1. Reversible reactions
A reversible reaction can proceed in both the forward and reverse directions. The forward reaction converts reactants into products, while the reverse reaction converts products back into reactants.
General reversible reaction
The double arrow indicates that both directions are possible. The letters represent chemical species, and the lowercase letters represent stoichiometric coefficients.
Irreversible and reversible reactions
Irreversible approximation
The reaction proceeds predominantly in one direction under the specified conditions.
Reversible reaction
Both forward and reverse reactions can occur, and the system may establish equilibrium in a closed system.
2. Dynamic equilibrium
Chemical equilibrium is reached when the forward and reverse reaction rates become equal. The concentrations of reactants and products then remain constant over time, provided external conditions do not change.
- The forward and reverse reactions continue to occur.
- The two reaction rates are equal at equilibrium.
- Concentrations are constant, but they are not necessarily equal.
- Equilibrium in this sense is established in a closed system.
Important: Dynamic equilibrium does not mean that reactions have stopped. It means that the two opposing processes occur at equal rates.
What happens as equilibrium approaches?
Imagine starting with reactants and almost no products. Initially, the forward reaction may be fast while the reverse reaction is slow. As products accumulate, the reverse reaction becomes faster. Eventually, the two rates become equal.
3. The equilibrium constant Kc
For a reaction at a fixed temperature, the equilibrium constant expresses the relationship between equilibrium concentrations of products and reactants.
Equilibrium constant in concentration form
For aA + bB ⇌ cC + dD, the square brackets represent equilibrium concentrations. Exponents are the coefficients in the balanced chemical equation.
Example
Consider the reaction:
If the equilibrium concentrations are [HI] = 0.80 mol L⁻¹, [H₂] = 0.20 mol L⁻¹ and [I₂] = 0.20 mol L⁻¹, then:
- Square the HI concentration: 0.80² = 0.64.
- Multiply the reactant concentrations: 0.20 × 0.20 = 0.04.
- Divide: Kc = 0.64 / 0.04 = 16.
For this example, Kc = 16 using the stated concentration convention.
Remember: Pure solids and pure liquids are omitted from the usual equilibrium-constant expression because their activities are treated as constant.
4. The equilibrium constant Kp
For gaseous reactions, an equilibrium constant can also be expressed using partial pressures. This form is called Kp.
Equilibrium constant in pressure form
Each P represents the equilibrium partial pressure of the corresponding gas.
Relationship between Kp and Kc
For ideal gases, Δn(g) equals total gaseous product coefficients minus total gaseous reactant coefficients. Use compatible pressure, concentration, gas-constant and temperature units.
For H₂(g) + I₂(g) ⇌ 2HI(g), Δn(g) = 2 − (1 + 1) = 0. Therefore, Kp = Kc for this reaction under the ideal-gas relation.
5. What does the value of K tell us?
The size of an equilibrium constant indicates whether products or reactants are favoured in the equilibrium mixture. It does not tell us how quickly equilibrium will be reached.
| Value | General interpretation |
|---|---|
| K ≫ 1 | Products are strongly favoured at equilibrium. |
| K ≪ 1 | Reactants are strongly favoured at equilibrium. |
| K around 1 | Neither side is overwhelmingly favoured solely by the value of K. |
The equilibrium constant applies to a particular balanced reaction at a particular temperature. If the reaction equation is reversed, its equilibrium constant becomes 1/K.
6. Reaction quotient Q
The reaction quotient has the same mathematical form as the equilibrium constant, but it uses the concentrations or partial pressures at the current moment, which may not be equilibrium values.
Compare Q with K
These comparisons predict the direction in which the system must shift to approach equilibrium under unchanged conditions.
Worked reasoning
Suppose a reaction has Kc = 10 and the current reaction quotient is Qc = 2. Since Qc is smaller than Kc, the mixture has relatively too little product compared with its equilibrium composition. The net reaction proceeds in the forward direction.
7. Le Chatelier's principle
When a system at equilibrium experiences a change in concentration, pressure or temperature, its composition shifts in a direction that tends to oppose the imposed change.
Changing concentration
- Adding a reactant generally favours the forward direction.
- Removing a reactant generally favours the reverse direction.
- Adding a product generally favours the reverse direction.
- Removing a product generally favours the forward direction.
Changing pressure or volume
For gaseous equilibria, decreasing volume increases pressure. The system tends to shift toward the side with fewer moles of gas. Increasing volume tends to favour the side with more moles of gas.
If both sides have the same total number of gaseous stoichiometric units, changing pressure by changing volume does not shift the ideal-gas equilibrium position.
Changing temperature
Temperature changes can alter the equilibrium constant itself. Treat heat as a reactant for an endothermic forward reaction and as a product for an exothermic forward reaction.
| Forward reaction | Temperature increase | Equilibrium constant |
|---|---|---|
| Endothermic (ΔH > 0) | Favours the forward direction | K increases |
| Exothermic (ΔH < 0) | Favours the reverse direction | K decreases |
What about a catalyst?
A catalyst increases the rates of both forward and reverse reactions. It helps the system reach equilibrium faster but does not change the equilibrium constant or the equilibrium composition at a fixed temperature.
8. Industrial example: the Haber process
Ammonia is manufactured industrially by reacting nitrogen with hydrogen.
- The forward reaction is exothermic.
- Higher pressure favours ammonia because the product side has fewer moles of gas.
- Lower temperature favours the exothermic forward direction at equilibrium.
- Industry uses a compromise temperature to obtain a useful reaction rate.
- A catalyst helps the system reach equilibrium faster.
- Removing ammonia as it forms can encourage further product formation.
Key distinction: Conditions that favour the equilibrium yield are not always the conditions that produce the fastest reaction. Industrial processes balance both.
9. Key ideas to remember
- At dynamic equilibrium, forward and reverse reaction rates are equal.
- Kc uses equilibrium concentrations; Kp uses equilibrium partial pressures.
- The exponents in equilibrium expressions come from the balanced equation.
- The reaction quotient Q predicts the direction of net change by comparison with K.
- Concentration and pressure changes can shift equilibrium, but do not change K at fixed temperature.
- Temperature changes can change K.
- A catalyst speeds up the approach to equilibrium without changing its position.