydrogen peroxide.
Which curve represents the decomposition of hydrogen peroxide with manganese (IV) oxide? Explain       ( 2 mks)
5. 
State and explain how the rate of reaction between zinc granules and steam can be increased.        ( 2 mks)
6. 
The table below gives three experiments on the reaction of excess sulphuric acid and 0.5g of zinc done under different condition. In each case the volume of gas was recorded at different time internals.
Experiment Term of zinc Conclusion of sulphuric acid
I Powder 0.8m
II Powder 1.0m
III Granules 0.8m
On the same axis draw and label the three curves that could be obtained from such results         ( 3 mks)
7. 
During the production of hydrogen iodide, hydrogen reacts with iodine according to the equation
H2(g) + I2(g)   2 Hl(g) ∆H = + 52.0 kj
Explain how the following would affect the yield of hydrogen iodide
(a)  Increase in temperature     ( 1 mk)
(b)  Decrease in pressure      ( 1 mk)

Ammonia can be converted to nitrogen (II) oxide as shown in the equation below
4NH3 (g) + 5O2   4NO2 + 6H2O (g)

The energy level diagram for the reaches is given above
(a)  Explain how an increase in temperature would affect the yield of Nitrogen (II) oxide       ( 2 mks)
9. 
The decomposition of calcium carbonate can be represented by the equation
CaCo3(s)   CaO(s) + Co2(g)
Explain how an increase in pressure would affect the equilibrium position
         ( 2 mks)
10. 
In the Haber process, the optimum yield of ammonia obtained when a temperature of 4500C a pressure of 200 atmospheres and iron catalyst are used
N2 (g) + 3H2 (g)  2NH3(g) ∆ H = -92kj
(a)  How would the yield of ammonia be affected if the temperature raised to 6000C.        ( 2 mks)
(b)  Give one use of ammonia     ( 1 mk)
11. 
The reaction between a piece of magnesium ribbon with excess 2m hydrochloric acid was investigated at 250C by measuring the volume of hydrogen gas produced as the reaction progressed. The sketch below represents the graph that was obtained

(a)  Name one piece of apparatus that may be used to measure the volume of hydrogen gas produced     ( 1 mk)
(b)  On the same diagram. Sketch the curve that would be obtained if the experiment was repeated at 350C.    ( 2 mks)
12. 
The sketch below shows the rate at which substance “H” is converted to “J Study it and answer the question that follows

Why do the two curves become horizontal after some time  ( 1 mk)
13. 
(a) What conditions is necessary for an equilibrium to be established? ( 1 mk)
(b) When calcium carbonate is heated, the equilibrium shown below is established
CaCO3(s)  CaO(s) + CO2 (g)
How would be the position of the equilibrium be affected if a small amount of dilute potassium hydroxide is added to the equilibrium mixture? Explain ( 2 mks)
14. 
Equal volume of 1m monobasic acids I and “M” were each reacted with excess magnesium turnings. The table below shows the volumes of the gas produced after one minutes.
Acid Volume of gas (cm3)
L 40
M 100
Explain the difference in the volumes of the gas produced  ( 2 mks)
15. 
In a closed system, aqueous iron (III) chloride reacts with hydrogen sulphide gas as shown in the equation below
2FeCL3(aq) + H2S(g)     2FeCl2(aq) + 2HCl(aq) + S (s)
State and explain the observation that would be made if dilute hydrochloric acid is added to the system at equilibrium    ( 2 mks)

A certain mass of a metal E1 reacted with excess dilute hydrochloric acid at 250C. The volume of hydrogen gas liberated was measured after every 30 seconds. The results were presented as shown in the graph below

(a)  Name one piece of apparatus that may have been used to measure the volume of the gas liberated.     ( 1 mk)
(b)  (i)  On the same axis, sketch the curve that would be  obtained if the
experiment was repeated at 350C   ( 1 mk)
(ii)  Explain the shape of your curve in b(i) above ( 1 mk)
17. 
Sodium thiosulphate reacts with dilute hydrochloric acid according to the following equation
S2O3-2 + 2H+ (aq) → H2O (l) + SO2(g) + S(s)
In an experiment to study how the rate of reaction varies with concentration, 10cm3 of 0.4M sodium thiosulphate was mixed with 10 cm3 of 2M hydrochloric acid in a flask. The flask was then placed on white paper marked with a cross (x). The time taken for the cross (x) to become invisible when viewed from above was noted and recorded in the table below. The experiment was repeated three times at the same temperature using the volumes in the table below.
Experiment Volume (in cm3 of 0.4m  thiosulphate Volume of water cm3 Volume of 2mHcl Time in seconds
1 10.0 0 10 16
2 7.5 2.5 10 23
3 5.0 5.0 10 32
4 2.5 7.5 10 72
(a)  (i)  Plot a graph of the volume of thiosulphate (vertical axis) against
time taken for the cross (x)  to become invisible ( 3 mks)
(ii)  From the graph, determine how long it would take for the cross to
become invisible if the experiment was done
I.  Using 6cm3 of the 0.4m thiosulphate solution ( 1 mk)
II.  Using 6cm3 of 0.2 m thiosulphate solution. Explain ( 1 mk)
(b)  (i)  Using the values for experiment 1. Calculate
(I)  Moles of thiosulphate used
(II)  Moles of hydrochloric acid used
(ii) (Which of the time reactants in experiment 1 controlled the rate of
the reaction? Explain
(c)  Give two precautions which should be taken in the experiments above to ensure that constant results are obtained.   ( 2 mks)
18. 
The table below gives the volumes of the gas produced when different volumes of 2m hydrochloric acid were reacted with 0.6g of magnesium powder at room temperature.
Volume of hydrochloric acid Volume of gas cm3
0 0
10 240
20 480
30 600
40 600
50 600
(a)  Write an equation for the reaction between magnesium and hydrochloric acid         (1mk)
(b)  On the grid provided plot a graph of the volume of gas produced (vertical axis) against the volume of acid added. Note that before the reaction produced is directly proportional to the volume of acid added  (3mks)
(c)  From the graph, determine:
(i)  The volume of the gas produced if 12.5 cm3 of 2M hydrochloric acid had been used
(ii)  The volume of 2M hydrochloric acid which reacted completely with 0.6 g of magnesium powder.    (1mk)
(c)  (i)  State and explain the effect on the rate of production of the gas if
0.6g of magnesium ribbon were used instead of magnesium powder        (2mks)
(ii)  3M hydrochloric acid was used instead of 2M hydrochloric acid
 (d)  Given that one mole of the gas occupies 24000 cm3 at room temperature, calculate the relative atomic mass of magnesium   (3mks) 
19. 
In an experiment to study the rate of reaction between duralumin (alloy of aluminium, magnesium and copper) and hydrochloric acid 0.5g of the alloy were reacted with excess 4M hydrochloric acid. The data in the table below were recorded; use it to answer the question that follows:
Time ( minutes) Total volume  of gas cm3)
0 0
1 220
2 410
3 540
4 620
5 640
6 640
7 640
(a)  (i)  From the graph determine the volume of gas  produced at the end
of 2 ½ minutes      ( 1 mk)
(b)  Determine the rate of reaction between 3rd and 4th minutes ( 1 mk)
(c)  Give a reason why some solid remained at the end of the experiment
( 2 mks)
(d)  Given that 2.5m3 of the total volume of the gas was magnesium and aqueous hydrochloric acid, calculate the percentage mass of aluminium present in 0.5g of an alloy. (Al = 27) (H=1)
(e)  State the properties of duralumin that make it more suitable than pure aluminum in aeroplane construction.    ( 2 mks)
20. 
Excess marble chips (calcium carbonate) was put in a beaker containing 100 cm3 of dilute hydrochloric acid. The beaker was then placed on a balance and the total loss in mass recorded after every two minutes as shown in the table below
Time (minutes) 0 2 4 6 8 10
Total loss in mass (g) 0 1.8 2.45 2.95 3.2 3.3
(a)  Why was there less in mass     ( 1 mk)
(b)  Calculate the average rate of loss in mass between
(i))  0 and 2 minutes      ( 1 mk)
(ii)  6 and 8 minutes      ( 1 mk)
(iii)  Explain the difference in the average rates of reaction in (b) (i) and (ii) above        ( 2 mks)
(c)  Write the equation for the reaction which takes place in the beaker  ( 1 mk)
(d)  State three ways in which the rate of the reaction above could be increased.       ( 3 mks)
(e)  The solution in the beaker was evaporated to dryness explain what would happen if the beaker and its contacts were left in the laboratory overnight
         ( 2 mks)
(f)  Finally some water was added to the contents of the beaker when aqueous sodium sulphate was added to the content of the beaker a white precipitate was formed
(i)  State one use of the substances identified in (f)(i) above
21. 
The table below shows the volumes of nitrogen (IV) oxide gas produced when different volumes of 1m nitric acid were each reacted with 2.07g of lead at room temperature.
Volume of 1m nitric acid Volume of nitrogen (IV) oxide gas cm3
5 60
15 180
25 300
35 450
45 420
55 480
(a)  Give a reason why nitric acid is not used to prepare hydrogen gas ( 1 mk)
(b)  Explain how the rate of reaction between lead and nitric acid would be affected if the affected if the temperature of the reaction mixture is raised         ( 2 mks)
(c)  On the grid provided below plot a graph of the volume of the gas produced vertical axis against the volume of acid   ( 3 mks)
(d)  Using the graph, determine the volume of
(i)  Nitrogen (IV) oxide produced when 30 cm3 of 1M nitric acid were acted with 2.07g of lead
(ii)   1M nitric acid which would react completely with 2.07g of lead
(1mk)
(e)   Using the answer in d (ii) above determine
(i)  The volume of 1 m nitric acid that would react with one mole of Pb/lead (Pb = 207)     ( 2 mks)
(ii)  The volume of nitrogen (IV) oxide gas produced when one mole of lead reacts with excess 1M nitric at room temperature
(f)  Calculate the number of moles of
(i)  1M Nitric acid that reacted with one mole of lead
(ii)  Nitrogen (IV) oxide produced when one molar of lead were reacted with excess nitric acid.  Molar gas volume = 24000 cm3 ( 1 mk)
(iii)  Using the answer in (21) (i) and (ii) above write the equation for the reaction  between  lead and nitric acid given that one mole of lead nitrate and two moles  of water were  also produced. ( 1 mk)
22. 
(a)  Methanol is manufactured from carbon (IV) oxide and hydrogen gas according to the equation
CO2(g) + 3H2     CH3OH(g) + H2O(g)
The reaction is carried out in the presence of a chromium catalyst at 700K and 30kpa. Under these conditions an equilibrium is reached when 2% of the carbon (IV) oxide is converted to methanol
(i)  How does the rate of the forward reaction compare with that of the reverse reaction when 2% of the carbon (IV) oxide is converted to methanol?          ( 1mk)
(ii)  Explain how each of the following would affect the yield of methanol
• Reduction in pressure      ( 2 mks)
• Using a more efficient catalyst      (2 mks)
(iii)  If the reaction is carried out at 500K and 30 kpa, the percentage of carbon     (IV) oxide converted to methanol is higher than 2%
(I) What is the sign of ∆ H for the reaction? Give a reason (2mks)
(II)  Explain why in practice the reaction is carried out at 700J but NOT at 500K        (1 mk)
(b)  Hydrogen peroxide decomposes according to the following equation
2H2O2(aq) → 2 H2O(l) + O2(g)
In an experiment, the rate of decomposition of hydrogen peroxide was found to be 6.0 x 10-8 mol dm-3S-1
(i)  Calculate the number of moles per dm3 of hydrogen peroxide that had decomposed within the first 2 minutes   ( 2 mks)
(ii)  In another experiment the rate of decomposition was found to be 1.8 x 10-7 mol dm-3S-1. The difference in the two rates could have been caused by addition of a catalyst, State giving reasons one other factor that may have caused the difference in the two rates of decomposition. ( 2 mks)
23. 
(a)  (i)  State the Le chatelier’s principle   ( 1 mk)
 (ii)  Carbon (II) oxide gas reacts with steam according to the reaction;
  CO(g) + H2O(g)    H2(g) + C02(g)
What would be the effect of increasing the pressure of the system at equilibrium? Explain    ( 2 mks)
(iii)  When the reaction in (ii) above was carried out at lower temperature, the yields of hydrogen and carbon (IV) oxide increased. What is the sign of ∆H for the reaction? Explain (2mks)
(b)  The table below gives the volume of oxygen gas produced at different times when hydrogen peroxide decomposed in the presence of a catalyst.
Time (sec) 0 10 20 30 40 50 60
Volume of oxygen (cm3) 0 66 98 110 119 120 120
(i)  Name the catalyst used for this reaction   ( 1 mk)
(ii)  On the grid provided. Draw the graph of volume of oxygen gas produced (vertical axis) against time.
(iii)  Using the graph determine the rate of decomposition of hydrogen peroxide after 24 seconds    ( 2 mks)
(iv)  Give a reason why the total volume of oxygen gas produced after 50 seconds remains constant    ( 1 mk)
24. Define the term rate of reactions      (1 mk)
25. State two methods used to measure rate of reactions    (2mks)
26. When a metal oxide of element “w” react  with hydrogen, the equation for the equation for the reaction is
WO3(s) + 3 H2(g) → W(s) + 3H2 (I)O
Comment on the reactivity of element “W” with respect to hydrogen (2mks)
27. 7.5g of calcium carbonate was placed in a conical flask containing 50cm3 of dilute hydrochloric acid. The flask kept at constant temperature and the volume of carbon (IV) oxide gas evolved was measured at 20 minutes intervals.
Not all the calcium carbonate was used up during the reaction the results were recorded in the table below
Time from start of reaction ( minutes) Volume of Co2 evolved cm3
0 0
20 555
40 810
60 695
80 1000
120 1020
(a)  Write an equation for the reaction between calcium carbonate and hydrochloric acid      ( 1 mk)
(b)  Plot a graph volume of carbon (IV) oxide produced against time (minutes)
         ( 3 mks)
(c)  What volume of carbon (IV) oxide were evolved during the 20th minutes intervals (20- 40) minutes     ( 1 mk)
(d)  Why was there no increase in volume of the gas evolved after 100 minutes?       ( 1 mk)
(e)  Calculate the mass of 11. 2 cm3 of carbon (IV) oxide gas evolved at stp: molar gas volume = 22.4 dm3
(f) Determine the mass of calcium carbonate which had reacted after 120 minutes
         ( 1 mk)
28. Consider the equilibrium  reaction  below
2SO4 (g) + O2 (g)                    2SO3 (g) ∆H = -ve
Which of the following will increase the yield of sulphur (vi) oxide
• Addition of catalyst
• Increase in pressure
• Increase in temperature
• Doubling the volume of the system     ( 1 mk)
29. (a)  Why does the rate of reaction
(i)  Increase with increase in temperature    ( 1 mk)
(ii)  Increase with use of a suitable catalyst   ( 1 mk)
(b)  The equation for gaseous reaction is
A(g) + 2B (g) → C(g) + D (g)
State the effect of the following on rate of reaction
(i)  The pressure of “B” is doubled but of A is the same  ( 1 mk)
(ii)  The amount pressure of both A and B are doubled  ( 1 mk)
(iii) The amount of A and B remain unchanged but an inert gas is added to double the over all pressure
30. Below is a graph  of the volume  of oxygen collected (cm3) against time when  powdered  and lump  of manganese (IV) oxide were used to decompose hydrogen peroxide

Which one of the curves correspond to the results obtained by using powdered manganese (IV) oxide. Give reasons     ( 2 mks)
31. Consider the following reaction
nCH2 = CH2(g)                  (CH2 – CH2-)n ∆H = -ve
What conditions favours the process     ( 2 mks)
32. Consider the reaction
A(g) + B)g)    AB(g) ∆H = + ve
Draw, an energy level diagram for this reaction, when un-catalyzed and when catalyzed        ( 2 mks)
33. For the following gaseous reaction
E(g) + F (g)       G(g) + H(g) ∆H = +ve
What is the effect on the rate if the?
(a)  Volume of the total reactants is doubled   ( 1 mk)
(b)  Temperature is doubled     ( 1 mk)
 
TOPIC 3
ELECTROCHEMISTRY 1 AND 2
1. 
A student set up an experiment as shown in the diagram below
(a)  Draw an arrow on the diagram to indicate the direction of the electron flow. Explain your answer     ( 2 mks)
(b)  What would be observed on the voltmeter (v) if both rods were Zinc rods?
2. 
Write an equation for the process that takes place at the anode during electrolysis of aqueous sodium sulphate solution using platinum electrodes ( 1mk)
3. 
3.8g of metal M were deposited when a molten salt of M was electrolyzed by passing a current of 0.6 amps for 90 minutes. Relative atomic mass of M= 226:
1. Faraday = 96500 coulombs)
(a)  Calculate the amount of electricity in coulomb
(i)  Needed to deposit 3.8g of metal M    (1mk)
(ii)  Needed to deposit 3.8g of metal M    (1mk)
(iii)  Deduced the charge on the ion of M    (1 mk)
4. 
Study the set up below and answer the questions that flows

State and explain the observations that would be made when the circuit is completed       (3 mks)
5. 
Explain the following observation
A chloride dissolves in water to form an electrolyte while the same chloride dissolves in methyl benzene to form non- electrolyte ( 2 mks)
6. 
Explain why it is not advisable to use aqueous sodium chloride solution as the salt bridge in electrochemical cell formed between half cells.
Pb2+ / Pbq+ = 0.13V and Cu2+/ Cu Ece = 0.34v  ( 2 mks)
7. 
Aqueous potassium sulphate was electrolyzed using platinum electrodes in a cell
(a)  Name the products formed at the cathode and anode ( 1 mk)
(b)  How does the concentration of electrolyte change during electrolysis?
(c)  Why would it not be advisable to electrolyte aqueous potassium sulphate using metal electrodes?
8. 
Use the information below to answer the questions that follows
Zn2+ + 2e → Zn (s) – 0.76
Al3+ + 3e Al   -1.66
Fe2+ + 2e → Fe(s) – 0.44
(a)  Calculate the EQ value for the electrochemical cell represented below Al(s)/Al3+ //Fe2+ (aq) / Fe + (s)    ( 1 mk)
(b)  Give a reason why aluminium metal would protect iron from rusting better than zinc metal      ( 1 mk)
9. 
The set up below was used to electrolyze aqueous copper (II) sulphate
(a)  Explain why the bulb light brightly at the beginning of the experiment and become dim after sometime.     ( 2 mks)
(b)  Write an ionic equation for the reaction that took place at the cathode
         ( 1 mk)
10. 
Use the cell representation below to answer the questions that follow
(a)  Write the equation for the cell reaction   ( 1 mk)
(b)  If the e.m.f of the cell is + 0.30v and Eq value of Fe(a)2+ / Fe(s) is 0.44V. Calculate the EQ value of Cr(a)3+ /CV(s)   ( 2 mks)
11. 
When amount of 1.5 amperes was passed through a cell containing M3+ ions of metal M for minutes the mass of the cathode increased by 0.26g. (Faraday = 96500 coulombs)
(a)  Calculate the quantity of electricity used   ( 1 mk)
(b)  Determine the relative atomic mass of metal “m”  ( 2 mks)
12. 
An element “P” has a relative atomic mass of 88. When a current of “P” for 32 minutes and 10 seconds 0.44g of “p” were deposited at the cathode.
Determine the change on an ion of “p”
13. 
The set up below was used to electroplate a metallic spoon. Study it and answer the question that follows
(a) Write an ionic equation for the reaction that occurred at the cathode ( 1 mk)
(b) State and explain what happen to the anode    ( 1 mk)

14. 
During purification of copper by electrolysis 1.48g of copper were deposited when a current was passed through aqueous for 2 ½ hrs   
Calculate the amount of current that was passed   ( 3 mks)
(CU = 63:5) (1 Faraday = 96,500 Coulombs)
15. 
A strip of metal “Q” was dipped into a solution of copper (II) sulphate and allowed to stand overnight. Given that
Cu2+ + 2s → CuEQ = + 0.34V
Q2+ (aq) + 2 e → Q(s) EQ = -0.13
(a) State the observations which were made    ( 2 mks)
(b) Give a reason for your answer in 19 (i) above   (2 mks)
16. 
The diagram below represent the set ups that were used to a study the effect of an electric current on pure water and dilute sulphuric acid.

State and explain the observations made when each experiment was started 
         ( 3 mks
17. 
In an experiment to investigate the conductivity of substance a student used the set up shown below

Student noted that the bulb did not light
(a)  What had been omitted in the set up    ( 1 mk)
(b)  Explain why the bulb  light when the omission is occurred ( 2 mks)
18. 
When a current of 0.82A was passed for 5 hours through an aqueous solution of metal “Z” 2.65g of metal were deposited
Determine the change on the ions of metal (A faraday = 96,500 coulomb) relative atomic mass of Z = 52       ( 2 mks)
19. 
Study the standard reduction potential given below and answer the questions that follow. The letters are not actual symbols of the elements
Actual symbols of elements   EQ values
M (aq) 2+ + 2e → M(s)    -0.76V
M (aq) 2+ + 2e →M(s)    - 2.36V
P (aq) + + e →P(s)    -0.80V
Q (aq)2+ + 2e → Co + (s)→   - 0.14V
(a)  The standard reduction potential for Fe2+ is 0.44V. Select the element which would best protect iron from rusting    ( 1 mk)
(b)  Calculate the EQ value for the cell M(s) / M2+ (a) // P+(aq) P(s)  (2mks)

20. 
(a)  Use the information given bellow to draw a labeled diagram of an electrochemical cell that can be constructed to measure the electromotive force between G and J.
G2+ (aq) + +2e → G (s); E0 =- 0.74v              (2mks)
J2+ (aq) + +2e   → J (s); E0 = - 0.14v
(b)  Calculate the E0 value for the cell constructed in (a) above.   (1mk)
21. 
(a)  When brine is electrolyzed using inert electrodes, chlorine gas is liberated at the anode instead of oxygen. Explain this observation.     (2mks)
(b)  Name the product formed at the cathode.                                       (1mk)
22. 
During the electrolysis of aqueous silver nitrate, a current of 0.5A was passed through the electrolyte for 3 hours.
(a)  Write the equation for the reaction which took place at the anode.  (1mk)
(b)  Calculate the mass of silver deposited (Ag = 108; 1F = 965000)       (2mks)
23. 
(a)  The following are half-cell reaction and their reduction potentials,
                                                                  E0 (Volts)
Zn2+ (aq) + 2e- → Zn(s)                       -0.76
Pb2+ (aq) + 2e -→Pb(s)                           -0.13
Ag+(aq) + e- → Ag(s)                             +0.80
Cu2+ (aq) + 2e-→ Cu(s) _                                      +0.30
(b)  Write the cell representation for the electrochemical cell that would give the highest Eθ                                                          (1mk)
(c)  State and explain the observations made when a copper rod is placed in a beaker containing silver nitrate solution.                                    (2mks)
24. 
The diagram below represents an experiment that was set up to investigate movement of ions during electrolysis.

When the circuit was completed, it was noticed that a blue colour spread towards the right
(a)  Explain this observation      (2mks)
(b)  Write the equation for the reaction that occurred at the anode (1 mk)
25. 
(a)  The table below gives reduction potentials obtained when the half cells for each of the metals represented by J, K, L, M and N were connected to a copper half of cells as the reference electrodes.
Metals Reduction potential
(vol/s)
J -1.10
K -0.47
L -0.00
M + 0.45
N 1.16
(i)  What is the metal “L” likely to be? Give a reason  ( 1 mk)
(ii)  Which of the metals cannot be displaced from solution of its salt by any other metal in the table give a reason   (2mks)

(iii)  Metal “K” and “M” were connected to form a cell as shown the diagram
below

(i)  Write the equation for the half cell reaction that occur at metal K electrode       ( 1 mk)
(ii)  If the slat bridge is filled with saturated sodium nitrate solution, how does it help to complete the circuit   (2mks)
(b)  When electric current is passed through copper (II) sulphate solution for several hours as shown in the diagram, a gas that relights a glowing splint is produced at electrode “C”

(i)  Which of the electrode is the cathode? Give a reason (2mks)
(ii)  Write an equation for the formation of the gas at electrode “D” 
(iii)  State and explain the observations that would be made
 I.  At electrode “D”     (1 mk)
 II.  In the copper (II) sulphate solution   (1 mk)
26. 
The extraction of aluminium from its ore takes place in two stages, purification stage and electrolysis stage. Below shows the set up for the electrolysis stage
(a)  (i)  Name the ore from which aluminum is extracted  (1 mk)
 (ii)  Name one impurity which is removed at the purification stage           (1 mk)
(b)  (i)  Label  on the diagram each of the following
 I.  Anode
 II.  Cathode
 III.  Region containing electrolyte
(ii)  The melting point of aluminium oxide is 20540C, but the electrolysis is carried out at between 800 C and 9000C
I.  Why is not carried out at 20500C   (2mks)
II.  What is done to lower the temperature  (1 mk)
(iii)  The aluminium which is produced is tapped off as a liquid. What does this suggest about its melting point?
(c)  A typical electrolysis cell uses a current of 40,000 amperes. Calculate the mass (in kg) of aluminium produced in one hour (Al = 27) Faraday = 96,500 coulombs)       (3mks)
27. 
Use the standard electrode potential for A, B, C, D and F given below to answer the questions that follows. The letters do not represent the actual symbols of the elements
    EQ volts
 A (aq) 2+ + 2e → A(s)  - 2.90V
 B (aq) 2+ 2e →B(s)  - 2.38V
 C (aq) + + e→ ½ C2  - 0.00V
 D (aq) +2 + 2e → D(s)  + 0.34V
 ½ Fe2 + e → F (aq)  + 2.87V
(i)  Which element is likely to be hydrogen? Give a reason for your answer
(2mks)
(ii)  What is EQ value for the strongest reducing agent?   (1 mk)
(iii)  In the space provide, draw a labeled diagram of the electrochemical cell that would be obtained when a half cells of element “B” and “D are combined        (3mks)
(iv)  Calculate the EQ value of the electrochemical cell constructed in (iii) above         (1 mk)
28. 
 The diagram below shows the extraction of sodium metal using the down cell
 Study it and answer the questions that follows
(i)  Explain why in this process the sodium chloride is mixed with calcium chloride       ( 2 mks)
(ii)  Why is the anode made of graphite and not steel?  ( 1 mk)
(iii) State two properties of sodium metal that make it possible for it to be collected as shown in the diagram    ( 2 mks)
(iv)  What is the function of steel gauze cylinder?   ( 1 mk
(v)  Write ionic equation for the reactions which take place at
 I.  Cathode      ( 1 mk)
 II.  Anode       ( 1 mk)
(vi)   Give one industrial use of sodium metal   ( 1 mk)
29. 
The set up below was used during the electrolysis of aqueous magnesium sulphate using inert electrodes

(i)  Name a suitable pair of electrodes for this experiment  (1 mk)
(ii)  Identify the anions and cations present in the solution  (2mks)
(iii)  On the diagram label the cathode     (1 mk)
(iv)  Write ionic equation for the reaction that took place at the
I:  Anode        (1 mk)
II.  Cathode       (1 mk)
30. 
(a)  The diagram below represents a mercury cathode cell that can be used in the industrial manufacture of sodium hydroxide. Study it and answer the question that follows

i.  Name the
 I:  Raw material introduced at “2”    (2mks)
 II.  Another substance, that can be used in the call instead of graphite           (1mk)
ii.  Identify the by products that come out at I    (1 mk)
iii.  Give
 1.  One use of sodium hydroxide     (1 mk)
 2.  Two reasons why mercury recycled    (1 mk)
(b)  A current of 1000 amperes was passed through the cell for five (5) hours
i.  Write equation for
 I.  The reaction that occurred at the mercury cathode  (1 mk)
 II.  The reaction in which sodium hydroxide was produced (1 mk)
ii.  Calculate the mass of sodium hydroxide that was produced (Na= 23) (O = 16) (H=1.0) Faraday = 96500 coulombs    (4mks)
31. 
(a) Study the standard electrode potentials for the half cells given below and answer the questions that follows. The letters do not represent the actual symbols of the elements
     E volts
N+(av) + e- → N  -2.92
J+(av) + e → J  +0.52
K+(aq) + e → ½ Kg  0.00
½ G(g) + e → G¬-(ag)  +1.36
M2+ (g) + 2 e →m(g)  -0.44
i.  Identify the strongest oxidizing agent: Give a reason for your answer
ii.  Which two half cells would produce the highest potential differences when combined?        (1 mk)
iii.  Explain whether the reaction represents below can take place (2mks)
 2M- + N → 2N + M2+
 (av) (s)   (s)      (aq)
(b)  100 cm3 of 2m sulphuric acid was electrolyzed using the set up represented by the diagram below
i.  Write an equation for the reaction that produce gas “L” ( 1 mk)
ii.  Describe how gas “k” can be identified   ( 1 mk)
iii.  Explain the differences in
(a)  The volume of gases produced at the electrodes
(b)  Brightness of the bulb if 100 cm3 of 2m ethanoic acid was used in place of sulphic acid    (2mks)

32. 
The table below gives the standard electrode potentials for the metals represented by letters D, E, F & G. study it and answer the questions that follows
Metals Standard electrical potential (volts)
D -0.13
E + 0.85
F + 0. 34
G - 0. 76
(a)  Which metal can be displaced from a solution of its salt by all the other metals in the table? Give a reason
 (b)  Metal “F” and “G” was connected to form a cell as shown in the diagram

 i.  Write the equation for the reactions that occur at the electrode F
and G
ii.  On the diagram indicate with an arrow the direction in which electrons would flow
iii.  What is the function of the salt bridge?   (1 mk)
 (c)  An electric current was passed though concentrated solution of copper (ii) chloride as shown in the diagram below.

i.  Explain the observation that would be made on the electrolyte as the experiment progress     (2mks)
ii.  After sometime test tube “H” was found to contain a mixture of two gases. Explain this observation     (3mks)
iii.  Which of the electrodes is the anode? Explain  (2mks)
33. 
The diagram below is a cross- section of a dry cell. Study it and answer the questions that follows

 (i)  On the diagram, show with a (+ve) sign the +ve (positive terminal) (1 mk)
 (ii)  Write the equation for the reaction in which electrons are produced (1 mk)
(iii)  The zinc can is line with ammonium chloride and zinc chloride paste. What would happen if the mixture was to become dry? Give a reason 
( 2 mks)
 (iv)  Give one advantage and one disadvantage of dry cell ( 2 mks)
(b)  The setup up below was used to electrolyze molten lead (ii) Iodide
i. State the observation that was made at the anode during the electrolysis. Give a reason for your answer.
ii.  A current of 0.5A was passed for two hours. Calculate the mass of load that was deposited (Pb= 207) (1 faraday = 96500c)   (3mks)
34. 
(a)  Brine usually contains soluble calcium and magnesium salts. Explain how sodium carbonate is used to purify brine    (2mks)
 
(b)  The diagram below represents a diagram cell used to electrolyte pure brim
i.  Write the equations for the reactions that take place at   (2mks)
I.  Cathode
II.  Anode
ii.  Name
I.  Products U:       (1 mk)
II.  Another material that can be used instead of titanium (1 mk)
III.  The impurity present in the product U   (1 mk)
iii.  State two functions of the porous diaphragm    (2mks)
(c)  Give one industrial use of the product “U”    (1 mk)
35. 
(a) The equations below shows the standard reduction potential for four half cell. Study it and answer the questions that follows. Letters are not actual symbols of the element.
       EQ Volts
F2 (aq) + 2e-  →2F-(av)  + 0.54
G2+ + 2e  →G(s)   -0.44
H+2 (aq) + 2 e   → H(s)  + 0.34
2J+ + 2eJ2(g)  →J2(g)   0.00
 i.  Identify the strongest reducing agent    (1 mk)
ii.  Write the equation for the reaction which takes place when solid “G” is added to a solution containing H2+ (ions)  (2mks)
iii.  Calculate the EQ value for the reaction in (ii) above  (1 mk)
 
(b)  The diagram below shows the apparatus used to electrolyze acidified water to obtain hydrogen and oxygen gases. Study it and answer the questions that follows?

i.   Identify the electrode at which oxidation takes place  (1 mk)
ii.  Give a reason why it is necessary to acidify the water (1 mk)
iii.  Explain why hydrochloric acid is not used to acidify the water          (2mks)
(c)  During electrolysis of aqueous copper (II) sulphate 144750 columbus of electricity were used. Calculate the mass of copper metal that was obtained (CU= 64) (1 Faraday = 96500 Columbus)   (3mks)

36. 
(a)  Below is a simplified diagram of the down’s cell used for the manufacture of sodium. Study it and answer the questions that follows.

i.  What material is the anode made of? Give a reason  (2mks)
ii.  What precaution is taken to prevent chlorine and sodium from re- combining?       (1 mk)
iii.  Write an ionic equation for the reaction in which chlorine gas is formed        (1 mk)
(b)  In the down’s cell (Used for manufacture of sodium) a certain salt is added to lower the melting point of sodium chloride from about 8000C to 6000C
i. Name the salt that is added
ii. State why is necessary to lower the temperature  (1 mk)
(c) Explain why aqueous sodium chloride is not suitable as an electrolyte for the manufacture of sodium in the down’s cell- process  (2mks)
(d)  Sodium metal reacts with air to form two oxides. Give the formula of the two oxides        (2mks)
(e)  State two uses of sodium metal     (2mks)
37. 
 (a)  What is an electrolyte       (1 mk)
 (b)  State how the following substances conduct electricity
 (i)  Molten calcium chloride
 (ii)  Graphite
(c)  The diagram below shows a set up that was used to electrolyze aqueous magnesium sulphate
(i)  On the diagram above, using an arrow, show the direction of the flow of electrons      (1 mk)
(ii)  Identify the syringe which hydrogen gas would be collected. Explain       (1 mk)
(d)  Explain why the concentration of magnesium sulphate was found to have increased at the end of the experiment.    (2mks)
(e)  During electrolysis a current of 0. 72A was passed through the electrolyte for 15 minutes. Calculate the volume of gas produced at the anode. I Faraday = 96500 Columbus. Molar gas volume is 24000 at room temperature        (4mks)

The diagram below represents asset up that can be used to electrolyze aqueous copper (II) sulphate
(a)  (i)  Describe how oxygen gas  is produced during the electrolysis           (2mks)
(ii)  Explain why copper electrodes are not suitable for this electrolysis
(2mks)
 (b)  Impure copper is purified by an electrolytic process
 (i)  Name one ore from which copper is obtained  (1 mk)
(ii)  Write the equation for the reaction that occurs at the cathode during the purification of copper    (1 mk)
(iii)  In an experiment to electroplate a copper spoon with silver, a current of 0.5A was passed for 18 minutes. Calculate the amount of silver deposited on the spoon.  (1F = 96500 coulombs, Ag) = 108)
(iv)  Give two reasons why some metals are electroplated (2mks)
39. The following tables give the standard electrode potential for a number of half reactions.
      E-volts
 Mg2+ (aq) + 2e →mg(s)   -2.3
          
 Mn2+ (aq) + 2e →mn (s)   -1.18
        
 Cd2+ (aq) + 2e → Cd   -0.402
 
 2H+ (aq) + 2e → H2   0.00
 
 Ag+ (aq) + e →Ag (s)    +0.799
     
 Ce+4 + E → CC3+   +1.61
(a)  Which one of the substance is the strongest oxidizing agent  (1 mk)
(b)  Which one if the substance is the strongest reducing agent  (1 mk)
(c)  Select one of the substances from the table that could be used to oxidize silver ions and write the equation for the reaction.   (2mks)
(d)   Given the two half reactions
 Cd2+ (aq) + 2e- →Cd(s)
 
 Mg2+ (aq) + 2e → Mg(s)
 
 (i)  Write the cell representation made up of these two half reactions (2mks)
(ii)  Write down the over all call reaction for the cell formed by these two half reactions      (2mks)
 (iii)  Calculate the EQ value of this cell    (2mks)
40. The diagram below shows a setup used to pass to electric current on molten lead (ii) bromide

(a)  (i)  What does the bulb show before the solid lead bromide is heated?           (1 mk)
 (ii)  Give a reason for your answer    (1 mk)
(b) Why was lead (ii) bromide in the molten state?   (1 mk)
(c) What observation is made at the cathode and anode respectively (2mks)
(d) Write equations for the reactions at both electrodes   (2mk)
41.  (i)  If the same arrangement was used to electrolyze aqueous potassium
iodide. Iodine vapor would be collected at the anode and hydrogen gas at the cathode instead of potassium. Explain why   (2mks)
(ii)  In an experiment chromium (iii) chloride is electrolyzed using the chromium electrodes. A current of 0.2A flows for 5788 seconds. The increase in mass of the electrode is 0.208g. Calculate the charge on the electrons. (Cr = 52) / Faraday = 96500C    (3mks)
42. Consider the call
 Mn(s) / Mn2+ (aq) // Cd2+ (aq) /Cd(s)
 Eq for the manganese electrode is – 0.40v calculate the e.m.f of the cell (1 mk)
43. Write cell reaction for the following electrochemical
ZN(s) /Znz+// Fe3+(aq)/Fe/pt      ( 2 mks)
44. Given the following standard electrode potential, Eq = -0.76V
 Zn2+(aq) + 2e →Zn(s) EQ = 0.76v
 Cl2(g) + 2e →2Cl-(aq) EQ = + 1.36V
 Calculate the EQ value for the cell     ( 1 mk)
45. Two incomplete half cells are given below

(a)  Complete the diagram to show how the two half cells are connected to give an electrochemical cell.      (2mks)
(b)  Using arrows show the direction of the electron flow  (1 mk)
(c)  Indicate the direction of current flow
(d)  Write the equation for the half cell/ reaction taking place at the electrodes
          (2mks)
(e)  Write the overall cell reaction     
(f)  How many moles of electrons are transferred?
(g)  Calculate the electronic charge transferred during reaction (F= 96,500 coulombs)        (1 mk)
46. Magnesium reacts rapidly with copper (II) ions as follows
 Mg(s) + Cu2+(aq) →Mg2+ + Cu (s)
 Give the half reaction for this reaction     (1 mk)
47. (a) Explain the changes that takes place in solution and at the electrodes in the
electrolysis of
 (i)  Aqueous Sodium sulphate with invert electrodes  (2mks)
 (ii)  Concentrated Sodium Chloride with carbon anode and mercury
cathode       (2mks)
(b)  Two electrolytic cells for solutions in a (i) and (ii) respectively were connected in series. A current of 1.5 A was passed for 600 seconds. The first cell contained aqueous copper (II) sulphate and had copper electrodes. The anode showed a loss in mass of 0.296 g but there was no change in the appearance of the electrolyte. The sodium chloride with little sodium hydroxide had copper electrodes and a reddish brown precipitate formed.
(i)  Why was there no change in the appearance of the electrolyte in the first cell
(ii)  Why was a small amount of sodium hydroxide added to aqueous sodium chloride in the second cell?
(iii)  Name the reddish- brown precipitate formed   (1 mk)
(iv)  Write an ionic equation for the formation of substance in (iii)
(v)  Calculate the value of Faraday constant   (1 mk)
48. Given that the standard electrode potential EQ are
 Mg2+(aq) + 2e →Mg(s) EQ = -2.38V
 Cl2(g) + 2e →2Cl(aq)¬ EQ = + 1.36V
Find the e.m.f of the cell
 
TOPIC 4
METALS
1. 
When magnesium metal is burnt in air it reacts with both oxygen and Nitrogen gas giving a white ash like substance. Write two equations for the two reactions that takes place.
2.
When excess Carbon (II) Oxide is passed over lead oxide in a combustion tube, lead (II) oxide is reduced.
 (a)  Write an equation for the reaction which took place   ( 1 mk)
(b)  What observations was made in the combustion tube when the reaction was complete        ( 1 mk)
(c)  Name another gas which would be used to reduce lead (II) oxide  ( 1 mk)
3. 
When the oxide of element “H” was heated with powdered carbon, the mixture glowed and carbon (IV) oxide gas was formed. When the experiment was repeated using oxide of “J” there was no apparent reaction
 (a)  Suggest one method that can be used to extract element J from its oxide
           ( 1 mk)
 (b)  Arrange element H, J and carbon in the order of their decreasing reactivity
           ( 1 mk)
4.
 Study the flow chart below and answer the question that follows

 (a)  State the conditions necessary for the reaction in step 2 to occur  ( 1 mk)
 (b)  Name
 (i)  Gas P        ( 1 mk)
 (ii)  One use of Zinc      ( 1 mk)

5. 
 The set up below was used to obtain a sample of iron

 Write two equations for the reactions which occur in the combustion tube
          ( 2 mks)
6. 
Dry carbon (II) oxide gas react with heated lead (II) oxide as shown in the equation below
 (a)  Name the process undergone by the lead (II) Oxide  ( 2 mks)
 (b)  Give a reason for your answer (a) above
(c)  Name another gas that can be used to perform the same function as carbon gas in the above reaction     ( 1 mk)
7. 
In the industrial extraction of lead metal, the ore is first roasted in a furnace. The solid mixture obtained is then fed into another furnace together with coke limestone and scrape iron. State the functions of each of the following in this process.
  (a)  Coke        ( 1 mk)
 (b)  Scrape iron       ( 1 mk)
 (c)  Limestone       ( 1 mk)
8. 
 Study the flowchart and answer the questions that follows
 Identify
 (a)  Solution K
 (b)  Solid
 (c)  Gas M
9. 
 The flow chart below shows steps used in the extraction of zinc form one of its
Ores.

(a)  Name the process that is used in step 2 to concentrate the ore. ( 1 mk)
(b)  Write an equation for the reaction which takes place in step 3 ( 1 mk)
 (c)  Name one use of zinc other than galvanizing    ( 1 mk)
10. 
During the extraction of aluminium from its ores; the ore is first purified to obtain alumina. The flow chart below shows the stages in the extraction of aluminium from alumina.
 (a)  Name
 (i)  Substance C1       ( 1 mk)
 (ii)  Process D1       ( 1 mk)
(b)  Give two reasons why aluminium is used extensively in making of cooking pans        ( 1 mk)
11. 
The flow chart below outlines some of the process involved in extraction of copper from pyrites. Study it and answer the questions that follows
 

 (a)
 (i)  Name gas “k”
 (ii)  Write an equation for the reaction that take place in the 1st roasting furnace
           ( 1 mk)
 (iii)  Write the formula of the cations present in the slag “M”
 (iv)  Identify gas “P”
(v)  What name is given to the reaction that take place in chamber N. Give a reason for your answer?
(b)  The copper obtained “M” is not pure. Draw a labeled diagram to show the set up you would use to refine the copper by electrolysis.  (2mks)
(c)  Given that the mass of copper obtained from the above extraction was 210 kg. Determine percentage purity of the ore (Copper pyrite) if 810 kg of it was fed to the 1st roasting furnace     (4mks)
  (Cu= 63.5) (Fe = 56) (S= 32)
(d) Give two effects that this process could have on the environment (2mks)
12. 
The flow chart below illustrates the industrial extraction of lead metal. Study it and answer the questions that follows
 (a)  (i)  Name the ore that is commonly – used in this process ( 1 mk)
  (ii)  Explain what take place in the roasting furnace  ( 1 mk)
  (iii)  Identify gas “p”      ( 1 mk)
(iv)  Write the equation for the main reaction that takes place in the smelting furnace      ( 1 mk)
(v)   Give two environmental hazards likely to be associated with extraction of lead
(vi)  What is the purpose of adding iron in the smelting furnace? (1 mk)
 (b) Explain why hard water flowing in lead pipes may be safer for drinking them soft water flowing in the same pipes    (3mks) 
(c) State one use of lead other than making lead pipes   (1 mk)
13.
 The raw material for extraction of aluminum is bauxite.
 (a)  Name the method that is used to extract aluminium from bauxite (1 mk)
 (b)  Write the chemical formula for the major components of bauxite (1 mk)
 (c)  (i)  Name the major impurities sin bauxite   (3mks)
  (ii)  Explain how the impurities in bauxite are removed  (3mks)
(d)  Crayolite is used in the extraction of aluminium from bauxite. State its function        (1 mk)
(e)  Describe how carbon (IV) oxide is formed during the extraction of aluminium        (2mks)
(f)  Aluminum is a reactive metal yet utensils made from aluminium do not corrode easily. Explain this observation

14.
The extraction of iron from its ore takes place in the blast furnace. Below is a simplified diagram of a blast furnace. Study it and answer the questions that follow.
(a) (i)  One of the substances in the slag    (1 mk)
(ii)  Another iron ore material used in the blast furnace  (1 mk)
 (One gas which is recycled)     (1 mk)
(b)  Describe the process which leads to the formation of iron in the blast furnace
(c)  State the purpose of limestone in the blast furnace    (1 mk)
(d)  Give a reason why the melting point of iron obtained from the blast furnace is 12000 while that of pure iron is 15350C   (1 mk)
(e)  State two uses of steel       (2mks)
15. The flow chart below shows a sequence of chemical reactions starting with copper. Study it and answer the questions that follow
 (a)  In step 1, excess 3M nitric acid was added to 0.5 of copper powder
 (i)  State two observations which were made when the reaction was in
progress       (2mks)
  (ii)  Explain why dilute hydrochloric acid cannot be used in step 1            (1 mk)
  (iii)  I.  Write the equation for the reaction that took place in step 1            (1mk)
II.  Calculate the  volume of 3M nitric acid that was needed  to react completely  with 0.5g  of copper  powder (Cu= 63.5)        (3mks)
(b)  Give names if the type of reactions that took place in steps 4 and 5 (1 mk)
(c)  Apart from the good conductivity of electricity, state two other properties that make it possible for copper to be extensively used in the electrical industry        (2mks)
16. Study the flow chart below and answer the questions that follow
(i)  Suggest a purpose for the industry process represented by the flow chart
           (1 mk)
 (ii)  Explain how process T is carried out     (2mks)
 (iii)  Explain why it is necessary to heat aluminum oxide before electrolysis is
carried out        (1mk)
(iv)  Suggest a reason to why carbon is not used for reduction of aluminium 
Oxide         (1 mk)
(v)  What properties of aluminum and the alloy make them suitable for use indicated?        (2mks)

17.  The flow chart illustrates the extraction of zinc and preparation of Zinc (II) sulphate crystals. Study it and answer the questions that follow
(a)
(i)  Name
 I.   Gas Q         (1 mk)
 II.  Liquid R        (1 mk)
 (ii)  Write an equation for the reaction that takes place in
- Chamber  I        (1 mk)
- The Roster        (1 mk)
- Chamber II        (1 mk)
 (iii)  Given that the zinc sulphide ore contain 45% of Zinc sulphate by mass calculate
I.  The mass in grains of Zinc sulphide that would be obtained from 250 kg of the ore         (1 mk)
II.  The volume of sulphur (IV) oxide (So2) that would be obtained from the mass of zinc        (1 mk)
III.  Sulphide obtained in 1 above at room temperature and pressure (Zn = 65.4) (S = 32.0) molar gas volume = 24 dm3
(b)  In such an experiment sulphur (IV) Oxide may keep escaping to the atmosphere. Explain how this could affect the environment.    (2mks)
(c)  Suggest one other man manufacturing plant that could be set up near Zinc extraction plant. Give a reason for your answer
18.  Iron Pyrites was heated in air to give Iron (III) oxide and a gas X: This is also when a yellow powder is burned in limited amount of air.
 (i)  Identify the yellow powder      (1 mk)
 (ii)  Identify gas X        (1 mk)
(iii)  Write a chemical equation to show the reaction between gas X and aqueous Sodium Hydroxide      (1 mk)
19. Hydrogen was passed over heated iron (III) oxide, but no reaction occurred. Iron (III) oxide was heated with carbon, Iron was formed and after separation it was dissolved in dilute sulphuric acid. A gas “Y” was evolved.
(a) (i)  Is the reaction between hydrogen and iron (III) oxide physical or
chemical explain      (2mks)
(ii)  Explain why carbon reached with iron (III) oxide while hydrogen did not        (2mks)
 (iii)  Identify gas Y       (1 mk)
(b)  Iron window frames corrode quickly unless carefully protected but aluminum window frames are resistant to corrosion
 (i)  Give the chemical name of the substance formed when iron rust
 (ii)  Why does aluminium items does not corrode as quickly as iron          (1 mk)
(iii)  Explain why galvanized iron is resistant to corrosion even when the protective layer of zinc is broken    (2mks)
20. Study the table below of oxides and sulphides formed by different elements and answer the questions that follow.
 
Elements Oxides Sulphides
Copper CUO, CU2¬O CuS, Cu2S
Hydrogen H2O H2S
With reference to the periodic table, what is the relationship between oxygen and sulphur         (1 mk)
21.  Two metals “A” and “B” have close packed and body centered cubic respectively. Which metal has the highest melting point     ( 1 mk)
22. Aluminium metal is a good conductor and is used for over head cables. State any two other properties that make aluminium suitable for this use.
23. The table below shows the properties of substances K, L, M and N
Substance Reaction with oxygen Melting point Conductivity
   Solid Molten
K Unreactive High Good Good
L Reactive Low Poor Poor
M Unreactive High Good Good
N Unreactive Low Good Good
  Select the substance which is likely to be
 (a)  Copper metal        (1 mk)
 (b)  Magnesium chloride
24. (a)  An ore is suspected  to containing mainly iron. Describe a method that can
be used to confirm the presence of iron in the ore   (4mks)
 (b)  Excess Carbon (II) oxide was passed over a heated sample of an oxide of iron as shown in the diagram below. Study the diagram and the data below it to answer the question that follows
 Mass of empty dish   10.98g
 Mass of empty dish + oxide of iron 13.30g
 Mass of empty dish + residue  12.66g
(i)  Determine the formula of the oxide of iron. Relative mass of oxide of iron is 232, Fe= 56, O = 16     (2mks)
(ii)  Write equation for the reaction which took place in the dish (1 mk)
 (c)  Corrosion is a destructive process in which iron is converted into hydrated (III) Oxide. State
(i) Two conditions necessary for rusting to occur  (1 mk)
(ii)  One method used to protect iron from rusting  (1 mk)
(d)  Explain why it is not advisable to wash vehicles using sea water  (2mks)
25. Lithium metal react with water less vigorously than sodium metal explain (1 mk)
 
TOPIC 5
ORGANIC CHEMISTRY II
1.  A compound where structure is shown below is found in detergent
 CH3(CH2)nCH     SO-3  Na+
With reference to the structure, explain how the detergent removes grease during washing         (2mks)
2. 
Complete the table below by inserting the missing information in the spaces provided         (4mks)
 
Name of polymer Name of monomer Use of polymer
Polystyrene  
 Vinyl chloride 
3. 
 The structure below represent five cleaning agents
 
 R – COO Na+  R  - OSO3Na+
      B
Which cleansing agent would be more suitable for washing in water containing magnesium sulphate? Explain       (2mks)
4. 
 (a)  Draw the structure of ethanol and propanoic acid   (2mks)
 (b)  Give the name of the organic compound formed when ethanol and
propanoic acid react in presence of concentrated sulphuric acid (1 mk)
5. 
 The structure below represent a portion of a polymer
 CH3  CH3  CH3
     |  |  |
 ―C―CH2―  C―CH2―――C―CH―
     |  |  | |
 COOCH3 COOCH3  COOCH3
 
 Give
 (a)  The name of the polymer      (1 mk)
 (b)  On industrial use of the polymer     (1 mk)
6. 
An organic compound with the formula C4H10O react with potassium metal to give hydrogen gas and a white solid.
(a) Write the structure formula of the compound    (1 mk)
(b) To which homologous series does the compound belong  (1 mk)
(c) Write the equation for the reaction between the compound and potassium metal         (1 mk)
7.
 Study the information in the table below and answer questions that follow
   
Alcohol’s Heat of combustion KJ/M
Methanol 715
Ethanol 1371
Prepanal 2010
Bufanal 2673
Give a reason why the differences in heat of combustion between successive alcohol are close
8. 
 Study the below chart and answer the questions that follows
 (a)  Identify N and P       (2mks)
 (b)  What name is given to the type of halogenation/ chlorinating reaction
given in step 2        (1 mk)
9. 
Name the process that takes place when crystals of Zinc Nitrate change into solution when exposed to air.       (1 mk)
10. 2007: PP 1 Q. 23
The table below shows the relative molecular masses and the boiling points of pentane and propane -1 –ol
Relative molecular mass Boiling point (0 C)
Pentane 72 36
Propan – 1-ol 60 97
 
Explain why the boiling point of propan –l –ol is higher than that of pentane
           (2mks)

11. 
The table below gives the information of some carboxylic acids and then draw points
 
Acid Boiling point (0C)
HCOOH 101
CH3COOH 118
CH3CH2COOH 141
CH3CH2CH2CH2COOH 187
CH3CH2CH2CH2CH2COOH 205
 (a)
(i)  Give the name of the acid whose formula
  CH3CH2CH2CH2COOH      (1 mk)
 (ii)  What is the empirical formula of CH3CH2CH2CH2CH2COOH (1 mk)
 (iii)  Plot the graph of boiling point against number of a ions of the carboxylic acids         (3mks)
I. From the graph determine the boiling point of the acid
CH3CH2CH2COOH      (2mks)
  (iv)  Explain giving reasons the shape of the graph   (2mks)
(b)  Explain the observation which would be made if NaHCO3 is added to an aqueous solution containing HCOOH    (2mks)
(c)  Calculate the volume of 0.2M sodium hydroxide solution which would be required to react completely with a solution containing 3.0 g of CH3COOH. (C= 12) (H= 1.0) (O= 16)    (3mks)
12.
 The formula given below represent a portion of a polymer
  H H H H
   |  |  |  |
         ―C     ―C      ―C     ―C―
   |  |  |  |

 (a)  Give the name of the polymer      (1 mk)
 (b)  One disadvantage of the continued use of this polymer  (1 mk)
13.
(a) When organ compound “Y” is reacted with aqueous sodium – carbonate. It produces carbon (IV) oxide. “Y” reacts with propanol to form a sweet smelling compound “Z” whose formula is.
               O
     |
CH3 ― CH2 ― C ― O ― CH2 ― CH2 ― CH3
 (i)  Name and draw the structural formula of compound “Y”  (1 mk)
 (ii)  What is the name of the group of compound to which “Z” belong  (1 mk)
(b)  In an experiment, excess ethanol is warmed with acidified potassium dichromate for about 20 minutes. State and explain the observations that was made at the end of the experiment
(c)  The scheme below was used to prepare a cleansing agent.  Study it and answer the questions that follow

(i)  What name is given to the type of cleansing agent prepared by the method shown in the scheme       (1 mk)
(ii)  Name one chemical substance in the scheme    (1 mk)
(iii)  What is the purpose of adding the chemical substance named in C (ii) above?         (1 mk)
(iv)  Name one other suitable substance that can be used in step 1 (1 mk)
(v)  Explain how an aqueous solution of the cleansing agent removed oil from utensils during washing      (3mks)
14. 
 (a)  Write the formula of
 (i)  Methanol       (1 mk)
 (ii)  Methodic acid       (1 mk)
(b)  Write the equation for the reaction between methanoic acid and the aqueous sodium        (1 mk)
(c)  (i)  Name the product formed when methanol react with methane acid
          (1 mk)
(ii)  State one condition necessary for the reaction in (c) (I) above to take place       (1 mk)
(iii)  Hydrogen gas reacts with hoxene form hexane. Calculate the volume as hydrogen as required to convert 42g of hexane to hexane at S.T.P (C= 12) (H=1) Molar gas volume at STP = 22.4dm3       (4mks)
 
15. 
The flow chart below shows a series of reactions starting with ethanol. Study it and answer the questions that follows.
(i) Name  I.  Process A      (1 mk)
  II.  Substance “B” and “C”    (1 mk)
(ii) Write the equation for the combustion of ethanol   (1 mk)
(iii) Explain why it is necessary to use high pressure to change B into polymer
          (1 mk)
(iv) State one use of methane      (1 mk)
16.
(a)  The list below gives the formula of some organic compounds. Use it to answer the questions that follow.
V1 CH3¬CH2CH2CH2OH
V2 CH3CH2CH3
V3 CH3CH2CH2C – OH
V4 CH3CH2CH= CH¬2
V5 CH3CH2CH2CH3
(i)  Select two compounds which
 I.  Are not hydrocarbons      (1mk)
 II.  Belong to the same homologous series   (1mk)
(ii)  Identify the compound that is likely to undergo polymerization.
 Give a reason for your answer    (2mks)
(b)  The structure below represents two cleansing agents
 R- COO – Na+
 R – OSO3 – Na+
In the table below give one advantage and one disadvantage using each of them
 Advantage Disadvantage
R- COO- Na+  
R-OSO3 – Na+  
 
(c)  Under certain conditions, Ethanoic acid C2H4O2 and ethanol reacts to form a sweet smelling compound
(i)  What is the general name of the compounds to which the sweet smelling compound belong       (1 mk)
(ii)  Write the formula of the sweet smelling compound   (1 mk)
(iii)  Give one use of ethanoic acid other than the formation of the sweet smelling compounds       (1 mk)
(iv)  Write an equation between dilute Ethanoic acid and solid potassium carbonate        (1 mk)
(d)  Fibres are either synthetic or natural. Give one
 (i)  Example of natural fibre     (1 mk)
 (ii)  Advantages synthetic fibres have over natural fibres  (1 mk)
17.
 (a) Give the systematic names of the following compounds
 (i) CH2 =   C   –  CH3

   CH3        (1 mk)
 (ii) CH3CH2CH2C ≡CH       (1mk)
(b)  State the observations made when propan-1-ol reacts with:
 (i)  Acidified potassium dichromate (VI) solution  (1 mk)
 (c)  Ethanol obtained from glucose can be converted to ethane as shown below
  C6¬ H12 O6   C2H5OH  CH2 = CH2
  Name and describe the processed that take place in steps I and II (3mks)
(d)  Compound A and B have the same molecule formula C3H6O2. Compound A liberates carbon (IV) oxide on addition of aqueous sodium carbonate while compound B does not. Compound B has a sweet smell. Draw the possible structures of.
- Compound A       (1 mk)
- Compound B        (1 mk)
(e)  Give two reasons why the disposal of polymers such as polychloethane by burning pollutes the environment.     (2mks)
 
18.
(a)  Alkanes, alkenes and alkynes can be obtained from crude oil. Draw the structures of the second member of the alkyne homologous series  (1 mk)
 
(b)  Study the flow chart below and answer the questions that follows

(i)  State the conditions for the reaction in step I to occur  (1 mk)
(ii)  Identify substance H       (1 mk)
(iii)  Give
 I. One disadvantage of the continued use of substances such as J
          (1 mk)
 II. The name of the process that takes place in step III  (1 mk)
 III. The name and the formula of substance K.   (2mks)
  Name………………………………
  Formula……………………………
(iv)   The relative molecular mass of J is 16,800 calculate the number of monomers that make up J      (2mks)
(c)  The table below gives the formula of four compounds L, M,  N and P
 
Compound Formula
L C2H6O
M C3H6
N C3H6O2
P C3H8
Giving a reason in each case, select the letter which represents a compound that:
(i)  Decolourise bromine in the absence of UV light  (2mks)
(ii)  Gives effervescence when reacted with aqueous sodium carbonate.
          (2mks)
19. The following is formula of monosaccharide (glucose)
  H H H H H H
   |  |  |  |  |  |
 H ――C―― C ―― C ――C ―― C ――C O
   |  |  |  |  |  |
  OH OH OH OH OH OH
 (i)  What is meant by monosaccharide     (1 mk)
 (ii)  How would glucose be converted into cellulose   (2mks)

20. Consider the following
 (i)     O
     
H2N ― CH2 ― C ― OH
 
 (ii)
    CH3          O
     |   
H2N ― C ―C
    
         OH
 (iii)
              O
        
NH2NCH ― C
    
         OH
(i)  What is the name of this class of compounds   ( 1 mk)
(ii)  What do ii and iii have in common?    ( 2 mks)
(iii)  Give the conditions of the reaction and name the products formed when compound i react with ethanol.
21. 2.635g of chloro propanoic acid (CLCH2CH2COOH), were dissolved into 250 cm3 of solution.  25 cm of the acid required, 25 cm3 of 0.1m potassium hydroxide solution for complete neutralization.
(i)  Write an equation for the reaction between potassium hydrate and chloropropanoic acid.       (1 mk)
 (ii)  Calculate the number of moles of chloropropanoic acid per dm3 (2mks)
 (iii)  Calculate the number of moles of
 (i)  Potassium hydroxide used     (1 mk)
(ii)  Chloropropanoic acid that would react with the number of moles of potassium hydroxide in 1 above    (2mks)
 
22. Below is a scheme of some reactions of ethanol. Study it and answer the questions that follow

 (i)  State the conditions and the reagents required in steps I, II, III and IV           (4mks)
 (ii)  Name the major products “A” and “B”    (2mks)
23.  A form (IV) student is interested in marking Tery lene for his project. He needs your advice on how to go about it.
 (a)  Explain to him what type of polymer is tery lene.                        (2mks)
 (b)  Given that tery lene is synthesized from ethane -1, 2-diol
  CH2CH2(OH)2   benzene -1, 4-dicarboxalic acid CH2  (COOH)2
(i)  Draw the polymer unit of tery lene consisting of two monomeric units.         (2mks)
(ii)  Name the product eliminated                                       (1 mk)
(c)  Give two
 (i)  Properties of tery lene                                                           (2mks)
 (ii)  Uses of tery lene                                                                           
(d)  (i)  Give two examples of natural polymer below.                   (2mks)
    (ii)  What is vulcanization?                          (2mks)
(e)  (i) Draw the monomer of the polymer below                  (1 mk)
           CH3        
              |   
― CH2 ― C ― = CH ― CH― n
 (ii)  Name the monomer                                                         (1mk)
                                       
24. Complete the following reaction
 CH 3CH 2OH (1)     Excess Con
                                                         
                                            H2 SO4
170                                                                            (1 mk)
25. Consider the following compounds
 (a) CH3 CH2 CH2COOH
 (b) CH3 CH2 COOCH2 CH2
 (c) HOOCCH2 CH2  COOH
 (d)CH3 CH(OH) CH3
Which of these compounds is
 (i)  Diabasic acid
 (ii)  An Ester  
26. How would each of the following compounds be chemically distinguished 
 CH3COOH and CH3CH2CH
27.  Name the regents and state the condition of the reaction necessary to affect the changes given below
 (a) C2H4 → C2H6        (1 mk)
 (b) C2H4¬ → C2H2        (1 mk)
 (c) C2H4 →CH3COOH       (1 mk)
28. The formula below represents the active ingredients in a detergent and in a soap respectively.
   H
    |
 CH3 (CH2)4 C   ― SO3- Na+
    |
CH3
 CH3(CH2)16COO-Na+
 (a)  What is a detergent?       (1 mk)
 (b)  Give two advantages and two disadvantages of using detergents as
cleansing agent       (2mks)
 (c)  Explain briefly the mode of action of soap during cleansing  (3mks)
 (d)  Give a reason for adding polyphosphate to the detergents  (1 mk)
 (e)  Explain briefly how the soap given above may be manufactured (3mks)
 
TOPIC 6
RADIOACTIVITY
1.  
 Complete the following equation
 (a) 14   14
      N + ?    →         C
   7   6      (1 mk)
 (b)  Give one use of radioactive elements     (1 mk)
2. 
 The table below gives the rate of decay for radioactive element Y
 
Number of days Mass (g)
0 348
270 48
 
Calculate the half – life of the radioactive element “Y”   (1 mk)
3. 
    233
100g of radioactive       Pa was reduced to 12.5g after 81 days
              91
 (a) Determine the half life of “Pa”
      233
 (b)  Pa decay by beta emission, what is the mass number and atomic number of
        91   the element formed       (1 mk)
4. 
 Complete the diagram below to show how ∂ and β particles from radioactive can be distinguished from each other. Label your diagram clearly. ( 3 mks)
 Source of radiation    Paper  Metal foil
5. 
 M grammes of radioactive is isotope decayed to 5.0g in 80g. The half life of the isolate is 25 days
 (a)  What is meant by half life      (1 mk)
 (b)  Calculate the initial mass “m” of radioactive isotope   (2mks)

6. 
    234
An isotope of uranium     U, decay by emission of an alpha particle to thorium (Th)
     94
 (a)  Write the equation for the nuclear reaction undergone by the isotope           (1 mk)
(b)  Explain why it is not safe to store radioactive substance in conditions made from aluminum sheet      (1 mk)
7. 
 The graph below shows the mass of a radioactive isotope plotted against time

 (a)  Using the graph determine, the half life of the isotope  (1 mk)
 (c)  Calculate the mass of the isotope present after 32 days  (2mks)
8. 
 A radioactive isotope X2 decay by emitting two alpha particles and one β particles to form 214
   β1
    83
 (a)  What is the atomic number of X2     (1 mk)
 (b)  After 112 days 1/16 of mass of X2¬ remained. Find the half life of X  (2mks)
9. 
 Study the nuclear reactions given in the scheme below and answer the questions that follows
 12  14  14
    C Step I    C Step II    C
   6    6    7
 
  12   14
 (a)     C and     C are isotopes. What is meant by the term isotopes?
    6     6
 (b)  Write an equation for the nuclear reaction in step II   (1 mk)
     14
(c)  Give one use of     C
           16
10. 
The graph below represents a radio active decay series for isotope “H”, study it and answer the equations that follows
(a)  Name the type of radiation emitted when isotope it changes to isotope “Y”
          (1 mk)
(b)  Write  an equation  for the nuclear reaction that occurs when “J” changes to isotope “K”        (1 mk)
(c)  Identify a pair of isotope of an element in the decay series  (1 mk)
11. 
100g of radioactive substance was reduced to 12.5 g within 15.6 years. Calculate the half life of the substance       (2mks)
12. 
 (a)  Complete the number equation below
  37  37
      A    +    B
  18  17
 (b)  State one
  (i)  Use of radioisotope in agriculture
  (ii)  Dangers associated with expose to human being to radioisotopes
           (1 mk)
13. 
 (a)  Distinguish between nuclear fission and nuclear fusion
 (b)  Describe how solid wastes containing radioactive substances should be
disposed of        (1 mk)
14. 
 (a) A radioactive substance emits three different particles.
  Give the symbol of the particles with the highest mass  (1 mk)
 (b)  (i)  Find the values of Z1 and Z2 in the nuclear equation below
   Z1   1   94    140  1
     U   +  n   Sr   +    xe +2 n
   92  0 38 Z0 0 
  (ii)  What type of nuclear reaction is represented in b (i) above? (1 mk)
15. 
 (a)  State the difference between chemical and nuclear reactions  (2mks)
 (b)  Below is a radioactive decay series starting from
  214   206
       B1 and ending at       Pb. Study it and answer the questions
    83     82
  that follows
  214  219  210  210  210  206
       BI Step I      tI Step II      pb Step III      BI Step IV      PO Step V      pb
    83    84    82    83    84    82
 (i)  Identify the particle emitted in step I and III.    (2mks)
 (ii)  Write the nuclear equation for the reaction which takes place in step V 
           (1 mk)
(c)  The table below gives the percentage of radioactive isotope of Bismuth that remains after decaying at different times.
Time (mm) 0 6 12 22 38 62 100
Percentage of Bismuth 100 81 65 46 29 12 3
(i)  On the grid provided plot a graph of the percentage of bismuth remaining (vertical axis) against time      (3mks)
(ii)  Use the graph, determine the
 I.  Half life the Bismith      (1 mk)
II.  Original mass of bismuth isotope given that the mass remained after 70 minutes was 0.16g     (2mks)
 d. Give one use of radioactive isotope in medicine   (1 mk)
16. Copper 64 has a half life of 12.8 his
 (a)  What is meant by half life?      (1 mk)
(b)  Draw a graph to show the decay of copper 64 from an initial activity to 64 counts per minute to four percent minutes    (4mks)
17. Complete the following nuclear equations
 (a) 55  55  
     Cr →    Mn + _________
  24  25 

 (b) 1 235  143           1
    n   +      U →      La + 3 n +_________
  0   92    57           0

18. A quality of 11X” was mentioned with a G.M tube scalar. The following results were obtained over a period of 70 minutes.
 
Time Cents per minute
0 800
10 560
20 427
30 305
40 225
50 165
60 122
70 85
(a)  Plot a graph of time against the counts per minutes   (4mks)
 (b)  Determine the half life of 44X      (3mks)
 (c)  Starting with 32g, of 44X how much of the isotope would be remaining after 110 minute?       (3mks)
19. Study the nuclear reaction and answer the questions that follows
 238
      U13   → X   → 13Y → 13  → Z
   92
 Determine the mass number and atomic numbers of X, Y and Z
20. (a)  When a stream of low energy  particles is directed towards a thin  of
aluminium, the following observation are made
 (i)   Most of particles pass straight the foil
(ii)  The remaining ones are either deflected or emerge from the same as they originally entered     (4mks)
(iii)  If the energy of the particles is increased, some are absorbed by the aluminium foil comments on this observation.  (4mks)
   31
21. The isotope     X  has a half life of 2.5 hours
           14
 Calculate the % (percentage) of a given mass of the isotope left after 7.5 hours
           (1 mk)
 
22. Below is a diagram of a deflection and penetrating powers of three radiations from a radioactive source

 (a)  Name the radiations labeled X, Y and Z    (3mks)
 (b)  Why are radiation X stepped by a thin piece of paper
23. Complete and balance the following nuclear reaction   (3mks)
  238  12  
 (i)      U  +    C → -96CF + _______
    92    6
  
28  16  1
 (ii)    U +    O →   n + -100FM
  __    8  0
  ―  241
 (iii)    Pu →      Am + _______
  94  95
ANSWERS TO TOPICAL QUESTIONS
TOPIC 1
SIMPLE CLASSIFICATION OF SUBSTANCES
1.
  W X Y K

2. To the mixture of sugar, camphor and alum, add either camphor dissolves leaving behind alum and sugar. Filter the mixture to obtain sugar and alum a residue. Add ethanol to this residue sugar will dissolve leaving behind alum as a residue. Filter the mixture, sugar will be in a solution of ethanol (filtrate) allow the filtrate to evaporate and solid sugar will be left behind.
3. (a)  In both cases the energy/ heat added  is used  to separate/ split/ weaken the
bonds holding the particles together.  We call this energy latent heat of fusion.
 (b)  CdCl2(s) → Cd(l)2+ + 2Cl-(l)
This is because CdCl2 is an ionic compound where the particles (ions) are held together by strong electrostatic force of attraction – compared to weak Vanderwaal forces and hydrogen bonds holding the molecules of water together.
4. (a)  Pass the mixture  of gas “D” and  “E” through  sulphuric acid. Gas “D”
will react to form salt- leaving behind gas “E” Collect gas E by downward delivery/ upward displacement of air since  it is heavier than air.
 OR
Pass a mixture of gas “D” and “E” over sodium hydroxide. Gas “D” will dissolve but gas “E” will not be affected. Collect gas “E” by downward delivery
 (b) Ammonia gas (HN3)
Ammonia is lighter that air.  It reacts with acids to form salt since itself is basic. It does react with sodium hydroxide since both are basic but will dissolve in it without any reaction.
5. Compress and cool the mixture to a temperature below 1960C i.e. (-2000C) to form liquid air. Allow the mixture to expand and warm. Nitrogen will vaporize first since it is more volatile. Oxygen will start to vaporize when a temperature of -1830C is attained.
6. (a) –  The thermometer is touching the  mixture
       -  Direction of flow of water in the liebigs condenser reversed
       -  Naked flame used to heat organic compound and yet they are flammable.
 (b)  –  Test the boiling point or
 -  Test the freezing point or
 -  Test its density or
 -  Its refractive index
7. Add water to the mixture and stir, sodium chloride will dissolve leaving behind copper (II) oxide. Filter the resulting mixture the filtrate will contain dissolved sodium chloride. Evaporate the filtrate to dryness to obtain solid sodium chloride.
8. (a)  Liebig’s condenser
(b)  Determine the point at which one of the liquids in a mixture has evaporated completely. Temperature tends to remain constant when one liquid in a mixture is evaporating.
 (c)  Liquid “C” since it is more volatile
9. (a) 3mg(s) + N2 (g)→ Mg3N2(s)
 (b) Argon
  Helium
  Krepton
  Xenon
 The above mentioned are rare/ un-reactive gases and do not combine with other substances easily
10. (a)  Reduction process
 (b)  Oxygen is removed from lead (II) oxide it’s reduced into lead metal
 (c)  – Hydrogen gas
  - Ammonia gas
11. C- Unburnt gas  D- Luminous yellow flame
12. (a)  G  (b)  A1
13. (a)  Cooling (b)  Latent heat of fusion
14. (a)  A black mass of substance which is spongy will be formed. A lot of heat is
given out. This is a chemical reaction. The formula of new substances are C for carbon
  H¬2 O for water vapour
 (b)  A purple vapour is formed that condense at the cooler part of the test tube as grey crystal. This is a physical change. No new compound is formed
(c)  A brownish gas is produced another gas lights a glowing splint. A black substance is left in the tube. This is a chemical change. The formula of new substances
 Copper (II0 Oxide  Cuo
 Nitrogen (IV) Oxide No2
 Oxygen gas  O2
 (d)  The pallets melts forming a colourless solution.
  Type of reaction physical or chemical
  Formula Na2CO3 and H2O
 Ps Sodium hydroxide is deliquescent it can also react with CO2 in solution to give sodium carbonate and water.
15. (a)  Fractional distillation
(b)  (i)  Add water to the mixture. Stir. Sodium chloride being ionic
dissolves. Filter the mixture to remove sulphur as a residue. Evaporate the filtrate to obtain solid sodium chloride
  (ii)  Determine the melting point, pure sulphur melts at 1140C OR
   Pure sulphur will have constant/ sharp boiling point
 (c)  (i)  Potassium bromide/ KBr
  (ii)  60 – 55 = 5g (units a  must)
  (iii)  Fractional crystallization
  (iv)  Separation of components of trona from lake Magadi
- Manufacture of Na2CO3
- Manufacture of NaCl
- Extraction
- Production
16. Pass the air through a filter to remove dust, then bubble it through potash solution to remove CO2; cool and compress the remaining air to get liquid air. Warm and allow it to expand. Nitrogen b.p – 1960C vaporizes first.
17. (a)  Fractional Distillation
 (b)  Paper chromatography
 (c)  Sublimation
 (d)  Use of a magnet
18. (a)  (i)  Over water
  (ii)  Upward delivery/ downward displacement of air
  (iii)  Downward delivery/ upward displacement of air
 (b) (i)  Fractional distillation
  (ii)  Upward delivery: It is less dense than air
  (iii)  Downward delivery: it is denser than air
19. (a)  Fractional distillation
(b) Round bottom flask: Fractionating column, Liebig’s condenser, thermometer, means of heating.
(c) Not to heat the mixture in open/ naked flame since the liquids are flammable. Use water bath
20. (a)  Carbon (IV) Oxide is removed in step I and oxygen removed  in step II
 (b) Step I – concentrated sodium/ potassium hydroxide
  Step II – Heated metal e.g. copper
21. Heat the mixture naphthalene will sublime leaving behind common salt. Cool the sublimate to get solid naphthalene.
22. (a) (i)  The solution was saturated
(ii)  The remaining solid will dissolves. This is because increase in
       temperature increases the solubility of potassium nitrate.
(iii)  Crystals will be formed
(b) (i)  Copper Nitrate and Sodium Sulphate/ soluble salt of copper and
soluble sulphate salt.
(ii) Cu2+(aq) + SO4-2 (aq) → CuSo4(aq)

(iii)  The solid will change from white to blue crystals. Heat will be produced. A chemical reaction will occur and anhydrous copper (II) sulphate will be hydrated.
23. (a)  Heat water steadily
  Thermometer should not touch the beaker
  Stir the naphthalene continuously
 (b) (i)  Determine the temperature
  (ii)  Stir the naphthalene so as to distribute heat evenly
  (iii)  Transfer heat to naphthalene so as to melt it.
 (c) - Presence of impurities
  - Experimental errors
  - Heat loss to the surrounding
24. (a)  Lime water
 (b)  White precipitate
 (c)  Co2(g) + Ca(OH)2 (aq) → CaCO3(s) + H2O(l
25. - To protect potassium from moisture and dry oxygen with which they react
 - Phosphorous reacts with dry oxygen not moist oxygen
26. Dissolve the moisture in cold water and stir R dissolve. Filter to get solid “S” and “V” as residue. Evaporate the filtrate to get R. Put “S” and V in hot water and stir. V dissolve filter to get S as a residue. Evaporate filtrate to get V.

TOPIC 2
ACIDS, BASES AND INDICATORS
1. White precipitate, which dissolve in excess of sodium hydroxide to give a clear/colourless solution.
2. Concentrated sulphuric acid is a covalent compound. Dilute sulphuric acid is an ionic compound. It ionizes fully producing more hydrogen ions (H+)
3. The evolution of carbon (IV) oxide increases with time then remains constant. Initially there were many particles reacting together. After 20 seconds all calcium carbonate were used up and the reaction came to a completion.
 4. M Zn(NH3)4+2
 N Zn (OH)4-2
5. Nitrous acid ionizes more compared to hypochlorous acid. Hypochlorous acids is a very weak acid. It ionizes partially producing few hydrogen ions(H+)
6. B Strong acidic
 C Weak acidic
 D Strong basic
7. Carbonate reacts with acid producing carbon (IV) Oxide which is weaker acidic in presence of water. It changes litmus to pink. Sulphite on the other hand reacts with dilute acid producing sulphur (IV). Oxide which is a strong acid in presence of water. It changes litmus to red.
8. NH4+ acts as an acid
 It donates protein (H+) to H20(l) and converts it to hydroxonium ion (H3O+). H3O+ act as an acid in the backward reaction as it donates proton to NH3(g) and convert it to NH4 + (aq).
9. React Lead Carbonate with dilute Nitric acid to get a solution of lead Nitrate.
 PbCO3(s) + 2HNO3 → Pb(NO3)(aq) + H20(l) + CO2(g)
 Dissolve potassium sulphate in water to get its solution. Mix potassium sulphate solution with Lead Nitrate Solution to obtain Lead Sulphate as a precipitate.
 Pb(NO3)2 + K2SO4(aq) → PbSO4(s) + 2KNO3(aq)
 Filter the resulting mixture to obtain Lead sulphate as a residue. Wash it with distilled water and dry it.
10. Strong acid is the one which ionizes fully producing more hydrogen ions when in solution with water e.g.
 - Hydrochloric acid
 - Nitric acid
 - Sulphuric acid
 Weak acid is the one which ionizes partially in solution of water producing few hydrogen ions e.g.
 - Ethanoic acid  - Propanoic acid
11. Add excess Lead (II) Carbonate to Nitric acid. Wait for the reaction to be completed. Filter the resulting solution mixture. To the filtrate (Lead Nitrate) add excess dilute hydrochloric acid. Filter the mixture to get lead (II) chloride.
 PbCO3(s) + 2HNO3(aq) →Pb(NO3¬)2(aq) + H2O(l) + CO2(g)
 
 Pb(NO3)2(aq) + 2HCl(aq) →PbCl2(s) + 2HNO3(aq)
12. Sting from the bee contains Histamine which is acidic. This causes irritation. Sodium hydrogen carbonate being alkaline/ basic neutralizes the acid to remove the irritation.
13. The blue crystal change to a white powder. Conc sulphuric acid is a dehydrating agent. It removes water of crystallization from hydrated copper (II) sulphate.
 
CuSO4:5H2O Conc H2SO4    +  CUSO4 +5H2O(l)
Blue crystals    White powder
14. Moles of HNO3 = Molarity x Vol = 2 x 50 = 0.1 moles
    1000  1000
 Moles of KOH in 50cm3 = 0.1 moles
 Moles of KOH in 100cm3 = 0.1 x 2 = 0.2 moles
 Mass of D = 0.2 x 56 = 11.2g
15. (a)  Brown ring where the layers of acid meets the layer of the  nitrate and
sulphate.
 (b)  2KNO3(s)      heat  2KNO2(aq) + O2 (g)
16. React with sodium hydrogen carbonate to form carbon (IV) Oxide which causes the dough to rise as it tries to escape.
17. - To neutralize soil acidity
 - Add Ca2+ ions to the soil which is needed by plants i.e. it acts as a fertilizer.
18. (a)  H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + H2O(l)
 (b)  Blue litmus paper change to red. The red litmus remained red.
 (c)  The acid used was in excess i.e
  Moles of both acid and bases are
  30 x 0.1 = 0.003 moles
   1000
 But NaOH: H2SO4 reacts in the ratio of 2:1
 Hence we expect 0.003 moles of NaOH to react with 0.0015 moles of H2SO4. The acid was in excess by 0.0015 moles.
19. (a)  Pb2+
 (b)  Zn2+
 (c)  Co3-2(a) + Zn2+(aq) → ZnCO3(s)
20. Hydrochloric acid is a strong acid. It ionizes fully in solution of water. Therefore there are more hydrogen ions to be displaced by magnesium. Ethanoic acid is a weak acid. It ionizes partially in solution of water. It contains few hydrogen ions to be displaced by magnesium.
21. (a)  Ammonia gas reacts  with water producing ammonia solution
  NH3(g) + H2O(l) →NH4OH(aq)
  Ammonia solution is a weak alkali. It ionizes partially producing hydroxyl ions [OH-]. The [oh-] ions changes red litmus to blue.
(b)  The funnel prevents the sucking back of water as ammonia is very soluble in water.
22. (i) ZnO(s) + H2SO4(aq) →ZnSO4(aq) + H2O(l)
 (ii) znO(s) + 2 NaOH(aq) →Na2ZnO2(aq) + H2O(l)
 (b) Zinc oxide is amphoteric in nature
23. Acid “L” is a weak acid. It contains few hydrogen ions to be displaced by magnesium. Acid “M” is a strong acid. It ionizes fully. There are more hydrogen ions (H+) to be displaced by magnesium.
24. (a)  Copper (II) Hydroxide [Cu(OH)2]
 (b)  Tetra – amine copper (II) ions [Cu (NH3)4+2]
25. The product from nettle plant is acidic aqueous ammonia solution being basic neutralize the acidic product.
26. (a)  (i) Colour change from green to brown
  (ii) Reddish brown precipitate
 (b) Fe3+(aq) 3OH-(aq) →Fe (OH)3(s)
27. (a)  O-2
 (b)  [Zn(OH)4]-2
28. Zn(s) + H2SO4(aq) →ZnSO4(aq) + 42(g)
 Zn(s) + 2H2SO4(l) →ZnSO4(aq) +SO2(g) + 2H2O(l)
29. Amphotric
30. (a)  Neutralization
 (b) (i)  Calcium hydrogen carbonate
  (ii)  Drying agent, extraction of sodium
31. (a)   (i)  Hygroscopy
  (ii)  Deliquescence
  (iii)  Efflorescence
 (b) (i)  Zn(OH)4-2
  (ii)  NH4(NH3)4+2
 (c) (i)
Elements Fe S O H2O
% by mass
RAM
Moles

Ratio 20.2
56
20.2
 56
0.36
0.36
1 11.5
32
11.5
 32
0.36
0.36
1 23.0
16
23.0
16
1.44
0.36
4 45.3
18
45.3
18
2.56
0.36
7
FeSO4:7H2O
(iii)  Moles of salts = mass =  6.95 = 0.025 moles
     RMM    278
Molarity = moles x 1000 = 0.025 x 1000
  Volume     250
= 0.1M
32. (a)
 
(b)  (i)  Conc of Br2 after 3 minutes 5.3 x 103 x 103 mol/dm3  - 0.1
  (ii)  Change in concentration
   Change in time
 (9.6 – 5.0) x 103
      3-0
 = 1.53 x 103 mol/dm3
(c) At high concentration the rate of reaction is high because the particles in the solution collide at a high frequency or more particles collide more often.
 (d)  At a lower temperature the particles have less kinetic energy hence frequency of collision is reduced or few particles have activation energy.
33. (a) Mg(s) + 2 HCl (aq) → MgCl2(aq) + H2(g)
(c) (i)  300cm3  10cm3 depending on the scale
 (ii)  The volume of 2M HCl which reacted completely with 0.6g of Mg powder is 30cm3
(d) (i)  The rate of reaction will be  lowered. Magnesium ribbon has a
small surface area  that the powder. Hence the collision of particle between magnesium particles in a ribbon and hydrochloric acid particles will be reduced.
(ii)  Rate is increased since the number of particles of HCL hydrochloric acid will be higher/ concentration is increased. Hence particles collide more frequently.
(e) Moles of hydrogen gas produced =    600cm3
      240000cm3
        = 0.25 moles
 Moles of mg = 0.25 moles
 RAM of mg = 0.6 = 24
     0.25
34. (a)  -   Malachine
      -  Copper pyrite
  -  Chalcosite
  -  Cuprite
 (b) (i)  Hydrogen Sulphide / H2S(g)
-Soluble carbonate i.e sodium carbonate Na2CO3/ potassium 
 carbonate/ K2CO3- including their bicarbonate KHCO3/NaHCO3
-Copper (II) Oxide/ CuO
(ii)  CuCO3   heat CuO(s) + CO2
  (iii)  Step 4
- Green Solid dissolves to form blue solution
- Effervescence and bubble  of  colourless gas which forms precipitate with lime water are produced
Step 7
Black solid dissolves to form a blue solution.
(c)  (i)  Tin/Sn
(ii)  Ornaments/ medals/ bearing metals in machines/ coinage/ gear               wheels
Earings/ door handles/ electrical contact.
35. (a) (i) Put soil in water in a beaker. To the mixture add universal
   indicator. Compare the colour change to the Ph Chart.
  (ii) Addition nitrogenous fertilizer which are acidic.
36. (i)  Q
 (ii)  Pink/Red
37. (a)  Number  of hydrogen ions (H+) which can be displaced by a metal or
ammonium radicals to form salts
(b)  Ethanoic acid had a basicity of one (i) since one one hydrogen ion in the carboxalate group (-COOH) can be displaced
38. (i)  Yellow in acidic medium: The H+ ions of the acid  react with OH- from
indicator producing more H2O. The equilibrium shift to the right side.
(ii)  Blue in alkaline medium. The OH- ions/ radicals from alkaline solution increases the concentration to the right. Equilibrium shift to the left side.
39. K+ and CO3-2
 Na+ and CO3-2
40. (i)  C
 (ii)  D
 (iii)  B
 (iv)  A
41. (a)  C
 (b)  A
 (c)  D
42. (a)  Dirty green  precipitate is formed
  Observations
  Dirty green precipitate changed to give a reddish brown precipitate
 (b)  (ii)  Explanation
Iron (II0 hydroxide which is green is oxidized to iron (III) hydroxide by oxygen in the air
43. Strong acid is the one which ionizes fully while in solution with water
 Weak acid ionizes partially while in solution with water
44. NaOH(aq) →Solution D
 CH3COOH(aq) → Solution C
 HCI (aq) → Solution B
 NH3(aq) →Solution A
45. (a)  Fe)s) + H2SO4(aq) →FeSO4(aq) + H2(g)
 (b) (i)  Dirty green  precipitate formed
  (ii)  Fe2+ (aq) + 2OH(aq) → Fe(OH)2(s)
 (c) (i)  2Fe-2 (aq) + 4H + (aq) + 2NO-3 → 2Fe3+(aq) + 2NO2 + 2H2O(l)
(ii)  Oxidizing agent: It oxidizes Iron (II) (Fe2+) to iron (III) compound (Fe3+)
(d) (i)  Green  solution will be formed
(ii)  Zinc acted as reducing agent. It reduces Iron (III) (Fe3+) to iron (II) Compound (Fe2+) which is green.
  (iii)  2 Fe3+ (aq) + Zn(s) →2Fe2+(aq) + Zn2+(aq)
 
TOPIC 3
AIR AND COMBUSTION
1. (a)  The blue litmus paper would turn pink/ red. Red litmus paper remains red.
The carbon (IV) oxide produced when the candle burns dissolves in water to form  a solution of weak carbonic acid.
 (b) x- y   x   100%
         x
2. Observation: At No rusting takes place
 Explanation: Zinc is more reactive than iron.  It reacts with oxygen in presence to iron hence preventing it from rusting. It acts as a sacrificial metal
 Observation at B
 The nail is covered by reddish brown substance/coating/rust
 Explanation: Copper is less reactive than iron. Iron combines first with oxygen in presence of moisture and rust.
3. (a)  To remove the layer of oxide  on their surfaces which  could inhibit the
reaction
 (b)  Q, R,P
4. 2Mg(s) + O¬2(g) → 2MgO(s)
 3 Mg(s) + N2(g) →Mg3N2(s)
5. Oxide: Highest oxidation number
 P2O5 (+5)
 Cl2O7 (+7)
6. CO(g) + PbO(s) →Pb(s) + CO2(g)
 Observations
7. -Iron will be covered by a reddish brown substance/coating/rust
 -Water in test tube rise and water in a beaker drops
 Explanation:
 Iron Combines with oxygen  in a presence of moisture to  form hydrated Iron (III) oxide / rust water rises up to occupy the  space which was occupied by oxygen  in the tube.
8. Al2O3 (Aluminium Oxide)
9. Change was greatest with Magnesium. Both react with oxygen gas to form oxides, but magnesium also reacts with nitrogen to form magnesium nitrate (Mg3N2)
10. (i)   Mass increase: Oxygen combines with copper metal to form copper (II)
Oxide.
(ii)  Mass decrease: copper Nitrate decomposes to give gases that escape leaving behind copper (II) oxide.
11. Magnesium is above iron in the reactivity series. It supply electrons to the iron bar hence   prevent it from rusting/ cathode protection.
12. Magnesium produces a lot of heat/ energy when burning. This splint sulphur (IV) oxide into sulphur and Oxygen. Magnesium burns in the oxygen produced.  Burning splint produces less energy which is not enough to break sulphur (IV) oxide.
13. (a)  Manganese (IV) Oxide/ MnO2(s)
 (b)  2 H2O2 (aq)   MnO2   2 H2¬O(l) + O2
 (c)  - Respiratory aids from patients suffering from respiratory diseases /
  during surgery.
- High mountain climbers and deep see divers
- Helps in combustion of rocket fuel
- Welding together with other gases such as hydrogen/ oxygen (hydrogen  
  flame) acetylene/ oxyacetylene flame.
14. Nitrogen (II) Oxide is oxidized by oxygen in air to form nitrogen (IV) oxide. This gas is acidic when dissolved in water. May cause acidic rain.  If inhaled by animals/ man may corrode respiratory surfaces exposing them to disease causing agents.
15. 2C(s) +  O2 (g) →2CO(g)
 Fe2O3 + 3CO(g) →2 Fe(s)  + 3 Co2(g)
16. (i)  SO2/ sulphur(IV) Oxide
 (ii)  2CuFeS2 + 402(g) → 2FeO(s) + 3S)2(g) + Cu2S(s)
 (iii)  Fe2+
 (iv)  Carbon (IV) Oxide or carbon (II) Oxide
(v)  Reduction/ oxidation = Redox since Cu2O is reduced to Cu and CO oxidized  to Co2
(b)

(c)  Mole ratio of CU in CuFeS2 = 1.1
 Moles of Cu produced = 210 = 3.3 moles
        63.5
RFM of CuFeS2 = 63.5 + 56 + 64 = 183.5
Mass of Cu in CuFeS2 = 3.3 x 183. 5 = 605. 6 kg
% purity = 605 x 100 = 74. 76%
  810
(d)  - Formation of acidic rain due to presence of sulphur (IV) oxide
 - Sulphur (IV) oxide is poisonous
 - Carbon (II) is poisonous
 - Global warming due to presence of carbon (IV) oxide
 - Dumping of wastes like slag prevents growth of vegetation
 - Soil erosion due to the excavation of the ores
17. (i)  Bitumen: It has the highest  boiling  point
(ii)  Fractional distillation: they have different boiling points, petrol boils out first
(iii)  Each component is a mixture of hydrocarbons/ impure or there is presence of isomes in each component.
(iv)  Methane → CH4 all alkane gases up to C = 4
(b)  Burning in limited air will produce carbon (II) oxide which is poisonous
(c)  - Manufacture of tar used in tarmac road/ surface of roads
- Amending leaking roofs.
18. (a) (i)
  (ii)  Sodium peroxide Na2O2
(b)  (i)  4P(s) + 5O2(g) → 2P2O5(g)
(ii)  Phosphorous (V) oxide dissolves in water to form an acid (Phosphoric acid)
(c)  A firm oxide (aluminium Oxide) is formed on the surface of the metal. This oxide protect aluminium from further attack
(d)  (i)  A reaction which proceeds  by production of heat i.e heat is lost to
the surroundings.
(ii)  The yield be lowered: through by Le- Chateliers principle, the yield is expected to increase. But lower temperatures will result into fewer particles attaining activation energy.
(iii)  RMM of SO3 = 32 + 48 = 80
  Moles of SO3 used = 350 = 4.38 moles
     80
 Moles of H2S2O7 = 4.38 moles
 RMM of H2S2O7 = 2 + 64 + 112 = 178
 Mass of H2S2O7 = 4.38 x 178 = 779.6 kg
19. (a)  (i)  Potassium Hydroxide or sodium hydroxide
(ii)  Air allowed to expand and warm up. Nitrogen gas vaporizes first since it is more volatile. On further heating- oxygen vaporizes.
 (b) (i)  Hydrogen gas
  (ii)  -     For the complete oxidation of ammonia gas
- To increase the yield  of  nitrogen  (II) Oxide
- To reduce the cost
(iii)  Nitrogen gas
(iv)  NH3(g) + HNO3(aq) →NH4NO3(aq)
(c) Brown gas (Nitrogen (IV) Oxide gas) and an acidic gas (sulphur (IV) oxide) formed
Nitric acid reduced into nitrogen (IV) oxide, water and oxygen. Sulphur is oxidized into sulphur (IV) oxide which dissolves in water forming sulphuric acid.
20. (a)  Carbon and  hydrogen
 (b)  (i)  The candle will go off/ extinguished since carbon (IV) oxide and
water vapour accumulate around the candle carbon (IV) oxide does not support burning.
OR The supply of oxygen will be supported and candle goes off
(ii)  Mass increase
Water combines with calcium oxide to form calcium hydroxide solution. This combine with carbon (IV) oxide to form calcium carbonate.
(iii)  - Carbon (IV) oxide
   - Carbon (II) oxide
(iv)  Protect calcium from obtaining water from the atmosphere
(v)  -Concentrated sulphuric acid
 -Calcium chloride
21. Iron metal is corroded by rust in presence of water and oxygen
22. There will be formation of a white precipitate. Candle burns producing carbon (IV) oxide.
23. Air contains carbon (IV) Oxide which dissolve in water producing a weak carbonic acid
24. Na + ions
25. 3Mg(s) + N2(g) →Mg3N2(s)
 Mg3N2(s) + 6H2O(l) → 3Mg(OH)2(aq) + 2NH3
26. (a)  Beaker A: No soot at the bottom
  Beaker B: A lot of black soot at the pattern
(b) Sample A: Non luminous flame produces a lot of heat.
(c)
Luminous Non Luminous
- Produce a lot of light
- Very sooty
- Large and wary
- Burns quietly - Produces less light
- Not Sooty
- Short and steady
- Burns with roaring noise
27. (a)  CO3-2 is an oxidizing agent. It removes hydrogen from water (H2O) and
oxidizes it to OH.
(b)  Fe2+ is a reducing agent. It adds electrons to Cl2 and reduces it to 2CL-
28. (a)  To allow all oxygen to be  used  up and also to allow the gas to contract/
cater for any expansion of gases
 (b)  To absorb carbon (IV) oxide which was produced by the burning candle
 (c)  % of oxygen 90 – 70 x 100 = 22.2%
    90
29. (a)  Curve B: Pure substances has  sharp/ fixed constant  melting and boiling
points
(b)  Impurities rises the boiling point pressure rises the boiling point i.e when pressure is high b.p is very high.
 
TOPIC 4
WATER AND HYDROGEN
1. i) If ignited immediately explosion would occur because it would still be
mixed with air.
 ii) 2H2(g)  + O2(g)             2H2O (g)
2. 
Metals Aqueous solution containing ions of metals
 P R T
P X X X
R  X 
T  X X
3. a) Sample II: because the volume of soap used is less i.e. 3.0 cm3 and
remains the same after boiling.
b) Sample II is temporary had water because after boiling it became soft.  Volume of soap change from 10.6 to 3.0 cm3
4. 2HCl(aq)      +   Zn(s)           ZnCl2(aq)    +   H2
 2H2(g)    +  O2(g)     2H2O (g) 
5. a) Moles of Zn = 1.96 = 0.03 moles
    65.4
  Moles of HCL: 100 x 0.2 = 0.02 moles
      1000
  Expected moles ratio of Zn: HCl
        1:2
  Moles reacting 0.01: 0.02
  Moles of Zn were in excess by
  0.03-0.01 = 0.02 moles
 b) Moles of H2 produced = 0.01 moles
  Volume = 22.4 x 0.01 = 0.224dm3
     OR     0.224cm3
6. a) 2Li(s) + H2O(g)             Li2O(s)   +H2(g)
b) Potassium is very reactive and the reaction is likely to be explosive/violent.
7. a) to generate steam which will push air out.
b) Oxygen in air would oxidize zinc to zinc Oxide and no gas/Hydrogen would be produced.
c) It is less dense than air,
8. a) SO4-2 and NH4
 b) From ammonium and sulphates based ferterlizers.
  NH4 can also be from humus- when they decay.
9 a) The Ca2+ and Mg2+ ions in the permutit
 b) By passing a solution of concentrated sodium chloride/ brine through the
permutit.
 c) Provide Ca+2 ions necessary for bone and teeth formation.
-When passed through lead pipe the lead sulphate coat the inside as it is insoluble.  This prevents chances of lead poisoning.
10. a) Cations: Al3+
 b) Anions: SO4-2
  Ba2+(aq)   + SO4-2   BaSo4(s)
 C) a) H2O(g) + mg(s)             MgO(s)   + H2(g)
11. a)     H2O(g) + Mg(s)         MgO(s)     +     H2(g)
 b) It is insoluble in water.
12 a) Effervescence and bubbles of colourless gas were liberated.
b) Copper turnings will settle at the bottom.  There will be no reaction since copper does not react with an acid unless the acid is an oxidizing agent.
13. a) Presence of Ca(HCO3)2  and  Mg(HCO3)2 salts which aresoluble.
b) During distillation pure water is evaporated and then condesed leaving behind solids CaCO3(s)    and MgCO3(s) as their hydrogen carbonates decompose during the process.
14. It has one electron in its outermost energy level which it can lose to form H+ showing oxidation state of +1 or gain one electron to form H- showing oxidation state of  -1.
15. H(g) + e-  H(g)
   ∆H = -ve
 H(g) → e- + H+(g) ∆H = +ve
16. a) H2O(g)   + C(s)             H2(g)   +  CO(s)
17. a) SO4-2 ions
 b) Ba2+(aq) + SO4   Tetraamine zinc  (II) ions
 c) Zn (NH3)2 and Ca (HCO3)2 decomposes producing CO2 when heated.
18. a) Carbon (IV) Oxide gas
 b) Mg(HCO3) and Ca(HCO3)2 decomposes producing CO2 when heated.
19. a) No change in volume since the number of moles of acid is equal in both  
cases.
 b) It is less dense and does not burn like hydrogen.
20. Changes anhydrous copper (II) sulphate from whit to blue.  Or changes colbalt chloride paper from blue to pink.
21. a) i) Add one drop of liquid to anhydrous compper (II) sulphate it will
turn blue from white.
 OR
Add one drop to anhydrous cobalt chloride; it will turn pink from blue.
 b) i) Large suspended particles e.g leaves, stones, sand, gravel/grit.
  ii) Sedimentation or precipitation
iii) (a) Causes the small suspended particles to settle/precipitate.  (b)  Destroy micro-organism
iv) a) Permanent hardness
c) Addition of washing soda Na2CO3 which precipitate g2+(aq) as gCO3(s).
 
22. a)

 b) 2H2(g)  +  O2(g)              2H2O(g)
 c) Moles of H2 produced: 1.2   = 0.05 moles
       24
  Moles of Zn = 0.05 moles
  Ram of Zn; 3.27 = 65.4
    0.05 
       d) -Hydrogenation of fat
  -Weather balloons
  -Welding when mixed with other gases e.g oxygen to give oxyhydrogen.
23. i) Hydrogen
 ii) Calcium hydroxide produced ionizes partially producing few (OH-) ions
 iii) Test the presence of carbon (IV) Oxide.
24. a) 2Na(s) + 2H2O(l)  2NaOH(aq)   + H2(g)
b) Sodium melts to form a silvery ball.  Float on the surface and dart about.  Hissing sound produced.
25. Na2CO3   X H2O
 % mass 36.8  63.2
 RMM  10.6  18
Moles  0.347  3.5
Ratio  1  10
X = 10: Na2CO3: 10H2O
26. a) -Lead oxide changes from yellow to brown when heated and finally grey
 shiny solid is formed.
-Anhydrous cobalt chloride changes from blue to pink.
b) H2(g) + PbO(s)   Pb(s) + H2O(g)
27. C, E, B, D
28. i) 2Li(s)   + H2O(g)  LiO(s)  + H2(g)
 ii) Potassium is very reactive and the reaction may be violent/explosive.
29. (i)  If the hydrogen gas is not removed from the system it will reduce the
oxide of iron.
 (ii) Weather balloon
  Welding
  Rocket fuel together with oxygen.
30. Hydrogen is more reactive than metal W since it is able to displace “W” from its oxide.
31. a) i) Sodium
  ii) Copper
 b) i) sodium hydroxide/alkaline solution
  ii) 2Na (s) +   2H2O (I)       2NaOH (aq)   + H2 (g)
 c) Sodium hydroxide is a strong alkali with a pH of 14.
This is because it ionizes completely in solution of water producing more hydroxyl (OH-) ions.
d) Pottassium and Rubidium
e) Burns with a “pop” sound.
32. Deliquescent aborbs water from the atmosephere and dissolves.
Hygroscopic absorbs water from the atmosphere and becomes fissed i.e.  it will float and helps in spreading of fire.

 
FORM TWO
TOPIC 1
STRUCTURES OF THE ATOMS AND THE PERIODIC TABLE
1. Proton = 27
Neutrons =32
Electrons = 27
2. a) X= 2:8::3
  Y= 2:8:6
 b) X2Y3
3. Hydrogen can gain one electron when combined with electronegative element to form H.  Hence behave like group seven elements can also lose one electron to form H+ i.e, behave like group one element.
4. a) Period 3
 b) Y-3
c) The ionic radius of Y is greater than its atomic radius Y reacts by gaining three electrons.  The electrons added increases the repulsion / screening effect between the adjacent energy levels.
 5. 
6. a) i) F,  (ii) i
 b) J is in-group VI, period 3
7. a) K+ has many electrons thus many energy levels.  Na+ has few number of
electrons and thus few energy levels.
b) Mg2+ contain large number of protons compared to Na+ i.e the effective nuclear charge of Mg+2 ions is high, thus results into strong force of attraction between the nucleus and the electrons in their energy levels.  Hence they are pulled close to the nucleus.
8. H(g) + e   H(g)       ∆H -ve
 H e + H+      ∆H = +ve
9. -Coinage, ornaments, soldering
 -Making, plumbing joints/musical instruments casing for bullets and bombs.
10. a) C and E contain equal numbers of protons/ atomic numbers.
 b) (I) Neutrons in b = 4
(II)   First ionization energy decreases with increase in atomic radius. When atomic radius increases the outermost electrons get further from the nucleus, less energy is thus required to remove it.
11 RAM = (62.93 x 69.09) + (64.93x 30.91) = 63.54
    100
12. Across the period there is a gradual increase in number of proteins in the nucleus.
 This increases the force as attraction between the nucleus and the electrons.
13. a) They are both metals and need to lose electrons to be stable 
 b) RCO3 (s)     RO (s)   +        CO2 (g)
 c) Q -3
14. Atoms of the same element s with the same atomic numbers but different mass
numbers.
15. a)
          
        R S 
N Q     V   T U
P          
          
 b)     U
 c) Q(s)   +   T2(g)       QT2(s) Or Mg(s)  + Cl2(g)     MgCl2(s)
16. a) T 2:8:2
  U 2:8:3
  V 2:8:4
  W 2:8:5
  X 2:8:6
  Y 2:8:7
 b) Perion 3, they all contain three energy levels,
 c) X has a small atomic radius compared to V. X has more protons so
nuclear charge is higher hence attract outmost electrons more strongly.
 d) U W
e) Ionic bond/ electrovalent bond “T” will react with “Y” by donating its outer most electrons to the atoms of “Y”
f) T2+, T+2, T
 (G) X -2 because it has a stable electronic arrangement of 2:8:8 X +2 has unsuitable electronic arrangement o (2:8:4)
h) i) Acidic oxide VO2, W2O 3    XO
 i) Basic Oxide TO
17. a) C =6
  H=1
  Na=11
  Ne= 20
 b) Ca2+  2:8:8
  P-3  2:8:8
 c) -259 + 273 = 14k
 d) Red phosphorous because it has a higher melting point.
 e) The one atomic number 24, because it is closer to the relative atomic mass
(24.3), that means that it contribute to RAM more than the other two.
 f) Al4C3
18. (i) Alkaline earth metals. 
 (ii) A: It has a stable electronic arrangement (duplet)
 (iii) Covalent bond.  This because electrons are shared between B and E.
 (iv) G belong to group V, period 3
19. a) i) Alkaline metals
  ii) Energy required to remove an electron from an atom
iii) “P” has the smallest ionic radius therefore, the outermost electrons are most strongly attracted to the nucleus, hence more energy is required to remove this electron.
iv) Melts because the reaction is exothermic.  Hissing sound because of the production of hydrogen gas.  Float because it is less dense that water.  Moves about due to propelling effect of escaping hydrogen.
b) A strong base ionizes completely in water producing more OH- ions e.g KOH and NaOH.  A weak base ionizes slightly producing few OH ions e.g NH4 OH, Ca (OH)2 and Mg (OH)2
c) i) Reaction between H+  ions from the acid and OH ions from bases
to form 1 mole  of water.
H+(aq) + OH (aq)   H2O(l)
              ii) Add 200cm3 of nitric acid to 200cm3 of 2m sodium hydroxide. 
Heat the mixture so as to make it saturated /concentrated.  Allow the mixture to cool for crystals to appear.  Filter/decant to obtain the crystals to appear.  Filter /decant to obtain the crystals.
  iii) NaNo3(s)       NaNO2(aq) + O2(g)
20. a)  i)

  ii)  Non metals
 b) i) KA// KBr//KI any one
  ii) Ionic/ electrovalent: “K” loss of electron to form K+ ions.  “A”
gains electrons to form A ions.  Two ions combine to give KA.
 c) Add strong alkaline solution KOH //NaOH to Magnesium Sulphate
solution to precipitate Mg (OH)2.(s).   Filter the filtrate to remove water.  The residue is magnesium Hydroxide.  Heat the hydroxide to remove water.
Or
Add soluble carbonate or hydrogen carbonate to the mixture.  Magnesium carbonate will be formed.   Heat the carbonate to get magnesium oxide.
d) Al (OH)3(aq)    + 3H+(aq)  Al3+(aq) + 3H2O(i)
  Al (OH)3 (s)     + OH-(aq)    Al(OH)  4(aq)
21. Add aqueous sodium carbonate to precipitate calcium carbonate and magnesium carbonate and then filter to obtain pure brine.
22. a) Na+ ions contain few electrons compared to K+ which has large number of
electrons.  Na+ has few energy levels.
b) The ionic radius decreases from Na+ Alst  .  This is because there is gradual increase in numbers of protons in the nucleus.  The added proton increases the attraction force between the nucleus and electrons.
23. a) W =Fe
  X =Na
  Y =Mg 
  Z =Ca
 b) X, Z, Y, W
24.

25. a) “G” it requires less ionization energy to pull out first electrons.
 b) Metallic group: atomic radius is large that the ionic radius.
26. a) i)     “T” gain either react by gaining or loosing electrons depending on
the electro negativity of the element it is reacted with.
  ii) Alkali metals
 iii) “Y” is unreactive because it has stable electron arrangement i.e
octet structure.
 b) i)    “Y” has a small atomic radius compared to X.  Y has many number
of  protons in its nucleus hence attract electrons very strongly towards the nucleus.
ii) “V” has a small atomic radius compare to “W”.  It can pull electrons to be gained very strongly i.e it has more electronegative.  W can only react by sharing electrons.
c) i) 
ii) WT4 is non polar molecule hence cannot dissolve in wate.  It exist in form of simple molecular structure hence melting point is low.
 
d) i)

 ii) (35x3) + (37x1)     = 35.5
   4
27. a) Isotopes refer to atoms of the same element with the same atomic numbers
but different mass numbers.
  (36 x0.34) + (38x0.06) + (40x 99.6) = 39.88
    100
28. a) 33-18=15
 b) Z(g) + 3e      Z-3(g)
29. a) i) Hydrogen has (H2)
  ii) Iron (II) Sulphide (Fes)
  iii) Hydrogen Sulphate (H2S)
    b) i) Burns with a pop sound
  ii) Darken the paper which is soaked in lead acetate: (forms black
precipitate with lead (Pb2+) salts).
30. a) E= 2,8,5
 b) The chloride of E is in form of a simple molecular structure.  The force
holding the molecules together is weak van der waals forces.
31. (10x18.7)+ (11 x 81.3) = 10.81
  100

 
TOPIC 2
CHEMICAL FAMILIES: PATTERNS IN THE PERIODIC TABLES
1. a) X: it has a stable electronic arrangement i.e Octet structure.
 b) i) “W” and “Y”
  ii) YW
2. IV, II, I, III
3. a) T(s) + X+(aq)  T2+ (aq)  + X(s)
 b) S, X, T, U
4. a) i) B  ii) C  b) D
5. G 3, it has the smallest atomic size, therefore outermost electrons are strongly
attracted by the nucleus.  A lot of energy is required to remove the outermost electron.
6. Element A= Sulphur, carbon, nitrogen
 Element B = Sodium, potassium, lithium
7. a) F2O5 = O: 2F + - 10 = 0     : 2F = 0+10
  F= +10      = +5
           2
 b) Group V
8. The yellow liquid is pcl3.  It is hydrolised in air to form Hcl which fumes since it absorbs water vapour from the atmosphere.
9. Group (VII) elements react by gaining the electrons Flourine has the smallest atomic radius in this group hence it attract electrons very strongly hence it gain electrons very easily making it to be more reactive.  Ease of electron gain decreases down the group.
10. a) Solid CD does not conduct electricity since the ions are not free to move. 
The ions are held together by electrostatic force of attraction.
 b) Aqueous CD is a strong electrolyte since the ions are free to move.
11. a) The outermost electrons in mg and Al are delocalized and free to move
hence allow the flow of electric current.
 b) Alluminium forms a protective coating and prevents further corrosion.
12. a) “K” and “N” they are in the same group or same number of valency
electron/or they loose two electrons.
 b) L2O
 c) “L” it has 7 electrons in its outermost energy level/ react by gaining one
electron.  Its ionic radius is bigger than atomic radius.
 d) M; It has highest tendency to loose electrons.
 e) The ions of “N” have many protons in its nucleus compared to M.  The
protons in N nucleus pulls the electrons very close to its nucleus.
f) “L” gains electrons to form L- ion, the added electron increases the repulsion/screening effect between electrons in the adjacent energy levels.
13. a) i)  “S” and “W”
 b) i)   “V” it is the only element whose boiling point is below 298 kj at
room temperature.
ii) V has stable electron arrangement
            c) i) T (NO3)3
  ii) 2S(s) + U(s)      S2U(s)
 d) Ionic or electrovalent bond “T” is a metal while “U” is a non metal,
therefore T loses electrons to “U”
 e) i) Cathode- Hydrogen gas
  ii) Anode – Oxygen gas.
14. a) i) Greenish yellow gas
  ii) Slightly soluble
  iii) Grey/ black solid.
 b) (i)   4HCl(aq) + MnO2(s)   MnCl2(aq) + 2HO(l)  + Cl2(g)
  ii) Oxidizing agent.  It oxidizes the chloride ions to chlorine gas.
 c) i) Iron (iii) Chloride
  ii) Mass of chlorine used
 = 8.06 – 6.30 = 1.76g
    RM M of Cl2 = 71
 Moles of cl2 =   1.77 = 0.0248 moles
        71
 Volume of chlorine
 = 0.0248 x 24000 = 595.2 cm3
           d)
   H H
   │ │
  H   — C   — C   — H
   │ │
   Cl Cl
  1, 2 – dichloroethane
15. i) “A” It is in group (VI) and gaining two electrons .
ii) Giant ionic structure: C2O3 is an ionic compound.  This is a very strong force of attraction (electrostatic force) between the ions.
iii) “E” is more reactive than H.  “E” has a small atomic radii and gains electrons very easily compared to H.
iv) (I) B (s) + Cl2 (g)     BCl2(s)
            (II) Moles of Cl2 = 1.21 = 0.054  moles
                     22.4
 Moles  of B = 0.054 moles
 RAM of      = 1.3    = 24
          0.054
v) “G” has a small atomic radius compared to F. G has many protons and hence attracts electrons very easily to its nuleus.
 b) i) The oxide of B is alkaline in mature with a PH greater than (8.0). 
B is a metal and forms basic oxide.  D is a non metal and forms acidic oxide with a PH less than 5.0
  ii)   i)  I U
              II W
      ii) I X 
    II Y
16. i) Potassium – sodium – Lithium
ii) Lithium has a small atomic radius compared to the others.  The outermost electrons are attracted very strongly by the nucleus charges.
A lot of energy is required to pull out the outer most electrons. Atomic radius decreases from potassium to lithium.
17. a) 3 Mg(s)   +   N2(g)     →       Mg3N2
  Moles of magnesium = 8 = 0.333 moles
      24
  Moles of N2 = 0.333 = 0.111 moles
       3
  Volume of N2 = 0.111 x 22.4 = 2.488dm3
 b) Mg3N2(s) + 6H2O(l) 3mg (OH)2(aq) + 2MH3(g)
  Moles of Mg3N2 =0.111 moles
  Moles of NH3 = 0.111 x 2 = 0.222 moles.
  Volume of NH3 = 0.222 x 22.4 = 4.97dm3
18. a) Ca(s) + 2H2O(l)  Ca (OH)2 (aq)  + H2(g)
 b) Moles of Ca = 2 = 0.05 moles
    40
  Moles of H2   0.05 moles
  Volume of H2 =0.05 x 24000 =1,200 cm3
 c) Ca (OH)2 is slightly soluble in water
d) Sodium reacts with water very vigorously.  Reaction may end being explosive since sodium is very reactive.
e) 2H2(g) + O2(g)  2H2O
Moles of H2 = 0.05 moles
 Moles of H2O  0.05 moles
 RMM of H2O =18
 Mass of H2 = 0.005 x 18 = 0.9g
f) Calcium is a metal and the outer most electrons are delocalized/ free to move.
19. a) “W” it has the largest atomic radius.  The outermost electrons are loosely
held by the nucleus.  Less energy is required to remove this electron.
 b) V+ X V + X
  V + Na NaV
20. 2Mg(s)   +   N2(g)   2MgO(s)
 3Mg(s)   +   N2(g)   Mg3N2(s)
21. a) Grey precipitate of iodine will be observed.  Chlorine is more reactive
than iodine and it displaces it from its solution of sodium iodide.
  Cl2(s)   +   2I(aq)    I2(s)  +2Cl (aq)
 (b) Covalent bond both chlorine and iodine are non metals and react by
sharing electrons.
22. Elements in group (VIII) which have a big atomic radius can react under special
condition by losing electrons e.g xenon and fluorine- can react to give xenon hexafluoride
23. i) Reddish brown liwuid
ii) No change
iii) Chlorine is more reactive than bromine and can displace it from its salts solution, but chlorine can not displace it self.
24. The group (VIII) element reacts by gaining electrons.  The atomic radius
decreases down the group.  Atoms with small ionic radius gains electrons very easily.  Hence gain electrons (electro negativity) decreases with an increase in atomic radius.
 
TOPIC 3
STRUCTURES AND BONDING
1. M: Metallic bonding
             N: ionic/ electrovalent bonding
2. i) NH3
 ii)   NH4

b)  NH3 posses one pair of electrons which can be shared with H+ ion which has no electrons to be stable.
3. An hydrous aluminum chloride is a covalent compound while magnesium chloride is an ionic compound.
4. CO2
 H2O
5. a) D      b)E
6. a) In a diamond all the carbon atoms are joined together by strong covalent
bonds, a three dimensional structure and therefore it is very hard.
 b) The carbon graphite atoms are bonded in layers.  The layers are held
together by weak van der waals forces of attraction.  The layers therefore slide over each other very easily.
7.  H H H H
  │ │ │ │
 H   — C   — C   — C   — C   — H Butane
  │ │ │ │
  H H H H 
8. Ionic bond.  It involves elections transfer.
9. HCl(g) is covalent, it dissolves in methyl benzene but does not ionize. Addition of water causes HCl(g) to ionize since it is polar. H+ ions are liberated which react with carbonate to produce CO2
10. In solution, molten of fussed since the ions are free.
11. PCl3 has a simple molecular structure.  Molecules are held together by weak van der waals forces. MgCl2 has giant ionic structure. Ions are held together by strong electrostatic force of attraction/strong ionic bond.
12. Neon is inert and will prevent oxidation of the filaments.
13. Covalent bond exists between two iodine atoms in an iodine molecule.  It involves sharing of the electrons. Van der waals forces exist between two or more molecules of iodine.  It is a weak force while covalent is a strong bond.
14. I. Conduct
 II. Ionic
 III. Covalent
15. a) The amount of heat absorbed by a mole of substance to change from liquid
state to gaseous state without changing the temperature of the surrounding
b) Boiling points increases with increase in molecular mass or increase in number of carbon atoms.
16 a)

 b)
17. Each carbon atom is bonded to other atom by covalent bond to form hexagonal
layers.  The layers are held together by weak van der waals forces.  The layers can slide over each other easily.
18 a) Covalent bond involves sharing of electrons between two or more atoms. 
Each atom contributes equal number of electrons to be share.  In co-ordinate bonding, the shared electrons are contributed by one partical in a molecule.  The products of covalent bonding are neutral molecules but in co-ordinate bonding by the products are charged.
ii) NH4

b)     or
19. Ethanol is a polar molecule; two forces van der Waals and hydrogen bonding
holds the molecules together.  Hexane is non-polar and only weak van der waals forces hold these molecules together.
20. a) Group (VII) elements
 b) Chlorine molecule is smaller and the strength of vanderwaals forces
between molecules of chlorine is weak as compared to iodine.
21. a) Metallic bonding
b) “C” it can gain electrons very easily since it has a small atomic radius.  It is very electronegative.
22. a) Ionic bonding/ Electrovalent bond
b) “C” it can gain electrons very easily since it has a small atomic radius.  It is very electronegative.
23. a) M 2:8
  C 2:8:8
 b) i) C   ii) N and C
 c)   Period 4
d)    “R” has a large atomic radius that “L”. The outermost electrons in “R” are not held tightly its nucleus.
e)

24. i) M= Graphite
  N = Diamond
 ii) Jewellary: drilling rocks, glass cutters
 iii) M/Graphite, the fourth electron is delocalized – in each carbon atom
25 a) i)

  ii)   The ions are not free at 250C since the salt is in solid state but
between but between 80 10C and 14130C the ions are free since electrostatic forces between the ions is overcome.
b) Ammonia reacts with eater to form ammonia solution.
c) Dative/co-ordinate bond
26. (i) period “2” it has two energy levels.
 (ii) I.  Across the period atomic radius decreases.  In A2 there is more
positive nuclear charge than in A1 hence elections are more pulled to the nuclear hence reduced size.
III.    A4 reacts by gaining electrons.  Then added electrons increases the
repulsion effect between the energy levels.
 (iii)

27. The energy used in bond breaking is higher than energy released when new bonds
are formed.
28. Water is a polar compound, two forces i.e van der waals forces and hydrogen bond held the molecules of water together.  Hydrogen sulphide molecule is non polar and the molecules are held together by weak van der waals forces.
29.
30. i) Electrons are transferred
 ii) Electrons are shared equally
31. i) sodium metal atoms has delocalized electrons in its outermost energy
level.  No ions in sodium solid metal
 ii) Iodine is a covalent substance, no free electrons or ions.
 iii) Sodium solid iodide has no free ions in solid state but when in solution
the ions are free.
32. Giant ionic structure: The compounds contain ions which are held together by strong electrostatic forces of attraction.
33. a) Cacl2: It has high melting point and requires a lot of energy to vaporize it.
 b) Simple molecular structure
 c) Ethanoal is polar with two forces van der Waals and hydrogen bonds
holding the molecules.  Carbon disulphide is non polar, only van der waals forces holds its molecules together.
34. a) i) U: conduct both in solid and liquid state
  ii) W
 b) i) V      (ii) Y
35. The molecule of methane is small hence the van der waals forces between molecule is weak. Hexane molecule is bulky with strong van der waals forces between molecules.
36. a)  G  (b) E
 
TOPIC 4
SALTS
1. a) 190C to 19.5 0C
 b) Place 80g of KNO3 in 100g of water and heat up to 500C.
 c) All the solid would dissolve because the solubility of calcium ethanoate
increases with decrease in temperature.
2. W: Because its solubility decreases with increases in temperature.
3. Dissolve K2SO4 in water; dissolve pbCO3(s) in nitric acid.  Mix the two solution
and filter to remove solid PbSo4
4. Add water to the solid mixture and stir.  “A” ddissolves while “B” does not. 
Filter the mixture and evaporate to dryness.
5. a) i) Dilute Nitric acid
ii) Lead (II) Sulphate (PbSO4)
 b) Pb (OH)2(s)  + 2OH-(aq)   Pb(HO)4-2(aq)
6. Crystals will be formed.  This is because the solubility of this substance decreases
with increase in temperature.
7. Crystals of KClO3. Cooling causes crystallization.  All solution is not yet
saturated in the solution because at 400C the solution is not yet saturated with
KNO3.
8. a) Potassium chloride
 b) Calcium chloride
 c) Lead (II) nitrate
9. - Making baking powder.
 - Treatment of stomach acidity
 - Health salts
 - Laxatives
 - Fire extinguishers.
 - Soft drinks
10. a) React MgO with Nitric acid to get Mg(NO3)2(aq). Add strong alkaline
solution e.g KOH / NaOH to precipitate Mg (OH)2.   Filter the mixture to get solid Mg (OH)2
 b) - In toothpaste
  - Neutralize acid in stomach (anti acid).
11. a) Cone sulphuric acid
 b) Cooling the concentrated solution to get crystals
 c) Anhydrous copper (II) Sulphate.
12. a) i) Deliquescency
13. React sodium with water to get sodium hydroxide. Bubbles into this solution
excess carbon (IV) oxide to get sodium hydrogen carbonate.
14. React copper with conc nitric acid to get copper nitrate solution.  Heat the
solution to dryness. Cu(NO3) decompose to give CuO. React CUO with dilute HCl to get CuCl2.  Filter and concentrate the solution to get crystals.s
 
15. a) i) Heating
  ii)

  iii) Zn2+(aq) + 4NH3(g)   zn(NH3)4+2(aq)
  iv) Brown gas/Fumes
v) Addition of anhydrous copper (II) sulphate.  It changes from white to blue or odd drops to anhydrous cobalt (II) chloride.  It changes from blue to pink.
b) i)  One of the salt R is insoluble in water because a residue is formed
when water is added.
 ii)  CO3 -2+ it react with acid to give CO2
 iii) Pb2+
c) Zinc nitrate and lead carbonate.
16. a) i) Hygroscopy/ hygroscopic
 ii) Deliquexscence
 iii) Efflorescence
b) i) Zn (OH) 4 -2
 ii) Cu(NH3) 4+2
c) i) Fe  S  O H2O
  20  11.5      23.0 45.3
  56  32  16 18
   0.36  0.36  1.44 2.52
   0.36  0.36  0.36 0.36
   1  1  4 7
  ii) FeSO4 : 7H2O
  iii) Moles of salt = 6.95 = 0.025 moles
       278
 Cone in moles /dm3 = 0.25 x 1000 = 0.1 M
          250
 17. i) Cu(s)      2e- + Cu2+  (aq)
             ii) Q= It = 0.2 x 5 x 60 x 60 = 3,600 c
  Loss in mass Cu  = 3600 x 64 = 1.19g
                965000 x2
 
18. i)

 ii) Moles of CuSO4 deposited = (0.185- 0.12) = 0.065
  Mass of CuSO4:  5H2O : RMM = 250
  = 0.065 x 250 = 16.25
 b) i) Moles of AgNO3 in 250cm 3  of solution
   = 0.1 x 24.1 = 0.0241 moles
         1000
         ii) Moles of NaCl in 25cm3
   Mole ratio 1:1
   = 0.00241 moles.
  iii) Moles of sodidum chloride in 250cm3
   = 0.00241 x 250 =0.00241 moles
    25
 iv) RMM of NaCl= 58.5
  Mass = 0.0241 x 58.5 = 1.41
 v) Mass of H2O = 5.35 - 1.41= 3.94g
 vi) 3.94g of water contain 1.41g of Nacl 100g of water containe 
  1.4 x 100= 35.79g = 35.79g
 3.94
 
19. a) Solution which cannot dissolve any more solute at aparticular temperature
 b) i)      

  ii) I) I 25g/100g of water
   II) Mass dissolved = 62g
Mass undissolved = 80-62 = 18g
20. a) i)
  ii) 71g/100g of water.
 iii) I.  A solution which has dissolved a lot of solute till it can
dissolve no more.
 II. Mass of solution at 250C = 100 + 71= 171g
  Mass in (g) = 100 x 71 = 41.52(g)
       171
21. a) i) Zn2+(aq)
  ii) Zn(OH)2(s)
 b) It is amphoteric
22. a) i) Iron (II) Suphide
  ii) Hydrogen Sulphide
 b) Darker paper soaked in lead acetate
23. a) BaSO3
 b) 2Hcl(aq) + BaSO3(s)   BaCl2(aq)  + SO2(g)   + H2O(l)
24. a) Pb2+(aq)  + SO4-2(aq)  PbSO4(s)
  Moles of Pb2+ salts  0.63 = 0.003 moles
     207
  RAM of PbSO4 = 303
  Mass of PbSO4 =303 x 0.003 = 0.91g.
25. a) Fe3+(aq)
 b) Oxidizing agent
 c) 2Fe(OH)3(s)         heat         Fe2O3(s)    + 3H2 O(l)
26. a) Zn
 b) Zn(NH3)+2
27. Dissolve lead carbonate in dilute nitric acid. React the mixture with dilute hydrochloric acid. Filter to get lead (II) chloride.
28. Sodium hydroxide is deliquescent. It absorbs water vapour from atmosphere and dissolves the solution formed (NaOH) absorbs CO2 to form Na2CO3 and H2O.  H2O.  H2O evaporate to leave a white solid of Na2 CO3.
29. Colour change from blue to white powder. Vapour which changes anhydrous cobalt chloride from blue to pick produced.
30. a) Brown precipitate of iron (III) hydroxide. Chlorine Fe3+ to Fe3+ ions.
b) It will dissolve to form a clear solution. Ammonia reacts with silver chlorine to give a complex salt.
 
TOPIC 5
CARBON AND SOME OF ITS COMPOUNDS
1. - Dense than air
 - Does not burn
 - Put off burning flame
2. Ca(OH)2 produces CaCO3 which is insoluble . NaOH forms Na2Co3 which is soluble.
3. SO2;   It is an acidic gas and react with Ca (OH)2 which is basic.
4. Equilibrium shift to the left to reduce the pressure.
5. Add water and stir.  Sodium carbonate will dissolve.  Filter to get lead carbonate as a residue.
6. Kerosene is less dense and float spreading the fire.  CO2 is more dense and covers the fire preventing oxygen reaching the fire.
7. K+/ Na+    and CO3-2
8. i) C(s)  + O2(s)    CO2(g)
 ii) CO2(g) + C(s)    2CO(s)
9. a) PBO(s)  + CO(s)  Pb(s) + CO2(g)
b) Colour of PbO change from yellow when cold, brown when hot, Finally grey.
c) Hydrogen gas
10. a) ammonia gas
 b) Filtration/precipitation/crystallization
 c) 2NaHCO3(s)     Na2CO3(s)  + CO2(g)  + H2O(l)
11. a) H = CaCO3
  J= CaO
 b) - Fertilizer/for liming/ making motar
  - Rising soil PH
12. Luminous   Non luminous
 -Sooty flame   Non-sooty
 -Produce more light  Less light
 -Less heat   very hot
 -Weavy flame   Stead flame
13. a) colour of solid change from black to reddish brown.
 b) CUO(s) + CO(g)     →       Cu(s)  + CO2(g)
 c) CO is poisonous gas
14. a) CO2(g)   Ca (OH)2(aq)     →        CaCO3(s) + H2O(l)
 b) White precipitate dissolves because Ca (HCO3)2 formed is soluble
15. Moles of Hcl = 20    = 0.02 moles
              1000
 Moles of GCO3    1   = 100
        0.01
 RAM of G = 100 – 60 = 40
16. a) To reduce PbO to Pb
 b) To remove silica as slag
 c) To reduce unreacted PbO to Pb
17. Equilibrium shift to the right to replace CO2 which is removed.
18. C(s)    +   O2(g)   → CO2(g)
 CO2(g)   +   C(s) → 2CO(g)
 FeO2(s)   +   CO(g)         → CO2(g)  + Fe(s)
19. a) Reduction;  Oxygen is removed
 b) Oxygen is removed/ oxidation state of Pb change from +2 to O.
 c) Ammonia gas/ Hydrogen gas
20. a) 2H2SO4(aq) + C(s)      →        CO2(g)          + 2SO2(g)   + 2H2O(l) 
b) Oxidation No: of S in SO2
+2   +   S    +    8=0
S = + 8 -2 = +6
Oxidation No:  of S in SO2
SO2 = 0
S + -4 =0
S= -+4
Change in oxidation from + 6    → +4 (reduction)
21. Sublimation
22. a) Cone:  Sulphuric acid and Ethanoic acid.
 b) C2H2O4(aq)           H2SO4 CO2(g)  +  CO(g)  + H2O(g)
 c) It is colourless and odourless.
23. a) Carbon (IV)Oxide
 b) Blue flame, carbon (II) oxide is burning
24. It is more dense than air
 It will react with calcium oxide since CO2 is acidic and CaO is basic.
25. a) The calcium and magnesium compounds in this water can not be
decomposed by heating i.e. Cacl2, CaSo4, MgSO4  MgSo4 and MgCL2
26. a) I. Pb2+   II. CO 3-2
 b) PbO(s) + 2H+ (aq)    Pb(s) + H2O(l)
27. a) i) Galenas
  ii) Some of the substance/ sulphide is converted with PbO or SO2.
  iii) Carbon (II) Oxide
  iv) PbO(s) + C(s)                  Pb(s)  + CO(g)
  v) SO2 is poisonous/ cause acidic rain or CO poisonous/ pb2+ also
poisonous.
b) Hard water contain Mg2+/ Ca2+ which form pbSO4 insoluble and form a protedine layer/soft water does not form these deposits/
c) - Radio active shielding
 - Alloys
28. i)  C(s) + O2(s)   →  CO2(g)
 ii)   KOH
 iii) Pass the gases through Ca(OH)2.  CO2 forms white precipitate but CO
does not..
 iv) Fuel in water gas and producer gas/ extraction of metals.
29. i) Step 2      →    CO2(g)
  Step 4 Dilute HCl.
 ii) Ca(HCO3)                      CaCO3 + H2O(l)  + CO2(g)
iii) Add H2SO4, add NaSO4/ K2SO4 filter to obtain CaSO4 as a residue. Heat the residue to dryness.
30. a) i) Allotropes
  ii) Add salt to methylbenzene, fullerence dissolves. Filter the mixture
to remove the residue.  Heat the filtrate to make it concentrated cool the solution slowly to get crystals
 iii) 12n= 720: n= 720=60
    12
  M.F = C60
31. Petrol is less dense float and stread fire
32. CUO(s) + CO(g)   →  CO2(g)  +   H2O(l)
33. a) i) - Decomposition of CaCO3 in S
   - Filtration
  ii) - Drying agent
  iii) NH3(g) + H2O(l)  + NaCl(aq) + NaHCO3(s)  + NH4Cl(aq)
 b) CaCO4(aq)   + Na2CO3(aq)        CaCO(3)(s)  + Na2 SO4(aq)
 c) Making baking powder.
 d) i) Na2CO3  + 2HCl           CO2 + 2NaCl  + H2O
  ii) Moles of CO2 = 672    = 0.03
       22400
   Moles of HCl = 0.03 x2 = 0.006 moles
   Conc of HCL = 0.006 x 1000= 1.0 m
      30
                             Value of x moles of Na2Co3   0.03
   
   X(mass) = 0.006x 1000 = 1.0
       30
  iii)    Value of x moles of Na2CO3 =106
   RMM of Na2CO3 =106
   X(mass) = 0.03 x 106 = 3.18g
34. a) Hardness caused by soluble Ca2+ + /mg2+
  HCO2- salts can be removed by warming.
 b) Hardeness caused by soluble CaSO4 cannot be removed by warming.
    i) 

  ii) Contain Ca2+ (aq)     + Na+p(s)     →     Na +(aq) + Cap (s)
35. CuCO3
 CuCO3(s)                 CuO (s)      +   CO2(g)
36. i) Burns with blue flame to give a gas which form white ppt with lime water.
 ii)  Forms white ppt with lime water.
37. Forms a coat of CaSO4 which prevent further reaction CaCl2 is soluble.
38. i) Combustion/ decay
 ii) Photosynthesis/ marine animals/ dissolve in water.
39.   C   H  O
 40   6.67  46.67
 12     1    16
 3.33   6.67  2.91
 3.33   3.33  3.33
 1     2    1
      CH2O
40. i) N= CO2
 ii) N is slightly soluble in water.
  PH decreases/ acidic NO2 dissolves in water to for HNO2.

 
FORM THREE WORK
TOPIC 1
GAS LAWS
1. Kinetic energy of the gas increases, and gas molecules moves faster.  The space
between them increases.
2. “Q” it diffuses more slowly i.e, it covered a short distance
3. Hydrogen; it is less dense than Coz and diffuses faster
4. Air is less dense than carbon (IV) oxide and so it enters the porous pot faster than
carbon (IV) oxide and so it enters the porous pot faster than carbon (IV) oxide leave out of it.
This creates a high pressure in the pot and the level of eater rises up as shown.
5. V1  = V2
 T1 T2
V2 = V1T2
    T1
T2  =250 x 315 = 262.5 cm 3
        300
6. P1V1 =   P2V2    T2= T1,P2V2
    T1        T2                   P1V1
            T2    =  500x 0.5 x 100 = 62.5K
   1x 400
7. P1   =  P2     P2 =  P1 T2
 T1 T2  T1
     P2 = 760 x 373  = 1038.39 mmHg
            273
8. Rmm of O3  = 16x 3 = 48
 Rmm of CO2 = 12+ 36 =44
 TCO2   =     44
  96                48
           Tcoz = 96x 6.63 =  91.9 seconds
               6.92
9. The entire solution turns pink/purple.
 - potassium permanganate/potassium manganate (VII) particles diffused into the water molecules.
10. a) The volume of a fixed mass of gas is inversely proportional to the 
pressure at constant temperature.
 b) P1V1 = P2V2
  V2 = P1V1:  V2 = 3 x 1 = 1.5 lts
     P2                   2
 
11. Mass due to C = 12   x   4.2  =  1.145(g)
 Mass due to H =  2  x 1.171  = 0.1899
       18
Elements C H
Mass      1.145
       12   0.1899
     1
Ram  
Moles  
Mole ratio  1 2
EF   =CH2
12 TO2  =   RMM O2
                          RMMSO2
 TSO2   =  50 x 64     =  TO2   x 70.7  seconds
                                    32
13. a) The volume of a fixed mass of gas is directly proportional to its
temperature in Kelvin.
 b) P1V1    =   P2V2
               T1          T2
      T2  =  291 x (1.0 x 105 )   x 2.8 x 10-2 =  2328k
                     (1.0 x 10 5 ) x (3.5 x 10 -2)
14. Purple/ pink particles spread to form a uniform solution; particles of water have
k.e they collide and disintegrate the particles of KMNO4 .  Diffusion takes place.
15. a) rate of diffusion is directly proportional to the square root of the density.
 b) Row =  RMMX    =   12 =    RMMA
  ROX     RMMO2       X        16
  X = 12 x 4   =  7.2365cms-1
             6.633
16. TA  =  RMMMA     24     =    RMMA
 TO2     RMMO2           20              32
       RMMA = 46
17. TCO2  =  RMMco2   200  =       44
 THCL     RMMCCL     THCL      36.5
                 THCL= 36.5 x 200 = 18.2 158 secs
18.       Moles of CO+2+ = 11  = 0.25 moles
                     44
 Moles of butane =   2x0.25 = 0.0625 moles
 Volume of butane =  0.0625 x 24 = 1.5 litres
19. P is less dense than air, so it diffuse into the porous pot fast compared to the rate
at which air moves out of the pot.  This increases the pressure in the porous pot and water rises as shown.  Q is more dense than air, hence a lot of air diffuses out of the porous pot compared to the amount of Q moving in.  This reduces the pressure inside the porous pot and atmospheric pressure forces water to vacuum left in the porous pot.
20. i) White deposit/ white slid/white fumes
 ii) Position of formation; Nearer the HCL side since NH3 is less dense and
diffuse faster compared to HCl
            iii)   NH3(g)  + HCl (g)    → NH4CL(s)
21. Rate “K”     =  RMMH2
 Rate  +H2”  RMM K
 Rate of K =  88  = 2.2 cm3/ sec
           40
 Rate H2 = 50  =  10cm3/sec
 
RAMK =(  10 x  2)
                      2.2
Rmm ‘K’ = 2x100 = 41.322
        (2.2)2
 22. NH3  +  O2   →    NO  + H2O
 Vol:  300 250          200      300
 Ratio : 4:    5:        4:     6:
Equation:
4NH3(g)     + 5H2O(g)      →       4NO(g)   +   6 H 2O(l)
23. 

Pressure
  Volume
24. V1   =   V2        T2   =    V2  T1
             T1        T2                           V1
 T2 =  300x 298(k)    = 447k
     200
25. P1V1    =   P2V2
   T1              T2
But P1 = P2
 V1 =  V2   V2 =  V1x T2
 T1       T2                  T1
       V2=    200x243  =   196.6cm3
  298
26.         VI   =  V2    T2   = V2T1
       T1        T2                   V1
T2 =   160x 298   =   238  238. 4k
 2000
Temp in 0C =  238.4- 237  =  -34.60C
27. P1V1 = P2V2  T2 = P2V2
 T1          T2                  P1V1
           T2= 800 x 190x 301k =  1.82. 4k
 760 x 330 
 
TOPIC 2
THE MOLES FORMLAE AND CHEMICAL EQUATIONS
1.      Mass of H2O = 34.8 -15.9 =  18.9 9(g)
 Components      Na2CO3:   X  H2O
 Mass    15.9g  18.9g
 Rmm    106  18
 Moles     15.9  18.9
     106  18
 Moles    15.9  18.9
     106  18
     0.15  1
 Ratio                                1        : 7
     X= 7
2.        2 moles of H2 react with 1 mole of O2
            100cm3 of H2 willreact with 50cm3 of O2
             O2 is in excess by 50cm3
3. 1 mole of CaCo3  react with 2 mole of Hcl
 0.1 mole caCo3 react with 0.2mole of Hcl
 Rmm CaCo3  = 40 + 12 + 48 = 100
 Moles of CaCO3 =  15 =0.15 moles
                                                       100
 Excess moles of CaCO3 = 0.15 – 0.1 = 0.05 moles
 Excess mass of caCO3     = 0.05 x 100= 5g
4. a) (C3H6O)n = 116
  (3x12) + (6x1) + 16) n = 116
  58n=  116:  n =2
  MF= (C3H6O)2   →  C6H12O2
 b) 12x6x100=  62.07 %
   116
5.  a) H2S(g) :- It adds (H) to Cl2  and reduce it to HCl . or the oxidation number
of cl2 reduced from O to -1
           b) Theoretical yield of H2S= 100 x 2.4 = 3.2g
                                                                    75
   Moles of H2S =  moles of S: 3.2 = 0.1 moles
     32
6. i) (C2H3) n = 54
  27n    =  54
                                      n=   2  :   MF =  (C2H3)2  -C4H6
ii)  H H   H   H
  │ │   │   │
C   ≡    C   — C   — C   — H   or H   — C   — C   ≡ C   — C   — H
 │  │ │   │   │
 H  H H   H   H
H   H
  │   │
 or  C   ≡    C   —   C   ≡    C
  │ │ │ │
  H H H H
iii) Alkyne if it has – C  C- or akene if it has – C= C- depending with
structural formula.
b) i)  Ca(OH)2(aq)  +   2CaCo3(s)   + 2H2O(i)
  ii)   90x0.01 =  0.0009
                                       1000
c) It will form “scum” initially then produce lather after adding a lot of soap solution.  All the ca2+ ions must be precipitated before soap lathers. 
7. Moles of H2 = 10= 5 moles
               2
 Moles of No2  = 5moles
 RMM of NO2 = 46
 5 moles of NO2 = 5 x 46 = 230g
8. Mass of H =  12 x 3.52 = 0.96g
            44
            Mass of H  2 x 1.44 = 0.16g
             18
 Elements C    H
 Moles  0.96  = 0.08  0.16   = 0.16
     12                                 1
Mole ratio 1  : 2
       EF   = CH2
(Ch2)n = 56:  14n = 56:  n =  56 =4
          14
          MF=  (CH2)4                      C4H8
9. Molarity of NaOH =  4=  0.1m/dm3
                                                            40
Moles of NaoH in 20cm3   = 0.1x20  = 0.002 moles
     1000
            Mole ratio 2:1
 Moles of H2SO4  =  0.002 =0.001 moles
    2
 Molarity of H2 SO4  =  0.001 x 1000  = 0.125m
                                                           8
10. RMM of H2O = 2 + 16 = 18
 RMM NCl2CO3 = 46+12+48 =106
 Moles of H2O =  14.5=  0.805 moles
         18
 Moles of Ncl2 CO3  85.5 = 0.886
                                             100
           Mole ratio    1:1 N = 1: Na2CO3: H2O
11. a) H2SO4(aq)  + 2NaOHcl(aq)      →        Na2SO2(aq)     + 2H2O(l)
 b) Blue litmus paper turns red while red litmus remains red
 c) The acid is in excess
12. Na2SO3(s) + 2Hcl(aq)                  2Nacl(aq)     + 2H2O(aq)
             Moles  of SO2    =  960    =   0.04 moles
                                           2400
Mole ratio 1:1
Moles of Na2SO3 = 126
Mass of Na2SO3 = 0.04 x 126 = 5.04 =g
13. a)   Mg(No3)2 + (NH4)2Co3 → MgCo3 + 2NH4No3
  Mg2+ + Co3 → Mg Co3
 b) RMM of Mg Co3 = 84
  Moles of Mg Co3 = 8.4 = 0.1 moles
             84
  Mole ratio 1:1
  Moles of Mg(No3)2  in x cm3 = 0.1 moles
 X=  1000 x 0.1 = 200cm3
  0.5
14. Moles of HCl = 20x1 = 0.02 moles
      1000
             Moles of GCO3 =  0.02 moles
             1
 RMM of G =  1x1 = 100
    0.01
 G = 100 - 60 = 40 RAM of G = 40
15. Mass of water =94.5-51.3= 43.2
 RMM  of Ba (OH)2 = 171: RMM of H2O = 18
 Moles of Ba(OH)2  = 51.3 =0.3
    171
 Moles of H2O = 43.2 = 2.4
       18
 Moles of ratio is 1:8
 n = 8
 E.F =  Ba(OH)2 : 8H2O
16. Mass in 500cm3+ = 15x1.05= 15.75g
 Mass in 100cm3  =  15.75x2= 31.5g
 Molarity =  315   = 0.103m
          60
17. a)
Elements C H O
% 64.9
  12 21.6
  16 13.5
   1
Moles 5.41 1.35 13.5
Ratio 4 1 10
E.F = C4H9OH
b) H H H H
 │ │ │ │
H   — C   — C   — C   — C   — O   — H
 │ │ │ │
 H H H H
18. AL2 (SO4)3    3 SO4-2    +  2AL3+
 Moles a2Al2(SO4)3 = 6.84 = 0.02
             342
 Moles a2 SO4-2 = 0.02 x 3 = 0.06
19.  CaSO4   H2O
 2.485   0.33
  136           48
 0.0183   0.0183  CaSO4:  H2O
 1   1
20. Moles of Ca3(PO4) I  =  115  =  0.37096
       310
 Moles of H3PO4 = 0.37096 x 2 = 0.74192
 Mass = 0.74192 x 98 = 72.71kg
21. a) Mg(OH)(l) + 2HCl  MgCl2  + 2H2O
  Mole of HCl = 23x0.1 = 0.0023 moles
     1000
 Moles of mg(OH)I = 0.0023   =  0.00115
      2
 Mass = 0.00115 x 58 = 0.00667g
 b) 0.00667 =  100= 13.34%
       0.5
22. a)  Brine (sodium Chloride)
    b) i)   2NaOH + H2SO4   Na2SO4 + 2H2O
  ii) No of moles of  H2SO4 =   40 x 0.5 = 0.02 mole
                    1000
  iii) I. 1 mole of NaOH= 0.02 x 2= 0.04 moles
    1000 x 0.04= 0.4 moles/dm3
        100
 II. RMMofNaOH in 1 ltre = 40x 0.4 = 16g
  Mass of NaOH in 1 litre = 40 x 0.4 = 16g
  Unreacted substance (NaCl) = 17.6 – 16= 1.6g
23. i) Elements  Fe:  S:  O:  H2O
  % Mass  20:2  11.5  23.0  45.3    RAM      56  32  16  18    Moles  20.2             11.5  23.0  45.3      56     32     16    18
     0.3607  0.359  1.438  2.5167
Mole ratio 1     1        1      1
Formula FeSo4 : 7H2O = 278
   Molarity = g/dm3 =  27.8  =  0.1M
                            Rmm       278
24. MgCl2(aq)  + 2 AgNO3(aq)   2Agcl(s)  + mg(N)3)2(aq)
 Moles  of Mgcl2  = 1.9  = 0.02 moles
            95
 Moles of AgNo3 =  0.02 x 2 = 0.04 moles
 RMM of AgCO3= 170
 Mass of AgNO3 = 0.04 x 170 = 6.8g
25. Elements  Fe  0
 Mass   8.4  3.6
 RAM   56  16
 Moles   8.4  3.6
    56  16
 Mole ratio 0.15 0.225
       0.15             0.15
         1 : 1.5
 X2         2          :       3
     Formular = Fe2O3
26.   2KOH(aq)   +  H2SO4(aq)  +  H2O(aq)
        Moles of KOH = 24x0.1 = 0.0024 moles
   1000
      Moles of H2SO4= 0.0012x 100= 0.04m
      30 
27. i) 
25.0 25.0 25.0
 ii) Average 25.0+25.0= 25.0cm3 
                                              3
 iii) a) Moles of acid   0.1x 25 =  0.0025 moles
                   1000
       b) Moles of X2CO3     1x 13.8  =  138g/dm3
    0.1m
iv) Molarity of carbonate = 0.0025 x 1000 = 0.1
       25
v) Formula mass X2CO3    1x13.8= 138g/dm3
                   0.1M
vi) X2CO3  = 138
 2x + 12+ 48 = 138
 X= 138.60 =39
           2
28 i) Mass of iron = 12.66.10.98 = 1.68g
             Mass of oxygen 13.30-12.60= 0.64
  Elements Fe  O
        1.68 0.64
          56   16
 Moles  0.03  0.04
 Ratio   3    4
  Formula Fe3O4
 ii) Fe3O4(s)   +  4CO(g)  3fe + 4CO2(s)
29 a) Ma2CO(aq) + H2SO4(aq)  M2SO4(aq)   + CO2 + HsO(l)
 b) Molarity of the acid =  9.8  = 0.1 mole/dm3
                         98
 c) Moarity of carbonate
  Moles a2   acid reacting =   0.1x12.3 = 0.00123
      1000
  Moles of carbonating reacting = 0.00123 moles by mole ratio1:1
  Molarity = 0.00123 x 1000 = 0.0984m
      12.5
 d) Molar mass of the carbonate =     g/dm3+
           Molarity
 =  13.8    = 140
     0.0984
          e) M2CO3 = 140
  2M= 140-(12+ 48)
  2M =80
  M= 80/2= 40
  RAM of “M” = 40
30. 2Pb(NO3)2(s)    2 PbO(s)  + 4 NO2(g)  + O2(g)
  RMM Pb =223
 Therefore moles of PbO = 22.3  = 0.1 moles
            223
 Moles of Pb(NO3)2    = 0.1 mole from mole ratio
 Rmm Pb (NO3)2  = 331 x 0.1 mole from mole ratio
 Rmm Pb (NO3)2 =33
 Mass pf Pb(NO3)2
 Mass of Pb(NO3)2 = 331 x 0.1= 33.1g
31. i) Average volume of B
  24.1 + 24.0 + 24.0   =  24.03cm
   3
 ii) Moles of A in 20cm3
  Molarity = 48=1.2m
          40
  Therefore moles = 20 x1.2  = 0.024 moles
           1000
 iii) Moles of acid B = 0.024      = 0.012   moles
              2
  Molarity of B = 0.012 x 100 = 0.499M
        24.03
        iv) Formula mass of (COO)2 :  nH2O
       =  63 x 1  =126g
                0.499
  v) Value of n
   (COO02  nH2O  = 126
   24 + 64 + 4 + 18n = 126
   
n=   126-90     = 2
     18                     
                                               n= 2
32.   CUCO3   CuO + CO2
        1 mole CUCO3 gives 1 mole of CO2(g)
          1 mole of CO2 at stp occupies 22.4 dm3
33.    Moles of N2 gas =  360
          24000
 No:  of molecules = 360   x 6.0 x 1023
         24000
    = 9.0 x 10 21 molecules
34. i) Mono atomic gas: These are gases which exist as single independent
atoms e.g Helium (He), Neo (Ne) Argon (Ar)
 ii) Diatomic gas:  gases which exist as combined atoms where two atoms are
combined together to form a molecule e.g oxygen (O2) Chlorine (Cl2) Hydrogen  (H2)
 iii) Atomicity of element: number of atoms in one molecule of it e.g zone (O3)
has atomicity of three.  The molecule formed is a triatomic or triatomic.
 
TOPIC 3
ORGANIC CHEMESTRY I
1. a) Substitution Chlorination/Halogenation
 b) U.V light /sunlight
2. a) sulphur
 b) To harden it /make it tough /to strengthen it.
3. a) (RCOO)2Ca and (RC6H5SO3)2 is better since it is not affected by hard
water.
4. a) Butanol
   H H H H 
   │ │ │ │
  H   — C   — C   — C   — C   — OH
   │ │ │ │
   H H H H
 b) C4H9OH(aq) + 6O2(g)  4CO2 (g)   + 5H2 O(l)
5. a) Sisal/ cotton/wood/silk/jute/hemp/fur/hair
 b) -Their strength can be varied to make them stronger
  - Not easily affected by chemicals
  - They last longer
6. a) 2220- 1560 = 660
  1560-890= 670
  -2220 + -650 =-2870kj
 b) ∆ Hc of Alkanes is an exothermic process since the values are negative i.e
heat is released from reaction.
7. a)
   H H H H 
   │ │ │ │
  H   — C   — C   — C   — C   — OH
   │ │ │ │
   H H H H
  Butanol/Butan-1-ol
 b) 2C4H9OH(ag) + 2K(s)  2C4H9OK(ag)   +  H2(g)
8. Add solid NaHCO3, to both, CH3COOH produces effervescence and a colourless
gas which give white precipitate with lime water is produce No  reaction with CH3CH2 CH2OH.
9. Reaction 1: Carbon is oxidized fully to it highest oxidation state in Co2.
10. Monomer CH2 = CH
         │
       CN
 Rmm of monomer = (12x3) + 1x3) + 1x14) =53
 53n= 5194
 n= 5194  =98
         53
11. Pentane:  It is non poler and will not react with sodium Hydroxide solution which is an ionic compound.
12. Tetrachloro methane
   CL
   │
  CL — C   — CL
   │
   C
13. -In pentane there will be no reaction
 -In pentanol, three will be effervescence and a colourless gas which burn with a
“pop” sound produced solution last is alkaline.
14 a)       
   H H H H H 
   │ │ │ │ │
  H   — C   — C   — C   — C   ═ C Pentene
   │ │ │ │ │
   H H H H H
   H H H H H 
   │ │ │ │ │
  H   — C   — C   ═ C   — C   ═ C Pent – 2 – ene
   │ │ │ │ │
   H  H H H
 b) C2H2(g)    + 2 Cl2(g)   CHClCHCl2(aq)
15. Methane/CH4(g)  
 CH4(g)  + 202(g)       CO2(g)  + 2H2O (i)
16. a) U.V. light /sunlight
 b) Bonds broken C-H and Br-Br
  Bonds formed C-Br and H- Br
17. a) 
  H  H
  │  │
  C   ═ C   — C   — H
  │ │ │
  H H H
 b) Propene
 c) Petroleum/crude oil/natural gas
18. Add water to the mixture in a separating funnel, ethanol being polar dissolves
while pentane does not.  Allow the mixture to separate into two layers.  Open the tap to drain the lower layer which contain ethanol.  Distill the aqueous layer to get ethanol.
19. a) Reaction which one or more hydrogen in alkaene molecule is/are replaced
by halogens.
 b)  H H
   │ │
  H   — C   — C   — CL
   │ │
   H H
20. a) Butane
 b) Hardening of oil in manufacturer of margarine
21. Butene/but – 1- ene
22. a) Isomerism is the occurrence of tow or more compounds with the same
molecula formula, but different molecular structure/structural formular.
 b)
  H  H H  
  │  │ │ 
  C   ═ C   — C   — C   — H But – 1 – ene
  │ │ │ │ 
  H H H H 
   H   H  
   │   │ 
  H   — C   — C   ═ C   — C    But – 2 – ene
   │ │ │ │ 
   H H H H 
   H  H  
   │  │ 
  H   — C   — C   ═ C     2 Methyl propene
   │ │ │ 
   H CH3 H 
23. Thermal cracking
24.  H H H H 
  │ │ │ │
H   — C   — C   — C   — C   — H
  │ │ │ │
  H H H H 
 Butane
H H H H 
  │ │ │ │
H   — C   — C   — C   — H
  │ │ │ 
  H CH3 H 
 2 Methyl propene
25.     (a)  C13 H 27COO- Na
(b)  Soapy detergent
 (c)  (CH3) (CH2)12  COO)2 Ca2+
       (CH3) (CH2)12 COO )2Mg2+
26.  (i)  C2 H4 O2 it melting point is  higher than 100 C
       (ii) CH14  and    C5 H12
C6 H14  has a higher melting point since it is more bulky compared to  C5 H12; hence the vanderwaals force between the molecules of C6 H14  is abit strong.
iii) C3H8O is more soluble in water than C5H12: because it forms hydrogen bonds with water molecules i.e it is polar due to the presence of (-OH) group.
      b) i) C4H8
  ii)   C4H8(g)    +   6O2(g)     →         4O2(g)         +   4 H2O(l)
c) i)
   H H H H H H  
   │ │ │ │ │ │
  H   — C   — C   — C   — C   — C   — C   — OH
   │ │ │ │ │ │
   H H H H H H 
  Reagents
  ii).   –  Concentrated sulphuric acid
- Al2O3 or phosphoric acid (Catalyst)
Conditions
Heat (160-1800C)
  d) i) Saponification/Hydrolysis
  ii) Fats/ ester
27. a) i) Butan-1 01
  ii) Propanoic acid
  iii) C5H10
   H H H H H  
   │ │ │ │ │ 
  H   — C   — C   — C   — C   — C   — H
   │ │ │ │ │ 
   H H H H H  
  or
   H   H H  
   │   │ │ 
  H   — C   — C   ═ C   — C   — C   — H
   │ │ │ │ │ 
   H H H H H  
28. a) i) Additional polymerization
  ii) Substitution reaction/chlorination
 b) i) Fractional distillation
  ii) Sink to the botton: effervescence/fizzing sound as hydrogen gas is
produced
  iii) -In thermometers
   -Fuel
   -Mild disinfectant
   -Solvent
 
 
c) i) C4H8
   H H H H   
   │ │ │ │  
   C   ═ C   — C   — C   — H  But-1-ene
   │ │ │ │  
   H H H H   
  Or
   H H H H   
   │ │ │ │  
  H   — C   — C   ═ C   — C   — H  But-2-ene
   │ │ │ │  
   H H H H   
  ii) Bromine water is decolourised because “X” is unsaturated or has a
(-C=C-) Double bond
  iii) C3H8(g)   +   5O2(g)                      3CO2(g)   +  4H2 O(l)
29. a) i) Pent-2-ene
  ii) Butanoic acid
    b) i) Substitution
  ii) Additional
c) i) C4H10(g)    +   13/2 O2(g)                              4 CO2(g)   + 5 H2O(l)
  2C4H10(g)   +    1302(g)                        8CO2  + 10H2O(l)
ii) The carbon (IV) oxide gas which is produced is acidic.  It dissolves in “K” water to form weak acid: carbonic acid.
d) i) Process whereby menometer (small molecules) join together to
form large molecules (Polymers)
  ii)
    F F
    │ │
      — C   — C   —
    │ │
    F F n
e) - Cheaper
 -  More durable
 -  Stronger
 -  Can be recycled
 - Not attacked by many chemicals
 - Corrosion resistant
30. a) i) Alkalyne
  ii) Carboxalic acid/Alkanoic acid
 b) i) vulcanization
  ii) To harden rubber/make it tough and stronger
 c) i) 2C3H2OH (aq)  +   2k(s)              2C3H2OK(aq) + H2
            ii) Process I: Dehydration
 iii) Additional hydrogenation
     A= 1,2 – Dibromoprepane
     B=Ethene/ C2H4
 iv) Nickel/platinum/palladium/platimin
 v) H H
  │ │
       — C   — C
  │ │
   CH3      H
 d) - Fuel/ source of fuel
  - Production of hydrogen gas
  - Production of   i) CCl4
       ii) Trichlomethane
      iii) Methanol
31. a) Ethane burns with a non luminous flame blue in colour whereas ethyne
burns with a luminous (yellow)flame which is very sooty- Ethane is saturated while ethyne is unsaturated with high percentage of carbon- particles.
 b)
   H   H   
   │   │  
  H   — C   — C   ═ C   — C   — H        │   │  
   H   H   
  Or
     H H   
     │ │  
      C   ═ C   — C   — C   — H  
   │  │ │  
   H  H H   
 c) i) A   = Oxidation “B is Ethene substance “C” sodium ethanoate
  ii) C2H5OH(g)   + O2(g)                  2CO2(g) 3H2O(l)
  iii) To bring reacting monomens into close contact.
  iv) -As a fuel
-Carbon black
-Manufacture of methanol
-Manufacture of di, tri and tetrachloromethene
32. a) i) Fractional distillation
  ii) boiling point
   molecular mass/ density
 b) i) C3H6
   Shake a sample with bromine, C3H8 does not decolorize it, but-
c3H6 decolourreses it.
Or
Use acidifical potassium magnate (VII) C3H6 decourise acidified potassium chromate (vi) C3 H6  Change it from orange to green while C3H6 burns with a smokey luminous falme.
 Alternative
Burn a sample of C3H8; it burns with a non luminous flame. C3H6 burns with a smoky luminous flame
c) P1
  H   — C   ═ C   — Cl        
   │ │    
   H H     
  P2
   H H H    
   │ │ │   
  H   — C   — C   — C   — Cl     
   │ │ │   
   H H H    
 d)  i) Ethanol
  ii) Slightly soluble in water
 e) Name: polythene/polythene
  Disadvantages of polythene
  -Non biodegradable
  -Pollute the environment by producing poisonous gases when burnt
33. a) Hydrocarbons
 b) i) fractional distillation
  ii) Fuel/component of =petrol/to drive small machines.
 c) i) CaC2  /Calcium distillation
  ii) phosphoric acid is the catalyst
  iii) H          C    C         H
  iv) Hydration
  v) I. - Wire insulation coat
    - Water prove seat covers
    - Motor cars seat covers
    - Shoes
    - Suitcase covers
           II.  Hardening of oil in manufacturing of margarine
 d) i) NaOH(aq)   +   CH3COOH                 CH3COONa(aq) + H2O(l)
  ii) Hydrochloric acid is a strong acid with many hydrogen ions to
react with the carbonate.  It is fully ionized in water.  Ethenoic acid is a weak acid with few Hydrogen ions.  It is partially ionized in water.
34. a) i) 2- Methy – prop – i-ene
  ii) pent –I – yne
35. a) i) Methane is a gas which is flammable in presence of oxygen.
  ii) Pass the mixture through a solution of calcium hydroxide to
remove CO2.  Then determine the volume of the gas left using a syringe.
 b)  i) Mass of methane   = 35.2  x   5 = 1.76kg  = 1760g
                                                            100
   
Moles = 176 =  0.11 moles
        16
  ii) CH4(g)  + 2O2(g)    CO2(g)    +  2H2O(g)
   Volume = 0.11 x 24 = 2.64dm3 = 2640cm3
 c) i) CO2 causes global warming
   -No causes acidic rain
   -Trichlorofluromethane destroy ozone layer
                         ii) I. Exhaust from vehicles
   II. Aerosal sprays.
36. i) Compounds containing carbon and hydrogen only.
 ii) A family of compounds having the same functional group and shows
similar chemical   characteristics.
iii) A hychocarbonic that contain maximum number of hydrogen atones possible banded to carbon atoms.
 Existence of different compounds with the same molecular formula but different structural formula.
37. i) But ¬¬¬____ 2 ____ ere
 ii) 2, Methylbutene
38. i) Step I reagents: Acidified potassium magnate (VIII)
  - Potassium dichloromate
  Conditions: -room temperature and pressure
  Step II reagents:   -Hydrogen gas
              Conditions  - Nickel catalyst/heat
 iii) A: [— CH2 — CH2 ¬¬¬—] n
  B: CH2CH3Br
  C: CH3CH3Br
  D: CH3CH2HSO4
39. i)
   H CH3 H    
   │ │ │   
  H   — H   — C   — C   — H     
   │ │ │   
   H CH3 H    
 ii)
   H H H    
   │ │ │   
  H   — C   — C   — C   — H     
   │ │ │   
   H Cl H    
40. a) Increase from “A” to “E”
 b) C15 — C25 ¬¬¬¬¬— D
  C4 — C12 — B
  C20 — upwards — E
  C9— C16 — C
  C1 — C4— A
41. Boiling point increases with increases in number of carbon atoms.  Pentane
molecules are big /large/bulky and the vander waals forces between these molecules is stronger compared to others.
42. i) C5 — C10
 ii) Carbon (ii) oxide / sulphure (iv) oxide/ nitrogen (iv)oxide
43. Sunlight energy split the halogen molecules into free radicle /atoms which are
very reactive i.e U.V act as a photocalolyst.
44. i) alkanes
 ii) Name: Propane:
   H H H    
   │ │ │   
  H   — C   — C   ═ C     
   │  │   
   H  H    
      iii) CH3CH(g)    =    CH2  + HBr(g) → CH3CHBrCH3(aq)
45. i) R:  Sodium hydroxide
 ii) T: tetrachloro methane/ carbon tetrachloride
 iii) CH3COONas(s)   + NaOH(aq)  → CH4(g) +  Na2CO3(aq)
46. i) Polyethene /polythene
 ii) (CH2 —  CH2 —)n  =  42000
  28n = 42000
  n = 42000  =1,500
        28
47. a) (C2H3)n = 54
  27n  = 54  = 2
     27
  (C2H3)2   C4H6   MF= C4H6
 b)   H H    
    │ │   
  C   ≡ C   — C   — C   — H     
  │  │ │   OR 
  H  H H 
   H   H    
   │   │   
  H   — C   — C   ≡ C   — C   — H     
   │   │   
   H   H    
  H   H    
  │   │   
  C   ═ C   — C   ═ C     
  │   │   
  H   H    
 c) Alkyne if it has [¬¬— C       C  —] or
  Alkene if it has [—  C   =    C   —]
    │ │
 
TOPIC 4
NITROGEN AND ITS COMPOUNDS
1. - Funnel has no tap/ does not dip into the reactant
 - Ammonia should not be collected over water as it is very soluble.
2. - Cracking/ descrpitating sound
 - Brown gas produced
 - Gas which relight a glown splint produced
 - Solid change from white to brown when hot and yellow when cold
3. a) i) NO-33                 N+ (-2 x 3) = -1
     N=  -1 + 6
      N= +5
  ii) NO   N + -2 =0
   N = 0 + 2
   N = +2
 b) Reduction:  because the Nitrogen in NO-3 ion gains electrons to form No
i.e. the oxidation number reduced from + 5 to +2/ oxygen is removed.
4.

5. Ammonium chloride and calcium hydroxide
6. RMM of (NH2)2CO =  22 + 4 + 12 + 16 = 60
 RMM of NH3 = 14 + 3 = 17
 Moles of NH3 =680  = 40 moles
       17
 Moles of Urea (NH2)2CO = 20x60 = 1200 kg
7. a) Zinc /Zc
 b) Zn (NH3)4 -2
8. a) NaOH or KOH
b) At first, light blue precipitate was formed.  In excess the precipitate dissolve to form a deep blue solution.
9. NH4Cl decomposes to give ammonia and hydrogen chloride gas.  Ammonia
diffuses faster than hydrogen chloride since it is less dense.  Ammonia is basic and Hcl is acidic in presence of moisture.
10. a) Oxygen gas
b) Thermal decomposition
11. Chemical test
Insert a blightly glowing splint it relight
Physical test
- Invert a gas jar of No. if it turns brown it is not N2O.
- Invert gas jar of “G” over cold water if the level rises it is N=2O
- Has a sweet sickly smell
12. a) The solution contained (OH) ions which change litmus to blue/Ammonia
is basic in presence of water.
 b) Prevent sucking back of water if the reacing vessel as ammonia is very
soluble.
13. a) Nitrogen gas
 b) Withdraw delivery tube from water.  This prevent sucking back of water.
14. a) Nitric acid is more volatile than concentrated sulphuric acid or Nitric acid
has a lower boiling point then concentrated sulphuric acid.  It therefore evaporate readily.
 b) NaNO3/ Sodium nitrate
 c) - Making  ammonium fertilizers
  - Making dye
  - Making explosions
  - Making synthetic fibres/nylon
  - Purification of metal/ gold
15. a) Platinized rhodium /gauze
 b) 2NH3(g) + 5/2O(g)                        2NO(g)    +  6H2O(l)
                       Or
  4NH3(g)  + 5O2(g)                     4NO(g) + 6H2O(l)
 c) - Nitrogenous fertilizers
  - Make explosive
16. White flames produces, ammonia react with Chlorine producing hydrogen
chloride gas which react with excess ammonia to give ammonium chloride.
17. White solid contain MgO and Mg=3N2 (magnesium nitride)   which react with
water to give ammonia gas.
18. a) An alkali is a  base that dissolves in water to give hydroxide ions(OH)
 b) i) Ammonia gas is very soluble in water thus it will dissolve in water
instead of being collected.
  ii) Ammonia is less dense than air and would therefore not displace
air in the collecting jar.
 c) Hydroxyl ions (OH)
 d) Moles of NH3 = 120  =0.005 moles
      24000
 e) i) The solution of Ammonium phosphate is heated slowly to about
half the volume so as to concentrate/saturate it.  It is then allowed to cool slowly to form crystals, then filtered.
  ii) From equation 3 moles of ammonia produces 1 mole of
Ammonium phosphate ration 3:1
 Noles of (NH4)3 PO4 = 0.005 = 0.0017 moles
         3
 RMM (NH4)3 PO4 =  (14 x 12) + 31+ 64 = 149
 Mass = 0.0017 x 149= 0.253g
19 a) i) Water
  ii) Black Copper (II) Oxide will change to brown copper metal
  iii) 2NH3(g)  + 3CU(s)  → 3HO2(l)   + N(2)(g)
  iv) (I)   Moles ratio of NH3 : N2 = 2:1
    i.e 2 mole NH3 gives   1 mole N2
    320cm3 NH3 will give 320 =  160cm3
                 2
   (II) Moles of NH3  =  320  =  0.0133 Moles
        24000
 Moles of CUO = 0.0133 x 3  = 0.02 moles
           2
 RMM CUO = 63.5 + 16 = 79.5
 Mass of CUO = 0.02 x 79.5 = 1.59g
(III)    Excess ammonia dissolve in water to form basic ammonia solution.
     b) The burning splint will be extinguished.
    c) - The method is cheaper
- Nitrogen will be pure i.e it will not be contaminated by other chemical as  
   is the case when obtained from ammonia.
20. i) Fusses calcium chloride/Cao (Quick lime)
 ii) To remove Carbon (IV) Oxide
 iii) 4Fe(s) + 3O2(g)                3Fe2O3(s)
Or
  3Fe(s) + 2O2(g)  Fe2O4(s)
 iv) Argon/helim/ Neon/ Krepton
 v) Provide very low temperature so that the semen does not decompose/ is
not destroyed.
 b) i) concentrated sulphuric acid
  ii) NaNO3(s)    +  H2SO4(l)    → NaHSO4(aq)   + HNO3(aq)
   Or
   2NaNO3(s)   + H2SO4(l)       →  Na2SO4+2HNO3
  iii) (I)     To avoid decomposition of Nitric acid by Sunlight/ Light
(II) Copper reacts with 50% nitric acid to give Nitrogen (II) oxide which is colourless.  Air oxides niteogen (II) oxide to Nitrogen (IV) oxide which is brown.
 c) NH3(g)  + HNO3(g)        → NH4NO3(s)
 Rmm of NH4NO=3 = 80
 Moles  of NH4NO3 ==  4800 = 60 moles
       80
From moles ratio 1:1 moles of NH3  60 moles
RMM of NH=3 = 17
Mass of NH3 = 60 moles
Rmm of NH3 = 60x 17 = 1020kg
21 a) I. Fractional distillation of air
  II. Neutralization
 b) electolysis of brine/water gas or cracking of alkane.
c) High pressure brings the molecules closer/increases the concentration of gas molecules/leads to more collusion.
Or
High pressure shift the equilibrium to the right hence the yield of more ammonia gas,
d) 2NH3(g) + H2SO4(aq)  → (NH4)2SO4(aq)
e) Catalyst :  platinum Rhodium/gauze
 Reagent : water and Oxygen
f) Ammonium nitrate
g) A fertilizer/as a fertilizer.
22 a) i) Heat
  ii) (I)   Soluble carbonate Na2CO3/H2CO3
(II) : Oxygen gas
(III) R = HNO3 Nitric (V) acid
S -   HNO2 Nitric (III) acid
  iii) I: Pb(OH)+-24
   II: Pbo(s)    +    H2(g) → H2O(l)¬     +  Pb(s)
 b) i) -Cheap
   -Corrosion resisistant
  ii) LEAD IS POISONOUS/ harmful/ affect nervous system/brain
 c) i) The reaction produce insoluble lead (II) sulphate which coats the
surface of Pb(NO3)2 preventing further contact.
  ii) Potassium Nitrate or Sodium Nitrate
23. i) 2KNO3(s)  → 2KNO2(aq)=O2
 ii) 2AgNO3(s) →        2Ag(s)  + 2 NO2(g)  +O2(g)
24 i) Nitrogen (II) oxide (NO)
 ii) NH3 (Oxidation NO of N = N + 3 = 0 = -3
  NO2 (Oxidation NO of N = N -4 = 0 = +4
  Oxidation No: of N increase from -3 to +4
 iii) NH4NO3(s)  Heat  N2O(g) + 2H2O(g)
 iv) - Fertilizers
- To make explosives
c) i) G
 ii) E2+(aq)   +  2OH- (aq)                      E(OH)2(s)
25. i) Nitric acid attack, rubber, cork wood and metals
 ii) Due to the presence of Nitrogen (IV) oxide formed by thermal
decomposition of Nitric (V) acid.
 iii) By bubbling in air which will make Nitrgen (IV) Oxide to combine with
water to given Nitric  (V) acid.
26. a) X – Ammonia
  Y- Water
  Z- Nitrogen gas
 b) NH3(g)  + CUO(s)                 N2(g)    + CU(s)   + H2O(l)
27. a) In precess, p, the mixture is passed through KOH to absorb – Carbon (IV)
oxide.  While in Q it is passed through concentrated sulphuric acid or fussed calcium chloride to absorb water vapour.
 b) By fractional distillation
28.  i)  N2(g) +3 H2(g)                2NH3(g)
  ii)   Platinum/platinized asbestos/vanadium (V) oxide
  iii)   Ammonium Sulphate
29. a) i) Dinitrogen tetra oxide (N2O4)
  ii) Nitrogen (IV) oxide (NO2)
 b) 2NO3(g)  + H2O(l)              HNO3(aq) + HNO2(aq)
30. a) Due to presence of dissolved Nitrogen (IV) Oxide (NO2)
 b) Nitrogen (IV) oxide (NO2)
 c) Oxygen gas
 d) Glass wood is to soak up Nitric acid.  It also conducts heat to the acid. 
Sand prevents direct heating to the acid, which might explode i.e prevent bumping which may cause cracking of glass.
            e) 4HNO3(aq)                      4NO2(g)  + 2H2O(l)  + O2(g)
31. a) The reaction is highly exothermic and the resultant heat causes the glow.
 b) Brown fume formed when the resultant gas (Nitrogen (II) Oxide combine
with oxygen in air to form Nitrogen (IV) oxide.
32. a) Haber process
 b) Finely divided iron catalyst
 c) Reaction between ammonia and oxygen in presence of platinum gauze
catalyst is exothermic.  Brown fumes are due to NO2(g).  Initially there is formation of NO(g) which is then oxidized in presence of oxygen to form to form brown gas (NO2)
33. a) Fe3+(g) and Cl-1(g)
 b) Sold Q fe2O3
  Brown precipitate FeCl=3
 c) i) Fe3+(g)   + 3OH(g)                 Fe(OH)3(s)
  ii) 2Fe(OH)3(s)             Fe2O3(s)  + 3H2O(l)
34. a) Impurities/ dust may poison the catalyst
 b) A- Oxygen /air
             B- Ammonia gas
 c) D- catalytic chamber
  E- Oxidation chamber
  F – Absorption chamber
 d) i) NH3(g)  +  O2(g)                NO(g)      + H2O(l)
  ii) 2NO2  +  H2O(l)           HNO3(aq)  + HNO2(aq)
 e) Platinum Rhodium /gauze/catalyst
 f) - Distillation
  - Oxidation of HNO2 by blowing in air.
 g) -  Manufacture of fertilizers
  - Manufacture of dyes
  - Refining precions metals/ gold
  - Manufacture of plastic /nylon
  - Manufacture of explosive/dynamites
 h) Concentrated Nitric acid is an oxidizes copper to copper ions and it self is
reduced into Nitrogen (II)Oxide which is colourless and water.  Nitrogen (II) Oxide is oxidized by oxygen in air to Nitrogen (IV) Oxide which is brown.
 
TOPIC 5
SULPHUR AND ITS COMPOUNDS
1. V =  Barium sulphite / BaSo3
 W=    Sulphure (IV) Oxide
2. a) Tube  I molten sulphure and water: Tube II super heated water.
 b) To force the molten sulphur out
3. Effervescemce.  Colourless gas with rotten egg smell. The gas darken the paper
soaked in lead acetate.
4. a) T as iron  (II) Sulphide
  U is hydrogen sulphide gas
 b) pass through soluble salt of lead e.g lead (II) nitrate and a black precipitate
of pbs is formed.
5. i) SO2 : s + -2 x 2 = 0
  S= +4
  SO3 : S + -2 x3 =0
  S= +6
  Oxidation number of sulphur increases from + 4 to + 6.
This is oxidation number of nitrogen decreases from + 4 to + 2. This is reduction.
 ii) SO2(g)     Sulphur (IV) oxide.
6. Due to formation of insoluble barium sulphate which “Coat” the reacting sulphite
and stops the reaction.
7. Sulphur is made up of poly atomic molecule (S8 ring).  The rings are held together
by weak vander waals forces.  On slightly heating the Vander walls forces are over come and the rings slid over each other.  On further heating, the rings open up to form chains of suphur atoms (S8) which then entangles making it viscous and dark.
8. SO2 which is poisonous is released in the air.  Acid rain which may cause corrosion will be formed.
9. Add dilute acid HCl or H2SO4 to each substance separately.  If it is sodium sulphide  (Na2S) a colourless gas with rotten eggs smell will be produced.  If it is sodium carbonate Na2CO3 effervescence and a colourless gas that forms white precipitate with lime water is produced.
10. Black precipitate formed
11. a) C = Fes or Zns
b) Hydrogen sulphide is very soluble in cold water but insoluble in warm water.
c) Black precipitate formed.
12. Concentrated nitric is a strong oxidizing agent.  It oxidizes iron (II) sulphate and
itself is reduced into nitrogen (IV) oxide gas which is brown and water.
13 a) 2NaOH(aq)  + H2SO4(aq)             Na2SO4(aq)   + 2H2O(l)
 b) - Blue – litmus paper turns red
  - Red litmus paper remains red
 c) The acid was in excess
14. a) Rhombic or monoclinic
 b) - Vulcanization of rubber
  - Preparation of calcium hydrogen sulphite which is a bleaching agent
  - Manufacture of sulphuric acid
  - Gun powder
  - Drugs/ointments
15. a) Sulphur (IV) oxide
 b) i) The gas escaped through the thistle funnel
  ii) The gas delivery tube was immersed in the reagents: Gas escape
.through the thistle funnel.
16. a) Sulphur (IV) oxide
b) i) The gas escaped through the thistle funnel.
  ii) The gas delivery tube was immersed in the reagents:  Gas escape
through the thistle funnel
17. a) H2S: (+ 1 x 2) + S = O
   S = O + 2
   S = + 2/2 = + 1
   S = + 1
18. a) solution from yellow/ orange to green
 b) 2, Fecl3(aq) + H2S(g)      →         2Fecl(aq) + 2Hcl(aq)  + S(s)
 c) Oxidation since hydrogen is removed oxidation number increase from -2
to 0
19. a) Concentrated sulphuric acid.
 b) Solution of blue solution is heated gently till it is half way its volume so as
to concentrate it.  It is then cooled slowly to obtain the crystals.
 c) An hydrous copper (II) sulphate.
20. The molecules which were in form of a ring open up to give chained molecule
(S8). This entangles each other reducing the flow of molten sulphur in increases
its viscosity.
21. a) A black solid is formed
 b) FeS(s)  + 2HCl(aq)               FeCl2(aq)   + H2S(g)
c) Iron powder has a very big surface area, hence high chance of parcels combining together.
22. Combustion of fuel produces sulphur (IV) oxide (SO2) which when dissolved in
water (rain) cause acidic rain which corrodes buildings and affect plants and
animals.
23. i) :I: 180C
  II:  at 100C solubility = 153 in 1000cm3
  In 15 litres / 1500cm3 maximum of
  SO2 153 = 153x 15000 = 2,295g
        1000
 ii) Solubility at 230C solubility is 98g/100cm3
  Moles of SO2 = 98 = 1.53 moles
      64
  Moles of NaOH = 2x 1.53 = 3.06 moles
  Volume of NaOH
  = 3.06 x 1000  = 1.53ocm3
   2
     b) i) I:  2 Fes(s) + 7/2 O2(g)        →            Fe2O3(s)  + 2SO2(g)
   Or
    4FeS(s)  +7O2(g)       →           2 Fe2O3(s)    + 4SO2(g)
   II:    SO3(l)  + H2SO4(l)       →        H=2S2O7(l)
   III:   H2S2O7(l)  + H2O(l)      →       2H2sO4(l)
ii) I:   To shift equilibrium position to the right and increase the
yield of  SO3(g) / Complete oxidation of SO2(g)
II: Vanadium (V) oxide platinum/platinized asbestos
24. (i) A reaction where heat is lost to the surroundings
(ii) The yield will lower: though by le-chatliers principles the yield is expected
to increase, the rate of reaction is lower because the reacting molecules have lower kinetic energy.
iii) RMM SO3 = 32 + 916 x 3 = 80
 Moles – of = SO3 = 350 = 4.375
             80
 RMMH2S2O7 = 2 + (32 x 2) + (16 x 7) = 178
 Mass of oleum = 4.375 x 178 = 778.75kg
25 a) Malachite (CUCO3: CU (OH2)
 b) i) Gas p is hydrogen sulphide reagent i.e is Na(2) CO3/K2CO3
   solid  R is CUO /copper (II) Oxide
  ii) CUCO3(s)                  CUO(s) CO2(g)
  iii) Step 4
   (i) – Green solid dissolves to form blue solution
        -  There is effervescence
  Step 7
- Black solid dissolves to form blue solution
c) i) Tin/Sn
 ii) Making
- Ornaments
- Medals
- Coins
- Gear wheels
- Clock springs
- Rims
- Metal bearings
- Jewellery/decorations
26       a) Super heated water – tube III/ Outer most/ widest pipe.
 b)    i) Platinum / vanadium (V) oxide
        ii) I:  The yield decreases. The high temperature decompose SO3
or the forward reaction is excothermic hence equilibrium will shift to the left
II:   Yield increases: there is increase in pressure: This will make equilibrium to shift to the right
iii) SO3 is dissolved into concentrated sulphuric acid to form oleum.  The oleum is diluted with water to make sulphuric acid.
c) 
i) 2NH3(g)  + H2SO4(aq)             (NH4)2(s) SO4
ii) Rmm of H2SO4 = 98
 Rmm of (NH4)2SO4 = 132
 Moles of (NH4)2SO4 =25  = 0.189 moles
    132
 Moles of H2(NH4)SO4 = 0.189 moles
 Mass of (NH4)2 = 0.189 x 98 = 18.6kg
27. a) 
 i) 2PbS(s)   + 3O2(s)                    2PbO(s)  + 2SO2(g)
 ii) 
- Pure so as not to poison the catalyst
- Dry so as not to interfere with collectin of SO3 which is very soluble. The H2SO4 formed may destroy catalyst.
iii) SO3 reacts with concentrated sulphuric acid to form oleum
 SO3(g) + H2SO4
 SO3 (g) + H2SO4(aq)                        H2S2O7(l)
iv) 
- Sulphur (IV) oxide
- Dissolves in rain water causing acidic  rain
vi) High pressure will increase the cost of production/even if the pressure is increased more than 3 atmospheres, the yield is not increased
b) i) Iron fillingss
 - Effervescemce
 - bubbles f colourless gas
 - Greenish solution
Crystals of white sugar
- Black spongy mass foams off
- Heat produced
- Vapour produced
ii) I.  Sulphuric acid is a strong acid
   Fe(s) + H2SO4(aq)      →         FeSO4(aq)  + H2(g)
  II. Concentrated sulphuric acid is dehydrating agent
c) Ammonium sulphate     
d) BaSO4 is insoluble in water hence the paint pigment will not be removed/ washed by water.
28 i) S8 = 256
 ii) Plastic sulphur
 iii) The rings of 8 atoms open up as the moelten sulphur is heated strongly the
long chins entangles and make the liquid sulphur to be viscous
29 a) The purple KMNO(aq)  is decolourized
  -Yellow solid is formed
 b) KMNO4 is reduced to colourless Mn2+ compounds the H2S is oxised to
yellow sulphur
 c) 2MNO4(aq) + 5H2S(g)                        6 H+(aq)¬   - 2Mn2(aq)  +  8H2O(l) + 5S(s)
30. a) i) Sulphure
  ii) Vanadium (V) oxide/ platinum
 b) i) Forward reaction is favoured hence more yield of sulphur (IV)    oxide
  ii) Low yield of SO3 since backward reaction is favoured because the
   reaction is exothermic.
 c) i) SO3(g)  + H2SO4(l)                       H2S2O7(l)
  ii) H2S2O7(l)  + H2O                      2H2SO4(l)
31. - Solution turns from yellow to green and yellow deposit of sulphur formed
 - Hydrogen sulphide is a reducing agent: it reduces Fe3 + to Fe2+ and itself is
   oxidized to sulphur .
32. - Sugar changes to brown and then black sugar charred off to give a black spongy
   mass of carbon.
 - Vapour produced
 - A lot of heat given out.
33.     a) i)       Sulphur (IV) oxide  - A
  ii) Oxygen – B
  iii) Platinum/platinized asbestos/ vanadium (V) oxide p
     b) 2SO2(g)  + O2(g)         →               2SO3(g)
34. Cl2(g)  + SO2(g)   + 2H2O(l)   →  2Hcl(aq)     + SO-2+4(aq) +4H+ (aq) or
            Cl2(g) +SO2(g)    + 2H2O      →    2Hcl(aq) + H2SO4(aq)
35. a) P- Barium sulphite
 b) 2HNa2SO3
 c) BaSO3(s)   + 2HNO3(aq)     →        Ba(NO3)(2)(aq)  + SO2(g)  + H2O(l)
36. i) Existence of a substance in a more than one form in the same physical
  state
 ii) Carbon
 iii) Rhombic and monoclinic
 iv) -Manufacture of sulphuric acid
  -Vulcanization of rubber
  -Fungicide
37. ZnS(s)   + H2SO4(aq)      →     H2S(g)  + Zn SO4(aq)
 Moles of H2SO4  =  0.2 x 100 = 0.02 moles
    100
 RMM a= ZnS = 97
 Moles of = ZnS = 9/97 = 0.09 moles
 ZnS is in excess by 0.09-0.02 = 0.07 moles
 
TOPIC 6
Chlorine and its compounds
1. a) Cl2(g)  + 2NaOH(aq)                    Nacl(aq)  + NaOcl(aq)  + H2O(l)
 b) NaOcl; decomposes to give oxygen atom that bleaches the dye/bleaches
  by oxidation.
2. a) additional chlorination
 b) CH3CH = CH2(g) + Cl2(g)                   CH3 CHClCH3(aq)
3. a) Sunlight U.V
 b)  H H H
   │ │ │
  H   — C   — C   —  C   — CL
   │ │ │       
                                 H        H         H
4. a) Yellow deposit of sulphur.
 b) H2S(g)  + Cl2(g)           →            2Hcl(g)    + S(s)
 c) In the fume cupboard since cl2   is poisonous or in the open air.
5. a) i) Concentrated Hcl(aq)  hydrochloric acid
  ii) Concentrated sulphuric acid H2SO4
 b) More dense than air
6. a) Chloric acid/ hypochlorous acid decompose to form atomic oxygen which
  oxidizes the dye and bleaches it
 b) 2HOCl(1q)      →          2HCl(aq)  + O2(g)
7. a) Iron (II) Chloride
 b) The solution was basic PH 14.  Excess Hcl neutralized the alkali and the
  solution became acidic as Hcl is acidic .
8. a) Cao is basic will Hcl is acidic.  They will react to form salt and water.
 b) Silical gel/ conc.  H2SO4/ used Cacl2
9. a) 2NaOH(aq)  + Cl2(g)            →                  NaOCl(aq) + H2O(l)
 b) - Bleaching agent.
  - Oxidising agent
10. Add silver nitrate solution white precipitate is formed which change to violet when exposed to light.
 White precipitate on adding lead nitrate.  The precipitate dissolves on warming.
11. a) It is drying agent
 b) 2HCl(g)  + Fe(s)            →            FeCl2(aq) + H2(g)
 c) -Picking of metals 
  -Making dye, drugs
12. 
No. Gas Test Observation
I Chlorine Put a moist red litmus paper into the gas. Turns red than white/bleached
II Sulphur(IV) oxide Potassium dichromate Paper turns green
III Butane Add a drop of bromine water Colourless solution
13. a) O2(Oxygen gas)
 b) PH drops: HOCL decompose to give HCL, which is a strong acid.
14. 

15. a) i) 
  ii) Remove HCl(g) sprays
             iii) In MnO2 manganase (Mn) is reduced.  Mn is (mnO2) has
   oxidation number +4 but in MnCl2  it has oxidation number +2
16. a) To remove oxygen / air which would react with the element to form an
  oxide.
 b) To absorb excess/unreacted chlorine
  To absorb moisture from the atmosphere
 c) Sodium chloride has  a high boiling point and the burner’s temperature is
  not able to vaporize the sodium chloride
 d) Calcium oxide /quick lime
 e) 2P(s) + 3cl2(g)     →        2PCl3(g)
             P4(s)   + 6Cl2(g)   →        4PCl3(g)
 f) Heat the mixture aluminium chloride sublimes, cool the vapour to obtain
  aluminium chloride.  Sodium chloride is left in the heated vessel.
17 i)

 ii) NaCl(s)    +  H2SO4             NaHSO4(aq) + HCl(g)
 iii) Concentrated sulphuric acid/used calcium chloride or silica gel.
 iv) A white precipitate is produced.  HCl(g) in water ionize to give H+ and Cl
  ions. The Cl- combines with pb2+ ions to form pbcl2.
 v) HCL is not an oxidizing agent, it only reacts and removes the oxides hence clearing the surface.  HNO3   is a strong oxidizing agent.  It re oxidizes the cleaned surface.
 b) i) HCl(aq)  + NaOH(aq)     →         NaCl(aq)   + H2O(l)
   Modes of NaOH = Modes of HCl
   = 46x11=   0.506 modes
       1000 
                       ii) Moles of Hcl in 250cm3
   = 0.506 x  250     =  5.06 modes
    25
   Rmm Hcl      =   36.5
   Mass = 36.5 x 4.06 = 184.69g
18. a) - potassium manganate (vii)
  - Lead (IV) oxide
  - Manganese (IV) oxide
  - Calcium chlorate (Caocl2)
 b) i) to remove all the oxygen which would form iron (iii) oxide instead
   of iron (iii) chloride.
  ii) CaO can absorb both cl2 and moisture, CaCl2 can only absorb moisture.
  iii) Cao(s)     +  Cl2                CaOCL2(s)
   Cao(s)   +  H2O(l)               Ca(OH) 2(aq)
   Ca(OH)2(s)  + Cl2(g)                CaOcl2(s)    : H2O
  iv) RMM Fecl3  = 162.5
   Moles  of fecl3     =     0.5    =  0.003
                         162.5
   Moles of Cl2  = 3 x 0.003 = 0.0045
   Vol of cl2 =  0.0045 x 24000 = 110.8cm3
 c) Fe3+  is reduced to Fe2+; H2S is oxidized to sulphur
 d) Turns, red then white because chlorine is acidic and a bleaching agent
                          inpresence of water.
19. a) i) Sodium hydroxide solution      →         A
   Ethane      →        B
 b) Additional polymerization
 c) - Making water prove pipes
  - making electric insulators
  - making water pipes.
 d. 2Cl -  + 2e(aq)    →        Cl2
 e) Dark brown solid is formed.  Chlorine is more reactive then iodine.  It
  displace it from solution.
 f) i) NaOH  + Cl2       →        NaCl  + NaOCl  + H2O
  ii) a) Moles = 2x 15000 = 30 moles
                        1,000
   b) Rmm NaOCl   =  74.5
    Moles of NaOCl = 30/2 = 15
    Mass of NaOCl = 15 x 74.5  =  1117.5g
       = 1.1175kg
20. a) i) Greenish yellow gas
  ii) Slightly soluble
  iii) Black/Grey solid
 b) i) 4HCL(aq) + MnO2(s)        →        Mncl2(s)  + 2H2O(l)
  ii) To osidise the chloride ions to chlorine gas/ oxidizing agent
 c) i) Iron         ii) Chloride       →        E
  iii) Mass of chlorine = 8.06 – 6.30= 1.76
   Rmm of Cl2 = 71
   Moles of Cl2 = 1.76   = 0.0248
                  71
            Vol  =  0.0248 x 24000 = 595.2 am3
    H          H
    │ │
 d)       CL  — C   —  C   —  CL         1,2 Dichloriethane
    │ │
                                    H         H
 e) - Manufacture of HCL
  - Manufacture of PVC, DDT
  - Manufacture of antiseptic.
21. a) - Carry experiment in a fume cupboard
  - Chlorine should not be allowed to escape to the atmosphere.
 b) Mno2 or K2Cl2O7
 c) General chlorine and drive out air which may combine with heat
  aluminium foil.
 d) Aluminium chloride sublime when heated.
 e) i) 2AL(s)  + 3Cl2(g)    →        2ALCL3(s)
   Moles of AL = 1.08 = 0.04
                    27  
  Moles of Cl2 = 0.04 x 3 = 0.06
  Mass of CL2 = 0.06 x 71= 4.26g
  ii)  3.47 x 100= 81.45%
                                         4.26
         f) Pass the vapour of phosphorous trichloride through a liebigs condenser to
  condense it.
22. i) They react to form a yellow solution of sodium hypochlorite and sodium
  chloride.
 ii) 2 NaOH(aq)  +  Cl2(g)                      NaOCL(aq)  + NaCL(aq)  + H2O(l)
23. i) FeCL3 /iron (III) Chloride
 ii) Reddish brown precipitate
 iii) 2NaOH + FeCL3(aq)                   Fe(OU)3(s) + 5NaCl(aq)
24. HCL(g) in water ionized while HCL(g) in methyl benzene dissolves as a molecule. HCl in water is acidic due to (H+) ions
25. a) Potassium  manganant (vii)
 b) Chlorine gas reacts with ammonium gas to produce white fumes of
  ammonium chloride. 
26. a) W Dry hydrogen gas
  Y Dry chlorine gas 
 b) To increase the surface area for absorption of HCL(g)/ Hydrogen chloride
  gas.
 c) H2(g)   +  CL2(g)                   2HCL(g)
 d) Due to the presence of dissolved chlorine gas.
27. i) lack of inverted tunnel/ dissolution through a delivery tube.
 ii) HCL(g)  is a molecular / covalent compound lacking free ions while hydrochloric  acid is ion; the free ions facilitate the reaction.
28. a) A               Hydrogen chloride
  D  Chloride gas
  B  Hydrochloric acid
  E  Iron (ii) Chloride
  F  Iron(iii) chloride
  J  Hydrogen gas
  Q  Zinc
  K  Zinc (ii) Chloride
b) Solution B: Turns blue litmus paper to red.
 Solution C: No effect on the litmus paper
c) E  Gree precipitate
 F Brown precipitate
d) i. - Excess chlorine
  - Chlorine is an oxidizing agent
 ii. Potassium margent VII or Manganese (IV)  Oxide.
e) Heat
29. a) - Effervescence
  - Green yellow gas
 b) - Use concentrated sulphuric acid as a drying agent
  - Heat the reactant
 c) To remove HCL(g) sprays
 d) i) 4HCL(aq) + MNO2              CL2(g) + MNCL2(aq)  + 2H2O(l)
  ii) To oxidize HCL to form chlorine
 e) Mole of HCL = 40 x 11 =  0.44 modes
      1000
  Moles of CL2   =   .44 x 1   ==   0.11 modes
           1000
  RMM of CL2  =  71
 Mass = 71x 0.11 = 7.81g
f) Fe   Cl
 0.28   0.53
 56   35.5

 0.005    0.0149
 0.005   0.005
 1   3
 FeCl3   Empirical formula
 g) Hydrogen and water
30. a) Concentrated Hydrochloric acid and potassium manganate (VII) or
manganese  (IV) oxide.
 b) Prevent formation of tri-iron oxide (Fe2O4) which will coat the iron
preventing reaction with chlorine. 
 c) It sublimes
 d) Calcium oxide; to absorb excess chlorine gas and water vapour.
 e) fume cupboard/open field; chlorine is poisonous
 f) 2Fe(s)  + 3CL2(g)                       2FeCL3(s)
 g) Yellow solid / sulphur.
31. a) hydrogen chloride
 b) NaCL(s)  + H2SO4(aq)                        NaHSO4(aq)  + HCL(g)
 c) Dense than air.
 d) Concentrated sulphuric acid
 e) i) -Increase the surface area for dissolution of gas
   - Prevent water sucking back
  ii) Silver chloride
   Ag+(aq)   + CL(aq)                 AgCL(s)
  iii) Hydrochloric acid
 
FORM 4 WORK
TOPIC 1
ENERGY CHANGES
1. a)  100- 389 = 289 kj/ male
 b) Exothermic: Energy in the reactant is higher than that of the products.
2. a) mg + Fe2+ + Fe2+                Mg2+  +  Fe(s)
 b) Heat change = 100g x 6.0 x 4.2 = 2520j
  Mole of Fe2t+ = 100 x 05  =  0.05 moles
        1000
  Molar heat = 2520 = 50400 = 50.4kj/moles
             0.05
3. Enthalpy of neutralization between CH3COOH and NaOH(aq)   is low than  that between HCL and NaOH because CH3COOH is a weak acid which does not dissociate fully in water.  HCL is a strong acid.  Some of the energy produced is used to dissociate CH3COOH so as to produce more (H+)
4. a) The energy change that takes place when one mole of a compound is
            formed from its constituents elements in their standard states.
b) (3x286) + (2x 394) – (-277)
 -853-788 + 277 = - 136kj/mole
5. H   = 500x9x4.2 = 18900 joules
             18900 J are produced by 0.06 of J
   38000 J are produced by     0.6   x     38000 = 12
                     18900
6 a) ∆H Activation energy
  ∆H3: Heat of reaction`
  ∆H3  = DH1 + DH2
  H      H
  │      │
7. H   —  C   —   H   +   CL –  CL       → H   —  C   — CL   +  HCl
  │      │
  H                          H
  BBE    BFE
  C — H = 414   C — CL = 326
  CL — CL =  + 244  H— CL = + 431
                757
  Total        +  658
  ∆H = BBE – BFE = 658 – 757  =  99KJ
8.     HF(products) – Hf(reagents
-1207- (-394- 635)= -1207 + 394 – 635 = -175kJ
 
9. a) 

 b) 538  =  269 KIJ
    1
10. It reacts with MaHCO3(s) to form CO2 which causes the dough to rise.
11. a) Hl= Lattice energy
 b) Let the heat be ∆H3   H3+ - 70 =15
12. ∆H1 __ H2      = H3
 — 1673.6  - (836) =    H1
 ∆H1=  836.2 JK/male
13. a) The heat absorbed by a mole of a substance to change from liquid state to
gases state at constant temperature.
 b) Boiling point increases with increase in molecular mass. This is due to
increase in strength of vander waals forces.
14. Moles of CuSO4 = 900  =  0.9  moles
         1000
 Moles g = BaCL2  =   600 =  0.6 mooles
    1000
 Heat change when 0.6 moles BaCL2 are used
 = 17.7 x 0.6  =  10.62KJ
    1500 x 4.2 x DT = (10.62 x 1000) J.
     ∆T =  10.6 x 1000 = 1.70C
     1500  x 4.2
15. There is a constant increase in mass caused by constant addition of - CH2 group.
 a)

b) ∆ Endothermic reaction.  The products are at a higher energy levels than the reactants.
16. First ionization energy decreases with increase in atom radius.  When the atom radius increases the uppermost electrons get further from the nucleus, less energy is thus required to remove it.
17. a) Latent heat of fusion
 b) Negative: particles are losing energy
19. a) Pale yellow liquid produced.  The equilibrium moves/shift to the right so
as to raise the temperature.  The forward reacton is exothermic and will be favoured by low temperature.
 b) Brown fume; reducing the volume of gases mixture will lower the
pressure hence; equilibrium shift to the left so as to raise the pressure.
20 a) Particles gains more kinetic energy and move very fast.
 b) X Y
 c) The heat added at the point helps to overcome the force of attraction
between water molecules i.e latent had of vaporization
21. a) ∆H1 = ∆H lattice/ latent heat of dissolution
  ∆H2 Heat of hydration
 b) ∆H3 = ∆H1 + H2
22. i) H2 + O2(g)                H2O2    ∆H    = -1333KJ mol -1
 ii) H2O(l)                H2(g)   + O2(g)  ∆HF= + 188k J mol +-1
 iii) H2O(l)                H2 O2(g)  ∆H   =   + 55 KJ mol -1

23. 

24. “J” It is very soluble in water at a very low temperature.  Its solubility decreases
with increase in temperature.
25. a) 2CH3OH(l)  +  3O2(g)             2CO2(g)  + 4H2O(l)
 b) i) Mass of methanol = 22.98g-22.11 = 0.87g
   RMM of CH3OH =32
   Moles = 0.87  = 0.027 males
        32
            ii) ∆T= 27-20 = 70C
    H = 500x 7 x 4.2 = 14700J
          iii) ∆HC =  14700 x1 =  544.4kj
      0.027x 1000
 c) -Heat lose to the surrounding from the apparatus
  - Incomplete combustion of methanol
 
 
d)

26. a) i) To get uniform mixture hence uniform distribution of heat.
   For complete neutralization
  ii) H+(aq)   + OH(aq)                H2O(l)
  iii) I.   Significance of Y2- Neutralization /end point neutralization
point.
II. Y1 and Y2: Neutralization is taking place producing heat.
III. Y2 and Y3:  reaction has come to an end and the products
are coding / losing heat to the surrounding.
  iv) I: ∆H = MCDT
    ∆T = 30.9 -24.5 = 6.4 0C  M =200g
    ∆H = 200 x 6.4 x 4.2 = 53765 = 5.376KJ
    Mole of NaOH = 100x 1 = 0.1 moles
            1000
    ∆HNt = 5376  = 53760J =    53.76KJ
       0.1
   Mole heat of neutralization = 53.76 KJ/mal
  v) It will be low since ethanoic acid is a weak acid and it is partially
inized in water, a lot of energy will be used to ionize the molecule further.  HCL is a strong acid fully ionized.
 b) 

27. a) Exothermic: heat energy is given out to the surrounding.
  Endothermic; Heat energy is absorbed from the surrounding.
 b) i) Vaporization /melting
  ii) Condensation /freezing
c) The water is undergoing change of state.  The heat supplied is used in breaking the inter particles forces between molecules of water.
d) i) Heat  of formation of FeCL3 (∆H1)
 ii) ∆H3 = ∆H1 + ∆H2
e) Butane:  because more bonds are formed on combustion of butane hence more heat is released.  Butane has the higher percentage of carbon.
28. a) i)   -There is is a redish brown deposit of copper
   -Blue colour of solution fade/become colourless
   -Grey solid of magnesium dissolve
  ii) ∆H= MC∆T
   ∆T = 43-25 = 180C
   ∆H=25x4.2 x 18 = 1890J
  iii) Moles of mg   = 0.15  = 0.00625 moles
         24
   Moles of CUCL2 = 25x2  = 0.05 moles
              100
   ∆Hppt = 1891x1 = 302400J
             0.00625
   Molar heat of displacement = -302.4 KJ
  iv) Mg(s)  + CU 2+ (aq)              Mg2+(aq)  + CU(s)
 b)

29. a) -the type of flame it produces.
  -amount of heat energy produced
 b) i) Heat produced = MC∆T
   ∆T = 46.5 – 25 = 21.50C
   ∆H  =  450 x 4.2 x 21.5 = 40635 joules
  ii) Moles of ethanol = 1.5 = 0.0326
             46
   Molar heat =    40635 = 1246472.392 joules
     0.0326
 c) C2H5OH (aq)   + 3O2(g)                 2CO2(g)  + 3H2O(l)
 d) -Heat less by radiation, conduction and convectional current.
  -Experimental errors when reading thermometer.
30. C(s)  + O2(g)                CO2(g)
  ∆H   = -360KJ/male
  1 mole of C produces 360 KJ
 30KJ will be produces by    1x30 =  0.83 moles
     360
 Mass of C= 0.083 x 12 = 0.99g of C
31. C3H7OH   + 5O2                3CO2 + 4H2O
 RMM for C3H7OH = 60
 5g C3H7OH produces 167 KJ
 60g C3H7OH  will produce
 60 x 167 = 2004kj
      5
 Molar heat of combustion = 2004 kj/male
32. i) C2H5OH(g) + 7/2O2              2CO2(g)   + 3H2O(g)
 ii) Heat produced by ethanol = heat gained by water
  ∆H = MC∆T
  500 x 4.2 x 60 =126, 00J = -126KJ
 iii) RMM of C2H5OH = 46
  Males of C2H5OH  = 5/46 = 0.1087 males
  0.11 mole produces 126kJ
  1 mole will produce 126 x1 = 1159.2 kj/mole
                                                                             0.1087
  Molar heat of combustion = 1159.2 kj/mole
33. Moles of CUSO4   = 100x0.1  = 0.01 Mole
    1000
 Heat produced = 100 x 4.2 x 4 = 1680
 ∆HDI(s)  = 1680 x 1 = 168000J = 168KJ
   0.01
 Molar heat of reaction = -168KJ/mole
34.
  H      H
  │      │
 H   — C   — H   +  CL — CL   → H   — ¬¬ C   — CL  +  CL — CL
  │      │
   H      H 
  HCL
  Bonds broken   Bonds formed
  C-H  = +444   C-CL   = -326
  CL-CL + 244   H-CLL  =-431
  Total energy +688(E1 Total -757 (EZ)
  ∆H = E1 + E2
  = + 688 + 758 = -70KJ:
35. a) JK: The molecules gain kinetic energy vibrate more and more
 b) KL: Change of state; solid naphlthatain melts.  The temperature remain
constant.
36. i) Exotherme reaction: the products are at a lower energy level compared to
the reactants.
 ii) ∆H is (-ve) negative: Heat is given out/exothermal reaction.
37. i) Heat liberated when 0.25 mole of CU is formed.
  = -526 x 0.25 = -131.5kJ
 ii) Heat liberated when o.5 mole of CU is formed = -63 x 0.5 = -31.5KJ
38. 2C(g) + 2H2(g)  H3  H3 C2(l)H(4)
  H1  =  -1356 H2 =     -1432
  2CO2       +  2H2O
 ∆H3 = ∆H1 - ∆H2
 = 1346 + 1532 = + 76KJ/molr
39. Weak acid is slightly ionized some heat is absorbed during ionization.
40. a) This is the heat change realized when one mole of a substance is formed
from its constituent elements under standard conditions.
 b) i) - Molar heat of combustion of hydrogen
   - Molar heat of formation of water vapour.
  ii) 
  iii) 2C(s) + 3H2(g)            C2H6
   2CO2(g) + 3H2O(l)             C2H6(g) + 7/2O2(g)
   2C(s)  + 2O2(g)         2CO2(g)
   3H2(g)  + 3/2O2(g)   3H2O(l)
   ∆H = -858KJ  ∆H = -86KJmol-1
  iv) I. E = MxC x O
        = 500g x 4.2 x 21.5
        = 45150J
        = 45.15KJ
   II.   C2H6(g)  + 7/2 O2(g)   →       2CO2(g)    +  3H2 O(l)
    ∆H = -156-KJmol-1
    1560Kj produced by 30g of Ethane.
    45.15kJ produced by 30g of Ethane.
    45.15KJ produced by   30  x 45.5g of Ethane
        1560
    = 0.8683g of Ethane 
 
TOPIC 2
RATE OF REACTION
1. - Effect: reaction will be faster
 - Explaination: powdered zinc offers a large surface area.
 - Heat increases the rate since particles collide more.
2. a) In both cases temperature remain constant because the heat energy is
being used to break up forces of attraction in the solid structure/ latent heat.
 b) CDCL2(s)     →      CD2+(l)      + 2CL-(l)
  This is because CdCL2 is an ionic compound which is held together by
electrostatic force athat are stronger than vanderwaals forces and hydrogen bonds holding the H2O  molecules together.  In water there is only one change (liquefaction) but in CdCL2 there are two changes ionization and liquefaction.
3. i) Curve (i)
 ii) Concentration of F increases with time.
iii) After time (t) concentration does not change because equilibrium has been established.
ii) Menganese (IV) oxide is a catalyst and increases the rate of decomposition of the hydrogen peroxide.
4. Curve (i)
i) Manganese (iv) oxide is a catalyst and increases the rate of decomposition of the hydrogen peroxide.
5. Use zinc power which has a large surface area.
6.
7. a) Yield would increase since H is positive.  Thus increase in temperature
shift the equilibrium to the right.
 b) No effect; The number of molecules / volume of gases is the same both to
the left and right side of reaction.
8. Increase in temperature would lower the yield of Nitrogen (ii) Oxide, this is
because the reaction is exothermic and equilibrium will shift to the left.
9. Increase in pressure would shift the equilibrium to the left since increase in
pressure followers the reaction which produces less volume of gas/ products/particles
10. a) The yield of ammonia qould decrease.
  - At high temperature ammonia decomposes
  - i.e Equilibrium moves to the left.
 b) -Manufacture of fertilizer, sodium carbonate
  - Smelling salts
  -As a refrigerants
  - Soften temporally hard water.
11. a) Gas syringe.
 b)
12. Equilibrium has been established or forwarded reaction equal to backward
reaction.
13. a) Reaction must be carried out in a closed vessel/system
 b) Equilibrium shift to the right or forward reaction, because CO2 will be
removed from the system by potassium hydroxide. 
14. Acid “M” is a strong acid than acid “L” it is fully ionized producing more (H+)
ions which react with magnesium turnings.
15. Brown solution produce; Equilibrium shift to the left so as to reduce the amount of HCL added.
16. a) Syringe barrel/graduated gas jar
b)i) particles gain more kinetic energy and collides very fast making reaction faster.
 
17 a) i)

  ii) I.  27 to 28 secons/read graph
   II.  28 x2 = 56 seconds.  The concentrating of [H+] ions is half
/read graph at 3 cm3.
 b) i) Moles of the sulphate = 10 x 0.4 = 0.004 moles
          1000
  ii) Moles of HCL = 10x2 = 0.02 moles
         1000
  iii) Thro sulphate: hydrochloric acid is in excess.  (1mk)
 c) Some cross should be used in each experiment.
  - The cross should be viewed from the same position.
 
18. a) Mg(s)  + 2HL(aq)    + 2HCL(aq)                       + H2(g)
 b)
 c) a)  i.   300cm3+
         ii. 26. 27cm3    0.5cm3
 d) i) Rate is lowered, because magnesium ribbon has a small surface
area then the powder/ collision between magnesium and hydrochloric acid is reduce.
  ii) Rate is increase: number of particles of HCL is higher or
concentration is increased.
   Modes of H2  = 600 = 0.025 modes
     24000
   Rmm mg =  0.6  =24
           0.025
 
19. a) i)
  ii) 480cm3 + 5.0 cm3+
 b) 620-540  = 1.33cm3 / second
      60
 c) Solid is due to presence of copper which had not reacted.  Copper is below
hydrogen in the reactivity series.
 d) Volume of gas H2 from (AL)
  =640- 2.5 = 637.5 cm3
  Moles of H2 = 637.5 = 0.0266moles
     24,000
  Moles of AL = 0.0266 x 2/3 = 0.0177 moles
  Mass of AL = 0.0177 x 27 = 0.4g
  Percentage of AL   = .48 x 100 = 95.6%
          0.5
 e) -Stronger than pure aluminum
  -Higher lensile strength
  -Harder than aluminum/laugher
  - More durable/more resistant ot corrosion/ rusting.
20. a) Carbon (Iv) oxide gas was lost
 b) i.   1.8 - 0 = 0.9g/ minute
          2 - 0
            ii. 3.2 – 2.95 = 0.12g/minute
         8-6
iii. The average rate of reaction in b(i) is higher than that in b(ii).    There are more particles between “O” and z minutes that between 6 and 8 minutes hence the frequency of collision in b(i) are higher than b(ii).
 c) CaCO3(s) + 2HCL(aq)  → CaCL2(aq)  + H2O(l)     + CO2(g)
 d) - Heating
  - Increase of concentration of HCL(aq)
  - Crushing the marble chips to increase the surface area.
 e) It turns dump/wet/increase in mass.  The caCl2 is hygroscope.  It absobs
water vapor from the atmosephere.
 f) i. Calcium sulphate.
  ii. -Making plaster for building/plaster of Paris. Cement/sulphur (IV)
 oxide/ aluminium
-Sulphate
-As filler material for paper (white out)
21. a) Nitric acid is an oxidizing agent and will oxidize hydrogen into water and
it self reduced to Nitrogen(iv) oxide and wate.
 b) Reaction rate will increase since the rate of particles collision will be
higher.
 c) 
 d) i. 370Cm3 + 0.5cm3
  ii. 45cm3 + 2cm3
 e) i.  2.07g of Pb react with 45cm3  of 1MHNO3
       207g of Pb will react with
   207 x 45 = 4500cm+2   4.5dm3
ii. From the graph: 45cm3+ of 1m HNO3  produces 480cm3 of NO2
4500cm3 a = 1MHNO3  produces
 4500 x 480  = 48,000cm3
       45
 f) i)    Moles of nitric acid to react with one mode of Pb
   = 4,500 x 1 = 4.5m:
           1000
  ii) Moles of NO3 produced by one mode of pb
   = 48000 = 2mole
      24000
 g) 4HNO3 + 2Pb                    Pb(NO3)2,              2NO2  + 2H2O
22. a) i) forward reaction is faster than the reverse reaction.
  ii) I. Production will reduce since equilibrium will shift backward so
as to raise the pressure.
II.  No change in amount of methanol since a catalyst will help reaction to come to equilibrium.
  iii. I. Negative:  The reaction is exothermic since it require low
temperature to be fast.
II.  To ensure that the reacting parcels possess more activation
energy.
  b) i)  No of seconds = 2x 60 = 120 sec
       moles of H2O2 decomposed
       = 120 x 6.0 x 10-8 = 7.20 x 10-6
ii) Concentration of H2O2 may be higher since concentration increase the rate of reaction.
23. a) i)  when a stress is introduced to a system in equilibriumshifts in such
a way as to minimize the effect of the stress.
  ii) No effect.  There are equal number of moles on both side of the
equition, therefore change of pressure does affect the equilibrium.
  iii) Negative.  The forward reaction is exothermic since it is favoured
by low temperature.
 
 
b) i) Manganese(IV) Oxide.
  iii)  Rate of O2 production = 14cm3  =    1.4cm3/sec
                                                                      10sec
24. Rate of reaction indicated the velocity of chemical reaction.  It is a measure of the
reactants consumed of products formed per unit  time.
25. - Measure a mount of product formed per unit time.
 - Measure the amount of reactant consumed against time.
 - Measure a mount of heat produced or consumed against time.
26. It is less reactive than hydrogen hence w is displaced by hydrogen from WO3.
 
27. a) CaCO3(s)  + 2HCL(g)   → CaCL2(aq) + CO2(g) + H2O(l)
 b)
 c) 810 – 55 = 255cm3
 d) All the acid was used up.
 e) Moles g2  CO2 = 11.2  =   0.0005
      22400
  Mass of CO2 = 0.0005moles x 44 = 0.022g
 f) Moles of g2 CO2 = 1020
           22400 = 0.0455
Moles of CaCO3 = 0.0455 moles
  RMM CaCO3  = 100
 Mass CaCO3 = 0.0455 x 100 = 4.55g
28. -Addition of catalyst
 -Increasing the pressure.
29. i) Increase in temperature increases the kinetic energy of the particles hence
Increases the rate of collusion
            ii)      Lowers the activation energy
 b) i. Increase
  ii. Increase
  iii. Unaltured
 c) M2(g)  + 3H2(g)           2MH3(g)
  D(-)  = -92Ks/Mole
- Temperature 4500/low
- High pressure 200-400 atmosphere.
- Catalyst iron prevented with AL2O2.
30. Curve (II) the reaction rate is higher because of bigger surface area.
31. Low temperature and high pressure.
32.
33. a) The rate of reaction is doubled.
 b) The rate of reaction increases.
 
TOPIC 3
ELECTRO CHEMISTRY I AND II
1. a) Arrow from zinc to copper rod: zinc is more reactive than copper. 
Zinc donate electrons more readily.
 b) No deflection
2. 4OH(aq)        →                4e + 2H2O(l)  + O2(g)
3. i. Q= 0.6x 90 x 60 = 3240 columbs
 ii. 3240 x 226 = 192695 columbus
         3.8
 iii. Charge = 192695   = 2
        96500
   Charge = +2
4. - Bulb will light since the current flow.
 - Grey metal of lead form at the cathode
 - Brown fumes of bromine at the anode.
5. Chloride ionizes in water since water is polar.  The same chloride dissolve in
methylbenzene as a molecule since the methylbenzene is non polar.
6. Cl ions will remove Pb2+ ions from electrolyte by farming insoluble pcCL2
7. CL ions will remove Pb2+ ions from electrolyte by farming insoluble PcCL2.
8. a) Cathode :  Hydrogen
  Anode:   Oxygen
 b) Increases: Since H2O is decomposed
 c) There would be an explosion because potassium is very reactive.
9. a) E reduced – Exudation = + 0.44 + 1.66 = + 1.22V
 b) Aluminum is more electropositive than Zn: hence react by losing
electropositive than Zn; hence react by losing electron ready.
10. a) Because the concentration of Cu +2 ions is high at the beginning and
decreases as the ions are discharged during electrolysis.
    (b) CU2+(aq)  + 2e            CU(s)
11. a) 2Cr(s) + 3Fe2+(g)             2Cr3+(g)   + 3Fe(s)
 b) 0.44 -E= 0.30v
  EQ =  -O.74v
12. a) Q= 1.5 x 15 x60 = 1350 ccolumbus
 b) 1350c gives 0.6g a= m
  3x96500C give 0.126 x 3 x 96500 = 55.76
               1350 
13. T= 32x 60 + 10 = 1930sec
 Q= 1930 x 0.5 = 965C
 0.44g produced by 965C
 88g = 965 x 88 =  193000C
  0.44
 Charge =  193000 = +2
        96500
14. a) Ag(aq)  +    +e-            Ag(g)
 b) Anode dissolves since it is active.
15. 63.5g requires  2x 96500C
 1.48g requires 1.48 x 2 x 96500 = 4498.3C
    63.5
 Q = it 1 = Q
           T
 T= (2x60x60) + 30 x 60 = 9,000 see
 1 =   4498.3   = 0.4998A
           9000
16. a) The colour of solution fades and Q disappears .
  - Brown solid was deposited at the bottom.
 b) Metal Q is more reactive than copper, therefore it displaces copper from
its solution.
17. i. Bulb did not light:  No ions are present in water.
 ii. Bulb light bubbles of colourless gas H2SO4 is an electrolyte.
18. a) No heating
 b) The solid melt, the ions become mobile.
19.  Q = it = 0.82 x 5 x60 x 60 = 14760  columbus
  No. of Faradays   =    14760  = 0.15F
                                                          96500
  Moles of Z =  2.65 = 0.05 moles
               52
  Change of Z = 0.15   = +3
     0.05
20. a) element “N” its more reactive
 b) EMF  = EQ reduced – Qoxidized.
  = + 0.80 + 0.76 = + 1.57v
21.

 b) E0  cell  = E0  reduced -  E0 oxidezed.
  = -0.14V- -0.74V = + 0.6V
22. a) Chloride ions in brime are in high concentration compare to oxide  ions in
solution
b) Hydrogen gas
23. a) Ag(a)       + e-    Ag(s)
 b) Ce = 1t = 5.0  x 3 x 60 x 60 = 54000C
  Mass of silver deposited
  =   108 x 54000     = 60.44
   96500
24. a) Zn(s) / Zn2+ // 2Ag+(aq) / 2A(s)
b) Greyish shinning solid deposited round copper.  Copper being more reactive displaces Ag from Ag2+ blue solution formed due to presence of CU2+ in solution.
25. a) CU2+ migrate toward the cathode
  CU2+ give solution a blue colour.
 b) 4OH(aq)           4e + 2H2O(l)  + O2(g)
26. a) i.   Copper: It is used as a standard electrode in this cell.
     The two electrodes have the same reduction potential.
  ii.   “J” because it has the most negative reduction potential.   Is easily
oxidized.
  iii.   I.   K(s)   →   2e(aq) + K+2
           2m(aq)  + 2e   →   2M(s)
  
   II.   By allowing  ions move into the two beakers.  Na+ ions
         -pass into the metal M electrode beaker and NO3 ions pass
into metal K electrode beaker.
 b) i. “D” Because oxygen gas is given out at electrode “C” thus “C” is
an anode
  ii. 4OH(aq)           4e-       +  2H2O(l)       +   O2(g)
  iii. I.   Brown substance /solid at electrod “D” This is because
CU2+  ions in solution gains electron at “d” to form Cu(s)
   II.     The solution will remain blue since the electrodes used are
copper and the anode will dissolve to replace the  CU2+ ions which are discharged.
27 a) i. Bauxite AL2O32H2O
  ii. Iron (iii) oxide, silica
 b) i. 

I      It is expensive / a lot of energy will be used
   II.   The ore is dissolve in cryolite (NaALF6)
   III. Its melting point is less than 8000
 c) Q=  40,000 x 60 x 60 = 144, 000, 000C
  Mass of AL = 144,000,000 x 27 = 13.43kg
         3x 96500
28. i. C2:  Hydrogen is used as a reference electrode whose E0 value is 0.00v
 ii. -240v
 iii.

 iv. EMF E0 red- E0 oxidized
  =  + 2.38 + 0.34 = + 2.27v
29. i) To lower the melting point from 800-6000C . Hence reduce the cost of
production.
 ii) Steal will react with chlorine while graphite will not
 iii) -Its melting point is lower than that of the electrolyte
  -It is less dense than the electrolyte
 iv) To prevent th products from coming into contact
 v) i. Cathode Na+(aq) + e → Na(s)
  ii. Anode 2Cl-(aq) → 2e(g) + Cl2
 vi) -Manufacture of Na2O2/NaCN
  -Liquid sodium is used as a coolant in nuclear reactor.
  -Sodium vapour is used in street lamps
  -Extraction of metals e.g Lithium and Aluminum in termite process.
30. i) Platinum /graphite/carbon
 ii) Cation Mg2+ and H+ anions SO42 and OH
 iii) To the left
  I. Anode: 4OH(aq)     4e + 2H2O(l) + O2(g)
  II.   Cathode 2H(aq) + +2e       H2(g)
iv) The concentration of a queous magnesium sulphate increase because water
molecules are broken down into hydrogen and oxygen.
31. i)  I .  Distilled water
     II.   Titanium /platinum
 ii) Chlorine gas
 iii)  -    Paper industry
- Glass industry
- Making soap/ detergents
- Extraction of aluminium
- Manufacture of drugs
b) i.  I.. Hg/Na+(aq)   + e                 Na/Hg(s)
     II.  2Na/Hg(s)  +  2H2O(l)             2Na OH(aq)  + H2(g) + Hg(s)
 ii. Q=it
  = 100x 5x 60 x 60 = 1,800,000C
  1 Faraday form 1 mode of Na
  1, Mole of Na forms 1 mode NaOH
  Rmm NaOH =40
  180000C  forms 40 x 1800000  = 746.1(g)
      96500C
32. i) “G” it has the highest +ve potential Eovalue
 ii) ½  G(g)  + e             G(aq)  and
     M+(aq)  + e              M(s)
 iii) Reaction can not take place from left to right “M” cannot displace
“N” from its solution.  “M” is more electropositive or the E0 value is –ve
 b) i. 40H(aq)             4e + 2H2O(l)  + O2(g)  
  ii. Insert a burning split in gas jar of gas K.  The gas burns
with a pop sound to show it is hydrogen.
   iii.     a)  Hydrogen is monovalent oxygen is divalent.  The
same amount of electricity liberate twice as much hydrogen.
        b) The bulb is brigher with sulphuric acid.  The acid is
a strong acid which is fully ionized.  Ethanoic acid is a weak acid partially ionized hence bulb will be dim.
33. a)  E
 b) i.  F2+(aq)   + 2e  →   F(s)
   G   →         2e + G2+(g)
       ii.    → V   →   From “G” to “F”
      iii.   -To complete the circuit
   -To compensate for the ions used or added to the
  electrolyte.
 c)    i) Bluish/green blue colour of the solution fades CU2+ are
removed from the solution.
ii) Chlorine gas and oxygen initially the concentration of chloride ions was high hence discharged.  With time the concentration of CL – ions decreased and [OH]  ions were discharged in preference to CL- ions.
iii)   “J:  The anions are –ve (negative) and are attracted at the anode.
34 i)

 ii) Zn(s)            Ze + Zn2+ (aq)
 iii) The cell would not produce any current ions are not mobile since the solid
is a non electrolyte.
iv) Advantage
 -Portable
 -Cheap
 Disadvantages
- Not rechargeable
- Cannot produce continous supply of electricity
- Causes environment pollution
 b)    i.  Purple /violet fumes produced since iodine vapour is produced.
        ii.   Q =  0.5 x 2 x 60 x 60 =  3600c
  Mass of Pb = 3600 x 207   =  3.86g
      2x 96500
35. a) Add aqueous sodium carbonate to precipitate calcium carbonate and
magnesium carbonate and filler.
  b) i.   Cathode:  2H+(aq)  +  2e              H2(g)
        Anode : 2CL(aq)              2e + CL2(g)
  ii. U I.    Sodium hydroxide
      II.  Graphite /platinum
      III.  Sodium chloride
     iii. To prevent mixing of chlorine gas with sodium hydroxide but
allow tree movement of ions
       c) - In paper industries
  - Manufacture of soap/detergents
  - Making bleaching agents
  - Purification of bauxite.
36. i) G
 ii) G(s)  + H2+(aq)             G2+(aq) + H(s)
 iii)      EMF = E0 red – E oxide
  + 0.34 + 0.44 = + 0.78v
 b) i.   H
  ii.   Pure water does not contain ions, acid is added to make water
ionize.
  iii.   HCl(aq)  is not used because the chloride ions will react with the
electrodes due to its high reactivity.
 c) 144750 Columbus = 144750  Taradays   = 1.5F
     96500
  2 faradays gives 64g of copper
  1.5 faradays give 1.5 x 64  = 48g
     2
37. i) Graphite/titanium: They do not react with chlorine.
 ii) A steel diagram is suspended between the electrodes
 iii) 2CL(aq)             2CL2(g) + 2e
 b) i.   Calcium chloride
  ii.   It is economical /reduce cost of production
 c) Hydrogen is preferentially discharged at the expense of sodium at the
cathode.  At the anode OH will be discharge in expense of CL.
 d) Na2O2
  Na2O
 e) -Making NaCN (Sodium cyanide used in extraction of gold.
  -Making sodium lead alloy used as antiknock in petol
  -Content in nuclear reactor.
38. a) Substamce which when molten fussed or in aqueous solution conduct
electricity and is decomposed.
 b) i. Conduct electricity when solution through the flow of the ions.
  ii. Graphite has a decolorized electrons which conduct electricity.
 c)                                 Electron flow

  ii. Syringe 1: H+ ions are positively changed and are discharged at
the cathode.
  a) During the process the water molecules are decomposed to give
hydrogen and water.
  b) Q = 0.72 x 15 x 60 = 648 Columbs 
   1 mole of gas (O2 requires 4 faraday i.e
   40H(aq)               4e(g)   + 2H2(l)  + O2(g)
             
   680 Columbus will liberate  648 x 1 = 0.001679 moles
                4 x 96500
   Volume of O2 = 2400 x 0.001679
    = 40.29cm3
39. a) (i) both SO42-  and OH migrate to the anode.  OH being lower on the
electrochemical series is preferentially discharged by losing electrons to form water and oxygen.
  ii) The anode would dissolve in water and move to the cathode as
copper(II) ions.  This would discharge the products of the electrolysis.       
 b) i) - Copper pyrites
   - Copper iron disulphide
     - Basic copper carbonate
  ii) CU 2+(aq)   + 2e       CU(s)
  iii) Q= It = 0.5 x 18 x 60 = 540C
   96000C  deposit 108g of Ag
   540C deposit   108 x 540g  of Ag
     96500 
  iv) - To prevent rusting/ carrion
   - For beauty
40. a)  Ce4+
 b) Mg(s)
 c) Ce+4 ions  Ce+4 (aq)   +   Ag(s)            Ag(aq)   +  Ce+3(aq)
 d) i) Mg(s) /Mg2++ //Cd2-(aq)+/Cd(s)
  ii) Mg(s)   +  Cd2+(aq)           Mg2+(aq)  / Cd(s)
  iii) E value = E0   oxidized
   = -0.402 + 2.37 = + 1.968V
41. a) i) The bulb does not light since solid bromide is a non electrolyte
  ii) Solid Lead (II) Bromide does not contain free ions
 b) To provide mobile ions
 c) Anode:  Brown gas evalued (Br2)
  Cathode: Grey solid (Pb) deposited
 d) Anode 2Br-(g)     2e + Br2(g)
  Cathode Pb2+(aq)  +  2 e   Pb(s)
42. i) Hydrogen ions are discharged in preference to potassium ion.  H+ are
below potassium in the preferential discharge series.    ii) Iodine is given off as a dark brown violet vapour.
iii) Q = 0.2 x 5788 = 1157.6 coulombs
 0.208 g of or requires 1157.6 coulombs
 52g = Cr requires 52 x 1157.6 = 289400 coulombs
    0.208 
Change of Cr = 289400 = +3
    96500 
43. Emf = E(-) produced – E(-)   oxidized
 = -0.40 + 1.19 = + 0.79 V
44. Zn(s) + 2 Fe3+(aq)      zn2+ + 2 Fe2+(aq)
45. EQ value = E0 reduced – E0 oxidized
 =  + 1.36 + 0.76 = + 2.12 V
46. a)  Current flow 
   Electron flow

 b)                                            Electron flow
     Current flow
 c) (see diagram)
 d) zn(s)                  2e   +  zn 2+ (aq)
  Cu2+ + 2e             Cu(s)
 e) zn(s)   + CU2+ (aq)   + CU(s)
 f) 2 moles of electrons
 g) 2 x 96500 = 19300 Coulombs
47. Mg(s)                 2e  +  Mg 2+ (aq)
 CU2+(aq)    + 2e                 CU(s)
48.   a)  i) Oxygen gas evolved at anode Hydrogen gas evolved at the cathode
OH- and H+ ions are discharged in preference of Na++ ions SO4-2 ions. 
  ii. Chlorine liberated at anode sodium discharged at the cathode to
form sodium Amalgem.  There is high concentration of Chloride ions in Brine.
High over voltage effect at the mercury cathode by hydrogen.  This sodium is discharged instead of Hydrogen. The resulting solution is an alkali.
 b) i) Copper ions were discharged and at the same time, the copper
anode dissolves to form Copper (II) ions.
  ii) To increase the concentration of OH ions (II) ions.
  iii) Copper
  iv) Cu2+(aq)  + 2e             Cu(s)
  v) Q= ie = 1.5 x 600 = 900 c
   900 C gives 0.296g of = Cu
   Hence 63.5 g of Cu produced by
   63.5 x 900 = 193074 C
                                        0.296
   Farady constant = 193074 = 96537C
      2
49. E.M.F = E0 Reduced – E0 Oxidised
 = + 1.36 + 2.38 = + 3.74 v
 
TOPIC 4
METALS
1. 2mg(s) + O2(g)               2MgO(s)
 3 mg(s)   + N2(g)                 Mg3N2(s)
2. a) PbO(s)  + CO(g)            Pb(s)  + CO2(g)
 b) silver white / grey metallic deposite of Lead
 c) Hydrogen gas / ammonia gas
3. a) Electrolysis of fused or molten oxide
 b) J- Carbon – H
4. a) Heat
 b) i) D= Sulphur (IV) oxide
  ii)  - Battery casing
- Galvanising ion
- Electroplating
5. Fe2O3(s)   + 3CO(s)            3CO2(s)   +  2Fe(s)
  2C(s)    + O2(g)               2CO(g)
                  2CO(g)  + O(g)                2CO2(g)
6. a) Reduction
 b) - Oxidation state of pb in pbo is reduced from +2 to O (zero)
  - Removal of oxygen
 c) Ammonia gas /Hydrogen gas
7. Coke: to reduce Pbo to Pb
 Limestone:  to remove silica as slag
 Scrap iron: To reduce unleaded Pbs to pb
8. a) dilute Nitric Acid
 b) Silver metal
 c) Oxygen
9. a) Froth floatation
 b) znCO3               znO(s) +  CO2(g)
 c) Manufacture of dry cells.  Zinc casing forms the anode of dry cells.
10. a) i) Cryolite NaAiF6
  ii) Electrolysis
 b) - Good conductor of heat
  - Resistant to corrosion
  - High melting point
  - Meleable
11. i) sulphuric (Iv) oxide
 ii) 2CUFeS2(s)  + 4O2(s)      2FeO(s)  + 3SO(s)  + 3SO2  + Cu2S(s)
 iii) Fe3+
 iv) Carbon (II)  Oxide or Carbon (IV) Oxide
 v) Redox or reduction & oxidation because Cu2O is reduced to CU(s)  and
CO oxidized to CO2
 
 
b)
 c) Moles of Cu = 210  = 3.3 moles
    63.5
  Rmm of CuFeS2 = 183.5
  Moles of CuFeS2  3.3 mole
  Mass of pure ore = 3.3 mole
  Mass of pure ore = 3.3 x 183.5 = 605.5 kg
  % purity = 605.5 x 100 = 74.85
         810
 d) - Formation of acidic rain due to SO2
  - Sulphur ((IV) oxide is poisonous
  - Carbon (II) oxide is poisonous
  - Green house effect due to CO2
  - Dumping of waste like slag prevents growth of plants.
  - Soil erosion due to extraction of ores from the ground
12. a) i) Galena /pbs
  ii) Some of the sulphide is converted into oxide (pbo or SO2)
  iii) Carbon (ii) oxide or carbon (IV) Oxide
  iv) Pbo + C(s) →   Pb(s)  +  Co(g)
  v) - So2 (g)  is poisonous
   - SO2 causes aacidic rain
   -- CO(g)  poisonous
   - Pb / pb2+ is poisonous / affect the nervous system.
  vi) To reduce the unreacted Pbs to Pb Lead
 b) Hard water cpmataom ca2+  and Mg2+ .  These ions form a protective layer
of CaCO3(s) on the lead.  This prevent Lead from dissolving hence no Lead poisoining.  Soft does not form these deposite.
 c) - Radioactive shilding
  - Lead/acid accumulators
  - Making alloy soldering wire
  - Making of anti-knock additive
  - Manufacture of paint
  - Manufacture of bullets
  - Manufacture of ball bearings.
13. a) Electrolysis /Hall / Hertoult cell
 b) Al2O3: 2H2O
 c) i) Iron (III) Oxide /Silica
  ii) Add hot concentrated NaOH(aq) /KOH(aq)  silica and Al2O3 oxide
dissolves.
Carbon (IV) oxide then add water and finally add Al(OH)3 to the filtrate to precipitate Al(OH)3(s).  Filter the Al(OH)3(s) and Silica will remain in the solution.
 d) Tolower the melting of Aluminium oxide from 2015 to 8500C/ also act as
an electrolyte.
 e) Oxygen gas produced at the graphite anode.  Carbon anode react with the
oxygen to form Carbon (IV) Oxide.
 f) Aluminium react with Oxygen to form Aluminium oxide which protect
aluminium from further corrosion.
14 a) i) Calcium silicate/Calcium aluminate
  ii) - Magnetite Fe3O4
- Siderite feCO3
- Pyrite Fes
iii) Carbon (Iv) Oxide
 b) Hot compressed air oxidizes coke Co2
  C(g)  + O2(g)                CO2(g)
  CO2(g)  + C(s)              2CO(g)
  Co / Carbon (II) Oxide reduces Fe2O3 to (iron)
  3CO2(g)   + Fe2O3(s)                    2Fe(s)  + 3CO2(g)
 c) Decompose to give CaO / calcium oxide which combine with silica and
Aluminium oxides/ impunities to remove them as slag.
 d) It contains many impurities such as carbon, and manganese.
 e) - Construction of bridges/ ship/ buildings
- Car bodies, nail, railway lines pipes, spoons, pressure cookers.
- Horse shoe magnet.
15. a) i) - Effervescence and brownish gas produced.
          - Blue solution formed
  ii) Dilute HCL is not an oxiding agent.
  iii) I.  Cu(s) + 4HNO3(aq)                  CU(NO3)4(aq)  +2NO2(g)  +2H2O(l)
   II.    Moles of CU = 0.5 = 0.007874
            63.5
   Moles of HNO3 = 0.007874 x 4 = 0.31496
   Volume of HNO3 = 0.00314 x 1000 = 10.49cm3
                                                   3
 b) Step 4:  Neautralization
  Step 5:  Displacement
 c) - Resistant to corrosion
  - It is tough, / strong metal
16. i) Extraction of Aluminium
 ii) Adding hot half concentrated sodium Hydroxide,
 iii) To melt it, so so as to make the ions mobile/make it an electrolyte.
 iv) Al2O3 os a stable ions compound which can onlybe reduced by
electrolysis.  Aluminium is more reactive than carbon.  
v) -Light metal
-Strong and durable
-Not easily corroded
17. i) I)  Carbon (II) Oxide / Carbon (IV) Oxide
  II) Dilute Sulphuric acid
  Chamber I
 ii) Zno(s)   +   C(s)                     CO(g) + zn(s)
  Roaster
  2Zns(s)   + 3O2(g)         CO(s)  +  Zn(s)
   Chamber II
  Zn(s)  + H2SO4(aq)                    ZnSO4(aq) + H2(g)
 iii) I:   Mass of zns  = 45  x250  =  112.5g
                       100
  II:   22 zn S(s)  +  3O2(g)                2SO2(g)  + 2ZnO(s)
   Moles of ZnS  = 112.5  =  1.16 moles
        97.4
   Volume of SO2     1.16 Moles
   Volume of So2 = 1.16 x 24 = 24.72 dm3
 b) - Cause acidic rain
  - SO2 is poisonous
 c) Contact process: SO2 (by product) can be used to manufacture sulphuric
18. i) Sulphur
 ii) Sulphur (IV) oxide
 iii) SO2(g)  + 2NaON(aq)                    NCl2SO3(aq)  +  H2O(l)
19. i) Physical change:  because there is no change in mass of iron (III) oxide.
 ii) Iron is more electropositive than hydrogen and less than carbon.
        iii) Hydrogen gas
 b) i) Hydrated iron (III) oxide
  ii) Aluminium form a coating of an oxide (AL2O3) which prevent
further corrosion.
  iii) Zinc is more reactive than iron so it loses it’s electrons more easily
than iron.  Hence zinc corrode before iron.
20. They are in the same group.
21. A
22. -Resistant to corrosion
 -Light metal
23. a) “M”  b) “L”
24. a) Dissolve the ore in Nitric acid.  To the filtrate add sodium Hydroxide
dropwise till in excess;  Brown / iron(III) ions (Fe3+ or add Ammodia solution till in excess again to obtain reddish brown precipitate of iron (III) ions Fe3+.
 b) i) Mass of iron Oxide
   = 13.30 – 10.98 = 2.32g
   Mass of iron/residue  = 12.66 – 10.98 = 2.32 g
   Mass of oxygen =  2.32- 1.68 = 0.64g
Elements Fe O
Mass 1.68 0.64
Moles 1.68  =0.03
56 0.64 = 0.04
16
Ratio 0.03 = 1
0.03 0.04  = 1.3
0.03
X 3 4
   Fe3O4)n  =  232
   (232)n  =  232:  n = 1 MF = Fe3O4
            ii) Fe3O4(s)  + 4CO(g)            3Fe(s)  + 4 CO2(g)
 c) i) - Moisture
   - Oxygen
  ii) - Galvanising
   - Painting/greasing
   - Plastic coating
   - Alloying
 d) Salt accelerate the rate of rusting /corrosion
25. Sodium atom has a large atomic radius and losses electrons very easily compared to Lithium which has a small atomic radius/Lithium outer most electrons are strongly attracted by the nucleus protons hence not easily removed.
 
TOPIC 5
ORGNIC CHEMESTRY 2
1. The ionic “head” lowers the surface tension of water faciliatating mixing of water and grease.  The non polar “tail” mix with grease, dislodging it from the fabric.
2. 
Name of polymer Name of monomer Use of polymer
Polystyrine Styrene
Phenythene Insulation, plastic pipes, biros, artificial rubber
Polyvinyl
Chloride
Polychloro
Ethane Vinyl chloride
Chloroethene Insulation of electric cables plastics tanks
3. “B”:  “B” does not form scum
4. a) Ethanol H H
    │ │
   H   — C   — C   — OH
    │ │
    H H
    H H  O
    │ │
Propanoic acid  H   — C   — C   — C   — OH
    │ │
    H H
 b) Alkanols / alcohols
5. a) Perspex/polymethyl/methacrylate
 b) As a substitute for glass in manufacture of
  - Safety screen
  - Plastic lens
  - Wind screens
6. a)  H H H H  
   │ │ │ │
  H   — C   — C   — C   — C   — OH
   │ │ │ │
   H H H H
 b) Alkanols/alcohols
 c) 2C4HqOH(l) + 2K(s)              2C4HqOK(s) + H2(g)
7. There is a constant increase in mass caused by constant addition of –CH2
8. a) N - Sodium ethanoate/CH3COO Na/sodium acetate.
  P - Methane/CH4
 (b) Substitution
9. Esterification
10. Penten -1-al is polar.  There are two forces, vanderwaals and hydrogen bonds holding its molecule together. Pentane is none polar.
11. a) CH3(CH2)12  COOONa
 b) Soapy detergent
 c) (CH3 (CH2)12 COO)2Ca/(CH3(CH2) 12COO)2mg
12. Butanoic acid and propanaol
13. i) Pentanoic acid
 ii) C3 H6O
 iii)

  I.  166 +-0.60C
iv) The boiling point increases with increase in mass.  The molecular mass increase by –CH2 Unit (14 units) this causes an increase in intermolecular forces between molecules. Hence more heat is required to bread the bonds in complex molecules.
 b) Effervescence /colourless gas is given off.  This is CO2 and it forms white
precipitate with lime water.
 c) CH3COOH + NaOH                CH3COONa + H2O
  Moles of CH3COOH    = 30      =  0.05 moles
          60
  Moles of NaOH = 0.05 moles
  Volume of Naoh = 0.05 x 1000 = 250cm3
     0.2
14. a) Pysteyrene/polythenyl ethane
 b) Cause pollution since it is non biogredable.
15. a) i) propanoic acid
    H H  O
    │ │
   H   — C   — C   — C   — OH
    │ │
    H H
  ii) Ester
 b) The colour of solution chante from orange/yellow to green because Cr + 6
is reduced  to reduced to Cr 3 + and ethanol is oxidized to ethanoic acid.
 c) i) Soap/soapy detergent
  ii) Sodium chloride
  iii) To make soap float
  iv) Potassium hydroxide/KOH
  v) A molecule of cleansing agent has polar and non polar parts.  Non
polar parts dissolves in oil and polar parts dissolves in water.  When the mixture is agitated the oil droplets coagulate and can be washed away with water.
16. a) i) CH3 OH
  ii) CH3COOH
 b) HCOOH (aq)  +  NaOH(aq)               HCOONa(aq) + H2O(l)
 c) i) Methyl methanoate /HCOOCH3
  ii) -Heat
   -Concentrated sulphuric acid
 d) i) Use of bromine water or acidified potassium manganate (VII).
   Hexane decolourises both at room temperature but hexane does
not.
  ii) -Fuel
   -Solvent
   -Manufacture of Hexanol and Hexanoic acid.
  iii) C6H12+ H2              C6H14
   Rmm of C6H12 = 84
   Moles of C6H2  = 42  =  0.5 moles
          84
   Moles of H2 = 0.5 moles
   Volume = 0.5 x 22.4 = 11.2dm3
17. i) I:  Oxidation
  II:  B = ethane
        C= sodium ethanoate
 ii) To bring the reacting monomers into close contact.
 iv) -As a fuel
  -In making carbon black
  -Manufacture of methanol
  -Manufacture of hydrogen cyanide
18. a) i)  I:  V1 and V3
      II:   V2 and V5
ii) V4:  It is unsaturated compound and during polymerization the double bond is broken to allow another monomer to combine.
 b) 
 Advantage Disadvantage
R-COO-
Na+ They are biogradable do not cause pollution Forms scum with Ca2+ and Mg2+
R- OSO3
Na They do not form scum with Ca2+ and Mg2+ (aq) They pollute the environment since they are non biogradable.
 c) i) Ester
  ii) CH3COOC2H3
  iii) -Used as solvent
   -Manufacture of drugs and chemicals
   -In flavouring and preservation of food
   -In manufacture of synthetic fibres
  iv) 2CH3COOH (aq)  + K2CO3(s)            2CH3COOK(aq) + H2O(l)
 d) i) Natural fibres include
   Rubber, Cellulose, Wool, starch, silk
  ii) Advantages of synthetic fibres.
- Can be made into complicated shapes more easily.
- Less expensive
- Resistant to corrosion
- Less dense and stronger
19. a) i) Change from orange to green.
  ii) Effervescence and a colourless gas which burn with a “Pop” sound
Produced.
 b) Step I:  Fermentation:  Glucose solution is mixed with yeast.  The
Enzymes from yeast convert glucose to ethanol.
Step II: fermentation: Dehydration: Ethanol is mixed with concentrated   sulphuric acid and heated in presence of AL2O3 as a catalyst.
e) i)  H H  O
    │ │
   H   — C   — C   — C   — O   — H
    │ │
    H H
  ii)  H  O H  
    │   │
   H   — C   — C   — O   — C   — H
    │   │
    H   H
 f) -Produces acidic – compounds which causes Global warming
-Produces acidic compounds which causes acidic rain.
20 a)
    H
    │
  C   ≡ C   — C   — H
  │  │
  H  H
 b) i) -High temperature (7000C) or
   -Produces acidic – compounds which causes acidic rain.
  ii) Ethane / C2H6
  iii) I. Polluting the environment / they are non biodegradable
   II. Hydrolysis
   III. Ethypropanoate
  iv)  H H
    │ │
      — C   — C   —
    │ │
    H H n = 16,800
    Monomer = 16800 = 600 monomers
      28
 c) i) “M” it is unsaturated with a double bond.  Its an alkene.
  ii) “N” It is an organic acid and will react with carbonate to give CO2.
21. i) Monomers of carbohydrates
ii) Condensation in which a molecule of water is eliminated between two monosacchararide.
22. i) Amino acids/ proteins
 ii) The carbon chain is linear
 iii) -Ester and water
  -Condition is –heat
  -Concentrated sulphuric acid/catalyst
23. i) ClCLH2COOH(aq)  + KOH(aq)  ---- CLCH2CH2COOK(aq)  + H2O(l)
 ii) Molarity = 2.635 x 1000 = 0.969 moles/dm3
             250
 iii) Moles of KOH  = 25x 0.1 = 0.0025
           250
 iv) Moles of acid = 0.0025 since ratio is 1:1
24. i) I:  -Concentrated sulphuric acid
   -Heat
  II:   Excess acidified potassium manganate (VImanganate (VII)
  III:  Sodium metal
  IV:  -Sulphuric acid
         -Ethanoic acid
 ii) -Textile /clothing
  -To make ropes
  -Safety bolts
  -Lents
  -Sails
 d) i) -Rubber
   -Cellulose, Wool, silk, Starch, Protein
  ii) Heating rubber with sulphur so as to make it strong hard and tough
 e) i)  CH3
    │
   CH2  ═ C   — CH  ═  CH2
  ii) 2, melty but -1, 3 – diene
25. CH3CH2OH(l)   Conc H2SO4  CH2 = CH2(g)  + H2O(l)
    1700C
26. i) C  ii) B
27. Ethanoic acid (CH3COOH) reacts with sodium Carbonate to liberate Carbon (IV)
Oxide while Ethanoic does not.
28. a) Reagent : Hydrogen gas
  Conditions :  Heat
    Nickel catalyst
 b) Catalytic cracking using asbestos as a catalyst and heat
 c) -Ozonised oxygen at 00C
  -Water
  -Acidified potassium Dichromate
29. a) A substance that improve the cleasing power of water.
  Advantages
 b) -Forms lather easily in both soft and hard water
  -Not alkaline or acidic
  Disadvantaged
- Non biodegradable
- Environmental pollution
- Eutrophication in water.
c) Polar end (_C00-) dissolves in water to form micable.  Non polar end
(CH(CH2)-) attract the greese /dirt.  The grease is then carried off while attracted to the non polar end linked to water to the polar end as a co agulant.
 d) To avoid scum formation in hand water by complexing with calcium and
magnesium ions.
 e) Add a little fat/oil to aqueous Sodium Hydroxide and boil for some time. 
Add saturated sodium, Chloride to precipitate out soap (salting out) filter and dry to obtainin solid soap which can then be made into flakes.
 
TOPIC 6
RADIOACTIVITY
1. a) 14  -0  14
            7        1                     6
b) - Nuclear reacter
  - Atomic bombs
  - Detecting leakage
  - Studying photosynthesis
 - Security measurement
 - Treatment of cancer
 - Sterilize surgical instruments
 - Dating
 - Killing bacteria
2. 384    t ½     192  t ½    192   t ½   96    t ½   48
    
  4t ½  = 270
  T ½  = 270  =  67.5 days
     3
3. a) 100   t ½    50      t ½    23    t ½     12.5
    3t ½ = 81
  T ½ = 81     = 27 days
   3
b) Mass number 233
 Atomic number 92
4. 
a  = Alpha particle
 b  = Gamma Ray
 c  = Beta particles
5. a. Time taken for a given mass of radioactive isotope to reduce to half.
 b. 100 = 4 half litres
   25
  80   40 20 10 5
  Original mass = 80g
6. a) 234  4      +2 230
       U                  He            +         Th
  94                   2                       92
 b) Some rays e.g gamma will penetrate through aluminium and may cause
the biological damage to the organisms.
7. a) t ½ 8   0.5 day
 b) 10         5   2.5   -   1.25        0.625
  32
   6
  Mass remaining = 0.625g
8. a) 222
       X
   86
 b) 1            ½               ¼              1/8             1/16
  4t ½ = 112
     T ½ = 112  = 28 days
      4
9. a) Atoms of the same element which have the same atomic numbe but
different mass numbers.
   0           14
 b) 146C                    C   +      N
         7
 c) -Dating young fossils
  -Isotopic tracer
  -tracking of biological process
10. a) Alpha
 b) 210      210           0
       J        K    +          e
81       82  -1
 c) “K” and “M”
11. a) 100   T½  50 T½ 25 T½ 12.5
   
  3 half   = 15.6 years
  T ½ = 15.6 years
    3
12. a) 37   0  37
      A                e        B
  18  -1               17
 b) i) Radioactive traces
  ii) - Causes cancer
- Cell mutation
13. a) Nuclear fusion is where two light nuclei cobine to give a heavy nucleus
with release of energy while nuclear fission is where a large nucleus splits into smaller nuclei with the release of enormous amount or energy.
 b) Wrap with aluminium or lead foil and bury them deep underground
14. a) 4 He +2   or  4He
  2                 2
 b) i) z1  →    235
   z1  →      54
  ii)    Nuclear fission
15 a)
Nuclear reactions Chemical reactions
Inolves protons and neutrons Involve valency electrons
Reaction rate not affected by element changes Reaction rate is influenced by element changes
Involve huge amount of energy Involve little amount of energy
There is change in mass No change in mass
 b) i) 1:  Alpha  II: beta
  ii) 210              206  4
        PO                   Pb     +      He 
   84   82       2
 c) i)
  ii) 1 2o minutes
   II % value at 70 minutes = 9%  2
   Mass = 0.16 x 100 = 1.778(g)
        9
d) - Treatment of cancer
  - Sterlization of surgical equipment
  - Treatment of leation of goiter
  - Regulate heat pace maker
  - Detection of blood circulation disorders
  - Measure of uptake of iodine.
16. a) Time take by radioactive isotopes to decay to half its mass.
 b) 

17. a) 54  55   0
        Cr                    Mn+             e
  24                   25                 -1
         235                143            1       90
 b) 10n +   u  La + 3   n+       Br
                                92        57             0      35
18. a)
 b) Half life = 22minutes
 c) 110  = 5 half lifes
   22
      32            16 8 4 2 1
                             1g will remain
19. 234  234  234
      X,       Y,                   Z
 90  91  92
20 i) The particles go through the inter atomic space in the metal foil because of
their small size.
 ii) Since the particles are positively changed, there which approach the
nucleus are repelled.
 iii) Those with low energy cannot over come the repulsive forces hence are
abosorbed.
21. Days 2.5  5 7.5
 Percentage ion  50   25%
  25% remains
22. a) X : Alpha
  Y : Gamma
  Z : Beta
 b) They are very heavy/less
  Penetrative /have large mass
23. i) 239  1   ii)    238            16
       CF      +        n           U    +           O
  38           O             92           8
 iii) 241  241        O
          PU      AM +        e
  94  95                -1

 

 

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