Chemistry - Moles, Bonding and Reactions
The chemistry paper covers atomic structure and electron configuration, periodic trends, chemical bonding, the mole concept and stoichiometry, solutions and concentration, acids, bases and pH, electrochemistry, and the organic families (alkanes, alkenes, alcohols, carboxylic acids). Numerical questions dominate, and almost all of them pass through the mole.
Core rules
- Mole relationships: n = m/M; n = V/22.4 for a gas at STP with V in dm³; n = C × V for a solution with V in dm³.
- One mole contains 6.02 × 10²³ particles (Avogadro's constant).
- An equation must be balanced before its coefficients can be used as mole ratios.
- pH = −log[H⁺]; a strong acid is assumed to ionise completely.
- Periodic trends: atomic radius decreases across a period and increases down a group; electronegativity increases across a period.
- Bonding: ionic involves transfer (metal + non-metal); covalent involves sharing; metallic involves delocalised electrons.
Worked example
How many molecules and how many grams of carbon are in 4.4 g of CO₂? (M: C = 12, O = 16; N_A = 6.02 × 10²³) Reasoning: M(CO₂) = 12 + 2(16) = 44 g/mol → n = 4.4/44 = 0.10 mol. Molecules = 0.10 × 6.02 × 10²³ = 6.02 × 10²². Each formula unit contains one carbon atom, so carbon mass = 0.10 × 12 = 1.2 g. Note the chain: mass → moles → particle count, and mass → moles → mass of an element. Every stoichiometry question follows this pattern.
Traps that cost marks
- Using mole ratios from an unbalanced equation.
- Applying 22.4 to a volume given in cm³ instead of dm³.
- Using the atomic mass of O for O₂, or of H for H₂.
- Confusing empirical formula (simplest ratio) with molecular formula (actual atoms).
- Sign and ordering errors when assigning oxidation numbers.