Cameroon GCE A-Level Chemistry cheat sheet

Atomic number Z = protons; mass number A = protons + neutrons; isotopes share Z but differ in neutron count.

Ionisation energy increases across a period (greater effective nuclear charge, smaller radius) and decreases down a group (outer electrons farther out and shielded).

Successive ionisation energies show large jumps at the point where a new quantum shell is breached; jump position reveals group number.

Electron configuration follows Aufbau order 1s 2s 2p 3s 3p 4s 3d 4p; 4s fills before 3d and empties before 3d in cations.

Period 3 trends: atomic radius decreases, electronegativity increases, metallic character decreases left to right.

Group trends: down Group 1 ionisation energy and electronegativity fall; down Group 17 oxidising power falls, boiling point rises.

Anomalies: Cr is [Ar]3d5 4s1 and Cu is [Ar]3d10 4s1 because half-filled and full d subshells are more stable.

First electron affinity is usually exothermic; second electron affinity is endothermic because the anion repels the incoming electron.

Ionic: lattice energy magnitude rises with higher ionic charge and smaller ionic radius (Born–Landé proportionality to q+ q− / r).

Covalent bond strength: bond order increases bond enthalpy (C≡C > C=C > C–C); bond length shortens as bond order rises.

VSEPR: electron pairs repel in order lone pair–lone pair > lone pair–bond pair > bond pair–bond pair; CH4 109.5°, NH3 107°, H2O 104.5°.

Shapes: 2 linear, 3 trigonal planar, 4 tetrahedral, 5 trigonal bipyramidal, 6 octahedral; lone pairs reduce angle by about 2.5° each.

Polarity: a molecule is polar only if bond dipoles do not cancel; CO2 non-polar (linear), H2O polar (bent), CCl4 non-polar (tetrahedral).

Intermolecular forces increase: London dispersion < dipole–dipole < hydrogen bonding; H-bonding needs H bonded to N, O or F.

Giant covalent (diamond, SiO2) melts high; simple molecular melts low; metals melt high due to delocalised electron attraction.

Graphite conducts because each carbon uses three bonds, leaving a delocalised p electron; diamond does not because all four valence electrons are bonded.

Phase changes: melting and boiling are endothermic; freezing and condensing are exothermic; enthalpy of sublimation = fusion + vaporisation.

Heating curve plateaus occur at melting and boiling points where added heat changes potential, not kinetic, energy.

Vapour pressure rises with temperature; a liquid boils when its vapour pressure equals external pressure.

Boiling point rises with stronger intermolecular forces and larger molar mass (more electrons, stronger London forces).

Cooling curves and solubility data locate the eutectic; impurities lower melting point and broaden the melting range.

Solid types: ionic (high m.p., conducts when molten), metallic, giant covalent, simple molecular, and amorphous.

Enthalpy of vaporisation is always larger than enthalpy of fusion for the same substance because all intermolecular forces must be broken.

Ideal gas assumptions fail near condensation: particle volume and intermolecular attraction become significant.

Ideal gas equation: pV = nRT with R = 8.314 J K⁻¹ mol⁻¹; T must be in kelvin (K = °C + 273).

Combined gas law: p1V1/T1 = p2V2/T2 for a fixed mass of gas.

Kinetic theory: average kinetic energy ∝ T (KE = 3/2 kT per molecule); pV = 1/3 N m c²rms.

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

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

Molar volume at s.t.p. (273 K, 1 atm) is 22.4 dm³ mol⁻¹; at r.t.p. (298 K, 1 atm) about 24.0 dm³ mol⁻¹.

Real gases deviate most at high pressure and low temperature, where attractions and molecular volume matter.

Avogadro's law: equal volumes of gases at the same T and p contain equal numbers of molecules.

Hess's law: total enthalpy change is independent of route; use cycle sums so ΔH_reaction = ΣΔHf(products) − ΣΔHf(reactants).

Bond enthalpy route: ΔH = Σ bonds broken − Σ bonds formed (both positive magnitudes).

Exothermic reactions have negative ΔH; endothermic have positive ΔH; activation energy is always positive.

Calorimetry: q = mcΔT; for a burner, ΔH = −q/n where n is moles burned; account for heat loss by comparison with data-book values.

Standard conditions for ΔHf and ΔHc: 298 K, 1 atm (100 kPa), substances in standard states; ΔHf of an element in its standard state is zero.

Mean bond enthalpies give approximate ΔH because values are averaged across different molecules.

Lattice enthalpy is exothermic to form; dissolving involves lattice enthalpy (endothermic) plus hydration enthalpy (exothermic).

Entropy increases with more gas moles, higher temperature, and greater disorder; ΔG = ΔH − TΔS, spontaneous when ΔG < 0.

Kc expression: for aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d / ([A]^a[B]^b); pure solids and liquids are omitted.

Kp uses partial pressures; p_i = mole fraction × total pressure.

Le Chatelier: raising pressure shifts equilibrium toward fewer gas moles; raising temperature shifts in the endothermic direction.

A catalyst speeds attainment of equilibrium but does not change K or the equilibrium position.

K depends only on temperature; concentration and pressure changes shift position but leave K unchanged.

If Q < K the forward reaction is favoured; if Q > K the reverse reaction is favoured; Q = K means equilibrium.

Haber process: N2 + 3H2 ⇌ 2NH3, ΔH negative; high pressure and moderate temperature (about 450 °C) with iron catalyst.

Equilibrium constant magnitude indicates extent: K >> 1 favours products, K << 1 favours reactants.

Rate = change in concentration per unit time; average rate uses the gradient of a concentration–time graph.

Rate equation: rate = k[A]^m[B]^n; orders are found experimentally, not from stoichiometry.

First order has constant half-life; zero order has constant rate independent of concentration; second order rate ∝ [A]².

Arrhenius equation: k = A e^(−Ea/RT); ln k = ln A − Ea/RT, so a plot of ln k against 1/T has gradient −Ea/R.

Collision theory: reaction needs collisions with energy ≥ Ea and correct orientation; catalysts lower Ea.

Oxidation is loss of electrons (increase in oxidation number); reduction is gain of electrons.

Oxidation number rules: O usually −2, H usually +1, sum equals overall charge; in oxoanions like MnO4⁻ manganese is +7.

Redox balancing: balance atoms, then oxygen with H2O, hydrogen with H⁺ (acidic) or OH⁻ (basic), then charge with electrons.

pH = −log[H⁺]; [H⁺] = 10^(−pH); pOH = 14 − pH at 25 °C.

Strong acid: [H⁺] = concentration × basicity (e.g. H2SO4 gives 2H⁺); weak acid uses Ka = [H⁺]²/[HA].

Ka × Kb = Kw = 1.0 × 10⁻¹⁴ at 25 °C; pKa + pKb = 14.

Buffer (weak acid + conjugate base): pH = pKa + log([salt]/[acid]); Henderson–Hasselbalch.

Buffer capacity is greatest when [salt] ≈ [acid], i.e. pH ≈ pKa; buffers resist pH change by neutralising added acid or base.

Titration: strong acid–strong base equivalence pH = 7; weak acid–strong base equivalence pH > 7; weak base–strong acid equivalence pH < 7.

Indicator choice: indicator pKa should lie within the steep part of the titration curve (e.g. phenolphthalein 8.3–10.0).

Salt hydrolysis: salt of weak acid and strong base gives alkaline solution; salt of strong acid and weak base gives acidic solution.

Group 1 metals react with water giving hydroxide and hydrogen; reactivity increases down the group.

Group 2: oxides are basic, hydroxides are alkaline; thermal stability of carbonates and nitrates increases down the group.

Solubility of Group 2 sulfates decreases down the group; hydroxide solubility increases down the group.

Group 17: oxidising power decreases down the group (Cl2 > Br2 > I2); displacement reactions confirm the order.

Hydrogen halides dissolve to give strong acids; HF is weak due to strong H–F bond.

Group 4/14: carbon forms CO2 (acidic oxide); silicon forms SiO2 (giant covalent, acidic oxide); lead(IV) oxide is amphoteric.

Nitrogen and phosphorus: NH3 is a weak base; nitric acid is a strong oxidising acid; phosphorus forms P4O10 and H3PO4.

Sulfur: SO2 is a reducing agent and acidic oxide; SO3 with water gives H2SO4; contact process uses V2O5 catalyst.

Transition metals have partially filled d subshells; they show variable oxidation states, coloured compounds, catalytic activity and complex formation.

Ligands donate electron pairs; coordination number is the number of ligand donor atoms (usually 4 or 6).

Complex shapes: tetrahedral (e.g. [CoCl4]²⁻) and octahedral (e.g. [Cu(H2O)6]²⁺); square planar occurs in Pt(II) and Ni(II) complexes.

Colour arises from d–d transitions splitting d orbitals; different ligands split the d orbitals differently, changing the colour.

Ligand substitution: adding excess NH3 or Cl⁻ to hexaaqua complexes replaces water ligands and changes colour.

Fe²⁺ gives a pale green solution; Fe³⁺ gives yellow-brown; adding SCN⁻ to Fe³⁺ gives a blood-red complex.

MnO4⁻ in acidic solution is a strong oxidising agent, reduced from Mn(VII) to Mn(II) (purple to colourless).

Catalytic action: Fe in Haber process, V2O5 in contact process, and transition metals in catalytic converters.

Alkanes: general formula CnH2n+2; substitution with halogens under UV light; combustion gives CO2 and H2O.

Alkenes: CnH2n; electrophilic addition with HBr, Br2 and H2O; bromine water is decolourised (test for unsaturation).

Alkynes: CnH2n−2; undergo addition reactions; terminal alkynes are weakly acidic.

Alcohols: primary oxidise to aldehydes then carboxylic acids; secondary to ketones; tertiary resist oxidation.

Aldehydes reduce Tollens' reagent to a silver mirror and Fehling's solution to a red precipitate; ketones do not.

Carboxylic acids are weak acids; react with alcohols to form esters (Fischer esterification, acid catalyst).

Amines are basic; primary amines react with nitrous acid to give alcohols and nitrogen; amides hydrolyse to acids and amines.

Functional group tests: bromine water (alkene), sodium (alcohol), NaHCO3 (carboxylic acid), 2,4-DNP (carbonyl).

Free-radical substitution: initiation (UV splits Cl2), propagation (Cl• + CH4), termination (radicals combine).

Electrophilic addition to alkenes: alkene attacks electrophile, carbocation intermediate forms; Markovnikov's rule applies.

Nucleophilic substitution: primary haloalkanes favour SN2 (one step, inversion); tertiary favour SN1 (carbocation, racemisation).

Elimination competes with substitution; hot concentrated alcoholic KOH favours elimination, aqueous KOH favours substitution.

Electrophilic substitution of benzene: nitration, halogenation, alkylation; catalyst needed (HNO3/H2SO4 for nitration).

Synthesis planning: convert functional groups in order (e.g. alkene → alcohol → aldehyde → acid), noting reagents and conditions.

Oxidation and reduction in organic synthesis: K2Cr2O7/H2SO4 oxidises alcohols; LiAlH4 reduces aldehydes and ketones.

Polymerisation: addition (ethene → poly(ethene)) and condensation (diol + diacid → polyester, releasing water).

Titration: rinse burette with the titrant, pipette with the solution being measured; read the meniscus at eye level.

Concordant titres agree within 0.10 cm³; average only concordant values, and discard the rough titre.

Moles = concentration (mol dm⁻³) × volume (dm³); convert cm³ to dm³ by dividing by 1000.

Percentage yield = (actual/theoretical) × 100; percentage error = (error/measured value) × 100.

Significant figures: report results to the precision of the least precise measurement; record burette readings to 0.05 cm³.

Qualitative tests: flame tests (Li red, Na yellow, K lilac, Ca brick-red, Cu blue-green); AgNO3 for halides (Cl white, Br cream, I yellow).

Gas tests: H2 pops, O2 relights a glowing splint, CO2 turns limewater milky, NH3 turns damp red litmus blue.

Data analysis: plot the best-fit line, use its gradient, and identify anomalous points rather than forcing them onto the line.