Equilibrium

The One Thing

Equilibrium is a dynamic state in which forward and reverse reactions continue at equal rates, so the composition of a closed system remains constant. Equilibrium constants, reaction quotients, and Le Chatelier’s principle describe the composition and response of chemical and ionic equilibria.

Reactions, Processes and Experiments

Key Terms

• Reversible reaction: A reaction that can proceed in both forward and reverse directions under suitable conditions.
• Dynamic equilibrium: A condition in a closed system where the forward and reverse reaction rates are equal, although both reactions continue.
• Law of mass action: At constant temperature, the rate of a reaction is proportional to the product of the active masses or concentrations of the reacting species.
• Equilibrium constant, Kc: For aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d/[A]^a[B]^b at a fixed temperature.
• Equilibrium constant, Kp: For gaseous reactions, Kp = (pC)^c(pD)^d/(pA)^a(pB)^b, where p represents the equilibrium partial pressure.
• Reaction quotient, Q: The expression obtained using current concentrations or pressures; comparing Q with K predicts the direction in which a reaction will proceed.
• Homogeneous equilibrium: An equilibrium in which all reacting species are in the same physical state.
• Heterogeneous equilibrium: An equilibrium involving substances in more than one physical state; pure solids and pure liquids are omitted from the equilibrium expression.
• Le Chatelier’s principle: When a system at equilibrium is disturbed, it shifts in the direction that reduces the effect of the disturbance.
• Arrhenius acid and base: An acid produces H+ or H3O+ in water, while a base produces OH− in water.
• Bronsted-Lowry acid and base: An acid donates a proton, and a base accepts a proton.
• Conjugate acid-base pair: Two species that differ by one proton, such as NH4+ and NH3 or H2CO3 and HCO3−.
• Strong electrolyte: A substance that ionizes or dissociates almost completely in aqueous solution.
• Weak electrolyte: A substance that ionizes only partially in aqueous solution and establishes ionic equilibrium.
• Ionic product of water, Kw: Kw = [H3O+][OH−] = 1.0 × 10^−14 at 25 °C in dilute aqueous solution.
• pH and pOH: pH = −log[H3O+] and pOH = −log[OH−]; at 25 °C, pH + pOH = 14.
• Acid dissociation constant, Ka: For HA ⇌ H+ + A−, Ka = [H+][A−]/[HA]. A larger Ka indicates a stronger weak acid.
• Base dissociation constant, Kb: For BOH ⇌ B+ + OH− or B + H2O ⇌ BH+ + OH−, Kb measures the extent of base ionization.
• Relationship between Ka and Kb: For a conjugate acid-base pair at 25 °C, Ka × Kb = Kw.
• Degree of ionization, α: The fraction of the original electrolyte that ionizes; for a weak electrolyte, α = ionized concentration/original concentration.
• Ostwald’s dilution law: For a weak electrolyte, its degree of ionization increases on dilution; for a weak acid, Ka = Cα^2/(1 − α), approximately Ka = Cα^2 when α is small.
• Common-ion effect: The ionization of a weak electrolyte is suppressed when a strong electrolyte containing a common ion is added.
• Buffer solution: A solution that resists significant change in pH when small amounts of acid or base are added.
• Henderson-Hasselbalch equation: For an acidic buffer, pH = pKa + log([salt]/[acid]); for a basic buffer, pOH = pKb + log([salt]/[base]).
• Solubility product, Ksp: The equilibrium constant for dissolution of a sparingly soluble ionic solid, written using the equilibrium concentrations of its ions.
• Precipitation condition: If the ionic product is less than Ksp, no precipitation occurs; if equal to Ksp, the solution is saturated; if greater than Ksp, precipitation is expected.

Easily Confused

• Dynamic equilibrium and static equilibrium: Dynamic equilibrium involves continuing forward and reverse reactions at equal rates; it is not a condition in which particles stop reacting.
• K and Q: K is calculated using equilibrium concentrations or pressures; Q uses current concentrations or pressures and predicts the direction of adjustment.
• K and reaction speed: K indicates whether products or reactants are favored at equilibrium; it does not indicate the speed of a reaction.
• Concentration or pressure changes and temperature changes: Concentration, pressure, volume, and catalysts do not change K at constant temperature; temperature changes K.
• Homogeneous and heterogeneous equilibrium: Homogeneous equilibrium contains one physical state; heterogeneous equilibrium contains more than one, with pure solids and pure liquids omitted from the expression.
• Strong and weak electrolytes: Strong electrolytes ionize or dissociate almost completely; weak electrolytes ionize partially and establish ionic equilibrium.
• Arrhenius and Brønsted–Lowry definitions: Arrhenius acids and bases are defined by H+/H3O+ and OH− production in water; Brønsted–Lowry acids and bases are defined by proton donation and acceptance.
• Ka and Kb: Ka measures weak-acid ionization, whereas Kb measures weak-base ionization; for conjugate pairs at 25 °C, Ka × Kb = Kw.
• Dilution and common-ion addition: Dilution generally increases weak-electrolyte ionization; common-ion addition suppresses it.
• Acidic and basic buffers: An acidic buffer uses a weak acid and its salt and applies pH = pKa + log([salt]/[acid]); a basic buffer uses a weak base and its salt and applies pOH = pKb + log([salt]/[base]).
• Saturated solution and precipitation: A solution is saturated when the ionic product equals Ksp; precipitation is expected when the ionic product is greater than Ksp.
• Kc and Kp: Kc uses equilibrium concentrations, while Kp uses equilibrium partial pressures of gases.

What Gets Asked

• Writing equilibrium expressions: Questions may require Kc or Kp for a balanced equation. Marks are lost by omitting stoichiometric coefficients as exponents or including pure solids and pure liquids.
• Manipulating equilibrium constants: Questions may reverse an equation or multiply all coefficients. The required changes are 1/K for reversal and K^n when coefficients are multiplied by n.
• Predicting equilibrium shifts: Questions may apply concentration, pressure, volume, temperature, or catalyst changes. The key slips are treating a catalyst as changing K, applying pressure effects when gaseous mole numbers are equal, or confusing endothermic and exothermic responses.
• Using Q and K: Questions may ask for the direction of reaction. The required comparisons are Q < K for the forward direction, Q > K for the reverse direction, and Q = K for equilibrium.
• Weak-acid, weak-base, and pH calculations: Questions may require Ka, Kb, α, Ostwald’s dilution law, pH, or pOH. Marks are lost by treating weak electrolytes as completely ionized or using the approximations without the stated small-ionization condition.
• Buffers, solubility, and precipitation: Questions may require the Henderson-Hasselbalch equation, Ksp, molar solubility, or an ionic-product comparison. The specific slips are interchanging acidic and basic buffer equations, using the wrong dissolution stoichiometry, or predicting precipitation when the ionic product is less than Ksp.