Synthesis, electrochemical, spectrophotometric and potentiometric studies of two azo-compounds derived from 4-amino-2-methylquinoline in ethanolic-aqueous buffered solutions

Article abstract of DOI:10.1590/S0103-50532012005000013

Two azo-compounds, 2-methyl-4-(5-amino-2-hydroxy-phenylazo)-quinoline (2) and 2-methyl- 4-(2-hydroxy-5-nitrophenylazo)-quinoline, derived from 4-amino-2-methylquinoline were synthesized. Their chemical structures were characterized and confirmed by means

Full text of DOI:10.1590/S0103-50532012005000013

J. Braz. Chem. Soc., Vol. 23, No. 8, 1523-1535, 2012.
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Article
A
Synthesis, Electrochemical, Spectrophotometric and Potentiometric Studies of
Two Azo-Compounds Derived from 4-Amino-2-Methylquinoline in
Ethanolic-Aqueous Buffered Solutions
Mona A. El-Attar,*,a Iqbal M. Ismailb and Mohamed M. Ghoneima
aAnalytical Chemistry Research Unit, Chemistry Department, Faculty of Science,
Tanta University, 31527 Tanta, Egypt
bChemistry Department, Faculty of Science, King Abdul Aziz University,
21589 Jeddah, Saudi Arabia
Dois compostos-azo, 2-metil-4-(5-amino-2-hidróxi-fenilazo)-quinolina e 2-metil-4-(5-amino-
2-hidróxi-nitrofenilazo)-quinolina, derivados da 4-amino-2-metilquinolina foram sintetizados. Suas
estruturas químicas foram caracterizadas e confirmadas através de análise elementar, espectroscopia
no infravermelho (IR), ressonância magnética nuclear (RMN) de 1H e espectrometria de massas
(MS). O comportamento eletroquímico do composto de partida (4-amino-2-metilquinolina) e dos
dois azo-derivativos sintetizados foi estudado com um eletrodo de mercúrio em solução tampão
universal B-R em diferentes valores de pH (2-11,5) contendo etanol 40% (v/v) usando polarografia
dc, voltametria cíclica e coulometria com potencial controlado. Os caminhos de reação dos
compostos no eletrodo foram elucidados e são discutidos. As constantes de dissociação (pKa) dos
compostos examinados, constantes de estabilidade e estequiometria dos complexos formados em
soluções dos compostos com alguns íons de metais de transição (Co(II), Ni(II), Cu(II), La(III) and
UO22+) foram determinadas.
Two azo-compounds, 2-methyl-4-(5-amino-2-hydroxy-phenylazo)-quinoline (2) and 2-methyl-
4-(2-hydroxy-5-nitrophenylazo)-quinoline, derived from 4-amino-2-methylquinoline were
synthesized. Their chemical structures were characterized and confirmed by means of elemental
1
chemical analysis, infrared (IR) spectroscopy, H nuclear magnetic resonance (NMR) and
mass spectrometry (MS). The electrochemical behavior of the starting compound (4-amino-
2-methylquinoline) and of the two synthesized azo-derivatives was studied at the mercury
electrode in the B-R universal buffer at various pH values (2-11.5) containing 40% (v/v) ethanol
using dc-polarography, cyclic voltammetry and controlled-potential coulometry. Their electrode
reaction pathways were elucidated and discussed. The dissociation constants (pKa) of the examined
compounds, stability constants and stoichiometry of their complexes in solution with some transition
metal ions (Co(II), Ni(II), Cu(II), La(III) and UO22+) were determined.
Keywords: 4-amino-2-methylquinoline, azo-derivatives, voltammetry, spectrophotometry,
potentiometry
Introduction
electrode reaction pathways of numerous azo-dyes in various
supporting electrolytes were studied and discussed.18-25
Although the electrochemical behavior of some Schiff base
compounds derived from 4-amino-2-methyquinoline has
been reported,26 no studies concerning the electrochemical
behavior of the latter one or its azo-derivatives are reported
in the literature to date.
In the present work, two azo-derivatives from 4-amino-
2-methylquinoline were synthesized and characterized.
Their electrochemical behavior was investigated at mercury
electrodes. Besides, the dissociation constants (pKa)
Synthetic azo-dyes are among the most explored classes
of organic compounds. Azo-dyes are widely used in many
practical applications such as photochromic materials,
colorants, non-linear optics, sensors and indicators.1-12 The
synthesis and spectral properties of several azo-dyes as well
as of their transition metal complexes have been reported
in the literature.13-17 The electrochemical behavior and
*e-mail: maema.2011@yahoo.com

1524
Synthesis, Electrochemical, Spectrophotometric and Potentiometric Studies of Two Azo-Compounds
J. Braz. Chem. Soc.
of the investigated compounds, stability constant and
stoichiometry of their metal complexes in solution with
some transition metal ions were determined.
IR (KBr) nmax/cm-1 3403 (OH), 1650 (N=N), 2975 (CH3);
1H NMR (DMSO-d6, 298 K, TMS) d/ppm 2.41 (s, 3H,
CH3); 4.23 (s, 2H, NH2), 5.11 (s, OH), 7.14 (s, pyridyl H),
6.55-7.32 (m, 3H subs. phenyl), 7.42-8.12 (m, 4H fused
phenyl). The diazo-coupling reaction27 for synthesis of
compound (2) is illustrated in Scheme 1.
Experimental
Synthesis of 2-methyl-4-(5-amino-2-hydroxyl-phenylazo)-
quinoline (2)
Synthesis of 2-methyl-4-(2-hydroxy-5-nitrophenylazo)-
quinoline (3)
According to the general procedure for synthesis of
azo-compounds reported in the literature,27 0.01 mol
of the starting compound, 4-amino-2-methylquinoline
(1) (Aldrich), was converted to 4-(diazonium salt)-
2-methylquinoline by dissolving in a mixture consisting
of 4 mL water and 8 mL of 9.8 mol L-1 HCl. The
reaction mixture was stirred until a clear solution was
obtained; cooled over crushed ice at 0-5 ºC, then 5 mL
water containing 0.01 mol of sodium nitrite were added
portion wise under gentle stirring for 30 min to complete
the reaction.27 The resulting diazonium salt was slowly
added under moderate stirring to a pre-cooled solution of
p-aminophenol (Aldrich) (0.01 mol) dissolved in 30 mL of
10% (v/v) NaOH aqueous solution and stirred for 30 min,
then allowed to reach room temperature. The final product
was obtained as a colored azo-dye (2-methyl-4-(5-amino-
2-hydroxy-phenylazo)-quinoline), which was separated,
filtered and several times washed with bidistilled water,
dried under vacuum and recrystallized from ethanol to give
2-Methyl-4-(2-hydroxy-5-nitrophenylazo)-quinoline
(3) was synthesized in a similar fashion as compound (2) but
using p-nitrophenol (Aldrich) instead of p-aminophenol to
give yellow crystals.Yield 64%; mp 112 oC; anal. calc. for
C16H12N4O3 (308.09 g mol-1) C, 62.33; H, 3.92; N, 18.17;
found C, 62.85; H, 4.15; N, 18.64; IR (KBr) nmax/cm-1
3327 (OH), 1662 (N=N), 2948 (aliphatic H), 1390 (NO2);
1H NMR (DMSO-d6, 298 K, TMS) d/ppm 2.52 (s, 3H, CH3),
6.34 (s, pyridyl H), 7.3-7.6 (m, 3H, subs. phenyl), 7.38-7.96
(m, 4H, fused phenyl). The diazocoupling reaction27 for
synthesis of compound (3) is illustrated in Scheme 1.
Solutions
Standard stock solutions (5 × 10-3 mol L-1) of each of
the investigated compounds (1, 2 and 3) were prepared in
appropriate volume of ethanol (Merck). Desired diluted
solutions were prepared by accurate dilution of the standard
stock solutions with ethanol. Solutions (2 × 10-3 mol L-1)
of each of the metal salts (CoCl2.6H2O, NiCl2.6H2O,
CuCl2.2H2O, La(NO3)3.6H2O and UO2(Ac)2.2H2O) were
o
brown crystals of compound (2). Yield 67%; mp 98 C;
anal. calcd. for C16H14N4O (278.12 g mol-1) C, 69.05;
H, 5.07; N, 20.13; found C, 69.73; H, 5.65; N, 20.52;
Scheme 1. Diazo-coupling reaction for synthesis of the two azo-compounds (2 and 3).

Vol. 23, No. 8, 2012
El-Attar et al.
1525
also prepared by dissolving the accurate weight of the
metal salt in appropriate volumes of deionized water.Also,
0.01 mol L-1 HCl, 1 mol L-1 KCl and 0.02 mol L-1 NaOH
solutions were prepared in deionized water.
A set of the Britton-Robinson (B-R) universal buffer
of pH values 2-11.5 was prepared in bidistilled water and
used as a supporting electrolyte in the presence of 40%
(v/v) ethanol. All chemicals used were of analytical grade
(BDH or Merck) and were used without further purification.
Mercury, gold, platinum and various carbonaceous
materials are the most frequently used working electrodes
in electrochemical studies. The use of mercury electrodes
has been recently discouraged mainly due to environmental
reasons (Hg is a toxic element). Mercury has a high
hydrogen overvoltage that greatly extends the cathodic
potential window up to about -2 V in aqueous solution.
Due to its wide cathodic potential window, mercury was
frequently used as working electrodes (e.g., DME and
HMDE) for electroreduction of substances of very negative
reduction potential in aqueous electrolyte. Besides, it was
successfully used as a working electrode for substance that
strongly adsorbs onto the surface of solid electrodes. The
mercury electrode is far superior to other solid electrodes
in this regard since great care must be exercised with such
solid electrodes to ensure that their surfaces are not changed
by the electrochemical reactions. On the other side, mercury
electrode (e.g., DME) also possesses a highly reproducible,
readily renewable and smooth surface.
The appropriate concentration of each of the examined
compounds in 10 mL B-R universal buffer in the presence of
40% (v/v) ethanol was introduced into the electrochemical
cell, and then deoxygenated with pure nitrogen gas for
about 10 min and for 30 s in each successive cycle, while
a stream of nitrogen gas was kept over the solution during
the measurements.
A potentiostat/galvanostat model 173-PAR incorporated
with a digital coulometer model 179-PAR (Princeton
Applied Research, Oak Ridge, TN, USA) was used
for the controlled-potential coulometric measurements
at a mercury pool cathode. A micro-coulometric cell
incorporated with a Pt wire sealed through the cell bottom
for contact with the mercury pool as a working electrode, a
reference SCE and a platinum gauze as a counter electrode
was used.
Controlled-potential electrolysis of solution
(2.5 × 10-4 mol L-1) of each of the investigated compounds
was performed in the B-R universal buffer of various pH
values containing 40% (v/v) ethanol to determine the total
number of electrons consumed in the overall electrode
reaction. The applied potential was adjusted to be around
the half-wave (E1/2) potential ( 0.1 V vs. SCE) or at the
plateau of the limiting current of each of the recorded
reduction waves. Prior to measurements, the electrolyzed
solutions were deoxygenated by bubbling with pure
nitrogen gas. During the measurements, a constant
stream of nitrogen gas was passed over the surface of
the electrolysis solution. The total charge Q (Coulombs)
consumed during the complete electrolysis of the reactant
was calculated by electronically integrating the current,
after subtracting the background current. The number
Apparatus and procedures
Physical measurements
Elemental analysis of the synthesized azo-compounds
(2 and 3) was carried out using Perkin Elmer 2400 elemental
analyzer (Central Laboratory for Searching & Microanalysis,
Tanta University, Egypt). Their melting points were measured
using a Gallenhamp apparatus and were uncorrected.
Infrared (IR) spectra of the solid compounds were recorded
on a Jasco FT/IR-4100-A spectrophotometer within
the range 4000-400 cm-1 as KBr discs (Microanalytical
1
unit, Kafr El-Sheikh University, Egypt). H nuclear
magnetic resonance (NMR) spectra were recorded on a
Varian EM 390-90 NMR spectrometer (Micro Analysis
Center, Cairo University, Egypt) using DMSO-d6 (Merck) as
a solvent and tetramethylsilane (TMS) as an internal standard
at room temperature. The 1H NMR chemical shifts (d, ppm)
are given relative to residual solvent peak. Mass spectral
measurements of the synthesized compounds were made
on a DI Analysis Shimadzu Qp-2010 Plus (Micro Analysis
Center, Cairo University, Egypt).
Electrochemical measurements
A Sargent-Welch Polarograph model 4001 (Fisher,
USA) was used in the dc-polarographic measurements.
The electrochemical cell used was as described by Meites.28
Characteristics of the capillary used as dropping mercury
electrode (DME) were m of 1.2 mg s-1 and t of 3.5 s in a
solution of 0.1 mol L-1 KCl at open circuit and a mercury
height of 60 cm. A saturated calomel electrode (SCE) was
used as a reference electrode.
A computer-controlled Potentiostat/Galvanostat model
273A-PAR (Princeton Applied Research, Oak Ridge, TN,
USA) and the electrode assembly model 303A-PAR with
the software package 270/250-PAR were used in cyclic
voltammetric measurements. A micro-electrochemical
cell incorporated of a three-electrode configuration system
comprising of a hanging mercury drop electrode (HMDE)
as a working electrode (surface area of 0.026 cm2), an
Ag/AgCl/KCls reference electrode and a platinum wire
auxiliary electrode was used.

1526
Synthesis, Electrochemical, Spectrophotometric and Potentiometric Studies of Two Azo-Compounds
J. Braz. Chem. Soc.
of electrons (n) transferred per electrolyzed molecule
at various pH values was estimated using Faraday’s
equation: N = Q/nF where N is the number of moles of
substance being electrolyzed and F the Faraday’s constant
(96,485 C mol-1).
A pH-meter model HI8014 (Hanna Instruments,
Italy) accurate to 0.01 pH units was used for the pH
measurements. A Mettler balance (Toledo-AB104,
Greifensee, Switzerland) was used for weighing the
solid materials. Deionized water was obtained from a
Purite-Still Plus deionizer connected to an AquaMatic
double-distillation water system (Hamilton Laboratory
Glass Ltd., Kent, UK).
metal ion (Co(II), Ni(II), Cu(II), La(III) or UO22+) + 15 mL
EtOH + 15 mL H2O.
A constructed QuickBasic language-PC program was
used in computing the data resulted from potentiometric
measurements to estimate the dissociation constants (pKa)
of examined azo-derivatives and the stability constants and
stoichiometry of their metal-complexes in solution.
Results and Discussion
Characterization of the synthesized azo-derivatives (2 and 3)
As described in the Experimental section, the results
provided by the different techniques are in good agreement
with the chemical formulae proposed for the synthesized
compounds 2 and 3. IR spectra of compounds 2 and 3
are shown in the Supplementary Information (SI) section
(Figure S1). In this section, the mass spectra of compounds
2 and 3 (Figure S2) and their fragmentation patterns
(Figures S3 and S4) are presented. The IR spectral bands
of compounds 2 and 3 at 1650-1662 cm-1 were assigned
for the N=N group. The stretching vibration bands, u(OH),
of the hydroxyl group were found at 3327-3403 cm-1.
The bands at 2948-2975 cm-1 characterize the u(CH3)
stretching vibration. In compound 3, the band at 1390 cm-1
characterizes the nitro group (Figure S1 in the SI section).
Spectrophotometric measurements
UV-Visible absorbance spectra of the examined
compounds were recorded within the wavelength range
200-800 nm at room temperature using a Shimadzu UV-Vis
spectrophotometer model 160A (Kyoto, Japan) with a
quartz spectrometric cell (1 cm bath length). The absorbance
spectra were scanned for the following mixtures: 5 mL of
B-R universal buffer + 4.8 mL H2O + 0.2 mL EtOH as
blank, while for the sample 5 mL of B-R universal buffer +
4.8 mL H2O + 0.2 mL of 2.5 × 10-3 mol L-1 azo-compound
dissolved in ethanol.
1
Potentiometric measurements
The H NMR spectrum of compound (2) exhibited
Potentiometric measurements were performed using a
pH-meter model HI8014 (Hanna Instruments, Italy) accurate
to 0.01 pH units. The electrode was standardized before and
checked after each titration with standard buffer solution
(Fisher, New Jersey, USA). A standard 0.02 mol L-1 NaOH
aqueous-ethanolic solution (40% v/v ethanol) was added from
a 5 mL total volume micro burette accurate to 0.01 mL and
the contents of the titration vessel were stirred using
a magnetic stirrer (Sargent-Welch, USA). All titration
measurements were carried out at 298 K by circulating
water from an Ultra-thermostat (JULABO F10, Seelbach,
Germany) through the annular space of a double-walled
Pyrex titration cell of 50 mL capacity. Each of the following
mixtures was prepared and potentiometrically titrated
against a standard 0.02 mol L-1 NaOH aqueous-ethanolic
solution (40% v/v ethanol) at 298 K. The volume was made
up to 50 mL using bidistilled water before the titration,
keeping ethanol content at 40% (v/v) in all titrated solution
mixtures: (i) 5 mL of 0.01 mol L-1 HCl + 5 mL of 1 mol L-1
KCl + 20 mL EtOH + 20 mL H2O; (ii) 5 mL of 0.01 mol L-1
HCl + 5 mL of 1 mol L-1 KCl + 5 mL of 5 × 10-3 mol L-1
azo compound + 15 mL EtOH + 20 mL H2O; and (iii) 5 mL
of 0.01 mol L-1 HCl + 5 mL of 1 mol L-1 KCl + 5 mL of
5 × 10-3 mol L-1 azo-compound + 5 mL of 2 × 10-3 mol L-1
singlet for 3 protons of (CH3) group at d 2.41, singlet for
(NH2) at d 4.23, singlet for (OH) at d 5.11, singlet for
pyridyl proton attached to C3 at d 7.14, multiplet for 3
aromatic protons at d 6.55-7.32 and multiplet for 4 protons
of fused phenyl at d 7.42-8.12.
1
For compound (3), the H NMR spectrum exhibited
a singlet for 3 protons of (CH3) group at d 2.52, singlet
for pyridyl proton attached to C3 at d 6.34, multiplet for
3 aromatic protons of substituted phenyl at d 7.3-7.6 and
multiplet for 4 protons of fused phenyl at d 7.38-7.96.
The mass spectral assignment of compound (2) showed
a strong molecular ion peak [M+] at m/z (278, 48%)
(Figure S2 in the SI section). Decomposition of this ion
may take place through two pathways. In the first path, the
compound loses CH3 radical forming ion at m/z (263, 14%),
then it loses three H radicals, producing ion with a base peak
at m/z (260, 100%). The second path, [M+] lose C6H6ON
radical, forming ion at m/z (170, 56%), this decomposition
step was confirmed by the appearance of peak at m/z (108,
74%), characterized for C6H6ON radical. The fragment
ion at m/z (170, 56%) may lose nitrogen gas followed by
five H radicals forming ions at m/z (142, 34%) and (137,
1%), respectively (Figure S3 in the SI section). The mass
spectral assignment of compound (3) showed an intense

Vol. 23, No. 8, 2012
El-Attar et al.
1527
molecular ion peak [M]+ at m/z (308, 100%) corresponding
to the formula C16H12O3N4 (Figure S2 in the SI section).
The [M+] may decompose through several fragmentation
paths. In the first, the molecule loses C10H8N3 moiety,
forming ion with a peak at m/z 138, 37%. The decomposed
ion (C10H8N3) with peak at m/z (170, 43%) may lose CH3
group and N2 gas forming an ion at m/z 127, which may
take hydrogen radical, forming ion with peak at m/z (128,
72%). In the second path, [M+] lose nitro group forming
ion with peak at m/z (262, 59%). The third path, [M+] may
lose CH3 group and hydroxyl radical, forming ion at m/z
276, in which it takes three hydrogen radicals, forming ion
with peak at m/z (279, 64%) (Figure S4 in the SI section).
Electrochemical studies
DC-polarographic and cyclic voltammetric behaviors
of 4-amino-2-methylquinoline (1) and its two synthesized
azo-derivatives (2 and 3) of 2.5 × 10-4 and 1 × 10-4 mol L-1,
respectively, were studied in the B-R universal buffer of
various pH values (2-11.5) containing 40% (v/v) ethanol.
The voltammograms of 4-amino-2-methylquinoline (1)
exhibited a single 2-electron irreversible step in solutions
of pH values (7-11) at E1/2 ca. -1.52 to -1.71 V (Figure 1,
curves a, b and c, and Figure 2, curves a and b). For
2-methyl-4-(5-amino-2-hydroxy-phenylazo)-quinoline (2),
the voltammograms exhibited two main 2-electron
irreversible cathodic steps (E1/2 ca. -0.18 to -0.71 V and
-0.75 to -1.26 V) for the 1st and 2nd waves, respectively,
over the entire pH range (2-11.4) (Figure 1, curves d,
e and f, and Figure 2, curves c and d). Besides, an additional
irreversible cathodic step was obtained at more negative
potentials (E1/2 ca. -1.1 to -1.43V) in solutions of pH values
4-9 (Figure 1, curve e, and Figure 2, curve c); its limiting
or peak current was gradually decreased upon the increase
of pH of the medium until complete disappearance at pH
higher than 9 (Figure 1, curve f, and Figure 2, curve d).
For 2-methyl-4-(2-hydroxy-5-nitrophenylazo]-
quinoline (3), the voltammograms showed a main 8-electron
irreversible cathodic step (E1/2 ca. -0.34 to -0.97 V) in
solutions of pH values lower than 10 (Figure 1, 1st wave,
curves g, h and i, and Figure 2, curve e), which splits into
two steps in solutions of pH values higher than 10 (Figure 1,
curve j ,and Figure 2, curve g). An additional irreversible
cathodic step was also obtained at more negative potentials
(E1/2 ca. -1.4 V) in solutions of pH values 4-9 (Figure 1,
curves h and i, and Figure 2, curve f), its limiting current
gradually decreased upon the increase of pH of the medium
until complete disappearance at pH values higher than 9
(Figure 1, curve j, and Figure 2, curve g). The half-wave
potentials (E1/2) or the peak potentials (Ep) shifted to more
Figure 1. DC-polarograms of the examined compounds (2.5 × 10-4 mol L-1)
in the B-R universal buffer of various pH values containing 40% (v/v)
ethanol, compound (1): (a) 6.2, (b) 7.2 and (c) 11.0, compound (2): (d) 2.0,
(e) 7.1 and (f) 11.4, and compound (3): (g) 2.4, (h) 6.1, (i) 7.7 and (j) 11.4.
negative values on the increase of either the pH of the
medium and the scan rate n, indicating the involvement
of proton in the electrode processes29 and the irreversible
nature of the reduction processes,29,30 respectively.
The reduction waves of the examined compounds at
the DME were analyzed using the fundamental equation
for the irreversible polarographic waves:28
Ed.e. = E1/2 - (0.0591/αna) log (i/(il – i))
(1)
Plots of Ed.e vs. log (i/(il - i)) for the polarographic
waves of investigated compounds at various pH values were
straight lines with slope values S1 (S1, mV = 59.1/αna)28
reported in Table 1. Values of αna (α is the symmetry
transfer coefficient and na is the number of electrons
involved in the rate-determining step) at various pH values

1528
Synthesis, Electrochemical, Spectrophotometric and Potentiometric Studies of Two Azo-Compounds
J. Braz. Chem. Soc.
were determined from slope S1 (Table 1).Also the E1/2 vs. pH
plots for the polarographic waves (of pH-dependent
E1/2) of the examined compounds were straight lines
with slope values S2 reported in Table 1. The number of
protons (p) participating in the rate-determining step of
the reduction process of the electro-active centers of the
investigated compounds was determined by applying the
relationship:29,31
d E1/2/d pH = S2 mV = (59.1/αna) p
p = S2 (αna/59.1) = S2/S1
(2)
(3)
The data reported in Table 1 indicated that the rate-
determining step of the reduction processes involved the
consumption of one proton (i.e., p = 1) over the entire pH
range. The ratio (p/na) may have the value 1 (when value of
p and na are equal) or 0.5 (when value of na is double that
of p). The most probable values obtained for α-parameter
(0.40-0.55) indicated that the ratio (p/na) equals 0.5, which
suggested that the number of electrons na involved in the
rate-determining step of each of the electro-active centers
should be double that of the involved protons (p), i.e.,
na = 2 and p = 1.
On the other side, linear Ep-pH plots of slope values
(58-60 mV) at scan rate of 200 mV s-1 for the examined
compounds were obtained. From slope values (13-15 mV)
of the obtained linear plots of Ep versus ln v (slope,
mV = 12.85/αna),32 values of αna (0.99-0.86) and
α (0.43-0.49) were also estimated at na = 2, confirming
again the irreversible nature32 of the electrode processes
of the examined compounds. Moreover, values of αna and
α were also determined at various pH values using the
equation: Ep - Ep/2 = 0.048/αna V,30 where (Ep - Ep/2) is
the difference between cathodic peak potential Ep and the
half-peak Ep/2 at half-height of peak current as a function
Figure 2. Cyclic voltammograms of the examined compounds
(1 × 10-4 mol L-1) in the B-R universal buffer of various pH values
containing 40% (v/v) ethanol, compound (1): (a) 7.4 and (b) 9.2,
compound (2): (c) 4.1 and (d) 11.2, and compound (3): (e) 3.3, (f) 7.4 and
(g) 11.1; scan rate 500 mV s-1.
Table 1. Dc-polarographic data for the investigated compounds in ethanolic-aqueous B-R universal buffer solutions containing 40% (v/v) ethanol at 25 oC
pH range
Compound (1) - reduction of quinoline moiety in a single step at various pH values:
7.0-11.0 64.5-66.7 0.91-0.88
S1 / mV
αna
S2 / mV
53
p (S2/S1)
0.82-0.80
p ca. 1
Compound (2) - reduction of N=N group in two steps (a) and (b), and reduction of quinoline moiety in a single step (c) at various pH values:
2.0-11.4 (a)
2.0-11.4 (b)
4.0-9.0 (c)
68.9-69.0
66.5-70.0
64.4-70.3
0.86
56.5
59.3
74.9
0.82
p ca. 1
0.89-0.84
0.92-0.84
0.89-0.85
p ca. 1
1.16-1.07
p ca. 1
Compound (3) - reduction of both N=N azo and NO2 groups in a single step:
2.4-11.4 78.5-83.4 0.75-0.71
74
0.94-0.89
p ca. 1

Vol. 23, No. 8, 2012
El-Attar et al.
1529
of scan rate v. The obtained αna values were found to equal
(0.96-1.06) and consequently values of the symmetry
transfer coefficient (α) over the pH range were found to
equal (0.48-0.53) at na equals 2. This provided additional
support of the irreversible nature32 of the electrode
processes of the examined compounds.
On the other hand, the UV-Vis absorbance spectra of
solutions(1×10-4 molL-1)oftheinvestigatedazo-derivatives
in ethanolic-aqueous B-R universal buffer (e.g., pH 7) before
controlled potential electrolysis showed a characteristic
band at lmax 400 nm due to the n-p* transition of -N=N-
group (e.g., Figure 3; spectrum (a) for compound 3).
This band disappeared completely after controlled-potential
electrolysis of the examined solutions (e.g., Figure 3;
spectrum (b) for compound (3)). It is important to notice
that the spectra of solutions of both derivatives (2) and (3)
after controlled-potential electrolysis resemble that of
the starting compound (1) (Figure 3; spectrum c). This
behavior confirmed the cleavage of the -N=N- bond in
both derivatives by controlled-potential electrolysis.
Controlled-potential electrolysis and TLC studies
As shown in Table 2, the total number of electrons
(n) transferred per molecule in the reduction process at
various pH values for 2.5 × 10-4 mol L-1 solution of each
of the examined compounds (1, 2 and 3) were determined
by means of controlled-potential electrolysis as discussed
in the Experimental section.
Table 2. Results of controlled-potential coulometry measurements for
the examined compounds
Number of electrons (n) transferred/molecule
Compound
pH
1st wave 2nd wave 3rd wave
Total
1
8.0
11.0
2.8
2
2
2
2
2
8
8
4
-
-
2
2
2
-
2
4
-
-
-
2
-
-
-
-
2
2
2
3
4
6.5
6
11.0
2.4
4
8
6.1
10
8
11.0
Thin layer chromatographic (TLC) experiments were
carried out for monitoring the products of complete
electrolysis of the solution of each of the investigated
azo-compounds (2 and 3) using benzene/acetone (70:30
v/v) as eluent. This was performed by concentration of the
completely electrolyzed solution of the each of examined
azo-compounds, and then followed by extraction of the
buffer ingredient with ether. The TLC-experiments showed
two clear spots compared to a single spot for the starting
compound (1), indicating that two products were formed
due to the reduction cleavage of the -N=N- double bond.
The two products were suggested to be the corresponding
primary aromatic amines. The presence of the latter one
was confirmed by carrying out diazo-coupling reaction27,33
on the completely electrolyzed solutions of each of the
investigated azo-compounds. Recovery of the characteristic
colored azo-dye27,33 in solution confirmed the presence of
aromatic amines in the completely electrolyzed solution as
a result of the reduction cleavage of -N=N- center in each
of the examined compounds (2 and 3), which confirmed
our suggested reduction mechanisms of the investigated
compounds.
Figure 3. UV-Vis spectra of the azo-compound (3) (1 × 10-4 mol L-1)
in ethanolic-aqueous B-R universal buffer solution of pH 7: (a) before
electrolysis and (b) after complete electrolysis at controlled-potential,
compared to that of 1 × 10-4 mol L-1 solution of the starting compound
(1) (c).
Electrode reaction pathways
Electrode reaction of compound (1)
4-Amino-2-methylquinoline (1) was reduced at the
mercury electrode in buffered solutions of pH values 7-11
in a single 2-electron irreversible step which is assigned to
the reduction-saturation of the -C=N- double bond of its
quinoline moiety,34 as illustrated in Scheme 2.
Electrode reaction of the azo-derivative (2)
The synthesized 2-methyl-4-(5-amino-2-hydroxy-
phenylazo)-quinoline (2) was reduced at the mercury
electrode in buffered solution of pH 2-11.5 in two 2-electron
steps (1st and 2nd steps), which were attributed to the
reduction of the -N=N- double bond to the amine stage as
illustrated in Scheme 3. The 3rd reduction step that appeared

1530
Synthesis, Electrochemical, Spectrophotometric and Potentiometric Studies of Two Azo-Compounds
J. Braz. Chem. Soc.
Scheme 2. Electrode reaction pathway of the parent compound (1).
Scheme 3. Electrode reaction pathway of the azo-derivative (2).
at more negative potential in solutions of pH values 4-9 is
due to reduction of the -C=N- double bond of the quinoline
moiety,34 as illustrated in Scheme 2 of the electrode reaction
of the starting compound (1).
The reduction step that appeared at more negative
potential in solutions of pH values 4-9 is due to reduction
of the -C=N- double bond of the quinoline moiety34 as
illustrated in Scheme 2 of the electrode reaction of the
starting compound (1).
Electrode reaction of azo-derivative (3)
The synthesized 2-methyl-4-(2-hydroxy-
5-nitrophenylazo)-quinoline (3) was also reduced at the
mercury electrode in buffered solution of pH values 2-10 in
a single 8-electron step, which is attributed to reduction of
either the -N=N- double bond to the amine stage and the
-NO2 group to the hydroxylamine stage, via the consumption
of four electrons for each centre,24,35 (Scheme 4, reaction 4).
At pH values higher than 10, this reduction step splits
into two steps. The 1st and 2nd steps were attributed to the
reduction-cleavage of the -N=N- double bond to the amine
stage and reduction of the NO2 group to the hydroxylamine
stage,35 respectively, via the consumption of four electrons
for each center (Scheme 4, reactions 5 and 6, respectively).
Spectrophotometric studies
UV-Vis absorption spectra of 5 × 10-5 mol L-1 of each of
compounds (2) and (3) were recorded in the B-R universal
buffer of various pH values (2-12) within the wavelength
range 200-800 nm. Compound 2 showed two absorption
bands at lmax 438 and 480 nm (n-p ), with an isosbestic point
at lmax 462 nm (Figure 4). While compound 3 showed two
absorption bands at lmax 315 and 400 nm (n-p ), with an
isosbestic point at lmax 345 nm (Figure 4).
The first band may be attributed to the non-ionized form,
whereas the second one that develops as the pH increases
is due to the ionized species of each of the examined
*
*

Vol. 23, No. 8, 2012
El-Attar et al.
1531
Scheme 4. Electrode reaction pathway of azo-derivative (3).
compounds as a result of proton dissociation of the hydroxyl
group. The isosbestic point indicated the presence of an
acid-base equilibrium between non-ionized and ionized
species. The ionic form absorbs at longer wavelength
indicating that the interamolecular charge transfer is easier
in the ionic form than in the non-ionic one, confirming that
the interamolecular charge transfer is influenced by the OH
group. For 4-amino-2-methylquinoline (1), its absorbance
spectra was also recorded under the same experimental
conditions and was found to exhibit absorption band at
pH curves at two different wavelengths were typical Z or
S-shaped (Figure 5) which indicated the transformation of
the molecule from one form into another one.
Values of the dissociation constant (pKa) of the
investigated compounds were determined at two different
wavelengths (except compound 1) using different
methods namely: (i) half-height, (ii) limiting absorbance,
(iii) modified limiting absorbance, (iv) isosbestic point and
(v) modified isosbestic point.36-38 The obtained pKa values
are reported in Table 3.
l
max 330 nm at pH values lower than 8, whereas, at higher
The observed high values of pKa of the synthesized
azo-derivatives (2 and 3) may be attributed to the location
of the ionizable group (-OH) in an ortho-position where
it has lower ability to release H+ ion.39,40
pH values, this band was blue shifted to lmax 300 nm. This
behavior may be due to that protonation of species takes
place on the nitrogen of quinoline moiety in solutions of pH
lower than 8. However, with the increase of the pH values
of the medium, deprotonation takes place. The changes in
the absorption spectra on varying the pH of solutions were
used for determination of the dissociation constant (pKa)
of the examined compounds (1, 2 and 3). The absorbance-
Potentiometric studies
The average number of protons associated with the
—
reagent molecule nA at different pH values41 was calculated

1532
Synthesis, Electrochemical, Spectrophotometric and Potentiometric Studies of Two Azo-Compounds
J. Braz. Chem. Soc.
Figure 4. UV-Vis spectra of the examined compounds (5 × 10-4 mol L-1)
in B-R universal buffer of various pH values, compound (2): (a) 2,
(b) 4, (c) 6, (d) 9, (e) 10 and (f) 11, and compound (3): (g) 3.0, (h) 6.3,
(i) 7.3 and (j) 9.2.
from the titration curves of HCl solution in the absence and
in presence of the examined azo-derivative (2 or 3) against
0.02 mol L-1 NaOH aqueous-ethanolic solution (40%
v/v ethanol) at 298 K (Figure 6). Thus, the formation
Figure 5. Plots of absorbance vs. pH of the investigated compounds
(5 × 10-4 mol L-1) in B-R universal buffer of different pH values,
compound (1): lmax at 330 nm (a), compound (2): lmax at 438 (b) and
480 nm (c), and compound (3): lmax at 315 (d) and 400 nm (e).
—
curves (nA vs. pH) for the proton-ligand systems42 were
constructed and found to extend between 0 and 1.0 in
—
the nA scale for both compounds (2 and 3). This means
that each of the examined compounds 2 and 3 has one
dissociable proton from the enolized hydrogen ion of the
hydroxyl group (Table 4). This indicates that the acid-base
equilibrium based on the molecular form HL is in complete
accordance with the following equilibrium:
As shown in Table 4, compound (2) has a lower acidic
character (pK1H = 8.75) than compound (3) (pK1H = 7.5).
This is quite reasonable because the presence of NH2 group
enhances the electron density by its high electrondonating
mesomeric and inductive effects, respectively; thereby
stronger O-H bond is formed. The presence of electron
withdrawing group NO2 leads to the opposite effect,39,40
thus the pK1H values for the examined compounds follow
(4)
Table 3. Dissociation constants (pKa) of the compounds under investigation in ethanolic-aqueous B-R universal buffer using: the half-height (i), limiting-
absorbance (ii), modified limiting-absorbance (iii), isosbestic point (iv) and modified isosbestic point (v) methods
Dissociation constants (pKa)
l
max / nm
(n-p )
Compound
Mean value
*
(i)
(ii)
10.1
11.05
6.35
8.20
6.95
(iii)
9.75
9.35
8.40
7.35
7.00
(iv)
-
(v)
-
1
2
330
438
480
315
400
9.8
9.4
8.6
7.4
6.9
9.88
9.53
7.78
7.45
6.95
9.55
8.30
3
7.25
7.05

Vol. 23, No. 8, 2012
El-Attar et al.
1533
Table 4. Stepwise stability constants (ML and ML2 chelates) of the synthesized azo-derivatives (2 and 3) in 40% (v/v) ethanol-water solution containing
1 mol L-1 KCl at 298 K
Compound (2)
Compound (3)
M+n
pK1H
log K1
log K2
pK1H
log K1
log K2
H+
8.75 0.09
7.5 0.1
Co2+
Ni2+
Cu2+
La3+
UO22+
5.87 0.07
6.15 0.07
6.48 0.06
6.30 0.10
6.62 0.09
4.02 0.06
4.29 0.09
4.55 0.08
4.43 0.09
4.75 0.10
6.32 0.06
6.95 0.10
8.45 0.11
7.54 0.08
9.26 0.08
5.11 0.08
5.83 0.12
6.30 0.07
6.10 0.09
6.80 0.10
Each value is the mean of three readings standard deviation.
complexes of the two azo-derivative (2 and 3) with the
selected metal ions (Co(II), Ni(II), Cu(II), La(III) or
UO22+) were obtained by plotting the average number
–
of ligands attached per metal ion (n) vs. the free ligand
exponent (pL).42
By analysis of the formation curves,43,44 the successive
stability constant log K1 and log K2 of the studied metal
complexes ML and ML2, respectively, were determined
(Table 4) using a constructed QuickBasic language-PC
program. As shown in Table 4, the differences observed
in the standard deviations (SD) of the obtained data under
the same experimental conditions were insignificant,
confirming reproducibility and precision of the results.
The data obtained can be pointed out based on the
following: (i) no precipitate was observed in the titration
vessel, indicating that the formation of metal hydroxide is
–
Figure 6. Potentiometric titration curves against 0.02 mol L-1 NaOH
ethanolic-queous 40% (v/v) solution at 298 K for the solution mixtures:
(a) 5 mL of 0.01 mol L-1 HCl + 5 mL of 1 mol L-1 KCl + 20 mL EtOH +
20 mL H2O, (b) 5 mL of 0.01 mol L-1 HCl + 5 mL of 1 mol L-1 KCl + 5 mL
of 5 × 10-3 mol L-1 azo-derivative (2) + 15 mL EtOH + 20 mL H2O, (c) 5 mL
of 0.01 mol L-1 HCl + 5 mL of 1 mol L-1 KCl + 5 mL of 5 × 10-3 mol L-1
azo-derivative (2) + 5 mL of 2 x 10-3 mol L-1 Co(II) + 15 mL EtOH +
15 mL H2O, (d) the same solution mixture + 2 × 10-3 mol L-1 Ni(II), and
(e) the same solution mixture + 2 × 10-3 mol L-1 Cu(II).
excluded; (ii) the maximum n values for all investigated
metal complexes were found to be ca. 2, revealing that
both ML and ML2 types of complexes are formed in
solution; (iii) for all complexes formed log K1 were always
found higher than those of log K2 (Table 4), because the
vacant sites of the metal ions more freely available for
binding of a first ligand than for a second one;40,45 (iv) the
order of stability constants of the metal complexes of
the azo-compounds (2 and 3) at 298 K were found as:
UO22+ > La(III) < Cu(II) > Ni(II) > Co(II). The sequence
of stability of complexes of compounds (2) and (3) with
Cu(II) > Ni(II) > Co(II) are in agreement with that reported
by Irving and Williams.46,47 The UO22+ has higher stability
than those of the other metal complexes. This may be
attributed to the bonded oxygen atoms which increase
the electrostatic attraction between the metal ion and the
coordinated ligands and overcome any steric hindrance
offered by the oxygen of the oxygenated cations.40
the sequence 2 > 3. The results obtained are of the same
order when compared to those spectrophotometrically
obtained (Table 3).
The average number of reagent molecules attached per
–
metal ion n, at different pH values, was calculated from the
titration curves of HCl solutions in the absence and the
presence of the examined azo-derivative (2 or 3) and the
selected metal ions (Co(II), Ni(II), Cu(II), La(III) or UO22+)
against 0.02 mol L-1 NaOH aqueous-ethanolic solution
(40% v/v ethanol) at 298 K (Figure 6).
The metal-titration curves (curves c-e) were well
separated from that of free ligand (curve b), along the axis
of added volume of NaOH solution, which is attributed
to the release of H+ ions as a result of the complexation
process (Figure 6). The formation curves for the metal
Conclusions
Two azo-compounds (2-methyl-4-(5-amino-2-hydroxy-
phenylazo)-quinoline (2) and 2-methyl-4-(2-hydroxy-5-

1534
Synthesis, Electrochemical, Spectrophotometric and Potentiometric Studies of Two Azo-Compounds
J. Braz. Chem. Soc.
nitrophenylazo)-quinoline (3)) derived from 4-amino-2-
methylquinoline (1) were synthesized and their chemical
structures confirmed by various techniques. Their electrode
reaction pathways at the DME and HMDE electrodes in
buffered solutions of pH values (2-11.5) were elucidated.
The dissociation constants (pKa) of the investigated
compounds were determined using spectrophotometric and
potentiometric methods. Besides, stability constant and
stoichiometry of their metal complexes in solution
with some transition metal ions (Co(II), Ni(II), Cu(II),
La(III) and UO22+) were determined potentiometrically.
The results indicated the existence of ML and ML2 types
of complexes in solution.
15. Nejati, K.; Rezvani, Z.; Seyedahmadian, M.; Dyes Pigm. 2009,
83, 304.
16. Ahn, S. B.; Kim, J. P.; Shim, W. S.; Fibers Polym. 2009, 10,
643.
17. Albayrak, C.; Gümrükçüoglu, I. E.; Odabasoglu, M.; Iskeleli,
N. O.; Agar, E.; J. Mol. Struct. 2009, 932, 4354.
18. Byabartta, P.; Russ. J. Coord. Chem. 2009, 35, 687.
19. Ghoneim, M. M.; El-Desoky, H. S.; Amer, S. A.; Rizk, H. F.;
Habazy, A. D.; Dyes Pigm. 2008, 77, 493.
20. Chandra, U.; Gilbert, O.; Swamy, B. E. K.; Bodke, Y. D.;
Sherigara, B. S.; Int. J. Electrochem. Sci. 2008, 3, 1044.
21. Menek, N.; Karaman, Y.; Dyes Pigm. 2006, 68, 101.
22. Menek, N.; Karaman, Y.; Dyes Pigm. 2005, 67, 9.
23. Menek, N.; Basaran, S.; Turgut, G.; Odabasoglu, M.; Dyes
Pigm. 2004, 61, 85.
Supplementary Information
24. La-Scalea, M. A.; Serrano, S. H. P.; Gutz, I. G. R.; J. Braz.
Chem. Soc. 1999, 10, 127.
Supplementary data of IR, and MS spectra of
the two synthesized azo-compounds (2-methyl-
4-(5-amino-2-hydroxy-phenylazo)-quinoline (2) and
2-methyl-4-(2-hydroxy-5-nitrophenylazo)-quinoline
(3)) (Figures S1-S4) are available free of charge at
http://jbcs.sbq.org.br as PDF file.
25. Goyal, R. N.; Verma, M. S.; Singhal, N. K.; Croat. Chem. Acta
1998, 71, 715.
26. Moharram, Y. I.; El-Hallag, I. S.; El-Attar, M. A.; Ghoneim,
M. M.; Bull. Electrochem. 2005, 21, 297.
27. Vogel, A. I.; A Text Book of Practical Organic Chemistry,
3rd ed.; Longman Publishers: New York, 1961.
28. Meites, L.; Polarographic Techniques, 2nd ed.; Interscience
Publishers: New York, 1965, p. 232-248.
References
1. Raposo, M. M; Fonseca, A. C.; Castro, M. R.; Belsley, M.;
Cardoso, M. S.; Carvalho, L. M.; Coelho, P. J.; Dyes Pigm.
2011, 91, 62.
29. Zuman, P.; The Elucidation of organic Electrode Process;
Academic Press: New York, 1969. p. 22-24.
30 Greef, R.; Peat, R.; Peter, L. M.; Pletcher, D.; Robinson, J.;
Instrumental Methods in Electrochemistry, Southampton
Electrochemistry Group; Ellis Horwood Limited, John Wiley
& Sons: New York, 1985, p. 187.
2. LiHong, N.; Cheng, Z.; ZiHui, C.; Zhi, Z.; Zhong, Yu, L.;
FuShi, Z.; YingWu, T.; Chin. Sci. Bull. 2009, 54, 1169.
3. Sacarescu, L.; Sacarescu, G.; Simionescu, M.; Hurduc, N.;
J. Inorg. Organomet. Polym. 2009, 19, 348.
31. Ghoneim, M. M.;Ashy, M.A.A.; Can. J. Chem. 1979, 57, 1249.
32. Galus, Z.; Fundamentals of Electrochemical Analysis; John
Wiley: New York, 1976, p. 238.
4. Dixit, B. C.; Patel, H. M.; Desai, D. J.; J. Serb. Chem. Soc.
2007, 72, 119.
5. Khosravi, A.; Moradian, S.; Gharanjig, K.; Taromi, F. A.; Iran.
Polym. J. 2005, 14, 667.
33. Pavia, D. L.; Lampman, G. M.; Kriz, G. S.; Engel, R. G.;
Introduction to Organic Laboratory Techniques - Small-
Scale Approach, 1st ed.; Harcourt Brace: Florida, 1998,
p. 313-322.
6. Raposo, M. M.; Castro, M. R.; Fonseca,A. C.; Schellenberg, P.;
Belsley, M.; Tetrahedron 2011, 67, 5198.
7. Raposo, M. M.; Castro, M. R.; Belsley, M.; Fonseca, A. C.;
Dyes Pigm. 2011, 91, 454.
34. Stradyn,Y. P.; Kadysh, V. P.; Giller, S. A. I.; Khim. Geterotsikl.
Soedin. 1973, 12, 1587.
8. He, T. C.; Wang, C. S.; Pan, X.; Zhang, C. Z.; Lu, G. Y.; Appl.
Phys. B: Lasers Opt. 2009, 94, 653.
35. Zuman, P.; Microchem. J. 1997, 57, 4-51.
36. Issa, R. M.; El-Ezaby, M. S.; Zewail, A. H.; Z. Phys. Chem.
(Leipzig) 1970, 244, 155.
9. Maruo, Y. Y.; Akaoka, K.; Nakamura, J.; Sens. Actuators, B
2010, 143, 487.
37. Issa, I. M.; Issa, R. M.; El-Ezabey, M. S.;Ahmed,Y. Z.; Z. Phys.
Chem. (Leipzig) 1969, 242, 169.
10. El-Safty, S. A.; Adsorption 2009, 15, 227.
11. Sharma, A. K.; Singh, I.; Food Anal. Methods 2009, 2, 221.
12. Sheremetev, S. V.; Kuznetsov, V. V.; J. Anal. Chem. 2007, 62,
319.
38. Abdelhal, F. M.; Issa, R. M.; El-Ezabey, M. S.; Hasanein,
A. A.; Z. Phys. Chem. (Frankfurt) 1970, 73, 59.
39. McMurry, J.; Organic Chemistry, 5th ed.; Brooks/Cole: USA,
2000.
13. Li, X.; Wu, Y.; Gu, D.; Gan, F; Dyes Pigm. 2010, 86, 182.
14. Sujamol, M. S.; Athira C. J.; Sindhu,Y.; Mohanan, K;
Spectrochim. Acta, Part A 2010, 75, 106.
40. Shehatta, I.; Kenawy, I.; Askalany, A. H.; Hassan, A. A.; Can.
J. Chem. 2001, 79, 42.

Vol. 23, No. 8, 2012
El-Attar et al.
1535
41. Fronaeus, S.; Complex System of Copper; Gleer Upska
Universitets-Bokhandeln: Lund, 1948.
45. Ghoneim, M. M.; El-Hallag, I. S.; El-Baradie, K.Y.; El-Desoky,
H. S.; El-Attar, M. A.; Monatsh. Chem. 2006, 137, 285.
46. Irving, H. M.; Williams, R. J. P.; Nature 1948,162, 746.
47. Irving, H. M.; Williams, R. J. P.; Analyst 1952, 77, 813.
42. Irving, H. M.; Rossotti H. S.; J. Chem. Soc. 1954, 76, 2904.
43. Rossotti, F. J. C.; Rossotti, H. S.; Acta Chem. Scand. 1955, 9,
1166.
44. Beck, M. T.; Nagybal, I.; Chemistry of Complex Equilibria;
Wiley: New York, 1990.
Submitted: April 3, 2012
Published online: July 31, 2012