temperature dependence of equilibrium
Of all the things you can do to a chemical equilibrium — add reactant, squeeze it, dilute it — temperature is the special one. Concentration and pressure only slide the mixture to a new spot that still satisfies the same K. Temperature is different: it actually changes K itself, the constant that defines where equilibrium lies. Heating or cooling does not just move the seesaw; it rebuilds the seesaw.
The direction of the change follows from whether the reaction releases or absorbs heat. Treat heat as a participant. For an exothermic reaction, heat is effectively a product, so adding heat (raising temperature) pushes the equilibrium backward and lowers K. For an endothermic reaction, heat is effectively a reactant, so adding heat pushes it forward and raises K. The van 't Hoff equation puts exact numbers on this trend.
This is why temperature is the one true lever on K and why industrial chemists agonise over it. The honest tension is that the temperature best for the equilibrium yield is often the worst for reaction speed: an exothermic reaction gives a higher yield when cool, yet runs hopelessly slowly there. Real processes therefore compromise, often with a catalyst to buy back the lost speed.
The brown-to-colourless conversion 2NO2 ⇌ N2O4 is exothermic. Warm the tube in hot water and it darkens — the equilibrium retreats toward brown NO2 and K falls. Chill it in ice and it pales — the equilibrium advances toward N2O4 and K rises.
Temperature is the only change that alters K, not just the mixture.
Treating heat as a reactant or product is a memory aid for direction only. It correctly predicts which way K moves, but the precise size of the change comes from the van 't Hoff equation, not from this picture.