Equilibrium and Le Chatelier’s Principle

Full Notes

What is Chemical Equilibrium?

Reversible reactions can go forward (reactants → products) and backward (products → reactants).

Dynamic equilibrium is reached in a closed system when:

• The rate of the forward reaction = rate of the reverse reaction.
• The concentrations of reactants and products remain constant (but are not necessarily equal).

It is dynamic because both reactions continue, but there is no overall change in the amounts.

Example: The Haber Process (ammonia production)

Forward: N₂ + H₂ → NH₃
Reverse: NH₃ → N₂ + H₂

At equilibrium, forward and reverse reactions occur at the same rate, so concentrations remain constant.

Matt’s exam tip

Remember that reactions are still occuring at equilibrium, it is just that they both happen at the same rate. Also, closed system means no matter (reactants or products) can get in or out of the system.

Understanding Equilibrium Through Graphs

Graphs can show how a chemical system changes over time as it approaches equilibrium. Whether we’re looking at concentration, rate of reaction, or partial pressure, the overall pattern is the same: dynamic changes at first, followed by a steady state.

• Concentration vs. Time (Left graph): The concentration of reactants (green) decreases, while the concentration of products (blue) increases. Once equilibrium is reached, both levels flatten — they remain constant, though the reactions are still happening in both directions.
• Rate of Reaction vs. Time (Middle graph): Initially, the forward reaction is faster than the reverse. Over time, the reverse reaction speeds up as more product forms. At equilibrium, both rates become equal, meaning no net change occurs — the system is dynamic, but balanced.
• Partial Pressure vs. Time (Right graph): For gaseous reactions, partial pressures behave just like concentrations (see K_p). The graph shows reactants’ pressures falling and products’ rising until both level out. This is another way to observe equilibrium being established.

Le Chatelier’s Principle

“If a system at equilibrium is subjected to a change, the position of equilibrium will shift to oppose that change.”

There are several factors that can affect equilibrium position:

1. Changing Concentration

Increasing reactant concentration shifts equilibrium right → more products.

Increasing product concentration shifts equilibrium left → more reactants.

2. Changing Pressure (for Gaseous Equilibria)

Increasing pressure shifts equilibrium towards the side with fewer gas molecules.

Decreasing pressure shifts equilibrium towards the side with more gas molecules.

No effect if the number of gas molecules is the same on both sides.

Example: Haber Process

Increasing pressure shifts equilibrium right → increases NH₃ yield.

3. Changing Temperature

Increasing temperature favours the endothermic direction (+ΔH).
Decreasing temperature favours the exothermic direction (−ΔH).

Example: Haber Process

Forward reaction is exothermic (−ΔH).
Increasing temperature shifts equilibrium left → reduces NH₃ yield.
Decreasing temperature shifts equilibrium right → increases NH₃ yield.

4. Effect of a Catalyst on Equilibrium

Catalysts do not affect the equilibrium position.

They increase the rate of both forward and reverse reactions equally, so equilibrium is reached faster, but yield is unchanged.

Example: Iron catalyst in the Haber Process speeds up NH₃ production but does not change equilibrium yield.

Summary

• Dynamic equilibrium occurs when forward and reverse reactions occur at the same rate and concentrations remain constant.
• Le Chatelier’s Principle predicts how equilibrium shifts to oppose changes in conditions.
• Concentration: Increasing reactants shifts right; increasing products shifts left.
• Pressure: Higher pressure favours the side with fewer moles of gas.
• Temperature: Higher T favours endothermic; lower T favours exothermic.
• Catalysts: Do not shift equilibrium but reach it faster.