Spectral properties of some metal complexes derived from uracil-thiouracil and citrazinic acid compounds

Article abstract of DOI:10.1016/j.saa.2006.07.046

The reaction of FeCl₃ with uracil (H₂L¹), citrazinic acid (H₂L⁶), 5-(phenylazo)citrazinic acid (H₂L⁷), 5-(m-hydroxyphenylazo)citrazinic acid (H₂L⁸) and 5-(m-nitrophenylazo)citrazinic acid (H₂L⁹) leads to the formation of complexes with the empirical formula Fe(HL)₃·nH₂O (n = 1-3). All of the prepared complexes have octahedral complexation geometry where the azo group is not the reactive site for complexation. Thiouracil (H₂L²) and the 5-(substituted phenylazo)thiouracil (H₂L³-H₂L⁵) ligands are bidentates on complexation with Co(II), Ni(II) and Cu(II). The complexes have been characterized by elemental analyses, IR, electronic spectra, magnetic susceptibility, DTA, electron spin resonance (copper complexes) and Moessbauer spectra (iron complexes). The coordination bond lengths between the metal ion and the active centers for complexation were calculated.

Full text of DOI:10.1016/j.saa.2006.07.046

Spectrochimica Acta Part A 67 (2007) 662–668
Spectral properties of some metal complexes derived from
uracil–thiouracil and citrazinic acid compounds
Mamdouh S. Masouda,∗, Amany A. Ibrahimb, Ekram A. Khalila, Adel El-Marghanyc
a Chemistry Department, Faculty of Science, Alexandria University, Alexandria, Egypt
b Chemistry Department, Faculty of Science, Ain Shams University, Cairo, Egypt
c Chemistry Department, Faculty of Education, Suez Canal University, Suez, Egypt
Received 14 March 2006; accepted 25 July 2006
Abstract
The reaction of FeCl3 with uracil (H2L1), citrazinic acid (H2L6), 5-(phenylazo)citrazinic acid (H2L7), 5-(m-hydroxyphenylazo)citrazinic acid
(H2L8) and 5-(m-nitrophenylazo)citrazinic acid (H2L9) leads to the formation of complexes with the empirical formula Fe(HL)3·nH2O (n = 1–3). All
of the prepared complexes have octahedral complexation geometry where the azo group is not the reactive site for complexation. Thiouracil (H2L2)
and the 5-(substituted phenylazo)thiouracil (H2L3–H2L5) ligands are bidentates on complexation with Co(II), Ni(II) and Cu(II). The complexes
have been characterized by elemental analyses, IR, electronic spectra, magnetic susceptibility, DTA, electron spin resonance (copper complexes)
and Mo¨ssbauer spectra (iron complexes). The coordination bond lengths between the metal ion and the active centers for complexation were
calculated.
© 2006 Elsevier B.V. All rights reserved.
1. Introduction
2. Experimental
Pyrimidine derivatives constitute an important class of com-
pounds because they are components of the biologically impor-
tant nucleic acid [1,2]. Compounds containing nitrogen and
sulphur as donor atoms have an important role to play as anti-
cancer and antiviral agents [3,4]. The Pt–uracil complexes have
high antitumor, antibacterial and antiviral activities and have a
low level of renal toxicity [5–9].
Citrazinic acid azo dyes were prepared to dye cellulosic and
nylon fibers [10]. Citrazinic acid is used in the field of pho-
tography [11–13] as an inhibitor-removing wash bath for direct
positive colour photographic development [12].
All the azo compounds (H2L3 to H2L5 and H2L7 to H2L9)
were prepared in a similar way by the usual diazotization process
[14]. An example H2L3 is given here: (0.1 mol, 9.3 g) of aniline
was dissolved in 5 mL (0.2 mol) HCl and 25 mL distilled water.
The hydrochloride compound was diazotized below 5 ◦C with
a solution of NaNO2 (0.1 mol, 6.9 g) in 20 mL distilled water.
The diazonium chloride was coupled with an alkaline solution
of H2L2 (0.1 mol, 12.8 g) in 30 mL distilled water. The crude
dye was filtered and crystallized from 75% ethanol–water (v/v),
then dried in a vacuum desiccator over CaCl2.
The iron complexes were prepared by dissolving 30 mmol
(3.36 g) of the ligand (H2L1) in 150 mL warm ethanol at 70 ◦C,
and mixed with 50 mL of an alcoholic solution containing
15 mmol (4.05 g) of iron(III) chloride for 1 h. The reaction mix-
ture was left to cool at room temperature where the complexes
precipitated.
The complexes of cobalt(II), nickel(II) and copper(II)
(10 mmol) were prepared in 10 mL of aqueous ammonia solu-
tion (25%) and 10 mL ethanol by refluxing 10 mmol of the
metal(II) chlorides with 20 mmol of the organic ligand dissolved
in 100 mL of ethanol. The reaction mixture was refluxed for 1 h
at 70 ◦C. The precipitated complexes then were separated by
filtration and washed with H2O followed by drying in a desic-
In this paper, an attempt is made to assign the centers of
complexation. The structures of complexes are based on vari-
¨
ous spectroscopic methods (UV, vis, IR, ESR and Mossbauer),
magneticsusceptibilityandDTA. Thecoordinationbondlengths
were calculated. The structures of the ligands are shown in
Fig. 1.
∗
Corresponding author. Tel.: +20 3 583 22 15; fax: +20 3 391 17 94.
E-mail address: drmsmasoud@yahoo.com (M.S. Masoud).
1386-1425/$ – see front matter © 2006 Elsevier B.V. All rights reserved.
doi:10.1016/j.saa.2006.07.046

M.S. Masoud et al. / Spectrochimica Acta Part A 67 (2007) 662–668
663
Table 1
Analytical data of the complexes
Complex
MW
mp (◦C)
% yield
Colour
Empirical formula
Calculated/Found (%)
M
C
H
N
Fe(HL1)3·3H2O
Fe(HL6)3·3H2O
Fe(HL7)3·2H2O
Fe(HL8)3·3H2O
Fe(HL9)3·H2O
Co(HL2)2·3H2O
Ni(HL2)2·3H2O
Cu(HL2)2·H2O
Co(HL3)2·3H2O
Ni(HL3)2·3H2O
Cu(HL3)2·2H2O
Co(HL4)2·3H2O
Ni(HL4)2·4H2O
Cu(HL4)2·2H2O
Co(HL5)2·4H2O
Cu(HL5)2·3H2O
443
572
866
932
320
325
340
360
365
315
319
325
318
320
333
320
321
335
322
340
85
79
77
82
80
85
73
76
82
85
82
75
76
80
81
82
Yellowish brown
Black
Brownish yellow
Brown
Brownish yellow
Pale brown
Green
Pale blue
Orange
Green
Green
Deep brown
Light green
Blue
C12H15FeN6O9
12.6 (12.8)
9.8 (9.6)
6.5 (6.4)
6.0 (6.1)
5.7 (5.7)
15.8 (15.5)
15.8 (15.7)
18.9 (18.7)
9.8 (10.0)
9.7 (9.8)
11.3 (11.2)
9.5 (9.3)
9.3 (9.2)
32.5 (32.6)
37.6 (37.5)
49.9 (49.8)
46.4 (46.4)
43.9 (43.8)
26.4 (26.2)
26.2 (26.3)
28.6 (28.5)
41.9 (41.8)
41.9 (41.8)
42.7 (42.6)
39.7 (39.8)
38.5 (38.5)
40.4 (40.5)
38.5 (38.5)
39.2 (39.1)
3.3 (3.4)
3.1 (3.2)
3.2 (3.1)
3.2 (3.2)
2.3 (2.2)
3.3 (3.2)
3.3 (3.2)
2.4 (2.4)
3.5 (3.4)
3.5 (3.5)
3.2 (3.1)
3.3 (3.3)
3.5 (3.3)
3.0 (3.1)
3.5 (3.5)
3.3 (3.2)
18.9 (19.1)
7.3 (7.0)
C18H18FeN3O15
C36H28FeN9O14
C36H23FeN9O18
C36H28FeN12O19
C8H12CoN4O5S2
C8H12NiN4O5S2
C8H8CuN4O3S2
C20H20CoN8O5S2
C20H20NiN8O5S2
C20H18CuN8O4S2
C20H20CoN8O7S2
C20H22NiN8O8S2
C20H18CuN8O6S2
C20H22CoN8O8S2
C20H20CuN8O7S2
14.5 (14.6)
13.5 (13.7)
17.1 (17.2)
15.3 (15.0)
15.3 (15.1)
16.7 (16.9)
19.5 (19.3)
19.5 (19.3)
19.9 (19.8)
18.5 (18.4)
18.0 (17.8)
18.9 (19.0)
18.0 (18.0)
18.3 (18.1)
983
365.9
365.7
335.5
373.9
373.7
561.5
605.9
623.7
593.5
623.9
611.5
10.7 (10.5)
9.2 (9.1)
10.4 (10.3)
Brown
Deep green
3.2. Iron–uracil complex
The IR spectra of this complex show the following features:
(1) The broad bands at 3410 and 3370 cm−1 in uracil and its iron
complex, respectively, are due to ν(OH) of water molecules
and hydrogen bonds of the type N–H· · ·O [14,16].
(2) The uracil ν(NH) bands at 3111 and 3041 cm−1 are strongly
affected on complexation to iron(III) with the presence of
only one band at 3100 cm−1 due to Fe–N interaction [16].
(3) Both ν(C O) bands of uracil at 1740 and 1714 cm−1 are
affected on complexation to iron(III) with different degrees.
The first band is shifted to 1727 cm−1, while the second one
disappeared. Such finding typified Fe–O bonding [14].
(4) The two amide bands of uracil at 1666 and 1643 cm−1 are
affected on complexation. The first band is affected very
slightly and appeared at 1663 cm−1 while the second band
is absent. Thus, uracil acts as a bidentate ligand through oxy-
gen and nitrogen atoms on reaction with iron. In the lower-
frequency spectral region bands for ν(Fe–O) and ν(Fe–N)
are seen at 548 and 428 cm−1, respectively [17].
Fig. 1. Structures of the ligands.
cator over anhydrous CaCl2. The metal content was determined
by usual complexometric titration procedures [15]. The analyt-
ical data and some of the physical properties are collected in
Table 1.
The Nujol mull electronic absorption spectra of the yel-
lowish brown–iron complex [Fe(HL1)3]·3H2O (Table 2) gave
charge transfer bands at 42,105 and 27,826 cm−1. The visible
d–d electronic spectral band at 20,121 cm−1 and the effective
room temperature magnetic moment (μeff = 5.9 BM) typified the
existence of an octahedral high-spin state [18]. The structure in
Fig. 2 is proposed for the Fe(III)–uracil complex.
3. Results and discussion
3.1. Iron complexes
3.3. Iron–citrazinic acid complex
All of the iron complexes are of 1:3 stoichiometry. A gener-
alized equation is given to represent these reactions:
The infrared spectrum of the iron–citrazinic acid complex
[Fe(HL6)3]·3H2O, shows a band at 3404 cm−1, which is due
to ν(O–H) of the citrazinate moiety or to H2O molecules in
the structure of the complex. The band at 3214 cm−1 is due
to ν(N–H) of the citrazinate moiety. The bands at 1705 and
FeCl3 + 3H2L → Fe(HL)3·nH2O + 3HCl
L stands for any ligand used (L1, L6, L7, L8, L9) and n is 3, 3,
2, 3, 1 in the sequence of these ligands.

664
M.S. Masoud et al. / Spectrochimica Acta Part A 67 (2007) 662–668
Table 2
Nujol mull electronic absorption spectra and room temperature (25 ◦C) magnetic
moment values of the metal complexes
Complex
λ (cm−1
)
μeff (298 K) (BM)
Fe(HL1)3·3H2O
Fe(HL6)3·3H2O
Fe(HL7)3·2H2O
Fe(HL8)3·3H2O
Fe(HL9)3·H2O
Co(HL2)2·3H2O
Ni(HL2)2·3H2O
Cu(HL2)2·H2O
Co(HL3)2·3H2O
Ni(HL3)2·3H2O
Cu(HL3)2·2H2O
Co(HL4)2·3H2O
Ni(HL4)2·4H2O
Cu(HL4)2·2H2O
Co(HL5)2·4H2O
Cu(HL5)2·3H2O
42,105; 27,826; 20,121
42,105; 37,175; 21,739
36,866 23,426; 20,779
36,866; 23,520; 20,000
5.9
6.1
5.96
5.8
5.85
5.1
3.1
1.8
5.2
3.2
1.9
5.1
3.1
1.82
5.15
1.83
36,866; 31,546; 23,529; 20,000
45,662; 37,736; 22,727; 17,857
47,170; 35,842; 22,988; 19,685
36,900; 27,397; 19,841
39,683; 31,746; 23,810; 19,569
34,483; 29,499; 22,727; 20,000
46,512; 37,037; 21,798
47,170; 38,462; 22,727; 18,762
38,168; 29,412; 22,890; 19,048
37,736; 30,581; 22,222
42,194; 30,769; 23,809; 19,231
38,462; 30,770; 20,619
Fig. 3. Suggested structure of the iron–citrazinic acid complex.
363 cm−1 are due to ν(Fe ← O) and ν(Fe–O) of the carboxy-
late group, respectively [19].The Nujol mull electronic absorp-
tion spectra of the brownish yellow [Fe(HL7)3]·2H2O iron–5-
(phenylazo)citrazinic acid complex (Table 2) gave three bands at
36,866, 23,426 and 20,779 cm−1. Its effective room temperature
magnetic moment value (μ = 5.96 BM) typified the existence of
an octahedral spatial structure of high-spin state [18]. Thus, the
structure of this complex is suggested as shown in Fig. 4.
1685 cm−1 are due to ν(C O) of the keto group of the citraz-
inate moiety. The band at 1613 cm−1 is due to ν(C N). The
two bands at 1460 and 1394 cm−1 are due to νas(COO−) and
νs(COO−), respectively, and confirm that the bonding occurred
via the bidentate carboxylate group. The presence of the low-
frequency bands at 432 and 365 cm−1 are due to ν(Fe ← O) and
ν(Fe–O) of the carboxylate group [19], respectively.
3.5. Mo¨ssbauer spectroscopy of iron complexes
The obtained isomer shift values for the iron–uracil and
iron–5-(m-nitrophenylazo)citrazinic acid complexes (δ = 0.3616
and 0.3488 mm/s), respectively, are less than those reported for
high-spin Fe(III) complexes (δ = 0.5–0.7 mm/s) [20], probably
due to the increase in the electron density at the nucleus or
due to the increase in covalency character of the bond between
Fe(III) and the oxygen atoms of the ligand. Such covalency
promotes d-electron transfer from the non-␴-bonding orbitals
of the central atom to ␲-acceptor orbitals of the ligand with
The Nujol mull electronic absorption spectra of the black
iron–citrazinic acid complex [Fe(HL6)3]·3H2O (Table 2) gave
charge transfer bands at 42,105 and 37,175 cm−1, and a visible
band at 21,739 cm−1 as well as the room temperature magnetic
moment value μ = 6.1 BM, which suggests an octahedral struc-
ture of high-spin state [18] as shown in Fig. 3.
3.4. Iron–5-(phenylazo)citrazinic acid complex
The IR spectra of the iron–5-(phenylazo)citrazinic acid com-
plex gave a well-defined band at 3460 cm−1 due to ν(OH) of
the citrazinic ring or to water molecules [14,15]. The bands at
3164, 3097, 1740, 1645, 1525, 1455 and 1410 cm−1 are due to
ν(N–H) of the citrazinate nucleus, hydrazo group, ν(C O) of
the keto group, ν(C N), ν(N N), νas(COO−) and νs(COO−),
respectively. The appearance of the two bands at 428 and
Fig. 4. Suggested structure of the iron–5-(arylazo)citrazinic acid complexes.
Fig. 2. Suggested structure of the iron(III)–uracil complex.

M.S. Masoud et al. / Spectrochimica Acta Part A 67 (2007) 662–668
665
the appropriate symmetry. The symmetry of the charge dis-
tribution around the iron nucleus results in a decrease in the
electronic field gradient at the position of the nucleus. This
leads to a decrease in the quadrupole splitting (ꢀEQ = 0.6624,
0.4384 and 0.7008 mm/s, respectively) for the iron–uracil,
iron–citrazinic acid and iron–m-(nitrophenylazo)citrazinic acid
complexes, respectively.
3.6. IR and electronic spectra of cobalt, nickel and copper
thiouracil complexes
The following equation is given to explain the mechanism of
complexation for 1:2 complexes:
M
2+ + 2H2L → M(HL)2 + 2H+,
Fig. 5. Suggested structure of metal–H2L2 complexes.
M = Co, Ni or Cu, L = L2, L3, L4 or L5
The IR spectra show the following characteristic features:
equals 1.8 BM. All of the data suggest a square-planar geometry
for such complexes.
Thus, the bidentate nature of thiouracil is suggested and its
complexation to the metal is given through oxygen and nitrogen
atoms as shown in Fig. 5.
(1) The ν(OH) band at 3526 cm−1 of thiouracil is shifted to
3324, 3370 and 3402 cm−1 in the corresponding Co(II),
Ni(II) and Cu(II) complexes, respectively, probably due
to the presence of water molecules in these complexes
[21].
(2) The doublet ν(N–H) bands of thiouracil at 3130 and
3042 cm−1 are absent in the cobalt complex. These are seen
at 3080 and 3160 cm−1 in the Ni(II) and Cu(II) complexes,
respectively, to identify M–N interaction [22].
3.7. IR and electronic spectra of 5-(phenylazo)thiouracil
and its metal complexes
The IR spectra show the following characteristic features:
(3) The ν(C O) and ν(C N) bands of thiouracil at 1696 and
1620 cm−1, respectively, are affected on complexation. In
the Co(II) complex, the first band is absent and the second
band is shifted to 1585 cm−1. However, in the Ni(II) com-
plex, both bands appeared at 1645 and 1580 cm−1. For the
(1) The broad ν(OH) band at 3470 cm−1 of the free ligand is
shifted to 3440, 3425 and 3420 cm−1 in the corresponding
Co(II), Ni(II) and Cu(II) complexes, respectively, probably
due to the presence of water molecules in these complexes
[21].
Cu(II) complex, the two bands shift to 1645 and 1595 cm−1
respectively. These data verify M–O interaction.
,
(2) The doublet ν(N–H) bands of the free ligand at 3110 and
3070 cm−1 shift to 3165 and 3050 cm−1 for the Co(II) com-
plex. However, both shift to 3155 and 3065 cm−1 for the
Ni(II) complex and for the Cu(II) complex to 3140 and
3050 cm−1. Thus, the imino nitrogen acts as a chelation
site [26,27].
(4) A new δ(OH) band appeared for the Co(II) complex at
1340 cm−1. This band is at 1350 and 1335 cm−1 in Ni(II)
and Cu(II) complexes, respectively [23].
The Nujol mull electronic absorption spectral bands of the
cobalt–thiouracil complex [Co(HL)2]·3H2O (Table 2), gave
charge transfer bands at 45,662 and 37,736 cm−1, and visible
d–d electronic spectral bands at 22,727 and 17,857 cm−1, and
a magnetic moment value of 5.1 BM, typical of an octahedral
spatial configuration of high-spin state [23,24].
The Nujol mull electronic absorption spectra of the green
Ni(II)–thiouracil complex [Ni(HL2)2]·3H2O (Table 2), gave
four bands at 47,170, 35,842, 22,988 and 19,685 cm−1, probably
due to octahedral spatial configuration. The [Ni(HL2)2]·3H2O
complex has an effective room temperature magnetic moment
value of 3.1 BM, reflecting spin–orbit coupling in an octahedral
configuration.
(3) The ν(C O) band of the free ligand at 1700 cm−1 is red-
shifted to 1660, 1675, and 1675 cm−1 in the corresponding
Co(II), Ni(II) and Cu(II) complexes, respectively, due to
complexation with the metal ion [27].
(4) The ν(N N) band of the ligand at 1510 cm−1 is shifted
slightly on complexation to 1520, 1530 and 1530 cm−1 in
the Co(II), Ni(II) and Cu(II) complexes, respectively, i.e. no
M-azo bonding exists [26,27].
(5) The δ(OH), γ(OH), δ(NH) and ν(C–N) modes of vibrations
of the free ligand are affected on coordination with the metal
ions to suggest that the oxygen and nitrogen N(3) atoms are
considered as centers for complexation [27].
The electronic absorption spectrum for the pale blue 1:2
copper(II)–thiouracil complex [Cu(HL2)2]·H2O, gave three
bands at 36,900, 27,397 and 19,763 cm−1, which may be
assigned to square-planar D2h symmetry of the Cu(II) complex
[25]. The effective room temperature magnetic moment value
The Nujol mull electronic absorption spectral bands
of the orange cobalt–5-(phenylazo)thiouracil complex
[Co(HL3)2(H2O)3] (Table 2), gave charge transfer bands at
39,683 and 31,746 cm−1. The visible d–d electronic spectral

666
M.S. Masoud et al. / Spectrochimica Acta Part A 67 (2007) 662–668
Fig. 6. Suggested structure of metal complexes 5-(arylazo)thiouracil.
bands observed at 23,810 and 19,569 cm−1 and the room
temperature magnetic moment value of 5.2 BM prove that this
complex exists in an octahedral complexation. The ligand is
bonded to Co(II) via the N(3) nitrogen adjacent to the carbonyl
group and via the oxygen atom. The electronic spectral data of
the Ni(II) complex [Ni(HL3)2(H2O)3] (Table 2), gave electronic
spectral bands at 34,483 and 29,499 cm−1 due to charge transfer
electronic transitions of the free ligand. The bands at 22,727
and 20,000 cm−1 and the effective room temperature magnetic
moment value of 3.2 BM support the octahedral complexation
for this complex [28]. The ligand is bonded to Ni(II) ion through
the oxygen and nitrogen N(3) atoms of the pyrimidine ring. The
5-(phenylazo)thiouracil–copper complex [Cu(HL3)2]·2H2O
gave bands at 46,512, 37,037 and 21,798 cm−1. The room
temperature magnetic moment for the copper complex is
1.9 BM. All of these data identified square-planar geometry
[27].
So, the bidentate nature of 5-(arylazo)thiouracil is suggested
and its complexation to the metal occurs through oxygen and
nitrogen atoms as shown in Fig. 6.
3.8. Electron spin resonance of copper complexes
The room temperature polycrystalline X-band ESR spec-
tra of Cu(II) (HL2)2·H2O, Cu(II) (HL3)2·2H2O and Cu(II)
(HL3)2·2H2O gave similar spectral patterns. All complexes are
of broad signals due to polymeric property. The data are of
isotropic nature with gs values at 2.14, 2.13 and 2.16 and A
values at 225, 250 and 200, respectively. [29]
3.9. Calculation of the coordination bond length
The calculations are based on recording the change in the
frequency of the band positions of the free ligand with respect
to its complexes. If it is suggests that the C N or C O groups
obtained by tautomerism are bonded to the metal ion, the coordi-
nation bond length (r) can be determined from the relation [30]
of (1).
ꢂ
ꢄ
ꢃ
ꢀ
ꢁ ꢀ
ꢁ
32πα
a2
νx=y − νx−y
2r
ꢀν =
exp −zπ
(1)
a

M.S. Masoud et al. / Spectrochimica Acta Part A 67 (2007) 662–668
667
Table 4
Thermal properties of iron–uracil, iron–citrazinic acid and cobalt–thiouracil complexes
Complex
Tm (◦C)
n
ꢀE (kJ mol−1
)
αm
Z (s−1
)
ꢀS* (kJ K−1 mol−1
)
ꢀH* (kJ mol−1
)
Weight of the
product (mg)
105
340
410
1.03
0.83
1.18
43.28
90.18
128.62
0.627
0.659
0.600
4.76
6.07
7.81
0.253
0.250
0.254
95.66
156.35
173.48
Fe(HL1)3·3H2O
18.1
90
220
365
440
1.02
0.93
0.82
1.18
35.93
33.896
98.73
0.626
0.735
0.664
0.602
4.10
2.80
6.39
0.254
0.260
0.225
0.247
92.19
128.02
162.67
176.32
Fe(HL6)3·3H2O
Co(HL2)2·3H2O
13.9
10.3
295.0
17.79
175
265
410
423
0.98
0.92
1.11
0.996
218.74
97.69
133.56
209.75
0.636
0.650
0.614
0.630
22.32
7.58
9.20
0.242
0.252
0.252
0.250
108.24
135.64
152.93
173.17
12.79
The Tm, n, ꢀE, αm, Z, ꢀS* and ꢀH* values are assigned to sample temperature at which the peak DTA deflection occurs, reaction order, activation energy, value
decomposed substance fraction, collision factor, entropies of activation and enthalpy of activation, respectively.
ꢀν is the shift in the oscillator frequency (νligand − νcomplex),
and α is the bond polarizability. The α values were computed
based on published data for complexes with a similar pyrimidine
compound to our complexes [31]. These are 48.61, 46.59 and
42.50 for cobalt, nickel and copper, respectively; a is the lattice
constant of the metal salt [a = 3.5, 3.524 and 3.614 for cobalt,
nickel and copper, respectively] [32]; νx=y the frequency of the
oscillator with a double bond [ν(C N) and ν(C O), respec-
tively]; νx−y the frequency of the oscillator with a single bond
[ν(C–N) and ν(C–O), respectively]; is the length of the oscil-
lator coordination to the metal ion. A literature survey for uracil
and pyrimidine complexes shows that the most prominent fea-
showed well-defined strong endothermic peaks. These peaks are
at 105, 340 and 410 ◦C for the [Fe(HL1)3]·3H2O complex, while
the [Fe(HL6)3]·3H2O complex gave four peaks at 90, 220, 365
and 440 ◦C. The peaks at 90 and 105 ◦C are due to dehydration
of water molecules in the outer sphere of the two complexes.
However, the second peak at 220 ◦C in the thermogram of the
iron–citrazinic acid complex may be due to the formation of
some decomposition products. On the other hand, the last two
peaks of both complexes represent their decomposition to Fe2O3
as a final decomposition product [17]. Calculation of the thermal
analysis data, ln ꢀt against 1000/TK plots, gives best-fit straight
lines for all peaks. The thermal parameters were calculated and
both complexes decompose via first order reactions.
˚
ture is that the N(3)–C(4) distance is 1.364 A and the O–C(4)
distance is 1.257 A [33].
The octahedral cobalt–thiouracil complex [Co(HL2)2·2H2O]
H2O gave an endothermic peak at 175 ◦C probably due to dehy-
dration of water molecules in the inner sphere of the complex.
The second peak at 265 ◦C may be due to the formation of some
decomposition products, while the last two peaks at 410 and
423 ◦C indicate the formation of CoO as the final product [22].
The ln ꢀt against 1000/T plots give straight lines for all peaks.
All the thermodynamic parameters were evaluated as shown in
Table 4.
˚
In the present work, one can use these values as the param-
eter in the calculation [33]. From such data, the r-values for the
bond between the metal and the nitrogen atom of the C N group
and the oxygen atom of the C–O group are computed [34].
Table 3 collects the values of the frequency shift, ꢀν
˚
(νligand − νcomplex), νx−y and r (A). For the thiouracil complexes,
the values of the calculated coordination bond length of the tran-
sition metal complexes decrease in the sequence either to the
oxygen or the nitrogen atom:
Ni < Co < Cu
References
For the octahedral Co(II) and Ni(II) complexes, the coordination
bond lengths are smaller than that of the square-planar Cu(II)
complex. In general, the M–N bond length is longer than that
of M–O. This is attributed to the electronegativity difference
between the oxygen and nitrogen atoms. Also, square-planar
complexes (e.g., copper) have longer bond lengths than octa-
hedral ones (e.g., cobalt and nickel). Thus, the geometry of the
complex compound plays a major factor for controlling such
trends. In turn, the flexibility factors assist such behaviour.
[1] E.R. Tucci, N.C. Li, J. Inorg. Nucl. Chem. 25 (1963) 17.
[2] H.R. Lindsay, J.C. Ramine, M.Y. Wong, Arch. Biochem. Biophys. 126
(1968) 812;
H.R. Lindsay, J.C. Ramine, M.Y. Wong, Chem. Abstr. 69 (1968) 94735t.
[3] A. Ali, S.E. Livingstone, Coord. Chem. Rev. 13 (1974) 102.
[4] J. Bauerova, K. Sebesta, F. Sorm, Z. Sormova, Collec. Czech. Chem. Com-
mun. 25 (1960) 2906;
J. Bauerova, K. Sebesta, F. Sorm, Z. Sormova, Chem. Abstr. 55 (1961)
4672a.
[5] T. Shimura, T. Tomohiro, H. Okuno, Inorg. Chim. Acta 155 (1989) 21.
[6] S. Sakamoto, M. Matsubara, H. Kudo, S. Suzuki, L. Liu, T. Nakayama, T.
Mitamura, H. Nagasawa, Anticancer Res. 14 (1994) 2515;
S. Sakamoto, M. Matsubara, H. Kudo, S. Suzuki, L. Liu, T. Nakayama, T.
Mitamura, H. Nagasawa, Chem. Abstr. 122 (1995) 281582r.
[7] L. Novotny, A. Vachalkova, Neoplasma 38 (1991) 223;
L. Novotny, A. Vachalkova, Chem. Abstr. 115 (1991) 105588u.
3.10. Differential thermal analysis (DTA)
The DTA curves of the octahedral iron complexes obtained
from H2L1 and H2L6 of the stoichiometry [Fe(HL)3]·3H2O

668
M.S. Masoud et al. / Spectrochimica Acta Part A 67 (2007) 662–668
[8] A. Vachalkova, L. Novotony, Neoplasma 37 (1990) 555;
[20] S.J. Lippard, J.M. Berg, Principles of Bioinorganic Chemistry, University
Science Books, Mill Valley, CA, 1994, p. 91.
[21] G.S. Pandemy, P.C. Nigum, V. Agrwala, Inorg. Chem. 39 (1977)
1877.
A. Vachalkova, L. Novotony, Chem. Abstr. 114 (1991) 223757a.
[9] A. Tsunoda, M. Shibusawa, N. Yasuda, K. Nakao, M. Kusamo, Anticancer
Res. 14 (1994) 2637;
A. Tsunoda, M. Shibusawa, N. Yasuda, K. Nakao, M. Kusamo, Chem.
Abstr. 122 (1995) 281259r.
[22] M.S. Masoud, S.A. Abou El Enein, H.M. Kamel, Ind. J. Chem. A 40 (2002)
297.
[10] A. Lee, J. Dvovak, V. Chmatal, P. Kuthan, J. Vencl, Ger. Offen. 755 (1978)
683;
[23] K. Jorgensen, Absorption Spectra and Chemical Bonding in Complexes,
Pergamon Press, 1962.
A. Lee, J. Dvovak, V. Chmatal, P. Kuthan, J. Vencl, Chem. Abstr. 89 (1978)
112366t.
[24] F. Albert Cotton, Advanced Inorganic Chemistry, 3rd ed., Wiley Eastern
Limited, London, 1979.
[11] K. Satorus, K. Shigeharu, K. Masayuki, H. Mooko, Japanese Patent, Kon-
ishiroku Photo Industry Co., Ltd., 28 January, 1988;
K. Satorus, K. Shigeharu, K. Masayuki, H. Mooko, Chem. Abstr. 111
(1989) 184060i.
[12] Y. Kazuyoshim I. Yoshio, Japanese Patent, Fuji Photo Film Co. Ltd., Jpn.
Kokai Tokkyo Koho, 30 June, 1992;
Y. Kazuyoshi, I. Yoshio, Appl. Chem. Abstr. 118 (1993) 4979j.
[13] K. Hiroaki, N. Shinsaku, K. Tsugio, K. Shigeharu, Japanese Patent, Konica
Co., Jpn., Kokai Tokkyo Koho, 13 October, 1988;
K. Hiroaki, N. Shinsaku, K. Tsugio, K. Shigeharu, Chem. Abstr. 114 (1991)
111793c.
[25] L. Dubick, R.L. Martin, Inorg. Chem. 5 (1966) 2203.
[26] M.S. Masoud, O.H. Abdel-Hamid, Z.M. Zaki, Trans. Met. Chem. 19 (1994)
21.
[27] M.S. Masoud, S.S. Haggag, Z.M. Zaki, M. El-Shabasy, Spectrosc. Lett. 27
(1994) 775.
[28] M.S. Masoud, A.K. Ghonaim, S.A. Abou El Enein, A.A. Mahmoud, J.
Coord. Chem. 55 (2002) 79.
[29] M.S. Masoud, E.A. Khalil, A.M. Hafez, A.F. El-Husseiny, Spectrochim.
Acta A 61 (2004) 989.
[30] G. Karagonins, O. Peter, Z. Electrochem. Ber. Bunsenges. Phys. Chem. 63
(1959) 1170;
[14] M.S. Masoud, Z.M. Zaki, Trans. Met. Chem. 13 (1988) 321.
[15] G. Schwarzenbach, in: H. Irving (Ed.), Complexometric Titration, Methnen
Co., London, 1957.
G. Karagonins, O. Peter, Chem. Abstr. 54 (1960) 5241F.
[31] C.H. Macgillavry, G.D. Riech, in:K. Lansdale(Ed.), PhysicalandChemical
Tables, vol. 3, 1962.
[16] M.S. Masoud, E.A. Khalil, A.M. Hindawy, A.M. Ramadan, Can. J. Anal.
Sci. Spectrosc. 50 (4) (2005) 175.
[32] R.K. Parashar, R.C. Sharma, A. Kuma, A. Kumor, G. Mohan, Inorg. Chim.
Acta 151 (1988) 201.
[17] M.S. Masoud, S.S. Haggag, E.A. Khalil, Nucleosides, Nucleosides
Nucleotides Nucleic Acids 25 (2006) 73.
[33] H. Rauter, E. Hillgeris, E. Andren, B. Lippert, J. Am. Chem. Soc. 116
(1994) 616.
[18] D.W. Barnum, J. Inorg. Nucl. Chem. 22 (1961) 183.
[19] M.S. Masoud, A.M. Ibrahim, E.A. Khalil, A. El-Marghany, Bull. Fac. Sci.
Assist. Univ. 25 (3-B) (1996) 7.
[34] M.M. Abou Sekkina, S.M. El-Hebawy, Proc. Indian Natl. Sci. A 51 (1985)
559;
M.M. Abou Sekkina, S.M. El-Hebawy, Chem. Abstr. 102 (1985) 124377n.