Why don't all reactions reach equilibrium?

Not all reactions reach equilibrium because some reactions go to completion, where all reactants are fully converted into products.

In more detail, a chemical reaction reaches equilibrium when the rate of the forward reaction equals the rate of the reverse reaction, resulting in a constant concentration of reactants and products. However, not all reactions behave this way. Some reactions are irreversible, meaning they proceed in one direction until all of the reactants are converted into products. These are known as reactions that go to completion.

The nature of the reactants and products, as well as the conditions under which the reaction takes place, can determine whether a reaction will reach equilibrium or go to completion. For instance, if a reaction produces a gas that is allowed to escape, the reaction is likely to go to completion because the removal of a product shifts the balance, according to Le Chatelier's principle, driving the reaction forward.

Similarly, if a reaction results in a solid precipitate, the reaction often goes to completion. The precipitate is removed from the reaction mixture, again shifting the balance and driving the reaction to continue until all reactants are used up.

In contrast, reactions that occur in a closed system, where neither reactants nor products can escape, are more likely to reach equilibrium. Here, the forward and reverse reactions occur at the same rate, maintaining a constant concentration of reactants and products.

In summary, whether a reaction reaches equilibrium or goes to completion depends on the nature of the reactants and products, and the conditions under which the reaction takes place.