Synthesis, structural characterization and antimicrobial activity evaluation of metal complexes of sparfloxacin

Article abstract of DOI:10.1016/j.saa.2011.07.072

The synthesis and characterization of binary Cu(II)- (1), Co(II)- (2), Ni(II)- (3), Mn(II)- (4), Cr(III)- (5), Fe(III)- (6), La(III)- (7), UO ₂(VI)- (8) complexes with sparfloxacin (HL₁) and ternary Cu(II)- (9), Co(II)- (10), Ni(II)- (11), Mn(II)- (12), Cr(III)- (13), Fe(III)- (14), La(III)- (15), UO₂(VI)- (16) complexes with sparfloxacin (HL₁) and dl-alanine (H₂L₂) complexes are reported using elemental analysis, molar conductance, magnetic susceptibility, IR, UV-Vis, thermal analysis and ¹H-NMR spectral studies. The molar conductance measurements of all the complexes in DMF solution correspond to non-electrolytic nature. All complexes were of the high-spin type and found to have six-coordinate octahedral geometry except the Cu(II) complexes which were four coordinate, square planar and U- and La-atoms in the uranyl and lanthanide have a pentagonal bipyramidal coordination sphere. The antimicrobial activity of these complexes has been screened against two Gram-positive and two Gram-negative bacteria. Antifungal activity against two different fungi has been evaluated and compared with reference drug sparfloxacin. All the binary and ternary complexes showed remarkable potential antimicrobial activity higher than the recommended standard agents. Ni(II)- and Mn(II) complexes exhibited higher potency as compared to the parent drug against Gram-negative bacteria.

Full text of DOI:10.1016/j.saa.2011.07.072

Spectrochimica Acta Part A 82 (2011) 414–423
Contents lists available at ScienceDirect
Spectrochimica Acta Part A: Molecular and
Biomolecular Spectroscopy
journal homepage: www.elsevier.com/locate/saa
Synthesis, structural characterization and antimicrobial activity evaluation of
metal complexes of sparfloxacin
Nadia E.A. El-Gamel∗, M.A. Zayed
Chemistry Department, Faculty of Science, Cairo University, Gamma Street 12613, Giza, Egypt
a r t i c l e i n f o
a b s t r a c t
Article history:
Received 4 June 2011
Received in revised form 17 July 2011
Accepted 20 July 2011
The synthesis and characterization of binary Cu(II)- (1), Co(II)- (2), Ni(II)- (3), Mn(II)- (4), Cr(III)- (5), Fe(III)-
(6), La(III)- (7), UO2(VI)- (8) complexes with sparfloxacin (HL1) and ternary Cu(II)- (9), Co(II)- (10), Ni(II)-
(11), Mn(II)- (12), Cr(III)- (13), Fe(III)- (14), La(III)- (15), UO2(VI)- (16) complexes with sparfloxacin (HL1)
and dl-alanine (H2L2) complexes are reported using elemental analysis, molar conductance, magnetic
susceptibility, IR, UV–Vis, thermal analysis and 1H-NMR spectral studies.
Keywords:
The molar conductance measurements of all the complexes in DMF solution correspond to non-
electrolytic nature.
Sparfloxacin (HL1)
dl-Alanine (H2L2)
Binary and ternary complexes
Spectral studies
Thermal analysis
Antimicrobial activity
All complexes were of the high-spin type and found to have six-coordinate octahedral geometry except
the Cu(II) complexes which were four coordinate, square planar and U- and La-atoms in the uranyl
and lanthanide have a pentagonal bipyramidal coordination sphere. The antimicrobial activity of these
complexes has been screened against two Gram-positive and two Gram-negative bacteria. Antifungal
activity against two different fungi has been evaluated and compared with reference drug sparfloxacin.
All the binary and ternary complexes showed remarkable potential antimicrobial activity higher than
the recommended standard agents. Ni(II)- and Mn(II) complexes exhibited higher potency as compared
to the parent drug against Gram-negative bacteria.
© 2011 Elsevier B.V. All rights reserved.
1. Introduction
Metal coordination to biologically active molecules can be used
in order to enhance their biological activity; therefore, numerous
The development of metal complexes as artificial nucleases is
an area of interest. Floroquinolone derivatives possess a broad
spectrum of activity against various pathogenic microorganisms,
which are resistant to aminoglycosides, penicillins, cephalosporins,
tetracyclines and other antibiotics. This class of compounds, when
compared to existing bactericidal drugs, shows improved phar-
macokinetic properties and a broad spectrum of activity against
parasites, bacteria, and mycobacteria, including resistant strains;
in addition to that they displayed significant in vitro antibacterial
activity against many Gram-positive and Gram-negative bacteria
through inhibition of their DNA gyrase [1].
Sparfloxacin (HL1, 5-amino-1-cyclopropyl-7-(3,5-dimethyl-
1-piperazinyl)-6,8-difluoro-1,4-dihydro-4-oxo-3-quinoline-
carboxylic acid) (Fig. 1) is widely used in the treatment of urinary
tract infections and is reported to exhibit higher activity against
the major respiratory pathogens and typical pathogens that cause
pneumonia [2].
studies regarding the interaction between quinolones with sev-
eral metallic cations have reported in the literature. Some of these
metal complexes represented potential antibacterial activity [3],
for example Co(II) and Cu(II) complexes with sparfloxacin exhib-
ited a significant enhancement against several pathogenic bacteria
than sparfloxacin [4].
The antibacterial and antifungal properties of palladium and
platinum complexes with sparfloxacin have also been reported,
where the complexes showed activity against Mycobacterium tuber-
culosis strain [5].
Fluoroquinolones behaved as bidentate ligands binding to the
metallic ion through the carboxylate and carbonyl oxygens as
observed in case of complexes with metal ions such as Co(II), Ni(II),
Zn(II), Cd(II), Mn(II) and Cu(II) [6]. Nevertheless, it was mentioned
that the platinum complexes were chelated with fluoroquinolones
via piperazine nitrogen atoms which is much less common [5a].
Recently some transition and earth metal complexes of sparfloxacin
were prepared and characterized and the antifungal activity were
evaluated where Fe(II), and Cd(II) complexes showed remarkable
activity [7].
In this context, the present work describes the synthesis of
binary and ternary complexes of sparfloxacin with dl-alanine
∗
Corresponding author. Tel.: +20 20112363028.
E-mail address: nadinealy@hotmail.com (N.E.A. El-Gamel).
1386-1425/$ – see front matter © 2011 Elsevier B.V. All rights reserved.
doi:10.1016/j.saa.2011.07.072

N.E.A. El-Gamel, M.A. Zayed / Spectrochimica Acta Part A 82 (2011) 414–423
415
2.3. Antimicrobial activity
CH3
Antimicrobial activity of the tested samples was determined
using a modified Kirby-Bauer disc diffusion method [8]. 100 ␮l of
the test bacteria/fungi were grown in 10 ml of fresh media until
they reached a count of approximately 108 cells/ml for bacteria
or 105 cells/ml for fungi [9]. 100 ␮l of microbial suspension was
spread onto agar plates. Plates inoculated with filamentous fungi
as Aspergillus flavus, Candida albicans at 25 ◦C for 48 h; Gram (+) bac-
teria as Staphylococcus aureus; Gram (−) bacteria as Escherichia coli;
they were incubated at 35–37 ◦C for 24–48 h and yeast as C. albi-
cans incubated at 30 ◦C for 24–48 h and, then the diameters of the
inhibition zones were measured in millimeters [9]. Standard discs
of Tetracycline (antibacterial agent), Amphotericin B (antifungal
agent) served as positive controls, while filter discs impregnated
with 10 ␮l of solvent (distilled water, chloroform, DMSO) were used
as a negative control.
HN
F
8
1
N
F
N
8a
4a
7
2
H3C
6
3
5
4
COOH
NH2
O
Fig. 1. The structure of sparfloxacin (HL1).
2.4. Synthesis of the complexes
with various metal ions like Cu(II), Co(II), Ni(II), Mn(II), Cr(III),
Fe(III), La(III) and UO2(VI), in order to continue the exploration of
the coordination behavior adopted by fluoroquinolones. For the
characterization of the compounds the following spectroscopic
and analytical techniques were employed: elemental analyses, IR
and 1H-NMR spectroscopy, UV–Vis, solid reflectance, magnetic
moment, ESR and thermogravimetric analysis. The biological activ-
ity of the parent ligand and its metal complexes is evaluated where
high antimicrobial is recorded.
Metal complexes were synthesized by addition of
a hot
water–ethanolic solution (60 ◦C) of the metal chlorides for Cu(II),
Co(II) Ni(II), Mn(II), Cr(III) and Fe(III), nitrate for UO2(VI) and
La(III) (25 ml, 0.1 mmol) to a hot ethanolic solution of HL1 (25 ml,
0.2 mmol in case of binary or 0.1 mmol in case of ternary complexes)
and dl-alanine (0.1 mmol). The resulting mixture was stirred under
reflux for 2 h and left to cool, whereby the complexes precipitated
as fine powders. The solid complexes were filtered, washed with
ethanol, then with diethyl ether, and dried in a vacuum desiccator
over anhydrous calcium chloride.
2. Experimental
3. Results and discussion
2.1. Materials and reagents
The found and calculated percentages of elemental analysis
(CHN) data are in well agreement with each other and prove the
suggested molecular formulas of the complexes (Tables 1 and 2).
The molar conductance measurements of all the complexes in DMF
solution were in the range of 18.10–11.99 ꢀ2 cm−1 mol−1. These
relatively low values indicate non-electrolytic nature of the com-
plexes [10].
All chemicals used were of the analytical reagent grade (AR)
and of highest purity available. They included sparfloxacin, dl-
alanine, Cu(II)Cl2·2H2O, La(NO3)3·6H2O, UO2(NO3)2·6H2O, and
MnCl2 (Sigma), Co(II)Cl2·6H2O, CrCl3·6H2O and Ni(II)Cl2·6H2O
(BDH); FeCl3·6H2O (Prolabo) were used. Absolute ethyl alcohol,
diethyl ether (Adwic), yeast extract and agar (Sigma) were also
used. De-ionized water collected from all glass equipments was
usually used in all preparations.
3.1. IR spectral studies
2.2. Instruments
All the characteristic absorption bands of the complexes are
represented in Table 3. The pyridone stretch ꢁ(C O)p in HL1
appears at 1641 cm−1 and the asymmetric and symmetric stretch-
ing ꢁ(COO)carb appeared at 1716 and 1371 cm−1, respectively.
The pyridone band is shifted to lower frequency (3–9 cm−1) and
(3–11 cm−1) in binary and ternary complexes, which indicates the
involvement in the chelation [3a]. The participation of the carboxy-
late O atom in the binary and ternary complexes is confirmed by
low shift in position of these bands to (1613–1555 cm−1) and high
shift to (1400–1380 cm−1) for asymmetric and symmetric bands,
respectively [3a].
FTIR spectra were obtained from dispersions in KBr using a
Perkin-Elmer FT-IR type 1650 spectrophotometer. The spectra were
collected in the range from 200 to 4000 cm−1 with a resolution of
2 cm−1. UV–Vis spectrophotometric measurements were carried
out using automated spectrophotometer UV-Vis Thermo Fischer
Scientific Model Evolution 60 ranged from 200 to 900 nm. The molar
magnetic susceptibility was measured on powdered samples using
the Faraday method. The diamagnetic corrections were made by
Pascal’s constant and Hg[Co(SCN)4] was used as a calibrant. The
solid reflectance spectra were performed on a Shimadzu 3101pc
spectrophotometer. The ESR spectra were recorded on JES FA-300
spectrophotometer at room temperature. Molar conductance was
measured on a ELICO(CM82T) conductivity bridge. 1H-NMR spectra
were recorded on a BRUKER DPX 400 instrument at room temper-
ature (d6-DMSO solution with TMS as internal standard). Thermal
analyses of the complexes were carried out using a Shimadzu
TGA-50H and DTA-50H thermogravimetric analyzer in a dynamic
nitrogen atmosphere (flow rate 20 ml min−1) with a heating rate of
10 ◦C min−1. The percentage weight loss was measured from ambi-
ent temperature to 1000 ◦C. Highly sintered ␣-Al2O3 was used as
reference.
IR spectrum of H2L2 shows sharp bands at 1594 and 1410 cm−1
assigned to the asymmetric and symmetric stretching vibrations
of the carboxylate moiety, respectively. These bands are shifted to
lower (3–63 cm−1) or higher (6–22 cm−1) frequencies, indicating
that H2L2 coordinates to the metal ions via deprotonated carboxy-
late [11].
The IR spectra of all complexes exhibit a broad split band
between 3400 and 3150 cm−1 which the N–H stretching vibration
of the piperazinyl moiety [12,13], therefore it is quite difficult to
detect the coordination vibration band of NH2 of H2L2; therefore we
focused on in-plane bending, (NH2) vibration and the shift of this
band from 1521 cm−1 to 1502–1517 cm−1 indicates participation in

416
N.E.A. El-Gamel, M.A. Zayed / Spectrochimica Acta Part A 82 (2011) 414–423
Table 1
Analytical and physical data of binary metal complexes.
Compound
Color (% yield)
m.p (◦C)
Found (calculated) (%)
ꢂeff (B.M.)
ꢃm (ꢀ2 mol−1 cm−1
)
C
H
N
Sparfloxacin (HL1)
C19H22F2N4O3
Yellow
240
57.95
(58.16)
51.50
(51.73)
52.16
(52.00)
52.41
(52.01)
51.94
(52.24)
50.00
(50.36)
43.35
(43.98)
50.32
5.55
(5.65)
5.45
(5.25)
5.92
(5.28)
5.22
(5.28)
5.59
(5.31)
5.10
(5.12)
4.55
(4.66)
5.20
14.20
(14.28)
12.53
(12.70)
12.03
(12.77)
12.63
(12.77)
12.55
(12.82)
12.42
(12.36)
12.69
(12.15)
12.02
–
–
[Cu(II)(L1)2]·2H2O
Green
(75)
Olive green
(85)
Yellowish green
(75)
Mustard
(88)
Dark green
(85)
Golden yellow
(87)
Brown
(72)
290
1.74
18.10
15.26
12.29
14.25
13.95
13.32
17.24
15.14
C
38H46CuF4N4O8
[Co(II)(L1)2·2H2O]
C38H46CoF4N8O8
[Ni(II)(L1)2·2H2O]
C38H46F4N8NiO8
[Mn(II)(L1)2·2H2O]
>300
>300
286
4.9
3.23
5.73
C
38H46F4MnN8O8
[Cr(III)(L1)2Cl·H2O]·H2O
>300
270
3.87
C
38H46ClCrF4N8O8
[La(III)(L1)2·NO3·H2O]·2H2O
Diamagnetic
5.90
C38H48F4LaN9O12
[Fe(III)(L1)2·Cl·H2O]·H2O
C38H46ClF4FeN8O8
289
(50.15)
42.50
(42.62)
(5.09)
4.20
(4.14)
(12.31)
10.02
(10.46)
[UO2(VI)(L1)2·H2O]
Orange
(65)
279
Diamagnetic
C38H44F4N8O9U1
Table 2
Analytical and physical data of ternary metal complexes.
Compound
Color (% yield)
m.p (◦C)
Found (calculated) (%)
ꢂeff (B.M.)
ꢃm (ꢀ2 mol−1 cm−1
)
C
H
N
[Cu(II)(L1)(HL2)]·2H2O
Green
>300
290
45.95
(45.63)
46.23
(46.00)
46.16
(46.02)
46.41
(46.32)
42.94
(43.82)
36.00
(36.88)
43.35
(43.55)
43.64
5.25
(5.40)
5.45
(5.44)
5.42
(5.44)
5.28
(5.48)
5.29
(5.18)
4.19
(4.36)
5.55
(5.15)
3.29
12.12
(12.09)
12.53
(12.19)
12.03
(12.20)
12.63
(12.28)
11.55
(11.62)
11.42
(11.73)
11.69
(11.54)
10.52
1.81
–
C22
H31CuF2N5O7
[Co(II)(L1)(HL2)·2H2O]
Green
(75)
Light green
(65)
Mustard yellow
(78)
Dark green
(85)
Dark yellow
(95)
Reddish brown
(89)
Orange
(87)
5.10
17.66
16.94
11.99
14.95
13.95
17.78
15.24
C22H31CoF2N5O7
[Ni(II)(L1)(HL2)·2H2O]
C22H31F2N5NiO7
>300
>300
>300
>300
>300
>300
3.36
[Mn(II)(L1)(HL2)·2H2O]
5.41
C22H31F2MnN5O7
[Cr(III)(L1)(HL2)Cl·H2O]·H2O
3.85
C
22H31ClCrF2N5O7
[La(III)(L1)(HL2)·NO3·H2O]·H2O
22H31F2LaN6O10
[Fe(III)(L1)(HL2)Cl·H2O]·H2O
22H31ClF2FeN5O7
[UO2(VI)(L1)(HL2)·H2O]
22H29F2N6O7U1
Diamagnetic
5.92
C
C
Diamagnetic
C
(43.52)
(3.82)
(10.98)
complex formation [14]. New bands in the complexes at 599–571
and 560–518 cm−1 can be assigned to the (M–O) stretching vibra-
tions of the pyridone and carboxylate groups, respectively [14]. In
ternary complexes bands in the 543–524 cm−1 and 430–415 cm−1
regions are attributed to the (M–O)carb and (M–N)amino stretch-
ing vibration of the carboxylate and amino groups [14]. There is
a paucity of such assignments in the literature concerning lan-
thanide complexes, new band at 620 cm−1 and the new shoulder
at 524 cm−1 and 361 cm−1 can be due to the La–O, La–N interac-
tions, respectively [15]. Several bands are observed which are at
Table 3
IR spectral data (4000–400 cm−1) of binary and ternary complexes.
Compound
␯(C O)p
ꢁ(COO) (asym)
ꢁ(COO) (sym.)
ꢄNH2
ꢁ(COO) amino
ꢁ(COO) amino
ꢁ(M–O)carb
ꢁ(M–O)p
ꢁ(M–O)carb
ꢁ(M–N)
(sym.)
(asym)
HL1
1
2
3
4
5
6
7
8
1641
1632
1638
1635
1634
1634
1636
1638
1638
1634
1635
1632
1633
1633
1630
1638
1636
1716
1555
1609
1612
1613
1611
1610
1608
1611
1612
1611
1606
1607
1608
1607
1610
1610
1371
1384
1384
1384
1390
1384
1385
1388
1385
1400
1387
1387
1383
1389
1383
1380
1383
–
–
–
–
–
–
–
–
–
–
–
–
–
–
–
–
–
––
1578
1580
1575
1555
1560
1540
1583
1583
–
–
––
–
–
–
–
–
–
1419
1420
1419
1408
1418
1402
1512
1422
–
529
525
526
528
523
528
525
518
520
520
540
533
530
533
520
528
–
587
597
598
588
593
571
580
586
588
592
599
590
585
590
588
589
–
–
–
–
–
–
–
–
–
540
530
534
532
543
548
–
–
–
–
–
–
–
–
–
420
425
419
415
418
418
426
430
9
1515
1517
1509
1508
1505
1508
1512
1502
10
11
12
13
14
15
16
524, 620
543

N.E.A. El-Gamel, M.A. Zayed / Spectrochimica Acta Part A 82 (2011) 414–423
417
1105, 1490 and 1300 cm−1 can be assigned for symmetric bending,
symmetric and asymmetric stretching of bidentate nitrate anions,
which confirm that the NO3 anion is used as bi-dentate in lan-
thanide complexes [16].
In uranyl chelates two intense sharp bands (920, 917 cm−1);
(775, 776 cm−1) were reliably assigned to asymmetric and symmet-
ric stretching vibrations of O
U O in binary and ternary chelates,
respectively [17]. In Cr(III) and Fe(III) complexes, medium bands
below 300 cm−1 are observed due to ꢁ(MCl) vibrations and fur-
ther substantiate the mode of coordination as suggested by NMR
spectral studies.
3.2. Electron spin resonance spectra
The ESR spectra for copper complex at room temperature gave
the corresponding parameters of spin-Hamiltonian (g⊥ = 2.090,
2.071; g11 = 2.406, 2.342) for binary and ternary complexes;respec-
tively. This trend, g11 > g⊥ observed for the copper complexes
Fig. 2. UV–Vis absorption spectra (2.5 × 10−4 M) of binary complexes.
suggests that the unpaired electron lies in the dx
orbitals giving
2
2
−y
2B1g as the ground state with g11 > g⊥ > 2; the ESR parameters of the
complexes coincide well with related systems which suggest that
the complexes have square-planar geometry and the system is axi-
ally symmetric [18]. According to Hathaway [19] G = gII − 2/g⊥ − 2;
if the value of G is larger than four, exchange interaction is negligi-
ble because the local tetragonal axes are misaligned. The g values
are 4.5 and 4.8 for the binary and ternary complexes, respec-
tively, which suggests that the local tetragonal axis is aligned
the range (13,260–19,415 cm−1), (24,930–27,390 cm−1
)
and
4
4
4
4
37,035–38,025 cm−1 due to A2g(F) → T2g(F), A2g(F) → T1g(F)
4
4
and A2g(F) → T1g(P) d–d transitions, respectively. These bands
suggest an octahedral geometry for the Cr(III) complexes [22].
Iron(III) complexes showed magnetic moment 5.90 and 5.92 B.M.
for 6 and 14, respectively corresponding to high spin octahedral
geometry [17]. The reflectance spectra exhibited a band at the
6
6
range of 21,295–20.904 cm−1, assigned to the A1g → T2g(G)
parallel or slightly misaligned and consistent with dx
state.
ground
2
2
transition in octahedral geometry of the complex [36]. The
−y
6A1g → T1g transition split into two bands at the range of
6
17,543–18.201 cm−1 and 14,995–12.141 cm−1
. These data are
consistent with octahedral geometry [22]. The spectra showed
bands at the range of 30.665–31.680 cm−1 attributed to ligand to
metal charge transfer [22]. La(III) and UO2(VI) and complexes are
diamagnetic.
3.3. Magnetic susceptibility and electronic spectra measurements
The magnetic moment values are calculated and reported
in Tables 1 and 2. Cu(II) complexes have ꢂeff values of 1.94
and 1.92 B.M., for 1 and 9, respectively, indicating a square
planar geometry; this was confirmed by the presence of one
band in the spectrum at 13.385 cm−1 with two shoulders on
either side at 18.860 and 11.900 cm−1. These are assigned to
3.4. UV–Vis absorption spectra
The UV–Vis absorption spectra of the binary and ternary com-
plexes (2.5 × 10−4 M) in ethanol–water mixture are shown in
Figs. 2 and 3, respectively. The absorption spectrum of HL1 showed
three absorption peak at 224, 292 and 366 nm and its molar absorp-
tivities were in range of (1.36–4.25) × 103 L mol−1 cm−1; these
peaks could be assigned to ␲–␲* and n–␲* transition within the
organic ligand [24]. In all binary complexes the maximum absorp-
tion peaks had relatively high bathochromic shifts (ꢅꢆ was in
2
2
2
2
2B1g → A1g
,
2B1g → B2g and B1g → E2g transitions, respec-
tively [20]. Cobalt(II) complexes have values of 4.9 and 5.1 B.M.
for 2 and 10, respectively, which indicate a high-spin octahe-
dral configuration [17,21]. These complexes displaced three
bands at 15.690–15.365, 17,467–18.019 and 21,930–22.120 cm−1
assigned to 4T1g(F) → T2g(F)(ꢁ1), 4T1g(F) → A2g(F)(ꢁ2) and
4
4
4T1g(F) → T2g(P)(ꢁ3) transitions, respectively, confirming the
4
octahedral geometry. Bands at 24.580–24.650 cm−1 could be
attributed to charge transfer bands [22]. Nickel(II) complexes
have a magnetic momentum of 3.23 and 3.36 B.M. for 3 and
11, respectively indicating a high spin octahedral configuration
[17,21]. The electronic spectra give three bands at ꢁ1 = 16.095 cm−1
,
ꢁ2 = 18.219 cm−1 and ꢁ3 = 21.410 cm−1; these bands are assigned
3
3
3
4
3
to A2g → T2g
,
3A2g(F) → T1g(F) and A2g → T1g(P) transitions,
respectively. Bands at 24.550–24.680 cm−1 were observed, which
could be attributed to LMCT [22]. Mn(II) complexes have ꢂeff
values of 5.73 and 5.84 B.M., for 4 and 12, respectively indicating
a high spin octahedral [22]. The electronic spectra of Mn(II) com-
plexes show adsorption bands in the range (16,940–19,530 cm−1),
(22,218–24,268 cm−1) and (26,310–27,765 cm−1). These adsorp-
6
4
6
4
tion bands are assigned to A1g → T1g(4G), A1g → T2g(4G)
6
4
4
and A1g → Eg, A1g(4G) transitions, respectively. These bands
suggest the Mn(II) complexes posses an octahedral geometry
[22]. ꢂeff values for Cr(III) complexes have been reported to be
3.87 and 3.85 for 5 and 13, respectively, this is in accordance
to high spin octahedral geometry [22,23]. The solid reflectance
spectra of Cr(III) complexes exhibited absorption bands in
Fig. 3. UV–Vis absorption spectra (2.5 × 10−4 M) of ternary complexes.

418
N.E.A. El-Gamel, M.A. Zayed / Spectrochimica Acta Part A 82 (2011) 414–423
the range of 1–16 nm) and their molar absorptivities were in the
range of (0.25–4.18) × 103 L mol−1 cm−1, this can be attributed to
the reaction of the metal ions with the ketonic oxygen and car-
boxylic oxygen in the conjugated system of the drug, which directly
resulted in the changes of absorption spectra and bathochromic
shifts of the maximum absorption peaks. Whereas in the case of
ternary complexes new absorption peaks appeared in 340–352 nm
and all the other absorption peaks are bathochromic shifted in
the range of (1–11 nm) and their molar absorptivities were in the
range of (0.59–3.61) × 103 L mol−1 cm−1, The shift of ꢆmax to higher
values (bathochromic shift) and the decrease of the absorbance
intensities in all of the mentioned complexes are indication to
the complexation behavior towards metal ions. It was observed
that in the case of lanthanides the energy of f–f transitions in
the complexes is slightly reduced, this may be due to covalent
interaction of 4f orbitals with vacant ligand orbitals or increased
nuclear shielding of f-orbitals due to slight L–M electron transfer
[25].
position of the ligand between 300 and 650 ◦C leaving NiO and CoO
as a residue. This decomposition is connected with endothermic
process. In the second and third steps in cobalt and nickel ternary
complexes the decomposition of the ligand took place in the range
of 200–750 ◦C leaving NiO and CoO as a residue. This decomposition
is connected with exothermic process. In the last steps of decom-
position, loss of the ligands took place leaving residual NiO and CoO
between 200 and 750 ◦C. DTA shows three successive exothermic
signals.
On the other hand the decomposition of manganese complexes
exhibited three successive decomposition steps, the first step in
the temperature range of 50–300 ◦C in which the complexes lose
the coordinated water with an endothermic effect. The subsequent
second and third steps (300–650 ◦C) correspond to the removal
of the organic part of the ligand leaving metal oxide (MnO) as a
residue. These steps are accompanied by two successive endo- and
exothermic process.
TG of Cr(III)- and Fe(III) complexes showed liberation of
hydrated and coordinated water and hydrated water molecules in
the first and second steps with an endothermic effects. The third
fourth and fifth steps are accounted for decomposition of ligands
and loss of 0.5Cl2 leaving Cr2O3 and Fe2O3 residues, these three
steps are accompanied by two endothermic and three exothermic
DTA signals.
The TG of binary La(III) complex showed that the decomposi-
tion occurred in three closed steps within the temperature range
of 50–650 ◦C. The first step in the range of 50–150 ◦C may be due
to the liberation of hydrated water molecules accompanied with
an endothermic effect. The second and third decomposition steps
at 150–650 ◦C correspond to loss of coordinated water, liberation
of NO2 and 0.5O2 and decomposition of the ligand and forma-
tion of La2O3 as a final residue. These steps are connected with
two endothermic effects. The thermal decomposition of ternary
lanthanide complex showed similar behavior in four successive
steps within the range of 50–630 ◦C, the first and second steps are
responsible for removal of coordinated and hydrated water with
an exothermic process. The third and fourth steps are accompanied
with liberation of NO2 and 0.5O2 and decomposition of ligands in
the range of 240–630 ◦C. DTA represented an exothermic signal.
Uranyl binary and ternary complexes thermally decomposed
in three and fifth decomposition steps, respectively. The first step
in the range of 50–300 ◦C can be assigned to liberation of coordi-
nated water with endothermic effects. The second till last mass loss
steps in the range of 300–650 ◦C are accounted for the decompo-
sition of both ligands molecules and formation of UO2 as a final
residue, this decomposition is followed by two endothermic and
two exothermic signals.
3.5. 1H-NMR spectra
1H-NMR spectra of uranyl and lanthanide complexes are
recorded in DMSO-d6. The ligand HL1 showed signals at 8.76 ppm
due to NH2 proton, which did not show any appreciable change
in the complexes, confirming non-coordination of NH2 moiety.
The signals for the aliphatic piperazine and the methyl protons
are practically unchanged since they lie far from the binding site
of the ligand, also the cyclopropane protons remain practically
unchanged as they were not involved in coordination. HL1 dis-
played a signal at 12.26 ppm characteristic of OH of carboxylic
which was shifted downfield to about 1.24–1.64 ppm, indicative
of coordination by the carboxylic group with deprotonation. In
ternary complexes the carboxylic and amino protons are shifted
downfield in comparison with alanine to about 0.52–1.39 ppm,
indicating that the nitrogen and oxygen atoms are the probable
coordination sites of alanine to the metal. A higherfield shift is
observed for CH and methyl groups of dl-alanine on comparison
with the spectra of the free amino acid; this may be due to forma-
tion of hydrogen bonding to the sulfoxide oxygen. In the complexes
a broad peak is observed at 2.89 ppm, due to the presence of water
molecules.
3.6. Thermal analysis
Thermal studies of the complexes have been done success-
fully by studying their simultaneous thermal analysis (TG/DTA).
The decomposition data for the complexes is presented in
Tables 4 and 5. The complexes are decomposed by losing water
molecules and ligands in parts in different heating steps and ulti-
mately in the final step metal oxide are found at high temperature.
The thermal decomposition of copper complexes are similar
where the loss of hydrated water molecules are observed in the
first step in case of ternary complex and second mass loss step in
the case of binary complex between 50 and 250 ◦C. This process is
strongly endothermic. A third mass loss step occurs in the range
of 250–650 ◦C, the decomposition of the ligand is observed leaving
CuO as a residue. This degradation process is accompanied by three
successive exothermic peaks.
The same trend is observed in the case of the cobalt and nickel
complexes, where the coordinated water is eliminated with an
endothermic effect within the range of 50–400 ◦C in the first and
second steps in case of cobalt and nickel binary complexes. While
in case of ternary complexes the coordinated water is eliminated in
the first step and connected with the exothermic processes within
the temperature range of 50–200 ◦C. The third and fourth steps in
cobalt and nickel binary chelates are accompanied with the decom-
3.7. Structures of the complexes
The structures of the binary Cu(II)- (1), Co(II)- (2), Ni(II)-
(3), Mn(II)- (4), Cr(III)- (5), Fe(III)- (6), La(III)- (7), UO2(VI)-
(8) complexes with sparfloxacin (HL1) and ternary Cu(II)- (9),
Co(II)- (10), Ni(II)- (11), Mn(II)- (12), Cr(III)- (13), Fe(III)- (14),
La(III)- (15), UO2(VI)- (16) complexes with sparfloxacin (HL1)
and dl-alanine (H2L2) complexes are confirmed by the elemen-
tal analyses, IR, molar conductance, magnetic, UV–Vis, ESR, NMR
and simultaneous thermal analysis. IR spectra revealed that HL1
drug behaves as a deprotonated bidentate ligand coordinated
to the metal ions via pyridone oxygen and carboxylate oxygen.
dl-Alanine is a uninegative bidentate ligand coordinated to the
metal ions via the deprotonated carboxylate-O and amino-N. The
molar conductance data, it is found that the complexes are non-
electrolytes.
1H-NMR spectra of HL1 displayed a signal at 12.26 ppm char-
acteristic of OH of carboxylic, this signal shifted downfield by

N.E.A. El-Gamel, M.A. Zayed / Spectrochimica Acta Part A 82 (2011) 414–423
419
Table 4
Thermal analysis data for binary complexes.
Complex
TG range, (◦C)
DTG, (◦C)
n
Mass loss
Total mass loss
Assignment
Metallic residue
Found (calculated) (%)
04.33 (04.07)
1
2
3
4
5
30–200
200–720
50–400
400–650
50–300
300–600
50–280
280–650
50–240
193
623
1
1
2
2
2
2
1
2
2
Loss of hydration water.
86.55 (86.93)
04.51 (04.10)
87.05 (87.37)
04.07 (04.10)
87.50 (87.39)
04.61 (04.12)
86.94 (87.78)
03.05 (03.97)
90.88 (91.00)
Decomposition of the HL1 ligand.
Loss of coordinated water.
Decomposition of the HL1 ligand.
Loss of coordinated water.
Decomposition of the HL1 ligand.
Loss of coordinated water.
Decomposition of the HL1 ligand
Loss of hydrated and coordinated
water.
CuO
CoO
NiO
80,306
446,570
78, 292
409,465
215
91.56 (91.47)
91.57 (91.94)
91.55 (91.90)
332,553
90, 167
MnO
240–600
50–200
301, 89,580
98
3
1
2
79.67 (79.26)
02.08 (01.98)
80.00 (80.49)
82.72 (83.23)
Loss of 0.5Cl2 and decomposition of the
HL1 ligand
Loss of hydrated and coordinated
water.
Loss of 0.5Cl2 and decomposition of the
HL1 ligand
Loss of hydration water.
Loss of coordinated water, liberation of
NO2 and 0.5O2 and decomposition of
the HL1 ligand.
Loss of coordinated water.
Decomposition of the HL1 ligand.
Cr2O3
6
7
8
200–650
279, 502
82.08 (82.47)
67.52 (68.61)
Fe2O3
La2O3
UO2
50–150
150–650
82
302,521
1
2
03.44 (03.47)
64.08 (65.14)
50–300
300–650
200
273, 553
1
2
01.80 (01.68)
72.27 (73.10)
74.07 (47.78)
n, number of decomposition steps.
the effect of complexation, In ternary complexes the carboxylic
and amino protons are shifted downfield in comparison with ala-
nine, indicating that the nitrogen and oxygen atoms are involved
in the coordination. On the basis of the above observations and
from the magnetic measurements, most of the complexes are
six-coordinated with distorted octahedral geometry while Cu(II)
complexes adopted square planar structure on the other hand U-
and La-atoms have a pentagonal bipyramidal coordination sphere.
The structures of the metal complexes can be given as shown below
(Fig. 4).
3.8. Anti microbial studies
From a medical point of view, the discovery of new therapeutical
targets and the development of new antibacterial, antiparasiti-
cal and antiviral drugs are urgently required, that is due to the
emergence and alarming spread of bacterial, parasitical and viral
strains that are resistant against the drugs used at present in
clinics. Metal ions play a vital role in a vast number of widely
different biological processes. The interaction of these ions with
biologically active ligands, for example in drugs, is a subject
Table 5
Thermal analysis data for ternary complexes.
Complex
TG range, (◦C)
DTG, (◦C)
n
Mass loss
Total mass loss
Assignment
Metallic residue
Found (calculated) (%)
9
10
11
12
13
50–200
200–650
168
210, 306, 579
1
3
05.95 (06.21)
80.57 (80.06)
Loss of hydration water.
Decomposition of the HL1 and H2L2
ligands.
Loss of coordinated water.
Decomposition of the HL1 and H2L2
ligands
Loss of coordinated water.
Decomposition of the HL1 and H2L2
ligands
Loss of coordinated water.
Decomposition of the HL1 and H2L2
ligands
Loss of hydration and coordinated
water.
Loss of 0.5Cl2, decomposition of the
HL1 and H2L2 ligands
Loss of hydration and coordinated
water.
Loss of 0.5Cl2, decomposition of the
HL1 and H2L2 ligands
Loss of hydrated and coordinated
water.
liberation of NO2 and 0.5O2 and
decomposition of the HL1 and H2L2
ligands
86.52 (86.27)
86.80 (86.96)
86.60 (86.99)
87.36 (87.57)
CuO
CoO
NiO
50–200
200–750
80
334, 561
1
2
06.40 (06.27)
80.40 (80.69)
50–200
200–600
70
1
3
06.22 (06.27)
80.38 (80.72)
285, 322, 489
60–300
300–600
105
323, 487, 530
1
2
06.25 (06.31)
81.11 (81.26)
MnO
50–300
300–650
50–200
103
395, 503, 600
80, 145
1
3
1
3
2
2
05.43 (05.97)
Cr2O3
68.99 (68.82)
74.42 (74.79)
14
15
03.25 (02.97)
200–600
50–240
240, 403, 587
118, 208
70.17 (70.73) 73.42 (73.70)
05.11 (05.02)
Fe2O3
240–630
397, 540
50.03 (49.41)
55.14 (54.43)
64.03 (64.73)
La2O3
UO2
16
50–300
300–600
228
1
4
02.10 (02.35)
61.93 (62.38)
Loss of coordinated water.
Decomposition of the HL1 and H2L2
ligands
280,365,480, 574
n, number of decomposition steps.

420
N.E.A. El-Gamel, M.A. Zayed / Spectrochimica Acta Part A 82 (2011) 414–423
CH3
CH3
F
NH
NH
F
N
N
F
N
N
F
CH3
CH3
O
O
O
O
O
NH2
NH2
O
O
O
NH2
O
O
O
U
Y nH2O
X
M
O
O
O
NH2
OH2
F
F
CH3
CH3
N
N
N
N
F
NH
F
NH
CH3
CH3
UO2(VI) binary complex
M = Cu(II), X = Y = 0, n = 2
M = Co(II),Ni(II), Mn(II), X = Y = H2O, n = 0
M = Cr(III), Fe(III), X = Cl, Y = H2O, n = 1
M = La(III), X = NO3, Y = H2O, n = 2
O
C
CH3
H
H
H3C
C
C
O
H2N
O
H2N
O
O
nH2O
Y
M
X
O
U
NH2
O
O
O
O
NH2
F
OH2
O
F
O
CH3
N
N
CH3
N
N
NH
F
F
NH
CH3
CH3
M = Cr(III), Fe(III), X = Cl, Y = H2O, n = 1
UO2(VI) ternary complex
M = Cu(II), Co(II), Ni(II), Mn(II), X = Y = H2O, n = 0
M = La(III), X = NO3, Y = H2O, n = 1
Fig. 4. Proposed chemical structures of the metal complexes.
of considerable interest. Some of the biologically active com-
pounds act via chelation [26], but for most of them is known
about how metal binding influences their activity. Our target is
to produce and evaluate some potential antimicrobial compounds
without causing any side effects on the patients; therefore we
have been interested in studying the complexing ability of biolog-
ically active ligands like sparfloxacin and discuss the coordination
chemistry behavior of some metal ion and/or in presence of dl-
alanine.
increase the chance of identify and observe the biological effi-
ciency of the tested materials. The studies were carried out
on, S. aureus (G+), E. coli (G−) species in bacteria and A. flavus
and C. Albicams among fungi using assay plates disc method
on appropriate nutrient medium. The results are included in
Figs. 5 and 6, the inhibition zone diameters were measured
after 48 h at 37 ◦C of incubation for all the mentioned complexes
(Table 6). On comparing the biological activity of the sparfloxacin
and its metal complexes with the standard (tetracycline, antibac-
terial) and (Amphotericin B, antifungal), the following results are
obtained:
In evaluating the antimicrobial activity of the mentioned
complexes we used more than one test organism in order to

N.E.A. El-Gamel, M.A. Zayed / Spectrochimica Acta Part A 82 (2011) 414–423
421
Fig. 5. Biological activity of HL1 and its complexes against (A, G+), (B, G−) bacteria.
Fig. 6. Fungal activity of HL1 and its complexes against Aspergillus flavus and Candida albicans fungus.

422
N.E.A. El-Gamel, M.A. Zayed / Spectrochimica Acta Part A 82 (2011) 414–423
Table 6
Antimicrobial activity of metal complexes and the standard effects of antibacterial and antifungal agents.
Diameter of zone of
inhibition (mm)
Complexes
Stapylococcus
aureus (G+)
Tetracycline
(antibacterial
agent)
Escherichia coli (G−
)
Tetracycline
(antibacterial
agent)
Aspergillus Amphoterician B
Candida
albicans
Amphoterician B
(antifungal agent)
flavus
(antifungal agent)
Distilled water,
chloroform, DMSO
HL1
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
0
0
0
0
0
0
0
0
0
0
32
35
35
37
33
37
35
32
32
33
39
35
34
33
38
32
32
31
31
31
31
31
31
31
31
31
31
31
31
31
31
31
31
31
36
34
34
37
36
36
36
36
36
36
36
38
37
35
35
36
36
33
33
33
33
33
33
33
33
33
33
33
33
33
33
33
33
33
16
16
16
16
16
16
16
16
16
16
16
16
16
16
16
16
16
19
19
19
19
19
19
19
19
19
19
19
19
19
19
19
19
19
15
18
13
14
0
0
0
0
0
0
14
13
0
0
0
12
12
0
0
0
0
0
0
17
12
0
0
0
0
0
0
0
3.8.1. Staphylococcus aureus (G+)
titative microbiological method examine potent and useful drugs
in both clinical and experimental tumour chemotherapy [33].
The free HL1 ligand showed higher antibacterial effect that
the standard tetracycline, furthermore, the biological activity
of binary and ternary complexes of Cu(II), Co(II), Ni(II), Mn(II),
Cr(III), and Fe(III) complexes is higher than that of the free HL1
ligand and standard tetracycline. While, UO2(VI) and La(III)
binary and ternary complexes have the same biological activity
of HL1 ligand and higher than that of tetracyclin. (Fig. 5), the
biological activity of the complexes are found to follow the order
10 > 15 > 3 = 5 > 4 > 6 = 1 = 2 = 11 > 12 > 9 > HL1 = 15 = 16 > 13 > stand-
ard tetracycline.
3.8.3. Using Aspergillus flavus and Candida albicans (Fungus)
HL1 and most of binary and ternary complexes did not repre-
sent any fungal activity comparing to the standard Amphoterician
B, nevertheless binary and ternary Co(II) and Ni(II) as well as ternary
Mn(II) complexes showed some fungal activity, this activity is rela-
tively less than the standard Amphoterician B (Fig. 6). The order
of activity can be arranged as follow: standard Amphoterician
B > 1 > 3 = 10 > 2 = 11 (in case of A. flavus) while in case of C. albicans
the order can be as follows Amphoterician B > 10 > 2 = 3 = 11.
By comparing the evaluation of the antimicrobial activity of
these complexes with the previous published quinolones com-
plexes, similar effects have been reported, very recently Sadeek
et al. represented, that uranium, zirconium and yttrium com-
plexes with sparfloxacin and gatifloxacin showed a good activity
against Gram −ve and Gram +ve than drugs alone, while no anti-
fungal activity is observed for both drugs and metal complexes
[34], on the other hand Ni(II) and Ag(I) enoxacin complexes rep-
resented slight activity against Gram +ve and Gram −ve bacteria
[35], while Cu–ciprofloxacin showed high standard antibacterial
activity against Gram +ve and Gram −ve [36], the same effect is
reported in the case of Fe–ciprofloxacin, where also no significant
inhibition towards the growth of fungal strains are reported [37].
3.8.2. Using Escherichia coli (G−)
The biological activity of the HL1 ligand is found to be higher
than that of the tetracycline standard, whereas binary [Mn(II),
Cr(III), Fe(III), La(III) and UO2 (VI)] and ternary [Cu(II), Co(II),
La(III) and UO2(VI)] complexes represented an equal activity
to HL1 ligand, while binary Ni(II) and ternary Ni(II) and Mn(II)
complexes showed a higher biological activity than both HL1
ligand and standard tetracycline. On the other hand binary
[Cu(II) and Co(II)] and ternary [Cr(III) and Fe(III)] complexes
represented less activity than HL1 but higher than standard
tetracycline (Fig. 5). The order of activity can be arranged as follow
11 > 12 = 3 > HL1 > 4 > 5 > 6 > 7 > 8 > 9 > 10 > 15 > 16 > 13 = 14 > 1 > 2>
standard tetracycline.
The significance of these results could be applied completely
in the treatment of some diseases caused by E. coli like, Gastroen-
teritis, Septicaemia, Urinary tract infections and hospital required
infections [27,28]. Many clinical research have been interested
in finding new anti-tumours depends mainly on attacking Gram-
negative bacteria [29–31], and since there are certain organisms
difficult to attack as most of them are Gram-negative rods, there-
fore the complexes which showed higher biological activity against
the Gram-negative strains may represent powerful attacking to the
barrier function as mentioned by Brown [31] and Nikaido et al.
[29]; therefore, the synthesis of these complexes might be recom-
mended in order to establish a new branch of research dealing with
the anti-tumour effect of many synthetic compounds as described
by Hodnett et al. [31] and Hickman [32]. Thus we are claimed that
these complexes may have a possible anti-tumour effect depending
on the fact that, Gram-negative bacteria can be considered a quan-
4. Conclusion
In this article, we describe the synthesis and characterization
of binary and ternary complexes of sparfloxacin (HL1) with dl-
alanine (H2L2). Elemental analysis of C, H and N data obtained
were in agreement with the predicted formula. Data of the infrared
and 1H-NMR spectroscopy represented that sparfloxacin acts as a
bidentate deprotonated ligand bound to the metal through pyri-
done oxygen and carboxylate oxygen. dl-Alanine is a uninegative
bidentate ligand coordinated to the metal ions via the deprotonated
carboxylate-O and amino-N. All the complexes are six-coordinated
with distorted octahedral geometry while Cu(II) complexes which
have square planar structure from ESR parameters, magnetic and
solid reflectance measurements. On the other hand uranyl and
lanthanide complexes have a pentagonal bipyramidal coordina-

N.E.A. El-Gamel, M.A. Zayed / Spectrochimica Acta Part A 82 (2011) 414–423
423
tion sphere. It has been observed that the metal complex have a
high activity than ligand against same organisms under the iden-
tical experimental condition. The comparison of inhibition zone
values for the newer metal complexes (Table 6) with another pre-
vious quinolones complexes reveals that the antimicrobial activity
is quite high which could be an indication to generate useful infor-
mation that contributes to the differentiation of new biomedical
branches.
[9] M.A. Pfaller, L. Burmeister, M.A. Bartlett, M.G. Rinaldi, Multicenter evaluation
of four methods of yeast inoculum preparation, J. Clin. Microbiol. 26 (1988)
1437–1441.
[10] J.A. Dean (Ed.), Lange’s Handbook of Chemistry, 14th ed., McGraw-Hill, New
York, 1992, Table 8.35.
[11] E. Santi, M.H. Torre, E. Kremer, S.B. Etcheverry, E.J. Baran, Vib. Spectrosc. 5
(1993) 285–293.
[12] E. Prestch, T. Clerc, J. Seibl, W. Simon, Tablas de Determinacion Estruc-
tural por Metodos Espectroscopicos, Springer-Verlag Iberica, Barcelona,
1998.
[13] N. Jimenez-Garrido, L. Perello, R. Ortiz, G. Alzuet, M. Gonzalez-Alvarez, E. Can-
ton, M. Liu-Gonzalez, S. Garcia-Granda, M. Perez-Priede, J. Inorg. Biochem. 99
(2005) 677–689.
[14] (a) M.A. Zayed, F.A. Nour El-Dien, G.G. Mohamed, N.E.A. El-Gamel, Spectrochim.
Acta 64A (2006) 216–232;
References
[1] (a) B.B. Lohray, V.B. Lohray, B.K. Srivastava, P. Kapadnis, P.P. Pandya, Bioorg.
Med. Chem. 12 (2004) 4557–4564;
(b) M.A. Zayed, F.A. Nour El-Dien, G.G. Mohamed, N.E.A. El-Gamel, Spectrochim.
Acta 60A (2004) 2843–2852;
(c) S. El-Khateeb, S. Abdel Fattah, S. Abdel Razeg, M. Tawakkol, Anal. Lett. 22
(1989) 101–115;
(d) W. Radecka-Paryzek, E. Luks, Monatsh. Chem. 126 (1995) 795–798;
(e) J. Costamagna, F. Carruso, M. Rossi, M. Campos, J. Canales, J. Ramirez, J. Coord.
Chem. 54 (2001) 247–259;
(f) G.G. Mohamed, N.E.A. El-Gamel, Spectrochim. Acta 60A (2004) 3141–3151;
(g) G.G. Mohamed, N.E.A. El-Gamel, Vib. Spectrosc. 36 (2004) 97–
104.
(b) D.E. King, R. Malone, S.H. Lilley, Am. Fam. Physician 61 (2000) 2741–2748;
(c) C.H. Gross, J.D. Parsons, T.H. Grossman, P.S. Charifson, S. Bellon, J. Jernee,
M. Dwyer, S.P. Chambers, W. Markland, M. Botfield, S.A. Raybuck, Antimicrob.
Agents Chemother. 47 (2003) 1037–1046;
(d) J.E.F. Reynolds, The Extra Pharmacopeia, 30th ed., The Pharmaceutical Press,
London, 1993, pp. 145–147;
(f) M.V.N. de Souza, Mini-Rev. Med. Chem. 5 (2005) 1009–1017;
(g) G. Anquetin, J. Greiner, N. Mahmoud, M. Santillana-Hayat, R. Gozalbes, K.
Farhati, F. Derouin, A. Aubry, E. Cambau, P. Vierling, Eur. J. Med. Chem. 41 (2006)
1478–1493.
[15] (a) R. Wysokinˇıski, B. Morzyk-Ociepa, T. Głowiak, D. Michalska, J. Mol. Struct.
606 (2002) 241–251;
(b) F.R. Allen, O. Kennard, D.G. Watson, L. Brammer, A.G. Orpen, R. Taylor, J.
Chem. Soc. Perkin Trans. 2 (1987) S1–S19;
[2] (a) D.T.W. Chu, P.B. Fernandes, in: B. Testa (Ed.), Advances in Drug Research,
vol. 21, Academic Press, London, 1991, pp. 39–144;
(b) D.E. King, R. Malone, S.H. Lilley, Am. Fam. Physician 61 (2000) 2741;
(c) Sparfloxacin, in: J.K. Aronson (Ed.), Meyler’s Side Effects of Drugs: The Inter-
national Encyclopedia of Adverse Drug Reactions and Interactions, Elsevier,
2006, pp. 3172–3174.
(c) E.J. Baran, R.C. Mercader, F. Hueso-Uren˜a, M.N. Moreno-Carretero, M. Quiros-
Olozabal, J.M. Salas-Peregrin, Polyhedron 15 (1996) 1717–1721;
(d) S. Lencioni, A. Pellerito, T. Fiore, A.M. Giuliani, L. Pellerito, M.T. Cambria, C.
Mansueto, Appl. Organomet. Chem. 13 (1999) 145–157.
[16] Z. Ji, Y. Chen, X. Cheng, G. Chen, J. Chem. Crystallogr. 32 (2002) 141–145.
[17] N.E.A. El-Gamel, D. Gerlach, J. Coord. Chem. 61 (2008) 2246–2265.
[18] (a) G.M.R. Larin, J. Coord. Chem. 25 (1999) 804–810;
(b) S. Djebbar-Sid, O. Benali-Baitich, J.-P. Deloume, Polyhedron 16 (1997)
2175–2182.
[3] (a) I. Turel, Coord. Chem. Rev. 232 (2002) 27–47;
(b) K.C. Skyrianou, C.P. Raptopoulou, V. Psycharis, D.P. Kessissoglou, G. Psomas,
Polyhedron 28 (2009) 3265–3271;
(c) M. Ruiz, L. Perelloˇı, J. Server-Carrioˇı, R. Ortiz, S. Garcıˇıa-Granda, M.R. Diaz, E.
Canto’n, J. Inorg. Biochem. 69 (1998) 231–239;
(d) B. Macias, M.V. Villa, M. Sastre, A. Castin˜eiras, J. Borraˇıs, J. Pharm. Sci. 91
[19] B.J. Hathaway, A.A.G. Tomlinson, Coord. Chem. Rev. 5 (1970) 1–43.
[20] S.-F. Tan, K.-P. Ang, Trans. Met. Chem. 13 (1998) 64–68.
[21] J. Manonmani, R. Thirumuruhan, M. Kandaswamy, V. Narayanan, S. Shanmuga,
S. Raj, M.N. Ponnus Wamy, G. Shanmugan, H.K. Fun, Polyhedron 20 (2001)
3039–3048.
[22] F.A. Cotton, G. Wilkinson, C.A. Murillo, M. Bochmann, Advanced Inorganic
Chemistry, 6th ed., Wiley, New York, 1999.
[23] A.B.P. Lever, Inorganic Electronic Spectroscopy, 1st ed., Elsevier, Amsterdam,
1968.
[24] C.-H. Song, H.W. Ryu, J.K. Park, T.-S. Ko, Bull. Kor. Chem. Soc. 20 (1999) 727–
730.
[25] K.B. Gudasi, R.V. Shenoy, R.S. Vadavi, M.S. Patil, S.A. Patil, R.R. Hanchinal, S.A.
Desai, H. Lohithaswa, Bioinorg. Chem. Appl. (2006) 1–8.
[26] A. Albert, Selective Toxicity. The Physico-Chemical Basis of Therapy, 6th ed.,
Chapman and Hall, London, 1979.
[27] E. Jawetz, J.L. Melnick, E.A. Adelberg, Review of Medical Microbiology, 16th ed.,
Lang Medical Publications, Los Anglos, CA, 1979.
[28] W.H. Hughes, H.C. Stewart, Concise Antibiotic Treatment, Butter Worth, Lon-
don, 1970.
(2002) 2416–2423;
(e) Z.H. Chohan, C.T. Supuran, A. Scozzafava, J. Enzyme, Inhib. Med. Chem. 20
(2005) 303–307;
(f) Z.H. Chohan, H. Pervez, A. Rauf, C.T. Supuran, Met. Based Drugs 8 (2001)
263–267;
(g) M.P. Loˇıpez-Gresa, R. Ortiz, L. Perelloˇı, J. Latorre, M. Liu-Gonzaˇı lez, S. Garcıˇıa-
Granda, M. Peˇırez-Priede, E. Cantoˇı n, J. Inorg. Biochem. 92 (2002) 65–74;
(h) J.R. Anacona, C. Toledo, Trans. Met. Chem. 26 (2001) 228;
(i) I. Turel, L. Golic, P. Bukovec, M. Gubina, J. Inorg. Biochem. 71 (1998) 53–60;
(j) I. Turel, I. Leban, N. Bukovec, J. Inorg. Biochem. 66 (1997) 241–245;
(k) F. Gao, P. Yang, J. Xie, H. Wang, J. Inorg. Biochem. 60 (1995) 61–67.
[4] S. Jain, N.K. Jain, K.S. Pitre, J. Pharm. Biomed. Anal. 29 (2002) 795–801;
(b) E.K. Efthimiadou, Y. Sanakis, C.P. Raptopoulou, A. Karaliota, N. Katsarosa, G.
Psomasa, Bioorg. Med. Chem. Lett. 16 (2006) 3864–3867.
[5] (a) L.M.M. Vieira, M.V. de Almeida, H.A. de Abreu, H.A. Duarte, R.M. Grazul, A.P.S.
Fontes, Inorg. Chim. Acta 362 (2009) 2060–2064;
(b) L.M.M. Vieira, M.V. de Almeida, M.C.S. Lourenc¸ o, F.A.F.M. Bezerra, A.P.S.
Fontes, Eur. J. Med. Chem. 44 (2009) 4107–4111;
(c) M. Ruiz, R. Ortiz, L. Perelloˇı, S. Garcıˇıa-Granda, M.R. Dıˇıaz, Inorg. Chim. Acta
[29] H. Nikaido, T. Nakae, Adv. Microbiol. Phys. 20 (1979) 163–250.
[30] N.R.W. Brown, Resistance of Pseudomonas aeruginosa, vol. 71, John Wiley, 1975.
[31] E.M. Hodnett, A.W. Wu, F.A. French, Eur. J. Med. Chem. Chem. Ther. 13 (1987)
577–579.
217 (1994) 149–154.
[6] H. Beraldo, D. Gambino, Mini-Rev. Med. Chem. 4 (2004) 31–39;
(b) J.R. Anacona, C. Toledo, Trans. Met. Chem. 26 (2001) 228–231;
(c) C. Chulvi, M.C. Mun˜oz, L. Perelloˇı, R. Ortiz, M.C. Mun˜oz, M.I. Arriortua, J. Via,
K. Urtiaga, J.M. Amigoˇı, L.E. Ochando, J. Inorg. Biochem. 42 (1991) 133–138;
(d) M. Ruiz, R. Ortiz, L. Perelloˇı, A. Castin˜eiras, M. Quiroˇıs, Inorg. Chim. Acta 211
(1993) 133–139.
[7] N. Sultana, M. Saeed, S. Gul, S. Shamim, J. Mol. Struct. 975 (2010) 285–291.
[8] (a) A.W. Bauer, W.M. Kirby, C. Sherris, M. Turck, Antibiotic susceptibility testing
by a standardized single disk method, Am. J. Clin. Pathol. 45 (1966) 493–496;
(b) National Committee for Clinical Laboratory Standards. (2002). Reference
Method for Broth Dilution Antifungal Susceptibility Testing of Conidium-
Forming Filamentous Fungi: Proposed Standard M38-A. NCCLS, Wayne, PA,
USA.
[32] J.A. Hickman, Biochemie 60 (1987) 997.
[33] T. Inoue, Y. Yamashita, M. Nishihara, S. Sugiyama, Y. Sonoda, T. Kumabe, M.
Yokoyama, T. Tominaga, Neuro-Oncology 11 (2009) 151–157.
[34] (a) S.A. Sadeek, W.H. El-Shwininy, J. Mol. Struct. 977 (2010) 243–253;
(b) S.A. Sadeek, W.H. El-Shwininy, J. Coord. Chem. 63 (2010) 3471–3482.
[35] Y. Cheng Liu, X.-J. Luo, C. Barta, H. Liang, J. Coord. Chem. 63 (2010) 3146–
3154.
[36] Y. Wang, G.-W. Lina, J. Honga, L. Lia, Y.-M. Yanga, T. Lu, J. Coord. Chem. 63 (2010)
3662–3675.
[37] P.B. Pansuriya, M.N. Patel, J. Enzyme Inhib. Med. Chem. 23 (2008) 230–239.