BC Chemistry 12 Study Guide: Equilibrium, Acids and Bases, and Electrochemistry for BC Graduation

BC Chemistry 12 is the capstone chemistry course for British Columbia secondary students heading to university science, engineering, health sciences, or related programs. The course focuses heavily on quantitative chemistry — writing and solving equilibrium expressions, calculating pH values, and working through multi-step redox problems — and these calculation skills are directly tested in both the school-based assessments and the university admission requirements that use Chemistry 12 marks.

The course builds on Grade 11 chemistry foundations but introduces significantly more mathematical rigour, particularly in the equilibrium and acid-base units. Students who invest in understanding the conceptual basis for each calculation type — not just memorising formulas — find the later units far more manageable.

Reaction kinetics: factors and rates

Rate and collision theory: Reaction rate increases with: higher concentration (more collisions per second), higher temperature (more molecules have sufficient activation energy E_a), increased surface area (for heterogeneous reactions), use of a catalyst (alternative pathway with lower E_a).

Rate laws: For a reaction A + B → products, the rate law is rate = k[A]^m[B]^n where m and n are the orders with respect to each reactant (determined experimentally, not from the balanced equation). Overall order = m + n. The rate constant k depends on temperature (Arrhenius equation: k = Ae^(−Ea/RT)).

Reaction mechanisms: The rate-determining step (the slowest step) determines the overall rate law. The rate law for the overall reaction matches the rate law for the rate-determining step (which may not be the stoichiometry of the overall reaction).

Chemical equilibrium

Equilibrium constant expression: Kc = [products]^stoich / [reactants]^stoich. Pure solids and pure liquids are excluded (their concentrations are constants absorbed into K). For gas-phase reactions, Kp uses partial pressures; Kp = Kc(RT)^Δn where Δn = moles of gaseous products − moles of gaseous reactants.

Le Chatelier's Principle: The key insight is that changing a concentration shifts the equilibrium position but not the value of K (at constant temperature). Changing the temperature changes the value of K. For exothermic reactions: increasing temperature decreases K (shifts toward reactants). For endothermic reactions: increasing temperature increases K.

ICE table method:

1. Write the balanced equation
2. Set up columns for each species
3. Row 1 (Initial): given initial concentrations (or zero for products if reaction starts from all reactants)
4. Row 2 (Change): −ax, +cx, etc. using stoichiometric coefficients
5. Row 3 (Equilibrium): sum of rows 1 and 2
6. Write the Keq expression and substitute equilibrium values
7. Solve for x

Use the Cornell Notes Tool to set up an ICE table template with the common simplification condition and verification step.

Acids, bases, and pH

Strong acids and bases: Complete dissociation means [H⁺] = [HCl], [OH⁻] = [NaOH], etc. pH = −log[H⁺], pOH = −log[OH⁻], pH + pOH = 14.

Weak acid/base calculations: Use ICE tables as described in the FAQ. The approximation √(KaC) works when Ka/C < 0.01 (less than 1% error). The full quadratic: Ka = x²/(C − x) → Kax − Kx − KaC = 0. No wait: x² + Ka·x − Ka·C = 0.

Actually: Ka = x²/(C−x), rearranging: x² + Kax − KaC = 0. Use quadratic formula: x = (−Ka + √(Ka² + 4KaC))/2 (positive root).

Buffers: A buffer contains a weak acid (HA) and its conjugate base (A⁻) — typically as a weak acid plus salt of its conjugate base. The buffer resists pH change because added H⁺ is consumed by A⁻ and added OH⁻ is consumed by HA. Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA]). Buffer capacity is highest when [HA] = [A⁻] (pH = pKa at this point).

Titration: Strong acid-strong base: equivalence point at pH 7 (salt of strong acid and strong base is neutral). Weak acid-strong base: equivalence point at pH > 7 (conjugate base of weak acid is basic). Indicators: change colour near their own pKa — choose an indicator whose pKa is close to the expected equivalence point pH.

Electrochemistry

Standard reduction potentials: The standard electrode potential E° measures the tendency of a half-reaction to proceed as a reduction. Higher E° = stronger oxidising agent. In a galvanic cell: the more positive half-reaction occurs as reduction (cathode); the less positive (more negative) as oxidation (anode). Cell voltage: E°cell = E°(cathode) − E°(anode). Positive E°cell → spontaneous.

Nernst equation: Under non-standard conditions: E = E° − (0.0592/n)logQ (at 25°C) where Q is the reaction quotient and n is the number of electrons transferred. At equilibrium, E = 0 and Q = K — this allows you to find K from E°.

Electrolysis and Faraday's laws: The number of moles of substance produced/consumed in electrolysis = Q/(nF) = It/(nF) where I = current (A), t = time (s), n = electrons per ion, F = 96,485 C/mol. Convert moles to mass using molar mass.

The Spaced Repetition Flashcard Tool works well for the half-reaction potentials and acid dissociation constants. Use the Pomodoro Timer for timed calculation practice. If you are also studying BC Physics 12, see the BC Pre-Calculus 12 study guide for the mathematical foundations that support both sciences.