Chemical Equilibrium Explained
Many chemical reactions do not go to completion. Instead they reach a state of dynamic equilibrium in which reactants are being converted to products at exactly the same rate as products are being converted back to reactants. Understanding equilibrium — and how to shift it — is central to industrial chemistry, from the manufacture of ammonia to the production of sulfuric acid.
What Is Dynamic Equilibrium?
Consider the reaction between hydrogen and iodine in a sealed container: H2(g) + I2(g) ⇌ 2HI(g). At first, only H2 and I2 are present, so only the forward reaction (making HI) occurs. As HI accumulates, the reverse reaction (HI breaking back into H2 and I2) begins. Over time, both rates adjust until the forward rate equals the reverse rate — this is dynamic equilibrium. The word “dynamic” is crucial: the reactions have not stopped. Molecules continue to convert in both directions; it is the rates that are equal, not the concentrations of reactants and products.
Dynamic equilibrium can only be established in a closed system where no substances can enter or leave. An open bottle of soda loses CO2 to the atmosphere and never reaches equilibrium; a sealed bottle eventually does.
The Equilibrium Constant Kc
For a general equilibrium aA + bB ⇌ cC + dD, the equilibrium constant in terms of concentration is defined as:
Kc = [C]c[D]d / ([A]a[B]b)
where [X] denotes the molar concentration of species X at equilibrium, and the exponents are the stoichiometric coefficients from the balanced equation. Kc has no units unless concentrations are expressed in non-standard ways; its numerical value tells you the relative proportions of products and reactants at equilibrium at a given temperature. A large Kc (much greater than 1) means the equilibrium lies to the right — mostly products. A small Kc (much less than 1) means the equilibrium lies to the left — mostly reactants. A Kc near 1 means significant amounts of both are present.
Crucially, Kc changes only when temperature changes. Adding a catalyst, changing concentrations, or changing pressure for gas reactions does not alter Kc.
Le Chatelier’s Principle
The French chemist Henri Louis Le Chatelier stated a principle in 1884: if a system at equilibrium is subjected to a change, the system will respond in the direction that partially opposes the change. This is sometimes called the “system fights back” rule. It allows us to predict qualitatively how equilibrium will shift in response to outside stresses.
| Change applied | Direction of shift | Effect on Kc |
|---|---|---|
| Increase concentration of reactant | Forward (toward products) | No change |
| Decrease concentration of reactant | Reverse (toward reactants) | No change |
| Increase concentration of product | Reverse | No change |
| Increase pressure (gas reaction) | Toward side with fewer moles of gas | No change |
| Decrease pressure (gas reaction) | Toward side with more moles of gas | No change |
| Increase temperature (exothermic reaction) | Reverse (endothermic direction) | Decreases |
| Increase temperature (endothermic reaction) | Forward | Increases |
| Add a catalyst | No shift | No change |
Concentration Changes
Adding more of a reactant increases the rate of the forward reaction momentarily. The system responds by producing more product until a new equilibrium is reached. At the new equilibrium, the concentration of all species is different from before, but Kc has the same value (because temperature is unchanged). Removing a product has the same effect as adding a reactant — the system shifts forward to replace the removed product. Industrial processes often continuously remove products to drive the equilibrium as far forward as possible.
Pressure Changes (Gas Equilibria)
Increasing pressure on a gas-phase equilibrium favours the side with fewer moles of gas. For the Haber process (N2(g) + 3H2(g) ⇌ 2NH3(g)), there are 4 moles of gas on the left and 2 on the right. Increasing pressure therefore shifts the equilibrium to the right, producing more ammonia. If both sides have equal moles of gas (as in H2 + I2 ⇌ 2HI), pressure changes do not shift the equilibrium at all.
Temperature Changes
Temperature is the one stress that actually changes the value of Kc. If the forward reaction is exothermic (releases heat, negative ΔH), treating heat as a product: raising temperature is like adding a product, shifting equilibrium to the left and decreasing Kc. Lowering temperature shifts it to the right, increasing Kc. For an endothermic forward reaction (positive ΔH), raising temperature favours more products and increases Kc.
This creates an industrial dilemma: for an exothermic reaction like the Haber process, a low temperature gives a high Kc (more ammonia at equilibrium) but a very slow rate. A compromise temperature of around 450°C and an iron catalyst are used — accepting a lower equilibrium yield in exchange for a commercially viable rate of production. Removing ammonia as it forms, and recycling unreacted gases, pushes conversion higher over time.
Summary
Dynamic equilibrium is reached in a closed system when the forward and reverse reaction rates are equal; concentrations become constant but are not necessarily equal. Kc quantifies the ratio of products to reactants at equilibrium and changes only with temperature. Le Chatelier’s principle predicts how concentration, pressure, and temperature changes shift the equilibrium position. These concepts are fundamental to understanding reaction types, industrial synthesis, and buffer systems in biology.