Modern Research in Catalysis, 2013, 2, 157-163
http://dx.doi.org/10.4236/mrc.2013.24021 Published Online October 2013 (http://www.scirp.org/journal/mrc)
Kinetic and Mechanistic Study of Oxidation of
Piperazines by Bromamine-T in Acidic Medium
Chandrashekar1,2, B. M. Venkatesha1*, S. Ananda3, Netkal M. Made Gowda4*
1Department of Chemistry, Yuvaraja’s College, University of Mysore, Mysore, India
2Department of Chemistry, PES College of Engineering, Mandya, India
3Department of Studies in Chemistry, Manasagangothri, University of Mysore, Mysore, India
4Department of Chemistry, Western Illinois University, Macomb, USA
Received December 25, 2012; revised April 24, 2013; accepted September 14, 2013
Copyright © 2013 Chandrashekar et al. This is an open access article distributed under the Creative Commons Attribution License,
which permits unrestricted use, distribution, and reproduction in any medium, provided the original work is properly cited.
ABSTRACT
Oxidations of piperazine, 1-methylpiperazine and 1-ethylpiperazine by bromamine-T (BAT) in buffered acidic medium
have been kinetically studied at 303 K. The reaction shows a first-order dependence of the rate each on [BAT]0 and
[piperazine]0, and an inverse fractional-order dependence on [H+]. The additions of halide ions and the reduction prod-
uct of BAT, p-toluenesulfonamide, have no effect on the reaction rate. The variation of ionic strength of the solvent
medium has no influence on the rate. Activation parameters have been evaluated from the Arrhenius and Eyring plots.
A common mechanism consistent with the kinetic data has been proposed for all piperazines. The protonation constants
of substrates have been evaluated. The Hammett linear free-energy relationship has been observed for the reaction with
ρ = −0.5 indicating that the electron-donating groups enhance the reaction rate by stabilizing the transition state. An
isokinetic relationship observed shows β = 368 K indicating the dominance of enthalpy factors on the reaction rate.
Keywords: Piperazines; Oxidation Kinetics; Mechanism; Bromamine-T; Buffer
1. Introduction
Piperazine is a heterocyclic nitrogenous compound [1]
that has chemical similarity with piperidine as it has two
opposing nitrogen atoms in the ring. In animals, includ-
ing man, piperazine and its salts are known to be highly
effective as anthelmintics [2]. It is used in the treatment
of gout and is an excellent solvent for uric acid [2,3].
Many uses of piperazine derivatives have been suggested
[4]. Their more important use is as intermediates for tran-
quilizing agents and antihistamines, insecticides, fungi-
cides, bactericides, analgesics, antispasmodics, filaricides
and anthelmintics. Some piperazines have been investi-
gated for the treatment of cancer [5,6], radiation sickness
[7], and anginapectoris [8]. The literature survey shows
that the kinetic investigations of reactions of piperazines
with iron (II) and cobalt (III) have been reported by
Aravindakshan et al. [9]. Aromatic N-halosulfonamides
are mild oxidants containing a strongly polarized N-
halogen bond where the halogen is in its +1 oxidation
state. The prominent member of this group, chloramine-T
(CAT), is a well-known analytical reagent and the me-
chanistic aspects of many of its reactions have been
documented [10,11]. Its bromine analogue, bromamine-T
(BAT), is a better oxidizing agent than CAT and chlor-
amine-B. However, meager information exists in the lit-
erature on BAT-piperazines reactions [12-15]. Hence, the
oxidation kinetics of piperazines adds much to the
knowledge of redox chemistry. These facts prompted us
to undertake the study of kinetics of oxidation of pipera-
zines by BAT in acidic buffer medium with a view to
elucidate the reaction mechanism.
2. Experimental
The oxidant BAT was prepared and purified using the
method of Nair and Indrasenan [16]. Its purity was
checked by iodometric and spectroscopic data [16,17].
Aqueous solutions of BAT were prepared, standardized
by the iodometric method, and preserved in amber-col-
ored bottles until use, to prevent its photochemical dete-
rioration. Piperazines (Spectrochem Co.) of acceptable
grades of purity were used without further purification.
Fresh aqueous solutions of piperazines were prepared
whenever required. All other chemicals used were of ac-
*Corresponding authors.
C
opyright © 2013 SciRes. MRC

CHANDRASHEKAR ET AL. 158
ceptable grades of purity. A constant ionic strength of the
reaction mixture was maintained at 0.1 mol·dm−3 by
adding a concentrated NaClO4 solution. Triply distilled
water was employed for preparing aqueous solutions. A
pH 4.0 buffer solution of acetic acid and sodium acetate
was prepared [18] and its pH value checked with a pH
meter.
2.1. Kinetic Measurement
The kinetic runs were performed under pseudo-first-or-
der conditions of [piperazine]0 >> [BAT]0 at 303 K. For
each run, requisite amounts of solutions of the piperazine,
NaClO4 and buffer of known pH were taken in stoppered
Pyrex glass tube whose outer surfaces were coated black
to eliminate photochemical effects. A required amount of
pH 4.0 acetate buffer solution was added to maintain a
constant total volume for all runs. The tube was thermo-
stated in a water bath set at a given temperature. To this
solution was added a measured amount of preequilibrated
BAT solution to give a known overall concentration. The
reaction mixture was periodically shaken for uniform
concentration. The progress of the reaction was moni-
tored by withdrawing aliquots of the reaction mixture at
regular time intervals and by iodometrically titrating the
unreacted BAT for over two half-lives. The pseudo-first-
order rate constants k' calculated were reproducible with-
in ±3.0%. The regression analysis of experimental data
was carried out on an Origin 5.0 HP computer to obtain
the regression coefficient, r.
2.2. Stoichiometry
Varying ratios of the oxidant-to-piperazine in pH 4.0
buffer were equilibrated at 303 K for 24 h. The unreacted
BAT in the reaction mixture determined iodometrically
showed that one mole of piperazine reacted with one
mole of BAT to give the corresponding N-oxide, which
is stoichiometrically represented as in Equation (1).
Here R = H for piperazine, CH3 for 1-methylpipera-
zine and C2H5 for 1-ethylpiperazine;
Ar = p-Me-C6H4.
2.3. Product Analysis
The reaction mixture in the stoichiometric ratio in the
presence of buffer medium was allowed to progress for
24 hr at 303 K. After completion of the reaction (moni-
tored by TLC), the reaction mixture was neutralized and
the products were extracted with diethyl ether. The or-
ganic products were subjected to spot tests and TLC
analysis. The N-oxide products corresponding to pipera-
zine oxide, 1-methylpiperazine oxide and 1-ethylpipera-
zine oxide were confirmed by GC-MS analysis. Mass
spectral data for the N-oxide products were obtained on a
17A Shimadzu gas chromatograph with LCMS—2010A
Shimadzu mass spectrometer. For example, the mass
spectrum showed a parent molecular ion peak at 102 amu
(Figure 1) confirming the formation of piperazine N-
oxide in the reaction mixture of piperazine and BAT. The
reaction product, p-toulenesulphonamide (ArSO2NH2),
was detected by paper chromatography [19]. Benzyl al-
cohol saturated with water was used as the solvent with
0.5% vanillin in 1% HCl in ethanol as the developing
reagent (Rf = 0.905).
3. Results and Discussion
The stiochiometry of BAT oxidation of piperazine and its
derivatives was found to be of 1:1 mole ratio. Oxidations
of piperazines by BAT were kinetically investigated,
under pseudo-first order conditions of [piperazine]0 >>
[BAT]0, at several initial concentrations of reactants in
pH 4.0 buffer medium. Under comparable experimental
50 60 70 80 90 100 110 120 130 140 m/z
55 63 92 112134 140
102
100
BG Mode: Peak Start 0.527(32)
72
Figure 1. Mass spectrum of piperazine-N-oxide with its molecular ion peak at 102 amu.
N
+
O
−
N
R
H
49 222
C H RN+ArSO NBrNa+H O22
+ArSONH+ NaBr


 (1)

CHANDRASHEKAR ET AL. 159
conditions, a similar oxidation kinetic behavior was ob-
served for all piperazines.
3.1. Effect of Varying Reactant Concentrations
on the Rate
The reaction performed in pH 4.0 buffer medium gave
linear plots of ln[BAT] vs time (r > 0.9948). The linear-
ity of these plots together with the constancy of the slope
for various [BAT]0 indicated a first-order dependence of
the reaction rate on [BAT]. The pseudo-first-order rate
constants, k', obtained at 303 K are listed in Table 1.
Under the same experimental conditions, an increase in
[piperazine]0 increased the rate. Plots of lnk' vs
ln[piperazine]0 were linear (r > 0.9948) with slopes of
0.992, (piperazine) 0.995, (1-methylpiperazine) and 1.00
(1-ethylpiperazine), which indicated a general first-order
dependence of the rate on the [substrate].
3.2. Effect of Acid on the Rate
The reaction rate increased with increasing pH (Table 2)
and plots of lnk' vs ln[H+] were linear (r > 0.9961) with
negative fractional slopes (−0.67 to −0.75) showing an
inverse fractional order dependence of the rate on [H+]
for each piperazine reaction.
3.3. Effets of Ionic Strength and ArSO2NH2 on
the Rate
The addition of Cl− ions in the form of NaCl at constant
pH and ionic strength, did not affect the rate. Hence the
dependence of the rate on pH reflected the net effect of
[H+]. The variation of ionic strength of the reaction me-
dium effected using NaClO4 (0.10 - 0.50 mol dm−3),
while keeping all other experimental conditions the same,
had no effect on the rate. Furthermore, the addition of
p-toluenesulfonamide or ArSO2NH2 (3.0 × 10−4 - 7.0 ×
10−4 mol·dm−3) had no effect on the rate indicating that it
is not involved in a pre-equilibrium to the rate determin-
ing step.
3.4. Effect of Temperature on the Rate
The reaction was studied at different temperatures in the
range, 298 K to 313 K, while keeping the concentrations
of reactants and other experimental conditions constant.
The rate constants are presented in Table 3. The activa-
tion parameters (Table 3) were calculated from the
slopes and intercepts of Arrhenius and Eyring plots of
logk' vs 1/T and logk'/T vs 1/T (>r − 0.9826), respec-
tively.
3.5. Test for Free Radicals
Addition of the reaction mixtures to aqueous acryl amide
monomer solutions, in the dark, did not initiate polym-
erization/precipitation, indicating the absence of in situ
formation of free radical species in the reaction sequence.
3.6. Mechanism
Bromamine-T (ArSO2NBrNa·3H2O), like its chlorine
analog chloramine-T (CAT), behaves as an electrolyte in
aqueous solutions [20] dissociating to furnish an anion
(Equation (2)). This anion undergoes protonation in acid
medium to form the free acid, ArSO2NHBr, as in Equa-
tion (3). Although the free acid has not been isolated, the
conductometric studies of CAT have provided ample
evidence of its formation [20,21]. In acid medium, the
Table 1. Effect of varying reactant concentrations on the reaction rate pH = 4.0; temperature = 303 K.
k' × 104 (s−1)
104 [BAT] (mol·dm−3) 102 [Piperazine]0 (mol·dm−3)
Piperazine 1-Methylpiperazine 1-Ethylpiperazine
3.00 1.00 2.23 3.96 6.15
4.00 1.00 2.20 3.98 6.25
5.00 1.00 2.18 3.90 6.00
6.00 1.00 2.16 3.97 6.10
7.00 1.00 2.25 3.88 6.20
5.00 5.00 1.20 2.04 3.14
5.00 7.00 1.52 2.63 4.26
5.00 10.00 2.18 3.90 6.00
5.00 12.0 2.57 4.57 7.23
5.00 15.0 3.65 5.83 9.43
5.00 20.0 4.47 8.03 12.7

CHANDRASHEKAR ET AL. 160
Table 2. Effect of varying pH on the reaction rate [pipera-
zine]0 = 1.00 × 10−2 mol·dm−3; [BAT]0 = 5.00 × 10−4
mol·dm−3; temperature = 303 K.
k' × 104 (s−1)
pH 105 [H+]
(mol·dm−3) Piperazine 1-Methylpiperazine 1-Ethylpiperazine
3.6 25.1 1.14 1.90 2.95
3.8 15.9 1.62 2.63 3.99
4.0 10.0 2.18 3.90 6.00
4.2 6.30 3.02 5.88 8.22
4.4 3.98 4.16 8.51 10.96
4.6 2.51 4.90 12.3 14.80
free acid undergoes a reaction to form dibromamine-T or
DBT (ArSO2NBr2) and p-toluenesulfonamide
(ArSO2NH2) (Equation (4)). ArSO2NHBr hydrolyzes to
give HOBr as one of the products [Equation (5)]. In high
acid concentrations, HOBr can get protonated to H2OBr+
as in Equation (6).
22
ArSONBrNaArSONBrNa


22

(2)
22
ArSO NBrHArSO NHBr

 (3)
22
2ArSO NHBrArSO NBr+ArSO NH (4)
22 22
ArSONHBrHOArSONH+ HOBr (5)
2
HOBr +HHOBr

 (6)
In acidic solutions, the probable oxidizing species are
the free acid (ArSO2NHBr), ArSO2NBr2 and HOBr. The
involvement of ArSO2NBr2 in the mechanism leads to a
second-order rate law, which is contrary to the experi-
mental observations, as Equation (4) indicates. If a slow
hydrolysis of ArSO2NHBr occurred as in Equation (5),
leading to HOBr as the primary oxidizing species, a first-
order retardation of the rate by the added ArSO2NH2
would be expected. This is contrary to the experimental
results. Hardy and Johnston [22], who have studied the
pH-dependence of relative concentrations of the species
present in acidified chloramine-T solutions of compara-
ble molarities, have shown that ArSO2NHBr is the likely
oxidizing species in acid medium.
Furthermore, ultraviolet spectral measurements showed
that the aqueous piperazine solutions have a sharp ab-
sorption band at 235 nm, while the BAT solution exhibits
a peak around 287 nm, both in pH 4.0 buffer solutions
and water. A mixture of BAT and piperazine shows a
max around 330 nm which indicates no direct reaction
between BAB and piperazines and no deprotanation from
BAB. However, piperazine in pH 4.0 buffer exhibits a
max at 380 nm, which shows a longer shift indicating the
formation of intermediate S' due to deprotanation of the
substrate. Based on the preceding discussion, Scheme 1
below is proposed for the reaction.
A detailed common mode of oxidation of piperazines
by BAT in acidic buffer medium along with structures of
intermediates is depicted in Scheme 2.
3.7. Rate Law Derivation
From the slow step of Scheme 1,




2
RatekS' BAT (7)
From step (1),
0
1
S' H
SK







 (8)
Also,




0
SSS
t

 '
(9)
Combination of Equations (8) and (9) leads to Equa-
tion (10)

1
0
SS'H
Piperazine
K


i) fast

2
k
22
complex
S'ArSO NHBrXArSO NH  ii) slow

2
222
XH OProducts
ArSO NHHArSO NH



iii) fast
Here Ar = p-Me-C
6
H
4
Scheme 1. General mechanism for the oxidation of pipera-
zines by BAT in pH 4 buffer.
Table 3. Effect of varying temperature and activation parameters for the oxidation of piperazines by BAT in acidic buffer
[piperazine]0 = 1.00 × 10−2 mol·dm−3; [BAT]0 = 5.00 × 10−4 mol·dm−3; pH = 4.0.
k' 104 (s−1) H≠ S≠ G≠ Ea
Substrate
298 303 308 313 (kJ mol−1) (JK−1 mol−1) (kJ mol−1) (kJ mol−1)
Piperazine 1.47 2.18 3.38 7.41 79.0 −55.6 95.4 81.6
1-Methylpiperazine 2.61 3.90 6.45 10.7 71.0 −76.0 94.16 73.5
1-Ethylpiperazine 3.71 6.0 8.77 12.6 60.2 −106.6 92.7 62.7

CHANDRASHEKAR ET AL. 161
Scheme 2. Detailed reaction pathway for the oxidation of
piperazines by BAT in acidic medium.
  


1
11
S' H
S' H
SS'
t
K
KK


  





1
1
S
S' H
t
K
K




(10)
Substitution for S’ in Equation (7) leads to,




21
1
SBAT
rate H
t
KK
K




(11)
The rate law (Equation (11)) obtained from Scheme 1
is in good agreement with the experimental results,
where the rate has a first-order dependence each on
[BAT]0 and [substrate piperazine]0 and an inverse frac-
tional-order on [H+].
Since, rate = [BAT], Equation (11) can be trans-
formed as,
k'


21
1
kS
k' H
t
K
K




Or
 
21 2
H
1
k'kSk S
tt
K




1
(12)
Based on Equation (12), plots of 1/k' vs [H+] at con-
stant [BAT]0, [piperazine]0 and temperature have been
found to be linear (Figure 2, r > 0.9951) for all piperazi-
nes. The deprotonation constants (K1) and protanation
0510 15 20 25
0
1000
2000
3000
4000
5000
6000
7000
8000
9000
Piperazine
Methyl piperazine
Ethyl piperazi ne
104/k'
105[H+]
Figure 2. Reciprocal plots of 1/ vs [H+] Experimental
conditions are as in Table 2.
k'
constants (KP) of the substrate and the reaction constant
(k2) were calculated from the slope and intercept of these
plots for the standard runs with [BAT]0 = 5.00 × 10−4 mol
dm−3, [piperazine]0 = 1.00 × 10−2 mol dm−3, and [H+] =
1.00 × 10−4 mol dm−3 at 303 K. Furthermore, the values
of protonation constant of the substrate (KP = 1/K1) de-
termined are presented in Table 4.
3.8. Structure-Reactivity Correlation
Structural modification of a reactant molecule may in-
fluence the rate or equilibrium constant of a reaction
through inductive, polar, steric and resonance effects,
which can be used to probe into the reaction mechanism.
Out of a number of empherical models proposed in de-
scribing the relationship between structure and reactivity,
the most successful and extensively investigated is the
linear free energy relationship [23] with Hammett equa-
tion as the most prominent example. Hammett treatment
describes the substituent effects on the rate and equilibria
of aromatic molecules. In the present system, structure-
reactivity relationship is ascertained by utilizing different
groups (H, CH3, C2H5) at one of the two nitrogens (N1)
of the piperazine ring and tested to fit results into the
Hammett equation [24]. The Hammett plot of log vs 
is reasonably linear (r = −0.9957). From such a plot, the
value of the reaction constant ρ is found to be −0.52 sig-
nifying that the electron releasing groups in the pipera-
zine ring enhance the rate. The positive inductive effect
of the substituent increases the electron density on nitro-
gen of the piperazine ring system and subsequently the
lone electron pair on nitrogen nucleophilically attacks the
positive bromine end of the reactive oxidizing species,
ArSO2NHBr, to form the N-bromopiperazine transition
state (Scheme 2). In the next fast step, the bromopipera-
zine intermediate undergoes hydrolysis to yield pipera-
zine-N-oxide as the end product. Furthermore, the posi-
k'

CHANDRASHEKAR ET AL. 162
Table 4. Protonation and deprotonation constants for the oxidation of piperazines by BAT in acidic medium at 303 K.
Substrate 10−5 K1 (mol·dm−3) 104 KP (dm3·mol−1) k2 (dm3·mol−1s−1)
Piperazine 4.65 2.14 0.072
1-Methylpiperazine 2.21 4.51 0.226
1-Ethylpiperazine 3.64 2.74 0.226
tive inductive effect of the substituent in the piperazine
ring system increases in the order: H < CH3 < C2H5,
which justifies the observed reactivity trend of piperazine
< 1-methylpiperazine < 1-ethylpiperazine.
3.9. Isokinetic Relationship
The largest activation energy for the slowest reaction
(Table 3) indicates that the reaction is enthalpy con-
trolled, within the reaction series. The variation in the
rate may be caused by changes in either the enthalpy or
entropy of activation or both. In this study, enthalpy and
entropy of activation are correlated by H≠ = H≠
0 +
 S≠, which is called the isokinetic relationship where β
is the isokinetic temperature. When the experimental
temperature T < β, the reaction rate is controlled mainly
by the enthalpy change. In the present case, the piperazi-
nes oxidations are linearly related by plotting H≠ vs S≠
(Figure 3, r = 0.9995). From the slope, the value of
isokinetic temperature (β) is computed to be 368 K. The
determined β value of 368 K being higher than the ex-
perimental temperature of 303 K, suggests that the reac-
tion is enthalpy controlled. The existence of isokinetic
relationship is very valuable to the mechanistic chemist
as this can be used as a supportive evidence for the me-
chanism along with other data. The large negative value
of S≠ indicates a more ordered associative transition
state with less degree of freedom. The near constant G≠
values show an identical common mechanistic pathway
in the oxidation of all the piperazines studied. Further-
-110-100-90 -80 -70 -60 -50
60
65
70
75
80
H
≠
(kJ mol
−1
)
S
≠
(JK
−1
mol
−1
)
Figure 3. Isokinetic plot of H≠ vs S≠.
more, the independent nature of the rate towards the ad-
dition of p-toulenesulfonamide, halide ion and neutral
salts supports the proposed mechanism and the rate law
derived.
4. Conclusion
Kinetics of oxidation of three piperazines using Broma-
mine-T as the oxidant was carried out in acid medium. A
mechanism has been proposed and the rate law has been
derived. The Hammett correlation of the substituent ef-
fect shows a linear free energy relationship with  =
−0.52 indicating that electron-donating centers enhance
the rate of reaction. An isokinetic study indicates that
enthalpy rather than entropy factors controls the reaction
rate.
5. Acknowledgements
Chandrashekar thanks the management of PES College
of Engineering, Mandya, Karnataka for granting permis-
sion to undertake this study and for encouragement.
REFERENCES
[1] W. R. Maynard Jr., “Gravimetric and Infra-Red Spectro-
photometric Determination of Piperazine,” Journal of the
Association of Official Analytical Chemists, Vol. 42, 1961,
pp. 610-612.
[2] G. L. Crenshaw and R. J. Shaver, “Down to Earth,” Dow
Chemical Company, Michigan, 1956.
[3] L. S. Goodman and A. Gilman, “The Pharmacological
Basis of Therapeutics,” 2nd Edition, Macillan Company,
New York, 1956.
[4] R. M. Ioset, “Personal Communication,” Dow Chemical
Company, Midland, Michigan, 1961.
[5] T. J Mc Nair, F. A. Wibin, E. T. Hoppe, J. L. Schmidt and
F. A. de Peyster, “Antitumor Action of Several New
Piperazine Derivatives Compared to Certain Standard
Anticancer Agents,” Journal of Surgical Research, Vol. 3,
No. 3, 1963, pp. 130-136.
http://dx.doi.org/10.1016/S0022-4804(63)80014-1
[6] V. A. Mikhalev, M. I. Dorokhova and N. E. Smolina,
“Prospidine and Some Other Derivatives of Steropoly
Piperazine,” Meditsinskaia promyshlennost’ SSSR, Vol.
17, No. 1, 1963, pp. 17-20.
[7] W. O. Foye and D. H. Kay, “Antiradiation Compounds
III. N-2-Mercaptoethylpiperazines” Journal of Pharma-

CHANDRASHEKAR ET AL. 163
ceutical Sciences, Vol. 51, No. 11, 1962, pp. 1098-1101.
http://dx.doi.org/10.1002/jps.2600511120
[8] A. Jouve, A. Grass and R. Benyamine, “Hemodynamic
Studies during Angina Pectoris,” Vie Med., Vol. 44, 1963,
pp. 115-120.
[9] K. K Aravindakshan and K. Muraleedharan, “Kinetics of
Non-Isothermal Decomposition of Polymeric Complexes
of N,N’-Bis(dithiocarboxy)piperazine with Iron(III) and
Cobalt(III)” Thermochimica Acta, Vol. 159, 1990, pp.
101-107.
http://dx.doi.org/10.1016/0040-6031(90)80098-J
[10] M. M. Campbell and G Johnson, “Chloramine T and Re-
lated N-Halogeno-N-Metallo Reagents,” Chemical Re-
views, Vol. 78, No. 1, 1978, pp. 65-79.
http://dx.doi.org/10.1021/cr60311a005
[11] K. K. Banerji, B. Jayaram and D. S. Mahadevappa,
“Mechanistic Aspects of Oxidations by N-Metallo-N-
Haloarylsulfonamides,” Journal of Scientific & Industrial
Research, Vol. 46, 1987, pp. 65-76.
[12] F. Ruff and A. F. Kucsman, “Oxidation of Dialkylsul-
hides with BAT in Alkaline Buffer Solutions,” Journal of
the Chemical Society, Perkin Transactions II, 1990, p.
1075.
[13] D. S. Mahadevappa and Puttaswamy, “Kinetics of Oxida-
tion of Aliphatic Ketones with Bromamine-T in Acid Me-
dium,” Bulletin of the Chemical Society of Japan, Vol. 61,
No. 2, 1988, pp. 543-547.
http://dx.doi.org/10.1246/bcsj.61.543
[14] K. S. Rangappa, H. Ramachandra, D. S. Mahadevappa
and N. M. Made Gowda, “Osmium (VIII) Catalyzed Ki-
netics and Mechanism of Indoles Oxidation with Aryl-N-
Haloamines in Alkaline Medium,” International Journal
of Chemical Kinetics, Vol. 28, No. 4, 1996, pp. 265-274.
http://dx.doi.org/10.1002/(SICI)1097-4601(1996)28:4<26
5::AID-KIN4>3.0.CO;2-T
[15] Puttaswamy and R. Ramachndrappa, “Ruthenium(III)-
Catalyzed Oxidation of Substituted Ethanols by Sodium
N-Bromo-p-Toluenesulfonamide in Hydrochloric Acid
Medium,” Transition Metal Chemistry, Vol. 24, No. 3,
1999, pp. 326-332.
http://dx.doi.org/10.1023/A:1006931803005
[16] C. G. R Nair and P Indrasenan, “New Redox Titrants in
Non-Aqueous or Partially Aqueous Media—VI1: Poten-
tiometric Determinations Using Dibromamine-T and Some
Further Applications of Dichloramine-T,” Talanta, Vol.
23, No. 3, 1976, pp. 239-241.
http://dx.doi.org/10.1016/0039-9140(76)80178-6
[17] M. S. Ahmed and D. S. Mahadevappa, “Bromamine-B as
a New Oxidimetric Titrant,” Talanta, Vol. 27, No. 8,
1980, pp. 669-670.
http://dx.doi.org/10.1016/0039-9140(80)80207-4
[18] R. M. C. Dawson, D. C. Elliott, W. Elliott and K. M.
Jones, “Data for Biochemical Research,” 3rd Edition,
Oxford Science Publications, 1986.
[19] D. S. Mahadeveppa and N. M. Made Gowda, “Estimation
of Glutathione with Chloramine-T and Dichloramine-T,”
Talanta, Vol. 22, No. 9, 1975, pp. 771-773.
http://dx.doi.org/10.1016/0039-9140(75)80226-8
[20] E. Bishop and V. J. Jennings, “Titrimetric analysis with
chloramine-T—I: The status of chloramine-T as a ti-
trimetric reagent,” Talanta, Vol. 1, No. 3, 1958, pp. 197-
212. http://dx.doi.org/10.1016/0039-9140(58)80034-X
[21] D. S. Mahadevappa and Rangaswamy, “Physico Chemi-
cal Properties of Chloramine-T. II. Conductometric study
of the Interaction of Chloramine-T with Chromium(III),
Aluminium(III) and Iron(III) Solutions,” Revue Roumaine
de Chimie, Vol. 22, 1977, pp. 1233-1241.
[22] F. F Hardy and J. P. J Johnston, “The Interactions of N-
Bromo-N-Sodiobenzenesulphonamide (Bromamine-B)
with p-Nitrophenoxide Ion,” Journal of the Chemical
Society, Perkin Transactions II, No. 6, 1973, pp. 742-746.
[23] R. D Gilliom, “Introduction to Physical Organic Chemis-
try”, Addison Wesley, London, 1970, p. 144.
[24] L. P. Hammett, “The Effect of Structure upon the Reac-
tions of Organic Compounds. Benzene Derivatives,” Jour-
nal of the American Chemical Society, Vol. 59, No. 1,
1937, pp. 96-103. http://dx.doi.org/10.1021/ja01280a022