Review and Capstone

What to recall. The whole quantitative core hangs on the mole. Convert with mol = mass ÷ molar mass; read mole ratios from the balanced equation to bridge any two species; track significant figures through the calculation. For solutions, molarity = mol solute ÷ L solution, dilutions follow M₁V₁ = M₂V₂, and an ion’s concentration is the formula-unit molarity times its subscript on dissolution. Finish with limiting-reagent and density conversions when a problem mixes mass, volume, and amount.

Review before this strand: Units & Measurement, Moles & Composition, Predicting & Balancing, Stoichiometry, Solutions & Molarity.

What to recall. Heat tracks with q = mcΔT, and calorimetry is conservation: q_rxn = −q_surroundings (a bomb gives q_rxn = −C_calΔT). Build reaction enthalpies by Hess’s law or ΔH°_rxn = ∑nΔH°_f(products) − ∑nΔH°_f(reactants); entropies the same way with ΔS°_rxn = ∑nS°(products) − ∑nS°(reactants) (watch J vs kJ). Spontaneity is ΔG = ΔH − TΔS: negative-ΔH/positive-ΔS always favorable, the reverse never, and the two mixed cases switch at T = ΔH ÷ ΔS.

Review before this strand: Thermochemistry, IMFs & Phases, Thermodynamics.

What to recall. Get rate-law orders from the ratio of initial rates when one concentration changes at a time; first-order decay obeys ln([A]₀/[A]) = kt with t_1/2 = ln 2 ÷ k. Write the equilibrium expression as products over reactants, each raised to its coefficient; compare Q to K for direction; solve concentrations with an ICE table. Shift predictions come from Le Châtelier — concentration, pressure/volume through moles of gas, and temperature through endo/exo. Gas mixtures partition by Dalton’s partial pressures.

Review before this strand: Gas Laws, Kinetics, Equilibrium.

What to recall. Strong acids and bases give pH straight from concentration. A weak acid needs an ICE table: K_a = x² ÷ (C − x) — drop the −x only when ionization is small (a few percent), otherwise solve the quadratic. Buffers follow Henderson–Hasselbalch: pH = pK_a + log([A⁻] ÷ [HA]). Get K_sp from molar solubility (mind the stoichiometric powers), salt pH from hydrolysis with K_a·K_b = K_w, and titration results from neutralization stoichiometry to the equivalence point.

Review before this strand: Acids & Bases, Buffers, Solubility Equilibria.

What to recall. Balance redox by half-reactions — atoms first, then O with H₂O, H with H⁺, charge with e⁻; in base, neutralize H⁺ with OH⁻. A galvanic cell’s standard potential is E°_cell = E°_cathode − E°_anode, spontaneous when E°_cell > 0 (since ΔG° = −nFE°); the Nernst equation handles nonstandard concentrations. Electrolysis is bookkeeping with Faraday: moles e⁻ = (I·t) ÷ F. The stronger oxidizing agent has the more positive E°.

Review before this strand: Redox & Balancing, Electrochemistry.

What to recall. Write ground-state electron configurations, including the Cr/Cu exceptions and the remove-4s-before-3d rule for transition-metal ions; count valence and d electrons from them. Predict VSEPR geometry from electron domains and the predominant intermolecular force (hydrogen bonding > dipole–dipole > dispersion, scaled by size). Periodic trends rank radius, ionization energy, and electronegativity. A cubic unit cell gives density by ρ = (Z · M) ÷ (N_A · a³).

Review before this strand: Atomic Structure, Electron Structure, Molecular Shape, IMFs & Phases.

What to recall. Name complexes by IUPAC rules — ligand prefixes in alphabetical order, the −ate suffix for an anionic complex, and the metal’s oxidation state in Roman numerals. The coordination number is the count of donor atoms; the metal’s oxidation number comes from charge balance against the ligands and counter-ions. From the d-electron count and the field strength, crystal-field theory fixes the spin state — strong field → low spin, weak field → high spin — which sets the number of unpaired electrons and whether the complex is paramagnetic or diamagnetic.

Review before this strand: Coordination Chemistry.