Revision notes for AQA GCSE Chemistry Conservation of mass and balanced chemical equations. Open the guide for explanations and worked examples. Written against the AQA GCSE Chemistry (8462) specification, so the content matches what's examinable rather than general Chemistry background.
Conservation of mass and balanced chemical equations
What you'll learn
- Why chemical reactions do not lose or create atoms.
- How conservation of mass links to balanced equations.
- The difference between subscripts inside formulae and multipliers before formulae.
- A reliable method for balancing symbol equations.
The big idea: reactions rearrange atoms
A chemical reaction is a change where atoms are rearranged to make new substances. Bonds may break and new bonds may form, but the atoms themselves are still there.
An atom is the smallest particle of an element that keeps that element’s chemical identity. An element contains only one type of atom. A compound contains atoms of two or more different elements chemically bonded together.
So when a reaction happens, you should imagine atoms being reorganised — not disappearing.
Reactants and products
In a chemical equation, the starting substances go on the left of the arrow, and the substances made go on the right.
Definition
Reactants and products
A reactant is a starting substance in a chemical reaction. A product is a substance formed by the reaction.
For example, in the equation:
2Mg(s) + O₂(g) → 2MgO(s)
magnesium and oxygen are reactants. Magnesium oxide is the product.
The arrow means “react to form”. The letters in brackets are state symbols:
- (s) means solid
- (l) means liquid
- (g) means gas
- (aq) means aqueous, meaning dissolved in water
State symbols are useful, but they do not change the atom counting.
Conservation of mass
Definition
Law of conservation of mass
The law of conservation of mass states that no atoms are lost or made during a chemical reaction, so the total mass of the products equals the total mass of the reactants.
This works because atoms have mass. If the same atoms are present before and after the reaction, the total mass must stay the same.
In a balanced equation, the atom counts match on both sides. This is how the equation shows conservation of mass.
Key Idea
The key conservation idea
Chemical reactions rearrange atoms. They do not create new atoms or destroy existing atoms.
Example
Using conservation of mass
In a closed container, 6.0 g of carbon reacts with oxygen to form 22.0 g of carbon dioxide. What mass of oxygen reacted?
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The reaction is C(s) + O₂(g) → CO₂(g), so the reactants are carbon and oxygen, and the product is carbon dioxide. In a closed system, total reactant mass equals total product mass.
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Set up the mass relationship: mcarbon+moxygen=mcarbon dioxidem_{\text{carbon}} + m_{\text{oxygen}} = m_{\text{carbon dioxide}}mcarbon+moxygen=mcarbon dioxide, so 6.0 g+moxygen=22.0 g6.0\text{ g} + m_{\text{oxygen}} = 22.0\text{ g}6.0 g+moxygen=22.0 g.
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Rearrange to find the missing mass: moxygen=22.0 g−6.0 g=16.0 gm_{\text{oxygen}} = 22.0\text{ g} - 6.0\text{ g} = 16.0\text{ g}moxygen=22.0 g−6.0 g=16.0 g. The mass of oxygen that reacted was 16.0 g.
Common Mistake
Open containers can mislead you
If a gas escapes from an open container, the measured mass may seem to decrease. If a gas from the air reacts, the measured mass may seem to increase. The law of conservation of mass still applies if you include all reactants and products.
Formulae, subscripts and multipliers
A chemical formula shows which elements are present in a substance and how many atoms of each element are in one particle or unit of that substance.
For example, H₂O(l) means each water molecule contains:
- 2 hydrogen atoms
- 1 oxygen atom
The small ₂ in H₂O(l) is not optional decoration — it is part of the formula.
Definition
Subscripts and multipliers
A subscript is the small lowered number inside a formula, such as the ₂ in H₂O(l). A multiplier or coefficient is the normal-sized number before a formula, such as the 2 in 2H₂O(l), and it multiplies the whole formula.
So:
- H₂O(l) means one water molecule.
- 2H₂O(l) means two water molecules.
- The 2 in front multiplies everything in H₂O(l).
Tip
Invisible ones
If there is no coefficient, treat it as 1. If an element symbol has no subscript, treat it as 1 atom in that formula.
Example
Counting atoms in a formula
How many aluminium atoms and oxygen atoms are shown by 2Al₂O₃(s)?
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The coefficient 2 means there are two units of Al₂O₃(s), so every atom count inside the formula is doubled.
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One unit of Al₂O₃(s) contains 2 aluminium atoms and 3 oxygen atoms, using the subscripts ₂ and ₃.
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Multiply by the coefficient: aluminium atoms = 2×2=42 \times 2 = 42×2=4; oxygen atoms = 2×3=62 \times 3 = 62×3=6.
Common Mistake
Changing subscripts to balance
Never change a subscript when balancing an equation. Changing H₂O(l) to H₂O₂(l) changes the substance from water to hydrogen peroxide.
What makes an equation balanced?
A symbol equation uses chemical formulae instead of just words. It is more precise than a word equation because it shows the number of atoms involved.
Definition
Balanced symbol equation
A balanced symbol equation is a chemical equation where each element has the same number of atoms on the reactant side and the product side.
Balanced does not mean there are the same number of molecules on each side. It means the same number of atoms of each element.
For example:
2H₂(g) + O₂(g) → 2H₂O(l)
Atom count:
- Reactants: 4 hydrogen atoms and 2 oxygen atoms
- Products: 4 hydrogen atoms and 2 oxygen atoms
So the equation is balanced.
How to balance equations
When you balance an equation, you are only allowed to add or change coefficients in front of formulae. You must not alter the formulae themselves.
A good method is:
- Write the correct formulae for the reactants and products.
- Count the atoms of each element on both sides.
- Add coefficients to make one element balance.
- Recount and adjust until every element balances.
- Check the coefficients are the smallest whole numbers that work.
Example
Balancing hydrogen and oxygen
Balance this equation:
H₂(g) + O₂(g) → H₂O(l)
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Count the atoms before balancing. On the left there are 2 hydrogen atoms and 2 oxygen atoms. On the right there are 2 hydrogen atoms and 1 oxygen atom.
-
Balance oxygen first by putting 2 in front of H₂O(l): H₂(g) + O₂(g) → 2H₂O(l). Now the right side has 2 oxygen atoms, but it also has 4 hydrogen atoms.
-
Balance hydrogen by putting 2 in front of H₂(g): 2H₂(g) + O₂(g) → 2H₂O(l). Now both sides have 4 hydrogen atoms and 2 oxygen atoms.
Example
Balancing aluminium reacting with oxygen
Balance this equation:
Al(s) + O₂(g) → Al₂O₃(s)
-
Oxygen is awkward because it appears as O₂(g) on the left and O₃ inside Al₂O₃(s) on the right. The smallest number both 2 and 3 go into is 6, so use 3O₂(g) and 2Al₂O₃(s).
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The equation is now Al(s) + 3O₂(g) → 2Al₂O₃(s). Count aluminium on the right: 2 units of Al₂O₃(s) contain 2×2=42 \times 2 = 42×2=4 aluminium atoms.
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Put 4 in front of Al(s): 4Al(s) + 3O₂(g) → 2Al₂O₃(s). Check oxygen: left has 6 oxygen atoms and right has 2×3=62 \times 3 = 62×3=6 oxygen atoms, so it is balanced.
Common Mistake
Balancing only one element then stopping
After changing a coefficient, always recount every element. Fixing one element can unbalance another one.
Why this matters
Balanced equations are the link between particle diagrams and chemical measurements. If the atom count is equal on both sides, the total mass is conserved.
Later, you will use balanced equations to work out reacting masses and amounts in moles. For now, the essential skill is careful atom counting.
Exam technique
In the exam
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Start by checking the formulae and state symbols, then count atoms on both sides of the arrow.
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Balance by changing coefficients only; never change subscripts inside formulae.
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Finish by checking every element has the same atom count on both sides, and that the coefficients are the smallest whole numbers possible.
Self review
Check yourself
- In 3CO₂(g), how many carbon atoms and oxygen atoms are shown?
- Why can a reaction in an open beaker seem to lose mass?
- Balance this equation: Mg(s) + O₂(g) → MgO(s).
Chemical measurements, conservation of mass and the quantitative interpretation of chemical equations
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