Development of a C3-symmetric benzohydroxamate tripod: Trimetallic complexation with Fe(III), Cr(III) and Al(III)

Article abstract of DOI:10.1016/j.saa.2016.02.033

The design, synthesis and physicochemical characterization of a C₃-symmetry Benzene-1,3,5-tricarbonylhydroxamate tripod, noted here as BTHA, are described. The chelator was built from a benzene as an anchor, symmetrically extended by three hydroxamate as ligating moieties, each bearing O, O donor sites. A combination of absorption spectrophotometry, potentiometry and theoretical investigations are used to explore the complexation behavior of the ligand with some trivalent metal ions: Fe(III), Cr(III), and Al(III). Three protonation constants were calculated for the ligand in a pH range of 2-11 in a highly aqueous medium (9:1 H₂O: DMSO). A high rigidity in the molecular structure restricts the formation of 1:1 (M/L) metal encapsulation but shows a high binding efficiency for a 3:1 metal ligand stoichiometry giving formation constant (in β unit) 28.73, 26.13 and 19.69 for [M₃L]; M=Fe(III), Al(III) and Cr(III) respectively, and may be considered as an efficient Fe-carrier. The spectrophotometric study reveals of interesting electronic transitions occurred during the complexation. BTHA exhibits a peak at 238 nm in acidic pH and with the increase of pH, a new peak appeared at 270 nm. A substantial shifting in both of the peaks in presence of the metal ions implicates a s coordination between ligand and metal ions. Moreover, complexation of BTHA with iron shows three distinct colors, violet, reddish orange and yellow in different pH, enables the ligand to be considered for the use as colorimetric sensor.

Full text of DOI:10.1016/j.saa.2016.02.033

Development of a C3-symmetric benzohydroxamate tripod:Trimetallic com-
plexation with Fe(III), Cr(III) and Al(III)
Minati Baral, Amit Gupta, B.K. Kanungo
PII:
DOI:
Reference:
S1386-1425(16)30090-7
doi: 10.1016/j.saa.2016.02.033
SAA 14299
To appear in:
Received date:
Revised date:
Accepted date:
14 July 2015
26 January 2016
28 February 2016
Please cite this article as: Minati Baral, Amit Gupta, B.K. Kanungo, Development of a
C3-symmetric benzohydroxamate tripod:Trimetallic complexation with Fe(III), Cr(III)
and Al(III), (2016), doi: 10.1016/j.saa.2016.02.033
This is a PDF file of an unedited manuscript that has been accepted for publication.
As a service to our customers we are providing this early version of the manuscript.
The manuscript will undergo copyediting, typesetting, and review of the resulting proof
before it is published in its final form. Please note that during the production process
errors may be discovered which could affect the content, and all legal disclaimers that
apply to the journal pertain.

ACCEPTED MANUSCRIPT
Development of a C3-symmetric benzohydroxamate tripod:Trimetallic
complexation with Fe(III), Cr(III) and Al(III)
Minati Baral1*, Amit Gupta1 and B.K. Kanungo2
1Department of Chemistry, National Institute of Engineering and Technology Kurukshetra,
Haryana-136119, India
2Department of Chemistry, Sant Longowal Institute of Engineering and Technology,
Longowal, Punjab-148106, India
Email: minatibnitkkr@gmail.com, Telephone:+91(1744)233544
Abstract
The design, synthesis and physicochemical characterization of a C3-symmetry Benzene-1,3,5 -
tricarbonylhydroxamate tripod, noted here as BTHA, are described. The chelator was built from
a benzene as an anchor, symmetrically extended by three hydroxamate as ligating moieties, each
bearing O,O donor sites. A combination of absorption spectrophotometry, potentiometry and
theoretical investigations are used to explore the complexation behavior of the ligand with some
trivalent metal ions: Fe(III), Cr(III), and Al(III). Three protonation constants were calculated for
the ligand in a pH range of 2-11 in a highly aqueous medium (9:1 H2O: DMSO). A high rigidity
in the molecular structure restricts the formation of 1:1 (M/L) metal encapsulation but shows a
high binding efficiency for a 3:1 metal ligand stoichiometry giving formation constant (in  unit)
28.73, 26.13 and 19.69 for [M3L]; M= Fe(III), Al(III) and Cr(III) respectively, and may be
considered as an efficient Fe-carrier. The spectrophotometric study reveals of interesting
electronic transitions occurred during the complexation. BTHA exhibits a peak at 238 nm in
acidic pH and with the increase of pH, a new peak appeared at 270 nm. A substantial shifting in
both of the peaks in presence of the metal ions implicates a s coordination between ligand and
metal ions. Moreover, complexation of BTHA with iron shows three distinct colors, violet,
reddish orange and yellow in different pH, enables the ligand to be considered for the use as
colorimetric sensor.
Keywords: Spectrophotometric, potentiometric, hydroxamate, formation constants, protonation
constants
1

ACCEPTED MANUSCRIPT
1. Introduction
Hydroxamic acids, R1C(O)N(R2)OH, where R1=alkyl/aryl and R2= alkyl/aryl or H are
ubiquitously known ligands in coordination chemistry for their high binding affinity for wide
range of transition metal ions, and have received much attention in biological, analytical,
chemical, medical and industrial field offering broad applications, particularly as clinical iron
removing agents [1-3]. They form stable chelates with various metal ions attributing to their
inhibitory action against enzymes having metalloprotein as their functional group [2-4]. The
solutions of these organic molecules produce intense color complexes contributing widely in
analytical chemistry as well as spectrophotometry for determination of metals [5-7]. Moreover, a
significant adsorption of metals (mainly iron) by hydroxamate may play a beneficial role in
recovery of metals from ground/waste water. Such as, excess metals dissolved in water caused
toxic concentration as in acid mine drainage, posing multiple health problems such as high
accumulation of iron in body, causing hemochromatosis [8]. Moreover, Al(III) a non-essential
element, is involved in causing dialysis dementia in patients and affecting the central nervous
system that causes different disorders such as amyotrophic lateral sclerosis [9-10]. The only
approved drug, desferrioxamine B (iron chelator) for hemochromatosis is posing several
drawbacks and side effects [11]. Although Cr+3 is absorbed by biological system, still, a high
instance of occupational exposure to this trivalent metal pertained with respiratory problems like
coughing, wheezing and dyspnea. Dermal exposure to chromium has also been demonstrated to
produce irritant and allergic contact dermatitis [12]. Moreover, industrial effluents also have high
concentration of dissolved chromium which causes chromium toxicity in nearby water resources.
So, to deal with such situations, some molecules are under necessity to reduce overload of
Fe(III), Cr(III) and Al(III) ions in an eco-friendly manner.
Despite of ongoing several researches for elucidating hydroxamate behavior, they remain poorly
characterized. Possibilities of several conformations depending on concentration, temperature
and the nature of the solvent attributes to the biggest challenge in hydroxamic acid
characterization. A detailed comparative study is thus required to evaluate the capability of the
ligand to serve as a potential chelator. High Stability constants of hydroxamic acids with
transition metal ions can be considered as an important tool for elucidating chelating (polarizing)
power of the ligands. The preferential binding mode in the metal hydroxamic acid complexes is
2

ACCEPTED MANUSCRIPT
through O, O (bidentate) donors in which the ligand is either monobasic (hydroxamato) or
dibasic (hydroximato) [13-15].
Hard metals with higher oxidation states preferentially form more stable hydroxamate
complexes. Benzohydroxamic acid (Bha), acetohydroxamic acid (Aha), salicylhydroxamic acid
and desferrixomine family (derived from Streptomyces species) are few hydroxamic acids
dominating the coordination chemistry and chemical biology. The Bha is extensively used as an
excellent spectrophotometric reagent for determination of metal ions such as Fe, Mo, V, U, and
Mn, and could be served as a potential colorimetric reagent [16]. Synthesis and studies of a
number of mono benzohydroxamates with various divalent and trivalent metal ions have been
documented [17-18]. Trimesic acid coated alumina has been reported as a potent adsorbent for
iron [8]. Considering the implications of Bha in analytical chemistry, there arises a burgeoning
interest in increasing the substitution and exploring the coordination behavior in relation with the
enhanced metal selectivity, efficiency and affectivity.
However, realizing the efficiency of hydroxamic acids as an efficient metal binder, it was
thought for a higher substitution on benzohydroxamate which may offer an increased metal
chelation in quantity, hence may be served as a potential metal scavenger. Keeping in view of the
above, in the present communication we describe the development and complexation studies of a
multidentate polyfunctional benzene-1,3,5-tricarbonylhydroxamic acid (BTHA). The
potentiometric and spectrophotometric studies also explore the coordination behavior of the
ligand with Fe, Al, and Cr trivalent metal ions in highly aqueous medium, H2O: DMSO = 9:1.
Due to a rigid molecular framework and the coordination ability of three equally spaced
hydroxamate groups, BTHA can serve as a strong metal-chelator and may find its use as a potent
in applications in the areas such as acid mine drainage carrying excess dissolved iron at low pH.
2. Experimental
2.1 Materials and Measurements
Ultrapure grade reagents such as trimesic acid, methylchloroformate, N-methylmorpholine were
purchased from Sigma Aldrich for the synthesis as well as titration purposes and used directly
3

ACCEPTED MANUSCRIPT
without further purification unless otherwise mentioned. Analytical grade solvents:
dimethylformamide, methanol, tetrahydrofuran and dimethylsulphoxide were purchased from
Qualigens and SD Fine Chemicals Limited and distilled over appropriate drying agents as per
standard procedures prior to use. Melting points were determined on MICROSIL digital melting
point apparatus and were uncorrected.
The progress of the reaction and the purity were monitored by thin layer chromatography (TLC),
performed on silica gel glass plates using 8:2 hexane and ethylacetate mixture and were
visualized in iodine vapors. CHN elemental analyses were performed on Euro EA 3000, 60Hz-
1200W elemental analyzer. UV spectrophotometric titrations were done on Thermofisher UV-
VISIBLE spectrophotometer. The potentiometric and spectrophotometric studies were carried
out in 9:1 :: H2O: DMSO mixture. The ionic strength was adjusted to 0.1 M KCl and the
temperature was maintained at 25 ± 1 ºC. The stock solutions were prepared in Millipore grade
deionized water and for weighing the appropriate amount, high precision electronic balance
CAS-CAUW220D was used that measures up to a level of four digits. KOH and HCl were
standardized by acid-base titration method using 0.1M potassium hydrogen phthalate as primary
standard and the exact concentrations were determined. 0.1M solutions were prepared freshly
immediately before use in ultrapure deoxygenated water that was flushed in continuity with inert
N2 to remove CO2 and O2.
1
13
The H and C NMR spectra of the ligand were obtained in d6DMSO using Bruker Avance II
400 NMR spectrometer. The chemical shifts were quantified in δ values (ppm).
Tetramethylsilane (TMS) was used as an internal reference. The FT-IR spectra were obtained in
the transmission mode as KBr pallet with Perkin- Elmer FT- IR spectrometer SPECTRUM 1000
in Mid-IR region range (4000- 400 cm-1).
2.2 Synthesis of Benzene-1,3,5- tricarbonylhydroxamic acid (BTHA)
In the synthesis of BTHA, the solvents used were freshly distilled and dried as per standard
method [19]. The synthesis is illustrated below as scheme 1. A solution of 2.83 g of
methylchloroformate (30.0 mmol) in 20 ml dimethylformamide was maintained at nearly -5ºC
using ice/CaCl2 mixture and was stirred till a clear solution is obtained in step 1. 2.10 g (10.0
4

ACCEPTED MANUSCRIPT
mmol) of trimesic acid dissolved in 10 ml of dimethylformamide was added to the above
solution followed by the addition of 3.03g (30.0 mmol) N-methylmorpholine. The mixture was
allowed to stir for 15 minutes. Precipitate of morpholine chloride was filtered out and the pale
yellow colored filtrate was kept at ~ 4 °C. In step 2, 2.1 g (30.0 mmol) hydroxylamine
hydrochloride in methanol (15.0 ml) was added to the methanolic solution of KOH (15.0 ml).
The mixture was stirred for 15 minute at 0 ºC. White precipitate of KCl was removed by
filtration and the filtrate containing free hydroxylamine was maintained at low temperature (0
°C). In step 3, cold solution of pale yellow colored filtrate of step 1 was added drop wise in
hydroxylamine with continuous stirring for 60 minutes at room temperature. The solution was
concentrated under vacuum to which 20 ml of distilled water was added and extracted with ethyl
acetate followed by evaporation under vacuum. White-grey colored residue thus obtained was
further washed with distilled water and re-crystallized in ethylacetate/methanol mixture giving a
white residue (yield: 1.83g, 72%,); m.p.: 210-214 C; FT-IR (KBr pellet, cm-1): 3370, 3168, 3019,
1649, 1586,1048, 910; 1H NMR (400mHz, DMSO-d6, δ ppm) 10.8-11.8 (s, 3H), 8.6-9.6 (s, 3H),
13
8.2-8.5 (m, 3H); C NMR (DMSO-d6, δ ppm): 166, 129, 127; and CHN analysis experimental
(calculated) for C9H9O6N3 found to be C = 42.12% (42.45%), N = 16.62% (16.54%), H = 3.41%
(3.64%).
2.3 Potentiometric and spectrophotometric titration procedure
Potentiometric and spectrophotometric methods were employed to determine the ligand
protonation constants and binding constants with the trivalent metal ions. Both the titrations were
carried out in a double wall glass jacketed titration cell and the temperature was maintained at
25±1 ºC by using a constant temperature circulating bath. The potentiometric titrations were
done on HACH Sension-2 potentiometer using glass electrode that recorded pH as –log[H+].
The electrode was duly calibrated in suitable buffers prepared by H2O-DMSO 9:1 mixture
following standard chemical method [19]. Borosilicate glass micro burette with least count
0.01ml was used for all titrations by using 9:1 H2O: DMSO mixture. The ligand solution (2 ml, 1
x 10-2M) was acidified with 2.0 ml of 0.1M HCl and titrated against standardized 0.1M KOH in
10% DMSO-water mixture at µ= 0.1M KCl at 25±1 ºC to determine the protonation constants in
pH range 2.0 to 11.0. Complexation behavior of BTHA towards Fe(III), Cr(III) and Al(III) were
studied by titrating 1:1 (2 ml, 1 x 10-2 M and 2 ml, 1x 10-2 M) and 3:1 (6 ml, 1 x 10-2 M and 2 ml,
5

ACCEPTED MANUSCRIPT
1x 10-2 M) metal-ligand mixtures against KOH, and formation constants (log β) of metal
complexes were determined. Final concentration of ligand and metal were maintained as 1 x10-3
Scheme 1. Schematic representation of synthesis of tripodal ligand BTHA
and 3 x 10-3 M respectively for the titrations by keeping the total volume of 20 ml. The titration
data was refined by the non-linear least square refinement program Hyperquad 2006 [20]. The
formation of protonation and metal-ligand equilibria and their formation constants (β) are
defined in equation 1.
pM + qH + rL
MpHqLr
(1)
β
pqr = [MpHqLr ] [M]p[H]qL]r
/
6

ACCEPTED MANUSCRIPT
The species distribution curves at different pH were plotted with HYSS 2009 [21].
2.4 Computational methods
All calculations were carried out on a Pentium R (IV) 3.20 GHz machine in Window XP
environment using HyperChem version 7.5 [22]. The initial geometry of BTHA molecule
leading to minimum strain energy was achieved through molecular mechanics calculation using
MM+ force field. The periodical search to global minimum energy conformer of the ligand and
its metal complexes were achieved using molecular dynamics simulation up to 2000 K followed
by the MM+ calculation. The minimum strain energy structures were further re-optimized
through semi-empirical (PM3) self-consistent fields (SCF) method, at the Restricted Hartree-
FOCK (RHF) level. The geometry optimizations were done by the application of the steepest
descent method followed by the Polak-Ribere method with convergence limit of 0.0001 kcal/mol
and RMS gradient of 0.001kcal/mol. For all the possible metal chelates, six coordinated
structures were drawn by adding appropriate number of water molecules with the metal ions
considering oxidation state of the metal ions. The electronic spectra of the various protonated
and deprotonated species of the ligand were calculated through semi-empirical methods applying
ZINDO/S Hamiltonian by using CAChe Work System Pro version 6.11 [23]. The full vibrational
analysis was carried out for the ligand to obtain free energy in values by applying semi-empirical
PM3 Hamiltonian. Prior to property calculations all the global minimum structures of the ligand
and the coordinated compounds were checked for imaginary frequency.
3. Results and Discussion
3.1 Synthesis and characterization of ligand
The polydentate chelator BTHA was obtained as white powder with a yield of 72% (pure) in
three steps as given in scheme 1. The compound is completely soluble in DMSO, partially
soluble in water and insoluble in methanol, ethanol, chloroform and tetrahydrofuran. It is stable
to air and light. Structural characterization of the molecule was done by analyzing the chemical
shifts (δ) in the 1H and 13C NMR spectra, FT-IR transmission data and CHN elemental analysis.
7

ACCEPTED MANUSCRIPT
BTHA exists in two forms: E and Z amide form (Fig.1). Both of the isomers can be identified
by characteristics infrared frequencies. The FT-IR spectrum of BTHA indicated several peaks
Fig. 1. Schematic representation of Z and E-amide isomers of hydroxamic acid
corresponding to various groups present in the molecular structure (Fig. 2) and the transmission
values were analyzed and compared with the identified values in the literature [24]. The sharp
band at 3168 cm-1 corresponds to amide type –NH stretch while a broad band obtained at 3370
cm-1 attributing for –OH stretch were obtained. The plausible explanation behind the prominent
band at 3019 cm-1 observed in the IR spectra is due to the Fermi-resonance overtone of CNH
bend + C-N stretch. The transmission value at 1640 cm-1 and 1048 cm-1 indicated the presence of
–C=O and –N-O group respectively in the substituted arms. The shifting in the transmission
value of free –OH stretch from 3600 to 3370 cm-1 and carbonyl group from 1690 to 1640 cm-1
implicates intra-molecular H-bonding between –OH and nearly oriented carbonyl group, which
is possible in the Z form of BTHA. Moreover, the appearance of the peak at 1586 cm-1 also
supports the predominance of Z isomeric form over E-form. As reported in the literature, CNH
bend shows peak at 1586 cm-1 corresponding to a characteristic peak for the Z isomer while the
peak at 1490-1440 cm-l is characteristic for the E-isomer [24-25]. The vibrational spectrum
obtained by semi empirical method using PM3 parameter showed closest agreement in terms of
pattern of peaks with the experimental one and confirms the assignments. Predominance
existence of Z-isomer of the ligand with a configuration of higher stability over E-isomer was
8

ACCEPTED MANUSCRIPT
well supported, and, also helped in predicting the non-bonded interaction in the molecule. The
approximate distance between the carbonyl oxygen and hydroxyl hydrogen in Z-amide form was
found to be 2.682 Å clearly indicates the presence of intra molecular H-bonding [22], which is
also evident in the experimental IR spectrum of the ligand.
Fig. 2. FT-IR spectrum of BTHA
The 1H and 13C NMR spectra of the compound under investigation with their characteristic peaks
are given in fig. 3. A singlet in the region 10.8 – 11.8 ppm attributes to the –NH group which
may be due to quadrupole broading as observed in case of nitrogen [26]. Another broad band
appearing at ca. 8.6 – 9.6 ppm indicates presence of –OH group. A multiplet was obtained for
aromatic hydrogens in the region 8.3 – 8.5 ppm. The 13C NMR spectrum illustrates the
absorption signal for the carbonyl carbon at 163 ppm while for aromatic carbons gave chemical
shift at 127 and 129 ppm.
a
9

ACCEPTED MANUSCRIPT
b
Fig. 3. (a) 1H NMR and (b) 13C spectra of BTHA in DMSO-d6 obtained on Bruker Avance II 400
NMR spectrometer
The ligand in 10% DMSO-water mixture shows only one peak at 238 nm which is due to π
π* transition, associated with the aromatic ring [27]. The simulated electronic spectra
obtained from semi-empirical/ZINDO/s method on PM3 optimized geometry also showed
similarity in the peaks pattern as shown in fig. 4a. The pictorial representation of corresponding
HOMO-LUMO molecular orbitals of the ligand causing transitions at aforesaid wavelength are
given in fig. 4b.
Fig. 4. (a) Electronic spectra of BTHA: (i) Experimental; (ii)Theoretical; (b) major transitions
HOMO-7 → LUMO+1 and HOMO-1 → LUMO+2, for two peaks at (I) and (II) respectively,
calculated by Semi-empirical/ZINDO/s
10

ACCEPTED MANUSCRIPT
3.2 Ligand protonation constants
The protonation constants of the ligand BTHA, which may be considered as triprotic LH3, were
determined by potentiometric and spectrophotometric methods. Due to the solubility problem of
the ligand, the potentiometric titrations were carried out in 10% DMSO-water mixture, at µ = 0.1
M KCl and 25 ± 1 ºC. The potentiometric titration curves of pH versus ‘a’, where a= the moles
of the base added per mole of the ligand, is presented in fig. 5. Inflection at a = 0 corresponds to
neutralization of the excess acid added whereas the other inflection at a = 3 implied the
consumption of base for the release of three moles of protons from the triprotic ligand. Analysis
of potentiometric data gave three protonation constants as best fit calculated from Hyperquad
using the equilibrium reactions as shown by the equation 2.
[LHn-1] + [H]
LHn,
K1n
=
(2)
Fig.5. Potentiometric titration curves of the ligand between pH vs. ‘a’, in the absence and
presence of the metal ions; M = Fe(III), Cr(III) and Al(III) in 3:1 metal-ligand ratio, [M] = 3x10-
3M, [L] = 1x10-3M at µ = 0.1 M KCl and T= 25 ± 1 ºC where ‘a’= moles of base added per mole
of the ligand
11

ACCEPTED MANUSCRIPT
BTHA has six dissociable protons and behaves as hexaprotic acid. However, in the present study
only three protonation constants for three –OH groups could be determined. Normally, N-acidity
or O-acidity of hydroxamic acids depends upon the solvents used. In the presence of non-protic
solvent (DMSO), hydroxamic acids favor N-acidity while aqueous polar solvent supports O-
acidity [26]. Further, the amide protons require a high pH condition for the deprotonation due to
poor acidity [28]. Hence, the three protonation constants due to the –NH groups could not be
determined under the present experimental conditions. The simulated curve did not match with
the experimental curve if the species LH1, LH2 and LH3 were considered only in the refinement
process. Inclusion of the hydrolyzed species LH-1 and LH-3 gave the best fit. Further, above pH ~
9, nature of curve indicated no steep change in pH value which could be explained due to the
possibility of hydrolysis in the hydroxamic acids. The mechanism of hydrolysis reactions is cited
in fig. 6.
Fig. 6. Plausible mechanism of hydrolysis of BTHA in alkaline medium( shown only at one arm)
The hydrolysis of hydroxamic acids in the basic medium is well-documented [29]. The
mechanism accounted for the hydrolysis of BTHA at higher pH, involves attack of OH- ion at the
carbonyl group in the presence of polar aqueous solvent forming salt of carboxylic acid and
hydroxyl amine eventually. The different species of the chelator formed are LH1, LH2, LH3, LH-
1, LH-3 in the adopted pH range. The species distribution curve is given in fig. 7 (a). LH3, the fully
protonated form, initially exists in 100% (at pH<5.0). As the pH increases, deprotonation starts
with the formation of LH2 and LH species with the maximum formation of 70 % and 60 % at pH
7.8 and 9.0, respectively. For further confirmation of the species formed and the corresponding
protonation constants obtained by the potentiometric method, spectrophotometric titrations were
carried out in a highly aqueous-DMSO mixture (9:1) at 1 × 10-5 M ligand concentration, 25 ±
1°C and constant ionic strength 0.1 M KCl. The measurements included absorbance with respect
to pH of the solution. The electronic spectra of the ligand were recorded from 200-550 nm in a
12

ACCEPTED MANUSCRIPT
wide pH range of 2.0-11.0 and are given in fig. 7(b). Protonation constants (log K) as well as
hydrolysis constants of the ligand were obtained by the analysis of the experimental data from
best fit curve using HypSpec simulating software [30] as depicted in table 1.
(a)
(b)
Fig. 7 (a). pH dependent distribution of the species of BTHA (b) Electronic spectra of BTHA as
a function of pH (2.0 -11.0) during a spectrophotometric titration where [BTHA] = 1x10-5 M, at
µ= 0.1M KCl and 25 ± 1ºC
Table 1: Protonation constants (log K) and hydrolysis constants of BTHA at 25 ± 1ºC, [L] =
0.1mol dm-3 and µ = 0.1 M KCl in H2O /DMSO (9:1) in pH 2-11 obtained by potentiometric and
spectrophotometric methods
Species
Potentiometry
Spectrophotometry
9.48 ± 0.04
8.56 ± 0.03
7.2 ± 0.05
9.67 ± 0.05
8.74 ± 0.03
7.4 ± 0.05
[LH]/[L] [H]
[LH2]/ [L][H]2
[LH3]/ [L][H]3
-9.84 ± 0.02
-21.89 ± 0.04
- 9.97± 0.02
-22.12 ± 0.04
[LH-1][H]/[L]
[LH-3][H]3/[LH-1]
No noticeable changes in electronic spectra were observed below pH 5.0 and above pH 9.0. The
equilibrium states between protonated and non-protonated forms of the ligand were studied by
examining the shifts in the ligands’ peaks. The protonation constants obtained by
13

ACCEPTED MANUSCRIPT
spectrophotometric method are in agreement with the results obtained by potentiometric method.
The characteristic electronic absorbance spectra of benzohydroxamic acid similar mono
substituted ligand have been reported [27].
Protonation constants of hydroxamic acids, R1C(O)N(R2)OH ( R1 = alkyl/aryl, R2= alkyl/aryl/H)
depend upon the nature of R1 and R2. The protonation behavior of some mono hydroxamic acids
have been studied both in aqueous and solvent mixture media [31]. Acetohydroxamic acid (Aha)
has higher protonation constant value due to the presence of electron donating group i.e., R1 = -
CH3 when compared to Bha, where R1= - C6H5 causes –I effect. Similarly in case of BTHA,
substituted aromatic ring also influences the protonation constants (log K) as shown in table 2.
Table 2: Protonation constants (log K) of mono hydroxamic acids and tri-hydroxamic acids,
BTHA in aqueous medium; T =25 °C, I =0.2 mol/dm3 KCl at µ= 0.1mol/dm3 KCl, 25±1 °C
Ligand
Aha
Pha
Hha
Bha
PhAha
MAha
iPAha
BTHA
R1
R2
-CH3
H
-C2H5 -(CH2)4CH3
-C6H5
H
-CH3
-C6H5
8.47
-CH3
-CH3
8.70
-CH3
-CH(CH3)2
9.26
H
H
log K
9.27
10.57*
9.33
9.36
8.69
9.93*
9.36
8.54
7.24
* in 50-50% (m/m) H2O: DMSO
Literature studies reveal that benzohydroxamic acid has maximum absorbance at 228 nm in
acidic and neutral pH. However, in alkaline condition it shows a new peak at 264 nm in addition
to the earlier one due to stabilization of the excited state by charge delocalization which
corroborates well of the existence of various forms in the basic medium as depicted in fig 8. The
tautomer II fetches the stability through charge delocalization that brings the higher excited state
π* closer to the ground π state and reduces energy of the transition state. Similar pattern was also
observed in the tri-hydroxamic acid substituted benzene (BTHA), a peak was observed at 238
nm in acidic medium and an additional broad peak appeared in the alkaline medium with λmax at
270nm. Deprotonation of the fully protonated ligand LH3 causes formation of a new peak with
-1
concomitant rise in absorbance intensity due to the formation of LH2 , LH-2 and L-3 species. The
14

ACCEPTED MANUSCRIPT
appearance of two peaks in the electronic spectra of the ligand justified by π
π* transitions
of two tautomeric forms of BTHA existing in equilibrium [27].
-
-
(I)
(II)
Fig. 8. Tautomeric forms of mono-substituted benzene hydroxamic acid (Bha)
3.3 Metal complex formation:
Complexation studies of BTHA with Fe, Cr and Al trivalent metal ions were conceded by
potentiometric and spectrophotometric methods in 1:1 and 3:1 metal-ligand molar ratios at ionic
strength µ= 0.1M KCl and temperature 25 ± 1°C. The titration curves of all the metal ions with
the ligand in 3:1 ratio are shown in fig. 5. The symbols represent data collected when no
turbidity or solid phase was there, whereas dashed lines represent points collected when solution
becomes turbid. The precipitation occurred for Fe(III) and Al(III) in the pH range of 5.3-6.0,
whereas in case of Cr (III), no precipitation was occurred. The volume of base used up to ‘a’ = 3,
indicates the release of three moles of protons from BTHA upon complexation with the metal
ions. The buffer region of the ligand shifted towards lower pH in the presence of metal ions with
the release of dissociable protons from the ligand. The larger deviation in the Fe-L curve from
the ligand only curve was more towards lower pH relative to the Cr-L and Al-L, indicates strong
complex formation between iron and the ligand. Hyperquad 2006 program was used for
calculating formation constants of the species using log  values of the ligand as input data. The
results are summarized in table 3. The best fit model obtained after simulation substantiated the
presence of M3LH2, M3L species for Fe(III) and Al(III) metal ions, while M3LH and M3L species
were formed in case of Cr(III) metal ions.
It has been well-established that metal ions exist as hexacoordinated [M(H2O)6] in aqueous
solution [32]. BTHA occupies two coordination sites of each metal ion through O, O- donor
15

ACCEPTED MANUSCRIPT
atoms, while rest sites were occupied by four water molecules. In the presence of Fe (III) and
Table 3: The overall formation constants (log β) of the metal complexes at 25 ± 1º C and µ = 0.1
M KCl (A= Potentiometry and B = Spectrophotometry)
Fe(III)
Cr(III)
Al(III)
A
B
A
B
A
B
[M3LH2]/[M]3[L][H]2
[M3LH1]/[M]3[L][H]
[M3L]/[M]3[L]
34. 35 ± 0.03
34.08 ± 0.08
-------------
-------------
33.98 ± 0.03
34.29 ± 0.08
-------------
-------------
28.98 ± 0.03
21.89 ± 0.08
16.87 ± 0.04
-1.45 ± 0.04
23.55 ± 0.09
19.69 ± 0.08
8.28 ± 0.04
2.40 ± 0.03
23.81 ± 0.09
19.83 ± 0.05
8.36 ± 0.04
2.86 ± 0.04
19.89± 0.04
-------------
-------------
26.19 ± 0.06
16.84 ± 0.04
-------------
28.73 ± 0.07
21.69 ± 0.03
16.54 ± 0.05
- 1.22 ± 0.03
26.13 ± 0.08
16.52 ± 0.03
-------------
[M3LH-2] [H+]2/[M]3[L]
[M3LH-3] [H+]3/[M]3[L]
[M3LH-6] [H+]6/[M]3[L]
-19.76 ± 0.03
- 4.15 ± 0.05
- 4.65 ± 0.03
Al(III) ions, precipitation occurred between pH 5.3-6.0, either due to hydrolysis of metal ions or
by the formation of mixed metal-ligand hydrolyzed species. No precipitation was observed in
case of Cr (III) till pH 11.0. Data above pH ~ 8.5 was not included in the calculation due to the
hydrolysis of the ligand as discussed earlier. Inclusion of hydrolyzed species M3LH-2, M3LH-3,
MLH-6 for Fe(III) and Cr(III); and M3LH-2, M3LH-6 for Al(III) in the model, gave complete
match of the theoretical and experimental curves. Possibility of formation of 1:1 ML
encapsulated complex or mono/dimetallic-complex were discarded as their inclusion in data
either worsen the refinement process or did not bring any improvement in the curves. Various
species can be represented by the following equilibria:
3M + L + 2H
3M + L + H
3M + L
M3LH2
M3LH
log β312 = [M3LH2] [M]3[L][H]2
(3)
(4)
(5)
(6)
(7)
(8)
/
log β311 =[M3LH] [M]3[L][H]
/
M3L
log β310 = [M3L] [M]3[L]
/
3M + L
M3LH-2 + 2H
M3LH-3 + 3H
M3LH-6 + 6H
log β31-2 = [M3LH-2][H+]2 [M]3[L]
/
3M + L
log β31-3 = [M3LH-3][H+]3 [M]3[L]
/
3M + L
log β31-6 [M3LH-6][H+]6 [M]3[L]
=
/
16

ACCEPTED MANUSCRIPT
In order to explore the nature of the precipitates obtained during the titrations of Fe (III) and
Al(III), the precipitates were filtered, washed with distilled water, dried and analyzed through IR
spectroscopic studies (Fig.9). The IR spectra of the precipitates were found to contain peaks of
of BTHA with a negative shift of ~24-37 cm-1 corresponding to C=O, might be due to the
shifting of charge density due to complex formation with the metal ions. This confirms, the
precipitates obtained during experiment above pH 5.3 are due to the M-L complexes, and the
ligand coordinates to the metal ions through two oxygen atoms of each hydroxamate arm.
(a)
(b)
Fig. 9. IR spectra of precipitates obtained for: a) Fe-BTHA; b) Al-BTHA
Alston et al., also claimed the formation of five membered ring complex through O, O-
coordination exhibiting violet color at low pH (Fig. 10) [33].
R
+3
Fig. 10. Coordination of iron with hydroxamate at low pH forms five member ring
The species distribution curves (Fig. 11) show that in an acidic medium (pH ~ 2.0) ligand forms
protonated metal complex M3LH2 with 80 % formation for both Fe and Al metal ions but, no
17

ACCEPTED MANUSCRIPT
such protonated species exists except M3LH, for the Cr ion with 17% formation only. With the
increase in pH, deprotonation of M3LH2 led to the formation of trimetallic M3L species with
maximum formation of 85% at pH ~3.0 and 78 % at pH~ 4.2 for Fe(III) and Al(III) metal ions
respectively. For Cr (III), M3L species was formed with 60% at pH ~ 5.2.
(a)
(b)
(c)
Fig. 11. pH-dependent distribution curves of the species of BTHA metal complexes (a) Fe(III),
(b) Cr(III) and (c) Al(III)
The interaction of transition metal ions with the ligand BTHA, was also studied by UV-Visible
spectrophotometric method in 3:1 metal-ligand stoichiometry ratio. The concentrations of the
ligand, [L] = 1.0 × 10-5 M, and metal ion, [M] = 3 × 10-5M were used to study the complexation
behavior in the wavelength range of 200-550 nm with the increasing of pH to elaborate different
possible forms of the species formed in the solution. The ligand metal interaction accompanied
with major spectral changes, has been depicted in fig. 12.
The hyperchromic shift in the ligand peaks in the presence of the metal ions implied the
coordination of the chromophoric group of the ligand with the metal ions. In case of Fe-BTHA,
the intensities of both the bands were higher as compared to the bands of BTHA, whereas, in
case of Cr(III)-BTHA and Al(III)-BTHA the intensities were lower. Complex formation between
BTHA and Fe(III) ion at pH ~2, was further evidenced by the appearance of a dark violet
coloration during the titration. The broad band at 430 nm is attributed to L
transfer transition.
M charge
18

ACCEPTED MANUSCRIPT
Fig. 12. Electronic spectra of BTHA as a function of pH (2-11) during spectrophotometric
titration with (a) Fe(III), (b) Cr(III) and (c) Al(III), [M]=3×10-5M, [L]=1×10-5M in H2O/DMSO
(9:1) mixture at µ = 0.1 M KCl at T=25 ± 1º C
Above pH 2.5, a reddish orange color along with a hyperchromic effect in the absorption spectra
was observed which may be due to the formation of M3L species, as evident from the species
distribution curve. Further increase in pH (> 5.0) caused formation of a yellow color. The
significant color variation in case of Fe-BTHA complex in a largely aqueous medium is due to
the allowance of a ligand to metal charge transfer transition involving protonated and de-
protonated species formed at varied pH. Fig. 13 represents distinct colour variations in BTHA-Fe
as a function of pH. Although complexation of the ligand with Cr(III) and Al(III) occurred, no
coloration was observed due to a negligible or no charge transfer transition.
Fig. 13. BTHA metal complexes in H2O/DMSO (9:1) solution exhibits color at different pH. Fe-
BTHA complex in test tube A: violet, pH 1.8-2.5; B: reddish orange, pH 2.5-5.0, C: yellow, pH
> 5.0; D and E represent Cr-BTHA and Al-BTHA complex respectively showing no noticeable
color
19

ACCEPTED MANUSCRIPT
The complexation ability of a ligand towards a metal ion at particular pH is better represented by
pM values. Metal ions compete with the protons during complexation in the solution, which
mainly depends upon pH and pKa of the ligand. The corresponding plots of pFe, pAl and pCr
versus pH over a range of 2-9, [L] = 10-3M and [M] = 3x10-3M were shown in fig.14. The pM
+
+3
+3
3
value for Fe (5.25) at pH 5.5 is relatively high as compared to Al and Cr (4.5 and 3.5)
indicating maximum complexation of BTHA with iron followed by Al and then Cr trivalent ions.
The strong coordination ability of the trisubstituted hydroxamate tripod, BTHA, in acidic pH and
subsequent precipitation formation of the complexes near pH 5.5 enables it to be considered as
effective sequestering agent for both Fe+3 and Al+3 ions.
Fig. 14. Plot of pM versus pH for BTHA, pM = -log[M], calculated for [M+3] = 3 x10-3M
and [L] = 1x10-3M
To verify the validity and agreement of the stability constants determined, the thermodynamic
properties of the species involved in the chemical equilibria were calculated by semi-
empirical/PM3 method. Change in the free energy for the metal complexes [L(M(H2O)4)3] where
M=Fe(III), Cr(III) and Al(III) were evaluated as -5434.58, -4564.29, and -4859.75 kJ/mol
respectively. The stability order: Fe3+ > Al3+ > Cr3+, corroborates well with the experimental
studies carried out by potentiometric and spectrophotometric methods and change in free energy
of formation of metal complexes. The enhancement in the stability of Fe(III)-BTHA as compared
20

ACCEPTED MANUSCRIPT
to chromium can be justified on the basis of Irving William stability order for transition metals
[34], according to which the stability of some trivalent transition metal ions, given below,
maintain the trend Co+3(low spin) > Mn+3 > Fe+3 > Cr+3 irrespective of particular ligand
involved. The trend observed by Irving William is largely determined by electrostatic effects and
can be related to the ionic potential (Z/r; Z= ionic charge and r= radius) as well as ionization
potential (i.e. M→ M+n) of metal ions.
Conclusions: The newly synthesized tripodal polydentate ligand BTHA, with the arms of three
hydroxamate binding moieties, shows greater binding efficiency towards Fe(III) as compared to
Cr(III) and Al(III). High formation constants of 3:1 metal-ligand complexes, M3L, can be
explained due to coordination of three pairs of O, O donors in the ligand that show higher
binding affinity towards hard metals, besides forming stable five membered rings chelates. As
seen from the log β values, the ligand BTHA was found to be excellent chelator towards Fe(III)
in acidic range and may be beneficial in developing of a more efficient adsorbent for Fe(III)
metal ions than the existing ones. Thus, the current ligand has the future prospect in reducing the
effect of water pollution containing iron in toxic concentration. The tripod exists predominantly
in Z-amide form which is favorable for intramolecular H-bonding. In the alkaline pH, BTHA
deprotonates giving rises to tautomers which are evident by the appearance of a new additional
peak in the electronic spectra of the ligand. Moreover, Fe-BTHA complex shows distinct colour
variation (violet-reddish orange-yellow) at different pH due to L
M charge transfer; this
enlightens its role as effective colorimetric agent and as potential iron sensor.
References
[1] P. V. Bernhardt, Dalton Trans (2007) 3214-3220
[2] R. Codd, Coord. Chem. Rev 252 (2008) 1387-1408
[3] C. J. Marmion, D. Griffith, K.B. Nolan, Eur. J. Inorg. Chem 2004 (2004) 3003-3016.
[4] I. M. Adcock, Curr. Med. Chem 7 (2006) 966-973
[5] B. Kurzak, H. Kozlowski, E. Farkas, Coord. Chem. Rev 114 (1992) 169-200
[6] Y. K. Agrawal and R. D. Roshania, Bull. Sot. Chim Belg 89 (1980) 159
[7] S. G. Tandon and S. C. Bhattacharya, J Indian Chem. Soc No 47 (1970) 583
[8] B. Saha, S. Chakraborty, G Das, J. Colloid Interface Sci 323 (2008) 26–32
[9] M. Kawahara, M. Kato-Negishi, Int. J. Alzheimers Dis 2011 (2011) 276393
[10] C. A. Shaw, M. S. Petrik, J. Inorg. Biochem 103 (2009) 1555
[11] S. Simon, P. A. Athanasiov, R. Jain, G. Raymond, J. S. Gilhotra, Indian J. Opthalmol 60
(2012) 315-317
21

ACCEPTED MANUSCRIPT
[12] a) L. Polak, In: D. Burrows (Eds) Chromium: metabolism and toxicity, FL: CRC Press,:
51-135
b) H.S. Novey, M. Habib et al., J. Allergy Clin. Immunol 72 (1983) 407- 412.
[13] R. Kakkar, R. Grover, P. Chadha, Org. Biomol Chem 1 (2003) 2200.
[14] T. W. Failes and T.W. Hambley, Dalton Trans (2006) 1995-1901
[15] C. J. Marmion, D. Griffith and K.B. Nolan, Eur. J. Inorg. Chem (2004) 3003-3016
[16] R. Arora, R. Kakkar, Int. Rev. Biophys. Chem (IREBIC) 3 (2012) 212-233
[17] M.M. Khalil and A.E.Fazary, Monatshefte fur Chemie 135 (2004) 1455-1474
[18] E. Farkas and O.Szabo, Inorg. Chim. Acta 392 (2012) 354-361
[19] B. S. Furniss, A. J. Hannaford, P.W.G. Smith, A.R. Tatchell, Vogel’S Textbook of
Practical Organic Chemistry, Prentice Hall, 5th Ed, 2008
[20] P. Gan, A. Sabatini, A. Vacca, Talanta,53 (1996) 1739
[21] a) L. Alderighia, P. Gans, A.Ineco, D. peters, A. Sabatini, A. Vacca, Coord. Chem. Rev
184 (1999) 311
[22] HyperChem Version 7.5, Hypercube Inc., 419 Phillip Street, Waterloo, Ontario, Canada,
N2L, 3X2
[23] CAChe version 6.11 (2003) Fujitsu Limited
[24] a) F. S. Higgins, L. G. Magliocco, N. B. Colthup, Appl. Spectrosc. 60 (2008) 279-287
b) A. Kaczor, J. Szczepanski, M. Vala, H. Kozlowski, L. M. Preoniewicz, Phys. Chem.
Chem. Phys 5 (2003) 2337- 2343
[25] B. Garcia, S. Ibeas, J.M. Leal, F. Secco, M. Venturini, M.L. Senent, A. Nino, C. Munoz,
Inorg. Chem 44 (2005) 2908-2919
[26] J. Schraml, M. Tkadlecova, S. Pataridis, L. Soukupova, V. Blechta, J. Roithova, O.
Exner, Magn. Reson. Chem 43 (2005) 535–542
[27] R. E. Plapinger, J. Org. Chem 24 (1959) 802–804
[28] B.K. Kanungo, M. Baral, S.K. Sahoo, S.E. Muthu Spectrochimica Acta A: Molecular
and Biomolecular Spectroscopy 74 (2009) 544–552.
[29] a) A. Ahmad, J. Socha and M. Večeřa Collect. Czech. Chem. Commun 39 (1974) 3293-
3303
b) K. K. Gosh, K. K. Krishnani, S. Ghosh, Indian J. Chem 38B (1999) 337-342
[30] P. Gans, A. Sabatini, A. Vacca, Ann. Chim. (Rome) 89 (1999) 45
[31] E. Farkas, E. Kozma, M. Petho, K. M. Herlihy, Micera G, Polyhedron 17 (1998) 3331-
3342
[32] S. K. Sahoo, R.K. Bera, M. Baral, B. K. Kanungo, J. Photochem. Photobiol., A 188
(2007) 298-310.
[33] W. C. Alston II, K. Ng, Forensic Sci. Int. 126 (2003), 114-117
[34] A. Ciferri, and A. Perico, Ionic interactions in natural and synthetic macromolecules,
John-Wiley & Sons, 2012
22

ACCEPTED MANUSCRIPT
Graphical Abstract
Highlights
 A tri substituted hydroxamate tripodal ligand developed, forms thermodynamically stable
complexes with high log β and pM values with trivalent metal ions : Fe, Cr and Al.
 The ligand with short arms of hydroxamic acid renders restriction for a 1:1 ligand- metal
encapsulation showing a large molecular strain energy
 Preferentially the ligand forms trimetallic complex of the type,[L(M(H2O)4)3], hence, can
be applied as potent adsorbent of the metal ions in solution present in toxic concentration.
 Fe(III)-L exhibits three distinct colors (violet-reddish orange-yellow) at different pH
implying the significant role of the chelator for the application as colorimetric sensor.
 The theoretical investigations confirm the experimental findings.
23