Push a system at equilibrium, and it pushes back. That is the essence of Le Chatelier's principle. If you change the conditions of a reversible reaction at equilibrium, the system shifts in the direction that counteracts the change. Add a reactant, and the system consumes it by making more product. Remove a product, and the system makes more to replace it. Increase the pressure, and the system shifts toward the side with fewer gas molecules. Raise the temperature, and the system shifts in the endothermic direction.

The principle is qualitative, but it is powerful. It guides the optimization of industrial reactions. The Haber-Bosch process for ammonia synthesis runs at high pressure and moderate temperature to favor the product. The contact process for sulfuric acid uses excess oxygen and removes sulfur trioxide as it forms. Both are designed around Le Chatelier's principle.

The principle applies to concentration, pressure, temperature, and volume changes. It does not apply to catalysts, which speed up the approach to equilibrium without changing its position. It also does not apply to adding an inert gas at constant volume, which does not change partial pressures and therefore does not shift the equilibrium. These nuances matter in practice.