Pletcher et Fleichmann Organic electrochmeistry 1969

Prévia do material em texto

ORGANIC E LECTRQC H EM ISTRY 
M. Fleischmann, B.Sc., Ph.D., A.R.C.S., and D. Pletcher, BSc., Ph.D., A.R.I.C. 
Chemistry Department, The University, Southampton SO9 5N H 
The special characteristics of electrode processes . . . . . . 
Reactive intermediates in electrode reactions . . .. . . . . 
Methods of investigation, 90 
Uncharged radicals, 95 
Carbanions and carbanion radicals, 97 
Carbonium ions and carbonium ion radicals, 98 
Other intermediates, 100 
Correlations between electrochemical data and other physical pro- 
perties . . . . .. . . . . . . . . . . . . 
Some further examples of electrochemical reactions of organic 
compounds . . .. . . . . . . . . . . . . 
The control of electrode reactions, 106 
Oxidations, 109 
Reductions, 1 12 
References . . . . . . . . .. . . .. . . .. 
85 
94 
103 
106 
114 
At the end of the last and the beginning of this century there was considerable 
interest in the possibility of synthesizing organic compounds by electro- 
chemical methods. Much of this work was summarized by Fichter.1 In some 
respects this is a depressing summary, as the early hopes of achieving selective 
reactions by electrochemical methods were clearly not realized, mainly 
because the electrolyses were carried out under conditions where the electrode 
potential, pH etc. were not controlled. 
Since the middle 1920s many of the advances in organic electrochemistry 
have been made by polarography, a technique in which a dropping mercury 
electrode is used to provide a highly reproducible surface in the construction 
of currentlpotential curves. However, polarographers have generally been 
reluctant to prepare and isolate the products of the electrode reaction, a 
factor which contributed to failure in communication between electro- 
chemistry and organic chemistry. 
In recent years, the forecast of a falling cost of electricity due to the advent 
of nuclear power and a realization that large scale industrial electrosynthesis 
is a real possibility have led to a resurgence of interest in organic electro- 
chemistry and research has been directed to the development of new selective 
routes and the improvement of existing electrosynthetic methods. At the 
same time it has become feasible to study the electrochemical oxidation and 
reduction of organic compounds in greater detail, partly because of the 
development of electronic control systems known as potentiostats, partly 
Fleischmann and Pletcher 87 
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because of the recognition and characterization of reaction intermediates and 
partly because of the considerable development in the understanding of 
electrode kinetics which has followed the development of numerous new 
techniques for studying electrode reactions. In fact, it is'now possible to 
study complex reaction sequences in aetail. 
The first section of this review (p. 88) deals briefly with the special 
characteristics of electrode reactions and with the most common methods of 
investigation, while the second (p. 94) illustrates the generation, nature and 
manner of reaction of intermediates of electrode reactions. The third (p. 103) 
and fourth (p. 106) sections deal with the correlation of electrochemical data 
with other physical properties and with some further examples of electro- 
chemical reduction, oxidation and substitution reactions which illustrate 
features of electrosynthesis additional to those described in the second section. 
For obvious reasons, this review is intended to be illustrative rather than 
comprehensive, and for fuller expositions of various aspects of these subjects, 
the reader is referred to reviews of electrode mechanisms72J organic polaro- 
graphy,4-6 synthesis,7-llJl reductions712J3 oxidations,l4-16 substitntion,l7 
and the Kolbe synthesis.18-20 Text-books on relevant topics of electrode 
processes include references 22-27. 
THE SPECIAL CHARACTERISTICS OF ELECTRODE PROCESSES 
It is well known that the standard free energy change of a reaction such as 
Rd + 0, - o x 
O x + H, - Rd 
is given by 
AGO = -zFEO 
where z is the number of electrons transferred in the process. If we refer 
immediately to a scale of free energies or electrode potentials (Fig. 1) it is 
evident that spontaneous reactions using oxygen or air as the oxidant, or 
hydrogen as a reducing agent, are only possible (in electrochemical terms) 
within the potential range limited by the reduction of oxygen and the oxidation 
of hydrogen. This driving force therefore amounts to roughly 0.5 eV or 
42 kJ mol-1. By contrast, it is possible to carry out electrochemical reactions 
between +3.5 V and -2.5 V, even in aqueous solutions, if suitable electro- 
lytes are chosen (e.g. perchlorates for oxidations and quaternary ammonium 
salts for reductions). The large driving force for electrochemical processes is 
therefore of the order of 3 eV or 250 kJ mol-1. 
Strictly, it is not valid to compare the available driving force to the standard 
free energy change of any chosen reaction since most processes are irrever- 
sible; we must therefore estimate the rate constants instead. At the surface 
between the metal and an electrolyte solution, there is a difference in electrical 
potential, 4, and in transferring a charge zFmol-1 a maximum amount of 
work ( Y Z ~ F can be done in reaching the activated complex, where a(O< (Y < 1) 
is the transfer coefficient. The rate of an electrochemical reaction is therefore 
88 R.I.C. Reviews 
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O 3 + 3 H 2 0 + 6 e - - 6OH- 
ClOh + e - -ClO, 
Ag2++ e- - Ag+ 
Z 1 0 - + H 2 0 + 2 e - - C I - + 2 0 H - 
0 , + 2 H 2 0 + 4 e - - 4 0 H - 
2 H + + 2 e - - H, 
Zn2++ 2e- - Z n 
Na++ e- - N a 
Solvated e lec t rons 
- + 4.0 
- + 3.0 
- + 2.0 
carbon i u m 
I ion 
I - e S 
-H+ I 
- + I .o radical ca t ion - radical 
1 
I - e - 
- ov subs t ra te 1 - e - 
+H + 
t 
-- radical an ion - radical 
-- 2.0 1 +e+ 
-- 3.0 
- - 4.0 
+ 
carbanion 
Fig. I . Representative electrode potentials 
Therefore, current 
azF 
JRT or log I = const. + -.$ 
at constant concentrations, where the product determines the slow step, vi is 
the order of the reaction with respect to species i which is present at a con- 
centration ci (for more exact expressions of the effect of the potential difference 
on the rates of electrode processes see reference 26). Equation 2 shows that 
we can modify the energy of activation, Ex, of electrochemical reactions to 
such an extent that they will take place at room temperature even if Ex is 
large. 
Fleischmann and Pletcher 89 
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Pot en t ios t a t F a - 
1 
Counter 
electrode 
Reference 
Potentiometers 
Osci I I bsco pe 
---I- t 
Osci I lator 
- - - 
Fig. 2. Circuit diagram for a potentiostat 
We can summarize this first point in the following way: 
reactions enable one to introduce a considerable amount 
t 
/ 
m 
m 
Y W 
I 
--L- 
Pulse profiles 
electrochemical 
of energy intomolecules at low temperatures. The order of magnitude of this energy is in 
fact comparable to that of the strength of chemical bonds and this explains 
why many of the chemical changes which are observed can take place. It 
is not surprising that many 'high energy' chemicals which are used as oxidants 
or reductants in synthesis are made electrochemically and it is clearly of some 
interest to avoid such intermediate steps in synthesis. 
Methods of investigation 
Many methods of investigating electrode reactions which have been developed 
during the past decade are based on the regulation of the potential of the 
working electrode with respect to an unpolarized reference electrode using a 
feedback amplifier (Fig. 2), the current being passed between the working 
and a subsidiary electrode. Such feedback amplifiers are known as potentio- 
stats. The potential to be applied between the working and reference elec- 
trodes is set at the second input of the amplifier. A succession of constant 
potentials may be chosen to construct a current/potential curve ; alternatively, 
a slow linear sweep of potential with time will give this curve (potentio- 
dynamic method). When a dropping mercury electrode is used, the current/ 
potential curve is known as a polarogram. 
Figure 2 shows that other potential/time profiles may be imposed on the 
working electrode. In the case of a triangular potential/time profile with a 
high sweep rate [0.1-1000 V s-11, the current/potential plot may be recorded 
oscillographically. This technique is known as cyclic voltammetry and is 
discussed later. A sine wave of small amplitude applied to the input gives 
Lissajou figures on the oscilloscope from which the electrode impedance may 
be calculated; alternatively the impedance may be found using a bridge. 
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I .5 
Fig. 3. Polarogram of anthracene in acetonitrile containing 0. I M Et,NCIO, 
3.0 
When a potential step is applied to the input, a current/time transient is 
obtained; Fig. 2 shows that square pulses may be also applied. 
Electrode reactions may also be studied by applying a constant current and 
measuring the potential/time curve or the potential in the steady state. While 
this method requires only simple apparatus, controlled potential methods 
have been increasingly used in the last few years. The advantages as com- 
pared to constant current techniques will be apparent, since the rates of 
electrode processes are controlled by the potential, as may be seen from 
equations 2 and 3. In addition, in preparative work a measure of selectivity 
in the reactions may be achieved (equation 1). 
Equation 3 shows that the current would increase indefinitely with potential; 
in practice, the diffusion of the electroactive species to the electrode, and 
possibly also of the products away from the electrode, will become rate 
determining and the current reaches a limiting value. A typical current/ 
potential curve obtained with a dropping mercury electrode, that is by 
polarography, is shown in Fig. 3. The portion AC illustrates the addition of 
one electron to an aromatic hydrocarbon and in the region BC the electrode 
process is diffusion controlled. The potential at which the current is half the 
diffusion current is known as the half-wave potential, E,. Polarographic 
waves are classified as being reversible or irreversible. Jn the first case, the 
rate of the electrode reaction is fast compared with diffusion so that the shape 
of the wave is controlled by mass transfer of the reagent and product and the 
equilibrium at the surface 
O x + ze-- Rd 
RT aox Therefore E E, + ~ In ~ 
zF a R d ( 5 ) 
91 Fleischmann and Pletcher 
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where aox and a R d are the activities of the oxidized and reduced forms at the 
electrode surface. It will be seen that in this case E, h EO. In the second case, 
the rate of the electrode process is slow compared with diffusion and E, is 
no longer equal to EO. However, with increasing potential, the current again 
becomes diffusion controlled. The decision as to whether the electrode 
process is reversible or irreversible may be made by considering the shape of' 
the wave. In the second case the wave is drawn out along the potential axis 
compared to a reversible wake (for exact criteria, see reference 25). In the 
example shown, and for many other hydrocarbons, the wave AC is reversible 
but that for the addition of a second electron CD is irreversible. 
In electrode reactions, it is frequently found that a chemical process 
precedes electron transfer and only one species is electroactive in a potential 
region. For example,28 
CH,(OH), HCHO + H,O ( a ) 
HCHO + 2e- + 2HS - CH,OH ( b ) 
A limiting current (which is controlled by the rate of the preceding reaction) 
may therefore be realized when b becomes fast compared to a. Such waves 
are known as kinetic waves and their distinction from diffusion controlled 
waves is described el~ewhere.~5 Other reaction sequences can also be investi- 
gated (e.g. following reaction^).^^ 
The steady-state current/potential data can further be analysed by plotting 
the initial part of the region AB in Fig. 3, according to equation 4. The linear 
relation between log I and + is known as a Tafel plot. The variation in 
position of these plots with the concentration of electroactive species gives 
information directly about the slow step of the electrode reaction. 
For measurements under non-steady state conditions, such as in cyclic 
voltammetry, all the information which has been described above may again 
be obtained23925 and further data of the reaction sequence may be derived. 
In the first place, the rate of diffusion in the non-steady state is greater than 
for steady polarization and the current can therefore reach higher values, as 
in section AB in Fig. 4. (This increase in current allows the study of the 
kinetics of faster electrode reactions than is possible by steady-state methods.) 
Non-steady state diffusion then becomes rate controlling and the current 
falls (section BC). On reversal of the potential sweep, the starting material is 
regenerated in the case of simple reactions and a peak, DEF, is observed in 
the oxidation current. Secondly, for a sequence such as 
R + e- - R - 
R ' - + X f A RX' - 3 (RX) 2 
some of the intermediate is removed by the chemical reaction and the reoxida- 
tion peak for 
R'-- R + e - 
is therefore diminished.30 In some cases peaks for the oxidation of RX' and 
(RX)2 may also be observed. In general, the shape of the peaks, and in 
particular their height and separation and their dependence on the sweep 
rate, is directly related to the overall kinetics of the electrode reaction. It is 
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B 
E 
Fig. 4. Cyclic voltammogram of anthracene in acetonitrileZ9 containing 0. I M Et,NCIO,, 
sweep rate 0. I5 V s-l 
of considerable importance that in this way the kinetics of fast reactions 
succeeding a slow electron transfer may be measured. 
Other pulse techniques give essentially similar informationand the choice 
of method will be governed largely by convenience although the interpretation 
of potential step experiments will usually require fewer assumptions than the 
interpretation of cyclic voltammograms and may therefore be preferable. 
Thus, in a single-step experiment, the initial part of the current/time transient 
gives the rate of the electron-transfer step and the whole transient the rate of 
preceding steps.27 Using a double step, say first to the cathodic and then to 
the anodic direction, the kinetics of a following reaction can again be deter- 
mined. For example, the rate constant, k, for the rearrangement of hydrazo- 
benzene to benzidine can be determined in this way.31 
Further information about the kinetics of electrode processes may be 
obtained using various translating electrodes which increase the rate of 
mass transfer. The most widely used of these is the rotating disc electr0de.3~ 
Information concerning chemical reactions succeeding electron transfer may 
be obtained using the ring-disc electrode.33 This consists of a disc electrode 
closely encircled by a ring electrode. The intermediate is generated at the 
disc and the fraction of the intermediate which is removed at the ring is 
related to the rate of removal of intermediate by chemical reaction during its 
transit from the disc to the ring. 
The illustrations which have been given have assumed that the inter- 
mediates are not adsorbed and that the coupled reactions take place in 
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solution. Frequently, this will not be the case but it is possible to investigate 
purely heterogeneous steps in precisely similar ways and indeed simultaneous 
heterogeneous and homogeneous pathways.34 
It will be apparent that very detailed information on the reaction mecha- 
nisms (i.e. the rate of preceding steps, the rate of electron transfer, the rate of 
succeeding reactions and data concerning adsorption) can now be obtained 
by electrochemical methods. As yet, such investigations have been carried 
out only on a very small proportion of known electrode reactions and although, 
where this information is available, the intermediates are fully consistent with 
those postulated in the next section and are in line with those postulated in 
physical organic chemistry, in other cases the existence of intermediates must 
be deduced from non-kinetic data such as the nature of products, esr spectra 
etc. 
REACTIVE INTERMEDIATES IN ELECTRODE REACTIONS 
Early workers in the field of mechanistic organic electrochemistry believed 
that electrochemical reductions took place via the production of ‘active 
hydrogen’ at the cathode and that the potential at which an organic com- 
pound was reduced was dependent on the potential energy required by the 
‘active hydrogen’ in order to react with the organic compound. Similarly, 
electrochemical oxidations were believed to proceed via the production of 
‘active oxygen or hydroxyl radicals’ at the anode. In recent years it has 
become clear that the intermediates produced during an electrode reaction 
are frequently the same as those commonly encountered in other fields of 
organic chemistry-carbanions, carbanion radicals, radicals, carbonium ion 
radicals, carbonium ions and, less commonly, intermediates such as biradicals, 
solvated electrons, transition-metal ions, surface oxide layers and inorganic 
radicals. 
Three factors have been mainly responsible for the recognition of these 
intermediate species. Firstly, the growing use of aprotic solvents, such as 
acetonitrile, dimethyl formamide, dimethyl sulphoxide, and propylene 
carbonate, in which the lifetime of the intermediates is much longer than in 
aqueous solutions ; secondly, the development of improved analytical 
techniques which allow the identification of minor as well as major products 
and, lastly, the development of a number of techniques, both electrochemical 
and non-electrochemical, for the study of transient species. 
Of the electrochemical techniques, cyclic voltammetry has been the most 
widely exploited although a number of linear sweep and pulse techniques for 
the study of short-lived species have been described. Electron spin resonance 
is by far the most common non-electrochemical technique since it allows the 
detection and identification of very dilute solutions of free radicals.35 Short- 
lived radicals have been studied using electrochemical cells in the resonance 
~avi ty .~6 The use of ir, visible and uv spectroscopy has been facilitated by the 
development of optically transparent electrodes37 which allow the electro- 
chemical cell to be placed directly in the light path. This technique permits 
the study of species close to the electrode while actual surface layers may be 
studied by total reflectance spectroscopy.38 Electrochemiluminescence39 is 
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another spectroscopic technique which has been used for the study of anion 
and cation radicals. The species Rf and R- are produced together, either by 
the use of alternating current or by placing two electrodes in close proximity, 
and are allowed to react to produce excited molecules, R*, which can then 
emit light which is characteristic of the system. 
In this section some of the electrochemical methods of producing inter- 
mediates are described and some illustrations of their reactions are given. 
It is important to realize that the behaviour of the intermediates will be 
different when they are adsorbed on the surface of the electrode from when 
they are free in solution; this is one of the factors which causes the electrode 
material to be important in controlling the products of an electrode reaction. 
Thus the Kolbe reaction 
p R - R 
RCOOH - RCOO' + e- + H+ - CO, + R' 
R++ e--- product 
has been shown to give widely differing products on platinumls and 
carbon40-42 electrodes. Moreover, although the Kolbe synthesis is believed to 
proceed via radical intermediates, no esr spectra have been obtained for 
them, presumably because the radicals are strongly adsorbed on the electrode. 
Uncharged radicals 
Radicals are intermediates in many electrode reactions, both oxidations and 
reductions, and a number of examples are cited as equations below. (R = 
alkyl, 4 = phenyl.) 
RCOOH - R'+ CO, + H+ + e- 
A l (CH,), - CH; + Al(CH,), + e- 43 
OH 0' 
0 0' 
45 
0 OH 
-CH (COOEt) , - CH(COOEt),+ e - 46 
+ 2 C = 0 + H++ e - - q5,C'OH 47 
RI + e----- R'+ I - 49 
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radical + RH RH + alkene 
I I 
I I 
R - C - C ' metal al kyl 
\ 
Fig. 5. Main modes of decomposition of uncharged radicals 
With the exception of a few radicals such as the diphenylpicrylhydrazyl 
radical, these intermediates are very reactive species and can react in a number 
of ways. The main modes of disappearance are summarized in Fig. 5. The 
most common mode of disappearance is dimerization as illustrated by the 
normal Kolbe synthesis, the formation of pinacol from the reduction of 
acetone47 and the formation of diphenyl from the reductionof tetraphenyl- 
ammonium ions.48 However, under suitable conditions, high yields of pro- 
ducts other than the dimer can be obtained, for example, 
4RMgl 4e- + 4Mg2++ 41-+ PbR, 50 
RCOOH .-- e- + H+ + CO, + R' C H 2 = C H - c H = C H 2 = _ 
RCH,-CCH=CH-CH;- dimer 51 
although by-products formed by competing reactions are almost always 
observed. At the same time, it is important to realize that the products and 
reaction routes will depend on the concentration of radicals present, i.e. will 
be dependent on the electrode potential, concentration of electroactive 
species and other electrolysis parameters. For example, a high concentration 
of radicals will favour dimerization. 
Other typical radical processes which have been demonstrated to occur 
during electrode reactions are the initiation of p~lymerization~~ and chain 
reactions. An example of a chain reaction is the reduction of benzyl iodide on 
mercury in aqueous ethanol53 where several times the coulombic yield of the 
product, benzylmercuric iodide, was obtained. The direct chemical reaction 
is very slow. Thus the mechanism postulated is 
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Carbanions and carbanion radicals 
In aprotic solvents, aromatic hydrocarbons may be shown (e.g. by polaro- 
graphy or cyclic voltammetry) to reduce in two one-electron steps to the 
carbanion radical and di~arbanion.5~ In the absence of proton donors or 
other addition agents, the radical anions are relatively stable. On the time 
scale of an electrochemical experiment, which is determined by the transition 
of the species through the boundary layer in contact with the electrode 
(approx. 0.1 s), the process may be reversed by changing the potential and 
the parent hydrocarbon will be regenerated.29 On a longer time scale, dimeri- 
zation or disproportionation will take place; the dianions formed will 
abstract protons relatively easily even from aprotic solvents so that dihydro- 
aromatic or dihydrodiaromatic compounds are formed at the potential at 
which the first electron is added. Thus anthracene gives 9, lo-dihydroanthra- 
cene55 and phenanthrene gives 9,9', lO,lO'-dihydrodiphenanthrene.56 The 
dicarbanion produced directly at more negative potentials, naturally, also 
protonates readily; in this case, reversal of the polarization in cyclic volt- 
ammetry can lead to the radical 4H*57 so that the protonation clearly takes 
place in two steps. 
+-- +-a +H - +H - +H ' 
I-.- 
+H- solvent- +H2 
+2- So'Vent 
t +- +- +e- 
+ + + 2 - 
When a proton donor [e.g. water or benzoic acid154 is added to the aprotic 
media, the first reduction wave increases at the expense of the second wave 
until the two waves merge to form a single, two-electron wave. The reduction 
scheme 
H'+e- - +H- 
#H-+H+- w 2 
has been postulated, the single wave being obtained because the species 
4H- has a higher electron affinity (i.e. it is more readily reduced) than the 
parent hydrocarbon. A similar scheme may be written for the reduction of the 
hydrocarbon in the presence of carbon dioxide55 
Fleischmann and Pletcher 97 
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+ + e--- +-- 
+ .- + CO, - $COO - 
+ 'COO- + e- - +-coo- 
+ -coo-+co, - + ( c o o - ) 2 
and when naphthalene is reduced in the presence of carbon dioxide, the 
product is 1,4-dihydro- 1,4-dicarboxynaphthalene. It is also possible to react 
these aromatic hydrocarbon carbanions with methyl iodide to give the 
dime thy1 di hydroderivative .55 
The carbon dioxide addition is a general reaction of carbanions. For 
example, when stilbene or benzyl chloride is reduced in the presence of 
carbon dioxide the products are 1,2-diphenylsuccinic acid55 and phenylacetic 
acid5* respectively. 
Aliphatic carbaiiions are also common intermediates and they usually occur 
at more negative potentials in reactions producing radicals. For example, 
RCI + 2e--- R-+CI- 49 
+CH = CH++ 2e-- +C-H -C-H+ 55 
CH CH3\ 
\C = + 2e- + H+ -- C--OH 59 
CH,' CH 3' 
and in most cases the final products arise by proton abstraction from the 
solvent. Thus, the reduction of ethyl chloride and acetone at very negative 
potentials gives ethane and isopropanol respectively. 
Compounds, such as acrylonitrile, which contain an activated alkenic 
group60 may be reduced to form a dianion 
CH, = CHCN + 2e-- C-H, - C-HCN 
CH, = CHCN + C-H,-C-HCN + 2Hf- NCCH2CH2CH2CH2CN 
which will react with another molecule of acrylonitrile to form the hydrodimer, 
adiponitrile. This reaction is the basis for a new industrial plant. 1,3-Buta- 
diene61 and c+unsaturated acids62 are other examples of the many alkenic 
compounds which reduce in the same way and it is also possible to form 
crossed hydrodimers from mixtures of activated alkenes.63 
Carboniunz ions and carbonium ion radicals 
It might be expected that in an aprotic solvent, aromatic hydrocarbons would 
also oxidize in two one-electron steps to give a carbonium ion radical and a 
carbonium ion 
+H - $H'++ e - 
+H'+ - +++ H++ e- 
or +H.+- + H ~ + + e - 
98 R.I.C. Reviews 
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However, the cation species are much more unstable than the corresponding 
anion species and this simple behaviour is found only for relatively few 
aromatic hydrocarbons such as 9, IO-diphenylanthra~ene~~.~~ and 1,8- 
dithionaphthalene.65 The behaviour found for most hydrocarbons consists 
of an ece mechanism, i.e. the first electron transfer is followed by a very fast 
chemical reaction, possibly dimerization or reaction between the cation 
radical and the solvent, to give a species which is oxidized further. The first 
electron transfer has been shown to be reversible by using very rapid cyclic 
techniques where the rate of potential change is fast compared with the 
following chemical reaction.66 The products which have been reported for 
the oxidation of hydrocarbons such as anthracene are many and varied and 
perhaps this is not surprising in view of the reactivity of the cation inter- 
mediates. However, in some cases such as hexamethylbenzene oxidized in 
acetonitrile good yields of a substituted amide have been reported67 
+ 
C H 3 ) 3 C H 3 CH 3 CH3@H3 CH, 
+ H+ + 2e- 
\ 
CH3 CH3 CH3 CH3 
CH,CN 
CH, CH, 3 3 4 . 3 CH3 \\ 
CH 3 
CH CO. NH.CH vlH3 - Hydrolysis CH3C+= N -CH2 
CH3 CH3 CH3 CH, 
Moreover, there now seems general agreement that the electrochemical 
substitution of aromatic species takes place via carbonium ion intermediates.68 
Thus, whether the reaction studied is halogenation, cyanation, methoxylation 
or acetoxylation the mechanism seems to be 
#I++ H++ 2e- x- +x +H - 
and not a mechanism involving formation of radicals followed by radical 
substitution, since substitution only takes place at potentials where the 
hydrocarbon is oxidized, even when the substituting anion is oxidized to 
radicals at lower potential^.^^ Furthermore, the distribution of product 
isomers coincides more closely with a reaction between a carbonium ion and 
an anion rather than attack on an aromatic system by a radical.70 
Aliphatic carbonium ions may be prepared by the oxidation of carboxylicacids,71 alkyl iodides,72 aliphatic hydrocarbons73 and primary amines.74 
RCOOH --- R + + C 0 2 + 2 e - + H + 
CH,I - C H ~ + + I+ + 2e- 
CH2 = CH- CH2R -CH2 =CH - CfHR + H++ 2e- 
RCH2NH2 - RCH2N'+H2 + e - - RC+H2 +NH2 
Fleischmann and Pletcher 99 
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ROCOCH ROMe 
-H+ CH CN H O alkene - R + 3 CH c += N RL CH ,CON H R 
R 4 RX 
Fig. 6. Probable reactions of aliphatic carbonium ions in common solvents 
A summary of the probable reactions of aliphatic carbonium ions in 
common solvents is shown in Fig. 6. 
Carbonium ions formed in electrode reactions show the usual carbonium 
ion rearrangement. Thus, when neopentyl iodide is oxidized in acetonitrile, 
N-tert-pentyl acetamide can be isolated as well as N-neopentyl acetamide. 72 
CH, 
+ 
CH,-C=N 
l cHlcN 
CH3- -C=N -CH,-CMe3 
CH3- C-CH, 
I 
C2H5 
It is also interesting to note that non-classical carbonium ions can be 
prepared electrochemically. For example, the anodic oxidation of exo- 
norbornene-2-carboxylic acid gives the non-classical carbonium ion.75 
Other intermediates 
Two interesting biradicals which have been prepared electrochemically are 
benzyne76 and dichl~rocarbene.~~ Since both are extremely reactive, their 
preparation is inferred by products from reactions with trapping agents. 
The small yield of product may be due to low efficiency in the preparation of 
100 R. I.C. Reviews 
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the biradicals but is also likely to be due to competing trapping reactions, e.g. 
reaction with the solvent. The biradicals were prepared by reduction in 
acetonitrile, dichlorocarbene from carbon tetrachloride and benzyne from 
o-dibromobenzene. The trapping agents were tetramethylethylene and furan 
respectively. 
/Me - ‘c-c Me2C=CMe2 Me CC 14 + 2e-- 2C I- + CC 1; 
Me/ ‘c’ ‘Me 
There is a number of oxidations where inorganic radicals have been 
postulated as intermediates. Thus, in acetonitrile, the anodic limit has been 
found to be strongly dependent on the inert e le~t ro ly te ,~~ which suggests 
that the initial reaction is oxidation of the anion present. The following 
mechanism, with perchlorate as the inert electrolyte, has been proposed :78 
ClO, - C l O i + e- 
C104+CH3CN - CH2CN+ HC104 
Similarly, the methoxylation of dimethyl formamide in a nitrate base 
electrolyte is believed to involve the nitrate radical :79 
N O - NO; + e- 
/ CH3 /CH3 
NO; + HCON - HCON + HNO3 
‘CH3 ‘CHj 
Clearly, the intermediates in the oxidation and reduction of substituted 
hydrocarbons will not always be carbanions and carbonium ions. For 
example, the intermediates from nitrogen containing organic compounds 
will frequently have the charge situated on the nitrogen. However, their 
reactions and properties do not differ greatly from the carbon species. 
Examples of such charged nitrogen species are shown below. 
The solvated electron is a probable intermediate in the electrochemical 
reduction of ben~ene.8~98~ It is only possible to reduce benzene in solvents 
such as ammonia, hexamethylphosphoramide, ethylenediamine and amines, 
Fleischmann and Pletcher 
8 
101 
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R 3 N - e - + R 3 N + s R3N+H 80 
CH,CONHR - e-+ CH,CON+HR SOIVenf_ CH,CON+H,R 81 
+ e-+ H+ 82 
NH - NH NH2 NH 
in which solvated electrons may be formed. Therefore, it seems likely that in 
the presence of Group I metal ions the reduction mechanism is 
w 
Li ( s h ) + e- - Li+(sol.) + e-(soh.) - products 
Metal ions in lower or higher valence states may also be used as inter- 
mediates in reductions and oxidations. For example, M g ( p and A1(1)8~ 
produced by the anodic oxidations of the metals will reduce 2-methoxy- 
phenyl mesityl ketone, and nitrosobenzene, azoxybenzene and azobenzene 
respectively. 
H H 
l l 
AI - AI++ e- 
3AI++2C6H5N0 + 4 H 2 0 - 60H-+ 3A13++ c 6 H 5 N - NC6H5 
or CO(III) will oxidize alkenes or aromatic compounds87 in a sequence 
such as 
cO2+ - co3+fe- 
and in this case the metal ion acts as a catalyst regenerated at the surface of 
the electrode. Similar oxidations may be carried out with metal oxides pro- 
duced on the electrode surface.87 
Mn2++ 2 H 2 0 - Mn02 +4HS+ 2e- 
M n 0 2 + Rd - Mn2++ Ox + 2e- 
An intermediate which is likely to be of considerable importance is the 
superoxide ion generated by the reduction of oxygen in non-aqueous soh- 
t i0ns.8~7~~ 
0; 0 2 + e- - 
This species will react with organic substrates leading, for example, to the 
102 R.I.C. Reviews 
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formation of organic peroxides in a reductive process. Electrochemical auto- 
catalysis can also be observed.g0 
The reactions which have been listed in this section have been formulated 
as though the intermediates are dissolved in solution. In fact, in many cases 
these intermediates will be adsorbed on the electrode surface and this will be 
particularly true for radical intermediates, for example as shown by the cis 
addition of methyl radicals, generated in the Kolbe reaction, to butadiene.91 
Again the formation of methanol by oxidation of acetate ions in basic solu- 
tions is best explained by the biradical reactiong2 
For the complete oxidation of hydrocarbons to carbon dioxide on platinum 
metal catalysed fuel-cell electrodes, the participation of OH radicals adsorbed 
on the surface has also been postulatedg3 and in these mechanisms one can 
indeed see an application of the original ideas of the mechanism of electro- 
oxidation. Again, it is possible to cite reduction reactions when using catalytic- 
ally active electrode materials such as platinum black where the original 
mechanisms involving hydrogen would be more appropriate (cf. hydro- 
genat ion). 94 
CORRELATIONS BETWEEN ELECTROCHEMICAL DATA AND OTHER PHYSICAL 
PROPERTIES 
For any electrode reaction, a thermodynamic cycle, such as that shown in 
Fig. 7 for the reduction of R, may be drawn up. From the cycle, it may be 
seen that 
where AGRd is the overall free energy change for the reduction R + e-+R-, 
AGEI. and AGE;. are the free energies of solvation of R and R- respectively and 
A is the electron affinity of R. Since the formal electrode potential, EO, (the 
potential on the scale of the normal hydrogen electrode) is given by 
EO = - AGRd/F and for a reversible reduction Eo E= EFd, it may be con- 
cluded that 
AGRd h - -/- AG:;. - FA (6) 
1 
F Ey = A + - [AGE,. - AGri.1 
Similarly for an anodic reaction we can derive the expression 
Fig. 7. Thermodynamic cycle for the reduction of compound R 
(7) 
Fleischmann and Pletcher 103 
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whereI is the ionization potential of R and A G S . is the free energy of solva- 
tion of R+. 
It is important to note that these expressions only hold for reversible 
electrode reactions in the absence of fast following chemical reactions 
removing R+ or R- and in the absence of complex formation between R+ or 
R- and the solvent, or ion-pair formation between R+ or R- and the base 
electrolyte ions. With these reservations in mind, equations 7 and 8 may be 
used directly to correlate Ey with ionization potentials95 and EFd with 
electron affinities.95 The results of such correlations are shown in Figs 8 and 9 
for a series of aromatic hydrocarbons-anthracene, 172-benzanthracene, 
pyrene, chrysene, phenanthrene and triphenylene. Both EFd and Ey were 
obtained in acetonitrile and the Ey were obtained at high voltage sweep 
rates in order to eliminate the following reaction which R+ would 
undergo. 
To within experimental error, linear plots are obtained for both correlations. 
Therefore, it must be concluded that the free energy terms in equations 7 and 
8 are either reasonably constant or vary linearly with potential. It would 
not be surprising if these free energy terms remain constant for series of 
similar large molecules. 
As the electron affinities and ionization potentials may be calculated for 
aromatic systems, for example by simple Hiickel theory, the correlations may 
also be made between E, and the orbital energy, i.e. with the level of the 
lowest unfilled or highest filled molecular orbital.96~97 For alternate hydro- 
carbons these levels are symmetrically placed with respect to the non-bonding 
orbital and this fact also allows correlation with the frequency for the first 
electronic t r an~ i t ion .~~ It has also been shown that EFd and Ey for the 
discharge of aromatic hydrocarbons may be correlated with the frequency of 
the charge-transfer transition of the molecules in the presence of donor or 
acceptor molecules respectively (e.g. hexamethylbenzene99 and tri- 
fluorenoneloo). These correlations are also based on a relation between the 
frequency of the transition and I or A . 
A number of such correlations for a variety of compounds has appeared 
in the literature and, despite the fact that in some cases half-wavepotentials 
for irreversible processes or processes including a coupled reaction have 
been applied, good correlations have been obtained. 
For irreversible cases, the term AGEi. - FA determines the energy of the 
final state of the reaction while AGEI. - F+ determines the energy of the 
initial state. There will therefore be a linear relation (in the first approxima- 
tion) between these energy terms and the free energy of activation, i.e. a 
‘linear free energy correlation’. Indeed, equation 4 relating log I to azF+/JRT 
is an example of such a relationship since + changes the free energy of activa- 
tion in a linear manner. Correlations between Ey and A will therefore still be 
observed provided AGEi. and AGE,. remain constant, the reactions have 
similar pathways and the heats of adsorption of R and R- are also independent 
104 R . I . C. Reviews 
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+ 1.2 I .4 I .6 I .8 
E px (V) 
Fig. 8 (upper) and Fig. 9 (lower). Correlations of €yx with ionization potential and Ey with 
electron affinity for a series of aromatic hydrocarbons 
of the nature of R, since these heat terms also determine the energies of the 
initial and final states. 
Similar linear free energy relationships can be obtained between half-wave 
potentials for compounds RAX and the structure of the compounds as 
expressed by the Hammett-Taft po or p*o* terms, i.e. 
(AE+)X == (E+)X - (E&)H = fn,R'X 
(AE+)x = (E+)x - (E+)H = P*~,RQ*X 
(9) 
(10) 
for a benzenoid series ( i e . A is an aromatic series) or 
Fleischmann and Pletcher 105 
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for a non-benzenoid series, where (E,)x and (E& are the half-wave poten- 
tials of the substituted and unsubstituted molecule respectively (substituent 
X) and p n , ~ , the reaction constants, measure the susceptibility of R to the 
effect of X and are independent of the nature of X, while OX, the substituent 
constants, depend on the position and nature of X but are independent of R. 
These substituent constants are determined by solution phase reactions. 
Linear correlations of the type 9 and 10 therefore relate the electrode process 
to other known reactions and indicate uniformity of reaction mechanism; 
their uses have been reviewed elsewhere.lo1 
SOME FURTHER EXAMPLES OF ELECTROCHEMICAL REACTIONS OF ORGANIC 
COMPOUNDS 
This section is divided into two parts; the first illustrates the importance of 
controlling the electrode potential and solution composition to give selective 
reactions and the second illustrates further reaction types, many of which 
could be applied to synthesis. 
The control of electrode reactions 
As early as 1898 it was shown that nitrobenzene could be reduced to phenyl- 
hydroxylamine at low negative and to aniline at more negative potentials. 
102 
As has already been explained (p. go), the development of electronic 
potentiostats allows the systematic application of this idea, for example, 
CH=NOH CH= NH CHl-NH, 
Q"" 2e-+2H+- VH + H,O 2e- t 2Hi- , @OH I 04 
/ / 
OH OH OH 
I05 
I 06 
R.I.C. Reviews 
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It is interesting to note that the stereochemistry of an electrode reaction 
can also depend on the potential. For example, in the reduction of benzil the 
ratio of cis to trans stilbenediol formed is potential dependent.107 
Furthermore, the course of a reaction may be changed by applying a pulse 
profile such that several different reactions take place successively and 
repetitively. An example is the alternate generation of carbanions and carbo- 
nium ions leading to electrochemiluminescence. An application to synthesis 
is the variation in the reduction mechanism of nitrobenzene caused by 
reoxidizing the intermediate, phenylhydroxylamine, before it has time to 
N O / 
0 
1 (y="o azoxybenzene 
t 
H H (y-yJ 
hydrazobenzene 
H2N m N H 2 benzidine '08 
diffuse away from the electrode surface. Thus, by choosing the potential for 
the reduction step it is possible to accumulate the coupling product, azoxy- 
benzene, or its reduction product. In effect, by applying a suitable square- 
wave potential/time profile, benzidine may be prepared directly from nitro- 
benzenelog and it is clear that many other products could similarly be made 
directly from nitrobenzene by choosing suitable pulse profiles. It will be 
apparent that the duration of the pulses is an additional control variable. 
A question of key importance for achieving a selective reaction is whether 
the desired product is itself electroactive at the potential at which the reaction 
is carried out. The half-wave potential for the substrate and product under 
the experimental conditions will be a useful guide. For example, it is clear 
that it will always be difficult to find selective oxidation routes for the forma- 
tion of phenols or acetates since these are usually more easily oxidizedthan 
the parent aromatic hydrocarbon. On the other hand, cyanation of aromatic 
compounds is a favourable reaction since the products have more positive 
half-wave potentials. Polarographic data will not necessarily be an accurate 
guide for the course of a reaction since large scale preparations are carried 
out on a longer time scale and can, therefore, be affected by disproportiona- 
Fleischmann and Pletcher 107 
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tions. Thus, dihydroaromatic compounds may be formed at the potential at 
which only one electron is added to the h y d r ~ c a r b o n . ~ ~ 
O M e in absence 
of pyridine 
The control of electrode reactions by the correct choice of the solvent has 
already been implied, for example in the discussion of Fig. 6 and in the 
discussion of the protonation of radical anions in aprotic solvents. In aqueous 
solution, the pH is of key importance and determines whether protonated 
or unprotonated species will be electroactive. These effects are also apparent 
in aprotic media where the presence of base (pyridine) can have a marked 
effect on the electrode mechanism. Examples which have been observed 
include the simple, two-electron oxidation of anthracene in the presence of 
pyridine to give a stable product 
M e o a c ' y H M e 
CH ,CHMe 
e - ' y ? ' ' C H M e formation dimer -Meon 3::; & M e 0 
M e 0 
= I M e 0 
Q O M e 
O M e O M e 
108 R.I.C. Reviews 
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Oxidat ions 
In addition to the reactions listed earlier, other oxidation processes can be 
formulated as proceeding through carbonium ion intermediates, for example, 
ring closures, expansions and contractions 
CH,CH,CH,COO- -2e- -'02- CH,CH,CH; CH2- CH, 4- other products 
40 
\ / 
CH2 
75 
and the methoxylation of furans may be formulated as 
r 1 
l 
Me0 O : M e M e 0 R 
On the other hand, many other oxidation reactions may be formulated as 
proceeding through radical intermediates. For example, the Kolbe and the 
related Brown-Walker reactions have in recent years been used for several 
Fleischmann and Pletcher 109 
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interesting syntheses which 
COOH 
proceed via radicals. Examples are 
2 EtOOC (CH,),COO- -Ze--2Cok EtOOC(CH 2)GCOOEt 114 
Other examples include the polymerization of vinyl acetate, vinyl chloride 
and methyl acrylate initiated by radicals formed by the Kolbe reaction.116 
Radicals may be prepared by the anodic oxidation of the sodium derivative 
of diethyl malonate,46 ethyl acetoacetate,46 ethyl phenylacetate46 and nitro- 
paraffins117 as well as less common anionic species. In the absence of reactive 
substrates, the radicals will dimerize, but clearly such radicals are particularly 
useful in synthesis, as in their reactions with alkenic bonds. A reaction scheme 
for the oxidation of diethyl malonate in the presence of vinyl ethyl ether118 is 
OEt OEt 
I I 
(EtOOC),CH - 
cH(COOEt), a 
- CH, - C H (COOE t)2 
formation 
dH(COOEt), CH2=CHoEt- (EtOOC),CH-CH,-tH - OEt 
I I-'- + 
(E t OOC), - CH - CH ,-CH (OEt), - So'Vent (EtOOC)2CH-CH2-CH- OEt 
I 
EtO' EtO' 
EtOOC 
EtO G O E t 
OEt 
110 R. I . C. Reviews 
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This scheme illustrates that radicals may dimerize, hydrogen abstract or be 
oxidized further to give carbonium ions which may react with the solvent, 
cyclize, or lose a proton, and is an example of the build up of complex mole- 
cules from simple starting materials. It is possible to use other alkenes such 
as cyclohexene, styrene or butadiene and the other carbanion species listed 
above .I18 
Several examples of elimination reactions have been reported during 
electrochemical oxidations. Examples include the oxidation of p-substituted 
phenols. 
O H 0 
I I9 
A number of indirect oxidations, where inorganic intermediates generated 
at the electrode react with substrates [cf. Mg+, Al+ in reductions; MnOz, 
Co3+ in oxidations], have been reported. Examples include the electro- 
chemical modification of the Clauson-Kaas methoxylation of furan where 
the bromine is generated at the anode,l21 
r- 1 
- HBr MeOH 
121 
1 
M e 0 O k M e 
the oxidation of propylene where the mercury(I1) is regenerated electro- 
chemically, 122 
CH,-CH=CH2+4Hg2++ H2O- CH,=CH-CHO+4H++ 2 H g r 
the oxidation of propylene to propylene oxide by electrochemically generated 
hyp~chlor i te l~~ 
U 
and the oxidative reaction of alkenes with carbon monoxide and methanol 
Fleischmnnn and Plefcher 11 1 
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in the presence of platinum carbonyls to give methyl esters of a,,B unsaturated 
acids1Z4 
H+ 
C6HSCR =CH, +CO +-OMe - C6H,CR= CHCOOMe 
I29 
Another example of an anodic reaction which is extensively used industrially 
is the perfluorination of aliphatic hydrocarbons in anhydrous hydrogen 
fl~oride.1259~26 The mechanism of these reactions, which clearly must be 
complex, is not fully understood. 
Reductions 
Processes analogous to those described for the reduction of aromatic hydro- 
carbons are also observed in heterocyclic series. Thus the reduction of 
pyrimidine is explained by the scheme127 
H 
112 R.Z. C. Reviews 
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In heterocyclic series, ring expansion has also been observed131 
and several different 
131 
ring closures have been reported, amongst which are 
I32 
OH 
H 
H 
I33 I I 
4e: i2+2 mo 
CH = C(COOH), COOH COO H 
CH3-C == CH -C-CH, 6e-+5H+ CH3-CH-CH2-CH-CH3 ,++ CH, -CH-CH,-CH-CH, 
c L I 
N' H2 HONH NH2 NH ___ 
I I -H,O I I I I 
HONH HN'OH 
I34 
Non-heterocyclic ring closures which have been observed include 
I33 
'OH 
which presumably arises from intramolecular radical coupling. An example 
which may be attributed to carbanionic attack on an activated alkenic bond 
is the intramolecular hydrodimerization, 
I35 
,CH 
CH2 'CH =CHCOOEt 
I 
CH2 
\ /CH =CHCOOEt 
CH2 
m CH2CooEt , CH.! CH2 CH = CHCOOEt 2H+ - ? = + I ___L 
CH2COOEt 
CHI ,C-H - C-HCOOEt 
\ 
CH2 
113 Fleischmann and Pletcher 
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The reactions of activatedalkenes, such as hydrodimerizations, may be 
used for a wide variety of organic syntheses many of which would lead to 
industrially useful products. What is more, such reactions can often be carried 
out in aqueous media using tetraalkylammonium p-toluenesulphonate 
(McKee’s salts) as electrolyte and solubilizing agent. 
As the work up and extraction of products from electrosynthesis will 
frequently be simple compared to conventional processes because of the 
absence of reagents, it is likely that selective electro-organic syntheses will 
find increasing application as additional and alternative reaction routes. 
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