Consider this reaction at equilibrium: $$\mathrm{C}(s)+\mathrm{H}_{2} \mathrm{O}(g) \rightleftharpoons \mathrm{CO}(g)+\mathrm{H}_{2}(g)$$ Predict whether the reaction will shift left, shift right, or remain un- changed after each disturbance. \begin{equation}\begin{array}{l}{\text { a. } C \text { is added to the reaction mixture. }} \\ {\text { b. } \mathrm{H}_{2} \mathrm{O} \text { is condensed and removed from the reaction mixture. }} \\ {\text { c. } \mathrm{CO} \text { is added to the reaction mixture. }} \\ {\text { d. } \mathrm{H}_{2} \text { is removed trom the reaction mixture. }}\end{array}\end{equation}

When a chemical reaction occurs, eventually it may reach a state where the rates of the forward and reverse reactions are equal. This dynamic state of balance is known as chemical equilibrium. At this point, the amounts of reactants and products remain constant over time, not because the reactions have stopped, but because they are occurring at the same rate.

Understanding chemical equilibrium is crucial because it allows us to predict how a reaction might behave when subjected to different conditions. In the given exercise, the solid carbon, C, does not alter the equilibrium when added because in such a balanced state, the addition of a solid doesn't change the reaction rates.

The reaction quotient, denoted Q, is a measure that compares the relative amounts of products and reactants present during a reaction at any given time. It is determined by the same expression as the equilibrium constant, K, but isn't limited to reactions at equilibrium. The reaction quotient helps in predicting the direction a reaction will shift to reach equilibrium.

To calculate Q, we use the concentrations of the gaseous reactants and products raised to the power of their coefficients in the balanced chemical equation. When Q is not equal to the equilibrium constant K, it indicates that the reaction will shift to reestablish equilibrium, either to the right towards more products or to the left towards more reactants.

Equilibrium shift refers to the direction a reaction moves to reattain equilibrium after a disturbance. Le Chatelier's principle predicts this behavior—if a system at equilibrium experiences a change in concentration, temperature, or pressure, the equilibrium will shift to counteract the imposed change and restore balance.

For instance, when CO is added to the reaction mixture in our exercise, the equilibrium shifts to the left, utilizing the excess CO to recreate the equilibrium by producing more reactants. On the other hand, removing a reactant or product, like H2O or H2 from the mixture, shifts the reaction towards the side from which the component was removed to counterbalance the change.

Changes in concentration of the reactants or products can significantly influence the position of equilibrium. According to Le Chatelier's principle, increasing the concentration of a reactant or a product will usually cause the equilibrium to shift away from the added substance, and decreasing the concentration will shift it towards the side with the lesser amount.

In the exercise provided, removing H2O condensate shifts the equilibrium to the left, inducing the formation of more H2O from the available CO and H2. Likewise, adding CO shifts the equilibrium left to reduce the concentration of the added product by turning some of it back into reactants.