Mercury(I) oxide decomposes into elemental mercury and elemental oxygen: \(2 \mathrm{Hg}_{2} \mathrm{O}(s) \rightleftharpoons 4 \mathrm{Hg}(l)+\mathrm{O}_{2}(g) .\) Write the equilibrium-constant expression for this reaction in terms of partial pressures. (b) Suppose you run this reaction in a solvent that dissolves elemental mercury and elemental oxygen. Rewrite the equilibrium- constant expression in terms of molarities for the reaction, using (solv) to indicate solvation.

Chemical equilibrium is a state in which the rate of the forward reaction equals the rate of the reverse reaction, resulting in no overall change in the concentrations of reactants and products. It's important to realize that equilibrium does not mean the reactants and products are present in equal amounts, but that their rates of formation are equal, leading to a constant ratio of concentrations that can be described by the equilibrium constant (K).

When a reaction reaches this stable balance, it is said to be at equilibrium. The equilibrium constant expression quantitatively describes the ratio of the concentrations of products to the concentrations of reactants, each raised to the power of their respective coefficients in the balanced equation. For the decomposition of Mercury(I) oxide, the equilibrium constant (called K or Kp when using partial pressures and Kc when using concentrations) provides an essential snapshot of the reaction's status at equilibrium.

Partial pressures are used when dealing with equilibrium involving gases. The partial pressure of a gas is the pressure that it would exert if it were alone in the container. In a mixture of gases, each gas contributes to the total pressure of the system in proportion to its concentration, known as its partial pressure.

The equilibrium constant expression can be written in terms of partial pressures (Kp). For gases, Kp is particularly useful because it directly relates to the gas's behavior under various pressures according to Dalton's Law of Partial Pressures. In the equilibrium constant expression for the decomposition of Mercury(I) oxide, the partial pressure of oxygen (O₂) is used, as it is the only gas present in the balanced reaction.

Molarity, denoted by M, is a measure of the concentration of a solute in a solution. It is defined as the number of moles of solute per liter of solution. This is a key concept when describing reactions that occur in solution, as the reaction rates and equilibria depend on the concentration of the reactants and products in the solution.

When the equilibrium involves substances that have been dissolved in a solvent, the equilibrium constant expression is written in terms of molarity (Kc). This version of the constant is based on the molar concentration of solvated reactants and products. In our exercise, after considering solvation, the equilibrium constant expression for the system changes to include the molarities of the solvated mercury and oxygen.

Reversible reactions are those that can proceed in both the forward and reverse directions. In such reactions, reactants can combine to form products, which can then break back down to yield the original reactants. Equilibrium is a dynamic process that occurs in reversible reactions; the system will constantly shuffle between reactants and products while maintaining a stable proportion that is defined by the equilibrium constant.

Understanding that reversible reactions reach a point where the rate of the forward process equals the rate of the reverse process is crucial for appreciating how equilibrium is achieved. The decomposition of Mercury(I) oxide is an example of a reversible reaction, where Mercury(I) oxide can decompose into mercury and oxygen and then recombine under the right conditions.

Solvation is the process by which solvent molecules surround and interact with solute ions or molecules when a solute is dissolved in a solvent. This interaction is crucial in changing the nature of the solute particles and their behavior in the solution, affecting properties such as conductivity, reactivity, and osmotic pressure.

In the context of chemical equilibrium, solvation can affect the equilibrium constant expression, as it changes when reactants and products are present in solvated forms. As seen in the reaction example, solvation of mercury and oxygen alters the equilibrium constant expression to account for their solvated states (indicated as [Hg(solv)] and [O2(solv)]), requiring the use of concentrations rather than partial pressures to describe the system at equilibrium.