Spectra in the Lab

Every chemical element has a unique `signature' which can be revealed by analyzing the light it gives off. This is done by spreading the light out into a spectrum — basically, a rainbow.

It may seem remarkable that we know anything about the composition of distant stars. In fact, we know a great deal, not only about the chemical elements present, but also about physical conditions. The key is to spread a star's light out, producing a spectrum like the one shown in Fig. 1.

Fig. 1. A spectrum. The light — in this case, from an ordinary light bulb — has been spread out into different colors. The scales above and below the spectrum will be explained below.


The nature of matter was debated for thousands of years. Suppose you have a chunk of gold, for example, and you start cutting it into smaller and smaller pieces. Can you always cut any piece, even a very small one, into two smaller pieces of gold? Or is there some minimum size a piece of gold can have? We now know the answer — the smallest possible piece is one atom of gold. Atoms are the building blocks of matter. There are about one hundred different kinds of atoms in the universe — these are known as the chemical elements.

The nature of light posed a very similar question: Is light composed of waves or of particles? If light is waves, then one can always reduce the amount of light by making the waves weaker, while if light is particles, there's a minimum amount of light you can have — a single `particle' of light. In 1905, Einstein found the answer: light is both! In some situations it behaves like waves, while in others it behaves like particles. This may seem strange and mystical, but it describes the nature of light very well.

A wave of light has a wavelength, defined as the distance from one crest of the wave to the next, and written using the symbol λ. The wavelengths of visible light are quite small: between 400 nm and 650 nm, where 1 nm = 10-9 m is a nanometer — one billionth of a meter. In Fig. 1, the scale on the bottom shows wavelengths in nanometers; as you can see, red light has long wavelengths, while blue light has short wavelengths.

A particle of light, known as a photon, has an energy E. The energy of a single photon of visible light is tiny, barely enough to disturb one atom; we use units of `electron-volts', abbreviated as eV, to measure the energy of photons. In Fig. 1, the scale on the top shows energies in electron-volts; photons of red light have low energies, while photons of blue light have high energies.

The relationship between energy E and wavelength λ is one of the most basic equations of quantum physics:

λ =  h c


Here c is the speed of light, and h is known as Planck's constant. Both c and h are constants of nature; they never change. From our point of view, the significance of this equation is that energy E and wavelength λ are inversely proportional to each other, and the relationship between them is the same in a laboratory on Earth and in the most distant stars and galaxies.


As quantum physics developed, physicists began to understand another puzzle. The light given off by atoms in a hot dilute gas does not form a spectrum of all colors as in Fig. 1; instead, only some colors are present, and each element produces a unique pattern, as shown in Fig. 2. Why do hot atoms behave this way? The answer involves two key ideas: first, each atom contains one or more electrons orbiting a central nucleus; second, in atoms of any given element, only certain orbits are allowed, and a very specific amount of energy is involved when an electron jumps from one orbit to another.

Hydrogen: a simple atom with a simple spectrum. Besides the three lines shown here, you may be able to see another in the blue near 410 nm. Spectra of hydrogen.
Helium: slightly more complex than hydrogen, with one yellow line and a number in the blue. Spectral of helium.
Neon: a very large number of lines in the red give neon signs their distinctive pink colors, but notice the two green lines. Spectra of neon.
Argon: the pastel color of argon is due to a wide range of lines throughout the spectrum. Spectra of argon.
Mercury: the strongest line, at 546 nm, gives mercury a greenish color. Spectra of mercury.
Fig. 2. When heated in a electric discharge tube, each element produces a unique pattern of spectral `lines'.

Fig. 3 illustrates this for hydrogen, which has only one electron. The allowed orbits of an electron in a hydrogen atom can be numbered using the symbol n, with n = 1 for the orbit closest to the nucleus, n = 2 for the next one out, and so on. For orbit n, the amount of energy required to completely separate the electron from the nucleus is

En = 13.6 eV


This quantity En is the energy level of orbit n. For example, an electron in orbit n = 2 requires energy E2 = 3.4 eV to be separated from the nucleus, while an electron in orbit n = 3 requires only E3 = 1.51 eV; we say that orbit n = 3 is less tightly bound to the nucleus than orbit n = 2. When an electron jumps down from orbit n = 3 to orbit n = 2, it gives off energy E = E2 − E3 = 1.89 eV. This energy can be carried away by a photon of light; the red line of hydrogen in Fig. 2 is due to photons with exactly this energy. Likewise, electrons jumping from orbit n = 4 to orbit n = 2 produce the blue-green line, and electrons jumping from orbit n = 5 to orbit n = 2 produce the deep blue line.

Energy levels and transitions of hydrogen.
Fig. 3. Energy levels (horizontal lines), and downward jumps (arrows) of hydrogen. The wiggly arrows in color represent the photons produced when an electron jumps down from one orbit to another. To save space, the lowest level (n = 1) is not shown.

What happens when an electron in a hydrogen atom jumps up to a higher orbit? This takes energy, which has to come from somewhere. One way to supply the energy is with a photon, but the photon has to have exactly the right amount of energy — no more, and no less. For example, an electron in orbit n = 2 can jump up to orbit n = 3 if it absorbs a photon with energy E = E2 − E3 = 1.89 eV.

Similar processes of emission and absorption happen in atoms of other elements. For atoms with more than one electron, the physics becomes much more complex, but the basic idea that electrons have only certain allowed orbits still holds. Each element has a different set of allowed orbits, so each element emits or absorbs photons with different energies — and therefore, different wavelengths. This is just what we see in Fig. 2!

Molecules also produce spectral lines, but their spectra are much more complex than the spectra of single atoms, and typically show broad bands instead of narrow lines, as in Fig. 4.

A spectrum of air.
Fig. 4. A spectrum of air. The bright bands are due to molecular oxygen (O2), molecular nitrogen (N2), and other molecules.


Examining different kinds of light with a spectroscope reveals a wide variety of spectra. The appearance of a spectrum tells us something about the physical conditions which produce the light.

For example, a continuous spectrum, like the one at the top of Fig. 5, is a featureless rainbow of color. This kind of spectrum is the hallmark of `black-body' radiation (so-called because a black object, heated until it glows, emits this kind of light). A hot solid, liquid, or very dense gas produces a continuous spectrum; while a wide range of wavelengths are always present, the overall color of the light depends on the temperature. For example, a bar of iron heated in a fire glows dull red; if heated more it glows orange, and if heated well beyond its melting point it shines with a brilliant blue-white light.

In contrast, an emission spectrum, like the one in the middle of Fig. 5, consists of bright lines or bands on a dark background. Emission spectra are produced when atoms of a dilute gas are `excited' — in effect, heated — by an electrical current, ultraviolet radiation, or some other source of energy. Excited atoms have electrons in high orbits, and these emit photons with specific wavelengths when they jump back down to lower orbits (as explained above). Neon signs produce emission spectra; so do objects like the Lagoon Nebula (M8) and the Ring Nebula (M57).

Continuous spectrum: a smooth gradation of color, with no distinct features. Continuous spectrum.
Emission spectrum: bright lines on dark background. Emission spectrum.
Absorption spectrum: dark lines superimposed on a continuous spectrum (shown in black & white because the subtle dark lines are hard to reproduce in color). Absorption spectrum
Fig. 5. Basic types of spectra.

Finally, an absorption spectrum, like the spectrum of sunlight shown in the bottom of Fig. 5, consists of dark lines or bands on top of a continuous spectrum. Absorption spectra are produced when light from a hot object travels through a cooler, dilute gas. When a photon with exactly the right wavelength encounters an atom of the cool gas, it is absorbed and its energy used to kick an electron into a higher orbit; if enough atoms of gas are present, all the photons of that wavelengths are absorbed, while photons with other wavelengths get through. The atmospheres of stars produce absorption spectra.

An element produces bright and dark lines with the same wavelengths. For example, hydrogen has three prominent lines with wavelengths of 434 nm, 486 nm, and 656 nm; these appear dark if the hydrogen is absorbing light, and bright if it is emitting light, but the same three wavelengths are seen in either case.

In some situations, we find spectra which mix different kinds of features: for example, a continuous spectrum with bright emission lines superimposed. Some older stars have spectra with dark absorption lines and bright emission lines; this is usually a sign that the star is ejecting gas in a stellar wind.


In the lab, we will explain how to use spectroscope, and how to adjust it so you can measure wavelengths accurately. You will then view and sketch different types of spectra.

We will set up several different discharge tubes, in which elements are excited electrically. You will be asked to identify these elements by looking at the light they produce using your spectroscope. The elements in question will be among those in Fig. 2.

We will also set up a light source which produces a bright spectral line, and ask you to measure the wavelength λ of this line. Once you've done this, you can identify the element involved by looking at the table printed on your spectrometer.

Finally, we will have you sketch and classify several additional light sources, possibly including a fluorescent light, a street-light, and a LED flashlight.



Joshua E. Barnes      (barnes at ifa.hawaii.edu)
Updated: 25 September 2011
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