Oxidation of NIII and N-I by an {Mn4O 6}4+ core in aqueous media: Proton-coupled electron transfer

Article abstract of DOI:10.1002/ejic.200700363

[Mn₄(μ-O)₆(bipy)₆]⁴⁺ (1 ⁴⁺; bipy = 2,2′-bipyridine) and its conjugate acid [Mn ₄(μ-O)₅(μ-OH)(bipy)₆]⁵⁺ (1H⁵⁺) quantitatively oxidise N^III (HNO₂ and NO₂^-) and N^-1 (NH₃OH⁺ and NH₂OH) to N^V (nitrate) and N^I (nitrous oxide), respectively, in aqueous solution (pH 2.0-6.0), with 1H⁵⁺ reacting much faster than 1⁴⁺. An uncommon feature of these reactions is the kinetic superiority of HNO₂ over its conjugate base NO ₂^-. NH₂OH, however, behaves normally - the conjugate acid NH₃OH⁺ is less reactive than NH ₂OH. These reactions show remarkable kinetic isotope effects: the observed rate of N^III oxidation increases in D₂O media whereas the N^-I oxidation rate slows down in media enriched with D₂O. A search of the available data on the redox kinetics of multinuclear oxidants suggests that the title {Mn₄O₆} ⁴⁺ reduction by N^III, the rate of which is accelerated in D₂O, is the only one established so far. A hydrogen atom transfer (HAT) mechanism (1e, 1H⁺; electroprotic) is proposed. Wiley-VCH Verlag GmbH & Co. KGaA, 2007.

Full text of DOI:10.1002/ejic.200700363

FULL PAPER
DOI: 10.1002/ejic.200700363
Oxidation of NIII and N–I by an {Mn4O6}4+ Core in Aqueous Media: Proton-
Coupled Electron Transfer
Suranjana Das[a] and Subrata Mukhopadhyay*[a]
Keywords: Kinetics / Mechanisms / Manganese / Oxidation / Nitrogen
[Mn4(µ-O)6(bipy)6]4+ (14+; bipy = 2,2Ј-bipyridine) and its con-
jugate acid [Mn4(µ-O)5(µ-OH)(bipy)6]5+ (1H5+) quantitatively
oxidise NIII (HNO2 and NO2–) and N–I (NH3OH+ and NH2OH)
to NV (nitrate) and NI (nitrous oxide), respectively, in aqueous
solution (pH 2.0–6.0), with 1H5+ reacting much faster than
14+. An uncommon feature of these reactions is the kinetic
superiority of HNO2 over its conjugate base NO2–. NH2OH,
however, behaves normally Ϫ the conjugate acid NH3OH+ is
less reactive than NH2OH. These reactions show remarkable
kinetic isotope effects: the observed rate of NIII oxidation in-
creases in D2O media whereas the N–I oxidation rate slows
down in media enriched with D2O. A search of the available
data on the redox kinetics of multinuclear oxidants suggests
that the title {Mn4O6}4+ reduction by NIII, the rate of which is
accelerated in D2O, is the only one established so far. A hy-
drogen atom transfer (HAT) mechanism (1e, 1H+; electro-
protic) is proposed.
(© Wiley-VCH Verlag GmbH & Co. KGaA, 69451 Weinheim,
Germany, 2007)
is stable in aqueous solution over a wide acidity range
(pH 2.0–6.0). Its one-electron-reduced, mixed-valent
MnIV3MnIII form is an EPR spectroscopic model for the S2
state,[14] which means that the electron-transfer reactivity of
the title Mn tetramer will thus be instructive.
Introduction
The oxygen-evolving complex (OEC)[1–5] oxidises water
to molecular dioxygen in photosystem II (PS II) at the re-
dox catalytic centre on the lumenal side of PS II, which
includes four manganese ions, a calcium ion and a chloride
ion.[6,7] The OEC cycles over five redox states S0–S4, where
the index refers to the number of oxidising equivalents
stored;[8,9] dioxygen is evolved in the S4ǞS0 transition.
Successive changes in the redox states of the Mn atoms in
the OEC during this cycle are associated with protonation/
deprotonation of the oxo bridges[10] and of the nearby
amino acid residues as well as the simultaneous movements
of proton and electrons within the OEC.[11]
The acid/base chemistry resulting from oxo-bridge pro-
tonations in synthetic multinuclear Mn complexes has been
studied mostly in non-aqueous media.[12] It has also been
demonstrated that oxo-bridge protonation sometimes in-
creases the reactivity of a binuclear Mn oxidant[12d,12e] and
completely quenches the catalase activity of an MnIV di-
mer.[12c] How such protonations affect the chemical behav-
iour of higher-valent synthetic Mn complexes is not known,
despite the relevance of such knowledge to PS II. This is
particularly true in water, the medium where PS II acts.
The title complex [Mn4(µ-O)6(bipy)6]4+ (14+; Figure 1)
formally corresponds to the fully oxidised S3 or S4 state of
OEC[2] [although involvement of an even higher state
(MnV) is feasible in the water oxidation process],[13] which
Figure 1. Schematic drawing of [Mn4(µ-O)6(bipy)6]4+. N-N = 2,2Ј-
bipyridine.
We demonstrated recently that 14+ oxidises glyoxylic and
pyruvic acid through an electroprotic mechanism and that
the oxo-bridged protonated species 1H5+ is a kinetically su-
perior oxidant to 14+ [15]
It therefore seemed worthwhile to
.
investigate whether the relative acidities of the reducing
agents influence the proton-coupled electron transfer
mechanism as well as the reactivity of the oxo-bridged pro-
tonated oxidant. The present investigation was thus
planned with two selected reducing agents, namely HNO2
and NH3OH+ (and their conjugate bases), which have a
sufficient separation in their pKa values.[16,17] We further
note that hydroxylamine[18] reduces the Mn ions in PS II
whereas nitrite does not.[19]
[a] Department of Chemistry, Jadavpur University,
Calcutta 700 032, India
E-mail: smukhopadhyay@chemistry.jdvu.ac.in
Supporting information for this article is available on the
WWW under http://www.eurjic.org or from the author.
4500
© 2007 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Eur. J. Inorg. Chem. 2007, 4500–4507

Proton-Coupled Electron Transfer
Table 1. Stoichiometry of reduction of the MnIV complex by NIII
Results and Discussion
4
and N–I (T = 25.0 °C, Cbipy = [Hbipy+] + [bipy] = 3.0 m).
[MnIV
]
4
pH
[NIII
]
[N–I]T
∆[MnIV4]/∆[TR]
Equilibrium Measurements
T
[m₎
(= TR) [m₎ (= TR) [m₎
The measured pKa values of these reducing acids in a
95% D2O medium are 3.80Ϯ0.10[20,21] and 6.40Ϯ0.10 for
nitrous acid and hydroxylammonium cation respectively.
The reported pKa values used in this investigation for
HNO2 and NH3OH+ in an H2O medium are 3.00[16] and
6.00,[17] respectively. All the pKa values are at 25.0 °C, I =
1.0  (NaNO3). The pH-metric titration of the title tetranu-
clear Mn oxidant, however, did not result in any ionisation
constant in the pH interval 2.0–7.0.
0.20
0.40
0.60
1.50
0.10
3.3
3.6
5.6
5.5
3.5
1.50
2.00
3.50
10.0
0.20
0.27
0.26
0.23
0.26
0.24
average = 0.25Ϯ0.02
0.20
0.50
0.80
1.30
0.20
2.5
4.5
5.0
5.3
5.1
2.00
4.00
6.00
10.0
0.60
0.23
0.24
0.27
0.23
0.23
average = 0.24Ϯ0.02
Stoichiometric Analyses
–
Nitrate (NO3 ) was detected as the oxidation product of
0.30 m). Averages of k0 values from at least three measure-
ments were taken and the average coefficients of variation
(CV)[23] for these measurements were within 3%. The ob-
served proton-dependences of k0 in these two redox reac-
tions is interesting (Table 2). For nitrite oxidation, k0 values
increase with an increase in acidity of the reaction medium
whereas they decrease with increasing acidity for hydroxyl-
amine oxidation until the very low pH region, where k0 in-
creases with a decrease in pH. We can suppose a superior
nitrite whereas no red colouration was observed when the
product solutions of the redox reaction of hydroxylamine
were tested for nitrite or nitrate. The results of several stoi-
chiometric experiments, under both kinetic (TR
Ͼ
4[MnIV4]) and non-kinetic conditions (TR Ͻ 4[MnIV4]), are
summarised in Table 1, which establishes a 1:4 stoichiome-
try (oxidant:reductant) for both the redox reactions, where
TR is the analytical concentration of the reducing agent de-
–
fined as TR = [HNO2] + [NO2 ] for [NIII
]
and TR =
–
kinetic reactivity of HNO2 over NO2 from these trends
T
[NH3OH+] + [NH2OH] for [N–I]T. The title Mn tetramer is
not a powerful oxidant (see below) and it is clear from the
ratio ∆[MnIV4]/∆[TR] (Table 1) that N2O is the sole oxi-
dation product of hydroxylamine Ϫ no further oxidation of
N2O to nitrite or nitrate was observed. The results of
EDTA titrations of the product solutions revealed quantita-
tive formation of MnII as the sole reduction product of
MnIV. The likely MnII species under these reaction condi-
tions is an MnII-bipy complex.[22] Optical spectra for the
product solutions were superimposable on those of mix-
tures of Mn(NO3)2 and 2,2Ј-bipyridine under similar exper-
imental conditions. Equations (1) and (2) are thus closest to
our above observations.
whereas NH2OH is superior to NH3OH+. However, the ti-
tle complex (14+) may suffer protonation at very low pH,
as was evident from its reaction with α-keto acids.[15] The
species formed (1H5+) may act as an additional oxidant in
the low pH region despite being present in only a low con-
centration. Protonated oxidants always react quicker than
their conjugate bases[24] within a system, and thus 1H5+ is
expected to react faster than 14+. We measured several k0
values in the presence of varying amounts of added bipy
over the entire pH range for both the redox reactions. The
lack of influence of added bipy on the observed rate of
either reaction at any acidity rules out the possibility of
bipy dissociation from the parent Mn complex in solution.
We observed a linear variation of k0 with TR in the entire
pH range studied and found no indication of rate saturation
even at the highest TR studied. Table 2 collects some repre-
sentative k0 values.
A strong pre-equilibrium (K) binding of the reducing
species with the title oxidant followed by reaction of the
intermediate thus formed with a second reducing species
would lead to the observed kinetics (linear dependence of
k0 on TR) provided that KTR is much greater than 1, which
requires a K value as high as 104 (lowest TR = 5.0ϫ10–3 ).
However, the tetranuclear Mn oxidant is coordinatively sat-
urated at each Mn centre and an inner-sphere attachment
Kinetics
The decay in absorbance vs. time plots could be fitted of reducing species seems unlikely. Moreover, d3 MnIV is
with standard first-order decay curves at least up to three substitution inert and the oxidant retains its bipy ligands in
half-lives. The observed first-order rate constants (k0) do aqueous solution in the entire pH range studied (see above),
not depend on the wavelength chosen for the measurements which means there is no scope for the formation of the aqua
(380–500 nm) where only the oxidant absorbs sufficiently, species that could lead to an inner-sphere mechanism in
the presence or absence of dissolved oxygen in the reaction solution. We thus disregard this kind of pathway in the ki-
medium or on the concentration of the Mn oxidant (0.05– netics of the title redox reaction. The observed proton-
Eur. J. Inorg. Chem. 2007, 4500–4507
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4501

S. Das, S. Mukhopadhyay
FULL PAPER
Table 2. Some representative first-order rate constants for the oxi-
dation of nitrite and hydroxylamine by the title complex (0.10 m)
at T = 25.0 °C, I = 1.0  (NaNO3). The values in parentheses are
calculated from Equation (9) using the rate constants reported in
Table 3.
pH
TR []
Cbipy [m]
104k0 [s–1]
Nitrite
2.50
2.82
3.14
3.78
4.18
4.75
5.15
5.56
4.69
4.69
4.67
4.15
4.13
4.13
3.22
3.24
3.23
4.71
4.70
2.51
2.53
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.025
0.075
0.10
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
20
491 (500)
197 (202)
74.0 (76.6)
11.2 (10.4)
3.44 (3.34)
0.74 (0.78)
0.28 (0.30)
0.11 (0.11)
0.43 (0.45)
1.39 (1.35)
1.81(1.89)
3.66 (3.62)
3.81 (3.82)
4.03 (3.82)
57.8 (59.6)
58.9 (56.0)
59.2 (57.8)
1.06[a]
Scheme 1. (RH = HNO2 or NH3OH+).
tion (11)]. A reaction path involving NO would require an
[HNO2]2 term in the rate law and NO+ requires a third-
order term [MnIV4][HNO2][H+]. The equilibrium constants
of these reactions are small, which might be the main
reason for not obtaining the reactivities, if any, of these spe-
50
80
20
50
–
cies. We therefore used [NIII
]
= [HNO2] + [NO2 ]. Except
T
for NH2OH and NH3OH+, no other pH-dependent N–I
species are known.
80
3.0
3.0
3.0
3.0
1.39[b]
496 [a]
499 [b]
Hydroxylamine
Plots of the left-hand side of Equation (9) resulted in a
good polynomial fit of [H+] for both redox reactions (Fig-
ure 2, r Ͼ 0.98), from which we obtain k1K1, (k2K1Ka + k3)
2.04
2.37
3.02
3.20
3.42
4.00
4.41
4.84
5.35
6.00
3.61
3.62
3.61
4.11
4.10
4.13
4.12
4.12
5.95
5.93
2.11
2.12
3.60
3.62
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.02
0.04
0.10
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
1.0
10
20
40
60
5.0
5.0
5.0
5.0
5.0
5.0
3.60 (3.58)
1.95 (1.92)
0.79 (0.84)
0.73 (0.76)
0.74 (0.74)
1.19 (1.12)
1.98 (2.10)
4.52 (4.69)
12.2 (12.4)
32.0 (32.2)
1.51 (1.58)
3.23 (3.18)
7.79 (7.90)
1.21 (1.30)
1.23 (1.28)
1.29 (1.34)
1.37 (1.32)
1.42 (1.32)
30.1[a]
30.5[b]
2.61[a]
2.08[b]
1.22[a]
1.05[b]
[a] I = 0.5  (NaNO3). [b] I = 0.1  (NaNO3).
dependence of k0 on [H+] could not be fitted to any type of
reaction scheme with only one protic equilibrium, which
means that Equations (3) and (4) need to be considered
simultaneously, as shown in Scheme 1 for both the reactions
under investigation.
Scheme 1 leads to the rate law in Equation (9), assuming
that K1[H+] is much less than 1. This scheme does not take
[25]
into account any protonation equilibrium of HNO2
to
Figure 2. Plot of the left-hand side of Equation (9) vs. [H+]. [com-
generate NO+ [Equation (10)] or a disproportionation equi-
plex] = 0.10 m, T = 25.0 °C, I = 1.0 . A: NIII, Cbipy = 3.0 m;
[26]
librium of HNO2
to generate NO and NO2 [Equa- B: N–I, Cbipy = 5.0 m.
4502
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Proton-Coupled Electron Transfer
Table 3. Rate constants for the reduction of 14+ by nitrite and hydroxylamine at T = 25.0 °C, I = 1.0  (NaNO3).
k1K1
k2K1Ka + k3
[–1 s–1]
k4
[–2 s–1]
[–1 s–1]
H2O medium NIII
H2O medium
N–I
390Ϯ20
3.43Ϯ0.20
(8.2Ϯ0.5)ϫ10–2
0
(4.4Ϯ0.3)ϫ10–3
(6.6Ϯ0.4)ϫ10–1
95% D2O medium
760Ϯ45
(5.0Ϯ0.3)ϫ10–2
(1.9Ϯ0.1)ϫ10–4
0
NIII
95% D2O medium
N–I
0.65Ϯ0.04
(5.5Ϯ0.3)ϫ10–2
and k4 (see Table 3). The non-existence of the k4 path for pated (1H5+ is a stronger oxidant than 14+ and the k4 path
the nitrite redox reaction, as determined by the above sec- is absent in NIII oxidation), while at pH 5–6 the value of k0
ond-order polynomial, was further confirmed by plotting remains unaltered in the range I = 1.0–0.1  (NaNO3) for
k0(Ka + [H+])/(TR[H+]) against [H+] (see electronic support- the N–I redox reaction (Table 2), in line with the maximum
ing information), which resulted in a good straight line (r contribution of the uncharged NH2OH path (k4) towards
Ͼ 0.98) with slope k1K1 and intercept (k2K1Ka + k3), in the overall rate at the highest pH extreme. Around pH 2.5,
–
close agreement (within 5%) with the values reported in where [HNO2] is much greater than [NO2 ], the rate of the
Table 3. On the other hand, the k4 path was found to be nitrite redox reaction was found to be independent of any
significant for hydroxylamine redox above a pH of about 4, variation in the ionic strength of the medium (Table 2),
and a plot of TR/k0 against [H+] yielded a good straight whereas at pH 2.1 or 3.6, where [NH3OH+] is much greater
line with slope 1/k4 and intercept 1/(k4Ka) at pH Ն 5.3 (see than [NH2OH], the rate of hydroxylamine oxidation de-
electronic supporting information). The value of k4 deter- creases significantly with decreasing ionic strength of the
mined in this way matches well with the values reported in medium (Table 2). These observations qualitatively support
Table 3.
the involvement of neutral HNO2 and cationic NH3OH+ as
It cannot be concluded, however, that, as 14+ does not reactive species with cationic 14+ or 1H5+
.
–
react with NO2 (the k4 path does not contribute to the
We can estimate K1 to be between 10–9 and 10 from the
–
overall rate), 1H5+ would not oxidise NO2 . Protonated fact that K1[H+] is much less than 1 and the k1K1 values
metal complexes are superior oxidants than their deproton- listed in Table 3. The fact that K1[H+] is much less than 1
ated analogues[20,24] and a finite contribution of k2 path sets an upper limit for the pre-equilibrium binding of nearly
may therefore be anticipated, although kinetic resolution of 10 as the maximum [H+] used in this study is 10–2 . As-
the proton-ambiguous k2 and k3 paths is not possible. We suming the diffusion-controlled limit is 1011 –1 s–1, the

also note that the k1 path starts contributing towards the lower limit of K1 derived from the k1K1 values in Table 3 is
observed rate at pH Յ 5.0 for the NIII redox reaction around 10–9. For comparison, we note that the reported
whereas for the N–I redox reaction a more acidic environ- pKa of the analogous complex [MnIV4(µ-O)5(µ-OH)-
ment is needed to ensure a finite contribution of the k1 path (bpea)4]5+ [bpea = N,NЈ-bis(2-pyridylmethyl)ethylamine] is
(pH Յ 4.0), as judged by the values of k1K1[H+]2 and –6,[12h] which shows that the bridging oxo group is much
(k2K1Ka + k3)[H+] for both redox reactions at these pHs. less basic. The strong acidity of the protonated oxo-bridge
Below pH 3.0, hydroxylamine is almost quantitatively pres- may thus be inferred. The Mn tetramer used here does not
ent as cationic NH3OH+, whose oxidation by 1H5+ is elec- undergo any spectral change in the pH range 2.0–6.0 there-
trostatically more unfavourable than the oxidation of fore spectral evaluation of its basicity was not possible.
NH2OH. Oxidation of cationic NH3OH+ requires higher
It appears from Table 3 that the maximum value of k3 is
concentrations of 1H5+, which are produced at much lower 4.4ϫ10–3
pH than that required for the oxidation of neutral HNO2 8.2ϫ10–2


–1 s–1 for N–I oxidation while it is
–1 s–1 for NIII oxidation. The reactivity of
purely on charge grounds.
NH2OH is higher than NH3OH+ as the rate increases with
It should be noted here that a scheme comprising only pH while it decreases with an increase in pH for the NIII
HNO2 as the active reducing species reacting with 1H5+ (k1) redox reaction. The maximum value of k2K1 for this latter
and 14+ (k1Ј) is an alternative to Scheme 1 and results in reaction is 82 (much lower than k1K1), which clearly indi-
k1K1 and k1Ј values identical to those reported in Table 3 cates that HNO2 reacts much quicker with 1H5+ or 14+ than
–
(k1Ј is similar to k3 in Scheme 1, assuming the k2 path is NO2 . The overwhelming dominance of the reactivity of
absent). However, this alternative cannot confirm the ab- 1H5+ is clear in both redox reactions.
sence of the k2 path in the nitrite redox reaction. The rate
constants measured in presence of varying amounts of
added NaNO3 might be instructive here. At pH 4.7 (where
[NO2 ] ϾϾ [HNO2]), k0 increases substantially with a de-
Mechanism
–
crease in ionic strength of the reaction medium (Table 2),
The electrochemistry of [Mn4(µ-O)6(bipy)6]4+ has been
which indicates a reaction between opposite charges, there- extensively studied by Dunand-Sauthier et al.[27] The cyclic
fore a finite contribution of the k2 path may thus be antici- voltammogram of [Mn4(µ-O)6(bipy)6]4+ in bipy/Hbipy+
Eur. J. Inorg. Chem. 2007, 4500–4507
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4503

S. Das, S. Mukhopadhyay
FULL PAPER
buffer shows only an irreversible reduction peak near
0.40 V (vs. Ag/AgCl) which leads to the formation of
[Mn(bipy)3]2+ species.[27] Attempts at electrolytic reduction
resulted in decomposition of the tetranuclear complex
into four MnII ions. Standard chemical reduction in solu-
tion also failed to achieve the monoreduced form
[MnIV3MnIII(µ-O)6(bipy)6]3+ [14]
Exhaustive reduction of
.
[Mn4(µ-O)6(bipy)6]4+ at 0.20 V (vs. Ag/AgCl) consumes al-
most eight equivalent of electrons per mole of the tetranu-
clear oxidant and leads to MnII.[27a] These authors have
also shown an intermediate generation of the mixed-valent
[MnIII,IV(µ-O)2(bipy)4]3+ species during the conversion to
MnII. This mixed-valent species could not, however, be gen-
erated selectively and quantitatively as the reduction of the
tetranuclear and the mixed-valent species occurs at a very
similar potential (approx. 0.4 V vs. Ag/AgCl).[27] We did not
observe formation of any such spectrally characterisable bi-
nuclear intermediate during the chemical reduction of the
title oxidant when reducing agents were used in either ex-
cess, deficit or stoichiometric amounts, see Equations (1)
and (2).
We can conclude from the known electrochemical behav-
iour of the title tetranuclear Mn comlex that this species is
–
Figure 3. Effect of the amount of D2O (mol-%) on k0. [complex] =
not a powerful oxidant. The one-electron oxidation of NO2
0.10 m, T = 25.0 °C, I = 1.0 . A: NIII, [NIII
] = 0.05 , Cbipy =
T
or HNO2 and NH2OH or NH3OH+ is, however, highly en-
3.0 m; B: N–I, [N–I]T = 0.01 , Cbipy = 5.0 m.
–
dothermic. The one-electron oxidation potentials for NO2 /
NO2, HNO2/NO2, H+, NH2OH/NH2OH+ and NH3OH+/
NH2OH+, H+ couples are –1.04,[26] –1.22,[28] –0.91[29] and
An initial one-electron transfer rate-limiting step in the
–1.30 V respectively.[28] An initial one-electron exchange be- overall eight electron transactions [Equations (5)–(8)] is
tween [Mn4(µ-O)6(bipy)6]4+ and these reducing species is proposed to result in the intermediate mixed-valent Mn spe-
thus energetically unfavourable although we observe that cies MnIV3MnIII, which is likely to be reduced quickly in the
these reducing agents smoothly reduce the MnIV oxidant to following steps by the excess reducing agents or the radicals
MnII. An inner-sphere attachment of the reducing species generated at the rate-determining step. We did not observe
to the MnIV centres that could provide an alternative path- any polymerisation in either redox reaction when the reac-
way is not likely as the d3 MnIV atoms in the coordinatively tions were carried out in 6% v/v acrylonitrile, although this
saturated oxidant are substitution inert. However, it can be does not rule out the production of radicals as they may
concluded from the known electrochemical behaviour of the react faster with the one-electron reduced mixed-valent Mn
tetranuclear Mn oxidant that the intermediate mixed-valent species, which is itself very reactive.[30] The one-electron re-
forms are too reactive. In fact, the very high reactivity of duced species of 14+ or 1H5+ would have much more basic
the most likely one-electron reduced species produced oxo bridges that immediately accept a proton from the reac-
by cryogenic radiolytic reduction of 14+
[MnIV3MnIII(µ-O)6(bipy)6]3+ has been documented.[30] ics barriers to further reduction.[33] A large increase in the
This species is very unstable even at low temperature basicity of oxo bridges after one-electron reduction of the
(190 K).[30]
bridge carrying higher valent Mn atoms in di- or polynu-
,
namely tion medium, which could remove thermodynamic or kinet-
,
In order to further explore the electron-transfer mecha- clear Mn complexes is well documented[34] and is a key fea-
nism we measured k0 values in media enriched with D2O ture in the redox steps of the Kok cycle in PS II.[1c,35] Radi-
and computed the second-order rate constants or their cals produced from one-electron oxidation of nitrite and
composites (Table 3) in the same way as in pure H2O media; hydroxylamine, namely NO2 and NH2OH+, are well-docu-
they showed kinetic isotope effects. Interestingly, we also mented in the literature and pathways leading to the forma-
note that the k1K1 path in NIII oxidation increases substan- tion of nitrate[16b,36] and nitrous oxide[37] are well-estab-
tially in D2O media whereas all the other rate constants or lished facets of nitrogen chemistry.
their composites decrease. A lowering of the rate in D2O
It has been known for a long time that rate increases
media is clearly indicative of an electroprotic mechanism.[31] significantly with D2O content in an H2O/D2O mixture if
Notably, we also found that the rate enhancement and rate acid/base pre-equilibrium steps are involved prior to the
decrease in H2O/D2O mixtures for NIII and N–I oxidations, rate-determining step.[38a–38c] D2O has a smaller autoproto-
respectively, are linearly related (r Ͼ 0.98) to the D2O con- lysis constant than H2O by a factor of about 5 and thus it
tent (Figure 3), which suggests transfer of a single proton is less basic than water,[38a] which means that the oxo brid-
in the slow steps.[32]
ges in the title oxidant will be able to compete with the
4504
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Eur. J. Inorg. Chem. 2007, 4500–4507

Proton-Coupled Electron Transfer
solvent for the deuteron in D2O more effectively than for around the OEC[13a,13b,40] cluster, thereby providing an in-
the proton in H2O. An increase in the K1 value in D2O is sight into the movement of protons and water in the site.
thus expected. In general, weak acids in H2O become much
weaker in D2O,[38] and we found that reducing acids like
glyoxylic, pyruvic, nitrous and hydroxylammonium cation Experimental Section
have considerably higher pKa’s in D2O.[15,20,21] We can thus
Materials: The complex salt hydrate [Mn4(µ-O)6(bipy)6][ClO4]4·
reasonably expect a higher concentration of 1H5+ in D2O
2H2O was synthesised following the literature procedure.[14] One of
and the observed reversal in the rate constants k1K1 in the
two redox reactions might be a result of competitive in-
crease in the K1 value and a decrease in the k1 value (elec-
troprotic mechanism).
its water molecules is easily lost[14] therefore the elemental analyses
of the material closely resemble the values expected for the mono-
hydrate. The tetramer used here thus appears to be the monohy-
drate. C60H50Cl4Mn4N12O23: calcd. C 43.16, H 2.99, N 10.07;
found C 43.39, H 3.03, N 10.00.
It has been shown recently[39] that an inverse deuterium
isotope effect (kH Ͻ kD) is observed during charge recombi-
nation between the fully reduced FADH– anion (FAD =
flavin adenine dinucleotide) and the neutral (deprotonated)
radical of the solvent-exposed tryptophan W306. This can
be interpreted by a model of fast protonation equilibrium
for the W306 radical that involves an increase in the pKa of
the radical in D2O, which is 0.56-times that in H2O buffer.
The authors were able to kinetically separate the rate con-
stants, and the rate constant for recombination via depro-
tonated radical was found to be higher in H2O than in D2O
(1.4 times).
The preparation, standardisation and storage of sodium nitrite, hy-
droxylamine nitrate and sodium nitrate have been described pre-
viously.[20,41] 2,2Ј-Bipyridine was used as received from Sigma. D2O
(99.9 atom-% D) and DNO3 (99+ atom-% D) were also purchased
from Sigma. All other chemicals were of reagent grade and used
as received. Doubly distilled, deionised and then freshly boiled
water was used throughout.
Equilibrium Measurements: The acid dissociation constants of the
reducing acids HNO2 and NH3OH+ cation in 95% D2O medium
were determined by pH-metric titration using a Metrohm (736 GP
Titrino) autotitrator at 25.0(Ϯ0.1) °C, I = 1.0  (NaNO3), as de-
scribed previously.[20,41,42] To avoid nitrous acid decomposition in
acidic media, solutions of NaNO2 were titrated with DNO3 while
aliquots of hydroxylamine solutions (initially at pH 2 made by add-
ing HNO3) were titrated with carbonate-free NaOH. The reported
values[16,17] of these ionisation constants in H2O media were used.
A similar pH-metric titration of [Mn4(µ-O)6(bipy)6]4+ was per-
formed in the pH range 2.0–7.0 to examine the basicity of its oxo
bridges.
It seems likely that the second-order rate constants de-
crease to such a great extent that this overcompensates the
effect of an increase in the K1 value in D2O for the N–I
redox reaction whereas it is unable to slow down the NIII
redox reaction. The rate retardations in D2O for the oxi-
dation of glyoxylic and pyruvic acid by the title Mn oxidant
are much smaller[15] than those observed in the oxidation
of N–I, which might indicate that the effect of an increase
in K1 is almost cancelled out by the decrease in rate con-
stants in D2O for the oxidations of these α-keto acids. A
substantial decrease of Ka in D2O (compared to that in
H2O) also contributes significantly to a remarkable de-
crease of the composite (k2K1Ka + k3) for the NIII redox
reaction in D2O.
Stoichiometry Measurements: The stoichiometries of the reactions
were measured with both an excess and a deficit of reducing agent.
Reaction mixtures containing 6–15-times NIII over the Mn oxidant
(0.05–1.50 m) were allowed to react until the solutions became
colourless (absorbance of less than 0.01 at 420 nm). Unreacted NIII
in these product mixtures was determined colourimetrically. After
appropriate dilution, the product solutions were treated with α-
naphthylamine and sulfanilic acid. The red dye thus formed was
estimated at 520 nm (ε = 4.0ϫ104 –1 cm–1).[43] We verified that
The contribution of reaction (8) in Scheme 1 (for R– =
–
NO2 ) towards the overall rate is statistically insignificant
–
bipyridine, MnII and ClO4 do not interfere in this coupling reac-
(Table 3). Moreover, (k2K1)max is much lower than k1K1 for
NIII oxidation, therefore we may reasonably expect that the
k3 path contributes significantly to the measured rate of
tion. NO3 as the NIII oxidation product was qualitatively detected
–
by the above-mentioned coupling reaction after adding Zn dust.
For this purpose reaction mixtures were prepared with com-
plex:NIII ratios of between 1:2 and 1:4. The Mn tetramer did not
interfere in this colour reaction. Unreacted hydroxylamine was de-
termined by titrating with standard KBrO3 in the usual manner[44]
–
NIII oxidation. The superior reactivity of HNO2 over NO2
in reducing 14+ or 1H5+ is thus clearly established. Both
NH2OH and NH3OH+, on the other hand, react with the
tetrameric Mn oxidant. It appears that the presence of an
H-atom in both these reducing agents is responsible for the
observed reactivity and thus a hydrogen atom transfer
(HAT) from the reducing acids may be mechanistically con-
when MnIV (0.5–1.50 m) was treated with (5–10 m) N–I. For-
4
mation of NO2 or NO3 as the N–I oxidation product was also
–
–
tested by the coupling reaction described above. Under non-kinetic
conditions, 0.10–0.50 m of the MnIV complex was treated with a
4
less than stoichiometric amount of the reducing agents. Unreacted
·
ceived to produce the radicals NO2 or NHOH· along with
MnIV complex was measured at 420 nm (ε = 7.5ϫ103 –1 cm–1)
4
the protonated oxo bridge in the mixed-valent MnIV3MnIII
species. The kinetic superiority of 1H5+ may be due, at least
in part, to its more oxidising nature than 14+ as well as a
more facile movement of hydrogen or hydrogen ion through
the H-bonding framework formed from the oxo bridges of
1H5+ and the oxygen and hydrogens of the reducing species.
The molecular mechanics computed structure of the OEC
when the reactions reached equilibrium. MnII in the product solu-
tions was estimated by EDTA titrations using EBT as indicator.[45]
For this purpose, excess reducing agent was treated with the Mn
oxidant until the absorbance at 420 nm became less than 0.01 and
–
–
then MnII was estimated. We verified that NO2 , NO3 , NH2OH
or bipy do not interfere in this complexometric titration.
Physical Measurements and Kinetics: Absorbance and electronic
reveals a network of organised hydrogen-bonded species spectra were recorded with a Shimadzu (1601 PC) spectrophotome-
Eur. J. Inorg. Chem. 2007, 4500–4507
© 2007 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
www.eurjic.org
4505

S. Das, S. Mukhopadhyay
FULL PAPER
J. Biesiadka, B. Loll, K. Sauer, J. Messinger, A. Zouni, V. Yach-
andra, Proc. Natl. Acad. Sci. USA 2005, 102, 12047–12052.
E. M. Sproviero, J. A. Gascón, J. P. McEvoy, G. W. Brudvig,
V. S. Batista, J. Chem. Theory Comput. 2006, 2, 1119–1134.
P. Joliot, G. Barbieri, R. Chabaud, Photochem. Photobiol. 1969,
10, 309–315.
B. Kok, B. Forbush, M. McGloin, Photochem. Photobiol. 1970,
11, 457–475. For a recent description of structure and oxi-
dation state changes of PS II Mn complexes in these states, see:
M. Haumann, C. Müller, P. Liebisch, L. Iuzzolino, J. Dittmer,
M. Grabolle, T. Neisius, W. Meyer-Klaucke, H. Dau, Biochem-
istry 2005, 44, 1894–1908.
a) J. B. Vincent, G. Christou, Inorg. Chim. Acta 1987, 136, L41–
L43; b) C. Meinke, V. A. Solé, P. Pospisil, H. Dau, Biochemistry
2000, 39, 7033–7040; c) P. E. M. Siegbahn, Inorg. Chem. 2000,
39, 2923–2935.
a) J. P. McEvoy, J. A. Gascon, V. S. Batista, G. W. Brudvig,
Photochem. Photobiol. Sci. 2005, 4, 40–949; b) L. I. Krishtalik,
Biochim. Biophys. Acta 1986, 849, 162–171; c) C. W. Hoganson,
N. Lydakis-Simantiris, X. S. Tang, C. Tommos, K. Warncke,
G. T. Babcock, B. A. Diner, J. McCracken, S. Styring, Pho-
tosynth. Res. 1995, 46, 177–184; d) C. W. Hoganson, G. T. Bab-
cock, Science 1997, 277, 1953–1956.
a) M. J. Baldwin, T. M. Stemmler, P. J. Riggs-Gelasco, M. L.
Kirk, J. E. Penner-Hahn, V. L. Pecoraro, J. Am. Chem. Soc.
1994, 116, 11349–11356; b) K. S. Hagen, T. D. Westmoreland,
M. J. Scott, W. H. Armstrong, J. Am. Chem. Soc. 1989, 111,
1907–1909; c) E. J. Larson, P. J. Riggs, M. L. Kirk, J. E.
Penner-Hahn, V. L. Pecoraro, J. Chem. Soc. Chem. Commun.
1992, 102–103; d) K. Wieghardt, U. Bossek, B. Nuber, J. Weiss,
J. Bonvoisin, M. Corbella, S. E. Vitols, J. J. Girerd, J. Am.
Chem. Soc. 1988, 110, 7398–7411; e) H. Chen, J. W. Faller,
R. H. Crabtree, G. W. Brudvig, J. Am. Chem. Soc. 2004, 126,
7345–7349; f) C. E. Dubé, D. W. Wright, W. H. Armstrong, J.
Am. Chem. Soc. 1996, 118, 10910–10911; g) J. M. Carroll, J. R.
Norton, J. Am. Chem. Soc. 1993, 115, 12639; h) C. E. Dubé,
D. W. Wright, S. Pal, P. J. Bonitatebus Jr, W. H. Armstrong, J.
Am. Chem. Soc. 1998, 120, 3704–3716; i) S. Mahapatra, T. K.
Lal, R. Mukherjee, Inorg. Chem. 1994, 33, 1579–1980; j) T. K.
Lal, R. Mukherjee, Inorg. Chem. 1998, 37, 2373–2382.
Several authors have proposed the involvement of a unique,
highly oxidised MnV=O species (in S4 state) in the O–O bond-
formation step (S4ǞS0). See, for example: a) J. Limburg, V. A.
Szalai, G. W. Brudvig, J. Chem. Soc. Dalton Trans. 1999, 1353–
1362; b) J. P. McEvoy, G. W. Brudvig, Phys. Chem. Chem. Phys.
2004, 6, 4754–4763; c) V. L. Pecoraro, M. J. Baldwin, M. T.
Candle, W. Y. Hsieh, N. A. Law, Pure Appl. Chem. 1998, 70,
925–929; d) P. E. M. Siegbahn, R. H. Crabtree, J. Am. Chem.
Soc. 1999, 121, 117–127; e) P. E. M. Siegbahn, Curr. Opin.
Struct. Biol. 2002, 6, 227–235; f) V. K. Yachandra, K. Sauer,
M. P. Klein, Chem. Rev. 1996, 96, 2927–2950.
ter using 1.00-cm quartz cells. The kinetics was monitored in situ
with the instrument in “kinetic mode” at 420 nm in an electrically
controlled thermostatted 25.0(Ϯ0.1) °C cell housing (CPS-240A).
Reactions with t1/2 less than 30 seconds were monitored using a
Biologic SFM-03/QS stopped-flow spectrophotometer with DT-
2801 data transmission system. The ionic strength was normally
maintained at 1.0  with NaNO3. The excess 2,2Ј-bipyridine con-
centration, Cbipy = [(Hbipy+) + (bipy)] was in the range 1–80 m
and acted as a good buffer in the pH range studied (2.0–6.0), which
was measured with an Orion-Ross Combined electrode system
(Model 81–02) before and after the reaction. Excess bipyridine
present in the reaction medium buffers the reaction within 0.05 pH
units in the entire pH range studied for nitrite oxidation while for
hydroxylamine oxidation a maximum pH increase of 0.15 units was
seen when working at a pH below 3.0 but above pH 3.0 the pH
drift was around 0.05 pH units. It is possible that HNO2 in the low
pH range and both bipyridine and hydroxylamine in the higher pH
range act as good buffers.
[7]
[8]
[9]
[10]
[11]
The electrode calibration was described earlier to read –log10[H+]
directly from a calibration curve constructed by plotting the pH-
meter reading against –log10[H+].[46] For reactions in D2O, pD was
calculated as pHmeasured + 0.40.[47] An excess of reducing agent was
maintained in all the kinetic runs. The log10(absorbance) vs. time
plots were found to be good straight lines for up to at least 90%
completion of reactions and the observed rate constants (k0) were
calculated in the usual way.
[12]
Supporting Information (see also the footnote on the first page of
this article): Linearisation of Equation (9) for the NIII redox reac-
tion and evaluation of k4 for the N–I redox reaction from a linear
plot at pH Ն 5.3.
Acknowledgments
The award of a Senior Research Fellowship (SRF) to Suranjana
Das by the Council of Scientific and Industrial Research, New
Delhi, India is gratefully acknowledged. We also thank Professor
Soumen Basak, Head, Chemical Sciences Division, Saha Institute
of Nuclear Physics (SINP), Salt Lake, Calcutta 700 064 and Dr.
Kalyan Giri of SINP for permitting us to use the stopped-flow
spectrophotometer at SINP and providing kind help.
[13]
[1] a) J. P. McEvoy, G. W. Brudvig, Chem. Rev. 2006, 106, 4455–
4483; b) J. Yano, J. Kern, K. Sauer, M. J. Latimer, Y. Pushkar,
J. Biesiadka, B. Loll, W. Saenger, J. Messinger, A. Zouni, V. K.
Yachandra, Science 2006, 314, 821; c) S. Mukhopadhyay, S. K.
Mandal, S. Bhaduri, W. H. Armstrong, Chem. Rev. 2004, 104,
3981–4026.
[2] J. W. Murray, J. Barber, Biochemistry 2006, 45, 4128–4130.
[3] Photosynthetic water oxidation: J. Nugent, Biochim. Biophys.
Acta 2001, 1503, 1 (special dedicated issue).
[4] Manganese Red-Ox Enzymes (Ed: V. L. Pecoraro), VCH Pub-
lishers, New York, 1992.
[5] J. Yano, Y. Pushkar, P. Glatzel, A. Lewis, K. Sauer, J. Mess-
inger, U. Bergmann, V. Yachandra, J. Am. Chem. Soc. 2005,
127, 14974–14975.
[6] K. N. Ferreira, T. M. Iverson, K. Maghlaoui, J. Barber, S.
Iwata, Science 2004, 303, 1831–1838; this work shows that the
OEC of PS II contains a cubane-like Mn3CaO4 cluster linked
to a fourth, distinct Mn by a mono µ-oxo bridge. However,
the conditions used for the structure determination by X-ray
crystallography may cause serious damage specifically to the
metal site structure. See, for example:J. Yano, J. Kern, K.-D.
Irrgang, M. K. Latimer, U. Bergmann, P. Glatzel, Y. Pushkar,
[14]
[15]
[16]
C. Philouze, G. Blondin, J.-J. Girerd, J. Guilhem, C. Pascard,
D. Lexa, J. Am. Chem. Soc. 1994, 116, 8557–8565.
S. Das, J. Bhattacharyya, S. Mukhopadhyay, Helv. Chim. Acta
2006, 89, 1947–1958.
a) P. Lumme, J. Tummavuori, Acta Chem. Scand. 1965, 19,
617–621; b) V. Zang, R. van Eldik, Inorg. Chem. 1990, 29,
4462–4468; c) S. Mukhopadhyay, R. Banerjee, J. Chem. Soc.
Dalton Trans. 1994, 1349–1353.
[17]
[18]
P. Lumme, P. Lahermo, J. Tummavuori, Acta Chem. Scand.
1965, 19, 2175–2188.
a) W. F. Beck, G. W. Brudvig, Biochemistry 1987, 26, 8285–
8295; b) R. D. Guiles, V. K. Yachandra, A. E. McDermott,
J. L. Cole, S. L. Dexheimer, R. D. Britt, K. Sauer, M. P. Klein,
Biochemistry 1990, 27, 486–496; c) W. F. Beck, G. W. Brudvig,
J. Am. Chem. Soc. 1988, 110, 1517–1523.
a) A. Sahay, A. Jajoo, S. Bharti, Luminescence 2006, 21, 143–
147; b) H. Spiller, P. Boger, Photochem. Photobiol. 1971, 26,
397–402; c) A. Stemler, J. B. Murphy, Plant Physiol. 1985, 77,
974–977.
[19]
4506
www.eurjic.org
© 2007 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Eur. J. Inorg. Chem. 2007, 4500–4507

Proton-Coupled Electron Transfer
[20]
[21]
[22]
[23]
[35] G. Renger, Biochim. Biophys. Acta 2004, 1655, 195–204.
[36] a) G. Davies, K. Kustin, Inorg. Chem. 1969, 8, 484–490; b) S. K.
Ghosh, R. N. Bose, E. S. Gould, Inorg. Chem. 1987, 26, 2684–
2687; c) D. W. Margerum, L. M. Schurter, J. Hobson, E. E.
Moore, Environ. Sci. Technol. 1994, 28, 331–337; d) D. W. John-
son, D. W. Margerum, Inorg. Chem. 1991, 30, 4845–4851.
[37] a) M. Simic, E. Hayon, J. Am. Chem. Soc. 1971, 93, 5982–
5986; b) J. N. Coopert, D. W. Margerum, Inorg. Chem. 1993,
32, 5905–5910; c) R. M. Liu, M. R. McDonald, D. W.
Margerum, Inorg. Chem. 1995, 34, 6093–6099.
[38] a) K. B. Wiberg, Chem. Rev. 1955, 55, 713–743; b) D. Ghosh,
A. D. Shukla, R. Banerjee, A. Das, J. Chem. Soc. Dalton Trans.
2002, 1220–1225; c) E. L. Purlee, J. Am. Chem. Soc. 1959, 81,
263–272; d) R. P. Bell, in Acid-Base Catalysis, Oxford Univer-
sity Press, London, 1941, p. 145; e) R. P. Bell, in The Proton in
Chemistry, Chapman and Hall, London, 1973; f) D. M. Quinn,
L. D. Sutton, in Enzyme Mechanism from Isotope Effects (Ed.:
P. F. Cook), CRC Press, Boca Raton, 1991, chapter 3, pp. 73–
126; g) K. B. Schowen, R. L. Schowen, in Methods in Enzy-
mology (Ed.: D. L. Purich), Academic Press, Inc., New York,
1982, vol. 87, pp. 551– 606.
J. Bhattacharyya, S. Mukhopadhyay, Helv. Chim. Acta 2005,
88, 2661–2674.
A. Das, S. Mukhopadhyay, Helv. Chim. Acta 2005, 88, 2561–
2572.
R. M. Smith, A. E. Martell, in Critical Stability Constants, Ple-
num, New York, 1976, vol. 2, p. 235.
CV was measured using the relation: CV = sdϫ100/x, where
sd is the standard deviation of the measurements of k0 and x
is the average k0. See: G. H. Jeffery, J. Bassett, J. Mendham and
R. C. Denney, in Vogel’s Textbook of Quantitative Chemical
Analysis, Longman Scientific and Technical, England, 5th ed.,
1989, p. 135.
a) P. Bandyopadhyay, B. B. Dhar, J. Bhattacharyya, S. Mukho-
padhyay, Eur. J. Inorg. Chem. 2003, 4308–4312; b) K. Lemma,
A. M. Sargeson, L. I. Elding, J. Chem. Soc. Dalton Trans. 2000,
2267–2272; c) A. Das, S. Mukhopadhyay, Polyhedron 2004, 23,
895–901.
[24]
[25]
[26]
N. S. Bayliss, R. Dingle, D. W. Watts, R. J. Wilkie, Aust. J.
Chem. 1963, 16, 933–942.
M. S. Ram, D. M. Stanbury, Inorg. Chem. 1985, 24, 2954–2962.
[27] a) M.-N. C. Dunand-Sauthier, A. Deronzier, A. Piron, X. Pra-
don, S. Ménage, J. Am. Chem. Soc. 1998, 120, 5373–5380; b)
M.-N. C. Dunand-Sauthier, A. Deronzier, A. Piron, J. Elec-
troanal. Chem. 1999, 463, 119–122; c) M.-N. C. Dunand-Sauth-
ier, A. Deronzier, A. Piron, X. Pradon, J. Am. Chem. Soc. 1997,
119, 3173–3174; d) C. Baffert, S. Romain, A. Richardot, J.-C.
Lepretre, B. Lefebvre, A. Deronzier, M.-N. C. Dunand-Sauth-
ier, J. Am. Chem. Soc. 2005, 127, 13694–13704.
[39] M. Byrdin, V. Sartora, A. P. M. Eker, M. H. Vos, C. Aubert,
K. Brettel, P. Mathis, Biochim. Biophys. Acta 2004, 1655, 64–
70.
[40] a) B. A. Diner, D. A. Force, D. W. Randall, R. D. Britt, Bio-
chemistry 1998, 37, 17931–17934; b) C. Jeans, M. J. Schilstra,
N. Ray, S. Husain, J. Minagawa, J. H. A. Nugent, D. R. Klug,
Biochemistry 2002, 41, 15754–15761; c) M. Haumann, A. Mul-
kidjanian, W. Junge, Biochemistry 1999, 38, 1258–1267; d) C.
Berthomieu, R. Hienerwadel, A. Boussac, J. Breton, B. A. Di-
ner, Biochemistry 1998, 37, 10547–10554.
[28] The oxidation potentials for HNO2 Ǟ NO2 + H+ + e– and for
NH3OH+ Ǟ NH2OH+ + H+ + e– were calculated from the
–
respective potentials for the NO2 /NO2 and NH2OH/NH2OH+ [41] S. Banerjee, U. Roy Choudhury, R. Banerjee, S. Mukho-
couples and the pKa values of HNO2 and NH3OH+.
[29] M.-L. Hung, M. L. McKee, D. M. Stanbury, Inorg. Chem.
2001, 40, 2206.
padhyay, J. Chem. Soc. Dalton Trans. 2002, 2047–2052.
[42] J. Bhattacharyya, K. Dutta, S. Mukhopadhyay, Dalton Trans.
2004, 2910–2917.
[30] G. Blondin, R. Davydov, C. Philouze, M. F. Charlot, S. Styr-
ing, B. Åkermark, J. J. Girerd, A. Boussac, J. Chem. Soc. Dal-
ton Trans. 1997, 4069–4074.
[43] a) G. Charlot, in Colorimetric Determination of Elements: Prin-
ciples and Methods, Elsevier Publishing Company, New York,
1964, p. 324; b) N. G. Bunton, N. T. Crosey, S. J. Patterson,
Analyst 1969, 94, 585–588.
[44] G. H. Jeffery, J. Bassett, J. Mendham and R. C. Denney, in Vo-
gel’s Textbook of Quantitative Chemical Analysis, Longman
Scientific and Technical, England, 5th edn., 1989, p. 408.
[45] Reference 44, p. 334.
[46] a) S. Banerjee, U. Roy Choudhury, B. C. Ray, R. Banerjee, S.
Mukhopadhyay, Anal. Lett. 2001, 34, 2797–2815; b) H. N. Ir-
ving, M. G. Miles, L. D. Pettit, Anal. Chim. Acta 1967, 38, 475–
488.
[31] a) W. H. Albery, in Proton Transfer Reactions (Eds.: E. Caldin,
V. Gold), Wiley, New York, 1975, ch. 9; b) J. A. Gilbert, S. W.
Gersten, T. J. Meyer, J. Am. Chem. Soc. 1982, 104, 6872–6873;
c) R. A. Binstead, M. E. McGuire, A. Dovletoglou, W. K.
Seok, L. E. Roecker, T. J. Meyer, J. Am. Chem. Soc. 1992, 114,
173–186; d) I. J. Rhile, J. M. Mayer, J. Am. Chem. Soc. 2004,
126, 12718–12719; e) M. Sjödin, S. Styring, H. Wolpher, Y. Xu,
L. Sun, L. Hammarström, J. Am. Chem. Soc. 2005, 127, 3855–
3863.
[32] W. H. Albery, M. H. Davis, J. Chem. Soc. Faraday Trans. 1972,
[47] a) P. K. Glasoe, F. A. Long, J. Phys. Chem. 1960, 64, 188–190;
b) P. Salomaa, L. L. Schaleger, F. A. Long, J. Am. Chem. Soc.
1964, 86, 1–7; c) A. K. Covington, R. A. Robinson, R. G.
Bates, J. Phys. Chem. 1966, 70, 3820–3824.
68, 167–181.
[33] D. M. Kurtz Jr, Chem. Rev. 1990, 90, 585–606.
[34] a) M. J. Baldwin, V. L. Pecoraro, J. Am. Chem. Soc. 1996, 118,
11325–11326; b) J. S. Vrettos, J. Limburg, G. W. Brudvig, Bi-
ochim. Biophys. Acta 2001, 1503, 229–245.
Received: April 2, 2007
Published Online: August 10, 2007
Eur. J. Inorg. Chem. 2007, 4500–4507
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