Consider the following equilibrium process: $$\mathrm{PCl}_{5}(g) \rightleftharpoons \mathrm{PCl}_{3}(g)+\mathrm{Cl}_{2}(g) \quad \Delta H^{\circ}=92.5 \mathrm{~kJ} / \mathrm{mol}$$ Predict the direction of the shift in equilibrium when (a) the temperature is raised; (b) more chlorine gas is added to the reaction mixture; (c) some \(\mathrm{PCl}_{3}\) is removed from the mixture; (d) the pressure on the gases is increased; (e) a catalyst is added to the reaction mixture.

When a chemical reaction reaches a state where the rate of the forward reaction equals the rate of the reverse reaction, we say the reaction is at equilibrium. However, various changes to the reaction conditions can 'disturb' this balance, leading to what is known as an equilibrium shift. This shift refers to the change in concentrations of reactants and products until a new equilibrium is established.

Applying Le Chatelier's principle helps us to predict the direction of an equilibrium shift. If a change is made to the conditions of a reaction at equilibrium (such as temperature, concentration, or pressure), the system will adjust in a way that counteracts this change. For instance, increasing the temperature of an endothermic reaction will shift the equilibrium to the right, favoring the formation of products to absorb the additional heat. Similarly, adding a reactant or removing a product shifts the equilibrium towards the products to counter the change.

The essence of a chemical reaction is the transformation of one set of chemical substances into another. Reactants are the starting substances that undergo change during the reaction and are converted into products, the new substances formed as a result of the reaction.

Chemical reactions can be represented by a chemical equation, which shows the reactants on one side of an arrow and the products on the other. In the case of reversible reactions, like our example with phosphorus pentachloride, the equation uses a double-headed arrow to show that the reaction can proceed in both forward and reverse directions. Equilibrium is achieved when the rates of these two opposing reactions are equal, meaning the concentration of reactants and products remains constant over time unless the system is disturbed.

The reaction rate is the speed at which a chemical reaction proceeds. It is determined by how quickly reactants are consumed and products are formed. Factors including temperature, concentration of reactants, surface area, pressure, and the presence of a catalyst can influence reaction rates.

In our equilibrium example, higher temperatures increase the rate of the endothermic reaction. A catalyst, while it increases the reaction rate for both the forward and reverse reactions, uniquely does not affect the equilibrium position because it speeds up both processes equally. Understanding reaction rates is crucial in predicting how changes in conditions will affect the rate at which equilibrium is reached, but not the composition of the equilibrium mixture.

In an endothermic reaction, the system absorbs heat from its surroundings. This input of energy is necessary for the reaction to occur and is reflected in a positive heat change \text(ΔH). The reaction between \text(\(\text{PCl}_{5}(g)\)) and its decomposition into \text(\(\text{PCl}_{3}(g)\)) and \text(\(\text{Cl}_{2}(g)\)) with a \text(ΔH^{\text{o}}) of 92.5 kJ/mol is an example of an endothermic process.

Understanding the endothermic nature of a reaction is vital in predicting how applying heat will affect it. For our example, increasing the temperature would drive the equilibrium position to the right, favouring the formation of more products, which is consistent with Le Chatelier's principle.