Le Chatelier's Principle: Factors Affecting Equilibrium

An equilibrium mixture is not frozen forever — if the conditions under which it was established are

disturbed, the system responds by shifting toward a new equilibrium position. Le Chatelier's principle provides a simple, purely qualitative rule for predicting the direction of that shift,

without needing to solve the equilibrium-constant expression at all: if a system at equilibrium is subjected to a change in concentration, pressure, volume, or temperature, the equilibrium shifts in the direction that tends to counteract, or partially undo, the effect of that change.

Effect of concentration change. Adding more of a reactant increases its concentration; the system

responds by consuming some of that excess, shifting the equilibrium forward (toward products) to

partially reduce the added reactant's concentration back down. Conversely, removing a product as it

forms (for instance, distilling it off) continually shifts the equilibrium forward, since the system

is perpetually trying to replace the product being taken away.

Effect of pressure and volume change. For a gaseous equilibrium in which the two sides have

different numbers of moles of gas, decreasing the volume (which raises the total pressure) shifts the

equilibrium toward whichever side has fewer moles of gas, since that side occupies less volume and

so partially relieves the increase in pressure. In ammonia synthesis,

$\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)$, the product side has only 2 moles

of gas against 4 moles on the reactant side, so compressing the mixture shifts the equilibrium

forward, favouring $\text{NH}_3$ — this is exactly why the industrial Haber process is run at high

pressure.

Effect of inert gas addition. Adding an inert gas such as argon at constant volume does not

change any of the reacting species' own partial pressures or concentrations, so it has no effect on

the equilibrium position at all. Adding it at constant total pressure, however, requires the vessel

to expand, which dilutes every reacting species and lowers its partial pressure — the equilibrium

then shifts toward the side with the greater number of moles of gas, exactly as if the pressure had

been lowered directly.

Effect of temperature change. Unlike concentration and pressure, changing the temperature

actually changes the numerical value of $K$ itself, not merely the position of a fixed equilibrium.

For an exothermic reaction (heat is a "product"), raising the temperature shifts equilibrium

backward, decreasing $K$ and lowering the equilibrium yield of product; for an endothermic reaction

the opposite occurs. This is why, even though ammonia synthesis is exothermic and would in principle

give the highest yield at low temperature, industrial plants operate at a moderately elevated …