Equilibrium

The One Thing

Equilibrium is a dynamic state in which opposing processes continue at equal rates, so measurable properties remain constant. Equilibrium constants, reaction quotients, Le Chatelier’s principle, and ionic-equilibrium relationships allow the composition, direction, pH, and precipitation behavior of chemical systems to be predicted.

Reactions, Processes and Experiments

Key Terms

• Physical equilibrium: A reversible physical process in which the forward and reverse changes occur at equal rates, such as water evaporating and condensing in a closed container.
• Chemical equilibrium: The dynamic state of a reversible chemical reaction in which the rates of the forward and reverse reactions are equal.
• Dynamic equilibrium: A state where reactions continue microscopically, but concentrations or other observable properties remain constant macroscopically.
• Reversible reaction: A reaction that can proceed in both forward and backward directions under suitable conditions.
• 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 reactants.
• Equilibrium constant, Kc: For aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d/[A]^a[B]^b at equilibrium; pure solids and pure liquids are omitted.
• Equilibrium constant, Kp: For gaseous reactions, Kp = (pC)^c(pD)^d/(pA)^a(pB)^b, using partial pressures.
• Reaction quotient, Q: The same expression as the equilibrium constant, but using concentrations or pressures at any instant.
• Le Chatelier’s principle: When an equilibrium system is disturbed by changing concentration, pressure, or temperature, it shifts in the direction that reduces the disturbance.
• Homogeneous equilibrium: An equilibrium in which all reactants and products are in the same physical phase.
• Heterogeneous equilibrium: An equilibrium involving substances in more than one physical phase.
• Degree of dissociation, alpha: The fraction of the initial amount that dissociates or reacts; alpha = amount dissociated/initial amount.
• Electrolyte: A substance that produces ions in aqueous solution or molten state and conducts electricity.
• Strong electrolyte: An electrolyte that ionizes or dissociates almost completely, such as HCl, NaOH, or NaCl.
• Weak electrolyte: An electrolyte that ionizes only partially and establishes ionic equilibrium, such as CH3COOH or NH3.
• Arrhenius acid: A substance that increases the concentration of H3O+ ions in aqueous solution.
• Arrhenius base: A substance that increases the concentration of OH- ions in aqueous solution.
• Bronsted-Lowry acid: A proton donor.
• Bronsted-Lowry base: A proton acceptor.
• Conjugate acid-base pair: Two species that differ by one proton, such as NH4+ and NH3, or H2O and OH-.
• Lewis acid: A species that accepts an electron pair.
• Lewis base: A species that donates an electron pair.
• Acid ionization constant, Ka: For HA + H2O ⇌ H3O+ + A-, Ka = [H3O+][A-]/[HA], indicating weak-acid strength.
• Base ionization constant, Kb: For B + H2O ⇌ BH+ + OH-, Kb = [BH+][OH-]/[B], indicating weak-base strength.
• Ionic product of water, Kw: Kw = [H3O+][OH-] = 1.0 × 10^-14 at 25 °C.
• pH: A measure of hydrogen-ion concentration: pH = -log[H3O+].
• pOH: A measure of hydroxide-ion concentration: pOH = -log[OH-].
• Acid strength: The extent to which an acid ionizes in water; larger Ka and smaller pKa generally indicate a stronger acid.
• Common-ion effect: Suppression of weak-electrolyte ionization when a strong electrolyte containing a common ion is added.
• Buffer solution: A solution that resists significant pH change when small amounts of acid or base are added.
• Acidic buffer: A mixture of a weak acid and its salt with a strong base, such as CH3COOH and CH3COONa.
• Basic buffer: A mixture of a weak base and its salt with a strong acid, such as NH3 and NH4Cl.
• Henderson-Hasselbalch equation: For an acidic buffer, pH = pKa + log([salt]/[acid]); for a basic buffer, pOH = pKb + log([salt]/[base]).
• Salt hydrolysis: The reaction of ions from a salt with water, making the solution acidic, basic, or nearly neutral.
• Solubility product, Ksp: The equilibrium constant for dissolution of a sparingly soluble salt; for MX ⇌ M+ + X-, Ksp = [M+][X-].
• Molar solubility: The number of moles of a sparingly soluble substance dissolved per litre of saturated solution.
• Precipitation: Formation of a solid when the ionic product exceeds the solubility product.

Easily Confused

• Physical equilibrium and chemical equilibrium: Physical equilibrium involves changes of physical state, such as evaporation and condensation; chemical equilibrium involves reversible chemical reactions.
• Dynamic equilibrium and static equilibrium: Dynamic equilibrium involves continuing forward and reverse processes at equal rates; the system is not microscopically static.
• K and Q: K uses equilibrium concentrations or pressures, whereas Q uses values at any instant.
• K and reaction rate: K indicates the extent or position of equilibrium, not how quickly the reaction occurs.
• Homogeneous and heterogeneous equilibrium: Homogeneous equilibrium has all species in one phase; heterogeneous equilibrium contains more than one phase.
• Strong and weak electrolytes: Strong electrolytes ionize or dissociate almost completely; weak electrolytes ionize only partially.
• Arrhenius and Bronsted-Lowry definitions: Arrhenius acids and bases are defined by producing H3O+ or OH- in aqueous solution; Bronsted-Lowry acids and bases are proton donors and acceptors.
• Bronsted-Lowry and Lewis definitions: Bronsted-Lowry theory concerns proton transfer; Lewis theory concerns electron-pair donation and acceptance.
• Ka and Kb: Ka measures weak-acid ionization, whereas Kb measures weak-base ionization.
• pH and pOH: pH measures hydronium concentration; pOH measures hydroxide concentration.
• Acidic and basic buffers: An acidic buffer contains a weak acid and its conjugate base; a basic buffer contains a weak base and its conjugate acid.
• Common-ion effect and buffer action: The common-ion effect suppresses weak-electrolyte ionization; buffer action involves consumption of added acid or base by a conjugate acid-base pair.
• Ksp and Qsp: Ksp describes a saturated solubility equilibrium; Qsp is calculated at any instant to determine whether precipitation will occur.
• Precipitation and dissolution: Precipitation forms a solid when Qsp > Ksp; dissolution transfers solid into solution until equilibrium is reached.

What Gets Asked

• Writing equilibrium expressions: Questions may require Kc for aA + bB ⇌ cC + dD; marks are lost by including pure solids or pure liquids.
• Predicting reaction direction: Compare Q with K; Q < K means forward reaction, Q > K means reverse reaction, and Q = K means equilibrium.
• Applying Le Chatelier’s principle: State the effect of concentration, pressure, or temperature changes; marks are lost by claiming that a catalyst changes K or the equilibrium composition.
• Calculating weak-electrolyte ionization: Use Ka = C alpha^2/(1 - alpha) or Kb approximately C alpha^2; marks are lost by applying the small-alpha approximation without the stated condition that ionization is small.
• Calculating pH, pOH, and buffer pH: Use pH = -log[H3O+], pOH = -log[OH-], and the appropriate Henderson-Hasselbalch form; marks are lost by using the acidic-buffer equation for a basic buffer or omitting pH = 14 - pOH at 25 °C.
• Predicting salt hydrolysis and precipitation: Identify whether a salt solution is acidic, basic, or neutral, and compare Qsp with Ksp; marks are lost by treating all salts as neutral or ignoring stoichiometric coefficients in expressions such as Ksp = (2s)^2(3s)^3 = 108s^5.