The Chemical Reaction That Cannot Make Up Its Mind — The Briggs–Rauscher Oscillating Reaction

Most chemical reactions seem reassuringly predictable.

Mix an acid with an alkali and the pH moves towards neutral. Burn a fuel and the fuel is gradually consumed. Add magnesium to hydrochloric acid and hydrogen is produced until one of the reactants runs out.

There is usually a clear sense of before, during and after.

Then there is the Briggs–Rauscher reaction.

At first, the mixture may appear almost colourless. Then it develops a yellow or amber colour. Suddenly it turns an intense blue-black. The darkness disappears. The amber colour returns. Then blue-black again.

And it repeats.

Not because anyone is adding more chemicals. Not because someone is switching the reaction on and off.

The chemistry itself is oscillating.

The Briggs–Rauscher reaction, first reported by Thomas Briggs and Warren Rauscher in 1973, became famous because these repeated colour changes make an extraordinarily complicated piece of chemical kinetics visible to the naked eye.

For students accustomed to thinking that reactions simply proceed from reactants to products, it raises a wonderful question:

How can a chemical reaction apparently change direction again and again?

And the answer takes us beyond much of ordinary GCSE and A-level Chemistry into feedback, autocatalysis, nonlinear systems and the chemistry of substances that are far from thermodynamic equilibrium.

A Reaction That Looks Almost Alive

There are some demonstrations that students understand intellectually but do not necessarily remember years later.

Then there are demonstrations where people simply stop talking and watch.

I think the Briggs–Rauscher reaction belongs firmly in the second category.

You can explain beforehand that the colours will oscillate.

You can tell someone:

clear → amber → blue-black → clear → amber → blue-black

But knowing what is going to happen does not really prepare you for seeing it.

There is something deeply counter-intuitive about watching a chemical mixture change colour, apparently finish changing, and then apparently decide to do it all over again.

A typical demonstration oscillates several times before the behaviour gradually dies away as the available reactants are consumed. The precise number, timing and appearance of the oscillations depend on the formulation, temperature and experimental conditions.

It immediately looks less like the simple chemistry found in introductory textbooks and more like a living system.

That comparison is worth exploring.

So What Is an Oscillating Chemical Reaction?

An oscillating reaction is one in which the concentrations of important intermediate chemical species repeatedly rise and fall.

The important word here is intermediate.

It is not normally the case that the whole reaction simply proceeds forwards, reverses completely, proceeds forwards again and then reverses again.

Instead, there is a network of competing reactions.

One pathway produces a substance.

Another pathway consumes it.

Those pathways influence one another.

Eventually one dominates.

Then conditions change sufficiently for the other pathway to dominate.

Then the first becomes favourable again.

The result is a chemical feedback loop.

In the Briggs–Rauscher reaction, changes involving iodine, iodide and several reactive intermediates produce the spectacular visible oscillations. The complete mechanism is complex enough that the system has been the subject of substantial kinetic research rather than being reducible to one simple school-level equation.

What Do We Actually Put Into the Reaction?

There are several formulations of the Briggs–Rauscher reaction. A Royal Society of Chemistry demonstration uses separately prepared solutions containing the principal components before they are mixed for the demonstration.

The chemistry involves:

hydrogen peroxide, H2O2;
potassium iodate, KIO3;
sulfuric acid, H2SO4;
malonic acid, CH2(COOH)2;
manganese(II) ions, usually supplied using manganese sulfate;
starch;
deionised water.

These ingredients each have very different jobs.

Hydrogen peroxide and iodate provide strongly oxidising chemistry.

Malonic acid participates in the competing iodine chemistry.

Manganese(II) acts as a catalyst within the reaction network.

Starch gives us the extraordinarily visible dark colour associated with iodine/iodide species bound within the starch structure.

And the acidic conditions are important to the overall kinetics.

It is therefore much more than an ordinary colour indicator experiment.

The starch allows us to see a kinetic battle taking place inside the solution.

Why Three Colours?

The colour sequence gives us clues about what is happening chemically.

Almost colourless

During parts of the cycle there is relatively little molecular iodine available to produce the characteristic iodine colour.

The solution consequently becomes pale or almost colourless.

Amber or yellow-brown

As molecular iodine, I2, accumulates, the characteristic yellow-brown or amber colour becomes visible.

It is rather like watching iodine appear in a conventional iodine reaction — except that here it will subsequently disappear again.

Blue-black

This is the spectacular stage.

Students are often taught simply:

iodine + starch → blue-black

That is useful at GCSE level, but the chemistry is more subtle.

The coloured starch complex involves iodine and iodide chemistry, including polyiodide species associated with the helical structure of amylose in starch. In the oscillating system, changes in iodine and iodide concentrations therefore produce the sudden dramatic darkening.

So even the familiar school starch test turns out to contain more chemistry than it first appears.

But Why Does It Keep Repeating?

This is the fascinating part.

A simplified way of imagining the reaction is to think of two competing chemical systems.

One favours the production of iodine.

Another eventually causes iodine to be consumed and iodide concentration to change.

Those changes then influence the first system.

So we can picture the cycle approximately as:

low iodide concentration

↓

a rapid reaction pathway becomes important

↓

iodine is produced

↓

iodine and related species accumulate

↓

iodide concentration changes

↓

the first pathway becomes inhibited

↓

iodine is consumed

↓

iodide eventually falls sufficiently

↓

the first pathway becomes important again

↓

the cycle repeats

This is not a complete mechanism — the real reaction involves several intermediates and radical processes — but it captures the crucial idea of feedback.

And feedback is one of the most important ideas extending beyond school chemistry.

Positive Feedback and Negative Feedback

Imagine turning on a microphone beside a loudspeaker.

A tiny sound enters the microphone.

The amplifier makes it louder.

The loudspeaker produces the louder sound.

The microphone hears that sound.

It gets amplified again.

Very quickly we hear the familiar howl of audio feedback.

That is an example of positive feedback.

A change encourages more of the same change.

Chemical reactions can behave similarly through autocatalysis, where the progress of a reaction creates something that helps the reaction proceed still faster.

But an oscillator also requires something that eventually restrains that process.

That is negative feedback.

One process accelerates.

Another process eventually catches up.

The first process is suppressed.

Conditions then gradually reset.

And the cycle starts again.

That combination of:

amplification + inhibition + delay

is a recipe for oscillation in many kinds of system.

The Briggs–Rauscher reaction gives us the unusual opportunity to watch that idea happening chemically.

Isn't This Breaking the Second Law of Thermodynamics?

This is perhaps the most interesting misconception the reaction can provoke.

If reactions naturally move towards equilibrium, shouldn't concentrations simply settle down smoothly?

How can something repeatedly move away from its previous state?

The key is that the oscillating mixture is not at thermodynamic equilibrium.

It begins containing chemical substances capable of reacting and releasing chemical free energy.

The system is therefore temporarily maintained far from equilibrium.

While that chemical "fuel" remains available, concentrations of intermediate species can fluctuate dramatically.

But the oscillations do not continue forever.

Eventually the reactants that sustain them become depleted.

The amplitude changes.

The oscillations cease.

The mixture settles towards its final state.

So thermodynamics has not been defeated at all.

The reaction simply demonstrates something important:

Thermodynamics tells us where a system can ultimately go. Kinetics tells us the route it may take getting there.

And that route can be astonishingly complicated.

Equilibrium Does Not Mean Nothing Is Happening

This also provides an opportunity to revisit another common misconception.

Students sometimes imagine chemical equilibrium as the moment when a reaction "stops".

It does not.

At dynamic equilibrium, forward and reverse processes continue, but their rates are equal so macroscopic concentrations remain constant.

The Briggs–Rauscher reaction is quite different.

During its oscillating phase, the concentrations of important species are not constant.

The system is therefore not sitting at equilibrium.

Visible concentrations are changing dramatically with time.

That distinction between:

dynamic equilibrium

and

dynamic behaviour far from equilibrium

is subtle but scientifically important.

Why Does the Reaction Eventually Stop?

Because there is no chemical perpetual-motion machine hiding in the flask.

Every oscillation consumes some of the chemical resources that sustain the reaction.

Eventually the conditions required for the feedback cycle can no longer be maintained.

One or more important reactants become sufficiently depleted.

The system changes permanently.

The oscillations disappear.

This makes the reaction rather like a wind-up mechanical clock.

The hands may move periodically, but only because stored energy is available.

Eventually that energy runs out.

The periodic behaviour was real.

It simply was not free.

A Wonderful Bridge from GCSE to University Chemistry

One reason I particularly like experiments such as this is that they can be appreciated at several different levels.

At GCSE

A student might investigate:

colour changes;
oxidation and reduction;
catalysts;
reaction rates;
hydrogen peroxide;
iodine chemistry;
starch as an indicator.

That alone makes an impressive demonstration.

At A-level

We can start discussing:

reaction mechanisms;
intermediates;
catalysts;
rate equations;
activation energy;
redox chemistry;
equilibrium;
reaction pathways.

Beyond A-level

Suddenly we reach:

nonlinear kinetics;
autocatalysis;
feedback loops;
radical mechanisms;
mathematical modelling;
bifurcation;
chemical oscillators;
self-organisation;
nonequilibrium thermodynamics.

The same beaker has taken us from GCSE observations into subjects normally encountered much later.

That is exactly why I enjoy demonstrations that go beyond the syllabus.

They show students that the syllabus is not the boundary of science.

It is merely the entrance.

Turning the Demonstration Into an Investigation

Rather than simply performing the reaction and watching it, there are several interesting ways to extract quantitative science from it.

Importantly, because this demonstration involves oxidising chemicals, acid and iodine-containing mixtures, modifications should be performed only in a suitably equipped laboratory using an established demonstration protocol and appropriate risk assessment. The RSC publishes a technician-supported formulation suitable for educational demonstration.

1. Measure the oscillation period

Record the experiment on video.

Measure the time between successive blue-black transitions.

Does the period remain constant?

Often it does not.

That immediately tells us that the underlying chemical conditions are evolving even when the visible sequence appears repetitive.

A graph could show:

Time / s

against

Oscillation number

or perhaps:

Oscillation period / s

against

Oscillation number.

Suddenly the spectacular demonstration has become experimental data.

2. Analyse the Colour Digitally

A camera can turn the colour change into a surprisingly sophisticated experiment.

Select one region of the flask in each video frame and measure its brightness or RGB value.

Plot something such as:

Blue intensity

against

Time.

Instead of merely saying:

"it changes colour repeatedly",

we obtain a waveform.

The chemistry begins to resemble a physics experiment investigating an oscillator.

One could measure:

period;
frequency;
maximum intensity;
minimum intensity;
changing amplitude;
number of cycles before oscillation stops.

It is a lovely example of different sciences merging together.

Chemistry creates the phenomenon.

Physics gives us ways of measuring light.

Computing processes the video.

Mathematics describes the resulting pattern.

3. What Does Temperature Do?

Reaction rates generally increase with temperature, so temperature provides an especially interesting variable when studying an oscillator.

Research on the Briggs–Rauscher system shows that temperature changes the timing and behaviour of the oscillations rather than merely making an identical sequence run slightly faster.

That raises a deeper question.

If several competing reactions all respond differently to temperature, what happens to the entire feedback system?

It is no longer enough simply to say:

higher temperature = faster reaction.

We have several reactions occurring simultaneously.

Some may become more important relative to others.

That is much closer to real chemical kinetics.

4. Does Stirring Matter?

It does.

Without sufficiently uniform mixing, different parts of the solution can experience somewhat different local chemical conditions.

Instead of the whole flask changing colour simultaneously, spatial patterns may develop.

That is fascinating because it introduces another major scientific idea:

reaction-diffusion behaviour.

Now chemistry is changing not simply through time but potentially through space.

The mixture can begin displaying fronts, waves and patterns rather than behaving as one perfectly mixed system. Studies and demonstrations of the Briggs–Rauscher system have deliberately examined such spatial as well as temporal behaviour.

From a Flashing Flask to Mathematical Modelling

Imagine plotting the concentration of one important intermediate against time.

A conventional reaction might produce a smooth curve.

An oscillator produces something much more interesting:

high → low → high → low → high...

But the heights and intervals may themselves gradually change.

To model that properly we may need several simultaneous differential equations describing different reaction rates.

Change one parameter slightly and the entire behaviour may change.

That leads us naturally towards another area of science:

nonlinear systems.

Small changes do not necessarily produce proportionally small results.

And chemistry begins to overlap with some of the same mathematics used to describe ecosystems, electronics, populations and other complex systems.

Chemical Clocks Are Not Quite the Same Thing

Students may already have encountered an iodine clock reaction.

That produces a spectacular sudden colour change after a delay.

But normally it happens once.

Wait...

wait...

BLUE.

An oscillating reaction goes significantly further.

It effectively contains repeated switching behaviour:

wait...

BLUE...

fade...

wait...

BLUE...

fade...

again and again.

So a useful progression might be:

Ordinary reaction

continuous change.

Clock reaction

delayed sudden change.

Oscillating reaction

repeated changes.

That progression makes the Briggs–Rauscher reaction feel much less mysterious.

Why Deionised Water Matters

One apparently mundane part of the demonstration is actually another chemistry lesson.

The quality of the water matters.

Chloride ions can interfere with the Briggs–Rauscher oscillator, and published work has specifically investigated chloride inhibition of the reaction. Educational protocols therefore use deionised water rather than relying on ordinary tap water.

That makes another valuable point for students:

purity sometimes matters enormously.

Tap water may look completely colourless.

Yet dissolved ions present at relatively low concentrations can alter sensitive chemical systems.

It is a reminder that "water" in chemistry rarely means simply H2O.

Safety — This Is a Demonstration, Not Kitchen Chemistry

The spectacular appearance can make the Briggs–Rauscher reaction look deceptively like one of the many safe colour-change experiments that can be tried using household materials.

It is not.

Established formulations involve hydrogen peroxide, acidic iodate solutions and iodine-producing chemistry. Appropriate eye protection, laboratory clothing, gloves where indicated by the risk assessment, suitable ventilation and correct handling and disposal procedures are required.

I would therefore treat this as a teacher or experienced demonstrator experiment performed in a properly equipped laboratory, rather than provide pupils with reagents and ask them simply to experiment.

For preparation, concentrations and disposal, I would use a recognised published protocol such as the Royal Society of Chemistry demonstration rather than an improvised internet recipe.

That does not reduce its educational value.

Quite the opposite.

Part of becoming a scientist is understanding that good practical chemistry means knowing not only what will react, but also how the reaction can be investigated responsibly.

What Should Students Predict Before Seeing It?

Before starting the demonstration, I would resist telling students exactly what happens.

I might show them the separate solutions and ask:

"Once these are mixed, what do you expect the colour to do?"

Most will probably predict something like:

colourless → brown

or:

colourless → blue-black.

Very few would predict:

colourless → amber → blue-black → colourless → amber → blue-black...

Then perform the reaction.

That moment when their prediction fails is valuable.

Science advances because observations sometimes refuse to behave according to our simple expectations.

The correct response is not:

"That shouldn't happen."

It is:

"Why did that happen?"

The Bigger Scientific Lesson: Feedback Is Everywhere

The Briggs–Rauscher reaction matters because feedback systems are not peculiarities confined to one flask.

Feedback is one of the organising ideas of modern science.

We encounter it in:

biochemical networks;
regulation inside cells;
ecological populations;
hormone systems;
electronic circuits;
climate processes;
engineering control systems;
neural systems.

The details are completely different, but the underlying question is often similar:

How does one changing quantity affect another quantity which then feeds back and alters the first?

Once students begin thinking in those terms, they start moving away from simple chains of cause and effect.

A causes B.

B causes C.

Towards interconnected systems:

A changes B.

B changes C.

C changes A.

Now genuinely complicated behaviour becomes possible.

Chemistry Is Still Studying It

Perhaps surprisingly for a reaction first reported more than half a century ago, the Briggs–Rauscher system has not become merely an old classroom curiosity.

Researchers continue to use modern analytical techniques to investigate its intermediates and changing chemical behaviour. A 2026 study, for example, reported real-time resonance Raman spectroscopic monitoring of the reaction.

I find that particularly appealing.

The same flask can make a school student say:

"Wow — why did it change colour again?"

while researchers can ask highly sophisticated questions about transient species, kinetics and mechanism.

That is a wonderful illustration of what science education should do.

A demonstration does not have to simplify science until all the interesting parts have disappeared.

Sometimes we should show students something difficult precisely because it gives them something to wonder about.

From Watching Chemistry to Thinking Like a Scientist

If I were using the Briggs–Rauscher reaction with students, the most important part would probably begin after the colours had stopped.

I would ask:

Why did it oscillate?

Why did it eventually stop?

Where did the energy come from?

Was the system at equilibrium while it was oscillating?

What caused the dark colour?

Why did stirring matter?

Would temperature alter the frequency?

Could we measure the oscillations rather than merely watch them?

Could we build a mathematical model?

And perhaps most importantly:

What other systems behave like this?

At that point the experiment has stopped being merely a spectacular chemistry demonstration.

It has become a doorway.

Conclusion — A Flask Full of Complexity

Most introductory chemistry necessarily teaches us to simplify.

One reactant reacts with another.

Products form.

Concentrations decrease.

Equilibrium is reached.

And those models are immensely useful.

But nature is under no obligation to remain simple.

The Briggs–Rauscher reaction shows what can happen when several reactions interact through positive feedback, inhibition, catalysis and changing intermediate concentrations.

The result is almost hypnotic:

clear... amber... blue-black... clear... amber... blue-black...

For a few minutes, the flask appears to have acquired a rhythm of its own.

Of course, it has not become alive.

It is following chemical laws throughout.

And perhaps that is the most fascinating thing about it.

The extraordinary behaviour is not occurring because chemistry has stopped obeying the rules.

It is occurring because the chemistry is obeying those rules in a system complicated enough for something unexpected to emerge.

That, for me, is one of the great attractions of science beyond the syllabus.

Sometimes the next level of understanding begins when an experiment does something that looks impossible.

21 August 2026