Kinetics of Oxidation of Azide by [Ethylenebis(biguanide)]silver(III) in Aqueous Acidic Media

Article abstract of DOI:10.1002/ejic.200300300

In aqueous acidic media sodium azide and the orange-yellow title complex react quantitatively generating N₂ gas along with a colourless solution of Ag⁺ ion and the ligand, ethylenebis(biguanide). The parent complex, [Ag(H₂L)]³⁺, and one of its conjugate bases, [Ag(HL)]²⁺, participate in the reaction with N ₃^- as the kinetically reactive reductant. Ag⁺ has no catalytic effect. At 25.0 °C, I = 1.0 M (NaNO₃), rate constants for the reactions [Ag(H₂L)]³⁺ + N ₃^- and that for [Ag(HL)]²⁺ + N₃ ^- are (12.3±0.5) and (3.8±0.6) M^-1· S^-1, respectively. The high rate constants along with the very low equilibrium constant (10^-15 M^-1) calculated for the first reaction strongly suggest inner-sphere redox mechanism. Wiley-VCH Verlag GmbH & Co. KGaA, 69451 Weinheim, Germany, 2003.

Full text of DOI:10.1002/ejic.200300300

FULL PAPER
Kinetics of Oxidation of Azide by [Ethylenebis(biguanide)]silver(III) in Aqueous
Acidic Media
Prabir Bandyopadhyay,[a] Basab Bijayi Dhar,[a] Jhimli Bhattacharyya,[a] and
Subrata Mukhopadhyay*[a]
Keywords: Kinetics / Mechanisms / Redox chemistry / Silver / Azides
−
In aqueous acidic media sodium azide and the orange-yellow
title complex react quantitatively generating N2 gas along
with a colourless solution of Ag+ ion and the ligand, ethyl-
enebis(biguanide). The parent complex, [Ag(H2L)]3+, and
stants for the reactions [Ag(H2L)]3+ + N3 and that for
[Ag(HL)]2+ + N3− are (12.3 0.5) and (3.8 0.6)
M
−1·s−1, respect-
ively. The high rate constants along with the very low equi-
librium constant (10−15
M
−1) calculated for the first reaction
one of its conjugate bases, [Ag(HL)]2+, participate in the reac- strongly suggest inner-sphere redox mechanism.
−
tion with N3 as the kinetically reactive reductant. Ag+ has
( Wiley-VCH Verlag GmbH & Co. KGaA, 69451 Weinheim,
no catalytic effect. At 25.0 °C, I = 1.0
M
(NaNO3), rate con-
Germany, 2003)
Introduction
equilibria.[14] Also, the AgIII-bound ligand, H2L may form
hydrogen bonds with a suitable reducing agent, like
Ϫ
Ethylenebis(biguanide)silver() (Figure 1) represents the
simplest cationic complex[1Ϫ3] of AgIII stable in aqueous
solution over a wide pH range against dechelation and in-
tramolecular redox decomposition.[4] The square-planar
acyclic tetraaza ligand imparts fascinating thermodynamic
stability[5] to the AgIII state and renders the title complex a
milder oxidant[6] than Agϩ. It is a potential inner-sphere
oxidant,[7Ϫ11] which may bind a reducing ligand in its vac-
ant axial coordination sites. But in this strong-field d8 sys-
tem such ligand binding is likely to be weak or at best mod-
erate. Hence, outer-sphere reactions may be energetically
more favourable[12,13] for weakly basic but strongly reducing
ligands with the AgIII complex. In redox reactions the com-
plex may act as either a one-[7] or a two-electron[10] oxidant.
In solution the title complex undergoes ligand-based protic
HC2O4 and lead to a dead-end adduct in the electron-
transfer pathway.[15] The mechanistic versatility of oxi-
dations of various organic and inorganic species by the title
AgIII complex is an outcome of such properties as dis-
cussed above.
Ϫ
Kirschenbaum et al. studied[16] the oxidation of N3 by
Ϫ
Ag(OH)4 in strongly alkaline media. They found that the
Ϫ
N3Ϫ ion replaced a coordinated hydroxide ion in Ag(OH)4
and the azido complex so produced underwent subsequent
Ϫ
redox decay with another N3 ion. Sengupta et al. pro-
Ϫ
posed an analogous mechanism[17] for the oxidation of N3
by [Ag(H2TeO6)2]5Ϫ, again in strongly alkaline media. We
are not aware of any mechanistic study of azide oxidation
by higher valent silver species in acidic media. The excep-
tional stability of the title complex towards redox decompo-
sition in acidic media affords a good chance to examine for
Ϫ
the first time the reactivity of HN3 and N3 towards a cat-
ionic AgIII complex in acidic solutions. Moreover, the li-
gand H2L in [Ag(H2L)]3ϩ is substitution-inert so that AgIII
never comes free in solution to open up extra, complicating
pathways in its kinetics.
Results and Discussion
Preliminary Observations
In the absence of any external buffer, the media pH be-
fore and after the reactions did not change by more than
0.05 units. Autodecomposition of the AgIII complex was in-
significant under the experimental conditions of pH and
temperature. No Agϩ catalysis (studied up to 0.01 ) was
found in the reactions. We did not extend the reaction pH
Figure 1. Structure of [Ag(H2L)]3ϩ; the gray-shaded area represents
Ϫ
a weakly bonded ClO4 ion at the axial site
[a]
Department of Chemistry, Jadavpur University,
Calcutta 700032, India
E-mail: subrataju@vsnl.net
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Kinetics of Oxidation of Azide by [Ethylenebis(biguanide)]silver(III) in Aqueous Acidic Media
FULL PAPER
Ͼ ca. 4.05, as reaction mixtures above pH ഠ 4.2 turn turbid
probably due to formation of insoluble silver azide after
ca. 30% decrease in their initial absorbances, thus seriously
interfering with the kinetic measurements.
Table 1. Some representative first-order rate constants for the oxi-
dation of azide by the AgIII complex
pH
TR [m]
104k0 [sϪ1 [a]
]
CV in%[b]
1.00
1.50
2.00
2.65
2.90
3.35
3.70
4.05
1.01
1.50
1.50
2.00
2.02
2.05
2.05
2.50
2.52
2.90
2.90
2.90
3.05
3.05
3.10
3.10
3.35
3.35
3.55
3.55
3.85
3.85
3.0
0.24 (0.24)
0.76 (0.76)
2.35 (2.34)
10.0 (9.9)
3.0
3.0
2.5
2.6
2.1
2.3
2.3
3.0
2.8
3.0
2.9
2.2
2.6
2.6
2.7
2.0
2.1
3.0
2.1
1.8
1.9
1.9
2.0
2.0
2.8
1.9
2.9
1.8
2.2
1.9
2.8
Stoichiometry and Reaction Products
Stoichiometric measurements yielded ∆[AgIII]T/∆TR
0.52Ϯ0.05. The principal conversion can thus be rep-
resented by Equation (1).
ϭ
16.8 (16.7)
39.2 (38.4)
60.0 (62.8)
84.0[c] (86.8)
2.29 (2.44)
1.56 (1.51)
3.10 (3.02)
9.20 (9.36)
5.00 (4.93)
22.0 (21.1)
25.8 (26.4)
34.2 (36.0)
70.0 (74.9)
8.30 (8.35)
33.0 (33.4)
65.5 (66.8)
44.5 (44.9)
88.0 (89.8)
12.1 (12.4)
24.2 (24.8)
37.3 (38.4)
76.0 (76.8)
25.4 (26.0)
107 (104)
AgIII ϩ 2 N3 Ǟ AgI ϩ 3 N2
Ϫ
(1)
30.0
6.0
12.0
12.0
6.0
24.0
30.0
15.0
30.0
1.5
6.0
12.0
6.0
12.0
1.5
3.0
3.0
6.0
1.5
6.0
1.5
6.0
1.5
The oxidation of azide to dinitrogen was qualitatively
confirmed by GC analysis. Immediately after completion of
the reaction (absorbance Ͻ 0.01 at 380 nm), the solution
pH was raised to ca. 5 by adding NaOH solutions when all
Agϩ was precipitated as AgN3 and removed by filtration.
From the filtrate, the free ligand ethylenebis(biguanide) was
quantitatively isolated[18] as the sparingly soluble [ethylene-
bis(biguanide)]copper() sulfate, and the amount of CuII
thus held was estimated iodometrically after decomposing
the CuII complex. The results showed that more than 95%
of the free ligand (H2L) was recovered.
Kinetics
No immediate spectral changes were observed on mixing
the reactants over the entire experimental range of pH. But
the absorbance of the reaction mixtures at λ Ն 380 nm
36.2 (37.0)
143 (148)
42.5 (43.4)
gradually decreased to less than 0.01 and the log10(absorb- 4.05
ance)/time plots were found to be linear to 90% completion
of the reactions. The first-order rate constants, kobs, defined
by Equation (2) were obtained from least-squares slopes of
these plots.
[a]
[complex] ϭ 0.20 m, T ϭ 25.0 °C, I ϭ 1.0  (NaNO3). Calcu-
[b]
lated kobs values are given in parentheses.
CV was measured
using the relation: CV ϭ sd ϫ 100/x, where sd ϭ standard devi-
ation of the measurements of kobs and x ϭ average kobs. See ref.[19]
,
p. 1123. 104 kobs [sϪ1] values are 91.0, 97.0 and 106 at I ϭ 0.5,
[c]
0.2 and 0.1  respectively.
Ϫd[AgIII]/dt ϭ kobs[AgIII
]
(2)
Averages of the kobs values (Table 1) from at least three
independent runs were taken and the average coefficients of
variation (CV) for these measurements were within 3%.
The reaction was found to be first-order in TR (ϭ [HN3]
ϩ [N3Ϫ]) and there is no TR-independent term. The follow-
ing changes of the reaction conditions had, within the limits
of experimental uncertainties, no influence on the values of
kobs: a fivefold variation in [AgIII] from 0.10 to 0.50 m,
addition of AgNO3 up to 0.01 , addition of ethylenebis(bi-
guanide) up to 1.0 m, occasional shaking of the spectro-
photometer cell, presence or absence of dissolved oxygen
and diffused light, and a variation in the monitoring wave-
length in the range 380Ϫ480 nm.
No polymer formed when 6% (v/v) acrylonitrile was ad-
ded to the reaction mixture and the entire reaction course
was EPR-silent. These experiments therefore indicate that
free radicals or other paramagnetic species like AgII do not
accumulate in the reaction media in concentrations detect-
able by either of these two methods.
The value of kobs sharply decreases with increasing [Hϩ].
Figure 2. Plot of kobs versus [Hϩ]Ϫ1 at TR ϭ 3.0 m; the solid line
A plot of kobs versus [Hϩ]Ϫ1 at a fixed TR (Figure 2) is a represents the fit of Equation (7)
Eur. J. Inorg. Chem. 2003, 4308Ϫ4312
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P. Bandyopadhyay, B. B. Dhar, J. Bhattacharyya, S. Mukhopadhyay
FULL PAPER
good straight line for pH Յ 2.9 ([Hϩ] Ն 1.26·10Ϫ3 ) with
zero intercept but droops at higher [Hϩ]Ϫ1. Such [Hϩ] de-
pendence of kobs indicates that the conjugate base form of
one of the reactants is more reactive while that of the other
Mechanism
The redox potentials for the couple [Ag(H2L)]3ϩ/2ϩ and
that for N3Ϫ/N3 are 0.43 V[27,6] and 1.33 V,[28] respectively.
The self-exchange rate constant[28] for N3Ϫ/N3 is 4·104
Ϫ

Ϫ1·sϪ1 and the upper limit for the self-exchange rate for
is less reactive than the parent acid. We infer that N3 is
much more reactive than HN3 while [Ag(H2L)]3ϩ is more
reactive than its conjugate base, [Ag(HL)]2ϩ. This is ex-
pected from known kinetic behaviour[11,12,19,20] of these
species and the fact[21,22] that protonation generally in-
creases the reactivity of metal complex oxidants whereas
the anionic forms of protolytic reductants are much more
reactive than the uncharged molecular compounds.[23Ϫ25]
The near-zero value for kobs at high [Hϩ] further suggests
that kinetic activity of HN3 is negligible for the present sys-
tem. The scheme in Equations (3) to (6) most simply ac-
commodates the kinetic observations and the known acid/
base equilibria Ka1 and Ka.
[Ag(H2L)]3ϩ/2ϩ has been estimated[27] to be 106.6. With
these values and using the Marcus cross-relation,[29] k12
ϭ
(k11k22K)1/2, the calculated maximum value of the outer-
Ϫ
sphere oxidation rate of N3 by [Ag(H2L)]3ϩ is 10Ϫ4

Ϫ1·sϪ1, considerable less (105 times) than k1 and we there-
fore propose an inner-sphere mechanism for the k1 path.
An inner-sphere mechanism appears very likely for the k2
path, too. Indeed, the very high thermodynamic price for
Ϫ
the oxidation of N3 to N3 essentially requires strong oxi-
dants[30,31] like [Ni(bipy)3]3ϩ or [CoW12O40]5Ϫ for outer-
sphere reactions to occur.
Coordination of the azide ion to the metal
centers[16,17,32Ϫ34] in [AgIII(OH)4]Ϫ, [AgIII(H2TeO6)2]5Ϫ
,
[CuIII(H2TeO6)2]5Ϫ
,
[MnIII(bipy)2(OH)]2ϩ
,
[MnIII(edta)-
3ϩ
(H2O)]Ϫ, or in Mnaq are well documented with the for-
mation constants of azido complexes lying in the range
10Ϫ1 to 102. A similar azide binding is likely for the AgIII
complex under investigation. X-ray photoelectron spectro-
scopic measurements[35] have shown that the silver() cen-
tre in the title complex has a high degree of positive charge
resulting in a high electron affinity of the AgIII ion. Expect-
edly, therefore, structural[36Ϫ38] and kinetic studies[8,9,11]
indicated axial coordination by AgIII with low molecular
weight alcohols, H2O2, SO42Ϫ, NO3Ϫ, or even ClO4Ϫ. The
kinetic steps (5) and (6) are therefore most likely to proceed
The derived rate-law for this scheme is Equation (7).
kobs(Ka1 ϩ [Hϩ])(Ka ϩ [Hϩ])/TR ϭ k1Ka[Hϩ] ϩ k2Ka1Ka
(7)
A plot of the left hand side of Equation (7) versus [Hϩ]
is a good straight line (Figure 3, r ϭ 0.99) with slope ϭ
(8.0Ϯ0.3) ϫ10Ϫ4 sϪ1 and intercept ϭ (3.9Ϯ0.6) ϫ10Ϫ8
·sϪ1. Using known Ka (ϭ 6.5·10Ϫ5 )[26] and Ka1 (ϭ
1.585·10Ϫ4 )[14] values, these lead to k1 ϭ (12.3Ϯ0.5)
Ϫ
via intermediate adduct formations between AgIII and N3
followed by intramolecular redox decomposition to form
the final products. The present results cannot straightaway
eliminate the possibility of proton-ambiguous paths like

Ϫ1·sϪ1 and k2 ϭ (3.8 Ϯ 0.6) Ϫ1·sϪ1. These rate constants
[Ag(H2L)]3ϩ ϩ HN3 ǟ {[Ag(H2L)]3ϩ·N3Ϫ} ϩ Hϩ;
{[Ag(H2L)]3ϩ·N3Ϫ} Ǟ Products, etc.
Ǟ
reproduced all experimental kobs values satisfactorily
(Table 1).
Nevertheless, the observed increase in reaction rate with
decreased ionic strength of the reaction media does not fit
into this alternative reaction scheme, rather it directs the
reaction between oppositely charged ions. Also, it has been
Ϫ
established that N3 is kinetically far more reactive than
HN3 and may exhibit overwhelming kinetic dominance over
[30,39]
HN3
even for redox reactions in 1  acid.
In the lower pH range (pH Յ 2.9), where [Ag(H2L)]3ϩ is
the only active and sole oxidant, the rate law simplifies to
kobs/TR ϭ kKa/[Hϩ]. The observed linearity of the kobs ver-
sus [Hϩ]Ϫ1 plot at higher [Hϩ] demands a maximum value
of the order of 101 Ϫ1 for preequilibrium formation of
any adduct between [Ag(H2L)]3ϩ and N3Ϫ. [Ag(HL)]2ϩ, a
possible weaker Lewis acid than its conjugate acid should
Ϫ
bind N3 even less tightly.
The present kinetic observations cannot unambiguously
settle the question of one- or two-electron transfer. The im-
mediate one-electron oxidation product[28,40Ϫ44] of azide is
N3, as firmly established by pulse radiolysis, flash photoly-
sis or by molecular beam studies. The azidyl radical (N3)
Figure 3. Plot of kobs(Ka1 ϩ [Hϩ])(Ka ϩ [Hϩ])/TR versus [Hϩ]; the
solid line represents the fit of Equation (7)
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Kinetics of Oxidation of Azide by [Ethylenebis(biguanide)]silver(III) in Aqueous Acidic Media
FULL PAPER
may decay to N2 in several ways, like by fragmentation[45] of cis-[Pt(N3)2(PPh3)2] has provided distinct spectral evi-
to N2 and N, the process, although spin-forbidden, is close dence for the intermediate formation[55] of neutral cyclic
to thermoneutral[46] with a rate constant[47,48] around 106 hexazine, N6.
s
Ϫ1. In yet another possibility, N3 can combine with
Ϫ
Ϫ
N3 forming[49Ϫ51] N6 (however, attempts to observe the
Ϫ
radical dimer N6 using pulse radiolysis met with no
Experimental Section
success,[48b] but nanosecond laser flash photolysis[50,51] and
kinetic evidences[16] suggest involvement of N6Ϫ in some in-
Materials: The silver() complex, [Ag(H2L)]3ϩ, was prepared by a
known procedure.[3,56] Sodium azide (G. R., E. Merck) solutions
were standardised by CeIV titration.[57] NaNO3 (G. R., E. Merck)
solutions were standardised by passing an aliquot through a Dowex
50W X-8 strong cation exchange resin in the Hϩ form and titrating
the liberated acid with standard NaOH to a phenolphthalein end-
point. To avoid any possible interaction of azide with oxides of
nitrogen,[58] the stock HNO3 solutions were purged with argon be-
fore use. All measurements were made at 25.0 °C and I ϭ 1.0 
(NaNO3). All solutions were prepared in water which was deion-
ised and then doubly distilled.
Ϫ
stances) and then rapid reaction of N6 with the oxidant
species or a rapid direct reaction of two mol of N3 radical
could produce three mol of N2 (2 N3 Ǟ 3 N2, 2k ϭ 9·109

Ϫ1·sϪ1).[28] The absence of EPR evidence for AgII or N3
radicals in our study cannot conclusively eliminate the for-
mation of such species as it may only indicate that such
species are formed but decayed too rapidly before they can
diffuse into the bulk solvent sufficiently quickly to permit
detection. For example, Equations (8) to (12) are one con-
ceivable expansion for the k1 path.
Physical Measurements and Kinetics: UV/Vis spectra were recorded
with a Shimadzu (1601 PC) spectrophotometer using 1.00-cm
quartz cells. Kinetics were generally monitored at 380 nm in situ in
the thermostatted (25.0Ϯ0.1 °C) cell-housing (CPS, 240A). For fas-
ter reactions sodium azide solutions (adjusted to the desired pH)
were directly injected into the spectrophotometer cell containing
other components of the reaction mixture at the same pH. The
desired concentrations of the complex and the reducing agent were
achieved after mixing which was immediately followed by auto-
matic monitoring of the change in absorbance with time. We used
capped quartz cells with minimum vacant space left and thus avo-
ided escape of volatile HN3. The solution pH (1.00Ϫ4.05) was
measured with a Systronics (India) pH meter (model 335) with elec-
trodes calibrated as described earlier.[59] All such pH measurements
were done quickly to ensure minimum evaporation of volatile HN3.
Spectrophotometric determination[60] of azide in blank experiments
indicated a near quantitative (Ն 97%) recovery of total azide under
our experimental conditions. Excess reducing agent, TR (ϭ [HN3]
ϩ [N3Ϫ]), 1.5Ϫ30.0 m over the complex (0.20 m) was used in all
kinetic runs.
Reaction (9) requires a very high thermodynamic price
and is as such unfavourable. Yet, the very high protonation
constants[52] of the ligand H2L drag the overall reaction to
completion in acidic media.
An alternative for the assigned fate of the azidyl radical
could be the dimerisation to produce N6 that ultimately
produces N2. But, this appears unlikely here as, if it were
the actual pathway, then sufficient accumulation of N3 to
cause appreciable reversal of reaction (9) and hence rate
retardation (deviation from simple first-order kinetics)
could be anticipated.[48b,53] No such retardation could be
detected and thus the dimerisation pathway is neglected in
the present study.
Stoichiometry and Reaction Products: The reaction stoichiometry
was determined under kinetic conditions in the pH range 3.5Ϫ4.0
by estimating unchanged azide spectrophotometrically. Excess az-
ide was mixed with the complex solutions and stored until the reac-
tion mixtures turned colourless indicating complete reduction of
AgIII to AgI. Fe(NO3)3 solution in 0.5  HNO3 was then added to
the product mixtures[60] to achieve [FeIII]/TR in the final solutions
within the range 1.5Ϫ2.5. The pH of the solutions was then ad-
justed[60] to 1.5Ϫ2.0 using HNO3 or NaOH, as needed. A red com-
plex, [Fe(H2O)5N3]2ϩ, generated within minutes. Its absorbance at
460 nm yielded the concentration of unchanged azide from a cali-
bration curve constructed under comparable conditions (ε ϭ 3570
Though the silver() complex under investigation is far
more stable than [Ag(OH)4]Ϫ or [Ag(TeO4)2(OH)4]5Ϫ, it
Ϫ
oxidises N3 at a faster rate. This can be attributed to a
difference in mechanism operating in these oxidations, for
example, two-electron transfer in [Ag(OH)3N3]Ϫ or in
[Ag(TeO4)2(OH)3(N3)]5Ϫ and the proposed one-electron
transfer in our inner-sphere adducts. A reactive species am-
biguity (HN3 versus N3Ϫ) may also be a reason for ob-
taining a higher rate in the present investigation. However,
of these two alternatives, the first appears more logical but
the proton-ambiguity situation is unresolvable. The tetraaza
ligand is expected to stabilise AgII much more than that by
hydroxide or tellurate.

Ϫ1·cmϪ1, literature value,[60] 3680 Ϫ1·cmϪ1). We verified that
Agϩ or the ligand, ethylenebis(biguanide), do not interfere with
the measurement of the azide concentration. The gaseous product
formed during the reactions was collected over a saturated sodium
chloride solution as described earlier[61] and analysed by gas chro-
matography using a Chemito (India) GCHT 8610 equipped with
TCD, a Porapack Q and molecular sieves 13 X column. Splitless
mode injection of 500 µL of sample was used. The oven tempera-
ture, injector and detector temperature were kept at 40, 55, and
100 °C, respectively. The carrier hydrogen gas flow was adjusted to
Direct two-electron changes in these redox steps to form
ϩ
Ϫ
N3 which reacts with N3 forming N6 and ultimately de-
cays to N2 cannot be ruled out in principle. Formation of
ϩ
N3 has been proposed for the thermal decomposition[54]
of [Mn(N3)2]2ϩ and photoinduced reductive decomposition 30 mL/min.
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