Kinetics and mechanism of reduction of a coordinated superoxide with hydroxylamine derivatives

Article abstract of DOI:10.1016/j.poly.2010.07.006

Rates for the uncatalyzed reactions of the coordinated superoxide in μ-superoxo bis[pentaminecobalt(III)]⁵⁺ complex (1) with hydroxylamine-N-monosulfonate, HONH(SO₃)^- (HMS) and hydroxylamine-N,N′-disulfonate, HON(SO3)22- (HDS) have been determined. Successive replacement of the -NH₂ protons with SO₃H group (electron withdrawing) increased the reaction rate from HMS to HDS but replacement of the O-H hydrogen halted the reaction. HMS and HDS are oxidized to a mixture of N₂O (20%), NO (80%) and SO₄^2-. The reactions are first-order in [1] and [reductant]. Reaction rate increased with pH though neither HMS nor HDS are involved in protic equilibria within the investigated pH range. We propose 1 is in protic equilibrium with its conjugate acid 1H which is kinetically a dead-end species. An electroprotic reaction between 1 and the -OH hydrogen in HMS is proposed whereas a simple electron-transfer mechanism has been proposed for the HDS oxidation.

Full text of DOI:10.1016/j.poly.2010.07.006

Polyhedron 29 (2010) 2833–2836
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Polyhedron
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Kinetics and mechanism of reduction of a coordinated superoxide
with hydroxylamine derivatives
a,
b
Kaustab Mandal a, Subrata Mukhopadhyay a, , Rupendranath Banerjee
, Aloke Ghosh
*
**
a Department of Chemistry, Jadavpur University, Kolkata 700 032, India
b Department of Chemistry, Burdwan University, Burdwan 713 104, India
a r t i c l e i n f o
a b s t r a c t
Article history:
Rates for the uncatalyzed reactions of the coordinated superoxide in
l
-superoxo bis[pentamineco-
balt(III)]5+ complex (1) with hydroxylamine-N-monosulfonate, HONH(SO3)ꢀ (HMS) and hydroxyl-
Received 14 April 2010
Accepted 2 July 2010
Available online 6 August 2010
2ꢀ
amine-N,N0-disulfonate, HONðSO3Þ2 (HDS) have been determined. Successive replacement of the –NH2
protons with SO3H group (electron withdrawing) increased the reaction rate from HMS to HDS but
replacement of the O–H hydrogen halted the reaction. HMS and HDS are oxidized to a mixture of N2O
(20%), NO (80%) and SO42ꢀ. The reactions are first-order in [1] and [reductant]. Reaction rate increased
with pH though neither HMS nor HDS are involved in protic equilibria within the investigated pH range.
We propose 1 is in protic equilibrium with its conjugate acid 1H which is kinetically a dead-end species.
An electroprotic reaction between 1 and the –OH hydrogen in HMS is proposed whereas a simple elec-
tron-transfer mechanism has been proposed for the HDS oxidation.
Keywords:
Hydroxylamine-N-monosulfonate
Hydroxylamine-N,N0-disulfonate
Superoxide
Cobalt(III)
Kinetics
Ó 2010 Elsevier Ltd. All rights reserved.
1. Introduction
2. Experimental
The superoxide anion, O2ꢀ, is known to be involved in many
biological and catalytic oxidations [1–5] like the superoxide radical
anion has been implicated as a by-product of the functioning of
aerobic organisms [6]. Free superoxide ions are known to be extre-
mely reactive [7,8]. However, their reactivity can be controlled
when bonded to a metal center as in the selected superoxo com-
2.1. Materials and reagents
The pentachloride salt of the superoxo complex 1 was synthe-
sized following a literature process [11]. The chloride salt was con-
verted to the corresponding perchlorate salt [12] and recrystallized
from 10% HClO4 (ꢀ
, Mꢀ1 cmꢀ1 at 670 nm = 834: literature value: 838
plex
l
-superoxo bis[pentaminecobalt(III)]5+ ion (1) [9]. Coordina-
2-mode blocks several of its
[13]). HDS [14,15], HMS [16] and HTS [17] were prepared by re-
ported procedures. Their purity were checked by H, N microanaly-
ses (Anal. Calc. in % for HON (SO3K)2ꢁ2H2O: H, 1.63; N, 4.50. Found:
H, 1.83; N, 4.37%. Anal. Calc. for HO–NH(SO3K)ꢁH2O: H, 2.36; N, 8.28.
Found: H, 2.29; N, 8.37%. Anal. Calc. for K3[(SO3)2NOSO3]ꢁ2H2O: H,
0.95; N, 3.31. Found: H, 0.98; N, 3.40%.). All other reagents including
the hydroxylamine-O-sulfonic acid (Aldrich) and the O-methyl
hydroxylamine (Alfa-Aesar) were of sufficient purity and used as
received. Doubly distilled water was used all through.
tion of the superoxo ligand in the
l
side reactions known in the free state [10] and thus helps in unam-
biguous mechanistic interpretations on its reactions. We here re-
port the kinetics and mechanism of the redox reactions of 1 with
the hydroxylamine-N-monosulfonate (HMS) and hydroxylamine-
N,N0-disulfonate (HDS). The kinetics with hydroxylamine itself is
too complicated to analyze. Reactivities of O-methyl hydroxyl-
amine, CH3–ONH2 (O-MeHA), hydroxylamine trisulfonate (HTS)
and hydroxylamine-O-sulfonic acid, HSO3–ONH2 have been also
presented for a comparison.
The present work has been carried in the range pH 3.5–5.5 in
acetate buffer. In this range, the complex suffers practically no
self-decomposition within the reaction period and the –SO3H
group in the reductants, present in the fully deprotonated form,
are not involved in any other parallel protic equilibria.
2.2. Stoichiometry and reaction products
The reaction stoichiometry with excess reducing agent was
determined iodometrically [18] quantifying the unused reducing
agent. After completion of the reactions, argon gas was passed
through the reaction mixtures to fully purge out any gas that might
have generated. NO was qualitatively tested in the ensuing gas by
passing it through a fresh FeSO4 solution acidified with sulfuric
acid. Presence of sulfate ion in the product mixture was qualita-
tively detected with the help of BaCl2 solution in the presence of
HCl.
* Corresponding author.
** Corresponding author. Tel.: +91 9432095474.
E-mail addresses: smukhopadhyay@chemistry.jdvu.ac.in (S. Mukhopadhyay),
rupenju@yahoo.com (R. Banerjee).
0277-5387/$ - see front matter Ó 2010 Elsevier Ltd. All rights reserved.
doi:10.1016/j.poly.2010.07.006

2834
K. Mandal et al. / Polyhedron 29 (2010) 2833–2836
2.3. Kinetic measurements
Kinetics was monitored spectrophotometrically at 670 nm,
where only 1 has an absorption maximum. All kinetic runs were
carried out in the presence of a large excess of the reducing agent
concentration over [1] at (25 0.1) °C in an electrically controlled
thermostated cell-housing of the instrument (Jasco V-1700). The
ionic strength (I) was maintained at 1.0 M (NaCl), unless men-
tioned otherwise. All kinetics was studied in the range pH 3.5–
5.5 maintained with acetate buffer (TOAc = 0.10 M). All pHs were
determined with a pH meter (Toshniwal CL-54, India); its elec-
trodes were calibrated using standard buffers as usual. All the reac-
tions were carried out under an inert argon atmosphere. The first-
order rate-constants ko were evaluated from the absorbance-time
data by standard non-linear least-squares programs.
0.12
0.08
0.04
0.00
0
50
100
150
t, s
3. Results and discussion
Fig. 1. First-order fit for the oxidation of
[1] = 0.20 mM, pH 4.5, I = 1.0 M (NaCl), TOAc = 0.10 M, T = 25.0 ( 0.1) °C.
1 with HDS. [HDS] = 10 mM,
3.1. Stoichiometry
Each mole of HMS and HDS consumed 2.8 mol of the complex
(Table 1). A brown color was seen when the product gas was
passed through acidified FeSO4 solution. Also, the reaction prod-
ucts formed a heavy, white precipitate with BaCl2 solution acidified
with HCl. The stoichiometric results may therefore be interpreted
as part oxidation (20%) to N2O and part oxidation to (80%) to NO.
Table 2
Some representative rate-constants for the reaction of HDS with 1. [1] = 0.2 mM,
I = 1.0 M (NaCl), TOAc = 0.10 M, T = 25.0 ( 0.1) °C.
103 [HDS] (M)
103 ko (sꢀ1) at pH
3.4 3.7
2.8 3.1
4.5
7.8
15.0
28.5
41.0
50.0
5.0
30.0
53.0
87.6.
159.0
177.0
5.5
95.0
2.0
5.0
10.0
15.0
20.0
2HONðSO3Þ2ꢀ þ 3H2O ! N2O þ 4SO2ꢀ þ Hþ þ 4e ð20%Þ
ð1Þ
6.6
13.9
20.0
25.0
11.0
14.6
21.0
27.0
169.0
2
4
a
a
a
HONðSO3Þ2ꢀ þ 2H2O ! NO þ 2SO2ꢀ þ 5Hþ þ 3e ð80%Þ
ð2Þ
2
4
a
The reduced form of 1 has been severally suggested to be 2 [13],
the hydroperoxo derivative of 1 and rapidly decomposes [19] to
The reactions under the condition are very fast and could not be followed by
conventional spectrophotometer.
CoII, NH4 ion and O2.
+
3.2. Reaction with hydroxylamine disulfonate (HDS)
The reaction obeyed first-order kinetics with respect to [1]
(Fig. 1), [HDS] (Table 2) and basicity (Fig. 2). None of the added me-
tal ions, Cu2+, Ni2+, Fe2+ and Ag+ exhibited any catalytic effect.
Above pH 5.7, the reaction velocity becomes too high to measure
with conventional spectrophotometry. On the other side HDS rap-
idly hydrolyzes [20] to HMS in solution at pH <3. So the kinetic
determinations were limited within the range pH 3.5–5.5. The
first-order rate-constants (ko) decreases with increasing ionic
strength (Table 3) as expected for reactions between two oppo-
sitely charged species.
0.16
0.12
0.08
0.04
3.3. Reaction with hydroxylamine monosulfonate (HMS)
0.00
0.00
Cu2+ dramatically catalyzed the reaction such that even impu-
rity level of Cu2+ (2.6 ꢂ 10ꢀ6 M, determined with Atomic Absorp-
tion Spectrometry) interferes with the kinetics (Fig. 3a) and
addition of 0.10 mM Cu2+ completed the reaction instantly.
0.10
0.20
0.30
10-5/ [H+]
Fig. 2. Inverse acid dependence of ko for the oxidation of HDS. [HDS] = 5.0 mM,
[1] = 0.20 mM, I = 1 .0 M (NaCl), TOAc = 0.10 M, T = 25.0 ( 0.1) °C.
Table 1
Stoichiometric results for the oxidation HMS and HDS with 1. TOAc = 0.10 M, I = 1.0 M
(NaCl), T = 25.0 ( 0.1) °C.
Table 3
105 [Reductant]
(M)
105 [1] (M) 105
D
[Reductant]
[1]/
The first-order rate-constants for the oxidation of HDS and HMS with 1 at different
media ionic strength. [1] = 0.20 mM, pH 4.5, TOAc = 0.10 M, T = 25.0 ( 0.1) °C.
(M)
D
[Reductant]
2.8a
2.8b
2.8a
2.8b
Reductant
103 ko (sꢀ1) at I (M)
[HMS], 23.8
[HMS], 13.3
[HDS], 17.4
[HDS], 13.3
8.0
6.65
5.8
2.84
2.4
2.1
2.35
1.0
1.5
2.0
2.5
HDS,10 mM
285.0
11.0
250.0
4.0
130.0
2.0
78.0
1.2
6.65
HMS, 10 mMa
a
pH 4.5.
pH 4.0.
a
b
[dipic] = 0.4 mM.

K. Mandal et al. / Polyhedron 29 (2010) 2833–2836
2835
5.5x10-2
5.0x10-2
4.5x10-2
4.0x10-2
3.5x10-2
3.0x10-2
2.5x10-2
2.0x10-2
1.5x10-2
1.0x10-2
5.0x10-3
0.16
0.12
0.08
0.04
0.00
HDS
HMS
3b
3a
0.0
0
100
200
300
time, s
400
500
600
0
2
4
6
8 10 12 14 16 18 20 22
Concentration (mM)
Fig. 3. Effect of a sequestering agent (dipicolinic acid) on the absorbance-time data
for the reaction of HMS. [1] = 0.20 mM, [HMS] = 10 mM, pH 5.0, I = 1.0 M,
TOAc = 0.10 M, T = 25.0 ( 0.1) °C (a). Under the same condition, but in presence of
0.4 mM dipicolinic acid (b) resulted.
Fig. 4. A comparison of first-order rate constant (ko, sꢀ1) for the oxidation of HDS
and HMS. [1] = 0.2 mM, pH 4.5, TOAc = 0.10 M, I = 1.0 M (NaCl), T = 25.0 ( 0.1) °C and
[dipic] = 0.40 mM for HMS.
Table 4
Effect of pH on the oxidation of HDS and HMS with 1.
[1] = 0.20 mM, TOAc = 0.10 M, I = 1.0 M (NaCl), T = 25.0
( 0.1) °C. Additionally, for the oxidation of HMS dipicol-
inic acid (0.40 mM) was added.
1.2x10-3
1.1x10-3
1.0x10-3
9.0x10-4
8.0x10-4
7.0x10-4
6.0x10-4
5.0x10-4
pH
[Reductant] (mM)
104 ko (sꢀ1
)
3.4
3.7
4.0
4.5
5.0
5.5
3.5
4.0
4.5
5.0
5.5
5.0a
5.0a
66
110
130
147
530
1690
1
3
11
47
149
5.0a
5.0a
5.0a
5.0a
10.0b
10.0b
10.0b
10.0b
10.0b
a
HDS.
HMS.
0
20
40
60
80
100
b
% of D2O
Dipicolinic acid (P0.40 mM) (Table S1) completely halted the cat-
alyzed path and only the direct redox reaction with HMS could be
seen (Fig. 3b). Kinetics of the direct reaction grossly followed ki-
netic pattern similar to those followed between 1 and HDS. Thus,
the direct reaction obeyed first-order kinetics with respect to [1],
[HMS] and basicity (Table 4). Increase in ionic strength decreased
the rate. Under comparable conditions, HDS reacts faster than
HMS (Fig. 4).
Fig. 5. Effect of mol% of D2O on the observed first-order rate-constants of the
oxidation of HMS. [1] = 0.2 mM, [HMS] = 10 mM, [dipic] = 0.40 mM, pH 4.5,
TOAc = 0.10 M, I = 1.0 M (NaCl), T = 25.0 ( 0.1) °C.
Thus we conclude that the reactions of hydroxylamine utilize its
–OH end but not the NH2 end.
4. Mechanism
3.4. Reaction with O-substituted derivatives
The reactions show inverse proton dependence but none of the
reducing agents suffer any kind of deprotonation under the reac-
tion condition. Further, presence of the –O–H proton seems essen-
tial and electron withdrawing groups on the N of the –NH2 moiety
accelerate the rate. Based on these kinetic observations, it appears
Substitution at the O-end of hydroxylamine halted the reactions
with [1], for example, the O-methyl hydroxylamine (O-MeHA),
hydroxylamine O-sulfonic acid (HSO3–ONH2) and hydroxylamine
trisulfonate (HTS) did not consume [1].
K
[(NH3)5CoO2HCo(NH3)5]6+
[(NH3)5CoO2Co(NH3)5]5+
H+
+
1
1H
k
ROx + [(NH3)5CoO2HCo(NH3)5]5+
1 + R
2
fast
ROx
Products, N2O and NO
fast
Products, CoII, NH4+ and O2
2
Scheme 1. R = HMS and Rox = one electron oxidized HMS.

2836
K. Mandal et al. / Polyhedron 29 (2010) 2833–2836
1 to be the active oxidant to which the reductants transfer their O–
H proton, and an electron (either simultaneously, viz. for HMS re-
dox or separately, viz. for HDS redox, vide infra), thus converting 1
to its hydroperoxo derivative 2, which undergoes rapid decompo-
sition to Co2+. Increased proton concentration partly scavenges 1
at a protic equilibrium (K) forming 1H and hence decreases the
reaction rate. Scheme 1 matches with the kinetic observations as
stated, for example, for HMS oxidation.
The HMS reaction showed prominent solvent kinetic isotope ef-
fect. A linear relation (Fig. 5) was observed when ko was plotted
against the mol% D2O in the H2O–D2O mixed media suggesting
transfer of a single proton at the rate step [21]. However, HDS
did not show any solvent kinetic isotope effect. On this basis, an
electroprotic mechanism for HMS oxidation but electron-transfer
separated from proton transfer for HDS oxidation is proposed.
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The work has been carried out with the financial assistance
from DST (Project No. SR/S1/IC-40/2007), UGC-CAS to the Depart-
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acknowledged.
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Appendix A. Supplementary data
Supplementary data associated with this article can be found, in
the online version, at doi:10.1016/j.poly.2010.07.006.