Short Answer

Le Chatelier’s principle states that when a system at equilibrium is disturbed by a change in temperature, pressure, or concentration, the system adjusts itself to counteract that change and restore a new equilibrium. This principle helps predict how equilibrium shifts when conditions change.

For example, adding more reactant shifts equilibrium toward the products, while increasing temperature favours the endothermic direction. Le Chatelier’s principle is widely used in chemistry to understand reaction behaviour and to optimise industrial processes.

Detailed Explanation :

Le Chatelier’s Principle

Le Chatelier’s principle is a fundamental rule in chemistry that explains how an equilibrium system responds to changes in external conditions. When a reversible reaction reaches equilibrium, the forward and backward reactions occur at equal rates. However, if a disturbance occurs—such as a change in temperature, pressure, or concentration—the system temporarily moves out of equilibrium.

According to Le Chatelier’s principle, the system will shift in a direction that reduces or opposes the effect of the disturbance. This shift continues until a new equilibrium is established. The principle is extremely useful for predicting how equilibrium will respond to changes and is widely applied in industrial chemistry to maximize product yield.

1. Effect of Change in Concentration

Changing the concentration of reactants or products disturbs equilibrium.

1. a) Adding Reactants

• System has too many reactant molecules.
• To reduce the effect, equilibrium shifts toward products.
• Forward reaction rate increases.

1. b) Removing Reactants

• System shifts toward reactants to replace the removed amount.

1. c) Adding Products

• System shifts toward reactants to reduce extra products.

1. d) Removing Products

• System shifts toward products to replace what was removed.

These shifts help the system regain stability.

2. Effect of Change in Temperature

Temperature affects equilibrium depending on whether the reaction is exothermic or endothermic.

1. a) Increasing Temperature

• System shifts to the endothermic side (absorbs heat).
• Helps reduce the added heat.

1. b) Decreasing Temperature

• System shifts to the exothermic side (releases heat).
• Helps replace lost heat.

Temperature changes also affect the value of the equilibrium constant (K), unlike concentration and pressure changes.

3. Effect of Change in Pressure

(Only applies to gaseous reactions)

1. a) Increasing Pressure

• System shifts to the side with fewer gas molecules.
• This reduces pressure.

1. b) Decreasing Pressure

• System shifts to the side with more gas molecules.
• This restores higher pressure.

Example:
N₂ + 3H₂ ⇌ 2NH₃
Left side: 4 gas molecules → Right side: 2 gas molecules
Increasing pressure favours ammonia formation.

4. Effect of Volume Change

Volume and pressure are inversely related.

• Decreasing volume increases pressure → equilibrium shifts to fewer gas molecules
• Increasing volume decreases pressure → equilibrium shifts to more gas molecules

This is another form of pressure influence.

5. Effect of Adding an Inert Gas

At Constant Volume

• No change in concentrations
• No shift in equilibrium position

At Constant Pressure

• Volume increases
• Concentrations decrease
• Equilibrium shifts toward the side with more gas molecules

6. Effect of Catalyst

A catalyst speeds up both the forward and backward reactions equally.

Important Point

• Catalyst does NOT shift the equilibrium position
• It only helps the system reach equilibrium faster

Thus, catalysts improve efficiency without changing the balance.

7. Real-life Examples of Le Chatelier’s Principle
8. a) Soda Bottle

When opened, pressure decreases → equilibrium shifts → dissolved CO₂ escapes as bubbles.

1. b) Haber Process

High pressure and moderate temperature favour ammonia formation according to Le Chatelier’s principle.

1. c) Saturated Sugar Solution

Adding more sugar (solid) does not shift equilibrium because solids do not affect equilibrium.

8. Industrial Applications

Industries use Le Chatelier’s principle to maximise product yield.

Examples

• Increasing pressure in ammonia production
• Increasing temperature in endothermic reactions
• Removing products continuously to shift equilibrium forward

This principle helps make industrial processes efficient and economical.

Conclusion

Le Chatelier’s principle explains how an equilibrium system responds to external changes. When concentration, temperature, or pressure is altered, the system shifts in a direction that reduces the effect of that change and restores equilibrium. This principle is extremely helpful in predicting reaction behaviour and optimizing industrial production. It remains one of the most important concepts in chemical equilibrium and reaction control.