chevron_left A heterogeneous equilibrium exists when substances in a reaction mixture are in different phases chevron_right

Heterogeneous Equilibrium

What is Chemical Equilibrium?

Before diving into heterogeneous equilibrium, let's first understand the general idea of chemical equilibrium. Imagine a busy playground where children are constantly moving between the swings and the slide. At first, many children run to the swings, but soon the swings are full. As some children get off the swings to go to the slide, others from the slide come to take their place. After a while, even though individual children are always moving, the number of children on the swings and the slide stays roughly constant. This is a state of balance, or equilibrium.

In chemistry, a reversible reaction is one that can go both forward and backward. Chemical equilibrium is the state reached when the rates of the forward and reverse reactions become equal, and the concentrations of the reactants and products remain constant over time. It is important to remember that this is a dynamic equilibrium; the reactions haven't stopped, they are just happening at the same speed.

Homogeneous vs. Heterogeneous Equilibrium

Not all equilibria are the same. Scientists classify them based on the physical states of the substances involved.

The Special Rule: Pure Solids and Liquids

The most important rule for understanding heterogeneous equilibrium concerns pure solids and pure liquids. For any pure substance (one that is not mixed with anything else), its concentration is directly related to its density. Since density is constant at a given temperature, the concentration of a pure solid or liquid is also constant.

Think of it like this: if you have a large cube of ice in a drink, the amount of solid ice (its concentration) doesn't change as it melts and refreezes at equilibrium; only the amount of liquid water changes. Therefore, we don't need to include the solid ice in our calculation of the equilibrium state.

Writing the Equilibrium Constant Expression (K)

The equilibrium constant, K, is a number that tells us the position of equilibrium. For a general reversible reaction: $ aA + bB \rightleftharpoons cC + dD $ the equilibrium constant expression is: $ K = \frac{[C]^{c}[D]^{d}}{[A]^{a}[B]^{b}} $ where the square brackets [ ] represent concentration in moles per liter (M).

For heterogeneous reactions, we apply the special rule and exclude pure solids and liquids. Let's see how this works with a classic example.

A Classic Example: Decomposition of Limestone

The thermal decomposition of calcium carbonate (limestone) is a perfect illustration of a heterogeneous equilibrium involving a solid and a gas.

The chemical equation is: $ CaCO_{3}(s) \rightleftharpoons CaO(s) + CO_{2}(g) $

To write the equilibrium constant expression, we first write it as if all species were included: $ K = \frac{[CaO][CO_{2}]}{[CaCO_{3}]} $

Now, we apply the rule. Both $ CaCO_{3} $ and $ CaO $ are pure solids. Their concentrations are constant and are therefore incorporated into K. We remove them from the expression. This leaves us with: $ K = [CO_{2}] $

This simplified expression is so important it gets its own special name. For reactions where a gas is produced from a solid, the equilibrium constant is called the equilibrium vapor pressure or, more specifically in this case, K_p (where "p" stands for pressure). Since the concentration of a gas is proportional to its partial pressure, we can write: $ K_{p} = P_{CO_{2}} $

This means that at a given temperature, the pressure of carbon dioxide gas in equilibrium with solid calcium carbonate and solid calcium oxide is a fixed value. This pressure is a constant property of the reaction at that temperature, just like the density of a substance.

Equilibrium in Everyday Life and Industry

Heterogeneous equilibrium is not just a textbook concept; it explains many natural phenomena and industrial processes.

1. The Formation of Stalactites and Stalagmites: In caves, carbon dioxide dissolved in groundwater forms weak carbonic acid. This acid reacts with limestone ($ CaCO_{3} $) underground, dissolving it into calcium and bicarbonate ions that are carried by water. When this water drips from the cave ceiling, the $ CO_{2} $ escapes, and the equilibrium shifts, causing the limestone to re-deposit as solid stalactites (from the ceiling down) and stalagmites (from the ground up). This is a slow-motion heterogeneous equilibrium in action.

2. The Vapor Pressure of Water: Consider a closed container partially filled with water. The system will establish a heterogeneous equilibrium between liquid water and water vapor: $ H_{2}O(l) \rightleftharpoons H_{2}O(g) $ The equilibrium constant expression, omitting the pure liquid water, is $ K_{p} = P_{H_{2}O} $. This P_H2O is the vapor pressure of water at that temperature. On a humid day, the air has a high partial pressure of water vapor, close to this equilibrium vapor pressure.

3. The Haber Process for Ammonia Synthesis: While the main reaction $ N_{2}(g) + 3H_{2}(g) \rightleftharpoons 2NH_{3}(g) $ is homogeneous, the industrial process uses a solid iron catalyst. The reaction then occurs on the surface of the solid catalyst, creating a temporary heterogeneous system where gas molecules interact with solid surface sites, a key concept in surface chemistry.

Important Questions

Footnote

¹ K: The equilibrium constant, a number that expresses the relationship between the amounts of products and reactants present at equilibrium.

² K_p: The equilibrium constant calculated from the partial pressures of gases. For a reaction $ aA(g) + bB(g) \rightleftharpoons cC(g) + dD(g) $, $ K_{p} = \frac{(P_C)^c (P_D)^d}{(P_A)^a (P_B)^b} $.

³ Concentration: The amount of a substance (solute) in a given volume of solution, typically measured in moles per liter (M).

⁴ Reversible Reaction: A chemical reaction that can proceed in both the forward and reverse directions.