What is Chemical Equilibrium?
When chemists study reactions, they often assume that reactions go to completion — that is, all the reactants are converted into products. But in reality, many reactions never fully complete. Instead, they reach a special state called chemical equilibrium.
Think about what happens when you open a bottle of fizzy drink. Carbon dioxide dissolved in the liquid constantly escapes into the gas above, while CO₂ from the gas constantly dissolves back into the liquid. Eventually, the rate of escaping and dissolving balance out — this is equilibrium in everyday life!
In chemistry, equilibrium occurs in closed systems — systems where no substances can enter or leave. This is crucial: if products escape, the reaction cannot reach equilibrium.
The state of a reversible reaction in which the rate of the forward reaction equals the rate of the reverse reaction, so the concentrations of reactants and products remain constant over time.
Equilibrium can only be established in a closed system. If the system is open (e.g., a gas escapes into the atmosphere), equilibrium cannot be reached because products are continuously removed.
Reversible Reactions: The Foundation of Equilibrium
Before we can understand equilibrium, we need to understand reversible reactions. Most reactions you have studied are written with a single arrow (→), suggesting they only go in one direction. However, many reactions can proceed in both directions.
A reaction in which the products can react together to reform the original reactants. Reversible reactions are written using a double arrow (⇌).
The two directions of a reversible reaction are:
- Forward reaction: Reactants → Products
- Reverse reaction: Products → Reactants
A classic example is the reaction between hydrogen gas and iodine gas:
At the start, only H₂ and I₂ are present, so the reverse reaction rate is effectively zero — no HI exists yet to react in reverse. As HI builds up, the reverse reaction begins. Over time, both reactions happen simultaneously until equilibrium is reached.
Imagine a crowded hallway with people walking in both directions. At first, all the people walk from left to right (forward reaction). As the right side fills up, more people start walking back to the left (reverse reaction). Eventually, the same number of people cross in each direction per minute — equilibrium!
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