Free QCE Chemistry Flashcard Review

A reaction that can proceed in both the forward and reverse directions simultaneously. Represented with the symbol ⇌. At equilibrium, the rate of the forward reaction equals the rate of the reverse reaction, though both continue to occur.

A state in which the rates of the forward and reverse reactions are equal, so the concentrations of reactants and products remain constant over time. The system appears static macroscopically but reactions continue at the molecular level.

1. Write products over reactants. 2. Each substance concentration is raised to the power of its stoichiometric coefficient. 3. Pure solids and pure liquids are omitted. 4. Kc = [products]/[reactants] with all concentrations in mol/L at equilibrium.

Equilibrium position: shifts right (towards products) to counteract the change. Kc: remains constant at that temperature. Kc is independent of concentration; it is a constant for a given reaction at a given temperature.

Equilibrium position: shifts towards the side with fewer moles of gas to counteract increased pressure. Kc: remains constant at that temperature. Pressure changes affect position only if there is a change in moles of gas.

Equilibrium position: shifts left (towards reactants) as temperature is removed from the system. Kc: decreases because the equilibrium constant is temperature-dependent. For endothermic reactions, the position shifts right and Kc increases.

Bronsted-Lowry acid: a proton (H⁺) donor. Bronsted-Lowry base: a proton (H⁺) acceptor. This definition applies in any solvent and explains why some substances act as bases despite not containing OH^-.

A conjugate acid-base pair consists of two species that differ by one proton (H⁺). Example: NH₃ and NH₄⁺ differ by one H⁺. When a base accepts a proton, it forms its conjugate acid; when an acid donates a proton, it forms its conjugate base.

Strong acid: completely ionises in water (e.g. HCl, H₂SO₄, HNO₃). Weak acid: partially ionises in water, establishing equilibrium between molecular and ionic forms (e.g. CH₃COOH). Strong acids are strong electrolytes; weak acids are weak electrolytes.

Strong base: completely ionises in water (e.g. NaOH, KOH). Weak base: partially ionises in water, establishing equilibrium (e.g. NH₃). Strong bases are strong electrolytes; weak bases are weak electrolytes.

pH + pOH = 14. Kw = [H⁺][OH^-] = 1.0 × 10^-14 at 25°C. pH = -log[H⁺] and pOH = -log[OH^-].

[H⁺] = 10^-ᵖᴴ = 10^-3·² = 6.3 × 10^-4 mol/L. Alternatively, use a calculator: antilog(−3.2) = 6.3 × 10^-4 mol/L.

ICE stands for Initial, Change, Equilibrium. 1. Write the equilibrium equation. 2. List initial concentrations of all species. 3. Determine concentration changes as ±x. 4. Calculate equilibrium concentrations (Initial + Change). 5. Substitute into Ka expression and solve for x.

ICE stands for Initial, Change, Equilibrium. 1. Write the equilibrium equation. 2. List initial concentrations (base concentration, [H⁺] or [OH^-] typically ≈ 0). 3. Determine changes as ±x. 4. Calculate equilibrium concentrations. 5. Substitute into Kb expression and solve for x.

A buffer solution resists large changes in pH when small amounts of acid or base are added. Composition: a weak acid and its conjugate base salt (e.g. CH₃COOH + CH₃COONa) or a weak base and its conjugate acid salt (e.g. NH₃ + NH₄Cl).

When H⁺ is added, the conjugate base anion (e.g. CH₃COO^-) neutralises it: CH₃COO^- + H⁺ → CH₃COOH. When OH^- is added, the weak acid (e.g. CH₃COOH) neutralises it: CH₃COOH + OH^- → CH₃COO^- + H₂O. This minimises pH change.

n = cV, where n = number of moles, c = concentration in mol/L, V = volume in L (or c and V in matching units). For titre calculations: n(acid)/n(base) = stoichiometric ratio from the balanced equation.

The equivalence point is the point in a titration where the moles of titrant added exactly equal the moles of analyte present, based on the stoichiometry of the reaction. At this point, one reagent is completely consumed. The pH at equivalence point depends on the acid-base nature of the products.

1. Elements in their elemental form have oxidation state 0. 2. Monatomic ions have oxidation state equal to their charge. 3. Oxygen is typically −2 (except −1 in peroxides). 4. Hydrogen is typically +1 (except −1 in hydrides). 5. Halogens are −1 (except when bonded to O or other halogens). 6. Algebraic sum in a molecule is 0; in a polyatomic ion is the charge.

An oxidising agent (oxidant) is a substance that accepts electrons and is reduced in a redox reaction. It increases the oxidation state of another species. Example: O₂ in the reaction 2Mg + O₂ → 2MgO. O₂ is reduced from 0 to −2.