In heterogeneous reaction systems, what types of substances do not appear in the equilibrium constant expression? Why?

In chemical reactions, you often see substances in different physical states taking part in the process. A reaction where reactants and products exist in different phases is called a heterogeneous reaction. For instance, you can have a reaction involving a solid and a gas. This is different from homogeneous reactions, where all reactants and products are in the same phase.

Let's consider an example of a heterogeneous reaction:
CaCO₃(s) ⇌ CaO(s) + CO₂(g)

Notice that calcium carbonate (CaCO₃) and calcium oxide (CaO) are solids, while carbon dioxide (CO₂) is a gas.

Due to the involvement of different phases (solid and gas), this reaction is considered heterogeneous. This concept is important because it affects how we write the equilibrium constant expression.

In the equilibrium constant expression for a chemical reaction, you may have noticed that pure solids and pure liquids are not included. Why is that?

Pure solids and pure liquids have constant compositions. Their concentrations do not change during the reaction. Because their activity (a measure of effective concentration) is always 1, they do not affect the equilibrium position.

So, when writing the equilibrium expression, we omit pure solids and pure liquids. Only the concentrations of gaseous and aqueous species are included.

For example, look at the reaction:
CaCO₃(s) ⇌ CaO(s) + CO₂(g)

Here’s the right equilibrium constant expression for this reaction:
\[ K = [CO_2] \]
Both CaCO₃ and CaO are solids and their concentrations are assumed constant and thus, not included in the expression.

When a chemical reaction reaches equilibrium, the rates of the forward and reverse reactions are equal, and the concentrations of the reactants and products remain constant over time. The state of balance is called a chemical equilibrium.

To describe this equilibrium quantitatively, we use the equilibrium constant (K), which is calculated using the equilibrium constant expression. This expression is derived from the balanced chemical equation.

For a reaction:
\[ aA + bB ⇌ cC + dD \]
the equilibrium constant expression is:
\[ K = \frac{[C]^c[D]^d}{[A]^a[B]^b} \]
Here, the concentrations of the reactants (A and B) and products (C and D) are raised to the power of their coefficients in the balanced equation. This expression allows us to understand the relationship between the concentrations of the substances involved at equilibrium.

In heterogeneous reactions, we exclude pure solids and liquids from the K expression, focusing only on the gaseous and aqueous phases because their concentrations can vary, affecting the overall equilibrium.