What you'll learn
Reversible reactions and equilibrium explain how some reactions can go both forwards and backwards, and how conditions can be changed to favour the products we want. In this guide you will learn what a reversible reaction is, what dynamic equilibrium means, how energy changes in reversible reactions, and how Le Chatelier's principle predicts the effect of changing concentration, temperature and pressure. These ideas are essential for understanding industrial processes such as the Haber process.
Key terms and definitions
Reversible reaction — a reaction that can proceed both forwards and backwards, shown with the ⇌ symbol.
Dynamic equilibrium — when the forward and backward reactions occur at the same rate in a closed system, so concentrations stay constant.
Closed system — one where no reactants or products can escape.
Le Chatelier's principle — if a change is made to a system at equilibrium, the position of equilibrium shifts to oppose that change.
Exothermic — a reaction that transfers energy to the surroundings.
Endothermic — a reaction that takes in energy from the surroundings.
Core concepts
Reversible reactions
In a reversible reaction, the products can react to reform the reactants. We write it with a double arrow (⇌): for example, ammonium chloride ⇌ ammonia + hydrogen chloride. Heating ammonium chloride breaks it down; cooling the products reforms it. The direction depends on the conditions.
Energy changes in reversible reactions
If a reversible reaction is exothermic in one direction, it is endothermic in the other, and the energy transferred is equal in both directions. For example, hydrated copper sulfate (blue) heated gives anhydrous copper sulfate (white) and water — endothermic; adding water reverses it — exothermic, releasing the same amount of energy.
Dynamic equilibrium
In a closed system, a reversible reaction reaches equilibrium when the forward and backward reactions happen at the same rate. The concentrations of reactants and products then stay constant (though not necessarily equal). It is called dynamic because both reactions continue — they just balance out.
Le Chatelier's principle
Le Chatelier's principle states that if you change the conditions of a system at equilibrium, the equilibrium position shifts to oppose (counteract) that change. This lets us predict how to increase yield.
Effect of concentration
If you increase the concentration of a reactant, the equilibrium shifts to the right (towards products) to use it up. If you remove a product, the equilibrium also shifts right to replace it. Decreasing a reactant or adding product shifts it left.
Effect of temperature
If you increase the temperature, the equilibrium shifts in the endothermic direction (to take in the extra energy); decreasing temperature shifts it in the exothermic direction. So for an exothermic forward reaction, a higher temperature reduces the yield of product.
Effect of pressure (gases)
For reactions involving gases, increasing the pressure shifts the equilibrium towards the side with fewer molecules of gas (to reduce the pressure); decreasing pressure shifts it towards the side with more gas molecules. Count the moles of gas on each side to decide.
Worked examples
Example 1: Recognising reversibility
What does the ⇌ symbol tell you?
It shows the reaction is reversible: it can go forwards and backwards, and can reach a dynamic equilibrium in a closed system.
Example 2: Temperature change
A forward reaction is exothermic. What happens to the yield if temperature is increased?
Increasing temperature shifts the equilibrium in the endothermic (backward) direction, so the yield of product decreases.
Example 3: Pressure change
In N₂ + 3H₂ ⇌ 2NH₃, what is the effect of increasing pressure?
The left has 4 gas molecules, the right has 2. Increasing pressure shifts equilibrium to the side with fewer molecules — the right — increasing the yield of ammonia.
Common mistakes and how to avoid them
Thinking equilibrium means equal concentrations. It means equal rates; concentrations are constant but usually unequal.
Forgetting the system must be closed. If products escape, equilibrium cannot be reached.
Getting the temperature shift backwards. Increasing temperature favours the endothermic direction.
Ignoring moles of gas for pressure. Count gas molecules on each side; only gases are affected by pressure.
Confusing rate and yield. Le Chatelier predicts the position of equilibrium (yield), not how fast it is reached.
Exam technique for Reversible Reactions and Equilibrium
Use the ⇌ symbol and state that the reaction is reversible.
Define dynamic equilibrium as equal forward and backward rates in a closed system.
Apply Le Chatelier's principle systematically for concentration, temperature and pressure.
Count moles of gas to decide the effect of pressure.
Match exothermic/endothermic directions when predicting temperature effects.
Quick revision summary
A reversible reaction can go both ways, shown by ⇌; if exothermic one way it is endothermic the other, transferring equal energy in each direction. In a closed system it reaches dynamic equilibrium, where the forward and backward reactions occur at the same rate and concentrations stay constant (not necessarily equal). Le Chatelier's principle predicts how the equilibrium shifts to oppose a change: increasing a reactant's concentration (or removing product) shifts it towards products; increasing temperature shifts it in the endothermic direction (lowering yield of an exothermic product); and for gases, increasing pressure shifts it towards the side with fewer gas molecules. To answer questions, identify whether the change is concentration, temperature or pressure, recall which direction is endothermic, count gas moles on each side, and remember that equilibrium concerns the position/yield, not the rate. These principles directly explain the compromise conditions used in the Haber process.