Predict the effect of increasing the container volume on the amounts of each reactant and product in the following reactions: (a) \(\mathrm{CH}_{3} \mathrm{OH}(l) \rightleftharpoons \mathrm{CH}_{3} \mathrm{OH}(g)\) (b) \(\mathrm{CH}_{4}(g)+\mathrm{NH}_{3}(g) \rightleftharpoons \mathrm{HCN}(g)+3 \mathrm{H}_{2}(g)\)

When a chemical reaction is at equilibrium, it means the rate of the forward reaction equals the rate of the reverse reaction. This balance can be tipped in favor of the forward or reverse reaction by changing the conditions. This is known as an equilibrium shift. Le Chatelier's Principle is a helpful guide to predict the direction of this shift when conditions such as pressure, temperature, or concentration are altered. For example, in reaction (a) \(\text{CH}_3\text{OH}(l) \rightleftharpoons \text{CH}_3\text{OH}(g)\), if the volume of the container is increased, the equilibrium will shift towards the side with more moles of gas, which is the gaseous \(\text{CH}_3\text{OH}(g)\) side.

Pressure changes can significantly influence the position of equilibrium in a reaction involving gases. According to Le Chatelier's Principle, if the pressure of a system at equilibrium is decreased, the system will shift to increase the pressure again. This usually means moving towards the side with more moles of gas. Conversely, if the pressure is increased, the system will shift towards the side with fewer moles of gas. In reaction (b) \(\text{CH}_4(g) + \text{NH}_3(g) \rightleftharpoons \text{HCN}(g) + 3 \text{H}_2(g)\), increasing the volume of the container (which decreases the pressure) causes the equilibrium to shift to the right, where there are more moles of gas.

Reaction dynamics involve understanding how different variables affect the speed and direction of a reaction. Le Chatelier's Principle is key here, because it helps predict how a system will respond to changes and return to equilibrium. By knowing whether a system will shift left or right, you can predict the concentrations of reactants and products. For instance, in our example reaction (b), knowing that increasing the container's volume decreases pressure allows you to predict an increase in the amounts of \(\text{HCN}(g)\) and \(\text{H}_2(g)\).

In a gaseous reaction, the number of moles of gas on each side of the equation plays a crucial role in determining how the equilibrium shifts in response to pressure changes. In reaction (a) \(\text{CH}_3\text{OH}(l) \rightleftharpoons \text{CH}_3\text{OH}(g)\), the equilibrium will shift towards the side with more gas moles (gaseous \(\text{CH}_3\text{OH}(g)\)) when the container volume is increased. Similarly, in reaction (b) \(\text{CH}_4(g) + \text{NH}_3(g) \rightleftharpoons \text{HCN}(g) + 3 \text{H}_2(g)\), the right side has 4 moles of gas compared to 2 moles on the left side, prompting the shift towards the right when the volume is increased.