Explain why the following equilibria are heterogencous and write the reaction quoticnt \(Q\) for each one. (a) \(\mathrm{NH}_{4} \mathrm{Cl}(\mathrm{s})=\mathrm{NH}_{3}(\mathrm{~g})+\mathrm{HCl}(\mathrm{g})\) (b) \(\mathrm{Na}_{2} \mathrm{CO}_{3} \cdot \mathrm{OH}_{2} \mathrm{O}\) (s) \(\rightleftharpoons \mathrm{Na}_{2} \mathrm{CO}_{3}(\mathrm{~s})+10 \mathrm{H}_{2} \mathrm{O}(\mathrm{g})\) (c) \(2 \mathrm{KCO}_{3}\) (s) \(\Rightarrow 2 \mathrm{KCl}\) (s) \(+3 \mathrm{O}_{2}\) (g)

Understanding the reaction quotient, Q, is a critical aspect of studying chemical reactions, especially when they don't exist exclusively in one phase state, known as heterogeneous equilibria. In these reactions, the substances involved can exist as solids, liquids, or gases.

Imagine a snapshot of a reaction in-progress, where reactants and products are present in various amounts. This snapshot is described quantitatively by the reaction quotient, Q, which is expressed similarly to the equilibrium constant, Keq, but is applicable at any moment before equilibrium is reached.

The formula for Q is composed of the concentrations of the gaseous and aqueous species raised to the power of their coefficients in the balanced chemical equation. Solid and liquid pure substances are not included as their concentrations are considered constant and do not affect the ratio.

For example, in the reaction \(NH_4Cl(s) \rightarrow NH_3(g) + HCl(g)\), Q is expressed by the equation \(Q = \frac{[NH_3][HCl]}{1}\). This formula takes into account both gaseous ammonia and hydrogen chloride. It's important to monitor Q because it indicates the direction in which the reaction is likely to proceed to reach equilibrium.

Chemical equilibrium represents a state in a chemical reaction where the rates of the forward and reverse reactions are equal, leading to no net change in the concentrations of reactants and products over time. It's a dynamic process where the reactants form products at the same rate as products revert to reactants.

At equilibrium, the reaction quotient Q equals the equilibrium constant Keq. This equality does not mean that the amounts of reactants and products are equal, but rather that their ratios remain constant. When Q and Keq are compared, if Q < Keq, the reaction will proceed forward, and if Q > Keq, the reaction will proceed in reverse to achieve equilibrium.

Understanding the concept of equilibrium is key to predict the behavior of a system under given conditions and to determine the optimal conditions for a reaction to proceed, which can be pivotal in industrial applications and in understanding biological systems.

The presence of different phase states in a reaction is characteristic of heterogeneous equilibria. All reactions involve various intermolecular forces and energy changes, which are significantly influenced by the phase states of the reacting species.

For example, in the dehydration of hydrated sodium carbonate, \(Na_2CO_3\cdot10H_2O(s) \rightleftharpoons Na_2CO_3(s) + 10H_2O(g)\), the reactant is a solid, and it decomposes into another solid and a gas when heated. The phase change introduces complexity to the reaction kinetics and thermodynamics.

Phase states are crucial in determining the form of the equilibrium expression for a reaction. Solids and liquids, due to their constant densities, do not appear in the equilibrium expression. Only the concentrations (or partial pressures) of gases and the molar concentrations of solutes in solution are included. In short, phase states have significant implications on the chemical properties and behaviors of substances in a reaction.