Le Chatelier's Principle: MDCAT Chemistry notes
Le Chatelier's Principle MDCAT notes: effects of concentration, pressure, temperature and catalyst on equilibrium, with BiCl3, SO3 and ammonia examples.
The principle
Le Chatelier's principle: if a stress (a change in concentration, pressure or temperature) is applied to a system at equilibrium, the system shifts in the direction that nullifies (opposes) the effect of the stress as far as possible. It lets us predict the direction in which an equilibrium will shift.
Effect of concentration
- Adding a reactant or removing a product shifts the equilibrium forward (towards products).
- Removing a reactant or adding a product shifts it backward.
- $K_c$ itself does not change; only the position shifts.
Example: $\mathrm{H_2 + CO_2 \rightleftharpoons H_2O + CO}$. Removing $\mathrm{CO_2}$ shifts the equilibrium to the left. In $\mathrm{2SO_2 + O_2 \rightleftharpoons 2SO_3}$, adding more $\mathrm{SO_2}$ increases the yield of $\mathrm{SO_3}$.
Bismuth chloride hydrolysis
$$\mathrm{BiCl_3 + H_2O \rightleftharpoons BiOCl + 2HCl}$$
BiOCl is a white insoluble solid, so the solution looks cloudy. Adding HCl shifts the equilibrium to the left and the solution becomes clear. Adding water shifts it right again and the cloudiness returns.
Effect of pressure (gaseous reactions)
- Increasing pressure shifts the equilibrium towards the side with fewer moles of gas, i.e. the direction of decreased volume.
- Decreasing pressure favours the side with more gas moles.
- If moles of gaseous reactants equal moles of gaseous products ($\Delta n = 0$), as in $\mathrm{H_2 + I_2 \rightleftharpoons 2HI}$, pressure has no effect.
Example: in $\mathrm{N_2 + 3H_2 \rightleftharpoons 2NH_3}$ (4 moles to 2), the same amounts of $\mathrm{N_2}$ and $\mathrm{H_2}$ in a smaller vessel are at higher pressure, so more ammonia forms.
Effect of temperature
Temperature is the only factor that changes the value of $K$.
| Forward reaction | Raise temperature | Lower temperature |
|---|---|---|
| Exothermic ($\Delta H$ negative) | Shifts backward; $K$ decreases | Shifts forward; $K$ increases |
| Endothermic ($\Delta H$ positive) | Shifts forward; $K$ increases | Shifts backward; $K$ decreases |
Examples: $\mathrm{N_2 + O_2 \rightleftharpoons 2NO}$ is endothermic, so its $K$ rises with temperature. Combustion of methane, neutralization, $\mathrm{2SO_2 + O_2 \rightleftharpoons 2SO_3}$ and ammonia synthesis are exothermic, so their $K$ falls with rising temperature.
Effect of a catalyst
A catalyst speeds up the forward and reverse reactions equally. Equilibrium is reached sooner, but its position and $K$ are unchanged.
Key formulas
- $\Delta n = n_{\text{gaseous products}} - n_{\text{gaseous reactants}}$; pressure matters only if $\Delta n \neq 0$.
- $K_p = K_c(RT)^{\Delta n}$
Common MDCAT traps
- Removing a product shifts the equilibrium towards the product side (forward), to replace it.
- For an endothermic reaction, decreasing temperature shifts it in reverse.
- Pressure does not affect equilibria with equal gas moles on both sides.
- Adding a catalyst never increases yield.
- For $K$ to increase with temperature, pick the endothermic reaction. Check the sign of $\Delta H$ rather than guessing.
Quick revision
- A system at equilibrium opposes any imposed change.
- High pressure favours fewer gas molecules.
- Heating favours the endothermic direction.
- Only temperature changes the equilibrium constant.
- HCl clears the cloudy $\mathrm{BiCl_3}$ solution.