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7 Light in Biology and Medicine
3. Metal band transitions: Many of the valence electrons in metals are delocalized
over the entire substance, occupying a dense set of possible quantum states
determined by the closely packed metal atoms. Metals at a finite temperature
have an ongoing dance of electrons moving up and down in energy near the
Fermi level. Exposing the surface of the metal to light gives extra energy
to electrons near the conduction band. The electrons quickly release excess
energy, causing a scattered wave. If the metal surface is smooth compared to the
wavelength of light, the scattered waves will form a reflected wave. The colors
reflected depend on the density of the energy levels near the Fermi level. The
valence electrons of copper and gold include atomic d-waves (i.e. electron wave
functions which carry two units of angular momentum). These d-wave electrons
form a ‘d-band’ in the quantum energies available to electrons in copper and
gold. Incoming photons need about 2 eV to excite electrons from the d-band
into the conduction band. This corresponds to photons in the blue part of the
spectrum. Some energy carried by these ‘blue photons’ is not returned as blue
light, but rather gets transferred to other electrons. This makes copper and gold
have an orange-yellowish color. Silver metal also has electrons in a d-band, but
they require ultraviolet photons to release them to the conduction band.
4. Semiconductor band transitions: Semiconductors have a very small population
of electrons in the conduction band, with most of the electrons tied to individual
atoms. Those that are least tied to atoms fill the valence band close to the
conduction band, and can be excited into the conduction band by incoming
light in the infrared and visible range of frequencies. Doping the silicon or
germanium with phosphorus or arsenic adds excess electrons, while doping
with boron causes unfilled electron energy levels (called “holes”). The dopants
create a new “donor” or “acceptor” band of quantum states between the
conduction and valence bands. The position of the bands can be selected to
match the photon energies which the material will absorb.
5. Transitions between unfilled atomic inner shell states in transition metals:
Transitions of electrons in the unfilled “d” and “f” atomic orbitals taking part in
ligand bonding can resonantly absorb visible light, giving color to such metal
ions and metal complexes.
6. Photoinduced charge transfer: Light-sensitive chemicals change chemical state
by exposure to light. In the process, photons absorbed from the light give
sufficient energy to cause electrons to jump from one atom to another. Silver
bromide in photographic film becomes silver and bromine by exposure to light.
Chlorophyll a and b absorb light energy, causing a series of charge transfer
reactions.
7. Color centers: Impurities and dislocations in some otherwise clear crystals can
cause them to exhibit color, as electrons localized near those impurities (or
the absence of electrons, i.e. ‘holes’) may have less binding energy than the
electrons taking part in the covalent or ionic binding of the crystal. Exciting
one of these electrons or holes can often be done with visible light, rather than
ultraviolet.
7 Light in Biology and Medicine
3. Metal band transitions: Many of the valence electrons in metals are delocalized
over the entire substance, occupying a dense set of possible quantum states
determined by the closely packed metal atoms. Metals at a finite temperature
have an ongoing dance of electrons moving up and down in energy near the
Fermi level. Exposing the surface of the metal to light gives extra energy
to electrons near the conduction band. The electrons quickly release excess
energy, causing a scattered wave. If the metal surface is smooth compared to the
wavelength of light, the scattered waves will form a reflected wave. The colors
reflected depend on the density of the energy levels near the Fermi level. The
valence electrons of copper and gold include atomic d-waves (i.e. electron wave
functions which carry two units of angular momentum). These d-wave electrons
form a ‘d-band’ in the quantum energies available to electrons in copper and
gold. Incoming photons need about 2 eV to excite electrons from the d-band
into the conduction band. This corresponds to photons in the blue part of the
spectrum. Some energy carried by these ‘blue photons’ is not returned as blue
light, but rather gets transferred to other electrons. This makes copper and gold
have an orange-yellowish color. Silver metal also has electrons in a d-band, but
they require ultraviolet photons to release them to the conduction band.
4. Semiconductor band transitions: Semiconductors have a very small population
of electrons in the conduction band, with most of the electrons tied to individual
atoms. Those that are least tied to atoms fill the valence band close to the
conduction band, and can be excited into the conduction band by incoming
light in the infrared and visible range of frequencies. Doping the silicon or
germanium with phosphorus or arsenic adds excess electrons, while doping
with boron causes unfilled electron energy levels (called “holes”). The dopants
create a new “donor” or “acceptor” band of quantum states between the
conduction and valence bands. The position of the bands can be selected to
match the photon energies which the material will absorb.
5. Transitions between unfilled atomic inner shell states in transition metals:
Transitions of electrons in the unfilled “d” and “f” atomic orbitals taking part in
ligand bonding can resonantly absorb visible light, giving color to such metal
ions and metal complexes.
6. Photoinduced charge transfer: Light-sensitive chemicals change chemical state
by exposure to light. In the process, photons absorbed from the light give
sufficient energy to cause electrons to jump from one atom to another. Silver
bromide in photographic film becomes silver and bromine by exposure to light.
Chlorophyll a and b absorb light energy, causing a series of charge transfer
reactions.
7. Color centers: Impurities and dislocations in some otherwise clear crystals can
cause them to exhibit color, as electrons localized near those impurities (or
the absence of electrons, i.e. ‘holes’) may have less binding energy than the
electrons taking part in the covalent or ionic binding of the crystal. Exciting
one of these electrons or holes can often be done with visible light, rather than
ultraviolet.