Conservation of Mass and Balancing Equations

This lesson covers the law of conservation of mass and how to balance chemical equations. These concepts are fundamental to all quantitative chemistry and are required by the Edexcel GCSE Chemistry specification (1CH0).

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The Law of Conservation of Mass

The law of conservation of mass states:

No atoms are created or destroyed during a chemical reaction.

This means the total mass of the reactants (the substances you start with) always equals the total mass of the products (the substances formed).

This law was established by Antoine Lavoisier in the 18th century and is one of the most fundamental principles in chemistry.

Why Is Mass Conserved?

In a chemical reaction, the atoms are rearranged — they break apart from one combination and join together in a new combination. The same atoms are present before and after the reaction; they have simply been reorganised.

Since the atoms are the same (and each atom has the same mass), the total mass cannot change.

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Apparent Changes in Mass

Sometimes, when you carry out a reaction in an open container, the mass of the container appears to change. This is not because mass has been created or destroyed — it is because gases have entered or left the container.

Mass Appears to Decrease

When a reaction produces a gas that escapes into the atmosphere:

Example: Heating calcium carbonate

CaCO₃ → CaO + CO₂

If carried out in an open crucible, the CO₂ gas escapes. The mass of the crucible and its contents decreases because the CO₂ has left.

Mass Appears to Increase

When a reaction involves a reactant from the air combining with a solid:

Example: Burning magnesium in air

2Mg + O₂ → 2MgO

The magnesium reacts with oxygen from the air. The mass of the product (MgO) is greater than the mass of magnesium you started with because oxygen atoms have been added.

Exam Tip: If asked to explain why mass appears to change, always state that mass is conserved overall, then explain which gas has entered or escaped. Never say mass is lost or gained — say it "appears to" change.

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Balancing Chemical Equations

Because atoms are conserved, every chemical equation must be balanced — there must be the same number of each type of atom on both sides of the equation.

Rules for Balancing

1. Write the word equation first (if not given).
2. Write the correct formulae for all reactants and products.
3. Count the atoms of each element on both sides.
4. Add coefficients (big numbers in front of formulae) to balance. Never change the subscript numbers in a formula.
5. Check that all atoms balance.

Exam Tip: You can only change the big numbers in front of formulae. Changing subscripts would change the substance itself (e.g., changing H₂O to H₂O₂ changes water to hydrogen peroxide).

flowchart TD
    A[Word equation] --> B[Write correct formulae for reactants and products]
    B --> C[Count atoms of each element on both sides]
    C --> D{Balanced?}
    D -- No --> E["Adjust coefficients (big numbers)<br/>Never change subscripts"]
    E --> C
    D -- Yes --> F[Final check: same atoms LHS = RHS]
    F --> G[Total mass of reactants = total mass of products]

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Worked Examples: Balancing Equations

Example 1: Hydrogen + Oxygen → Water

Unbalanced: H₂ + O₂ → H₂O

Count atoms:

• Left: H = 2, O = 2
• Right: H = 2, O = 1 ← oxygen is unbalanced

Put a 2 in front of H₂O:

H₂ + O₂ → 2H₂O

Count again:

• Left: H = 2, O = 2
• Right: H = 4, O = 2 ← hydrogen is now unbalanced

Put a 2 in front of H₂:

2H₂ + O₂ → 2H₂O

Final check:

• Left: H = 4, O = 2 ✓
• Right: H = 4, O = 2 ✓

Balanced equation: 2H₂ + O₂ → 2H₂O

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Example 2: Methane + Oxygen → Carbon Dioxide + Water

Unbalanced: CH₄ + O₂ → CO₂ + H₂O

Count atoms:

• Left: C = 1, H = 4, O = 2
• Right: C = 1, H = 2, O = 3

Balance hydrogen — put 2 in front of H₂O:

CH₄ + O₂ → CO₂ + 2H₂O

Count:

• Left: C = 1, H = 4, O = 2
• Right: C = 1, H = 4, O = 4

Balance oxygen — put 2 in front of O₂:

CH₄ + 2O₂ → CO₂ + 2H₂O

Final check:

• Left: C = 1, H = 4, O = 4 ✓
• Right: C = 1, H = 4, O = 4 ✓

Balanced equation: CH₄ + 2O₂ → CO₂ + 2H₂O

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Example 3: Iron + Oxygen → Iron(III) Oxide

Unbalanced: Fe + O₂ → Fe₂O₃

Count:

• Left: Fe = 1, O = 2
• Right: Fe = 2, O = 3

This requires more steps. Balance iron first by putting 2 in front of Fe:

2Fe + O₂ → Fe₂O₃ (O still unbalanced: 2 vs 3)

Try using 4Fe + 3O₂ → 2Fe₂O₃:

Check:

• Left: Fe = 4, O = 6 ✓
• Right: Fe = 4, O = 6 ✓

Balanced equation: 4Fe + 3O₂ → 2Fe₂O₃

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Example 4: Sodium + Water → Sodium Hydroxide + Hydrogen

Unbalanced: Na + H₂O → NaOH + H₂

Count:

• Left: Na = 1, H = 2, O = 1
• Right: Na = 1, O = 1, H = 3

Put 2 in front of Na, 2 in front of H₂O, and 2 in front of NaOH:

2Na + 2H₂O → 2NaOH + H₂

Check:

• Left: Na = 2, H = 4, O = 2 ✓
• Right: Na = 2, O = 2, H = 4 (2 from NaOH + 2 from H₂) ✓

Balanced equation: 2Na + 2H₂O → 2NaOH + H₂

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Example 5: Balancing with Brackets

Unbalanced: Ca(OH)₂ + HCl → CaCl₂ + H₂O

Count:

• Left: Ca = 1, O = 2, H = 3, Cl = 1
• Right: Ca = 1, Cl = 2, H = 2, O = 1

Balance Cl by putting 2 in front of HCl:

Ca(OH)₂ + 2HCl → CaCl₂ + H₂O