Temperature Changes

Temperature changes are fundamentally different from concentration and pressure changes. Changing concentration or pressure shifts the equilibrium, but K stays the same. Changing temperature actually changes the value of K itself.

This is the single most important fact in this section: temperature is the ONLY factor that changes K.

The Heat-as-a-Species Trick

The easiest way to predict the effect of temperature is to treat heat as if it were a chemical species:

• Exothermic reaction (delta-H < 0): heat is a “product”
  • A + B ⇌ C + D + heat
• Endothermic reaction (delta-H > 0): heat is a “reactant”
  • heat + A + B ⇌ C + D

Now apply Le Chatelier’s principle as if heat were a concentration:

Summary Table

Why K Changes (Unlike With Concentration/Pressure)

When you change concentration, Q changes but K stays fixed - the system shifts to bring Q back to K. But when you change temperature, K itself changes to a new value. The system then shifts so that the current Q moves toward this NEW K.

This happens because K depends on the rate constants $k_{f}$ and $k_{r}$ (recall K = $k_{f}$/$k_{r}$). Temperature changes both rate constants according to the Arrhenius equation, but it changes them by different amounts depending on the activation energies. The net result is that K itself changes.

Example: The Haber Process

N₂(g) + 3H₂(g) ⇌ 2NH₃(g), delta-H = -92 kJ/mol (exothermic)

• Increasing temperature: Shifts left (fewer products), K decreases
• Decreasing temperature: Shifts right (more products), K increases

This creates a dilemma in industrial chemistry: low temperature gives a better K (more product), but the reaction is too slow. The compromise is moderate temperature (~450 degrees C) with a catalyst to speed things up.