Equilibrium and Structural Study of
m-Methyl Red in Aqueous Solutions:
Distribution Diagram Construction
Sa’ib J. Khouri, Ibrahim A. AbdelRahim, Ehab M. Alshamaileh & Abdel
Mnim Altwaiq
Journal of Solution Chemistry
ISSN 0095-9782
J Solution Chem
DOI 10.1007/s10953-013-0068-9
1 23
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1 23
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J Solution Chem
DOI 10.1007/s10953-013-0068-9
Equilibrium and Structural Study of m-Methyl Red
in Aqueous Solutions: Distribution Diagram
Construction
Sa’ib J. Khouri • Ibrahim A. Abdel-Rahim • Ehab M. Alshamaileh
Abdel Mnim Altwaiq
•
Received: 20 August 2012 / Accepted: 30 April 2013
Ó Springer Science+Business Media New York 2013
Abstract The UV/Vis spectra of the m-methyl red (m-MR) ({3-[4-(dimethyl-amino)
phenylazo] benzoic acid}) were examined in aqueous solutions at various acidities. These
were characterized by the overlap of the different bands of m-MR. The thermodynamic
acid dissociation constant, Ka2, of the equilibrium between m-H2MR? (diprotic form) and
m-HMR (monoprotic form) was determined as 1.02 9 10-2 at 25 °C, and that for the
equilibrium between m-HMR and MR- (basic form), Ka3, was determined as 4.94 9 10-5
at 25 °C. Based on the two observed Ka values, the distribution diagram of the three forms
of m-MR in water was constructed. The neutral monoprotic form (m-HMR) has a maximum fraction of 0.883 at pH = 3.14.
Keywords Azo-dye m-Methyl red UV/Vis spectra Acid dissociation
constant Distribution diagram
1 Introduction
m-Methyl red (m-MR), 3-[4-(dimethyl-amino)phenylazo] benzoic acid, is one of the three
isomers of methyl red that is one of a large number of indicators based upon aminoazobenzene, which is still a useful volumetric indicator. Methyl red isomers have been the
subject of several investigations, such as photochromism [1], inclusion complexes with
cyclodextrins [2, 3], pK determination [4–6], and structure [7]. m-MR has many
S. J. Khouri (&)
Department of Chemistry, American University of Madaba (AUM), Madaba, Jordan
e-mail: sbkhouri@yahoo.com; s.khouri@aum.edu.jo
I. A. Abdel-Rahim E. M. Alshamaileh
Department of Chemistry, University of Jordan, Amman, Jordan
A. M. Altwaiq
Department of Chemistry, Petra University, Amman, Jordan
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applications in many fields such as the paper and textile industries [8], ink-jet printing [9],
and as acid–base indicators [10].
o-Methyl red (o-MR, the main isomer) has been used by several authors as an example
for determining the pK value of an acid–base indicator by a spectrophotometric method
[11–13]. This indicator has four possible basic centers to receive added protons with
increasing of the acidity in aqueous solutions. Similarly m-MR has the same four basic
centers with only a change in the position of one group. These centers are the –COOgroup, the a- and b- nitrogen of the azo linkage, and the nitrogen of the dimethylamino
group, Fig. 1. The nonionic structure of m-methyl red (m-HMR, neutral form, structure B
in Fig. 1) forms from m-MR- (the basic form, structure A in Fig. 1), where the first proton
adds to the carboxyl anion, followed with the diprotic form m-H2MR? (structures C, D,
and E, Fig. 1), and finally with the triprotic form m-H3MR2? (structure F, Fig. 1) [14].
The determination of Ka values of organic dyes in water using several developed
methods has great importance for many chemical practical applications and scientific
research areas, such as titration, solvent extraction, electrophoresis, chromatography, drug
synthesis, and many others [6]. Many researchers studied two equilibria of relevant dyes to
m-MR concerning the values of the acid dissociation constants (as Ka2 and Ka3 in Fig. 1),
H3C
N
α N
β N
H3C
CH3
N
HN+
H+
N
H+
Ka3
N
Ka2
CO2-
CH3
H3C
CH3
N
CO2H
CO2H
B
A
Kt
N
C
CH3
H3C
+
CH3
H
N N
NH
N+
CH3
HO2C
D
H+
Ka1
H
N+
N
N
N
H N+
CH3
CH3
HO2C
CO2H
E
F
Fig. 1 The acid–base equilibria of m-methyl red in aqueous acidic solutions
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and their results indicate that these equilibria are overlapping, which means that the dye
cannot be present solely in the monoprotonated form in solutions, which is the common
form in the two overlapping equilibria [14].
In this work, we examine spectrophotometrically the different m-MR forms and
structures in aqueous solutions at different acidity values. The values of Ka2 and Ka3 at
25 °C were evaluated. From the calculated Ka values, the distribution diagram of basic,
monoprotic, and diprotic forms of m-MR at 36 pH values was constructed. The absence of
such information in the literature on m-MR motivated us to carry out this study.
2 Experimental
The highly purified neutral form of m-MR was purchased from BDH chemicals, UK and
was used as received. Other chemicals used in this study were reagent grade. Typical stock
solutions of m-MR were prepared in the presence of dilute aqueous NaOH. A typical stock
solution had a concentration of 2.04 9 10-4 moldm-3 and a pH of 10.5. The pH of the
test solution was adjusted by adding an appropriate amount of either NaOH or HCl
solutions. The pH value of each prepared solution was measured with a pH meter
immediately after each absorption measurement by the UV/Vis spectrophotometer.
The UV/Vis spectra of the test solutions were recorded using a double-beam spectrophotometer (Carry 100 Bio Varian) and a quartz cell with optical path length of 1.00 cm.
The cell holder in the spectrophotometer was connected to a constant temperature water
bath thermostat, where its temperature was controllable to ±0.1 °C. The pH-meter (Hanna
Instruments pH 211 microprocessor pH-meter) was calibrated with different standard
buffer solutions. The UV/Vis instrument was connected to a personal computer for data
collection in ASCII-file format.
3 Results and Discussion
3.1 UV/Vis Spectra of m-MR in Acidic and Basic Solutions
Figure 2 shows three UV/Vis spectra of 1.537 9 10-5 moldm-3 m-MR at the pH values 7.27,
3.61, and 2.51 representing three different forms of the dye; the basic (m-MR-), neutral
(m-HMR), and di-protonated (m-H2MR?) forms, respectively. Figure 3 illustrates the transfer
of m-MR from the basic form to the neutral one, and Fig. 4 illustrates the transfer from the
neutral form to the di-protonated one by gradually lowering the pH value in each case.
In Fig. 3, the intensities of the main bands in the visible and UV regions decrease with
lowering pH, as a result of the gradual transformation of m-MR from its basic form with a
strong yellow color to its red colored neutral form. The neutral form of azo dyes generally
has a weaker absorption band and hence less intense color [15]. Another important
observation is that a band around 218 nm starts to appear as a result of the protonation of
the carboxylate ion to form the carboxyl group, which absorbs generally around 220 nm
[16]. The presence of an isosbestic point at 508 nm in Fig. 3 is an indication of the
presence of the first equilibrium between the basic and the neutral forms of the m-MR [17],
where a shoulder at 530 nm grows with decreasing pH, resulted from the formation of the
neutral form.
As the acidity of the m-MR solution increases, the intensities of the three bands around
505, 291 and 222 nm that are attributed to the diprotic form (m-H2MR?) increase, in addition
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Fig. 2 The UV/Vis spectra of
the basic, neutral, and diprotonated forms of
1.537 9 10-5 moldm-3 m-MR
Spectrum #
0.5
1
2
3
0.4
Absorbance
pH
Form
--------------------------------------7.27
3.61
2.51
basic
neutral
diprotonated
3
1
0.3
0.2
2
0.1
0.0
200
300
400
500
600
Wavelength (nm)
Fig. 3 The effect of acidity
(pH = 7.27–3.61) on the UV/Vis
spectra of 1.537 9 10-5
moldm-3 m-MR
0.5
Absorbance
0.4
0.3
Spectrum # pH λmax/nm
---------------------------------------1
7.27
449
2
4.90
453
3
4.79
454
4
4.62
455
5
4.25
458
6
3.88
465
7
3.61
471
1
7
0.2
0.1
0.0
200
508 nm
300
400
500
600
Wavelength (nm)
to formation of a shoulder centered at 400 nm, as illustrated in Fig. 4. All the bands undergo
red shifts accompanied by an increase in the intensity. The di-protonated species
(m-H2MR?) has two different tautomeric structures: ammonium tautomer (structure C,
Fig. 1) and two zwitterionic azonium tautomers (structures D and E, Fig. 1), hence
m-H2MR? is the most stable among the other forms (m-MR-, m-HMR and m-H3MR2?). The
shoulder centered at 400 nm is attributed to the ammonium tautomer, while the absorption at
505 nm is attributed to the azonium tautomer of an azobenzene dye [18]. The spectra in
Fig. 4 indicate that the band of the azonium tautomer increases while that of the ammonium
tautomer decreases as the HCl concentration is increased. This pattern is due to a shift in the
position of the tautomeric equilibrium of m-MR as the acidity of the solution is increased, a
phenomenon generally observed for the tautomerism of azobenzene dyes [18].
Another observation in Fig. 4 is the broadening of the n ? p* transition of the amino
group at 291 nm, which can be attributed to the resonance hybrid structures of m-H2MR?
where the lone pair on the amino group is delocalized and contributes to the resonance
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Fig. 4 The effect of acidity
(pH = 3.32–8.0 moldm-3
H2SO4) on the UV/Vis spectra of
1.537 9 10-5 moldm-3 m-MR
0.7
0.6
Absorbance
0.5
0.4
Spectrum # pH λ max/nm
---------------------------------------1
3.32
491
2
3.02
500
3
2.85
503
4
2.73
503
5
2.62
504
6
2.56
504
7
2.51
505
8
8.0M H 2SO4 506
9
3.7M H 2SO4 506
9
1
0.3
0.2
0.1
0.0
200
300
400
500
600
Wavelength (nm)
hybrid structures of the azonium tautomers (structures D and E, Fig. 1). The formation of
the azido group (C=N–) with delocalized p electrons will cause broadening of the n ? p*
peak towards the red shift since it will suffer more delocalized p ? p* character. Some
researchers reported that the maximum electronic absorption of the azido group is located
around 190 nm, but when it is coupled with ethylene group, the maximum undergoes a red
shift with increasing intensity and then is located around 220 nm [16]. The last observation
on Fig. 4 is the presence of the second isosbestic point centered at 464 nm, which represents the presence of the second equilibrium between the neutral and di-protonated forms
of m-MR.
When H2SO4 is used to attain higher acidities, the intensities of the absorption bands of
the di-protonated form of m-MR (m-H2MR?) decrease as the concentration of H2SO4 is
increased, and a new absorption maximum appears at 408 nm that represents the tri-protonated form of m-MR (m-H3MR2?) as indicated in Fig. 5. The solution becomes yellow in
color in 16 moldm-3 H2SO4. In this high H2SO4 concentration, the protonation does not
stop on the azo group, but the amino group is also protonated and the final structure possesses
two positive centers [6, 19]. In this case, a dramatic blue shift is noticeable for p ? p*
transitions from 505 to 407 nm, while a red shift occurs in the n ? p* transitions from 222 to
228 nm. When the second protonation at the amino group takes place in the harsh acidic
media, all of the resonance structures and tautomeric structures bear no more lone pairs of
electrons on the amino group, which is replaced by the ammonium group with positive
charge. This will cause the disappearance of the n ? p* peak of the amino group, as is seen
in Fig. 5. The last observation on Fig. 5 is the expected third isosbestic point centered on
443 nm which represents the presence of the third equilibrium between the di-protonated
form (m-H2MR?) and the tri-protonated form (m-H3MR2?).
3.2 Calculation of Acid Dissociation Constants of Acid–Base Equilibria of m-MR
The thermodynamic acid dissociation constants Ka2 and Ka3 of the equilibria between the
different m-MR forms in aqueous solution were determined based on the following
equations:
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Fig. 5 The effect of H2SO4 on
the UV/Vis spectra of
1.537 9 10-5 moldm-3 m-MR
0.8
0.7
Absorbance
0.6
Spectrum # [H2SO4 ] / M λmax/nm
---------------------------------------------------1
3.7
506
2
8.0
506
3
12.0
408
4
16.0
408
0.5
1
2
4
3
0.4
0.3
0.2
443 nm
0.1
0.0
200
300
400
500
600
Wavelength (nm)
Ka2
H2 MRþ Hþ þ HMR
½Hþ ½HMR cHþ cHMR
Ka2 ¼
½H2 MRþ cH2 MRþ
Ka2
HMR Hþ þ MR
½Hþ ½MR cHþ cMR
Ka3 ¼
½HMR]
cHMR
ð1Þ
ð2Þ
ð3Þ
ð4Þ
where H2MR?, HMR, and MR- represent the diprotic, neutral, and basic form of m-MR,
respectively, and ci represents the molar activity coefficient. The value of product,
½Hþ cHþ , is the activity of H? (aHþ ) and can be calculated from the pH value of a
solution. The activity coefficient of the diprotic form (cH2 MRþ ) was calculated according to
the Guggenheim extension of the Debye–Hückel equation for singly charged organic ions
in water at 25 °C [20]:
pffiffi
0:512 I
pffiffi þ 0:20 I
log10 cH2 MRþ ¼ ð5Þ
1þ I
where I represents the ionic strength of the solution. The activity coefficient of the basic
form (cMR ) was calculated at low ionic strength from the Debye–Hückel limiting law
pffiffi
(log10 cMR ¼ 0:512 I in water at 25 °C), and the activity coefficient of the neutral form
(cHMR) was assumed to be unity.
The equilibria of Eqs. 1 and 3 are overlapping, since the ratio ½Hþ ½HMR]/[H2 MRþ :
þ
½H ½MR =½HMR] is about 250 [4]. Equation 2 can be rewritten in terms of absorbance as
in the following form:
A Aa
aHþ cHMR
Ka2 ¼
ð6Þ
Ab A
cH2 MRþ
where A, Aa and Ab represent the absorbances of isomolar solutions of a mixture of H2MR?
and HMR, a solution of H2MR?, and a solution of HMR, respectively. Likewise, Eq. 4 can
be written in terms of absorbances as follows:
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Ka3 ¼
A0 A0a
A0b A0
aHþ cMR
cHMR
ð7Þ
where A’, A0a and A0b are the absorbances of isomolar solutions of a mixture of HMR and
MR-, a solution of HMR, and a solution of MR-, respectively. The overlapping in the
equilibria of Eqs. 1 and 2 make it impossible to determine directly the values of Ab and A0a :
This difficulty can be solved, using a linear graphical method after rearranging Eqs. 6 and
7 to evaluate Ka2 and Ka3 as in the following forms:
aHþ cHMR
¼ K a2 Ab Ka2 A
ð8Þ
ðA Aa Þ
cH2 MRþ
cHMR
0
0
0
þ A0a
A ¼ Ka3 ðAb A Þ
ð9Þ
aHþ cMR
A plot of A against the LHS of Eq. 8 gives a slope equal to –Ka2 and intercept equal to
Ka2Ab. Likewise in Eq. 9, a plot of ðA0b A0 ÞðcHMR =aHþ cMR Þ against A0 gives a slope
equal to Ka3 and an intercept equal to A0a . Figure 6 is a plot according to Eq. 8 at
k = 500 nm that contains typical data for evaluating Ka2. After making three sets of data to
evaluate Ka2 by Eq. 8, the average value of Ka2 in water at 25 °C is (1.02 ± 0.03) 9 10-2
moldm-3, with a pKa2 value of 1.99 ± 0.01. The value of Ka3 was obtained from a plot of
Eq. 9 at three different wavelengths. Figure 7 contains a typical data set at k = 450 nm.
The average value of Ka3 in water at 25 °C was calculated to be (4.94 ± 0.02) 9 10-5
moldm-3, with a pKa3 value of 4.31 ± 0.01. The ratio Ka2/Ka3 as calculated from our
results is 206, and seems to be consistent with the assumption that the two equilibria given
in Eqs. 1 and 2 are overlapping equilibria [4].
3.3 Distribution Diagram of m-MR Species
Figure 1 shows three acid–base equilibria, with acid dissociation constants Ka1, Ka2, and
Ka3. The constants Ka2 and Ka3 belong to overlapping equilibria in view of their relative
magnitude compared with that of Ka1. In this study we considered only the distribution of
the species participating in these overlapping equilibria. These species, as in Fig. 1, are the
-0.0004
-0.0005
(A-Aa )(aH+.γHMR / γH2MR+ )
Fig. 6 Plot of Eq. 8 at 500 nm.
The absorbance of the diprotic
form (Aa) is 0.677 at
3.7 moldm-3 H2SO4 and
1.537 9 10-5 moldm-3
m-methyl red
-0.0006
-0.0007
-0.0008
-0.0009
-0.0010
0.44
0.45
0.46
0.47
0.48
0.49
A
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Fig. 7 Plot of Eq. 9 at 450 nm.
The absorbance of the basic form
(A0b ) is 0.363 at pH = 10.5 and
1.537 9 10-5 moldm-3
m-methyl red
0.35
A'
0.30
0.25
0.20
0.15
500 1000 1500 2000 2500 3000 3500 4000 4500 5000
(A' b -A)(γHMR /a H+.γMR- )
basic form (structure A), the monoprotic form (structure B), and the diprotic form (the
tautomeric mixture of structure C and the resonance hybrid of structures D and E). The
expressions of the fraction, Fi, of the species participating in the overlapping equilibria
were written depending on the thermodynamic expressions of Ka2 and Ka3, and are given as
follows:
FMR ¼ Ka2 Ka3 = D
þ
ð10Þ
FHMR ¼ Ka2 ½H cHþ ðcMR = cHMR Þ = D
ð11Þ
FH2 MRþ ¼ ½Hþ 2 c2Hþ ðcMR = cH2 MRþ Þ = D
ð12Þ
where the denominator, (D) is defined by Eq. 13:
D ¼ Ka2 Ka3 þ Ka2 ½Hþ cHþ ðcMR =cHMR Þ þ ½Hþ 2 c2Hþ ðcMR =cH2 MRþ Þ
ð13Þ
The calculation of Fi using Eqs. 10–12 was based on some considerations: (a) Ionic
strength, I, does not depend on the contribution of m-MR due to its very low concentration,
and its value was calculated directly from the pH of the solution. (b) The activity coefficients cMR and cH2 MRþ of the singly charged ions were assumed equal. (c) The product
½Hþ cHþ , was calculated from the pH of the solution, aHþ ¼ 10pH , which is the activity of
the H? species, aHþ . (d) The activity coefficients cMR and cH2 MRþ were calculated
pffiffi
according to the Debye–Hückel limiting law: log10 ci ¼ 0:512 I , except in the pH range
0 pH 2; the activity coefficient cH2 MRþ was calculated according to Eq. 5. (e) In the pH
range 0–9, there is no contribution from the triprotic form (structure F in Fig. 1) to the total
concentration of m-MR, since Ka1 of m-MR is expected to be much larger than Ka2 or Ka3.
Figure 8 shows the fractions given by Eqs. 10–12, which was established by using 35
different values of pH in the range 0–9. The figure indicates that the fraction of monoprotic
form (curve # 2) is always less than one at any pH value. The maximum fraction of the
monoprotic form corresponds to a pH of 3.14, at which the distribution curve of the
monoprotic form appears to be symmetrical. The fraction of the basic and the diprotic
forms are nearly equal at this pH, which is the isoelectric point of m-MR [4]. The fractions
obtained at pH = 3.14 are 0.053 for the basic form, 0.883 for the monoprotic form and
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Fig. 8 The distribution diagram
of m-methyl red species in water
at 25 °C
1.0
1
2
3
Fraction
0.8
Curve #
0.6
Form
-----------------------1
2
3
0.4
diprotic
monoprotic
basic
0.2
0.0
0
2
4
6
8
10
pH
0.064 for the diprotic form. It was found that m-MR is completely in the basic form at pH
values [7, and completely the diprotic form at pH = 0. It was also found that the diprotic–
monoprotic equilibrium dominates in the pH range 0 \ pH \ 3, while the fraction of the
diprotic form is nil in the monoprotic-basic equilibrium in the pH range 4.25 B pH B 7.
Acknowledgments This work was supported by Petra University (Project 2009/1/7). The experimental
measurements were done at University of Jordan.
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