Concerted proton-electron transfer in the oxidation of hydrogen-bonded phenols

Article abstract of DOI:10.1021/ja054167+

Three phenols with pendant, hydrogen-bonded bases (HOAr-B) have been oxidized in MeCN with various one-electron oxidants. The bases are a primary amine (-CPh₂NH₂), an imidazole, and a pyridine. The product of chemical and quasi-reversible electrochemical oxidations in each case is the phenoxyl radical in which the phenolic proton has transferred to the base, ^.OAr-BH⁺, a proton-coupled electron transfer (PCET) process. The redox potentials for these oxidations are lower than for other phenols, predominately from the driving force for proton movement. One-electron oxidation of the phenols occurs by a concerted proton-electron transfer (CPET) mechanism, based on thermochemical arguments, isotope effects, and ΔΔG^?/ΔΔG^o. The data rule out stepwise paths involving initial electron transfer to form the phenol radical cations [^.+HOAr-B] or initial proton transfer to give the zwitterions [^-OAr-BH⁺]. The rate constant for heterogeneous electron transfer from HOAr-NH₂ to a platinum electrode has been derived from electrochemical measurements. For oxidations of HOAr-NH₂, the dependence of the solution rate constants on driving force, on temperature, and on the nature of the oxidant, and the correspondence between the homogeneous and heterogeneous rate constants, are all consistent with the application of adiabatic Marcus theory. The CPET reorganization energies, λ = 23-56 kcal mol^-1, are large in comparison with those for electron transfer reactions of aromatic compounds. The reactions are not highly non-adiabatic, based on minimum values of H_rp derived from the temperature dependence of the rate constants. These are among the first detailed analyses of CPET reactions where the proton and electron move to different sites.

Full text of DOI:10.1021/ja054167+

Published on Web 04/19/2006
Concerted Proton-Electron Transfer in the Oxidation of
Hydrogen-Bonded Phenols
Ian J. Rhile, Todd F. Markle, Hirotaka Nagao, Antonio G. DiPasquale, Oanh P. Lam,
Mark A. Lockwood, Katrina Rotter, and James M. Mayer*
Contribution from the Department of Chemistry, Campus Box 351700, UniVersity of Washington,
Seattle, Washington 98195-1700
Received June 23, 2005; Revised Manuscript Received March 12, 2006; E-mail: mayer@chem.washington.edu
Abstract: Three phenols with pendant, hydrogen-bonded bases (HOAr-B) have been oxidized in MeCN
with various one-electron oxidants. The bases are a primary amine (-CPh2NH2), an imidazole, and a
pyridine. The product of chemical and quasi-reversible electrochemical oxidations in each case is the
•
phenoxyl radical in which the phenolic proton has transferred to the base, OAr-BH+, a proton-coupled
electron transfer (PCET) process. The redox potentials for these oxidations are lower than for other phenols,
predominately from the driving force for proton movement. One-electron oxidation of the phenols occurs
by a concerted proton-electron transfer (CPET) mechanism, based on thermochemical arguments, isotope
effects, and ∆∆Gq/∆∆G°. The data rule out stepwise paths involving initial electron transfer to form the
phenol radical cations [•+HOAr-B] or initial proton transfer to give the zwitterions [-OAr-BH+]. The rate
constant for heterogeneous electron transfer from HOAr-NH2 to a platinum electrode has been derived
from electrochemical measurements. For oxidations of HOAr-NH2, the dependence of the solution rate
constants on driving force, on temperature, and on the nature of the oxidant, and the correspondence
between the homogeneous and heterogeneous rate constants, are all consistent with the application of
adiabatic Marcus theory. The CPET reorganization energies, λ ) 23-56 kcal mol-1, are large in comparison
with those for electron transfer reactions of aromatic compounds. The reactions are not highly non-adiabatic,
based on minimum values of Hrp derived from the temperature dependence of the rate constants. These
are among the first detailed analyses of CPET reactions where the proton and electron move to different
sites.
Scheme 1
Proton-coupled electron transfer (PCET) is of much current
interest, as it is important in a variety of chemical and biological
processes.1,2 Such reactions can occur by concerted or stepwise
mechanisms. The stepwise possibilities include initial transfer
of the proton followed by electron transfer (PT-ET), sometimes
termed proton-gated electron transfer,3 and ET followed by PT
(ET-PT). Reactions in which the proton and electron transfers
occur in one single kinetic step have recently been termed
“concerted proton-electron transfer” (CPET).4,5 CPET encom-
passes a range of processes that involve the transfer of an
electron and a proton, including hydrogen atom transfer (HAT),6
and non-HAT processes where the e- and H+ are separated in
the reactants, products, and/or at the transition structure.7-11
While HAT reactions continue to be the subject of extensive
(1) Cukier, R. I.; Nocera, D. G. Annu. ReV. Phys. Chem. 1998, 49, 337-369.
(2) Mayer, J. M. Annu. ReV. Phys. Chem. 2004, 55, 363-390.
study in organic radical chemistry, the second class of CPET
has received less attention. This report describes studies of a
set of reactions of the latter class: oxidations of intramolecularly
hydrogen-bonded phenols (Scheme 1). Removal of an electron
from these compounds results in transfer of the phenolic proton
to the base. These reactions involve movement of both e- and
H+ but cannot be described as HAT.
(3) For an example, see: Chen, K.; Hirst, J.; Camba, R.; Bonagura, C. A.;
Stout, C. D.; Burgess, B. K.; Armstrong, F. A. Nature 2000, 405, 814-
817.
(4) Costentin, C.; Evans, D. H.; Robert, M.; Save´ant, J.-M.; Singh, P. S. J.
Am. Chem. Soc. 2005, 127, 12490-12491.
(5) The term PCET has been variously used to refer to all processes involving
transfers of H+ and e-, specificially only to concerted processes, or to
proton-coupled electron transfers that are not HAT. We support the recent
suggestion4 that CPET be used to specifically refer to concerted processes.
(6) (a) Free Radicals; Kochi, J. K., Ed.; Wiley: New York, 1973. (b) Mayer,
J. M. Acc. Chem. Res. 1998, 31, 441-450.
PCET oxidations of phenols to phenoxyl radicals are of
particular importance in biological systems because of the
(7) (a) Biczo´k, L.; Gupta, N.; Linschitz, H. J. Am. Chem. Soc. 1997, 119, 12601.
(b) Gupta, N.; Linschitz, H. J. Am. Chem. Soc. 1997, 119, 6384.
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J. AM. CHEM. SOC. 2006, 128, 6075-6088
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A R T I C L E S
Rhile et al.
widespread involvement of tyrosyl radicals in enzymatic
processes.12 They have been implicated as intermediates in class
I ribonucleotide reductases,13 photosystem II,14 prostaglandin
H synthases 1 and 2,15 cytochrome c oxidase,16 galactose
oxidase,17 amine oxidases,18 and other systems.12 In many cases,
the phenoxyl radical is generated from the phenol by outer-
sphere electron transfer, with release of the proton to a nearby
residue (histidine, arginine, lysine, etc.) or to a hydrogen-bonded
network.12 An interesting example is the oxidation of tyrosine
160 of the D2 subunit (YZ) in Photosystem II by long-range
proton transfer (PT), in some cases to bulk solution, while the
HOAr-B compounds reported here have an intramolecular PT
in aprotic media. The use of aprotic media and a strong initial
hydrogen bond provides the advantage of being able to keep
track of the proton but may limit the generality of the
conclusions. More studies are required to model biological and
chemical systems with weaker hydrogen-bonding interactions
and systems in which the formation of charged intermediates
is more facile (perhaps with a higher local effective dielectric
constant). Our studies and the model systems mentioned above
all conclude that concerted proton-electron transfer is the
dominant pathway under most conditions, but Hammarstro¨m
and co-workers have shown that a proton-first mechanism takes
over at high pH, where deprotonation of tyrosine is energetically
accessible.8 Similarly, elegant work by Okamura and others has
indicated stepwise mechanisms for quinone reduction in pho-
tosystem I.20
electron transfer to the light-induced chlorophyll radical cation
+ 19
P680
.
The phenolic proton of YZ is likely transferred to a
hydrogen-bonded histidine (His190 of subunit D1). This tyrosyl
radical then is involved in the oxidation of the manganese cluster
and eventually the conversion of water to O2.
The HOAr-B systems examined here were designed to model
such phenol oxidations with concomitant proton transfer. Related
model studies include oxidation of tyrosine by a pendant
photogenerated [Ru(bpy)3]3+ or a photoexcited ReI center8,9 and
electron transfer from phenol-pyridine adducts to photoexcited
C60.7 These previous studies have all involved intermolecular
The motif of a tyrosine hydrogen-bonded to a base may be
viewed as a biological redox cofactor. A variety of other electron
transfer cofactors, such as iron-sulfur clusters, hemes, and
quinones, have been studied and understood on the basis of the
Marcus-Hush theory of electron transfer.21 We have previously
shown that rate constants for hydrogen atom transfer reactions
are in many cases well predicted by the Marcus cross relation.22
This report shows that Marcus theory can also be applied to
non-HAT CPET reactions, and it describes the characteristics
of the HOAr-B compounds as electron transfer reagents,
highlighting the influence of the PT on the thermodynamics
and kinetics of electron transfer. The results are also discussed
in light of the more recent and more sophisticated theoretical
models of CPET.23 A preliminary report has described the
oxidation of one of the phenols, HOAr-NH2.24
(8) (a) Magnuson, A.; Berglund, H.; Korall. P.; Hammarstro¨m, L.; Åkermark,
B.; Styring, S.; Sun, L. J. Am. Chem. Soc. 1997, 119, 10720-5. (b) Sjodin,
M.; Styring, S.; Åkermark, B.; Sun, L.; Hammarstro¨m, L. J. Am. Chem.
Soc. 2000, 122, 3932-3936. (c) Sjo¨din, M.; Styring, S.; Wolpher, H.; Xu,
Y.; Sun, L.; Hammarstro¨m, L. J. Am. Chem. Soc. 2005, 127, 3855-3863.
(d) Sjo¨din, M.; Ghanem, R.; Polivka, T.; Pan, J.; Styring, S.; Sun, L.;
Sundstro¨m, V.; Hammarstro¨m, L. Phys. Chem. Chem. Phys. 2004, 6, 4851-
4858.
(9) Reece, S. Y.; Nocera, D. G. J. Am. Chem. Soc. 2005, 127, 9448-9458.
(10) (a) Shukla, D.; Young, R. H.; Farid, S. J. Phys. Chem. A 2004, 108, 10386-
10394. (b) Turro, C.; Chang, C. K.; Leroi, G. E.; Cukier, R. I.; Nocera, D.
G. J. Am. Chem. Soc. 1992, 114, 4013. (c) Chang, M. C. Y.; Yee, C. S.;
Nocera, D. G.; Stubbe, J. J. Am. Chem. Soc. 2004, 126, 16702-16703. (d)
Lehmann, M. W.; Evans, D. H. J. Phys. Chem. B 2001, 105, 8877-8884.
(e) Mayer, J. M.; Hrovat, D.; Thomas, J. L.; Borden, W. T. J. Am. Chem.
Soc. 2002, 124, 11142-11147. (f) Anglada, J. M. J. Am. Chem. Soc. 2004,
126, 9809-9820. (g) Weatherly, S. C.; Yang, I. V.; Armistead, P. A.; Thorp,
H. H. J. Phys. Chem. B. 2003, 107, 372-378. (h) Stubbe, J.; Nocera, D.
G.; Yee, C. S.; Chang, M. C. Y. Chem. ReV. 2003, 103, 2167-2201. (i)
DiLabio, G. A.; Ingold, K. U. J. Am. Chem. Soc. 2005, 127, 6693-6699.
(j) Huynh, M. H. V.; Meyer, T. J. Proc. Natl. Acad. Sci. U.S.A. 2004, 101,
13138-13141. (k) Meyer, T. J.; Huynh, M. H. V. Inorg. Chem. 2003, 42,
8140-8160.
(11) (a) Kojima, T.; Sakamoto, T.; Matsuda, Y.; Ohkubo, K.; Fukuzumi, S.
Angew. Chem., Int. Ed. 2003, 42, 4951. (b) Haddox, R. M.; Finklea, H. O.
J. Electroanal. Chem. 2003, 550-551, 351.
(12) (a) Stubbe, J.; van der Donk, W. A. Chem. ReV. 1998, 98, 705-762. (b)
Pesavento, R. P.; van der Donk, W. A. AdV. Protein Chem. 2001, 58, 317-
385.
(13) (a) Ehrenberg, A.; Reichard, P. J. Biol. Chem. 1972, 247, 3485-3488. (b)
Sjo¨berg, B.-M.; Reichard, P.; Gra¨slund, A.; Ehrenberg, A. J. Biol. Chem.
1978, 253, 6863-6865. (c) Sahlin, M.; Gra¨slund, A.; Ehrenberg, A.;
Sjo¨berg, B.-M. J. Biol. Chem. 1982, 257, 366-369. (d) Griepenburg, U.;
Lassmann, G.; Auling, G. Free Radical Res. 1996, 26, 473-481.
(14) (a) Barry, B. A.; El-Deeb, M. K.; Sandusky, P. O.; Babcock, G. J. Biol.
Chem. 1990, 265, 20139-20143. (b) Barry, B. A.; Babcock, G. T. Proc.
Natl. Acad. Sci. U.S.A. 1987, 84, 7099-7103.
(15) (a) Tsai, A.-L.; Kulmacz, R. J.; Palmer, G. J. Biol. Chem. 1995, 270,
10503-10508. (b) Tsai, A.-L.; Palmer, G.; Kulmacz, R. J. J. Biol. Chem.
1992, 276, 17753-17759. (c) Tsai, A. L.; Palmer, G.; Xiao, G.; Swinney,
D. C.; Kulmacz, R. J. J. Biol. Chem. 1998, 273, 3888. (d) Hsi, L. C.;
Hoganson, C. W.; Babcock, G. T.; Smith, W. L. Biochem. Biophys. Res.
Commun. 1994, 202, 1592-1598.
Results
1. Syntheses and Characterization of Compounds. The
phenol-amine HOAr-NH2 was synthesized as outlined in
(19) (a) Tommos, C.; Babcock, G. T. Biochim. Biophys. Acta 2000, 1458, 199-
219. (b) Vrettos; J. S.; Limburg, J.; Brudvig, G. W. Biochim. Biophys.
Acta 2001, 1503, 229-245. (c) Renger, G. Biochim. Biophys. Acta 2004,
1655, 195-204. (d) Rappaport, F.; Lavergne, J. Biochim. Biophys. Acta
2001, 1503, 246-259. (e) Nugent, J. H. A.; Rich, A. M.; Evans, M. C. W.
Biochim. Biophys. Acta 2001, 1503, 138-146. (f) Kuhne, H.; Brudvig, G.
W. J. Phys. Chem. B 2002, 106, 8189-8196. (g) Faller, P.; Goussias, C.;
Rutherford, A. W.; Un, S. Proc. Natl. Acad. Sci. U.S.A. 2003, 100, 8732-
8735. (h) Ferreira, K. N.; Iverson, T. M.; Maghlaoui, K.; Barber, J.; Iwata,
S. Science 2004, 303, 1831-1838. (i) Zouni, A.; Witt, H.-T.; Kern, J.;
Fromme, P.; Krauss, N.; Saenger, W.; Orth, P. Nature 2001, 409, 739-
743. (j) Rhee, K.-H.; Morris, E. P.; Barber, J.; Kuhlbrandt, W. Nature 1998,
396, 283-286. (k) Haumann, M.; Mulkidjanian, A.; Junge, W. Biochemistry
1999, 38, 1258-1267. (l) Ka´lma´n, L.; LoBrutto, R.; Allen, J. P.; Williams,
J. C. Nature 1999, 402, 696-699. (m) Petrouleas, V.; Koulougliotis, D.;
Ionnidis, N. Biochemistry 2005, 44, 6723-6728.
(20) Graige, M. S.; Paddock, M. L.; Bruce, J. M.; Feher, G.; Okamura, M. Y.
J. Am. Chem. Soc. 1996, 118, 9005-9016.
(16) (a) Ferguson-Miller, S.; Babcock, G. T. Chem. ReV. 1996, 96, 2889-2907.
(b) Gamelin, D. R.; Randall D. W.; Hay, M. T.; Houser, R. P.; Mulder, T.
C.; Canters, G. W.; de Vries, S.; Tolman, W. B.; Lu, Y.; Solomon, E. I. J.
Am. Chem. Soc. 1998, 120, 5246-5263 and references therein. (c)
Proshlyakakov, D. A.; Pressler, M. A.; DeMaso, C.; Leykam, J. F.; DeWitt,
D. L.; Babcock, G. T. Science 2000, 290,1588-1591.
(21) (a) Page, C. C.; Moser, C. C.; Chen, C.; Dutton, L. Nature 1999, 402,
47-5217. (b) Barbara, P. F.; Meyer, J. T.; Ratner, M. A. J. Phys. Chem.
1996, 100, 13148-13168.
(22) Roth, J. P.; Yoder, J. C.; Won, T.-J.; Mayer, J. M. Science 2001, 294,
2524-2526.
(17) Whittaker, M. M.; Whittaker, J. W. J. Biol. Chem. 1990, 265, 9610-9613.
(18) (a) Janes, S. M.; Mu, D.; Wemmer, D.; Smith, A. J.; Kaur, S.; Maltby, D.;
Burlingame, A. L.; Klinman, J. P. Science 1990, 248, 981-987. (b) Janes,
S. M.; Palcic, M. M.; Scaman, C. H.; Smith, A. J.; Brown, D. E.; Dooley,
D. M.; Mure, M.; Klinman, J. P. Biochemistry 1992, 31, 12147-12154.
(c) Cooper, R. A.; Knowles, P. F.; Brown, D. E.; McGuirl, M. A.; Dooley,
D. M. Biochem. J. 1992, 288, 337-340. (d) Brown, D. E.; McGuirl, M.
A.; Dooley, D. M.; Janes, S. M.; Mu, D.; Klinman, J. P. J. Biol. Chem.
1991, 266, 4049-4051. (e) Mu, D.; Janes, S. M.; Smith, A. J.; Brown, D.
E.; Dooley, D. M.; Klinman, J. P. J. Biol. Chem. 1992, 267, 7979-7982.
(23) (a) Swalina, C.; Pak, M. V.; Hammes-Schiffer, S. Chem. Phys. Lett. 2005,
404, 394-399. (b) Hammes-Schiffer, S.; Iordanova, N. Biochim. Biophys.
Acta 2004, 1655, 29-36. (c) Hammes-Schiffer, S. Acc. Chem. Res. 2001,
34, 273-281. (d) Hammes-Schiffer, S. ChemPhysChem 2002, 33-42. (e)
Cukier, R. I. J. Phys. Chem. B 2002, 106, 1746-1757. (f) Georgievskii,
Y.; Stuchebrukhov, A. A. J. Chem. Phys. 2000, 113, 10438-10450. (g)
Kuznetsov, A. M.; Ulstrup, J. Can. J. Chem. 1999, 77, 1085-1096. (h)
Krishtalik, L. I. Biochim. Biophys. Acta 2000, 1458, 6-27. (i) References
1 and 2.
(24) Rhile, I. J.; Mayer, J. M. J. Am. Chem. Soc. 2004, 126, 12718-12719.
9
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Concerted Proton−Electron Transfer in Phenols
A R T I C L E S
Scheme 2. Synthesis of HOAr-NH2
Scheme 2, following literature precedents.25 The tertiary
t
-CPh2NH2 and Bu substituents in the 2, 4, and 6 positions
confer stability on the derived phenoxyl radical; 2,4,6-
tBu3C6H2O•, for instance, is stable in solution.26 Recently, 2,4-
di-tert-butyl-6-(N-methyl-2-pyrrolidyl)phenol was reported to
give a persistent oxidized form, decaying over 30 min after bulk
electrolysis.27 Related compounds with a -CH2- spacer
between the amine and the phenol are readily available via the
Mannich reaction (phenol + formaldehyde + amine),28 but such
compounds are susceptible to radical attack at the benzylic
hydrogens (and at other C-H bonds R to the amine).29 The
Mannich procedure cannot be used to make tertiary substituents
because of the decreased reactivity of the ketone-derived
iminium cation.28 HOAr-NH2 was therefore synthesized by
addition of benzophenone to the lithiated phenol, leading to the
gem-diphenyl substituents.25a Subsequent trityl chemistry leads
to products.25b,c The corresponding chemistry with gem-methyl
groups is diverted by elimination from the HOArCMe2OH
intermediate under the mild acidic conditions.
Figure 1. ORTEP drawings of (a) HOAr-NH2, (b) HOAr-im, and (c)
HOAr-py.
Table 1. Structural, Spectroscopic, and Electrochemical Data for
Phenol-Base Compounds
δO
E1 2 (V)
/
[∆E
p (mV)]b
-
H
phenol (HOAr-B)
dO N (Å)
‚‚‚
(ppm)a
HOAr-NH2
HOAr-im
HOAr-py
2.550(2), 2.613(3)c
2.646(2)
2.561(3), 2.567(3),
2.573(3)f
12.32
13.42
14.83
0.37 [143]d
0.42 [105]e
0.58 [100]d
a 1H NMR data in CD3CN. b E vs Cp2Fe+/0
.
c Two independent molecules
The related phenol with a 4,5-bis(4-anisyl)-2-imidazolyl
substituent, HOAr-im, has been reported by Benisvy,30 and the
pyridyl compound HOAr-py has been prepared by Fujita.31 In
each case, the authors explored the compounds’ properties as
ligands to metals. Both compounds fluoresce under ultraviolet
light due to excited-state intramolecular proton transfer (ES-
IPT).32
in the unit cell. d Scan rate ) 200 mV s-1
.
e Scan rate ) 100 mV s-1 f Three
.
independent molecules in the unit cell.
11.9-15.4° for the three crystallographically independent
molecules of HOAr-py. In the two independent molecules of
HOAr-NH2, the NCCC torsion angles are 33.4 and 42.0°.
Similar structures have been observed for related molecules.33
The O‚‚‚N distances across the hydrogen bond vary between
2.550(3) and 2.646(2) Å (Table 1), which are in the shorter
portion of the known range for OH‚‚‚N hydrogen bonds.33,34
Crystal packing forces appear to play a significant role in these
distances, as the two independent molecules of HOAr-NH2 in
the unit cell have O‚‚‚N distances that differ by 0.063 (4) Å;
for the three molecules of HOAr-py, the O‚‚‚N distances vary
by 0.012(3) Å. The imidazole derivative crystallizes with a
molecule of methanol that is hydrogen-bonded to the imidazole
NH hydrogen.
The X-ray crystal structures of HOAr-NH2, HOAr-py, and
HOAr-im (Figure 1) all show molecules with intramolecular
hydrogen bonds from the phenol to the nitrogen base. There is
some twisting between the phenol and pyridyl or imidazolyl
rings, with inter-ring torsion angles of 22.6° for HOAr-im and
(25) (a) Talley J. J.; Evans, I. A. J. Org. Chem. 1984, 49, 5267-5269. (b)
Gomberg, N.; Nishida, D. J. Chem. Soc. 1922, 190-207. (c) Mandell, L.;
Piper, J. V.; Pesterfield, C. E. J. Org. Chem. 1963, 28, 574-575.
(26) Altwicker, E. R. Chem. ReV. 1967, 67, 475-531.
(27) Maki, T.; Araki, Y.; Ishida, Y.; Onomura, O.; Matsumura, Y. J. Am. Chem.
Soc. 2001, 123, 3371-3372.
(28) (a) Tramontini, M.; Angiolini, L. Tetrahedron 1990, 46, 1791-1837. (b)
Gevorgyan, G. A.; Agababyan, A. G.; Mndzhoyan, O. L. Usp. Khim. 1984,
53, 971-1013. (c) Tramontini, M. Synthesis 1973, 703-775. (d) House,
H. O. Modern Synthetic Reactions, 2nd ed.; W. A. Benjamin: New York,
1972; p 654. (e) See also refs 33a and 39.
(29) Sparfel, D.; Baranne-Lafont, J.; Cuong, N. K.; Capdevielle, P.; Maumy,
M. Tetrahedron 1990, 46, 803-814.
(30) (a) Benisvy, L.; Bill, E.; Blake, A. J.; Collison, D.; Davies, E. S.; Garner,
C. D.; Guindy, C. I.; McInnes, E. J. L.; McArdle, G.; McMaster, J.; Wilson,
C.; Wolowska, J. Dalton Trans. 2004, 3647-3653. (b) Benisvy, L.; Bittl,
R.; Bothe, E.; Garner, C. D.; McMaster, J.; Ross, S.; Teutloff, C.; Neese,
F. Angew. Chem., Int. Ed. 2005, 44, 5314-5317.
(31) Inoue, Y.; Nakano, T.; Tanaka, H.; Kashiwa, N.; Fujita, T. Chem. Lett.
2001, 1060-1061.
(32) (a) Stolow, A. Annu. ReV. Phys. Chem. 2003, 54, 89-119. (b) LeGourrierec,
D.; Kharlanov, V.; Brown, R. G.; Rettig, W. J. Photochem. Photobiol. A
1998, 117, 209-216. (c) Braeuer, M.; Mosquera, M.; Perez-Lustres, J. L.;
Rodriguez-Prieto, F. J. Phys. Chem. A 1998, 102, 10736-10745.
NMR spectra of the phenols in dry CD3CN all show sharp
downfield resonances for the phenolic proton, e.g., 12.32 ppm
(33) (a) In Mannich bases, the range of O- - -N distances across the hydrogen
bond is 2.56-2.71 Å: Koll, A.; Wolschann, P. Monatsh. Chem. 1996, 127,
475-486. (b) Imidazoles, 2.55-2.60 Å: Foces-Foces, C.; Llames-Saiz,
A. L.; Claramunt, R. M.; Cabildo, P.; Elguero, J. J. Mol. Struct. 1998,
440, 193-202. Benisvy, L.; Blake, A. J.; Collison, D.; Davies, E. S.; Garner,
C. D.; McInnes, E. J. L.; McMaster, J.; Whittaker, G.; Wilson, C. J. Chem.
Soc., Dalton Trans. 2003, 1975-1985. (c) Pyridines, 2.54-2.56 Å: Shu,
Wenmaio; Valiyaveettil, S. Chem. Commun. 2002, 1350-1351. Kaczmarek,
L.; Balicki, R.; Lipowski, J.; Borowicz, P.; Grabowska, A. J. Chem. Soc.,
Perkin Trans. 2 1994, 1603-1610. (d) Most Mannich bases HOAr-CH2-
NR2 are nonplanar due to steric pressure from the R groups.33a-c,75
(34) (a) Pimental, G. C.; McClellan, A. L. The Hydrogen Bond; Freeman: New
York, 1960. (b) Frey, P. A. Magn. Reson. Chem. 2001, 39, S190-S198.
(c) Gilli, P.; Bertolasi, V.; Gilli, G. J. Am. Chem. Soc. 2000, 122, 10405-
10417.
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A R T I C L E S
Rhile et al.
HOAr-py shows no peak above 385 nm that would be
characteristic of the proton-transferred structure.40
2. Cyclic Voltammetry and Chemical Oxidations. Oxida-
tions of HOAr-NH2, HOAr-im, and HOAr-py with near-
stoichiometric amounts of [N(p-C6H4Br)3]•+ yield the corre-
sponding phenoxyl radical (Scheme 1 above), as confirmed by
both UV-vis and 1H NMR spectroscopies (at ∼10 µM and ∼1
mM concentrations, respectively). The reactions are marked by
a rapid decrease of the intense absorption of the blue aminium
ion at 699 nm (40 000 M-1 cm-1). These reactions, and most
of the solution measurements in this report, were done in MeCN.
Oxidation of HOAr-im with [N(p-C6H4Br)3]•+ yields a blue
solution of OAr-imH+ with an absorption at 695 nm (8300)
•
Figure 2. UV-vis spectra of (a) phenol HOAr-NH2 and (b) phenoxide
-OAr-NH2 in MeCN. The inset of spectrum (a) is the spectrum of a
saturated solution of HOAr-NH2 in a 10 cm path length cell.
which decays to ∼33% intensity over 1.5 h. Phenoxyl radicals
typically have absorptions between 420 and 720 nm, with higher
intensity for the more conjugated radicals [λmax, nm (ꢀ, M-1
cm-1)]: 2,4,6-tri-tert-butylphenoxyl radical (tBu3ArO•), 630
(400), benzene;41 2,6-tBu2-4-Ph-C6H2OH, 488 (2780); 2,6-tBu2-
4-(Me2NC6H4)C6H2OH, 650 (6000).26 Treating •OAr-imH+ in
MeCN with triethylamine or excess pyridine (pKa ) 18, 12,
respectively38) produces a purple solution with λmax ) 544 nm
(approximately 6100) due to the deprotonated phenoxyl radical
•OAr-im (Figure 3, eq 2). This species likely still has an
for HOAr-NH2, typical of intramolecularly hydrogen-bonded
phenols.35a The chemical shifts for HOAr-py (14.83 ppm) and
HOAr-im (13.42 ppm) are farther downfield, as has been
previously observed for related compounds36 that have “resonance-
assisted hydrogen bonds” due to the conjugation between the
phenol and the basic site.34c,37 The two p-anisyl groups in
HOAr-im are inequivalent, indicating that intermolecular proton
transfer between imidazole nitrogen atoms is slow on the NMR
time scale, presumably due in part to the strong OH‚‚‚N
hydrogen bond.
The UV-vis spectrum of HOAr-NH2 contains absorptions
typical of aromatic compounds35b at 207 (40 000) and 287 nm
(3600) (Figure 2a; the ꢀ value is stated parenthetically after each
λmax in M-1 cm-1) The deprotonated phenol (-OAr-NH2) is
generated in MeCN by addition of excess di(tetra-n-butylam-
monium) succinate.38 -OAr-NH2 has additional absorptions at
259 (6900) and 327 nm (4700) (Figure 2b), low-energy
absorptions that are typical of phenoxides.35 A UV-vis spectrum
of a saturated (16.0 mM) MeCN solution of HOAr-NH2 in a
10.00 cm quartz cell shows no absorption maximum in the
phenoxide region (inset of Figure 2a). These optical spectra
provide an estimate of the equilibrium constant KPT2 for
formation of the zwitterion -OAr-NH3+ (eq 1). Mannich bases
with strongly acidic phenols can exist in this tautomeric form,
and the optical spectrum of the phenoxide (e.g., -OAr-NH2)
has often been taken as a model for the low-energy part of the
spectrum of the zwitterion.39 With this assumption, the lack of
intramolecular hydrogen bond from the imidazole hydrogen to
the oxyl radical. OAr-im was prepared independently by
•
heterogeneous PbO2 oxidation of HOAr-im in MeCN or
•
DMSO.42 This isolated OAr-im had an absorption at 544 nm
and contained some HOAr-im by 1H NMR. Addition of 1 equiv
•
of triflic acid to solutions of OAr-im formed the 695 nm
•
absorption characteristic of OAr-imH+ (eq 2).
Oxidation of HOAr-py by [N(p-C6H4Br)3]•+ gives a yellow
solution with λmax ) 481 nm (1600), which fades with t1/2 ≈ 6
(39) (a) Koll, A.; Wolschann, P. Monatsh. Chem. 1999, 130, 983-1001. (b)
Przeslawska, M.; Koll, A.; Witanowski, M. J. Phys. Org. Chem. 1999, 12,
486-492. (c) Teitelbaum, A. B.; Derstuganova, K. A.; Shishkina, N. A.;
Kudryavtseva, L. A.; Bel’skii, V. E.; Ivanov, B. E. Bull. Acad. Sci. USSR,
DiV. Chem Sci. (Engl. Transl.) 1980, 558-562; IzV. Akad. Nauk SSSR,
Ser. Khim. 1980, 803-808.
(40) (a) Kro´l-Starzomska, I.; Filarowski, A.; Rospenk, M.; Koll, A. J. Phys.
Chem. A 2004, 108, 2131-2138. This work states that zwitterionic
o-hydroxy Schiff bases have 385 nm < λmax < 430 nm. (b) Popp, G. J.
Org. Chem. 1972, 37, 3058-3062. (c) Lapachev, V.; Stekhova, S.; Mamaev,
V. Monatsh. Chem. 1987, 118, 669-670.
(41) (a) Woon, T. C.; Dicken, C. M.; Bruice, T. C. J. Am. Chem. Soc. 1986,
108, 7990-7995. (b) Pokhedenko, V. D.; Khizhny, V. A.; Koshechko, V.
G.; Samarskii, V. A. Theor. Exp. Chem. (Engl. Transl.) 1975, 11, 489-
493; Teor. Eksper. Khim. 1975, 11, 579-584. This work reports tBu3ArO•
λmax (nm, MeCN) ) 316, 628 nm (but no ꢀ’s).
(42) (a) Xie, C.; Lahti, P. M. Tetrahedron Lett. 1999, 40, 4305-4308. (b) Xie,
C.; Lahti, P. M.; George, C. Org. Lett. 2000, 2, 3417-3420.
an absorption maximum at 327 nm (ꢀHOAr-NH (327) ) 1.1 M-1
2
cm-1) implies that essentially no zwitterion is present in MeCN
solution, that KPT2 < 10-4. Similarly, the UV-vis spectrum of
(35) Silverstein, R. M.; Bassler, G. C.; Morrill, T. C. Spectrometric Identification
of Organic Compounds, 5th ed.; Wiley: New York, 1991; (a) p 184 and
(b) pp 306-311.
(36) Rozwadowski, Z.; Majeewski, E.; Dziembowska, T.; Hansen, P. J. Chem.
Soc., Perkin Trans. 2 1999, 2809-2817.
(37) Gilli, G.; Belluci, F.; Ferretti, V.; Bertolasi, V. J. Am. Chem. Soc. 1989,
111, 1023-1028.
(38) Izutsu, K. Acid-Base Dissociation Constants in Dipolar Aprotic SolVents;
IUPAC Chemical Data Series No. 35; Blackwell Scientific Publications:
Boston, MA, 1990; pp 17-35. The pKa2 of succinic acid in MeCN is 29.0.
9
6078 J. AM. CHEM. SOC. VOL. 128, NO. 18, 2006

Concerted Proton−Electron Transfer in Phenols
A R T I C L E S
•
•
Figure 3. Visible spectra of (a) OAr-imH+ and (b) OAr-im in MeCN.
h. Reactions of HOAr-NH2 with [N(p-C6H4Br)3]•+ show no
absorptions above 400 nm at 100 µM, indicating a colorless
radical product. A complex EPR spectrum was recorded for one
of the oxidation mixtures of HOAr-NH2 in CH2Cl2 (see
Supporting Information of ref 24). 1H NMR monitoring of
reactions of HOAr-NH2 with substoichiometric amounts of
[N(p-C6H4Br)3]•+ in MeCN showed reduced signals for HOAr-
NH2 and the appearance of N(p-C6H4Br)3. With excess [N(p-
C6H4Br)3]•+, the amine is not observed because there is rapid
exchange between NAr3 and [NAr3]•+ by electron transfer.43
The cyclic voltammograms of the three phenols in 0.1 M
nBu4NPF6/MeCN (Table 1, Figures 4a and S20 in Supporting
Information) are quasi-reversible, with almost equal anodic and
cathodic currents but with peak separations (∆Ep) larger than
the theoretical 59 mV. The rate constant for heterogeneous
electron transfer (kel) for HOAr-NH2 has been determined by
analysis of the CV data at different scan rates ν (Figure 4a).44
kel is related to ∆Ep and ν by eqs 3 and 4,
Figure 4. (a) Cyclic voltammograms of HOAr-NH2 with ν ) 0.05, 0.1,
0.2, and 0.5 V s-1. Inset: An overlay of the simulated (blue dots) and
experimental CVs at 100 mV s-1. (b) Plot of ψ vs ν-1/2 for HOAr-NH2
(see eqs 3 and 4 in text).
have recently reported a much slower heterogeneous rate
constant of (9 ( 5) × 10-7 cm s-1 for CPET reduction of a
water-superoxide complex, which exhibits a much more
distorted cyclic voltammogram.4
Similar quasi-reversible voltammograms have been reported
for other phenols with intra- or intermolecularly hydrogen-
bonded amine or pyridine bases.7,27,30,47 In contrast, electro-
chemical oxidations of phenols without an attached base are
irreversible in dried aprotic solvents, occurring via an EC
mechanism. The chemical step (“C”) is typically proton transfer
into the bulk solution, which is effectively irreversible in aprotic
media.48 Thus, oxidation of 2,4,6-tBu3ArOH is irreversible in
dry MeCN, even though the radical 2,4,6-tBu3ArO• is stable.48a
Matsumura et al. have shown27 that moving the attached base
from the ortho to the para position changes the oxidation from
quasi-reversible to irreversible. Oxidation of the methyl ether
MeOAr-NH2 is irreversible (Ep,a ) 1.2 V; all potentials in this
report are vs Cp2Fe0/+ in MeCN), probably because without the
stabilizing proton transfer the high-energy anisyl or aminium
radical cation decays rapidly. The related phenol-alcohol
1/2
R/2
πDOFν
DR
k0el ) ψ
(3)
(
) ( )
RT
DO
ln ψ ) 3.69 - 1.16 ln(∆Ep - 59)
(4)
where DO and DR are the diffusion constants (cm2 s-1) for the
oxidized and reduced forms of the analyte, R is the transfer
coefficient (taken to be 0.545), and R and T have their standard
meanings. A, DO, and DR were determined using chronoamper-
ometry (see Supporting Information). kel was found to be (3 (
1) × 10-3 cm s-1 from the slope of a plot of ψ vs ν-1/2 (Figure
4b). To support our measurements, these parameters were used
to simulate the CVs using DigiSim46 with good results (inset,
Figure 4a). For comparison, Evans, Save´ant, and co-workers
(46) DigiSim software is a product of Bioanalytical Systems, Inc. (http://
www.bioanalytical.com/products/ec/digisim/index.html).
(43) Sorensen, S. P.; Bruning W. H. J. Am. Chem. Soc. 1973, 95, 2445-2451.
(44) (a) Swaddle, T. W. Chem. ReV. 2005, 105, 2573-2608. (b) Bard, A. J.;
Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications,
2nd ed.; John Wiley and Sons Inc.: New York, 2001.
(45) The value of ψ is nearly independent of R (0.3 < R < 0.7) (Nicholson, R.
S. Anal. Chem. 1965, 37, 1351). When ipc/ipa ) 1, as is the case here, the
value of R is typically close to 0.5. Simulated CVs using R ) 0.4, 0.5, and
0.6 showed no significant difference.
(47) (a) Thomas, F.; Jarjayes, O.; Jamet, H.; Hamman, S.; Saint-Aman, E.;
Duboc, C.; Pierre, J.-L. Angew. Chem. 2004, 116, 604-607; Angew. Chem.,
Int. Ed. 2004, 43, 594-597. (b) Rhile, I. J.; Mayer, J. M. Angew. Chem.
2005, 105, 1624-1625; Angew. Chem., Int. Ed. 2005, 44, 1598-1599.
(48) (a) Bordwell, F. G.; Cheng, J.-P. J. Am. Chem. Soc. 1991, 113, 1736-
1743. (b) Cf.: Williams, L. L.; Webster, R. D. J. Am. Chem. Soc. 2004,
126, 12441-12450. (c) For aqueous electrochemistry, see: Li, C.; Hoffman,
M. Z. J. Phys. Chem. B 1999, 103, 6653-6656.
9
J. AM. CHEM. SOC. VOL. 128, NO. 18, 2006 6079

A R T I C L E S
Rhile et al.
a
Table 2. Potentials of Oxidants (in MeCN, V vs Cp2Fe+/0
)
HOAr-OH [(2-CPh2OH)(4,6-tBu2)C6H2OH] also shows ir-
reversible electrochemistry (Ep,a ) 1.1 V), possibly because
proton transfer to the weakly basic primary alcohol is not
favored and the proton is lost to the bulk solution. CV of the
oxidant
E1
/2
[N(p-C6H4Br)3]•+
0.67
0.48
0.38
0.32
0.16
0.70
0.58
0.53
0.46
0.32
[N(p-C6H4OMe)(p-C6H4Br)2]•+
[N(tol)3]•+ b
n
phenoxide -OAr-NH2, as the Bu4N+ salt, shows a reversible
[N(p-C6H4OMe)2(p-C6H4Br)]•+
[N(p-C6H4OMe)3]•+
[Fe(bpy)3]3+
oxidation wave centered at -0.57 V, essentially equal to the
E1/2 for 2,4,6-tri-tert-butylphenoxide, -0.572 V.48a,53
[Fe(5,5′-Me2bpy)3]3+
[Fe(4,7-Me2phen)3]3+
[Fe(3,4,7,8-Me4phen)3]3+
[MPT]•+ c
The average of the anodic and cathodic peaks for HOAr-
NH2, HOAr-im, and HOAr-py is taken as the E1/2 for the
coupled proton-electron transfer (CPET), the potential for
transfer of both an electron to the electrode and the phenolic
proton to the amine (Scheme 1). This is the interpretation of
most of the previous electrochemical studies of phenol-base
systems.7,27,30 One recent paper has interpreted the large ∆Ep
for oxidation of a phenol-amine as indicating a stepwise EC
(ET-PT) mechanism,47a but this is, in our view, inappropriate.47b
The assignment of E1/2 as the energetics of CPET is supported
by the thermodynamic discussion below and by the following
equilibration experiment.
a See Experimental Section for conditions. b Tri-p-tolylaminium. c 10-
Methylphenothiazinium.
phenanthroline derivative), and the 10-methylphenothiazinium
ion [MPT]•+. All displayed reversible cyclic voltammograms
(Table 2). The [Fe(N-N)3]3+/2+ potentials in MeCN vary
substantially with ionic strength due to differences in ion-pairing
between the FeII and FeIII forms.50 For [Fe(5,5′-Me2bpy)3]3+/2+
,
the potential changes by -40 ( 4 mV/log(i). Kinetic studies
using [Fe(R2bpy)3]3+ and [Fe(Mexphen)3]3+ were done at 0.1
M ionic strength to match the electrochemical conditions. For
the singly charged [N(tol)3]•+, the change in potential with ionic
strength was found to be minimal [3 ( 1 mV/log(i)].
3. Kinetics. The rates of oxidation of the phenols have been
monitored by stopped-flow kinetics, following the disappearance
of the oxidant in reactions of [NAr3]•+ or the appearance of
[Fe(N-N)3]2+ in reactions with iron oxidants. Reactions of the
iron complexes were performed in MeCN containing 0.1 M
nBu4NPF6 to match the electrochemical conditions (see above).
When possible, reactions were performed with a large excess
of phenol relative to oxidant (>5 equiv).
The time sequences of optical spectra were globally analyzed
to derive rate constants using SPECFIT software51 (or, in one
instance,49 Microsoft Excel) (Table 3). For thermodynamically
favorable reactions (Keq . 1) run under pseudo-first-order
conditions, the second-order rate constant was taken as the slope
of a plot of kobs vs [HOAr-B]. Particularly fast reactions were
Oxidation of HOAr-NH2 by [N(tol)3]•+ yields an equilibrium
mixture with the phenoxyl radical and the tri-p-tolylamine, with
equilibrium constant K5 (eq 5). Addition of N(tol)3 to the
K5 ) 2.0 ( 0.5
HOAr-NH2 + [N(tol)3]•+ y
z •OAr-NH3
+
+
N(tol)3 (5)
reaction mixture causes an increase in the optical absorbance
due to [N(tol)3]•+, yielding K5 ) 2.4. Alternatively, addition of
aliquots of triflic acid to solutions containing large excesses of
N(tol)3 and HOAr-NH2 versus [N(tol)3]•+ quantitatively pro-
tonates HOAr-NH2 and therefore shifts the equilibrium toward
[N(tol)3]•+. This experiment afforded K5 ) 1.5, and also
established that a second proton transfer equilibrium, eq 6, is
not significant (K6 , 1).49 It should be noted that these
•OAr-NH3+ + HOAr-NH2 &
•OAr-NH2 +
HOAr-NH3 (6)
F
+
analyzed with second-order kinetics, and reactions with Keq
j
1 were analyzed as opposing second-order reactions. In each
case, the rate constant was derived from approximately 25
kinetic runs, at five different concentrations. The temperature
dependence of the rate constants was measured over 30-47 K
ranges (Figure 5), yielding the Eyring parameters52 in Table 4
(below). Variations in driving force over the appropriate
temperature ranges for the reactions of HOAr-NH2 + [N(tol)3]•+
and HOAr-py + [Fe(5,5′-Me2bpy)3]3+ were evaluated by cyclic
voltammetry of the individual reagents. The difference between
the half-reaction potentials measured at 2 and 46 °C (for HOAr-
py and [Fe(5,5′-Me2bpy)3]3+) or 49 °C (for HOAr-NH2 and
[N(tol)3]•+) were found to be within the propagated experimental
error ((30 mV).
equilibrium experiments are possible only because of the
stability of the phenoxyl radical on the chemical time scale.
Together, the equilibria establish an overall equilibrium constant
K5 ) 2.0 ( 0.5, which implies a difference in redox potential
between E(HOAr-NH2+/0) and E([N(tol)3]•+/0) of 18 ( 8 mV.
This is in excellent agreement with the 20 ( 30 mV difference
in the electrochemical E1/2 values: 0.36 ( 0.02 V for HOAr-
+/0
NH2
and 0.38 ( 0.02 V for [N(tol)3]•+/0. The agreement
validates the assignment of the phenol E1/2 values as E(CPET).
The oxidants used in this study include variously substituted
triarylaminium ions [N(p-C6H4X)3]•+, iron(III) tris-polypyridyl
complexes [Fe(N-N)3]3+ (N-N ) 2,2′-bipyridine or 1,10-
To determine the kinetic isotope effects, MeCN solutions of
HOAr-NH2 and HOAr-py were prepared with 0.5-1% v/v
CH3OD and the kinetics performed otherwise as above. The
large molar excess of CH3OD provided high isotopic enrichment
at the exchangeable OH and NH2 positions (the rate constants
were corrected for the residual proton content in the CH3OD).
Control experiments showed that addition of 1% v/v protio-
methanol (CH3OH) does not affect the rate constant for these
reactions. kHOAr-NH /kDOAr-ND ranges from 1.6 ( 0.2 to 2.6 (
(49) See the Supporting Information of ref 24.
(50) (a) Noel, M.; Vasu, K. I. Cyclic Voltammetry and the Frontiers of
Electrochemistry; Aspect: London, 1990; pp 141-143. (b) Braga, T. G.;
Wahl, A. C. J. Phys. Chem. 1985, 89, 5822-5828. (c) Chan, M.-S.; Wahl,
A. C. J. Phys. Chem. 1978, 82, 2542-2549.
(51) Binstead, R. A.; Zuberbu¨hler, A. D.; Jung, B. Specfit, version 3.0.36 (32-
bit Windows); Spectrum Software Associates: Chapel Hill, NC, 2004.
(52) Eyring equation: k ) κ(kBT/h) exp[(T∆Sq - ∆Hq)/RT]; κ is assumed to be
1.
(53) (a) From ref 48a, E(2,4,6-tBu3ArOH) ) 1.85 V in MeCN vs Ag/AgI in
+
+
MeCN and EAg/Ag ) ESCE + 0.365 V. Using EFc/Fc ) ESCE - 0.40 V (ref
t
53b), the potential for Bu3ArOH in MeCN is 1.09 V vs Cp2Fe+/0. (b)
Connelly, N. G.; Geiger, W. E. Chem. ReV. 1996, 96, 877-910.
2
2
9
6080 J. AM. CHEM. SOC. VOL. 128, NO. 18, 2006

Concerted Proton−Electron Transfer in Phenols
A R T I C L E S
Table 3. Rate Constants for Phenol Oxidations (295 ( 2 K, MeCN)
1
phenol
oxidanta
k (M-1 s-
)
Erxnb (V)
HOAr-NH2
[Fe(bpy)3]3+
(4 ( 1) × 106
(1.5 ( 0.2) × 105
(3.8 ( 0.4) × 105
(8 ( 1) × 105
0.34
0.31
0.22
0.16
0.12
0.09
0.02
-0.20
-0.04
-0.04
0.22
[N(p-C6H4Br)3]•+
(4 ( 2) × 107
[Fe(5,5′-Me2bpy)3]3+
[Fe(4,7-Me2phen)3]3+
[N(p-C6H4OMe)(p-C6H4Br)2]•+
[Fe(3,4,7,8-Me4phen)3]3+
[N(tol)3]•+
(3.0 ( 0.3) × 104
(1.1 ( 0.2) × 105
(1.1 ( 0.1) × 103
(2.7 ( 0.3) × 104
(3.2 ( 0.3) × 104
(5.8 ( 0.6) × 104
kH/kD ) 2.6 ( 0.4d
(4.3 ( 0.4) × 104
kH/kD ) 2.5 ( 0.3d
(6.9 ( 0.7) × 102
kH/kD )1.6 ( 0.2d
(1.1 ( 0.1) × 104
(5.2 ( 0.8) × 106
(5.8 ( 0.9) × 105
(1.9 ( 0.4) × 106
(3.3 ( 0.6) × 105
(1.5 ( 0.2) × 106
kH/kD ) 2.8 ( 0.6
(2.3 ( 0.4) × 105
kH/kD ) 2.5 ( 0.6
[N(p-C6H4OMe)3]•+
[N(p-C6H4OMe)2(p-C6H4Br)]•+
[MPT]•+
c
DOAr-ND2
[Fe(5,5′-Me2bpy)3]3+
[N(tol)3]•+
0.02
[N(p-C6H4OMe)3]•+
-0.20
HOAr-im
HOAr-py
[N(p-C6H4OMe)3]•+
[Fe(bpy)3]3+
-0.26
0.12
0.00
-0.05
-0.12
0.12
[Fe(5,5′-Me2bpy)3]3+
[Fe(4,7-Me2phen)3]3+
[Fe(3,4,7,8-Me4phen)3]3+
[Fe(bpy)3]3+
DOAr-pyc
[Fe(5,5′-Me2bpy)]3]3+
0.00
b
a Reactions with [Fe(N-N)3]3+ were performed in 0.1 M Bu4NPF6/MeCN. Erxn ) E1/2(oxidant) - E1/2(phenol). c Reactions with deuterated substrates
were performed in 0.5-1% v/v CH3OD in MeCN. Rate constants are corrected for residual proton content using kexpt ) kD(1 - fH) + fHkH, where fH is the
d
fraction protonated. kHOAr-NH /kDOAr-ND
.
2
2
Scheme 3. Thermochemical Cycle Indicating the Effect of
Hydrogen Bonding on Redox Potentials
Figure 5. Temperature dependence of the rate constants for HOAr-py +
[Fe(5,5′-Me2bpy)3]3+ (2), HOAr-NH2 + [N(tol)3]•+ (b), and HOAr-im
+ [N(C6H4OMe)3]•+ (9). The curve fits are to the non-adiabatic form of
the Marcus equation (eq 15; see Discussion).
0.4 and kHOAr-py/kDOAr-py ) 2.5 ( 0.6 or 2.8 ( 0.6, depending
on the oxidant (Table 3).
Discussion
Phenols with an intramolecular hydrogen bond react with one-
electron oxidants to generate phenoxyl radicals in which the
proton has transferred. The phenols and phenoxyl radicals have
been characterized by spectroscopy, cyclic voltammetry, and
chemical reactivity. The kinetics of oxidation have been
examined for a number of phenol-oxidant pairs, in one case
electrochemically. We discuss first the thermochemistry of outer-
sphere oxidation of these phenols, and then the mechanistic data
that implicate a concerted proton-electron transfer (CPET)
pathway for the reactions. Finally, analysis of the CPET rate
constants indicates that the classical Marcus theory is an
excellent starting point to understand these processes.
has Ep,a(tBu3ArOH•+/0) ) 1.09 V.53 Such large shifts have been
suggested to be due to hydrogen-bonding effects, but the analysis
below shows that the shifts must be due to proton transfer.47
Consider the oxidation of a phenol hydrogen-bonded to a base
B by electron transfer without proton transfer. The effect of
the hydrogen bonding is illustrated by the thermochemical cycle
in Scheme 3a, which compares the potentials for the H-bonded
+/0
+/0
and non-H-bonded forms (EArOH-B , EArOH ). The difference
between these two potentials is equal to the difference in the
strengths of the hydrogen bonds in the reduced and oxidized
forms (∆GHB/red - ∆GHB/ox) (eq 7). Note that the absolute
hydrogen bond strengths are not important, only the change upon
1. Phenol Potentials. The potentials for HOAr-B (0.36-
0.58 V vs Cp2Fe+/0 in MeCN, Table 1) are significantly lower
than reported values for one-electron oxidation of phenols
without a pendant base. 2,4,6-Tri-tert-butylphenol, for example,
oxidation. Since most hydrogen bonds are 3-8 kcal mol-1 54
,
shifts of more than ca. 5 kcal mol-1 (0.2 V) would be quite
9
J. AM. CHEM. SOC. VOL. 128, NO. 18, 2006 6081

A R T I C L E S
Rhile et al.
Scheme 4. Three Possible Mechanisms for Oxidation of
Phenol-Base Compounds
unusual. A very recent experimental study in general supports
these thermochemical arguments.55
The effect of the hydrogen bonding on a CPET redox
potential is similar. Progressing around the thermochemical cycle
in Scheme 3b, (1) the hydrogen bond in HOAr-B is broken
(-∆GHB/red); (2) the non-hydrogen-bonded phenol is oxidized
+/0
(EArOH ); (3) the proton is transferred (-RT∆pKa); and (4) the
hydrogen bond of •OAr-BH+ is formed (∆GHB/ox). The sum of
these four steps is equal to the overall potential (eq 8). For the
+/0
sterically crowded phenols discussed here, EArOH is probably
t
well approximated by the Bu3ArOH•+/0 potential.56 With this
•
approximation, the difference between the potential for OAr-
t
HB+/HOAr-B and that for Bu3ArOH•+/0 is the energetics of
the proton-transfer step (-RT∆pKa) plus the difference in
hydrogen bond strengths. As in the pure electron transfer case
of Scheme 3a, it is the change in H-bond strengths rather than
their absolute value that is important.
In sum, the difference of 0.5-0.7 V in redox potentials for
HOAr-B vs tBu3ArOH•+/0 is too large to be due to changes in
hydrogen bond strength. This difference is primarily due to the
proton transfer from the phenol radical cation to the base, step
3 in Scheme 3b. In MeCN, 2,4,6-tri-tert-butylphenol radical
cation has a pKa of ca. 060 and protonated benzylamine has a
pKa of 17,38 yielding a ∆pKa of 17. This provides a crude
prediction of a shift of 1 V (∆E ) 0.059 V × ∆pKa), somewhat
larger than the observed 0.73 V difference between E(•ArO-
NH3+/HOAr-NH2) vs E(tBu3ArOH•+/0).
2. Mechanistic Analysis. There are three reasonable mech-
anisms for the one-electron oxidation of the hydrogen-bonded
phenols (Scheme 4). Rate-limiting outer-sphere electron transfer
could yield the phenolic radical cation (•+HOAr-B), which
would be followed by fast proton transfer to give the product
(ET1-PT1). Radical cations of simple phenols are well-
established transients, particularly in photochemical processes.61
Alternatively, pre-equilibrium proton transfer to yield the
zwitterion (-OAr-BH+) could be followed by electron transfer
(PT2-ET2). Zwitterions such as -OAr-BH+ are well known
in phenol-base chemistry, particularly when the phenolic
portion is highly acidic, as in p-nitrophenols.39,40,62 Rate-limiting
proton transfer is ruled out because different oxidants react at
different rates and because PT between electronegative elements
in general occurs at very fast rates.63 The defining characteristic
of the stepwise mechanisms is the formation of an intermediate
with a finite lifetime. The third mechanism is the concerted
transfer of both particles, CPET, defined by the absence of an
intermediate along the reaction coordinate. “Concerted” implies
that both particles move in a single kinetic step but does not
imply synchronous movement of the proton and electron.
There are three experimental markers indicating the oxidation
mechanism as CPET. First, isotope effects on the oxidation of
DOAr-ND2 and DOAr-py (1.6-2.8 depending on oxidant and
The change in hydrogen bond strength and the attendant shift
of the redox potential is likely to be quite small. As noted above,
the hydrogen-bonded phenoxide -OAr-NH2 has the same
potential as the non-H-bonded 2,4,6-tBu3ArO-. The potentials
for tBu3ArOH (+1.09 V), the anisole MeOAr-NH2 (∼1.2 V),
and the hydroxyphenol HOAr-OH (∼1.1 V) are quite similar,
despite what are likely very different H-bonds (OH‚‚‚NCMe,
O‚‚‚HN, and OH‚‚‚OH). Hammarstro¨m and co-workers attribute
0.10 V (2 kcal mol-1) to the change in hydrogen bonding for
the tyrosine-histidine pair in their model system.57 Phenoxyl
radicals are known to make strong hydrogen bonds in some
systems,58 so the H-bond in •OAr-BH+ could be stronger than
that in HOAr-B, but this effect is usually small. The H-bond
strengthening upon oxidation of catechols to oxyl radicals has
been variously estimated as between ∼4 and <1 kcal mol-1 58
.
The strengthening in catechols and 1,8-naphthalenediols is
particularly large because oxidation yields hydrogen bonds in
which PT is degenerate; one report describes a ∼7 kcal mol-1
(0.3 eV) strengthening for 1,8-naphthalenediols.58d For the case
of HOAr-NH2, the hydrogen bond could be even stronger in
the neutral phenol because of the much larger pKa mismatch
between donor and acceptor in the radial cation.59 This would
shift the potential in the opposite direction.
(54) March, J. AdVanced Organic Chemistry, 4th ed.; Wiley: New York, 1992;
p 76.
(55) Kanamori, D.; Furukawa, A.; Okamura, T.; Yamamoto, H.; Ueyama, N.
Org. Biomol. Chem. 2005, 3, 1453-1459.
(56) For sterically encumbered phenols such as tBu3ArOH, hydrogen bonding
to solvent is a small effect. For instance, 2,6-di-tert-butyl-4-methylphenol
(BHT) is only 14% hydrogen-bonded in MeCN: Wren, J. J.; Lenthen, P.
M. J. Chem. Soc. 1961, 2557-2560.
(57) Sjo¨din, M.; Styring, S.; Åkermark, B.; Sun, L.; Hammarstro¨m, L. Philos.
Trans. R. Soc. London B 2002, 357, 1471-1479.
(58) (a) Lucarini, M.; Pedulli, G. F.; Guerra, M. Chem. Eur. J. 2004, 10, 933-
939. (b) Lucarini, M.; Mugnaini, V.; Pedulli, G. F.; Guerra, M. J. Am.
Chem. Soc. 2003, 125, 8318-8329. (c) Amorati, R.; Lucarini, M.;
Mugnaini, V.; Pedulli, G. F. J. Org. Chem. 2003, 68, 5198-5204. (d) DFT
calculations suggest an H-bond strengthening of 8.6 kcal mol-1 for
4-methoxy-1,8-naphthalenediol, but experimental results indicate that this
is overestimated by as much as 2 kcal mol-1. Foti, M. C.; Barclay, L. R.
C.; Ingold, K. U. J. Am. Chem. Soc. 2002, 124, 12881-12888.
(59) Hydrogen bond strengths have been shown to increase with decreasing
difference in pKa of the hydrogen bond donor and acceptor (∆pKa: Shan,
S.-O.; Herschlag, D. Proc. Natl. Acad. Sci. U.S.A. 1996, 93, 14474-14479).
This value is smaller in the phenol-amine (pKa(phenol) - pKa(amine) ≈ 27 -
(60) (a) Reference 48a. (b) The difference between the phenol pKa values in
DMSO and MeCN is taken as 9.5 units according to the following:
Chantooni, M. K., Jr.; Kolthoff, I. M. J. Phys. Chem. 1976, 80, 1306-
1310.
(61) (a) Brede, O.; Hermann, R.; Karakostas, N.; Naumov, S. Phys. Chem. Chem.
Phys. 2004, 6, 5184-5188. (b) Ganapathi, M. R.; Hermann, R.; Naumov,
S.; Brede, O. Phys. Chem. Chem. Phys. 2000, 2, 4947-4955.
(62) (a) Rospenk, M.; Fritsch, J.; Zundel, G. J. Phys. Chem. 1984, 88, 321-
323. (b) Koll, A.; Rospenk, M.; Sobczyk, L. J. Chem. Soc., Faraday Trans.
1 1981, 77, 2309-2314. (c) Rospenk, M. J. Mol. Struct. 1990, 221, 109-
114.
+
17 ) 10, vs the radical cation pKa(PhOH ) - pKa(amine) ≈ 17 - 0 ) 17. The
case is opposite for HOAr-py, where the relevant pKa’s are 27 (PhOH),
12 (pyridine), and 0 (PhOH+) (MeCN pKa data from refs 38 and 60).
(63) Bell, R. P. The Proton in Chemistry; Cornell University Press: Ithaca, NY,
1973.
9
6082 J. AM. CHEM. SOC. VOL. 128, NO. 18, 2006

Concerted Proton
−
Electron Transfer in Phenols
A R T I C L E S
Scheme 5. Putative Kinetic Scheme for ET-PT Involving
phenol) can only be explained through CPET. In rate-limiting
electron transfer (ET1), no bond is made or broken in the ET
step and, like other electron transfers,64a,65 there would be only
a small secondary isotope effect. The proton-first pathway PT2-
ET2 would have an equilibrium PT isotope effect, which would
also be small.
Precursor and Successor Complexes, Assuming Kprecursor
)
Ksuccessor ) 1 M-1 and k-ET1 ) kPT1 ) 1013 s-1 as the Best-Case
Scenario
Second, the intermediates in the stepwise mechanisms appear
to be too high in energy to be involved in the reactions. In the
ET1-PT1 mechanism for HOAr-NH2 + [N(tol)3]•+, the ET1
step (eq 9) is estimated to have EET1 ) -0.71 V (∆GET1 ) 16
kcal mol-1, KET1 ) 10-12). This estimate uses E(•+HOAr-NH2/
kET1
HOAr-NH2 + Ntol3
8 •+HOAr-NH2 +
Ntol3 EET1 ) -0.71 V (9)
•+
librium constants Kprecursor ) Ksuccessor ) 1 M-1, as is commonly
done.67 No evidence for a precursor or successor complex is
evident in optical spectra of reaction mixtures. The ET1 step
generates the successor complex in which the proton has not
transferred, [•+HOAr-NH2|Ntol3]. In the scenario most favorable
to the ET1-PT1 pathway, this intermediate partitions equally
HOAr-NH2))1.09V,takentobethesameasE(tBu3PhOH+/0 53
)
and E(HOAr-OH+/0) and 0.11 V below E(MeOAr-NH2+/0) (as
noted above, hydrogen-bonding effects are likely to be small).
The estimated value of ∆GET1 is significantly higher than the
observed Eyring barrier for this reaction, ∆Gq ) 11 kcal mol-1
(from k ) 1.1 × 105 M-1 s-1 and the Eyring equation with κ
) 1;52 with smaller prefactors or κ < 1 the discrepancy would
be larger66,67). For •+HOAr-NH2 to be a viable intermediate
(eq 9), our estimate of the potential would have to be in error
by more than 0.2 V. (This would also predict a dependence on
driving force different than what is observed, as described
below.) This analysis can equivalently be framed in terms of
rate constants instead of barriers. The values KET1 ) kET1/k-ET1
) 10-12 (see above) and kET1 ) kobs ) 105 M-1 s-1 would imply
an impossible k-ET1 ) 1017 M-1 s-1, much faster than the
diffusion limit in MeCN.68 Similar arguments hold for the
HOAr-py and HOAr-im systems.
between back electron transfer to the precursor complex (k-ET1
)
and forward proton transfer (kPT1) to [•OAr-NH3 |Ntol3], with
both occurring at the fastest possible rate, ∼1013 s-1. With these
assumptions, kET1 would have to be 2 × 105 M-1 s-1, based on
the experimentally observed kobs ) 1 × 105 M-1 s-1 and K5 )
2 (eq 5, above). This requires KET1 for ET within the precursor
complex (kET1/k-ET1) to be 2 × 10-8, more than 104 larger than
the KET1 based on the estimated redox potentials (see above).
And even in this best-case scenario, [•+HOAr-NH2|Ntol3] is
barely an intermediate, since the 1013 s-1 rate constants imply
a half-life of only 35 fs, roughly one vibrational period for a
1000 cm-1 mode. The reaction with [N(p-C6H4OCH3)3]•+
provides even tighter constraints, because ET1 becomes an
additional 0.22 V uphill (KET1 is less favorable by 5 × 10-3),
but kobs only changes by 10-2. The constraints on KET1 are also
more stringent if the back electron or proton transfers have any
+
A more complete analysis of the ET1-PT1 pathway includes
precursor and successor complexes, as illustrated for the HOAr-
NH2 + [N(tol)3]•+ reaction in Scheme 5. The formation of the
precursor and successor complexes is assumed to have equi-
barrier and are slower than the maximal 1013 s-1
.
It should be added that the most recent computational and
experimental reports conclude that similar intermolecularly
hydrogen-bonded [PhOH|base]•+ species are not minima in the
gas phase (proton transfer from O to the base proceeds without
barrier).69 If this is also the case for the solution ArOH-B•+
species discussed here, they cannot, by definition, be intermedi-
ates.
The PT2-ET2 pathway is also very unlikely, based on
thermochemical arguments similar to those above. The rate
constant for PT2-ET2 is the product of the equilibrium constant
for initial proton transfer (KPT2, eq 1, above) and the ET rate
constant from the zwitterion (kET2, eq 10). KPT2 might be
(64) Eberson, L. Electron Transfer Reactions in Organic Chemistry; Springer-
Verlag: New York, 1987; (a) pp 77-79, (b) pp 51-52, (c) p 27, and (d)
pp 30-34 (using the radii in ref 71).
(65) (a) Buhks, E.; Bixon, M.; Jortner, J. J. Phys. Chem. 1981, 85, 3763. (b)
Smaller secondary isotope effects have been observed for related electron
transfers, e.g. Gould, I. R.; Farid, S. J. Am. Chem. Soc. 1988, 110, 7883-
5.
(66) Pre-exponential factors from 1011 to 1013 M-1 s-1 have been used for
bimolecular adiabatic electron transfer reactions, values described as a
collision frequency or a rate of crossing at the transition state64,66a,67 (at
298 K, the Eyring prefactor kBT/h is 6 × 1012 s-1). One recent paper66a
describes the choice of prefactor as related to whether solvent or inner-
sphere motions are dominant in the reorganization energy, a level of detail
that is not yet available for CPET processes. In this mechanistic section,
the analysis uses the Eyring equation as is typical in mechanistic chemistry;
in the Marcus theory section that follows, the most typical67 Z ) 1011 M-1
s-1 is used. Using Z ) 1011 M-1 s-1 in this mechanistic context would
give ∆GqZ)1011 ) 8 kcal mol-1 and would strengthen the argument against
the PT1-ET1 mechanism; the estimate of ∆GET1 would have to be off by
more than 0.3 V. Thus, the use of the Eyring equation here is a conservative
(worst-case) choice for this argument. (a) Hamann, T. W.; Gstrein, F.;
Brunschwig, B. S.; Lewis, N. S. J. Am. Chem. Soc. 2005, 127, 13949-
13954.
KPT2
kET2
HOAr-NH2 y z •OAr-NH3
8 •OAr-NH3 (10)
+
+
expected to be ∼10-9 on the basis of the difference in pKa of
amines (∼18) and phenols (∼27) in MeCN.38 Since this ignores
potential electrostatic interactions in the zwitterion, we instead
(67) The Marcus equation is often applied to bimolecular reactions using 1011
M-1 s-1 as the prefactor (rather than the Eyring kBT/h): k ) (1011 M-1
s-1) exp(-∆Gq/RT)67a-d; KP is typically assumed to be ∼1 M-1, with work
terms added where appropriate.67a-d,71 (a) Marcus, R. A.; Sutin, N. Biochim.
Biophys. Acta 1985, 811, 265-322. (b) Sutin, N. Prog. Inorg. Chem. 1983,
30, 441-499. (c) Sutin, N. Acc. Chem. Res. 1982, 15, 275-282. (d) Nelsen,
S. F.; Pladziewicz, J. R. Acc. Chem. Res. 2002, 35, 247-254.
(68) Cf.: (a) McClelland, R. A.; Kanagasabapathy, J. V. M.; Banait, N. S.;
Steenken, S. J. Am. Chem. Soc. 1992, 114, 1816-1823. (b) de Carvalho,
I. M. M.; Gehlen, M. H. J. Photochem. Photobiol. A: Chem. 1999, 122,
109-113. (c) Kikuchi, K.; Sato, C.; Watabe, M.; Ikeda, H.; Takahashi,
Y.; Miyashi, T. J. Am. Chem. Soc. 1993, 115, 5180-5184.
(69) (a) Fang, Y.; Liu, L.; Feng, Y.; Li, X.; Guo, Q. J. Phys. Chem. A 2002,
106, 4669-4678. (b) Feng, Y.; Liu, L.; Fang, Y.; Guo, Q. J. Phys. Chem.
A 2002, 106, 11518-11525. (c) Wang, Y.; Eriksson, L. A. Int. J. Quantum
Chem. 2001, 83, 220-229. (d) O’Malley, P. J. J. Am. Chem. Soc. 1998,
120, 11732-11737. (e) Kim, H.; Green, R. J.; Qian, J.; Anderson, S. L. J.
Chem. Phys. 2000, 112, 5717-5721. (f) A minimum for a [HOAr-im]•+
species in solution has been calculated in ref 30b.
9
J. AM. CHEM. SOC. VOL. 128, NO. 18, 2006 6083

A R T I C L E S
Rhile et al.
use the more conservative experimental value, KPT2 < 10-4
,
derived from UV-vis spectra (see above). The value kobs
>
107 M-1 s-1 for HOAr-NH2 + [N(p-C6H6Br)3]•+ then implies
kET2 > 1011, faster than the diffusion limit in MeCN.68 An
analogous argument can be made with barrier heights and with
precursor and successor complexes.
The third argument for the CPET mechanism is the depen-
dence of barrier on driving force. For HOAr-NH2 + [NAr3]•+,
a plot of ∆∆Gq vs ∆∆G° has a slope 0.53 [∆∆G° ) nF(E1 -
En) and ∆∆Gq ) RT ln(k1/kn)]. Following the discussion above,
the stepwise path with rate-limiting ET1 would require ∆Gq
ET1
≈ ∆G°ET1 and therefore that ∆∆Gq/∆∆G° = 1 (∆G° ≈ λ in
the Marcus picture; see below).67 The initial PT2 mechanism
requires that ∆Gq
≈ 0, so ∆∆Gq/∆∆G° = 0 (-∆G° ≈ λ).
ET2
As discussed in the next section, ∆∆Gq/∆∆G° ) 0.53 is close
to the value of 0.5 predicted by Marcus theory for the concerted
process in this |∆G°CPET| , 2λ situation. The dependence of
barrier on driving force has previously been used by Okamura
et al. to discuss stepwise vs concerted PCET pathways.20
In sum, the isotope effects, the thermochemistry, and the
dependence of the rate constants on driving force are consistent
only with a concerted mechanism. These conclusions are
consistent with the findings of Linschitz, Hammarstro¨m, and
Nocera, who have all found CPET mechanisms for their
systems, which include both aqueous and nonaqueous media.7-9
Figure 6. log(k) vs Erxn for oxidations (a) of HOAr-NH2 by [NAr3]•+ (b)
and (b) of HOAr-py by [Fe(Mexphen)3]3+ (2) and [Fe(R2bpy)3]3+ (9) and
HOAr-NH2 by [Fe(Mexphen)3]3+ (1) and [Fe(R2bpy)3]3+ ([). The curves
are fits to eq 12.
3. Analysis Using Marcus Theory. From one perspective,
the phenol-bases HOAr-B are simply outer-sphere electron
transfer reagents, so Marcus theory may be appropriate to
analyze these reactions. However, the inner-sphere reorganiza-
tions for HOAr-B/•OAr-BH+ are unusual because they involve
not only small shifts in equilibrium bond distances, as in the
standard Marcus picture, but also movement of a proton across
an OH‚‚‚N hydrogen bond. The proton can be thought of as
transferring ∼0.7 Å between two minima on an adiabatic
potential energy surface.70 This would not seem to fit easily
into the standard Marcus model, where a single parabolic
surface, defined by the reorganization energy λ, describes all
of the solvent and inner-sphere reorganizations. Current, more
refined theoretical formulations of CPET treat the proton transfer
explicitly but are more complicated and require more parameters
than are readily determined experimentally.23 So, despite its
simplifications, the adiabatic Marcus equation is still the logical
starting point because it predicts barriers and rate constants using
only the two parameters λ and ∆G° (eqs 11 and 12).67,71
λ
4
∆G° 2
λ
∆Gq ) 1 +
(11)
(
)
rxn/λ)]2
k ) [1011 M-1 s-1] e-(λ/4RT)[1+(-FE
(12)
(13)
1
λ12 ) (λ11 + λ22)
2
The kinetic and thermochemical data for HOAr-B + oxidant
reactions (Table 3) can be fit by eq 12, as shown in Figure 6.72
The data sets are also well fit by straight lines; as noted above,
the slope ∆∆Gq/∆∆G° for the HOAr-NH2 + [NAr3]•+ reactions
is 0.53. The limited range of experimentally accessible driving
forces (0.51 V) does not provide a test of the predicted parabolic
dependence of log(k) on ECPET
.
The reorganization energies for the oxidations of HOAr-NH2
derived from eq 12 are 34 ( 1 kcal mol-1 (1.5 eV) for the
reactions with [NAr3]•+, 38 ( 2 kcal mol-1 (1.9 eV) for the
reactions with [Fe(N-N)3]3+ (N-N ) R2bpy, Mexphen), and
35 ( 1 kcal mol-1 (1.5 eV) from the single rate constant
oxidation by 10-methylphenothiazinium (MPT•+) (Table 4).
Following the additivity postulate (eq 13), each of these cross
reaction values (λ12) is the average of the λ’s for the individual
self-exchange reactions A+ + A f A + A+. The λ11 for N(tol)3/
[N(tol)3]•+ self-exchange in MeCN is 12 kcal mol-1 (0.5 eV),43
which is taken as characteristic of the series of NAr3•+ oxidants
used here. λ11 for MPT•+/0 is similar (9 kcal mol-1, 0.4 eV);73
that for [Fe(bpy)3]3+/2+ in MeCN is twice as large (24 kcal
mol-1, 1.0 eV).74 Using eq 13, these values yield reorganization
energies for HOAr-NH2/•OAr-NH3+ self-exchange of 53 ( 3,
(70) (a) O‚‚‚N distances in these and related structures are in the range 2.53-
2.65 Å.33 With typical O-H and N-H distances of 0.96 and 1.01 Å,70a
the distance between proton positions in OH‚‚‚N vs O‚‚‚HN tautomers
should be 0.56-0.68 Å in a linear hydrogen bond. In the bent structures
likely for HOAr-B, the distances should be ∼0.7-0.8 Å. (b) CRC
Handbook of Chemistry and Physics, 54th ed.; Weast, R. C., Ed.; CRC
Press: Cleveland, OH, 1973; pp F198-F199.
(71) (a) An electrostatic correction to ∆G° for the Marcus analysis64c has been
included for the HOAr-B + [Fe(N-N)3]3+ reactions: ∆G°′ - ∆G° ) (Z1
- Z2 - 1)(331.2f)/(Dr12) ) 0.76 kcal mol-1, or 0.03 eV; Z1 ) 0; Z2 ) +3;
f ) 0.60 (for 0.1 M ionic strength);64c D ) 37.5; r12 ) 13.9 Å. The radius
of HOAr-B was calculated from the lowest crystallographic volume ∼1400
Å3 and V ) (4/3)πr3, yielding r ) 6.9 Å. The radius of [Fe(bpy)3]3+ is
assumed to be 7.0 Å.71b For the HOAr-B + [NAr3]•+ reactions, one of the
species on each side of the equation is uncharged; therefore, the correction
is zero. (b) Schlesener, C. J.; Amatore, C.; Kochi, J. K. J. Am. Chem. Soc.
1984, 106, 3567-3577.
(72) Fitting of k vs T data to eqs 12 and 15 implicitly assumes that λ is constant
with temperature. For λ to change significantly over such a small
temperature range (30-50 K) would require λ to have a very large entropic
component (∆λ ) ∆[∆Hλ°] - ∆[T∆Sλ°] ≈ ∆T[∆Sλ°], which would be
e1 kcal mol-1, even if ∆Sλ° were 20 cal K-1 mol-1
.
9
6084 J. AM. CHEM. SOC. VOL. 128, NO. 18, 2006

Concerted Proton−Electron Transfer in Phenols
A R T I C L E S
Table 4. Activation Parameters, Adiabatic Reorganization Energies, and Apparent Non-adiabatic Reorganization Energies and Hrp Values
for Phenol-Base Oxidationsa
c
d
e
e
reaction
λ
12(E)b
λ
12(T)b
λ11
[
λ
12(non-ad)]d
[Hrp
]
∆
Hq
∆
Sq
HOAr-NH2 + [N(Ar)3]•+ f
HOAr-NH2 + [Fe(N-N)3]3+ g
HOAr-NH2 + [MPT]•+
34 ( 1
38 ( 2
35 ( 1i
27 ( 1
22 ( 1
25 ( 2i
33 ( 1
53 ( 3
52 ( 4
58 ( 3i
30 ( 3
23 ( 3
36 ( 4i
[29.6 ( 1.6]
[10 ( 4]
6.3 ( 0.4
-14.1 ( 1.3
h
h
h
h
h
h
h
h
h
h
HOAr-py + [Fe(R2bpy)3]3+ g,j
HOAr-py + [Fe(Mexphen)3]3+ g
HOAr-im + [N(anisyl)3]•+ k
27 ( 1
h
25 ( 2
[23.2 ( 1.6]
h
[17 ( 3]
[6 ( 2]
h
[4 ( 3]
4.7 ( 0.4
h
7.0 ( 0.7
-16.0 ( 1.2
h
-17 ( 3
a ∆Hq and λ in kcal mol-1, ∆Sq in cal K-1 mol-1 (eu), and Hrp in cm-1
.
b λ12(E) and λ12(T) are the adiabatic reorganization energies calculated from the
c
dependence of k on either E° or T using eq 12 (Figures 5 and 6). λ11 is the adiabatic reorganization energy for HOAr-B/•OAr-HB+ self-exchange from
d
eq 13 [using the average of λ12(E) and λ12(T)]. λ12(non-adiabatic) and Hrp from eq 15, which may not be appropriate; see text. e Determined by fitting kobs
vs T data (Figure 5) to the Eyring equation.52 f Temperature-dependent results for HOAr-NH2 + [N(tol)3]•+, 280 e T e 327 K. g Corrected for work terms
following refs 64 and 71. h Not determined. i From a single phenol-oxidant pair. j Temperature-dependent results for HOAr-py + [Fe(Me2bpy)3]3+, 279 e
T e 318 K. k Anisyl ) -C6H4OMe; 279 e T e 309 K.
53 ( 4, and 58 ( 3 kcal mol-1 (2.4 ( 0.2 eV). These are the
same within experimental error, which is an indication that it
is appropriate to use the adiabatic Marcus equation to analyze
these reactions. The single rate constant for HOAr-im + [N(p-
C6H4OMe)3]•+ gives λ11 ) 36 ( 3 kcal mol-1 (1.4 eV) for
CPET self-exchange.
made. This agreement, between two different kinds of analysis
and involving mostly independent data sets, supports the use
of the Marcus equation for these CPET reactions. Another
correct prediction of the Marcus treatment is that ∆∆Gq/∆∆G°
) 0.5 + ∆G°/2λ. Using the values in Tables 3 and 4 for the
five HOAr-NH2 + NAr3•+ reactions, 0.5 + ∆G°/2λ ranges from
0.39 to 0.57, with an average value of 0.49. This is in good
agreement with the experimental linear fit, ∆∆Gq/∆∆G° ) 0.53.
The electrochemical rate constant kel provides an additional
test of the applicability of the adiabatic Marcus treatment. While
rigorous comparison of heterogeneous and homogeneous elec-
tron transfer kinetics is complex, there is often a good
correspondence between kel and the homogeneous self-exchange
rate constant k11, via eq 14.44a,77 Equation 14 follows from the
The rate constants for oxidation of HOAr-py fall on different
lines for oxidants [Fe(R2bpy)3]3+ (R ) H, Me) vs [Fe(Mex-
phen)3]3+ (Mex ) 4,7-Me2, 3,4,7,8-Me4). The λ12 values are 27
( 1 kcal mol-1 (R2bpy oxidants) and 22 ( 1 kcal mol-1 (Mex-
phen oxidants). The distinction is surprising because of the
similarity of these oxidants. The self-exchange rate constants
for these species are similar, with the phenanthroline derivatives
reacting ca. 3 times faster, although this comparison is com-
plicated by scatter among the different derivatives, counterion
and ionic strength effects, etc.lc Taking λ11 = 24 and 21 kcal
mol-1 for [Fe(R2bpy)3]3+/2+ and [Fe(Mexphen)3]3+/2+, respec-
tively, yields apparent λ11(HOAr-py) values of 30 ( 3 and 23
( 3 kcal mol-1 for the different oxidants. These values are
different, as indicated by the distinct lines in Figure 6. This
discrepancy suggests a deviation from the adiabatic Marcus
treatment, perhaps due to non-adiabaticity or to ion-pairing
issues, as will be probed in future work.75 The oxidations of
HOAr-NH2 do not show such a distinction between reactions
with [Fe(R2bpy)3]3+ vs [Fe(Mexphen)3]3+, although there is some
scatter in the rate constants (bottom curve of Figure 6b). Fitting
these rate constants separately yields cross reaction λ12 and
HOAr-NH2 self-exchange λ11 values that agree within error:
k11
kel
)
(14)
1011 M-1 s-1 104 cm-1 s-1
x
assumption that a given reagent has similar intrinsic barriers
for homogeneous and heterogeneous electron transfer. The rate
constants are divided by the different collision frequencies, and
the self-exchange k11 appears as a square root because it involves
two molecules and therefore two intrinsic barriers. k11 for
HOAr-NH2 has not been directly determined but is calculated
to be 8 M-1 s-1 using the adiabatic Marcus equation (eq 12)
with λ11 ) 55 kcal mol-1 (the average of the three experimen-
tally derived values in Table 4). Then (k11/1011)1/2 ) 9 × 10-6
,
a factor of 20 larger than (kel/104) ) 3 × 10-7. This is good
agreement given the approximate nature of eq 14 and that the
kel of 3 × 10-3 cm s-1 lies on the cusp of the conditions where
eq 14 holds s according to Swaddle, only for kel e 10-2 cm
[Fe(R2bpy)3]3+, λ12 ) 39 ( 2, λ11 ) 53 ( 4; [Fe(Mexphen)3]3+
λ12 ) 37 ( 2, λ11 ) 53 ( 4 kcal mol-1
,
.
Reorganization energies can also be derived from the tem-
perature dependence of the rate constants using eq 12.76 These
are given as λ12(T) in Table 4, to distinguish them from the
s-1 77b
.
The results reported here are among the first confirmations
that the adiabatic Marcus equation is applicable to this class of
CPET reactions, in which the proton and electron are clearly
separated in the reactants or products. We and others have used
Marcus theory for CPET reactions, assuming its applicability.8,24
The tests described here are the equivalence of the intrinsic
barriers derived from the dependence on driving force, from
the dependence on temperature, and from different reagents,
and the agreement between electrochemical and solution rate
constants. It should be noted that these tests are not especially
stringent and that there is the possibility of a deviation in the
difference λ’s derived for HOAr-py with the different iron
oxidants, which will be explored in more detail in future work.75
Hammarstro¨m and co-workers have previously shown that
reorganization energies derived from the dependence on ECPET
λ12(E). The λ12(E) and λ12(T) values are the same within
experimental error for the three cases where comparisons are
,
(73) Kowert, B. A.; Marcoux, L.; Bard, A. J. J. Am. Chem. Soc. 1972, 94, 5538-
5550.
(74) (a) λ22([Fe(bpy)3]3+) was calculated from the self-exchange ratelc using the
adiabatic Marcus equation with Z ) 1011, taking λ ) 0.25∆Gq and a
[Fe(bpy)3]3+ self-exchange rate. The self-exchange rate for [Fe(41,41 ′-Me2-
bpy)3]3+ is 2 times faster than that for the bipy derivative,lc so it is likely
that [Fe(5,5′-Me2bpy)3]3+ would also be slightly faster.
(75) Markle, T. F.; Rhile, I. J.; Nagao, H.; DiPasquale, A. G.; Mayer J. M.,
manuscript in preparation and work in progress.
(76) The calculation of an adiabatic λ from the temperature-dependent rate data
assumes that the equilibrium constants for formation of the precursor and
successor complexes are 1 M-1 at all temperatures; see, however, the
discussion of non-adiabatic character below.
(77) (a) Marcus, R. A. J. Phys. Chem. 1963, 67, 853-857. (b) Fu, Y.; Swaddle,
T. W. J. Am. Chem. Soc. 1997, 119, 7137-7144.
9
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A R T I C L E S
Rhile et al.
other words, initial pure ET or PT is disfavored because of the
high energy of the intermediate that would be formed. A similar
argument has been advanced by Hammarstro¨m et al. for aqueous
CPET reactions.8
Extrapolation of these conclusions to a specific biological
system requires caution because typically the driving forces for
CPET, pure ET, and pure PT are not known. The local dielectric
constant and nearby protein residues can substantially affect
these values (and the intrinsic barriers). When initial ET or PT
is energetically competitive with CPET, as found for instance
for quinone reductions in PS I, stepwise pathways are favored.20
However, the concerted mechanism likely occurs in many
situations when pure ET and pure PT are high in energy.
4. Adiabatic vs Non-adiabatic Electron Transfer. The
discussion above has utilized the adiabatic Marcus equation,
but many electron transfer reactions are non-adiabatic. Current
theoretical descriptions of CPET use non-adiabatic formalisms.23
In a non-adiabatic reaction, there is a low probability of crossing
from the reactant to product diabatic surfaces when the system
reaches the transition structure (transmission coefficient κ ,
1). In adiabatic reactions, the system is well described by a
ground-state potential energy surface with κ ≈ 1. The matrix
element Hrp is a measure of the interaction of the surfaces and
appears in the pre-exponential of the non-adiabatic Marcus
equation (eq 15). Values of Hrp less than ∼200 cm-1 (∼kBT)
normally indicate a non-adiabatic reaction.78
Figure 7. Marcus potential energy surfaces for (a) CPET, with a favorable
∆G°CPET but a large λCPET, in contrast to (b) initial ET (to be followed by
PT), with an unfavorable ∆G°ET but a smaller λET
.
similar λ’s are derived from driving force and temperature-
dependent measurements.8
The 56 kcal mol-1 (2.4 eV) reorganization energy for HOAr-
+
NH2/•OAr-NH3 self-exchange in MeCN is quite a large
value.64b,67d HOAr-NH2 is fairly close in size to N(tol)3 and
has the same charge, yet the phenol-amine CPET λ11 is 4.7
times that of the triarylamine: 56 vs 12 kcal mol-1. The 12
kcal mol-1 value for N(tol)3/[N(tol)3]•+ is typical of λ’s for outer-
sphere electron transfer by aromatic organic molecules, usually
k )
4π2Hrp
2
-{([λ(non-ad)]+∆G°)2/4[λ(non-ad)]kBT}
KP
e
h 4π[λ(non-ad)]k T
x
B
(15)
e20 kcal mol-1 64b
HOAr-NH2 has a much higher intrinsic
.
Hrp is best determined through measurements in the region
where -∆G° = λ, but such measurements are not possible with
this system. An alternative though problematic approach fits k
as a function of T to eq 15, to obtain Hrp and the non-adiabatic
reorganization energy λ12(non-ad).72 This analysis requires the
assumption that the equilibrium constants for forming the
precursor complexes KP are 1 at all temperatures, in the absence
of electrostatic work (note ∆G° has been found to be roughly
constant with temperature). The apparent values for Hrp, 10 (
4, 6 ( 2, and 4 ( 3 cm-1, and λ12(non-ad) are given in Table
4. These Hrp values would normally indicate a non-adiabatic
reaction. However, KP is likely to be smaller than 1 (two
standard estimating approaches give KP = 0.8664d and 0.0279)
and is likely to have a temperature dependence, becoming
smaller at higher temperatures due to an unfavorable entropy.
Including either KP < 1 or such a temperature dependence would
increase the value of Hrp. Thus, the apparent Hrp values
calculated from eq 15 with KP ) 1 are lower limits.80 For
instance, if the entropy of forming the precursor complex, ∆S°P,
were -10 cal K-1 mol-1 and KP ) 0.02 at 298 K, the derived
barrier because ET is coupled to transfer of the proton.
Hammarstro¨m and co-workers have reached the same conclusion
in their studies, that CPET oxidations of tethered phenol and
indole groups in water have much higher intrinsic barriers than
the pure electron transfers from the same reagents.8 The
similarities of our conclusions are striking in light of the
differences in our systems, Hammarstro¨m measuring rate
constants in water for intramolecular ET coupled to proton
transfer to the bulk aqueous solution. The λ11’s for HOAr-im
and HOAr-py are smaller than that of HOAr-NH2 but still
larger than those for aromatic organic molecules. The significant
differences in intrinsic barriers for the amino vs the pyridyl and
imidazolyl derivatives will be discussed in a future report.75
The large intrinsic barriers to CPET s indicating that it is
inherently difficult s would suggest that this concerted reaction
would be disfavored relative to the stepwise ET-PT mechanism.
If the driving forces for the competing rate-limiting steps were
identical, this would indeed be the case. However, in this system
E°(CPET) is substantially more favorable than E°(initial ET),
which leads to the lower barrier for the CPET (Figure 7). In
Hrp would be 130 cm-1 81
.
In sum, the CPET reactions described here appear to be at
most mildly non-adiabatic. The slowness of the electron transfer
reactions of HOAr-NH2 is due not to substantial non-adiabatic
character but rather to large reorganization energies. For
example, HOAr-NH2 + [N(tol)3]•+ is ∼105 slower than
[N(tol)3]•+/0 self-exchange; both processes have ∆G° = 0. This
is because λ12 for HOAr-NH2 + [N(tol)3]•+ is substantially
(78) Newton, M. D.; Sutin, N. Annu. ReV. Phys. Chem. 1984, 35, 437-480
(esp. p 451).
(79) Sutin, N. Prog. Inorg. Chem. 1983, 30, 441-498.
(80) The Eyring parameters for these reactions (Table 4) also suggest adiabatic
processes. Non-adiabatic reactions should be marked by large negative ∆Sq
to reflect the low probability of reaction, but the values observed, -14 to
-17 eu, are modest for a bimolecular process.
(81) KP ) 0.1 (at 298 K) and ∆SP ) -10 eu imply ∆HP ) -1.6 kcal mol-1
.
Using these assumptions, estimated values of Hrp and λ can be derived by
fitting kET (where kET ) kobs/KP) vs T data to eq 15.
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6086 J. AM. CHEM. SOC. VOL. 128, NO. 18, 2006

Concerted Proton−Electron Transfer in Phenols
A R T I C L E S
• +/0
larger than the λ11 of 12 kcal mol-1 for [N(tol)3 ]
self-
A deviation from the adiabatic Marcus equation may have been
observed in the different λ11 values obtained for oxidations of
HOAr-py with [Fe(R2bpy)3]3+ (30 ( 3 kcal mol-1) vs [Fe(Mex-
phen)3]3+ (23 ( 3 kcal mol-1).
The CPET rate constants are slower than the rate constants
of pure ET reactions of comparable organic reagents, especially
for HOAr-NH2. This is a result of the large intrinsic barriers
for CPET. A small part of this rate difference could be that the
CPET reactions are more non-adiabatic, but the data do not
support highly non-adiabatic CPET. The 27-56 kcal mol-1
adiabatic reorganization energies for these reactions are sig-
nificantly larger than typical λ’s for ET reactions of aromatic
organic compounds. For instance, [N(tol)3]•+/0 self-exchange has
exchange,43 whether one uses the adiabatic λ12 ) 34 kcal mol-1
(from eq 12) or the non-adiabatic λ12(non-ad) ) 30 kcal mol-1
(from eq 15 with KP ) 1). HOAr-py and HOAr-im have
intrinsic barriers that are also large but are smaller than that for
HOAr-NH2. The origin of these barriers and the differences
among these structurally similar phenol-bases will be discussed
in a future publication.75
Conclusions
One-electron oxidation of phenols hydrogen-bonded to a
pendant base, HOAr-B, yields radical cations in which the
phenolic proton has transferred to the base, OAr-BH+. Three
•
λ11 ) 12 kcal mol-1 43
. The pyridyl and imidazolyl compounds
have significantly lower intrinsic barriers than the amino
derivative. Future work will probe the origins of these barriers
and the differences among the different phenols.75
cases are reported here, with amino, pyridyl, and imidazolyl
bases. These systems serve as models for hydrogen-bonded
tyrosine residues in proteins, and more generally as an archetype
for a class of coupled proton-electron transfer reactions where
the electron and proton travel to different sites. The redox
potentials of these phenols are lower than those of simple
phenols, reflecting the favorable transfer of the proton to the
hydrogen-bonded base.
Experimental Section
General. Unless otherwise noted, reagents were purchased from
Aldrich, solvents from Fischer, and deuterated solvents from Cambridge
Isotope. MeCN was used as obtained from Burdick and Jackson (low-
water brand) and stored in an argon-pressurized stainless steel drum
Reactions of HOAr-B with [NAr3]•+ or [Fe(N-N)3]3+
oxidants in MeCN follow simple bimolecular kinetics. The
mechanism of oxidation involves concerted transfer of the proton
and electron (CPET). Three arguments rule out the alternative
stepwise mechanisms of initial proton and subsequent electron
transfer, or initial electron and subsequent proton transfer. First,
the primary kinetic isotope effects (1.6-2.8) are inconsistent
with the stepwise pathways. Second, the rates of oxidation are
too fast to involve the high-energy intermediates of the stepwise
pathways (the observed barriers are lower than the estimated
free energies of [-OAr-BH+] and [•+HOAr-B]). Third, the
dependence of the rate on driving force for the reaction HOAr-
NH2 + [NAr3]•+, ∆∆Gq/∆∆G°) 0.53, is consistent only with
the |∆G°| , 2λ situation found for CPET. Based on this work
and related model systems,7-9 CPET is likely a common (albeit
underappreciated) mechanism for phenol oxidations. It is favored
when the phenol is hydrogen-bonded to a base and when the
intermediates in the stepwise paths are high in energy. These
conditions probably occur often in biological systems, although
the local protein environment can have a substantial influence
on the relevant energetics.
The CPET reactions are in general well described by the
adiabatic Marcus equation (eq 12). Fitting the variation in k
with ECPET for a series of oxidants yields self-exchange
reorganization energies λ11 ) 56 ( 3, 27 ( 4, and 36 ( 3 kcal
mol-1 for HOAr-NH2/•OAr-NH3+, HOAr-py/•OAr-pyH+, and
HOAr-im/•OAr-imH+, respectively. For HOAr-NH2, the same
λ11 is found for three different oxidants, as required by the
Marcus treatment. For each of the phenols, the temperature
dependence of the rate constants gives the same λ11 as found
from k vs ECPET. The ∆∆GqCPET/∆∆G°CPET ) 0.53 for HOAr-
NH2 + [NAr3]•+ reactions is very close to the value of 0.5
predicted by the Marcus equation for this |∆G°CPET| , 2λ case.
The electrochemical rate constant for HOAr-NH2 correlates well
with the calculated HOAr-NH2 self-exchange rate constant.
These results support the use of simple Marcus theory for such
CPET systems, although these are not particularly stringent tests.
n
plumbed directly into a glovebox. Bu4NPF6 was recrystallized three
times from EtOH and dried in vacuo for 2 days at 110 °C prior to use.
1H NMR and 13C NMR spectra were recorded on Bruker AF300,
AV300, AV301, DRX499, or AV500 spectrometers at ambient tem-
peratures; chemical shifts are reported relative to TMS in ppm by
referencing to the residual solvent signals. The UV-vis spectra were
obtained on a Hewlett-Packard 8453 diode array spectrophotometer
and are reported as λmax in nm (ꢀ, M-1 cm-1), except for the long-
pathlength spectrum that was obtained on a CARY-500 instrument.
The EPR spectrum was recorded on a Bruker EPX CW-EPR spec-
trometer operating at X-band frequency at room temperature.
The synthesis of HOAr-NH2, equilibration experiments, crystal-
lographic data, and the EPR spectrum of •OAr-NH3+ are given in the
Supporting Information of this paper and of ref 24. Preparation of
MeOAr-NH2 followed the procedure in Scheme 2, starting from the
methyl bromoaryl ether C6H2(OMe)(2-Br)(4,6-tBu2) (see Supporting
Information). HOAr-im was prepared as described by Benisvy30 by
the condensation reaction: aldehyde + ammonium acetate + 4,4′-
dimethoxybenzil. The preparation of HOAr-py used the Ni(dppe)2
coupling of the phenol-derived Grignard and 2-bromopyridine described
by Fujita,31 except with a BBr3 deprotection of the methyl ether.82
Electrochemistry. Cyclic voltammograms were taken on an E2
Epsilon electrochemical analyzer (Bioanalytical Systems) at ca. 5 mM
substrate in anaerobic 0.1 M nBu4NPF6/acetonitrile solution, unless
otherwise specified. The electrodes were as follows: working, platinum
disk (unless noted otherwise); auxiliary, platinum wire; and reference,
Ag/AgNO3 (0.01 M) in electrolyte solution. All potentials are reported
vs a Cp2Fe+/0 internal standard. Errors are estimated to be (0.02 V.
Representative CVs are included as Figure S17 of the Supporting
Information.
For cyclic voltammetry at elevated and depressed temperatures, a
single solution was used for each analyte. Typically, the electrochemical
cell was prepared and degassed and CVs were collected, yielding E1/2
vs Ag/AgNO3. The entire cell was then placed in a warm water bath
and allowed to come to thermal equilibrium before CVs were collected.
This process was repeated using an ice bath. The cell was allowed to
return to ambient temperature, and a final series of CVs were collected;
in all cases, the E1/2 was found to be within 5 mV of the initial room-
temperature measurements. Last, ferrocene was added as an internal
standard and another CV was obtained. The potential of ferrocene vs
Ag/AgNO3 was found to vary by less than 5 mV over the temperatures
(82) Zhang, H.; Kwong, F. Y.; Tian, Y.; Chan, K. S. J. Org. Chem. 1998, 63,
6886-6890.
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A R T I C L E S
Rhile et al.
studied, so room-temperature Cp2Fe+/0 was used as a reference for all
of the CVs. A glassy carbon working electrode (φ ) 3 mm) was used
for HOAr-NH2 and HOAr-py in these experiments. The driving force
for the HOAr-NH2 + [N(tol)3]+• and HOAr-py + [Fe(Me2-bpy)3]3+
reactions was taken as the difference of E1/2(oxidant) - E1/2(HOAr-
B); see Table S21 in Supporting Information.
In the determination of the heterogeneous kel for HOAr-NH2, the
platinum disk working electrode (φ ) 1.6 mm) was polished before
each scan with commercial alumina solution and rinsed with water,
dilute HNO3(aq), ethanol, and acetonitrile before use. The uncompen-
sated resistance Ru of the electrochemical cell was measured at each
experiment and was found to be on the order of 500 Ω, which is
expected to have a negligible effect on the measured potential E (<5
mV). Simulated CVs were produced with DigiSim version 3.03,46 using
the experimentally measured values for E0, Eint, Erev, Eend, V, electrode
area (planar), DO, DR, and kel. The mechanism model used was B + e-
f A. The parameter R was taken to be 0.5.45
with a large excess of benchtop CH3OD (1% v/v for HOAr-NH2 or
0.5% v/v for HOAr-py). Control experiments showed that aerobic
addition of an equivalent amount of CH3OH to MeCN solutions did
not affect the rate. The isotope effects were corrected for the OH content
in the CH3OD (determined via 1H NMR), 7% for the experiments with
DOAr-ND2 and 4% for those with DOAr-py. Rate constants and data
analyses are given in the Supporting Information of this paper and of
ref 24.
X-ray crystallographic data and experimental descriptions are in
the Supporting Information.
Acknowledgment. We gratefully acknowledge support from
the U.S. National Institutes of Health (grant 2 R01 GM50422).
Special thanks to Prof. Arnold Rheingold for crystallographic
assistance with the HOAr-py crystal structure.
Supporting Information Available: Syntheses and cyclic
voltammograms, and kinetic, chronoamperometry, and X-ray
crystallographic data (PDF, CIF). This material is available free
of charge via the Internet at http://pubs.acs.org.
Kinetics. Kinetics experiments were performed on an OLIS RSM-
1000 stopped-flow apparatus in anaerobic MeCN. The data were
analyzed with SpecFit global analysis software.51 Kinetics were fit to
pseudo-first-order, second-order, or opposing second-order kinetics as
appropriate. To determine the isotope effects, solutions were prepared
JA054167+
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6088 J. AM. CHEM. SOC. VOL. 128, NO. 18, 2006