shifting an equilibrium

Once a reaction has settled at equilibrium, it sits there contentedly — until you change something. Then it stirs and slides to a new resting point. Shifting an equilibrium means deliberately disturbing a settled reaction so that the mixture moves toward more product or more reactant. It is how a chemist coaxes a reluctant reaction to give a better yield, instead of meekly accepting whatever balance it first found.

Mechanically, every shift can be understood through the reaction quotient Q and the constant K. Disturb the system and Q is momentarily knocked away from K; the reaction then runs in whichever direction brings Q back to K. Adding reactant lowers Q, so the system runs forward; removing product also lowers Q, again running forward; compressing a gas mixture changes the partial pressures and pushes toward fewer gas molecules. Le Chatelier's principle is the quick way to call the direction.

The deep point worth holding onto is that there are two genuinely different kinds of shift. Changing concentration, pressure, or volume moves the equilibrium position to a new spot but leaves K unchanged. Changing temperature is the special case: it changes K itself, and so moves the position by redefining where equilibrium lives. Pulling product out as it forms — as the Haber process does — is the workhorse trick of industrial chemistry, continually shifting the equilibrium forward and dragging a reaction toward completion.