Spectroscopy Basics
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Spectroscopy takes advantage of the fact that all atoms and molecules absorb and emit light at
certain wavelengths. To understand why, you must understand how atoms are structured. You can
read about atomic structure in How Atoms Work, but a quick recap here will be helpful. In 1913, a
Danish scientist by the name of Niels Bohr took Ernest Rutherford's model of the atom -- a dense
nucleus surrounded by a cloud of electrons -- and made some slight improvements that better fit with
experimental data. In Bohr's model, the electrons surrounding the nucleus existed in discrete orbits,
much like planets orbiting the sun. In fact, the classic visual image we all have of atoms, such as the
one on the right, is modeled after Bohr's concept. (Scientists have since moved away from some of
Bohr's conclusions, including the idea of electrons moving around the nucleus in fixed paths, instead
envisioning electrons congregating around the nucleus in a cloud.)
In the Bohr atom, an electron in a particular orbit is associated with a specific amount of energy.
Unlike planets, which remain fixed in their orbits, electrons can hop from one orbit to another. An
electron in its default orbit is in its ground state. To move from the ground state to an orbit farther
away from the nucleus, an electron must absorb energy. When this happens, chemists say the
electron is in an excited state. Electrons generally can't remain in an excited state indefinitely.
Instead, they jump back down to the ground state, a move that requires the release of the same
energy that enabled them to become excited in the first place. This energy takes the form of a
photon -- the tiniest particle of light -- at a certain wavelength and, because wavelength and color are
related, at a certain color.
An atom absorbs energy in the form of heat, light or electricity. Electrons may move from a lowerenergy orbit to a higher-energy orbit.
Each element on the periodic table has a unique set of Bohr orbits that no other element shares. In
other words, the electrons of one element exist in slightly different orbits than the electrons of another
element. Because the internal structures of the elements are unique, they emit different wavelengths
of light when their electrons get excited. In essence, every element has a unique atomic "fingerprint"
that takes the form of a set of wavelengths, or a spectrum.
William Wollaston and Joseph von Fraunhofer developed the first spectrometer to see the spectral
fingerprints of elements. A spectrometer is an instrument that both spreads out light and displays it for
study. Light enters a narrow slit and passes through a lens that creates a beam of parallel rays.
These rays travel through a prism, which bends the light. Each wavelength is bent a slightly different
amount, so a series of colored bands is produced. A second lens focuses the light on an exit slit,
which allows one color of light to pass through at a time. Scientists often use a small telescope,
mounted on a turntable, to observe the color exiting through the slit more easily. Then, the scientist
rotates either the telescope or the prism to bring another color into view. By noting the angle of the
prism or the telescope, the wavelength of the exiting light can be determined. Using a spectroscope
to analyze a sample may take several minutes, but it can reveal much about the light source. Some
spectrometers, known as spectrographs, are set up to photograph the spectrum.
As we have noted in the section on the Bohr atom, isolated atoms can absorb and emit packets of
electromagnetic radiation having discrete energies dictated by the detailed atomic structure of the
atoms. When the corresponding light is passed through a prism or spectrograph it is separated
spatially according to wavelength, as illustrated in the following image.
Separation of light by a prism according to wavelength
Continuum, Emission, and Absorption Spectra
The corresponding spectrum may exhibit a continuum, or may have superposed on the continuum bright lines
(an emission spectrum) or dark lines (an absorption spectrum), as illustrated in the following figure.
Continuous, emission, and absorption spectra
Origin of Continuum, Emission, and Absorption Spectra
The origins of these three types of spectra are illustrated in the following figure.
Sources of continuous, emission, and absorption spectra
Thus, emission spectra are produced by thin gases in which the atoms do not experience many
collisions (because of the low density). The emission lines correspond to photons of discrete energies
that are emitted when excited atomic states in the gas make transitions back to lower-lying levels.
A continuum spectrum results when the gas pressures are higher. Generally, solids, liquids, or dense
gases emit light at all wavelengths when heated.
An absorption spectrum occurs when light passes through a cold, dilute gas and atoms in the gas
absorb at characteristic frequencies; since the re-emitted light is unlikely to be emitted in the same
direction as the absorbed photon, this gives rise to dark lines (absence of light) in the spectrum.