Why Does Salt Turn a Flame Orange? Electrons Falling Back

Heat in a flame supplies enough energy to knock electrons in the metal atoms into higher energy arrangements than they normally occupy. Those arrangements are unstable, and the electron falls back within a very short time, releasing the difference in energy as a particle of light. The energy of that light determines its colour, and because the available arrangements in any given element are fixed and specific to it, the energies released are fixed and specific too. Sodium produces an intense yellow orange, potassium a pale lilac, copper a blue green, lithium a crimson, barium a pale green and calcium a brick red. The colours are the direct visible consequence of the internal structure of each kind of atom.

The procedure is simple and the details matter:

• •A wire loop of platinum or nichrome is cleaned by dipping in acid and heating
• •Cleaning is repeated until the wire no longer colours the flame at all
• •The loop is moistened and dipped in the sample to pick up a small quantity
• •It is held in the hot edge of a non-luminous flame rather than in the yellow part
• •The colour is observed immediately, since some fade within a second or two
• •Blue glass or a prism separates the contribution of sodium, which swamps everything

Contamination by sodium is the practical difficulty that dominates the technique. Sodium compounds are everywhere, in glass, in dust, on skin and in tap water, and the emission from sodium is unusually intense, so a trace produces a strong yellow that masks the weaker colours of everything else. A sample containing potassium alongside any sodium at all appears yellow rather than lilac. The standard remedy is to view the flame through cobalt blue glass, which absorbs the yellow strongly while transmitting the violet and red, revealing the potassium behind it. Scrupulous cleaning of the wire between samples is the other requirement, and a wire that still colours the flame after cleaning must be discarded rather than used.

The colours appear in several familiar places once they are recognised. Fireworks are the obvious case, with the colours produced by exactly these metal compounds, strontium for red, barium for green, copper for blue and sodium for yellow, and blue remains the hardest because the copper compounds that produce it break down at the temperatures involved. Street lamps of the older orange type contained sodium vapour and emitted essentially the same light as a sodium sample in a flame, which is why they rendered everything in shades of grey and orange. Salt spilled onto a gas hob yellows the flame instantly. And the light from stars carries the same signatures, which is how their composition is known and how the technique escaped the laboratory entirely.

The schoolroom demonstration is the visible end of a technique that transformed several sciences. Robert Bunsen and Gustav Kirchhoff in 1860 combined a clean hot flame with a prism, separating the emitted light into precise lines rather than an impression of colour, which made identification unambiguous and sensitive to tiny quantities. They used it to discover two new elements within two years, each named for the colour of its characteristic lines. Kirchhoff then established that the dark lines in sunlight correspond to the bright lines of known elements, which meant the composition of the sun could be determined from Earth and which founded astrophysics as a subject. Modern instruments applying the same principle measure metal concentrations routinely in water, soil, blood and industrial samples.

Heat pushes electrons into higher arrangements, and they fall back releasing light whose energy is fixed by the structure of that element, which is why each metal gives its own colour. A clean wire and a non-luminous flame are essential, and sodium contamination swamps everything unless viewed through cobalt blue glass. Adding a prism to the same idea in 1860 produced two new elements and the ability to analyse the sun.