Reversible reactions and dynamic equilibrium

Many reactions run to completion: you mix the reactants, they are used up, and that is the end of it.

Some do not. As soon as products form, they start turning back into reactants. The reaction is reversible, and chemists show it with a half-headed arrow:

$\text{N}_2\text{O}_4(g) \rightleftharpoons 2\text{NO}_2(g)$

The arrow pointing right is the forward reaction; the arrow pointing left is the reverse reaction.

Heating blue hydrated copper(II) sulfate drives off water and leaves white anhydrous copper(II) sulfate:

$\text{CuSO}_4\cdot 5\text{H}_2\text{O}(s) \rightleftharpoons \text{CuSO}_4(s) + 5\text{H}_2\text{O}(g)$

Adding a drop of water to the white solid turns it blue again, producing the hydrated salt once more.

This is the same reaction read in the opposite direction — the everyday test for water is the reverse of the everyday test for "water of crystallisation".

A static equilibrium would be a reaction that has stopped. Dynamic equilibrium is the opposite: the reaction has not stopped, it has become a two-way traffic with equal flows.

Imagine a room with two doors and people walking through both at the same rate. The number of people inside stays the same, yet people are constantly moving.

In the same way, individual molecules keep reacting in both directions. Because the two rates match, the amount of each substance on the bench stays put.

Equilibrium requires no matter to enter or leave. If a product escapes — a gas fizzing away from an open beaker — the reverse reaction loses its raw material and can never catch up with the forward reaction.

So the forward reaction keeps winning, and the reaction behaves as if it went to completion. That is why equilibrium experiments are done in sealed tubes or stoppered flasks.

Energy may still be exchanged with the surroundings. Closed means closed to matter, not insulated.