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Spectroscopy and Spectrometry

Splitting light and identifying chemical elements from the spectrum

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Spectroscopy and Spectrometry

970 words, estimated reading time 5 minutes.

Spectroscopy is the study of the interaction between energy and matter as a function of wavelength. Spectrometry is a technique used for the identification of elements through analysis of a spectrum.
Introduction to Astronomy Series
  1. Introduction to Astronomy
  2. The Celestial Sphere - Right Ascension and Declination
  3. What is Angular Size?
  4. What is the Milky Way?
  5. The Magnitude Scale
  6. Sidereal Time, Civil Time and Solar Time
  7. Equinoxes and Solstices
  8. Parallax, Distance and Parsecs
  9. Flux
  10. Luminosity of Stars
  11. Apparent Magnitude, Absolute Magnitude and Distance
  12. Variable Stars
  13. Spectroscopy and Spectrometry
  14. Redshift and Blueshift
  15. Spectral Classification of Stars
  16. Hertzsprung-Russell Diagram
  17. Kepler's Laws of Planetary Motion
  18. The Lagrange Points
  19. What is an Exoplanet?
  20. Glossary of Astronomy & Photographic Terms

Before we take a look at spectroscopy, lets first have a look at what light is. Light, as we know, is only a small part of a larger Electro-Magnetic (EM) spectrum.

A diagram of the EM spectrum

For the purposes of this article we will focus on the visible wavelengths only, and we will start with looking at Spectrometry.


Animation of the dispersion of light as it travels through a triangular prism

Way back in 40AD, Seneca observed the light scattering properties of glass prisms but it wasn't until 1666 when Newton observed his own spectra that the idea of light being made of colours became popular. In 1802 the English chemist William Hyde Wollaston observed dark lines (absorption lines) in spectrum from a glass prism. Later in 1814 German physicist Joseph von Fraunhofer independently rediscovered the lines and began a systematic study and careful measurement of the wavelength of these features. In all, he mapped over 570 lines, and designated the principal features with the letters A through K, and weaker lines with other letters.

Hydrogen absorption lines in the visible spectrum

If you were to observe the Sun's spectrum through a prism you may be able to see dark lines as shown in the diagram above. These are called Fraunhofer lines or absorption lines.

In 1859 Gustav Robert Kirchhoff and Robert Bunsen showed that each chemical element has a unique "signature" of emission lines, and deduced that the dark lines in the solar spectrum were caused by absorption by those elements in the upper layers of the sun. Some of the observed features are also caused by absorption in oxygen molecules in the Earth's atmosphere.

They did this by looking at the spectrum emitted by elements as they burn. In this example, Hydrogen will be burnt. By viewing hydrogen as it burns through a spectrometer a different set of lines will be observed - emission lines.

Emission lines of Hydrogen

As you can see, the emission lines from hydrogen match the absorption lines from the Suns spectrum. Kirchhoff and Bunsen deduced that the Sun's upper atmosphere must be absorbing these wavelengths due to the presence of Hydrogen. Their experiment involved looking at the spectra of the Sun as it passes through a hot gas (from the Bunsen burner) and comparing it with the spectra emitted by heating different elements.
It was during the process of developing spectroscopy that the Bunsen burner came into being.

The Kirchhoff-Bunsen Experiment

In all there are over 1000 Fraunhofer lines observable in the Sun's spectrum and because each element has its own signature, we can deduce the chemical composition of the Sun, or any unknown object by analysing the spectral lines.

What causes these lines?

Electron levels in the Bohr model

Atoms consist of protons, neutrons, and electrons. Protons are positively charged, electrons are negative, and neutrons have no charge (electrically neutral). Danish physicist Niels Bohr devised a model of the atom which helps explain absorption and emission lines. In his model, protons and neutrons are in the nucleus, the electrons orbit the nucleus. In the Bohr model electrons are only allowed to orbit at certain distances from the nucleus, much in the same way as planets can only orbit the sun at certain distances. The further away from the nucleus the more energy is needed. Each of these "distances" is called an energy level. Electrons can move between energy levels, but it does require an exchange of energy.

When we discuss the energy of a photon we can also talk about the wavelength since the two are related. The energy required is determined by the energy difference between the two levels and is different for every energy level and every element. Combining elements into molecules also changes the energy requirements.

The energy (E) of a photon (in Joules) is given by the formula:

Equation 29 - Energy of a Photon

Where h is the Planck constant (6.624 x 10-34 joule-sec) and the frequency (f) is a function of wavelength (λ). Frequency is given by the formula below:

Equation 30 - Frequency of Light

Where c is the speed of light (3x108 ms-1) and λ it's wavelength in hertz.

For an electron to move to a higher energy level it must gain energy. One way is to absorb a photon having the right amount of energy. When the electron absorbs the photon the corresponding wavelength appears to be missing from the spectrum because it has been absorbed.

When an electron moves to a lower energy level it releases the same amount of energy. This causes an emission line.

Energy levels are generally noted as n, the first energy level being n = 2 (n = 1 for the nucleus). A jump from n = 2 to n = 3 requires an absorption of energy, while moving from n = 3 to n = 2 releases it.

Going back to our hydrogen example, when it gains energy from a photon in the sun an electron makes the jump from n = 2 to n = 3 and an absorption line is formed. In this case light of 656.3nm (red). When we heat hydrogen in a burner we actually excite the electron with energy, then it releases it again. As the electron returns to n = 2 it emits the same amount of energy and we see an emission at 656.3nm.

Electrons can jump from n = 2 to n = 3, or to n = 4, 5 and so on. The amount of energy required is summarised in the table below for hydrogen. This is also known as the Balmer Series.

Transition of n3→4→25→26→27→28→29→2∞→2
Wavelength (nm)656.3486.1434.1410.2397.0388.9383.5364.6

Each different element has it's own unique energy levels and when an elemental atom is combined in a molecule the energy levels again change. Because of this we can use spectroscopy to identify almost any element or compound.

Last updated on: Friday 8th September 2017

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