When an external change in concentration, pressure, or temperature is applied to a system at equilibrium, the equilibrium shifts in the direction that counteracts the change. This is Le Chatelier's principle (the principle of equilibrium shift).
Basic Knowledge
Factors and equilibrium shifts:
① Concentration: Increasing reactants shifts equilibrium in the forward direction; increasing products shifts it in the reverse direction
② Pressure: Increasing pressure shifts equilibrium toward the side with fewer gas molecules (no effect for liquids and solids or when the numbers of molecules on both sides are equal)
③ Temperature: Increasing temperature shifts equilibrium in the endothermic direction
④ Catalysts: Do not change the equilibrium position (they only shorten the time required to reach equilibrium)
⑤ Addition of an inert gas: No effect at constant volume; at constant pressure, expansion shifts equilibrium toward the side with more molecules
Le Chatelier's principle（Equilibrium shifts in the direction that reduces the effect of a change）
Concentration change（Shifts in the direction that decreases the increased component）
Pressure change（Shifts toward the side with fewer moles）
Temperature change（Shifts in the endothermic direction to reduce the effect）
Haber process（High pressure and low temperature are favorable, but a catalyst is used as a compromise for rate）
Contact process（V2O5 converts SO2→SO3. Low temperature is favorable because it is exothermic, but approximately 450℃ is used to maintain the rate）
Further Exploration (Principles and Applications)
For the Haber–Bosch process(N2+3H2⇌2NH3, ΔH<0), low temperature and high pressure are favorable from an equilibrium perspective, but because the rate is too slow at low temperatures, a compromise of around 500℃, approximately 200 atm, and an Fe catalyst is used.
The Contact process also involves SO2+½O2⇌SO3(ΔH<0): low temperature favors equilibrium, while high temperature favors the rate. It has been industrialized using a V2O5 catalyst system at approximately 450℃. The “trade-off between equilibrium and rate” is a key point in industrial chemistry.
- Principle: shifts in the direction that reduces the change
- Concentration↑ ⇒ direction of consumption
- Pressure↑ ⇒ direction with fewer molecules
- Temperature↑ ⇒ endothermic direction
- Catalysts have no effect on the equilibrium position
- Haber process=high pressure and moderate temperature
- Contact process=moderate temperature and catalyst
Points to Note (Common Confusions and Frequent Mistakes)
① For pressure changes, consider the difference in moles of gases. There is no effect if the numbers on both sides are equal. ② Temperature changes K itself, whereas concentration and pressure changes do not change K (they only change the concentration ratio). ③ Addition of an inert gas: no effect at constant volume. ④ A catalyst changes the rate but not the equilibrium position.
Practice
- For the equilibrium N2+3H2⇌2NH3(ΔH<0), how does the equilibrium shift in each case: (a) pressure↑, (b) temperature↑, (c) addition of an Fe catalyst?
- Give one example of a reaction in which a pressure change does not affect equilibrium.
- Explain why a temperature of 500℃ is chosen for the Haber process from both equilibrium and rate perspectives.