Emission Spectra

An Emission Spectrum is the range of colours given out by a light source. To view spectra, the light must be dispersed in some way for the range of colours to be seen. This can be seen naturally (a rainbow) or artificially using a spectroscope / spectrometer.

There are two main types of Emission Spectra:-

1. Continuous Spectra.

2. Line Spectra.

Continuous Spectra

If a beam of white light is passed through a prism or a grating a thermal continuous spectrum will be seen.

For visible light

Violet Light ~ 400nm (short Wavelength)

Red Light ~ 700nm (long Wavelength)

Line Spectra

When the source of light is from a specific element such as a Sodium vapour lamp, a continuous spectra is not observed. The observed spectra is a series of Lines, corresponding to specific frequencies of radiation being emitted.

The below diagram shows the Line Spectra for several different elements:-

Absorption Spectra

By observing first our own sun and then other stars we see not a line spectrum as expected, instead we observe a cross between an Continuous Spectrum and a Line Spectrum, which we call a Absorption Spectrum:-

The black bands in the above heliograph correspond exactly with where we would expect to find Line Emissions for individual elements:-

Quantum Theory of Spectra

In order to explain the formation of emission and absorption spectra, we must understand how the light is emitted or absorbed at an atomic level.

The below diagram shows two alternate ways to demonstrate the inner workings of an atom:-

The left hand diagram shows the traditional Bohr model of an atom, while the right hand diagram shows an Energy Level diagram of the same atom. In order to understand how emission and absorption spectra are formed, we must use the Energy Level diagram.

Note - At Higher Level, we will only discuss the energy levels of Hydrogen (1 proton / 1 electron).

Energy Levels and Emission Spectra

Each energy level is capable of containing an Electron, but for a stable atom, the electron will be found in the lowest energy level - the Ground State.

If an Electron is given enough energy it can escape the atom entirely. We call the energy required to do this the Ionisation Energy.

Note - By convention we refer to the electron at Ionisation Energy as having an energy value of zero. This means that the energies within the atom are "negative" energy values.

Absorption Spectra

An Electron is able to move to a higher energy state within the atom by absorbing a Photon of the correct energy to make the jump. We call the higher energy state an Excited State.

If a continuous spectrum is passed through a gas, Photons of the correct Energy and therefore the correct Frequency are absorbed, leaving dark bands - an Absorption spectrum.

Emission Spectra

Once in an excited state, an electron is able to return to the ground state by giving out Energy in the form of a Photon. As the Energy change is between the excited and ground states, the Photon that is emitted has exactly the same Energy as the absorbed Photon.

This is why in the above Hydrogen Spectrum diagram, the patterns of both the Emission and Absorption spectra match.

The electron can also be excited to higher than just the first excited state and transitions between multiple levels can occur:-

By applying the formula E = hf, we can see that large transitions ( W3 to W0 ) give high Energy and Frequency, therefore short Wavelength, Photons.