Investigating Light with Spectroscopy and SpectrometryA look at what light is together with how we use spectroscopy and spectrometry to better understand the universe and composition of stars.
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Spectroscopy studies the interaction between energy and matter as a function of wavelength. Spectrometry is a technique used to identify elements through a spectrum analysis. Before we look at spectroscopy, let's first look at what light is. Light, as we know, is only a small part of a larger Electromagnetic (EM) spectrum.
For this article, we will focus only on the visible wavelengths.
Spectroscopy and Spectrometry
Way back in 40AD, Seneca observed the light scattering properties of glass prisms, but it wasn't until 1666 when Newton observed his spectra, that the idea of light is made of colours became popular. In 1802, the English chemist William Hyde Wollaston observed dark lines (absorption lines) in the 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.
If you see the Sun's spectrum through a prism, see it 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 Sun's upper layers. Some of the observed features are also caused by the absorption of 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.
As you can see, the emission lines from hydrogen match the absorption lines from the Sun's spectrum. Kirchhoff and Bunsen deduced that the Sun's upper atmosphere must absorb these wavelengths due to the presence of Hydrogen. Their experiment involved looking at the spectra of the Sun as it passes through hot gas (from the Bunsen burner) and comparing it with the spectra emitted by heating different elements. During the process of developing spectroscopy, the Bunsen burner came into being.
In all, over 1000 Fraunhofer lines are observable in the Sun's spectrum, and because each element has its signature, we can deduce the chemical composition of the Sun or any unknown object by analysing the spectral lines.
What causes Fraunhofer lines in Spectroscopy?
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 like 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 an energy exchange is required.
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 formula gives the energy (E) of a photon (in Joules):
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 &lambda is wavelength in hertz.
For an electron to move to a higher energy level, it must gain energy. One way is to absorb a photon with the right amount of energy. When the electron absorbs the photon, the corresponding wavelength appears missing from the spectrum because it has been absorbed.
When electrons move to a lower energy level, they release the same 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 energy absorption 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 jumps from n = 2 to n = 3, forming an absorption line. In this case, the light is 656.3nm (red). When we heat Hydrogen in a burner, we 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 n = 4, 5, etc. The amount of energy required is summarised in the table below for Hydrogen. This is also known as the Balmer Series.
| Transition of n | 3→ | 4→2 | 5→2 | 6→2 | 7→2 | 8→2 | 9→2 | ∞→2 |
|---|---|---|---|---|---|---|---|---|
| Name | H-α | H-β | H-γ | H-δ | H-ε | H-ζ | H-η | |
| Wavelength (nm) | 656.3 | 486.1 | 434.1 | 410.2 | 397.0 | 388.9 | 383.5 | 364.6 |
| Colour | Red | Blue-green | Violet | Violet | Ultraviolet | Ultraviolet | Ultraviolet | Ultraviolet |
Each different element has its 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.
