Le Chatelier's principle helps predict how a chemical equilibrium responds when conditions change. It matters because many important reactions, from industrial ammonia production to blood oxygen transport, are reversible and can shift in either direction. At equilibrium, the forward and reverse reaction rates are equal, but the amounts of reactants and products do not have to be equal.
A stress such as adding reactant, changing volume, or heating the system causes the equilibrium position to shift to reduce that stress.
Understanding Chemistry: Le Chatelier's Principle Worked Examples
Worked examples begin with a balanced chemical equation. The coefficients matter because they show the proportions used when the system readjusts. Consider nitrogen gas reacting with three moles of hydrogen gas to form two moles of ammonia gas.
If extra hydrogen is injected into a closed container, the system has more hydrogen than its equilibrium mixture requires. Some nitrogen and hydrogen are then used to make more ammonia. The change continues only until the balance between the two reaction rates is restored.
Students should separate the immediate change from the later response. The concentration changes at once when hydrogen is added. The equilibrium shift happens afterward as the reaction proceeds.
The reaction quotient gives a more reliable method than relying on a verbal rule. It compares the mixture that exists now with the equilibrium constant for that temperature. Suppose a reaction makes one product from one reactant.
If the current product concentration is two and the current reactant concentration is four, the quotient is one half. If the equilibrium constant is two, the mixture contains too little product relative to reactant. The forward reaction becomes faster than the reverse reaction for a time.
More product forms until the quotient reaches two. In problems with coefficients, concentrations are multiplied by themselves the required number of times.
A coefficient of two means that concentration is squared. This is a common source of calculator mistakes.
Gas volume examples require students to count only moles of gas. For the ammonia reaction, four moles of gaseous reactants become two moles of gaseous product. Compressing the container raises the pressure.
The system can reduce that pressure effect by making more ammonia, because the product side has fewer gas particles. Solids and pure liquids are not included in this count. Adding more solid to a flask usually does not shift an equilibrium, provided some of that solid was already present.
An inert gas can be confusing too. At constant volume, adding inert gas raises total pressure but leaves the reacting gas concentrations unchanged, so the equilibrium position does not move.
Temperature needs different treatment because it changes the equilibrium constant itself. For an exothermic reaction, energy is released as product forms. Heating favors the reverse direction, while cooling favors the forward direction.
For an endothermic reaction, the pattern is reversed. This idea matters in industrial ammonia production. High pressure favors ammonia, but very low temperature can make the reaction too slow for practical use.
Factories choose a compromise, then remove ammonia as it forms to encourage further production. Catalysts are another important distinction.
A catalyst speeds both directions and helps equilibrium be reached sooner. It does not change the final equilibrium mixture at a fixed temperature.
Key Facts
- At equilibrium, rate forward = rate reverse.
- Reaction quotient: Q = [products]^coefficients / [reactants]^coefficients.
- If Q < K, the reaction shifts right toward products.
- If Q > K, the reaction shifts left toward reactants.
- For gases, decreasing volume increases pressure and shifts equilibrium toward the side with fewer moles of gas.
- For an exothermic reaction, heat is a product, so increasing temperature shifts left.
Vocabulary
- Dynamic equilibrium
- A state in which forward and reverse reactions continue at equal rates so concentrations remain constant.
- Le Chatelier's principle
- The rule that an equilibrium system shifts to oppose a change in concentration, pressure, volume, or temperature.
- Reaction quotient
- A value calculated like the equilibrium constant using current concentrations to predict the direction of shift.
- Equilibrium constant
- A constant value at a fixed temperature that compares product concentrations to reactant concentrations at equilibrium.
- Endothermic reaction
- A reaction that absorbs heat, so heat can be treated as a reactant in equilibrium predictions.
Common Mistakes to Avoid
- Saying the system shifts toward the side with more particles after compression is wrong because increased pressure favors the side with fewer moles of gas.
- Including pure solids or pure liquids in K or Q expressions is wrong because their activities are treated as constant and do not appear in the expression.
- Assuming a catalyst changes the equilibrium position is wrong because a catalyst speeds up both forward and reverse reactions equally and only helps the system reach equilibrium faster.
- Forgetting that temperature changes alter K is wrong because concentration and pressure changes shift position at the same K, but changing temperature changes the equilibrium constant itself.
Practice Questions
- 1 For N2(g) + 3H2(g) ⇌ 2NH3(g), a container at equilibrium has [N2] = 0.20 M, [H2] = 0.30 M, and [NH3] = 0.60 M. Calculate Q and compare it with K = 100 to predict the shift direction.
- 2 For 2SO2(g) + O2(g) ⇌ 2SO3(g), predict the shift when the volume of the container is cut in half. Then state which side has fewer gas moles.
- 3 For Co(H2O)6^2+(aq) + 4Cl−(aq) ⇌ CoCl4^2−(aq) + 6H2O(l), adding chloride ion makes the solution more blue because CoCl4^2− is blue. Explain the shift using Le Chatelier's principle.