Dynamic Equilibrium Notes

Dynamic Equilibrium
Definition

Dynamic Chemical Equilibrium: A reversible reaction where forward and reverse rates are equal, resulting in constant reactant and product concentrations.

Conditions for Establishment

  • Reversible Reactions: Products can revert to reactants.

  • Closed System: Conserves mass, allows energy exchange.

    • Open System: Exchanges both energy and matter.

  • Constant Conditions: Temperature (T), pressure (p), and concentration (c) are constant.

  • Catalysts: Affect forward and reverse rates equally; do not shift equilibrium.

Equilibrium Establishment

  • Forward reaction rate equals reverse reaction rate.

  • Reactants in equilibrium with products:Reactants \rightleftharpoons Products

  • Reactant and product quantities remain constant.

  • Temperature and rate factors are constant.

  • Forward and reverse reactions occur simultaneously.

Equilibrium Position

  • To the Right: Favours product formation (more products).

  • To the Left: Favours reactant formation (more reactants).

Le Chatelier’s Principle

Equilibrium shifts to counteract stress (pressure, temperature, concentration changes), establishing a new equilibrium.

Factors Affecting Equilibrium

  • Pressure (gases only)

  • Concentration

  • Temperature

  • Catalysts (no change in equilibrium position)

How to Approach Le Chatelier Questions

  1. Identify disturbance.

  2. Refer to Le Chatelier's principle.

  3. Determine the system's response and favoured reaction.

  4. Apply collision theory.

  5. Determine the effect and answer.

Cheat Sheet

Disturbance

Shift Direction

Effect on Products

Increase Reactants

Shifts Right (Forward Reaction)

Increases

Decrease Reactants

Shifts Left (Reverse Reaction)

Decreases

Increase Temperature

Shifts Away from Heat (Endothermic)

Depends on Reaction

Decrease Temperature

Shifts Towards Heat (Exothermic)

Depends on Reaction

Increase Pressure

Shifts to Fewer Gas Moles

Favors Side with Fewer Moles

Decrease Pressure

Shifts to More Gas Moles

Favors Side with More Moles

Add Catalyst

No Shift

Increases Rate Equally

Example:2 SO2 (g) + O2 (g) \rightleftharpoons 2 SO_3 (g) \Delta H < 0

  1. Increase reactants: shifts forward to consume.

  2. Decrease temperature: Shifts exothermic (forward) to produce heat.

  3. Volume halved (pressure increased): shifts to fewer gas moles (forward).

Other Factors

  • Precipitation: Insoluble salts affect equilibrium.

  • Common Ion Effect: Reduces salt solubility when a common ion is added.

Example

Cu(H2O)6^{2+} (aq) + 4 Cl^- (aq) \rightleftharpoons CuCl4^{2+} (aq)+ 6 H2O (l) \Delta H > 0

  1. Add HCl (aq): Shifts right (green).

  2. Add H_2O (l): Shifts left (blue).

  3. Increase temperature: shifts right (green).

  4. Add AgNO_3: Shifts left (blue).

Equilibrium Constant (K_c)

  • Predicts reaction yield at equilibrium.

  • Conditions altered to increase product yield.

  • K_c: Product to reactant concentration ratio at equilibrium.

  • K_c > 1: High product concentration (economically viable).

  • K_c < 1: Low product concentration (not economically viable).

Formula

For aA + bB \rightleftharpoons cC + dD :

K\frac \frac{[C]^c [D]^d} {[A]^a ⁣^b}

Factors Affecting K_c

  • Only temperature affects K_c.

  • Liquids and solids excluded (value = 1).

  • Gases and solutions included.

  • Use equilibrium concentrations to calculate.K_c.

  • K_c constant at a given temperature. Changes if temperature changes.

    • Increase K_c: Forward favoured.

    • Decrease K_c: Reverse favoured.

Examples of Equilibrium Constant Expressions

  1. 2 SO2 (g) + O2 (g) \rightleftharpoons 2 SO3 (g) Kc = \frac {[SO3]^2} {[SO2]^2 [O2]}

  2. CH3COOH (aq) + H2O (l) \rightleftharpoons CH*3COO^- (aq) + H3O^+ (aq) Kc = \frac {[CH3COO^-] [H3O^+]} {[CH3COOH]}

  3. CaCO3 (s) \rightleftharpoons CaO (s) + CO2 (g)
    Kc = [CO_2]

  4. MgO(s) + CO2(g) \rightleftharpoons MgCO3(s)
    Kc = \frac{1}{[CO_2]}

Example Calculation

For H2 (g) + I2 (g) \rightleftharpoons 2 HI (g) + energy:

  • [H_2] = 0.02 M

  • [I_2] = 0.02 M

  • [HI] = 0.16 M

Kc = \frac{[HI]^2} {[H2][I2]} = 64

RICE Tables

Method for calculating equilibrium concentrations.

Structure

Reactants

Products

Reaction

Initial

Change

Equilibrium

Modified Structure (Using Moles)

Reactants

Products

Reaction

Initial

Moles

Moles

Change

Moles

Moles

Equilibrium

Moles

Moles

Example 1

(NH4)2S(g) \rightleftharpoons H2S(g) + 2NH3(g)

Kc = \frac{[H2S][NH3]^2} {[(NH4)2S]} = 0.0057

Example 2

X (g) + Y (s) \rightleftharpoons Z (g) + R (g)

K_c = \frac {[Z][R]} {[X]}

[Z] = 0.0157 M

Equilibrium Graphs

Concentration vs. Time

Analyzing Concentration vs Time Graphs

2 SO2(g) + O2(g) \rightleftharpoons 2 SO_3(g) + energy

At equilibrium, temperature decreased: shifts right, products increase.

CO2(g) + energy \rightleftharpoons 2 CO(g)O₂(g) + O

At equilibrium, temperature decreased: shifts left, reactants increase.

Rate vs. Time

Description of oA with equal forward and reverse rates of rate graphs: reactants, volume, temperature, catalyst.