GCE A/L Chemistry cheat sheet

Atomic number Z = number of protons; mass number A = protons + neutrons; isotopes have same Z, different A.

Electron configuration: fill orbitals in order of increasing (n+l); for equal (n+l), lower n first. Aufbau order: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p.

Hund's rule: within a subshell, electrons occupy degenerate orbitals singly with parallel spins before pairing.

Pauli exclusion principle: no two electrons in an atom can have the same four quantum numbers.

Periodic trends: atomic radius decreases across a period (increasing Zeff) and increases down a group (added shells).

First ionisation energy generally increases across a period and decreases down a group; exceptions at group 2→13 (e.g., Mg→Al) and 15→16 (e.g., P→S) due to subshell/electron-pairing effects.

Electron affinity: energy released when an electron is added to a gaseous atom; generally more negative across a period, less negative down a group (F exception vs Cl).

Electronegativity (Pauling) increases across a period and decreases down a group; F is the most electronegative element.

Successive ionisation energies show large jumps when a new shell is breached, revealing group number.

Ionic bonding: electrostatic attraction between oppositely charged ions; lattice energy magnitude increases with higher charges and smaller ionic radii (Born–Landé).

Covalent bond: shared electron pair; bond order = (bonding electrons − antibonding electrons)/2 in MO theory.

VSEPR: electron domains (bonding + lone pairs) arrange to minimise repulsion; lone-pair repulsion > bond-pair repulsion, reducing bond angles (e.g., CH4 109.5°, NH3 107°, H2O 104.5°).

Hybridisation: steric number 2 → sp (linear), 3 → sp2 (trigonal planar), 4 → sp3 (tetrahedral); multiple bonds count as one domain.

Polarity: a molecule is polar if it has polar bonds and an asymmetric arrangement so bond dipoles do not cancel (e.g., CO2 nonpolar, H2O polar).

Intermolecular forces strength: ion–dipole > hydrogen bonding > dipole–dipole > London dispersion (for similar molar mass).

Hydrogen bonding occurs when H is bonded to N, O or F and interacts with a lone pair on N, O or F of another molecule; explains high boiling points of H2O, HF, NH3.

London dispersion forces increase with number of electrons (polarisability) and molecular surface area; branched isomers have weaker dispersion than straight-chain isomers.

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

Mole: n = m/M; Avogadro constant NA = 6.022 × 10^23 mol^−1; number of particles = n × NA.

Ideal gas equation: pV = nRT; R = 8.314 J K^−1 mol^−1; use p in Pa, V in m^3, T in K.

At STP (273.15 K, 1.00 × 10^5 Pa), molar volume of an ideal gas ≈ 22.7 dm^3 mol^−1; at 273 K and 1 atm ≈ 22.4 dm^3 mol^−1.

Dalton's law: total pressure = sum of partial pressures; partial pressure = mole fraction × total pressure.

Kinetic theory assumptions: negligible particle volume, no intermolecular forces, random motion, elastic collisions, average kinetic energy ∝ T (Kelvin).

Root-mean-square speed: v_rms = √(3RT/M); lighter gases diffuse faster (Graham's law: rate ∝ 1/√M).

Real gases deviate from ideality at high pressure and low temperature; van der Waals equation: (p + an^2/V^2)(V − nb) = nRT.

Empirical formula from percentage composition: divide by atomic masses, then divide by smallest ratio; molecular formula = empirical formula × n where n = M_r / empirical formula mass.

Concentration: c = n/V (mol dm^−3); dilution: c1V1 = c2V2.

Enthalpy change ΔH = H_products − H_reactants; exothermic ΔH < 0, endothermic ΔH > 0.

Standard conditions: 298 K, 1 bar (10^5 Pa), substances in standard states; standard enthalpy of formation ΔHf° of an element in its standard state is zero.

Hess's law: ΔH for a reaction is independent of route; ΔH°_rxn = ΣΔHf°(products) − ΣΔHf°(reactants).

Bond enthalpy: ΔH°_rxn = Σ(bonds broken) − Σ(bonds formed); mean bond enthalpies give approximate values.

Calorimetry: q = mcΔT; for a reaction in solution, ΔH = −q/n (n = moles of limiting reactant).

Coffee-cup calorimeter measures heat at constant pressure (ΔH); bomb calorimeter measures heat at constant volume (ΔU).

Lattice enthalpy: enthalpy change when one mole of an ionic solid forms from gaseous ions; Born–Haber cycle combines ionisation energy, electron affinity, atomisation and lattice enthalpy.

Entropy S increases with temperature and number of moles of gas; ΔS°_rxn = ΣS°(products) − ΣS°(reactants).

Gibbs free energy: ΔG° = ΔH° − TΔS°; reaction feasible when ΔG° < 0.

Equilibrium constant Kc = [products]^coefficients / [reactants]^coefficients (activities of pure solids/liquids = 1); Kp uses partial pressures.

Le Chatelier's principle: a system at equilibrium shifts to oppose a change in concentration, pressure or temperature; catalysts do not shift equilibrium.

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

Kw = [H+][OH−] = 1.0 × 10^−14 at 298 K; pH = −log[H+]; pOH = −log[OH−]; pH + pOH = 14 at 298 K.

Strong acid/base: fully dissociated; weak acid: Ka = [H+][A−]/[HA]; for a weak acid, [H+] ≈ √(Ka × c) when dissociation is small.

pKa = −log Ka; Henderson–Hasselbalch: pH = pKa + log([A−]/[HA]) for a buffer.

Buffer capacity is greatest when [A−] ≈ [HA] (pH ≈ pKa) and when concentrations are high.

Titration: at equivalence moles of acid = moles of base (adjusted for stoichiometry); pH at equivalence depends on salt hydrolysis (strong acid–weak base → acidic).

Solubility product Ksp = [cation]^a[anion]^b for a sparingly soluble salt; precipitation occurs when ionic product > Ksp.

Rate = −d[reactant]/dt = +d[product]/dt; rate law is determined experimentally, not from stoichiometry.

Order: rate = k[A]^m[B]^n; overall order = m + n; zero order → rate independent of concentration; first order → constant half-life t½ = ln2/k.

Arrhenius equation: k = A e^(−Ea/RT); ln(k2/k1) = (Ea/R)(1/T1 − 1/T2).

Activation energy Ea is the minimum energy needed for a collision to lead to reaction; catalysts lower Ea without changing ΔH.

Collision theory: rate depends on collision frequency, orientation and fraction of collisions with energy ≥ Ea (Boltzmann distribution).

Rate-determining step is the slowest step in a mechanism; the rate law reflects the species involved up to and including that step.

Catalysis: homogeneous (same phase) and heterogeneous (different phase, surface adsorption); enzymes are biological catalysts with high specificity.

Clock reactions and initial-rate methods (e.g., measuring gas volume, colour change, conductivity) are used to determine rate laws experimentally.

For first-order reactions, concentration decays exponentially: [A] = [A]0 e^(−kt); ln[A] = ln[A]0 − kt.

Oxidation is loss of electrons (OIL), reduction is gain (RIG); oxidising agent is reduced, reducing agent is oxidised.

Oxidation number rules: free element 0; monatomic ion = charge; O usually −2 (except peroxides −1, OF2 +2); H usually +1 (except metal hydrides −1); sum in neutral compound = 0, in polyatomic ion = charge.

Galvanic cell: anode is negative (oxidation), cathode is positive (reduction); electrons flow anode → cathode in external circuit.

Standard electrode potential E° is measured vs standard hydrogen electrode (SHE, E° = 0.00 V) at 298 K, 1 M, 1 bar.

Cell emf E°_cell = E°_cathode − E°_anode; positive E°_cell indicates a spontaneous reaction under standard conditions.

ΔG° = −nFE°_cell; F = 96 485 C mol^−1; n = number of moles of electrons transferred.

Nernst equation: E = E° − (RT/nF) ln Q; at 298 K, E = E° − (0.0592/n) log Q.

Electrolysis: oxidation at anode, reduction at cathode; products depend on relative electrode potentials and concentration (e.g., dilute vs concentrated NaCl).

Faraday's laws: mass deposited ∝ charge; charge Q = It; moles of electrons = Q/F.

IUPAC naming: identify longest carbon chain containing the principal functional group; number to give the principal group the lowest locant; name substituents alphabetically.

Homologous series: alkanes CnH2n+2, alkenes CnH2n, alkynes CnH2n−2, alcohols CnH2n+1OH, carboxylic acids CnH2n+1COOH.

Functional groups: hydroxyl (–OH), carbonyl (C=O), carboxyl (–COOH), ester (–COO–), amine (–NH2), amide (–CONH2), halogeno (–X), nitrile (–CN).

Isomerism: structural (chain, position, functional group) and stereoisomerism (geometric cis/trans, optical with chiral carbon bearing four different groups).

Alkenes show geometric isomerism when each carbon of the C=C has two different groups; restricted rotation about C=C.

Chirality: a carbon atom bonded to four different groups gives non-superimposable mirror images (enantiomers) that rotate plane-polarised light.

Primary, secondary, tertiary classification for alcohols and halogenoalkanes depends on the number of carbon atoms attached to the carbon bearing the functional group.

Aromaticity (Hückel): cyclic, planar, fully conjugated, with 4n + 2 π electrons (e.g., benzene, n = 1, six π electrons).

Boiling points within a homologous series increase with chain length (stronger London dispersion); branching lowers boiling point.

Nucleophilic substitution: SN1 (two steps, carbocation intermediate, racemisation, tertiary favoured) vs SN2 (one step, backside attack, inversion, primary favoured).

Elimination: E1 (carbocation intermediate) and E2 (concerted, anti-periplanar, strong base); competes with substitution.

Electrophilic addition to alkenes: Markovnikov's rule - H adds to the carbon with more hydrogens; carbocation stability tertiary > secondary > primary.

Free-radical substitution of alkanes: initiation (homolytic fission), propagation, termination; requires UV light.

Electrophilic aromatic substitution (nitration, halogenation) proceeds via a resonance-stabilised arenium ion; substituents direct ortho/para (activating) or meta (deactivating).

Oxidation: primary alcohols → aldehydes → carboxylic acids; secondary alcohols → ketones; tertiary alcohols resist oxidation. Reagents: acidified K2Cr2O7 or KMnO4.

Reduction: aldehydes/ketones → alcohols (NaBH4 or LiAlH4); carboxylic acids → primary alcohols (LiAlH4); nitrobenzene → phenylamine (Sn/HCl).

Esterification: carboxylic acid + alcohol ⇌ ester + water, acid catalyst, reversible; hydrolysis of esters under acid or base conditions.

Synthesis routes: convert functional groups (e.g., halogenoalkane → alcohol via NaOH(aq), → nitrile via KCN, → amine via excess NH3); use protecting groups when necessary.

Haber process: N2 + 3H2 ⇌ 2NH3, exothermic; conditions ~450 °C, 200 atm, iron catalyst; compromise between rate and equilibrium yield.

Contact process: SO2 + ½O2 ⇌ SO3, vanadium(V) oxide catalyst, ~450 °C, 1–2 atm; SO3 absorbed in concentrated H2SO4.

Ostwald process: NH3 oxidised to NO (Pt/Rh catalyst), NO to NO2, then NO2 + H2O → HNO3.

Greenhouse gases: CO2, CH4, N2O, water vapour, CFCs; absorb infrared radiation and re-emit it, warming the atmosphere.

Acid rain: SO2 and NOx emissions form H2SO4 and HNO3; mitigation via flue-gas desulfurisation (CaO/CaCO3) and catalytic converters.

Ozone depletion: CFCs release Cl radicals in the stratosphere which catalytically destroy O3; Montreal Protocol phased out CFCs.

Water treatment: sedimentation, filtration, chlorination (disinfection), and measurement of dissolved oxygen (BOD/COD) to assess pollution.

Polymers: addition polymers (polyethene, PVC) from alkenes; condensation polymers (nylon, PET) with loss of small molecule; recycling codes and biodegradability.

Fertilisers: NPK (nitrogen, phosphorus, potassium); eutrophication from excess nitrate/phosphate runoff.

Titration procedure: rinse burette with solution, read meniscus at eye level, use a white tile, add indicator; concordant titres within 0.10 cm^3.

Standard solution: accurately weighed primary standard (e.g., anhydrous Na2CO3, potassium hydrogen phthalate) dissolved and made up to a known volume in a volumetric flask.

Acid–base indicators: phenolphthalein (colourless → pink, pH 8.2–10.0) for strong base–weak acid; methyl orange (red → yellow, pH 3.1–4.4) for strong acid–weak base.

Qualitative cation tests: flame tests (Li crimson, Na yellow, K lilac, Ca brick-red, Cu blue-green); NaOH(aq) precipitates (Cu(OH)2 blue, Fe(OH)3 red-brown, Fe(OH)2 green, Al(OH)3 white soluble in excess).

Anion tests: carbonate (effervescence with acid, limewater milky), sulfate (white BaSO4 precipitate with BaCl2 then acid), halide (AgCl white, AgBr cream, AgI yellow; soluble in NH3 in order).

Gravimetric analysis: precipitate, filter, wash, dry to constant mass, weigh; calculate moles from formula mass.

Colorimetry: absorbance ∝ concentration (Beer–Lambert law); use a calibration curve of known standards.

Chromatography: separation by differential adsorption/partition; Rf = distance travelled by spot / distance travelled by solvent front.

Spectroscopy: infrared identifies functional groups (e.g., C=O ~1700 cm^−1, O–H broad ~2500–3300 cm^−1); mass spectrometry gives molecular ion and fragmentation; NMR gives environment of H/C atoms.