NCEA L2 Chemistry cheat sheet

Mole calculations: n = m/M, n = cV (V in L), n = V/22.4 L for gases at STP; M is molar mass in g mol⁻¹.

Avogadro's number N_A = 6.02 × 10²³ particles per mole; number of particles = n × N_A.

Concentration: c = n/V; dilution obeys c₁V₁ = c₂V₂ because moles of solute are conserved.

Percentage yield = (actual mass ÷ theoretical mass) × 100; theoretical mass comes from the limiting reagent.

Limiting reagent: divide moles of each reactant by its stoichiometric coefficient; the smallest quotient limits the reaction.

Empirical formula from percentage composition: divide each % by atomic mass, then divide by the smallest ratio; molecular formula = empirical formula × integer from M.

Ionic equations include only species that change; spectator ions are cancelled. Solubility rules: most nitrates, sodium, potassium and ammonium salts are soluble; most chlorides are soluble except AgCl and PbCl₂; most sulfates are soluble except BaSO₄, PbSO₄ and CaSO₄ (slightly).

Precipitation occurs when the product of ion concentrations exceeds Ks; a precipitate is identified by its colour and its solubility behaviour in excess reagent.

Ionic bonding: electrons transferred from metal to non-metal; lattice held by electrostatic attraction between cations and anions; high melting point, conducts when molten or dissolved.

Covalent molecular substances: discrete molecules with weak intermolecular forces; low melting point, do not conduct. Covalent network substances (diamond, SiO₂) have high melting points.

Metallic bonding: positive ions in a sea of delocalised electrons; explains conductivity, malleability and ductility.

Electronegativity difference guides bond type: ΔEN < 0.4 non-polar covalent, 0.4–1.7 polar covalent, > 1.7 largely ionic (guideline only).

Intermolecular forces increase in the order dispersion < dipole–dipole < hydrogen bonding; H-bonding requires H bonded to N, O or F.

Polarity: a molecule with polar bonds can be non-polar if bond dipoles cancel by symmetry (e.g. CO₂, CCl₄).

Shapes from VSEPR: 2 regions linear (180°), 3 trigonal planar (120°), 4 tetrahedral (109.5°), 4 with one lone pair trigonal pyramidal (~107°), 4 with two lone pairs bent (~104.5°).

Lewis structures: count valence electrons, form single bonds, complete octets, then place remaining electrons as lone pairs on the central atom; formal charge = valence electrons − (non-bonding electrons + ½ bonding electrons).

Types of reaction: combination, decomposition, displacement, precipitation, neutralisation (acid + base → salt + water) and redox.

Acid–base: acid donates H⁺, base accepts H⁺; strong acids (HCl, HNO₃, H₂SO₄) ionise fully, weak acids (CH₃COOH) ionise partially. pH = −log₁₀[H₃O⁺]; [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C.

Neutralisation titration: n(acid) × coefficient ratio = n(base); c_acid V_acid / coefficient = c_base V_base / coefficient.

Redox: oxidation is loss of electrons (OIL), reduction is gain (RIG); oxidant is reduced, reductant is oxidised. Oxidation number rules: O = −2 (except peroxides), H = +1 (except hydrides), free element = 0, sum equals overall charge.

Balancing redox in acidic solution: split into half-equations, balance atoms, balance O with H₂O, balance H with H⁺, balance charge with electrons, multiply to equal electrons, then add and cancel.

Enthalpy: ΔH is negative for exothermic, positive for endothermic; ΔH = Σ bonds broken − Σ bonds formed; bond breaking absorbs energy, bond making releases energy.

Calorimetry: q = mcΔT; ΔH = −q/n (kJ mol⁻¹), with m in grams, c = 4.18 J g⁻¹ °C⁻¹ for aqueous solutions.

Activation energy is the minimum energy for a successful collision; a catalyst lowers E_a by providing an alternative pathway without being consumed.

Rate increases with higher concentration, higher temperature, larger surface area and catalysts; temperature raises the proportion of collisions exceeding E_a, not just collision frequency.

Equilibrium is dynamic: forward and reverse rates are equal and concentrations remain constant; K_c = [products]^coefficients ÷ [reactants]^coefficients, pure solids and liquids omitted.

Le Chatelier's principle: a system at equilibrium shifts to oppose a change in concentration, pressure or temperature; a catalyst changes only the rate of approach, not the equilibrium position.

Organic homologous series: alkanes CₙH₂ₙ₊₂ (single bonds), alkenes CₙH₂ₙ (C=C), alkynes CₙH₂ₙ₋₂ (C≡C), alcohols CₙH₂ₙ₊₁OH, carboxylic acids CₙH₂ₙ₊₁COOH.

IUPAC naming: longest chain gives the parent, number from the end nearest the functional group, name substituents alphabetically with locants (e.g. 2-methylbutan-1-ol).

Isomerism: structural isomers differ in connectivity (chain, position, functional group); geometric (cis–trans) isomers require restricted rotation (C=C) and two different groups on each carbon.

Alkene addition reactions: with H₂/Ni gives alkane, with HCl gives haloalkane, with Br₂ water decolourises (test for unsaturation), with H₂O/H⁺ gives alcohol.

Functional group tests: bromine water decolourised by alkenes; acidified dichromate turns orange to green with primary/secondary alcohols; limewater turns milky with CO₂; carboxylic acids turn blue litmus red and react with carbonates releasing CO₂.