Chemical equilibrium: Class 11 Chemistry

Key points

A reversible reaction can proceed in both directions, shown by the sign ⇌. In a closed vessel the reactants form products and the products form reactants again. Chemical equilibrium is reached when the rate of the forward reaction equals the rate of the reverse reaction. After that the concentrations of all substances stay constant.

Equilibrium is dynamic: both reactions go on all the time, but there is no net change. It can be reached from either side, needs a closed system, and the same state results at a given temperature. A catalyst helps equilibrium to be reached faster but does not change its position.

The law of mass action (Guldberg and Waage) states that the rate of a chemical reaction is proportional to the product of the active masses (molar concentrations) of the reactants. For aA + bB ⇌ cC + dD at equilibrium, the equilibrium constant is Kc = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ. Concentrations are in mol dm⁻³, and the powers are the coefficients in the balanced equation.

Pure solids and pure liquids are left out of the expression because their concentrations are constant. For CaCO₃(s) ⇌ CaO(s) + CO₂(g), Kc = [CO₂]. The units of Kc depend on the reaction and are (mol dm⁻³) raised to the power Δn, where Δn is moles of gaseous products minus moles of gaseous reactants. Kc depends only on temperature.

For gas reactions the constant can be written with partial pressures as Kp. The two are related by Kp = Kc(RT)^Δn. When Δn = 0, as in H₂ + I₂ ⇌ 2HI, Kp = Kc and there are no units. A large K (above about 10³) means the reaction goes almost to completion, and a very small K means little product forms.

The reaction quotient Q has the same form as K but uses concentrations at any time. If Q < K, the reaction moves forward. If Q > K, it moves backward. If Q = K, the system is at equilibrium.

Le Chatelier's principle: if a stress is applied to a system at equilibrium, the system shifts in the direction that reduces the stress. The effects are:

• Concentration: adding a reactant (or removing a product) shifts the equilibrium to the right.
• Pressure: raising the pressure shifts the equilibrium toward the side with fewer gas moles. There is no effect if the moles are equal.
• Temperature: heating favours the endothermic direction and cooling favours the exothermic direction. This is the only factor that changes the value of K.
• Catalyst and inert gas at constant volume: no change in position.

Industrial applications use the principle. In the Haber process, N₂ + 3H₂ ⇌ 2NH₃ (ΔH negative), high pressure (about 200 atm) favours ammonia because 4 moles of gas make 2. A low temperature favours the yield, but a moderate temperature (about 450 °C) with an iron catalyst is used to get a fast rate. In the contact process, 2SO₂ + O₂ ⇌ 2SO₃ is run with vanadium(V) oxide as catalyst at about 450 °C and the SO₃ is removed continuously.

Water ionises slightly: H₂O ⇌ H⁺ + OH⁻. The ionic product of water is Kw = [H⁺][OH⁻] = 1 × 10⁻¹⁴ mol² dm⁻⁶ at 25 °C. In pure water [H⁺] = [OH⁻] = 1 × 10⁻⁷ mol dm⁻³. The pH scale is pH = −log[H⁺] and pOH = −log[OH⁻], and pH + pOH = 14 at 25 °C. A solution is acidic if pH < 7, neutral if pH = 7 and basic if pH > 7.

Strong acids and bases ionise completely. Weak acids such as CH₃COOH ionise partly, with an acid dissociation constant Ka = [H⁺][A⁻]/[HA]. A smaller Ka means a weaker acid. For weak bases the constant is Kb. The Ostwald dilution idea is that a weak electrolyte ionises more as it is diluted, although the ion concentration falls.

The common ion effect is the decrease in the ionisation of a weak electrolyte when a compound with a common ion is added. Adding CH₃COONa to CH₃COOH lowers [H⁺]. A buffer solution resists a change in pH when small amounts of acid or base are added. An acidic buffer is a weak acid with its salt (CH₃COOH and CH₃COONa) and a basic buffer is a weak base with its salt (NH₄OH and NH₄Cl). Blood is buffered at about pH 7.4.

For a sparingly soluble salt such as AgCl, the solubility product Ksp is the product of the ion concentrations in a saturated solution, each raised to its power. For AgCl, Ksp = [Ag⁺][Cl⁻]. For Ca(OH)₂, Ksp = [Ca²⁺][OH⁻]². A salt precipitates when the ionic product exceeds Ksp. Adding a common ion lowers the solubility.

Study tip: For every equilibrium question write the balanced equation, then the expression, then substitute with units. In Le Chatelier questions state the stress, the direction of shift and the reason in one sentence. Memorise Haber and Contact process conditions with the reason for each one.

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Common mistakes

• Including pure solids and liquids in the Kc expression.
• Thinking that a catalyst shifts the equilibrium or changes K. It does neither.
• Writing the products over the reactants the wrong way round, or forgetting to use coefficients as powers.
• Believing equilibrium means equal concentrations. It means equal rates.
• Forgetting that only a temperature change alters the value of K.

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