Chemical reactions: nomenclature, oxidation numbers and balancing redox equations
Oxidation numbers, names of inorganic compounds, types of reaction, balancing by inspection and with the half-reaction method, with worked examples.
A chemical reaction rearranges atoms: none is created and none is destroyed, which is why every equation must be balanced. This lesson covers the names of inorganic compounds, the types of reaction, oxidation numbers and balancing, including oxidation-reduction reactions, the basis of batteries and electrochemistry.
Oxidation numbers
The oxidation number is the charge an atom would have if all its bonds were ionic. Rules to assign it:
- in free elements (Fe, O2, Cl2) it is 0;
- in monatomic ions it equals the charge (Na+ +1, Cl− −1);
- oxygen is usually −2 (−1 in peroxides such as H2O2), hydrogen +1 (−1 in metal hydrides);
- the sum is 0 in a neutral molecule and equals the charge in an ion.
Example: in H2SO4 sulphur is x with 2(+1) + x + 4(−2) = 0, so x = +6. In K2Cr2O7 chromium is +6, in MnO4− manganese is +7.
Essential nomenclature
| Class | How it forms | Examples |
|---|---|---|
| Basic oxides | metal + oxygen | Na2O sodium oxide, Fe2O3 iron(III) oxide |
| Acidic oxides | non-metal + oxygen | CO2 carbon dioxide, SO3 sulphur trioxide |
| Hydroxides | basic oxide + water | NaOH sodium hydroxide, Ca(OH)2 |
| Oxoacids | acidic oxide + water | H2SO4 sulphuric acid, HNO3 nitric acid |
| Binary acids | hydrogen + non-metal | HCl hydrochloric acid, HF |
| Salts | acid + base | NaCl sodium chloride, CuSO4 copper(II) sulphate, KNO3 |
Modern IUPAC naming gives the oxidation number in Roman numerals (iron(III) oxide) or uses prefixes (diphosphorus trioxide). The traditional endings are still everywhere: -ous/-ic for acids (sulphurous, sulphuric) and -ite/-ate for the corresponding salts (sulphite, sulphate).
Types of reaction
- Synthesis: 2 H2 + O2 → 2 H2O.
- Decomposition: CaCO3 → CaO + CO2.
- Single displacement: Zn + CuSO4 → ZnSO4 + Cu.
- Double displacement, often with a precipitate: AgNO3 + NaCl → AgCl↓ + NaNO3.
- Neutralisation: HCl + NaOH → NaCl + H2O.
- Combustion: CH4 + 2 O2 → CO2 + 2 H2O.
Balancing an equation
For simple reactions, proceed by inspection: balance first the elements that appear in only one substance on each side, and hydrogen and oxygen last. Example, the combustion of propane: C3H8 + O2 → CO2 + H2O. Three carbons give 3 CO2, eight hydrogens give 4 H2O, so there are 6 + 4 = 10 oxygens on the right and 5 O2 are needed:
C3H8 + 5 O2 → 3 CO2 + 4 H2O
Oxidation-reduction (redox) reactions
A species is oxidised when its oxidation number increases (it loses electrons) and reduced when it decreases (it gains electrons). Both always happen together: the species oxidised is the reducing agent, the one reduced is the oxidising agent. To balance them, use the half-reaction method. Example in acidic solution: permanganate oxidises iron(II).
- Half-reactions: Fe2+ → Fe3+ (oxidation) and MnO4− → Mn2+ (reduction).
- Balance oxygen with H2O and hydrogen with H+: MnO4− + 8 H+ → Mn2+ + 4 H2O.
- Balance the charges with electrons: MnO4− + 8 H+ + 5 e− → Mn2+ + 4 H2O; Fe2+ → Fe3+ + e−.
- Multiply so the electrons match (the second by 5) and add:
MnO4− + 5 Fe2+ + 8 H+ → Mn2+ + 5 Fe3+ + 4 H2O
Final check: the same atoms on both sides and the same total charge (−1 + 10 + 8 = +17 on the left, +2 + 15 = +17 on the right). In basic solution you proceed the same way and at the end neutralise the H+ with OH−.
Redox reactions are everywhere in electronics: etching printed circuit boards with ferric chloride, corrosion of contacts, soldering with a flux that reduces oxides, and above all batteries, where the redox reaction is split into two compartments and the electrons travel through the external circuit.
