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*Revision Materials* 1 Atomic Structure 2 Atoms, molecules and stoichiometry 3 Chemical Bonding 4 States of matter 5 Chemical energetics 6 Electrochemistry 7 Equilibria 8 Reaction kinetics 9 The Periodic Table, chemical periodicity 10 Group 2 11 Group 17 12 Nitrogen and sulfur 13 Organic 14 Hydrocarbons 15 Halogen compounds 16 Hydroxy compounds 17 Carbonyl compounds 18 Carboxylic acids and derivatives 19 Nitrogen compounds 20 Polymerisation 21 Organic synthesis 22 Analytical techniques 23 Chemical energetics 24 Electrochemistry 25 Equilibria 26 Reaction kinetics 27 Group 2 28 Chemistry of transition elements 29 Organic 30 Hydrocarbons 31 Halogen compounds 32 Hydroxy compounds 33 Carboxylic acids and derivatives 34 Nitrogen compounds 35 Polymerisation 36 Organic synthesis 37 Analytical techniques

2 Atoms, molecules and stoichiometry

2.1 Relative masses of atoms and molecules 2.2 The mole and the Avogadro constant 2.3 Formulas 2.4 Reacting masses and volumes (of solutions and gases)

Formulas

Specification Reference Physical Chemistry: Atoms, molecules and stoichiometry 2.3

Quick Notes

  • Ionic charges can be predicted from Periodic Table group number.
  • You need to know these compound ions:
    • NO3−, CO32−, SO42−, OH−, NH4+, Zn2+, Ag+, HCO3−, PO43−
  • Charges can be used to construct balanced formulas for ionic compounds.
  • Ionic equations exclude spectator ions.
  • Empirical formula: simplest whole-number ratio of atoms.
  • Molecular formula: actual number of atoms of each element in a molecule.
  • Anhydrous = no water; hydrated = contains water; water of crystallisation = fixed water in crystal.

Full Notes

Writing Formulas from Charges and Oxidation Numbers

Predicting Ionic Charges

You can predict the charge of an ion based on its position in the Periodic Table:

CIE A-Level Chemistry periodic table diagram indicating typical ionic charges by group.
Group Likely Ion
1 +1
(e.g. Na+)
2 +2
(e.g. Mg2+)
3 +3
(e.g. Al3+)
5 −3
(e.g. N3−)
6 −2
(e.g. O2−)
7 −1
(e.g. Cl−)

Transition metals can form variable charges, often shown using Roman numerals.

Example Iron (II) = Fe2+ , Iron(III) = Fe3+

Common Ions to Memorise

Ion Formula
Nitrate NO3−
Carbonate CO32−
Sulfate SO42−
Hydroxide OH−
Ammonium NH4+
Zinc Zn2+
Silver Ag+
Hydrogencarbonate HCO3−
Phosphate PO43−

We can use these charges to write correct formulas for ionic compounds by balancing charges so the total charge is zero.

Examples
Ca2+ and OH− would combine to form Ca(OH)2
Al3+ and SO42− would combine to form Al2(SO4)3

Writing Balanced Equations (Including Ionic Equations)

Full and Ionic Equations

Start by writing the balanced chemical equation.

Then write the ionic equation, removing spectator ions (ions that don’t change).

CIE A-Level Chemistry diagram contrasting a full equation with the simplified ionic equation by removing spectator ions.

Full equation:
NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)

Ionic equation:
OH−(aq) + H+(aq) → H2O(l)

State Symbols

Always include state symbols when writing ionic equations:

Empirical and Molecular Formulas

The empirical formula shows the simplest whole number ratio of atoms in a compound and the molecular formula represents the actual number of atoms in a molecule.

For Example: Glucose (C6H12O6)
Empirical formula = CH2O (simplest ratio 1:2:1)
Molecular formula = C6H12O6 (actual composition)

Water of Crystallisation

Hydrated compounds contain water molecules as part of their crystal structure. The water molecules are trapped between particles in the solid.

Water of crystallisation refers to this water, shown in the formula as “•xH2O". Where x is the moles of water in the solid compared to moles of compound.

For Example:
A common hydrated form of copper (II) sulfate is CuSO4•5H2O (s)
This formula tells us that for every one mole of CuSO4 in the solid crystal, there are also 5 moles of H2O molecules.

CIE A-Level Chemistry diagram illustrating CuSO4·5H2O as a hydrated crystal with water of crystallisation.

Anhydrous compounds contain no water.

For Example:
CuSO4(s) is anhydrous as the formula contains no water of crystallisation.

CIE A-Level Chemistry diagram showing anhydrous copper(II) sulfate CuSO4 with no water of crystallisation.

Calculating Empirical and Molecular Formulas

To find the empirical formula from mass or percentage composition:

Worked Example Calculation

Find the empirical for the compound with a composition by mass of C 52.2%, H 13.0% and O 34.8%

CIE A-Level Chemistry worked example showing C 52.2%, H 13.0%, O 34.8% converted to moles and ratios to give empirical formula C2H6O.

Determining the Molecular Formula

To find the molecular formula, use:

n = Molecular mass / Empirical mass

Multiply the empirical formula by n to get the molecular formula.

Worked Example Calculation

The empirical formula of a compound is CH2O, and its molar mass is 180 g/mol. Find the molecular formula.

Calculate the empirical formula mass:

Find n:

Multiply empirical formula by n:

CIE A-Level Chemistry worked example showing CH2O empirical mass of 30.0, n = 180/30 = 6, giving molecular formula C6H12O6.

Summary