Reversible Reactions

Some reactions can go in both directions. They are called reversible reactions and are shown with the symbol ⇌, for example: $$\text{A} + \text{B} \rightleftharpoons \text{C} + \text{D}$$

Dynamic equilibrium in a closed system

In a closed system (nothing added or lost), a reversible reaction reaches equilibrium. At equilibrium, the rate of the forward reaction is equal to the rate of the reverse reaction, and the concentrations of reactants and products stay constant. The reaction has not stopped; it is like people going up and down an escalator at the same speed, so the total number on the escalator stays the same.

Changing conditions and the direction of change

• Concentration: Adding a reactant pushes the equilibrium toward the products. Removing a product also pushes forward. Removing a reactant or adding a product pulls it backward.
• Temperature: Heating favours the endothermic direction (takes in heat). Cooling favours the exothermic direction (releases heat). If the forward reaction has $\Delta H < 0$ (exothermic), heating shifts equilibrium backward.
• Pressure (gases only): Increasing pressure favours the side with fewer gas molecules; decreasing pressure favours the side with more gas molecules. Example: $$\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)$$ 4 gas molecules ⇌ 2 gas molecules, so higher pressure favours ammonia.
• Catalyst: Speeds up both forward and reverse reactions equally. It helps equilibrium happen faster but does not change the equilibrium position.

Familiar reversible examples

• Hydrated copper(II) sulfate: Blue crystals lose water on heating to form white anhydrous copper(II) sulfate; adding water reverses it. $$\text{CuSO}_4\cdot5\text{H}_2\text{O} \rightleftharpoons \text{CuSO}_4 + 5\text{H}_2\text{O}$$
• Cobalt(II) chloride: With more water present the pink form is favoured; with less water (dry, warm) the blue form is favoured.

Common misconceptions

• Equilibrium does not mean equal amounts; it means equal rates.
• At equilibrium, reactions continue both ways; they do not stop.
• A closed system is needed; if gases escape, equilibrium cannot be maintained.
• Pressure changes matter only for gaseous reactions.