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What is Le Chatelier's Principle?

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University Course Reader · STEM

Applying Le Chatelier’s principle, an increase in pressure shifts the equilibrium toward the side with fewer moles of gas.

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Overview

A reaction at equilibrium makes products and reactants at the same speed, so the amounts stay steady. Change something and one direction runs faster until the two speeds match again. The reaction does not return to its old amounts — it settles at a new balance.
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Overview

Poke a balanced reaction and it pokes back — add reactant, it makes more product; take something away, it makes more to replace it. But it never rewinds to the old amounts. It just finds a new balance somewhere else and calls it even. 😎

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Detail

Equilibrium does not mean stopped: both directions keep running at matched speeds, so the amounts hold still. Add reactant and collisions increase, the forward direction speeds up, and product builds until the reverse direction catches up and the speeds tie again. At that new balance there is more product than before and also more reactant, because only part of what you added was consumed. Removing product as it forms works the opposite way and pulls the reaction forward continuously, which is why ammonia plants drain product off as it appears. Squeezing a gas mixture shifts it toward the side with fewer gas molecules, since fewer molecules occupy less volume. Heat can be written into the equation as a participant: heating a reaction that releases heat shifts it back toward reactants, and cooling shifts it toward products. A catalyst changes none of this, speeding both directions equally so the same new balance arrives sooner.
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Detail

Equilibrium isn't a nap — both directions run full speed, tied, so the amounts sit still. Dump in reactant: more collisions, the forward direction pulls ahead, product piles up until the reverse catches up and they tie again. New tie, new numbers — more product than before AND more reactant, because only some of what you dumped got used. Steal the product as fast as it appears and the reaction chases that gap forever, which is exactly how factories keep one running. Squeeze a gas mixture and it runs to whichever side has fewer molecules, because fewer molecules need less room. Classic trap: catalysts shift NOTHING, both directions just get faster, same finish line, less waiting.

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Analogy

A concert lets out, hundreds of people want rides within three blocks, and prices surge — which pays drivers to leave their corners and come over. More drivers arrive and prices ease, but they never drop back to the normal fare, because the crowd is still standing there. The area settles at a new steady price with more drivers working it.
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Analogy

Two checkout lines, six people each, nobody switching because switching gains nothing. A tour bus empties into line A, people at the back drift over to B, and the drifting stops the second it stops being worth it: eleven and eleven. Balanced again — but you never get your six back, because the lines rebalanced against each other, not against the bus.
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Formal definition — The same term, explained the usual way

Le Chatelier's principle states that a system at chemical equilibrium subjected to a change in concentration, pressure, or temperature will adjust in the direction that partially offsets the imposed change, establishing a new equilibrium position. Catalysts alter the rate at which equilibrium is attained without displacing its position.

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