Studies on some salicylaldehyde Schiff base derivatives and their complexes with Cr(III), Mn(II), Fe(III), Ni(II) and Cu(II)

Article abstract of DOI:10.1016/j.saa.2006.09.013

The formation constants of some transition metal ions Cr(III), Mn(II), Fe(III), Ni(II) and Cu(II) binary complexes containing Schiff bases resulting from condensation of salicylaldehyde with aniline (I), 2-aminopyridine (II), 4-aminopyridine (III) and 2-aminopyrimidine (IV) were determined pH-metrically in ethanolic medium (80%, v/v). The formation constants were determined for all binary complexes. The important infrared (IR) spectral bands corresponding to the active groups in the four ligands and the solid complexes under investigation were studied. The solid complexes have been synthesized and studied by thermogravimetric analysis. The thermal dehydration and decomposition of these complexes were studied kinetically using the integral method applying the Coats-Redfern equation. It was found that the thermal decomposition of the complexes follow second order kinetics. The thermodynamic parameters of the decomposition are also reported. The electronic absorption spectra of the investigated ligands were carried out to determine the pK_a values spectrophotometrically.

Full text of DOI:10.1016/j.saa.2006.09.013

Spectrochimica Acta Part A 67 (2007) 950–957
Studies on some salicylaldehyde Schiff base derivatives and their
complexes with Cr(III), Mn(II), Fe(III), Ni(II) and Cu(II)
S.A. Abdel-Latifa,∗, H.B. Hassibb, Y.M. Issab
a Chemistry Department, Faculty of Science, Helwan University, Cairo, Egypt
b Chemistry Department, Faculty of Science, Cairo University, Giza, Egypt
Received 3 July 2006; received in revised form 11 September 2006; accepted 12 September 2006
Abstract
The formation constants of some transition metal ions Cr(III), Mn(II), Fe(III), Ni(II) and Cu(II) binary complexes containing Schiff bases
resulting from condensation of salicylaldehyde with aniline (I), 2-aminopyridine (II), 4-aminopyridine (III) and 2-aminopyrimidine (IV) were
determined pH-metrically in ethanolic medium (80%, v/v). The formation constants were determined for all binary complexes. The important
infrared (IR) spectral bands corresponding to the active groups in the four ligands and the solid complexes under investigation were studied. The
solid complexes have been synthesized and studied by thermogravimetric analysis. The thermal dehydration and decomposition of these complexes
were studied kinetically using the integral method applying the Coats–Redfern equation. It was found that the thermal decomposition of the
complexes follow second order kinetics. The thermodynamic parameters of the decomposition are also reported. The electronic absorption spectra
of the investigated ligands were carried out to determine the pKa values spectrophotometrically.
© 2006 Elsevier B.V. All rights reserved.
Keywords: Salicylaldehyde Schiff bases; Transition metal complexes; Absorption spectra; Thermal analysis; Potentiometry; Infrared
1. Introduction
compounds effective and stereospecific catalysts for oxidation,
reduction and hydrolysis and they show biological activity,
and other transformations of organic and inorganic chemistry
[25]. It is well known that some drugs have higher activity
when administered as metal complexes than as free ligands
[25].
The present study deals with the preparation of some Schiff
bases and their chelates with some transition metal ions, e.g.
Cr(III), Mn(II), Fe(III), Ni(II) and Cu(II). The solid complexes
have been synthesized and studied by elemental analyses, IR
and thermogravimetric analysis (TG). The thermal dehydration
and decomposition of these complexes were studied kineti-
cally using the integral method applying the Coats–Redfern
equation.
Schiff base ligands are considered “privileged ligands”
because they are easily prepared by the condensation between
aldehydes and imines. Stereogenic centers or other elements
of chirality (planes, axes) can be introduced in the synthetic
design. Schiff base ligands are able to coordinate many dif-
ferent metals [1–5], and to stabilize them in various oxidation
states. The Schiff base complexes have been used in catalytic
reactions [6] and as models for biological systems [7,8]. Many
copper complexes of Schiff bases were prepared [9–14]. It has
been reported that the structure of the substituent bonded to
the imino nitrogen affects the coordination geometry of the
complex [15]. During the past two decades, considerable atten-
tion has been paid to the chemistry of the metal complexes of
Schiff bases containing nitrogen and other donors [16–21]. This
may be attributed to their stability, biological activity [22] and
potential applications in many fields such as oxidation catal-
ysis [23], electrochemistry [24]. The complexes make these
2. Experimental
All chemicals used in the present investigation were of Analar
or the highest purity grade from BDH used as received without
further purification. The solvents used were methanol, ethanol,
isopropanol, n-hexane and chloroform, either of spectroscopic
pure from BDH or purified according to recommended proce-
dures [26].
∗
Corresponding author. Tel.: +20 100244590.
E-mail address: salatif 1@yahoo.com (S.A. Abdel-Latif).
1386-1425/$ – see front matter © 2006 Elsevier B.V. All rights reserved.
doi:10.1016/j.saa.2006.09.013

S.A. Abdel-Latif et al. / Spectrochimica Acta Part A 67 (2007) 950–957
951
2.1. Preparation of the Schiff bases I–IV
3. Apparatus and techniques
The Schiff bases were prepared by mixing equimolecular
amounts of salicylaldehyde and the appropriate amine in 10 ml
absolute ethanol in a round bottomed flask equipped with a
condenser. The mixture was brought into reflux for 4 h. The
products obtained after cooling were filtered off and crystallized
from absolute ethanol. The crystallized solids were dried under
vacuum and kept dry in a desiccator over anhydrous calcium
chloride. Melting points were measured and elemental analy-
sis for the prepared Schiff bases was done. The results obtained
were in good agreement with the calculated values. The prepared
Schiff bases have the following structural formulae:
The pH-measurements were carried out using a Jenway
3310 digital pH-meter at 25 ◦C and 0.1 M ionic strength. The
infrared spectra of all ligands and their complexes were recorded
(4000–200 cm−1) using a Perkin-Elmer FTIR 2000 spectopho-
tometer applying the KBr disc technique. The infrared spectral
vibration modes are listed in Table 2.
TheUV–visibleabsorptionspectrawererecordedusingJasco
V-350 recording spectrophotometer at room temperature. The
absorption spectra of the Schiff bases under investigation were
scanned over the wavelength range 280–500 nm in buffer solu-
tions of different pH values to determine their ionization con-
stants. Thermal analyses (TG and DTG) were obtained in a nitro-
gen atmosphere using a type TGA 50 Shimadzu derivatograph
thermal analyzer. The heating rates were suitably controlled at
10 ◦C min−1 under a nitrogen atmosphere and the w◦eight loss
∼
was measured from ambient temperature up to 800 C.
=
4. Results and discussion
4.1. Determination of the pKa values of the investigated
ligands spectrophotometrically
Table 1 shows the analytical data of the prepared Schiff bases.
The absorption spectra of 10−4 M of the Schiff bases
(I–IV) in 80% ethanolic buffer solutions of varying pH val-
ues were recorded to investigate the spectral properties of the
species liable to exist in such media and to determine the
ionization constants (pKa) values of the acidic group present.
Britton–Robinson universal buffers [28] were used to control the
pH over the range 2.00–12.00. The spectra in acidic solutions
of pH 2.00–6.00 are characterized by a strong band absorbing
maximally within the range 310–320 nm. These bands are due to
absorption of the nonionised form liable to exist in such media
2.2. Preparation of the solid complexes
The solid complexes were prepared by mixing hot saturated
alcoholic solutions of 0.001 M of metal salt solutions with the
requisite amount of the ligands under study (I–II) sufficient to
form 1:1 or 1:2 (M:L) complexes. The reaction mixture was
stirred on a water bath. If the solid complexes did not sep-
arate on standing, diluted ammonia solution 1:10 was added
to the complexes of Cr(III), Fe(III) and Ni(II) to maintain the
pH at 5–6. The solid complexes were filtered off and washed
several times with ethanol until the filtrate becomes colour-
less. The obtained complexes were kept in a vacuum desiccator.
The metal content of the prepared complexes was determined
[27].
*
and may be assigned to ␲–␲ electronic transition within the
Schiff base of aniline (I), 2-aminpyridine (II), 4-aminopyridine
(III) and 2-aminopyrimidine (IV) salicylaldehyde influenced by
intramolecular charge transfer. The spectra in alkaline solutions
Table 1
Analytical data of the prepared Schiff bases
Compound
Tentative formula
mp (◦C)
Colour
Molecular
weight
Yield %
C% calculated (found)
H% calculated (found)
N% calculated (found)
I
II
III
IV
C13H11NO
C12H10N2O
C12H10N2O
C11H9N3O
49
65
66
Orange
Orange
Orange
Orange
197
198
198
199
70
48
52
45
79.18 (79.4)
72.72 (72.7)
72.71 (72.3)
66.33 (65.4)
5.58 (5.4)
5.05 (4.9)
5.05 (5.3)
4.52 (4.5)
7.10 (7.2)
14.14 (13.8)
14.14 (13.9)
21.10 (20.7)
123
Table 2
IR band assignments of the investigated free ligands (I–IV)
Compound
νOH
νC–H asymmetry
νC–H symmetry
νC
N
ν(C–O)
νC–H (CH N)
I
II
III
IV
3447b
3421b
3365b
3419b
2884m
2556m
2361m
2831m
3055m
3051m
3224m
3052m
1650m
1645m
1647m
1643m
1275s
1275s
1275s
1275s
980w
993w
998w
990w
s: strong, b: broad, m: medium, w: weak.

952
S.A. Abdel-Latif et al. / Spectrochimica Acta Part A 67 (2007) 950–957
Table 3
are characterized by the presence of a strong band absorbing
maximally within the range 378–384 nm which may be assigned
to the absorption of the ionized form liable to exist at high pH
values as a result of acid–base equilibrium. The absorbance of
the bands assigned to the ionized form increases gradually by
increase of pH attaining a maximum value at the pH range 11–12
for the investigated ligands (I–IV).
The variation of absorbance with pH is used for the cal-
culation of the ionization constants of the investigated ligands
using (a) the half height method and (b) the modified limiting
absorbance method [29].
Collective data of log K1, log K2, log β2 and ꢀG values for binary complexes of
Cr(III), Mn(II), Fe(III), Ni(II) and Cu(II) with the investigated ligands (I–IV) at
25 ◦C and 0.1 M ionic strength
Complex
log K1
log K2
log β2
−ꢀG (kJ mol−1
)
Cr(I)
Mn(I)
Fe(I)
Ni(I)
Cu(I)
5.97
4.45
6.60
6.81
6.92
4.20
3.43
5.10
5.42
5.94
10.17
7.88
11.70
12.23
12.86
58.02
44.96
66.75
69.77
73.37
Cr(II)
Mn(II)
Fe(II)
Ni(II)
Cu(II)
6.19
5.02
6.62
6.82
6.86
4.93
3.81
5.21
5.52
6.10
11.12
8.83
11.83
12.34
12.86
63.45
50.38
67.50
70.41
73.37
The pKa values are 7.6, 7.9, 6.6 and 8.4 for the investigated
ligands I–IV, respectively.
Cr(III)
Mn(III)
Fe(III)
Ni(III)
Cu(III)
6.89
5.59
7.06
7.38
7.39
5.26
4.52
5.55
6.06
6.31
12.15
10.11
12.61
13.44
13.70
69.32
57.68
71.95
76.68
78.17
4.2. Determination of the protonation constants of the
investigated ligands potentiometrically
The titration curves show that there is one sharp jump indica-
tive of one neutralisation equilibrium corresponding to the
librated proton of the OH group of the aldehydic moiety. The
average numbers of protons associated with the ligand (n¯A) at
different values were calculated from acid and ligand titration
curves using Irving and Rossotti equations [30].
Cr(IV)
Mn(IV)
Fe(IV)
Ni(IV)
Cu(IV)
6.44
5.19
7.14
7.88
7.66
5.35
3.78
6.12
6.09
6.62
11.79
8.87
13.26
13.97
14.28
67.27
51.18
75.65
79.71
81.47
The data obtained shows that the proton–ligand ioniza-
tion constants increases in the order I (7.35) < II (7.99) < III
(8.04) < IV (8.55), respectively. These results reveal that the
value of pKa increases with increase of the basicity of the hetero
ring.
4.4. Infrared spectra of the free ligands and their metal
complexes
The metal complexes of ligands I and II with Cr(III), Mn(II),
Fe(III), Ni(II) and Cu(II) were prepared and subjected to many
analytical tools such as elemental analyses C, H, N and halogen
(Table 4), IR and TG analysis. The assignment of the impor-
tant bands was made and recorded in Table 5. In order to give
conclusive idea about the structure of the metal complexes, the
main IR bands were compared with those of the free ligands
(Table 2). The strong broad band observed at 3447–3135 cm−1
in the spectra of the free ligands is assigned to νOH. The OH
stretching frequency appearing in the spectra of all complexes
studied as a broad band in the range 3447–3148 cm−1 is due
to the presence of water of hydration and/or coordinated water.
The Schiff bases under investigation displays the C N band
within the wave number range 1643–1650 cm−1. Upon chela-
tion a shift to lower wave number has been observed. This
indicated that it may be considered as a center of chelation.
The phenolic C–O of the free ligands (I and II) is observed at
1275 cm−1 [31]. Upon chelation, this band is shifted to higher
wave number which means that the shifts are due to coordination
of ligand to metal atom by the azomethine nitrogen and pheno-
lic oxygen [31]. In the free ligands (I and II) a band observed at
980 and 993 cm−1 is ascribed to aldehydic νC–H (CH N) and
shifted to higher wave number on complexation. The new IR
bands appearing at 431–455 and 555–470 cm−1 are assigned to
ν(M–O) and ν(M ← N) vibrations [32], respectively. The most
significant conclusion drawn from the proceeding arguments is
that ligands I and II acts as monobasic bidentate ligands towards
the central metal ion to form cationic six- or four-coordinate
chelates.
4.3. Determination of the stability constants of the
complexes
The stability constants of the divalent Mn, Ni, Cu and triva-
lent Cr and Fe metal ions with the investigated Schiff bases
(I–IV) have been determined at 25 ◦C and 0.1 M ionic strength.
An examination of the titration curves indicates that the metal
titration curve is separated from the ligand titration curve. The
end points of the titration of the three solution increases in the
order acid < ligand < complex curves. This is indicative of the
formation of the complexes.
The formation curves of the complexes were obtained by
plotting the relationship between average number of ligands
attached per metal ion n¯ and free ligand exponent pL. The val-
ues of log β1(log K1) and log β2(log K1 + log K2) derived from
these data are collected in Table 3. These are obtained from the
formation curves, where log K1 is the pL value at n¯ = 0.5 and
log K2 is the value at n¯ = 1.5. Since the maximum value of n¯
in all cases was less than 2, no complexes of order higher than
1:2 are presumed to form. Also the values of the free energy of
formation calculated by the use of the relationship:
−ꢀG = 2.303RT log β2
It is well known that the larger the negative value of
ꢀG, the more stable is the complex formed. From Table 3,
it is observed that the stability decreases is the order
Cu(II) > Ni(II) > Fe(III) > Cr(III) > Mn(II).

Table 4
Analytical data of the investigated ligands I and II, and their Cr(III), Mn(II), Fe(III), Ni(II) and Cu(II) complexes
Complexes
Empirical formula
M:L
FW
% Yield
C% (calculated)
found
H% (calculated)
found
N% (calculated)
found
Cl% (calculated)
found
M% (calculated)
found
[Cr(I)·Cl·OH·2H2O]H2O
[Cr(I)2·Cl·H2O]
C13H17CrNO5Cl
C26H22CrN2O3Cl
1:1
1:2
1:1
1:2
1:1
1:2
1:1
1:2
1:1
1:2
1:1
1:2
1:1
1:2
1:1
1:2
1:1
1:2
1:1
1:2
354.5
497.5
340.4
482.9
358.3
537.3
344.2
486.7
330.5
491.5
355.5
535.5
341.4
484.9
359.3
503.3
345.1
488.6
350
80
75
75
65
75
80
80
75
80
70
75
80
70
75
80
70
85
75
70
75
(44.00) 43.8
(62.71) 62.4
(45.82) 45.5
(64.60) 64.1
(43.53) 43.73
(58.06) 58.2
(45.32) 45.4
(64.10) 64.3
(47.2) 47.4
(63.47) 63.5
(40.50) 40.8
(53.78) 53.6
(42.17) 42.4
(59.39) 59.6
(40.07) 40.1
(57.22) 57.5
(41.72) 41.3
(58.94) 59.0
(41.14) 41.3
(58.35) 58.4
(4.79) 4.7
(4.42) 4.9
(4.70) 5.1
(4.96) 4.9
(4.74) 4.5
(4.83) 5.1
(4.64) 4.5
(4.93) 4.5
(5.14) 5.0
(4.88) 4.6
(4.5) 4.4
(4.48) 4.6
(4.39) 4.3
(4.53) 4.7
(4.45) 4.6
(3.97) 4.2
(4.34) 4.1
(4.50) 4.3
(4.28) 4.4
(4.45) 4.3
(3.94) 3.8
(5.62) 5.3
(4.11) 3.9
(5.79) 5.3
(3.90) 3.6
(5.21) 5.5
(4.06) 4.2
(5.75) 5.7
(4.23) 4.1
(5.69) 5.0
(7.87) 7.9
(10.45) 10.5
(8.2) 8.1
(11.54) 11.3
(7.79) 7.7
(11.12) 10.9
(8.11) 8.2
(11.46) 11.6
(8.00) 8.2
(11.34) 11.2
(10.01) 10.1
(7.13) 7.3
(10.42) 9.17
–
(9.90) 9.7
(6.6) 6.4
(10.31) 10.5
–
(14.66) 14.0
(10.45) 10.8
(16.12) 16.9
(11.36) 10.9
(15.57) 15.1
(10.38) 10.4
(17.05) 17.4
(12.06) 12.5
(19.21) 19.8
(12.91) 12.8
(14.62) 14.7
(9.71) 9.6
(16.08) 16.8
(11.32) 11.7
(15.53) 15.4
(11.08) 11.6
(17.00) 17.4
(12.01) 12.3
(18.14) 18.5
(12.86) 12.7
[Mn(I)·Cl·3H2O]
[Mn(I)2·2H2O]
C
13H16MnNO4Cl
C26H24MnN2O4
C13 17FeNO5Cl
[Fe(I)·Cl·OH·2H2O]H2O
[Fe(I)2·Cl·H2O]2H2O
[Ni(I)·Cl·3H2O]
H
C26H26FeN2O5Cl
C13H16NiNO4Cl
C26H24NiN2O4
[Ni(I)2·2H2O]
[Cu(I)·OH·3H2O]
[Cu(I)2·2H2O]
C
13H17CuNO5
–
–
C26H24CuN2O4
[Cr(II)·Cl·OH·2H2O]H2O
[Cr(II)2·Cl·H2O]2H2O
[Mn(II)·Cl·3H2O]
[Mn(II)2·2H2O]
C12H16CrN2O5Cl
C24H24CrN4O5Cl
C12H15MnN2O4Cl
C24H22MnN4O4
C12H16FeN2O5Cl
C24H20FeN4O3Cl
C12H15NiN2O4Cl
(9.98) 9.7
(6.62) 6.7
(10.39) 10.5
–
(9.88) 9.9
(7.05) 7.2
(10.28) 10.7
–
[Fe(II)·Cl·OH·2H2O]H2O
[Fe(II)2·Cl·H2O]
[Ni(II)·Cl·3H2O]
[Ni(II)2·2H2O]
C24
C12H15CuN2O4Cl
C24 22CuN4O4
H22NiN4O4
[Cu(II)·Cl·3H2O]
[Cu(II)2·2H2O]
(10.14) 10.3
–
H
493.5

954
S.A. Abdel-Latif et al. / Spectrochimica Acta Part A 67 (2007) 950–957
Table 5
IR band assignment of Cr(III), Mn(II), Fe(III), Ni(II) and Cu(II) complexes with the investigated ligands I and II
Complex
M:L
ν(OH)
ν(C N)
δOH·H2O
ν(C–N) (CH N)
ν(C–O)
γOH·H2O
ν(M–O)
ν(M ← N)
[Cr(I)·Cl·OH·2H2O]H2O
[Cr(I)2·Cl·H2O]
1:1
1:2
1:1
1:2
1:1
1:2
1:1
1:2
1:1
1:2
1:1
1:2
1:1
1:2
1:1
1:2
1:1
1:2
1:1
1:2
3148b
3364b
3422b
3326b
3421b
3135b
3384b
3305b
3447b
3312b
3357b
3420b
3320b
3360b
3410b
3366b
3442b
3347b
3447b
3440b
1614m
1615m
1619m
1612m
1606m
1610m
1619s
1620m
1623m
1623m
1618m
1619m
1620m
1606m
1610m
1615m
1617m
1625m
1626m
1630m
1529s
1539s
1543s
1541s
1542s
1590s
1533s
1533s
1558s
1532s
1528s
1541s
1530s
1540s
1535s
1540s
1533s
1530s
1532s
1545s
1030w
983w
997w
1295s
1300s
1305s
1303s
1325s
1330s
1323s
1297s
1310s
1315s
1326s
1304s
1298s
1300s
1315s
1320s
1325s
1305s
1308s
1315s
829s
859s
840s
850w
835s
845s
826s
815s
820s
820s
840s
830s
817s
817s
840s
822s
814s
825s
819s
815s
435m
440m
433m
450m
430m
440m
434m
432s
445m
440m
430m
443m
455m
433m
440m
431m
450m
435m
440m
445m
541m
543m
480m
527m
483m
545m
469m
487m
555m
555m
511m
470m
519m
516m
479m
488m
487m
450m
470m
465m
[Mn(I)·Cl·3H2O]
[Mn(I)2·2H2O]
970w
[Fe(I)·Cl·OH·2H2O]H2O
[Fe(I)2·Cl·H2O]2H2O
[Ni(I)·Cl·2H2O]H2O
[Ni(I)2·2H2O]
1013w
1028w
1025w
1025w
1025w
1025w
1033w
1030w
1032w
1029w
1035w
1028w
938w
[Cu(I)·OH·3H2O]
[Cu(I)2·2H2O]
[Cr(II)·Cl·OH·2H2O]H2O
[Cr(II)2·Cl·H2O]2H2O
[Mn(II)·Cl·3H2O]
[Mn(II)2·2H2O]
[Fe(II)·Cl·OH·2H2O]H2O
[Fe(II)2·Cl·H2O]
[Ni(II)·Cl·3H2O]
[Ni(II)2·2H2O]
1030w
1025w
1030w
[Cu(II)·Cl·3H2O]
[Cu(II)2·2H2O]
4.5. Thermal analyses (TG and DTG)
[33] in the temperature range 100–316 ◦C. The organic part of
the complexes may decompose in one or more step with the
possibility of the formation of one or two intermediates. These
intermediates may include the metal ion with a part of the Schiff
base in case of 1:1 or 1:2 complexes. These intermediates may
finally decompose to stable metal oxides.
The decomposition of all complexes ended with the stoichio-
metric oxide formation. The metal content percentages of the
complexes are thus calculated and compared with those obtained
from the metal content analytical determination [27]. The aver-
age metal percentages were found to be in good agreement with
those calculated from the tentative formulae based on the ele-
mental analyses.
In the present investigation, the heating rates were suit-
ably controlled at 10 ◦C min−1 under nitrogen atmosphere and
the weight loss was measured from ambient temperature up
◦
∼
to 800 C. The TG and DTG curves are represented in
=
Figs. 1 and 2. The weight losses for each chelate were calcu-
lated for the corresponding temperature ranges and are shown in
Table 6. The metal percentages calculated from the metal oxide
residues were compared with those determined by the analytical
metal content determination [27]. The hydrated water molecules
are associated with the complex formation and found outside the
coordination sphere formed around the central metal ion. The
dehydration of this type of water takes place in the tempera-
ture range 25–220 ◦C, the weight loss corresponds to one water
molecule. On the other hand, the coordinated water molecules
are eliminated at higher temperatures than the water molecules
of hydration. The water of coordination is usually eliminated
The initial decomposition and inflection temperatures have
been used as an indication on the thermal stability of complexes
[34]. The relative lower temperature assigned to the first step
decomposition for (1:2) complexes may be attributed to the pres-
ence of two organic ligands around the metal ion.
Fig. 1. Thermogravimetric and derivative thermogavimetric analyses curves of
Cr(II) (1:1) complex.
Fig. 2. Thermogravimetric and derivative thermogavimetric analyses curves of
Mn(II) (1:1) complex.

S.A. Abdel-Latif et al. / Spectrochimica Acta Part A 67 (2007) 950–957
955
Table 6
Thermogravimetric results of Cr(III), Mn(II), Fe(III), Ni(II) and Cu(II) complexes with the investigated ligands I and II
Complex
Molecular weight
Temperature (◦C)
DTG temperature (◦C)
Calculated loss %
Found loss %
Assignment
[Cr(I)·Cl·OH·2H2O]H2O
[Cr(I)·Cl·OH·2H2O]
[Cr(I)·Cl·OH]
354.5
336.5
282.5
247
25–116
116–229
229–387
387–512
512–616
616–800
100–220
220–316
316–341
341–445
445–574
574–800
100–208
208–262
262–390
390–508
508–800
100–211
211–339
339–535
535–750
81–230
61
192
–
438
–
–
126
–
–
344
–
–
–
250
–
468
–
130
310
468
–
200
407
5.07
16.04
12.52
31.17
15.29
23.69
5.25
11.16
12.39
36.26
17.51
20.68
5.22
4.63
16.66
12.49
31.65
14.99
23.62
5.45
10.91
12.72
37.26
17.27
20.43
5.62
H2O
2H2O
Cl
C6H6
NC
CrO2
H2O
2H2O
Cl
C6H5N
CO
MnO
H2O
2H2O
Cl
NOCH
NiO
[Cr(I)·OH]
170
[Mn(I)·Cl·3H2O]
[Mn(I)·Cl·2H2O]
[Mn(I)·Cl]
340.4
322.4
286.4
250.9
159.9
[Mn(I)]
[Ni(I)·Cl·3H2O]
[Ni(I)·Cl·2H2O]
[Ni(I)·Cl]
344.2
326.2
290.2
254.7
11.03
12.23
16.88
21.70
3.61
11.25
12.50
16.25
21.30
3.20
[Ni(I)]
[Cr(I)2·Cl·H2O]
[Cr(I)2·Cl]
[Cr(I)2]
497.5
479.5
444
H2O
Cl
7.13
8.49
68.91
20.10
16.84
66.46
15.34
7.32
10.09
22.22
24.17
16.17
5.06
10.66
18.24
36.91
23.62
26.21
47.24
25.45
29–57
37.32
22.71
66.99
21.13
16.92
66.59
15.92
7.27
2C6H5, 2C6H4
CrO3
2H2O, NO2
3C6H5, N, 2C
NiO
2H2O
NO2
C6H5, N
C6H5
[Ni(I)2·2H2O]
486.7
404.7
230–557
[Cu(I)2·2H2O]
491.5
455.5
409.5
129–154
154–170
170–229
229–541
541–800
25–100
100–203
203–306
306–450
450–606
190–334
334–450
450–631
166–190
190–667
667–1000
–
–
177
–
621
55
222
396
–
[Cu(I)2]
9.99
22.72
24.59
17.82
4.54
CuO
H2O
[Cr(II)·Cl·OH·2H2O]H2O
[Cr(II)·Cl·OH·2H2O]
[Cr(II)·Cl·OH]
355.5
337.5
301.5
246.5
11.36
19.67
37.31
24.80
26.55
47.54
25.85
28.99
37.78
22.40
2H2O
NCHCO
C5H5N, C
CrO2
3H2O, Cl
C7H5NO
MnO2
3H2O, Cl, N
C5H4N2
CuO
–
[Mn(II)·Cl·3H2O]
341.4
251.9
132.9
350
303
375
–
126
–
[Mn(II)]
[Cu(II)·Cl·3H2O]
246.5
–
For Cr(I), Mn(I), Ni(I) (1:1) and Cr(I), Ni(I), Cu(I) (1:2) com-
plexes (Table 6), a mass loss occurred within the temperature
range 25–230 ◦C corresponding to the loss of hydrated water
molecule and at the temperature range 116–316 ◦C correspond-
ing to a loss of coordinated water molecules. In the temperature
range 229–390 ◦C a mass loss occured corresponding to a loss
of Cl atoms. At the temperature range 393–535 ◦C mass losses
occured due to the formation of intermediate species through
the decomposition of the organic moiety of the complexes and
this continues till a constant weight is obtained where a metallic
oxide residue is formed.
For Cr(II), Mn(II), Cu(II) (1:1) complexes, a mass loss
occurred within the temperature range 25–100 ◦C corresponding
to the loss of hydrated water molecule in case of Cr(II) complex.
At the temperature range 100–334 ◦C corresponding to a loss
of coordinated water molecules (Table 6). At higher tempera-
ture a mass loss corresponding to the formation of intermediate
species through the decomposition of the organic moiety of the
Fig. 3. Coats–Redfern plot for Cr(II) (1:1) complex.

956
S.A. Abdel-Latif et al. / Spectrochimica Acta Part A 67 (2007) 950–957
complexes and this continues till a constant weight is obtained
where a metallic oxide residue is formed.
Based on the results gained from elemental analysis, IR and
TG the structure of the formed complexes may be represented
as follows:
Fig. 4. Coats–Redfern plot for Mn(II) (1:1) complex.
was used for the determination of the values of the reaction order.
Here Ws stands for the weight remaining at a given temperature
Ts, i.e. the DTG peak temperature, W0 and W are the initial
∞
and final weights of the substance, respectively. It was reported
that the first order reaction has a Cs value of 0.368 and the
second order one has Cs value of 0.5 [35]. The values of Cs
for the thermal decomposition of Schiff base complexes under
investigation are in the range 0.49–0.93 which indicates that the
decomposition follows second order kinetics.
4.6.2. Integral method using the Coats–Redfern equation
For a second order process the Coats–Redfern equation [36]
may be written in the form:
ꢀ
ꢁ
ꢀ
ꢂ
ꢃꢁ
(Wf/(Wf − W))
AR
θE∗
2RT
E∗
E∗
log
= log
1 −
−
T2
2.303RT
4.6. Kinetics of thermal decomposition of the complexes
4.6.1. Determination of reaction order of decomposition
where Wf is the mass loss at the completion of the reaction,
W the mass loss up to temperature T; (Wr = Wf − W), R the gas
constant, E* the activation energy in J mol−1 and θ is the heating
The Horowitz and Metzger [35] equation, Cs = (n)1/1−n
,
rate. Since 1 − 2RT/E* 1, a plot of the left-hand side of the
∼
=
where Cs is the weight fraction of the substance present at the
DTG peak temperature, Ts, is given by
above equation against 1/T was constructed (Figs. 3 and 4) and
E* was calculated from the slope and A (Arrhenius constant)
was found from the intercept.
Ws − W
The activation entropy S*, the activation enthalpy H* and the
free energy of activation G* were calculated using the following
∞
Cs =
(1)
W0 − W
∞
Table 7
Kinetic data of the thermal decomposition of the Schiff bases 1:1 and 1:2 complexes
Complex
E* (kJ mol−1
)
R2
S* (J K−1 mol−1
)
H* (kJ mol−1
)
G* (kJ mol−1
)
Cs
[Cr(I)·Cl·OH·2H2O]H2O
[Mn(I)·Cl·3H2O]
[Ni(I)·Cl·3H2O]
37.43, 41.22
70.02, 99.71
54.03, 52.50
106.36
18.59, 81.68
58.66, 95.73
21.54, 170.95
218.20, 130.35
11.25
0.998, 0.996
0.987, 0.999
0.997, 0.995
0.999
0.996, 0.995
0.999, 0.998
0.962, 0.979
0.999, 0.998
0.997
−228.62, −248.87
−230.07, −230.95
−229.12, −232.55
−231.15
33.78, 35.73
65.27, 93.25
50.03, 47.22
100.55
15.17, 76.47
55.04, 88.82
17.97, 165.51
213.51, 124.16
7.56
106.34, 176.98
142.02, 185.09
119.87, 172.37
271.83
123.97, 233.31
158.23, 296.08
132.09, 318.95
345.01, 274.16
100.24
0.73, 0.64
0.61, 0.67
0.56, 0.49
0.93
0.77, 0.23
0.59, 0.72
0.49, 0.20
0.64, 0.35
–
[Cr(I)2·Cl·H2O]
[Ni(I)2·2H2O]
−230.03, −230.65
−229.30, −231.83
−230.55, −229.36
−228.29, −230.20
−232.28
[Cu(I)2·2H2O]
[Cr(II)·Cl·OH·2H2O]H2O
[Mn(II)·Cl·3H2O]
[Cu(II)·Cl·3H2O]
R2 = correlation coefficient of the fitted straight line.

S.A. Abdel-Latif et al. / Spectrochimica Acta Part A 67 (2007) 950–957
[3] S.A. Sallam, Transition Met. Chem. 31 (2006) 46.
957
equations:
[4] M. Cindric´, N. Strukan, V. Vrdoljak, T. Kajfezˇ, B. Kamenar, Croatica Chim.
Acta 76 (2003) 157.
[5] C. Sousa, C. Freire, B. de Castro, Molecules 8 (2003) 894.
[6] D.E. Hamilton, R.S. Drago, A. Zombeck, J. Am. Chem. Soc. 109 (1987)
374.
ꢂ
ꢃ
Ah
KT
S∗ = 2.303 log
G∗ = H∗ − TsS∗
R,
H∗ = E∗ − RT,
[7] D. Chen, A.E. Martel, Inorg. Chem. 26 (1987) 1026.
[8] J. Costamagna, J. Vargas, R. Latorre, A. Alvarado, G. Mena, Coord. Chem.
Rev. 119 (1992) 67.
[9] M.J. Samide, D.G. Peters, J. Electroanal. Chem. 443 (1998) 95.
[10] M. Kato, Y. Muto, Coord. Chem. Rev. 92 (1988) 45.
[11] J. Losada, I. Del Peso, L. Beyer, Inorg. Chem. Acta 301 (2001) 107.
[12] M.L.P. Santos, I.A. Bagatin, E.M. Pereina, A.M.C. Ferreira, J. Chem. Soc.
Dalton Trans. (2001) 838.
where K and h are the Boltzman’s and Plank’s constants, respec-
tively. The calculated values of E*, S*, H* and G* for the
dehydration and the decomposition steps are given in Table 7.
The activation energies of the decomposition were found to
be in the range 11–218 kJ mol−1. The high values of the activa-
tion energies reflect the thermal stability of the complexes. The
entropy of activation was found to have negative values in all of
the complexes, which indicates that the decomposition reactions
proceed with changes in order which may be disordered due to
the decomposition of the complexes.
[13] T.N. Rahmonui, S.D. Sid, N. Chenah, O.B. Baitich, Synt. React. Inorg.
Met. Org. Chem. 29 (1999) 79.
[14] M.A. Martinez, R. Aquiler-Saloma, N.M. Ruvalcaba, R.C. Rosado, A.N.
Vazquez, V.G. Vidales, A.Z. Dehesa, R.A. Toskano, S.H. Ortega, G.J.M.
Fernandez, J. Chem. Soc. Dalton Trans. (2001) 2346.
[15] L.T. Yildirm, K.C. Emregul, R. Kurtaran, O. Atakol, Cryst. Res. Technol.
37 (2002) 1344.
5. Conclusion
[16] S.S. Djebbar, B.O. Benali, J.P. Deloume, Polyhedron 16 (1997)
2175.
The spectrophotometric determination of the pKa suggests
the values 7.6, 7.9, 6.6 and 8.4 for the investigated ligands I–IV,
respectively. The ionization constants increase in the order I
(7.35) < II (7.99) < III (8.04) < IV (8.55), respectively. The sta-
bility constants of the divalent Mn, Ni, Cu and trivalent Cr and Fe
metal ions with the Schiff bases (I–IV) have been determined.
The metal complexes of ligands I and II with Cr(III), Mn(II),
Fe(III), Ni(II) and Cu(II) were prepared. The main IR bands of
the complexes were compared with those of the free ligands.
The TG and DTG curves are documented. The weight losses for
each chelate were calculated for the corresponding temperature
ranges. The values of Cs for the thermal decomposition of Schiff
base complexes under investigation were calculated indicating
that the decomposition follows second order kinetics. The acti-
vation energies of the decomposition were also calculated. The
high values of the activation energies reflect the thermal stability
of the complexes. The entropy of activation was found to have
negative values in all of the complexes, which indicates that the
decomposition reactions proceed with a change in order which
may be disordered due to the decomposition of the complexes.
[17] P. Bhattacharyya, J. Parr, A.T. Ross, J. Chem. Soc. Dalton (1998)
3149.
[18] L. He, S.H. Gou, Q.F. Shi, J. Chem. Crystallogr. 29 (1999) 207.
[19] J.C. Wu, N. Tang, W.S. Liu, M.Y. Tan, A.S.C. Chan, Chin. Chem. Lett. 12
(2001) 757.
[20] D. Chatterjee, A. Mitra, J. Coord. Chem. 57 (2004) 175.
[21] G. Minu, H. Bhowon, A. Li Kam Wah, M.O. Dosieah Ridana,
D. Ramalingum Lacour, Synth. React. Inorg. Metal-Org. Chem. 34
(2004) 1.
[22] C.M. Liu, R.G. Xiong, X.Z. You, Y.J. Liu, K.K. Cheung, Polyhedron 15
(1996) 4565.
[23] S.S. Djebbar, B.O. Benali, J.P. Deloume, Transition Met. Chem. 23 (1998)
443.
[24] Y.J. Hamada, IEEE Trans. Electron Devices 44 (1997) 1208.
[25] R. Ramesh, M. Sivagamasundari, Synth. React. Inorg. Met.-Org. Chem. 33
(2003) 899.
[26] A.I. Vogel, A Text Book of Practical Organic Chemistry, 4th ed., Longman,
London, 1978, p. 925.
[27] A.M.G. Macdonald, P. Sirichanya, Microchem. J. 14 (1969) 199.
[28] H.T.S. Britton, Hydrogen Ions, Chapman and Hall, London, 1955.
[29] R.M. Issa, A.H. Zewail, J. Chem. UAR 14 (1971) 461.
[30] H.M. Irving, H.S. Rossotti, J. Chem. Soc. 74 (1953) 3397.
[31] E. Canpolat, M. Kaya, Turk. J. Chem. 29 (2005) 409.
[32] X. Tai, X. Yin, Q. Chen, M. Tan, Molecules 8 (2003) 439.
[33] N.T. Abdel-Ghani, O.E. Sherif, Thermochim. Acta 156 (1989) 69.
[34] M. Nath, P. Arora, Synth. React. Inorg. Met-Org. Chem. 23 (1993)
1523.
References
[1] S.M.E. Khalil, Chem. Papers 54 (2000) 12.
[2] A.H. Osman, Transition Met. Chem. 31 (2006) 35.
[35] H.H. Horowitz, G. Metzger, J. Anal. Chem. 35 (1963) 1464.
[36] A.W. Coats, J.P. Redfern, Nature 201 (1964) 68.