WACE Chemistry ATAR cheat sheet

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

Electron configuration fills in order 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p; the 4s subshell fills before 3d.

Each orbital holds two electrons of opposite spin; s holds 2, p holds 6, d holds 10.

Period number equals the number of occupied electron shells; group number (main groups) equals outer-shell electron count.

Across a period, effective nuclear charge rises, atomic radius decreases and first ionisation energy generally increases.

Down a group, atomic radius and shielding increase, so first ionisation energy falls and metallic character rises.

Successive ionisation energies jump sharply at the electron removed from the next inner shell, revealing shell structure.

An emission spectrum shows discrete lines because electrons fall between quantised energy levels, emitting photons of energy E = hf.

Electronegativity increases across a period and decreases down a group (Pauling scale).

Ionic bonding: metal loses electrons, non-metal gains them; lattice energy increases with higher charges and smaller ionic radii.

Covalent bonding: shared electron pairs; bond order rises give shorter, stronger bonds (triple > double > single).

Bond polarity follows electronegativity difference; ΔEN roughly 0-0.4 nonpolar, 0.4-1.7 polar covalent, >1.7 largely ionic.

VSEPR: electron pairs repel and adopt positions of maximum separation; lone pairs repel more strongly than bonding pairs.

Electron-domain geometries: 2 linear, 3 trigonal planar, 4 tetrahedral, 5 trigonal bipyramidal, 6 octahedral.

Lone pairs reduce bond angles; CH4 109.5°, NH3 107°, H2O 104.5°.

Molecular shapes from domains: 4 domains with 1 lone pair trigonal pyramidal, 2 lone pairs bent; 6 domains with 2 lone pairs square planar.

Polar bonds plus asymmetric geometry gives a polar molecule; symmetric molecules (CO2, CCl4) are nonpolar overall.

Intermolecular forces: dispersion < dipole-dipole < hydrogen bonding; strength rises with molecular size and polarity.

n = m/M; n = cV (V in litres); n = V/24.79 at STP for gases in WACE data; molar volume may be given as 22.4 L at 0 °C, 1 atm.

Avogadro constant NA = 6.022 × 10^23 mol^-1; number of particles = n × NA.

Balance equations by conserving atoms and charge; coefficients give mole ratios for all species.

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

Percentage yield = (actual yield / theoretical yield) × 100; percentage purity = (mass pure / mass sample) × 100.

Percentage composition by mass = (mass of element in formula / molar mass of compound) × 100.

Concentration after dilution: c1V1 = c2V2 (moles conserved).

Empirical formula from percentage or mass data: divide by atomic mass, then by the smallest ratio; molecular formula = empirical × n.

Combustion analysis: all carbon goes to CO2 and all hydrogen to H2O, allowing mole ratios to be found.

Brønsted-Lowry: acid is a proton donor, base is a proton acceptor; conjugate pairs differ by one H+.

Strong acids (HCl, HNO3, H2SO4) ionise completely; weak acids ionise partially with Ka = [H+][A-]/[HA].

pH = -log10[H3O+]; [H3O+] = 10^-pH; at 25 °C pKw = 14 and pH + pOH = 14.

For a weak acid: Ka ≈ [H+]^2/c when ionisation is small, giving [H+] = sqrt(Ka × c).

Buffer of weak acid and its conjugate base: pH = pKa + log10([A-]/[HA]) (Henderson-Hasselbalch).

Neutralisation: acid + base → salt + water; acid + carbonate → salt + water + CO2.

Salt hydrolysis: salts of weak acid/strong base are basic; strong acid/weak base are acidic; strong/strong are neutral.

Titration equivalence: moles of H+ donated equal moles of H+ accepted; use n = cV with stoichiometric ratio.

An indicator changes colour over about pKa ± 1 and must bracket the equivalence pH.

Oxidation is loss of electrons (OIL), reduction is gain (RIG); oxidant is reduced, reductant is oxidised.

Oxidation number rules: free element 0, monatomic ion equals charge, O usually -2, H usually +1, sum equals overall charge.

Balance half-equations by atoms, then O with H2O, then H with H+, then charge with electrons; equalise electrons before adding.

In acidic solution use H+/H2O; in basic solution add OH- to both sides to neutralise H+.

Galvanic cell: oxidation at anode (negative), reduction at cathode (positive); electrons flow anode to cathode.

Standard cell potential E°cell = E°cathode - E°anode; a positive E°cell means the reaction is spontaneous.

Electrolytic cells use external power; anode is positive and oxidation still occurs there.

Faraday's laws: moles of electrons = It/F with F = 96 485 C mol^-1; charge = current × time.

Activity series predicts displacement: a more reactive metal displaces a less reactive metal from solution.

Enthalpy change q = mcΔT; for a reaction, ΔH = -q/n (per mole), with sign convention exothermic negative.

Exothermic reactions release energy (ΔH < 0); endothermic absorb energy (ΔH > 0).

Bond breaking absorbs energy, bond formation releases energy; ΔH = Σ bonds broken - Σ bonds formed.

Hess's law: ΔH for a route is independent of pathway; add equations and their ΔH values accordingly.

ΔH = Σ ΔHf(products) - Σ ΔHf(reactants), using standard formation enthalpies (element in standard state = 0).

Standard conditions: 100 kPa and stated temperature (usually 298 K), all species in standard states.

Calorimetry assumes dilute aqueous solutions have density 1.00 g mL^-1 and specific heat capacity 4.18 J g^-1 K^-1.

Activation energy is the minimum energy for reaction; catalysts lower it without changing ΔH.

Energy profile diagrams show reactants, products, activation energy and ΔH as the difference between reactant and product levels.

Rate = change in concentration of reactant or product per unit time; units mol L^-1 s^-1.

Rate increases with concentration, surface area, temperature and catalysts, and with pressure for gases.

Collision theory: reaction needs collisions with sufficient energy and correct orientation; only collisions exceeding Ea succeed.

Temperature raises rate mainly by increasing the fraction of collisions exceeding Ea, not by increasing collision frequency alone.

Equilibrium is dynamic: forward and reverse rates are equal and concentrations remain constant, not necessarily equal.

Le Chatelier: adding reactant shifts forward; increasing pressure shifts toward fewer gas moles; heating shifts in the endothermic direction.

Catalysts do not shift equilibrium position or K; they only shorten the time to reach it.

Kc = [products]^coefficients / [reactants]^coefficients for the balanced equation; K depends on temperature only.

For gases, Kp uses partial pressures; a large K means products predominate at equilibrium.

Alkanes CnH2n+2 are saturated; alkenes CnH2n and alkynes CnH2n-2 are unsaturated with C=C and C≡C.

Homologous series differ by CH2; IUPAC names use the longest chain and lowest locants for substituents.

Functional groups: alcohol -OH, aldehyde -CHO, ketone C=O (internal), carboxylic acid -COOH, ester -COO-, amine -NH2, amide -CONH2.

Primary alcohol oxidises to aldehyde then carboxylic acid; secondary alcohol to ketone; tertiary alcohol resists oxidation.

Carboxylic acid + alcohol (acid catalyst) gives ester + water; esterification is reversible.

Addition reactions of alkenes: hydrogenation, halogenation, hydrohalogenation and hydration; alkanes undergo substitution.

Bromine water decolourises with alkenes but not alkanes - a test for unsaturation.

Isomers share molecular formula but differ in structure; structural isomers have different physical and chemical properties.

Naming suffixes: -ane alkane, -ene alkene, -ol alcohol, -al aldehyde, -one ketone, -oic acid carboxylic acid.

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

Precipitation tests: Cl- with AgNO3 gives white AgCl; SO4^2- with BaCl2 gives white BaSO4; CO3^2- gives CO2 with acid.

Ammonium ion: warm with NaOH releases NH3 gas that turns damp red litmus blue.

Flame atomic absorption or emission spectroscopy quantifies metal ions by measuring light absorbed or emitted at characteristic wavelengths.

Chromatography separates by differential affinity between mobile and stationary phases; Rf = distance solute / distance solvent front.

Titration technique: rinse burette with titrant, read at eye level at the bottom of the meniscus, use a white tile and a dropwise end point.

Volumetric flask preparation: dissolve, transfer with rinsing, make to the calibration mark with the meniscus on the line.

Significant figures in results follow the least precise measurement; burette readings are recorded to 0.05 mL.

Gravimetric analysis: filter, wash, dry and weigh precipitate to constant mass before calculating moles.