Kinetics of oxidation of hydroxylamine by [ethylenebis(biguanide)]silver(III) in aqueous media

Article abstract of DOI:10.1016/S0277-5387(02)01101-4

The complex cation [ethylenebis(biguanide)]silver(III), [Ag(H₂L)]³⁺, and its conjugate bases, [Ag(HL)]²⁺ and [AgL]⁺ oxidise NH₂OH quantitatively to N₂O in the pH range 3.00-7.30, themselves

Full text of DOI:10.1016/S0277-5387(02)01101-4

Polyhedron 21 (2002) 1893ꢁ1898
/
www.elsevier.com/locate/poly
Kinetics of oxidation of hydroxylamine by
[ethylenebis(biguanide)]silver(III) in aqueous media
Prabir Bandyopadhyay, Subrata Mukhopadhyay *
Department of Chemistry, Jadavpur University, Calcutta 700 032, India
Received 21 January 2002; accepted 31 May 2002
Abstract
The complex cation [ethylenebis(biguanide)]silver(III), [Ag(H2L)]3ꢀ, and its conjugate bases, [Ag(HL)]2ꢀ and [AgL]ꢀ oxidise
NH2OH quantitatively to N2O in the pH range 3.00ꢁ7.30, themselves being reduced to Ag(I). Free ethylenebis(biguanide) was
recovered in near-quantitative yield. The reactions are first-order in both [complex] and total hydroxylamine concentration. The
reaction proceeds through three parallel paths: [Ag(H2L)]3ꢀ ꢀNH2OH (k1), [Ag(HL)]2ꢀ ꢀNH2OH (k2) and [AgL]ꢀ ꢀ
NH2OH (k3),
where k1, k2 and k3 are 20.1, 12.5 and 4.35 dm3 molꢂ1 sꢂ1, respectively, at 25.0 8C and Iꢃ1.0 mol dmꢂ3 (NaNO3). No silver(I)-
/
/
/
/
/
catalysed path was detected and the reactions appear to be inner-sphere in nature. # 2002 Elsevier Science Ltd. All rights reserved.
Keywords: Silver; Hydroxylamine; Kinetics; Mechanism; Redox; Catalysis
1. Introduction
Hydroxylamine, the presently chosen reductant, is
useful for probing reaction mechanisms of inorganic
complexes, and the mechanistic versatility noted in its
reactions is attractive in the way that it acts both as an
oxidant or a reductant [15] and can co-ordinate through
either the N- or the O- end [16].
The hypervalent silver species [ethylenebis(biguani-
de)]silver(III) (Fig. 1) is one of the earliest known
authentic [1,2] cationic complex of silver, strongly
stabilised [3] by a square-planar acyclic tetraaza ligand.
The in-plane ligand field stabilisation is so extensive that
the complex behaves as a milder oxidant (0.43 V for the
couple [Ag(H2L)]3ꢀ/2ꢀ) than Agꢀ [4]. It is also stable in
both strongly acidic and strongly alkaline media, a
property unknown in other silver(III) complexes, and is
freely soluble in aqueous media below pH approxi-
mately 4.5. Though to a limited extent, it is also soluble
at least up to pH 8, if alkali is added, carefully avoiding
local excesses. The study of reactivity of different
silver(III) species connected through protic equilibria is
thus possible because, in spite of this remarkable
thermodynamic stability, the silver(III) complex oxidises
We have now observed that hydroxylamine can
reduce the title silver(III) complex at a moderately fast
rate and report herein the kinetics of reactions of the
parent complex and two of its conjugate bases in
aqueous nitrate media and to our knowledge this is
the first detailed investigation on the mechanistic path-
ways of electron-transfer between hydroxylamine and a
silver(III) complex. However, it may be noted that
ꢂ
Ag(OH)4 oxidises NH2OH and NH2OCH3 in basic
media in a rather complicated manner [17].
2. Experimental
[5ꢁ14] various organic and inorganic molecules with
/
notable variation in the mechanistic pathways of
electron transfer.
2.1. Materials
The silver(III) complex, abbreviated as [Ag(H2L)]3ꢀ
hereafter, was prepared according to the known proce-
dure [18] with a slight modification in the crystallisation
process [5]. Hydroxylamine nitrate was prepared in
solution by double decomposition of hydroxylamine
* Corresponding author. Tel.: ꢀ
6266
E-mail address: subrataju@vsnl.net (S. Mukhopadhyay).
/
91-33-546-1508; fax: ꢀ91-33-414-
/
0277-5387/02/$ - see front matter # 2002 Elsevier Science Ltd. All rights reserved.
PII: S 0 2 7 7 - 5 3 8 7 ( 0 2 ) 0 1 1 0 1 - 4

1894
P. Bandyopadhyay, S. Mukhopadhyay / Polyhedron 21 (2002) 1893ꢁ1898
/
made to ꢀ
/
2. A solution of Fe2(SO4)3 at pHꢀ2 was
/
then added in each set and after a few minutes excess
2,2?-bipyridyl solution (at pHꢀ2) was added. The
/
resulting red [Fe(bipy)3]2ꢀ, generated by the quantita-
tive reduction of excess Fe(III) by hydroxylamine was
measured [21] at 523 nm (oꢃ )
8500 dm3 molꢂ1 cmꢂ1
after proper dilution.
/
Fig. 1. Structure of [Ag(enbbg)]3ꢀ. Bond lengths are in A. is weakly
˚
3. Results and discussions
ꢂ
bonded ClO4 ion at the axial site.
3.1. Preliminary observations
sulfate (G.R., E. Merck) with barium nitrate (G.R., E.
Merck) and standardised by oxidation with KBrO3 in
presence of dilute H2SO4 [19]. We found this solution to
be stable for at least 6 months when kept at low
Within the time span of the reactions the autodecom-
position [22] of the silver(III) complex is insignificant.
No silver(I) catalysis was found in the reactions below
temperature (ꢀ
/
5ꢁ10 8C) and in dark. Aqueous solu-
/
pHꢀ5.0. Above this pH, addition of AgNO3 in the
/
tions of NaNO3 (G.R., E. Merck) was standardised by
passing through a Dowex 50W X-8 strong cation
exchange resin in the Hꢀ form and titrating the
liberated acid with standard NaOH to a phenolphtha-
lein end-point. All measurements were made at 25 8C
reaction mixture produces a faint blackish suspension
probably due to metallic silver and/or silver oxide and
thus Agꢀ catalysis, if any, was examined below pHꢀ
5.0. No external buffer was used in the kinetic measure-
ments since buffering (by acetate for example) affects
the spectra of the complex; specific interaction of buffer
with the complex was thus indicated. Even in the
absence of an external buffer, the media pH before
and after the reactions did not change by more than 0.05
units.
/
and Iꢃ
1.0 mol dmꢂ3 (NaNO3). All solutions were
prepared in water which was deionised and then double
distilled.
/
2.2. Physical measurements and kinetics
All absorbance versus time data were recorded with a
Shimadzu (1601 PC) spectrophotometer using 1.00 cm
quartz cells. The kinetics was monitored in situ in the
‘kinetic mode’ of the instrument at 420 nm in the
3.2. Stoichiometry and reaction products
Stoichiometric experiments under kinetic conditions
indicate 1:1 reaction between the redox partners (Table
1). The principal conversion may then be represented by
Eq. (1).
electrically controlled thermostated (25.090.1 8C) cell-
/
housing (CPS-240A). In this mode the change in
absorbance is automatically and continuously recorded.
Hydroxylamine solution (adjusted to the desired pH)
was directly injected in to the spectrophotometer cell
containing other components of the reaction mixture
kept at the same pH. The desired concentration of the
complex and reducing agent was achieved after mixing.
The ionic strength (I) was maintained at 1.0 mol dmꢂ3
(NaNO3). The solution pH values were adjusted with
HClO4 or NaOH. Although the measured pH is usually
defined in terms of the activity of hydrogen ions, we
used the hydrogen ion concentration by calibrating the
pH electrode with analytically prepared solutions as
described earlier [20].
2Ag(III)ꢀ2NH2OH
0 2Ag(I)ꢀN2OꢀH2Oꢀ4Hꢀ
Immediately after completion
(absorbanceB
(1)
of
reaction
/
0.01 at 380 nm), all Agꢀ was precipitated
as AgCl and removed by filtration. From the filtrate, the
free ligand, ethylenebis(biguanide), was quantitatively
isolated [23] as the sparingly soluble [ethylenebis(bigua-
nide)]copper(II) sulfate, and the amount of Cu(II) thus
held was estimated iodometrically after decomposing
the Cu(II) complex. The results showed that more than
95% of the free ligand (H2L) was recovered.
The complicated redox chemistry of hydroxylamine in
solution has been studied extensively [15]. It was
observed that hydroxylamine is most commonly oxi-
2.3. Stoichiometric measurements
ꢂ
The stoichiometry was measured under kinetic con-
ditions by estimating excess hydroxylamine spectro-
photometrically. Excess hydroxylamine was mixed
with the complex in the pH range 4.0ꢁ6.0. After the
reaction mixtures turned colourless indicating complete
reduction of Ag(III) to Ag(I), pH of the mixtures were
dised to N2O [24ꢁ
/
26]. Oxidation to N2, NO2ꢂ, NO3
and NO are relatively rare [16,24ꢁ
/
26]. ESR spectro-
scopy has established the formation of the H2NO+
radical in the oxidation of NH2OH with one-electron
oxidants [27]. The ultimate oxidation product varies
depending on the mode of decay (Eq. (3)) of H2NO+
/

P. Bandyopadhyay, S. Mukhopadhyay / Polyhedron 21 (2002) 1893ꢁ
/1898
1895
Table 1
Stoichiometric results for the oxidation of N(ꢂI) by the silver (III) complex
a
[Ag(III)]T (mmol dmꢂ3
0.15
)
[N(ꢂI)]T (mmol dmꢂ3
)
pH [N(ꢂI)]left (mmol dmꢂ3
)
D[N(ꢂI)]T/D[Ag(III)]T
0.30
0.40
0.50
0.80
1.20
1.00
1.50
3.8 0.17
4.0 0.26
4.9 0.34
5.0 0.48
4.6 0.94
4.9 0.52
3.9 1.05
0.87
0.93
1.07
1.07
0.30
0.50
0.87
0.96
0.90
Averageꢃ0.9790.10.
a
Tꢃ25.0 8C; Iꢃ1.0 mol dmꢂ3 (NaNO3).
radical under different conditions [15]. Thus H2NO+
may decay to N2 (Eq. (2)) or may be further oxidised
(Eq. (3)) to HNO, which dimerises to cis- and trans-
hyponitrous acid [28]. The trans form is more stable
than the cis- form, decomposes into N2O [29]. With
strong oxidants such as Mn3ꢀ(aq) or Ag2ꢀ(aq), present
function of pH and total NH2OH (defined as [N(ꢂ
/
I)]T,
[NH3OHꢀ]) and some representa-
tive data are summarised in Table 2. Over the entire pH
which is [NH2OH]ꢀ
/
range studied the reaction shows a clear first-order
dependence on [N(ꢂ
/
I)]T and there is no [N(ꢂI)]T
/
independent term. The following changes in the reaction
conditions had, within the limits of experimental un-
certainties, no influence on the values of k0: a fivefold
ꢂ
in excess, the free radical is further oxidised to NO3
[30].
variation in [Ag(III)] from 0.1 to 0.5 mmol dmꢂ3
,
H2NO+ 0 0:5N2 ꢀH2O
(2)
(3)
(4)
addition of AgNO3 (studied up to 0.01 mol dmꢂ3),
variation of ionic strength maintained by NaNO3 or
H2NO+ 0 HNOꢀHꢀ ꢀeꢂ
2HNO 0 H2N2O2 0 N2OꢀH2O
NaClO4 in the range 0.1ꢁ
1.0 mol dmꢂ3, the occasional
shaking of the spectrophotometer cell, presence or
/
Thermodynamic oxidising strength of the title com-
plex under investigation is moderate [4,11]. Hence N(III)
or N(V) is not an expected oxidation product.
Table 2
First-order rate constants for the oxidation of N(ꢂI) by the Ag(III)
complex
a
3.3. Kinetics
pH
[N(ꢂI)]T (mmol dmꢂ3
)
104 k0 (sꢂ1
)
No immediate spectral change was observed on
mixing hydroxylamine with the Ag(III) complex over
the entire range of experimental pH. However, the
3.00
3.10
3.25
3.45
3.73
4.00
4.35
4.60
4.93
5.21
5.55
5.85
6.05
6.30
6.65
6.92
7.15
7.30
3.10
3.25
4.00
4.00
4.00
4.00
6.65
3.0
0.57
0.71
0.98
1.49
2.65
4.58
9.25
15.4
29.9
51.0
89.0
129
absorbance at l ]380 nm gradually decreases to less
/
than 0.01. Trace metal ions [31] in solution strongly
catalyse the oxidation of hydroxylamine and affects
both rate and the oxidation product. Ambient light
sometime also plays a definite role [32]. We measured
some k0 values in presence of added ethylenebis(bigua-
nide) [1.0 mmol dmꢂ3] and found to be identical with its
absence within experimental uncertainties, and we ob-
served no effect of diffused light in our system. The
log10(absorbance) versus time plots were found to be
linear for more than 90% of reactions and the pseudo
first-order rate constants, k0, defined by Eq. (5) were
obtained from the least-squares slopes of these plots.
152
169
168
159
150
145
6.0
12.0
6.0
1.40
4.05
9.25
18.0
26.8
47.0
82.2
ꢂd[Ag(III)]=dtꢃk0[Ag(III)]
Any initial drop in absorbance for faster reactions can
be computed from the time of mixing of solutions (ꢀ
(5)
12.0
18.0
30.0
1.5
/
4
s) and the observed first-order rate constants. Averages
of the k0 values from at least three runs were taken and
individual runs were reproducible within 3%. Detailed
data of the kinetics of the process were obtained as a
a
Tꢃ25.0 8C, Iꢃ1.0 mol dmꢂ3; [complex]ꢃ0.15 mmol dmꢂ3
.

1896
P. Bandyopadhyay, S. Mukhopadhyay / Polyhedron 21 (2002) 1893ꢁ1898
/
a2 ꢃKa=(Ka ꢀ[Hꢀ])
(14)
absence of dissolved oxygen, and a variation in the
monitoring wavelength in the range 380ꢁ480 nm.
/
Non-linear fitting of left-hand-side of Eq. (12) with
[Hꢀ] yielded k1, k2Ka and k3Ka Ka . Using known
values of Ka and Ka , the rate constants k2 and k3 were
evaluated. All these second-order rate constants are
presented in Table 3 which reproduced all k0 within 5%.
No polymer is formed when 6% (v/v) acrylonitrile is
added to the reaction mixture and whole reaction course
is epr silent. There is thus no evidence for the formation
of detectable amounts of free radicals or other para-
magnetic species during the reaction. However, one can
not conclude from these observations as to whether the
reactions of Ag(III) with hydroxylamine are one or two
electron processes, because the radicals produced from
one-electron oxidation of hydroxylamine may react in a
very fast manner with Ag(III) before they can escape
into the bulk solvent, thus evading detection.
1
1
2
1
2
It may also be noted that in the pH interval 3.00ꢁ
/
4.60, the second dissociation of the Ag(III) complex
(Ka ) can be neglected and Eqs. (9) and (10) are sufficient
to represent the observed first-order rate constants. On
2
the other hand, in the higher pH range 6.30ꢁ7.30 the
/
first dissociation of the Ag(III) complex (Ka ) can be
1
neglected and only Eqs. (10) and (11) may account for
the observed situation. The second-order rate constants
evaluated in this manner agreed nicely with the values
represented in Table 3 within 7%.
A plot of k0 versus pH has a bell shape (Fig. 2). Such a
pH dependence indicates the involvement of two or
more acidꢁ/base equilibria [33]. Considering the values
for the acid dissociation constants of the complex in
aqueous solution [34] (Ka and Ka , which are 1.58ꢄ
/
3.4. Mechanism
2
10ꢂ4 and 5.01ꢄ
(ꢃ1.0ꢄ
stant [35] of NH3OHꢀ, the scheme in Eqs. (3)ꢁ
/
10ꢂ7 mo1l dmꢂ3, respectively) and Ka
/
/
10ꢂ6 mol dmꢂ3), the acid dissociation con-
The observed sequence k1ꢁ
/
k2ꢁk3 indicates that
/
/(8)
protonation of the oxidant increases its kinetic reactivity
whereas the opposite is true for protonation of the
reductant. Majority of the other known redox systems
also follow a similar trend [36,37]. The trend was
observed earlier for the oxidation [12] of N(III) with
the title Ag(III) complex. In the present investigation we
found that the conjugate acid NH3OHꢀ is kinetically
inactive towards the Ag(III) complexes under investiga-
tion and its conjugate bases, whereas HNO2 was active,
provide a reasonable explanation for the kinetic ob-
servations.
Ka1
[Ag(H2L)]3ꢀ X[Ag(HL)2ꢀ ꢀHꢀ
(6)
Ka2
[Ag(HL)]2ꢀ X[AgL]ꢀ ꢀHꢀ
(7)
(8)
Ka
NH3OHꢀXNH2OHꢀHꢀ
k1
[Ag(H2L)]3ꢀ ꢀNH2OH 0 Products
(9)
ꢂ
though to a much lesser extent than NO2 towards the
k2
[Ag(HL)]2ꢀ ꢀNH2OH 0 Products
(10)
(11)
reduction of the Ag(III) complex. We anticipate that an
unfavourable charge interaction between the cationic
silver(III) complexes with NH3OHꢀ makes the situation
a different one. A molecule with a free electron pair
should have a greater affinity for cation than would
another positively charged ion.
k3
[AgL]ꢀ ꢀNH2OH 0 Products
From these, expression (12) can be derived where the
fraction of the complex present as [Ag(H2L)]3ꢀ is given
by Eq. (13) and the fraction of [N(ꢂ
/
I)]T present as
NH2OH is given by Eq. (14).
One may use Marcus cross relation [38], k12ꢃ
/
(k11k22K)1/2 to identify the nature of electron transfer.
The equilibrium constant K of the one-electron transfer
between [Ag(H2L)]3ꢀ and NH2OH may be calculated
from the known redox potentials of [Ag(H2L)]3ꢀ/2ꢀ
k0[Hꢀ]2=([N(ꢂI)]Ta1a2)
ꢃk1[Hꢀ]2 ꢀk2K [Hꢀ]ꢀk3K K
(12)
(13)
a1
a1 a2
a1 ꢃ[Hꢀ]2=([Hꢀ]2 ꢀKa [Hꢀ]ꢀKa Ka
1
1
2
couple (0.43 V) [4,11] and NH2OHꢁ
V) [39]. The self-exchange rate constant for the couple
NH2OHꢁ
NH2OHꢀ is 5ꢄ10ꢂ13 dm3 molꢂ1 sꢂ1 [39]
/
NH2OHꢀ (ꢂ
0.42
/
/
/
and the maximum possible self-exchange rate for
Table 3
Second-order rate constants (in dm3 molꢂ1 sꢂ1) for the oxidation of
a
N(ꢂI) by the Ag(III) complex
Reaction
Rate constant
[Ag(H2L)]3ꢀꢀNH2OH
[Ag(HL)]2ꢀꢀNH2OH
[AgL]ꢀꢀNH2OH
k1ꢃ20.191.0
k2ꢃ12.590.60
k3ꢃ4.3590.50
Fig. 2. Dependence of k0 on pH at 25.0 8C, [Ag(III)]Tꢃ
dmꢂ3, [N(ꢂ 3.0 mmol dmꢂ3, Iꢃ1.0 mol dmꢂ3 (NaNO3). The
I)]Tꢃ
solid line represents the fit of Eq. (12).
/
0.15 mmol
a
[Complex]ꢃ0.15 mmol dmꢂ3
,
Iꢃ1.0 mol dmꢂ3 (NaNO3),
/
/
/
Tꢃ25.0 8C.

P. Bandyopadhyay, S. Mukhopadhyay / Polyhedron 21 (2002) 1893ꢁ
/1898
1897
[Ag(H2L)]3ꢀ/2ꢀ couple is 106.6 [13]. Using these values,
Acknowledgements
the calculated maximum value of k12 becomes ꢀ
/
1.7ꢄ
/
10ꢂ3 dm3 molꢂ1 sꢂ1, much lower than the experimen-
tally observed value (20.1 dm3 molꢂ1 sꢂ1), which
refutes one-electron outer sphere pathway as the rate-
determining step. This is not unexpected in light of the
We are grateful to Professor L.J. Kirschenbaum,
University of Rhode Island, Kingston RI 02881, for
many suggestions. Financial assistance from Depart-
mental Special Assistance (DSA) Programme in Chem-
istry, Jadavpur University from University Grants
Commission (New Delhi) is thankfully acknowledged.
remarkably low NH2OHꢁ
NH2OHꢀ self-exchange rate
/
coupled with a near zero E8 value for the one electron
reaction.
It is well documented that hydroxylamine (NH2OH)
can co-ordinate to higher valent metal ions [40ꢁ42]. The
/
References
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high [43] as established by X-ray photoelectron spectro-
scopic measurements that the Ag(III) centre has a high
degree of positive charge and it is known from structural
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[44ꢁ46] and kinetic [6,7] studies that the complex
/
undergoes axial co-ordination with alcohols, H2O2,
ꢂ
[5] R. Banerjee, K. Das, A. Das, S. Dasgupta, Inorg. Chem. 28 (1989)
585.
SO42ꢂ, NO3 or even ClO4ꢂ. We thus anticipate a
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4. Conclusions
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Trans. (1993) 567.
Reduction of the complex cation [Ag(H2L)]3ꢀ may
proceed via a Agꢀ- catalysed path and/or an uncata-
lysed path depending on the nature of the reductant.
Only the uncatalysed path is detectable in the oxidation
[11] R. Banerjee, Proc. Indian Acad. Sci. (Chem. Sci.) 106 (1994) 655.
[12] S. Mukhopadhyay, R. Banerjee, J. Chem. Soc., Dalton Trans.
(1994) 1349.
[13] S.P. Ghosh, M.C. Ghosh, E.S. Gould, Inorg. Chim. Acta 225
(1994) 83.
of N(ꢂI), N(III) [12], alcohols [6], ascorbic acid [10] and
/
[14] S. Mukhopadhyay, S. Kundu, R. Banerjee, Proc. Indian Acad.
Sci. (Chem. Sci.) 107 (1995) 403.
Fe2ꢀ(aq) [11] by this complex. The catalysed path
ꢀ
involves a Agꢀ ꢁAg0 cycle (for HCO2H [5], N2H5 [8]
/
[15] K. Wieghardt, Advances in Inorganic and Bioinorganic Mechan-
isms, vol. 3, Academic Press, New York, 1984, p. 213.
[16] R.A. Scott, G.P. Haight, Jr., J.N. Cooper, J. Am. Chem. Soc. 96
(1974) 4136.
and H3PO2 [9]) or a Ag(I)ꢁAg(II) cycle (for H2O2 [7]).
/
Predominance of the uncatalysed or catalysed path
depends on a critical balance of thermodynamic and
also on kinetic factors, but grossly speaking when the
[17] L.J. Kirschenbaum, Personal communication.
[18] (a) P. Ray, K. Chakravarty, J. Indian Chem. Soc. 21 (1944) 47;
(b) P. Ray, Inorg. Synth. 6 (1960) 74.
ꢀ
reductant is powerful [47] like HCO2H, N2H5 or
H3PO2 or a good ligand for Ag(III), then catalysed
[19] A.I. Vogel, A Text Book of Quantitative Inorganic Analysis
including Elementary Instrumental Analysis, 3rd ed., The English
Language Book Society and Longmans, Green & Co, London,
1961, p. 391.
path dominates. We observed that below pHꢀ5.0
/
hydroxylamine does not react with Agꢀ over a long
period and we observed no Agꢀ catalysis in our
reactions. Below pH 3.0, hydroxylamine was found to
be kinetically inactive towards the reduction of the title
complex which indicates absence of any possible
[20] S. Banerjee, U. Roy Choudhury, B.C. Ray, R. Banerjee, S.
Mukhopadhyay, Anal. Lett. 34 (2001) 2797.
[21] R.L. Davies, K.W. Dunning, J. Chem. Soc. (1965) 4168.
[22] R. Banerjee, D. Banerjea, Indian J. Chem. 17A (1979) 246.
[23] P. Ray, Chem. Rev. 61 (1961) 313.
[Ag(H2L)]3ꢀ ꢀ
/
NH3OHꢀ path. We found that
[24] F.T. Bonner, N.Y. Wang, Inorg. Chem. 25 (1986) 1858.
[25] (a) R.M. Liu, M.R. McDonaldand, D.W. Margerum, Inorg.
Chem. 34 (1995) 6093;
NH3OHꢀ is kinetically inactive but the conjugate
base, NH2OH, a neutral species reacts with the charged
Ag(III) complexes. An indirect outcome of these types
of reacting species may be the innocent influence of
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