Category 1: Relative Masses and the Mole Concept (5 Questions)

• Relative Formula Mass (Mr): Calculate the Mr for:

  • (a) Ammonium Sulfate, (NH4)2SO4

  • (b) Hydrated Copper(II) Sulfate, CuSO4.5H2O

• Mass to Moles: How many moles are present in 10g of Calcium Carbonate (CaCO3)?

• Moles to Mass: Calculate the mass of 0.05 moles of Sodium Carbonate (Na2CO3).

• Avogadro’s Constant: How many atoms are present in 2.3g of Sodium metal? (Na)

• Percentage Composition: Calculate the percentage by mass of Nitrogen in Urea, CO(NH2)2.

Category 2: Empirical and Molecular Formulae (5 Questions)

• Empirical Formula (EF): A compound contains 40% Carbon, 6.7% Hydrogen, and 53.3% Oxygen by mass. Determine its EF.

• Molecular Formula (MF): A compound has the EF CH2 and a relative formula mass of 84. Find its MF.

• Water of Crystallisation: 5.00g of hydrated Barium Chloride (BaCl@.xH2O) was heated until all water was lost. The remaining anhydrous salt weighed 4.26g. Find the value of x.

• Reduction of Oxides: 1.6g of an oxide of Iron was reduced by Hydrogen to form 1.12g of Iron metal. Determine the formula of the oxide.

• Combustion Analysis: When 0.1 moles of a hydrocarbon CxHy is burned completely, it produces 0.3 moles of CO2 and 0.4 moles of H2O.​​ Find the formula of the hydrocarbon.

Category 3: Molar Volume of Gases (5 Questions)

(Assume Molar Volume = 22.4 dm3​​ at s.t.p.)
11. Volume to Moles: Calculate the number of moles in 480 cm3 of Oxygen gas at​​ s.t.p.
12. Mass to Volume: What volume (at s.t.p.) is occupied by 4.4g of Carbon Dioxide?
13. Reacting Volumes (Gay-Lussac): For the reaction N2​​ + 3H2​​ ⇌ 2NH3, what volume of Ammonia is produced from 15 cm3​​ of Nitrogen?
14. Density: Calculate the density of Chlorine gas (Cl2)​​ in g/dm3 at​​ s.t.p.
15. Unknown Gas: 0.28g of a gas occupies 240 cm3​​ at​​ s.t.p. Calculate its Mr​​ and identify the gas.

Category 4: Reacting Masses & Stoichiometry (5 Questions)

• Basic Mass-Mass: What mass of Magnesium Oxide is formed when 12g of Magnesium burns in excess Oxygen?

• • 2Mg + O2​​ → 2MgO

• Thermal Decomposition: What mass of CO2 is released when 20g of Calcium Carbonate is heated?

• • CaCO3​​ → CaO CO2

• Limiting Reagent: If 5g of Iron reacts with 5g of Sulfur, which is the limiting reagent?

• • Fe + S → FeS

• Percentage Yield: A student expected 10g of product but only obtained 8g. Calculate the percentage yield.

• Purity: 5.0g of impure Zinc reacted with excess HCl to produce 1.2 dm3 of H2 at r.t.p. Calculate the percentage purity of the Zinc.

Category 5: Solutions and Concentrations (5 Questions)

• Molarity (mol/dm3): Calculate the concentration of a solution containing 4g of NaOH dissolved in 250 cm3 of water.

• Mass in Solution: How many grams of Silver Nitrate (AgNO3) are needed to prepare 500 mL of a 0.1 mol dm-3​​ solution?

• Dilution: If 100 cm3​​ of 2.0 mol dm-3​​ HCl​​  is diluted to 500 cm3, what is the new concentration?

• Gram Concentration: Convert 0.25 mol dm-3 of H2SO4 into g dm-3.

• Ion Concentration: What is the concentration of Cl- ions in a 0.5 mol dm-3 solution of MgCl2?

Category 6: Titration Calculations (5 Questions)

• Acid-Base Titration: 25 cm3 of 0.1 mol dm-3 required 20.0 cm3 of H2SO4 for neutralization. Calculate the molarity of the acid.

• • 2NaOH + H2SO4​​ → Na2SO4​​ + 2H2O

• Back Titration (Advanced): An excess of 50 cm3 of 1.0 mol dm-3​​ HCl was added to 2g of impure CaCO3. The excess acid required 10cm3​​  of 1.0 mol dm-3​​ NaOH for neutralization. Calculate the % purity of the carbonate.

• Unknown Metal: A divalent metal M​​ reacts with HCl. ​​ 0.12g​​  of M​​  produced 120 cm3 of H2​​ at r.t.p. Find the Ar of M.

• • M + 2HCl → MCl2​​ + H2

• Gas from Solid: What volume of CO2 (at r.t.p.) is produced when 5.3g of Na2CO3 reacts with excess HCl?

• • Na2CO3​​ + 2HCl → 2NaCl + CO2​​ + H2O

Acids and Bases: Reactions

• Neutralization Reactions: Write balanced chemical equations (including state symbols) for the reaction between:

  • Dilute Hydrochloric acid and Sodium Hydroxide.

  • Sulfuric acid and Copper(II) oxide.

• Reactions with Carbonates: Describe the observations made when dilute Nitric acid is added to a beaker containing Calcium Carbonate chips. How would you test the gas evolved?

• Amphoteric Oxides: Define the term amphoteric. Name two oxides that exhibit this property and state how they behave when reacted with a strong alkali like Sodium Hydroxide.

• Acidity and Alkalinity: Distinguish between a strong acid and a concentrated acid. Use the concept of ionization in your answer.

• Proton Transfer: Using the Bronsted-Lowry theory, identify the acid and the base in the following reaction:
  

Salt Preparation

• Soluble Salts (Excess Base Method): Describe the steps required to prepare a pure, dry sample of Copper(II) Sulfate crystals starting from Copper(II) Oxide and Sulfuric acid. Explain why the oxide is added in excess.

• Soluble Salts (Titration): Why is the titration method used to prepare Sodium Chloride instead of the "excess base" method? Explain the role of the indicator in this process.

• Insoluble Salts (Precipitation): You are required to prepare a sample of Barium Sulfate.

  • Name two soluble starting reagents you could use.

  • Briefly outline the filtration and washing process needed to obtain the pure salt.

Hard Water

• Causes of Hardness: Identify the two specific metal ions responsible for causing hardness in water.

• Temporary vs. Permanent: Explain the chemical difference between temporary hardness and permanent hardness in terms of the anions present.

• Removal of Hardness: Write a chemical equation to show how boiling removes temporary hardness. Why does boiling fail to remove permanent hardness?

• Ion Exchange: Describe how an ion-exchange resin works to soften water. Which ions are being swapped?

• Economic Impact: State one advantage and one disadvantage of living in a hard water area (e.g., in the Maltese Islands).

Pollution

• Acid Rain: Identify the two main gases responsible for acid rain. State one man-made source for each gas.

• The Greenhouse Effect: Name two greenhouse gases. Explain how an increase in these gases leads to global warming.

• Photochemical Smog: Explain how nitrogen oxides (NOx) are formed in car engines and name the device fitted to cars to reduce these emissions.

• Water Pollution: Explain the process of eutrophication, starting from the runoff of nitrogenous fertilizers into a valley or pond.

• Carbon Monoxide: Why is Carbon Monoxide (CO) considered a toxic gas? Relate your answer to its effect on hemoglobin in the blood.

States of Matter & Kinetic Theory

• Particle Arrangement: Use the kinetic theory to contrast the arrangement and movement of particles in a solid versus a gas.

• Phase Changes: Define the terms sublimation and condensation. Give one example of a substance that sublimes at room temperature.

• Diffusion: Explain why a smell (like perfume) spreads faster in a warm room than in a cold room. Refer to the kinetic energy of particles.

• Gas Pressure: Using the kinetic theory, explain why the pressure inside a car tire increases when the car is driven for a long distance on a hot day.

• Cooling Curves: Sketch a cooling curve for a pure substance as it turns from a liquid to a solid. Label the melting point and explain why the temperature remains constant during the change of state.

Day 1  Day 2  Day 3  Day 4  Day 5  Day 6

Day 1

Question 1

Aqueous iron(III) chloride, FeCl3, reacts with aqueous potassium iodide, KI.

vFeCl3​​ + wKI​​ ​​ xFeCl2​​ + yKCl + I2

Which statements are correct?

• In the balanced equation, v, w, x and y have the same value.

• Potassium iodide is an oxidising agent.

• A dark brown solution is produced in the reaction.

• 1 and 2

• 1 and 3

• 2 only

• 2 and 3

Question 2

The relative atomic mass, Ar, of an element is the average mass of the isotopes of that element compared to another particle.

Which particle is used for this comparison?

• A proton

• An atom if​​ 12C

• An atom of​​ 40Ca

• An atom of​​ 1H

Day 2

Question 1

The equation for the combustion of methane is shown.

CH4​​ + 2O2​​ ​​ CO2​​ + 2H2O

Which mass of methane produces 36g of water?

• 16​​ g

• 18​​ g

• 32​​ g

• 64​​ g

Question 2

Calcium carbonate, CaCO3, reacts with dilute hydrochloric acid to produce carbon dioxide.

The equation for the reaction is shown. The relative formula​​ mass of calcium carbonate is 100.

CaCO3​​ + 2HCl​​ ​​ CaCl2​​ + H2O + CO2

10g of calcium carbonate is reacted with excess of dilute hydrochloric acid.

Which mass of carbon dioxide is produced?

• 2.2 g

• 2.8 g

• 4.4 g

• 44 g

Day 3

Question 1

The equation for the reaction between sodium carbonate and excess dilute hydrochloric acid shown.

Na2CO3​​ + 2HCl​​ ​​ 2NaCl + H2O + CO2

When 26.5 g of sodium carbonate reacts with excess dilute hydrochloric acid, what is the maximum volume of carbon dioxide produced?

• 6 dm3

• 12 dm3

• 18dm3

• 24dm3

Question 2

The equation shows the reaction between magnesium and sulfuric acid.

[Ar: H, 1; O, 16; Mg, 24; S, 32]

Mg + H2SO4​​ ​​ MgSO4​​ + H2

In this reaction, which mass of magnesium sulfate is formed when 6 g of magnesium react with excess sulfuric acid?

• 8 g

• 24 g

• 30 g

• 60 g

Day 4

Question 1

A compound contains 1.10 mol of K, 0.55 mol of C, and 1.65 mol of O. What is the empirical formula of this compound?

Question 2

An organic compound contains carbon, hydrogen, and oxygen. If elemental analysis reveals 38.71% carbon and 9.67% hydrogen, which is its empirical formula?

• CH3O

• CH2O

• CHO

• CH4O

Day 5

Question 1

Magnesium reacts with steam.

Mg + H2O​​ ​​ MgO + H2

When 2.43g of magnesium reacts with an excess of steam, the products are 4.03 g of magnesium oxide and 0.20 g of hydrogen.

What is produced when 7.29 g of magnesium is reacted with an excess of steam?

• 1.34 g of magnesium oxide and 0.07 g of hydrogen.

• 4.03 g of magnesium oxide and 0.20 g of hydrogen.

• 8.06 g of magnesium oxide and 0.40 g of hydrogen.

• 12.09 g of magnesium oxide and 0.60 g of hydrogen.

Day 6

Question 1

The equation for the thermal decomposition of sodium hydrogencarbonate is shown.

2NaHCO3​​ ​​ Na2CO3​​ + H2O + CO2

The Mr​​ of sodium hydrogencarbonate, NaHCO3, is 84.

The Mr​​ of sodium carbonate, Na2CO3, is 106.

In an experiment, 2.1 g of sodium hydrogencarbonate is heated but not all of it decomposes. All of the carbon dioxide is collected and measured at room temperature and pressure. The total volume of carbon dioxide produces id 0.21dm3.

The volume of 1 mole of gas at room temperature and pressure is 24dm3.

Which statement is correct?

• The mass of sodium carbonate produced is 0.93 g.

• The mass of sodium carbonate produced is 1.33 g.

• The percentage yield of carbon dioxide is 10%.

• The percentage yield of carbon dioxide is 35%.