The biology of light – part II

July 6, 2012

In previous posts I tried to explain why light is so important for biology and how its properties can be used in biology research. The process of absorption is especially important for understanding the role of light within the field of biology, which is obvious since light can only have an effect if it interacts with matter. Molecules can absorb light because their electrons become more “energetic” and rise form the so-called ground state to one of the excited states. These different states and their significance for biology will be explained in the following. Molecules that are able to absorb light of the biologically interesting spectrum with wavelengths form approximately 100 to 800 nm are also called chromophores. Such molecules often contain a delocalized π-electron system which means that these molecules can form bonds that occur between electron orbitals (outer electron cloud around the atomic nucleus) called π-orbitals. Many of such π-bonds result in an electron system within the molecule which is flexible once it absorbs the energy of photons which make up light. The electrons within such a flexible system can then become delocalized or spread out across the molecule. Therefore, molecules which contain a delocalized π-electron system are especially sensitive to the absorption of light. In practice a π-electron system can be found in molecules containing aromatic systems and/or a relatively high number of conjugated double bonds. Figure 1 displays how pH changes influence the protonation states of an anthocyanine molecule, a plant colorant which gives flowers and berries their distinct colors in many plants. Of course, one finally only observes the non-absorbed photons. The anthocyanine molecule is a good example for the interesting color properties of molecules that can change once the atomic make-up changes. This molecule is also a good example for a π-electron system (see the aromatic rings).

Figure 1

An electron excitation can, however, only occur if photon energy matches the energy difference between the ground and the excited state. When the electrons of single atoms fall from the excited state back to the ground state they emit the previously absorbed energy again in the form of a photon which matches the wavelength of the previously absorbed photon. Therefore, absorption and emission spectra of single atoms are identical. However, things are different for molecules which of course by definition always contain at least two atoms. Bound atoms can vibrationally interact with each other which costs energy. In a molecule every electron state can be subdivided vibrational states and each vibrational state can again be split into different rotational states. Still, the photon’s energy must match the energy difference between the ground and excited state in order to become absorbed. This results in the fact that molecules can absorb a range of wavelengths and due to vibrations emit longer wavelengths containing less energetic photons. While single atoms can only absorb and emit single spectral lines (Figure 2) molecules can absorb and emit broader band light spectra.

Figure 2

Hopefully, the basics of absorption are now a bit clearer. However, we have not discussed yet how an excited electron can lose its energy again. Depending on the excited state it is in the electrons can choose different paths to fall back into the ground state. As described, during these paths the electrons emit longer wavelengths photons which are responsible for biologically interesting processes such as fluorescence, bioluminescence, or phosphorescence. But also other processes  occur that are not directly visible by eye when the electrons fall back to the ground state. These include internal conversion, intersystem crossing, resonance energy transfer, and photochemical reactions. In the following, all of these effects will be described. A so-called Jablonski diagram (Figure 3) schematically displays the different electron states and their subdivision into vibrational and rotational states. A Jablonski diagram also helps to “visibly understand” what happens to excited and returning electrons. So let’s start.

Figure 3 (Source: Olympus)

First of all electrons become excited by photon absorptions which rises their energy level to the first or second singlet state depending on the available energy and molecular properties. This process is extremely fast is depicted by the green arrows.

Internal conversion (IC) always occurs in molecules when excited electrons fall back to the ground state. During IC the absorbed energy is converted into kinetic energy in the form of vibrations or rotations. No electromagnetic radiation occurs.  Yellow curly arrows indicate this process which logically occurs between the vibrational states, but also within one vibrational state containing more rotational states (not shown).

An observable process is fluorescence. When an electron falls back from the first singlet state into the ground state it emits electromagnetic radiation in the form of a photon. However the emission wavelength is longer because the electron has lost energy due to IC on its way from the second singlet state to the first singlet state or due to IC within just the first singlet state. Fluorescence is indicated by the red down facing arrows from S1 to S0. Bioluminescence is the process of fluorescence within biomolecules such as green fluorescent protein (GFP).

Another important feature of electron states is the process of intersystem crossing. During intersystem crossing an electron moves from the excited first singlet state into the first triplet state. A triplet state is a state in which an electron can only be found once its quantum mechanical spin reverses from -1/2 to +1/2. Quantum mechanically this very unlikely and therefore a triplet state occurs less often than the other two excited states. Intersystem crossing is indicated by the blue curly line.

The process of phosphorescence occurs one an electron has managed to enter the triplet state by intersystem crossing and falls back into the ground state. As in fluorescence a photon is emitted, but it contains less energy and it time delayed with a factor of about one million because a triplet state is not stable state. In research applications (such as the earlier described fluorescence correlation spectroscopy) where only the emitted fluorescence at a specific point in time is required, a certain percentage of phosphorescence signal therefore needs to be subtracted from the total signal. Phosphorescence is indicated by the red arrow facing to the lower left corner.

Zooming out and looking not only at one molecule, but for example two neighboring molecules, makes it possible to observe two other interesting electron effects. The first one is called resonance energy transfer. If the fluorescence spectrum of a donor molecule and the absorption spectrum of a acceptor molecule match the vibrations of the former molecule’s electron can excite the latter molecule’s electron. Then a longer wavelength fluorescent emission of the second molecule can be observed even though it has not been excited directly with photons. This process is the basis of Förster Resonance Energy Transfer (FRET) which can be used to determine whether two molecules are in close proximity. If two proteins are close two each other, the chromophore electrons of the first one therefore might vibrationally excite the electrons of the second one and fluorescence of a distinct wavelength can be observed. A second interesting photo effect between two neighboring molecules is a so-called photochemical reaction. Here, a very strong excitation removes an electron from its original orbital and the molecule therefore becomes oxidized. Another molecule which receives the electron is reduced. Oxidized and reduced molecules can then participate in “regular” chemical reactions. Photochemical reactions are the basis of photosynthesis where chlorophyll electrons power the electron transport chain in plants.

I hope that this short summary of light and biology serves its purpose of demonstrating that understanding a little bit of physics and a little bit of biology and joining both, can lead to some very interesting insights! For people who are interested more in photons and their effects on electrons I want to recommend an extremely appealing Java-animated tutorial created by Olympus.

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