Spectroscopic investigation and oxidation of the anticholinergic drug atropine sulfate monohydrate by hexacyanoferrate(III) in aqueous alkaline media: A mechanistic approach

Article abstract of DOI:10.3906/kim-1307-4

The oxidation of the anticholinergic drug atropine sulfate monohydrate by hexacyanoferrate(III) in aqueous alkaline media was investigated spectrophotometrically by monitoring the decrease in absorbance of hexacyanoferrate(III) (HCF(III)). Oxidation products were identied. The oxidation mechanism was proposed from kinetic studies. The reaction constants involved in the dierent steps of the mechanism were calculated. The eects of added products, ionic strength, and dielectric constant of the reaction were investigated. The polymerization test revealed that oxidation occurred with intervention of free radicals. The activation parameters were evaluated. TUeBITAK.

Full text of DOI:10.3906/kim-1307-4

Turk J Chem
Turkish Journal of Chemistry
(2014) 38: 477 – 487
¨
http://journals.tubitak.gov.tr/chem/
˙
c
⃝ TUBITAK
doi:10.3906/kim-1307-4
Research Article
Spectroscopic investigation and oxidation of the anticholinergic drug atropine
sulfate monohydrate by hexacyanoferrate(III) in aqueous alkaline media: a
mechanistic approach
Manjunath METI, Sharanappa NANDIBEWOOR, Shivamurti CHIMATADAR∗
P. G. Department of Studies in Chemistry, Karnatak University, Pavate Nagar, Dharwad, India
Received: 01.07.2013
•
Accepted: 22.11.2013
•
Published Online: 14.04.2014
•
Printed: 12.05.2014
Abstract: The oxidation of the anticholinergic drug atropine sulfate monohydrate by hexacyanoferrate(III) in aqueous
alkaline media was investigated spectrophotometrically by monitoring the decrease in absorbance of hexacyanoferrate(III)
(HCF(III)). Oxidation products were identified. The oxidation mechanism was proposed from kinetic studies. The
reaction constants involved in the different steps of the mechanism were calculated. The effects of added products,
ionic strength, and dielectric constant of the reaction were investigated. The polymerization test revealed that oxidation
occurred with intervention of free radicals. The activation parameters were evaluated.
Key words: Kinetics of oxidation, mechanism, hexacyanoferrate(III), atropine sulfate
1. Introduction
Hexacyanoferrate(III) has been widely used to oxidize numerous organic and inorganic compounds in alkaline
media.1,2 Many transition and nontransition metal ions in their complex form act as good oxidants in acidic,
basic, or neutral media. However, oxidation capacity depends on their redox potential. It is also known that the
redox potential of the couple is pH dependent. For instance, the redox potential3 of [Fe(CN)6 ]3− /[Fe(CN)6 ]4−
in acid medium is +0.36 V and in basic medium is +0.40 V. This indicates that hexacyanoferrate(III) is a good
oxidant in basic medium. It is a one-equivalent oxidant leading to its reduction to hexacyanoferrate(II), a stable
product.4 Oxidation by HCF(III) ion generally proceeds through an outer sphere electron transfer mechanism,
which depends not only on the nature of the substrate but also on the medium of the reaction.5
Tropane alkaloid (atropine) is extracted from deadly nightshade (Atropa belladonna), jimsonweed (Datura
stramonium), mandrake (Mandragora officinarum), and other plants of the family Solanaceaeare widely used as
parasympatholytic, anticholinergic, and antiemetic drugs.6 Atropine sulfate is (RS)-(1R,3r,5S)-3-tropoyloxytro-
panium sulfate monohydrate (Figure 1).7
Atropine sulfate, the monohydrate of (1R,3r,5S)-3-tropoyloxytropanium sulfate, can be used for suppress-
ing unstriated muscle, controlling glandular excretion, in small doses stimulating the central nervous system,
having the action of mydriasis in ophthalmology, etc.8 However, most alkaloids have special and significant
biological activity, and so caution is called for in establishing a sound method to detect atropine sulfate in a
clinical assay.9 Its degradation by microorganisms has been reported by several groups10,11 and, in cases for
∗Correspondence: schimatadar@gmail.com
477

METI et al./Turk J Chem
which the mechanism has been studied, hydrolysis of the ester linkage to give the 2 separate cyclic components
is the initial step.12
CH
3
N
. H SO . H O
2 4 2
H
C
O
C
O
CH OH
2
2
Figure 1. Chemical structure of Atropine sulfate monohydrate.
Although some work on oxidation of atropine sulfate monohydrate by various oxidants has been carried
out,12 there is a lack of literature on the oxidation of this drug by hexacyanoferrate(III). The objectives of the
present study were to (i) accumulate the kinetic data, (ii) elucidate plausible mechanisms, (iii) design kinetic
rate laws, (iv) ascertain the reactive species, (v) deduce thermodynamic parameters, and (vi) characterize the
products.
2. Results and discussion
2.1. Reaction orders
The order with respect to [ASM] and [alkali] were found from the graph of log kobs versus log(concentration)
plots.
2.2. Effect of [hexacyanoferrate(III)]
The concentration of hexacyanoferrate(III) was varied in the range 0.50 × 10−4 –5.0 × 10−4 mol dm−3 at
constant [ASM], [OH− ], ionic strength, and temperature. The fairly constant kobs value (Table 1) indicates that
order with respect to hexacyanoferrate(III) concentration was unity. This was also confirmed by the linearity
of the plot of log(absorbance) versus time up to 80% completion of the reaction.
2.3. Effect of [atropine sulfate]
The concentration of ASM was varied in the range 5.0 × 10−4 –5.0 × 10−3 mol dm−3 at a constant [HCF(III)],
[OH− ], ionic strength, and temperature. The rate of reaction increased with the increase in [atropine sulfate]
(Table 1). A plot of log kobs versus log [ASM] was linear and was found to be less than unity.
2.4. Effect of [alkali]
The concentration of OH− was varied in the range 0.10–1.0 mol dm−3 at constant [HCF(III)], [ASM], ionic
strength, and temperature. The rate of reaction increased with the increase in [alkali] (Table 1) and the order
was found to be less than unity, i.e. 0.60.
478

METI et al./Turk J Chem
Table 1. Effect of variation in hexacyanoferrate(III), atropine sulfate, and OH− on the oxidation of atropine sulfate by
hexacyanoferrate(III) at 25 ◦ C and I = 1.0 mol dm−3
.
[HCF] × 104
[ASM] × 103
[OH−]
(mol dm−3
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.10
0.30
0.50
0.70
1.0
kobs× 103
(mol dm−3
0.50
1.0
)
(mol dm−3
2.0
2.0
)
)
(s−1
1.09
1.03
1.05
1.07
1.08
0.31
0.65
1.05
1.56
2.08
0.45
0.81
1.05
1.48
1.71
)
2.0
3.0
5.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
0.50
1.0
2.0
3.0
5.0
2.0
2.0
2.0
2.0
2.0
2.0.
2.5. Effect of ionic strength (I) and dielectric constant (D)
The effect of ionic strength was varied by varying KNO3 concentration between 0.60, 0.80, 1.0, 1.20, and 1.40
mol dm−3 . The rate was found to increase with increase in ionic strength. A plot of log kobs versus I 1 was
linear with positive slope. The effect of dielectric constant was studied by varying the t-butyl alcohol-water
(v/v) percentage composition from 0% to 30%. It was found that as the percentage composition of t-butyl
alcohol increased in the reaction medium, the rate of reaction decreased and the plot of log kobs versus 1/D
was linear with negative slope.
/
2
2.6. Effect of initially added products
The initially added products, hexacyanoferrate(II), tropine, and benzaldehyde, did not have any significant
effect on the rate of reaction.
2.7. Test for free radicals (polymerization study)
The reaction mixture was mixed with a known quantity of acrylonitrile monomer and kept for 2 h under inert
atmosphere. On diluting with methanol, a white precipitate of polymer was formed, indicating the intervention
of free radicals in the reaction. The experiment of either Fe(CN)36− or ASM with acrylonitrile alone did not
induce polymerization under conditions similar to those induced with the reaction mixture. Initially added
acrylonitrile decreases the rate, also indicating the free radical intervention in the reaction.13
2.8. Effect of temperature
The rate of reaction was measured at different temperatures, 15, 25, 35, and 45 ◦ C, under varying concentrations
of [ASM] and [alkali], keeping the other conditions constant for the reaction. The rate constants (k) of the slow
step of Scheme 1 were obtained from the intercepts of the plots of 1/kobs versus 1/[ASM] at the 4 different
temperatures. The values are given in Table 2. The energy of activation for the rate determining step was
479

METI et al./Turk J Chem
obtained by the least square method of the plot of log k versus 1/T, and the other activation parameters were
calculated and are given in Table 2.
CH3
N
CH3
N
K1
OH-
+
H2O
+
H
O
C
C
O
C
O
-O
C
CH2OH
CH2OH
CH3
N
I
K2
[Fe(CN)6]3-
+
Complex(C)
O
C
II
-O
C
CH2OH
I
CH3
N
k
[Fe(CN)6]4-
Complex (C)
H2O
+
+
+
+
CO2
slow
.CH
CH2OH
H
OH
II
III
IV
V
fast
[Fe(CN)6]3-
H+
[Fe(CN)6]4-
+
+
H2O
+
+ CH3OH +
.CH
CHO
CH2OH
III
VI
VII
H+
H2O
OH-
+
Scheme 1. Stoichiometry of the oxidation of atropine by hexacyanoferrate(III) in alkaline medium.
480

METI et al./Turk J Chem
Table 2. Activation parameters and thermodynamic quantities for the oxidation of atropine sulfate by hexacyanofer-
rate(III) in alkaline medium with respect to the slow step of Scheme 1.
(A) Effect of temperature and activation parameters
Temperature (K) k× 102 (s−1
)
Parameters
Values
53 ꢀ 3
288
298
308
318
0.27 ꢀ 0.01
0.52 ꢀ 0.02
1.04 ꢀ 0.03
2.15 ꢀ 0.03
Ea (kJ mol−1
)
∆H# (kJ mol−1
∆S# (J K−1 mol−1
∆G# (kJ mol−1
log A
)
50 ꢀ 3
)
–120 ꢀ 10
85 ꢀ 3
)
7.0 ꢀ 0.1
(B) Effect of temperature on first and second equilibrium step of Scheme 1
Temperature (K) K1 (dm3 mol−1
)
K2× 10−2 (dm3 mol−1
4.98 ꢀ 0.20
)
288
298
308
318
0.78 ꢀ 0.03
1.82 ꢀ 0.05
3.44 ꢀ 0.10
5.81 ꢀ 0.25
3.03 ꢀ 0.20
2.15 ꢀ 0.10
0.98 ꢀ 0.06
(C) Thermodynamic quantities with respect to K1 and K2
Thermodynamic quantities Values from K1 Values from K2
∆H (kJ mol−1
∆S (J K−1 mol−1
∆G
(kJ mol−1
)
51 ꢀ 4
–39 ꢀ 3
–85 ꢀ 5
–14 ꢀ 0.8
)
)
174 ꢀ 12
–1.50 ꢀ 0.1
298
The variation in the concentrations of the oxidant, substrate, and alkali, while keeping the others
constant, showed that the reaction is first order in oxidant and less than the unit order in substrate and
alkali concentrations (Table 1). The reaction between atropine sulfate and Fe(CN)36− has a stoichiometry of
1:2. Based on the experimental results, a mechanism can be proposed for which all the observed orders in each
constituent such as [oxidant], [reductant], and [OH− ] may be well accommodated. Oxidation of atropine sulfate
by hexacyanoferrate(III) in alkaline media is a noncomplementary reaction with 2 moles of oxidant reacting with
1 mole of substrate.
In the present study, alkali combines first with atropine sulfate to give the anionic form of atropine(I)
in a prior equilibrium step, which is also supported by the observed fractional order in [OH− ] and [ASM].
The hexacyanoferrate(III) species reacts with the anionic form of atropine sulfate(I) to give a complex C(II),
which decomposes in a slow step to form an intermediate 2-phenyl ethanol free radical species(III) and a final
product of tropine(IV) derived from atropine sulfate and byproduct CO2 (V). This intermediate 2-phenyl ethanol
free radical species(III) further reacts with another mole of hexacyanoferrate(III) in a fast step to form final
products such as benzaldehyde(VI), methyl alcohol(VII), and Fe(CN)46− . All these results may be interpreted
in the detailed mechanistic Scheme 1 as shown below.
Since Scheme 1 is in accordance with the generally well-accepted principle of noncomplementary oxida-
tions taking place in the sequence of one-electron steps, the reaction between the substrate and oxidant would
afford a radical intermediate. A free radical scavenging experiment revealed such a possibility. Spectroscopic
evidence for complex formation between oxidant and substrate was obtained from the UV-Vis spectra of hex-
acyanoferrate(III) (2.0 × 10−4 mol dm−3), ASM (2.0 × 10−3 mol dm−3), [OH− ] (0.5 mol dm−3), and a
mixture of hexacyanoferrate(III), ASM, and alkali. A hypsochromic shift of about 4 nm from 263 to 259 nm in
the spectra of hexacyanoferrate(III) to a mixture of hexacyanoferrate(III) and ASM was observed. The plots of
481

METI et al./Turk J Chem
kobs versus [ASM] and kobs versus [OH− ] were nonlinear, whereas the linearity of the Michalis–Menten plots
proved the complex formation between oxidant and substrate, which also explains less than unit order in [ASM].
Scheme 1 leads to the rate law (1)
−d[Fe(CN)36−
kK1K2[ASM][OH−][Fe(CN)63−
1 + K1[OH−] + K1K2[ASM][OH−]
Rate =
=
(1)
(2)
dt
Rate
[Fe(CN)36−
kK1K2[ASM][OH−]
1 + K1[OH−] + K1K2[ASM][OH−]
kobs
=
=
Eq. (2) can be rearranged to the following form, which is suitable for verification:
1
1
1
1
=
+
+
(3)
kobs
kK1K2[ASM][OH−] kK2[ASM]
k
The increase in the rate with increasing ionic strength is contrary to a reaction between neutral and charged
species of reactants, as presented in Scheme 1. This might be due to the presence of different ions and the
high ionic strength of the reaction medium. The effect of solvent on the reaction rate has been described in
detail in the literature.14 For the limiting case of a zero angle approach between 2 dipoles or an anion–dipole
system, Amis15 has shown that a plot of log kobs versus 1/D gives a straight line, with a negative slope for
a reaction between a negative ion and a dipole or 2 dipoles, and with a positive slope for positive ion and
dipole interaction. In the present study, the plot observed had a negative slope as shown in, which supports the
involvement of negative ions as given in Scheme 1. The thermodynamic quantities for the different equilibrium
steps in Scheme 1 can be evaluated as follows. The [ASM] and [OH− ] (Table 1) were varied at 4 different
temperatures. According to Eq. (3), other conditions being constant, plots of 1/kobs versus 1/[ASM] and
1/kobs versus 1/[OH− ] should be linear and are found to be so (Figures 2A and 2B). From the slopes and
intercepts, the values of K1 , K2 , and k were calculated at different temperatures (Table 2). A van’t Hoff plot
was made for the variation in K1 and K2 with temperature (log K1 versus 1/T and log K2 versus 1/T). The
values of enthalpy of reaction ∆H, entropy of reaction ∆S, and free energy of reaction ∆G were calculated
for the first and second equilibrium steps of Scheme 2. These values are given in Table 2C. A comparison of
the ∆H value (51 kJ mol−1) from K1 of the first step with that of ∆H# (50 kJ mol−1) obtained for the
rate determining step shows that the reaction before the rate determining step is fairly fast as it involves low
activation energy.16 A high negative value of ∆S# (–112 J K−1 mol−1) suggests that the intermediate complex
(C) is more ordered than the reactants.17
In conclusion, the oxidation of atropine sulfate monohydrate by hexacyanoferrate(III) in aqueous alkaline
media was investigated. Based on the experimental observations, a mechanism was proposed via the formation
of an intermediate complex between atropine sulfate and hexacyanoferrate(III). The rate constant of the slow
step and other equilibrium constants involved in the mechanism were evaluated and activation parameters with
respect to the slow step of the reaction were computed. The overall sequence described here is consistent with
all experimental findings, including the product, and spectral, mechanistic and kinetic studies.
482

METI et al./Turk J Chem
Figure 2. (a) Plots of kobs versus [ASM] and 1/kobs versus 1/[ASM] and (b) kobs versus [OH− ] and 1/kobs versus
1/[OH− ] at 4 different temperatures (conditions as in Table 1).
CH3
CH3
N
N
2 [Fe(CN)6]3-
2 OH-
+
+
+
H
C
O
C
O
CH2OH
H
OH
4-
+ CH3OH +
+
2 [Fe(CN)6]
CO2
CHO
Scheme 2. Detailed scheme for the oxidation of atropine sulfate by alkaline hexacyanoferrate(III).
3. Experimental
3.1. Materials and reagents
All materials employed in the present work were of reagent grade. Atropine sulfate was obtained from s.d.
fine-Chem Ltd and K3 Fe(CN)6 were purchased from SISCO CHEM. The stock solution of atropine sulfate was
prepared by dissolving known amounts of samples in distilled water. Solutions of atropine sulfate were always
freshly prepared before use. A stock solution of the oxidant, hexacyanoferrate(III), was prepared by dissolving
K3 Fe(CN)6 in distilled water and the solution was standardized iodometrically.18 Hexacyanoferrate(II) solution
483

METI et al./Turk J Chem
was prepared by dissolving a known amount of K4 Fe(CN)6 (s.d. fine-Chem) in water. The required alkalinity
and ionic strength were maintained with KOH (Fisher Scientific) and KNO3 (Fisher Scientific), respectively,
in the reaction solutions. t-Butyl alcohol (SPECTROCHEM) was used to vary the dielectric constant of the
media.
3.2. Kinetic measurements
Reactions were carried out under pseudo-first-order conditions with known excess of [substrate] over [oxidant]
at constant temperature. The reaction was initiated by mixing required quantities hexacyanoferrate(III) with
the ASM solution, which also contained the required concentrations of KNO3 and KOH. The progress of
the reaction was monitored spectrophotometrically at 420 nm by measuring the decrease in absorbance of
hexacyanoferrate(III) (ε = 1070 ꢀ 10 dm3 mol−1 cm−1).
The spectral changes during the chemical reaction for the standard condition at 25 ◦ C are shown in
Figure 3. It is evident from the figure that the concentration of HCF decreases at 420 nm. It was verified
that there was almost no interference from other species in the reaction mixture at this wavelength (420 nm).
The first order rate constants (kobs) were obtained from the plots of log(absorbance) versus time plots. The
rate constants were reproducible within ꢀ5%. In view of the modest concentrations of alkali used in the
reaction medium, attention was also given to the effect of the surface of the reaction vessels on the kinetics.
Use of polythene/acrylic ware and quartz or polyacrylate cells gave the same results as glass vessels and cells,
indicating that the surfaces play no important role in the rate of reaction.
Figure 3. UV-Vis spectral changes during the oxidation of atropine sulfate by alkaline hexacyanoferrate(III) at 25 ◦ C;
[Fe(CN)36− ] = 2.0 × 10−4 ; [ASM] = 2.0 × 10−3 ; [OH− ] = 0.5; and I = 1.0 mol dm−3 with scanning time interval of
1.0 min.
3.3. Instruments used
(i) For kinetic measurements, a Peltier Accessory (temperature control) attached Varian CARY 50 Bio UV-Vis
spectrophotometer (Varian, Victoria-3170, Australia) was used. (ii) For product analysis, a QP-2010S Shimadzu
484

METI et al./Turk J Chem
gas chromatograph mass spectrometer, Nicolet 5700-FT-IR spectrometer (Thermo, USA), and 300 MHz 1 H
NMR spectrophotometer (Bruker, Switzerland) were used.
3.4. Stoichiometry and product analysis
Different sets of reaction mixtures with different concentrations of reactants were kept for 6 h at 25 ◦ C
under nitrogen atmosphere in a closed vessel. The remaining concentration of hexacyanoferrate(III) was
assayed spectrophotometrically by measuring the absorbance at 420 nm. The results indicated that 2 moles of
hexacyanoferrate(III) reacted with 1 mole of atropine as given in Eq. (4):
The product was extracted with ether and purified. The ethereal layer was subjected to column chro-
matography using hexane and ethyl acetate in 8:2 (v/v) and the fractions were subjected to spectral investiga-
tions. The main reaction products were identified as tropine and benzaldehyde. The products were characterized
by spectral studies as given below.
The IR spectrum of tropine showed a broad band at 3245 cm−1 assigned to –OH stretching. The 1 H
NMR spectral analysis of tropine exhibited a broad singlet for –OH at δ 12.34 ppm (D2 O exchangeable), and
a triplet in the range 3.89–3.95 ppm was observed for the protons present on the carbon carrying OH group. A
multiplet due to 2 methine protons (C–H) adjacent to nitrogen and 8 methylene protons appeared at around
3.54–3.66 ppm. Methyl protons attached to nitrogen appeared as a singlet at δ 3.36 (Figure 4). The GC-MS
Figure 4. 1 H NMR spectra of tropine, the oxidation product of atropine sulfate by hexacyanoferrate(III) in alkaline
medium.
485

METI et al./Turk J Chem
mass spectrum of tropine showed a base peak at 124 amu and a molecular ion peak at 141 amu (Figure 5).
Another product, benzaldehyde, was confirmed by its GC-MS spectrum, which showed a molecular ion peak at
106 amu and was also confirmed by its hydrazone derivative.19
Figure 5. GC-MS spectra of tropine showed molecular ion peak at m/z 141 amu and base peak at m/z 124 amu.
The byproducts were identified as methyl alcohol, which was confirmed by sodium test,14 and CO2 was
qualitatively detected by bubbling nitrogen gas through the acidified reaction mixture and passing the liberated
gas through a tube containing limewater. The reaction products did not undergo further oxidation under the
present kinetic conditions.
Acknowledgments
One of the authors (Manjunath D Meti) thanks INSPIRE (Innovation in Science Pursuit for Inspired Research)
Fellowship (IF110548) given by the Government of India, Ministry of Science and Technology, New Delhi.
References
1. Kelson, E. P.; Phengsy, P. P. Int. J. Chem. Kinet. 2000, 32, 760–770.
2. Vovk, A. I.; Muraveva, I. V.; Kukhar, V. P.; Baklan, V. F.; Russ. J. Gen. Chem. 2000, 70, 1108–1112.
3. Day, M. C.; Selbin, Theoretical Inorganic Chemistry, Reinhold, New York, NY, USA, 1964.
4. Sharanabasamma, K.; Mahantesh, A. A.; Suresh, M. T. The Open Catalysis Journal 2011, 4, 1–8.
5. Anjali, G.; Shivani, S. Transition. Met. Chem. 2010, 35, 549–557.
6. Barar, F. S. K. Essentials of Pharmacotherapeutics, 4th Edition, S. Chand and Co. Ltd., New Delhi, India, 2007,
246.
7. Arvadiya, A. C.; Dahivelker, P. P. Scientific Paper Chemical Industry and Chemical Engineering Quarterly 2013,
19, 333−337.
8. Chinese Pharmacopoeia, Chemical Engineering Press, Beijing, 4th edn. 1985, 2, 541–543 (in Chinese).
9. Yumei, L.; Lihong, L.; Weifeng, L.; Deliang, H.; Lihua, N.; Shouzhuo, Y. The Analyst 1999, 124, 1629–1634.
10. Berends, F.; R¨orsch, A.; Stevens, W. F. Research Institute National Defence, Stockholm, Sweden, 1967, 45–54.
11. R¨orsch, A.; Berends, F. A.; Bartlema, H. C.; Stevens, W. F.; Winsinck, F. Proc. Kon. Ned. Akad. Wetensch. Ser.
C 1971, 74, 132–152.
486

METI et al./Turk J Chem
12. Byadagi, K. S.; Hosahalli, R. V.; Nandibewoor, S. T.; Chimatadar, S. A. Z. Phys. Chem. 2012, 226, 233–249.
13. Kolthoff, I. M.; Meehan, E. J.; Carr, E. M. J. Am. Chem. Soc. 1953, 75, 1439–1441.
14. Moelwyn-Hughes, E. A. Physical Chemistry, 2nd ed.; Pergamon Press: New York, NY, USA, 1961.
15. Amis, E. S. Solvent Effects on Reaction Rates and Mechanisms; Academic Press: New York, NY, USA, 1966; 183.
16. Bilehal, D. C.; Kulkarni, R. M.; Nandibewoor, S. T. Can. J. Chem. 2001, 79, 1926–1933.
17. Weissberger, A.; Lewis, E. S. ed. Investigations of Rates and Mechanism of Reactions in Techniques of Chemistry.
Wiley, New York NY, USA. 1974, 4, 421.
18. Jeffery, G. H.; Bassett, J.; Mendham, J.; Denny, R. C. Vogel’s Textbook of Quantitative Chemical Analysis, ELBS,
5th ed.; Longman: Essex, UK, 1996, 339–345.
19. Furniss, B. S.; Hannaford, A. J.; Smith P. W. G.; Tatchell, A. R. Vogel’s Textbook of Practical Organic Chemistry,
5th ed.; Pearson Education Ltd: Upper Saddle River, NJ, USA, 2004.
487

Copyright of Turkish Journal of Chemistry is the property of Scientific and Technical
Research Council of Turkey and its content may not be copied or emailed to multiple sites or
posted to a listserv without the copyright holder's express written permission. However, users
may print, download, or email articles for individual use.