Free HSC Chemistry Flashcard Review

For a reaction aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d / [A]^a[B]^b. Kc is the ratio of product concentrations to reactant concentrations at equilibrium, each raised to the power of their stoichiometric coefficients. Kc is constant at a given temperature only.

A large Kc value (Kc >> 1) indicates that at equilibrium, the reaction has proceeded far to the right, with products significantly favoured over reactants. The equilibrium lies far towards the products.

A small Kc value (Kc << 1) indicates that at equilibrium, only a small proportion of reactants have been converted to products. The equilibrium lies far towards the reactants.

If a system at equilibrium is disturbed (by a change in concentration, pressure, temperature, or volume), the system will shift to counteract that change and re-establish equilibrium.

An increase in reactant concentration shifts equilibrium to the right (towards products) to consume the excess reactant and re-establish Kc. The system produces more products.

A decrease in pressure shifts equilibrium towards the side with the greater number of gas moles. If the forward reaction increases the number of moles, equilibrium shifts right; if it decreases moles, equilibrium shifts left.

An increase in temperature shifts an exothermic equilibrium to the left (towards reactants) because heat is a product. The system counteracts the addition of heat. Kc decreases.

N₂ + 3H₂ ⇌ 2NH₃. Industrial conditions use high pressure (150–300 atm), moderately high temperature (400–500°C), and an iron catalyst. High pressure favours products (fewer moles on right). Lower temperature favours products (exothermic) but slower reaction, so 450°C is a compromise for reasonable yield and rate.

2SO₂ + O₂ ⇌ 2SO₃. Industrial conditions use moderate pressure (1–2 atm), moderate temperature (400–500°C), and a vanadium(V) oxide catalyst. The reaction is exothermic; lower temperatures favour products but slow the reaction. About 99% conversion is achieved.

A Brønsted–Lowry acid is a proton (H⁺) donor. It donates H⁺ ions in aqueous solution (e.g., HCl → H⁺ + Cl^-).

A Brønsted–Lowry base is a proton (H⁺) acceptor. It accepts H⁺ ions in aqueous solution (e.g., NH₃ + H⁺ → NH₄⁺).

A conjugate acid–base pair differs by one proton (H⁺). If HX is an acid, X^- is its conjugate base. If NH₃ is a base, NH₄⁺ is its conjugate acid. They are related by: acid ⇌ base + H⁺.

Kw = [H⁺][OH^-] = 1.0 × 10^-14 at 25°C. In pure water, [H⁺] = [OH^-] = 1.0 × 10^-7 M, making pure water neutral (pH = 7).

pH = −log₁₀[H⁺]. pOH = −log₁₀[OH^-]. pH + pOH = 14 at 25°C. A solution is acidic if pH < 7, neutral if pH = 7, and basic if pH > 7.

For a weak acid HA, Ka = [H⁺][A^-] / [HA]. Ka is a measure of the strength of a weak acid; the larger the Ka, the stronger the acid (more dissociation). Ka is constant at a given temperature.

For a weak base B, Kb = [BH⁺][OH^-] / [B]. Kb is a measure of the strength of a weak base; the larger the Kb, the stronger the base (more dissociation). Kb is constant at a given temperature.

For a weak acid HA and its conjugate base A^-: Ka × Kb = Kw = 1.0 × 10^-14 at 25°C. This shows that the stronger the acid, the weaker its conjugate base.

A buffer is a solution that resists changes in pH when small amounts of acid or base are added. A typical buffer contains a weak acid and its conjugate base (or a weak base and its conjugate acid) in roughly equal amounts.

When H⁺ is added to an acetate buffer, the conjugate base (CH₃COO^-) reacts with it: H⁺ + CH₃COO^- → CH₃COOH. When OH^- is added, the weak acid (CH₃COOH) reacts: OH^- + CH₃COOH → CH₃COO^- + H₂O. This maintains pH.

pH = pKa + log₁₀([A^-] / [HA]). This relates the pH of a buffer to the pKa of the weak acid and the ratio of conjugate base to weak acid. When [A^-] = [HA], pH = pKa.