An equilibrium is balanced, but it is not stuck. Add more of a reactant, squeeze the container, or heat it, and the concentrations start changing again until a new equilibrium is reached. Chemists who make ammonia or sulfuric acid on an industrial scale deliberately choose conditions that push these equilibria toward the product they want. In this lesson you will learn two tools for predicting the direction of the shift: a calculation (the reaction quotient, Q) and a principle (Le Chatelier's).
The Reaction Quotient, Q
The reaction quotient has exactly the same form as the equilibrium constant expression, but you substitute the concentrations the system has right now, which may not be equilibrium values. Comparing Q with K tells you which way the reaction must run to reach equilibrium.
| Comparison | What it means | Shift |
|---|---|---|
| Too few products compared with equilibrium | Forward reaction is favoured: shifts right | |
| Already at equilibrium | No shift | |
| Too many products compared with equilibrium | Reverse reaction is favoured: shifts left |
Le Chatelier's Principle
Le Chatelier's principle states that when a system at equilibrium is disturbed by a change in conditions (a stress), the system shifts in the direction that partly counteracts that change. It never fully undoes the stress, but it moves to reduce it. Every prediction below can also be justified with Q, which is what makes the principle more than a rule of thumb.
- Adding a reactant or removing a product: Q becomes smaller than K, so the system shifts right, using up some of the added reactant.
- Adding a product or removing a reactant: Q becomes larger than K, so the system shifts left.
- Decreasing volume (increasing pressure) of a gas mixture: the system shifts toward the side with fewer moles of gas, which lowers the pressure. If both sides have the same number of moles of gas, there is no shift.
- Increasing volume (decreasing pressure): shifts toward the side with more moles of gas.
- Adding an inert gas at constant volume: no shift, because the concentrations (and partial pressures) of the reacting gases do not change.
- Adding a catalyst: no shift. A catalyst speeds up the forward and reverse reactions equally, so equilibrium is reached sooner but at the same position.
- Adding a pure solid or liquid: no shift, because it does not appear in Q or K.
You can see the temperature effect with the system from the last lesson. The forward reaction is endothermic, so warming a sealed tube shifts it right and the gas turns a darker brown; cooling it in ice water shifts it left and the colour fades. Another classroom example is the cobalt(II) chloride equilibrium: [\text{Co{(H}_2\text{O)}}_6]^{2+} (pink) (blue) . Adding chloride ions or heating (the forward reaction is endothermic) turns the solution blue.
The Haber Process
Ammonia is made industrially for fertilizers by the Haber process. Le Chatelier's principle explains the conditions. High pressure (typically around 200 atm) favours the right side, which has fewer moles of gas. The reaction is exothermic, so a low temperature would give a larger K and more ammonia at equilibrium, but it would also make the reaction far too slow. Plants compromise with a moderate temperature of roughly 400 to 450 °C and an iron-based catalyst, which gets the mixture to equilibrium quickly without changing K. Ammonia is also continuously removed by cooling it into a liquid, and the unreacted nitrogen and hydrogen are recycled, so the system keeps shifting right.