Le Chatelier Principle Calculator

Predict how equilibrium shifts when concentration, pressure, or temperature is changed.

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Why Equilibrium Shifts

Le Chatelier’s principle states that a system at equilibrium, when disturbed, responds so as to partially oppose the disturbance. It is a useful predictive shortcut, but the underlying mechanism is worth understanding because it explains the exceptions.

The real driver is the reaction quotient Q compared with the equilibrium constant K. At equilibrium Q = K. Disturb the system and Q moves away from K; the reaction then proceeds in whichever direction restores equality. If Q < K the forward reaction runs; if Q > K the reverse runs.

DisturbanceDirection of shiftDoes K change?
Add reactantToward productsNo
Remove productToward productsNo
Increase pressure (compress)Toward the side with fewer gas molesNo
Increase temperature, exothermicToward reactantsYes — K decreases
Increase temperature, endothermicToward productsYes — K increases
Add a catalystNo shiftNo
Add inert gas at constant volumeNo shiftNo

That column matters more than it looks. Only temperature changes K. Concentration and pressure changes move Q away from K and the system returns to the same K by shifting composition. Temperature changes K itself, which is why heating an exothermic reaction genuinely reduces the achievable yield rather than merely redistributing it.

This is the tension at the heart of the Haber process. Ammonia synthesis is exothermic, so low temperature favours yield — but low temperature makes the reaction impossibly slow. Industrial practice settles near 450°C with high pressure and a catalyst, accepting reduced equilibrium yield in exchange for a workable rate.

Worked Examples

Example 1: Add N2 to N2+3H2=2NH3
Shifts RIGHT, more NH3 formed
Result: Q becomes < K, forward reaction proceeds
Industrial: remove NH3 product continuously
Example 2: Increase T for exothermic reaction
Shifts LEFT, K decreases
Result: Less product formed at higher T
Haber process: compromise T for rate vs yield
Example 3: Adding argon at constant volume
N2 + 3H2 ⇄ 2NH3, inert gas added, volume fixed
Result: No shift
Total pressure rises but each partial pressure is unchanged, so Q still equals K. A frequent exam trap.
Example 4: Equal moles of gas on both sides
H2 + I2 ⇄ 2HI, pressure increased
Result: No shift
Two moles of gas on each side means compression affects both directions equally. Pressure changes only shift equilibria with unequal gas moles.
Example 5: Removing product continuously
Esterification with water distilled off as it forms
Result: Shifts strongly toward products
Continuously removing a product holds Q below K, driving the reaction far past its normal equilibrium position. This is the basis of Dean-Stark operation.

Common Mistakes

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Believing a catalyst improves yield

A catalyst lowers the activation barrier for forward and reverse reactions equally. It reaches equilibrium faster but does not change K or the equilibrium position at all.

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Including solids and pure liquids in pressure arguments

Only gas-phase moles count when assessing a pressure change. Solids and pure liquids have essentially fixed activity and do not appear in the equilibrium expression.

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Assuming adding inert gas causes a shift

At constant volume, adding argon raises total pressure but leaves every partial pressure unchanged, so Q is unaffected. Only at constant pressure, where the volume expands, does dilution cause a shift.

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Thinking concentration changes alter K

They do not. K depends only on temperature. Concentration and pressure changes move Q away from K, and the system shifts composition to restore the same K.

Frequently Asked Questions

Does catalyst affect equilibrium?
No. Catalyst lowers Ea equally for forward and reverse. Equilibrium position and K unchanged. Reaction reaches equilibrium faster only.
Pressure and liquids/solids?
Pressure changes only affect gases. Pure liquids and solids have activity=1 regardless of amount. Dissolved species in solution: pressure changes negligible.
Does a catalyst shift equilibrium?
No. It speeds up forward and reverse reactions equally, so equilibrium is reached sooner but at exactly the same position. K is unchanged.
Why does only temperature change K?
Because K is determined by the standard Gibbs free energy change, which is temperature dependent. Concentration and pressure changes alter Q, not K, so the system simply returns to the same K.
Do solids and liquids affect equilibrium shifts?
Pure solids and liquids have constant activity and are excluded from the equilibrium expression. Adding more solid does not shift equilibrium, and only gas moles matter for pressure effects.
What happens when inert gas is added?
At constant volume, nothing — partial pressures are unchanged. At constant pressure the mixture expands, partial pressures fall, and the equilibrium shifts toward more gas moles.
Why does the Haber process use high temperature if the reaction is exothermic?
Because low temperature favours yield but makes the rate impractically slow. Around 450°C is a compromise between an acceptable equilibrium yield and a usable reaction rate.

Formula Explorer connections

Interpretation: This formula describes how reactants, products, ions or phases distribute when opposing processes reach equilibrium. Assumption: Use equilibrium rather than initial concentrations, correct stoichiometric exponents, and the specified temperature; activities may replace concentrations in nonideal systems.

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