Le Chatelier's principle
/ luh-SHAT-lee-ay /
Imagine a hammock with a person resting in the middle. Push down on one side and the whole hammock shifts to take up the slack, settling into a new balanced shape that partly absorbs your push. A chemical system at equilibrium behaves the same way. Le Chatelier's principle says that if you disturb a system at equilibrium, the system responds by shifting in the direction that partly opposes — and so relieves — the disturbance.
In practice the 'disturbances' are familiar: add more of a substance and the system consumes some of it; remove a substance and the system makes more of it; raise the pressure on a gas mixture and the equilibrium shifts toward the side with fewer gas molecules; heat the mixture and it shifts in the direction that absorbs heat (the endothermic direction). Each time, the equilibrium moves to a new position that softens the change you imposed.
It is a wonderfully handy rule of thumb, but treat it as a qualitative guide, not a law of nature. It predicts the direction of a shift, not the exact amount, and it can mislead in tricky cases — for example, adding an inert gas at constant volume changes nothing, even though the total pressure rises. The principle is really a shorthand for what the equilibrium constant and reaction quotient would tell you if you did the full calculation.
In the equilibrium N2 + 3H2 ⇌ 2NH3, four gas molecules on the left become two on the right. Squeeze the mixture into a smaller volume and the equilibrium shifts toward ammonia — the side with fewer gas molecules — because that partly relieves the rise in pressure.
Disturb an equilibrium and it shifts to soften the disturbance.
A subtle trap: adding an inert gas at constant volume raises total pressure but does not shift the equilibrium, because it changes neither the partial pressures nor the concentrations of the reacting species.