A-Level Chemistry: Understanding Equilibrium and Le Chatelier's Principle

Le Chatelier's principle sounds abstract until you realize it's just the chemical world's way of saying "the system resists change." A reaction at equilibrium is like a tug-of-war that's evenly balanced. Pull one side, and the balance shifts back toward center. Change temperature, pressure, or concentration, and the equilibrium position moves to counteract that change. (This guide has been for the 2025-26 syllabus.)

Knowing the principle isn't enough. You need to predict which direction the equilibrium shifts when conditions change, calculate how the equilibrium constant responds to temperature changes, and explain what's actually happening at the molecular level. This guide walks you through all of it, with the focus on understanding why these shifts happen, not just memorizing which way to shift.

Equilibrium: The Starting Point

Before understanding Le Chatelier's principle, you need equilibrium clarity.

A reaction reaches equilibrium when the forward and reverse reaction rates are equal. At that point, the concentrations of reactants and products aren't changing — not because the reaction stopped, but because they're being produced and consumed at equal rates.

Key: Equilibrium is a dynamic state. Reactions are still happening in both directions; they're just balanced.

The equilibrium constant (K): For the reaction aA + bB ⇌ cC + dD, K = [C]^c[D]^d / [A]^a[B]^b. The equilibrium constant is a number that tells you the ratio of products to reactants at equilibrium. It's constant for a given reaction at a given temperature (emphasis on temperature — this matters).

What Does Le Chatelier's Principle Actually Mean?

If you change the conditions of a system at equilibrium, the system shifts to counteract that change and re-establish equilibrium.

Translation: If you increase the concentration of a product, the equilibrium shifts left (toward reactants) to decrease that product. If you increase pressure, the equilibrium shifts toward whichever side has fewer moles of gas. If you increase temperature, the equilibrium shifts in the endothermic direction.

The system isn't intelligent. It's not "trying" to resist. What's really happening is that when you change conditions, the quotient Q (the current ratio of products to reactants) no longer equals K. The system re-equilibrates by shifting in whichever direction brings Q back to K.

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The Three Main Stressors: Concentration, Pressure, and Temperature

Stressor 1: Changing Concentration

Rule: If you increase the concentration of a reactant or product, the equilibrium shifts away from that substance. If you decrease it, the equilibrium shifts toward that substance. Explore our detailed guide on chemistry career paths for more tips.

Why: When you increase a reactant, Q changes and no longer equals K. The system shifts right (toward products) to decrease the new reactant concentration and bring Q back to K. Vice versa for decreasing.

Example: For the equilibrium N₂ + 3H₂ ⇌ 2NH₃:

• If you add more N₂, the equilibrium shifts right (toward NH₃) to consume some of that N₂.
• If you remove some NH₃, the equilibrium shifts right (toward NH₃) to replace what you removed.
• The K doesn't change; only the concentrations change until a new equilibrium is reached.

Important: If you add an inert gas (one not in the equilibrium) at constant volume, nothing happens — the equilibrium doesn't shift. The partial pressures and concentrations of the reacting species don't change.

Stressor 2: Changing Pressure (or Volume)

Rule: If you increase pressure (by decreasing volume), the equilibrium shifts toward the side with fewer moles of gas. If you decrease pressure, it shifts toward the side with more moles of gas.

Why: Pressure increases when you compress the system (decrease volume). The system shifts to reduce the number of gas molecules, which lowers the pressure back down.

Example: For N₂ + 3H₂ ⇌ 2NH₃ (4 moles on left, 2 on right):

• Increase pressure → equilibrium shifts right (fewer moles of gas = lower pressure).
• Decrease pressure → equilibrium shifts left (more moles of gas = higher pressure).

Example: For 2NO₂ ⇌ N₂O₄ (2 moles on left, 1 on right):

• Increase pressure → equilibrium shifts right (1 mole is fewer than 2).
• Decrease pressure → equilibrium shifts left.

Exception: If the reaction has equal moles of gas on both sides, pressure changes don't shift the equilibrium. Example: H₂ + I₂ ⇌ 2HI (2 moles on each side). Changing pressure has no effect.

Stressor 3: Changing Temperature

Temperature is the only stressor that changes K itself. Concentration and pressure stressors shift the equilibrium without changing K; temperature stressors actually change K.

For exothermic reactions (ΔH < 0):

• Increase temperature → K decreases → equilibrium shifts left (toward reactants).
• Decrease temperature → K increases → equilibrium shifts right (toward products).

For endothermic reactions (ΔH > 0):

• Increase temperature → K increases → equilibrium shifts right (toward products).
• Decrease temperature → K decreases → equilibrium shifts left (toward reactants).

Why: Think of the reaction equation as if heat is a "chemical" on one side:

Exothermic: Reactants → Products + Heat. Adding heat is like adding a product. Equilibrium shifts left (toward reactants) to consume the added heat.

Endothermic: Reactants + Heat → Products. Adding heat is like adding a reactant. Equilibrium shifts right (toward products) to consume the added heat.

Example: N₂ + 3H₂ ⇌ 2NH₃ is exothermic (ΔH = -92 kJ/mol).

• Increase temperature → K decreases → equilibrium shifts left → less NH₃ produced.
• This is why the Haber process (making ammonia industrially) uses moderately high temperature (400°C) — too low and the reaction is too slow; too high and equilibrium favors reactants. Understanding redox equilibrium principles in industrial processes demonstrates how electron transfer drives industrial chemistry.

How to Predict Equilibrium Shifts: The Systematic Approach

Step 1: Write the equilibrium equation and identify whether it's exothermic or endothermic.

Step 2: Identify the stressor (concentration change, pressure change, or temperature change).

Step 3: For concentration and pressure stressors, use Le Chatelier's principle to predict which direction the equilibrium shifts. Don't ask "which side has more product?" Ask "which direction counteracts the change I made?"

Step 4: For temperature stressors, remember that exothermic reactions shift left when heated and right when cooled. Endothermic reactions are opposite.

Step 5: State the shift and what it means for the concentrations of reactants and products. "The equilibrium shifts right, so [NH₃] increases and [N₂] and [H₂] decrease."

Example problem: The reaction 2NO(g) + O₂(g) ⇌ 2NO₂(g) is exothermic. If you increase the volume of the container, what happens to the equilibrium?

1. Identify: Exothermic. 3 moles on left, 2 on right.
2. Stressor: Increase volume = decrease pressure.
3. Apply principle: Decreasing pressure shifts equilibrium toward the side with more moles of gas (left, 3 moles). Equilibrium shifts left.
4. Result: [NO₂] decreases, [NO] and [O₂] increase.

Calculating Equilibrium Concentrations After a Shift

After you've predicted the shift direction, you can calculate new equilibrium concentrations using an ICE table (Initial, Change, Equilibrium). When mastering these calculations, understanding strategic study approaches helps you apply these principles systematically across different problem types. Learn more in our guide on master the IB chemistry syllabus a perfect.

Example: For N₂ + 3H₂ ⇌ 2NH₃, suppose at equilibrium [N₂] = 0.5 M, [H₂] = 1.5 M, [NH₃] = 1.0 M. K = 1.6. If you add 0.2 M of N₂, what are the new equilibrium concentrations?

1. After adding N₂: [N₂] = 0.7 M (the others haven't changed yet).
2. Calculate Q: Q = [NH₃]²/([N₂][H₂]³) = 1.0²/(0.7·1.5³) ≈ 0.42. Q < K, so equilibrium shifts right.
3. Set up ICE table for the shift:

N₂	+	3H₂	⇌	2NH₃
0.7		1.5		1.0
-x		-3x		+2x
0.7 - x		1.5 - 3x		1.0 + 2x

4. Use K to solve for x: 1.6 = (1.0 + 2x)² / ((0.7 - x)(1.5 - 3x)³). Solve (algebra). Once you have x, you have the new equilibrium concentrations. For mastering calculation techniques in chemistry, laboratory reporting practices that emphasize precise calculations transfer to equilibrium problem-solving. For more on this, see our guide on mastering organic chemistry.

Need help setting up and solving ICE tables after a shift? Work with a chemistry tutor who can teach you the systematic approach to equilibrium calculations → You may also find our resource on use chemistry AI solver for effective IB helpful.

Common Mistakes and How to Avoid Them

1. Confusing which direction to shift for pressure changes. Remember: the equilibrium shifts toward the side with FEWER moles of gas when pressure increases.

2. Forgetting that only temperature changes affect K. Concentration and pressure changes shift the equilibrium without changing K. Temperature changes both.

3. Adding an inert gas at constant volume and expecting a shift. Inert gases don't participate and don't change the partial pressures of the reacting species. No shift.

4. Not identifying whether a reaction is exothermic or endothermic before predicting temperature effects. This is crucial. Get it backwards and your prediction is backwards.

5. Forgetting that equilibrium is dynamic. When you shift the equilibrium, reactions are still happening in both directions. Concentrations change as the system re-equilibrates, but it reaches a new equilibrium.

Master Equilibrium and Le Chatelier's Principle

Work with a chemistry tutor who can help you build intuition for how systems respond to stress and teach you to calculate new equilibrium positions accurately → Le Chatelier's principle is testable through multiple angles: predicting shifts, calculating new equilibria, designing industrial processes. Our tutors teach you the logic behind each principle, help you avoid the most common mistakes, and give you targeted feedback on your problem-solving approach. Whether you're building confidence on basics or refining your technique for exam questions, expert support accelerates your progress. For understanding how equilibrium principles apply to internal assessment investigations, connecting theory to practical work strengthens both.

FAQs

Why does only temperature change the equilibrium constant K?

K depends on the relative stability of products and reactants at a given temperature. Changing temperature changes that relative stability, so K changes. Concentration and pressure changes just shift how much of each substance is present, but they don't change which state (products vs reactants) is more stable.

If I increase pressure by adding an inert gas at constant volume, does equilibrium shift?

No. The total pressure increases, but the partial pressures (and therefore concentrations) of the reacting species don't change. Le Chatelier's principle responds to changes in partial pressure or concentration, not total pressure from inert gases.

What's the difference between how the equilibrium shifts and how K changes?

The equilibrium SHIFTS (concentrations change) in response to concentration, pressure, or temperature stressors. K (the ratio of products to reactants at equilibrium) only changes in response to temperature. For concentration and pressure stressors, K stays the same; the system just reaches a new set of concentrations that still have the same K.