Question

Consider an equilibrium mixture consisting of \(\mathrm{H}_{2} \mathrm{O}(g), \mathrm{CO}(g), \mathrm{H}_{2}(g),\) and \(\mathrm{CO}_{2}(g)\) reacting in a closed vessel according to the equation $$\mathrm{H}_{2} \mathrm{O}(g)+\mathrm{CO}(g) \rightleftharpoons \mathrm{H}_{2}(g)+\mathrm{CO}_{2}(g)$$ a. You add more \(\mathrm{H}_{2} \mathrm{O}\) to the flask. How does the new equilibrium concentration of each chemical com- pare to its original equilibrium concentration after equilibrium is reestablished? Justify your answer. b. You add more \(\mathrm{H}_{2}\) to the flask. How does the concentration of each chemical compare to its original concentration after equilibrium is reestablished? Justify your answer.

Short answer

a. After adding more H2O and reaching the new equilibrium: - H2O concentration is higher than before. - CO concentration is lower than before. - H2 concentration is higher than before. - CO2 concentration is higher than before. b. After adding more H2 and reaching the new equilibrium: - H2O concentration is higher than before. - CO concentration is higher than before. - H2 concentration is higher than before. - CO2 concentration is lower than before.

Step-by-step solution

Part a: Adding more H2O

We have the following reaction at equilibrium: \[\mathrm{H}_{2}\mathrm{O}(g) + \mathrm{CO}(g) \rightleftharpoons \mathrm{H}_{2}(g) + \mathrm{CO}_{2}(g)\] When we add more H2O to the flask, according to Le Chatelier's Principle, the equilibrium will shift to counteract the change, in this case by consuming some of the added H2O. The equilibrium will shift to the right, meaning the concentration of H2 and CO2 will increase, and the concentration of CO will decrease. Once the new equilibrium is established: - The concentration of H2O will be greater than before (because some of the added H2O remains). - The concentration of CO will be lower than before (due to the shift to the right). - The concentration of H2 will be higher than before (because more H2 is produced). - The concentration of CO2 will be higher than before (due to the shift to the right).

Part b: Adding more H2

We again have the equilibrium reaction: \[\mathrm{H}_{2}\mathrm{O}(g) + \mathrm{CO}(g) \rightleftharpoons \mathrm{H}_{2}(g) + \mathrm{CO}_{2}(g)\] When we add more H2 to the flask, according to Le Chatelier's Principle, the equilibrium will shift to counteract the change. In this case, the equilibrium will shift to the left, consuming some of the added H2. This means the concentration of H2O and CO will increase, while the concentration of CO2 will decrease. Once the new equilibrium is established: - The concentration of H2O will be higher than before (due to the shift to the left). - The concentration of CO will be higher than before (because more CO is produced). - The concentration of H2 will be higher than before (because some of the added H2 remains). - The concentration of CO2 will be lower than before (due to the shift to the left).

Key concepts

Le Chatelier's Principle

Le Chatelier's Principle is a fundamental concept in chemistry that describes how a reaction at equilibrium responds to changes in its conditions. In simpler terms, it states that if an external change is imposed on a system at equilibrium, the system will adjust itself to minimize that change. Think of it as the system's way of resisting the effects of the disturbance. It's like a chemical tug-of-war, where the reactants and products are pulling against each other. When something external, like temperature or pressure, or the concentration of one of the reactants or products changes, the system will react by shifting the equilibrium in a way that reduces the impact of this change.

For instance, if more reactants are added to the system, the equilibrium will shift to produce more products, consuming the added reactants in the process. Conversely, if some of the product is removed, the system will produce more product to replace it. The principle can be applied to changes in concentration, temperature, volume, or pressure, and it's a useful tool for predicting how a system will respond to such changes.

Equilibrium Shift

An equilibrium shift occurs when the balance between the reactants and products in a reversible chemical reaction is disturbed. This shift is described by Le Chatelier's Principle and can be triggered by changes in concentration, pressure, volume, or temperature. Understanding how an equilibrium shifts allows chemists to predict the direction in which a reaction will proceed when subjected to different conditions.

For example, in the original exercise, adding more H2O to the system caused the equilibrium to shift to the right, leading to more production of H2 and CO2. This process, simply put, is the system's way of reducing the increased concentration of H2O by using it up to create more products. It's like a dance where everyone has to find a new partner if someone new joins in — the existing pairs rearrange themselves to accommodate the change.

Reaction Concentrations

Reaction concentrations refer to the amounts of reactants and products present in a chemical reaction. In an equilibrium situation, these concentrations remain constant over time, which is different from saying that they are necessarily equal. In a closed system, where no substances can enter or leave, the concentrations dictate the position of equilibrium.

The change in reaction concentrations is an important aspect when considering equilibrium shifts. As seen in the textbook exercise, adding more H2 caused a shift which increased the concentrations of H2O and CO due to the reaction moving towards the reactants to re-establish equilibrium. This principle helps scientists and engineers to control reactions, ensuring they can optimize conditions for the desired outcome such as the maximum yield of product. By manipulating the concentrations, they can 'push' the reaction in one direction or another according to what is needed at the time.