Wavelength (l) = distance between peaks
Frequency (u) = waves/sec, # of wave crests passing a point per unit time
Speed of light (C) = 3 x 1010 cm s-1 = 186,000 miles s-1
C = lu
As a particle, light behaves as descret packet of energy:
Quantum of light, or a photon, equals a unit of electromagnetic radiation.
Quantum Energy (E) = hu
h= 1.58 x 10-34 cal s-1 photon-1 (Plank's constant)
If E = hu and C = lu
then u = C/l
and E = h (C/l)
According to this last equation, shorter wavelengths of light have higher quantum energy.
For biological systems, light is usually measured in terms of moles of photons:
fluence = moles / unit area (µmol m-2)
fluence rate = moles / unit area / unit time (µmol m-2 s-1)
1 mole of photons = 1 einstein
Stark-Einstein Law of light absorption: A molecule can absorb one photon at a time causing the excitation of one electron.
In a single atom, the displaced electron can fall back to the ground state and dissipate the energy by emitting light (fluorescence) of the nearly the wavelength or by emitting heat. Due to the discrete nature of atomic orbitals, atoms can absorb wavelengths that match the energy characteristics of each orbital. This results in very narrow absorption and emission bands. (other forms of energy than light can be used to excite atoms, as in fluorescent bulbs)
In molecules, electrons are shared by the atoms of the molecule. The resulting molecular orbitals are not as tightly defined as in an atom due to variations in bond behavior. This bond variation results in a broadening of the energy levels for each excited state and molecules have broad absorption bands. In addition, a complex molecule can have multiple excited states that result in light absorption at multiple wavelengths.
The light absorption characteristics of a molecule can be diagnostic of it's molecular properties and function. There is a "tuning" or resonance between light and the electron orbitals of a molecule such that if a molecule contains electrons that can resonate at a certain frequency, it can absorb light at those frequencies. The longer the "resonance pathway", the longer the wavelength of light that can be absorbed. This will be demonstrated in class.
The figure below shows on the left of the green and black squiggly lines, red and blue excitation levels for a pigment molecule like chlorophyll. Because blue light is more energetic, blue light can excite electrons to the second excited state whereas the energy in red light can only elevate an electron to the first excited state.
After light is absorbed and an electron is raised to an excited state, there are several possible outcomes. As with a single atom, the excited electron can return to the ground state and give off heat or fluorescence. In addition, if the pigment is near another pigment of similar electronic configuration, the energy can be transferred by "excitation energy transfer". In excitation energy transfer, the excited electron on one molecule drops to it ground state and the energy is transferred to the neighbor molecule rasing an electron on that molecule to an excited state. The electron doesn't change molecules. Photochemistry is another possible outcome. For photochemistry to occur, the excited electron has to move to another molecule. The physical transfer of an electron changes the chemical properties of both the electron donor and the electron acceptor. Controlled photochemistry is what makes photosynthesis work. The lifetime of an electron in an excited state is on the order of a nanosecond.
When a molecule is excited and emits fluorescence as the electron returns to the ground state, the wavelength of light is always of a lower energy than the light that was absorbed (there is always some energy loss for any process). Since the emitted light is of lower energy, it is of a longer wavelength that what was absorbed during excitation. Thus, fluorescence is always shifted towards the red end of the spectrum. This property is known as Stokes Shift. This processes will be demonstrated in class.