NSSCAS Chemistry (8224) cheat sheet

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

Electron configuration: fill order 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p; exceptions: Cr [Ar]3d⁵4s¹, Cu [Ar]3d¹⁰4s¹.

Ionisation energy increases across a period (greater nuclear charge, same shielding) and decreases down a group (extra shells, greater shielding, larger atomic radius).

Successive ionisation energies show large jumps when a new shell is broken; the number of electrons removed before the jump equals the group number.

Atomic radius decreases across a period and increases down a group; ionic radius: cations smaller than parent atom, anions larger.

Period 3 oxides: Na₂O and MgO basic, Al₂O₃ amphoteric, SiO₂ acidic (giant covalent), P₄O₁₀ and SO₃ acidic, Cl₂O₇ acidic.

First ionisation energy definition: energy needed to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.

Electronegativity increases across a period and decreases down a group; Pauling scale, F is most electronegative (4.0).

Ionic bonding: electrostatic attraction between oppositely charged ions; high melting points, brittle, conduct when molten or aqueous, soluble in polar solvents.

Covalent bonding: shared electron pair; giant covalent (diamond, graphite, SiO₂) have high melting points; simple molecular have weak intermolecular forces and low melting points.

Metallic bonding: lattice of positive ions in a sea of delocalised electrons; good conductivity, malleable, ductile.

VSEPR shapes: 2 electron domains linear (180°), 3 trigonal planar (120°), 4 tetrahedral (109.5°), 3+1 lone pair trigonal pyramidal (~107°), 2+2 bent (~104.5°).

Polarity: polar bonds plus asymmetric shape gives polar molecule; symmetrical shapes (CO₂, CCl₄, BF₃) are non-polar despite polar bonds.

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

Giant ionic lattice energy increases with higher charges and smaller ionic radii (Born–Landé); MgO has higher lattice energy than NaCl.

Graphite conducts electricity because each carbon has a delocalised electron in a layer of sp² hybridised atoms; diamond does not (sp³, all electrons localised).

Mole equations: n = m/M; n = V/24 dm³ at room temperature and pressure (r.t.p.); n = concentration × volume (in dm³).

Avogadro constant L = 6.022 × 10²³ mol⁻¹; molar gas volume = 24 dm³ mol⁻¹ at r.t.p. (25 °C, 1 atm).

Empirical formula from percentage composition: divide by atomic mass, divide by smallest ratio, multiply to whole numbers.

Limiting reagent: calculate moles of each reactant, divide by stoichiometric coefficient; smallest quotient is limiting.

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

Titration calculation: moles = concentration (mol dm⁻³) × volume (dm³); use stoichiometric ratio from balanced equation.

Concentration conversions: mol dm⁻³ = g dm⁻³ / molar mass; ppm = mg dm⁻³ for dilute aqueous solutions.

Balanced equations must conserve atoms and charge; state symbols (s, l, g, aq) are required in thermochemical and electrochemical equations.

Enthalpy change ΔH: negative for exothermic (heat released), positive for endothermic (heat absorbed); measured at constant pressure.

Standard conditions: 298 K, 100 kPa (1 bar); standard enthalpy of formation ΔHf° is for one mole of compound from elements in standard states.

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

Bond enthalpy: ΔHr = Σ(bonds broken) − Σ(bonds formed); mean bond enthalpies give approximate values because they are averaged over different compounds.

Calorimetry: q = mcΔT; for solutions use c ≈ 4.18 J g⁻¹ K⁻¹ and mass of solution; ΔH = −q/n (negative for exothermic).

Lattice energy (Born–Haber): formation of one mole of ionic solid from gaseous ions; exothermic; use cycle with ionisation, atomisation, electron affinity.

Entropy S increases with temperature, phase change (s → l → g), and more moles of gas; ΔS positive for increasing disorder.

Gibbs free energy: ΔG = ΔH − TΔS; reaction feasible when ΔG < 0; T = ΔH/ΔS at equilibrium boundary.

Rate factors: concentration, pressure (gases), surface area, temperature, catalyst; rate = change in concentration / time.

Collision theory: particles must collide with energy ≥ activation energy (Ea) and correct orientation; catalyst lowers Ea.

Maxwell–Boltzmann distribution: temperature increase shifts curve right and flattens it, increasing fraction of molecules with E ≥ Ea.

Rate equation: rate = k[A]^m[B]^n; orders determined experimentally; overall order = m + n; k is temperature-dependent.

Le Chatelier's principle: system shifts to oppose change in concentration, pressure, or temperature; catalyst does not shift equilibrium position.

Equilibrium constant Kc = [products]^stoich / [reactants]^stoich; only temperature changes Kc; Kc large means products favoured.

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

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

Brønsted–Lowry: acid is a proton donor, base is a proton acceptor; conjugate pair differs by one H⁺.

pH = −log₁₀[H⁺]; [H⁺] = 10^(−pH); Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 298 K; pH + pOH = 14.

Strong acid fully dissociates (HCl, HNO₃, H₂SO₄); weak acid partially dissociates (CH₃COOH, Ka small).

Ka = [H⁺][A⁻]/[HA]; pKa = −log Ka; for weak acid pH ≈ ½(pKa − log c).

Buffer: weak acid + conjugate base (or weak base + conjugate acid); resists pH change on adding small amounts of acid or base; pH ≈ pKa + log([A⁻]/[HA]).

Titration curve features: strong acid–strong base equivalence pH 7; weak acid–strong base equivalence pH > 7; indicator pKa near steep region.

Salt hydrolysis: salts of weak acid and strong base give basic solution; salts of strong acid and weak base give acidic solution; strong–strong neutral.

Solubility rules: all nitrates and group 1 salts soluble; most chlorides soluble except AgCl, PbCl₂; most sulfates soluble except BaSO₄, PbSO₄, CaSO₄ (sparingly).

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

Oxidation state rules: elements 0; monatomic ion = charge; O usually −2 (peroxide −1); H +1 (hydride −1); sum equals overall charge.

Balance redox half-equations: balance atoms, then O with H₂O, H with H⁺ (acidic) or OH⁻ (basic), then charge with electrons.

Electrochemical cell: anode oxidation (negative in galvanic cell), cathode reduction (positive); E°cell = E°cathode − E°anode; positive E°cell means spontaneous.

Standard hydrogen electrode (SHE): 2H⁺ + 2e⁻ ⇌ H₂, E° = 0.00 V; used to measure standard electrode potentials at 298 K, 1 mol dm⁻³, 100 kPa.

Electrolysis: anode positive (oxidation), cathode negative (reduction); products depend on ion discharge series and concentration.

Faraday's laws: moles of electrons = charge (C) / 96500; mass deposited = (Q × M) / (z × 96500).

Disproportionation: same element simultaneously oxidised and reduced (e.g., Cl₂ + 2OH⁻ → Cl⁻ + ClO⁻ + H₂O).

Group 2: reactivity increases down group (lower ionisation energy); oxides and hydroxides increasingly basic; solubility of hydroxides increases down group, sulfates decrease.

Group 17: oxidising power decreases down group (Cl₂ > Br₂ > I₂); displacement reactions: more reactive halogen displaces less reactive halide.

Group 17 tests: Cl₂ bleaches damp litmus; Br₂ orange/brown; I₂ purple vapour; AgCl white, AgBr cream, AgI yellow precipitates with AgNO₃.

Period 3: Na, Mg, Al metals; Si metalloid; P, S, Cl non-metals; Ar noble gas; melting points rise to Si then fall.

Transition metals: variable oxidation states, coloured compounds, catalytic activity, form complex ions; e.g., Fe²⁺/Fe³⁺, Cu²⁺/Cu⁺, MnO₄⁻ purple.

Nitrogen chemistry: NH₃ basic; HNO₃ strong acid and oxidising; NO and NO₂ are atmospheric pollutants; fertilisers from NH₃.

Sulfur chemistry: SO₂ acidic and reducing; SO₃ forms H₂SO₄; contact process uses V₂O₅ catalyst at ~450 °C.

Flame tests: Li crimson, Na yellow, K lilac, Ca brick red, Ba apple green, Cu blue-green.

Homologous series: same general formula and functional group, differ by CH₂; IUPAC naming: longest chain, lowest locants, suffix for functional group.

Alkanes: saturated, general formula CnH2n+2; free-radical substitution with halogens (initiation, propagation, termination).

Alkenes: CnH2n, C=C double bond, planar; electrophilic addition with HBr, Br₂, H₂O; test with bromine water decolourised.

Alcohols: CnH2n+1OH; primary oxidised to aldehyde then carboxylic acid; secondary to ketone; tertiary resistant; dehydration to alkene with hot H₂SO₄.

Aldehydes and ketones: carbonyl C=O; aldehydes reduce Tollens' reagent (silver mirror) and Fehling's solution (red precipitate); ketones do not.

Carboxylic acids: weak acids, react with carbonates to give CO₂; esters formed with alcohols and acid catalyst; esters hydrolyse.

Amines: basic due to lone pair on N; primary amines formed by reduction of nitriles or ammonia with haloalkanes.

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

Apparatus: burette (±0.05 cm³) for titrations, pipette (±0.1 cm³) for fixed volumes, measuring cylinder for approximate volumes, volumetric flask for accurate dilution.

Titration procedure: rinse burette with solution, fill and remove air bubble, record initial and final readings to 0.05 cm³, repeat for concordant results (within 0.1 cm³).

Filtration: gravity for simple separation, suction for faster; crystallisation: dissolve, filter hot, cool, filter crystals, wash with cold solvent, dry.

Distillation: simple for separating solvent from non-volatile solute; fractional for liquids with different boiling points.

Tests: CO₂ limewater milky; O₂ relights glowing splint; H₂ squeaky pop; Cl₂ bleaches damp litmus; NH₃ turns damp red litmus blue; water turns anhydrous CuSO₄ blue.

Purity: melting point sharp and at expected value; impurities lower and broaden melting point; chromatography Rf value = distance spot / distance solvent front.

Errors: systematic (e.g., zero error, calibration) affect accuracy; random (e.g., reading variation) affect precision; reduce random error by repeating and averaging.

Safety: goggles, lab coat, fume cupboard for toxic/vaporous substances; waste disposal according to local rules; never pipette by mouth.