Quick Summary
- Reversible reactions can proceed in both directions; products can react to reform reactants.
- Equilibrium is reached when forward and reverse reaction rates are equal in a closed system.
- Dynamic equilibrium means reactions continue at the molecular level, but concentrations stay constant.
- Le Chatelier's Principle predicts how equilibrium shifts when conditions change.
- Temperature, pressure, and concentration all shift the position of equilibrium.
- Catalysts speed up equilibrium without shifting its position.
- Industrial processes like Haber Process and Contact Process use compromise conditions to balance yield, rate, safety, and cost.
1. Reversible Reactions
What Is a Reversible Reaction?
Most reactions you've studied go to completion—reactants are completely used up to form products, and the reaction stops. Reversible reactions are different.
In a reversible reaction, the products can themselves react with each other or decompose, reforming the original reactants. The reaction can happen in both directions:
- Forward reaction: Reactants → Products
- Reverse reaction: Products → Reactants
We show this with the special symbol: ⇌ (not just →)
Key Property: Energy Transfers in Both Directions
Here's something crucial to remember:
If the forward reaction is exothermic, the reverse reaction is endothermic. The same amount of heat is transferred in both directions, but in opposite ways. This makes sense because energy is conserved.
Example: The Haber Process Equation
One of the most important reversible reactions in industry is:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
The forward reaction (making ammonia) is exothermic—energy is released. The reverse reaction (ammonia decomposing) is endothermic—energy is absorbed.
Example: Hydrated and Anhydrous Salts
A brilliant way to see reversible reactions in action is with salt hydration.
Water of Crystallisation
Water molecules that are chemically bonded into the structure of some salts during crystallisation. It's not just water sitting on top—it's part of the crystal structure itself.
Copper(II) Sulfate Example:
CuSO₄·5H₂O ⇌ CuSO₄ + 5H₂O
┌─────────────────────────────┬─────────────────────────┐
│ │ │
│ CuSO₄·5H₂O │ CuSO₄ │
│ (Hydrated) │ (Anhydrous) │
│ BLUE crystals │ WHITE powder │
│ Contains water │ No water │
│ │ │
│ HEAT (endothermic) │ │
│ ←─────────────────────→ │ │
│ ADD WATER (exothermic) │ │
│ │ │
└─────────────────────────────┴─────────────────────────┘
- When you heat hydrated copper sulfate, water is driven off → you get white anhydrous powder. This is endothermic.
- When you add water to white powder, the reaction is highly exothermic (releases lots of heat) → blue colour returns.
Exam Tip
- The colour change of CuSO₄ (white → blue) and CoCl₂ (blue → pink) are used as chemical tests for water.
- You must remember these colour changes for exam questions.
- CuSO₄·5H₂O = blue; anhydrous CuSO₄ = white.