Factors Affecting Equilibria

Full Notes

Introduction

Chemical equilibrium is a dynamic state where forward and reverse reactions occur at the same rate. When an external change is applied, the system adjusts to restore equilibrium.

This response is predicted by Le Chatelier’s Principle, while the reaction quotient (Q_c) helps compare the current state to the equilibrium condition, allowing us to predict how the system will respond.

How Are Q and K Related to Le Châtelier’s Principle?

Le Châtelier’s Principle explains how a system at equilibrium responds to stress: it shifts in a direction that opposes the disturbance and works to re-establish equilibrium.

We can predict the direction the equilibrium will shift in using Q (reaction quotient) and K (equilibrium constant, K_c or K_p).

Q vs. K: Predicting the Direction of Shift

K (for example K_c or K_p) is the ratio of product to reactant concentrations at equilibrium.

Q is the ratio of those concentrations at any moment, even if the system is not at equilibrium.

When a system is disturbed:

• If Q < K
  Too many reactants, not enough products  and the system shifts right to form more products.
• If Q > K
  Too many products, not enough reactants  and the system shifts left to form more reactants.
• If Q = K
  The system is already at equilibrium  and no shift occurs.

In other words,Q tells us where we are, and K tells us where we’re going.

How This Links to Le Châtelier’s Principle

Le Châtelier’s Principle says:
When a system at equilibrium is disturbed, it will shift in the direction that opposes the change and restores equilibrium.

This is exactly what happens when Q ≠ K. The system adjusts:

• Concentrations or pressures shift
• Q changes
• Until Q = K again — and the system reaches a new equilibrium

Examples of Disturbances

What About Temperature?

Temperature changes actually change the value of K, not just Q.

• For endothermic reactions (ΔH > 0):
   Increasing T means K increases = system shifts right
• For exothermic reactions (ΔH < 0):
   Increasing T means K decreases = system shifts left

After a temperature change:

• Q remains the same initially.
• But K has changed.
• The system is now out of equilibrium (Q ≠ K), and will shift to re-establish Q = K.

How Q and K Predict Reaction Direction

6.8.1 Effect of Concentration Change

If the concentration of a reactant or product is changed, the system will shift to oppose the change:

• Adding a reactant or product shifts the equilibrium to consume the added substance.
• Removing a component shifts the equilibrium to replace it.

Example: For the reaction:

H₂(g) + I₂(g) ⇌ 2HI(g)

• Adding H₂ or I₂ → shifts equilibrium to the right (produces more HI)
• Removing HI → shifts right (to make more HI)
• Adding HI → shifts left (to consume excess HI)

Experimental Illustration: In the reaction:

Fe³⁺(aq) + SCN⁻(aq) ⇌ Fe(SCN)²⁺(aq)

• Adding Fe³⁺ or SCN⁻ deepens the red colour → shift to the right
• Diluting the solution (adding water) lightens the colour → shift to the left

6.8.2 Effect of Pressure Change

This effect is relevant only for reactions involving gases.

• For reactions where the number of gas molecules changes (Δn ≠ 0):
  • Increasing pressure shifts the equilibrium to the side with fewer gas molecules
  • Decreasing pressure shifts it to the side with more gas molecules
• If Δn = 0 (equal number of moles on both sides), a change in pressure has no effect.

For Example: Haber Process

4 moles (N₂ + 3H₂) ⇌ 2 moles (NH₃).
Higher pressure shifts equilibrium right, increasing NH₃ yield.

Role of Q_c in Pressure Changes

The reaction quotient (Q_c) helps determine the direction of the shift:

• If Q_c < K_c, the reaction proceeds forward
• If Q_c > K_c, the reaction shifts backward
• If Q_c = K_c, the system is at equilibrium

When pressure increases:

• The concentrations of all gases increase
• If the increase disproportionately affects the side with more gas molecules, then Q_c > K_c → shift left
• If the increase favours the side with fewer gas molecules, then Q_c < K_c → shift right

6.8.3 Effect of Inert Gas Addition

• At constant volume: Adding an inert gas does not affect the partial pressures of the reacting gases → no effect on equilibrium.
• At constant pressure: Volume increases, so partial pressures decrease. The system shifts depending on the change in number of gas molecules (Δn):
  • If one side has more gas molecules, the equilibrium shifts in that direction to counteract the drop in pressure.

6.8.4 Effect of Temperature Change

Temperature affects both the position of equilibrium and the value of the equilibrium constant (K).

• For endothermic reactions (ΔH > 0): Increasing temperature shifts equilibrium to the right
• For exothermic reactions (ΔH < 0): Increasing temperature shifts equilibrium to the left

For Example: In the Haber Process:

N₂(g) + 3H₂(g) ⇌ 2NH₃(g); ΔH = – kJ/mol
Increasing temperature shifts equilibrium left, reducing NH₃ yield.
Decreasing temperature shifts equilibrium right, increasing NH₃ yield.

6.8.5 Effect of a Catalyst

• Catalysts:
  • Increase the rate of both forward and reverse reactions equally
  • Do not affect the position of equilibrium
  • Do not change the value of the equilibrium constant (K)
• Their role is simply to help the system reach equilibrium faster.

Summary

• Le Chatelier’s Principle predicts shifts that oppose external changes.
• Use Q versus K to determine the direction of shift.
• Pressure changes matter only when Δn is not zero.
• Temperature changes the position and the value of K.
• Catalysts speed up attainment of equilibrium without changing K.