QCE Chemistry cheat sheet

Empirical formula from percentage composition: divide each percentage by the molar mass of that element, then divide all results by the smallest value to obtain the simplest whole-number ratio.

Molecular formula = (empirical formula)n, where n = molar mass of compound ÷ mass of one empirical formula unit.

n = m ÷ M; c = n ÷ V (V in litres); for gases at STP, n = V ÷ 22.71 L mol⁻¹.

Percentage yield = (actual yield ÷ theoretical yield) × 100; atom economy = (mass of desired product ÷ total mass of reactants) × 100.

Limiting reagent: calculate n for each reactant, divide by its stoichiometric coefficient; the smallest quotient identifies the limiting reagent.

Ionisation energy increases across a period (greater effective nuclear charge, smaller radius) and decreases down a group (outer electrons further from nucleus, more shielding).

Atomic radius decreases across a period and increases down a group; ionic radius of a cation is smaller than its parent atom, an anion larger.

Bond type follows an electronegativity difference guide: ΔEN < 0.4 non-polar covalent, 0.4–1.7 polar covalent, > 1.7 predominantly ionic (a guide, not a strict rule).

Metals have high melting points, are malleable and conduct electricity due to delocalised valence electrons; ionic solids are brittle, high-melting and conduct only when molten or dissolved.

Intermolecular forces ranked by strength: ion–dipole > hydrogen bonding > dipole–dipole > dispersion; boiling point rises with the strongest force present and with molecular mass/shape.

Hydrogen bonding requires H bonded directly to N, O or F, plus a lone pair on a nearby N, O or F atom.

Dispersion forces increase with the number of electrons (larger molar mass) and with greater surface contact (linear molecules have more contact than compact branched isomers).

Like dissolves like: polar and ionic solutes dissolve in polar solvents; non-polar solutes dissolve in non-polar solvents, driven by entropy and solute–solvent interactions.

Rate of reaction factors: concentration (collision frequency), temperature (fraction of collisions exceeding Ea), surface area, catalyst (lowers Ea by providing an alternative pathway).

Rate = Δ[product] ÷ Δt or −Δ[reactant] ÷ Δt; average rate decreases over time as reactant concentration falls.

Collision theory requires collisions with sufficient energy (≥ activation energy) and correct orientation; only a small fraction of collisions are effective.

Enthalpy: ΔH = Σ bond energies broken − Σ bond energies formed; exothermic reactions have ΔH < 0 and release energy to surroundings.

Calorimetry: q = mcΔT, where c(water) = 4.18 J g⁻¹ K⁻¹; ΔH = −q ÷ n for the amount of substance reacting.

Equilibrium is dynamic: forward and reverse rates are equal and concentrations remain constant, but both reactions continue; it requires a closed system.

Kc for aA + bB ⇌ cC + dD equals ([C]^c [D]^d) ÷ ([A]^a [B]^b); pure solids and liquids are excluded.

Le Chatelier's principle: a system at equilibrium shifts to partially oppose a change in concentration, pressure, volume or temperature.

A catalyst does not change the equilibrium position or K; it only reduces the time taken to reach equilibrium.

For an endothermic reaction, increasing temperature increases K; for an exothermic reaction, increasing temperature decreases K.

pH = −log₁₀[H⁺]; [H⁺] = 10^(−pH); Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C; pH + pOH = 14.

Strong acids ionise completely (HCl, HNO₃, H₂SO₄); weak acids ionise partially (CH₃COOH) and have Ka expressions describing their extent of ionisation.

Neutralisation stoichiometry: n(H⁺) = n(OH⁻) at equivalence; c₁V₁ ÷ a = c₂V₂ ÷ b for an acid–base reaction with coefficients a and b.

Oxidation is loss of electrons (increase in oxidation number); reduction is gain of electrons; the oxidant is reduced and the reductant is oxidised.

Balance redox half-equations by adding H₂O to balance O and H⁺ to balance H in acidic solution, then equalise electron numbers before combining.

IUPAC naming: identify the longest carbon chain, number to give substituents the lowest locants, and name the principal functional group with its suffix.

Homologous series prefixes: meth-, eth-, prop-, but-, pent-, hex-; suffixes -ane (alkane), -ene (alkene), -yne (alkyne), -ol (alcohol), -oic acid (carboxylic acid).

Isomers share a molecular formula but differ in structure: chain (skeletal), position, functional group and geometric (cis–trans) isomerism.

Alkenes undergo addition reactions (with H₂, Br₂, HCl, H₂O) because the C=C double bond can open; alkanes undergo substitution with halogens under UV light.

Alcohols oxidise to aldehydes then carboxylic acids (primary), to ketones (secondary), and resist mild oxidation (tertiary).

Esters form from a carboxylic acid plus an alcohol with an acid catalyst (condensation, releasing water) and hydrolyse back in acid or base.

Addition polymers form from monomers containing C=C; condensation polymers form with loss of a small molecule such as water, and both can be recycled or redesigned for degradability.

Green chemistry design principles include atom economy, use of renewable feedstocks, fewer hazardous solvents, and design for degradation or recycling.

Infrared spectroscopy: O–H broad 3200–3600 cm⁻¹, C=O strong 1680–1750 cm⁻¹, C–H 2850–3100 cm⁻¹; used to identify functional groups.