Light Emission
The primary instrument for optical investigations into matter is the spectrograph. The
basic design has remained almost unchanged since the 1830’s. The principle components
are:
Slit – used to isolate the object and form images of the spectral lines
Collimating Lens – used to make the diverging light from the slit into parallel beams
Dispersive Element – separates the light into its component colors. This can be
either a prism or a diffraction grating.
Imaging Lens – used to focus the spectrum
Recording Device – to observe and record the spectrum. This has changed from the
eye to photographic film to digital imaging systems.
Nature provides only three types of spectra, each of which provide information of the
source.
Continuous spectra have all colors with no gaps and are produced by hot, dense
sources. The prototype is an incandescent light bulb.
Emission (discrete or bright line) spectra have patterns of bright lines, the pattern
varying with the element. They are produced by hot, tenuous sources. The
prototype is a neon sign
Absorption (dark line) spectra have the continuous background with dark lines of
gaps. They are produced by a hot, dense source that has an overlying cooler,
tenuous gas. The prototype is a star.
Light comes from fundamental atomic processes, and to understand light, we must
talk a bit about the structure of the atom. The study of atomic structure begins in 1896
with the discovery of the electron. You learned in grade school that atoms are composed
of electrons (negative charges), protons (positive charges), and neutrons (no charge). The
early experiments that led to a viable atomic model were performed at the beginning of
the 20th century by Ernst Rutherford. His experiments used helium nuclei as “bullets” to
probe the structure of the atom. The pattern of scattering of the helium nuclei could be
used as a judge of the structure of the atom. He concluded that atoms have a positively
charged and very tiny center (nucleus). Surrounding the nucleus was the negative charge
of the electrons.
Rutherford’s work became formalized into a “planetary model.” For the simplest
atom (hydrogen) with a single electron and a single proton, the electron was viewed as
orbiting the proton in a circle. Using Newton’s Laws on this situation, we can argue:
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The electron is kept in orbit because of a force (the electrical force).
A force implies that the electron is accelerating.
Accelerating charges radiate light.
Light carries energy, so that the electron must be losing energy. The electron
must move inward to places that require less energy. Bigger orbits require
more energy.
This basic argument can be applied time and again to show that the electron must
spiral into and collide with the proton, and that atoms should not be stable under
Newton’s Laws. This was the first failure of Newton’s Laws.
Bohr’s Model
Niels Bohr introduced a simple model that allowed atoms to exist and permitted an
explanation of the line spectra that we observed above. The model centers around two
hypotheses:
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A system of stable orbits exists surrounding every atom. Electrons can only
be on these stable orbits and the electron then has no tendency to lose energy.
Transitions are possible from one stable orbit to another. The electron must
end up with the energy of the new state.
Recall that the large orbits have large energy and small orbits have small energy. If
the electron drops to a smaller orbit, energy must be lost and a photon of light is emitted.
Similarly, if the electron makes a transition to a larger orbit, energy must be gained. One
source of this energy is to absorb light. Since the stable orbits come at precise energies,
the transition energies are also precise. Thus the photon has only a certain amount of
energy. Energy is related to its frequency, frequency is inversely related to the
wavelength, and wavelength is the same as color. Bohr’s hypotheses show that atoms
should emit emission spectra. Every one of the 92 naturally occurring elements has a
unique pattern of stable orbits, and thus a unique set of spectral lines.
If enough energy is given to the atom, one or more electrons can be dislodged from
the atom entirely. This process is called ionization. The element retains its chemical
identity, since the number of protons in the nucleus hasn’t changed. But the patterns of
stable orbits has shifted and along with it, the lines emitted in the spectrum. So not only
can we identify which element is present in a sample, we can also tell the state of
ionization.
Continuous Radiation
When we measure the intensity of the continuous spectrum, we find that one color is
the brightest. The intensities of the other wavelengths follow a well-defined curve called
the radiation curve, blackbody curve, or Planck curve. By measuring the area under
the curve or by measuring the wavelength of the peak of the curve, we can measure the
temperature of the source. The first idea is called Stefan-Boltzmann’s Law:
Area under BB Curve  Total energy emitted  T4
The size of the blackbody curve grows quickly as the temperature increases. The
other idea is called Wien’s Displacement Law.
 max 
1
T
As the temperature decreases the source emits longer wavelength light. Using Wien’s
Law we see that everything in nature must emit light. Hot objects emit short wavelengths
while cool objects emit long wavelengths.
Fluorescence and Phosphorescence
Consider an atom with only three atomic energy states. If the electron in the ground
state gains enough energy, it can jump to the top energy state. In normal material the
electron returns directly to the ground state from which it came. In fluorescent and
phosphorescent material, however, there is a strong preference to return to the ground
state in smaller steps, emitting longer wavelengths at each step. Thus UV may stimulate
the process, but visible light is emitted. In fluorescent material the intervening states are
normal atomic states of very short duration (10-8 s), so the light is only emitted while the
stimulating UV source is available. In phosphorescent material the intervening state is
metastable with lifetimes 100,000 times as long as normal states. These long lived states
cause the material to emit light well after the UV stimulation source is turned off.
Lasers (Light Amplification by the Stimulated Emission of Radiation)
Here we attempt to overpopulate the metastable state of an atom. A source of high
energy makes the electrons jump to an excited state. This is a normal, short-lived state
which quickly decays to the long-lived metastable state. Electrons from many atoms
accumulate in the metastable state. If a photon passes having the energy of the final
downward transition, it will stimulate an avalanche of other photons as all of the
electrons in the metastable state make the downward transition together.