2019 WASSCE Chemistry Theory — Questions & Solutions

2019 WASSCE Chemistry Theory — Questions & Solutions

1.(a) Explain briefly why 

$\mathrm{^4_2He}$ has a stable electron configuration compared to $\mathrm{^9_4Be}$.

Answer:
$\mathrm{^4_2He}$ has a full 1s shell with two electrons, which makes it very stable (noble gas configuration). $\mathrm{^9_4Be}$ has two electrons in the 1s shell and two in the 2s shell, which is less stable due to higher energy and electron repulsion.

(b) Consider the elements $^1_1H$ and $^3_3Li$.

(i) State the number of electrons that an atom of each element would have after forming an ionic bond.

Answer:

• $H$ atom would have 2 electrons (gains one to form H⁻ ion).

• $Li$ atom would have 0 electrons in its outer shell (loses one electron to form Li⁺ ion).

(ii) Give a reason for each of your answers stated in (b)(i).

Answer:

• Hydrogen gains one electron to attain the stable helium configuration.
• Lithium loses one electron to achieve the stable configuration of helium.

(c) State two factors that should be considered when siting a chemical industry.

Answer:

• Availability of raw materials.
• Accessibility to transport facilities.

(d) State two advantages of using a catalyst instead of high temperatures in chemical reactions.

Answer:

• Catalysts lower the activation energy, allowing the reaction to proceed at lower temperatures.
• Catalysts increase reaction rate without being consumed.

(e) Turpentine burns in chlorine according to the following equation:

$$\mathrm{C_{10}H_{16} (l) + 8 Cl_2 (g) \rightarrow 10 C(s) + 16 HCl(g)}$$

Calculate the mass of turpentine that would completely burn in 21.3 g of chlorine.

$$\text{Given: } M_{Cl_2} = 71 \text{ g/mol}, \quad M_{\text{turpentine}} = 136 \text{ g/mol}$$

Answer:

• Moles of $\mathrm{Cl_2}$ = $\frac{21.3}{71} = 0.3$ mol

• From equation, 8 moles of $\mathrm{Cl_2}$ react with 1 mole of turpentine, so moles turpentine = $\frac{0.3}{8} = 0.0375$ mol

• Mass turpentine = moles × molar mass = $0.0375 \times 136 = 5.1$ g

(f) What is cracking?

Answer:
Cracking is the process of breaking down long-chain hydrocarbons into shorter, more useful hydrocarbons, usually by heat or catalyst.

(g) State two factors that may influence the value of electron affinity.

Answer:

• Atomic size (smaller atoms generally have higher electron affinity).
• Nuclear charge (higher charge attracts electrons more strongly).

(h) What are carbohydrates?

Answer:
Carbohydrates are organic compounds composed of carbon, hydrogen, and oxygen, typically with a hydrogen:oxygen ratio of 2:1.

(i) State two differences between a simple sugar and starch.

Answer:

• Simple sugars are monomers (e.g., glucose); starch is a polymer.
• Simple sugars are soluble in water; starch is insoluble.

(j) Write an equation to show the dissociation of each of the following acids:

(i) $\mathrm{H_2CO_3}$

$$\mathrm{H_2CO_3 \rightleftharpoons H^+ + HCO_3^-}$$

(ii) $\mathrm{CH_3COOH}$

$$\mathrm{C}{\mathrm{H}}_3\mathrm{C}\mathrm{O}\mathrm{O}\mathrm{H}\rightleftharpoons {\mathrm{H}}^{+}+\mathrm{C}{\mathrm{H}}_3\mathrm{C}\mathrm{O}{\mathrm{O}}^{-}$$

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2. (a) (i) Name three different methods for preparing salts.

Answer:

• Neutralization
• Precipitation (double displacement)
• Reaction of metal with acid

(ii) Give one example of a balanced equation for each of the methods named in (a)(i). .

• Neutralization:

$$\mathrm{H}\mathrm{C}\mathrm{l}+\mathrm{N}\mathrm{a}\mathrm{O}\mathrm{H}\to \mathrm{N}\mathrm{a}\mathrm{C}\mathrm{l}+{\mathrm{H}}_2\mathrm{O}$$

• Precipitation:

$$\mathrm{A}\mathrm{g}\mathrm{N}{\mathrm{O}}_3+\mathrm{N}\mathrm{a}\mathrm{C}\mathrm{l}\to \mathrm{A}\mathrm{g}\mathrm{C}\mathrm{l}(\mathrm{s})+\mathrm{N}\mathrm{a}\mathrm{N}{\mathrm{O}}_{\mathrm{3}}$$

• Metal-acid reaction:

$$\mathrm{M}\mathrm{g}+2\mathrm{H}\mathrm{C}\mathrm{l}\to \mathrm{M}\mathrm{g}\mathrm{C}{\mathrm{l}}_2+{\mathrm{H}}_{\mathrm{2}}$$

(iii) State two uses of sodium trioxocarbonate(IV).

Answer:

• Used in water softening
• Used in fire extinguishers

(b) Describe how to obtain pure copper from impure copper by electrolysis.
Answer:
Use electrolytic refining with impure copper as the anode and pure copper as the cathode. Copper dissolves at the anode and deposits on the cathode, impurities fall as anode sludge.

(c) Given solubilities of sodium chloride and silver nitrate at 30°C and 100°C:

Salt	Solubility at 30°C (g/100 g water)	Solubility at 100°C (g/100 g water)
Sodium chloride	36.3	39.0
Silver nitrate	297.0	952.0

(i) Calculate the percentage of each salt deposited on cooling from 100°C to 30°C.

Answer:
Sodium chloride:
Amount deposited = 39.0 - 36.3 = 2.7 g
Percentage = $\frac{2.7}{39.0} \times 100 = 6.92\%$

Silver nitrate:
Amount deposited = 952.0 - 297.0 = 655.0 g
Percentage = $\frac{655.0}{952.0}\times 100=$

(ii) Which salt is purified more efficiently by crystallization?
Answer: Silver nitrate, as it deposits a higher percentage upon cooling.

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3. (a) (i) State the important points of Dalton’s atomic theory.

Answer:

• Elements are made of tiny indivisible atoms.
• Atoms of the same element are identical.
• Compounds form from atoms of different elements in fixed ratios.

(ii) How does this theory explain the law of multiple proportions?
Answer:
Atoms combine in simple whole number ratios, explaining why elements form multiple compounds with different ratios.

(b) (i) Name three useful products from chemical transformation of vegetable oils.

Answer:

• Soap
• Biodiesel
• Margarine

(ii) With the aid of equations, describe the chemical reaction involved in the transformations stated in (b)(i).

Answer:

• Saponification: oils react with alkali to form soap.
• Hydrogenation: addition of hydrogen to oils to form margarine.
• Transesterification: conversion of oils to biodiesel.

(c) Explain why Beryllium’s properties differ from Magnesium’s but resemble Aluminium’s.
Answer:
Beryllium has a small atomic size and high charge density, causing covalent bonding like Aluminium, unlike the more metallic Magnesium.

(d) Write balanced equation for reaction between iodine and aqueous sodium thiosulphate.

$$\mathrm{I_2 + 2Na_2S_2O_3 \rightarrow 2NaI + Na_2S_4O_6}$$

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4. (a) (i) What is meant by hardness of water?

(ii) Describe briefly how you would determine what proportion of hardness in a given sample of water is due to permanent hardness.

(iii) Give two reasons why hardness of water is an undesirable property.

(b) State the:

(i) reagents;

(ii) conditions for the laboratory preparation of trioxonitrate (V) acid.

(iii) How does concentrated trioxonitrate (V) acid react with:

(I) sulphur;

(II) aluminium.

(c) Name one amphoteric oxide.

Question 4 Answers

4. (a) (i) What is meant by hardness of water?
Hardness of water is the presence of dissolved calcium (Ca²⁺) and magnesium (Mg²⁺) salts, such as calcium bicarbonate (Ca(HCO₃)₂) and magnesium bicarbonate (Mg(HCO₃)₂), which prevent soap from forming a lather easily.

(ii) Describe briefly how you would determine what proportion of hardness in a given sample of water is due to permanent hardness.

• Measure the total hardness by titrating the water sample with a soap solution or EDTA.
• Boil the water sample to precipitate and remove temporary hardness (mainly bicarbonates).
• Retitrate the boiled sample.
• The hardness remaining after boiling is permanent hardness, while the difference between total hardness and permanent hardness is temporary hardness.

(iii) Give two reasons why hardness of water is an undesirable property.

• It causes scale formation (like CaCO₃ deposits) in boilers and pipes, reducing efficiency and damaging equipment.
• It reduces soap efficiency, forming scum (insoluble salts), increasing soap consumption.

(b) (i) State the reagents for the laboratory preparation of trioxonitrate (V) acid (nitric acid).
Reagents:

• Concentrated sulfuric acid, H₂SO₄
• Potassium nitrate, KNO₃ (or sodium nitrate, NaNO₃)

(ii) State the conditions for the laboratory preparation of trioxonitrate (V) acid.

• The mixture of KNO₃ and concentrated H₂SO₄ is gently heated.
• This produces nitrogen pentoxide (N₂O₅) gas, which dissolves in the acid to form nitric acid (HNO₃).

(iii) How does concentrated trioxonitrate (V) acid react with:

(I) Sulfur (S)
Concentrated HNO₃ oxidizes sulfur to sulfuric acid:

$$\mathrm{S} + 6\mathrm{HNO}_3 \rightarrow \mathrm{H}_2\mathrm{SO}_4 + 6\mathrm{NO}_2 + 2\mathrm{H}_2\mathrm{O}$$

(II) Aluminium (Al)
Aluminium reacts with concentrated HNO₃ to form a thin oxide layer that protects the metal (passivation), so no further reaction occurs.

(c) Name one amphoteric oxide.

• Aluminium oxide, $\mathrm{Al}_2\mathrm{O}_3$

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5. (a) (i) Write an equation for the reaction by which sulphur dioxide in solution could be converted to tetraoxosulphate (VI) acid.

(ii) State one test to confirm the conversion of sulphur dioxide to tetraoxosulphate (VI) acid.

(iii) State the reaction of concentrated tetraoxosulphate (VI) acid with:

(I) oxalic acid;

(II) copper.

(iv) What property of concentrated tetraoxosulphate (VI) acid does each of the reactions stated in (a)(iii) illustrate?

(b) (i) Explain briefly why water is referred to as a universal solvent.

(ii) Give one chemical test for water.

(c) (i) What is the major component of synthetic gas?

(ii) Give one reason why synthetic gas is not a major source of air pollution.

(d) Name one product of destructive distillation of coal that is:

(i) solid;

(ii) liquid;

(iii) gas.

(e) (i) Write a balanced chemical equation for the complete combustion of carbon.

(ii) State one:

(I) physical;

(II) chemical property of the products in (e)(i).

Question 5 Answers

5. (a) (i) Write an equation for the reaction by which sulphur dioxide in solution could be converted to tetraoxosulphate (VI) acid.

Sulphur dioxide dissolves in water to form sulphurous acid, which can be oxidized to sulphuric acid (tetraoxosulphate (VI) acid):

$${\mathrm{S}\mathrm{O}}_2+{\mathrm{H}}_2\mathrm{O}\to {\mathrm{H}}_2{\mathrm{S}\mathrm{O}}_3\text{(sulphurous acid)}$$

Then oxidation:

$$2{\mathrm{H}}_2{\mathrm{S}\mathrm{O}}_3+{\mathrm{O}}_2\to 2{\mathrm{H}}_2{\mathrm{S}\mathrm{O}}_4\text{(tetraoxosulphate (VI) acid)}$$

(ii) State one test to confirm the conversion of sulphur dioxide to tetraoxosulphate (VI) acid.

• Add barium chloride (BaCl₂) solution; formation of a white precipitate of barium sulphate (BaSO₄) confirms sulphuric acid presence.

(iii) State the reaction of concentrated tetraoxosulphate (VI) acid with:

(I) Oxalic acid (H₂C₂O₄)

$${\mathrm{H}}_2{\mathrm{C}}_2{\mathrm{O}}_4+{\mathrm{H}}_2{\mathrm{S}\mathrm{O}}_4\to {\mathrm{C}\mathrm{O}}_2+{\mathrm{H}}_2\mathrm{O}+\text{other products}$$

(Reaction involves oxidation of oxalic acid by concentrated H₂SO₄ producing carbon dioxide gas.)

(II) Copper (Cu)

$$\mathrm{C}\mathrm{u}+2{\mathrm{H}}_2{\mathrm{S}\mathrm{O}}_4\to {\mathrm{C}\mathrm{u}\mathrm{S}\mathrm{O}}_4+{\mathrm{S}\mathrm{O}}_2+2{\mathrm{H}}_2\mathrm{O}$$

(Concentrated sulfuric acid acts as an oxidizing agent, releasing sulfur dioxide gas.)

(iv) What property of concentrated tetraoxosulphate (VI) acid does each of the reactions stated in (a)(iii) illustrate?

• The ability of concentrated H₂SO₄ to act as a strong oxidizing agent.

(b) (i) Explain briefly why water is referred to as a universal solvent.

• Water dissolves a wide range of substances due to its polarity and ability to form hydrogen bonds.

(ii) Give one chemical test for water.

• Add anhydrous copper(II) sulfate (CuSO₄). It changes from white to blue in the presence of water.

(c) (i) What is the major component of synthetic gas?

• Carbon monoxide (CO).

(ii) Give one reason why synthetic gas is not a major source of air pollution.

• It burns cleanly to produce carbon dioxide and water, reducing harmful emissions compared to fossil fuels.

(d) Name one product of destructive distillation of coal that is:

(i) Solid: Coke (carbon-rich solid)

(ii) Liquid: Coal tar

(iii) Gas: Coal gas (a mixture including hydrogen, methane, and carbon monoxide)

(e) (i) Write a balanced chemical equation for the complete combustion of carbon.

$$\mathrm{C} + \mathrm{O}_2 \rightarrow \mathrm{CO}_2$$

(ii) State one:

(I) Physical property of the product $\mathrm{CO}_2$:

• It is a colorless, odorless gas.

(II) Chemical property of $\mathrm{CO}_2$:

• It turns lime water milky (forms CaCO₃ precipitate).

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6. (a) State the conditions necessary for the cracking of long-chain hydrocarbons to produce more gasoline.

(b) State two reasons why metallic objects are electroplated.

(c) (i) Explain briefly why calcium oxide cannot be used to dry hydrogen chloride gas.

(ii) State one drying agent for hydrogen chloride gas.

(d) Concentrated trioxonitrate (V) acid was added to a solution of iron (II) tetraoxosulphate (VI) and the mixture heated. The mixture turned from pale green to yellow with the evolution of a brown gas. Explain briefly these observations.

(e) (i) Write the equation for the reaction between zinc oxide and

1. dilute tetraoxosulphate (VI) acid,

2. sodium hydroxide solution.

(ii) State which property of zinc oxide is shown by the reaction in (e)(i).

(f) Two isotopes of chlorine are 3517Cl and 3717Cl. State one:

1. similarity;

2. difference between the isotopes.

(g) State the two products formed when chlorine water is exposed to sunlight.

(h) Consider the reaction represented by the following equation:

(State the:

1. species that is undergoing oxidation;

2. oxidizing agent.)

(i) What is meant by carbon-12 scale?

(j) State two properties of a chemical system in equilibrium.

Question 6 Answers

6. (a) State the conditions necessary for the cracking of long-chain hydrocarbons to produce more gasoline.

• High temperature (about 600–700°C)
• Presence of a catalyst such as silica-alumina or alumina

(b) State two reasons why metallic objects are electroplated.

• To prevent corrosion (protect the metal surface)
• To improve appearance (give a shiny or decorative finish)

(c) (i) Explain briefly why calcium oxide cannot be used to dry hydrogen chloride gas.

• Calcium oxide reacts chemically with hydrogen chloride gas to form calcium chloride and water, so it cannot be used as a drying agent:

$$\mathrm{CaO} + 2\mathrm{HCl} \rightarrow \mathrm{CaCl}_2 + \mathrm{H}_2\mathrm{O}$$

(ii) State one drying agent for hydrogen chloride gas.

• Concentrated sulfuric acid (H₂SO₄)

(d) Concentrated trioxonitrate (V) acid was added to a solution of iron (II) tetraoxosulphate (VI) and the mixture heated. The mixture turned from pale green to yellow with the evolution of a brown gas. Explain briefly these observations.

• Iron (II) ions (Fe²⁺, pale green) are oxidized to iron (III) ions (Fe³⁺, yellow) by concentrated nitric acid.
• The brown gas evolved is nitrogen dioxide (NO₂), produced due to the oxidizing action of concentrated nitric acid.

(e) (i) Write the equation for the reaction between zinc oxide and:

• Dilute tetraoxosulphate (VI) acid (H₂SO₄):

$$\mathrm{Z}\mathrm{n}\mathrm{O}+{\mathrm{H}}_2{\mathrm{S}\mathrm{O}}_4\to {\mathrm{Z}\mathrm{n}\mathrm{S}\mathrm{O}}_4+{\mathrm{H}}_2\mathrm{O}$$

• Sodium hydroxide solution (NaOH):

$$\mathrm{Z}\mathrm{n}\mathrm{O}+2\mathrm{N}\mathrm{a}\mathrm{O}\mathrm{H}+{\mathrm{H}}_2\mathrm{O}\to {\mathrm{N}\mathrm{a}}_2[\mathrm{Z}\mathrm{n}(\mathrm{O}\mathrm{H}{)}_4]$$

(Zincate ion formed in solution)

(ii) State which property of zinc oxide is shown by the reactions in (e)(i).

• Zinc oxide is amphoteric, reacting both as an acid and a base.

(f) Two isotopes of chlorine are $\mathrm{^{35}_{17}Cl}$ and $\mathrm{^{37}_{17}Cl}$. State one:

• Similarity:
• Both isotopes have the same number of protons (17) and therefore are chlorine atoms.
• Difference:

• They have different numbers of neutrons (18 neutrons in

  $\mathrm{^{35}Cl}$, 20 neutrons in $\mathrm{^{37}Cl}$), resulting in different atomic masses.

(g) State the two products formed when chlorine water is exposed to sunlight.

• Hydrochloric acid (HCl)
• Hypochlorous acid (HClO)

(h) Consider the reaction represented by the following equation:
(Note: The exact reaction is not shown in the excerpt. I will give general definitions.)

1. Species that is undergoing oxidation: The species that loses electrons (increase in oxidation state).
2. Oxidizing agent: The species that gains electrons (causes oxidation).

(i) What is meant by carbon-12 scale?

• A scale of atomic masses where carbon-12 ($\mathrm{^{12}C}$) isotope is assigned exactly 12 atomic mass units (amu), and other elements are measured relative to it.

(j) State two properties of a chemical system in equilibrium.

• The forward and reverse reactions occur at the same rate.
• The concentrations of reactants and products remain constant.

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7. (a) A hydrocarbon having the formula 

$\mathrm{C}_{10}\mathrm{H}_{22}$ was cracked to produce $\mathrm{C}_6\mathrm{H}_{14}$ and another hydrocarbon P.

(i) Give the molecular formula of P.

• The molecular formula of P is $\mathrm{C}_4\mathrm{H}_8$.
  (Because $\mathrm{C}_{10}\mathrm{H}_{22} \rightarrow \mathrm{C}_6\mathrm{H}_{14} + \mathrm{C}_4\mathrm{H}_8$)

(ii) Draw the structures of two isomers of P.

• Two isomers of $\mathrm{C}_4\mathrm{H}_8$ (an alkene) are:

1. 1-Butene:

$$\mathrm{C}{\mathrm{H}}_2=\mathrm{C}\mathrm{H}-\mathrm{C}{\mathrm{H}}_2-\mathrm{C}{\mathrm{H}}_{\mathrm{3}}$$

2. 2-Butene:

$$\mathrm{CH_3-CH=CH-CH_3}$$

(iii) Give a reason why P could be polymerized.

• P contains a carbon–carbon double bond (C=C) which allows it to undergo addition polymerization.

(b) State the guiding principles which are used to explain the way electrons of the atoms of the elements are arranged in atomic orbitals.

• Aufbau principle: Electrons fill the lowest energy orbitals first.
• Pauli exclusion principle: Each orbital can hold a maximum of two electrons with opposite spins.
• Hund’s rule: Electrons occupy degenerate orbitals singly first, with parallel spins, before pairing.

(c) Consider each of the following substances: NaH, H₂, H₂S, NH₄Cl.

(i) Describe the nature of the intermolecular forces holding the units or molecules together in the condensed (liquid or solid) state.

• NaH: Ionic bonding between Na⁺ and H⁻ ions.

• H₂: Weak Van der Waals (dispersion) forces between nonpolar molecules.

• H₂S: Dipole-dipole interactions due to polarity of H–S bond.

• NH₄Cl: Ionic bonding between $\mathrm{NH}_4^+$ and Cl⁻ ions.

(ii) Explain briefly what happens when a sample of each of the substances is added to water.

• NaH: Reacts with water to form sodium hydroxide and hydrogen gas:

$$\mathrm{N}\mathrm{a}\mathrm{H}+{\mathrm{H}}_2\mathrm{O}\to \mathrm{N}\mathrm{a}\mathrm{O}\mathrm{H}+{\mathrm{H}}_{\mathrm{2}}$$

• H₂: Does not react; insoluble in water.

• H₂S: Dissolves slightly; weak acid, partly ionizes:

$$\mathrm{H}_2\mathrm{S} \rightleftharpoons \mathrm{H}^+ + \mathrm{HS}^-$$

• NH₄Cl: Dissociates into ions:

$${\mathrm{N}\mathrm{H}}_4\mathrm{C}\mathrm{l}\to {\mathrm{N}\mathrm{H}}_4^{+}+{\mathrm{C}\mathrm{l}}^{-}$$

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(iii) Write the chemical equations of any reactions occurring or of any equilibria established.

• For NaH (reaction with water):

$$\mathrm{N}\mathrm{a}\mathrm{H}+{\mathrm{H}}_2\mathrm{O}\to \mathrm{N}\mathrm{a}\mathrm{O}\mathrm{H}+{\mathrm{H}}_{\mathrm{2}}$$

• For H₂S (ionization equilibrium):

$${\mathrm{H}}_2\mathrm{S}\rightleftharpoons {\mathrm{H}}^{+}+{\mathrm{H}\mathrm{S}}^{-}$$

• For NH₄Cl (dissociation):

$${\mathrm{N}\mathrm{H}}_4\mathrm{C}\mathrm{l}\to {\mathrm{N}\mathrm{H}}_4^{+}+{\mathrm{C}\mathrm{l}}^{-}$$

(d) Element J has the following electron configuration: $1s^2 2s^2 2p^6 3s^2$.

(i) How many unpaired electrons can be found in J?

• There are no unpaired electrons; all orbitals are fully filled.

(ii) State whether J would be a good oxidizing or reducing agent.

• J would be a reducing agent.

(iii) Give a reason for the answer in (d)(ii).

• J has electrons to donate (especially the 3s electrons), so it tends to lose electrons and reduce other substances.

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8. (a) In the Solvay process, explain briefly with equations the functions of the following substances:

• limestone;
• ammonia.

(b) (i) Write a chemical equation for the fermentation of glucose.

(ii) Explain briefly why a tightly-corked glass bottle filled to the brim with fresh palm-wine shatters on standing for some time.

(c) Consider the following metals: Na, Fe, K and Cu.

(i) Arrange the metals in order of increasing reactivity.

(ii) Which of the metals will react with cold water?

(iii) Which of the metals could form coloured salts?

(d) (i) What is a redox reaction?

(ii) Identify which of the following reaction equations are redox.

(I) 2Na + Cl2 → 2NaCl

(II) AgCl + 2NH3 → [Ag(NH3)2]Cl

(III) C2H2 + H2 → C2H4

(IV) HCl + KOH → KCl + H2O

(V) 2FeCl3 + 2KI → 2FeCl2 + 2KCl + I2

(iii) Give a reason for each of the answers in (d)(ii).

(iv) Write balanced equations of the half reactions for any two of the redox reactions in (d)(ii).

Question 8 Answers

8. (a) In the Solvay process, explain briefly with equations the functions of the following substances:

• Limestone (CaCO₃):
  Decomposes on heating to produce calcium oxide (CaO) and carbon dioxide (CO₂), which is used in the process.

$$\mathrm{CaCO}_3 \xrightarrow{\Delta} \mathrm{CaO} + \mathrm{CO}_2$$

• Ammonia (NH₃):
  React with carbon dioxide and water to form ammonium bicarbonate, which reacts with sodium chloride to produce sodium bicarbonate (precipitate).

$${\mathrm{N}\mathrm{H}}_3+{\mathrm{C}\mathrm{O}}_2+{\mathrm{H}}_2\mathrm{O}\to {\mathrm{N}\mathrm{H}}_4{\mathrm{H}\mathrm{C}\mathrm{O}}_{\mathrm{3}}$$

(b) (i) Write a chemical equation for the fermentation of glucose.

$$\mathrm{C}_6\mathrm{H}_{12}\mathrm{O}_6 \rightarrow 2\mathrm{C}_2\mathrm{H}_5\mathrm{OH} + 2\mathrm{CO}_2$$

(ii) Explain briefly why a tightly-corked glass bottle filled to the brim with fresh palm-wine shatters on standing for some time.

• Carbon dioxide gas produced during fermentation builds up pressure inside the sealed bottle, causing it to shatter.

(c) Consider the following metals: Na, Fe, K and Cu.

(i) Arrange the metals in order of increasing reactivity.

$$\mathrm{C}\mathrm{u}<\mathrm{F}\mathrm{e}<\mathrm{N}\mathrm{a}<\mathrm{K}$$

(ii) Which of the metals will react with cold water?

• Potassium (K) and Sodium (Na) react vigorously with cold water.

(iii) Which of the metals could form coloured salts?

• Iron (Fe) and Copper (Cu) form coloured salts.

(d) (i) What is a redox reaction?

• A reaction in which oxidation and reduction occur simultaneously, involving the transfer of electrons.

(ii) Identify which of the following reaction equations are redox.

(I) $2\mathrm{Na} + \mathrm{Cl}_2 \rightarrow 2\mathrm{NaCl}$ — Redox

(II) $\mathrm{AgCl} + 2\mathrm{NH}_3 \rightarrow [\mathrm{Ag}(\mathrm{NH}_3)_2]\mathrm{Cl}$ — Not redox

(III) $\mathrm{C}_2\mathrm{H}_2 + \mathrm{H}_2 \rightarrow \mathrm{C}_2\mathrm{H}_4$ — Redox

(IV) $\mathrm{HCl} + \mathrm{KOH} \rightarrow \mathrm{KCl} + \mathrm{H}_2\mathrm{O}$ — Not redox

(V) $2\mathrm{FeCl}_3 + 2\mathrm{KI} \rightarrow 2\mathrm{FeCl}_2 + 2\mathrm{KCl} + \mathrm{I}_2$ — Redox

(iii) Give a reason for each of the answers in (d)(ii).

• (I) Sodium is oxidized and chlorine is reduced.
• (II) No change in oxidation states; complex formation only.
• (III) Addition of hydrogen; hydrogen is oxidized, acetylene reduced.
• (IV) Neutralization reaction; no change in oxidation states.
• (V) Iron(III) is reduced to Iron(II) and iodide ion is oxidized to iodine.

(iv) Write balanced equations of the half reactions for any two of the redox reactions in (d)(ii).

• For (I):
  Oxidation:

$$2\mathrm{N}\mathrm{a}\to 2{\mathrm{N}\mathrm{a}}^{+}+2e^{-}$$

Reduction:

$${\mathrm{C}\mathrm{l}}_2+2e^{-}\to 2{\mathrm{C}\mathrm{l}}^{-}$$

• For (V):
  Oxidation:

$$2{\mathrm{I}}^{-}\to {\mathrm{I}}_2+2e^{-}$$

Reduction:

$$2{\mathrm{F}\mathrm{e}}^{3+}+2e^{-}\to 2{\mathrm{F}\mathrm{e}}^{2+}$$

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9. (a) The following reaction scheme is an illustration of the contact process. Study the scheme and answer the questions that follow.

(i) Name X and Y.

(ii) Write a balanced chemical equation for each of the processes I, II, III and IV.

(iii) Name the catalyst used in process II.

(iv) Using Le Chatelier's principle, explain briefly why increasing the temperature would not favour the reaction in II.

(v) State two uses of SO2.

(b) Consider the following equation: 2H2(g) + O2(g) → 2H2O(g)

Calculate the volume of unused oxygen gas when 40 cm3 of hydrogen gas is sparked with 30 cm3 of oxygen gas.

(c) Calcium carbonate of mass 1.0 g was heated until there was no further change.

• Write an equation for the reaction which took place.
• Calculate the mass of the residue.
• Calculate the volume of the gas evolved at s.t.p.
• What would be the volume of the gas measured at 15 °C and 760 mm Hg?

[C=12.0, O=16.0, Ca=40.0, molar volume of a gas at s.t.p. = 22.4 dm3]

Question 9 Answers

9. (a) The following reaction scheme is an illustration of the contact process. Study the scheme and answer the questions that follow.

(i) Name X and Y.

• X = Sulfur dioxide, $\mathrm{SO}_2$

• Y = Sulfur trioxide, $\mathrm{SO}_3$

(ii) Write a balanced chemical equation for each of the processes I, II, III and IV.

• Process I: Combustion of sulfur

$$\mathrm{S}+{\mathrm{O}}_2\to {\mathrm{S}\mathrm{O}}_{\mathrm{2}}$$

• Process II: Oxidation of sulfur dioxide to sulfur trioxide

$$2\mathrm{SO}_2 + \mathrm{O}_2 \xrightarrow{\text{V}_2\mathrm{O}_5} 2\mathrm{SO}_3$$

• Process III: Absorption of sulfur trioxide in concentrated sulfuric acid

$${\mathrm{S}\mathrm{O}}_3+{\mathrm{H}}_2{\mathrm{S}\mathrm{O}}_4\to {\mathrm{H}}_2{\mathrm{S}}_2{\mathrm{O}}_7\text{(oleum)}$$

• Process IV: Dilution of oleum to produce sulfuric acid

$${\mathrm{H}}_2{\mathrm{S}}_2{\mathrm{O}}_7+{\mathrm{H}}_2\mathrm{O}\to 2{\mathrm{H}}_2{\mathrm{S}\mathrm{O}}_{\mathrm{4}}$$

(iii) Name the catalyst used in process II.

• Vanadium(V) oxide, $\mathrm{V}_2\mathrm{O}_5$

(iv) Using Le Chatelier's principle, explain briefly why increasing the temperature would not favour the reaction in II.

• The oxidation of $\mathrm{SO}_2$ to $\mathrm{SO}_3$ is an exothermic reaction. Increasing temperature shifts the equilibrium to the left (reactants) to counteract the change, decreasing $\mathrm{SO}_3$ formation.

(v) State two uses of $\mathrm{SO}_2$.

• As a preservative in food and drinks
• In the manufacture of sulfuric acid

(b) Consider the following equation:

$$2{\mathrm{H}}_2(g)+{\mathrm{O}}_2(g)\to 2{\mathrm{H}}_2\mathrm{O}(g)$$

Calculate the volume of unused oxygen gas when 40 cm³ of hydrogen gas is sparked with 30 cm³ of oxygen gas.

• From the balanced equation, 2 volumes of $\mathrm{H}_2$ react with 1 volume of $\mathrm{O}_2$.

• Volume of $\mathrm{O}_2$ required for 40 cm³ $\mathrm{H}_2$ = $40÷2=$ cm³.

• Initial $\mathrm{O}_2$ = 30 cm³.

• Unused $\mathrm{O}_2 = 30 - 20 = 10$ cm³.

(c) Calcium carbonate of mass 1.0 g was heated until there was no further change.

• Write an equation for the reaction which took place:

$$\mathrm{CaCO}_3(s) \xrightarrow{\Delta} \mathrm{CaO}(s) + \mathrm{CO}_2(g)$$

• Calculate the mass of the residue:
  Molar masses:
  $\mathrm{CaCO}_3 = 100$ g/mol
  $\mathrm{CaO} = 56$ g/mol
  Mass of $\mathrm{CaO} = \frac{56}{100} \times 1.0 = 0.56 \text{ g}$

• Calculate the volume of the gas evolved at s.t.p.:
  Moles of $\mathrm{CaCO}_3 = \frac{1.0}{100} = 0.01$ mol
  Moles of $\mathrm{CO}_2 = 0.01$ mol (1:1 ratio)
  Volume at s.t.p. = $0.01 \times 22.4 = 0.224$ dm³ = 224 cm³

• What would be the volume of the gas measured at 15 °C and 760 mm Hg?
  Use the combined gas law:

$$\frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2}$$

Where:

• $P_1 = 760$ mm Hg, $V_1 = 224$ cm³, $T_1 = 273 \text{ K}$
  

• $P_2=\mathrm{760\; mm\; Hg, }$$T_2 = 15 + 273 = 288 \text{ K}$
  

• $V_2 = ?$
  

Rearranged:

$$V_2=V_1\times \frac{T_2}{T_1}=224\times \frac{288}{273}\approx 236.3{\text{ cm}}^3$$

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10. (a) (i) Draw and label a diagram to illustrate the preparation and collection of dry chlorine gas in the laboratory.

(ii) State two uses of chlorine.

(b) Describe the preparation of hydrogen from water gas.

(i) Name the chief ore of aluminium.

(ii) Why is the ore purified?

(iii) Name the electrode used in the electrolysis.

(iv) Give one reason why cryolite, NaAlF6, is added to the electrolyte.

(c) Name three products obtained directly from the destructive distillation of coal.

Question 10 Answers

10. (a) (i) Draw and label a diagram to illustrate the preparation and collection of dry chlorine gas in the laboratory.

• Use manganese(IV) oxide (MnO₂) as catalyst, with concentrated hydrochloric acid (HCl) added slowly to MnO₂ in a conical flask.
• Chlorine gas (Cl₂) is evolved and collected by downward displacement of air (because chlorine is denser than air).
• Include a drying tube containing concentrated sulfuric acid (H₂SO₄) to dry the chlorine gas before collection.

(a) (ii) State two uses of chlorine.

• Used for water purification (disinfectant) germicide/sterilization/antiseptic/insecticides 
• Used in the manufacture of hydrochloric acid and bleaching agents.
• Production of polymers/PVC/plastics 
• Production of chlorinated chemicals e.g. CHCl3, CCl4 

(b) Describe the preparation of hydrogen from water gas.

• Water gas (a mixture of CO and H₂) is passed over heated zinc oxide (ZnO).
• Zinc oxide reacts with carbon monoxide to produce zinc vapor and carbon dioxide.
• Hydrogen is separated by fractional distillation or absorption.

(i) Name the chief ore of aluminium.

• Bauxite ($\mathrm{Al}_2\mathrm{O}_3 \cdot 2\mathrm{H}_2\mathrm{O}$)

(ii) Why is the ore purified?

• To remove impurities such as iron oxides and silica which would interfere with electrolysis.

(iii) Name the electrode used in the electrolysis.

• Carbon (graphite) electrodes are used as the anode and cathode.

(iv) Give one reason why cryolite, Na₃AlF₆, is added to the electrolyte.

• To lower the melting point of alumina, reducing energy consumption during electrolysis.

(c) Name three products obtained directly from the destructive distillation of coal.

• Solid: Coke
• Liquid: Coal tar
• Gas: Coal gas (a mixture including hydrogen, methane, carbon monoxide)