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High Yield Topic

Equilibrium

Chemistry Unit 7
25 min read
IAT Advanced
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1. Equilibrium in Physical Processes

Equilibrium can exist between different phases of the same substance. It is dynamic—the process continues in both directions at equal rates.

Process Type Example Condition / Constant
Solid-Liquid Ice ↔ Water Occurs at Melting Point (P = 1 atm, T = 273K).
Liquid-Gas Water ↔ Steam Vapour pressure is constant at a given Temperature.
Solid-Gas I2(s) ↔ I2(v) Sublimation equilibrium in a closed vessel.
Dissolution Sugar(s) ↔ Sugar(aq) Occurs in a Saturated Solution.

General Characteristics

  • Equilibrium is possible only in Closed Systems.
  • Both forward and backward processes occur at the same rate.
  • Measurable properties (concentration, P, T) remain constant.

2. Chemical Equilibrium

In a reversible chemical reaction, equilibrium is reached when the concentrations of reactants and products stop changing.

Kc = [C]c[D]d / [A]a[B]b
Law of Chemical Equilibrium for aA + bB ↔ cC + dD.
Kp = Kc(RT)Δng
Δng = (moles of gaseous products) - (moles of gaseous reactants).
If Δng = 0, then Kp = Kc.

Types of Equilibria

  • Homogeneous: All reactants and products are in the same phase (e.g., all gases).
  • Heterogeneous: Components are in different phases. Crucial: Concentrations of Pure Solids and Pure Liquids are taken as unity (1) and omitted from K expressions.

3. Applications of Equilibrium Constant

K tells us the extent of a reaction and helps predict its direction.

Qc = [C]tc[D]td / [A]ta[B]tb
Reaction Quotient (Q): Calculated using instantaneous concentrations.
• Q < K: Net reaction moves Forward.
• Q > K: Net reaction moves Backward.
• Q = K: System is at Equilibrium.
ΔG = ΔG° + RT ln Q
At Equilibrium (ΔG = 0, Q = K):
ΔG° = -RT ln K = -2.303 RT log K

4. Le Chatelier’s Principle

If a system at equilibrium is subjected to a change in concentration, pressure, or temperature, the system shifts in a direction that tends to undo the effect of the change.

Factor Change Shift Direction
Concentration Add Reactant / Remove Product Forward (towards products)
Pressure Increase Pressure Side with Fewer gas moles
Temperature Increase (Exothermic, -ΔH) Backward (Towards reactants)
Temperature Increase (Endothermic, +ΔH) Forward (Towards products)
Inert Gas Added at Constant Volume No Effect

IAT Pro Tip: Catalysts

A catalyst does not shift the equilibrium. It only helps the system reach equilibrium faster by lowering the activation energy for both forward and backward reactions equally.

5. Ionic Equilibrium

Ionic equilibrium involves the balance between un-ionized molecules and ions in solution.

Theory Acid Definition Base Definition
Arrhenius Produces H+ in water Produces OH- in water
Brønsted-Lowry Proton (H+) Donor Proton (H+) Acceptor
Lewis Electron-pair Acceptor Electron-pair Donor
pH = -log[H3O+] | pKw = pH + pOH = 14
At 298 K, Kw = 1.0 × 10-14.

Ostwald’s Dilution Law

For a weak electrolyte (e.g., acetic acid): α = √(Ka / C)

As concentration (C) decreases (dilution increases), the degree of ionization (α) increases.

6. Buffers & Solubility Product

Buffer Solutions

Solutions that resist pH change upon addition of small amounts of acid or base.

  • Acidic Buffer: Weak Acid + Salt of its conjugate base (e.g., CH3COOH + CH3COONa).
  • Basic Buffer: Weak Base + Salt of its conjugate acid (e.g., NH4OH + NH4Cl).
pH = pKa + log([Salt] / [Acid])
Henderson-Hasselbalch Equation. Used for acidic buffers. For basic buffers: pOH = pKb + log([Salt]/[Base]).

Solubility Product (Ksp)

Ksp = [Mn+]x [Xm-]y
For salt MxXy. Precipitation occurs if Ionic Product (Qsp) > Ksp.

Common Ion Effect

The solubility of a sparingly soluble salt decreases in the presence of a common ion. (e.g., AgCl is less soluble in NaCl solution than in pure water).

Quick Revision Flashcards

Hover over a card to reveal the definition/formula.

Dynamic Equilibrium

State where forward and backward rates are equal. Macroscopic properties are constant, but microscopic activity continues.

Lewis Acid

A species that can accept an electron pair (e.g., BF3, AlCl3, H+).

Buffer Action

The ability of a buffer solution to resist changes in pH when acid or base is added.

7. Common Mistakes

  • Δng Calculation: Ignoring the state of matter (s, l, g). Only include gaseous moles when calculating Δng for Kp = Kc(RT)Δng.
  • Logarithm Errors: If [H+] = 2 × 10-4, pH = 4 - log(2) ≈ 3.7. Don't simply guess 4.
  • Pure Solids/Liquids: Forgetting to exclude them from Kc/Kp expressions. Their concentration remains constant and is taken as 1.

IAT Exam Focus Points

High-Yield Areas

  • Le Chatelier Numericals: Predict shifts when volume is decreased or inert gas is added at constant P vs constant V.
  • pH Calculation: Mixtures of strong acid/base (neutralization) vs weak acid/base (buffer).
  • Solubility Product: Calculating solubility (s) from Ksp for different salt types (1:1, 1:2, 2:3).

8. Practice Mock Test

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Equilibrium

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