Wait, What? When an equilibrium mixture changes colour, the “old” reaction has not stopped and the “new” reaction has not taken over completely.

At dynamic equilibrium, forward and reverse reactions continue. Their rates are equal, so the macroscopic composition remains constant. Disturb the conditions and the rates become temporarily unequal; the composition changes until a new equilibrium state is established. A colour shift is therefore evidence of a change in relative amounts of coloured species, not proof that one side of the equation has gone to 100% completion.

This practical’s RFE is to teach students how to connect a visible macroscopic signal to a reversible molecular system without turning Le Chatelier’s principle into a magic sentence that replaces mechanism.

The model: equilibrium is dynamic

For a reversible reaction:

A + B ⇌ C + D

equilibrium is reached when the forward and reverse reaction rates are equal. Concentrations then remain constant over time, but molecules continue reacting in both directions.

Royal Society of Chemistry guidance emphasises this distinction: equilibrium appears macroscopically static but is dynamically active at the particle level. See RSC: Understanding equilibrium—a delicate balance.

Colour is a concentration proxy

Some equilibrium systems contain species with different colours. If the relative concentration of one coloured species increases, the mixture’s observed colour shifts toward that species’ appearance.

The RSC cobalt-equilibrium demonstration uses two differently coloured cobalt species and shows that changing chloride concentration or temperature changes the observed colour. See the RSC cobalt equilibrium demonstration.

A visible shift is not the same as total conversion

Suppose an equilibrium mixture appears purple because pink and blue species coexist. Adding chloride causes the mixture to look more blue. The correct conclusion is that the equilibrium composition has shifted toward the blue-associated species.

It is usually wrong to say “all the pink species became blue.” Unless the equilibrium constant and concentrations justify near-complete conversion, both sides generally remain present.

Le Chatelier predicts direction, not mechanism

Le Chatelier’s principle is useful shorthand: when an equilibrium system is disturbed, the composition changes in a direction that partially opposes the imposed change.

But RSC guidance explicitly warns that the principle is a summary, not a full explanation. A deeper account considers how the disturbance changes forward and reverse rates or, at JC level, how the reaction quotient Q compares with the equilibrium constant K.

Concentration change: what happens immediately?

If a reactant concentration is suddenly increased, the reaction quotient changes immediately while the equilibrium constant remains fixed at constant temperature. The forward and reverse rates are no longer equal, so the system evolves until Q returns to K.

This is stronger reasoning than “the system wants to remove the added reactant.” Molecules do not possess intentions. The shift emerges from kinetics and thermodynamics.

Temperature is different because K changes

Changing concentration at constant temperature changes Q but not K. Changing temperature can change the equilibrium constant itself.

For an endothermic forward process, raising temperature typically increases the equilibrium constant for that direction; for an exothermic forward process, raising temperature tends to favour the reverse composition. The colour change therefore contains thermodynamic information, but only if the reaction enthalpy and species identities are known.

Why a control tube matters

Human colour memory is poor. A reference tube kept under the original condition gives a simultaneous visual comparison. RSC practical guidance for equilibrium demonstrations commonly retains a control sample for exactly this reason.

The control does not create the equilibrium explanation. It improves the reliability of the macroscopic observation.

Use safe, validated school procedures

Some classic equilibrium demonstrations use hazardous reagents or concentrated acids and are intended for teacher-led or microscale use under institutional laboratory procedures. Follow approved school risk assessments and authoritative practical instructions rather than improvising concentrations or chemical combinations.

The learning objective is equilibrium reasoning, not chemical bravado.

Quantitative window: reaction quotient

For a simplified equilibrium:

A(aq) + B(aq) ⇌ C(aq)

the concentration-form equilibrium expression might be:

K_c = [C]/([A][B])

If equilibrium concentrations initially give K_c = 4.0, then [A] is suddenly doubled while everything else is momentarily unchanged, the immediate reaction quotient becomes half its previous value:

Q < K

The net reaction proceeds forward until the concentration ratio again satisfies K at that temperature.

Colour intensity can become quantitative—with calibration

If one species absorbs strongly at a chosen wavelength, a colorimeter or spectrophotometer can measure absorbance. With an appropriate calibration and Beer–Lambert conditions, this can estimate concentration rather than relying on eye judgement.

But the measurement then inherits calibration, overlapping spectra, cuvette path length, wavelength selection and linear-range limits. Instrumental colour is stronger evidence only when the optical model is controlled.

Dilution can alter more than one concentration at once

Adding water does not simply “remove reactant.” It changes the concentrations of all dissolved species and can change solvent participation or complex-ion equilibria. Therefore equilibrium predictions from dilution should come from the equilibrium expression and stoichiometry, not a memorised rule that dilution always shifts toward “more particles.”

Catalysts do not shift the equilibrium position

A catalyst lowers kinetic barriers for both forward and reverse processes. It can help the system reach equilibrium faster, but it does not change K or the final equilibrium composition at a fixed temperature.

This is a valuable practical distinction: faster colour stabilisation does not necessarily mean a different equilibrium colour.

Observation versus inference

Observation: “After the concentration disturbance, the mixture changed from purple toward blue and then remained approximately constant in colour.”

Inference: “The relative concentration of the blue-associated equilibrium species increased and a new macroscopic steady composition was reached.”

Stronger inference: “At the stated temperature, the disturbance changed Q so net reaction proceeded until Q again matched K.”

Overclaim: “The reaction moved completely to the product side.” A colour shift alone does not establish complete conversion.

Failure modes that cap standards

• Colour memory used without a control: subtle shifts become subjective.
• Temperature changes unintentionally during concentration treatment: K may change as well as Q.
• Different sample volumes: colour intensity and thermal response become harder to compare.
• Le Chatelier treated as mechanism: prediction replaces explanation.
• Shift interpreted as completion: equilibrium composition is overclaimed.
• Catalyst said to increase yield at equilibrium: kinetics and equilibrium are confused.
• Hazardous demonstration improvised: method integrity and safety both fail.

Unfamiliar transfer: gas-phase equilibrium

For a gas equilibrium such as N₂O₄(g) ⇌ 2NO₂(g), pressure changes can alter equilibrium composition because the stoichiometric numbers of gas particles differ. The visible colour of NO₂ can make the shift observable. The same reasoning transfers, but pressure now changes gas concentrations through volume rather than adding a dissolved reagent.

Secondary → JC → deeper Chemistry

Secondary: recognise reversible reactions, describe colour changes accurately and understand that equilibrium does not mean reactions stop.

JC: use K and Q, separate concentration changes from temperature changes, explain catalyst effects and connect colour signals to changing species concentrations.

Deeper Chemistry: equilibrium extends to activities, coupled equilibria, van ’t Hoff relationships, speciation modelling, non-ideal solutions and kinetic relaxation methods.

Checkpoint 1: the colour shift

An equilibrium mixture becomes much bluer after a disturbance. Can you conclude the pink species concentration became zero?

Checkpoint 2: the catalyst

A catalyst makes the final colour appear sooner but the final colour is unchanged. What does that tell you?

Answer key and WHY reasoning

Checkpoint 1: no. The colour shift shows a change in relative composition, not complete disappearance of one equilibrium species.

Checkpoint 2: the catalyst changed the rate at which equilibrium was reached but not the equilibrium composition at the same temperature.

How to study this practical

Always write four lines: what was changed; what was observed; what happened to Q or reaction rates immediately; what new equilibrium composition is supported. This keeps macroscopic colour, particle-level dynamics and symbolic equilibrium reasoning connected.

Evidence boundaries

A colour-shift experiment demonstrates changes in equilibrium composition for a particular chemical system. Colour alone does not yield exact concentrations unless calibrated, and Le Chatelier’s principle predicts direction without replacing the thermodynamic explanation of why K has the value it does.

Authoritative next steps

• Royal Society of Chemistry: The equilibrium between two coloured cobalt species
• RSC: Teaching reversible reactions and equilibrium
• RSC: Understanding equilibrium—a delicate balance
• SEAB A-Level syllabus directory

Teaching Guide

Ask students to ban the phrase “the equilibrium shifts to oppose the change” for five minutes. They must instead explain what happened immediately to concentrations, forward/reverse rates or Q, and why the colour changed. Only then allow Le Chatelier back as a summary. This prevents the slogan from replacing the science.