Emission spectra are produced when photons are emitted from atoms as excited electrons return to a lower energy level. Whenever energized/excited, an electron moves to a higher energy level, in this case by jumping orbitals/shells. At the point in time when the electron eventually falls back to its regular 'ground level' (its original position) the light is emitted. There are numerous conceivable electron moves for every molecule, and every move has a particular energy level distinction. This gathering of various moves are what prompts distinctive emanated wavelengths that make up an emission spectrum. All the elements have a emission spectrum unique to them. If one wanted to, they could deduce the element by it's emission spectrum.
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For example, above you can see the absorption and emission spectrum of Cobalt (Co). The absorption spectrum is drawn showing all the visible light spectrum with black gaps/lines in the places that the color gets absorbed. To display an emission spectrum, the background is black and now where those gaps of color were on the visible light spectrum, there is the bands of color it previously absorbed.
The principle energy level or shell is given a whole number, n, and can hold, at most, number of electrons: 2n2. A more definite model of the molecule portrays the division of the fundamental energy levels into s, p, d and f sub-levels in the order of progressively higher energies. The makeup of sub-levels is a settled number of orbitals, areas where there is a high likelihood of finding an electron. Each orbital has a characterized energy state for a given electronic setup and substance environment and can hold two electrons of clockwise and anti-clockwise (respectively) direction.
Limit of convergence at high frequencies in an emission spectrum correspond to the first ionization energy
Refer to this website: www.chemguide.co.uk/atoms/properties/hspectrum.html