How does Le Chatelier's Principle apply to the Haber process?

The Haber process, $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$, $\Delta H = -92.4\,\text{kJ/mol}$, is a classic example. To maximize ammonia yield, we apply Le Chatelier's Principle: high pressure (shifts to fewer moles of gas, i.e., product side), low temperature (exothermic reaction, favors product formation), and continuous removal of ammonia (shifts to product side to replenish). However, a compromise temperature is used (around $400-450^{\circ}C$) because very low temperatures make the reaction too slow, even with a catalyst.

Why does adding an inert gas at constant volume not affect equilibrium?

When an inert gas is added to a reaction mixture at equilibrium in a fixed volume, the total pressure of the system increases. However, the partial pressures of the reacting gases remain unchanged because their concentrations (moles per unit volume) do not change. Since the equilibrium position is determined by the partial pressures (or concentrations) of the reacting species, and these haven't changed, the equilibrium remains undisturbed. Le Chatelier's Principle only applies to changes in the *partial* pressures of the reacting components.

Can a change in concentration affect the equilibrium constant (K)?

No, a change in concentration of reactants or products does not affect the value of the equilibrium constant, K. K is a constant for a given reaction at a specific temperature. While changing concentrations will cause the equilibrium to shift to re-establish balance, the ratio of products to reactants at the *new* equilibrium state will still be equal to the original K value. Only temperature changes can alter the numerical value of the equilibrium constant.

How do you determine the effect of pressure on reactions involving solids or liquids?

For reactions involving only solids and/or liquids, changes in external pressure have a negligible effect on the equilibrium position. This is because solids and liquids are virtually incompressible, meaning their volumes and thus their concentrations are hardly affected by pressure changes. Le Chatelier's Principle regarding pressure primarily applies to gaseous systems where changes in pressure significantly alter the concentrations (or partial pressures) of the reacting species due to volume changes.

What is the difference between an exothermic and endothermic reaction in terms of temperature effect?

For an exothermic reaction ($\Delta H 0$), heat is a reactant. Increasing temperature shifts the equilibrium to the right (product side) to consume the added heat, increasing K. Decreasing temperature shifts it to the left (reactant side) to produce heat, decreasing K.

Effect of Concentration, Pressure and Temperature

Chemistry

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Version 1Updated 22 Mar 2026

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Le Chatelier's Principle states that if a change of condition is applied to a system in equilibrium, the system will shift in a direction that relieves the stress. This fundamental principle is crucial for understanding how chemical equilibria respond to external perturbations, specifically changes in concentration of reactants or products, total pressure (for gaseous systems), and temperature. By…

Quick Summary

Le Chatelier's Principle is a fundamental concept in chemical equilibrium, stating that a system at equilibrium will counteract any applied stress to re-establish a new equilibrium. Three primary stresses are concentration, pressure, and temperature.

Increasing reactant concentration or decreasing product concentration shifts equilibrium towards products. Conversely, decreasing reactant concentration or increasing product concentration shifts it towards reactants.

For gaseous reactions, increasing pressure shifts equilibrium towards the side with fewer moles of gas, while decreasing pressure shifts it towards the side with more moles of gas. Temperature changes affect the equilibrium constant (K) itself.

For exothermic reactions, increasing temperature shifts equilibrium to reactants (decreasing K), while for endothermic reactions, increasing temperature shifts it to products (increasing K). Catalysts accelerate the attainment of equilibrium but do not alter its position.

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Key Concepts

Concentration Effect and Reaction Quotient

The reaction quotient, $Q_c = \frac{[Products]^x}{[Reactants]^y}$ , helps predict the direction of shift. If…

Pressure Effect and Moles of Gas

Pressure changes are significant only for gaseous reactions where the number of moles of gas changes. The…

Temperature Effect and Enthalpy Change

Temperature is the only factor that changes the equilibrium constant $K$ . For exothermic reactions ($\Delta H…

Concentration — Add reactant/remove product

\implies

shift right. Remove reactant/add product

\implies

shift left.\n- Pressure (gases only): Increase P

\implies

shift to fewer moles of gas. Decrease P

\implies

shift to more moles of gas. (No effect if

\Delta n_g = 0

).\n- Temperature: Exothermic (

\Delta H < 0

): Increase T

\implies

shift left (K decreases). Decrease T

\implies

shift right (K increases). Endothermic (

\Delta H > 0

): Increase T

\implies

shift right (K increases). Decrease T

\implies

shift left (K decreases).\n- Catalyst: No effect on equilibrium position or K; only speeds up attainment of equilibrium.\n- Inert Gas (constant V): No effect on equilibrium position.

To remember Le Chatelier's Principle effects: 'CPT'

Concentration: Consume what's Crowded, Create what's Clear.

Pressure: Push to People (moles) Poorer (fewer).

Temperature: Take Thermal (heat) side for Temp increase; Take Thermal side for Temp decrease (opposite of what's added/removed).