Effect of Concentration on Equilibrium (Higher Tier)

Le Chatelier’s Principle

Le Chatelier’s Principle states that if a system at dynamic equilibrium is subjected to a change in concentration, temperature, or pressure, the system will respond to oppose that change and restore a new equilibrium.

Dynamic equilibrium means the forward and reverse reactions continue at the same rate, so the concentrations of reactants and products remain constant.

When the concentration of a reactant or product changes, the position of equilibrium shifts to reduce that change, favouring either the forward or reverse reaction.

Effect of Concentration on Equilibrium

Changing the concentration of substances in a reversible reaction affects the position of equilibrium:

• Increasing the concentration of a reactant shifts the equilibrium to the right, producing more products.
• Increasing the concentration of a product shifts the equilibrium to the left, producing more reactants.
• Decreasing the concentration of a reactant shifts the equilibrium to the left, producing more reactants.
• Decreasing the concentration of a product shifts the equilibrium to the right, producing more products.

This shift happens because the system tries to restore balance by consuming the added substance or replacing the removed substance.

For example, consider the reversible reaction:

$$\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)$$

If the concentration of nitrogen gas ($\text{N}_2$) is increased, the equilibrium shifts to the right to produce more ammonia ($\text{NH}_3$). Conversely, if ammonia is removed, the equilibrium also shifts to the right to replace it.

This adjustment continues until a new equilibrium is established with different concentrations but the same rate of forward and reverse reactions.

For instance, if ammonia is added to the system, the equilibrium shifts left to reduce ammonia concentration by producing more nitrogen and hydrogen.

Example in main text: If the concentration of hydrogen gas in the reaction $\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3$ is increased, the system shifts right to produce more ammonia, using up the extra hydrogen.

Predicting Direction of Shift

To predict which way equilibrium will shift when concentration changes:

• Identify which substance's concentration has changed.
• If concentration of a reactant or product is increased, equilibrium shifts to consume it (opposite side).
• If concentration is decreased, equilibrium shifts to produce more of that substance (same side).

This principle helps predict how the system responds without needing detailed calculations.

For example, in the reaction:

$$\text{A} + \text{B} \rightleftharpoons \text{C} + \text{D}$$

If the concentration of C is increased, equilibrium shifts left to use up C and form more A and B.

If B is removed, equilibrium shifts left to produce more B by breaking down C and D.

This concept is crucial for understanding how to control chemical reactions in industry and laboratory settings.

Example: Consider the equilibrium reaction:

$$\text{H}_2(g) + \text{I}_2(g) \rightleftharpoons 2\text{HI}(g)$$

If the concentration of hydrogen gas ($\text{H}_2$) is increased, the system shifts right to produce more hydrogen iodide ($\text{HI}$).

Practical Applications

In industry, controlling concentrations is vital to maximise product yield in reversible reactions.

For example, the Haber process for ammonia production uses high concentrations of nitrogen and hydrogen to shift equilibrium towards ammonia.

By continually removing ammonia as it forms, the equilibrium shifts right, increasing yield.

Optimising concentrations reduces waste and improves efficiency, saving costs and energy.

Understanding how concentration affects equilibrium helps chemists design better processes and control reaction conditions.