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GCSE Chemistry · Topic 3

Quantitative chemistry

Quantitative chemistry is the maths of chemistry: how much of each substance reacts, and how much product you can make.

We will work with relative formula mass, moles, balanced equations and ionic equations, empirical formulae, concentration and gas volumes.

Then we use limiting reactants, percentage yield and atom economy to judge how efficient a reaction is.

In this topic

  1. 3.1Conservation of mass and the moleFree below · 59 questions
  2. 3.2Reacting masses, concentration and yieldOpen in Full Marks · 62 questions
  3. 3.3Relative formula mass, percentage composition and empirical and molecular formulaeOpen in Full Marks · 47 questions
  4. 3.4Writing and balancing chemical equations, state symbols and ionic equationsOpen in Full Marks · 43 questions
  5. 3.5Limiting reactants and using moles to balance equationsOpen in Full Marks · 52 questions
  6. 3.6Atom economyOpen in Full Marks · 56 questions
  7. 3.7Molar volume and calculations involving gasesOpen in Full Marks · 56 questions

3.1 · Free notes

Conservation of mass and the mole

Diagram: Conservation of mass and the mole

Key points

In more detail

Practice

Conservation of mass and the mole: 10 questions

Pick an answer to check it, or open the answer for the explanation.

1. What is relative formula mass?

  1. The mass of one mole in kg
  2. The sum of relative atomic masses in a formula
  3. The number of atoms in a compound
  4. The mass of a single atom
Show answer

B: The sum of relative atomic masses in a formula. Add up the relative atomic masses of every atom shown in the formula.

2. How many particles are in one mole of a substance?

  1. 6.02 × 10²³
  2. 1000
  3. 22.4
  4. 6.02 × 10⁻²³
Show answer

A: 6.02 × 10²³. This is Avogadro's constant.

3. What is the relative formula mass of water, H₂O? (H = 1, O = 16)

  1. 16
  2. 17
  3. 18
  4. 20
Show answer

C: 18. (2 × 1) + 16 = 18.

4. How many moles are in 22 g of CO₂? (Mr = 44)

  1. 0.5
  2. 1
  3. 2
  4. 22
Show answer

A: 0.5. moles = mass ÷ Mr = 22 ÷ 44 = 0.5.

5. Why does the measured mass appear to decrease when a gas is produced in an open flask?

  1. Mass is destroyed
  2. The gas escapes into the air
  3. Atoms disappear
  4. The reaction isn't real
Show answer

B: The gas escapes into the air. The mass of the escaped gas is no longer in the container, but no mass is actually lost overall.

6. What is the Mr of CO₂? (C = 12, O = 16)

  1. 28
  2. 32
  3. 44
  4. 48
Show answer

C: 44. 12 + (2 × 16) = 44.

7. What is the relative formula mass (Mr) of calcium hydroxide, Ca(OH)₂? (Ca = 40, O = 16, H = 1)

  1. 74
  2. 58
  3. 97
  4. 57
Show answer

A: 74. The bracket multiplies O and H by 2: 40 + 2×(16 + 1) = 74. Forgetting the bracket gives 57.

8. How many moles are in 36 g of water, H₂O? (Mr = 18)

  1. 0.5 mol
  2. 2 mol
  3. 18 mol
  4. 648 mol
Show answer

B: 2 mol. Moles = mass ÷ Mr = 36 ÷ 18 = 2 mol. Dividing the other way round gives 0.5 mol.

9. What mass of sodium hydroxide, NaOH, is 0.25 mol? (Mr = 40)

  1. 10 g
  2. 160 g
  3. 6.25 g
  4. 0.25 g
Show answer

A: 10 g. Mass = moles × Mr = 0.25 × 40 = 10 g. Dividing 40 by 0.25 wrongly gives 160 g.

10. How many molecules are there in 2 mol of oxygen gas, O₂?

  1. 2.41 × 10²⁴
  2. 3.01 × 10²³
  3. 1.20 × 10²⁴
  4. 6.02 × 10²³
Show answer

C: 1.20 × 10²⁴. Particles = moles × 6.02 × 10²³ = 2 × 6.02 × 10²³ = 1.20 × 10²⁴ molecules. The question asks about molecules, not atoms.

Keep going in Full Marks

Full Marks has 59 questions on Conservation of mass and the mole (49 more than here), set out in practice levels. You also get notes and 3D models for all 7 subtopics of Quantitative chemistry, plus games, timed exams and past papers.

Practise 3.1 in Full MarksOpen Quantitative chemistry

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