Does changing pressure change the equilibrium constant?
Short answer: Only temperature changes K. Pressure shifts the position toward fewer gas moles.
The trap
The wrong idea
“Changing pressure changes the equilibrium constant.”
It sounds reasonable, and a lot of JEE aspirants carry it into the exam. Here's a one-question check, then the explanation that makes it click.
Quick check
Try it before you peek
Compressing N₂ + 3H₂ ⇌ 2NH₃ at constant temperature:
- Changes K
- Shifts right, K unchanged
- Shifts left
Show the answer
B. Shifts right, K unchanged
Only temperature changes K. Pressure shifts the position toward fewer gas moles.
The click
For a given reaction, K depends only on temperature. Changing the pressure, by compressing or expanding the container, can change the position of equilibrium, meaning the amounts of each substance, but not K.
Take N₂ + 3H₂ ⇌ 2NH₃, with 4 moles of gas on the left and 2 on the right. Compress the mixture and every partial pressure rises. The reaction quotient Q = p(NH₃)² / (p(N₂)·p(H₂)³) has more pressure factors below the line than above, so Q immediately falls below K. The system responds by forming more NH₃ until Q equals K again. It shifts right, toward fewer moles of gas, as Le Chatelier's principle predicts, and K ends exactly where it started.
If both sides have the same number of gas moles, as in H₂ + I₂ ⇌ 2HI, a pressure change does not shift the position at all. Adding an inert gas at constant volume has no effect either, because the partial pressures of the reacting gases don't change.
Temperature is different: it changes K itself. For an exothermic reaction such as ammonia synthesis (ΔH < 0), raising the temperature lowers K. That is the key distinction for exams: changes in concentration and pressure change Q, and the system moves back to the same K; a change in temperature changes K.
Got this one? The Knot Scan finds the misconception you don't know you have, in 8 questions.