Equilibrium constants
Equilibrium constants
- The equilibrium constant links the amounts of products and reactants.
- For concentrations we use $K_c$; for gases, $K_p$.
- Only temperature changes its value.
$K_c$
For $a\text{A} + b\text{B} \rightleftharpoons c\text{C} + d\text{D}$:
$$K_c = \frac{[\text{C}]^c\,[\text{D}]^d}{[\text{A}]^a\,[\text{B}]^b}$$
- Square brackets mean concentration in $\dfrac{\text{mol}}{\text{dm}^3}$; the powers are the equation's numbers.
Practice
In the expression for Kc, the concentrations of the products are:
Kc = [products]^powers / [reactants]^powers, with the powers taken from the balanced equation.
$K_p$ and partial pressure
- For gases we use partial pressures ($K_p$).
- The partial pressure of a gas is its share of the total:
$$\text{partial pressure} = \text{mole fraction} \times \text{total pressure}$$
Practice
The partial pressure of a gas equals:
Partial pressure = mole fraction × total pressure; Kp is written using these.
Practice
Kp is written using partial pressures rather than concentrations.
For gas-phase equilibria, Kp uses the partial pressures of the gases.
What changes K
- Only temperature changes the value of $K_c$ or $K_p$.
- Changing concentration or pressure, or adding a catalyst, shifts the position but leaves $K$ unchanged.
Practice
Which change alters the value of the equilibrium constant?
Only temperature changes K; the others shift the position of equilibrium but leave K unchanged.
You've got it
Key idea
- $K_c = \dfrac{[\text{C}]^c[\text{D}]^d}{[\text{A}]^a[\text{B}]^b}$ (powers from the equation; $\dfrac{\text{mol}}{\text{dm}^3}$)
- $K_p$ uses partial pressures = mole fraction × total pressure
- only temperature changes $K$; concentration/pressure/catalyst shift the position only