What Are Emission Spectra?
When atoms absorb energy , from heat, electricity, or light , their electrons are promoted to higher energy levels (excited states). These excited states are unstable, so electrons quickly fall back to lower energy levels, releasing the absorbed energy as photons of light.
When this emitted light is passed through a prism or diffraction grating, it produces an emission spectrum: a series of bright, coloured lines against a dark background. Each line corresponds to a specific wavelength of light, which represents a precise energy transition inside the atom.
Emission spectrum: An emission spectrum is the pattern of discrete bright lines produced when excited atoms release energy as photons of specific wavelengths.
Ground state: The ground state is the lowest energy arrangement of electrons in an atom , the configuration an atom adopts under normal conditions. An excited state occurs when one or more electrons have been promoted to higher energy levels by absorbing energy.
The key equation linking a photon's energy to its frequency is:
where:
- = energy of the photon (J)
- = Planck's constant ( J s)
- = frequency of the radiation (Hz, or s)
Since the speed of light , this can also be written as:
where is the wavelength in metres and m s.
Energy is directly proportional to frequency and inversely proportional to wavelength. Shorter wavelengths (e.g. ultraviolet) carry more energy than longer wavelengths (e.g. infrared).
Continuous vs. Line Spectra
Not all light sources produce the same kind of spectrum. There are two fundamentally different types:
Continuous spectrum: A continuous spectrum contains all wavelengths of visible light, blending seamlessly from red to violet , like a rainbow. It is produced by incandescent objects such as the Sun or a filament bulb.
Line spectrum: A line spectrum consists of discrete lines at specific wavelengths against a dark background. It is produced by excited gaseous atoms of a particular element.
The existence of a line spectrum , rather than a continuous one , is direct evidence that energy levels in atoms are quantized. Electrons can only occupy certain fixed energy levels, so only specific energy differences (and therefore specific wavelengths) are possible.
Sodium street lamps glow with a characteristic yellow-orange colour. When viewed through a spectroscope, this resolves into two distinct lines at 589.0 nm and 589.6 nm , a unique spectral "fingerprint" for sodium. This is also why sodium is used in flame tests, producing a persistent yellow flame.
Don't confuse emission and absorption spectra. An emission spectrum shows bright lines on a dark background , produced when excited electrons fall to lower energy levels and release photons of specific wavelengths. An absorption spectrum shows dark lines on a continuous background , produced when electrons absorb photons of specific wavelengths and are promoted to higher energy levels. The lines appear at the same positions in both spectra for a given element, because the same energy differences are involved.
Note: absorption spectra are not a primary assessment focus at SL , understanding emission spectra is the core requirement.