Dynamic Equilibrium and le Chatelier’s Principle

Reversible Reactions: A reaction that takes place in both forward and
reverse directions if reversible. Equilibrium only in reversible
• Dynamic Equilibrium System: Equilibrium that exists in a closed system when rate of forward
reaction is equal to rate of reverse reaction is dynamic equilibrium. Concentrations of reactants
and products do not change. Both reactions taking place, it is whichever side is dominating. Where
pH would remain constant.
• How Reaches Dynamic Equilibrium: Rate of forward reaction slows down and rate of backward
reaction speeds up until rate of forward reaction is the same as the rate of the backward reaction.
• Closed System: For reaction to remain in equilibrium or to reach dynamic equilibrium, system
must be closed. So temperature, pressure and concentration or reactants/ products unaffected
by outside influences.

• Position of Equilibrium: The relative quantities of reactants and products, indicating the extent of
a reversible reaction at equilibrium.
• le Chatelier’s Principle: When a system in dynamic equilibrium is subjected to an external change
in conditions, the position of equilibrium will shift to minimise the effect of change.
• On a graph, system reaches dynamic equilibrium when concentration are constant.
• Investigating Concentration: Equilibrium between chromate ions CrO4
(yellow solution) and
dichromate ions Cr2O7
(orange solution) changes with acid conc so easy to see shift in
(aq) + 2H+
(aq) ⇌ Cr2O7
(aq) + H2O (l).
• Method: Add yellow potassium chromate K2CrO4. Add sulfuric acid H2SO4 until no further change.
Solution turns orange. Add aqueous sodium hydroxide NaOH(aq) until no further change.
Solutions turns back yellow.
• Increasing Reactant:
– When adding dilute sulfuric acid, increasing concentration of H+ (aq) ions.
– Increases rate of forward reaction.
– Causes position of equilibrium to shift to right minimise change in H+
(aq) concentration. New
position of equilibrium established towards products. Solution turns orange as Cr2O7
• Decreasing Reactant:
– When add NaOH (aq), added OH-
(aq) ions reacts with H+ (aq) ions to form H2O. This decreases
concentration of H+ (aq) ions, the reactant.
– Decreases rate of forward reaction.
– Causes position of equilibrium to shift to left to minimise change in concentration. New
position of equilibrium is established towards reactant. Solution turns yellow as CrO4
• Temperature: Forward and reverse reactions have same value for enthalpy change, but signs are
• Investigating Temperature Method: Dissolve cobalt chloride CoCl2. Add HCl. Place boiling tube in
iced water. Solution is pink colour. Transfer boiling tube into water bath. Solutions turns blue.
Changes pink when back in ice water.
• Concentration: HCl added to provide more Cl-
(aq) ions. Shifts equilibrium slightly towards right
to achieve colour changes.
• Increasing Temperature:
– Increase in temperature, shifts position of equilibrium in the forward direction in endothermic
– So to absorb energy and minimise increase in temperature.
– More products made. Solution turns blue.
• Decreasing Temperature:
– Decreasing temperature, shifts the position of equilibrium in the exothermic direction.
– More reactants made.
• Pressure: Gases only. 2NO2 (g) ⇌ N2O4 (g). Nitrogen dioxide has 2 moles and is brown. Dinitrogen
tetroxide has 1 mol and is colourless.
• Increasing Pressure:
– Increasing pressure will shift the position of equilibrium to the side of fewer moles of gas on
the reactant/ product side.

– Reducing number of gaseous moles to minimise increase in pressure- quote number of moles.
More colourless N2O4 (g) formed and brown colour fades.
• Decreasing Pressure: Brown colours deepens.
• H2 + Br2 ⇌ 2HBR. No change occurs to position of equilibrium when an increase in pressure as
same number of moles on both sides. PCl5 ⇌ PCl3 + Cl2, here product has more moles.
• Catalyst: Does not change position of equilibrium. It speeds up forward and reverse reactions
equally. Increases the rate at which an equilibrium is established. Catalysis take place at lower
temperatures with lower energy demand, so reduce CO2 emissions/ less fossil fuels burnt.
• Haber Process: N2 (g) + 3H2 ⇌ 2NH3 (g). le Chatelier’s principle used to predict best conditions to
produce maximum yield.
• A low temperature would produce a high yield of product, but can’t overcome activation energy
so slow rate of reaction.
• High pressure not only increases yield but also rate of reaction. However need large amounts of
energy so expensive and also safety risks as toxic ammonia could leak under pressure.
• Use a certain value for high enough pressure/ temperature for good rate of reaction without
compromising yield. Increased temperature or pressure, increases rate of reaction- see other
topic. Compromise- choose a higher temperature which creates reduced yield but in shorter space
of time.
• Unreacted nitrogen and hydrogen recycled to be converted into ammonia.