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FACS Newsletter 1/2000
Electron Transfer Reactions in Chemistry and Biology:
History and Future
Dr. R.A. Marcus
the Professor of California Institute of Technologyˇˇ
The field of electron transfer reactions has grown
considerably from the late 1940s and early 1950s. During that period the rates
of many of the simplest type of electron transfer reactions, isotopic exchange electron
transfers, were measured using radioactive isotopes, and the results provided
invaluable insight. Such studies led Libby to suggest the importance of the
Franck-Condon principle in controlling the rates of the electron transfers (a
suggestion which he attributed to James Franck). Several years later Dr.
Marcus showed that to implement Libby's suggestion and still satisfy conservation of energy during the transfer, it
was necessary for the solvent and any vibrational coordinates to "reorganize" prior to the
transfer. The reorganization amounts to thermal fluctuations which permit the system to
reach the intersection of two potential energy surfaces, that of the reactants and
environment and that of the products and environment. The Franck-Condon principle is
satisfied only at the intersection.
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The electron transfer field subsequently expanded considerably into diverse areas
such as inorganic reactions in solution, reactions across organic bridged systems, reactions
at interfaces (metal-liquid, semiconductor-liquid, liquid-liquid, liquid-polymer-metal, and
other modified electrodes), photoinduced electron transfers (solar energy conversion,
photosynthesis, charge transfer spectra of donor-acceptor systems), biological electron
transfers (such as in photosynthesis), long range electron transfer, and solvent dynamics of
charge transfer systems. Some of the ideas have been extended to electron transfers
accompanied by bond dissociation, as well as to various atom or group (e.g., methyl)
transfer reactions which do not involve electron transfer, and to SN2 reactions of the ET
type. Theory, which has helped link these fields, has led to a variety of predictions and
tests.
Examples of the expansion of the electron transfer field into different areas are given
in Figure 1, below.
The results there can usually be interpreted with an equation for the rate constant ket of the form in eq. 1, sometimes modified by nuclear tunneling, diffusion, or solvent dynamics effects, or by adaptation to the specific problem:
(1)
where A depends on the process considered (e.g., whether it occurs in the bulk phase or at
an interface), ¦¤Go is the standard free energy of reaction or its equivalent (e.g.,
"activation overpotential") for the transfer step, and l
is a "reorganization energy", expressible in terms of changes of bond lengths and the properties of the solvent (or other
environment) for the transfer step. One consequence of eq. 1 is the counter-intuitive so
called "inverted effect", whereby when the reaction is extremely downhill, i.e., has a very
negative¦¤Go (¦¤Go > l), the reaction actually goes more slowly the more downhill the
reaction is. For atom or group transfer reactions eq. 1 becomes for 0.2 <
(-Go/l) < 0.8 an
approximation to a different equation, which Dr. Marcus derived in 1968, one which has,
for atom or group transfers, no inverted effect. This difference is also understandable on
physical grounds.
The main features of electron transfer theory and examples of the interaction between
theory and experiment on several of these topics will be illustrated in this lecture. It is
interesting, for example, to see how Nature's photosynthetic reaction center is constructed
so as to yield a highly efficient separation of charge across a membrane, in a way
consistent with chemical and physical concepts inferred from simpler systems. Dr.
Marcus also described several other new directions in the field.
A few general references to the theory and experiments were given in his talk.
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