Chemistry: Mole Logic, Structure Rules, and Common Errors
What the chemistry section tests
Physical chemistry (mole concept, stoichiometry, gaseous state, thermodynamics, equilibrium, ionic equilibria, electrochemistry, kinetics), inorganic chemistry (periodic trends, chemical bonding, coordination compounds, p-block and d-block facts), and organic chemistry (hybridisation, isomerism, reaction mechanisms, named reactions, biomolecules). Expect numericals on moles and pH alongside structure-based organic questions.
Core rules
- One mole = 6.022×10²³ particles = molar mass in grams = 22.4 L of an ideal gas at STP.
- The limiting reagent decides product yield; compare mole ratios, never masses.
- pH = −log[H⁺]. Strong acids are treated as fully dissociated; weak acids need Ka.
- Hybridisation from steric number: 2 → sp, 3 → sp², 4 → sp³.
- Le Chatelier: raising pressure favours the side with fewer gas moles.
- Optical isomerism requires a chiral centre with four different groups; geometrical isomerism requires restricted rotation.
Worked example
How many grams of water form when 4 g of H₂ reacts with 32 g of O₂? Equation: 2H₂ + O₂ → 2H₂O. Moles: H₂ = 4/2 = 2 mol; O₂ = 32/32 = 1 mol. The required ratio is 2:1 and the available ratio is 2:1, so neither reagent is limiting. Water formed = 2 mol × 18 g/mol = 36 g.
Traps that cost marks
- Taking mole ratios from an unbalanced equation.
- Treating 22.4 L as the molar volume at any temperature.
- Confusing empirical and molecular formulae.
- Writing pH of a weak acid as −log(concentration) without using Ka.
- Forgetting that enantiomers rotate plane-polarised light in opposite directions.
- Counting a carbon as chiral when two attached groups are identical.