Dynamic Equilibrium
Simple Explanation
Dynamic equilibrium is the state a reversible reaction reaches when the forward and reverse reaction rates become exactly equal — the concentrations of reactants and products stop changing overall, even though both reactions are still actively happening.
Why Do We Need It?
The word "dynamic" is essential: equilibrium is NOT a static state where the reaction has stopped — it is a constant, balanced back-and-forth, which is why concentrations stay constant without either reaction ever actually ceasing.
Why Does This Work?
As a reversible reaction proceeds, the forward rate decreases as reactants are used up, while the reverse rate increases as products build up — eventually these two changing rates cross and become equal. Once equal, for every product molecule formed by the forward reaction, one product molecule is being converted back by the reverse reaction (on average), so net concentrations stop changing.
Real-Life Example
A saturated solution
Sugar stops dissolving once a solution becomes saturated, even though solid sugar remains at the bottom.
This is a dynamic equilibrium: sugar molecules are still constantly dissolving into solution AND crystallizing back out of solution at the same rate — the amount of dissolved sugar stops changing, but the process never actually stops.
Practice
At dynamic equilibrium, which statement is true?
MediumCommon mistake
Believing that equal concentrations of reactants and products indicate equilibrium — equilibrium means the RATES are equal, which usually happens at very unequal concentrations, depending on how far the reaction favours products or reactants.
Quick Review
- Dynamic equilibrium: forward rate = reverse rate.
- Concentrations stay constant, but both reactions keep happening.
- Equilibrium does NOT mean equal concentrations of reactants and products.