Green Fluorescent Protein (GFP) has existed for more than one hundred and sixty million years in one species of jellyfish, Aequorea victoria. The protein is found in the photoorgans of Aequorea.
The discovery of GFP
The jellyfish Aequorea victoria is bioluminescent, i.e. it produces light with the help of
chemical reactions that provide the energy for photon emission and emits green light as first
described by Davenport and Nicol (1955). In 1960, Osamu Shimomura joined the laboratory
of Frank Johnson at Princeton to clarify the molecular mechanism of the bioluminescence of
Aequorea victoria. Shimomura came from Nagoya University, where he had completed
extensive work on the bioluminescence of the small ostracod Cypridina, together with Prof.
Y. Hirata. Aequorea victoria jellyfish were collected during the following summers in Friday
Harbor in the Puget Sound of Washington state on the coast of the Pacific Ocean
(Shimomura, 2005).
The active component of the Aequorea bioluminescence was identified as a protein, named
aequorin, emitting blue light in a Ca2+- dependent manner (Shimomura et al., 1962). That the
light emission of purified aequorin peaked in the blue part of the visible spectrum came as a
surprise, since the bioluminescence of Aequorea victoria is distinctly green.
Structure
GFP has a typical beta barrel structure, consisting of one β-sheet with alpha helix(s) containing the chromophore running through the center.Inward facing sidechains of the barrel induce specific cyclization reactions in the tripeptide Ser65–Tyr66–Gly67 that lead to chromophore formation. This process of post-translational modification is referred to as maturation. The hydrogen bonding network and electron stacking interactions with these sidechains influence the color of wtGFP and its numerous derivatives. The tightly packed nature of the barrel excludes solvent molecules, protecting the chromophore fluorescence from quenching by water.
Use
The availability of GFP and its derivatives has thoroughly redefined fluorescence microscopy and the way it is used in cell biology and other biological disciplines.While most small fluorescent molecules such as FITC (fluorescein isothiocyanate) are strongly phototoxic when used in live cells, fluorescent proteins such as GFP are usually much less harmful when illuminated in living cells. This has triggered the development of highly automated live cell fluorescence microscopy systems which can be used to observe cells over time expressing one or more proteins tagged with fluorescent proteins. For example, GFP had been widely used in labelling the spermatozoa of various organisms for identification purposes as in Drosophila melanogaster, where expression of GFP can be used as a marker for a particular characteristic. GFP can also be expressed in different structures enabling morphological distinction. In such cases, the gene for the production of GFP is spliced into the genome of the organism in the region of the DNA which codes for the target proteins, and which is controlled by the same regulatory sequence; that is the gene's regulatory sequence now controls the production of GFP, in addition to the tagged protein(s). In cells where the gene is expressed, and the tagged proteins are produced, GFP is produced at the same time. Thus, only those cells in which the tagged gene is expressed, or the target proteins are produced, will fluoresce when observed under fluorescence microscopy. Analysis of such time lapse movies has redefined the understanding of many biological processes including protein folding, protein transport, and RNA dynamics, which in the past had been studied using fixed (i.e. dead) material. .
Another powerful use of GFP is to express the protein in small sets of specific cells. This allows researchers to optically detect specific types of cells in vitro (in a dish), or even in vivo (in the living organism. Genetically combining several spectral variants of GFP is a useful trick for the analysis of brain circuitry (Brainbow). Other interesting uses of fluorescent proteins in the literature include using FPs as sensors of neuron membrane potential, tracking of AMPA receptors on cell membranes , viral entry and the infection of individual influenza viruses and lentiviral viruses,etc.
GFP in nature
The purpose of both bioluminescence and GFP fluorescence in jellyfish is unknown. GFP is co-expressed with aequorin in small granules around the rim of the jellyfish bell. The secondary excitation peak (480nm) of GFP does absorb some of the blue emission of aequorin, giving the bioluminescence a more green hue. The serine 65 residue of the GFP chromophore is responsible for the dual peaked excitation spectra of wild type GFP. It is conserved in all three GFP isoforms originally cloned by Prasher. Nearly all mutations of this residue consolidate the excitation spectra to a single peak at either 395nm or 480nm. The precise mechanism of this sensitivity is complex, but probably involves donation of a hydrogen from serine 65 to glutamate 222, which influences chromophore ionization. Since a single mutation can dramatically enhance the 480nm excitation peak, making GFP a much more efficient partner of aequorin, A. victoria appears to evolutionarily prefer the less-efficient, dual peaked excitation spectrum. Roger Tsien has speculated that varying hydrostatic pressure with depth may effect serine 65's ability to donate a hydrogen to the chromophore and shift the ratio of the two excitation peaks. Thus the jellyfish may change the color of its bioluminescence with depth. Unfortunately, a collapse in the population of jellyfish in Friday Harbor, where GFP was originally discovered, has hampered further study of the role of GFP in the jellyfish's natural environment.
Saturday, July 25, 2009
Saturday, July 18, 2009
ORGANIC RADICALS
Background & Introduction
A radical is an atomic or molecular species having an unpaired, or odd, electron. Some radicals, such as nitric oxide (NO), are relatively stable, but most are so reactive that their isolation and long-term study is not possible under normal laboratory conditions. The electrons in most stable organic compounds are paired in atomic or molecular orbitals, so the total electron count is an even number. Molecular oxygen (O2) is a rare example of a stable biradical (two unpaired electrons having the same spin), with an even number of electrons.
Early chemists used the term "radical" for nomenclature purposes, much as we now use the term "group". Many doubted that such open-valenced species could exist, although there was circumstantial evidence for their participation in gas phase reactions. Credit for the first isolation and characterization of a "free radical" goes to Moses Gomberg, a young instructor at the University of Michigan. In 1900 Gomberg attempted a synthesis of hexaphenylethane by reacting triphenylmethyl chloride with finely divided metals such as silver and zinc. When air was excluded from the reaction, he obtained a yellow solution, the color of which darkened reversibly on heating and cooling. This solution yielded a colorless, crystalline C38H30 hydrocarbon which Gomberg assumed to be hexaphenylethane.
If the yellow solution was exposed to air (or oxygen) a C38H30O2 peroxide was obtained, and identified by reduction to the known alcohol, triphenylmethanol. In a similar fashion the yellow solution reacted with iodine to produce triphenylmethyl iodide.
Gomberg concluded that the colored solutions contained reactive triphenylmethyl free radicals, formed by thermal dissociation of their dimer (Keq = 2 • 10–4 at 25º C). The exceptional stability of this carbon radical is attributed to odd electron delocalization into the three phenyl rings. Discrete Kekule formulas demonstrate that this benzyl-like decocalization places the electron on ortho and para carbons, but not on meta carbons.
The resonance structures may give the impression that the triphenylmethyl radical is planar (flat). Actually the phenyl groups are turned by about 35º, producing a shape similar to a three bladed propellor. Despite this twist, the p-pi orbital overlap is still over 80%, so the electron delocalization is not seriously diminished.
More than fifty years later, the reactive dimer of triphenylmethyl radical was shown to be the para-coupled compound and not hexaphenylethane. The steric crowding of phenyl groups in the simple ethane dimer is apparently so severe that bonding between two 3º-carbon atoms is prohibited. Since the electron delocalization places radical character at the para carbons of the phenyl groups, bonding to this relatively unhindered location is preferred, although at the cost of one benzene ring's aromaticity. If the para-locations are themselves hindered by large meta substituents, then an unstable hexaarylethane may actually be formed.
Other relatively stable radicals, such as galvinoxyl have been prepared and studied. These species usually owe their stability to a combination of odd electron delocalization and steric hindrance to dimerization, as the ortho tert-butyl groups in galvinoxyl demonstrate. The term "free radical" is now loosely applied to all radical intermediates, stabilized or not.
Detection and Observation of Radicals
Only triphenymethyl and a few other stabilized radicals may be generated in concentrations suitable for examination by traditional laboratory methods. Evidence for the transient existence of more reactive radical species in chemical reactions usually requires special techniques, including low-temperature isolation in solids and high speed spectroscopic probes. However, an interesting chemical detection of the methyl radical was carried out by the Austrian chemist Fritz Paneth not long after Gomberg's preparation of triphenylmethyl radical. The Paneth experiment involved gas phase thermal decomposition of tetramethyllead to methyl radicals and lead atoms in a glass tube.Gaseous tetramethyllead is carried through the glass reactor tube in a stream of nitrogen.
Electron Paramagnetic Resonance
The same unpaired or odd electron that renders most radical intermediates unstable and highly reactive may be induced to leave a characteristic "calling card" by a magnetic resonance phenomenon called "electron spin resonance" (esr) or "electron paramagnetic resonance" (epr). Just as a proton (spin = 1/2) will occupy one of two energy states in a strong external magnetic field, giving rise to nmr spectroscopy; an electron (spin = 1/2) may also assume two energy states in an external field. Because the magnetic moment of an electron is roughly a thousand times larger than that of a proton, the energy difference between the spin states falls in the microwave region of the spectrum (assuming a moderate magnetic field strength). The lifetime of electron spin states is much shorter than nuclear spin states, so esr absorptions are much broader than nmr signals.
One way of improving the signal to noise ratio in esr spectra is to display them as first derivatives rather than absorptions. In practice, esr spectra may be quite complex.This complexity is the result of hyperfine splitting of the resonance signal by protons and other nuclear spins, an interaction similar to spin-spin splitting in nmr spectroscopy. For example, the esr signal from methyl radicals, generated by x-radiation of solid methyl iodide at -200º C, is a 1:3:3:1 quartet (predicted by the n + 1 rule). The magnitude of signal splitting is much larger than nmr coupling constants (MHz rather than Hz), and is usually reported in units of gauss. The complexity of the triphenylmethyl spectrum is due to three different hyperfine splittings: 3 para hydrogens, 6 ortho hydrogens & 6 meta hydrogens. Ideally this should produce 196 lines, but imperfect resolution reduces the number observed.
Methods of Generating Free Radicals
The homolytic cleavage of covalent bonds produces radicals, and since this is an endothermic process, it requires the introduction of energy from the surroundings. Heat serves this purpose by collisional interconversion of kinetic energy into vibrational energy, and the temperature required for bond homolysis will be proportional to the bond dissociation energy. Absorption of light may also lead to radical species by intra- or intermolecular conversion of the increased electronic energy into vibrational energy. As expected, weaker covalent bonds dissociate into radicals more readily than stronger covalent bonds.
A. Thermal Cracking
At temperatures greater than 500º C, and in the absence of oxygen, mixtures of high molecular weight alkanes break down into smaller alkane and alkene fragments. This cracking process is important in the refining of crude petroleum because of the demand for lower boiling gasoline fractions. Free radicals, produced by homolysis of C–C bonds, are known to be intermediates in these transformations. Studies of model alkanes have shown that highly substituted C–C bonds undergo homolysis more readily than do unbranched alkanes. In practice, catalysts are used to lower effective cracking temperatures.
In contrast to stronger C–C and C–H bonds, the very weak O–O bonds of peroxides are cleaved at relatively low temperatures ( 80 to 150 ºC ). The resulting oxy radicals may then initiate other reactions, or may decompose to carbon radicals.
Organic azo compounds (R–N=N–R) are also heat sensitive, decomposing to alkyl radicals and nitrogen. Azobisisobutyronitrile (AIBN) is the most widely used radical initiator of this kind, decomposing slightly faster than benzoyl peroxide at 70 to 80 ºC. The thermodynamic stability of nitrogen provides an overall driving force for this decomposition, but its favorable rate undoubtedly reflects weaker than normal C-N bonds.
C. Photolytic Bond Homolysis
Compounds having absorption bands in the visible or near ultraviolet spectrum may be electronically excited to such a degree that weak covalent bonds undergo homolysis. Examples include the halogens Cl2, Br2 & I2 (bond dissociation energies are 58, 46 & 36 kcal/mole respectively), alkyl hypochlorites, nitrite esters and ketones. . Ketones undergo n to π* electronic excitation near 300 nm. The resulting excited state is a diradical in which one of the odd electrons is localized on the oxygen atom. Cleavage of an alkyl group may then take place.
D. Electron Transfer
The action of inorganic oxidizing and reducing agents on organic compounds may involve electron transfers that produce radical or radical ionic species. Ferrous ion, for example, catalyzes the decomposition of hydrogen peroxide ( Fenton's reagent ) and organic peroxides. In some cases the radical intermediates formed in this manner are sufficiently stable to be studied in the absence of oxygen.
The alkali metals lithium, sodium and potassium reduce the carbonyl group of ketones to a deep blue radical anion called a "ketyl. A similar reduction of benzene and its derivatives also proceeds by way of radical anion intermediates.
E. Hydrogen and Halogen Atom Abstraction
If free radical reactions are to be useful to organic chemists, methods for transfering the reactivity of the simple radicals generated by the previously described homolysis reactions to specific sites in substrate molecules must be devised. The most direct way of doing this is by an atom abstraction, as shown here.
R–H + X• –––> R• + H–X
Indeed, when X is Cl or Br, this is a key step in the alkane halogenation chain reaction. Hydrogen abstraction reactions of this kind are sensitive to the nature of both the attacking radical ( X•) and the R–H bond. . Thus the rate of reaction of 1º C–H with Cl• is a thousand times faster than with Br•. However, the less reactive bromine atom shows much greater selectivity in discriminating between 1º, 2º and 3º C–H groups.
Certain C–H bonds are so susceptible to radical attack that they react with atmospheric oxygen (a diradical) to form peroxides. Typical groups that exhibit this trait are 3º-alkyl, 2º & 3º-benzyl and alkoxy groups in ethers.
R–H + O2 –––> R• + •O2H –––> R–O–O–H
The exceptional facility with which S–H and Sn–H react with alkyl radicals makes thiophenol and trialkyltin hydrides excellent radical quenching agents, when present in excess. At equimolar or lower concentration they function well as radical transfer agents..
Carbon halogen bonds, especially C–Br and C–I, are weaker than C–H bonds and react with alkyl and stannyl radicals to generate new alkyl radicals. This reaction has been put to practical use in a mild procedure for reducing alkyl halides to alkanes.
ELECTRONEUTRALITY
Nature seems to strongly discourage any process that would lead to an excess of positive or negative charge in matter. Suppose, for example, that we immerse a piece of zinc metal in pure water. A small number of zinc atoms go into solution as Zn ions, leaving their electrons behind in the metal:
Zn(s) → Zn2+ + 2e–
As this process goes on, the electrons which remain in the zinc cause a negative charge to build up within the metal which makes it increasingly difficult for additional positive ions to leave the metallic phase. A similar buildup of positive charge in the liquid phase adds to this inhibition. Very soon, therefore, the process comes to a halt, resulting in a solution in which the concentration of Zn2+ is still too low (around 10–10 M) to be detected by ordinary chemical means.
Transport of zinc ions from the metal to water; the build-up of negative charge in the metal (and positive charge in the solution) soon brings the process to a halt.
There would be no build-up of this opposing charge in the two phases if the excess electrons could be removed from the metal or the positive ions consumed as the electrode reaction proceedes. For example, we could drain off the electrons left behind in the zinc through an external circuit that forms part of a complete electrochemical cell; this we will describe later. Another way to remove these same electrons is to bring a good electron acceptor (that is, an oxidizing agent) into contact with the electrode. A suitable acceptor would be hydrogen ions; this is why acids attack many metals. For the very active metals such as sodium, water itself is a sufficiently good electron acceptor.
The degree of charge unbalance that is allowed produces differences in electric potential of no more than a few volts, and corresponds to unbalances in the concentrations of oppositely charged particles that are not chemically significant. There is nothing mysterious about this prohibition, known as the electroneutrality principle; it is a simple consequence of the thermodynamic work required to separate opposite charges, or to bring like charges into closer contact. The additional work raises the free energy change of the process, making it less spontaneous.
The only way we can get the oxidation of the metal to continue is to couple it with some other process that restores electroneutrality to the two phases. A simple way to accomplish this would be to immerse the zinc in a solution of copper sulfate instead of pure water. As you will recall if you have seen this commonly-performed experiment carried out, the zinc metal quickly becomes covered with a black coating of finely-divided metallic copper. The reaction is a simple oxidation-reduction process, a transfer of two electrons from the zinc to the copper:
Zn(s) → Zn2+ + 2e– Cu2+ + 2e– → Cu(s)
The dissolution of the zinc is no longer inhibited by a buildup of negative charge in the metal, because the excess electrons are removed from the zinc by copper ions that come into contact with it. At the same time, the solution remains electrically neutral, since for each Zn ion introduced to the solution, one Cu ion is removed. The net reaction
Zn(s) + Cu2+ → Zn2+ + Cu(s)
quickly goes to completion.
Potential differences at interfaces
The transition region between two phases consists of a region of charge unbalance known as the electric double layer. As its name implies, this consists of an inner monomolecular layer of adsorbed water molecules and ions, and an outer diffuse region that compensates for any local charge unbalance that gradually merges into the completely random arrangement of the bulk solution. In the case of a metal immersed in pure water, the electron fluid within the metal causes the polar water molecules to adsorb to the surface and orient themselves so as to create two thin planes of positive and negative charge. If the water contains dissolved ions, some of the larger (and more polarizable) anions will loosely bond (chemisorb) to the metal, creating a negative inner layer which is compensated by an excess of cations in the outer layer.
Electrochemistry is the study of reactions in which charged particles (ions or electrons) cross the interface between two phases of matter, typically a metallic phase (the electrode) and a conductive solution, or electrolyte. A process of this kind can always be represented as a chemical reaction and is known generally as an electrode process. Electrode processes (also called electrode reactions) take place within the double layer and produce a slight unbalance in the electric charges of the electrode and the solution. Much of the importance of electrochemistry lies in the ways that these potential differences can be related to the thermodynamics and kinetics of electrode reactions. In particular, manipulation of the interfacial potential difference affords an important way of exerting external control on an electrode reaction.
The interfacial potential differences which develop in electrode-solution systems are limited to only a few volts at most. This may not seem like very much until you consider that this potential difference spans a very small distance. In the case of an electrode immersed in a solution, this distance corresponds to the thin layer of water molecules and ions that attach themselves to the electrode surface, normally only a few atomic diameters. Thus a very small voltage can produce a very large potential gradient. For example, a potential difference of one volt across a typical 10–8 cm interfacial boundary amounts to a potential gradient of 100 million volts per centimeter— a very significant value indeed!
Interfacial potentials are not confined to metallic electrodes immersed in solutions; they can in fact exist between any two phases in contact, even in the absence of chemical reactions. In many forms of matter, they are the result of adsorption or ordered alignment of molecules caused by non-uniform forces in the interfacial region. Thus colloidal particles in aqueous suspensions selectively adsorb a given kind of ion, positive for some colloids, and negative for others. The resulting net electric charge prevents the particles from coming together and coalescing, which they would otherwise tend to do under the influence of ordinary van der Waals attractions.
Interfacial potential differences are not directly observable.The usual way of measuring a potential difference between two points is to bring the two leads of a voltmeter into contact with them. It's simple enough to touch one lead of the meter to a metallic electrode, but there is no way you can connect the other lead to the solution side of the interfacial region without introducing a second electrode with its own interfacial potential, so you would be measuring the sum of two potential differences. Thus single electrode potentials, as they are commonly known, are not directly observable.
What we can observe, and make much use of, are potential differences between pairs of electrodes in electrochemical cells.
Zn(s) → Zn2+ + 2e–
As this process goes on, the electrons which remain in the zinc cause a negative charge to build up within the metal which makes it increasingly difficult for additional positive ions to leave the metallic phase. A similar buildup of positive charge in the liquid phase adds to this inhibition. Very soon, therefore, the process comes to a halt, resulting in a solution in which the concentration of Zn2+ is still too low (around 10–10 M) to be detected by ordinary chemical means.
Transport of zinc ions from the metal to water; the build-up of negative charge in the metal (and positive charge in the solution) soon brings the process to a halt.
There would be no build-up of this opposing charge in the two phases if the excess electrons could be removed from the metal or the positive ions consumed as the electrode reaction proceedes. For example, we could drain off the electrons left behind in the zinc through an external circuit that forms part of a complete electrochemical cell; this we will describe later. Another way to remove these same electrons is to bring a good electron acceptor (that is, an oxidizing agent) into contact with the electrode. A suitable acceptor would be hydrogen ions; this is why acids attack many metals. For the very active metals such as sodium, water itself is a sufficiently good electron acceptor.
The degree of charge unbalance that is allowed produces differences in electric potential of no more than a few volts, and corresponds to unbalances in the concentrations of oppositely charged particles that are not chemically significant. There is nothing mysterious about this prohibition, known as the electroneutrality principle; it is a simple consequence of the thermodynamic work required to separate opposite charges, or to bring like charges into closer contact. The additional work raises the free energy change of the process, making it less spontaneous.
The only way we can get the oxidation of the metal to continue is to couple it with some other process that restores electroneutrality to the two phases. A simple way to accomplish this would be to immerse the zinc in a solution of copper sulfate instead of pure water. As you will recall if you have seen this commonly-performed experiment carried out, the zinc metal quickly becomes covered with a black coating of finely-divided metallic copper. The reaction is a simple oxidation-reduction process, a transfer of two electrons from the zinc to the copper:
Zn(s) → Zn2+ + 2e– Cu2+ + 2e– → Cu(s)
The dissolution of the zinc is no longer inhibited by a buildup of negative charge in the metal, because the excess electrons are removed from the zinc by copper ions that come into contact with it. At the same time, the solution remains electrically neutral, since for each Zn ion introduced to the solution, one Cu ion is removed. The net reaction
Zn(s) + Cu2+ → Zn2+ + Cu(s)
quickly goes to completion.
Potential differences at interfaces
The transition region between two phases consists of a region of charge unbalance known as the electric double layer. As its name implies, this consists of an inner monomolecular layer of adsorbed water molecules and ions, and an outer diffuse region that compensates for any local charge unbalance that gradually merges into the completely random arrangement of the bulk solution. In the case of a metal immersed in pure water, the electron fluid within the metal causes the polar water molecules to adsorb to the surface and orient themselves so as to create two thin planes of positive and negative charge. If the water contains dissolved ions, some of the larger (and more polarizable) anions will loosely bond (chemisorb) to the metal, creating a negative inner layer which is compensated by an excess of cations in the outer layer.
Electrochemistry is the study of reactions in which charged particles (ions or electrons) cross the interface between two phases of matter, typically a metallic phase (the electrode) and a conductive solution, or electrolyte. A process of this kind can always be represented as a chemical reaction and is known generally as an electrode process. Electrode processes (also called electrode reactions) take place within the double layer and produce a slight unbalance in the electric charges of the electrode and the solution. Much of the importance of electrochemistry lies in the ways that these potential differences can be related to the thermodynamics and kinetics of electrode reactions. In particular, manipulation of the interfacial potential difference affords an important way of exerting external control on an electrode reaction.
The interfacial potential differences which develop in electrode-solution systems are limited to only a few volts at most. This may not seem like very much until you consider that this potential difference spans a very small distance. In the case of an electrode immersed in a solution, this distance corresponds to the thin layer of water molecules and ions that attach themselves to the electrode surface, normally only a few atomic diameters. Thus a very small voltage can produce a very large potential gradient. For example, a potential difference of one volt across a typical 10–8 cm interfacial boundary amounts to a potential gradient of 100 million volts per centimeter— a very significant value indeed!
Interfacial potentials are not confined to metallic electrodes immersed in solutions; they can in fact exist between any two phases in contact, even in the absence of chemical reactions. In many forms of matter, they are the result of adsorption or ordered alignment of molecules caused by non-uniform forces in the interfacial region. Thus colloidal particles in aqueous suspensions selectively adsorb a given kind of ion, positive for some colloids, and negative for others. The resulting net electric charge prevents the particles from coming together and coalescing, which they would otherwise tend to do under the influence of ordinary van der Waals attractions.
Interfacial potential differences are not directly observable.The usual way of measuring a potential difference between two points is to bring the two leads of a voltmeter into contact with them. It's simple enough to touch one lead of the meter to a metallic electrode, but there is no way you can connect the other lead to the solution side of the interfacial region without introducing a second electrode with its own interfacial potential, so you would be measuring the sum of two potential differences. Thus single electrode potentials, as they are commonly known, are not directly observable.
What we can observe, and make much use of, are potential differences between pairs of electrodes in electrochemical cells.
- Summary
Electroneutrality principle - Bulk matter cannot have a chemically-significant unbalance of positive and negative ions.
Dissolution of a metal in water can proceed to a measurable extent only if some means is provided for removing the excess negative charge that remains. This can be by electron-acceptor ions in solution, or by drawing electrons out of the metal through an external circuit.
Interfacial potentials - these exist at all phase boundaries. In the case of a metal in contact with an electrolyte solution, the interfacial region consists of an electric double layer.
The potential difference between a metal and the solution is almost entirely located across the very thin double layer, leading to extremely large potential gradients in this region.
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