Calculate the reaction Gibbs free energy of \(\mathrm{H}_{2}(\mathrm{~g})+\) \(\mathrm{I}_{2}(\mathrm{~g}) \longrightarrow 2 \mathrm{HI}(\mathrm{g})\) at 700 . \(\mathrm{K}\) when the partial pressures of \(\mathrm{H}_{2}, \mathrm{I}_{2}\), and HI are \(0.35\) bar, \(0.18\) bar, and \(2.85\) bar, respectively. For this reaction, \(K=54\) at 700 . \(\mathrm{K}\). (b) Indicate whether this reac mixture is likely to form reactants, is likely to form produ is at equilibrium.

Understanding chemical equilibrium is fundamental to grasping how reactions behave over time. It occurs when the rates of the forward and reverse reactions are equal, leading to constant concentrations of the reactants and products. This state doesn’t mean the reactants and products are in equal concentration, but that their rates of formation are stabilized.

For the given chemical reaction (hydrogen gas reacting with iodine gas to produce hydrogen iodide), reaching equilibrium means that the amount of hydrogen, iodine, and hydrogen iodide no longer changes over time. This balanced state does not imply the reaction has stopped; molecules continue to react, but their concentrations remain unaltered due to the equal and opposite reaction rates.

When we say a reaction is 'at equilibrium,' it’s an implication that no net change can be observed. The system can still respond to changes, as predicted by Le Chatelier's principle, leading to shifts in equilibrium to counteract the disturbance.

The reaction quotient (Q) helps chemists determine the direction in which a reaction mixture is likely to proceed at any given point before reaching equilibrium. It’s a ratio that compares the concentrations (or partial pressures in the case of gases) of the products to the reactants, raised to the power of their coefficients from the balanced equation.

To calculate Q, you take a ‘snapshot’ of the system to assess the concentrations or partial pressures at that moment. This snapshot for the reaction between hydrogen and iodine to form hydrogen iodide could be taken at any time before the system reaches equilibrium. Q is pivotal because it predicts if the reaction will proceed forward to produce more products, reverse to yield more reactants, or remain at chemical equilibrium.

The equilibrium constant (K) is deeply connected to the concept of chemical equilibrium. It is defined at a particular temperature and only changes if the temperature is altered. K provides a ratio of the concentrations (or partial pressures) of the products to reactants at equilibrium. Unlike Q, which can be calculated at any point, K is specific to the equilibrium state.

For the reaction at hand, K has been given as 54 at 700 K. This number guides us in predicting whether the reaction favors the formation of products or reactants. A large K value implies a greater concentration of products at equilibrium, suggesting that the reaction strongly favors the formation of products under the given conditions.

When dealing with gases, partial pressures are a crucial consideration. In the context of chemical reactions, the partial pressure refers to the pressure that a specific gas in a mixture would exert if it were alone in the container. It's proportional to the concentration of the gas, making it an alternative way of expressing how much of that gas is present.

In the reaction between hydrogen and iodine, the partial pressures are essential inputs for calculating both Q and K. They represent how much of each gas is available for the reaction at a snapshot in time. Knowing that these pressures are 0.35 bar for hydrogen, 0.18 bar for iodine, and 2.85 bar for hydrogen iodide allows for the precise calculation of the reaction quotient, Q, and enables chemists to determine the reaction's direction relative to equilibrium.