What is the significance of chemical formulae in stoichiometry?

Chemical formulae are crucial in stoichiometry as they provide the exact ratio of atoms in a compound, which is essential for calculating the amounts of reactants and products in a chemical reaction. Understanding formulae allows students to balance equations and perform mole calculations, which are key components of the Cambridge IGCSE Chemistry curriculum.

How do you determine the empirical formula from experimental data?

To determine the empirical formula from experimental data, you need to find the simplest whole-number ratio of moles of each element in a compound. This involves converting the mass of each element to moles using their atomic masses, then dividing by the smallest number of moles to find the ratio. This process is a fundamental skill in stoichiometry taught in the Cambridge IGCSE Chemistry syllabus.

What is the difference between empirical and molecular formulae?

The empirical formula represents the simplest whole-number ratio of elements in a compound, while the molecular formula shows the actual number of each type of atom in a molecule. For example, the empirical formula of glucose is CH2O, whereas its molecular formula is C6H12O6. Understanding this distinction is important for Cambridge IGCSE students studying stoichiometry.

How can you use a chemical formula to calculate the molar mass of a compound?

To calculate the molar mass of a compound using its chemical formula, sum the atomic masses of all the atoms present in the formula. For instance, the molar mass of H2O is calculated by adding the masses of two hydrogen atoms and one oxygen atom. This calculation is essential for stoichiometry problems in the Cambridge IGCSE Chemistry curriculum.

Why is it important to balance chemical equations in stoichiometry?

Balancing chemical equations is crucial in stoichiometry because it ensures the law of conservation of mass is upheld, meaning the same number of each type of atom is present on both sides of the equation. This allows for accurate calculations of reactants and products, a key skill for Cambridge IGCSE Chemistry students.

Why is "Changing a chemical formula to balance an equation" a common mistake in Formulae?

Why it happens: Tempting shortcut. How to avoid it: Only change COEFFICIENTS in front of formulas. Never change subscripts. Source: 0620/42 — every series

Why is "Omitting state symbols" a common mistake in Formulae?

Why it happens: Speed. How to avoid it: Include (s), (l), (g) or (aq) for every species.

Why is "Charges not balanced in an ionic equation" a common mistake in Formulae?

Why it happens: Forgetting charge as well as atoms. How to avoid it: Total charge LHS = total charge RHS. Always check both atoms AND charges.

Why is "Writing $1$ before a formula" a common mistake in Formulae?

Why it happens: Trying to be explicit. How to avoid it: Coefficient of 1 is implied — don't write it.

StoichiometryCambridge IGCSE Chemistry (0620)

Formulae

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Formulae and Equations — Cambridge IGCSE 0620 Chemistry Extended (2026)

Writing chemical formulae from ion charges, balancing equations, and including state symbols. The base skill for everything stoichiometric.

What you’ll learn

Mapped to the Cambridge IGCSE 0620 syllabus (2026-2028).

3.1 — Write the formulae of simple compounds from ionic charges.
3.1 — Balance chemical equations.
3.1 — Include appropriate state symbols.
3.1 — Construct ionic equations from full equations (Extended).

Writing chemical formulae

Total positive charge = total negative charge. Use subscripts to balance.

Steps.

Write the symbol of each ion with its charge.
Find the lowest common multiple of the charges to balance.
Use subscripts to indicate how many of each ion. Drop the charges in the final formula.

Worked. Calcium chloride from Ca²⁺ and Cl⁻.

Need 2 Cl⁻ to balance 1 Ca²⁺.
Formula: $\text{CaCl}_{2}$ .

Worked. Aluminium oxide from Al³⁺ and O²⁻.

Need 2 Al³⁺ ( $+6$ total) to balance 3 O²⁻ ( $-6$ total).
Formula: $\text{Al}_{2}\text{O}_{3}$ .

Worked. Ammonium sulfate from NH₄⁺ and SO₄²⁻.

Need 2 NH₄⁺ to balance 1 SO₄²⁻.
Formula: $(\text{NH}_{4})_{2}\text{SO}_{4}$ . (Brackets when more than one polyatomic ion.)

Common ions to memorise.

Ion	Charge	Notes
Na⁺, K⁺, Li⁺	$+1$	Group 1 metals
Mg²⁺, Ca²⁺	$+2$	Group 2 metals
Al³⁺	$+3$	Group 13
NH₄⁺	$+1$	Ammonium
Cl⁻, Br⁻, I⁻	$-1$	Halogens
O²⁻, S²⁻	$-2$	Oxide, sulfide
OH⁻	$-1$	Hydroxide
NO₃⁻	$-1$	Nitrate
CO₃²⁻	$-2$	Carbonate
SO₄²⁻	$-2$	Sulfate
PO₄³⁻	$-3$	Phosphate

Transition metals have variable charges; Cambridge usually quotes the charge in the question (e.g. iron(II) = Fe²⁺, iron(III) = Fe³⁺, copper(II) = Cu²⁺).

Total + = total − in a formula.
Memorise common polyatomic ions.
Brackets for multiple polyatomic ions.
Roman numeral = ion charge for transition metals.

Balancing equations

Same number of each type of atom on both sides. Use coefficients only.

Conservation of mass. Atoms are neither created nor destroyed in a chemical reaction. So the count of each atom must be the same on both sides.

Strategy.

Write the unbalanced equation (correct formulae for reactants and products).
Count atoms of each element on both sides.
Add coefficients (numbers in front of formulae) to balance.
NEVER change subscripts within a formula — that changes what the substance IS.

Worked. Combustion of methane.

Unbalanced: $\text{CH}_{4} + \text{O}_{2} \to \text{CO}_{2} + \text{H}_{2}\text{O}$ .
Carbon: 1 = 1 ✓.
Hydrogen: 4 vs 2 → put 2 in front of $\text{H}_{2}\text{O}$ : $\text{CH}_{4} + \text{O}_{2} \to \text{CO}_{2} + 2\text{H}_{2}\text{O}$ .
Oxygen: 2 vs 4 → put 2 in front of $\text{O}_{2}$ : $\text{CH}_{4} + 2\text{O}_{2} \to \text{CO}_{2} + 2\text{H}_{2}\text{O}$ .
Final check: C 1=1, H 4=4, O 4=4 ✓.

Worked. Aluminium + chlorine.

Unbalanced: $\text{Al} + \text{Cl}_{2} \to \text{AlCl}_{3}$ .
Need 2 Al + 3 Cl₂ → 2 AlCl₃ to balance.
Final: $2\text{Al} + 3\text{Cl}_{2} \to 2\text{AlCl}_{3}$ .

Tip. When you have an odd-numbered atom on one side (like the H in NH₃) and even-numbered on the other, multiply both sides by 2 to find a common ground.

Atoms must balance (conservation of mass).
Coefficients only, never subscripts.
Balance one element at a time.
Re-check at the end.

State symbols and ionic equations

Use $(s), (l), (g), (aq)$ to show physical state. Ionic equations show only the ions that change.

State symbols.

$(s)$ — solid.
$(l)$ — pure liquid.
$(g)$ — gas.
$(aq)$ — aqueous (dissolved in water).

Always include them in Cambridge mark-scheme answers.

Worked. $\text{Mg}(s) + 2\text{HCl}(aq) \to \text{MgCl}_{2}(aq) + \text{H}_{2}(g)$ .

Ionic equations (Extended). Show only the ions that actually change. Spectator ions (those that appear unchanged on both sides) are removed.

Steps.

Write the full balanced equation with state symbols.
Split aqueous compounds into their ions.
Cancel any ions that appear identically on both sides (spectators).
What remains is the ionic equation.

Worked. Reaction of sodium hydroxide with hydrochloric acid.

Full: $\text{NaOH}(aq) + \text{HCl}(aq) \to \text{NaCl}(aq) + \text{H}_{2}\text{O}(l)$ .
Split aqueous: $\text{Na}^{+}(aq) + \text{OH}^{-}(aq) + \text{H}^{+}(aq) + \text{Cl}^{-}(aq) \to \text{Na}^{+}(aq) + \text{Cl}^{-}(aq) + \text{H}_{2}\text{O}(l)$ .
Cancel Na⁺ and Cl⁻ (spectators).
Ionic: $\text{H}^{+}(aq) + \text{OH}^{-}(aq) \to \text{H}_{2}\text{O}(l)$ .

This is the SAME ionic equation for ANY acid + alkali neutralisation — that's why neutralisation always involves H⁺ and OH⁻ combining to make water.

Charge must balance in ionic equations too. Sum of charges on left = sum on right.

$(s), (l), (g), (aq)$ .
Ionic equation: drop spectator ions.
Acid + alkali: $\text{H}^{+} + \text{OH}^{-} \to \text{H}_{2}\text{O}$ .
Charge AND atoms balance.

How it’s examined

Formulae and equations are foundational and appear EVERY paper. Paper 2: 3-5 marks for writing formulae and balancing. Paper 4: 5-7 marks combined with stoichiometric calculations. Examiner reports flag changing subscripts when balancing (you only adjust coefficients) and missing state symbols.

Sources: Cambridge IGCSE Chemistry 0620 syllabus 2026-2028 (3.1); 0620/42 Oct/Nov 2024 — Q7 (balancing, ionic equations); 0620 Examiner Reports 2022-2024. Last reviewed 2026-05-06.

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Step-by-step worked examples — Formulae

Step-by-step solutions to past-paper-style questions on formulae, written exactly the way a tutor would explain them at the board.

1Use ion charges to write formulae
Extended• Adapted from 0620/22 May/Jun 2024 Q4• formula
Question
Write the formulae of: aluminium oxide; calcium nitrate; ammonium sulfate.
Step-by-step solution
1. Step 1
  $\text{Al}^{3+}$ and $\text{O}^{2-}$ → swap charges → $\text{Al}_{2}\text{O}_{3}$ .
2. Step 2
  $\text{Ca}^{2+}$ and $\text{NO}_{3}^{-}$ → $\text{Ca(NO}_{3})_{2}$ .
3. Step 3
  $\text{NH}_{4}^{+}$ and $\text{SO}_{4}^{2-}$ → $(\text{NH}_{4})_{2}\text{SO}_{4}$ .
Answer
$\text{Al}_{2}\text{O}_{3}$ , $\text{Ca(NO}_{3})_{2}$ , $(\text{NH}_{4})_{2}\text{SO}_{4}$ .
2Balance an equation
Core• balance
Question
Balance: $\text{C}_{4}\text{H}_{10} + \text{O}_{2} \to \text{CO}_{2} + \text{H}_{2}\text{O}$ .
Step-by-step solution
1. Step 1
  Balance C and H first.
  $\text{C}_{4}\text{H}_{10} + \text{O}_{2} \to 4\text{CO}_{2} + 5\text{H}_{2}\text{O}$
2. Step 2
  RHS has $4 \times 2 + 5 = 13$ oxygen atoms. Use $13/2$ on the left.
  $\text{C}_{4}\text{H}_{10} + \dfrac{13}{2}\,\text{O}_{2} \to 4\text{CO}_{2} + 5\text{H}_{2}\text{O}$
3. Step 3
  Multiply through by 2 to clear the fraction.
  $2\text{C}_{4}\text{H}_{10} + 13\,\text{O}_{2} \to 8\text{CO}_{2} + 10\text{H}_{2}\text{O}$
Answer
$2\text{C}_{4}\text{H}_{10} + 13\,\text{O}_{2} \to 8\text{CO}_{2} + 10\text{H}_{2}\text{O}$
3Add state symbols
Extended• states
Question
Add state symbols: zinc + dilute hydrochloric acid → zinc chloride solution + hydrogen gas.
Step-by-step solution
1. Step 1
  Zinc: solid (s); HCl(aq); ZnCl $_{2}$ in solution; H $_{2}$ gas.
Answer
$\text{Zn(s)} + 2\text{HCl(aq)} \to \text{ZnCl}_{2}\text{(aq)} + \text{H}_{2}\text{(g)}$
4Write an ionic equation
Extended• ionic equation
Question
Silver nitrate solution reacts with sodium chloride solution to form a silver chloride precipitate. Write the ionic equation.
Step-by-step solution
1. Step 1
  Full: $\text{AgNO}_{3}(\text{aq}) + \text{NaCl}(\text{aq}) \to \text{AgCl}(\text{s}) + \text{NaNO}_{3}(\text{aq})$ .
2. Step 2
  Cancel spectator ions ( $\text{Na}^{+}$ , $\text{NO}_{3}^{-}$ ).
3. Step 3
  Net ionic equation.
  $\text{Ag}^{+}(\text{aq}) + \text{Cl}^{-}(\text{aq}) \to \text{AgCl}(\text{s})$
Answer
$\text{Ag}^{+}(\text{aq}) + \text{Cl}^{-}(\text{aq}) \to \text{AgCl}(\text{s})$

Key Formulae — Formulae

The formulae you need to memorise for formulae on the Cambridge IGCSE 0620 paper, with every variable defined in plain English and a note on when to use it.

Ion-charge swap (criss-cross)
$M^{a+} + X^{b-} \to M_{b}X_{a}$
When to use
Quickly write the formula of an ionic compound.

Key Definitions and Keywords — Formulae

Definitions to memorise and the exact keywords mark schemes credit for formulae answers — sharpened from recent examiner reports for the 2026 0620 sitting.

State symbols
Examiner keyword
(s) solid, (l) liquid, (g) gas, (aq) aqueous (dissolved in water).
Balanced equation
Examiner keyword
Equal numbers of each type of atom (and charge) on both sides.
Ionic equation
Examiner keyword
Shows only the ions/atoms that take part in the reaction (spectators removed).
Spectator ion
Examiner keyword
An ion that appears unchanged on both sides of the equation.

Common Mistakes and Misconceptions — Formulae

The traps other students keep falling into on formulae questions — taken from recent Cambridge IGCSE 0620 examiner reports and mark schemes — and how to avoid them.

✕Changing a chemical formula to balance an equation
0620/42 — every series
Why it happens
Tempting shortcut.
How to avoid it
Only change COEFFICIENTS in front of formulas. Never change subscripts.
✕Omitting state symbols
Why it happens
Speed.
How to avoid it
Include $(\text{s})$ , $(\text{l})$ , $(\text{g})$ or $(\text{aq})$ for every species.
✕Charges not balanced in an ionic equation
Why it happens
Forgetting charge as well as atoms.
How to avoid it
Total charge LHS = total charge RHS. Always check both atoms AND charges.
✕Writing $1$ before a formula
Why it happens
Trying to be explicit.
How to avoid it
Coefficient of $1$ is implied — don't write it.

Practice questions

Exam-style questions with step-by-step worked solutions. Try one before checking the method.

Past paper style quiz

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Video lesson

Short walkthrough of the concepts students most often get stuck on.

Formulae — frequently asked questions

The things students keep getting wrong in this sub-topic, answered.

Ion

Charge

Notes

Na⁺, K⁺, Li⁺

+1

Group 1 metals

Mg²⁺, Ca²⁺

+2

Group 2 metals

Al³⁺

+3

Group 13

NH₄⁺

+1

Ammonium

Cl⁻, Br⁻, I⁻

-1

Halogens

O²⁻, S²⁻

-2

Oxide, sulfide

OH⁻

-1

Hydroxide

NO₃⁻

-1

Nitrate

CO₃²⁻

-2

Carbonate

SO₄²⁻

-2

Sulfate

PO₄³⁻

-3

Phosphate

State symbols.

$(s)$ — solid.
$(l)$ — pure liquid.
$(g)$ — gas.
$(aq)$ — aqueous (dissolved in water).

Always include them in Cambridge mark-scheme answers.

Worked. $\text{Mg}(s) + 2\text{HCl}(aq) \to \text{MgCl}_{2}(aq) + \text{H}_{2}(g)$ .

Ionic equations (Extended). Show only the ions that actually change. Spectator ions (those that appear unchanged on both sides) are removed.

Steps.

Write the full balanced equation with state symbols.
Split aqueous compounds into their ions.
Cancel any ions that appear identically on both sides (spectators).
What remains is the ionic equation.

Worked. Reaction of sodium hydroxide with hydrochloric acid.

Full: $\text{NaOH}(aq) + \text{HCl}(aq) \to \text{NaCl}(aq) + \text{H}_{2}\text{O}(l)$ .
Split aqueous: $\text{Na}^{+}(aq) + \text{OH}^{-}(aq) + \text{H}^{+}(aq) + \text{Cl}^{-}(aq) \to \text{Na}^{+}(aq) + \text{Cl}^{-}(aq) + \text{H}_{2}\text{O}(l)$ .
Cancel Na⁺ and Cl⁻ (spectators).
Ionic: $\text{H}^{+}(aq) + \text{OH}^{-}(aq) \to \text{H}_{2}\text{O}(l)$ .

This is the SAME ionic equation for ANY acid + alkali neutralisation — that's why neutralisation always involves H⁺ and OH⁻ combining to make water.

Charge must balance in ionic equations too. Sum of charges on left = sum on right.

Formulae

Short Study Notes in the form of Slides

Detailed Study Notes

What you’ll learn

How it’s examined

1Use ion charges to write formulae

2Balance an equation

3Add state symbols

4Write an ionic equation

Ion-charge swap (criss-cross)

State symbols

Balanced equation

Ionic equation

Spectator ion

✕Changing a chemical formula to balance an equation

✕Omitting state symbols

✕Charges not balanced in an ionic equation

✕Writing 111 before a formula

Practice questions

Past paper style quiz

Assess your understanding

Video lesson

Formulae — frequently asked questions

Formulae

Short Study Notes in the form of Slides

Detailed Study Notes

What you’ll learn

How it’s examined

1Use ion charges to write formulae

2Balance an equation

3Add state symbols

4Write an ionic equation

Ion-charge swap (criss-cross)

State symbols

Balanced equation

Ionic equation

Spectator ion

✕Changing a chemical formula to balance an equation

✕Omitting state symbols

✕Charges not balanced in an ionic equation

✕Writing 111 before a formula

Practice questions

Past paper style quiz

Assess your understanding

Video lesson

Formulae — frequently asked questions

✕Writing $1$ before a formula

✕Writing $1$ before a formula