Questions – Chapter 2: Atomic Structure
🔹 Section A: Multiple Choice Questions (MCQs)
1. What is the relative charge of a neutron?
A) +1
B) -1
C) 0
D) +2
2. Which subatomic particle is found in the nucleus and has a positive charge?
A) Neutron
B) Electron
C) Proton
D) Positron
3. Which element has the electronic configuration 2,8,1?
A) Oxygen
B) Magnesium
C) Sodium
D) Potassium
4. What is the mass number of an atom with 11 protons, 12 neutrons, and 11
electrons?
A) 22
B) 11
C) 23
D) 12
🔹 Section B: Short Answer Questions
5. Define the following terms:
a) Atomic number
b) Mass number
c) Isotope
6. State the three subatomic particles of an atom and describe their relative mass and
charge.
7. How is the periodic table arranged in terms of atomic structure?
8. What information does the atomic number of an element provide?
9. Explain the difference between the mass number and the relative atomic mass (Ar).
10. Why do atoms have no overall electrical charge?
🔹 Section C: Electronic Structure
11. What is the electronic configuration of the following elements?
a) Oxygen (O)
b) Sodium (Na)
c) Magnesium (Mg)
d) Chlorine (Cl)
12. Draw the electronic configuration (shell diagram) for:
a) Carbon (Atomic number = 6)
b) Neon (Atomic number = 10)
13. Which group of the periodic table contains elements with a full outer shell of
electrons? What is the significance of this?
14. An element has an atomic number of 17 and a mass number of 35.
a) How many protons, neutrons, and electrons does it have?
b) Write its electronic configuration.
c) Identify the element.
🔹 Section D: Isotopes and Relative Atomic Mass
15. What are isotopes? Give two examples of isotopes of hydrogen and their uses.
16. The relative atomic mass of chlorine is 35.5. Explain why it is not a whole number.
17. Chlorine has two isotopes: Cl-35 (75%) and Cl-37 (25%).
a) Calculate the relative atomic mass (Ar) of chlorine.
18. Carbon exists mainly as two isotopes: carbon-12 and carbon-13.
a) Describe a use of carbon isotopes in science.
b) How does the existence of isotopes affect the mass number shown on the
periodic table?
🔹 Section E: Structure of the Atom
19. Describe the structure of an atom in terms of a central nucleus and surrounding
electrons.
20. How did Rutherford’s experiment lead to the discovery of the nuclear model of the
atom?
21. Compare the Bohr model of the atom with the plum pudding model.
🔹 Section F: Challenge / Exam-Style Questions
22. A sample of lithium contains two isotopes:

⁶Li (mass = 6, abundance = 7.5%)

⁷Li (mass = 7, abundance = 92.5%)
a) Calculate the relative atomic mass of lithium.
b) How many protons, neutrons, and electrons are in each isotope?
23. An unknown element X has the following properties:

It has 18 electrons in a neutral atom

Its mass number is 40
a) Identify element X
b) Calculate the number of protons, neutrons, and electrons
c) Write its electronic configuration
d) Predict its chemical properties based on its group
Questions – Chapter 3: Chemical Bonding
🔹 Section A: Multiple Choice Questions (MCQs)
1. What type of bonding occurs between metal and non-metal atoms?
A) Ionic bonding
B) Covalent bonding
C) Metallic bonding
D) Hydrogen bonding
2. Which of the following compounds contains only covalent bonds?
A) NaCl
B) MgO
C) H₂O
D) CaCl₂
3. What is the correct formula of the compound formed between calcium and
fluorine?
A) CaF
B) Ca₂F
C) CaF₂
D) Ca₂F₂
4. Which substance has a giant covalent structure?
A) Water
B) Diamond
C) Sodium
D) Carbon dioxide
🔹 Section B: Short Answer Questions
5. Define the following:
a) Ionic bond
b) Covalent bond
c) Metallic bond
6. State the type of bonding in the following substances:
a) Sodium chloride
b) Oxygen gas
c) Copper
d) Silicon dioxide
7. Explain how an ionic bond is formed between sodium and chlorine.
8. Why do ionic compounds have high melting and boiling points?
9. Describe how a single covalent bond forms between two hydrogen atoms.
10. What are lone pairs and bond pairs of electrons?
🔹 Section C: Drawing Diagrams
11. Draw dot-and-cross diagrams to show the bonding in:
a) Magnesium chloride (MgCl₂)
b) Oxygen molecule (O₂)
c) Ammonia (NH₃)
d) Water (H₂O)
12. Show the formation of sodium oxide using electron transfer diagrams.
13. Use a dot-and-cross diagram to show the bonding in a methane (CH₄) molecule.
🔹 Section D: Properties of Substances
14. Compare the physical properties of ionic compounds and covalent molecular
substances under the following headings:
a) Melting point
b) Electrical conductivity
c) Solubility in water
15. Explain why diamond is hard, but graphite is soft and slippery, even though both
are forms of carbon.
16. Why do metals conduct electricity?
17. Explain how metallic bonding gives metals their properties:
a) Malleability
b) Electrical conductivity
🔹 Section E: Structures and Types of Bonding
18. Describe the structure and bonding in:
a) Diamond
b) Graphite
c) Silicon dioxide (SiO₂)
19. What is meant by a giant ionic lattice? Give an example.
20. Why do substances with simple molecular structures (like CO₂ and H₂O) have low
melting and boiling points?
🔹 Section F: Challenge / Exam-Style Questions
21. Explain, with diagrams, the differences between:
a) A giant covalent structure (e.g., diamond)
b) A simple molecular structure (e.g., carbon dioxide)
22. Magnesium reacts with nitrogen to form magnesium nitride (Mg₃N₂).
a) Write the electron configurations of magnesium and nitrogen.
b) Describe how electrons are transferred to form Mg₃N₂.
c) Draw a dot-and-cross diagram for the compound.
23. Aluminum and oxygen form aluminum oxide (Al₂O₃).
a) Describe the ionic bonding in aluminum oxide.
b) Explain why aluminum oxide has a high melting point.
c) Predict whether aluminum oxide conducts electricity when molten and explain
your answer.
Questions – Chapter 4: Chemical Formulae and Equations
🔹 Section A: Multiple Choice Questions (MCQs)
1. What is the correct chemical formula for aluminum sulfate?
A) Al₂SO₄
B) AlSO₄
C) Al₂(SO₄)₃
D) Al₃(SO₄)₂
2. Which of the following compounds is correctly named?
A) KCl – potassium chlorate
B) MgO – magnesium oxide
C) Na₂O – sodium dioxide
D) Fe₂O₃ – iron(I) oxide
3. What is the formula of a compound formed between calcium (Ca²⁺) and nitrate
(NO₃⁻)?
A) CaNO₃
B) Ca(NO₃)₂
C) Ca₂NO₃
D) Ca₂(NO₃)₃
🔹 Section B: Short Answer Questions
4. Define the following terms:
a) Empirical formula
b) Molecular formula
c) Relative atomic mass (Ar)
5. Write the balanced symbol equation for the reaction of hydrochloric acid with
sodium hydroxide.
6. Explain why chemical equations must be balanced.
7. Identify the type of compound and write the correct name:
a) NH₄Cl
b) CaCO₃
c) SO₂
8. State the physical states (using state symbols) of the reactants and products in this
reaction:
HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(…)
🔹 Section C: Writing Formulae and Naming Compounds
9. Write the chemical formulae for the following compounds:
a) Magnesium chloride
b) Iron(III) oxide
c) Ammonium nitrate
d) Potassium permanganate
10. Write the names of the following compounds:
a) Na₂SO₄
b) Ca(OH)₂
c) FeCl₃
d) HNO₃
🔹 Section D: Balancing Equations
11. Balance the following chemical equations:
a) H₂ + O₂ → H₂O
b) C₃H₈ + O₂ → CO₂ + H₂O
c) Fe + HCl → FeCl₂ + H₂
12. Balance this equation and add state symbols:
Na + H₂O → NaOH + H₂
🔹 Section E: Relative Mass Calculations
13. Calculate the relative molecular mass (Mr) of:
a) H₂SO₄
b) Ca(OH)₂
c) C₂H₆O (ethanol)
14. A compound contains 40% carbon, 6.7% hydrogen, and 53.3% oxygen by mass.
a) Determine its empirical formula.
b) If its molecular mass is 180 g/mol, what is its molecular formula?
🔹 Section F: Water of Crystallization
15. A 4.98 g sample of hydrated magnesium sulfate (MgSO₄·xH₂O) is heated to constant
mass, leaving 2.44 g of anhydrous MgSO₄.
a) Calculate the mass of water lost.
b) Calculate the number of moles of MgSO₄ and H₂O.
c) Determine the value of x in the formula.
🔹 Section G: Challenge / Exam-Style Questions
16. A compound contains aluminum and chlorine only. 6.7 g of aluminum reacts with
chlorine to form 33.3 g of a compound. Determine its empirical formula.
17. Lead(II) nitrate reacts with potassium iodide to form lead(II) iodide and potassium
nitrate:
a) Write a balanced symbol equation.
b) Name the type of reaction.
c) What observation would you expect?
📘 Chapter 5: Chemical Calculations – Question Set
🔹 Section A: Basic Mole Calculations
1. Define the term "mole".
2. How many moles are there in:
o
a) 12 g of carbon (C)?
o
b) 36 g of water (H₂O)?
3. Calculate the number of molecules in 2 moles of oxygen gas (O₂).
4. Find the mass of:
o
a) 0.5 moles of NaCl
o
b) 3 moles of CO₂
🔹 Section B: Calculations Using Equations
5. Write a balanced chemical equation for the reaction between hydrogen and oxygen
to form water.
6. Given: 2Mg + O₂ → 2MgO, how many grams of magnesium are required to form 10 g
of MgO?
7. If 5 g of calcium reacts with water, what mass of hydrogen gas is produced?
🔹 Section C: Gases and Molar Volume
8. Calculate the volume occupied by:
o
a) 2 mol of hydrogen gas at room temperature and pressure (r.t.p.)
o
b) 0.25 mol of CO₂ gas at r.t.p.
9. If 6.0 dm³ of hydrogen reacts with chlorine, what volume of HCl is formed?
🔹 Section D: Solutions and Concentration
10. Calculate the concentration (in mol/dm³) of a solution containing:

a) 0.2 mol NaOH in 500 cm³ of solution

b) 0.5 mol H₂SO₄ in 250 cm³
11. A student dissolves 10 g of sodium carbonate (Na₂CO₃) in 200 cm³ of water. Calculate
the concentration of the solution in mol/dm³.
🔹 Section E: Empirical and Molecular Formulae
12. A compound contains 52.2% carbon, 13.0% hydrogen, and 34.8% oxygen. Find its
empirical formula.
13. If the empirical formula of a compound is CH₂ and its relative molecular mass is 56,
find its molecular formula.
14. A hydrocarbon contains 85.7% carbon and 14.3% hydrogen. Determine its empirical
and molecular formula if the relative molecular mass is 56.
🔹 Section F: Percentage Yield and Purity
15. Define:

a) Percentage yield

b) Percentage purity
16. In an experiment, 4.8 g of copper was obtained from 6.0 g of copper(II) oxide.
Calculate the percentage yield.
17. A student obtains 0.8 g of pure salt from 1.0 g of an impure sample. What is the
percentage purity?
🔹 Section G: Water of Crystallization
18. Define “water of crystallization”.
19. A hydrated salt has the formula CuSO₄·xH₂O. A 6.25 g sample is heated and leaves
behind 4.00 g of anhydrous CuSO₄. Find the value of x.
20. A 10.0 g sample of hydrated magnesium sulfate (MgSO₄·xH₂O) loses 4.0 g on heating.
Find the value of x.
🔹 Section H: Limiting Reactants
21. What is meant by the term limiting reactant?
22. In the reaction: 2H₂ + O₂ → 2H₂O, if 3 mol of H₂ react with 2 mol of O₂, which is the
limiting reactant?
23. If 2.0 g of zinc reacts with 10 cm³ of hydrochloric acid (2 mol/dm³), calculate the
limiting reactant and the mass of ZnCl₂ formed.
📝 Bonus: Mixed Concept Problems
24. A student reacts 0.1 mol of calcium carbonate with excess hydrochloric acid.
Calculate:

a) The mass of calcium chloride formed

b) The volume of carbon dioxide evolved at r.t.p.
25. Determine the empirical formula of a compound formed when 0.48 g of magnesium
reacts with 1.6 g of oxygen.
Chapter 6: Electrochemistry – Question Set
🔹 Section A: Basics of Electrochemistry
1. Define electrolysis.
2. What is the difference between electrical conductors and electrolytes?
3. Why do solid ionic compounds not conduct electricity, but molten or aqueous ones
do?
4. What are the roles of the anode and the cathode in electrolysis?
5. Distinguish between a strong and a weak electrolyte, with examples.
🔹 Section B: Electrolysis of Molten Compounds
6. Write the ionic half-equations for the electrolysis of molten sodium chloride.
7. In the electrolysis of molten lead(II) bromide:
o
a) What is formed at the cathode?
o
b) What is formed at the anode?
o
c) Write the half-equations for both electrodes.
8. Why must the electrolyte be molten or in aqueous solution for electrolysis to occur?
🔹 Section C: Electrolysis of Aqueous Solutions
9. What four ions are typically present in an aqueous sodium chloride solution?
10. During the electrolysis of dilute sulfuric acid, identify:

a) Products at the anode and cathode

b) Their respective half-equations
11. Describe and explain the products formed during the electrolysis of copper(II) sulfate
using:

a) Graphite electrodes

b) Copper electrodes
12. Why is hydrogen produced instead of sodium during the electrolysis of aqueous
sodium chloride?
🔹 Section D: Electroplating
13. What is electroplating, and what are its uses?
14. Describe the setup required to electroplate a metal spoon with silver.
15. In the electroplating process:

a) What should the anode be made of?

b) What should the cathode be made of?

c) What should the electrolyte contain?
🔹 Section E: Hydrogen Fuel Cells
16. Write the overall equation for the reaction that takes place in a hydrogen fuel cell.
17. What are the advantages of using hydrogen fuel cells over fossil fuels?
18. What challenges are associated with using hydrogen as a fuel?
19. Compare a hydrogen fuel cell with a conventional electrolysis setup.
20. Explain why hydrogen is described as a “clean” fuel.
🔹 Section F: Analytical and Predictive Questions
21. Predict the products formed during the electrolysis of:

a) Aqueous potassium nitrate

b) Concentrated hydrochloric acid
22. Identify the ions discharged at each electrode during the electrolysis of concentrated
sodium chloride solution.
23. A student electrolyzes copper(II) sulfate solution with carbon electrodes. After 10
minutes, the blue color of the solution fades. Explain why.
24. Why do mass changes occur at the electrodes during electrolysis with copper
electrodes?
🔹 Section G: Practical and Application Questions
25. Draw a labeled diagram of an electrolysis setup for molten lead(II) bromide.
26. How can you test the gases produced during electrolysis? Describe suitable tests for:

a) Hydrogen

b) Oxygen

c) Chlorine
27. A solution of copper(II) sulfate is electrolyzed with graphite electrodes. Predict and
explain what happens to the pH of the solution over time.
28. How does increasing the current affect the rate of electrolysis?
29. During the electrolysis of brine (concentrated NaCl solution), identify all three
products and describe how they can be collected
Chapter 7: Chemical Energetics – Question Set
Section A: Basic Concepts
1. Define an exothermic reaction. Give two examples.
2. Define an endothermic reaction. Give two examples.
3. Explain what happens to the temperature of the surroundings in an exothermic
reaction.
4. Why do endothermic reactions cause a decrease in temperature of the
surroundings?
Section B: Energy Level Diagrams
5. Draw an energy level diagram for an exothermic reaction and label the reactants,
products, and energy change.
6. Draw an energy level diagram for an endothermic reaction and explain the key
features.
Section C: Bond Energies and Calculations
7. Explain the difference between bond breaking and bond making in terms of energy
changes.
8. Using bond energies, calculate the enthalpy change for the reaction:
H₂ + Cl₂ → 2HCl
(Bond energies: H–H = 436 kJ/mol, Cl–Cl = 243 kJ/mol, H–Cl = 431 kJ/mol)
9. Why is the reaction in question 8 exothermic?
Section D: Enthalpy Changes
10. What is meant by standard enthalpy change of formation?
11. What is meant by standard enthalpy change of combustion?
12. Calculate the enthalpy change for the combustion of ethane (C₂H₆) given the
following enthalpies of formation:
o
C₂H₆ = –85 kJ/mol
o
CO₂ = –393 kJ/mol
o
H₂O = –286 kJ/mol
Reaction: C₂H₆ + 3.5 O₂ → 2 CO₂ + 3 H₂O
Section E: Fuels and Applications
13. Name two advantages and two disadvantages of hydrogen as a fuel.
14. Write the overall reaction that takes place in a hydrogen fuel cell.
15. Explain why hydrogen fuel cells are considered environmentally friendly.
Section F: Practical and Application Questions
16. Describe an experiment that can be used to investigate whether a reaction is
exothermic or endothermic.
17. How is the heat energy change calculated during a chemical reaction using
calorimetry? Include the formula and define each term.
18. Explain why dissolving ammonium nitrate in water causes a temperature decrease.
Chapter 8: Rates of Reaction – Question Set
Basic Understanding
1. What is meant by the rate of reaction?
2. Name two ways to measure the rate of a reaction involving a gas being produced.
3. How does increasing the concentration of a reactant affect the rate of reaction?
Explain why.
Collision Theory and Activation Energy
4. Explain collision theory in your own words.
5. What is activation energy and why is it important in chemical reactions?
6. Why do all collisions not lead to a reaction?
Factors Affecting Rate
7. How does increasing the surface area of a solid reactant affect the rate of reaction?
Give an example.
8. Describe what happens to the rate of reaction when temperature is increased. Why?
9. What role does a catalyst play in a chemical reaction?
Practical and Calculation-Based
10. A reaction between zinc and hydrochloric acid produces hydrogen gas. How would
you measure the rate of this reaction?
11. In an experiment, the time taken for a color change in an iodine clock reaction
decreases when the temperature is raised. What does this tell you about the effect of
temperature on reaction rate?
12. If the rate of reaction doubles when the temperature is increased by 10°C, what can
you infer about the relationship between temperature and reaction rate?
Extended Thinking
13. Why does powdered calcium carbonate react faster with acid than large chunks?
14. Why is it important to keep variables like temperature and concentration constant
during an experiment on reaction rates?
15. Describe an experiment you could do to investigate the effect of concentration on
the rate of reaction.
Chapter 9: Reversible Reactions and Equilibrium — Question Set
Basic Concepts
1. What is a reversible reaction? Give an example.
2. Explain what is meant by dynamic equilibrium in a closed system.
3. In a reversible reaction at equilibrium, what happens to the concentrations of
reactants and products?
Le Chatelier’s Principle
4. Describe what happens to the position of equilibrium if the concentration of a
reactant is increased.
5. How does increasing temperature affect an exothermic reversible reaction at
equilibrium?
6. Explain the effect of increasing pressure on the equilibrium position of the Haber
process.
Application and Examples
7. Write the chemical equation for the dehydration and rehydration of copper sulfate.
8. Why is the Haber process carried out at moderate temperature and high pressure?
9. What observable change occurs when the equilibrium between nitrogen dioxide and
dinitrogen tetroxide shifts?
Extended Thinking
10. A system at equilibrium is disturbed by decreasing the concentration of a product.
Describe how the system responds according to Le Chatelier’s principle.
11. Why can’t we obtain 100% yield of ammonia in the Haber process even if conditions
are optimized?
12. Explain why the color of a mixture of NO₂ and N₂O₄ changes when the temperature is
altered.
Practical Questions
13. Design a simple experiment to demonstrate the effect of temperature on the
position of equilibrium in a reversible reaction.
14. How would you expect the rate of the forward reaction to change if the pressure is
increased in a gaseous reversible reaction?
Chapter 10: Redox Reactions – Question Set
Basic Concepts
1. What is a redox reaction?
2. Define oxidation and reduction in terms of electrons.
3. Give an example of a reaction where oxygen is gained and explain why it is oxidation.
Identifying Redox Reactions
4. In the reaction between zinc and copper sulfate, which species is oxidized and which
is reduced?
5. What is the role of an oxidizing agent? Give an example.
6. What is the role of a reducing agent? Give an example.
Half-Equations and Electron Transfer
7. Write the oxidation half-equation for zinc reacting to form Zn²⁺ ions.
8. Write the reduction half-equation for copper ions forming copper metal.
9. Explain why redox reactions always involve both oxidation and reduction.
Real-Life Applications
10. Describe the redox changes occurring during the rusting of iron.
11. Explain how redox reactions are involved in the process of respiration.
12. Why is carbon used in the extraction of metals from their oxides?
Extended Thinking
13. Hydrogen peroxide can both oxidize and reduce substances in the same reaction
(disproportionation). Explain this statement with a chemical equation.
14. How can you tell from a chemical equation whether a redox reaction has taken
place?
15. In a displacement reaction, why does a more reactive metal displace a less reactive
metal from its compound?
Chapter 11: Acids and Bases – Question Set
Basic Understanding
1. Define an acid and a base in terms of proton transfer.
2. What is the difference between a base and an alkali?
3. Give two examples each of strong acids and strong alkalis.
4. What ions do acids and alkalis produce in aqueous solution?
Reactions and Properties
5. What happens to blue litmus paper in an acidic solution?
6. Write the general equation for the reaction between an acid and a metal carbonate.
7. What is meant by neutralization? Write a word equation to show it.
8. Explain why acids conduct electricity in aqueous solution.
The pH Scale and Strength
9. What does the pH scale measure?
10. A solution has a pH of 2. Is it a strong or weak acid? Explain.
11. Compare a strong acid and a weak acid with examples.
12. Why is universal indicator better than litmus in measuring pH?
Extended Questions
13. Explain what happens when hydrochloric acid reacts with calcium carbonate.
14. Describe how you could safely investigate the pH of a variety of household
substances.
15. Describe a test to distinguish between hydrochloric acid and sodium hydroxide
solution.
Chapter 12: Preparation of Salts – Question Set
General Understanding
1. What is a salt in chemistry?
2. What determines the name of the salt formed in a neutralization reaction?
3. Name the salt formed when:
a. Sulfuric acid reacts with sodium hydroxide
b. Hydrochloric acid reacts with calcium carbonate
c. Nitric acid reacts with potassium hydroxide
Methods of Salt Preparation
4. Outline a method for preparing a soluble salt from an acid and an insoluble base.
5. What safety precautions should be taken when preparing salts in the lab?
6. Why is filtration used in some salt preparations but not in others?
Reactions
7. Write a balanced chemical equation for the reaction between magnesium and
hydrochloric acid.
8. Describe the steps needed to prepare copper(II) sulfate crystals from copper(II) oxide
and sulfuric acid.
9. How can an insoluble salt like lead(II) sulfate be prepared in the laboratory?
Practical and Application
10. Explain how titration can be used to prepare sodium chloride.
11. Why is it important to add excess base or carbonate when reacting with acids?
12. After filtering a salt solution, how do you obtain dry crystals of the salt?
Extended/Challenge
13. A student wants to prepare zinc sulfate crystals. Which acid and base should they
use, and how?
14. Compare the precipitation method with the titration method of salt preparation.
15. Why can't you prepare sodium nitrate by reacting sodium with nitric acid directly in
school labs?
Chapter 13: The Periodic Table – Question Set
🔹 Section A: Multiple Choice / One-word Answer
1. What does the atomic number of an element represent?
2. Which group contains the noble gases?
3. What is the term for elements that have properties of both metals and non-metals?
4. Which side of the Periodic Table are non-metals generally found?
5. Which Group 1 element is more reactive: lithium or potassium?
🔹 Section B: Short Answer Questions
6. Why are elements in the same group said to have similar chemical properties?
7. State two physical properties that distinguish metals from non-metals.
8. Define "transition metals" and give two of their characteristics.
9. Why is helium placed in Group 0 even though it has only 2 electrons?
10. Compare the reactivity of chlorine and iodine. Which is more reactive and why?
🔹 Section C: Application & Reasoning
11. Explain why sodium reacts more vigorously with water than lithium.
12. Predict the formula of the compound formed when magnesium reacts with chlorine.
13. Why do noble gases not form compounds easily?
14. Describe how you can use the Periodic Table to predict whether an unknown
element is a metal or non-metal.
15. Identify two trends observed across Period 3 of the Periodic Table.
🔹 Section D: Higher-order / Challenge Questions
16. A student claims that all metals are hard and dense. Use knowledge of Group 1
metals to challenge this statement.
17. Fluorine displaces bromine from a solution of potassium bromide. What does this tell
you about the relative reactivities of halogens?
18. Describe the trend in ionization energy across a period and explain it in terms of
atomic structure.
19. Explain how the positions of elements in the Periodic Table relate to their electronic
configuration.
20. Iron (Fe) forms Fe²⁺ and Fe³⁺ ions, while sodium forms only Na⁺. Explain this
difference using knowledge of electronic structure and group positioning
Chapter 14 – Question Set: Metallic Elements and Alloys
🔹 Section A: Recall & Definitions (1-mark questions)
1. What is meant by the term malleable?
2. Define an alloy.
3. Name a metal that is liquid at room temperature.
4. Which metal is used in galvanising iron?
5. Which metal is the most reactive: iron, zinc, or potassium?
🔹 Section B: Short Answer (2–3 mark questions)
6. State three physical properties typical of metals.
7. What are the two main methods used to extract metals from their ores?
8. Explain why alloys are often harder than pure metals.
9. What is the composition and one common use of brass?
10. Describe a method to prevent iron from rusting.
🔹 Section C: Applications & Explanations (4–5 mark questions)
11. Explain why copper is suitable for use in electrical wiring.
12. Compare the reactivity of magnesium and iron with dilute hydrochloric acid, and
write balanced equations for both reactions.
13. Describe what happens during a displacement reaction between zinc and copper(II)
sulfate solution.
14. Iron is extracted from its ore using a blast furnace. Write a chemical equation
showing how carbon reduces iron(III) oxide to iron.
15. Discuss the environmental problems caused by metal extraction and how recycling
helps.
🔹 Section D: Higher-Order / Challenge Questions (5+ marks)
16. Design an experiment to compare the corrosion of three different metals (e.g. iron,
copper, and zinc) when exposed to air and water.
17. Why is aluminium used in aircraft construction even though it is more reactive than
iron?
18. Explain, with examples, how the properties of an alloy differ from the metals it is
made of.
19. The reactivity series helps predict which metal will displace another. Use this concept
to explain what would happen if aluminium metal is placed in a solution of iron(III)
chloride.
20. Stainless steel is used for surgical instruments. Explain the benefits of using an alloy
instead of pure iron for this application
Chapter 15 – Question Set: Reactivity of Metals
🔹 Section A: Recall & Basic Understanding (1–2 marks)
1. What is meant by the term "reactivity series"?
2. Name the most reactive metal in the reactivity series.
3. List three metals that do not react with cold water.
4. Which gas is produced when a reactive metal reacts with water or acid?
5. Write a word equation for the reaction of magnesium with steam.
6. What is observed when iron is added to copper(II) sulfate solution?
7. Which method is used to extract potassium from its ore?
8. State two physical properties that are common among reactive metals.
🔹 Section B: Application & Reasoning (3–4 marks)
9. Describe the reaction between zinc and hydrochloric acid. Write a balanced chemical
equation.
10. Explain why copper does not displace zinc from zinc sulfate solution.
11. Predict what will happen if aluminium is added to a solution of iron(III) chloride.
Write the balanced chemical equation.
12. Why do metals like gold and silver occur uncombined in nature?
13. Compare the reactions of calcium and magnesium with cold water.
14. A student placed three metals (A, B, C) into iron(II) sulfate solution. Only metal C
caused a reaction. Arrange the metals in order of increasing reactivity.
🔹 Section C: Extended Response & Analysis (5+ marks)
15. Explain how the reactivity of metals influences their method of extraction. Include at
least three examples.
16. Design a simple lab experiment to compare the reactivity of magnesium, zinc, and
copper using dilute hydrochloric acid.
17. What is a displacement reaction? Give two examples with balanced equations and
explain the observations.
18. The thermite reaction is used in rail welding. Write the chemical equation for the
reaction and explain why aluminium is used.
19. Discuss the redox changes in the following displacement reaction:
Zn+CuSO4→ZnSO4+CuZn + CuSO_4 → ZnSO_4 + CuZn+CuSO4→ZnSO4+Cu
Identify what is oxidised and what is reduced.
🔹 Section D: Real-World Application
20. How does the concept of reactivity play a role in the protection of underground
pipelines?
21. Why is magnesium used as a sacrificial metal in ship hulls?
22. Lithium is highly reactive and used in batteries. What precautions should be taken
when handling such reactive metals?
Chapter 16: Extraction and Corrosion of Metals — Detailed Notes
🔷 16.1 Introduction to Metal Extraction

Metals in the Earth’s crust are mostly found as ores — compounds of metals
combined with oxygen, sulfur, or other elements.

To use metals, they must be extracted (separated) from their ores.

The method of extraction depends largely on the reactivity of the metal (see
Chapter 15).
🔷 16.2 Methods of Extraction
1. Reduction with Carbon (Smelting)

Metals less reactive than carbon can be extracted by heating their oxides with
carbon.

Carbon removes oxygen from the metal oxide to leave the pure metal.
Example: Extraction of Iron in a Blast Furnace
Fe2O3+3C→2Fe+3COFe_2O_3 + 3C \rightarrow 2Fe + 3COFe2O3+3C→2Fe+3CO

Iron(III) oxide is reduced to iron.

Carbon acts as the reducing agent.
Note:

Carbon monoxide (CO) is often the actual reducing agent in blast furnaces.

Blast furnaces operate at high temperature, using coke (carbon), limestone, and iron
ore.
2. Electrolysis

Used for metals more reactive than carbon (e.g., aluminium, sodium, potassium).

Metal ores are melted or dissolved to allow electric current to flow.

Electrolysis breaks down compounds into metals and non-metals.
Example: Extraction of Aluminium from Bauxite

Bauxite (mainly Al₂O₃) is dissolved in molten cryolite to lower melting point.

Electrolysis produces aluminium metal at cathode and oxygen gas at anode.
3. Other Methods

Displacement reactions: Less common industrially but used in some cases.

Bioleaching and phytomining: Emerging methods using bacteria and plants to
extract metals (not usually in IGCSE syllabus but good to know!).
🔷 16.3 Corrosion of Metals

Corrosion is the slow chemical attack of metals by substances in their environment
(mainly oxygen and water).

Most commonly seen as rusting of iron.
Rusting of Iron

Iron reacts with oxygen and water to form hydrated iron(III) oxide, commonly called
rust.

Rust is porous and flaky, so it does not protect the metal underneath — rusting
continues until iron is completely destroyed.
Chemical Reaction:
4Fe+3O2+6H2O→4Fe(OH)34Fe + 3O_2 + 6H_2O \rightarrow 4Fe(OH)_34Fe+3O2+6H2
O→4Fe(OH)3

Iron(III) hydroxide dehydrates to form rust (Fe₂O₃·xH₂O).
🔷 16.4 Factors Affecting Rusting

Presence of water (including moist air) is essential.

Presence of oxygen in air.

Salt (like in seawater) speeds up rusting by increasing electrical conductivity of water.

Acidic conditions also accelerate corrosion.
🔷 16.5 Preventing Rusting (Corrosion Protection)
1. Barrier Methods

Painting, oiling, greasing: Physical barrier to block water and oxygen.

Plastic coatings: Modern, durable barrier layers.
2. Galvanising

Coating iron or steel with a layer of zinc.

Zinc acts as a sacrificial metal — it corrodes instead of iron because it is more
reactive.
3. Sacrificial Protection

Attach blocks of more reactive metals (e.g., magnesium or zinc) to iron structures.

These metals corrode preferentially, protecting the iron.
Used in: Ships’ hulls, underground pipelines.
4. Alloying

Stainless steel contains chromium which forms a thin, protective oxide layer.

Prevents rusting and corrosion even in wet conditions.
🔷 16.6 Extraction and Corrosion in Everyday Life

Iron: extracted in blast furnace, easily rusts.

Aluminium: extracted by electrolysis, resists corrosion due to protective oxide layer.

Copper: extracted by smelting, corrodes slowly forming a green layer of copper
carbonate (patina).
� Summary Table
Metal
Extraction
Method
Corrosion
Resistance
Protection Method
Iron
Blast furnace (C
reduction)
Rusts easily
Painting, galvanising,
sacrificial protection
Aluminium
Electrolysis
Forms protective
oxide layer
Usually no extra
protection needed
Copper
Smelting
Forms green patina
(corrosion)
Coating, alloying
Metal
Extraction
Method
Corrosion
Resistance
Protection Method
Zinc
Electrolysis
Corrodes slowly
Used in galvanising
💡 Examples & Applications

Galvanised steel roofs: Zinc layer protects steel underneath.

Oil pipelines: Sacrificial anodes attached to prevent corrosion.

Rust on cars: Must be removed and treated to prevent structural damage.

Copper roofs and statues: Develop patina that protects metal beneath