Why don't reactions always go to completion?

Reactions don't always go to completion due to factors like equilibrium, reaction reversibility, and insufficient reactants.

In more detail, one of the main reasons reactions don't always go to completion is the concept of equilibrium. In a chemical reaction, the reactants are transformed into products. However, at the same time, the products can also react to form the reactants. When the rate of these two processes becomes equal, the reaction is said to be in a state of equilibrium. At this point, the concentrations of the reactants and products remain constant, but the reaction has not necessarily gone to completion.

Another reason is the reversibility of reactions. Not all reactions are one-way processes. Many reactions are reversible, meaning the products can react together to produce the original reactants. In a reversible reaction, the reaction can proceed in both the forward and reverse directions. When the rate of the forward reaction equals the rate of the reverse reaction, the system is in dynamic equilibrium. This means that the reaction may appear to have stopped, but in reality, the forward and reverse reactions are still occurring, just at the same rate.

Lastly, the availability of reactants can also affect whether a reaction goes to completion. If there is not enough of one reactant to react with the other(s), the reaction will stop before all of the reactants have been used up. This is known as a limiting reactant problem. The reactant that is completely used up first limits the amount of products that can be formed and stops the reaction from going to completion.

In conclusion, while it might seem logical to assume that reactions should always go to completion, the reality is more complex. Factors such as equilibrium, reaction reversibility, and the availability of reactants all play a role in determining whether a reaction will go to completion or not.