Why does Le Chatelier's principle shift toward products

A system at equilibrium has forward and reverse rates equal. Any disturbance creates an imbalance, and the reaction moves in the direction that reduces the disturbance. This shift does not change the equilibrium constant; it only changes the concentrations until the original ratio is re‑established.

Haber‑process example

Consider the exothermic synthesis of ammonia: $$\text{N}_2(g)+3\text{H}_2(g)\rightleftharpoons2\text{NH}_3(g)$$ At 400 °C, $K_p\approx1.6\times10^{-4}$. Doubling the total pressure from 100 atm to 200 atm reduces the total volume, favoring the side with fewer gas moles (2 mol of NH₃), so the equilibrium shifts toward products, increasing $[\text{NH}_3]$ while $K_p$ stays the same.

Typical stresses that push the equilibrium toward products:

• Adding one or more reactants
• Removing one or more products
• Increasing pressure when Δn_{gas}<0
• Lowering temperature for an exothermic reaction

How to predict the direction of shift:

1. 1Write the balanced equation and identify Δn_{gas} (moles of gas products − reactants).
2. 2Determine whether the reaction is exothermic or endothermic (look up ΔH).
3. 3Apply the stress: concentration, pressure, or temperature change.
4. 4If the stress favors the side with lower total pressure or opposite temperature effect, the shift is toward that side (usually products).

Stress versus shift direction: