Fireworks and Electron Transitions: The Physics Behind the Colours

A firework shell bursts high in the sky and scatters stars of brilliant red, green and blue. Every one of those colours is chemistry and atomic physics working together: a fast, hot reaction supplies energy, and particular atoms and molecules turn that energy into light of very specific colours. This article explains where firework colours come from, which chemicals produce them, and why some colours are much harder to make than others.

Two ways to make light

Incandescence: glowing because it’s hot

Any object glows when it’s hot enough: first dull red, then orange, yellow and, at very high temperatures, white. This incandescence comes from the random thermal motion of hot particles and gives a continuous spread of colours. In fireworks, burning metal powders such as aluminium, magnesium and titanium produce brilliant white sparks and flashes by incandescence, while iron and charcoal give golden, orange sparks.

Luminescence: glowing at specific colours

Coloured stars rely mostly on luminescence: light emitted when excited electrons in atoms or molecules drop back to lower energy levels, releasing photons of particular energies and therefore particular colours. This is the same process that produces flame test colours and neon signs. See photons and energy levels.

For pure, vivid colours, firework makers want as much luminescence and as little white incandescence as possible, which means controlling the flame temperature carefully.

The chemistry inside a firework star

A coloured “star” (the small pellet that produces each point of light) typically contains:

1. An oxidiser, which supplies oxygen: often potassium perchlorate (KClO₄) or potassium nitrate (KNO₃). Historically, potassium chlorate was used, but it’s sensitive and dangerous.
2. A fuel, which burns rapidly: charcoal, sulfur, metal powders or organic resins.
3. A colourant: a compound of the metal that produces the colour.
4. A chlorine donor: often a chlorinated plastic such as PVC, or the perchlorate itself, which helps form coloured chloride species.
5. A binder to hold the pellet together.

When the star ignites, the fuel and oxidiser react very exothermically, producing temperatures of roughly 1,500–3,000 °C. That heat vaporises the colourant and excites its electrons. See oxidation and reduction and exothermic vs endothermic.

The colours and their sources

Colour	Typical colourant	Main light-emitting species
Red	strontium carbonate or nitrate	SrCl and SrOH molecules
Orange	calcium salts	CaCl and CaOH molecules
Yellow	sodium salts (e.g. sodium nitrate, cryolite)	sodium atoms (589 nm)
Green	barium nitrate or chlorate	BaCl molecules
Blue	copper compounds (e.g. copper carbonate, copper oxychloride)	CuCl molecules
Purple	strontium (red) + copper (blue) together	mixture
White / silver	aluminium, magnesium, titanium	incandescence, plus metal oxide emission
Gold	iron filings, charcoal	incandescent sparks

Atoms or molecules?

Here’s a detail textbooks often skip: many firework colours don’t come from free metal atoms at all, but from small, short-lived molecules formed in the flame:

• The red of strontium comes mainly from SrCl and SrOH.
• The green of barium comes mainly from BaCl.
• The blue of copper comes mainly from CuCl.

That’s why a chlorine donor is so important: it provides chlorine atoms to form these molecules. Without enough chlorine, barium gives a paler, whiter light. Molecules produce bands of closely spaced lines rather than single sharp lines, but they’re concentrated in the right colour region.

Sodium is the exception: its intense yellow comes from sodium atoms, emitting at 589.0 and 589.6 nm. It’s so strong that even a trace of sodium contamination can wash out other colours, so fireworks for pure colours must avoid sodium impurities.

Why blue is the hardest colour

Firework makers have long considered a deep, bright blue the most difficult colour:

• The blue-emitting molecule CuCl breaks down at temperatures above about 1,200 °C, so the flame must be relatively cool.
• But a cooler flame gives less light overall.
• Any white incandescence from hot particles easily washes out the blue.
• Chlorine donors and fuels must be carefully balanced.

Getting all this right in a composition that is also safe, stable and reliable is a real challenge, and blue stars are often dimmer than reds and greens. Chemists continue to research better blue formulations.

Energy and colour: a quick calculation

Different colours correspond to different photon energies. Using E = hc/λ:

• Strontium’s red emission, around 640 nm, carries about 1.9 eV per photon (about 187 kJ/mol).
• Barium’s green emission, around 520 nm, carries about 2.4 eV (about 230 kJ/mol).
• Copper’s blue emission, around 430–450 nm, carries about 2.8 eV (about 270 kJ/mol).

Blue photons need the most energy per photon, part of the reason blue is harder to produce efficiently. See wavelength, frequency and photon energy calculations.

Effects beyond colour

• Bangs: “flash powder” (a fine metal fuel with a strong oxidiser) burns extremely fast in a confined casing, producing a sharp report.
• Whistles: certain organic salts burning in a tube produce oscillating combustion that makes a whistling sound.
• Crackles: small granules that burn and then pop in rapid succession.
• Glitter and strobes: compositions that burn in pulses, flashing on and off.

Safety and the environment

Fireworks are explosives and must only be handled by trained professionals, following local laws. Beyond immediate safety, chemists are working to reduce their environmental impact:

• Perchlorates, common oxidisers, can contaminate water and affect the thyroid gland. Alternatives are being developed.
• Barium compounds are toxic; researchers have developed barium-free greens, for example using boron compounds.
• Chlorinated organics can produce harmful by-products when burned; chlorine-free formulations are an active research area.
• Smoke and particulates affect air quality after large displays.

See heavy metals and risk assessment in chemistry.

A short history

Fireworks began in China more than a thousand years ago, after the invention of black powder (gunpowder), a mixture of potassium nitrate, charcoal and sulfur. For centuries, fireworks were limited to the gold and white sparks of burning charcoal and iron. Bright reds, greens and blues only became possible in the nineteenth century, when chemists isolated strontium and barium compounds and introduced potassium chlorate as an oxidiser that burned hot enough to excite their colours.

A classroom connection

The flame test in a school lab is essentially a miniature, safe version of a firework star: heat a metal salt in a flame, excite its electrons, and watch the colour. Looking at the flame through a simple spectroscope shows that each colour is made up of specific lines or bands, the fingerprint of the element. See flame test colours.

Key takeaways

• Firework colours come mainly from luminescence: heat excites electrons, which emit light of specific colours as they drop back.
• A firework star combines an oxidiser, a fuel, a colourant, a chlorine donor and a binder.
• Strontium gives red, calcium orange, sodium yellow, barium green and copper blue; many colours come from molecules such as SrCl, BaCl and CuCl.
• Blue is hardest because CuCl breaks down at high temperatures and is easily washed out by white light.
• Modern research aims to make fireworks safer and less polluting, replacing perchlorates, barium and chlorinated compounds.