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10.15 Emission Lines and Bands

As an electron drops to a lower energy level, it loses energy by emitting a photon. The photon's energy equals the difference in energy between its energy level and the one below it. Electrons in an atom can experience only certain, fixed changes in energy level. For instance, if an electron starts in level 3, it can drop only to level 2 or 1. Atoms can emit only photons of energy corresponding to these differences in energy level. As a result, each element can emit only certain wavelengths of light. A familiar astronomical example is the radiation produced by a sample of hydrogen gas containing neutral (or uncharged) atoms with their electrons at different energy levels. An emission line results from the emitted photons and appears in a projected spectrum as a line or narrow bar of color.

Spectral emission lines are extraordinarily useful in astronomy. Each element has a unique number of protons in the nucleus and electrons orbiting the nucleus. This means in turn that each element has a unique set of electron energy levels that create a specific set of wavelengths represented by the emission lines. If you see a set of emission lines, you can match it with a single element and infer that atoms of that element are present and glowing in the distant object. If more than one element is present, the pattern is more complex because there are more lines, but the principle is the same. You can deduce the object's composition, even without having a sample! By contrast, the smooth spectrum of thermal radiation only gives information about the temperature of the object. At the same temperature, a lump of iron or a carbon rod or a cloud of hydrogen all emit the same thermal spectrum. Thus spectral lines are much more useful if we want to determine the chemical composition of a distant object.

 

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The Bohr model for the atom, showing the nucleus and different levels that an electron can occupy. Click here for original source URL



If the electrons in an atom are in their lowest possible energy level, that atom cannot produce an emission line. This is because the electrons cannot drop to any lower energy level. An atom in which all electrons are in the lowest possible energy level is said to be in its ground state. An atom in which one or more electrons are in energy levels higher than the lowest available ones is said to be in an excited state. Excited states usually last only a fraction of a second before the electrons decay to the lowest available energy level — trying to reach equilibrium. Atoms generally need to be disturbed to produce and maintain excited states. This can happen in two ways. Radiation will add energy to a gas and so cause electrons to raise their energy state. In a hot gas, the same role can be served by collisions of the atoms or molecules themselves. Heating a gas enclosed within a certain volume increases the velocity of atoms and so increases the probability that they will collide with each other.

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Transition levels and wavelengths for Hydrogen. Click here for original source URL.

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Emission line spectrum of hydrogen in the visible part of the spectrum. Click here for original source URL.

Where in the electromagnetic spectrum would we look for emission lines from atoms? Take hydrogen, the simplest element. The energy required to raise an electron from the ground state to be free of the atom is the largest amount of energy that can result in a spectral line. Therefore, it corresponds to the shortest wavelength feature we might see. This wavelength is about 90mm, which is in the ultraviolet too blue for our eyes to see. Other electron transitions in hydrogen have smaller energy differences, so they yield redder spectral lines. Heavier elements have more electron energy levels so they have more possible transitions and a denser thicket of emission lines. But for the most common elements like carbon and nitrogen and oxygen and silicon, the spectral lines fall in the same region of the electromagnetic spectrum. Most of the useful emission lines fall in the decade of wavelength from 100 nm to 1000 nm (or 0.1 micron to 1 micron). This spans the visible spectral range but also extends to ultraviolet and infrared wavelengths.
 

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Absorption spectrum of liquid water across a wide wavelength range. Strong Bands: 2900nm, 1950nm, 1450nm. Medium bands: 1200nm, 900nm. Weak bands: 820nm, 730nm. Click here for original source URL.

Molecules are grouping of atoms that can share their electrons. Therefore, the electron structure in a molecule is more complex than in an atom. The electron's path may take it around two or more nuclei. As a result, the emission line structure of a molecule can be complex. For example, a gas containing water molecules (H2O) has many more emission lines than a gas containing single H and O atoms. The molecule has various ways of responding to a disturbance in addition to having its electrons change energy levels. For example, it may vibrate like two balls linked with a spring, or it may rotate. As a result, the energy levels from a molecule are vastly more numerous, and the resulting emission lines blend together into a broader emission feature called an emission band. The rest of the story is the same. A given molecule (such as H2O) can produce only certain emission bands, allowing us to identify the molecule in a remote source.

In order for atoms and molecules to produce clear emission lines and bands, they must be detached from one another, as in gases. If the atoms were linked together, they would form molecules, and if the molecules were linked, they would form a solid or liquid. In most solid or liquid substances, the electron structure is so complex that emissions are not confined to one wavelength, but are smeared out. Therefore, emission features of solids and liquids are barely discernible. Most emission lines and bands arise from gases.
 

 

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Diagram of the Electromagnetic Spectrum. Click here for original source URL.

Where in the electromagnetic spectrum would we look for emission features from molecules? Some spectral features from molecules are seen in the visible spectrum. However, molecules have shapes that allow them to vibrate and oscillate in many different ways. Most of these modes involve less energy than a typical electron transition, so the spectral features are typically found at longer wavelengths. Many of the most important spectral features from molecules are found in the infrared or even the microwave part of the electromagnetic spectrum.