To determine the direction in which the equilibrium shifts when HCl is added to the reaction, we need to consider Le Chatelier’s principle.
Le Chatelier’s principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract the change.
Given the reaction:
[tex]\[ \text{H}_2(g) + \text{Cl}_2(g) \rightleftharpoons 2 \text{HCl}(g) + 184.6 \, \text{J} \][/tex]
When hydrochloric acid (HCl) is added to the system, we are increasing the concentration of one of the products of the reaction. According to Le Chatelier’s principle, if the concentration of a product is increased, the system will respond to partially counteract this change by shifting the equilibrium position to the left.
This shift to the left means that the reaction will favor the formation of the reactants ([tex]\(\text{H}_2\)[/tex] and [tex]\(\text{Cl}_2\)[/tex]) from the products, decreasing the concentration of HCl and thus partially offsetting the initial increase.
Therefore, when HCl is added to the reaction, the equilibrium will shift to the left, or towards the reactants.
