GCSE Chemistry Equations & mass questions by topic

Complete the sentences about chemical equations.

1. (a)The large numbers in front of formulae in an equation are called …[1]
2. (b)The state symbol (g) shows a substance is a …[1]
3. (c)Balancing an equation must never change a …[1]
4. (d)Write the formula of the diatomic element chlorine.[1]

Balance the following equations.

1. (a)Mg + O₂ → MgO[1]
2. (b)NaOH + HCl → NaCl + H₂O (already balanced — state this)[1]
3. (c)CaCO₃ → CaO + CO₂ (already balanced — state this)[1]
4. (d)Al + O₂ → Al₂O₃[1]

A student heats 1.00 g of calcium carbonate strongly in a crucible with a loose lid, and the mass of solid left is 0.56 g.

1. (a)Name the gas released during heating.[1]
2. (b)Calculate the mass of gas released.[1]
3. (c)Explain why this reaction does not break the law of conservation of mass.[2]

A piece of magnesium ribbon with mass 2.4 g is burned in a crucible and 4.0 g of white powder remains.

1. (a)Name the white powder.[1]
2. (b)Explain why the powder’s mass is greater than the ribbon’s.[2]
3. (c)State the mass of oxygen that combined with the magnesium.[1]

Give the relative formula mass of each compound (H 1, C 12, O 16, Mg 24).

1. (a)H₂O[1]
2. (b)CO₂[1]
3. (c)MgO[1]

A student mixes lead nitrate solution with potassium iodide solution in a glass beaker standing on a balance. A bright yellow precipitate forms at once and settles to the bottom.

1. (a)Name the yellow precipitate.[1]
2. (b)Predict what happens to the balance reading as the precipitate forms, and explain your answer.[2]
3. (c)The student repeats the reaction in a sealed flask on a second balance. State the reading change and explain why it is the same as or different from the beaker result.[2]

Ammonia, NH₃, is the starting material for nitrogen fertilisers (N 14, H 1).

1. (a)Calculate the relative formula mass of ammonia.[1]
2. (b)Calculate the percentage by mass of nitrogen in ammonia. Give your answer to 1 decimal place.[2]
3. (c)Suggest why fertiliser bags state the nitrogen content as a percentage by mass.[1]

A student writes the unbalanced equation H₂ + O₂ → H₂O and then “fixes” it by writing H₂ + O₂ → H₂O₂.

1. (a)Explain why H₂O₂ is not an acceptable correction.[2]
2. (b)Write the correctly balanced equation.[1]
3. (c)Add state symbols for hydrogen burning in oxygen at room conditions.[2]

Copper(II) sulfate solution reacts with sodium hydroxide solution, giving a blue precipitate of copper(II) hydroxide and sodium sulfate solution.

1. (a)Write the full balanced equation with state symbols.[2]
2. (b)Write the ionic equation for the reaction.[2]
3. (c)Name the spectator ions and explain what happens to them.[2]

A compound contains by mass 30.4% nitrogen and 69.6% oxygen (N 14, O 16).

1. (a)Calculate the empirical formula.[3]
2. (b)The compound’s Mr is 92. Deduce its molecular formula.[2]

A student measures the resolution and behaviour of a balance and a thermometer, then runs three titrations: 23.10 cm³, 23.30 cm³ and 23.20 cm³.

1. (a)State the uncertainty of a single reading on a balance with 0.1 g resolution.[1]
2. (b)Calculate the mean titre and the range.[2]
3. (c)Explain what the range tells you about the readings, and why repeating improves confidence.[3]

5.72 g of hydrated sodium carbonate crystals, Na₂CO₃·xH₂O, are heated to constant mass, leaving 2.12 g of anhydrous sodium carbonate (Na 23, C 12, O 16, H 1).

1. (a)Calculate the amount, in moles, of anhydrous Na₂CO₃ remaining.[1]
2. (b)Calculate the mass and the amount, in moles, of water driven off.[2]
3. (c)Deduce the value of x and hence the full formula of the crystals.[3]

A chemist confirms sulfate ions in a solution by adding dilute hydrochloric acid and then barium chloride solution; a dense white precipitate forms.

1. (a)Write the full balanced equation, with state symbols, for barium chloride solution reacting with sodium sulfate solution.[2]
2. (b)Write the ionic equation for the formation of the precipitate.[2]
3. (c)Explain why dilute hydrochloric acid is added first, naming one ion that would otherwise cause a false positive.[2]

12.8 g of copper oxide is fully reduced in a stream of hydrogen, leaving 10.2 g of copper (Cu 64, O 16).

1. (a)Calculate the mass of oxygen removed and its amount in moles.[2]
2. (b)Calculate the amount of copper in moles and hence the empirical formula of the oxide.[3]
3. (c)Write the balanced equation for the reduction with H₂.[1]

This question is about quantitative reading of equations.

1. (a)In 4Al + 3O₂ → 2Al₂O₃, state the ratio of aluminium atoms to oxygen molecules consumed.[1]
2. (b)Explain, in terms of atoms, why the total mass of products must equal the total mass of reactants.[2]
3. (c)A sealed flask contains 6.4 g of methane and excess oxygen. The methane burns completely. Using CH₄ + 2O₂ → CO₂ + 2H₂O, explain how you could predict the mass of CO₂ formed, and state it (C 12, H 1, O 16).[4]

A gaseous hydrocarbon X burns completely in oxygen: X + 5O₂ → 3CO₂ + 4H₂O.

1. (a)Deduce the formula of X, explaining your reasoning from the products.[3]
2. (b)Calculate the Mr of X and name the compound (C 12, H 1).[2]
3. (c)State why the coefficient 5 in front of O₂ confirms your formula.[1]

A required-practical style investigation: 0.30 g of magnesium ribbon is burned in a crucible and the magnesium oxide formed weighs 0.50 g (Mg 24, O 16).

1. (a)Explain why the product has a greater mass than the original ribbon.[2]
2. (b)Show by calculation that the results are consistent with the formula MgO.[2]
3. (c)Write the balanced equation for the reaction.[1]
4. (d)During heating the lid must be lifted carefully from time to time. Evaluate the effect on the calculated empirical formula if some white smoke escapes each time.[3]

This question is about the thermal decomposition of copper(II) carbonate: CuCO₃ → CuO + CO₂ (Cu 64, C 12, O 16). A 2.48 g sample is heated to constant mass in an open crucible.

1. (a)Calculate the amount, in moles, of copper(II) carbonate in 2.48 g.[1]
2. (b)Use the equation to calculate the mass of copper(II) oxide expected.[3]
3. (c)The student’s final residue actually weighs 1.55 g. Give two possible reasons for the difference from your answer to (b).[2]
4. (d)Calculate the mass of carbon dioxide released, and state the decrease in the crucible reading shown on a balance underneath.[2]