CHEMISTRY FORM THREE TOPICAL EXAMINATION

CHEMICAL EQUATIONS

TIME: 2½ HOURS

INSTRUCTIONS

1. This paper consists of Sections A and B.
2. Answer all questions.
3. Write all answers in your answer booklet.
4. Show working where calculations are required.
5. Use blue or black ink, except for diagrams, which should be drawn in pencil.

SECTION A (20 MARKS)

1. Choose the correct answer for each item.

(a) The half-equation X²⁺ + 2e⁻ → X represents:

1. Oxidation
2. Reduction
3. Displacement
4. Decomposition
5. Neutralisation

(b) The reaction between aqueous silver nitrate and aqueous sodium chloride to form silver chloride and sodium nitrate is:

1. Combination
2. Single displacement
3. Double displacement
4. Thermal decomposition
5. Combustion

(c) The correct net ionic equation when ammonium chloride is warmed with sodium hydroxide solution is:

1. H⁺(aq) + OH⁻(aq) → H₂O(l)
2. NH₄⁺(aq) + OH⁻(aq) → NH₃(g) + H₂O(l)
3. Na⁺(aq) + Cl⁻(aq) → NaCl(aq)
4. NH₄⁺(aq) + Cl⁻(aq) → NH₄Cl(s)
5. NH₄⁺(aq) → NH₃(g) + H⁺(aq)

(d) A compound contains 40.0% carbon, 6.7% hydrogen and 53.3% oxygen by mass. Its empirical formula is:

1. C₂H₄O
2. CH₂O
3. CH₄O
4. C₂H₂O
5. CHO

(e) What mass of water forms when 1.0 g of hydrogen reacts completely with excess oxygen?

1. 1.8 g
2. 9.0 g
3. 4.0 g
4. 18.0 g
5. 0.5 g

(f) The correct net ionic equation for the reaction of hydrochloric acid with sodium hydroxide is:

1. Na⁺(aq) + Cl⁻(aq) → NaCl(s)
2. Na⁺(aq) + OH⁻(aq) → NaOH(s)
3. H⁺(aq) + Cl⁻(aq) → HCl(aq)
4. H⁺(aq) + OH⁻(aq) → H₂O(l)
5. 2H⁺(aq) + OH⁻(aq) → H₂O(l)

2. Match each description in List A with the most appropriate response in List B.

SECTION B

3. Complete and balance the following chemical equations:

1. Zn(s) + Cl₂(g) →
2. CaO(s) + H₂O(l) →
3. Li(s) + N₂(g) → Li₃N(s)
4. K₂CO₃(aq) + HCl(aq) →
5. Pb(NO₃)₂(aq) + Na₂SO₄(aq) →

4. Write the balanced net ionic equations for the following reactions:

1. Magnesium reacts with dilute sulphuric acid.
2. Silver nitrate solution reacts with iron(II) chloride solution.
3. Dilute hydrochloric acid reacts with solid calcium carbonate.
4. Iron displaces copper from copper(II) sulphate solution.
5. Chlorine gas displaces iodine from potassium iodide solution.

5. (a) Define the following terms:

1. Ionic equation
2. Molecular equation

(b) Complete and balance the following chemical equations:

1. KClO₃(s) → [heat, MnO₂ catalyst]
2. Pb(NO₃)₂(s) → [heat]
3. H₂SO₄(aq) + KOH(aq) →
4. CuCO₃(s) → [heat]
5. Na(s) + H₂O(l) →

6. Explain what happens in each situation below, using a balanced chemical equation where a reaction occurs:

1. Potassium carbonate is heated strongly.
2. Concentrated sulphuric acid is added to dry sodium chloride and the mixture is warmed.
3. Sodium nitrate is heated strongly.
4. Concentrated sulphuric acid is added carefully to sugar.
5. Aqueous sodium hydroxide is added to aqueous copper(II) sulphate.
6. Carbon dioxide is bubbled through clear limewater.
7. Water is added to white anhydrous copper(II) sulphate.
8. A glowing wooden splint is introduced into separate jars of hydrogen and carbon dioxide.
9. Solid ammonium chloride is heated.
10. Dilute nitric acid reacts separately with (i) warm copper(II) oxide and (ii) zinc carbonate.

7. Write the products and balance the following chemical equations:

1. AgNO₃(aq) + NaCl(aq) →
2. Zn(s) + H₂SO₄(aq) →
3. MgCl₂(aq) + AgNO₃(aq) →
4. Na₂SO₄(aq) + BaCl₂(aq) →
5. ZnCO₃(s) + HCl(aq) →
6. Na(s) + H₂O(l) →

8. Complete and balance the following chemical equations:

1. Fe(OH)₃(s) + H₂SO₄(aq) →
2. MgCO₃(s) + HCl(aq) →
3. Na(s) + O₂(g) → Na₂O(s)
4. H₃PO₄(aq) + Al₂O₃(s) →

9. Explain each type of chemical reaction and give one relevant balanced chemical equation:

1. Synthesis (combination) reaction
2. Decomposition reaction
3. Precipitation reaction
4. Single displacement reaction
5. Neutralisation reaction

10. (a) Define the following terms:

1. Redox reaction
2. Oxidation reaction
3. Reduction reaction

(b) Identify oxidation and reduction in each of the following, with a brief reason:

1. S(s) + O₂(g) → SO₂(g)
2. N₂(g) + 3H₂(g) → 2NH₃(g)
3. Fe²⁺(aq) → Fe³⁺(aq) + e⁻
4. Fe³⁺(aq) + e⁻ → Fe²⁺(aq)

(c) Write balanced chemical equations for:

1. Hydrogen reacts with chlorine to form hydrogen chloride.
2. Magnesium oxide is reduced by carbon to form magnesium and carbon monoxide (at sufficiently high temperature).
3. Silver nitrate reacts with sodium chloride to form silver chloride and sodium nitrate.

11. (a) Explain the meaning of a net ionic equation.

(b) Write balanced net ionic equations for:

1. Carbon dioxide reacts with aqueous potassium hydroxide to form potassium carbonate and water.
2. Magnesium reacts with dilute hydrochloric acid.

CHEMISTRY FORM THREE TOPICAL EXAMINATION

HARDNESS OF WATER

TIME: 2½ HOURS

INSTRUCTIONS

1. This paper consists of Sections A, B and C.
2. Answer all questions.
3. Write answers in your answer booklet.
4. Show working where calculations are required.
5. Use blue or black ink, except for diagrams, which should be drawn in pencil.

SECTION A (20 MARKS)

1. Choose the correct answer for each item.

(a) Hard water that can be softened by boiling contains dissolved:

1. Calcium hydrogen sulphate
2. Magnesium chloride
3. Magnesium carbonate
4. Calcium sulphate
5. Calcium hydrogen carbonate

(b) Which compound can be used to remove hardness of water?

1. Calcium chloride
2. Sodium carbonate
3. Magnesium hydroxide
4. Potassium hydroxide
5. Zinc chloride

(c) Which dissolved substance causes permanent hardness of water?

1. Carbon dioxide
2. Calcium sulphate
3. Magnesium carbonate
4. Sodium carbonate

(d) One disadvantage of hard water is that it:

1. Always corrodes water pipes
2. Causes increased tooth decay
3. Requires more soap for washing
4. Always contains harmful minerals

(e) Hard water that can be softened by boiling contains dissolved:

1. Calcium carbonate
2. Calcium chloride
3. Sodium carbonate
4. Magnesium sulphate
5. Calcium hydrogen carbonate

(f) Which method removes both temporary and permanent hardness of water?

1. Ion exchange
2. Boiling
3. Adding calcium hydroxide only
4. Simple filtration

(g) Which of the following is NOT an advantage of hard water?

1. Provides calcium and magnesium ions
2. Can contribute to healthy bones and teeth
3. May be preferred for its mineral taste
4. Forms scum and wastes soap

(h) One advantage of soft water is that it:

1. Forms lather readily with soap
2. Always contains more calcium
3. Always has a better taste
4. Is used to treat ulcers

(i) Which method produces water free of dissolved hardness-causing salts when correctly carried out?

1. Boiling
2. Distillation
3. Adding washing soda
4. Adding calcium hydroxide

(j) The main natural reservoir of salt water on Earth is:

1. Seas and oceans
2. Freshwater springs
3. Glaciers
4. Freshwater wells

2. Match the descriptions in List A with the responses in List B.

SECTION B

3. Write balanced chemical equations showing how:

(a) Temporary hardness is removed by:

1. Boiling (one equation)
2. Adding a suitable chemical (two equations)

(b) Permanent hardness is removed by:

1. Adding washing soda (one equation)
2. Ion exchange (one representative equation)

4. Define the following terms:

1. Soft water
2. Hard water
3. Permanent hardness of water
4. Temporary hardness of water

5. (a) What is hardness of water?

(b) Briefly explain the types of water hardness.

(c) State the causes of each type of hardness.

(d) Explain how each type of hardness can be removed.

(e) State three advantages and three disadvantages of hard water.

6. Equal volumes of water from four sources were tested with soap solution. The volumes of soap required to form a lather persisting for 30 seconds are shown below.

(a) Which untreated sample is the hardest? Give a reason.

(b) Which sample behaves like distilled water? Explain.

(c) Name one compound that might cause the hardness of (i) sample A and (ii) sample B.

(d) Write a balanced equation showing one way of removing hardness from sample C.

7. (a) State two advantages of hard water.

(b) State two disadvantages of hard water.

(c) Give two methods of removing temporary hardness of water.

8. (a) Distinguish between:

1. An acidic salt and a basic salt
2. A hydrated salt and an anhydrous salt

(b) Vinegar, lemon juice and yoghurt taste sour. State three other properties you would expect from their acidic solutions.

SECTION C

9. (a) Explain the meaning of:

1. Basicity of an acid
2. The pH scale

(b) Name substances that cause:

1. Temporary hardness of water
2. Permanent hardness of water

(c) Use one balanced equation in each case to explain how:

1. Temporary hardness is removed by boiling
2. Permanent hardness is removed by a chemical method

10. (a) Explain why pure water has no effect on red or blue litmus paper.

(b) (i) What may happen to a tightly stoppered, completely filled glass bottle of water left in a freezer overnight? Explain.

(ii) Explain why untreated hard water can cause problems in steam pipes and boilers.

(c) (i) Name two compounds that cause temporary hardness and two ions mainly responsible for permanent hardness.

(ii) Write one balanced equation for removing temporary hardness by boiling and another for removing permanent hardness with sodium carbonate.

11. (a) Define:

1. Base
2. Alkali
3. Weak acid
4. Hygroscopic substance

(b) With the aid of balanced chemical equations, describe how temporary and permanent hardness of water can be removed.

CHEMISTRY FORM THREE TOPICAL EXAMINATION

ACIDS, BASES AND SALTS

TIME: 2½ HOURS

Name: ______________________________ Class: __________________

INSTRUCTIONS

1. This paper consists of Sections A, B and C.
2. Answer all questions.
3. Write all answers in your answer booklet.
4. Show all working for calculations.
5. Use blue or black ink, except for diagrams, which should be drawn in pencil.

SECTION A: OBJECTIVE QUESTIONS

1. Choose the correct answer for each item.

(i) Sea water contains various salts. Which salt is present in the largest proportion?

1. Magnesium sulphate
2. Sodium chloride
3. Calcium sulphate
4. Magnesium chloride

(ii) Which pair consists of substances commonly used as drying agents?

1. Calcium carbonate and lead(II) chloride
2. Anhydrous calcium chloride and concentrated sulphuric acid
3. Sodium chloride and calcium carbonate
4. Concentrated nitric acid and concentrated hydrochloric acid

(iii) A weak acid is best described as an acid that:

1. ionises only partially in aqueous solution
2. is always dilute
3. does not react with other substances
4. is harmless

(iv) What happens when aqueous calcium chloride and sodium carbonate are mixed?

1. A white precipitate of sodium chloride forms
2. A white precipitate of calcium carbonate forms
3. Only a colourless solution forms
4. Both sodium chloride and calcium carbonate precipitate

(v) An insoluble salt such as lead(II) chloride is usually prepared by:

1. crystallisation
2. precipitation
3. thermal decomposition
4. evaporation

(vi) A hygroscopic substance:

1. evaporates into air
2. loses water to air
3. adds water chemically to another compound
4. absorbs moisture from air
5. removes the elements of water from another compound

(vii) An efflorescent substance:

1. evaporates into air
2. loses water of crystallisation to air
3. absorbs moisture from air
4. removes elements of water from another compound
5. adds water to another compound

(viii) Which of the following is a weak acid?

1. Sulphuric acid
2. Hydrochloric acid
3. Ethanoic acid
4. Nitric acid

(ix) An aqueous solution with pH 1 is:

1. slightly acidic
2. slightly alkaline
3. strongly acidic
4. neutral

(x) Which substance is NOT a product when solid silver nitrate is strongly heated?

1. Silver oxide
2. Silver metal
3. Nitrogen dioxide
4. Oxygen

2. Match each description in List A with the correct response in List B.

SECTION B: STRUCTURED QUESTIONS

3. (a) Acids react with metal carbonates, hydrogen carbonates, metal oxides and metal hydroxides. Illustrate each reaction with a balanced chemical equation. Include an equation showing how carbon dioxide is tested using limewater.

(b) Explain how to prepare a flower-extract indicator and use it to distinguish acidic and alkaline solutions.

4. Distinguish between the following pairs of terms:

1. Hydrated and anhydrous salts
2. Efflorescence and deliquescence
3. Base and alkali
4. Acid and base
5. Acid salts and normal salts

5. (a) State the expected observation in each test.

5. (b) Complete the indicator colour table.

6. (a) Heating 488 g of BaCl₂·nH₂O produces 416 g of anhydrous BaCl₂. Calculate n. (Ba = 137, Cl = 35.5, H = 1, O = 16).

(b) Define (i) an acid and (ii) a 1.0 mol dm⁻³ acid solution. (iii) State three characteristics of acids.

7. Complete and balance the following thermal decomposition equations:

1. PbCO₃(s) →
2. Ca(NO₃)₂(s) →
3. (NH₄)₂CO₃(s) →
4. NH₄Cl(s) ⇌

8. (a) Distinguish between strong and weak acids. (b) Explain three applications of neutralisation.

9. A student pipetted 25.0 cm³ of potassium hydroxide solution into a conical flask, added phenolphthalein and titrated it with dilute hydrochloric acid. The endpoint was reached after 21.0 cm³ of acid was added.

(a) (i) Name the apparatus used to measure 25.0 cm³ of alkali accurately. (ii) State the initial indicator colour. (iii) State the colour at the endpoint.

(b) (i) Compare the molar concentrations of the acid and alkali, assuming the volumes stated are exact and the reaction is complete. (ii) Name the salt formed. (iii) Write the balanced equation. (iv) Classify the salt as normal or acid salt and explain.

(c) (i) Explain how pure crystals of the salt could be obtained using the titration results. (ii) State whether the salt is soluble or insoluble and give a reason.

10. A water sample forms scum with soap even after it has been boiled and cooled. (a) Identify the type of hardness. (b) Suggest one treatment to soften it.

SECTION C: APPLICATION AND EXTENDED QUESTIONS

11. (a) Define an acid–base indicator. (b) Give two uses of the pH scale. (c) Give two uses of salts.

12. (a) Define (i) basicity of an acid and (ii) the pH scale.

(b) Name dissolved substances that cause (i) temporary hardness and (ii) permanent hardness of water.

(c) Use balanced chemical equations to show (i) how boiling removes temporary hardness and (ii) how sodium carbonate removes permanent hardness.

(d) Distinguish between an acid salt and a basic salt. Classify the following as normal, acid or basic salts: PbSO₄, Mg(HSO₄)₂, Zn(OH)Cl, Mg(HCO₃)₂, NH₄HSO₄ and Ba(NO₃)₂.

13. (a) Define neutralisation and give a balanced equation. 

(b) Name and give the formulae of three common laboratory acids. 

(c) Define an indicator and give four examples. 

(d) State the usual products when an acid reacts with 

1. a reactive metal and 
2. a metal carbonate; give one balanced equation for each.

14. (a) Name the element present in all common acids. 

(b) Describe two tests for identifying an acidic aqueous solution, giving the expected result. 

(c) If dilute acid spills on clothing, which of the following is most suitable for cautious neutralisation: sodium hydroxide solution, sodium hydroxide pellets, concentrated ammonia solution or sodium hydrogen carbonate? Explain, and state an immediate first-aid action.

CHEMISTRY EXAMINATION – FORM THREE

TOPICAL EXAMINATION: ENERGY AND FUELS

Time: 2½ Hours

INSTRUCTIONS

1. This paper consists of Sections A, B and C.
2. Answer all questions.
3. Write your answers clearly in the answer booklet.
4. Use blue or black ink; use pencil for diagrams.

SECTION A: OBJECTIVE QUAESTIONS

1. Choose the correct answer for each of the following items.

i. Which of the following energy transformations produces hydroelectric power (HEP)?

1. Electrical energy to heat energy
2. Electrical energy to mechanical energy
3. Mechanical energy to electrical energy
4. Mechanical energy to chemical energy

ii. Which of the following statements about coke as a fuel, compared with coal, is incorrect?

1. It does not produce carbon dioxide gas
2. It does not produce poisonous gas
3. It has a higher heat content
4. It is clean and smokeless

iii. Which of the following is a characteristic of a good fuel?

1. It must be firewood
2. It is easily available
3. It must melt
4. It can burn houses into ashes

iv. What is kindling temperature?

1. A kind temperature
2. Temperature outside a burning material
3. The highest temperature obtained from a burning substance
4. The lowest temperature at which a combustible material catches fire

v. Which of the following is a renewable source of energy?

1. Coal
2. Petroleum
3. Biogas
4. Natural gas

vi. Which of the following is not a primary source of energy?

1. Wood
2. Kerosene
3. Crude oil
4. Natural gas

vii. Which of the following is not a characteristic of a good fuel?

1. It is readily available
2. It has a low energy value
3. It is affordable
4. It is easy to transport

viii. In hydroelectric power generation, which energy transformation takes place?

1. Electrical energy to heat energy
2. Electrical energy to mechanical energy
3. Mechanical energy to electrical energy
4. Chemical energy to electrical energy

ix. Which of the following is a source of chemical energy?

1. Natural gas
2. Biomass
3. Coal
4. Solar energy

x. Which of the following is not an advantage of biogas?

1. It can be a cheaper source of energy
2. It pollutes the environment
3. It is a renewable source of energy
4. It can create employment opportunities

2. Match the descriptions in List A with the correct terms in List B.

SECTION B: STRUCTURED QUAESTION

3. Electricity can be used to perform work and for heating. Explain how electricity is used in each case and give four practical examples for each use.

4.

1. What is fuel?
2. Mention three categories of fuel and give two examples for each category.
3. Give two pieces of advice to Tanzanian communities on how to reduce charcoal and firewood consumption through efficient use of fuel.

5.

1. Classify each of the following fuels as primary (P) or secondary (S): wood, petrol and coal gas.
2. Identify the gaseous fuel produced by each of the following processes:
   1. Destructive distillation of coal
   2. Destructive distillation of wood
   3. Passing air over red-hot coke at about 1000°C

6.

1. Write the chemical symbols of the following elements:
   1. Argon
   2. Lead
   3. Silver
   4. Manganese
   5. Barium
   6. Xenon
2. Write the chemical formulae of the following compounds:
   1. Potassium sulphate
   2. Copper(II) carbonate
   3. Aluminium oxide
   4. Magnesium hydroxide

7. Explain the major methods used to obtain artificial fuels from natural fuels in each of the three physical states: solid, liquid and gaseous.

SECTION C: EXTENDED RESPONSE QUESTIONS

8. With the aid of a well-labelled diagram, explain how destructive distillation of coal can be carried out experimentally.

9.

1. Define fuel.
2. Give two examples of each of the following categories of fuels:
   1. Solid fuels
   2. Liquid fuels
   3. Gaseous fuels
3. State any two qualities of a good fuel.

10.

1. Define the following terms:
   1. Greenhouse effect
   2. Global warming
   3. Acid rain
   4. Pollution
2. Give four examples of greenhouse gases.

11.

1. Define the following terms:
   1. Fuel
   2. Calorific value of a fuel
   3. Energy value of a fuel
2. Give two examples of each of the following:
   1. Solid fuels
   2. Liquid fuels
   3. Gaseous fuels
3. Name four characteristics of a good fuel.

CHEMISTRY EXAMINATION FORM THREE

TOPICAL EXAMINATION: MOLE CONCEPT

TIME: 2½ HOURS

NAME: __________________________ 

CLASS: __________________________

INSTRUCTIONS

1. This paper consists of Sections A, B and C.
2. Answer all questions.
3. Write your answers in your answer booklet.
4. Use blue or black ink, except for drawings, which should be done in pencil.

SECTION A: MULTIPLE CHOICE QUESTIONS

1. Choose the correct answer for each of the following items.

i. If 0.9 g of calcium metal is burnt in air, the mass of powder formed is:

1. 1.14 g
2. 1.18 g
3. 12 g
4. 1.26 g

ii. If 1 g of hydrogen is exploded in air, the mass of water formed is:

1. 1.8 g
2. 9 g
3. 4 g
4. 18 g

iii. What mass of pure sulphuric acid is found in 400 cm³ of its 0.1 M solution?

1. 2.45 g
2. 9.80 g
3. 3.92 g
4. 4.90 g

iv. The volume of 18 M concentrated sulphuric acid that must be diluted with distilled water to prepare 10 litres of 0.125 M sulphuric acid is:

1. 69.44 cm³
2. 22500 cm³
3. 225 cm³
4. 4440 cm³

v. If two jars labelled W and Z contain 22.4 cm³ of oxygen gas and 22.4 dm³ of nitrogen gas at STP respectively, then it is true that:

1. There were 6.02 × 10²³ oxygen molecules in jar W and 6.02 × 10²³ nitrogen molecules in jar Z.
2. 6.02 × 10²³ oxygen atoms were in jar W and 6.02 × 10²³ atoms of nitrogen in jar Z.
3. There were 12.4 × 10²³ molecules of oxygen and nitrogen in the gas jars W and Z.
4. 6.02 × 10²³ molecules of oxygen and nitrogen were in the two jars W and Z.

vi. Which of the following statements is true about one mole of CO₂?

1. It contains Avogadro’s number of molecules.
2. It contains 6.023 × 10²³ molecules at room temperature.
3. It contains 44 g of carbon and 16 g of O₂.
4. It contains 12 g of carbon and 16 g of O₂.

vii. The loss in mass when 100 g of marble is heated to a constant mass is:

1. 200 g
2. 88 g
3. 44 g
4. 22.4 g
5. 22400 g

viii. The number of atoms present in 0.025 moles of iron is:

1. 6.02 × 10²³ atoms
2. 56 atoms
3. 5.6 × 10²³ atoms
4. 22.4 × 10²³ atoms
5. 1.51 × 10²³ atoms

ix. How many atoms of zinc are there in 165 g of zinc?

1. 1.505 × 10²⁵
2. 1.0997 × 10²²
3. 2.2586 × 10²³
4. 9.63 × 10²⁵

2. Match the descriptions in List A with the correct terms in List B by writing the letter of the correct answer beside the corresponding item number.

SECTION B: STRUCTURED QUESTIONS

3. With the help of a chemical equation, calculate the volume of carbon (IV) oxide produced when 20 g of calcium carbonate is heated at standard temperature and pressure (STP). (Ca = 40, C = 12, O = 16)

4. What would be the volume of the gas produced at 20°C and 750 mmHg? (Molar gas volume at STP is 22.4 litres.)

5. 25.0 cm³ of sodium hydroxide solution containing 4.0 g/dm³ were required for complete neutralization of 0.18 g of a dibasic acid. Calculate the relative molecular mass (RMM) of the acid. (Na = 23, H = 1, O = 16)

6. Answer the following questions:

1. How many molar volumes of 132.0 g of CO₂ are there at STP?
2. Determine the number of molecules in 0.25 moles of lead (II) nitrate.

7. Answer the following questions:

1. What mass in grams of hydrated sodium carbonate (Na₂CO₃·10H₂O) is in 65 cm³ of 0.2 M solution?
2. What volume of carbon dioxide would be evolved at STP when 6.2 g of copper (II) carbonate is reacted with sulphuric acid?

8. When excess lead (II) nitrate solution was added to a solution containing sodium chloride, a white precipitate was formed and weighed 5.56 g. Determine the amount of sodium chloride in the solution. (Pb = 207, Cl = 35.5, Na = 23)

9. 32.5 g of iron (III) chloride was reduced by hydrogen to iron metal. Calculate the volume of hydrogen used in the reaction at room temperature and pressure. The equation is: 2FeCl₃(s) + 3H₂(g) → 2Fe(s) + 6HCl(g). (Fe = 56, Cl = 35.5; molar gas volume = 24 litres at RTP.)

SECTION C: ADDITIONAL CALCULATIONS

The following questions retain the order and wording of the source, including its repeated question numbers.

7. Answer the following questions:

1. How many moles of calcium are present in 20 g of the metal?
2. Determine the grams of magnesium that contain 1.8 × 10²⁴ atoms.
3. How many ions are there in 4.6 g of Na⁺ ions? (Ca = 40, Mg = 24, Na = 23)

8. Find the mass of 3.0 g of calcium carbonate.

1. 150 g of calcium carbonate were reacted with 0.2 moles of hydrochloric acid.

   1. Write the equation of the reaction.
   2. Calculate the amount of calcium carbonate that will remain unreacted.
2. 0.7 g of gas X occupies 560 cm³ at STP. Determine its relative molecular mass. (Molar gas volume = 22.4 L.)
3. A certain carbonate ACO₃ reacts with dilute hydrochloric acid according to the equation given below: ACO₃(s) + 2HCl(aq) → Al₂(aq) + CO(g) + H₂O(l). If 1 g of the carbonate reacts completely with 20 cm³ of 1 M HCl, calculate the relative atomic mass of G. (C = 12.0, O = 16.0)
4. When a hydrocarbon was completely burnt in oxygen, 4.2 g of carbon dioxide and 1.71 g of water were formed. Determine the empirical formula of the hydrocarbon. (H = 1.0, C = 12.0, O = 16.0)

9. Suppose you have 3.92 g of hydrated ammonium iron (II) sulphate. How many moles of the compound do you have? (H = 1, N = 14, Fe = 56, S = 32, O = 16)

10. How many grams are there in 3 moles of water?

11. Determine the number of molecules in:

1. 1.3 dm³ of hydrogen at STP.
2. 1.2 dm³ of sulphur (IV) oxide at RTP. (L = 6.0 × 10²³; molar gas volume at STP = 22.4 dm³.)

CHEMISTRY FORM THREE TOPICAL EXAMINATION

ELECTROLYSIS

2½ HOURS

INSTRUCTIONS

1. This paper consists of Sections A, B and C.
2. Answer all questions.
3. Write answers clearly.
4. Use blue or black ink; use pencil for diagrams.

SECTION A: OBJECTIVE QUESTIONS

1. Choose the correct answer for each of the following items.

i. During electrolysis of molten aluminium oxide, three Faradays are needed to deposit one mole of aluminium. The number of electrons involved is:

1. 6.02 × 10²³
2. 1.806 × 10²³
3. 18.06 × 10²³
4. 180.6 × 10²³
5. 1806 × 10²³

ii. If a steady current of 2 amperes is passed through an aqueous solution of iron, the mass deposited at the cathode will be:

1. 54 g
2. 56 g
3. 0.54 g
4. 28 g
5. 0.52 g

iii. The use of electricity to decompose molten sodium chloride into its component elements is an example of:

1. Electroplating
2. Galvanization
3. Hydrolysis
4. Precipitation
5. Electrolysis

iv. What is the meaning of electrolysis?

1. Decomposition of a chemical compound by electricity
2. Breaking up by using electric power
3. Oxidation-reduction process
4. Passing electricity through a conductor

v. Which is the correct property of an electrolyte?

1. It contains free charge carriers
2. It does not have free ions
3. It cannot be decomposed by an electric current
4. It contains mainly undissociated molecules

vi. Which of the following is not an electrolyte?

1. Aqueous sulphuric acid
2. Solid common salt
3. Molten calcium chloride
4. Alcohol in water

vii. When dilute sodium chloride solution is electrolysed using carbon electrodes, which substance is collected at the cathode?

1. Hydrogen
2. Sodium
3. Chloride
4. Oxygen

viii. What are cations?

1. Cathode-going ions
2. Anode-moving ions
3. Negatively charged ions
4. Unionized particles

ix. Which particles will be anions in the electrolysis of aqueous copper sulphate using platinum electrodes?

1. SO₄²⁻ and OH⁻
2. Cu²⁺ and H⁺
3. Cu²⁺ and SO₄²⁻
4. Platinum

x. What is the name given to the quantity of a substance deposited by one coulomb of electricity on an electrode?

1. Electrochemical equivalent
2. Chemical equivalent
3. One mole of the substance
4. One gram molecular volume at S.T.P.

xi. Which of the following is an application of electrolysis?

1. Electroplating
2. Isolation of elements
3. Purification of metals
4. All of the above

xii. What is the name of the vessel in which electrolysis is conducted?

1. Voltameter
2. Voltmeter
3. Avometer
4. Beaker

2. Match the items in List A with the correct responses from List B.

SECTION B: STRUCTURED QUESTIONS

3. (a) Define (i) electrochemical equivalent and (ii) chemical equivalent.

(b) An electric current was passed in series through solutions of calcium chloride and copper(II) sulphate using carbon electrodes. If 2.5 litres of chlorine gas at S.T.P. were produced, calculate the volume of oxygen gas and the mass of copper produced.

4. (a) Explain (i) electrolysis and (ii) electroplating.

(b) State Faraday's laws of electrolysis.

(c) (i) Write the chemical equations for reactions at each electrode during electrolysis of dilute sulphuric acid using platinum electrodes.

(ii) 19,300 coulombs of electricity were passed through copper(II) nitrate solution. Calculate the mass of copper deposited at the cathode.

5. (a) With a well-labelled circuit diagram, explain the electrolysis of dilute sodium chloride solution using graphite electrodes.

(b) Does this electrolysis have a special name? Explain.

(c) Will further electrode reactions occur if electrolysis continues until the electrolyte is nearly dry? Explain.

(d) Calculate the mass of magnesium produced during electrolysis of molten magnesium chloride if a current of 1.93 A flows for 16 minutes 40 seconds.

6. (a) State Faraday's first law of electrolysis.

(b) Calculate the mass of silver and the volume of oxygen at S.T.P. liberated by 9,650 coulombs of electricity.

7. (a) State Faraday's laws of electrolysis.

(b) Element P has relative atomic mass 88. When a current of 0.5 A passes through its fused chloride for 32 minutes 10 seconds, 0.44 g of P is deposited at the cathode. 

(i) Calculate the number of Faradays needed to liberate one mole of P. 

(ii) Write the formula of the P ion. 

(iii) Write the formula of the hydroxide of P.

8. (a) Distinguish between chemical equivalent and electrochemical equivalent of a substance.

(b) Describe the electrolysis of copper(II) sulphate using graphite electrodes. Include a labelled diagram.

9. (a) State Faraday's laws of electrolysis.

(b) Element Z has relative atomic mass 88. A current of 0.5 A passes through its fused chloride for 32 minutes 10 seconds, depositing 0.44 g of Z at the cathode. (i) Calculate the number of Faradays needed to liberate one mole of Z. (ii) Write the formula of its ion. (iii) Write the formula of its hydroxide.

SECTION C: APPLICATION QUESTIONS

10. (a) (i) List three factors affecting the selection of ions discharged at electrodes. 

(ii) Define an electrolyte.

(b) A blue aqueous copper sulphate solution is electrolysed using copper electrodes. 

(i) Write ionic equations for the reactions at the cathode and anode. 

(ii) Explain what happens to the blue colour as electrolysis continues.

(c) (i) Define electroplating. 

An iron knife is electroplated with silver using a silver anode and an electrolyte containing Ag⁺, K⁺ and CN⁻ ions. Refer to the apparatus in the original examination. 

(ii) Which electrode is the cathode? 

(iii) Is the anode process oxidation or reduction? 

(iv) Write ionic equations for each electrode. 

(v) What happens to the electrolyte?

11. (a) 0.02 moles of electrons are passed through sodium hydroxide solution using platinum electrodes. 

(i) Name the gases evolved at each electrode. 

(ii) Write ionic equations for the electrode reactions. 

(iii) Calculate the moles and volumes at S.T.P. of each gas.

(b) What mass of copper is liberated during electrolysis of copper sulphate solution by one Faraday of charge?

(c) Element X has relative atomic mass 88. A current of 0.5 A is passed through its fused chloride for 32 minutes 10 seconds, depositing 0.44 g of X at the cathode. 

(i) Calculate the number of Faradays needed to liberate one mole of X. 

(ii) Write the formula of its ion. 

(iii) Write the formula of its hydroxide.

CHEMISTRY FORM THREE – TOPICAL EXAMINATION

CHAPTER 06: VOLUMETRIC ANALYSIS

PART I: MULTIPLE CHOICE QUESTIONS

1. Choose the correct answer for each of the following questions.

i. What mass of potassium chloride is needed to prepare 250cm3 of 0.235M solution?

1. 4.38g
2. 9.23g
3. 15.60g
4. 31.3g
5. 1.0g

ii. Which of the following laboratory equipments can accurately measure volume of a titre that has reacted completely with an a liquid?

1. Pipette
2. Burette
3. Measuring cylinder
4. Beaker

iii. What will be the molarity of a solution which contains 26.5g of anhydrous sodium carbonate in 5dm3of solution?

1. 0.05 M
2. 0.25 M
3. 5.30 M
4. 0.025 M
5. 0.50 M

iv. What volume of 0.2M strongSO4 is required to neutralize completely 25cm3 of 0.05M KOH?

1. 0.656cm3
2. 6.125cm3
3. 3.225cm
4. 3.125cm3

v. The mass of sodium hydroxide contained in 25cm3 of 0.1M is .

1. 0.5,
2. 2.85g
3. 25.0g
4. 0.1g
5. 25g

vi. A solution of sodium carbonate was prepared in order to get a 2M solution. 200cm3 of this solution was used in a titration experiment. The number of moles present in 200cm3 of 2M solution used in the titration will be;

1. 4.0
2. 0.04
3. 0.40
4. 0.045

vii. Which indicator is suitably used for the titration of weak acid and a strong base?

1. Methyl orange
2. Phenolphthalein
3. Litmus
4. Bromothymol blue

PART II: STRUCTURED QUESTIONS

2. Given that in the chemistry  national four examination, chemistry practical (2A), the following data were obtained by titrating-0.1M HCl with 5.3g/dm³ of X2 C03 using methyl orange as an indicator. Pipette used was 20cm3

Questions:

(i) Complete the table above.

(ii) Calculate the mean titre volume.

(iii) Write a balanced chemical equation between X2CO3 and HCl.

(iv) Calculate the molarity of X2CO3.

(v) Identify and name element X.

3. What do we mean by volumetric analysis?

4. State the use of each of the following apparatus used in the volumetric analysis:

(a) Volumetric flask

(b) Burette

(c) Pipette

(d) Measuring cylinder

(e) Weighing balance f. Dropper.

5. An experiment was carried out to find the value of X in Na2CO3.Xstrong . O.12.4g of Na2CO3.XstrongO was dissolved in water to make 0.5dm3 of solution. 25cm3 of this solution was transferred to a conical flask and few drops of methlyorange(MO) was added.The solution in a conical flask was then titrated against 0.25M HCl and 20cm3 of acid were required to reach the end point. i. write balanced chemical equation for the reaction ii. Calculate concentration of Na2CO3.XstrongO in mol/dm³. iii. Calculate the value of X.

6. 6cm3HCL required 22 cm3Of 0.2 M NaOH for complete neutralization. Calculate the concentration of the acid in moles per time.

7. Solution Q was made by dissolving 2.65g of a X2CO3in water and diluting it to 250cm3. 25 cm³ of P was titrated with 0.25 MHCL solution. The results obtained were tabulated as shown.

(a) Calculate the average volume of HCL acid used.

(b) Calculate the concentration of solution Q in: i.  Mol /dm³ ii. g/ dm³

8. Calculate the molarities of the following solutions:

(a) 10.og of sodium hydroxide, NaOH, in 250cm3 of solution

(b) 3.5g of iron (II) Sulphate FeSO4 in 100 Cm³

9. Calculate the molarities of the following solutions:

(a) 10.0 g sodium hydroxide, NaOH, in 250 Cm³ of solution

(b) 3.5g of iron (II) sulphate FeSO4,  in 100cm3 solution.

10. Calculate the number of moles in:

(a) 20cm3 of 0.1M Nitric (v) acid.

(b) 150cm3of 0.5 Sodium carbonate.

11. (a) Explain the meaning of the following. i. Titration ii. Molar solution

(b) Given that 5.6g of KOH were dissolved in 1dm3 .25cm3 of this solution ( KOH) were titrated with 25cm3 of hydrochloric acid. Calculate;

(i) Molarity of acid

(ii) Concentrationof KOH

(iii) Name  suitable indicator for this titration.

12. (a) Define i. End point of neutralization ii. Indicator iii. Standard solution iv. Secondary standard solution

(b) A solution of sulphuric acid is made by dissolving 5.2g of its mass containing some impurities to make 1 dm³ of solution. If 25cm3 of this solution required exactly 25cm3 of 0.1M NaOH for complete  neutralization during the titration process. Calculate the percentage purity of sulphuric acid.

13. Define the following terms:

(a) Standard solution

(b) End points

(c) Titration

14. 190 cm³ of the 0.4M lithium carbonate solution was neutralized by 16.0Cm3 of 1MstrongSO4. Write an equation for the reaction.

15. In an experiment, 22.00cm3 of 1.0 NaOH  Was neutralized by  170.0 cm³ of nitric (v) Acid. What is the concentration of nitric (v) acid in mol/ dm³.

16. Calculate the molarities of the following solutions:

(a) 25.0g of sodium hydroxide, NaOH in 200cm3 of solution.

(b) 3.5 g of iron (II) sulphate, FeSO4, in 250cm3 of solution.

17. Calculate the volume of

(a) 0.1M  Na,CO3 solution that neutralizes compeletely 2400cm3 of 0.2M strongSO4 Solution.

(b) 2M nitric (V) acid required to neutralize 25.00 cm³ of 0.5M NaOH.

18. Xg of Sodium hydroxide were dissolved in distilled water to make 100cm3  of solution . 50 cm³ of the solution required 50cm3. 2M nitric (V) acid for complete neutralization. Caliculate neutralization. Calculate the mass (x) of Sodium hydroxide dissolved ( Na ═ 23 O ═ 16 H ═ 1)

19. Calculate the amount of calcium carbonate that would remain in 17.0g of calcium carbonate were reacted with 0.25 Moles of hydrochloric acid. The equation for the equation for the reaction is CaCO3(g) + 2HCL(aq) → CaCl2(aq) + CO2 + strongO(I) (C ═ 12 O ═ 16 Ca ═ 40).

20. 0.88gcm-3. Calculate the morality of the ammonia solution.

21. How would you prepare 0.15 L of 0.5 M NaOH,  Starting with 6.0M Solution.

22. How would you prepare 0.75 litre of a 0.5M NaoH Solution, starting with solid sodium hydroxide and water.

23. Calculate the volume of 1M sulphuric (vi) acid need to completely neutralize 100cm3 of sodium hydroxide 0.2M in concentration.

24. The reaction below shows, reaction between calcium hydroxide and nitric (iv) acid. Ca(OH)2(S) + 2HNO3(aq) → Ca(NO3)2(aq) + 2strongO(l) If 20 Cm³ of0.2 M nitric(v) acid was used, calculate the volume of Ca(OH)2 used, given its morality is 0.1M

25. 3M  nitric acid reacted with 0.5g copper powder as shown below CU(S) + 2HNO3 (aq) → Cu (NO3) 2 + strongO Given (CU = 63.5) Calculate the volume of 3M nitric acid that reacted with copper.

26. Calculate the mass of nitrogen present in 25kg bag of ammonium phosphates (NH4)2 HPO4. (N = 14, H= 1, P = 31, O = 16)

27. When 1.08g of aluminium foil were   in a stream of  chlorine gas, the mass of product formed was 3.47g. calculate the i. Maximum mass of product formed of chlorine was excess (AL 27, CL = 35.5) ii. iii. Percentage yield of the product formed.

28. 15.0 cm³ of ethanoic acid (CH3COOH) was dissolved in water to make 500cm3 of solution. Calculate the concentration of the solution in moles per litre. (C= 12.0 H = 10, O = 16.0) (Density= 1.05g/cm³ of ethanoic acid)

29. The relative formula mass of a mass of hydrocarbon is 58. Draw and  name two possible structure of the hydrocarbon

30. An alcohol has this composition H= 13.5% O2 = 64.9 %

(a) Determine the structure of alcohol if empirical formula and molecular are the same.

(b) Draw the structure of alcohol if empirical formula and molecular formula are the same.

31. 6.84g of aluminum sulphate were dissolved in 150cm3 of water calcululate molar concentration of the sulphate ions in the solution. In of aluminium sulphate is 342).

32. Work out the simplest formula of a  compound. Which has the follows composition • Magnesium 9.8g • Sulphur13.0g • Oxygen26.0g • Water of crystallization 51.2g

33. 150g of Calcium carbonate were reacted with 0.2 moles of hydrochloric acid.

(a) Write equation of reaction

(b) Calculate amount of calcium carbonate that will remain unreached.

34. 0.7 g of gas × occupies 560 cm³ at s.t.p. determine its relative molecular mass. Molar gas vol = 22.4

35. A certain carbonate aCO3 reacts with dilute hydrochloric acid according to the equation given below. ACO3(s) + 2HCl(aq) → al2aq + CO (g) + strongO (l) If 1g of the carbonate reacts completely with 20cm3 of the IM HCL.Calculate the relative atomic mass of G. (C = 12.0, O = 16.0)

36. When a hydrocarbon was completely bunted in oxygen, 4.2g of carbon dioxide has1.71g of water were formed. Determine the empirical formula of late hydrocarbon. ( H= 1.0, C = 12.0 , O = 16.0 )

37. (a) What is titration?

(b) 25cm3 of impure sulphuric acid contains 5.2g/dm³ reacted with 23cm3 of sodium hydroxide solution made by dissolving 4.0g of NaOH in distilled water to make 1.0dm3 solution

(c) Calculate the percentage of

(i) Purity of the acid and(ii) its impurity

(d) Mention two areas of application of volumetric analysis

38. 20.0cm3 of a 1.4M, Dibasic acid strongX solution was titrated against 1.0M NaOH solutionuntil neutralization was complete.

(a) Write down the equation for the reaction.

(b) Calculate the volume of sodium hydroxide solution used.

CHEMISTRY EXAMINATION FORM THREE

TOPICAL EXAMINATION: COMPOUNDS OF METALS

NAME: ................................................ 

CLASS: ........................ 

TIME: 2½ HOURS

INSTRUCTIONS

1. This paper consists of Sections A, B and C.
2. Answer all questions.
3. Write all answers in the spaces provided.
4. Use blue or black ink for writing and pencil for drawings.

SECTION A (20 MARKS)

1. Choose the correct answer for each of the following items.

(i) Which of the following properties generally increases down the group?

1. Ionization energy
2. Atomic size
3. Electronegativity
4. Sodium and zinc

(ii) When does a chemist fail to identify a compound between sulphur and iron?

1. A black solid is formed
2. Heat is used to join them up
3. Yellow colour of sulphur and silvery shiny appearance
4. The resulting mass is greater than the individual mass of the elements
5. Any of the above A–D does not take place

(iii) Which of the following sets of elements is arranged in order of increasing electronegativity?

1. Chlorine, fluorine, nitrogen, oxygen, carbon
2. Fluorine, chlorine, oxygen, nitrogen, carbon
3. Carbon, nitrogen, oxygen, chlorine, fluorine
4. Nitrogen, oxygen, carbon, fluorine, chlorine
5. Fluorine, nitrogen, oxygen, chlorine, carbon

(iv) Which method could be used to separate the products in the following equation? Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq)

1. Chromatography
2. Crystallisation
3. Distillation
4. Filtration
5. Condensation

(v) The metal nitrate which will NOT give a metal oxide on heating is

1. Calcium nitrate
2. Silver nitrate
3. Lead nitrate
4. Copper nitrate
5. Zinc nitrate

(vi) Which of the following pairs of compounds can be used in preparation of calcium sulphate?

1. Calcium carbonate and sodium sulphate
2. Calcium chloride and ammonium sulphate
3. Calcium hydroxide and barium sulphate
4. Calcium nitrate and lead (II) sulphate
5. Calcium chloride and barium sulphate

(vii) The ionic equation when aqueous ammonium chloride reacts with sodium hydroxide solution is represented as:

1. NH₄⁺(aq) + OH⁻(aq) → NH₃(g) + H₂O(l)
2. Na⁺(aq) + Cl⁻(aq) → NaCl(s)
3. H⁺(aq) + OH⁻(aq) → H₂O(l)
4. 2NH₄⁺(aq) + 2Cl⁻(aq) → 2NH₃(g) + 2HCl(g)

(viii) Three elements X, Y and Z are in the same period. The oxide of X is amphoteric, of Y is basic and of Z is acidic. Which order shows increasing atomic number?

1. X, Y, Z
2. Y, Z, Y
3. Z, X, Y
4. Y, X, Z
5. X, Z, Y

(ix) The only metal which does not react with dilute hydrochloric acid is:

1. Magnesium
2. Aluminium
3. Copper
4. Zinc
5. Sodium

(x) Brown ring test is a technique used to test the presence of:

1. Sulphate
2. Nitrate
3. Chloride
4. Potassium salts

2. Match the items in List A with their corresponding responses in List B.

SECTION B

3. (a) Give reasons for the following:

1. A solution of chlorine in water is acidic.
2. Yellow phosphorus is stored under water.

(b) Give an account of the following:

1. Anhydrous copper (II) sulphate becomes coloured when exposed to air for a long time.
2. Carbon dioxide can be collected by the downward delivery method.
3. Concentrated sulphuric acid is not used for drying hydrogen sulphide gas.
4. Sodium metal is kept in paraffin oil.

4. (a) With the aid of chemical equations, explain what will happen when aluminium chloride reacts with water.

(b) A student accidentally broke a beaker containing copper (II) sulphate crystals. He decided to separate the blue crystals from small pieces of glass by first dissolving the mixture and then filtering. What were the next steps?

5. (a) Asubuhi Njema’s child was sick. At hospital she was prescribed medicine including a bottle of syrup labelled “Shake before use”. What does this signify?

(b) Answer the following:

1. What is the first step when identifying the contents of a salt containing one anion and one cation?
2. In a solution of salt and water, identify the solute and solvent. Justify your answer.

(c) Sodium is solid while chlorine is gas at room temperature although both are in the same period. What causes the difference?

6. (a) Describe the effect of:

1. Strongly heating a piece of marble in a Bunsen burner flame.
2. Moistening the residue from (i) with water.

(b) Explain:

1. Why is slaked lime added to soil in gardening?
2. Why is concentrated sulphuric acid used as a drying agent?

(c) Why is zinc used as a coating for iron and not vice versa?

(d) What will happen when:

1. Yellow flowers are introduced into chlorine gas?
2. A burning splint is introduced into hydrogen gas?
3. A glass rod dipped in concentrated hydrochloric acid is introduced into ammonia gas?
4. Sulphur dioxide is bubbled through yellow acidified potassium dichromate solution?

7. (a) Name the ore commonly used in extraction of copper metal.

(b) Write a balanced chemical equation for each extraction step:

1. Roasting concentrated ore (CuFeS₂) in air.
2. Heating roasted ore with silica in absence of air.
3. Burning copper sulphide in a regulated supply of air.
4. Purifying copper by electrolysis using copper sulphate solution, pure copper cathode and impure copper anode.

(c) Explain with chemical equations what happens to:

1. An iron earring dropped into copper (II) sulphate solution.
2. A copper knife dipped into zinc nitrate solution.
3. Copper turnings dropped into dilute hydrochloric acid.

8. (a) Explain the following observations using equations:

1. White anhydrous copper (II) sulphate turns blue when water is added.
2. A small piece of sodium reacts vigorously with water.
3. Adding zinc to copper (II) sulphate solution decolourises the solution and deposits a brown solid.

9. (a) Sodium, magnesium, zinc, copper and silver are listed in decreasing order of reactivity. Which metal is:

1. Likely to tarnish most rapidly in air?
2. Most likely to be found free in nature?
3. Least likely to react with steam?

(b) Two metals above are extracted by electrolysis of their molten chlorides. Name them and give one reason.

(c) Answer about an electrolytic cell:

1. Name the positive and negative electrodes.
2. To which electrode do sodium ions migrate in dilute aqueous sodium chloride?
3. What other ions migrate to that electrode?
4. Which ions are discharged at that electrode?

10.

(a) Classify the following reactions:

1. Action of heat on lead nitrate.
2. Addition of dilute hydrochloric acid to silver nitrate.
3. Action of heat on calcium carbonate in a closed vessel.
4. Action of chlorine gas on iron (II) chloride.

11. Explain how you can separate a mixture of iodine and copper (II) oxide.

12. List three products formed when copper (II) nitrate is heated strongly in a test tube.

13. Aluminium has atomic number 13. Answer the following:

1. Write the formula of aluminium chloride.
2. State and explain the observation when burning magnesium ribbon is plunged into carbon (IV) oxide.
3. State why water is not a suitable extinguisher of an electrical fire.
4. State the name given to the reaction in which ethane is converted to ethane.
5. Calculate the volume of 0.09 M acid that can be prepared from 1 dm³ of 1.2 M acid.
6. Explain how to distinguish a sulphate using barium chloride solution and dilute hydrochloric acid.
7. When 10.64 g of metal X reacts with excess oxygen, 15.2 g of oxide forms. Determine its formula (X = 56, O = 16).
8. Name the products at cathode and anode during electrolysis of concentrated sodium chloride using carbon electrodes.

14. (a) Some iodine became contaminated with black copper (II) oxide. Explain how to separate the mixture.

(b) Name the chemical techniques that can be used to separate sodium chloride when it is electrolysed using carbon electrodes.

15. Name a substance which decomposes on heating to give two gases other than water vapour.

16. Name the two gases referred to in Question 15.

17. A compound J reacts with concentrated sulphuric acid (H₂SO₄) to give brown fumes of compound K. A solution of K in water reacts with copper metal to give a dark brown gas L. Name J, K and L.

CHEMISTRY EXAMINATION FORM THREE

TOPICAL EXAMINATION: EXTRACTION OF METALS

NAME: ........................................ CLASS: ................ TIME: 2½ HOURS

INSTRUCTIONS

1. This paper consists of Sections A, B and C.
2. Answer all questions.
3. Write answers clearly.
4. Use blue or black ink, except for drawings, which should be in pencil.

SECTION A: OBJECTIVE QUESTIONS

1. Choose the correct answer for each item.

(i) In the reaction 2FeCl₂ + Cl₂ → 2FeCl₃, chlorine may be regarded as:

1. A reducing agent
2. An oxidizing agent
3. A catalyst
4. Halogens

(ii) The best method to extract iron from its ore is:

1. Reduction using carbon
2. Electrolysis of its solution
3. Roasting of iron ore in air
4. Electrolysis of its molten form

(iii) In the blast furnace, iron ore is converted into iron metal. Which statement is NOT true?

1. The reactions in the blast furnace produce heat which keeps iron molten
2. Limestone reacts with impurities to form slag
3. Coke reacts to reduce iron oxide to iron
4. Coke combines with some impurities in the ore

(iv) A less electropositive metal may be extracted by:

1. Ore concentration and electrolytic reduction
2. Ore concentration and chemical reduction
3. Ore concentration and purification
4. Ore concentration, reduction and purification

(v) Carbon is placed in the metal reactivity series because:

1. It has both metallic and non-metallic properties
2. It is very electropositive
3. It acts as a reducing agent like metals
4. It is metallic in nature

(vi) Which of the following is the most electronegative element?

1. Sodium
2. Sulphur
3. Fluorine
4. Chlorine

(vii) Adding calcium chloride to sodium chloride during sodium extraction helps to:

1. Increase the solubility of sodium
2. Lower the boiling point of sodium
3. Increase the boiling point of sodium chloride
4. Make the solution more soluble

(viii) Which is NOT a method of purifying an ore after extraction?

1. Distillation
2. Oxidation
3. Concentration
4. Electro-refining

(ix) The common ore used in extraction of aluminium is:

1. Bauxite
2. Cryolite
3. Feldspar
4. Kaolin

(x) The following are stages followed during extraction of metals EXCEPT:

1. Ore concentration
2. Ore purification
3. Refining
4. Extraction

2. Match each item in List A with the correct response in List B.

SECTION B: STRUCTURED QUESTIONS

3. (a) (i) Name four important ores of iron. (ii) What is achieved when concentrated iron ore is roasted?

(b) Briefly describe the balanced chemical reactions taking place in the blast furnace during iron extraction.

4. (a) List four common stages in the extraction of less reactive metals such as zinc and copper.

(b) Name the ore commonly used in the extraction of iron.

(c) Consider the following reactions in the blast furnace (equations reproduced as printed in the source; verify before use):

1. C(s) + O₂ → CO₃ + heat
2. CO₂ + C → 2CO
3. Fe₂O₃ + 3C → 2Fe + 3CO
4. Fe₂O₃ + CO → Fe + 2CO₂

1. Indicate two reducing agents in the blast furnace.
2. Explain the importance of steps (1) to (3).
3. Name compound L, added to produce a substance that removes impurities as slag.
4. Write the complete chemical reactions undergone by compound L to form slag.

5. Study the following flow chart showing aluminium extraction and answer the questions.

1. Name the ore labelled T.
2. State why the ore is dissolved in excess sodium hydroxide.
3. Write the formula of the aluminium compound in the solution.
4. Name the process in Step II.
5. Write the ionic equation for Step IV.
6. Explain why the anode is replaced from time to time.
7. Name one compound in the residue.
8. Explain why aluminium alloys are preferred to pure aluminium for overhead power cables.

6. Study the simplified blast furnace diagram and answer the questions.

(a) State the functions of (i) coke, (ii) limestone and (iii) hot air blast.

(b) (i) State one important use of the pipe marked in the original diagram. (ii) Give the uses of slag inside the furnace and as a product of the process.

7. Study the simplified Downs cell diagram used in extraction of sodium.

1. What is a Downs cell?
2. Name substances X and Y.
3. Why is calcium chloride added to molten sodium chloride?
4. Write electrode equations for (i) the anode and (ii) the cathode.
5. Explain why sodium is collected at the top.
6. Mention one environmental effect of sodium extraction.
7. Explain the role of the steel gauze diaphragm.
8. How is sodium stored in the laboratory?
9. Mention uses of sodium.

SECTION C: EXTENDED RESPONSE QUESTIONS

8. (a) What is a blast furnace?

(b) Name the common ore used in extraction of iron.

(c) Explain two roles of limestone in a blast furnace.

(d) What is steel? Name two uses of steel.

(e) Describe an environmental impact of iron extraction.

9. (a) Define the term ore.

(b) Name the chief ores of (i) copper, (ii) iron, (iii) aluminium and (iv) zinc.

(c) Give factors determining the method used for extracting a metal.

(d) Explain why it took so long for chemists to extract aluminium from its ore.