VCE Chemistry cheat sheet

Mole relation: n = m/M, where M is molar mass in g mol⁻¹; n = N/N_A with N_A = 6.02 × 10²³ mol⁻¹; for gases at STP (0 °C, 100 kPa) n = V/22.7 L mol⁻¹.

Empirical formula: divide each element's mole amount by the smallest value; molecular formula = (empirical formula) × k, where k = M(molecular)/M(empirical).

Percentage yield = (actual yield ÷ theoretical yield) × 100; percentage atom economy = (M of desired product ÷ sum of M of all reactants) × 100.

Limiting reagent: calculate n of each reactant, divide by its stoichiometric coefficient; the smallest quotient is limiting and sets the maximum product amount.

Concentration: c = n/V (mol L⁻¹); dilution obeys c₁V₁ = c₂V₂ because moles of solute are conserved; ppm = mg solute per kg (or L) of solution.

Ionic equations: cancel spectator ions; strong acids and soluble salts are written as separated ions, weak acids and precipitates as intact formulas.

Solubility rules: all group 1 and ammonium salts are soluble; all nitrates and acetates are soluble; most chlorides are soluble except AgCl, PbCl₂ and Hg₂Cl₂; most carbonates and hydroxides are insoluble except those of group 1 and ammonium.

Redox: oxidation is loss of electrons (increase in oxidation number), reduction is gain; in a balanced half-equation, atoms and charge must both balance, with H⁺ and H₂O added in acidic solution.

Organic naming: find the longest carbon chain containing the principal functional group; number to give that group the lowest locant; name substituents alphabetically with locants.

Priority of functional groups for suffix naming: carboxylic acid > ester > amide > nitrile > aldehyde > ketone > alcohol > amine > alkene > alkyne > halogen/alkyl.

Esterification: carboxylic acid + alcohol ⇌ ester + water, acid-catalysed and reversible; the ester is named alkyl alkanoate, e.g. ethyl ethanoate from ethanol and ethanoic acid.

Addition polymers form from monomers containing C=C; condensation polymers (polyesters, polyamides) form with loss of a small molecule such as H₂O, and their monomers each need two reactive groups.

Calorimetry: q = mcΔT with c(water) = 4.18 J g⁻¹ K⁻¹; ΔH = −q/n (kJ mol⁻¹), negative for exothermic reactions; a bomb calorimeter uses constant volume.

Equilibrium: K = [products]^coefficients ÷ [reactants]^coefficients for homogeneous systems at constant temperature; K depends only on temperature, and Q compared with K predicts direction of shift.

Le Chatelier: increasing concentration of a reactant shifts right; increasing pressure shifts toward fewer gas moles; raising temperature shifts in the endothermic direction and changes K.

Acid–base: pH = −log₁₀[H₃O⁺]; Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C; strong acids ionise completely, weak acids only partially (Ka small, conjugate base stronger).

Volumetric analysis: at the equivalence point of a monoprotic acid–base titration, n(acid) = n(base); use c₁V₁/n₁ = c₂V₂/n₂ for other stoichiometries, and choose an indicator whose colour change brackets the equivalence pH.

Instrumental analysis: AAS measures absorbance of a specific wavelength by ground-state atoms to quantify an element via a calibration curve; IR identifies bonds by absorption wavenumber; ¹H NMR gives chemical shift, integration and splitting (n + 1 rule) for hydrogen environments; MS gives molecular ion mass and fragmentation pattern.

Bonding continuum: electronegativity difference > about 1.7 gives predominantly ionic bonding, intermediate values polar covalent, near-zero nonpolar covalent; metallic bonding involves a lattice of cations in a delocalised electron sea.

VSEPR: electron domains repel to maximum separation - 2 domains linear (180°), 3 trigonal planar (120°), 4 tetrahedral (109.5°); lone pairs compress bond angles, e.g. H₂O about 104.5°.

Intermolecular forces ranked by strength: dispersion (all molecules, increases with electron number and surface area) < dipole–dipole < hydrogen bonding (H bonded to N, O or F).

Properties follow structure: ionic solids have high melting points and conduct only when molten or dissolved; metals conduct as solids; molecular covalent substances have low melting points; giant covalent networks (diamond, SiO₂) are hard and high-melting.

Reaction types: precipitation (two soluble salts form an insoluble product), acid–base neutralisation (H⁺ + OH⁻ → H₂O), acid–metal (metal + acid → salt + H₂), acid–carbonate (→ salt + H₂O + CO₂), and combustion.

Exothermic reactions release energy (ΔH negative, products lower enthalpy); endothermic absorb energy (ΔH positive); bond breaking absorbs energy and bond forming releases it, so ΔH = energy in (bonds broken) − energy out (bonds formed).

Avogadro's law: equal volumes of gases at the same temperature and pressure contain equal numbers of molecules, so gas volume ratios equal mole ratios in a balanced equation.

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

Fossil fuel combustion: complete combustion of a hydrocarbon gives CO₂ + H₂O; incomplete combustion gives CO and/or soot, and CO binds haemoglobin, reducing oxygen transport.

Biofuel comparison: ethanol from fermentation (glucose → 2 ethanol + 2 CO₂, yeast, about 37 °C, anaerobic) versus ethanol from ethene hydration (H₃PO₄ catalyst, about 300 °C, high pressure); energy density and carbon neutrality differ.

Galvanic cells: oxidation occurs at the anode (negative electrode), reduction at the cathode (positive); electrons flow through the external circuit from anode to cathode, ions migrate in the electrolyte to complete the circuit.

Standard electrode potentials: a half-cell with more positive E° is reduced; E°(cell) = E°(reduction, cathode) − E°(reduction, anode); a positive E°(cell) means the reaction is spontaneous under standard conditions.

Rate of reaction = change in concentration of a reactant or product per unit time; factors increasing rate: higher concentration or pressure, higher temperature, greater surface area, and a catalyst.

Collision theory: reaction requires collisions with energy at least equal to the activation energy and correct orientation; a catalyst lowers activation energy by providing an alternative pathway.

Rate–time and concentration–time graphs: a steeper initial gradient means a faster rate; rate decreases over time as reactant concentration falls; a catalyst raises the initial gradient but not the final yield.

Organic reaction pathways: alkene + H₂O (H⁺ catalyst) → alcohol; alcohol + H₂SO₄/heat → alkene (dehydration); primary alcohol oxidised by acidified dichromate to aldehyde then carboxylic acid; secondary alcohol to ketone; haloalkane + NaOH(aq) → alcohol (substitution).