Le Chatelier's Principle

Le Chatelier's Principle describes how a system already at equilibrium responds when it's disturbed: it shifts to counteract the disturbance and re-establish equilibrium. This subtopic covers the three stresses that trigger a shift — changes in concentration, pressure (or volume), and temperature — and how to predict which direction the equilibrium moves in each case.

Key Takeaways

• Le Chatelier's Principle: a system at equilibrium shifts to counteract an applied stress and re-establish equilibrium.

• Concentration: adding reactant/removing product drops Q below Keq and shifts the system forward; removing reactant/adding product raises Q above Keq and shifts the system in reverse.

• Pressure/volume (gas-phase only): increasing pressure (decreasing volume) shifts toward fewer moles of gas; decreasing pressure (increasing volume) shifts toward more moles of gas. Solids and liquids are unaffected.

• Temperature is the one stress that changes Keq itself: for endothermic reactions, heat acts as a reactant, so increasing temperature favors products; for exothermic reactions, heat acts as a product, so increasing temperature favors reactants.

• Catalysts don't shift equilibrium or change Keq — they only speed up how fast equilibrium is reached.

What Le Chatelier's Principle Says

Le Chatelier's Principle states that if a stress is applied to a system at equilibrium, the system adjusts to counteract that stress and re-establish equilibrium. "Stress" here means a change in concentration, pressure/volume, or temperature. The principle applies to all three, but each one shifts the equilibrium through a different mechanism.

Concentration Changes

When the concentration of a reactant or product is changed, the system shifts to counteract the change and restore equilibrium. This case connects directly to the reaction quotient Q: changing a concentration moves Q away from Keq, and the system reacts by shifting until Q equals Keq again.

• Adding reactants or removing products: Q becomes less than Keq. The system shifts forward, producing more products, until Q rises back to equal Keq.

  • Example: for N₂(g) + 3H₂(g) ⇌ 2NH₃(g), adding more N₂ shifts the system right, producing more NH₃.

• Removing reactants or adding products: Q becomes greater than Keq. The system shifts in reverse, producing more reactants, until Q falls back to equal Keq.

  • Example: for the same reaction, adding NH₃ shifts the system left, producing more N₂ and H₂.

Pressure and Volume Changes

Changes in pressure (and volume) affect gas-phase equilibria by shifting the system toward whichever side has fewer or more moles of gas.

• Increased pressure (decreased volume): the system shifts toward the side with fewer moles of gas, which lowers the pressure back down.

  • Example: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) has 4 moles of gas on the reactant side and 2 moles on the product side. Increasing the pressure shifts the equilibrium right, toward the side with fewer moles of gas.

• Decreased pressure (increased volume): the system shifts toward the side with more moles of gas, which raises the pressure back up.

  • Example: for the same reaction, decreasing the pressure shifts the equilibrium left, toward the side with more moles of gas.

MCAT Callout — Gases Only: This stress applies only to gaseous species. Solids and liquids are essentially incompressible, so their concentrations don't change with pressure or volume changes — only the moles of gas on each side matter.

Temperature Changes

Temperature is different from the other two stresses: changing concentration or pressure shifts where the equilibrium sits without changing Keq itself, but changing temperature actually changes the value of Keq, because temperature is tied to the reaction's enthalpy change (ΔH). Which direction the shift goes depends on whether the reaction is endothermic or exothermic.

• Endothermic reactions (ΔH > 0) absorb heat, so heat can be treated as a reactant.

  • Increasing temperature (adding heat) shifts the equilibrium toward products.

  • Decreasing temperature (removing heat) shifts the equilibrium toward reactants.

  • Example: for the endothermic reaction N₂O₄(g) ⇌ 2NO₂(g), increasing temperature shifts the equilibrium right, producing more NO₂; decreasing temperature shifts it left, producing more N₂O₄.

• Exothermic reactions (ΔH < 0) release heat, so heat can be treated as a product.

  • Increasing temperature (adding heat) shifts the equilibrium toward reactants.

  • Decreasing temperature (removing heat) shifts the equilibrium toward products.

  • Example: for the exothermic reaction 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), increasing temperature shifts the equilibrium left, producing more SO₂ and O₂; decreasing temperature shifts it right, producing more SO₃.

MCAT Callout — Catalysts: a catalyst speeds up the forward and reverse reactions equally, so it shortens the time it takes to reach equilibrium — but it does not shift the equilibrium position or change the value of Keq. Only concentration, pressure/volume, and temperature changes trigger a Le Chatelier shift.

Common MCAT Mistakes

• Thinking a catalyst shifts equilibrium. A catalyst speeds up both the forward and reverse reactions equally — it changes how fast equilibrium is reached, never where the equilibrium position or Keq itself ends up.

• Applying the pressure/volume stress to solids or liquids. This stress only affects gas-phase species. Increasing pressure has no meaningful effect on the concentration of a solid or liquid, so only the balance of gas moles on each side matters.

• Forgetting that temperature is the one stress that changes Keq. Concentration and pressure/volume changes shift the position of equilibrium but leave Keq unchanged; temperature changes actually alter the numerical value of Keq itself.

• Mixing up which direction heat favors. Treat heat as a reactant for an endothermic reaction and as a product for an exothermic reaction — then apply the same logic used for any other reactant/product concentration change to figure out which way the shift goes.

MCAT-Style Concept Check

Question: For the exothermic reaction N₂(g) + 3H₂(g) ⇌ 2NH₃(g), which single change would shift the equilibrium toward more NH₃?

• A) Increasing the temperature

• B) Increasing the pressure (decreasing the volume)

• C) Removing N₂ from the system

• D) Adding NH₃ to the system

Answer: B

Explanation: The reactant side has 4 moles of gas (1 N₂ + 3 H₂) and the product side has 2 moles of gas (2 NH₃). Increasing the pressure shifts the equilibrium toward the side with fewer moles of gas — the product side — producing more NH₃. Increasing temperature would shift an exothermic reaction toward reactants (A is wrong), removing N₂ would shift the system toward reactants to replace it (C is wrong), and adding NH₃ would shift the system toward reactants to consume the excess (D is wrong).

FAQ

Does a catalyst shift equilibrium toward products?

No. A catalyst lowers the activation energy for both the forward and reverse reactions by the same amount, so it speeds up how quickly the system reaches equilibrium — but it doesn't change the equilibrium position or the value of Keq.

Why does pressure only matter for gases in Le Chatelier's Principle?

Pressure and volume changes shift equilibrium by changing the concentration of gas-phase species — compressing a gas raises its concentration in a way that changing pressure doesn't do for a solid or liquid, since solids and liquids are essentially incompressible.

How do I know which direction a temperature change shifts an equilibrium?

Treat heat as a reactant in an endothermic reaction (ΔH > 0) or as a product in an exothermic reaction (ΔH < 0), then apply the same logic as any other concentration change: adding "heat" to the reactant side shifts the system forward, adding it to the product side shifts the system in reverse.

Is Le Chatelier's Principle the same thing as the reaction quotient Q?

They're closely related but not identical. Q vs. Keq is the tool for predicting shift direction after any stress changes a concentration; Le Chatelier's Principle is the broader statement that a stressed system shifts to counteract the stress, covering concentration, pressure/volume, and temperature changes alike.