Spectroscopic studies on gallic acid and its azo derivatives and their iron(III) complexes

Article abstract of DOI:10.1016/j.saa.2013.10.054

Azo gallic derivatives and their iron(III) complexes were synthesized and characterized. The stereochemistry and the mode of bonding of the complexes were achieved based on elemental analysis, UV-Vis and IR. The thermal behaviors of the complexes were studied. The effect of pH on the electronic absorption spectra of gallic acid and its azo derivatives are discussed. Different spectroscopic methods (molar ratio, straight line method, continuous variation, slope ratio and successive method) are applied for determination of stoichiometry and pK values for the complex formation of gallic acid with iron(III) in aqueous media. Iron(III) complexes of gallic acid is formed with different ratio: 1:1, 1:2, 1:3 and 1:4 (M:L).

Full text of DOI:10.1016/j.saa.2013.10.054

Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 120 (2014) 505–511
Contents lists available at ScienceDirect
Spectrochimica Acta Part A: Molecular and
Biomolecular Spectroscopy
journal homepage: www.elsevier.com/locate/saa
Spectroscopic studies on gallic acid and its azo derivatives and their
iron(III) complexes
a
c
Mamdouh S. Masoud a, Alaa E. Ali b, , Sawsan S. Haggag , Nessma M. Nasr
⇑
a Chemistry Department, Faculty of Science, Alexandria University, Alexandria, Egypt
b Chemistry Department, Faculty of Science, Damanhour University, Damanhour, Egypt
c Faculty of Pharmacy, Alexandria University, Alexandria, Egypt
h i g h l i g h t s
g r a p h i c a l a b s t r a c t
Effect of pH on the electronic absorption spectra of 4 ꢁ 10ꢂ5 M of gallic acid.
ꢀ
The effect of pH on the absorption
spectra of gallic acid and its azo
derivatives is studied.
ꢀ
ꢀ
The dissociation constants of the azo
ligands are evaluated.
Different spectroscopic methods are
applied for determination of pK
values of gallic acid–iron(III) complex.
a r t i c l e i n f o
a b s t r a c t
Article history:
Received 7 May 2013
Received in revised form 8 October 2013
Accepted 10 October 2013
Available online 23 October 2013
Azo gallic derivatives and their iron(III) complexes were synthesized and characterized. The stereochem-
istry and the mode of bonding of the complexes were achieved based on elemental analysis, UV–Vis and
IR. The thermal behaviors of the complexes were studied. The effect of pH on the electronic absorption
spectra of gallic acid and its azo derivatives are discussed. Different spectroscopic methods (molar ratio,
straight line method, continuous variation, slope ratio and successive method) are applied for determina-
tion of stoichiometry and pK values for the complex formation of gallic acid with iron(III) in aqueous
media. Iron(III) complexes of gallic acid is formed with different ratio: 1:1, 1:2, 1:3 and 1:4 (M:L).
Ó 2013 Elsevier B.V. All rights reserved.
Keywords:
Gallic acid
Dissociation constants
Spectrophotometric methods
Synthesis
Complexes
DTA
Introduction
the complexation starts from pH = 3 and continuous to pH = 9.
The degree of chelation increases with increasing the pH. Iron is
Gallic acid (3,4,5trihydroxy benzoic acid) is a strong chelating
agent and forms complexes of high stability with iron [1], where
attached to gallic acid through two adjacent OH [2]. The kinetics
and mechanisms of the reactions of a number of pyrogallol-based
ligands with iron(III) have been investigated in aqueous solution
at 25 °C and an ionic strength 0:5MNaClOꢂ4 , where 1:1 complex
is formed initially, when the metal reacts with ligand subsequently
⇑
Corresponding author. Tel.: +20 1115799866; fax: +20 4553293011.
E-mail address: dralaae@yahoo.com (A.E. Ali).
1386-1425/$ - see front matter Ó 2013 Elsevier B.V. All rights reserved.
http://dx.doi.org/10.1016/j.saa.2013.10.054

506
M.S. Masoud et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 120 (2014) 505–511
through an electron transfer reaction, with an evidence for the for-
mation of 1:2 at higher pH-values [3]. When excess of iron(III) is
added to an acidic solution of either gallic-acid or its methyl ester,
a dark blue color appears initially and then rapidly fades, ascribed
to the formation of 1:1 complex, in which the metal ion is coordi-
nated to two adjacent hydroxyl groups of the ligand, while the
third hydroxyl group remains protonated due to its high pKa [4].
This is subsequently decomposed to Fe(II) and the corresponding
semiquinine and then reacts rapidly with another Fe(III) species
to form the quinine. Iron(III)–gallic acid complex shows, a maxima
at 415 and 690 nm. The quinine absorbs at 380 nm with a shoulder
at 445 nm .Gallic acid forms (1:3) complex with iron in the pH rang
4–6 [5]. Kinetic study on the complexation of gallic acid with fer-
rous sulfate was performed using UV–Vis absorption spectroscopy.
The stoichiometry composition of the formed complex is 1:1 [6]
and the absorption band of gallic acid at 263 nm undergoes a
bathochromic shift of 34 nm with the addition of ferrous ions,
due to the formation of gallic acid–Fe(II) complex with the
formation of a new absorption band at 570 nm, which is assigned
to the complex formation. In the present manuscript, it is aimed
to give more information about the chemistry of gallic acid and
three azo compounds of groups of different electronic characters
and some complexes especially Fe(III). The scope of the studies is
itemized as follows: (i) studies on the free ligands (effect of pH
on the electronic absorption spectra and evaluation of pK values),
(ii) studies on the transition metal complexes (interaction of Fe(III)
with gallic acid and evaluation of the possible species exist and
their stabilities, (iii) synthesis and characterization of some solid
complexes by elemental analysis, spectroscopy and magnetic
moments) and thermal analysis (DTA) of some ligands and their
complexes.
Results and discussion
Studies on the free ligands
Effect of pH on the electronic absorption spectra of gallic acid
Fig. 2 is taken as a representative example for the effect of pH
on the electronic absorption spectra of 4 ꢁ 10ꢂ5 solution of gallic
acid, which showed three well defined bands with kmax at 230,
280 and 425 nm. The first two bands are due to
p–
pꢃ transition,
while the last one is due to n–pꢃ transition. The acidity constants
of gallic acid are pK1 = 4.9, pK2 = 7.5, pK3 = 10.3 [1]. The ionization
mechanism of gallic acid proceeds in two steps. The first acidity
constant is associated with ionization of carboxylic group (1B).
The second one is likely to be associated with the OH group in
the 4-positions (1C), because this allows the delocalization of the
negative charge and the presence of two ortho-positioned OH
groups. This stabilizes the negative charge by intramolecular
hydrogen bonds (1E), which lies lower in energy than structure
(1F). In acidic medium [9], below pH = 3.4, gallic acid exists in its
neutral form with its absorbance maximum at 269 nm. By increas-
ing the pH gradually, the anion form of gallic acid exists, which
shifts the absorbance maximum to 257 nm. However, from pH
4.5 to 7.5, the shift of maximum is small. Further increase of
pH > 7, leads to the formation of a new additional peak with a max-
imum at 295 nm. At basic pH, gallic acid undergoes fast autooxida-
tion, which leads to colorization of the solution. The peak with a
maximum at 225 nm at pH = 2 gradually shifts to 231 nm at
pH = 7. However, H4L2 shows two bands at 225 and 380 nm. Mean-
while, H5L3 gave four bands at 225, 260, 350 and 450 nm. Finally,
H5L4 compound gave three bands at 225, 280 and 450 nm with
one isosbestic point at kmax ¼ 346 nm. Remarkable features are gi-
ven, where the compounds H4L1, H4L2 and H5L3 possess no isobes-
tic point, probably due to the overlapping of absorbing species. The
shorter wavelength region is due to the electronic transitions (up
Experimental
to ꢄ250 nm) mainly of the
p–
pꢃ type, while the longer wavelength
The UV–VIS spectra were measured with a Perkin Elmer Lamb-
da 4B spectrophotometer. Gallic acid and ferric chloride com-
pounds were purchased from Sigma and Aldrich. 0.01 M stock
solution of Fe(III) was prepared by dissolving the required weight
in distilled water. The solution was acidified by the addition of
hydrochloric acid to keep the Fe(III) soluble in water and the exact
concentration was achieved complexometrically using salicylic
acid as indicator. 0.01 M stock solution of gallic acid was prepared.
For every new experiment fresh solution was prepared because
gallic acid in aqueous solution does not persist for a long time. It
oxidizes to quinine. Universal buffer solution was prepared by tak-
ing 0.04 M each of H3BO3, H3PO4 and CH3COOH acids and adding
the required volume of 0.2 M NaOH to give the desired pH [7].
The pH was checked by using a Jenway 3015 pH-meter, previously
calibrated with standard buffer solutions of pH 4.00, 7.02 and 9.18.
0.10 M KCl solutions were prepared and used to adjust the ionic
strength of the solutions. 30 mmol of the ligand was dissolved in
alcohol mixed with 15 m-mole of iron(III) chloride, which was dis-
solved in water. The reaction mixture was refluxed for one hour,
and then left over-night, where the precipitated complexes were
separated by filtration and washed by H2O and dried in a desicca-
tor over anhydrous CaCl2. Azo compounds were prepared in a sim-
ilar way by the usual diazotization process [8]. The required
substituted amines (0.1 mol) were dissolved in (0.2 mol) HCl and
25 ml distilled water. The hydrochloride compounds were diazo-
tized below 5 °C with a solution of NaNO2 (0.1 mol) and 20 ml dis-
tilled water. The diazonium chloride was coupled with an alkaline
solution of gallic acid (0.1 mol/30 ml distilled water). The crude
dyes were filtered off and crystallized, then dried in a vacuum des-
iccator over P4O10. Fig. 1 illustrated the structures of gallic acid and
its azo derivatives.
side (>225 nm) can be argued to the electronic transitions mainly
of n–pꢃ type. The introduction of azo group affects the mode of ion-
ization to some extent. In general, the azo compounds undergo a
regular bathochromic shift on increasing the pH as a result of pro-
ton elimination. The aryl azo substituents exert an acid strength ef-
fect on the ligands (as will be seen later from the pK values).
Evaluation of pK-values
Different spectrophotometric methods deduced from As-pH
relations were applied, half height [10], modified limiting absorp-
tion [11] and collector methods [12]. The data are collected in
Table 1. Most of compounds under investigation give slopes = 1,
2 indicating one and two protons are ionized. In case of gallic acid,
the slope = 3 leads to be completely tautomerized.
Distribution diagrams for gallic acid and its azo derivatives at different
pH
From the distribution diagram plots between the fractions of an
acid species of the ligands plotted against pH, we can found that
the variation of these species are due to the acid dissociation is
shifting as pH changes [13].
For a diprotic acid : H2L ꢀ H2L þ HLꢂ þ Lꢂ2
For a triprotic acid : H3L ꢀ H3L þ H2Lꢂ þ HLꢂ2 þ Lꢂ3
For H4L1 compound only three ionizable protons are traced, if
the pH is 2.0 or less where the acid exists in the form of H3L. It
seems that the H2Lꢂ predominates in the pH range 4.5–6.5. How-
ever, on increasing the pH from 8.0–10.0, the HLꢂ2 species

M.S. Masoud et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 120 (2014) 505–511
507
HO
X
COOH
O
HO
OH
N
N
HO
Gallic acid (H4 L1)
HO
OH
OH
Compounds
X
H4 L2
H5 L3
OH
H5L4
H
COOH
Fig. 1. Structures of gallic acid and its azo derivatives.
similar trend, where only two protons are traced. However, in case
of H5L3, for pH = 4.0, the compound exists in the form of H2L, on
increasing the pH from 6.5 to 8.0, so the predominant species is
HLꢂ, and if the pH > 10.5, the Lꢂ2 exists. Meanwhile, in case of
H5L4, if the pH = 6.0, the predominant species is H2L, on increasing
the pH from 8.0 to 10.0, the predominant species is HLꢂ and if the
pH is greater than 12.0, the compound exists in the form of Lꢂ2. In
case of H5L3, the values of the slopes are 1.7 and 1.3. This suggests
that the mechanism of AOH ionization is of complicated equilibria
accompanied by the existence of an associated structure through
intermolecular hydrogen bonding through two hydroxy groups
and or intramolecular hydrogen bonding between the OH and
N@N groups. However, in case of H5L4, the values of the slopes
are 1.5 and 0.6, which may support the presence of such intramo-
lecular hydrogen bonding between the azo group and the
OACOOH.
Studies on the Fe(III)-complexes
Interaction of Fe(III)–gallic acid and evaluation of the possible species
exist and their stabilities
Molar-ratio method [14]. Fig. 3 represents the electronic absorption
spectra of Fe(III)–gallic acid complex .The concentration of Fe(III) is
kept constant at 4 ꢁ 10ꢂ4 M, while that of ligand is varied between
1 ꢁ 10ꢂ4–1.1 ꢁ 10ꢂ3 M. The stoichiometry of the reaction is ob-
tained by plotting the absorbance of the resulting solution against
the molar ratio of the variable concentration of the ligand with re-
spect to that of Fe(III). It seems that the complex reaction by this
procedure produced through 1:2 M ratios, Fig. 4.
Fig. 2. Effect of pH on the electronic absorption spectra of 4 ꢁ 10ꢂ5 M of H4L1.
Straight line method [15]. The following equation is applied:
ꢀ
ꢁ
1
v
CD
Kc As
1
D
Kc
¼
ꢅ
ꢂ
ð1Þ
predominates, on further increasing the pH up to 12.0 or higher,
the predominant species is Lꢂ3, while in case of H4L2, if the pH is
less than 3.2, HL exists, on further increasing the pH up to 5.0 or
higher, the L species exists. The compounds H5L3 and H5L4 are of
n
By plotting 1/vn against 1/As, where (n = 1, 2, 3), a straight line is
obtained at the proper stoichiometry of the reaction, Fig. 5. It is
Table 1
Dissociation constants of the organic compounds spectrophotometrically.
Compound
Half height
Modified limiting method
Collecter
Average pK
pK1
pK1
pK2
pK3
pK1
pK2
pK3
pK1
pK2
pK3
pK2
pK3
H4L1
H4L2
H5L3
H5L4
4.9
3.3
4.2
7.7
7.5
–
9.2
10.5
10.3
–
–
4.3(2)
7.1(3.2)
–
9.1(2.0)
10.1(0.6)
9.6(0.8)
4.6
3.3
4.4
7.6
6.9
–
9.3
10.3
10.2
–
–
4.6 0.3
3.2 0.15
4.5 0.3
7.6 0.05
7.1 0.30
–
9.2 0.1
10.2 0.2
10.1 0.37
3.0(1.2)
4.9(1.7)
7.7(1.5)
–
–
–
–
–
–
–
–
The value between brackets denotes the slope from the linear plots.

508
M.S. Masoud et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 120 (2014) 505–511
0.4
0.35
[Fe (III)] = 4x10-4
M
0.3
0.25
0.2
0.15
0.1
a
b
c
= 1x10-4
M
= 2x10-4 M
= 4x10-4 M
d = 1.1x10-3M
0.05
0
0.5
1
1.5
2
1+Log V
Fig. 5. Straight line logarithmic plot for iron(III)–gallic acid reaction at
kmax = 600 nm.
1
0.9
0.8
0.7
0.6
0.5
0.4
0.3
0.2
0.1
Fig. 3. Electronic absorption spectra of iron(III)–gallic acid in aqueous medium.
0.1
0.2
0.3
0.4
0.5
0.6
[Fe] / [Fe]+[L]
Fig. 6. Continuous variation method for iron(III)–gallic acid reaction at
kmax = 380 nm.
solutions by taking two sets of experiments, Fig. 7. (i) The concen-
tration of the ligand is kept constant in excess at (2 ꢁ 10ꢂ3 M),
while that of iron(III) is varied between (1 ꢁ 10ꢂ4–6 ꢁ 10ꢂ4 M),
(ii) the concentration of the iron(III) is kept constant in excess at
(2 ꢁ 10ꢂ3 M), while that of the ligand is varied between
(1 ꢁ 10ꢂ4–6 ꢁ 10ꢂ4 M). The absorbance of the two sets of solutions
was measured at kmax ¼ 400 nm. The limiting absorption method
[18] is based on the same mode of thinking given before for the
slope ratio method. On plotting the logarithmic values of the
absorbance versus the logarithmic values of the concentration of
the variable, Fig. 8. The stoichiometry of the complex reaction is
calculated based on the ratio of the slopes of the two straight lines
obtained. So, the traced complex is with the stoichiometry 1:1. It is
observed from, Fig. 7b, in which, a straight line is obtained passed
through the origin, indicating that the complexes formed obey
Fig. 4. Molar ratio plot for iron(III)–gallic acid reaction at kmax = 600 nm.
apparent that n = 1 is the predominant to assign the presence of a
complex with a stoichiometry 1:1. If the complex is strongly
dissociated:
Kc
D
1
v
CD
Kc
Beer’s law. The high
e value 762.4 is an advantage for applications,
ꢅ
>> 1 ) logAs ¼ log
þ n log
v
ð2Þ
n
where very low concentrations in the pH-scale could be obtained.
Plotting logAs versus logv, a straight line is obtained from which n
CD
c
and logðK Þ values are evaluated from the slope and the intercept,
The successive method [19]. This depends on the changing of both
the concentration of the ligand and the metal ion. The following
equation is applied, where As is the absorbance due to [ML], q is
the molar extinction coefficient of the formed complex.
respectively. So, the stoichiometry and the equilibrium constants
are calculated. The pK value is found to be 8.3.
The continuous variation method [16,17]. In this reaction, the total
molar concentration is kept constant at 2 ꢁ 10ꢂ3 M. The measure-
ment was done at kmax ¼ 386 nm, Fig. 6. The extrapolation of such
plot, gives 0.2, 0.35 and 0.45 mol fraction of Fe(III) to suggest the
presence of complexes of 1:4, 1:2 and 1:1, respectively.
ꢂ
CMCL
As
1
q 2
1
2
ꢂ
)
¼
þ
ðCM þ CL Þ
By plotting the left-hand side of the above equation versus
ꢂ
ðCM þ CL Þ, a straight line is obtained with a slope equals to 1/
e
The slope ratio and the limiting logarithmic methods [14]. These
and an intercept of 1/q
e
. The results are diagrammed in Fig. 9. The
methods are used to trace the complex formation in dilute
values of molecular extinction (
e
) and the formation constant (q)

M.S. Masoud et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 120 (2014) 505–511
509
Table 2
pK values and the stoichiometry for iron(III)–gallic acid complex by using different
spectroscopic methods.
Methods
Stoichiometry of iron(III)–gallic acid
complexes
pK
Molar ratio
Straight line
Continuous
variation
1:2
1:1
1:1
1:2
1:4
1:1
–
8.30
3.63
6.43
8.40
–
Slope ratio
Limiting logarithmic 1:1
–
Fig. 7. Slope ratio method for iron(III)–gallic acid reaction at kmax = 600 nm. (a) [Fe
(III)] = 2 ꢁ 10ꢂ3, [L] is variable, (b) [L] = 2 ꢁ 10ꢂ3, [Fe (III)] is variable.
gand gave a broad band at 3312 cmꢂ1 assigned to mOH [21–23],
where some changes occur on complexation. The data in such re-
gion are related either to that the AOH group is affected on coordi-
nation or due to the existence of water molecule. Hydrogen bonded
structures are expected due to interamolecular H-bonding be-
tween the azo group and the o-carboxy [24,25], H-bonding of the
type OAHꢅ ꢅ ꢅN between the AOH or ACOOH substituent and the
N@N group. The mC@O [26] of H4L1, H4L2, H5L3 and H5L4 ligands
are at 1666, 1670, 1681 and 1710 cmꢂ1, respectively, which is
shifted in case of [Fe (H4L3) (H5L3) Cl2ꢅ5H2O] and [Fe (H2L4)ꢅ6H2O]
complexes, suggesting that the carbonyl group is strongly affected
on coordination with transition metal ions either by direct interac-
tion between the lone pair of electrons of the carbonyl group with
metal ion or through tautomerization. Also, dC@O vibration band
reached to the same conclusion. H4L2, H5L3 and H5L4 compounds
gave bands at 1493, 1452 and 1491 cmꢂ1, respectively, identified
to azo vibration [27,28]. The shift of this band in all complexes
indicated that the azo group is involved in coordination. The pres-
Fig. 8. Limiting logarithmic method for iron(III)–gallic acid reaction at kmax = 400
nm. (a) [Fe+3] = 2 ꢁ 10ꢂ3, [L] is variable, (b) [L] = 2 ꢁ 10ꢂ3, [Fe (III)] is variable.
ence of cc@o, mc@o [29] modes of vibrations in the metal complexes,
are 11280 and 72900, respectively. The pK values obtained from dif-
ferent spectrophotometric methods are summarized in Table 2.
suggest that the oxygen atom is also a center of coordination. New
bands are appeared in all azo gallic acid complexes in the
frequency ranges 437–592 cmꢂ1 and 503–646 cmꢂ1, which are
Synthesis and characterization of some solid complexes by elemental
analysis, spectroscopy and magnetic moments
The infrared spectrum of the metal complexes, Table 3, com-
pared with those of the frequencies of coordinated functional
attributed to mMAN and
mMAO, respectively. From the analytical
data, electronic spectra and magnetic moment results, Table 4,
iron(III)-complexes are of Oh geometry [30,31]. The following
chemical equations suggested the formation of the complexes:
groups (e.g. mOH
depending on the strength of
the metal ion and -electrons of the functional groups. Some of
,
mC@O
,
mN@N) are affected with different degrees
FeCl3 þ 3H4L1 ! FeðH4L1Þ ꢅ Cl3 þ 4H2O;
p
-interaction occurring between
3
p
FeCl3 þ H4L2 ! FeðH2L2Þ ꢅ Cl þ 2HCl
the carbonyl bands are disappeared, and others are slightly shifted,
probably due to tautomerization to enol form followed by com-
plexation through deprotonation. The IR of gallic acid gave three
FeCl3 þ ðH5L3Þ ! FeðH4L3ÞðH5L3Þ ꢅ Cl2 þ 5H2O þ HCl
bands at 3494, 3279 and 2669 cmꢂ1, which are assigned to mOH
.
2
In case of [Fe (H4L1)3Cl3ꢅ4H2O] complex, the three bands are shifted
at 3410, 3341 and 2679 cmꢂ1, where phenolic OH groups are par-
ticipating in the formation of the complexes [20]. However, In case
of H4L2 and H5L4, three bands are appeared at 3452, 3218,
3033 cmꢂ1 and 3367, 3203, 3072 cmꢂ1, respectively, while H5L3 li-
FeCl3 þ H5L4 ! FeðH2L4Þ þ 4H2O þ 3HCl
DTA of some ligands and their Fe(III) complexes
Differential thermal analysis (DTA) was carried out, where the
rate of heating was 10 °C/min. The order of chemical reactions
(n) was calculated via the peak symmetry method [32]. The value
8
7
6
5
4
3
2
1
of the decomposed substance fraction, (am), at the moment of max-
imum development of reaction with T = Tm being determined from
1
1ꢂn
the relation: ð1 ꢂ amÞ ¼ n [33]. The values of collision factor, Z,
can be obtained from the following equation:
!
!
E
RTm
E
RT2m
kTm
h
D
S#
R
z ¼
/ exp
¼
exp
where
stant, / rate of heating (Ksꢂ1), k the Boltzmann constant and h
the Planck’s constant. The heat of transformation,
H#, can be cal-
culated from the DTA curves [34]. In general, the change in enthalpy
H#) for any phase transformation taking place at any peak
D
S# represents the entropy of activation, R molar gas con-
0.2
0.4
0.6
0.8
1
[Fe]+[L] x104
D
Fig. 9. Successive method for iron(III)–gallic acid reaction at kmax = 271 nm.
(D

510
M.S. Masoud et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 120 (2014) 505–511
Table 3
Fundamental infrared bands for organic ligands and their iron(III)-complexes.
Compound
H4L1
mOH
mC@O
mC@C and mCAC mN@N mCAOH and dOH
d C–OH
dC@O and c CO cOAH H. bonding OH in COOH mMAO mMAN
3494
3279
2669
3410
3341
2679
3452
3218
3033
–
1666
1612
1540
1426
1610
1538
1439
1595
1445
–
1318
1384
1266
1098
634
655
690
694
865
794
733
732
750
745
–
–
Fe(H4L1)3Cl3ꢅ4H2O
H4L2
1666
1670
–
–
1303
1383
1263
1097
875
796
471
–
–
1493
1510
1306
1410
1244
1068
835
840
–
Fe(H2L2) Clꢅ2H2O
1579
–
–
1257
499
603
3355
3053
3312
H5L3
1681
1693
1710
1606
1518
1602
1488
1601
1452
1452
1445
1491
1368
1344
1108
1108
663
669
688
818
833
879
753
755
758
–
–
Fe(H4L3) (H5L3) Cl2ꢅ5H2O
3212
530
–
607
–
H5L4
3367
3203
3072
3377
3188
1307
1379
1214
1116
Fe(H2L4)ꢅ6H2O
1730
1587
–
1481
1309
1379
1230
1147
684
887
761
445
530
Table 4
Analytical data, electronic absorption spectra and room temperature magnetic moment values.
Complex
% Calculated/(found)
k(nm)
leff
Electronic transition
M
C
H
N
–
X
Fe(H4L1)3Cl3ꢅ4H2O
Fe(H2L2) Clꢅ2H2O
Fe(H4L3) (H5L3) Cl2ꢅ5H2O
Fe(H2L4)ꢅ6H2O
7.3
33.3
(33.7)
39.1
(39.4)
39.2
2.9
(3.5)
3
(3.4)
3.6
13.9
(13.5)
8.7
(8.4)
9.0
300,390
450,510
300,389,675
5.9
5.9
5.9
5.9
CT (t2g
CT ( ? e.g.)
CT (t2g
CT ( ? e.g.)
CT (t2g
CT ( ? e.g.)
CT (t2g
CT ( ? e.g.)
?
pꢃ
pꢃ
pꢃ
pꢃ
)
(6.8)
13.8
(13.7)
7.0
p
7.0
(7.4)
7
?
)
p
386,314,350,460,510
400,485,510,565
?
)
(7.2)
11.6
(11.0)
(39.8)
35.3
(35.8)
(3.7)
3.9
(4.4)
(7.3)
5.8
(6.1)
(9.5)
–
p
?
)
p
Table 5
DTA ligands and their Fe(III) complexes.
Compound
H4L1
Type
Tm (°K)
D
E (kJ/molꢂ1
)
n
am
D
S– (kJ Kꢂ1molꢂ1
)
D
H– (kJ molꢂ1
)
103Z (sꢂ1
)
Endo
Endo
Exo
Endo
Exo
Exo
Exo
Exo
Exo
Exo
Exo
Exo
379.5
540.6
721.8
513.5
623.0
682.2
383.0
833.3
873.0
373.0
747.7
793.0
109.1
1.4
1.2
0.5
1.2
0.4
1.6
1.4
0.4
2.3
1.6
0.7
1.4
0.6
0.6
0.7
0.6
0.8
0.5
0.6
0.8
0.5
0.5
0.7
0.6
ꢂ0.29
ꢂ0.29
ꢂ0.31
ꢂ0.30
ꢂ0.29
ꢂ0.29
ꢂ0.29
ꢂ0.32
ꢂ0.31
ꢂ0.29
ꢂ0.27
ꢂ0.29
ꢂ0.32
ꢂ0.29
ꢂ0.27
ꢂ0.29
ꢂ0.32
ꢂ0.29
ꢂ0.27
ꢂ0.30
ꢂ0.30
ꢂ0.29
ꢂ0.29
ꢂ0.32
ꢂ117.2
ꢂ156.0
ꢂ225.0
ꢂ154.0
ꢂ181.0
ꢂ199.0
ꢂ111.0
ꢂ270.0
ꢂ267.0
ꢂ114.0
ꢂ219.0
ꢂ249.2
0.03
0.09
0.01
0.25
0.08
0.08
0.06
0.01
0.02
0.01
0.07
0.01
439.8
50.74
107.9
436.6
468.9
203.9
14.37
143.0
25.73
484.8
41.6
Fe(H4L1)3Cl3ꢅ4H2O
H4L2
Fe(H2L2) Clꢅ2H2O
H5L3
Exo
Exo
Exo
Exo
Exo
Exo
Exo
Exo
Exo
Endo
Exo
373.0
658.9
888.0
903.0
572.0
603.0
373.0
705.0
853.7
588.3
643.0
196.0
4230.0
1605.0
24.71
196.0
2006.0
20.5
0.5
1.5
1.5
1.6
0.7
0.5
0.5
0.7
0.7
0.7
0.6
0.7
0.5
0.6
0.5
0.7
0.5
0.5
0.7
0.7
0.7
0.6
ꢂ108.0
ꢂ180.0
ꢂ254.0
ꢂ290.0
ꢂ169.3
ꢂ167.3
ꢂ114.7
ꢂ216.4
ꢂ253.6
ꢂ169.0
ꢂ207.0
0.06
0.77
0.21
0.01
0.04
0.40
0.01
0.01
0.05
0.12
0.01
Fe(H4L3) (H5L3) Cl2ꢅ5H2O H5L4
79.2
383.2
597.2
10.41
Fe(H2L4)ꢅ6H2O
temperature Tm can be given by the following equation:
D
S#
=
D
H#/Tm.
D
S#, values for all complexes, are nearly of the same magnitude
The (DTA) for H4L1, H4L2, H5L3 and H5L4 compounds and their
Fe(III)-complexes are collected in Table 5. The change of entropy,
and lie within the range of (ꢂ0.27 to ꢂ0.32) kJ Kꢂ1 moleꢂ1. So, the
transition states are more ordered, the fraction appeared in the

M.S. Masoud et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 120 (2014) 505–511
511
[4] H. Kipton, J. Powell, M.C. Taylon, J. Aust. Chem. 35 (1982) 739.
ccalculated order of the thermal reaction, n, confirmed that the
reactions proceeded in the complicated mechanisms. The calculated
values of the collision parameters, Z, showed a direct relation to Ea
[35]. The values of the decomposed substance fraction (am), at max-
imum development of the reaction were calculated. It is nearly with
the same magnitude and lies within the range 0.5–0.8. Based on
least square calculations, the lnDT versus 1000/T plots for all com-
plexes gave straight lines from which, the activation energies were
calculated [36,37].
[5] M.S. Masoud, S.S. Hagagg, A.E. Ali, N.M. Nasr, J. Mol. Struct. 1014 (2012) 17.
[6] L.L. Lu, Y.H. Li, X.Y. Lu, Spectrochim. Acta 73A (2009) 829.
[7] T.S. Britton, Hydrogen Ions, Chapman. Hall Ltd, London, 1955.
[8] G. Schwarzenbach, Complexometric Titration (H. Irving, Trans), Methuen Co.,
London, 1957.
[9] K. Polewski, S. Kniat, D. Slawinska, Curr. Top Biophys. 26 (2002) 217.
[10] I.M. Issa, R.M. Issa, F.I. Taha, M.A. Khattab, Egypt. J. Chem. 14 (1971) 59.
[11] A.A. Muck, M.B. Pravico, Anal. Chim. Acta 45 (1969) 534.
[12] J.C. Colleter, Z. Ann. Chim. 5 (1960) 415.
[13] A.A. Shoukry, M.M. Shoukry, Spectrochim. Acta 70A (2008) 686.
[14] A.E. Harvey, H. Manning, J. Am. Chem. Soc. 72 (1950) 4488.
[15] E. Asmus, J. Anal. Chem. 178 (1960) 104.
[16] A. Holme, F.J. Langmyhr, Anal. Chim. Acta 28 (1963) 501.
[17] P. Jop, Ann. Chim. 9 (1928) 113.
Conclusion
[18] F.W. Fowler, A.R. Katritzky, R.J. Rutherford, J. Chem. Soc. B (1971) 460.
[19] M.S. Masoud, T.M. Salem, M.M. Amer, J. Rev. Roumaine Chim. 27 (1982) 235.
[20] L. Wang, J. Fan, J. Nano-scale Res. Lett. 5 (2010) 81.
[21] F. Billes, I.M. Ziegler, P.J. Bombicz, Vib. Spectrosc. 43 (2007) 193.
[22] I.M. Ziegler, F. Billes, J. Mol. Struct. 618 (2002) 259.
[23] D. Slawins, K. Polewski, P.R. Wski, J.S. Ski, J. Int. Agrophys. 21 (2007) 199.
[24] M.M. Aly, N.I. Al-Shatti, Trans. Met. Chem. 23 (1998) 361.
[25] L.J. Bellamy, The infrared spectra of complex molecules, Wiley, New York,
1958.
Gallic acid and it azo derivatives gave stable complexes with ir-
on(III). These complexes were found to have different metal: ligand
ratios. The thermal behavior of these complexes proved that the
mechanism of their thermal reactions is not simple where fractions
appeared in the calculated orders of these reactions. The calculated
values of the collision numbers showed a direct relation to the
reactions’ activation energies.
[26] A. Kula, J. Therm. Anal. Calorim. 75 (2004) 79.
[27] M.S. Masoud, E.A. Khalil, A.M. Hindawy, A.E. Ali, E.F. Mohamed, Spectrochim.
Acta 60A (2004) 2807.
Appendix A. Supplementary material
[28] M.S. Masoud, A.E. Ali, R.H. Mohamed, A.A. Mostafa, Spectrochim. Acta 62A
(2005) 1114.
[29] D. Hadzi, J. Chem. Soc. (1956) 2725.
Supplementary data associated with this article can be found, in
the online version, at http://dx.doi.org/10.1016/j.saa.2013.10.054.
[30] D.W. Barnum, J. Inorg. Nucl. Chem. 22 (1961) 183.
[31] N.L. Allinger, J.C. Tai, A.M. Miller, J. Am. Chem. Soc. 88 (1966) 4495.
[32] E. Kissinger, Anal. Chem. 29 (1957) 1702.
[33] M.S. Masoud, T.S. Kasem, M.A. Shaker, A.E. Ali, J. Therm. Anal. Calorim. 84
(2006) 549.
References
[34] M.L. Dhar, O.J. Singh, Therm. Anal. 37 (1991) 259.
[35] E. Koch, Thermochim. Acta 94 (1985) 43.
[36] M.S. Masoud, A.A. Soayed, A.E. Ali, Spectrochim. Acta 60A (2004) 1907.
[37] R. Iordanova, E. Lefterova, I. Uzunov, Y. Dimitriev, D. Klissurski, J. Therm. Anal.
Calorim. 70 (2002) 393.
[1] A.S. Li, B. Bandy, S.S. Tsang, J. Free Rad. Res. 33 (2000) 551.
[2] A. Avdeev, S.R. Sofen, T.L. Bregante, K.N. Raymond, J. Am. Chem. Soc. 100 (1978)
5362.
[3] M.J. Hynes, M. O’Coinceanainn, J. Inorg. Biochem. 85 (2001) 131.