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Why is a sodium street lamp orange? Why does copper turn a flame green, and strontium turn it red? Why do neon signs glow red, while “neon” signs of other colours use different gases? All of these come down to one idea: electrons in atoms can only have certain energies, and when they move between those energies, they absorb or emit light of very specific colours. This article connects the physics of photons with the colours you see every day.
Light comes in packets: photons
Light behaves both as a wave and as a stream of particles called photons. Each photon carries a fixed amount of energy that depends on its frequency (or wavelength):
E = hf = hc ÷ λ
- h = Planck’s constant, 6.626 × 10⁻³⁴ J s
- f = frequency (Hz)
- c = speed of light, 3.00 × 10⁸ m/s
- λ = wavelength (m)
Higher frequency and shorter wavelength mean more energy per photon. Across the visible spectrum:
| Colour | Approximate wavelength (nm) | Energy per photon (eV) |
|---|---|---|
| red | 620–750 | about 1.7–2.0 |
| orange | 590–620 | about 2.0–2.1 |
| yellow | 570–590 | about 2.1–2.2 |
| green | 495–570 | about 2.2–2.5 |
| blue | 450–495 | about 2.5–2.8 |
| violet | 380–450 | about 2.8–3.3 |
Ultraviolet photons carry even more energy (enough to damage skin and DNA); infrared photons carry less. For calculations, see wavelength, frequency and photon energy calculations.
The idea that light comes in packets was proposed by Max Planck (1900) and developed by Albert Einstein (1905) to explain the photoelectric effect.
Electrons have fixed energy levels
Electrons in atoms can’t have just any energy. They occupy quantised energy levels, like the rungs of a ladder: an electron can stand on a rung but never between rungs.
- The lowest-energy arrangement is the ground state.
- Arrangements with electrons in higher levels are excited states.
The exact energies depend on the element, because each element has a different nuclear charge and number of electrons. See the Bohr model and electron shells explained.
Absorbing and emitting photons
- Absorption: an electron can jump to a higher level by absorbing a photon whose energy exactly matches the gap.
- Emission: an excited electron falls to a lower level, releasing a photon whose energy exactly equals the gap.
Energy of photon = E(upper level) − E(lower level)
A photon with slightly the wrong energy simply isn’t absorbed. It’s all or nothing, which is why atomic spectra consist of sharp lines. See emission vs absorption spectra.
Why each element has its own colours
Because each element has a unique set of energy levels, it has a unique set of energy gaps, and therefore a unique set of emitted wavelengths. If some of those fall in the visible region, the element gives a characteristic colour.
| Source | Element | Main visible colour | Why |
|---|---|---|---|
| Sodium street lamp | sodium | orange-yellow (589 nm) | strong 3p → 3s transition |
| Red neon sign | neon | red-orange | several strong red lines |
| Firework red | strontium (and lithium) | red | red emission from strontium species |
| Firework green | barium (and copper) | green | green emission lines and bands |
| Firework blue | copper compounds | blue | copper(I) chloride emission |
| Flame test lilac | potassium | lilac | violet and red lines combined |
See flame test colours.
Where do firework colours come from?
When a firework explodes, the heat excites metal atoms and ions (or small molecules containing them). As the electrons fall back, they emit light of their characteristic colours. Chemists choose compounds to produce specific colours: strontium salts for red, sodium for yellow-orange, barium for green, copper for blue. Blue is famously the hardest colour to get bright and pure.
Energy per mole: connecting to chemistry
Chemists often convert photon energies to energy per mole, to compare with bond energies and ionisation energies:
- A mole of red photons (700 nm) carries about 171 kJ.
- A mole of violet photons (400 nm) carries about 299 kJ.
- A mole of UV photons at 250 nm carries about 479 kJ, more than enough to break many covalent bonds (C–C is about 350 kJ/mol).
This is why UV light can trigger chemical reactions (such as breaking chlorine molecules into atoms, or damaging DNA), while visible red light usually can’t.
Fluorescence and phosphorescence
Some substances absorb higher-energy light (often UV) and re-emit lower-energy visible light:
- Fluorescence: emission happens almost immediately (nanoseconds). Highlighter pens, tonic water under UV light (quinine glows blue) and “optical brighteners” in laundry detergents fluoresce. The emitted light has a longer wavelength than the absorbed light because some energy is lost as heat first.
- Phosphorescence: emission is delayed, continuing for seconds to hours. Glow-in-the-dark stars and some watch dials phosphoresce, because electrons are trapped temporarily in states from which returning is slow.
Lasers
A laser produces intense light of a single wavelength by stimulated emission: when a photon of exactly the right energy passes an excited atom, it can trigger the atom to emit an identical photon in the same direction. By keeping many atoms excited (a population inversion) and bouncing light between mirrors, a laser builds up a powerful beam of identical photons. The colour depends on the energy levels of the material: red helium–neon lasers, green lasers based on neodymium crystals, and many others.
Molecules and solids: broader colours
Atoms in a gas give sharp lines, but:
- Molecules have extra vibrational and rotational energy levels, so their absorption and emission form broad bands.
- Solids have huge numbers of interacting atoms, producing bands of energy levels. A hot solid (such as a filament or molten iron) glows with a continuous spectrum, whose colour depends on temperature: red at about 700 °C, orange-yellow hotter, and white at very high temperatures.
- Coloured compounds (such as copper sulfate or potassium permanganate) absorb some visible wavelengths through electron transitions and show the complementary colour. See transition metals.
LEDs
In a light-emitting diode, electrons in a semiconductor drop across an energy gap called the band gap, emitting photons. The size of the gap sets the colour: gallium arsenide phosphide for red, indium gallium nitride for blue. White LEDs usually combine a blue LED with a phosphor that converts some blue light to yellow. See semiconductors explained.
Colour you can predict
Once you know the energy gaps in an atom, you can predict its colour. For example, the gap between sodium’s 3p and 3s levels is about 2.1 eV. Using E (eV) ≈ 1240 ÷ λ (nm), λ ≈ 590 nm: yellow-orange, exactly the colour of a sodium flame.
Key takeaways
- Light consists of photons with energy E = hf = hc/λ; shorter wavelengths carry more energy.
- Electrons occupy quantised energy levels; they absorb or emit photons only of energies exactly matching the gaps.
- Each element’s unique energy levels give it unique colours, used in flame tests, street lamps, neon signs and fireworks.
- Fluorescence, phosphorescence, lasers and LEDs all rely on electrons moving between energy levels.
- Molecules and solids have many closely spaced levels, giving bands and continuous spectra rather than sharp lines.
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