Concept 01 / 08Equilibrium
Le Chatelier's Principle: Effect of Concentration, Pressure, Temperature and Catalysts
Le Chatelier's principle applied to changes in concentration, pressure or volume, temperature and addition of inert gas, the effect of a catalyst, why only temperature changes K, and industrial applications to the Haber and Contact processes.
Le Chatelier's principle. If a system at equilibrium is disturbed by a change in concentration, pressure or temperature, the equilibrium shifts in the direction that tends to counteract the change.
Concentration. Adding a reactant (or removing a product) makes ; the reaction proceeds forward until again. Adding a product shifts it backward. itself does not change. In the figure, adding to an ammonia equilibrium consumes some of it together with and forms more . Removing a product continuously (for example escaping from a lime kiln, or water removed in esterification) drives a reaction to completion.
Pressure (by changing volume). Compressing a gaseous equilibrium shifts it towards fewer moles of gas; expanding shifts it towards more moles. If , pressure has no effect (). For ( moles), high pressure favours ammonia. Solids and liquids are hardly affected.
Inert gas. Added at constant volume: partial pressures of the reacting gases are unchanged, so there is no shift. Added at constant total pressure: the volume increases, which acts like a decrease in pressure, and the equilibrium shifts towards more moles of gas.
Temperature. Temperature is the only factor that changes . Raising the temperature favours the endothermic direction; lowering it favours the exothermic direction. For an exothermic forward reaction (), decreases as temperature rises; for an endothermic one, increases. Quantitatively (van 't Hoff equation):
.
Example: (colourless) ⇌ (brown) is endothermic; the mixture turns darker brown on heating.
Catalyst. A catalyst speeds up the forward and backward reactions equally, by lowering the activation energy of both. Equilibrium is reached faster, but its position and are unchanged.
Industrial applications.
Haber process, , : high pressure (about ) favours ammonia; low temperature favours it too but makes the reaction too slow, so a compromise of about with an iron catalyst (molybdenum promoter) is used, and ammonia is removed by liquefaction.
Contact process, , exothermic: excess air, moderate temperature (about ), catalyst, pressure of to (conversion is already high).
Key Takeaways:
- Concentration and pressure changes shift the position but not ; temperature changes .
- Heating favours the endothermic direction.
- Compression favours fewer gas moles; no effect if .
- Inert gas at constant : no effect; at constant : shift towards more gas moles.
- Catalysts do not shift equilibrium.
- Pitfall: melting of ice () is favoured by pressure, which is why ice melts under pressure; the principle applies to physical equilibria too.
Brain-Teaser 1: For , , predict the effect of (a) raising temperature, (b) compressing the mixture, (c) adding , (d) adding argon at constant pressure.
Solution: (a) Forward (endothermic), increases. (b) Backward, towards fewer moles ( instead of ). (c) Backward, unchanged. (d) Volume increases, so forward, towards more moles.
Brain-Teaser 2: Why does the yield of ammonia fall at higher temperature, yet industrial plants do not operate at room temperature?
Solution: The reaction is exothermic, so decreases with temperature and the equilibrium yield falls. But at low temperatures the rate is impractically slow even with a catalyst. About is a compromise between yield and rate, compensated by high pressure and recycling of unreacted gases.
Brain-Teaser 3: of a reaction doubles when the temperature rises from to . Estimate .
Solution: . . (positive, since increased with ).
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