Kinetics of oxidation of hydrogen peroxide and ascorbic acid by a tribridged manganese(IV,IV) dimer in feebly acidic media

Article abstract of DOI:10.1139/v99-062

In weakly acidic, aqueous buffer (MeCO₂^-+ bipy), the complex ion [Mn₂^IV(μ-O)₂(μ-MeCO ₂)(bipy)₂(H₂O)₂]³⁺, 1 (bipy = 2,2′-bipyridine), coexists in rapid equilibrium with its hydrolytic derivatives, [Mn₂^IV(μ-O)₂(bipy)₂(H ₂O)₄]⁴⁺, 2, and [Mn₂^IV(μ-O)₂(μ-MeCO ₂)(bipy)(H₂O)₄]³⁺, 3. The solution quantitatively oxidizes hydrogen peroxide to oxygen and ascorbic acid to dehydroascorbic acid, itself being reduced to Mn^II. In the presence of excess reductant, the reactions follow simple first-order kinetics with no evidence for the accumulation of a significant amount of any intermediate manganese complex. The ascorbate anion shows overwhelming kinetic dominance over ascorbic acid, but no evidence is available for deprotonation of hydrogen peroxide. The preferred intimate mechanism for hydrogen peroxide is inner sphere but that for ascorbic acid is uncertain. For both reductants, increased extent of aquation leads to increased kinetic activity in the order: 1 < 2 < 3.

Full text of DOI:10.1139/v99-062

451
Kinetics of oxidation of hydrogen peroxide and
ascorbic acid by a tribridged manganese(IV,IV)
dimer in feebly acidic media
Anup Kumar Bhattacharya, Anath Bondhu Mondal, Anadi C. Dash,
G.S. Brahma, and Rupendranath Banerjee
–
Abstract: In weakly acidic, aqueous buffer (MeCO2 + bipy), the complex ion [Mn2IV(µ-O)2(µ-MeCO2)(bipy)2(H2O)2]3+,
1 (bipy = 2,2′-bipyridine), coexists in rapid equilibrium with its hydrolytic derivatives, [Mn2IV(µ-O)2(bipy)2(H2O)4]4+, 2,
and [Mn2IV(µ-O)2(µ-MeCO2)(bipy)(H2O)4]3+, 3. The solution quantitatively oxidizes hydrogen peroxide to oxygen and
ascorbic acid to dehydroascorbic acid, itself being reduced to MnII. In the presence of excess reductant, the reactions
follow simple first-order kinetics with no evidence for the accumulation of a significant amount of any intermediate
manganese complex. The ascorbate anion shows overwhelming kinetic dominance over ascorbic acid, but no evidence
is available for deprotonation of hydrogen peroxide. The preferred intimate mechanism for hydrogen peroxide is inner
sphere but that for ascorbic acid is uncertain. For both reductants, increased extent of aquation leads to increased
kinetic activity in the order: 1 < 2 < 3.
Key words: kinetics, manganese, ascorbic acid, hydrogen peroxide, 2,2′-bipyridine.
Résumé : En milieu faiblement acide, tampon aqueux (acétate + bipyridine), le complexe ionique [Mn(IV)2(µ-O)(µ-
MeCO2)(bipy)2(H2O)2]3+ (1 dans lequel bipy = 2,2’-bipyridine) coexiste en équilibre rapide avec ses dérivés
d’hydrolyse, [Mn(IV)2(µ-O)(bipy)2(H2O)4]4+ (2) et [Mn(IV)2(µ-O)2(µ-MeCO2)(bipy)2(H2O)4]3+ (3). La solution oxyde
quantitativement le peroxyde d’hydrogène en oxygène et l’acide ascorbique en acide déhydroascorbique alors que le
manganèse est réduit en Mn(II). En présence d’un excès de réducteur, les réactions suivent une cinétique du premier
ordre simple, sans accumulation de quantités significatives de complexes de manganèse intermédiaires. L’anion
ascorbate présente une dominance cinétique importante par rapport à l’acide ascorbique, mais on n’a pas pu mettre en
évidence de déprotonation du peroxyde d’hydrogène. La sphère interne correspond au mécanisme intime préféré pour
le peroxyde d’hydrogène; celui pour l’acide ascorbique est toutefois incertain. Pour les deux réducteurs, une
augmentation de l’aquation conduit à une activité cinétique accrue dans l’ordre 1 < 2 < 3.
Mots clés : manganèse, acide ascorbique, peroxyde d’hydrogène, 2,2′-bipyridine.
[Traduit par la Rédaction] Bhattacharya et al. 458
Introduction
In addition to their significance in light-driven oxidation
of water molecules to oxygen, these complexes are novel
oxidants useful for the mechanistic studies of electron trans-
fer reactions of higher-valent manganese in solution. We re-
cently investigated the kinetics of reduction of two dioxo-
bridged manganese(III,IV) dimers (7) and the tribridged
[Mn2IV(µ-O)2(µ-MeCO2)(bipy)2(H2O)2]3+ ion (1; bipy =
2,2′-bipyridine; see Fig. 1) in aqueous media (8–10). We
The oxygen-evolving complex (OEC) in photosystem II
(PS II) consists of a manganese cluster that binds substrate
water molecules and accumulates oxidizing equivalents in a
cyclic sequence of reactions known as the Kok or the S-state
cycle (1). The OEC is made up of four manganese ions (2),
held together by µ-oxo and carboxylate bridges, and a tetra-
meric structure was recently proposed on the basis of
EXAFS and EPR data (3). ESSEM 14N/15N labelling studies
(4) and 15N ENDOR spectroscopy (5) have shown that nitro-
gen-containing ligands are bound to manganese in PS II.
Oxo- and carboxylato-bridged manganese dimers (6) con-
taining N-donor ligands are therefore considered as “molecular
bricks” for the OEC. Investigation of the kinetics, equilib-
rium, and mechanism of the ligand substitution and redox
reactions of such complexes is intrinsically interesting.
IV
thus found that the Mn2 species in excess bipy generate a
MnIIIMnIV dimer, [(bipy)2MnIV(µ-O)2MnIII(bipy)2]3+, as in-
termediate while reacting with thiosulfate and hydrazine but
not with nitrite. From the observed proton dependence, the
intimate mechanism for hydrazine appeared to be inner
sphere. In this paper, we report the kinetics of reduction of 1
and its aqua derivatives by ascorbic acid (H2A) and hydro-
gen peroxide in an attempt to further explore the reactivity
pattern of higher-valent oxo-bridged manganese complexes.
Received July 6, 1998.
A.K. Bhattacharya, A.B. Mondal, and R. Banerjee.1 Department of Chemistry, Jadavpur University, Calcutta-700032, India.
A.C. Dash1 and G.S. Brahma. Department of Chemistry, Utkal University, Bhubaneswar-751004, India.
1Authors to whom correspondence may be addressed. e-mail (RB): rbju@yahoo.com; (ACD): dsthcr28@cal2.vsnl.net.in
Can. J. Chem. 77: 451–458 (1999)
© 1999 NRC Canada

452
Can. J. Chem. Vol. 77, 1999
IV
a
Fig. 1. Graphical structure for 1 drawn on the basis of its crystal
Table 1. Stoichiometry of the reduction of Mn2 complexes.
structure.
IV
[Mn2
]
pH
[R]
∆ [Mn2IV]/∆[R]
R = H2O2
0.8
1.4
1.8
2.4
0.2
0.6
0.8
1.0
1.2
4.53
4.79
4.85
4.96
5.12
0.48
0.52
0.49
0.47
0.51
2.8
Avg. = 0.495 ± 0.02
R = ascorbic acid
0.2
0.4
0.9
1.2
1.6
4.48
4.52
4.87
5.05
5.23
0.7
1.2
2
2.8
3.8
0.49
0.47
0.52
0.48
Experimental
Materials
0.51
The solutions of hydrogen peroxide were prepared by di-
lution of 30% (w/v) stabilizer-free hydrogen peroxide (G.R.,
E. Merck) and were standardized by iodometry using starch
indicator (11). L-Ascorbic acid (A.R., S.R.L) was used with-
out further purification. All other materials including the
complex ion 1 have been described earlier (8). Solutions
were prepared in doubly distilled water, generally just before
use.
Avg. = 0.496 ± 0.02
aCbipy = 35.0; all concentrations are in mmol dm–3 ; total volume, 100
mL.
most common oxidation product (14, 15) of H2A, is pro-
duced in the present experiments.
Solution equilibria
Physical measurements and kinetics
The crystal structure of the complex salt 1(ClO4)3·H2O
shows a bridging ethanoate ligand, a bis(µ-oxo) bridge, and
two water molecules, each directly coordinated to one man-
ganese(IV) ion (16). The ethanoate bridge and a bipy ligand
are labile and dissociate in aqueous solution according to
eqs. [1] and [2] (8).
Kinetics were measured with a stopped-flow spectropho-
tometer (HITECH SF-51), interfaced with an Apple-II GS
personal computer for data acquisition and analysis of
absorbance–time data for the first-order rate constants. Five
hundred and twelve data points were collected by the stopped-
flow apparatus for each experiment. Excess reducing agent
IV
maintained pseudo-first-order conditions in
a
mixed
3+ +2H2O
[Mn2 (µ −O)2(µ − MeCO2)(bipy)2(H2O)2 ]
(MeCO2H/MeCO2 + Hbipy+/bipy) aqueous buffer medium
in the pH range 4.2–5.5. Each reported first-order rate con-
stant is the average of three to five replicate measurements
for a given reaction mixture. We measured the solution pH with
an Elico (L 1120) pH-meter (8, 12). The linearity of the
electrode was established using pH 4, 7, and 9 buffers. The
electrode was calibrated to read –log [H+] directly using a
series of acid solutions at the ionic strength used for kinetic
measurements. Kinetics was measured at 25.0oC for hydro-
gen peroxide, but at 15.0oC for fast-reacting ascorbic acid.
Stoichiometries were determined under the kinetic condi-
tions. An excess of a reducing agent, hydrogen peroxide or
ascorbic acid, was added in a single portion to a solution of
–
[1]
1
K12
º
IV
−
[Mn2 (µ −O)2(bipy)2(H2O)4]4+ +MeCO2
2
K13
[2]
1 + 2 H2O
º
IV
3+ +bipy
[Mn2 (µ −O)2(µ − MeCO2)(bipy)(H2O)4]
3
Kea
º
−
[3]
[4]
MeCO2H
Hbipy+
MeCO2 + H+
IV
the Mn2 complex 1. The amount of hydrogen peroxide re-
maining after complete reaction was determined iodome-
trically(11). Ascorbate concentration in spent reaction
mixtures was determined iodimetrically using starch indica-
tor near the end point in a dilute sulphuric acid medium
(13). Standard iodine solution rapidly and quantitatively oxi-
dized ascorbic acid to dehydroascorbic acid (13, 14). Other
materials present in the mixture did not interfere.
KHbipy
H+ + bipy
º
The solution is otherwise stable under the present experi-
mental conditions.
A labile ethanotate bridge is found also in the dinuclear
Mn(III,IV) complex [Mn2(µ-O)2(µ-MeCO2)(fac-bpea)2]2+
(bpea is N,N-bis(2-pyridylmethyl)ethylamine) and in its diferric
analogue, [Fe2(µ-O)(µ-MeCO2)2(L)2]2+ (L is substituted bpea)
(17–19).
Results and discussion
The chelating ligands in the complexes [L′2MnIII(µ-
O)2MnIVL′2]3+ (L′ = 2,2′-bipyridine and 1,10-phenanthroline)
are also labile (7, 20). The binding sites in the S3-state (the
S-state notation (1) identifies the number of oxidizing equiv-
alents stored in the manganese cluster in the water oxidation
Stoichiometry and reaction products
Each mole of the complex consumed two moles of H2O2
and thus suggests quantitative oxidation of H2O2 to O2 (Ta-
ble 1). The 1:2 stoichiometry was also observed for ascorbic
acid (H2A), indicating that dehydroascorbic acid (A), the
© 1999 NRC Canada

Bhattacharya et al.
453
Fig. 2. Variation of first-order rate constants for H2O2 with pH at a fixed Cbipy(0.035) and fixed Cea(0.08) at 25.0oC (᭞). Variation of
first-order rate constants for ascorbic acid with pH at a fixed Cbipy (0.035) and Cea (0.1) at 15.0oC (᭺).
complex of PS II) are labile and exchange 18O-labeled water
in less than 1 s (21). However, the bridging oxo-ligands are
kinetically inert (22, 23). Interestingly, equilibrium as well
as kinetic studies (8, 10) suggest that the coordinated water
molecules in 1 and its aqua derivatives do not deprotonate,
at least up to pH 5.5. Weak acidity of coordinated water is
known also for [Mn3IV(µ-O)4(bipy)4(H2O)2]4+ (24) and
[Mn(pd)2(H2O)2]+ (pd is pentane-2,4-dione). The latter com-
plex has a pKa value of 7.3 (25–27). The weak acidity ob-
served is consistent with recent findings that the O—H bond
strength in water coordinated to higher-valent manganese is
comparable to the phenolic O—H bond strength in tyrosine
(28, 29).
ria [1]–[4], the deprotonation equilibrium of H2A, eq. [8]
(15), and an overwhelming kinetic dominance of the
ascorbate ion, HA–, over ascorbic acid, H2A. (Scheme 2,
eqs. [8]–[11]).
Scheme 1.
kA
→
[5]
1 + H2O2
Products (MnII, H+, O2)
k
B
[6]
[7]
2 + H2O2
3 + H2O2
→
Products
Products
kC
→
Kinetics and reaction schemes
All reactions in the presence of excess reducing agent
obeyed first-order kinetics for at least 90% reaction. The
first-order rate constants, k0 for H2O2 and k0′ for ascorbic
acid (see Table 2), increased with increasing concentration
of the reducing agent but decreased with increasing [bipy]
Scheme 2.
Ka
º
[8]
[9]
H2A
H+ + HA
–
and [MeCO2 ] at a fixed pH. However, pH change affects k0
′
kA
and k0′ differently. Increased pH decreases k0 but increases
k0′ (Fig. 2). The variations in k0 can be explained on the ba-
sis of only equilibria [1]–[4], provided the hydrolytic deriva-
tives 2 and 3 are kinetically more active than their parent, 1.
We therefore propose Scheme 1 as a simple model for the
reaction of H2O2.
1 + HA–
→
Products (MnII, H+, A)
Products
′
kB
[10] 2 + HA–
[11] 3 + HA–
→
′
kC
The pH dependence of the reactions of ascorbic acid can
be reasonably explained if one assumes, along with equilib-
→
Products
© 1999 NRC Canada

454
Can. J. Chem. Vol. 77, 1999
IV
a
Table 2. First-order rate constants for the reduction of Mn2 complexes.
–
pH
[bipy]
[MeCO2 ]
[R]
k0/s–1
Reducing agent (R), H2O2
4.25
4.50
4.63
4.79
5.06
5.23
5.34
5.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
14.1
19.1
21.6
24.5
28.5
30.3
31.2
32.3
2.73
5.46
8.19
9.55
10.9
13.65
16.4
19.1
19.1
19.1
19.1
19.1
19.1
19.1
19.1
19.1
19.1
19.1
19.1
19.1
19.1
19.4
29.0
34.8
42.1
53.9
60.3
63.8
68.0
29.0
29.0
29.0
29.0
29.0
29.0
29.0
1.81
3.62
5.44
7.26
9.07
12.9
36.3
54.4
29.0
29.0
29.0
29.0
29.0
29.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.5
2.0
2.5
3.0
4.0
5.0
1.16 (1.13)
1.01 (0.99)
0.93 (0.91)
0.84 (0.82)
0.74 (0.73)
0.71 (0.69)
0.69 (0.68)
0.67 (0.66)
2.81 (2.68)
1.81 (1.78)
1.37 (1.43)
1.34 (1.32)
1.23 (1.24)
1.12 (1.26)
1.06 (1.09)
1.75 (1.76)
1.59 (1.61)
1.49 (1.50)
1.39 (1.41)
1.34 (1.33)
1.20 (1.22)
0.92 (0.94)
0.87 (0.86)
1.52 (1.48)
2.04 (1.98)
2.41 (2.47)
3.02 (2.97)
3.89 (3.96)
4.88 (4.95)
Reducing agent, ascorbic acid
–
pH
[bipy]
14.1
19.1
22.4
23.8
27.7
29.3
30.5
32.3
2.73
5.46
8.19
9.55
13.65
17.2
19.1
19.1
19.1
19.1
19.1
19.1
19.1
19.1
19.1
19.1
19.1
[MeCO2 ]
[R]
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.5
2.0
2.5
3.0
4.0
10–1 k0′/s–1
3.07 (3.05)
3.54 (3.56)
3.81 (3.83)
3.99 (3.94)
4.25 (4.19)
4.35 (4.29)
4.39 (4.36)
4.41 (4.45)
5.18 (5.20)
4.40 (4.24)
3.99 (3.92)
3.84 (3.83)
3.65 (3.69)
3.62 (3.61)
3.99 (3.98)
3.91 (3.90)
3.87 (3.84)
3.68 (3.69)
3.57 (3.61)
3.43 (3.51)
5.21 (5.33)
7.21 (7.12)
8.96 (8.89)
10.5 (10.7)
14.3 (14.2)
4.25
4.50
4.67
4.75
5.00
5.13
5.25
5.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
4.50
24.2
36.3
45.8
50.3
64.3
70.8
76.2
85.1
36.3
36.3
36.3
36.3
36.3
36.3
1.81
3.63
5.44
13.8
23.6
54.4
36.3
36.3
36.3
36.3
36.3
© 1999 NRC Canada

Bhattacharya et al.
455
Table 2. (concluded).
–
pH
[bipy]
[MeCO2 ]
[R]
k0/s–1
4.50
4.50
19.1
19.1
36.3
36.3
4.5
5.0
16.2 (16.0)
17.8 (17.8)
aTemperature, 25.0oC for H2O2; 15.0oC for ascorbic acid. [Mn2IV] = 0.10; all concentrations are in mmol dm–3 except I = 1.0 mol dm–3 (NaNO3); λ, 420
nm for H2O2; 380 nm for ascorbic acid. Values in parentheses are calculated rate constants.
Schemes 1 and 2 lead to eqs. [12] and [13], respectively.
{kA[MeCO2− ][bipy] + kBK12[bipy] + kCK13[MeCO2− ]}[H2O2 ]
[12] k0 =
−
[MeCO2− ][bipy] + k12[bipy] + k13[MeCO2
]
{KakA′ [MeCO2− ][bipy] + KakB′ K12[bipy] + KaK13kC′ [MeCO2− ]}Ca
{[H+ ] + Ka}{[MeCO2− ][bipy] + K12[bipy] + K13[MeCO2− ]}
[13] k0′ =
where Ca (= [H2A] + [HA–]) is the total analytical concentration for ascorbic acid. Equation [12] leads to eq. [14] and eq. [13]
leads to eq. [15].
[14] k0([MeCO2−][bipy]+ K12[bipy] + K13[MeCO2−])/[H2O2]
−
= kA[MeCO2−] [bipy] + kBK12[bipy] + kCK13[MeCO2
]
]
[15] k0′([H+] + Ka)([MeCO2−][bipy] + K12[bipy] + K13[MeCO2−])/CaKa
= kA′ [MeCO2−] [bipy] + kB′ K12[bipy] + kC′ K13 [MeCO2
−
–
[MeCO2 ] = Cea Kea / ([H+] + Kea) and [bipy] = Cbipy KHbipy
/
(31) have shown that oxidation of ascorbic acid by
[Fe(H2O)6]3+ ion follows an outer-sphere path, and oxidation
by the [Fe(H2O)5(OH)]2+ ion proceeds by an H-atom transfer
from HA– although both H2A and HA– are known (32) to be
good ligands for Fe(III). Obviously, presence of a labile co-
ordination site and the ability of the reductant to bind to a
metal centre are insufficient information to decide in favour
of one of the two intimate mechanisms — inner sphere or
outer sphere. However, an H-atom transfer mechanism may
be discarded since the Mn2IV complexes have no OH group.
–
–
([H+]+KHbipy), where Cea = ([MeCO2 ] + [MeCO2 ]) and
C
bipy = ([Hbipy+] + [bipy]) are the total analytical concentra-
tion of ethanoate and 2,2′-bipyridine, respectively. Values for
the left-hand sides of eqs. [14] and [15] may be calculated
using known rate constants (Table 2) and equilibrium con-
stants (Table 3). The calculated left-hand sides of eqs. [14]
and [15] each yielded excellent straight lines (Figs. 3 and 4)
–
when plotted against [bipy] or [MeCO2 ], thus justifying the
proposed schemes. The slopes and intercepts of the straight
lines yielded the specific rate constants kA, kB, kC, kA′, kB′
and kC′ displayed in Table 4 along with known rate constants
for comparable reactions. The rate law derived for an equiv-
alent scheme admitting significant kinetic contribution from
H2A contains an additional term, first order in [H+]. Such a
rate law does not fit the experimental data.
The simple first-order kinetics observed by us, without
any indication of rate saturation, only implies the validity of
the inequality, Qx[HA–] << 1 for every individual value of
Qx, the formation quotient for adducts that may form be-
tween HA– and the MnIV complexes, 1, 2, 3. In our experi-
2
ments the maximum value for [HA–] is 3.6 × 10–3 mol dm–3.
The inequality, therefore, imposes an upper limit on Qx,
30 dm3 mol–1. The observed situation is similar to those for
reduction of [Mn2III,IV(µ-O)2(bipy)4]3+, 4, by hydroquinone
The K12 and K13 values are available only at 25.0oC and
were used in computing the kn′ values for ascorbic acid at
15.0oC. This difference in temperature is unlikely to intro-
duce serious error. We verified that a simultaneous increase
or decrease to the extent of 50% in K12 and K13 changes k1′
by only 12% even when k2′ and k3′ are kept fixed to see what
happens if the whole thrust of the change in the two equilib-
rium constants are put on k1′ alone. Similar were the cases
for k2′ and k3′.
–
and HSO3 . No rate saturation was observed in these two
reactions, as was also the case in the reaction between
HSO3– and 4, and that between HA– and the one-electron re-
–
duced MnIII form of 4 (33). However, the (MnIII2)–HSO3
adduct (Q =2450 dm3 mol–1), as well as that formed from 4
and ascorbic acid, is strong enough to permit their kinetic
detection (33). Apparently, the Q values may be sufficiently
different to evade rationalization even in closely related sys-
tems.
Intimate mechanism
Ascorbic acid can act both as an inner-sphere as well as
an outer-sphere reducing agent (30). In the present study, la-
bile coordination sites present in the metal complexes may
induce an inner-sphere mechanism but it is not a trivial mat-
ter to pinpoint the mechanism. For example, Bansch et al.
Oxidation of hydrogen peroxide is generally inner sphere
(34) excepting when the oxidant is very powerful and
coordinatively saturated, such as [NiIII(bipy)3] (35). For ex-
ample, the reactions of H2O2 with moderately oxidizing (36)
© 1999 NRC Canada

456
Can. J. Chem. Vol. 77, 1999
–
Fig. 3. Variation of k0 G / [H2O2] with [MeCO2 ] at a fixed pH (4.50) and fixed Cbipy (0.035) (᭿). Variation of k0 G / [H2O2] with
–
–
[bipy] at fixed pH (4.50) and fixed Cea (0.08) (᭹). G = [MeCO2 ][bipy] + K12[bipy] + K13[MeCO2 ]. Other conditions: [Mn2IV], 1.0 ×
10–4; I, 1.0; [H2O2], 1.0 × 10–3; temp., 25.0oC. All concentrations are in mol dm–3.
nated H2O2 and enhanced acidity. We recall that water mole-
cules coordinated to Mn(IV) in 1 and its hydrolytic
derivatives release no H+ in the experimental range of pH.
Similar behaviour is definitely expected for coordinated H2O2
and is not in contradiction with the proposed mechanism.
Table 3. Values of the equilibrium constants used.
o
Reaction
Temp./ C
I / mol dm–3
pK
2.0a
3.0a
4.55b
4.43c
4.49d
4.08e
Equation [1]
Equation [2]
Equation [3]
Equation [4]
25
25
20
25
15
15
1.0
1.0
1.0
0.33
0.1
1.0
IV
Table 4 shows that the kinetic activity of the Mn2 com-
plexes towards various reducing agents increases with in-
creased extent of aquation in the order: 1 < 2 < 3. Limburg
et al. (38) as well as Banerjee and co-workers (7) noticed
that, kinetically, 4 is a much inferior oxidant than its sol-
volyzed derivatives. Again, in a recent model for the mecha-
nism of the OEC in PS II, it has been assumed that only the
two terminal manganese atoms containing coordinated water
in a d-shaped tetranuclear cluster are redox active; the two
central Mn4+ ions without coordinated water are inactive
(23). A rationale for the greater kinetic activity of the aqua
complexes was presented earlier (9).
Equation [8]
aReference 8 in the text.
bReference 43. Change in temperature little affects the value; e.g, for
I → 0, the value over the temperature range 15–25oC is 4.756 ± 0.002
(44).
cReference 45. Effect of I on pK is small; ±0.01 unit for a change in I
from 0.33 to 0.1 mol dm–3; cf. ref. 46.
dReference 47. This value was used in calculations for ascorbic acid.
Comparison with footnote c shows that the effect of temperature on pK is
small.
eReference 15 in the text.
IV
2–
+
Reduction of the Mn2 complexes by S2O3 and N2H5
produced the MnIVMnIII complex, 4, as an intermediate, which
subsequently oxidized additional molecules of reductant and
thus decayed to MnII. No intermediate manganese complex
could be detected in the present experiments. This, and the
observed exponential profiles along with the absence of any
absorbance drops immediately after mixing, suggests that
generation of 4 and its subsequent reactions with ascorbic
acid and hydrogen peroxide are kinetically insignificant in
the present systems. It may be that the one-electron transfer
complexes such as [Mn(cydta)(H2O)]– (cydta is cyclohexyl
ethylenediaminetetraacetate) (37) and solvolyzed 4 follow
the inner-sphere path (7d, 36). Oxidation of potassium
peroxymonosulfate to O2 with solvolyzed 4 and its close
analogues also follows the inner-sphere mechanism (38),
and the same mechanism appears very likely for the reac-
tions under the present investigation. Earlier examples of the
IV
−
inner-sphere mechanism for the Mn2 complexes are their
products, the radicals HO2 and A (generated in the rate-
+
determining steps) rapidly react with the MnIVMnIII species.
This is a strong possibility if the reactions are inner sphere,
when rapid further reduction of the successor complex might
reactions with N2H5 (10).
In this study, no evidence is available for deprotonation of
H2O2, though an inner-sphere mechanism implies a coordi-
© 1999 NRC Canada

Bhattacharya et al.
457
–
Fig. 4. Variation of k0′ D {[H+] + Ka} / CaKa with [MeCO2 ] at a fixed pH (4.50) and fixed Cbipy (0.035) (ٗ). Variation of k0′ D {[H+]
–
–
+ Ka} / CaKa [bipy] at a fixed pH (4.50) and fixed Cea (0.1) (᭺). D = {[MeCO2 ][bipy] + K12[bipy] + K13[MeCO2 ]}. Other
conditions: [Mn2IV], 1.0 × 10–4; I, 1.0; Cea, 1.0 × 10–3; temp.,15.0oC. All concentrations are in mol dm–3.
IV
Table 4. Second-order rate constants (dm3 mol–1 s–1) for the reduction of different Mn2 species by different reducing agents (R).
Rate constant for
R
1
2
3
Ref.
–
NO2
(2.84 ± 0.13) × 10–3
—
2.0 ± 0.02
(1.70 ± 0.08) × 102
(4.22 ± 0.27) × 104
0.149 ± 0.007
2.07 ± 0.12
10
11
9
+
N2H5
(1.11 ± 0.08) × 10–3
40 ± 1
6.33 ± 0.2) × 10–3
(1.45 ± 0.04) × 102
(1.02 ± 0.07) × 104
(9.20 ± 0.45) × 104
2–
S2O3
H2O2
HA–
(1.97 ±0.06) × 103
(5.62 ± 0.31) × 104
This work
This work
occur before the radicals can escape into the bulk solvent
(kind suggestion of a reviewer; see also ref. 39).
Such a mechanism, where the radicals are generated and
consumed within the coordination sphere of the metal com-
plex, is, however, closely analogous to and rarely distin-
strate ligands according to the Hoganson and Babcock model
(23). It appears that a mechanism involving a pair of elec-
tron acceptors becomes necessary when the driving force for
electron transfer from a multielectron donor becomes too
small for free-radical production (41). In such situations,
binuclear complexes may provide a low-energy pathway for
two-electron transfer. Nevertheless, ascorbic acid rarely un-
dergoes two-electron transfer (42). One suggested example
involves a ferriporphyrin dimer(41).
guishable from
a
simultaneous two-electron transfer
mechanism, proposed for different multinuclear manganese
complexes. For example, Wieghardt et al. (40) isolated a
peroxo-bridged complex, [L2′Mn2IV(µ-O)2(µ-O2)]2+ (L′ = tri-
methyl-1,4,7-triazacyclononane), which releases molecular
oxygen rapidly in aqueous solution at 20oC by an intra-
2–
molecular two-electron oxidation of coordinated O2 with
Conclusions
IV
III
concomitant reduction of Mn2 to a binuclear Mn2 com-
plex. Peroxo-bridged Mn4 clusters have recently been pro-
posed to represent the S4 state, in which intramolecular two-
electron oxidation of the peroxo-bridge produces O2 and the
S0 state is regenerated (23). The S3 to S4 transition may also
be described as a probable two-electron oxidation of sub-
Rapid reduction of the complex ion [Mn2IV(µ-O)2(µ-
MeCO2)(bipy)2(H2O)2]3+ and its hydrolytic derivatives with
hydrogen peroxide seems to be inner sphere, but the intimate
mechanism for the reduction by ascorbic acid is uncertain.
An increased number of coordinated water molecules in-
© 1999 NRC Canada

458
Can. J. Chem. Vol. 77, 1999
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Acknowledgment
R.B. is grateful to the Department of Science and Tech-
nology (New Delhi) for sponsoring the project under which
the present work was done.
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