Question

Consider this reaction at equilibrium. $$ \mathrm{CO}(g)+\mathrm{Cl}_{2}(g) \rightleftharpoons \mathrm{COCl}_{2}(g) $$ Predict the effect (shift right, shift left, or no effect) of these changes. (a) adding \(\mathrm{Cl}_{2}\) to the reaction mixture (b) adding \(\mathrm{COCl}_{2}\) to the reaction mixture (c) adding \(\mathrm{CO}\) to the reaction mixture

Short answer

Adding Cl2 or CO causes a shift right. Adding COCl2 causes a shift left.

Step-by-step solution

Predict the Effect of Adding Cl2

According to Le Chatelier's Principle, if a stress is applied to a reaction mixture at equilibrium, the system will adjust itself to partially counteract the imposed change. In this case, adding more \(\mathrm{Cl}_{2}\) would increase the concentration of one of the reactants. The system will shift to the right, towards the formation of the products, to consume some of the excess \(\mathrm{Cl}_{2}\).

Predict the Effect of Adding COCl2

Similarly, by adding more \(\mathrm{COCl}_{2}\) we are increasing the concentration of the product. To counteract this change, the equilibrium will shift to the left to consume some of the added product and form more reactants.

Predict the Effect of Adding CO

Adding more \(\mathrm{CO}\) increases the concentration of this reactant. Therefore, according to Le Chatelier's Principle, the equilibrium will shift to the right in order to reduce the concentration of \(\mathrm{CO}\) by forming more \(\mathrm{COCl}_{2}\).

Key concepts

Chemical Equilibrium

Understanding chemical equilibrium is pivotal when dissecting chemical reactions. It is a state in which the rate of the forward reaction equals the rate of the reverse reaction. As a result, the concentrations of reactants and products remain constant over time, but not necessarily equal. At equilibrium, a particular ratio of concentrations – the equilibrium constant – is established.

For the reaction \(\text{CO}(g) + \text{Cl}_2(g) \rightleftharpoons \text{COCl}_2(g)\), equilibrium is reached when the rate at which \(\text{CO}\) and \(\text{Cl}_2\) combine to form \(\text{COCl}_2\) is equal to the rate at which \(\text{COCl}_2\) decomposes back into \(\text{CO}\) and \(\text{Cl}_2\). It's important to note that even though the macroscopic properties appear static, molecules continue to react at a microscopic level.

Reaction Quotient

When analysing a reaction's progress towards equilibrium, we use the reaction quotient, \(Q\). This dimensionless number is calculated similarly to the equilibrium constant, \(K\), but with the initial concentrations of the reactants and products rather than the concentrations at equilibrium.

The reaction quotient provides insight into the direction the reaction will take to achieve equilibrium. If \(Q < K\), the forward reaction is favored, and the system will shift to the right. Conversely, if \(Q > K\), the reverse reaction is favored, indicating a shift to the left until equilibrium is reestablished.

Equilibrium Shift

Le Chatelier's Principle is a key tool for predicting how a system at equilibrium will respond to changes. It states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract the change. For example, in the given exercise, when \(\text{Cl}_2\) is added, the system shifts to the right, increasing the production of \(\text{COCl}_2\) to reduce the \(\text{Cl}_2\) concentration.

Similarly, by adding \(\text{COCl}_2\), we essentially shift the equilibrium to the left, favoring the reverse reaction to decrease the concentration of the added product. It is fascinating how the principles of equilibrium provide a quantitative and predictive framework for understanding chemical reactions.

Concentration Changes

Variations in concentration can significantly influence the direction of the equilibrium shift. In the case of the exercises provided, adding more of a substance shifts equilibrium away from the added component. This is a direct application of Le Chatelier's Principle.

Increasing the concentration of \(\text{Cl}_2\) or \(\text{CO}\) causes a shift to the right, promoting product formation. Conversely, increasing \(\text{COCl}_2\) prompts a shift to the left, in favor of reactant production. Understanding these fundamental effects of concentration changes is essential for mastering chemical equilibrium concepts and solving related problems.