Difficulty: Intermediate | Prerequisites: Equilibrium fundamentals (K_eq, Q, equilibrium expressions). Read the companion notes on Equilibrium Fundamentals first.
Once you know what equilibrium is and how K and Q work, the next question is: what happens when something disturbs the system? Le Chatelier's Principle is the framework for answering that. It tells you how a system at equilibrium responds when you change concentration, pressure, or temperature.
This chapter also introduces an important distinction between kinetic and thermodynamic control of reactions, which explains why the product you get depends on the conditions you choose. These ideas carry forward into organic chemistry, biochemistry, and industrial process design.
Le Chatelier's Principle says that if you stress a system at equilibrium, it shifts to partially undo that stress. Adding reactants pushes the reaction forward; raising the temperature shifts it toward the endothermic direction; changing pressure affects gaseous equilibria depending on moles of gas on each side. Separately, low temperatures favour the fast (kinetic) product, while high temperatures favour the stable (thermodynamic) product.
Le Chatelier's Principle
If a stress is applied to a system at equilibrium, the system shifts in the direction that partially relieves that stress.
Think of it as the system pushing back against whatever you did to it.
Stress (on an equilibrium system)
Any change in concentration, pressure, volume, or temperature that moves Q away from K_eq.
Endothermic reaction
A reaction with ΔH > 0, meaning it absorbs heat from its surroundings. In Le Chatelier terms, heat acts as a reactant.
Exothermic reaction
A reaction with ΔH < 0, meaning it releases heat to its surroundings. In Le Chatelier terms, heat acts as a product.
Kinetic product
The product that forms faster because it requires less activation energy (lower energy barrier). Favoured at lower temperatures.
In simple terms, the "quick" product. It forms first but is less stable.
Thermodynamic product
The product that is more stable (lower free energy) but forms more slowly because it requires more activation energy. Favoured at higher temperatures.
In simple terms, the "better" product. It takes longer to form but once it does, it sits at a lower energy state.
The core idea is straightforward: disturb a system at equilibrium and it shifts to oppose the disturbance.
When you add reactants or remove products, Q drops below K_eq. The system responds by shifting forward (to the right), producing more products until Q = K_eq again.
When you remove reactants or add products, Q rises above K_eq. The system shifts backward (to the left), producing more reactants.
Practical application: to drive a reaction toward completion, start with a large excess of one reactant or continuously remove the product as it forms. Industrial processes routinely use both strategies.
This section applies only to reactions involving gases.
Compressing the system (decreasing volume) increases total and partial pressures. The system is no longer at equilibrium.
The system responds by shifting toward the side with fewer moles of gas, which reduces the total pressure.
Expanding the system (increasing volume) decreases pressures. The system shifts toward the side with more moles of gas to restore pressure.
If both sides have the same number of moles of gas, changing pressure has no effect on the equilibrium position.
Adding an inert gas at constant volume does not change partial pressures and therefore does not shift the equilibrium.
Temperature is the only common stress that changes the value of K_eq itself. Concentration and pressure changes shift Q relative to a fixed K; temperature changes K.
Concentrations and partial pressures do not change the instant you change temperature, so Q stays the same initially. But K_eq is now a different number, so Q ≠ K_eq.
For an endothermic reaction (ΔH > 0), treat heat as a reactant. Raising the temperature shifts the equilibrium toward products and increases K_eq.
For an exothermic reaction (ΔH < 0), treat heat as a product. Raising the temperature shifts the equilibrium toward reactants and decreases K_eq.
Lowering the temperature has the opposite effect in each case.
When a reaction can produce more than one product, the conditions determine which one you get.
Kinetic product (lower temperature)
Forms faster because it has a lower activation energy (smaller energy barrier to the transition state).
Less stable: sits at a higher free energy at the end of the reaction.
Dominates when the reaction is run at lower temperatures, where there is not enough energy to reach the higher transition state of the alternative pathway.
Thermodynamic product (higher temperature)
Forms more slowly because it requires more activation energy.
More stable: sits at a lower free energy (more negative ΔG).
Dominates when the reaction is run at higher temperatures, where enough energy is available to overcome the larger barrier and the system has time to reach the more stable outcome.
The key distinction is between speed of formation (kinetic) and stability of the result (thermodynamic). Low temperature traps the fast product; high temperature lets the system find the stable one.
Bicarbonate buffer system (Le Chatelier example)
CO₂(g) + H₂O(l) ⇌ H₂CO₃(aq) ⇌ H⁺(aq) + HCO₃⁻(aq)
When blood pH drops (too acidic), the body increases respiration to blow off CO₂, shifting the equilibrium to the left and raising pH.
When blood pH rises (too basic), the body slows respiration, trapping CO₂ and shifting the equilibrium to the right.
Kinetic vs thermodynamic product summary
Kinetic product: higher free energy at the end, lower activation energy, forms at lower temperatures.
Thermodynamic product: lower free energy at the end (more negative ΔG), higher activation energy, forms at higher temperatures.
Your blood's pH regulation is Le Chatelier in action: the bicarbonate buffer shifts equilibrium via CO₂ levels controlled by your breathing rate. In industry, the Haber process for ammonia uses high pressure (fewer moles of gas on the product side) and moderate temperature (a compromise between speed and equilibrium position) to maximise yield.
Students often think adding an inert gas at constant volume shifts the equilibrium. It does not, because partial pressures of the reacting gases are unchanged.
Students assume that raising the temperature always shifts the equilibrium toward products. The direction depends on whether the reaction is endothermic or exothermic.
Students confuse the effect of a catalyst with Le Chatelier's Principle. A catalyst speeds up both the forward and reverse reactions equally. It does not shift the equilibrium position or change K_eq.
Students sometimes think the kinetic product is always "worse." It is less stable, but it may be exactly the product you want, and controlling temperature lets you select it deliberately.
⚠️ Know how to predict the direction of shift for each type of stress: concentration, pressure/volume, and temperature.
⚠️ Temperature is the only stress that changes K_eq. This is a very common exam question.
⚠️ Be able to explain the bicarbonate buffer as a Le Chatelier example.
⚠️ Know the difference between kinetic and thermodynamic products, and which conditions favour each.
⚠️ Remember that catalysts do not shift equilibrium. If a multiple-choice question offers "add a catalyst" as a way to shift equilibrium, it is a trap.
True or false: Adding a catalyst shifts the equilibrium toward products.
False. A catalyst speeds both directions equally and does not change K_eq.
Fill in the blank: For an exothermic reaction, raising the temperature shifts equilibrium toward the ______.
Reactants (left).
True or false: Compressing a gaseous system always shifts equilibrium toward the side with fewer moles of gas.
True.
Fill in the blank: The kinetic product forms ______ but is ______ stable than the thermodynamic product.
Faster; less.
True or false: Changing temperature changes Q but not K_eq.
False. Temperature changes K_eq, not Q (Q stays the same initially because concentrations have not yet changed).
Q: For the reaction N₂(g) + 3 H₂(g) ⇌ 2 NH₃(g) (exothermic), what happens to the equilibrium position if you increase the pressure by decreasing the volume?
A: The equilibrium shifts to the right (toward NH₃), because the product side has fewer moles of gas (2) than the reactant side (4).
Q: The same reaction is exothermic. What happens to K_eq if you raise the temperature?
A: K_eq decreases, because for an exothermic reaction heat is effectively a product, and raising the temperature shifts equilibrium toward reactants.
Q: A student claims that adding an inert gas to a sealed container at constant volume will shift the equilibrium. Are they correct?
A: No. Adding an inert gas at constant volume does not change the partial pressures of the reacting gases, so Q is unchanged and no shift occurs.
Q: A reaction can produce product A (kinetic) or product B (thermodynamic). At low temperature, which product dominates, and why?
A: Product A dominates. At low temperature, there is not enough energy to overcome the higher activation barrier for product B, so the faster pathway to A wins.
Q: In the bicarbonate buffer system, what happens when blood pH drops?
A: The body increases respiration to expel CO₂. Removing CO₂ shifts the equilibrium to the left, consuming H⁺ ions and raising pH back toward normal.
Le Chatelier's Principle is the practical application layer of everything in the companion Equilibrium Fundamentals notes. If K and Q are the diagnostic tools, Le Chatelier tells you what the system will actually do.
These ideas extend directly into acid-base buffers (the Henderson-Hasselbalch equation is an equilibrium expression), solubility equilibria (the common-ion effect is a concentration stress), and organic chemistry (kinetic vs thermodynamic control governs product selection in reactions like enolate formation and Diels-Alder cycloadditions).
The temperature-dependence of K also connects to the van 't Hoff equation, which you may encounter in physical chemistry or advanced general chemistry.
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