# Reversible and irreversible reactions


It is a common observation that most of the reactions when carried out in closed vessels do not go to completion, under a given set of conditions of temperature and pressure. In fact in all such cases, in the initial state, only the reactants are present but as the reaction proceeds, the concentration of reactants decreases and that of products increases. Finally a stage is reached when no further change in concentration of the reactants and products is observed.

Example 1

If a mixture of gaseous hydrogen and iodine vapors is made to react at 717 k in a closed vessel for about 2 - 3 hours, gaseous hydrogen iodide is produced according to the following equation:

$H_{2\; (g)} + I_{2\; (g)} \rightarrow 2HI_{(g)}$

But along with gaseous hydrogen iodide, there will be some amount of unreacted gaseous hydrogen and gaseous iodine left. On the other hand if gaseous hydrogen iodide is kept at 717K in a closed vessel for about 2 - 3 hours it decomposes to give gaseous hydrogen and gaseous iodine.

$2HI_{(g)} \rightarrow H_{2\; (g)} + I_{2\; (g)}$

In this case also some amount of gaseous hydrogen iodide will be left unreacted. This means that the products of certain reactions can be converted back to the reactants. These types of reactions are called reversible reactions.

Thus, in reversible reactions the products can react with one another under suitable conditions to give back the reactants. In other words, in reversible reactions the reaction takes place in both the forward and backward directions. The reversible reaction may be expressed as:

$H_{2\; (g)} + I_{2\; (g)} \rightleftharpoons 2HI_{(g)}$

These reversible reactions never go to completion if performed in a closed container. For a reversible chemical reaction, an equilibrium state is attained when the rate at which a chemical reaction is proceeding in forward direction equals the rate at which the reverse reaction is proceeding.

## Contributors and Attributions

Binod Shrestha (University of Lorraine)

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