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(see Graphic)
William Henry Perkin observed in the late 19th century that the wonders in
his Oxford garden were produced by nature at the modest temperatures of an
English summer and without the use of concentrated sulphuric acid. He knew
from his pioneering work on organic dyestuffs that chemical synthesis in
the laboratory and factory was a much harsher affair. But some areas of
modern industrial chemistry are coming to have more in common with the
gentle ways of nature than Perkin ever dreamt possible.

The building blocks in many natural chemical reactions are highly reactive
molecules called radicals which are produced from more stable entities by
the addition or removal of electrons, usually by enzymes. Vital biochemical
processes, including respiration and photosynthesis, depend on this electron
transfer. Addition and removal of electrons is also the basis of
electrochemistry, which uses electric potentials as the driving force behind
chemical reactions.

Electrochemistry was used to make molecules as long ago as the middle of the
19th century. It grew in industrial importance in the 1960s when it was used
to make raw materials for the burgeoning polymer industry. Now the
importance of electrochemistry is growing again, particularly for
carbon-based (organic) chemistry, as chemicals industries shift from making
basic chemicals in enormous quantities to smaller amounts of more
specialised chemicals, and as tougher environmental laws begin to bite.
Electrochemistry is also the basis of new methods being developed for
destroying toxic organic pollutants.

Electrochemical reduction or oxidation has several advantages over
conventional methods of making reactive chemicals. First, it works at low
temperatures. In this respect it is similar to photosynthesis, where light
supplies the energy required to transfer electrons from one molecule to
another and so drive the reaction along. In electrochemistry this energy
comes from an electrical potential. Both types of reaction avoid the need
for thermal energy – heat – which ordinary chemical reactions often need in
order to go at a reasonable speed. These milder conditions mean lower energy
costs and more selective reactions.

Secondly, electrochemistry works without the use of harsh chemical reagents
– only the direct transfer of electrons is involved. The electron is clean,
relatively cheap, need not be stored, and is required in small quantities:
about half a milligram of electrons will do the same job as 23 grams of the
reducing metal sodium. Not that electrons come in bottles. In the chemicals
industry, zinc, tin and sulphur compounds are widely used as reductants, and
nitric oxides and lead and chromium salts are used as oxidants. These are
not environmentally friendly reagents: nitric oxide is an atmospheric
pollutant, lead salts are toxic, and chromium salts can cause dermatitis and
cancer. Controlling and disposing of them are growing problems for the
chemicals industry. Against this background, straightforward electrochemical
reactions, in which electrons are supplied to a chemical reaction by
positive and negative electrodes dipped into the reaction mixture, have
obvious attractions.

Electrochemical cells

In general, an electrochemical reaction involves applying an electrical
potential to a liquid called the electrolyte, into which two electrodes are
immersed; this arrangement forms an ‘electrochemical cell’. A simple example
is the electrolysis of water in which hydrogen is produced by reduction of
protons (H⁺) at the negative electrode or ‘cathode’, and oxygen by
oxidation of hydroxide ions (HO^–) at the positive electrode or ‘anode’. The
net transfer of electrons from the cathode to the anode means that an
electric current passes through the cell.FIG-mg18844502.GIF

In an organic electrochemical reaction, organic molecules are converted into
reactive ‘intermediates’ similar to those involved in normal chemical
reactions. Many of the intermediates have an odd number of electrons; these
odd-electron species are called radicals, or, if they also carry an
electrical charge, radical-ions. These intermediates might combine to form
new bonds, synthesising a larger molecule. Alternatively, the removal or
addition of electrons from stable molecules may generate an intermediate in
which certain bonds are weakened and ultimately break. Both pathways are
important in making new molecules.

As long ago as 1834, Michael Faraday found that electrochemical oxidation
and reduction were not confined to inorganic ions and molecules. When he
used an electrochemical cell to oxidise the organic acetate anion (CH₃COO)
of acetic acid, he detected hydrocarbons in the gaseous products. Chemists
now know that these hydrocarbons were methane (CH₄) and ethane (C₂H₆), and
the key intermediate was the methyl radical. Along the way, the stable
carbon dioxide (CO₂) molecule splits out, so even this first organic
electrosynthesis involved both bond formation and bond breaking. The whole
thing takes place in water at about 20 degreeC.

Ethane is not an impressive target for synthesis by modern organic chemists,
who have their sights set on more complex molecules which are biologically
and industrially important. But the key step is the same: combining radicals
that were formed electrochemically. In the past few years, Hans
Schaumlautfer’s group at the University of Muumlautnster in Germany has used
electrochemical methods to synthesise several natural products, including
brevicomin, the sex attractant of the western pine beetle. More recently
Schafer has used an electrochemical cell to generate radicals which join
together to make fatty acids with chemical formulae like
CF₃(CF₂)₇(CH₂)₉CO₂H. In this molecule, all the hydrogen atoms in a long
section of the hydrocarbon chain [CF₃(CF₂)₇-] are replaced by fluorine, and
so this part of the molecule resembles fluorocarbons such as Teflon, which
are well known for their special physical properties. In addition, other
parts of the chain [-(CH₂)₉-] are wax-like, and the end [-CO₂H] is water
soluble. This should give it some unusual and potentially useful lubricating
and surfactant properties.

In my own group at Queen Mary and Westfield College in London, we have been
concentrating on the reactions of radicals generated electrochemically from
quinones and related starting materials. These radicals, and similar ones
formed by oxidation of phenols (benzene rings with hydroxyl groups
attached), combine to form molecules similar to natural polymers such as
lignin or the related lignans. Podophyllotoxin, a naturally occurring
antitumour agent, is of this type. In nature, enzymes transfer the electrons
in the oxidation reactions.

Electrochemistry can also be used for ‘indirect’ electrolysis, in which
conventional oxidising and reducing chemicals are generated and recycled
electrochemically in a sealed system. For instance, salts of metal chromium
are highly toxic and are powerful oxidants when the chromium is in the high
oxidation state chromium (VI) – the Roman numeral refers to the number of
electrons that the chromium atom has lost. The reaction reduces the chromium
to a lower oxidation state, chromium (III), which in the past was discarded.
As chromium (III) is highly toxic, this poses a pollution problem. In the
1970s, Robert Clarke, an industrial chemist working with the British
dyestuffs firm L. B. Holliday, developed a simple electrochemical routine in
which the chromium salt was recycled within the system
(see Figure 1).FIG-mg18844501.GIF
This
was used successfully by Holliday to manufacture anthraquinone, an
intermediate molecule essential for the production of dyes.

Industrial revival

A similar approach now seems likely to revive the industrially important
Wacker process in which a group of hydrocarbons known as olefins, which are
obtained from the cracking of oil, are oxidised by a compound of the
precious metal palladium to form ketones and aldehydes. In the mid-1970s
this process accounted for 85 per cent of the 650 000 tonnes of acetaldehyde
(CH₃CHO) produced in the US each year. The original process used oxygen as
the oxidant and a copper salt as the catalyst to regenerate the expensive
palladium. But the mixture was corrosive, and the high pressure of oxygen
needed was potentially dangerous, so the Wacker process became obsolete. Now
D. G. Miller and Dan Wayner, working at the National Research Council of
Canada in Ottawa, have devised a way of replacing the copper salt by a
recyclable quinone catalyst and the oxygen by an anode, in an
electrochemical reactor.

A difficult challenge facing industrial chemists is to devise better and
cleaner ways of converting unreactive methane (CH₄) which is usually burnt
as a fuel, to methanol (CH₃OH), a much more valuable chemical. Yet some
bacteria and enzymes cheerfully consume hydrocarbons and at least one of
them usefully catalyses the oxidation of methane. Now Derek Barton and his
colleagues at the Natural Products Institute in Gif-sur-Yvette near Paris,
and later at the University of Texas A&M, College Station, have found an
indirect method of oxidising paraffin hydrocarbons such as methane, one
version of which is electrochemical. When dissolved in the common solvent
pyridine, the hydrocarbons are oxidised when small amounts of iron in
oxidation state II are present, along with an electron source to regenerate
it. This reaction is simple to carry out and is potentially very important
commercially. Not so simple is the sequence of reactions that take place:
the detailed reaction pathway is not yet completely understood, but it
certainly involves formation of a complex iron compound in a high oxidation
state, which becomes active in contact with the hydrocarbon.

The indirect technique has proved useful for destroying molecules as well as
for making them. At high enough temperatures, any organic chemicals will
burn (a process of oxidation) to form water and carbon dioxide, which is why
incineration is used to destroy noxious molecules such as benzene and nerve
gases. But before these molecules have been completely destroyed they may
turn into harmful gases which could escape. Electrochemistry offers an
alternative to incineration which avoids this problem: highly oxidising
metal ions generated electrochemically could destroy the unwanted chemicals
at lower temperatures, without releasing noxious gases.

For instance, last year Joseph Farmer and his co-workers at Lawrence
Livermore National Laboratory in California found that silver (II), the
metal’s highest oxidation state, will completely ‘combust’ common
hydrocarbon components of organic waste to form mainly water and carbon
dioxide. Silver salts are not cheap, but the silver (II) is one of the most
powerfully oxidisers known, and as Farmer discovered, in a closed
electrochemical cell the silver salt can be recycled. The engineering trick
which makes this possible is the use of a rotating cylindrical anode which
stirs the solution so well that the oxidising silver ions and the organic
waste are rapidly mixed at the surface of the anode. This also ensures the
silver (II) is regenerated so fast that relatively small quantities are
sufficient. This technique can be adapted to destroy benzene, a waste
product of the US Department of Energy’s Savannah River Plant which may be
contaminated with radioactive caesium-137, so it cannot be incinerated.
Destroying the benzene electrochemically leaves a relatively small quantity
of radioactive waste.

Destroying toxic waste

An alternative method for destroying organic waste materials and one which
has great potential for disposing of chemical warfare agents such as
organophosphorus nerve gases is being developed by Allen Bard, Marye Anne
Fox and their co-workers at the University of Texas at Austin. This combines
the power of the electrochemical cell with the commonest solvent, water –
though not at normal temperatures and pressures. In the region of 400K (127
degreeC) and at pressures of about 240 atmospheres water becomes
‘supercritical’, a state in which it behaves very differently from ordinary
water. Not only is it corrosive but it will also completely dissolve many
organic compounds, such as benzene, which are normally insoluble in water.
Bard and Fox have developed electrochemical cells which operate with
supercritical water as the electrolyte. The oxygen generated
electrochemically in their sealed cells will completely combust organic
waste materials.

Although the basic technology of electrosynthesis is well established,
variations of the technique are still opening up alternatives to
conventional synthesis. In 1990 a group of French chemists announced the
development of an unusual electrochemical cell for synthesising the
anti-inflammatory drug Fenoprofen, used for treating rheumatism and
arthritis. About $1.5 billion worth of this and related drugs are
sold every year. The conventional chemical synthesis consists of seven
steps, but chemists at the research centre of France’s nationalised
electricity industry in Paris, have perfected a three-step route. This
depends on an electrochemical reaction combined with an ingenious
electrochemical cell which supplies magnesium ions to the solution at
exactly the right rate, so that for every electron which leaves the cathode
one arrives at the anode.

The key to the French team’s ‘pencil sharpener’
(see Figure 3)
(see Figure 3)
is the use of
a magnesium anode, which is oxidised to the magnesium (Mg²⁺) cations that
are consumed in the reaction. The key synthetic step is the formation at the
steel cathode of an organic anion which reacts there with carbon dioxide;
the magnesium cations are needed to drive the reaction. They do this by
stabilising the carboxylate anion as soon as it has formed. This method has
several advantages. First, there is no need for an expensive membrane to
separate the electrolyte in the two halves of a conventional electrochemical
cell. Secondly, a relatively cheap and safe solvent is used. And finally,
the cathode is made of a harmless material, rather than the toxic mercury or
lead which are usually required and which could contaminate materials made
in the cell, which could be drugs or foodstuffs.

Zipping up molecules

During the past year or so electrochemistry has played a key role in the
synthesis and study of conducting organic polymers such as
poly(paraphenylenevinylene)s or PPVs. These materials combine the desirable
property of electrical conductivity with light weight and the ability to be
moulded or extruded. As was recently discovered by the physicist Richard
Friend and organic chemist Andrew Holmes, both of the University of
Cambridge, some PPV compounds also emit light when carrying current and can
be used to make light-emitting diodes (LEDs).

At Queen Mary and Westfield, we are using electrochemistry to synthesise
PPVs. In our reaction, PPV is formed from organic radicals containing eight
carbon atoms, which are formed in a cleavage reaction at the cathode of an
electrochemical cell. The radicals combine by ‘zipping up’ to form long
chains. This is exactly how some natural polymers, such as three-dimensional
lignins, form from naturally occurring phenols, and how they combine with
cellulose to make the woody materials called lignocellulosics which give
mechanical strength to many plants.

Academic and industrial chemists have moved on a long way since Perkin’s
day. Today they are in a good position to exploit the advantages of
nature’s chemistry – without harming the environment. 2:

James Utley is professor of organic chemistry at Queen Mary and Westfield
Collge, University of London and an active researcher in organic
electrochemistry.

* * *

1: Electrochemistry large and small

The importance to the molecular sciences of a theoretical description of
electron transfer processes was recognised by the award of the 1992 Nobel
Prize for Chemistry to Rudy Marcus of the California Institute of
Technology. Forty years ago, Canadian-born Marcus became curious about how
exactly electrons transferred between atoms and molecules. By 1956 he had
produced a relatively simple description which encompassed electron transfer
reactions in such diverse fields as photosynthesis, solar energy conversion
and, not least, electrochemistry. The transfer of an electron is the first
step in any electrochemical process.

The beauty of the Marcus theory turned out to be its counterintuitive
nature. Chemical reactions have a ‘driving force’, the free energy
difference between the reactants and products (δ³Ò). A roller skater
might view this as the drop from the top of a hill to the bottom. However,
for molecules a small hump must be climbed before they can ‘free-wheel’ down
the other side. The height of this hill is the ‘free energy of
activation’, (δ³Ò^#). This quantity, along with temperature (T), controls
the rate of a reaction.FIG-mg18844504.GIF

Electron transfer reactions are no exception. Their rates (k_ET) are governed
by an equation like this:

k_ET α e^{-δ³Ò#/°ì°Õ}

This equation tells us that small δ³Ò^# values and high temperatures
favour faster reactions. k is a constant called Boltzmann’s constant.

Marcus recognised that as well as being influenced by classical
electrostatic forces, molecules and atoms move through a sea of solvent
molecules which are clustered around the reactants. Once reactants have
become products, solvent molecules need to reorganise so that they can
cluster about these newly formed molecules.

This elaborate minuet of molecules also takes place following electron
transfers. The molecules involved here are massive compared to electrons, so
the change is relatively slow and requires ‘reorganisation energy’ which
helps to increase δ³Ò^#. Marcus brought all this into a version of the
equation:

δ³Ò^# = Electrostatic factors + &lgr;/4(1 + δ³Ò^o/&lgr;)²

Here, the ‘reorganisation energy’ is represented by &lgr; δ³Ò^o
is essentially the driving force for the reaction – the measure
of the energy required for two or more atoms or ions to exchange electrons
that is known in electrochemical terminology as the ‘standard redox
potential’. Its value can be measured.

The ‘electrostatic factors’ include the changes in charge on the species
exchanging electrons and their distance apart. The key feature of the Marcus
equation is that its second term is parabolic.

Before Marcus came up with his equation, molecular scientists had imagined
that rates of reactions would increase continuously, and exponentially, with
the ‘driving force’, rather as the roller skater expects only to roll
downhill and to roll faster, the steeper the slope. A parabolic plot of rate
versus δ³Ò^o predicts that at some point the rates (for
example, of electron transfer) would decrease as δ³Ò^o
increased.

Final proof of the theory came when physical chemists were able to observe
experimentally the parabolic nature of the relationship. One of the key
experiments was published in 1984 by John Miller and his colleagues at the
Argonne National Laboratory in the US. They measured the rates of electron
transfer between various organic ‘functional groups’ of atoms and one fixed
group held apart by a rigid molecular framework or ‘spacer’. Then they used
standard electrochemical methods to measure differences in redox potential
– that is, the free energy difference, the driving force for the reaction.
The graph that shows how these two quantities are related
(shown here on the left)
(shown here on the left)
is a parabola.

2: Electrochemistry large and small

In the early 1960s, the US chemicals company Monsanto introduced a very
large scale electrosynthetic process for the manufacture of adiponitrile,
one of the starting materials for the manufacture of nylon. Today Monsanto,
the German company BASF and Asahi in Japan produce several hundred thousand
tonnes of adiponitrile each year by this process.

Monsanto’s industrial success triggered much fundamental and applied
research. In particular, research by Lennart Eberson at Lund in Sweden,
Henning Lund at Aarhus in Denmark, Jean-Michel Saveant in Paris and Allen
Bard at the University of Texas at Austin in the US, among others, has led
to a detailed understanding of the different electrochemical pathways
molecules might take. The result has been numerous new applications of
electrosynthesis. Many of these research results have in turn had a profound
impact on the study of organic reaction mechanisms, especially those which
have radicals and radical-ions as intermediates. This month an EC-funded
research network was set up to link laboratories investigating the role of
electron transfer in activating organic molecules. The network involves
researchers in Aarhus, Belfast, Bonn, Braga, Copenhagen, London, Munster,
Paris and Southampton.

Some of the most effective developments in industrial-scale electrochemistry
have taken place in the developing world. India, for example, operates more
electrochemical processes than any other country, as it cannot afford
expensive chemical catalysts but has plenty of electricity. The Indian
processes are mainly small and medium-scale operations, which produce
organic compounds for India’s own dyestuffs and pharmaceuticals industries.