Chemical Equilibrium

Chemical Equilibrium – Complete Study Notes

Chemical equilibrium is a state in a reversible reaction where the rate of the forward reaction equals the rate of the reverse reaction. At this point, the concentrations of reactants and products remain constant (but not necessarily equal) over time. Equilibrium is dynamic — both reactions continue, but at equal rates, so no net change is observed.

1. Reversible and Irreversible Reactions

  • Irreversible reactions go to completion — all reactants convert to products. Example: Burning paper.
  • Reversible reactions can proceed in both directions. They reach a state of equilibrium. Example: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

2. The Equilibrium Constant (Kc)

For the general reaction aA + bB ⇌ cC + dD, the equilibrium constant is:

Kc = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ

  • Kc depends only on temperature. It does not change with concentration, pressure, or catalyst.
  • If Kc >> 1: Reaction favours products (equilibrium lies to the right).
  • If Kc << 1: Reaction favours reactants (equilibrium lies to the left).
  • Note: Pure solids and pure liquids are NOT included in the expression (their concentrations are constant).

3. Kp – Equilibrium Constant in Terms of Partial Pressures

For gaseous reactions, we can use partial pressures: Kp = Kc(RT)^Δn, where Δn = (moles of gaseous products) – (moles of gaseous reactants). When Δn = 0, Kp = Kc.

4. Le Chatelier's Principle

"If a system at equilibrium is subjected to a change in conditions, the equilibrium will shift in the direction that opposes the change." This allows us to predict how the equilibrium position changes when conditions are altered.

Effect of Concentration:

  • Adding a reactant → equilibrium shifts forward (towards products).
  • Removing a product → equilibrium shifts forward.
  • Adding a product → equilibrium shifts backward (towards reactants).

Effect of Pressure (Gaseous Reactions):

  • Increasing pressure → equilibrium shifts towards the side with fewer moles of gas.
  • In Haber Process: N₂ + 3H₂ ⇌ 2NH₃ (4 moles → 2 moles), so high pressure favours NH₃ production.
  • If Δn = 0 (equal moles both sides), pressure change has no effect on equilibrium position.

Effect of Temperature:

  • Increasing temperature → equilibrium shifts in the endothermic direction.
  • For exothermic reactions (like Haber process): increasing temperature shifts equilibrium backwards, reducing yield of NH₃. This is why a compromise temperature of 400–500°C is used.
  • Temperature DOES change the value of K (Kc or Kp).

Effect of Catalyst:

  • Catalyst does NOT shift the equilibrium position or change K.
  • It speeds up BOTH forward and reverse reactions equally, so equilibrium is reached faster.

5. Reaction Quotient (Q)

Q has the same formula as Kc, but is calculated using the current (non-equilibrium) concentrations:

  • If Q < Kc: System will shift forward (towards products).
  • If Q > Kc: System will shift backward (towards reactants).
  • If Q = Kc: System is already at equilibrium.

6. Industrial Applications

  • Haber Process (NH₃): High pressure (200 atm), moderate temperature (~400–500°C), Fe catalyst. Removes ammonia as it forms to shift equilibrium forward.
  • Contact Process (H₂SO₄): 2SO₂ + O₂ ⇌ 2SO₃, uses V₂O₅ at 450°C, 1–2 atm. High yield (~98%).

7. Ionic Equilibrium – Weak Acids and Bases

  • Strong electrolytes (HCl, NaOH, NaCl) dissociate completely in water.
  • Weak electrolytes (CH₃COOH, NH₃) partially dissociate — equilibrium is established.
  • Ka (Acid dissociation constant): Measures acid strength. Larger Ka = stronger acid.
  • Common Ion Effect: Adding a common ion suppresses the ionization of a weak acid or base (shifts equilibrium backward).
  • Buffer Solution: Resists changes in pH. Made from a weak acid and its conjugate base (or weak base + conjugate acid).
  • Solubility Product (Ksp): For sparingly soluble salts, Ksp = [Cation]ᵐ[Anion]ⁿ. Precipitation occurs when ionic product > Ksp.

Key Exam Tips

  • Temperature is the ONLY factor that changes the value of K.
  • Catalyst reaches equilibrium faster — same K, same equilibrium amounts.
  • Adding inert gas at constant volume → no effect on equilibrium (concentrations unchanged).
  • Adding inert gas at constant pressure → shifts towards more gas moles side (volume increases).
  • Active mass of pure solid or liquid = 1 (omitted from K expression).