Changing an Equilibrium

A disturbance changes what must balance

An equilibrium mixture can respond when its surroundings or composition change. Le Chatelier’s principle summarises this response: the net change tends to reduce the imposed disturbance. It is a qualitative guide, not a claim that the original state is restored completely. The quotient test provides a more precise way to justify many predictions.

Adding a reactant at fixed volume generally lowers the corresponding product-to-reactant quotient and promotes its consumption. Removing a product also promotes forward change. K remains the same at fixed temperature. After a reactant is added, some of it is consumed, but its new equilibrium amount need not fall all the way back to its old value. A shift is a net reaction during readjustment, not a permanent difference between the final opposing rates.

Colour changes make concentration effects visible

Mixing dilute iron(III) and thiocyanate solutions forms a deep-red complex, $\mathrm{FeSC}{\mathrm{N}}^{2+}$. At equilibrium its colour intensity becomes steady. Adding thiocyanate increases the red colour because more complex forms. Removing free iron(III) or thiocyanate encourages the complex to dissociate and its red colour fades.

Oxalate can bind iron(III), and mercury(II) can bind thiocyanate, removing free reactant ions from this equilibrium. These observations illustrate competing equilibria: a reagent can change the free concentration even when the element remains somewhere in solution. They are teacher-controlled demonstrations; the mercury reagent is not needed to reproduce the idea in independent practice. In industry, removing ammonia by condensation or removing CO2 from a lime kiln similarly promotes further product formation.

Compression depends on gas stoichiometry

Compressing an ideal gas mixture at constant temperature raises all its partial pressures and concentrations initially. If the two sides have different numbers of gaseous molecules, the quotient changes. The subsequent net reaction favours the side with fewer gas moles. Expansion favours the side with more gas moles. Pure solid and liquid coefficients do not enter this gas count.

For $\mathrm{C}(\mathrm{s})+\mathrm{C}{\mathrm{O}}_2(\mathrm{g})\rightleftharpoons 2CO(\mathrm{g})$, compression promotes the reverse direction because one gaseous mole is favoured over two. If the gas-mole change is zero, uniform compression leaves the ideal-gas quotient unchanged. Adding an inert gas at constant volume raises total pressure but leaves reacting gases’ partial pressures unchanged, so their equilibrium is not shifted. The constant-volume condition is essential; an expansion performed to maintain constant total pressure is a different disturbance.

Temperature changes the equilibrium constant

Heating does more than change the current quotient: it changes K itself. For an exothermic forward reaction, higher temperature favours the reverse, heat-absorbing direction and K decreases. For an endothermic forward reaction, higher temperature favours forward change and K increases. Temperature can also alter the reaction rates, so its effect on speed must be considered separately.

For $2N{\mathrm{O}}_2(\mathrm{g})\rightleftharpoons {\mathrm{N}}_2{\mathrm{O}}_4(\mathrm{g})$, the forward reaction is exothermic. Brown NO2 gives way to less-coloured N2O4 on cooling. In a demonstration, sealed samples placed in cold, room-temperature and warm baths show progressively different brown intensities. The gases and preparation reagents require appropriate laboratory controls; the useful evidence is the reversible colour response to temperature.

Another example is the pink hexaaquacobalt(II) complex exchanging ligands with chloride to form blue tetrachlorocobaltate(II): $[\mathrm{Co}({\mathrm{H}}_2\mathrm{O}{)}_6{]}^{2+}+4C{\mathrm{l}}^{-}\rightleftharpoons [\mathrm{CoC}{\mathrm{l}}_4{]}^{2-}+6{\mathrm{H}}_2\mathrm{O}$. The blue direction is endothermic under the described conditions, so warming favours blue and cooling favours pink. Both complexes contain cobalt in oxidation state +2.

A catalyst changes the route, not the destination

A catalyst provides a pathway with lower barriers and accelerates approach to equilibrium in both directions. It does not change the reaction’s overall standard Gibbs energy, K or equilibrium composition. A catalyst cannot make an equilibrium with negligible useful product become product-favoured. It can make a favourable but slow conversion practically accessible.