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Energized electrons must shed excess energy to settle into a lower, stable state. That release manifests as photons—light. Atomic emission spectra are therefore a map of electrons slipping back into lower energy levels. Quantum mechanics dictates that electrons can only absorb or emit specific, discrete energy quanta. Each element’s unique orbital configuration governs the wavelengths, and thus colors, of its emission lines.
While the macroscopic world follows continuous, deterministic laws, the microscopic realm is ruled by discrete states and probability. Electrons occupy distinct energy levels with no intermediate states. When excited, an electron jumps instantaneously to a higher level; when it relaxes, it emits a photon whose energy equals the gap between the two levels. Unlike a steadily burning fire, which emits energy gradually, an electron releases its energy all at once.
Energy from light exists in packets called photons. Photons have different energies that correspond to different wavelengths. Therefore, the color of emission lines reflects the amount of energy released by an electron. This energy changes depending on the orbital structure of the atom and the energy levels of its electrons. Higher energies correspond to wavelengths toward the shorter, blue end of the visible light spectrum.
When light passes through atoms, those atoms can absorb some of the light’s energy. An absorption spectrum shows us which wavelengths of light were absorbed by a particular gas. An absorption spectrum looks like a continuous spectrum, or rainbow, with some black lines. These black lines represent photon energies absorbed by electrons in the gas. When we view the emission spectrum for the corresponding gas, it will display the inverse; the emission spectrum will be black everywhere except for the photon energies that it previously absorbed.
Emission spectra can have a large number of lines. The number of lines does not equal the number of electrons in an atom. For example, hydrogen has one electron, but its emission spectrum shows many lines. Instead, each emission line represents a different jump in energy that an electron of an atom could make. When we expose a gas to photons of all wavelengths, each electron in the gas may absorb a photon with exactly the right energy to excite it into the next possible energy level. Hence, the photons of an emission spectrum represent a variety of possible energy levels.
