What Causes Emission Spectra, and Why Do Elements Produce Different-Colored Light?

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Here, $\Delta E$ is the energy difference between the levels, $h$ is Planck's constant, $f$ is frequency, $c$ is the speed of light, and $\lambda$ is wavelength. A larger energy difference produces a higher-frequency, shorter-wavelength photon.

Because electrons can occupy only discrete energy levels, only particular energy differences are possible. Therefore, an element produces separate bright lines rather than a continuous range of wavelengths. This is called a line emission spectrum.

Each element has a different nuclear charge and electron configuration, giving it a unique arrangement of energy levels. Its permitted electron transitions therefore release a characteristic set of photon energies and colors. This spectral pattern acts like an atomic fingerprint.

A common misconception is that electrons themselves have colors. They do not: the observed color comes from the wavelength of the emitted photons. Some emitted wavelengths may also lie outside the visible region.

Exam Technique

For an explain question, connect all three ideas: electrons occupy discrete levels, downward transitions release photons, and $\Delta E$ determines wavelength. Do not merely state that different elements have different colors; explain that their energy-level spacings are different.