13.3 Shifting Equilibria: Le Châtelier’s Principle

Chemical equilibrium is a dynamic balance between forward and reverse reactions. When conditions change, the system responds in a predictable way to restore balance. Le Châtelier's Principle captures this idea: if you disturb a system at equilibrium, it shifts to partially counteract that disturbance.

This topic covers the major factors that can shift equilibrium (concentration, temperature, and pressure), how to use the reaction quotient to predict which way a reaction will shift, and why catalysts don't actually change the equilibrium at all.

Factors Affecting Equilibrium

Dynamic Equilibrium and Reaction Quotient

At dynamic equilibrium, the forward and reverse reactions are still happening, but they proceed at equal rates. The concentrations of reactants and products stay constant, even though molecules are continuously reacting in both directions.

The reaction quotient (Q) has the same formula as the equilibrium constant $K$, but you can calculate it at any point during a reaction, not just at equilibrium. It tells you where the reaction currently stands compared to where it "wants" to be.

Comparing Q to K predicts which direction the reaction will shift:

• If $Q < K$: too few products relative to equilibrium. The reaction shifts right (toward products).
• If $Q > K$: too many products relative to equilibrium. The reaction shifts left (toward reactants).
• If $Q = K$: the system is already at equilibrium. No net shift occurs.

Effects of Concentration on Equilibrium

The core idea here is straightforward: adding more of a substance pushes the equilibrium away from it, and removing a substance pulls the equilibrium toward replacing it.

• Increasing the concentration of a reactant shifts equilibrium to the right (toward products). The forward reaction speeds up to consume the added reactant until a new equilibrium is reached.
• Decreasing the concentration of a reactant shifts equilibrium to the left (toward reactants). The reverse reaction speeds up to replenish what was removed.
• Increasing the concentration of a product shifts equilibrium to the left (toward reactants). The system works to consume the excess product.
• Decreasing the concentration of a product shifts equilibrium to the right (toward products). The system produces more of the product that was removed.

For example, consider the reaction $N_2 + 3H_2 \rightleftharpoons 2NH_3$. If you add extra $H_2$, the system shifts right and produces more $NH_3$. If you continuously remove $NH_3$ as it forms, the equilibrium keeps shifting right to replace it. This is actually how industrial ammonia production (the Haber process) maximizes yield.

Temperature Changes in Chemical Equilibria

Temperature is unique among these factors because it's the only one that actually changes the value of $K$ itself. The trick to predicting temperature effects is to treat heat as if it were a reactant or product.

Exothermic reactions ($\Delta H < 0$) release heat. Think of heat as a product:

$\text{Reactants} \rightleftharpoons \text{Products} + \text{heat}$

• Increasing temperature adds heat, so the system shifts left (toward reactants) to consume it. $K$ decreases.
• Decreasing temperature removes heat, so the system shifts right (toward products) to replace it. $K$ increases.

Endothermic reactions ($\Delta H > 0$) absorb heat. Think of heat as a reactant:

$\text{Reactants} + \text{heat} \rightleftharpoons \text{Products}$

• Increasing temperature adds heat, so the system shifts right (toward products). $K$ increases.
• Decreasing temperature removes heat, so the system shifts left (toward reactants). $K$ decreases.

Pressure Impacts on Gas-Phase Equilibrium

Pressure changes only matter for reactions involving gases, and only when the total moles of gas differ between the two sides of the equation.

When pressure increases (or volume decreases), the system shifts toward the side with fewer moles of gas. This makes sense because fewer gas molecules in the same space means lower pressure, which partially counteracts the change.

When pressure decreases (or volume increases), the system shifts toward the side with more moles of gas.

For example, in $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$, there are 4 moles of gas on the left and 2 on the right. Increasing pressure shifts the reaction to the right (toward $NH_3$) because that side has fewer gas molecules.

Pressure changes have no effect on equilibrium when:

• The total moles of gas are equal on both sides of the equation
• The reaction involves only solids or liquids (their volumes don't change meaningfully with pressure)

Catalysts and Equilibrium

A catalyst speeds up both the forward and reverse reactions by the same factor. Because both rates increase equally, the system reaches equilibrium faster, but the final ratio of products to reactants doesn't change.

Catalysts do not shift the position of equilibrium, and they do not change the value of $K$. They only affect how quickly equilibrium is reached.