Post on 16-Feb-2020
Chapter 3
Tris(aminomethyl)ethane-5-oxine
(TAME5OX) and its Interaction with
Trivalent Metal Ions
Chapter 3
Tris(aminomethyl)ethane-5-oxine and its
Interaction with Trivalent Metal Ions
3.1 Introduction
Bidentate 8-hydroxyquinoline (8HQ), is a quinoline derivative originating in plants
as well as from synthesis. Out of seven potential isomeric mono-hydroxyquinolines,
only 8-hydroxyquinoline can readily form a five-membered chelate. Owing to its
multifunctional properties, such as diverse bioactivities, therapeutic potentials,
magnificent photophysical applications like sensors, liquid crystalline, nanoscale
materials, redox-active switches etc. [1], this scaffold is employed as molecular
platform in coordination chemistry to construct many unique ligand systems for the
recognition of important metal ions especially main group, transition and rare-earth
metal ions. The chemistry of 8-hydroxyquinoline experiences a renaissance due to
the importance of its aluminum complex as emitter for organic light emitting
devices (OLEDs), resulted studies on synthesis and performances of many 8-
hydroxyquinoline derivative ligands and their different metal complexes [2]. In
addition, this moiety has been also employed as a functional receptor for metal ions
owing to its unique pyridyl -N and -OH structural characteristics [3]. Research
showed that the modification of an 8-hydroxyquinoline ligand by changing its
substituents at different positions have changed its selectivity and coordination
behaviour with different metal ions [4]. To enhance metal binding selectivity of 8-
HQ, introduction of additional moieties (central unit and connecting groups) at the
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C-2, C-3 and/or C-7 positions, adjacent to the original binding sites in 8-HQ, has
been extensively examined [5-6].
Moreover, a number of chelators that are based on triamines as central unit
with zero to three pendant donor groups have been studied. Of the fully aliphatic
tripodal amine ligands, the C3v symmetric ligands tren (tris(2-aminoethyl)amine),
trpn (tris(3-aminopropyl)amine) and tame (tris(1-aminomethyl)amine) in which all
three arms are of the same length, represent ideal fragments to construct variety of
multidentate ligands. A multidentate ligand offers two advantages: a higher
thermodynamic stability than its bi- or tridentate analogues, and a metal complex
stability indifferent to a dilution effect. Martell and Hancock have pointed out that
the close proximity of the donor atoms maximizes the number of 5- and 6-
membered chelate rings formed upon metal ion complexation, and this, in turn,
maximizes the entropic contribution to a high thermodynamic stability [7]. As a
result, a number of ligands have been prepared offering a variety of donor atom sets
[8-9] are shown in Figure 3.1. Accordingly, rich coordination chemistry has
developed, and there exists a body of structural and thermodynamic information for
a variety of metal complexes of this type. Detailed review on such ligands is
presented in chapter 2.
In the quest for new chelators with the intent of developing efficient binders
and fluorophores for trivalent metal ions, it is germane to employ a polyfunctional
multidentate ligand. The first hexadentate ligand O-Trensox in tripodal
orchestration [10], composed of three 8-HQ binding units linked to central amine
backbone of tris(2-aminoethyl)amine (Tren) has been reported by Serratrice et al.,
Following O-Trensox approved as strong chelator for Fe(III) over a wide range of
pH, the structural and aqueous coordination chemistry of several multidentate 8-
hydroxyquinoline ligands for iron and other metal ions have been determined
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(Figure 3.1). The ferric complex of COX200 [11] (an analog of Cox2000 in which
the 8-hydroxyquinoline subunits are connected to the tripodal spacer, based on a C-
pivot scaffold grafted with a polyoxyethylenic chain units) was studied by X-ray
crystallography and was the first example of a tripodal tris-hydroxyquinoline
structure with an 8-hydroxyquinolinate coordinating mode compared to the ferric
complex of O-TRENSOX [12] which was crystallized in acidic medium showed
salicylate coordinating mode. The structures of ferric complexes of these two model
ligands are presented in Figure 3.2. It was observed that grafting a polyoxyethylenic
chain does not change the structure of the coordination sphere and has a negligible
effect on the pFe. Indeed, the COX-2000 derivative exhibits a pFe value of 29.1,
which is very close to that of O-TRENSOX. In order to compete effectively with
transferrin for iron, a ligand should have a pM greater than 21.6, the pM of a human
serum transferrin (calculated with [HCO3-] = 0.23 M) [13].
Figure 3.1: Structures of Tripodal O-TRENOX, N-TRENOX (and their sulfonyl
derivatives), Csox, Cox200, Cox750 and Cox2000.
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Figure 3.2: X-Ray Structures of [Fe(COX-200)] (left) [11] and of [Fe(O-TRENSOX)]
(right) [12].
It is noted, however that these investigations hitherto reported are limited
mainly to 8-HQ derivatives with substituents on the C-2 or C-7 position. Derivation
at C-5 position of the 8-HQ has not been explored in detail. A possible reason could
be that due to the para position of pyridyl N and OH groups in this arrangement, not
much room remains for encapsulation of metal ions in the central cavity of ligand
which facilitate coordination with more strain. However, further exploration
revealed that substitution at the C-2, C-3 and/or C-7 positions in the 8-HQ moiety
mainly tunes the coordinating property of the pyridyl N group, while substitution at
the C-5 positions induces change in the binding ability of the OH group [14].
Keeping in view the scarcity of polydentate 5-linked 8-HQ based tripod, to compare
from theoretical standpoints with its -2 and -3 linked counter parts, ease of synthesis
and to mimic the C-3 symmetric tripodal siderophores, the present study was
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undertaken to develop a new C-3 symmetric tripodal hexadentate 8-HQ based
ligand linked at 5-position through amine (Figure 3.3) and to evaluate its
coordination and photophysical properties. The design of TAME5OX relies on a
short tris(aminomethyl)ethane (TAME) backbone as an anchor symmetrically fitted
with three individual 8-hydroxyquinoline (8HQ) moieties at the C-5 position. The
spacer bears a secondary amine function, and its length has been chosen to achieve
a tight orientation of binding units. The foremost rationale is to explore the pH
dependent binding and fluorescent behaviour of the edifice.
Figure 3.3: Structure of 5,5'-(2-(((8-hydroxyquinolin-5-yl) methylamino) methyl)-2-
methylpropane-1,3-diyl) bis(azanediyl) bis(methylene) diquinolin-8-ol (TAME5OX).
In section 3.2.1 of this chapter, synthesis, characterization, protonation (or
deprotonation) behaviour, photophysical and DFT studies of the new 8-
hydroxyquinoline based tripodal chelator, namely 5,5'-(2-(((8-hydroxyquinolin-5-
yl)methylamino)methyl)-2-methylpropane-1,3-diyl)bis(azanediyl)bis(methylene)
diquinolin-8-ol, (TAME5OX), (Scheme 3.1) are discussed.
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The second section (3.2.2) of the chapter presents the synthesis, solution
thermodynamic evaluation and photophysical properties of the coordination
compounds of the titled ligand with three trivalent metal ions viz., Fe3+
Al3+
and
Cr3+
. The coordination behaviour of the chelator in the corresponding coordination
compounds are evaluated using combination of absorption and emission
spectrophotometry, potentiometry, infra red, electrospray mass spectrometry and
theoretical investigations. The pM values for these metals ions, calculated at pH
7.4, indicating TAME5OX is a powerful synthetic metal chelator. The density
functional theory is employed for study of electronic behaviours like evaluation of
vibrational modes, NBO analysis, excitation and emission properties of the species
formed during solution studies to validate the experimental findings and elucidate
the proposed molecular structures of these complexes.
Following the theme established in section 3.2.1 and section 3.2.2, a systematic
study on the interactions of the TAME5OX with four lanthanide ions viz., La3+
,
Eu3+
, Tb3+
and Er3+
are explored in section 3.2.3. The thermodynamic stability and
aqueous coordination chemistry of the chelator with the said lanthanide ions have
been probed by same techniques as mentioned above, except theoretical studies by
semiempirical sparkle model method. The coordination geometries are optimized
using PM7 as sparkle/PM7 model and the theoretical spectrophotometric studies are
also carried out in order to validate the experimental findings, based on ZINDO/S
methodology at configuration interaction with single excitations (CIS) level.
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3.2 Results and Discussion
3.2.1 Tris(aminomethyl)ethane-5-8-hydroxyquinoline (TAME5OX): Design,
Synthesis, Charaterization, Solution thermodynamics, Photophysical and DFT
Studies
3.2.1.1 Design, Synthesis and Characterization
For an efficient biomimetic siderophore based on enterobactin, a metal chelator for
trivalent metal ions with high complexing ability requires a topology which has a
pivotal basic skeleton with suitable spacer length with attachment position of the
binding unit (cf. Figure 1.10 from Chapter 1); it is essential that the three electron
donating arms should arrange themself around the tripodal ligand to work
cooperatively for uptake of the cation. The molecular edifice containing tris
(aminomethyl) ethane as a central unit, anchored on a quaternary carbon, and
extended by three diverging arms bearing the oxine coordination sites, connected at
the C-5 position of the 8-hydroxyquinoline than commonly used C-2, C-3, or C-7
positions [14] was considered for the study. The whole synthesis of the TAME5OX
was performed through a multistep pathway by following the general plan as
depicted in Scheme 3.1, leading to a more reliable procedure.
Commercially available (Sigma-Aldrich) 1,1,1-tris(hydroxymethyl)ethane, 1,
the synthetic precursor bearing a methyl group at the bridgehead carbon, was
converted to its tritosylate 1a by a nucleophilic substitution reaction with tosyl
chloride in THF. Intermediate tritosylate 1a was isolated in high yield as white
crystalline solid which was further reacted with sodium azide in DMSO to afford
the corresponding triazide 1b. The crude azide was obtained as slightly yellow oil
and was used immediately after work-up to avoid explosion risks. Subsequent
reduction of the triazide with LiAlH gave the required triamine (1,1,1 tris amino
methyethane) 1c in good yield. Since the triamine is a malodorous air sensitive
liquid, it was converted to its corresponding salt as stable white solid by passing dry
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HCl gas into it. Triamine has been prepared by reported procedure [15] with slight
changes.
The derivatization of oxine starts from 8-hydroxyquinoline 2, a very cheap
precursor, which is easily transformed into a 5-chloromethyle derivative 2a by
chloromethylation of 8-hydroxyquinoline at 5-position [16] (Scheme 3.1.1). The
chloromethylation step is best done with formaldehyde (40%) and HCl (37%) [17].
It has been reported that the synthesis of the free base 2a could be carried out by the
reaction of its hydrochloride salt with sodium hydrogen carbonate (NaHCO3) [18];
however, 2a is unstable and readily undergo hydrolysis to form hydroxymethyl
quinolinol [17]. Hence, the hydrochloride salt of 2a was directly used for coupling.
Scheme 3.1: Reagents and conditions: (i) NaOH:H2O, TsCl, THF, 0°C, 6 hrs; (ii) NaN3,
DMSO, 137°C, 12 hrs, under N2; (iii) LiAlH4, dry THF, reflux, 18 hrs, under N2; (iv)
HCHO, 37% HCl, 50°C, 7 hrs, dry HCl gas;(v) K2CO3, acetone:water, reflux, 3 hrs.
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The synthesis of the final product, novel tripodal ligand based on C-3
symmetrical carbon anchor (TAME5OX 3), involve direct condensation of
hydrochloride salts of 1c and 2a. The reaction was performed by a conventional
method, refluxing in acetone water mixture (because of insolubility of
hydrochloride salt of 1c and 2a in organic solvents) in presence of K2CO3 for 3 hrs,
which led formation of the final product, TAME5OX, 3. The compound was
isolated as an air-stable greenish off-white solid, recrystalised from chloroform as
pure compound with 96% yield and was characterized by CHN analysis, FTIR, 1H,
and 13C NMR and mass spectrometry.
3.2.1.1.1 Infrared spectra
The absorption peak at 3399 cm-1
of TAME5OX is ascribed to the -OH stretching
of 8-hydroxyquinoline unit. The band observed at lower frequency than the normal
position of free oxine, 3730-3520 cm-1
, is attributed to presence of intramolecular
hydrogen bonding in the ligand, which can be also seen in 1H NMR of TAME5OX.
The appearance of bands at 2925 and 1469 cm-1
are due to N-H stretching and
bending vibrations, respectively of -CH2-NH-CH2- spacer, and no peak attributable
to unreacted amine or chloromethylene group was group was present. The
absorption peaks around 1599, 1556 and 1389 cm-1
are due to the aromatic
carboncarbon stretching vibrations due to 8hydroxyquinoline nucleus [17]. The
weak bands around 2112 and 1786 cm-1
may be due to asymmetric and symmetric
stretching vibrations of methylene groups (-CH2-), and near 1300 cm-1
and 1320
cm-1
are due to asymmetric and symmetric stretching vibrations of the -CH2-NH-
CH2-bridge.
In order to confirm the assignments of IR frequencies, theoretical calculations
were carried out by DFT method at B3LYP/6-31G* level. The calculated IR
frequencies of the TAME5OX as well as the experimental spectra are shown in
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Figure 3.4. For comparison, the theoretical and experimental values are also given
in Table 3.1. The calculated vibrational frequencies are obtained at higher values
than experimental ones, but their trend are corroborate with the experimental data
indicating that theoretical results can be implemented for the interpretation of
experimental results. The differences in calculated and experimental spectra may be
mainly due to, somewhat, to the fact that the experimental spectra were measured in
the solid phase and the theoretical spectra were obtained in the gas phase. However,
a linear correlation obtained between the theoretical and experimental data (Figure
3.5) gave an acceptable correlation.
Figure 3.4: Infra-red (IR) spectrum of TAME5OX (a) experimental (ATR) (b) theoretical
(calculated by B3LYP/6-31G*).
NH
NH
NH
N
N
N
OH
HO
HO
3
89
Figure 3.5: Correlation between the experimental and theoretical IR frequencies of
TAME5OX
3.2.1.1.2 NMR Spectra
The 1H NMR of TAME5OX presents set of signals given in experimental section
3.4 with the assignments proposed on the basis of the numbering scheme are given
in „a‟ of Figure 3.6. On proton NMR spectra of the ligand, signals were observed
characteristics to OH, CH, CH2 and CH3 protons. Signals due to CH of 8-
hydroxyquinoline were observed in the aromatic region of the spectrum in the range
of 8.87-6.9 ppm (inset of Figure 3.6) which showed two doublets at 8.4 and 8.8
ppm, one double doublet at 7.5 and one triplet around 6.9 ppm. Singlets at 1.1, 2.0
and 2.4 ppm are assigned to the protons of one methyl (-CH3) and two methylene (-
CH2) groups: one attached to carbon of centre unit, and the other to 8-
hydroxyquinoline unit, respectively. TAME5OX showed a signal at 4.61 ppm due
to NH proton and board peak at 9.74 ppm for hydroxyl proton of quinoline ring;
this weak and broad band of hydroxyl protons observed were most probably
resulted from intramolecular H-bonding of OH proton with N atom of quinoline
unit [19].
90
Figure 3.6:
1HNMR spectrum of TAME5OX: (a) experimental (DMSO, 300 MHz), inset
showing splitting of protons of aromatic part of ligand, and (b) calculated (applying
DFT/B3LYP/6-31G* method).
The 13
C NMR spectrum of TAME5OX, with the assignments proposed on the
basis of the numbering scheme, is given in Figure 3.7. The ligand showed bands at
34, 56 and 61 ppm due to the CH3 and CH2-NH-CH2 carbons, respectively. Signals
at 148, 139, 135, 127, 126, 122, 120 and 110 ppm attributed to the CH groups
belong to aromatic oxine unit, whereas the peak at 153 ppm corresponds to carbon
attached to -OH present in the oxine unit.
The theoretically calculated 1H and
13C NMR chemical shifts of TAME5OX
obtained through DFT/B3LYP method have been compared with the experimental
data in “b” of Figure 3.6 and Table 3.1. The methyl protons showed singlet peak at
2.472.47
2.47
1.16
2.0 2.02.0
3.81
6.80
7.06
8.44
7.60
8.86
3.81
7.06
6.80
8.86
7.60
8.44
3.817.06
6.80
8.86
7.60
8.44
9.83
9.83
9.83
NH
NH
NH
N
N
N
OH
HO
HO
(a)
(b)
91
Figure 3.7:
13C NMR spectrum of TAME5OX (DMSO, 300 MHz), insets showing chemical
shifts due to aromatic carbons (left), and aliphatic carbons (right).
1.2 ppm. The methylene protons of spacer showed signals in the spectrum at 2.05
ppm due -CH2 group attached to central unit carbon and at 2.47 ppm due -CH2
group attached to 8-hydroxyquinoline unit. Five peaks at 6.8, 7.2, 7.6, 8.4 and 8.9
were calculated for the aromatic -CH protons of 8-hydroxyquinoline unit in contrast
to experimental findings where only four signals were obtained for these protons.
The double doublet appeared at 6.94 ppm in the experimental 1HNMR spectrum of
TAME5OX due to 2 -CH protons were calculated theoretically as two doublets
separately at 6.8 and 7.2 ppm as shown in “b” of Figure 3.6. In addition the
experimentally observed peak for hydroxyl proton at 9.74 ppm, was calculated
theoretically at 9.78 ppm which was found merged with aromatic protons of oxine
group in both cases due to intramolecular hydrogen bonding. The calculated
chemical shifts gave qualitative trends although the values are not in quantitative
agreement with the experiment measurements, whereas the linearity in curves
between the calculated 1H and
13C NMR values with their respective experimental
values (Figure 3.8, a-b) gave a perfect correlation.
92
Figure 3.8: Correlation between the experimental and theoretical (a)
1HNMR, and (b)
13CNMR, data of TAME5OX.
Table 3.1: Calculated (B3LYP/6-31G*) and Experimental (DMSO, 300 MHz)
NMR (1H and
13C) Chemical Shifts (ppm), and IR Frequencies (ATR, cm
-1) for
TAME5OX
NMR IR
Proton B3LYP/6-31G* Experimental Carbon B3LYP/6-31G* Experimental B3LYP/6-31G* Experimental
H1 1.18 1.2 C1 36.78 37.01 3619 3470
H2 2.05 2.0 C2 41.23 42.65 2998 2932
H3 2.5 2.4 C3 55.99 57.45 1649 1594
H4 4.73 3.61 C4 61.72 62.66 1556 1454
H5 7.45 7.41 C5 111.1 111.44 1383 1396
H6 6.94 6.85 C6 116.65 117.88 1152 1239
H7 7.5 7.55 C7 122.04 121.93 1087 1172
H8 8.42 8.43 C8 128.31 127.57 1022 1080
H9 8.84 8.85 C9 135.82 135.97
H9 9.98 9.74 C10 137.86 138.34
C11 149.14 148.87
C12 154.21 152.51
C13 157.39 153.58
3.2.1.1.3 Mass Spectra
High-resolution electrospray ionization time of flight mass spectrometry TOF
MS(ES+) was used to identify the product by evaluation of the molecular ion peak
of TAME5OX (Figure 3.8). The spectrum of TAME5OX showed two peaks at m/Z
598.2(20%) and 599.3(10%) corresponding to the neutral [M]+ and protonated
[M+H]+ form of ligand, respectively.
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Figure 3.9: MS(ES+) mass spectrum of TAME5OX.
3.2.1.2 Ligand Properties: Protonation equilibrium, Solution thermodynamics
and Photophysical studies
3.2.1.2.1 Ligand Protonation Constants
In order to evaluate the competition between the protons and metal ions for the
coordination sites, it was necessary to determine first the acid-base properties of the
ligand. The deprotonation constants of TAME5OX were determined separately by
potentiometry, UV-visible and luminescence spectrophotometry methods. Owing to
insolubility of nonadentate ligand in water, it was dissolved in appropriate amount
of standard 0.1M HCl to make it water soluble by converting into its hydrochloride
salt. The neutral TAME5OX is addressed as LH3, whereas the fully protonated form
of TAME5OX is considered as a potentially nona-protic acid and is denoted as
(H9L)6+
. It possesses nine potential protonation sites: three phenolate oxygen atoms,
three pyridine nitrogens and three secondary amine nitrogen atoms of pendant arms.
The extracted pKa values (protonation equilibria) obtained from titrations in the pH
range 1.6-12.5, are defined by the following equations:
𝐻10−𝑖𝐿 7−𝑖 + ⇋ 𝐻9−𝑖𝐿
6−𝑖 + + 𝑖𝐻+ (3.1)
𝐾𝑎𝑖 = 𝐻9−𝑖𝐿 6−𝑖 + H+ i 𝐻10−𝑖𝐿
7−𝑖 + 𝑖 = 1 − 9 (3.2)
The potentiometric titration of the protonated form (H9L)6+
of the ligand was carried
out in the pH range of 1.6-12.5 in 0.1 M KCl with KOH at 25ºC. Due to its poor
solubility in water (as precipitation below pH 7.0 at concentration 1×10-3
M-1
×10-4
M),
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the final concentration was maintained to 1×10-5
M and data were analyzed by the
HYPERQUAD program [20]. Analysis of potentiometric curve (Figure 3.10) allowed
the determination of nine protonation constants (the number in parentheses
corresponds to the standard deviation in the last significant digit): log K1 = 9.89(4),
log K2 = 8.92(5), log K3= 8.42(4), log K4 = 6.71(3), log K5 = 6.38(2), log K6 =
5.88(1), log K7 = 4.60(2), log K8 = 4.12(1) and log K9 = 3.64(6).
Figure 3.10: Potentiometric titration curve of TAME5OX (1×10
-5M). Solvent H2O, I =
0.1M (KCl), T= 25(2)°C, and ‘a’ is the moles of base added per mole of TAME5OX present.
Symbols and solid lines represent the experimental and calculated data, respectively.
UV-visible spectrophotometric titrations were also carried out showing three buffer
regions in the pH range of 1.97-10.95: one, in the pH range of 1.97-6.69 („a‟ of
Figure 3.11; corresponding spectra exhibit three isosbestic points at 217, 250 and
272 nm), second region in the pH range 7.539.35 (“b” of Figure 3.11, spectra
exhibit three more isosbestic points at 277, 295 and 315 nm) and the third buffer
region in the pH range 9.6210.81 („b‟ of Figure 3.11, spectra overlayed exhibit
max at 325nm). Analysis of spectrophotometric data performed by the program
Hypspec [21] gave a model with same number of absorbing species (as observed by
potentiometric titration) involving three proton exchange in each pH range and also
nine deprotonation constants, which are in good agreement with that determined by
potentiometric method.
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Figure 3.11: UV-vis absorption spectra of TAME5OX (1.0×10
−5M) as a function of pH: (a)
pH = 1.97−6.69; (b) pH = 7.53−10.81. Solvent: H2O, I = 0.1M (KCl), T = 25.0(2)ºC.
In consideration to further support the number and nature of species obtained
by both potentiometric and spectrophotometric methods and the corresponding
deprotonation constants discussed above, luminescence titrations were also carried
out for 1×10-5
M solution of TAME5OX from 1.97-10.81 pH range. The typical
example of experimental data for the luminescence titration of the ligand is depicted
in Figure 3.12. The best refinement of the luminescence-pH data by the Hypspec
program corresponds to the presence of nine species, in concurrence with the
potentiometric and spectrophometric data; the values of Log Ks are returned in
Table 3.2 (the number in parentheses corresponds to the standard deviation in last
significant digit).
Figure 3.12: Dependence of the fluorescence emission spectra of aqueous solution of
TAME5OX (1.0 × 10-5
M) on the change of the pH value from acidic to basic range (1.97-
10.95 pH), T = 25.0(3)°C (excitation and emission slit widths of 5.0 nm) with an excitation
wavelength of 385 nm.
All constants are reported in Table 3.2 together with some similar ligands
containing 8-hydroxyquinoline or its sulphonate derivative (oxinobactin,
96
sulfoxinobactin) [22], Tsox, TsoxMe) [23], (OTRENSOX) [24] based on the
backbone of tris(2aminoethyl)amine), (COX2000 [11], based on a C-pivot scaffold
grafted with a polyoxyethylenic chain units). The structures of the reference ligands
have appeared elsewhere: O-Trensox and Cox2000 Figure 3.1, for Tsox, TsoxMe:
Figure 3.34 in section 3.2.3 of this chapter, Oxinobactin and Sulfoxinobactin:
Figure 5.1 in chapter 5. The three highest log K values found between 8.42 and
9.89 and three lowest values between 3.64 and 4.60 are assigned to the
protonation of the hydroxyl and pyridinium nitrogen moieties of the three
quinolinate groups, respectively, whereas the log K between 5.88 and 6.71 were
attributed to the secondary amines for these ligands. Inspite of the differences in
solvents or other experimental conditions, the log K values of TAME5OX and these
reported ligands can be compared: the average value for the log K of the oxygen
sites, 9.07 for TAME5OX, is much higher than the corresponding value (7.73 for
oxinobactin, 7.02 for sulfoxinobactin, 8.07 for O-TRENSOX and 8.65 for Tsox).
The same trend was observed for protonation of the nitrogen sites: the average
value is 4.13 for TAME5OX and (3.02 for oxinobactin, 2.07 for sulfoxinobactin,
2.46 for OTRENSOX, 3.15 for Cox2000 and 3.13 for Tsox). This differences
seems to be due to the absence of an electron withdrawing effects of sulphonate
group in case of oxinobactin and sulfoxinobactin [22]
and carbonyl group of amide
linkage at ortho position to hydroxyl moieties [23,24]. The higher values of
TAME5OX could be due to electron donating effect of CH2-NH-CH2 linkage
attached at 5-position (para to hydroxyl group) of 8-hydroxyquinoline. Moreover,
the protonation constant (log K) ranges for the nitrogen and oxygen atoms differ
from the statistical factor of (log 4 = 0.6), which points to cooperativity taking place
between the three arms of the tripodal ligand, probably via intra- or intermolecular
hydrogen bonds, similar to the situation for oxinobactin and Tsox, but in contrast to
O-Trensox.
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Table 3.2: Protonation constants[a]
(logK) of TAME5OX, (p = potentiometry, s = spectrophotometry and f = fluorometrically) and ligands
containing 8-hydroxyquinoline subunits.
Ligand log K1 log K2 log K3 average
log K1-3
log K4 log K5 log K6 average
log K4-6
log K7 log K8 log K9 average
log K7-9
TAME5OX[b]
9.89(4)p
9.82(5)s
9.79(3)f
8.92(5)p
8.97(4)s
8.87(6)f
8.42(4)p
8.39(3)s
8.36(7)f
9.07p
9.06s
9.00f
6.71(8)p
6.79(7)s
6.63(3)f
6.38(2)p
6.35(1)s
6.30(5)f
5.88(4)p
5.82(3)s
5.80(1)f
6.32
6.32
6.24
4.60(4)p
4.65(5)s
4.63(3)f
4.12(2)p
4.08(1)s
4.16(1)f
3.64(4)p
3.67(3)s
3.62(6)f
4.12
4.13
4.13
Oxinobactin[c]
8.51(3) 7.78(2) 6.90(4) 7.73 4.07(7) 2.75(9) 2.28(2) 3.02 - - - -
Sulfoxinobactin[c]
7.84(1) 7.03(1) 6.18(2) 7.02 2.78 ~2 1.42(3) 2.07 - - - -
Tsox[b,d]
9.46(3) 8.44(4) 6.4(1) 8.65 4.4(1) 3.2(1) 1.8(1) 3.13 - - - -
TsoxMeb,d
9.19(7) 8.20(9) 6.65(11) 8.01 3.94(14) - - - - -
OTRENSOX[b,e]
8.62(4) 8.18(5) 7.44(4) 8.07 3.10(4) 2.55(8) 1.83(1) 2.46 - - -
NTRENOX[b.e]
9.69(1) 8.73(1) 8.26(2) 8.98 6.99(2) 6.44(1) 3.53(7) - 3.10(6) - - -
COX2000b,f
10.38(1) 8.40(1) 6.80(2) 8.53 4.13(2) 2.89(2) 2.43(1) 3.15 - - - - aNumbers in parentheses represent the standard deviation in the last significant digit.
bIn water (I = 0.1M KCl).
cIn methanol/water (80/20 w/w, I
= 0.1M NaClO4). cReference [22].
dReference [23].
eReference [24].
fReference [11]
.
98
The corresponding distribution curves of TAME5OX obtained from the log K
values are presented in Figure 3.13 which indicate that in acidic solution
TAME5OX initially exists 100% in the fully protonated form as LH9 below pH 2.9.
As the pH is increased, deprotonation starts, the ligand loses two protons
subsequently from the Npyridyl of oxine unit with formation of LH8 and LH7
which exist between pH 0.8-4.0 and 1.2-4.8 with maximum concentration of 28%
and 81% at pH 2.8 and
~3.4 respectively. Further release of protons lead formation
of LH6, LH5 and LH4 successively in the pH range of 2.9-6.5 with maximum
concentration 42%, 49% and 63% at pH 4.1, 4.8 and 5.2, respectively. Further
deprotonation leads to the formation of LH3, the neutral form of TAME5OX, in pH
range of 4.9-7.4 with maximum concentration 57% at pH 6.3. Other species LH2
and LH were formed after pH 5.2 and 7.5, with maximum concentration of 82% and
61%, respectively. No isosbestic points in the electronic spectra („b„ Figure 3.11)
were observed above the pH 9.62 and the spectra overlapped with each other.The
fully deprotonated ligand (L) was predominant above pH 9.0with 100%
concentration.
Figure 3.13: Species distribution curves of TAME5OX containing species, computed from
the protonation constants given in Table 2. Calculated for [TAME5OX]tot = 1×10-5
M,
Solvent : H2O, I = 0.1M (KCl), T = 25.0(2)°C.
99
In order to validate the assignment of experimental protonation constants,
aqueous-phase free energy for the entire representative species of TAME5OX were
calculated by density functional theory by B3LYP/6-31G* quantum mechanical
approach. For an acid species AH, the pKa defined as the negative logarithm of the
dissociation constant of the reaction AH = A- + H
+, is given by the thermodynamics
relation pKa= ∆Gaq,AH/2.303 RT. The change in free energy in aqueous-phase is
obtained according to the following equation:
∆Gaq ,AH = Gaq ,A− + Gaq ,H+ − Gaq ,AH (3.3)
For which the proton free energy at 298 K and 1 atm is:
Gaq ,H+ = 2.5RTTS° =1.487.76
= 6.28 kcal=mol
The greater acidity belongs to the group in which its hydrogen release is easier
and there must be a decrease in G values in the deprotonation process. A clear
decrease in G was observed in case of TAME5OX (Figure 3.14) as the
deprotonation take place from protonated N of 8hydroxyquinoline from fully
protonated free ligand, LH96+
, followed by protonated amines of pendant arms and
hydroxyl groups of binding units. The validity of this theoretically calculated Gaq
for the different species of TAME5OX was compared with the calculated
experimental pKa, which resulted an acceptable correlation with R2 = 0.9973.
Figure 3.14: Correlation between the experimental log K and calculated ∆Gº of
TAME5OX.
100
3.2.1.2.2 Absorption spectra of ligand (TAME5OX) and its protonated species
The ligand TAME5OX shows two peaks 255 and 310 nm (ε = 110, 000 M−1
cm−1
and 60, 100 M−1
cm−1
, respectively) in the UV-visible region assigned to →* and
n→* transitions. Since, a variety of species are formed with change on pH, the
large spectral changes (Figure 3.11, a-b) observed in the spectrophotometric
titration of the TAME5OX as a function of pH is indicative of a change in the
protonation and deprotonation behaviour of the ligand. At pH ≤ 2.9, the electronic
spectrum of the TAME5OX in its protonated form is symbolized by a strong
absorption band at higher energies 263 nm, (ε = 110, 000 M−1
cm−1
) assigned for
→* transitions and a broad band at 380 nm, (ε = 5, 000 M−1
cm−1
) assigned for
n→* transition, respectively. Upon rising pH above 2.9, the deprotonation of the
pyridinium nitrogen results in hypochromic shifts (263 nm, ε = 72,000 M−1
cm−1
and
380 nm, ε = 2, 300 M−1
cm−1
) with the simultaneous ascent in absorbance towards
lower wavelength, 230 nm (ε =32, 000 M−1
cm−1
) which results in the formation of
three isosbestic points at 217, 250 and 272 nm as shown in Figure 3.11. The
isosbestic points assert the formation of three protonated species of the ligand in the
pH range of 1.9-6.69, consequently which can be well comprehended from species
distribution curves (Figure 3.13).
Above pH 6.7, the deprotonation of protonated secondary amines of
TAME5OX, is characterized by a main absorption band centered at 308 nm, (ε =
60, 100 M−1
cm−1
) with a shoulder at about 279 nm, (ε =38, 000 M−1
cm−1
) attributed
to n→* and →* transitions, respectively, and the broad band around 380 nm
get diminished in this pH range („b‟ of Figure 3.11). The deprotonation results in
the bathochromic and hypochromic shifts with the rise of pH from acidic to neutral
medium. The formation of three more isosbestic points at 276 nm, 295 nm and 315
101
nm, confirms the formation of LH6, LH5 and LH4 species, respectively. Similarly,
under basic conditions the electronic spectra of ligand exhibits two bands, one at
higher energies was red shifted (about 17 nm) to 325 nm, (ε = 28, 000 M−1
cm−1
)
and low energy band was also hyperchromically red shifted (about 14 nm) to 290
nm, (ε = 40, 000 M−1
cm−1
) (Figure 3.11), and evidenced the deprotonation of
hydroxyl function from neutral form (LH3) of TAME5OX to form LH2, LH and L
upon variation of pH upto 10.81. Bands at λ > 300 nm are attributed to n→* state
with chargetransfer character while the band around 285 nm concerns the
quinoline ring [24]; the band due to protonation of Npy disappears in this pH range.
No spectral change was observed on further pH variation upto 11.5.
To corroborate the above argument, for the deprotonation behavior of
ligand, the electronic spectra of different species of TAME5OX obtained were
calculated by time-dependent density functional theory (TD-DFT) at the B3LYP/6-
31G* level. The calculated spectra for the neutral (LH3), protonated and
deprotonated species (LH9, LH8, LH7, LH6, LH5, LH4 and LH2, LH) were obtained
by stepwise protonation and deprotonation of N-, NH- and OH- groups, respectively
(Figure 3.15). The experimental spectra of TAME5OX gave major variations for
the * transition; the peak of protonated (LH9, LH8 and LH7) at 262 nm („a‟ of
Figure 3.15) shifted bathochromically with formation of LH6, LH5 and LH4, show
transitions at 274, 276 and 278 nm, respectively. Similar variations were also
observed in the calculated electronic spectra („b‟ of Figure 3.15) and the
corresponding calculated values are reported in Table 3.3. The calculated -* and
n-* transitions for the LH6, LH5 and LH4 obtained at 245 nm and 280 nm, shifted
to 260 nm and 300 nm for the species LH3, LH2 and LH respectively. The simulated
electronic absorption spectra for corresponding predicted species are in good
102
concurrence in the contour, number of absorption bands and relative intensities with
the species except small red-shifts in LH3, LH2 and LH forms, which are due to the
impossibility of considering all the symmetrical orbitals in TD-DFT calculations.
Also, strong absorption peaks were obtained from theoretical calculations at 264 nm
for LH9 and at the same position with comparatively lower intensities for LH8 and
LH7 species, which corroborates well with the experimental results.
Figure 3.15: pH-dependent electronic spectra (absorption) of the nine species as a function
of molar absorptivity and wavelength (a) predicted from Hypspec using experimental data
and (b) calculated through TDDFT/B3LYP method by employing 631G* basis set, for
the protonation and deprotonation of ground state geometrical optimized neutral species
(LH3) of TAME5OX.
3.2.1.2.3 Fluorescence Spectra of ligand (TAME5OX)
In general, the binding moiety of 8-hydroxyquinoline serves as a fluorophore
because of the significant enhancement of fluorescence that is known to occur when
the excited state intramolecular proton transfer (ESIPT) between the hydroxyl
group and quinoline nitrogen is suppressed upon binding of a metal cation [25].
However, some HQ derivatives are poorly luminescent [26] due to an
intramolecular photoinduced proton transfer (PPT) process between the hydroxyl
group (which is relatively strong photoacid) and the nearby quinoline nitrogen
(which is, in turn, a strong photo base). In addition, in protic media, intermolecular
103
Table 3.3: Computed Optical Properties: Absorption-S0 and emission-S1; wavelengths, oscillator strengths (fcalc), electronic transitions and
energies calculated at the TD-DFT/6-31G* level of theory, using DFT/B3LYP and DFT/cam-B3LYP for ground state and excited state geometry,
respectively, for different species of the TAME5OX.
S0 S1
Species nm (abs) fcalc Nature Transition
(Excitation energy)
nm(em) fcalc Nature Transition
(Emission energy)
(H9L)6+
270.54 0.3281 HOMO-2→LUMO+1 110→114 (1.186 eV)
425.64 0.1091 HOMO-1→LUMO+1 111→114 (0.162 eV)
(H8L)5+
272.34 0.1794 HOMO-1→LUMO 111→113 (0.883 eV)
425.89 0.1581 HOMO→LUMO+1 112→114 (0.293 eV)
(H7L)4+
274.89 0.1563 HOMO-2→LUMO 110→113 (0.726 eV)
519.95 0.1674 HOMO-2→LUMO 110→112 (0.385 eV)
(H6L)3+
220.84
310.23
0.1068
0.0921 HOMO-1→LUMO
HOMO→LUMO
111→113 (0.590eV)
112→113 (0.513 eV)
522.83
0.1748
HOMO-1→LUMO 111→112 (0.571eV)
(H5L)2+
223.60
319.37
0.0821
0.0731 HOMO→LUMO+1
HOMO-1→LUMO
112→ 114 (0494 eV)
111→113 (0.489 eV)
526.64
0.1921 HOMO-1→LUMO 111→ 112 (0.836 eV)
(H4L)+ 225.30
337.42
0.0675
0.0598 HOMO-1→LUMO+2
HOMO-1→LUMO
111→115 (0.314 eV)
111→113 (0.379 eV)
528.93 0.2107 HOMO→LUMO+1 112→114 (1.243 eV)
(H3L) 229.91 350.42
0.0401 0.0463
HOMO-2→LUMO+1
HOMO→LUMO
110→114 (0.253 eV)
112→113 (0.288 eV)
530.37 0.2193 HOMO-1→LUMO+1 111→114 (1.391 eV)
(H2L)- 234.98
372.87
0.0374
0.0412
HOMO-1→LUMO+1
HOMO→LUMO
111→114 (0.192 eV)
112→113 (0.219 eV)
474.67 0.2116 HOMO→LUMO 112→113 (1.287 eV)
(HL)2-
244.26
376.05
0.0283
0.0403 HOMO-1→LUMO
HOMO→LUMO
111→113 (0.175 eV)
112→113 (0.209 eV)
470.23 0.1875 HOMO-1→LUMO 111→113 (0.723eV)
(L)3-
245.16
379.08
0.0138
0.0294 HOMO→LUMO+1
HOMO→LUMO
112→114 (0.159 eV)
112→113 (0.184 eV)
429.34 0.1845 HOMO-2→LUMO 110→113 (0.596 eV)
104
PPT process involving solvent molecules can also occur, further decreasing the
fluorescence quantum yield of this kind of fluorophore. The 8-hydroxyquinoline
ligands are indeed characterized by strong intramolecular hydrogen bonding that
may be photoinduced to undergo proton transfer in the excited state [27,28].
In the present tripodal ligand containing three units of 8-hydroxyquinoline,
upon excitation into the ligand transitions, a maximum of emission at about 425 nm
was observed (Figure 3.12). This fluorescence emission originates from an excited-
state intramolecular proton transfer (ESIPT), and enhanced fluorescence with
absolute quantum yield ≈0.091) may be attributed to presence of three units of 8-
hydroxyquinoline in TAME5OX.
3.2.1.2.3.1 pH-Dependent Fluorescence Spectra
From survey of pH-dependent photophysical properties of 8-hydroxyquinoline
compounds, it is seen that, the fluorescence intensity of these compounds changes
in a monotonous way, depending only on the change of pH, which gets increased or
decreased along with the change in the pH value in a consistent way due to the only
one photoinduced electron transfer mechanism between pyridyl -N and -OH group
of quinoline unit, resulting in the OFF-ON or ON-OFF type of fluorescent pH
sensors [29]. Moreover, there are reports on the pH sensors based on the opposite
photoinduced electron transfer processes between a receptor and signaling unit [3].
To explore the pH-dependent optical properties of TAME5OX with unique pyridyl-
N and -OH structural characteristics, the pH vs fluorescence titration experiment
was carried out in aqueous medium. An appropriate volume of 0.1M HCl and 0.1M
NaOH were used for protonation of pyridyl -N group and deprotonation of the -OH
group, respectively. As shown in Figure 3.11, the -* and n-* transitions in the
electronic absorption spectrum of TAME5OX exhibit the exemplary behaviour (red
and hypochromic shifts) along with changing the pH of the system from acidic to
105
basic (as discussed in protonation of TAME5OX). In good comparison, the changes
in the blue-green emission of the ligand maxima located at 425 nm which was
excited at 385 nm, were monitored as a function of pH ranging from 1.9710.81
(Figure 3.12) at room temperature. The emission intensity at 425 nm (with meagre
quantum yield) as well as maximum wavelength remains almost unchanged, no
substantial change in behaviour of the fluorophore was observed upon measuring
emission under acidic conditions within the pH range of 1.97-2.88. Most
interestingly, with an increase in the pH from 3.04 to 6.69, resulted the evolution of
emission peak at 520 nm upto5 fold (could be attributed to the competition
between the intramolecular photon induced proton transfer), while no change in
blue-green emission at 425 nm. Further rise of pH resulted the intense fluorescence
of TAME5OX gradually get blue shifted (about 50 nm) towards 470 nm, along with
quenching of fluorescence signal about 1, 2 and 3-fold at 7.53, 7.94 and 8.27 pH,
respectively. After addition of more base (0.1M NaOH) (pH > 8.5), the
deprotonation effect again evidenced in growth of fluorescence signal upto 3-fold
which promptly got blue shifted to original blue-green emission wavelength (420
nm). Surprisingly, with an increase in the pH from 8.9 to 10.9, the intense
fluorescence of the same compound under basic conditions also gets gradually
quenched in a similar manner as observed under acidic conditions, (cf. Figure 3.12).
It is worthnoting that more significant changes under physiological pH and less
significant changes under acidic and basic medium confirm that protonation of N
atom lowers the LUMO energy and decreases the vertical S1→S0 transition energy,
and the deprotonation of the hydroxyl function also destabilize the HOMO, thus
closing the gap giving rise to smaller transition, while in neutral form the HOMO-
LUMO gap increases which corresponds significant red shifts along with intense
106
fluorescence. Thus, among all the species formed during the pH range of 1.8-10.9,
the neutral species, LH3, was found to be more fluorescent that predominates under
physiological pH, while the more protonated and deprotonated species were found
to be less fluorescent.
These results divulge (reveal) the typical characteristics of a pH-fluorescent
probe through the photoinduced proton transfer (PPT) enhancement process and
photoinduced electron transfer (PET) quenching processes between pyridyl-N and -
OH group for this tripodal chelator [3]. In other words, these two processes between
the 8-HQ receptor units are accountable for the particular fluorescent properties in
TAME5OX ligand along with either decreasing the pH from 6.7 to 1.97 or
increasing the pH from 6.9 to 10.88. The unique pH-dependent fluorescent property
of this ligand indicates the potential application, as pH indicator under both acidic
as well as basic conditions, and, will act as novel OFF-ON-OFF type of fluorescent
pH sensor.
3.2.1.2.3.2 pH-Dependent PET and PPT Mechanisms
For the intention of cognizance the remarkable pH-dependent fluorescence
characteristics of TAME5OX, the following PET quenching and PPT enhancement
mechanisms determined by the orbital energy levels of the excited states of 8-HQ
unit of the ligand is presented at Figure 3.16. In principle, 8-HQ unit can serve as
both electron donor and acceptor upon photoexcitation; the photoinduced
electron can transfer from the excited pyridyl-N moiety to the lowest unoccupied
molecular orbital (LUMO) of the -OH group (donor-excited PET; d-PET) or just
reversely from the highest occupied molecular orbital (HOMO) of -OH to the
excited pyridyl-N moiety (acceptor excited PET; a-PET) [30]. When the N atom is
protonated in the form of NH+, the LUMO energy of electron deficient 8-HQ unit
becomes lower than that of the excited aromatic ring, (see left part of. Figure 3.16),
107
leading to a d-PET quenching process. In contrast, when the OH group of 8-HQ is
deprotonated into O-, the HOMO energy of aromatic ring gets increased and
becomes higher than that of excited N atom, (right part of Figure 3.16) which
induces the electron transfer from -OH moiety to pyridyl-N unit, giving rise to a
reductive a-PET. These two opposite photoinduced electron transfer processes are
in turn responsible for the special fluorescence characteristics of TAME5OX along
with the change in the pH value under acidic and basic conditions, resulting in a
novel OFF-ON-OFF type of pH fluorescent switch to acidic and basic conditions,
respectively.
Figure 3.16: Proposed photoinduced electron transfer (PET) and photoinduced proton
transfer (PPT) mechanisms between N-pyridyl and -OH groups of 8-HQ moieties of
TAME5OX in protonated quinolinium, neutral, and deprotonated quinolinate states,
respectively.
108
The above-mentioned two processes (PPT and PET) observed during pH
dependent fluorescence of TAME5OX are clearly validated by density functional
(DFT) calculation with the camB3LYP/6-31G* method based on the excited
state optimized molecular structure of the ligand. The calculated frontier
molecular orbital (MO) energies and surfaces of protonated, neutral and
deprotonated species of TAME5OX are shown in Figure 3.17. The highest occupied
orbital (HOMO) is mainly located on phenolate side and the lowest unoccupied
* orbital (LUMO) is mainly located on the pyridyl moiety. As can be found, both
the HOMO and LUMO of TAME5OX in the neutral state are located almost
completely on the aromatic side of 8-HQ moiety, while the HOMO-1 and LUMO+1
are located mainly on the -N pyridyl and hydroxyl group of 8-HQ unit, respectively. As a
result, the electronic transition between HOMO and LUMO, leading to fluorescence
from the system of this compound. When the N atom of 8HQ in TAME5OX is
protonated in the form of NH+, the LUMO energy of pyridyl moiety, 1.168eV,
becomes higher than that for the aromatic part in the molecule, 0.385eV, which in
Figure 3.17: Calculated CAM−B3LYP/6−31G* energy levels and surfaces of frontier
molecular orbitals (MOs) of TAME5OX in protonated, neutral, and deprotonated states,
respectively.
109
turn results in the electron transfer towards system in moiety under neutral
conditions as shown in Figure 3.17. When the OH group of the 8HQ moiety of
TAME5OX is deprotonated into O- in the basic system, change occurs in the
molecular structure of this compound. Along with structural change, the HOMO
energy of the 8-HQ moiety in the molecule gets increased to be close to those of the
LUMO, which in turn decreases largely from that of system of moiety, which
decreases the energy gap and quenches the fluorescence intensity. For a better
interpretation, the trends of the results are plotted in Figure 3.18. The computed
vertical S1→S0 emission energies (that is, fluorescence) of highly acidic (LH9 and
LH8) and basic (L) species were obtained with maxima at 425 nm. Other protonated
and neutral (LH3) forms are red-shifted by ca. 95 nm with significant enhancement
of fluorescence intensity along stepwise deprotonation from LH7 to LH3, while
deprotonated species LH2 and LH show blue shifts with respect to the
corresponding neutral emission value with decreasing intensity from LH2→L, in
good agreement with experimental results
Figure 3.18: TD-camB3LYP simulated emission spectra of TAME5OX as a function of
protonation and deprotonation of excited state geometrical optimized neutral species (LH3)
of by employing 6-31G* basis set.
110
3.2.2 Complexation, Evaluation of Thermodynamics and Photophysical
Properties of TAME5OX with Trivalent Fe, Al and Cr Ions
The coordination ability of TAME5OX with trivalent Fe, Cr and Al is being
reported through synthesis, solution as well as theoretical study in the current
chapter.
3.2.2.1 Synthesis and Characterization of [M(TAME5OX)]
For synthesis of the metal complexes non-aqueous solution was employed. The
reaction of TAME5OX with the metal ions in ethanol medium led to compounds
with general formula [M(TAME5OX)] with moderate yields (10-29%). The
complexes were characterized by elemental analysis, IR and ESI MS mass
spectrometry.
3.2.2.1.1 Infrared spectra
In the IR spectra of the metal complexes of TAME5OX, the absence of a broad
band at 3399 cm-1
confirms the coordination of ligand with metal ions through OH
groups. There are sharp absorption bands at 991, 1082 and 1103 cm-1
for
Fe(TAME5OX), Al(TAME5OX) and Cr(TAME5OX), respectively, which
associate with C-O vibrations at the C-O-M site [31]. The (C=N) quinoline bands
at 1633 cm-1
and 1081 cm-1
of the oxine sub-units of ligand are generally shifted to
lower wave numbers 1612cm-1
and 991cm-1
, respectively, than those in the free
ligand on coordination to the metal ion [32]. The IR spectra of the complexes
showed bands at ca. 603, 479 and 461, cm-1
for Fe3+
, Al3+
and Cr3+
, respectively,
slightly shifted to lower wave numbers than those in free ligand, indicating the
coordination of the oxine groups to the metal ions [33]. The experimental IR spectra
of ferric complex of TAME5OX is shown in Figure 3.19.
In order to confirm the assignments of IR frequencies, theoretical calculations
were carried out by DFT method at B3LYP/6-31G* level. For comparison, the
111
theoretical and experimental values are also given in Table 3.4. The calculated
vibrational frequencies are obtained at higher values than experimental ones, but
their trend are corroborate with the experimental data indicating that theoretical
results can be implemented for the interpretation of experimental results. The
differences in calculated and experimental spectra are mainly due to, somewhat, to
the fact that the experimental spectra were measured in the solid phase and the
theoretical spectra were obtained in the gas phase. However, a linear correlation
obtained between the theoretical and experimental data (Figure 3.20) gave an
acceptable correlation.
Figure 3.19: Infra-red (IR) spectrum of Fe(TAME5OX).
Table 3.4: Calculated (B3LYP/6-31G*) and experimental IR frequencies (KBr, cm-
1) for Fe(TAME5OX), Al(TAME5OX) and Cr(TAME5OX) complexes.
Fe(TAME5OX) Al(TAME5OX) Cr(TAME5OX)
B3LYP/631G* Experimental B3LYP/631G* Experimental B3LYP/631G* Experimental
3209 2932 2998 2930 3076 2927
1615 1612 1673 1668 1669 1652
1492 1432 1556 1543 1542 1510
1349 1345 1483 1456 1454 1437
1155 1136 1346 1392 1397 1386
1094 1085 1235 1223 1217 1209
995 991 1089 1082 1105 1103
615 603 971 965 613 603
475 461 644 638 675 461 - - 491 479 415 403
- - 421 417 - -
112
Figure 3.20: Correlation between the experimental and theoretical IR frequencies of
M(TAME5OX); [M= Fe, Al and Cr].
3.2.2.1.2 Mass Spectra
High-resolution electrospray ionization time of flight mass spectrometry TOF
MS(ES+) was used to identify the product by evaluation of the molecular ion peak
of metal complexes of TAME5OX. The mass spectrum of Fe(TAME5OX) (Figure
3.21) showed two peaks at m/Z 653.4 (35%) and 654.4 (22%) attributed to the
protonated complex fragments [Fe(M+2H)]+
and [Fe(M+3H)]+, respectively.
Similarly the mass spectrum of Al(TAME5OX) and Cr(TAME5OX) (not shown in
figure) display molecular ion peaks that are consistent with formulation of the
M(TAME5OX).
Figure 3.21: MS(ES+) mass spectrum of Fe(TAME5OX).
Due to the insolubility of the coordination compounds we could not isolate any
crystal, and could not establish the exact structure. However, on the basis of the
physico-chemical evidences the structure of the complexes [M(TAME5OX)] is
proposed to be distorted octahedral with N3O3 coordination bonded through three
sets of pyridyl and phenolate ions of oxine.
113
3.2.2.2 Complexation. Interaction of TAME5OX with Trivalent Metal (Fe3+
, Al3+
and Cr3+
) Ions
In contrast to the solid state study, in which the positions of the atoms of a molecule
are freezed during crystal formation, whereas the solution structure of molecule
differes due to the dynamic equilibrium where the free energy, entropy, enthalpy
play important role and the atoms rearrange themselves to attain lowest energy
confirmations.
3.2.2.2.1 Metal Complex Equilibria
In aqueous solution, the metal ions usually assume hexacoordinated structure with
type [M(H2O)6]. In presence of ligands there are many equilibrium reactions and
there is dynamic equilibrium between various species, which are dependent on pH.
Many species exist in solution. The ligand TAME5OX is a potential trinegative
nonadentate (N6O3) chelator and theoretically expected to satisfy nine-coordination
of a trivalent metal ion, whereas the most probable coordination number for Fe3+
,
Al3+
and Cr3+
is six. Molecular modeling suggests that the topology of the ligand is
such that six coordination to the same metal ion increase strain energy due to the
positioning of N,O-coordinate sites of each oxinate group opposite to the cetral unit
of ligand. But, the anchoring tetrahedral sp3 TAME on the pivot C atom facilitate
six-coordinate distorted octahedral geometry of M3+
. Keeping in view the dynamic
equilibrium in solution, many strain free structures were proposed by adding water
molecules to the coordination sphere by replacing some coordination sites of the
ligand turn wise. A coordination scan was performed using molecular mechanics
with varying coordinate water molecules (one to three, assigning maximum
coordination number 3) with SYBYL force field (the strain energy of the water in
this force field assumes zero [34]) suggest 6-coordination with three water
molecules with stoichiometry [M(TAME5OX)(H2O)3]. As mentioned earlier, with
variation of pH a variety of species form in solution, the nature and stability of the
114
species forming between the trivalent metal (Fe3+
, Al3+
and Cr3+
) ions and
TAME5OX have been probed at variable pH and in equimolar solutions of ligand
and metal ions by the combined use of potentiometric, spectrophotometric and
luminescence titrations in the pH range 1.0-12.7. The potentiometric titration
curves of 1:1 solutions for all the metal-ligand systems (Figure 3.22) show a pH
variation from free ligand (TAME5OX) indicating the release of protons upon M3+
coordination. The curves (ii-iv) deviate from the curve for the protonation of ligand
(i) alone, indicates the formation of species in which different number of protons
has been displaced from TAME5OX by the trivalent metal ions possibly from the
protonated N-pyridyl and hydroxyl groups. Keeping in view these preliminary
observations, many sets of possible models were tested in the minimization
programme and the best-fit model was obtained when formation of species MLH3,
MLH2, MLH and ML were considered for Fe(III) and Cr(III), where as MLH3 and
ML were considered for Al(III). The potentiometric data were refined by the
program Hyperquad to obtain the overall formation constants (log s) and the
results are summarized in Table 3.5. The equilibrium reactions for the overall
formation of the complex species and their cumulative formation constants, 11n, are
defined by the following equations (charges are omitted for clarity).
Figure 3.22: Potentiometric titration curves: (i) 1×10
-5M TAME5OX, (ii)
[TAME5OX]/[Fe3+
] = 1/1, 1×10-5
M, (iii) [TAME5OX]/[Al3+
] = 1/1, 5×10-5
M, (iv)
[TAME5OX]/[Cr3+
] = 1/1, 1×10-5
M. Solvent H2O, I = 0.1M (KCl), T= 25(2)°C, and ‘a’ is
the moles of base added per mole of M[TAME5OX] present. Symbols and solid lines
represent the experimental and calculated data, respectively.
115
𝑀 + 𝐿 ⇌ 𝑀𝐿, 𝛽110 = 𝑀𝐿
𝑀 𝐿 (3.4)
𝑀 + 𝐿 + 𝐻 ⇌ 𝑀𝐿𝐻, 𝛽111 = 𝑀𝐿𝐻
𝑀 𝐿 𝐻 (3.5)
𝑀 + 𝐿 + 2𝐻 ⇌ 𝑀𝐿𝐻2, 𝛽112 = 𝑀𝐿𝐻2
𝑀 𝐿 𝐻 2 (3.6)
𝑀 + 𝐿 + 3𝐻 ⇌ 𝑀𝐿𝐻3, 𝛽113 = 𝑀𝐿𝐻3
𝑀 𝐿 𝐻 3 (3.7)
In order to predict the reaction path, it is advantageous to represent the equilibrium
reactions in terms of stepwise formation of complexes. If MLH3 is regarded to be
first species formed by the reaction of M and LH93+
, and MLH2, MLH and ML are
formed due to dissociation of protons in steps from MLH3. Accordingly, the
stepwise dissociation and the derived successive formation constants (log K) can be
represented as Equations 3.8 − 3.11.
(Fe Al Cr)
𝑀 + 𝐿𝐻9 ⇌ 𝑀𝐿𝐻3 + 6𝐻, 𝑙𝑜𝑔 𝐾𝐿𝐻9
𝑀𝐿𝐻3 = 𝑀𝐿𝐻3 𝐻
6
𝑀 𝐿𝐻9 = 2.85 3.98 1.79 (3.8)
𝑀𝐿𝐻3 ⇌ 𝑀𝐿𝐻2 + 𝐻, 𝑙𝑜𝑔 𝐾𝑀𝐿𝐻3
𝑀𝐿𝐻2 = 𝑀𝐿𝐻2 𝐻
𝑀𝐿𝐻3 = 2.60 6.22 2.46 (3.9)
𝑀𝐿𝐻2 ⇌ 𝑀𝐿𝐻 + 𝐻, 𝑙𝑜𝑔 𝐾𝑀𝐿𝐻2
𝑀𝐿𝐻 = 𝑀𝐿𝐻 𝐻
𝑀𝐿𝐻2 = 3.08 − 4.43 (3.10)
𝑀𝐿𝐻 ⇌ 𝑀𝐿 + 𝐻, 𝑙𝑜𝑔 𝐾𝑀𝐿𝐻𝑀𝐿 =
𝑀𝐿 𝐻
𝑀𝐿𝐻 = 1.89 − 5.24 (3.11)
Fe[TAME5OX]:
116
Al[TAME5OX]:
Cr[TAME5OX]:
Figure 3.23: UVvis absorption spectra of 1:1 solution of M3+
and TAME5OX (1/1, 1×10-
5M) as a function of p[H]. For [TAME5OX]/[Fe
3+] (a) p[H] = 1.97-7.54 (b) p[H] = 8.34-
11.04; for [TAME5OX]/[Al3+
] (a) p[H] = 1.947.42 (b) p[H] = 9.2611.70; and for
[TAME5OX]/[Cr3+
] (a) p[H] = 1.897.98 (b) p[H] = 8.0410.85. Solvent H2O, I = 0.1M
(KCl), T= 25(2)°C.
The spectrophotometric titrations were also carried for study of formation of metal-
ligand complexes and the results of Fe(TAME5OX), Al(TAME5OX) and
Cr(TAME5OX) complexes are shown in Figure 3.23. The data obtained in two
different pH ranges were fitted separately. Between pH 1.9 and 7.5, the best fit of
absorbance data correspond to the formation of two protonated species for Fe3+
and
Cr3+
whereas only one protonated species, was obtained for complexation of Al(III).
Combining these data with that obtained from the potentiometric titrations allowed
to fit three proton-dependent equilibria, described by equations 3.8-, except for
117
Al3+
, in which only one species was evidanced. Above pH 8.0, both potentiometric
and spectrophotometric data revealed the presence of the single neutral (ML)
species in all the three ligandmetal systems. The calculated formation constants
defined by equations 3.4-3.7 are listed in Table 3.5.
To get further support the number and nature of species obtained by both
potentiometric and spectrophotometric methods and the corresponding formation
constants, luminescence titrations were carried out for 1:1 solutions of Fe3+
, Al3+
and Cr3+
ions with the ligand from 1.9-11.2 pH range. The experimental graphs for
the luminescence titrations for iron, aluminum and chromium complexes are
presented in Figure 3.24. For Fe[TAME5OX] and Cr[TAME5OX] systems, the best
refinement of the luminescence-pH data by the program Hypspec correspond to the
presence of four species each: three protonated, [M(H3L)]3+
, [M(H2L)]2+
and
[M(HL)]+, and one neutral species [(ML)] in concurrence with both potentiometric
and spectrophotometric results; but, for Al[TAME5OX] system, the best fit data
returned only two complexes in the pH range of 1.9-11.7, one protonated
[Al(H3L)]3+
[log = 29.62(5)] and one neutral (AlL) [log =
23.04(6)], in harmony
with both potentiometric and spectrophotometric data. It is worth to mention that all
species could not be detected by a single method, whereas they could be detected by
combination of potentiometry, UV-visible spectrophotometry and fluorometry. The
formation constants obtained by all the three techniques defined by equations
3.4−3.7, together with the average log values are reported in Table 3.5.
118
Fe[TAME5OX] ` Al[TAME5OX]
Cr[TAME5OX]
Figure 3.24: Dependence of the emission spectra of 1:1 aqueous solution of
M3+
:TAME5OX = (1.0 × 10-5
M) on the change of the pH value from acidic to basic range,
T = 25.0(3)°C (excitation and emission slit widths of 5.0 nm). (a) for Fe
3+/TAME5OX (1.89-
11.23 pH), (b) for Al3+
/TAME5OX (1.97-11.04 pH), and (c) for Cr3+
/TAME5OX (1.94-
11.70 pH), with an excitation wavelength of 385, 370 and 345 nm, respectively.
(a) (b)
(c)
119
Table 3.5: Equilibrium constantsa (log 11n), absorption, emission characteristics of Fe3+
, Al3+
and Cr3+
complexes of TAME5OX and
corresponding pM values.
Ligand Complex Log 11nb Log 11n
c Log
11nd
Average
Log 11n
ab (nm) ɛ(M-1cm-1) em(nm) Φ ePL pM3+ e
TAME5OX [Fe(H3L)]3+ 37.82(5) 37.74(3) 37.69(1) 37.75 267, 290, 383 108 000, 20 000, 8 000 425 0.01 31.16
[Fe(H2L)]2+ 35.07(6) 35.25(2) 35.13(4) 35.15 258, 272, 300 45 000, 18 000, 6 230 525, 470 0.06 [Fe(HL)]+ 32.01(1) 32.19(5) 32.03(3) 32.07 350 50 000 525, 500 0.09
[(FeL)] 30.20(2) 30.39(4) 29.95(2) 30.18 350 55 000 430, 500 0.15
[Al(H3L)]3+ 29.30(7) 29.57(2) 29.62(5) 29.49 255, 282 98 000, 40 000 425, 485 0.09 18.07
[(AlL)] 23.23(3) 23.54(3) 23.04(6) 23.27 245, 345, 370 40 100, 41 000, 50 000 435, 500 0.38
[Cr(H3L)]3+ 33 .33(4) 33 .25(1) 33 .17(2) 33 .25 275 96 500 420 0.01 18.12
[Cr(H2L)]2+ 31.82(2) 31.70(5) 31.55(5) 31.69 270, 280, 300 45 300, 30 000, 25 640 425, 468 0.03
[Cr(HL)]+ 27.41(3) 27.26(6) 27.11(7) 27.26 305, 325 26 000, 50 000 475, 500 0.04
[(CrL)] 22.09(1) 21.79(2) 22.18(2) 22.02 310, 330 25 200, 55 000 425 0.19
Oxinobactin[g] [Fe(LH3)]3+
[(FeL)]
36.75(5) 33.61(13)
450
450
38 000
48 000
- - 32.8
Sulfoxinobactin[g] [Fe(LH3)]
[(FeL)]
30.66(5)
26.73(9) 430
430
40 000
19 000
- - 27.1
O-TRENSOX[h] [Fe(LH5)]2+
[Fe(LH)]2-
[Fe(L)]3-
42.2(1)
36.5(1) 30.9(1)
435
443 443
82 000
52 000
54 000
-
- 29.5
Cox2000[i] [Fe(LH4)]4+
Fe(LH)]+
[(FeL)]
38.98(1)
34.24(5)
32.12(9)
440
440
450
4400
4600
4600
-
-
-
-
-
-
29.1 aNumbers in parentheses represent the standard deviation in the last significant digit. In water (I = 0.1M KCl, T = 25.0°C). bPotentiometric method, cUV-visible spectrophotometric method, dLuminescence spectrophotometric method, e0.1 M solution of quinine sulphate in 0.5 M H2SO4 as standard (Φ = 0.546), f pM3+ = −log [M3+] calculated for [M]tot = 10−5M, [L]tot = 10−5M, and
p[H] = 7.4, [g] Reference [22], [h] Reference [24], [i] Refrence [11].
120
To verify the assignments made for the experimental formation constants,
aqueous-phase free energy for the entire representative species for the Fe3+
, Al3+
and Cr3+
complexes of TAME5OX were also calculated through density functional
theory by B3LYP/631G* quantum mechanical approach. The calculated
experimental log K values were compared with the theoretically calculated ∆G°,
which resulted an acceptable correlation presented in Figure 3.25.
Figure 3.25: Correlation between the experimental log K and calculated ∆Gº (a)
Fe(TAME5OX) and (b) Cr(TAME5OX).
Figure 3.26: Species distribution curves of (a) Fe(TAME5OX), (b) Al(TAME5OX) and (c)
Cr(TAME5OX) containing species, computed from the formation constants given in Table.
Calculated for [TAME5OX]tot = [M3+
]tot = 1×10-5
M, Solvent : H2O, I = 0.1M (KCl), T =
25.0(2)°C.
The pH-dependent species distribution curves of Fe3+
, Al3+
and Cr3+
complexes
of TAME5OX, are shown in Figure 3.26. It has been observed that complexation
starts below pH 1.0 with the formation of Fe(LH3) as the major species (98.5%)
present in the range of 0.8-5.6 pH. After pH 5.9, Fe(LH3) diminishes with the
formation of new species Fe(LH2) which predominates at pH 7.2 (95%). Due to
(a) (b) (c)
121
further deprotonation, Fe(LH) was found as major species at pH 8.0 (67%) along
with the neutral complex FeL which remained predominantly present during the pH
range of 8-11. Similarly, in case of Al and Cr, Al(LH3) and Cr(LH3) complexes (cf.
Figure 3.26, b-c) were found major specis during pH range of 1.0-6.0 and 0.6-5.3,
respectively. An increase in pH led to diminish these species and after pH 5.3 and
2.9 with the formation of new species AlL and Cr(LH2), with maximum
concentration at pH 7.4-10.5 and 5.3, respectively. Cr(LH2) upon deprotonation at
higher 8.0 and 10.3 led the formation of Cr(LH) and CrL. Cr(LH) was the major
species for a wide range of pH 4.0-12.0. Further scrutiny of the distribution curves
showed a steep decrease of the free Fe3+
, Al3+
and Cr3+
concentration from pH 0.0-
2.2 before it completely disappears at pH > 5.8. These results suggest a change in
the coordination around the metal ion with variation of pH.
3.2.2.2.2 Solution coordination behaviour and selectivity
The fully protonated form of ligand, LH96+
, releases six protons to give neutral
uncharged coordinated ligand LH3 in MLH3 and obviously these protons must be
released from protonated N-pyridyl and from protonated secondary amine (-NH)
groups. The first species MLH3 presumably formed due to the coordination of three
N-pyridyl atoms to give cappedtype geometry. It is known that Fe, Al and Cr are
hard metal ions, and show strong preference towards hard donor atoms such as
negatively charged oxygen atoms and in aqueous solution these metal ions exists as
M(H2O)63+
[35]. In order to satisfy the coordination sphere in MLH3, the three
remaining coordination sites of metal may be considered to be occupied by three
water molecules. As the pH increases, as expected, coordinated water molecules are
replaced by the deprotonated quinolinate hydroxyl groups of L in three steps to
form MLH2, MLH and ML. Extrusion of protons from the complex MLH3 may take
place from hydroxyl groups of quinoline moieties, for which the protonation
122
constant values have been evaluated. There are two possibilties (as discussed
above) for the coordination of oxygen atoms to metal centre, either coordinated
water molecules or N-M bonds will be replaced by O-bonds. These two
probabilities lead to two distinct geometries, either nitrogen-oxygen encapsulated or
xygen-oxygen encapsulated. In the first case, upon coordination if the metal ion
facilitates to enter into the nitrogen-oxygen cavity with release of coordinated water
molecules, it may form the architecture which will not favor octahedral or distorted
octahedral geometry, as all the six coordination groups/sites are located on the
para position of anchoring unit which will coordinate metal ion on single side
only, lead to unstable complex geometry, verified by molecular modeling through
molecular mechanics (MM+) and semi-emprical (PM6) calculations. However, in
the second case of oxygen–oxygen encapsulation, the ligand coordinates through its
quinoline oxygens as anionic donors on one side of metal ion, by replacement of
nitrogen-metal center to oxygen-metal with three water molecules remains
coordinated on the other side forming distorted octahedral geometries of complexes,
and was found most stable complex species with least strain energy, calculated
theoretically by above mentioned methods. On the basis of these observations, we
can rationalize the preference of three hydrated tris(quinoline) bonding for M(III)
than the tris(aminoquinoline) or tris(hydroxyquinoline). Thus, nitrogen-oxygen
encapsulated structures can be suggested for species MLH3, MLH2 and MLH,
respectively, while oxygen-oxygen encapsulated structure for the neutral form (ML)
of complex. Consequently, the N-pyridal atoms remain coordinated in the
protonated complexes which get replaced with M-O coordination with the increase
in pH, as can be evidanced from the red shifts at higher pH observed from both
UV-visible and fluorescence spectra. So the coordination sphere around the metal
centre changes from M-N to M-O with change in pH range. As a consequence, a
quinolinate coordination is favoured for M(TAME5OX) in acidic medium since
123
hydrogen bonding between pyridal N and hydroxy oxygen “preorganizes” the
ligand for this mode of bonding with metal ion. Furthermore, in basic medium, the
coordination changes from a quinolinate to a salicylate mode of bonding upon
deprotonation of the complex.
In solution, the metal ion competes with the protons during complexation. The
proton competition mainly depends on the pH and pKa of the ligand and thus, the
pM (pM = −log[Mn+
]) [36] is a better consideration for the relative complexation
efficiency of the ligand under given conditions of pH, Mn+
and L concentrations.
The pM values for (M3+
= Fe3+
, Al3+
and Cr3+
) were calculated at pH 7.4, [L] total=
1 × 10−5
M and [M3+
] total= 1 × 10−5
M, and summarized in Table 3.5 along with
other tripodal or reference ligands, as computed from known formation constants. It
is evident from Table 3.5, that at physiological pH, the ligand binds Fe(III) more
selectively than Cr(III) followed by Al(III). According to these values, TAME5OX
is strong synthetic iron chelating agent over the pH range of 3-14; its efficiency is
higher than O-Trensox and only 1.34 pFe log unit less than oxinobactin. The
enhanced affinity of TAME5OX for ferric ion and the very similar coordination
geometry around other metal ions for its metal complexes support to the ability of
the 8-hydroxyquinoline moieties to be predisposed with the C pivot central unit.
The pH dependence of the selectivity was also examined owing to the fact that in
particular biological compartments or circumstances, pH values far from 10 can be
observed. It is important to see whether the selectivity of the ligand still remains in
a biological relevant pH range. The corresponding plots of pFe3+
, pAl3+
and pCr3+
versus pH for the TAME5OX over the pH range of 2.0-10 are presented in Figure
3.27, which inferred that the ligand for all the three metal ions is maintained over
this wide pH range. For TAME5OX, the pM values translates to 31.16, 18.0 and
18.12 for Fe, Al and Cr, respectively. Thus the C pivot building block is impressive
and the complexes based on TAME5OX appear to be sufficiently stable in water for
124
potential in vivo applications may be done to chelate effect of the 8-
hydroxyquinoline moieties with predisposed architecture.
Figure 3.27: Plot of pM versus p[H] for TAME5OX, pM = -log [M
3+], calculated for [M
3+]
= 10−5
M and [L] = 10−5
M. (a) p[Fe], (b) p[Al], and (c) p[Cr].
3.2.2.3 Photophysical properties of metal complexes
3.2.2.3.1 Absorption spectra: experimental and theoretical correlation from DFT
The ligand TAME5OX shows two peaks 255 and 310 nm (ε = 110, 000 M−1
cm−1
and 60, 100 M−1
cm−1
, respectively) in the UV-visible region assigned to →* and
n→* transitions, and on metal coordination, there is a considerable shift on both
transitions. Since, a variety of species are formed with change on pH, the large
spectral changes (Figure 3.23, for Fe(TAME5OX)) observed in the
spectrophotometric titrations of the metal-TAME5OX complexes in acidic medium
and at physiological pH is indicative of a change in the coordination of metals with
the ligand. At pH < 2.9, the electronic spectra of TAME5OX with Fe, Al and Cr
ions is symbolized by a strong absorption band at higher energies 267 nm, (ε = 108,
000 M−1
cm−1
), 255 nm, (ε = 98, 000 M−1
cm−1
) and 275 nm, (ε = 96, 500 M−1
cm−1
),
respectively, assigned for →* transitions. In case of ferric complex, the low
energy broad band observed near visible region at 383 nm (ε = 7, 640 M−1
cm−1
) is
attributed for n→* transition (cf. Figure 3.23 (a)), while no such peak at higher
125
wavelength was observed for Al and Cr complexes in this pH range. Upon rise of
pH above 2.9, the behaviour of TAME5OX showed similar variations with all the
three metal ions. The high energy bands at 267 nm, (ε = 108, 000 M−1
cm−1
), 255
nm, (ε = 98, 000 M−1
cm−1
) and 275 nm, (ε = 96, 500 M−1
cm−1
), declines and get
shifted towards lower wavelengths, at 258 nm (ε =50, 000 M−1
cm−1
), 245 nm, (ε =
40, 100 M−1
cm−1
) and 270 nm, (ε = 45, 300 M−1
cm−1
), with the concomitant ascent
in absorbance at about 282-310 nm, 270-300 nm and 290-315 nm for Fe, Al and Cr
complexes, respectively. The formations of these non-structured bands at higher
wavelengths were observed due to the deprotonation and complexation of Npyr
groups, but less significantly for Fe complex at 330-340 nm. Analogously, along
with the increment of pH upto 7.9 this low energy band at 335 nm gets red shifted
towards higher wavelength with the formation of broad band around 380-410 nm (ε
=8, 000 M−1
cm−1
). No useful information about the d-d transition of metal ions
could be obtained from the electronic spectra of the complexes in a solution. Since
the absorption coefficients of these bands are large for a d-d transition, they are
likely due to ligand-to-metal charge transfer (LMCT). However, the appearance of
band at higher wavelengths (390 nm) with lower intensity may be attributed to the
coordination of pyridine nitrogen atoms to the metal ion upon chelation that
corresponds to the formation of Fe(LH3) complex species. It can thus be assumed
that the band at 338 nm and 380-410 nm is of the Npyr→Fe(III) type. Similar results
have also been reported by ligands Tsox and TsoxMe [23]. The isosbestic points
assert the formation of different protonated complexes of Fe, Al and Cr with the
ligand in the pH range of 1.97-7.94, 1.94-7.97 and 1.89-7.98. The various species
found during complex formation are presented in the species distribution curves
(Figure 3.26).
126
In contrarary to acidic→neutral pH range, no significant spectral changes were
observed over the pH range of 7.9-11.5. Above pH 8, the deprotonation and
complexation of hydroxyl groups with Fe, Al and Cr ions, resulted in the
bathochromic and hyperchromic shifts (cf. Figure 3.23 for Fe(TAME5OX)) of n-
π* transitions as expected at 350 nm, (ε =43, 000 M−1
cm−1
), 345 nm, (ε =41, 000
M−1
cm−1
) and at 325 nm, (ε =50, 000 M−1
cm−1
), respectively. Similarly, upon
further increase in pH upto 11, these bands of M[TAME5OX] were again
hyperchromically red shifted to 370 nm, (ε =50, 000 M−1
cm−1
) and 330 nm, (ε =55,
000 M−1
cm−1
) in case of Al and Cr ions, respectively. While for Fe[TAME5OX],
the peak at 350 nm, (ε =50, 000 M−1
cm−1
) get only hyperchromically shifted with ε
=55, 000 M−1
cm−1
and low energy band in the visible region at 400 nm, (ε = 8,
000M−1
cm−1
was completely disappeared in this pH range („b‟ of Figure 3.23)
indicating the formation of neutral (FeL) complex. As observed in analogues
quinolinate [22,24] and phenolate complexes the bands at higher wave lengths are
assigned to charge transfer by O→M+3
. It should be noticed that the bands due to
Npyr→M disappears in the basic pH and the change in absorption spectrum can be
interpreted by a change in the coordination properties of the metal ion. It can thus
be assumed that TAME5OX is the efficient M3+
chelator in wide range of pH 1.7-
11.5, and form stable complexes not only with the fully deprotonated form of ligand
(L3-
) but also with the neutral (LH3), mono (LH2) and di (LH) anionic forms of
ligand.
In order to support the experimental results for formation of metal complexes,
further investigation for the electronic transitions and electronic structures of
complexes were performed from a theoretical stand point at DFT level. The main
goal was to see what influence the 8-hydroxyquinoline binding unit has on the
distribution in complexes with the present ligand and to get insights into the trends
127
of the basic electronic structure with the variation of metal ions, and also within
different protonated and neutral complexes. Geometry optimization, frequency and
NBO analysis of complexes with the metals Fe and Al were performed by Jaguar
7.6 [37] with the B3LYP exchange correlation functional using the 6-31G* basis
set. These conditions were used for similar metal complexes (e.g., Alq3) before
and have proved to be on a sufficient level of theory to describe the properties
of the complexes adequately [38]; for the Cr complex, the LACVP* basis set was
used instead [39]. Time-dependent density functional theory (TD-DFT) calculations
were carried out at the B3LYP/6-31G* approximation level through Gaussian 09
programme [40] to evaluate the absorption properties (excitation wavelength,
oscillator strengths, etc.) of the metal complexes. The computed optical data of
complexes obtained from these calculations are listed in Table 3.6. The simulated
electronic absorption spectra of predicted species of Fe(TAME5OX),
Al(TAME5OX) and Cr(TAME5OX) obtained via TD-DFT method are in good
agreement in the profile, number of absorption bands and relative intensities with
the species obtained from experiment [Figure 3.28]. The ground state optimized
structures of TAME5OX and its Fe, Al and Cr complexes are shown in Figure 3.29
(a-d), respectively. The optimized structures showed no unusual features except
slightly distortion. From an electronic point of view, for all form of, viz., neutral or
protonated, the second electronic transition basically corresponded to a
HOMO→LUMO excitation and orbitals had same character (namely and *,
respectively). On analysing the molecular orbitals, the HOMO is somehow
localized on the aromatic ring carrying the OH function while the antibonding
LUMO is delocalised over the nitrogen atom of the quinoline system; that, this
transition is expected to be relatively sensitive, in terms of intensity and position, to
overall protonation degree of the molecules. In particular, the significant
128
hypochromic shifts are computed going from protonated (acidic) to neutral species
for all above mentioned metal ions and small blue shifts were also predicted in
going from [M(H3L)]3+
to [M(H2L)]2+
. While the remarkable red shifts (of about 45
nm) were observed in going from dito mono protonated complexes and also about
30 nm red shift from mono protonated (MLH) to neutral(ML) forms.
Fe[TAME5OX]:
Al[TAME5OX]:
Cr[TAME5OX]:
Figure 3.28: pH-dependent electronic spectra (absorption) of the different metal complexes
as a function of molar absorptivity and wavelength (a) predicted from Hypspec using
experimental data and (b) calculated through TD-DFT/B3LYP method by employing 6-
31G* basis set, for Fe(TAME5OX), Al(TAME5OX) and Cr(TAME5OX).
129
Figure 3.29: The DFT/B3LYP optimized ground state geometries of (a) TAME5OX, (b)
Fe[TAME5OX], (c) Al[TAME5OX] and (d) Cr[TAME5OX] from 6−31G* level of
calculation.
A quantitative explanation of these variations in absorption transitions and
energies can be given by the inspection of the molecular orbitals computed for Fe,
Al and Cr complexes of TAME5OX. Indeed protonation of the N atom increases
the acceptor character of the quinoline ring. Thus, it is expected to raise the LUMO
energy and increases the vertical S0→S1transition energy, in agreement with the
above mentioned computed red shift on deprotonation of NH+ upon complexation
(cf. Figure 3.28). On the other hand, the deprotonation of the hydroxyl function will
destabilize the HOMO, thus closing the gap and giving rise to smaller transition.
Notably both the above mentioned effects (i.e; stabilization of the LUMO and
destabilization of the HOMO) are concomitant when going from M(H3L)]3+
to
(ML). Not surprisingly these species are computed to have the HOMO→LUMO
gap (2.984eV) and a significant red shift of the second transition (98 nm) is
consistently computed particularly going from Al(H3L)]3+
to (AlL) (cf. Figure 3.28).
The first excited state basically corresponds to a HOMO-3→LUMO+1 transition,
and it is the only intense transition in this spectral region for all compounds but
Fe(LH2) and FeL that show a second transition of relatively the same intensity at 278
nm and 435 nm, respectively (cf. “a” of Figure 3.28). This later corresponds, in both
cases, HOMO to LUMO+1 excitation, the LUMO+1 being a * orbital delocalized
over all quinoline rings with negligible contribution from the acidic functions.
(a) (b) (c) (d)
130
The electronic spectra and electronic structure of metal complexes are
exemplified in figure 3.30 and 3.31, respectively, for three neutral metal complexes
of TAME5OX (Metal = Fe, Al and Cr), for the sake of comparision. A few general
trends, can be seen comparing the three complexes (cf. Figure 3.31). These
complexes show considerable delocalization of the ligandbased orbitals over at
least two binding units, especially for the unoccupied orbitals (cf. Figure 3.31,
orbitals 130 = LUMO+2, 129 = LUMO+1, for Fecomplex, and orbitals 129 =
LUMO+2, 128 = LUMO+1, for Crcomplex). The exceptions are the three virtual
Figure 3.30: Simulated TDDFT spectra of neutral (ML, M = Fe, Al and Cr) complexes of
TAME5OX as a function of 6-31G* basis set employed in (a) UV/vis absorption by B3LYP
for excitation energy determination. (b) Emission by camB3LYP for first excited state
geometries of complexes.
131
Figure 3.31: Calculated (B3LYP/631G*) frontier orbitals for the metal complexes of
TAME5OX (Metal = Fe, Al and Cr. respectively).
132
orbitals in Al(TAME5OX) (orbitals 122 = HOMO-2, 123 = HOMO-1 and 124 =
HOMO), which are each localized on only one of the three equivalent 8HQ binding
units. For the Al(TAME5OX) complex, the HOMO-2 (-9.471 eV), HOMO-1 (-
9.340 eV) and HOMO (-9.314 eV) orbitals (cf. Figure 3.31) are very close in
energy. The same holds true for the LUMO, LUMO+1, and LUMO+2 orbitals.
Importantly, the metal centers do not participate in a significant way. The same
trend can be seen in the Fe for HOMO-1(-9.678 eV) and HOMO (-9.473 eV)
orbitals and for the LUMO (-9.338 eV), LUMO+1(-9.318 eV). Similarly, the
energy of unoccupied orbitals of Cr(TAME5OX) complex is close while little
variation in case of occupied orbitals (cf. Figure 3.31 orbitals: HOMO-2 (-11.169
eV), HOMO-1 (-10.452 eV) and HOMO (-9.683 eV)). The orbitals are almost
identical and show no dependence on the nature of the metal center.
The results of the TD-DFT calculations on the complexes are summarized in
Table 3.6. Two long-wave and two short-wave bands are found in each case. For all
three metal complexes these absorptions are predominantly composed of transitions
involving HOMO→LUMO (Fe,127→128; Al, 124→125; Cr, 126→127),
HOMO→LUMO+1 (Fe,127→129; Al, 124→126; Cr, 126→128), and
HOMO→LUMO+2 (Fe,127→130; Al, 124→127; Cr, 126→129). In every case,
the strongest transitions are predicted to be the long wave ones (with oscillator
strengths f ≈ 0.13-0.18). While metal-centered orbitals do not participate in the
frontier orbitals, but there is trend toward shorter wavelengths for these transitions
with heavier metals (e.g., Table 3.6: Fe→Al→Cr = 265.5 nm→260.3 nm→245.9
nm) most likely due to the concomitant elongation of the M-O bonds, resulting in a
small increase in the dimensions of the complexes with retained general shape.
133
3.2.2.3.2 Experimental and TD-DFT Studies: Luminescence Spectra
Since the tripodal ligand bears three 8hydroxyquinoline fluorescent probes
(absolute quantum yields1-3%), which act as metal binding units, it was worthy to
examine in detail the emission properties of corresponding metal complexes. To
examine the pH-dependent optical properties of metal complexes of TAME5OX
with unique pyridyl -N and -OH structural characteristics, an investigation of the
fluorescence emission of 1:1 solutions of M:ligand over a range of pH were carried
out. The pH vs fluorescence titration experiments (cf. Figure 3.24 for
Fe(TAME5OX) were carried out in aqueous medium with an appropriate volume of
0.1M HCl and 0.1M NaOH were used for protonation of pyridyl N-group and
deprotonation of the OH-group, respectively (as discussed in complexation section).
For TAME5OX to be suitable as a probe for the rapid monitoring of aqueous metal
ions in environmental or biological settings, the fluorescence emission should be
tolerant to changes in pH that may occur in unbuffered natural systems or in the
analysis of samples from strongly acidic and/or basic environment. The changes in
the blue-green emission maxima of M(TAME5OX) located at 425 nm were
monitered at solution pH values ranging from 1 to12 (cf. Figure 3.24). While
emission at 425 nm was greatly reduced in strongly acidic and basic media,
significant enhancement in the behavour of metal complexes were observed upon
measuring emission at a range of pH from 1 to 12. Inconsistant activity over such a
wide range of pH makes TAME5OX applicable for the analysis of environmental
samples that would occur well within this extended range of pH, or for use in
buffered medium.
134
Table 3.6: Computed Optical Properties: Absorption-S0 and emission-S1; wavelengths, oscillator strengths (fcalc) and transition composition of
excited-state functions, calculated at the TD-DFT/B3LYP/631-G* level of theory using DFT/B3LYP and DFT/cam-B3LYP for ground state and
excited state geometries, respectively, of the molecules.
S0 S1
Molecule (nm)
Abs
fcalc Nature Transition
(Excitation-energy) (nm)
Em
fcalc Nature Transition
(Excitation-energy)
Fe(TAME5OX) 405.91
370.34
329.55
264.84
0.1042
0.1038
0.0734
0.0347
HOMO-2→LUMO
HOMO→LUMO+1
HOMO-1→LUMO
HOMO→LUMO+2
HOMO→LUMO
HOMO→LUMO+3
125→128 (0.216 eV)
127→ eV
126→ eV
127→ eV
127→ eV
127→ eV
478.64
0.1594
HOMO→LUMO 127→ eV
127→ eV
Al(TAME5OX) 392.91
356.34
287.45
260.84
0.1094
0.1334
0.0767
0.0803
HOMO-2→LUMO
HOMO→LUMO+1
HOMO-1→LUMO
HOMO→LUMO+2
HOMO→LUMO
HOMO→LUMO+1
122→125 (0.306 eV)
124→ eV
123→ eV
124→ eV
124→eV
124→ eV
590.64
510.26
0.0194
0.2334
HOMO-1→LUMO
HOMO→LUMO
123→ eV
124→eV
123→eV
124→ eV
Cr(TAME5OX) 360.91
346.34
293.55
242.84
0.1164
0.1194
0.0267
0.0103
HOMO-2→LUMO+1
HOMO-2→LUMO+2
HOMO-1→LUMO+2
HOMO→LUMO
HOMO→LUMO+4
HOMO→LUMO+6
124→129 (0.256 eV)
126→ eV
126→ eV
126→ eV
126→ eV
126→ eV
421.64
0.1694
HOMO→LUMO
126→ eV
126→eV
Orbital numbers according to figure 3.31
135
With this (pH dependant emission) data in hand, we proceeded to explore the
fluorescence properties of TAME5OX in presence of metal ions at physiological pH
(7.4). Since the tripodal ligand shows efficient fluorescence (absolute quantum
yield ≈0.19) at this pH, it was worthy to examine in detail the emission properties of
TAME5OX corresponding to particular metal ions. The fluorescence intensities
were measured as a function x, where x is defined as the molar fraction of M(III)
versus ligand concentration. Typical fluorescence emission spectra recorded by the
titration of a 50μM aqueous solution of TAME5OX with increasing amounts of
Fe+3
and
Cr+3
solutions, showed continuous fluorescence quenching of the
bluegreen fluorescence (max = 475 nm and 425 nm, respectively) at pH 7.4,
shown in Figure 3.32 (a-b). An efficient quenching of the red shifted (of about 50
nm) and ligand-centered emission was observed (ϕ = 0.008 and 0.01), in the
presence of 1.2 and 1.4 equivalents of Fe+3
and Cr+3
, respectively. No further
quenching was observed upon the addition of molar excesses of Fe(III) and Cr(III)
to the systems. The quenching of fluorescence spectra of TAME5OX in the
prescence of these metal ions might be explained via energy or electron transfer or
the heavy atom effect. The minimum fluorescence intensity value of ligand reached
in the presence of excess iron(III) is about zero, but non-negligible (20%) for
Chromium(III). In the later case, the intramolecular quenching is less efficient, and
could be explained by the dissociation at pH 7.4 of a protonated form (LH) in
TAME5OX. The protonated bound 8-hydroxyquinoline moiety is responsible of the
blue-shifted residual emission (cf. Figure 3.32). The 52μM and 54μM of Fe(III) and
Cr(III) solutions that resulted from the titration had a pH of 5.2 which confirms that
none of the observed fluorescence quenching is due to the acidification of the
sample.
136
Figure 3.32: Quenching of the fluorescence emission spectrum of TAME5OX (1.0×10
-5M)
in water at pH 7.4, T = 25.0(3)ºC (excitation and emission slit widths of 5.0 nm), upon
increasing concentration of (a) Fe3+
from 0→1.2 equivalent (exc = 385 nm) (b) Cr3+
from
0→1.4 equivalent (exc = 345 nm).
The highly emissive nature of the Al(TAME5OX) complex enables the
monitoring of the reaction stoichiometry by emission titration. Emission for
TAME5OX is initially low (Figure 3.33) in very diluted (1×10-6
) aqueous solution
with emission slit widths of 2.5 nm at pH 7.4 but incremental addition of Al+3
to the
solution causes the emission intensity at 425 nm to decrease, concurrently forming a
new emission peak at 500 nm. In the presence of 1 equivalent of Al3+
, the mixture
showed an intense red fluorescence (cf. Figure 3.33) with a quantum yield of 0.31,
and an approximately 90fold enhancement in fluorescence signal at 500 nm was
estimated. The typical shift in emission peaks in higher wavelength region (500 nm)
relative to the emission of TAME5OX and significant enhancement of fluorescent
signal could be explained on the basis of fact that as in the free ligand fluorescence
quenching may be associated with the internal charge transfer, since the
fluorescence emission of the complex depends on the influence of metal cation on
this photoinduced charge transfer. The Al3+
cation, because of its large charge
density, should be theoretically impair reduce the charge transfer on excitation, thus
restoring the fluorescence emission.
(a) (b)
137
Figure 3.33: Evolution of the fluorescence emission spectrum (exc = 371 nm, excitation
and emission slit widths of 2.5 nm) of TAME5OX (1.0×10-6
M) upon increasing
concentration of Al3+
in ethanol from 0→1.0 equivalent at pH 7.4, T = 25.0(3)ºC.
To get the insight on mechanism of interesting fluorescence behaviour metal
complexes of TAME5OX, theoretical calculations for excited state DFT/cam-
B3LYP with 6-31G* optimized structures were carried out. The calculated emission
spectra of these complexes are shown in Figure „b‟ of 30. The computed emission
wavelengths (em) is in agreement well with experimental data, not only for the
peak values at 425, 475 and 500 nm for Cr, Fe and Al complexes, respectively, but
also the shape of the emission spectra. Computed vertical S1→S0 emission energies
(Table 3.6) are, as expected, all redshifted with respect to the corresponding
absorption values. This observation is in line with the very small structural changes
computed for each form when going from the ground to the excited state. Indeed,
the largest structural evolutions are related to a modification of the hydrogen bond
138
network at the excited state due to the photoacidity or photobasicity of the different
functional groups. In particular, the NH+/N pair seems to act as a photobase, the
nitrogen atoms interacting stronger with water‟s hydrogens at the excited state than
at the ground state. On the contrary, the OH function acts as a photoacid, its
hydrogen interacting stronger with water‟s oxygen at the excited state than at the
ground state. These results can justify the off-on-off behaviour as fluorescence
observed for these complexes in solution.
3.2.2.4 Natural Bond Orbital (NBO) analysis
As an adequate method for studying intra– and intermolecular bonding and
interaction among bonds, NBO analysis was performed on TAME5OX and its
metal complexes (Table 3.7). According to NBO analysis, the weaker hydrogen
bond interaction was formed between the nitrogen and hydroxyl of the quinoline
group with 3.76 kcal/mol of stabilization energy. It is evident that the electron
density (ED) in O–H antibonding ζ* orbital was significantly increased (0.00709e)
by the strong hydrogen bond between hydroxyl group (O–H) and nitrogen N) in the
molecule. The interaction between N and H–O leads to an increase in electron
density (ED) of O–H antibonding orbital. The increase of population in O–H
antibonding orbital weakens the O–H bond, which event in bond elongation and
arrive red shift of stretching frequency in IR spectrum.
The nature of metal–terminal oxygen and nitrogen interaction in TAME5OX–
M (M = Fe, Al and Cr) complexes were also interpretated by NBO analysis. Each
natural bond orbital (NBO) ζAB is written in terms of two directed valence hybrids
(HHOs) hA and hB on atoms A and B:
ζAB = cAhA + cBhB
139
where cA and cB are polarization coefficients. Each valence bonding NBO must in
turn be paired with a corresponding valence anti–bonding NBO:
ζ*AB = cBhA + cAhB
to complete the span of the valence space. The Lewis–type (donor) NBOs are
thereby complemented by the non–Lewis–type (acceptor) NBOs that are formally
empty in an idealized Lewis picture [35-36]. The interactions between „filled‟
Lewistype NBOs and „empty‟ non–Lewis NBOs impel to loss of occupancy from
the localized NBOs of the abstracted Lewis structure into the empty non–Lewis
orbitals, and are ascribed to as „delocalization‟ corrections to the zeroth–order
natural Lewis structure [41]. The results of NBO analysis of TAME5OX and
M[TAME5OX] complex are shown in Table 3.7. The detected natural M–O bond
orbitals are of p and d character for Fe and Cr complexes; they arise from overlap of
the three occupied oxygen and nitrogen porbitals with the empty dxy, dxz and dyz
metal orbitals, while for Al complexes d character was negligible. The oxine
subunit has lone pair orbitals, and electron density is strongly delocalized into non–
Lewis metal orbital. Accordingly, the metal–oxygen and metal–nitrogen
interactions in the investigated complexes can be described as a coordinate–
covalent bond.
140
Table 3.7: The Occupancy of the Calculated Natural Bond Orbitals (NBOs) Showing the Lewis and non–Lewis Orbitals of TAME5OX and NBOs
Between the Iron and the Oxine Unit for Fe[TAME5OX], Al[TAME5OX] and Cr[TAME5OX].
Bond(A-B) Occupancy BD s(%) p(%) d(%) Occupancy BD* s(%) p(%) d(%)
O46-H85 0.99701 0.8684(sp3.16)O
0.4959(s)H
23.98%
100.00%
75.83%
-
0.19% 0.00555 0.4959(sp3.16)O
-0.8684(s)H
23.98%
100.00%
75.83%
-
0.19%
O47-H83 0.99671 0.8794(sp3.05)O
0.4761(s)H
24.63%
100.00%
75.20%
-
0.17% 0.00840 0.4761(sp3.05)O
-0.8794(s)H
24.63%
100.00%
75.20%
-
0.17%
O48-H84 0.99803 0.8782(sp3.13)O
0.4783(s)H
27.65%
100.00%
72.23%
-
0.11% 0.00797 0.4783(sp3.13)O
-0.8782(s)H
27.65%
100.00%
72.23%
-
0.11%
Fe-O46 0.99303 0.5709(d1.27)Fe
0.8223(p)O
20.00%
16.00%
41.66%
54.90%
38.34%
29.10%
0.00445 0.8223(d1.27)Fe
-0.5709(p)O
16.00%
20.00%
54.90%
41.66%
29.10%
38.34%
Fe-O47 0.99281 0.5683(d1.65)Fe
0.8239(p)O
12.63%
11.71%
49.91%
59.98%
32.34%
28.32%
0.00238 0.8239(d1.65)Fe
-0.5683(p)O
11.71%
12.63%
59.98%
49.91%
28.32%
32.34%
Fe-O48 0.99326 0.5791(d1.39)Fe
0.8106(p)O
18.70%
13.00%
50.94%
59.98%
30.36%
29.02%
0.00530 0.8106(d1.39)Fe
-0.5791(p)O
13.00%
18.70%
59.98%
50.94%
29.02%
30.36%
Al-O46 0.98719 0.5107(p1.67)Al
0.8519(p)O
12.74%
27.27%
63.90%
70.98%
23.36%
05.56%
0.25625 0.8519(d1.67)Al
-0.5107(p)O
27.27%
12.74%
70.98%
63.90%
05.56%
23.36%
Al-O47 0.98707 0.5189(p1.77)Al
0.8530(p)O
29.03%
21.92%
48.36%
72.85%
22.61%
06.99%
0.25109 0.8530(d1.77)Al
-0.5189(p)O
21.92%
29.03%
72.85%
48.36%
06.99%
22.61%
Al-O48 0.98735 0.5264(p1.78)Al
0.8487(p)O
29.37%
27.02%
50.62%
66.85%
20.07%
09.65%
0.25351 0.8487(d1.78)Al
-0.5264p)O
27.02%
29.37%
66.85%
50.62%
09.65%
20.07%
Cr-O46 0.99412 0.5612(d1.02)Cr
0.8156(p)O
10.25%
36.00%
11.45%
57.90%
69.65%
06.10%
0.00310 0.8156(d1.02)Cr
-0.5612(p)O
36.00%
10.25%
57.90%
11.45%
06.10%
69.65%
Cr-O47 0.99437 0.5594(d1.12)Cr
0.8317(p)O
07.17%
31.71%
12.55%
59.98%
70.49%
08.32%
0.00314 0.8317(d1.12)Cr
-0.5594(p)O
31.71%
07.17%
59.98%
12.55%
08.32%
70.49%
Cr-O48 0.99279 0.5683(d1.44)Cr
0.8214(p)O
09.74%
23.00%
16.75%
69.98%
67.11%
09.02%
0.00319 0.8214(d1.44)Cr
-0.5683(p)O
23.00%
09.74%
69.98%
16.75%
09.02%
67.11%
141
3.2.3 Solution Thermodynamics and Photophysical Behaviour of
Tris(aminomethyl)ethane-5-oxine with La3+
, Eu3+
, Tb3+
and Er3+
Lanthanide chemistry has been experiencing an upsurge in research activities since
the late 1980s, gaining further momentum from Lehn‟s proposal in 1990 that
lanthanide complexes could be regarded as light conversion molecular devices [42].
Due to their escalating applications in chemical, biomedical, industrial, analytical
[43] and other fields, especially as fluorescent materials [44] and as ideal probes in
the studies of biological systems [45], has led to the discovery of numerous new
compounds. Also, owing to the lanthanides unique electronic structures, optical and
magnetic properties well-designed lanthanide complexes have been extensively
used in magnetic resonance image, biological assays, light emitting diodes,
telecommunications, optical fibers, light amplifiers, and lasers, etc. [46]. Due to the
effect of inner core shells, the f electrons do not involve in the chemistry of their
complexes. The f–f transitions are laporte forbidden [47], inducing low absorption
coefficients due which the photophysical properties of these ions are markedly
dependent on their environment [48].
Strongly absorbing chromophores are
generally employed to populate the lanthanide excited states, to attain efficient
emissions [49]. Such chromophores could furthermore shield the Ln3+
centre from
solvent molecules which can quench lanthanide emissions [50]. As a consequence,
a number of chromophoric ligands, especially the β-diketonate [51], carboxylate
[52], and 8-hydroxyquinolinate [53] ligands are used to effectively transfer light
energy to the metal for sensitized lanthanide emission because they display intense
absorptions in the UV-vis region. Especially 8-hydroxyquinoline (8-HQ) and its
derivatives are versatile coordination ligands towards a wide range of metal ions,
including lanthanide ions. In fact, Ln3+
complexes of 8-hydroxy quinolinates have
been considered as one of the most promising materials for the design of
electroluminescent devices [54] and are ideal candidates as light-harvesting
142
chromophores for sensitization of visible and NIR luminescence from Ln3+
ions
[55]. However, 8-hydroxyquinolines being bidentate monoanionic ligands and
hence cannot saturate coordination sphere of a lanthanide ion upon formation of a
charge-neutral tris-complex [56]. The lack of coordination control around the
lanthanide ion imposes a challenge to the formation of lanthanide supramolecular
architectures. To make the outcome of complexation more predictable, it is
advantageous to employ polydentate 8-hydroxyquinolines for coordination with
lanthanides and many chelates as exemplified at of Figure 3.34 have been
designed, incorporating quinoline moieties into podand structures for multidentate
coordination [57].
Figure 3.34: Tetrapodal hydroxyquinoline ligands for the coordination of rare earth ions.
3.2.3.1 Complexation. Interaction with La3+
, Eu3+
, Tb
3+ and Er
3+ Ions.
Keeping in view the importance of lanthanide 8-hydroxyquinolinate-tris chelates, it
was intended to study the interaction of La3+
, Eu3+
, Tb
3+ and Er
3+ ions
with new
synthesized
nona-dentate TAME5OX in solution by the combined use of
potentiometric, spectrophotometric and luminescence titrations. These ions were
chosen as standard example because La is the very first ion and Er lies in the second
half part of the Ln3+
series whereas Eu and Tb lounge approximately in the middle
part of the series. Not much difference in the stability of the resulting podates is
expected along the lanthanide series [58].
143
For lanthanide ion, the metal ion varies its coordination number from 6 to 12,
and TAME5OX a nonadentate N6O3 chelator is expected to satisfy the most
common nine-coordination. The four different above mentioned Ln(III) complexes
of TAME5OX examined with the coordination scan were found to have a hydration
number of 3. As TAME5OX facilitates 9-coordinate a square-face tricapped
geometry of Ln3+
with stoichiometry [M(TAME5OX)(H2O)3] with bonding six sites
of three oxinate groups leaving three coordination sites open for binding waters is
proposed. The molecular modeling parabola (Figure 3.35) clearly shows that when
water molecules are added or removed to the Ln3+
complex of the said ligand the
stability get distorted. At 9 coordination, the stability (as measured in strain energy, -
E) is well elevated than the 8 and 10 coordination state.
Figure 3.35: Coordination scan plots for the ligand (TAME5OX) with the varied
coordination number for Ln3+
complexes.
The method used for complexation studies is similar that employed for Fe3+
,
Al3+
and Cr3+
in the previous section. The potentiometric titration curves of a 1:1
solutions of La3+
, Eu3+
, Tb3+
and Er3+
ions and ligand TAME5OX are shown in
Figure 3.36. The curve ii, iii, iv and v for La3+
, Eu3+
, Tb3+
and Er3+
respectively,
showed deviation from the curve of ligand i alone. This indicates the formation of
144
complexes due to displacement of protons from TAME5OX by the lanthanide
ions. It may be noted that formation of hydroxo-complexes were also detected at
pH above 9.0 as expected for lanthanides. The potentiometric data were refined
using the Hyperquad and the formation constants obtained for La[TAME5OX],
Eu[TAME5OX], Tb[TAME5OX] and Er[TAME5OX] are reported in Table 3.8 as
expected for lanthanides. The formation of the complex species and their
cumulative formation constants, 11n, are defined by equations 3.12 and 3.13,
respectively.
𝐿𝑛𝑎𝑞 3+ + 𝐿3− + 𝑛𝐻+ ⇋ 𝐿𝑛 𝐻𝑛 𝐿 𝑛+ 3.12
𝛽11𝑛 = 𝐿𝑛 𝐻𝑛𝐿 𝑛+ 𝐿𝑛3+ 𝐿3− 𝐻+ 𝑛 (3.13)
Figure 3.36: Potentiometric titration curves: (i) 5×10
-5M TAME5OX, (ii)
[TAME5OX]/[La3+
] = 1/1, 5×10-5
M (iii) [TAME5OX]/[Eu3+
] = 1/1, 5×10-5
M, (iv)
[TAME5OX]/[Tb3+
] = 1/1, 5×10-5
M, and (v) [TAME5OX]/[Er3+
] = 1/1, 5×10-5
M,. Solvent
H2O, I = 0.1M (KCl), T= 25(2)°C, and ‘a’ is the moles of base added per mole of
Ln[TAME5OX] present. Symbols and solid lines represent the experimental and calculate
data, respectively.
The spectrophotometric titration results are shown in Figure 3.37. The data were
fitted separately in two different pH ranges. Between pH 1.9 and 7.5, the best fit of
absorbance data corresponds to the presence of four protonated complex species for
Tb3+
and three for Er3+
, but only two for La3+
and Eu3+
, mainly because several
protonated forms of the ligand coexist in this pH range. Combination of these data
145
with the data obtained from the potentiometric titrations allowed to fit four proton-
dependent equilibria described by equations 3.14 to .
𝐿𝑛𝑎𝑞 3+ + 𝐻6𝐿
3+ ⇋ 𝐿𝑛 𝐻5𝐿 5+ + 𝐻+ (3.14)
𝐿𝑛 𝐻5𝐿 5+ ⇋ 𝐿𝑛 𝐻4𝐿
4+ + 𝐻+ (3.15)
𝐿𝑛 𝐻4𝐿 4+ ⇋ 𝐿𝑛 𝐻3𝐿
3+ + 𝐻+ (3.16)
𝐿𝑛 𝐻3𝐿 3+ ⇋ 𝐿𝑛 𝐻2𝐿
2+ + 𝐻+ (3.17)
Above pH 7.3, presence of three more species were obtained La3+
, Eu3+
, Tb3+
whereas four species were obtained for Er3+
both from potentiometric and
spectrophotometric data. Their formation constants defined by equations 3.18 to
3.21 were refined by considering two successive deprotonation and the formation of
one hydroxo complex at pH higher than 9.0. The results are tabulated in Table 3.8.
𝐿𝑛 𝐻2𝐿 2+ ⇋ 𝐿𝑛 𝐻𝐿 + + 𝐻+ (3.18)
𝐿𝑛 𝐻𝐿 + ⇋ 𝐿𝑛𝐿 + 𝐻+ (3.19)
(𝐿𝑛𝐿) ⇋ 𝐿𝑛𝐿 𝑂𝐻 − + 𝐻+ (3.20)
(𝐿𝑛𝐿) + (𝐻6𝐿) 3+ ⇋ 𝐿𝑛𝐿2 + 3𝑂𝐻− (3.21)
La[TAME5OX]:
146
Eu[TAME5OX]:
Tb[TAME5OX]:
Er[TAME5OX]:
Figure 3.37: UV−vis absorption spectra of 1:1 solution of Ln3+
and TAME5OX as a
function of p[H], [TAME5OX] = [Ln3+
]tot = 5×10-5
M, Solvent: H2O, I = 0.1M(KCl), T =
25.0(2)°C. For La[TAME5OX]: (a) p[H] = 1.94-7.49 (b) p[H] = 7.78-10.25; for
Eu[TAME5OX]: (a) p[H] = 1.95-7.21 (b) p[H] = 7.51-10.82; for Tb[TAME5OX]: (a) p[H]
= 1.93-7.34 (b) p[H] = 7.49-10.88; and for Er[TAME5OX] (a) p[H] = 1.94-7.44; (b) p[H]
= 7.51-10.65.
The luminescence titrations are presented in Figure 3.38. For La(TAME5OX)
system, the best refinement of the luminescence-pH data by the Hypspec program
corresponds to the presence of four species, two protonated complex species that is,
147
[La(H2L)]3+
[log = 40.98(3)] in consistent with spectrophotometric data, and
[La(HL)]2+
[log = 36.78(5)] which was found in concert with both potentiometric
and spectrophotometric data. Along with the protonated species, one neutral species
[(LaL)] [log = 33.41(4)] and one hydroxo-complex [LaL(OH)]
- [log =9.92(1)]
were also formed that favourably consents the formation of these species by both
potentiometric and spectrophotometric method. For Eu(TAME5OX) system, the
luminescence-pH corresponds to the presence of five species, three protonated
complex species that is, [Eu(H5L)]5+
[log = 55.76(6)] and [Eu(H2L)]
2+ [log =
40.49(2)] in consistent with potentiometric and spectrophotometric data,
respectively, and [Eu(HL)]+
[log =
36.39(4)] in agreement with both
potentiometric and spectrophotometric data. One neutral species [(EuL)] [log =
33.41(4)] and one hydroxo-complex [EuL(OH)]- [log =9.92(1)]
were also formed
that favorably agree the formation of these species by both potentiometric and
spectrophotometric method. In case of Tb(TAME5OX) system in the pH range of
1.93-7.54, the best fit of the luminescence-pH data returned considering formation
of three protonated complexes, [Tb(H3L)]3+
species [log = 46.43(3)] concurred
favorably with spectrophotometric data, while [Tb(H2L)]2+
and [Tb(HL)]
+ species
{log = 39.76(3) and 35.37(5), respectively}, which were found in good agreement
with both potentiometric and spectrophotometric results. Above pH 7.4, the
luminescence data impart to the presence of neutral (TbL) [log = 32.23(5)],
(TbL2) [log = 29.43(6)], and hydroxo [TbL(OH)]
- [log =
8.24(4)] species, were
found in good concert with both potentiometric and spectrophotometric data. For
Er(TAME5OX) system in the pH range of 1.93-7.54, the best fit was attained
considering formation of three protonated complexes, Er(H2L)]2+
[log = 37.81(5)],
is in concurrence with only the spectrophometric value, while [Er(H3L)]3+
and
148
[Er(HL)]+
species {log = 45.78(1) and 33.54(3), respectively} were observed in
agreement with both potentiometric and spectrophotometric results. Above pH 7.6,
the luminescence data impart to the presence of neutral (ErL) [log = 35.81(5)],
(ErL2) [log = 37.32(3)], and hydroxo [ErL(OH)]
- [log =
7.98(3)] species in
consistent to both potentiometric and spectrophotometric data, which corresponds
to the presence of [(ErL2)] complex [log = 37.85(4) and 37.39(5), respectively]
along with (ErL) and [ErL(OH)]-. It is worth to mention that all species could not be
detected by a single method, whereas they could be detected by combination of
potentiometry, UV-visible spectrophotometry and fluorometry (spectro-
(a) La[TAME5OX] (b) Er[TAME5OX]
(c) Eu[TAME5OX] d) Tb[TAME5OX] Figure 3.38: Dependence of the emission spectra of 1:1 aqueous solution of Ln:TAME5OX
(5.0 × 10-5
M) on the change of the pH value from acidic to basic range, T = 25.0(3)°C
(excitation and emission slit widths of 5.0 nm). (a) for La[TAME5OX] (1.94-10.23 pH), (b)
for Er[TAME5OX] (1.94-10.65 pH), (c) for Eu[TAME5OX] (1.7-12.3 pH) and (d) for
Tb[TAME5OX] (1.9-12.5 pH); with an excitation wavelength of 325 nm for La and Eu and
380 nm for Er and Tb.
149
electrometric). The formation constants {(log, obtained by all the three techniques,
potentiometrically as well as spectrophotometrically (UV-visible and fluorescence
techniques)} of these complexes, which were defined above by equations 3.12 and
3.13, together with the average log values are reported in Table 3.8. The step-wise
formation constants (log K), which are direct indicator of stability, are also returned
in the table.
In order to verify the validity and assignments made for the experimental
formation constants, aqueous-phase free energy for the entire representative species
for all the four metal ions were calculated through semi-empirical sparkle/PM7
COSMO quantum mechanical approach.
For the metal complex MLH(n-1), the log K defined as the negative logarithm of
the formation constant of the reaction Lnaq + LHn = [Ln(LH)(n-1)]+ + H3O
+, is given
by the thermodynamics relation ∆G° = -RT (log K/2.303). The change in free
energy in aqueous-phase is obtained according to the following equation:
∆𝐺𝑎𝑞 ,[𝐿𝑛(𝐿𝐻)𝑛 ] = 𝐺𝑎𝑞 ,[𝐿𝑛 𝐿𝐻 (𝑛−1)] + 𝐺𝑎𝑞 ,𝐻3𝑂+ − 𝐺𝑎𝑞 ,𝐿𝑛 + 𝐺𝑎𝑞 ,𝐿𝐻𝑛 + 𝐺𝑎𝑞 ,𝐻2𝑂 (3.22)
Since the metal ions compete with the protons for the coordination sites, the
complexation of metal ion with a ligand occurs upon release of hydrogen and
greater acidity belongs to the group in which its hydrogen release is easier. A
clear decrease in ∆G° was observed for Ln(TAME5OX) (as shown in Figure 3.39)
as the deprotonation take place from fully protonated free ligand, LH96+
, upon
complexation with lanthanide metal ions. The validity of this calculated ∆G° for the
different species of La, Eu, Tb and Er complexes of TAME5OX were compared
with the calculated experimental log K, which resulted an acceptable correlation
with R2 = 0.9893, 0.99562, 0.98891 and 0.99041, respectively. It is known that, the
properties of chemical species match, if the theoretical geometry coincides with the
150
actual geometry of molecule. The theoretical electronic spectra of each species
matches with experimental observed spectra, suggest the authencity of the modelled
structures of protonated, neutral or hydroxy species.
Figure 3.39: Correlation between the Experimental Log K and Calculated ∆Gº of (a)
La(TAME5OX), (b) Eu(TAME5OX) (c) Tb(TAME5OX), and Er(TAME5OX).
To best of our knowledge no neutral lanthanide complexes (ML) of tripodal
chelator based on 8-hydroxyquinoline has been reported so far. For occupancy of
the same number of coordination sites of three bidentate oxine units of the
TAME5OX, the log (or log K) of M(TAME5OX) can be compared with 3 of
M(Ox)3 complex. Surprisingly, the chelation of TAME5OX caused tremendously
enhanced stability of about 8-9 orders of magnitude in log compared to Ln(ox)3:
(log 3=23.37) [59]. These surprising results led us to check the model several times
determined by each method by varying the log values. The best fit results were
obtained only at 34.02, 33.41, 32.23 and 33.15 for LaL, EuL, TbL and ErL,
correspondingly, and also only at 29.43 and 29.83 for TbL2 and ErL2, respectively.
The stability of these complexes may be viewed as described below: In the present
151
case, as seen from ΔGº value (Figure 3.39) there is a strong driving force for the
ligand with formation of thermodynamically favourable complexes. This can be
interpreted as dominated by a favourable entropy change associated with chelation
and also as a result of the anchoring effect [60]; after coordination of one unit of
ligand, attachment of the second and third oxine subunits of tripodal TAME5OX is
easier than oxine molecule from outside. Also, there are various other effects like
chelating effect, pre-organization that arise as a result of molecular shape that add
complications. The topology of C-pivot tripodal ligand TAME5OX is such that
the basic skelton containing the methyl group and three pendant arms at the
tetrahedral pivot carbon (sp3 hybrid) gives binding pendant arms a pyramid shape,
which cooperatively work in the coordination of electron donating binding
units towards metal ions. Also, the donor groups are pre-organized in more
appropriate positions and in close proximity („a‟ of Figures 3.40) for face-capped
binding to the lanthanide ions with three water molecules on opposite side in case
Figure 3.40: Sparkle/PM7 optimized ground state geometries in COSMO (water) of
complexes (a) TAME5OX, showing preorganized three oxine units as readiness for metal
complex formation (b) La[TAME5OX], (c) Er[TAME5OX], (d) Eu[TAME5OX] and (e)
Tb[TAME5OX].
of ML and with no water molecule attached in case of ML2 as the coordination
sphere of metal ion gets completely saturated. Whereas free, flexible oxine
molecules approaching three dimensionally must undergo significant
translational motion “stitch” themselves onto the metal ion, which is less favored.
Though similar enhanced stability for lanthanide chelates has not been documented
(a) (b) (c) (d) (e)
152
so far, but natural occurring iron chelates in siderophore show tremendous high
stability; for example, Fe(enterobactin) (log = 49 [61]) containing three
catecholate units is 18 times more stable than Fe(catechol)3 (log3 =31) [62].
The distribution diagrams of various species formed are shown in Figures 3.41.
Despite the complexity of the system, LaL is the major species (98.5%) present in
the range of 7.5-14.0 pH. Other species present in La(TAME5OX) system at acidic
pH is [La(H4L)]4+
(100% at pH ≤ 4.2); upon an increase in the pH, successive
deprotonation of La(LH4) species leads to the formation
of [La(H3L)]
3+, [La(H2L)]
2+
and [La(LH)]+ which exists at pH 5.2 (35%), 6.4 (64%) and 7.6 (98.4%),
respectively. For Eu(TAME5OX), EuL is the major species (98.5%) present in the
range of 7.1-14.0 pH. Additional species present at very acidic pH is [Eu(H5L)]5+
(100% at pH ≤ 2.8); upon an increase in the pH, successive deprotonation of
Eu(H5L) species leads to the formation of [Eu(H4L)]
4+, [Eu(H3L)]
3+, [Eu(H2L)]
2+
and [Eu(HL)]+ which exists at pH 4.3 (62%), 5.0 (38%), 5.8 (62%) and 7.8 (98.3%)
(98.3%), respectively. Likewise in case of Tb, [Tb(H5L)]5+
complex predominates at
pH 1.3; an increase in pH leads to diminish this species and new [Tb(H4L)]4+
species
starts forming with maximum concentration at pH 2.9 (55%). Due to further
deprotonation, [Tb(H3L)]3+
was found as major species during pH 2.0-6.1 (74%) along
with [Tb(H2L)]2+
at pH 5.0 (75%). [Tb(HL)]+
exists 42% at physiological pH and the
neutral complexes TbL and TbL2 were found during the pH range of 5.7-12.7 with
concentration 78% and 17%, respectively, with TbL2 remains dominant over higher pH
range. In case of Er, [Er(H4L)]4+
complex predominates at pH 1.8; an increase in pH
leads to diminish this species and new [Er(H3L)]3+
species starts forming with
maximum concentration at pH 3.9 (65%). Due to further deprotonation, [Er(H2L)]2+
was found as major species during pH 3.8-7.6 (59%) along with [Er(HL)]+ at pH
153
6.2 (35%). The neutral complexes ErL and ErL2 were found at physiological pH
which exists 42% and 7%, respectively. Further scrutiny of the distribution curves
shows that LaL and ErL were dominant over higher pH range. A steep increase in
concentration of the hydroxyl complexes of each Ln3+
ion was observed from pH
10.0 and 9.2 followed by decrease down to a pH of ~13.5 for Tb3+
and Er3+
ions
only.
La(TAME5OX) Er(TAME5OX)
Eu(TAME5OX) Tb(TAME5OX)
Figure 3.41: Species distribution curves of La[TAME5OX], Er[TAME5OX],
La[TAME5OX] and Er[TAME5OX], respectively, containing species, computed from the
formation constants given in Table. Calculated for [TAME5OX]tot = [Ln3+
]tot = 10-5
M,
Solvent : H2O, I = 0.1M (KCl), T = 25.0(2)°C.
154
The complexation efficiency of a metal ion Maq by a given ligand is evaluated
by the pM(= -log[Maq]) value calculated at physiological pH. The plots of pLn
3+
versus pH for the TAME5OX over the pH range of 3-12 for La, Eu, Tb and Er are
presented in Figure 3.42. For TAME5OX, this translates into pLn values of
14.6(La), 16.14(Eu), 19.48(Tb) and 19.8(Er) compared to pEu=15.6 for Tsox [23]
or 19.6 for [Eu(dtpa)]2-
(dtpa= diethylenetriaminopentaacetic acid), as computed
from known formation constants [60]. The chelate effect of the podand with respect
to the 8- hydroxyquinoline building block is impressive and the complexes based on
TAME5OX appear to be sufficiently stable in water for potential in vivo
applications.
Figure 3.42: Plot of pLn versus p[H] for TAME5OX, pLn = -log [Ln3+
], calculated for
[Ln3+
] = 10−5
M and [L] = 10−5
M, for (a) Ln= La3+
and Er3+
and (b) Ln= Eu3+
and Tb3+.
155
Table 3.8: Equilibrium Constantsa (log 11n), Absorption, Emission Characteristics of La
3+, Er
3+, Eu
3+ and Tb
3+ Complexes of TAME5OX
and Corresponding pLn Values.
Complexes
L=TAME5OX
Species Equilibrium: Log
Log K
Log 11nb
Log K11nb
Log 11nc
Log K11nc
Log 11nd
Log K11nd
Average
Log 11n
max (nm) ɛ(M-1cm-1) em(nm) ePL pLn3+ f
(LaL)
[La(H4L)]4+ [La(H4L)4+]/[La3+][L3-][H+]4
[La(H4L)4+]/[La(H3L)3+][H+]
53.35(4)
4.93 (4)
53.43(5)
5.01(5)
_
_
53.39
262, 285 15 200, 6 800 425 0.007 14.6
[La(H3L)]3+ [La(H3L)3+]/[La3+][L3-][H+]3
[La(H3L)3+]/[La(H2L)2+][H+]
48.42(3)
7.14 (3)
48.69(4)
7.41 (4)
_
_
48.51 262, 286, 320 13 000, 7 200, 3 600 425,
478
0.007,
0.018
[La(H2L)]2+ [La(H2L)2+]/[La3+][L3-][H+]2
[La(H2L)2+]/[La(HL)+][H+]
_
_
41.28(4)
4.74(4)
40.98(3)
4.20(3)
41.13
262, 288, 320 11 200, 6 600, 4 000 425,
495
0.007,
0.026
[La(HL)]+ [La(HL)+]/[La3+][L3-][H+]
[La(HL)+]/[LaL][H+]
36.24(2)
2.30(2)
36.54(3)
2.24(3)
36.78(5)
2.96(5)
36.52
268, 293, 325 9 000, 6 500, 4 200 500 0.051
[(LaL)] [LaL]/[La3+][L3-]
[LaL]/[La3+][L3-]
33.94(3)
33.94(3)
34.30(5)
34.30(5)
33.82(5)
33.82(5)
34.02
274, 305, 335 7 400, 6 600, 4 400 500 0.059
[LaL(OH)]- [(LaLOH)-][H+]/[La3+][L4-]
[(LaLOH)-][H+]/[LaL)]
10.34(2)
23.6(2)
10.66(1)
23.64(1)
10.61(4)
23.21(4)
10.53 280, 300 5 800, 5 400 428
0.019
(ErL)
[Er(H4L)]4+ [Er(H4L)4+]/[Er3+][L3-][H+]4
[Er(H4L)4+]/[Er(H3L)3+][H+]
54.45(4)
7.01(4)
54.43(4)
6.99(4)
_
_
54.44 282, 305, 349 43 600, 22 000, 4 000 425 0.006 19.8
[Er(H3L)]3+ [Er(H3L)3+]/[Er3+][L3-][H+]3
[Er(H3L)3+]/[Er(H2L)2+][H+]
_
_
47.44(2)
7.10(2)
47.35(3)
7.02(3)
47.38 282, 305, 325, 349, 400 38 000, 26 000, 4 100, 3 820, 3 200 425,
475
0.006,
0.017
[Er(H2L)]2+ [Er(H2L)2+]/[Er3+][L3-][H+]2
[Er(H2L)2+]/[Er(HL)+][H+]
40.35(5)
4.01(5)
40.34(4)
4.00(4)
40.33(3)
4.05(3)
40.34 282, 305, 325, 349, 400 34 000, 28 000, 4 190, 3 700, 3 340 425,
500
0.006,
0.028
[Er(HL)]+ [Er(HL)+]/[Er3+][L3-][H+]
[Er(HL)+]/[ErL][H+]
36.33(5)
3.23(5)
36.34(6)
3.13(6)
36.28(5)
3.14(5)
36.31 282, 305, 325, 315, 400 32 000, 31 910, 4 300, 3 550, 3 500 500 0.041
[(ErL)] [ErL]/[Er3+][L3-]
[ErL]/[Er3+][L3-]
33.10(7)
33.10(7)
33.21(6)
33.21(6)
33.14(5)
33.14(5)
33.15
293, 319, 355 27 000, 37 800, 7 800 508 0.057
[(ErL2)] [ErL2]/[Er3+][L3-]2
[ErL2]/[ErL][L3-]
29.84(6)
3.26(6)
29.79(6)
3.42(6)
29.86(6)
3.28(5)
29.83 320, 358 33 400, 7 600 442 0.055
[ErL(OH)]- [(ErLOH)-][H+]/[Er3+][L4-]
[(ErLOH)-][H+]/[ErL)]
9.50(6)
20.34(6)
9.72(5)
20.07(5)
9.48(4)
20.38(4)
9.56 315, 345 2 000, 6 000 425 0.016
156
Complexes
L=TAME5OX
Species Equilibrium: Log
Log K
Log 11nb
Log K11nb
Log 11nc
Log K11nc
Log 11nd
Log K11nd
Average
Log 11n
max (nm) ɛ(M-1cm-1) em(nm) ePL pLn3+ f
(EuL)
[Eu(H5L)]5+ [Eu(H5L)5+]/[Eu3+][L3-][H+]5
[Eu(H5L)5+]/[Eu(H3L)3+][H+]2
55.98(4)
3.29(4)
_
_
55.76(6)
2.94(6)
55.87
245, 267 38 000, 17 200 410 0.007 16.16
[Eu(H4L)]4+ [Eu(H4L)4+]/[Eu3+][L3-][H+]4
[Eu(H4L)4+]/[Eu(H3L)3+][H+]
52.69(5)
4.96 (5)
52.85(5)
5.12(5)
_
_
52.77
245, 267, 300 30 400, 16 000, 4 500 410,
470
0.008,
0.005
[Eu(H3L)]3+ [Eu(H3L)3+]/[Eu3+][L3-][H+]3
[Eu(H3L)3+]/[Eu(H2L)2+][H+]
47.73(5)
7.09 (5)
_
_
_
_
47.73 246, 268, 298 26 000, 15 400, 6 000 410,
496
0.008,
0.019
[Eu(H2L)]2+ [Eu(H2L)2+]/[Eu3+][L3-][H+]2
[Eu(H2L)2+]/[Eu(HL)+][H+]
_
_
40.64(3)
4.52(3)
40.49(2)
4.10(2)
40.56
248, 269, 300 22 600, 13 000, 5 700 410,
500
0.008,
0.027
[Eu(HL)]+ [Eu(HL)+]/[Eu3+][L3-][H+]
[Eu(HL)+]/[EuL][H+]
36.12(1)
2.65(1)
36.27(2)
2.62(2)
36.39(4)
2.98(2)
36.26
250, 271, 306 19 400, 10 400, 5 900 410,
500
0.008,
0.053
[(EuL)] [EuL]/[Eu3+][L3-]
[EuL]/[Eu3+][L3-]
33.47(2)
33.47(2)
33.65(5)
33.48(5)
33.41(4)
33.41(4)
33.51
255, 275, 310 15 000, 9 000, 5 970 410,
465
0.008,
0.045
[EuL(OH)]- [(EuLOH)-][H+]/[Eu3+][L4-]
[(EuLOH)-][H+]/[EuL)]
9.67(3)
23.8(3)
9.83(2)
23.65(2)
9.92(1)
23.49(1)
9.80 248, 268 11 600, 8 000 410,
435
0.010,
0.009
(TbL)
[Tb(H5L)]5+ [Tb(H5L)5+]/[Tb3+][L3-][H+]5
[Tb(H5L)5+]/[Tb(H3L)3+][H+]2
56.11(7)
3.32(7)
56.03(6)
3.17(6)
_
_
56.07
220, 260, 280, 310 45 500, 45 000, 24 000, 5 600 420 0.006 19.48
[Tb(H4L)]4+ [Tb(H4L)4+]/[Tb3+][L3-][H+]4
[Tb(H4L)4+]/[Tb(H3L)3+][H+]
52.79(4)
6.01(4)
52.86(4)
6.08(4)
_
_
52.82 222, 260, 280, 310, 345 45 700, 36 000, 23 900, 5 650, 4 300 420,
475
0.006,
0.007
[Tb(H3L)]3+ [Tb(H3L)3+]/[Tb3+][L3-][H+]3
[Tb(H3L)3+]/[Tb(H2L)2+][H+]
_
_
46.78(2)
7.09(2)
46.43(3)
6.67(3)
46.60 224, 260, 280, 310, 345 45 840, 31 600, 23 890, 5 820, 4 000 425,
480
0.006,
0.018
[Tb(H2L)]2+ [Tb(H2L)2+]/[Tb3+][L3-][H+]2
[Tb(H2L)2+]/[Tb(HL)+][H+]
39.78(5)
4.13(5)
39.69(4)
4.11(4)
39.76(3)
4.39(3)
39.74 226, 260, 280, 315, 350 45 900, 24 000, 23 000, 6 200, 3 700 425,
495
0.006,
0.027
[Tb(HL)]+ [Tb(HL)+]/[Tb3+][L3-][H+]
[Tb(HL)+]/[TbL][H+]
35.65(5)
3.43(5)
35.58(6)
3.27(6)
35.37(5)
3.14(5)
35.53 230, 268, 289, 315, 350 45 200, 23 000, 21 000, 5 500, 3 250 500 0.043
[(TbL)] [TbL]/[Tb3+][L3-]
[TbL]/[Tb3+][L3-]
32.22(7)
32.22(7)
32.31(6)
32.31(6)
32.23(5)
32.23(5)
32.16
233, 270, 292, 318 22 200, 21 500, 4 600, 2 700 505 0.065
[(TbL2)] [TbL2]/[Tb3+][L3-]2
[TbL2]/[TbL][L3-]
29.14(6)
3.08(6)
29.39(6)
2.92(6)
29.43(6)
2.80(5)
29.32 235, 270, 295, 320 19 400, 20 750, 3 700, 2 140 472 0.051
[TbL(OH)]- [(TbLOH)-][H+]/[Tb3+][L4-]
[(TbLOH)-][H+]/[TbL)]
8.25(6)
20.89(6)
8.36(5)
21.03(5)
8.24(4)
21.19(4)
8.28 238, 265, 285 14 800, 18 400, 1 600 415 0.014
aNumbers in parentheses represent the standard deviation in the last significant digit. In water (I = 0.1M KCl, T = 25.0°C).
bPotentiometric method,
cUV-visible
spectrophotometric method, dLuminescence spectrophotometric method,
e0.1 M solution of quinine sulphate in 0.5 M H2SO4 as standard ( = 0.546),
fpLn
3+ = −log
[Ln3+
] calculated for [Ln]tot = 10−6
M, [L]tot = 10−5
M, and p[H] = 7.4.
157
3.2.3.2 Photophysical Properties
3.2.3.2.1 UV-visible spectrophotometric Studies
The ligand TAME5OX shows two peaks 257 and 310 nm (ε = 110, 000 M−1
cm−1
and 60, 100 M−1
cm−1
, respectively) in the UV-visible region assigned to →* and
n→* transitions, and on metal coordination, there is a considerable shift on both
transitions. Since, a variety of species are formed with change in pH, the substantial
spectral changes (Figure 3.37) observed in the spectrophotometric titration of the
above titled lanthanide TAME5OX complexes in acidic medium and at
physiological pH is indicative of a change in the coordination of metals with the
ligand. The f→f transitions are not expected for complexes of La3+
(f 0), whereas for
Eu3+
(f 6) and Tb
3+ (f
8) absorption peaks are expected at 460 nm, 535 nm, 540 nm
and 490 nm, 625 nm, 700 nm, respectively [63]. For complexes of Er3+
(f11
) weak
peaks are expected at 995 nm and 1,486 nm which are usually unaffected in the
ligand environment. Absence of peaks in the region 460 nm to 700 nm and 995 nm
to 1,486 nm confirm, the observed spectra are due to ligand transitions or charge
transfer transitions. At pH ≤ 2, the absorption spectrum of the TAME5OX with La
ion is epitomized by a strong absorption band at higher energies 262 nm, (ε = 15,
000 M−1
cm−1
), assigned for →* transitions and comparatively low absorption
band at 285 nm (ε = 7, 800 M−1
cm−1
), assigned for n→* transitions.
Eu(TAME5OX) is symbolized by a strong absorption band at 255 nm, (ε = 38, 000
M−1
cm−1
) and relatively low absorption band at 265 nm (ε = 17, 200 M−1
cm−1
) and
for Tb complex, bands at 205 nm (ε = 39, 600 M−1
cm−1
) and 258 nm (ε = 46, 000
M−1
cm−1
), are assigned for →* transitions. In contrarily to La and Eu complexes,
the low absorption band at 270 nm (ε = 23, 600 M−1
cm−1
), assigned for →*
transition and a less intense, non-structured band around 320 nm (ε = 4, 000
158
M−1
cm−1
), assigned for n→* transitions, were also observed at the same pH for Tb
ion. In the same manner, electronic spectrum for Er ion in the similar conditions is
characterized by strong absorption band at 283 nm, (ε = 43, 600 M−1
cm−1
) due to
→* transition, while peak at 305 nm, (ε = 21, 400 M−1
cm−1
), and the less
significant non-structured band at 349 nm (ε = 3, 800 M−1
cm−1
), are assigned for
n→* transitions.
Upon rise of pH above 2, the behaviour of TAME5OX with Tb and Er is
inconsistent and more compelling from La and Eu. In case of La(TAME5OX), with
the rise in pH the intensity of high energy band at 262 nm declines (ε = 11, 000
M−1cm−1) and at 285 nm remains almost unchanged (ε = 7, 400 M−1
cm−1
), with
the formation of isosbestic points at 251 nm and 268 nm. For Eu(TAME5OX), the
intensity of bands at 255 nm (ε = 27, 000 M−1
cm−1
) and 265 nm (ε = 14, 600
M−1
cm−1
), declines with the formation of isosbestic points at 234 nm and 275 nm,
while for Tb(TAME5OX), the ligand peak at 205 nm, (ε = 39, 600 M−1
cm−1
) get red
shifted with slight ascent around 218 nm, and the high energy band at 258 nm, (ε
=46, 000 M−1
cm−1
) declines upto (ε =33, 600 M
−1cm
−1) at same wavelength forming
isosbestic point at 235 nm. Correspondingly for Er(TAME5OX), the peak at 283
nm (ε =43, 000 M−1
cm−1
) declines (ε =31, 600 M−1
cm−1
) with the concomitant rise
in absorbance band at 305 nm (ε = 31, 520 M−1
cm−1
) with the development of
isosbestic points at 278 nm, 295 nm and 312 nm, due to the deprotonation and
complexation of Npyr groups, whereas for La complex low absorption band
appeared at 310 nm (ε =4, 200 M−1
cm−1
). In case of Eu complex, the band at 265
nm (ε =17, 200 M−1
cm−1
) declines with the simultaneous growth in absorbance at
292 nm (ε =6, 000 M−1
cm−1
) due to the deprotonation and complexation of Npyr
groups, where as in case of Tb complex, very less significant non-structured band
159
appeared at 288 nm (ε =4, 000 M−1
cm−1
), assigned for n→* transition with no
significant change of the band at 270 nm (ε = 23, 600 M−1
cm−1
). Analogously, the
very low energy band at 320 nm (ε =3, 600 M−1
cm−1
) for Tb ion and the band at
349 nm for Er ion, shifted towards higher wavelength with formation of broad
bands around 365-405 nm (ε =2, 400 M−1
cm−1
) and 380-430 nm (ε =2, 800
M−1
cm−1
), respectively, along with formation of three isosbestic points at 278 nm,
308 nm and 345 nm in case of Tb, and blue shifted less intense band at 325 nm (ε
=5, 600 M−1
cm−1
) for Er. No such shifts were observed in case of La and Eu
complex with the increment of pH. Since the absorption coefficients of these bands
are large for a d-d or f-f transition, they are likely due to ligand–to–metal charge
transfer (LMCT). Appearance of band at higher wavelengths with lower intensity
may be ascribed to the coordination of pyridine nitrogen atoms to the metal ion
upon chelation. It can thus be assumed that the band at 292 nm or 365-405 nm is for
the Npyr→Eu3+
or Npyr→Tb3+
transition and the band at 310-320 nm or 385-435 nm
is for the Npyr→La3+
or Npyr→Er3+
type, respectively. The isosbestic points assert
the formation of four protonated complexes in the pH range of 1.9-7.5 for each
metal ions, which can also be comprehend from species distribution curves (Figure
3.41).
Above pH 7.3, the deprotonation and complexation of hydroxyl groups in
case of both Tb and Er ions result in the bathochromic and hypochromic shifts („b‟
of Figure 3.37). As expected the peaks at 262 nm (ε =30, 000 M−1cm−1), and at
283 nm (ε =23, 400 M−1
cm−1
), get shifted towards a broad band around 292 nm,
elongated upto 330 nm, (ε =5, 600 M−1
cm−1
) for Tb(TAME5OX), the bands at 302
nm (ε =30, 000 M−1
cm−1
), and at 320 nm (ε =23, 400 M−1
cm−1
) changed to a broad
band around 335 nm and elongated upto 350 nm (ε =5, 600 M−1
cm−1
) for
160
Er(TAME5OX). For La(TAME5OX) the bands at higher energies were red shifted
to 290 nm (ε = 24, 000M−1
cm−1
), and 310 nm (ε = 13, 600M−1
cm−1
), for
Eu(TAME5OX), the band at lower wave length was also red-shifted to 265 nm (ε =
24, 000M−1
cm−1
), and 284 nm (ε = 13, 600M−1
cm−1
) shifts hypochromically in both
cases (Figure 3.37). For La and Eu, the low energy bands were also shifted towards
higher region at 345 nm (ε = 6, 400M−1
cm−1
), and at 320 nm (ε = 6, 400M−1
cm−1
),
respectively, indicating the formation of neutral complexes. The appearance of
observed bands at 350 nm, 325 nm, 328 nm and at 345 nm can be assigned to
charge transfer by O→La+3
, O→Eu+3
, O→Tb+3
and O→Er+3
respectively, [24]
whereas the bands due to Npyr→Ln were disappeared in this pH range indicating the
change in the coordination sphere of the lanthanide ion. In contrary to acidic to
neutral pH range, the concurrent downslope of LMCT bands of Tb and Er
complexes were observed over the pH range of 7.410.8. The behaviour of ligand
with La and Eu ions in the basic medium was somehow similar as observed in
physiological pH with slight upslope at 320 nm and appearance of isosbestic point
at 298 nm. No noticeable spectral change was observed on further variation of pH
upto 10.8 for La, and 11.5 for Eu, Tb and Er complexes, indicating the formation
of hydroxo complexes for each metal ion.
Investigation for the electronic transitions and electronic structures of different
species formed for each corresponding Ln complexes in solution, were also
performed from a theoretical stand point. Though the ab-initio and DFT methods
are more accurate for lanthanide metal coordinate compounds, it has been
established that the sparkle model in conjugation with PM6/PM7 produce
comparable result with Hartree-fock (HF) or DFT method [64]. It has been proved
that the sparkle/PM7 is an accurate and statistically valid tool for the prediction of
161
the geometrical features of lanthanide coordination polyhedral and, by design, is
expected to perform best with ligands with nitrogen or oxygen as coordinating
atoms present in the vast majority of all coordination compounds of the trivalent
rare earth metals [65-66]. It has also been suggested [67] that the use of quantum
chemical methodology other than semiempirical sparkle model such as Hartree-
Fock (HF) or density functional theory (DFT) using an effective core potential
(ECP) to treat the Ln3+
ions is unfeasible owing to the high computational effort
needed. In this study the semiempirical sparkle model calculations were
performed for all possible species. The computed optical data of all the
PM7/sparkle optimized species are listed in Tables 3.9. The simulated electronic
absorption spectra of predicted complexes obtained via ZINDOS/CIS method
are in good agreement in the profile, number of absorption bands and relative
intensities with the species obtained from experiment (Figure 3.43) except small
red-shifts in Eu complexes, which are due to the impossibility of considering all the
symmetrical orbitals in CI calculation which was not considered in the calculation
and partly due to the parametric effect used for the calculations. The ZINDO
electronic structure model does not have diffuse functions; hence, excited states of a
primarily Rydberg nature are not reproduced well. Table 3.9 present the calculated
singlet→singlet electronic transitions and corresponding oscillator strengths of
different species of metal complexes of TAME5OX and their assignments to the
experimental bands. The ground state optimized structures of TAME5OX and its
La, Eu, Tb and Er complexes are shown in Figure 3.40 (b-e). From an electronic
point of view, for all the Ln[TAME5OX] acidic forms, the second electronic
transition basically corresponds to a HOMO→LUMO excitation and orbitals have
162
La[TAME5OX]
Eu[TAME5OX]
Tb[TAME5OX]
Er[TAME5OX]
Figure 3.43: pH-dependent electronic spectra of the different species as a function of
molar absorptivity and wavelength (a) predicted from Hypspec using experimental data
and (b) calculated by applying semi-empirical sparkle PM7/ZINDOS methods for
Ln[TAME5OX] (Ln= La, Eu, Tb and Er, respectively).
163
the same character (namely and *, respectively). On analysing the molecular
orbitals, the HOMO is somehow localized on the aromatic ring carrying the OH
function while the antibonding LUMO is delocalised over the nitrogen atom of the
quinoline system; we expect this transition to be relatively sensitive, in terms of
intensity and position, to overall protonation degree of the molecules. In particular,
the significant hypochromic shifts are computed going from protonated (acidic) to
neutral species for each lanthanide ion and remarkable red shifts were also
predicted in going from [La(H4L)]4+
to LaL and [Eu(H5L)]5+
to [Eu(HL)]+ while the
blue shifts were observed of the same magnitude from [LaL] to [LaL(OH)]
- and
[EuL] to [EuL(OH)]
-. Some remarkable red shifts of about 19-25 nm were observed
in going from di-protonated to neutral forms.
A quantitative explanation of these variations in absorption transitions and
energies can be given by the inspection of the HOMO and LUMO orbitals
computed for all species shown in Figure 3.44. Indeed, protonation of the N atom
increases the acceptor character of the quinoline ring. Thus, it is expected to raise
the LUMO energy and increases the vertical S0→S1 transition energy, in
agreement with the above mentioned computed red shift on deprotonation of
NH+ upon complexation. On the other hand, the deprotonation of the hydroxyl
function will destabilize the HOMO, thus closing the gap and giving rise to smaller
transition. Notably, both the above mentioned effects (i.e., stabilization of the
LUMO and destabilization of the HOMO) are concomitant when going from
Ln(H5L)]5+
to [LnL(OH)]-. Not surprisingly these species are computed to have the
HOMO→LUMO gap (1.935eV) and a significant red shift of the second transition
(37 nm and 35 nm) is consistently computed particularly going from Tb(H4L)]4+
to [TbL(OH)]- and
Er(H4L)]
4+ to [ErL(OH)]
-, respectively. The first excited state
164
Figure 3.44: Calculated energy gapes and surfaces of frontier molecular orbitals for the
ground state of Ln[TAME5OX] complexes of different protonated, neutral and hydroxo-
species, respectively, from ZINDO/S-CIS method
basically corresponds to a HOMO-3→LUMO+1 transition, and it is the only
intense transition in this spectral region for all compounds but EuL and [EuL(OH)]-
that show a second transition of relatively the same intensity at 346 nm and 352 nm,
165
and also ErL and [Er(L2)] show another transition at 355 nm and 400 nm,
respectively. This later corresponds, in both cases, HOMO to LUMO+1 excitation,
the LUMO+1 being a * orbital delocalized over all quinoline rings with negligible
contribution from the acidic functions.
All of these results precisely concede the hypothesis that the influence of the
lanthanide ion on the absorption and electronic properties (and consequently on the
ground state geometries) of the complexes is a small effect. A further important
point is that taking into account, at least partially; the core interaction between the
metal ion (+3 point charge) and its neighbour atoms appreciably improves the
quality of the theoretical results.
3.2.3.2.2 Luminescence Studies:
Before the pH dependent luminescence properties of the lanthanide complexes
of TAME5OX can be expounded, an insight on the pH dependent fluorescence
properties of the ligand alone, in which 8HQ moiety is utilized as fluorescent
receptor unit is required to be addressed. It is interesting to note that the fluorescent
intensity of TAME5OX changes in a monotonous way depending only on the
change of pH value, which gets increased or decreased along with the change
(either increase or decrease) in the pH value in a consistent way due to the only one
photoinduced electron transfer mechanism between pyridyl –N and OH group of
each quinoline unit of arm, resulting in the OFFON or ONOFF type of
fluorescent pH sensor.
Some reports have been documented on the pH sensors whose fluorescent
intensity changes in acidic and basic systems along with decrease and increase in
the pH value, respectively, based on the opposite photoinduced electron transfer
processes between a receptor and signaling unit [3].
166
Table 3.9: Computed Optical Properties of La(TAME5OX), Er(TAME5OX),
Eu(TAME5OX) and Tb(TAME5OX):: Absorption-S0 and Emission-S1;
Wavelengths, Oscillator Strengths (fcalc), Electronic Transitions and Energies
Calculated at the CIS-ZINDO/S Method, Using Sparkle/PM7 Ground State and
Excited State Geometry of the Complexes.
S0 S1
Complexes nm
(abs)
fcalc Nature Transition
(Excitation energy) nm(em) fcalc Nature Transition
(Emission energy)
La(H4L) 275.24,
580.37
0.1034
0.0512
HOMO-3→LUMO+2 HOMO→LUMO+1
125→131 (1.054 eV) 128→130 (0.422 eV)
430.39 0.0691 HOMO→LUMO 128→129 (0.1249 eV)
La(H3L) 278.38,
570.41
0.0953
0.0452
HOMO-2→LUMO+1 HOMO→LUMO
126→130 (0.514 eV) 126→129 (0.473 eV)
445.13,
497.76
0.0942
0.4034
HOMO-1→LUMO+1
HOMO-2→LUMO+1
127→130 (0.1884 eV)
126→130 (0.7069 eV)
La(H2L) 285.84,
545.64
0.0722
0.0323
HOMO-3→LUMO+1 HOMO→LUMO
125→130 (0.711 eV) 128→131 (0.331 eV)
446.27,
498.43
0.0994
0.4759
HOMO-1→LUMO+1
HOMO-2→LUMO+2
127→130 (0.615 eV)
126→131 (0.8417 eV)
La(HL) 305.83,
375.79
0.0534
0.0282
HOMO-2→LUMO+1 HOMO→LUMO
126→130 (0.650 eV) 128→129 (0.213 eV)
448.42
500.54
0.1724
0.6532
HOMO→LUMO
HOMO-2→LUMO+3
128→129 (0.1448eV)
126→132 (1.2064 eV)
LaL 330.65,
385.72
0.0260
0.0062
HOMO-1→LUMO+2 HOMO-1→LUMO
127→ 131 (0.252 eV) 127→129 (0.164 eV)
502.95 0.2243 HOMO-3→LUMO+3 125→132 (1.3626 eV)
LaL(OH) 300.54,
350.74
0.0284
0.0079
HOMO-1→LUMO+1 HOMO→LUMO
127→130 (0.311 eV) 126→129 (0.173 eV)
443.52 0.2615 HOMO→LUMO+1 128→130 (0.6453 eV)
Er(H4L) 479.32,
275.91
0.0314
2.1713
HOMO→LUMO HOMO-3→LUMO+2
128→129 (0.2332 eV) 125→131 (2.0222 eV)
425.39 0.0791 HOMO→LUMO 128→129 (0.1349 eV)
Er(H3L) 486.01,
283.14
0.0298
1.8712
HOMO→LUMO HOMO-2→LUMO+2
128→129 (0.2204 eV) 126→131 (1.8235 eV)
430.13,
490.76
0.1042
0.4134
HOMO-1→LUMO+1
HOMO-2→LUMO+1
127→130 (0.1984 eV)
126→130 (0.8169 eV)
Er(H2L) 248.08,
434.13
0.0263
1.5338
HOMO→LUMO HOMO-2→LUMO+1
128→129 (0.2032 eV) 126→130 (1.7710 eV)
446.27,
499.43
0.1094
0.4859
HOMO-1→LUMO+1
HOMO-2→LUMO+2
127→130 (0.725 eV)
126→131 (0.9517 eV)
Er(HL) 255.31,
282.23,
330.07
0.0219
1.3207
0.2225
HOMO→LUMO HOMO-2→LUMO+1 HOMO-1→LUMO
129→130 (0.1145 eV) 126→130 (1.1350 eV) 128→130 (0.2273 eV)
444.42
500.54
0.1824
0.6632
HOMO→LUMO
HOMO-2→LUMO+3
128→129 (0.2548eV)
126→132 (1.3164 eV)
ErL 278.83,
365.15,
492.23
0.0189
1.0821
0.1924
HOMO-1→LUMO HOMO-2→LUMO+3 HOMO→LUMO+1
127→ 129 (0.0852 eV) 126→ 132 (1.0114 eV) 128→130 (0.1835 eV)
440.95
508.37
0.2343
0.8013
HOMO→LUMO+2
HOMO-3→LUMO+3
128→131 (0.4586 eV)
125→132 (1.5726 eV)
ErL2 300.46,
428.51,
515.50
0.0162
1.1085
0.1572
HOMO→LUMO HOMO-3→LUMO+2 HOMO-1→LUMO
128→129 (0.0626 eV) 125→131 (1.0234 eV) 127→129 (0.1478 eV)
455.52 0.2715 HOMO→LUMO+1 128→130 (0.6553 eV)
ErL(OH) 310.74,
372.12
0.0151
0.7235
HOMO-1→LUMO+1 HOMO-2→LUMO+1
127→130 (0.0326 eV) 126→130 (0.8529 eV)
431.53 0.2142 HOMO→LUMO 128→129 (0.4184 eV)
Eu(H5L) 240.91,
352.65,
480.35
0.2141
0.1023
0.0214
HOMO-3→LUMO+2 HOMO-1→LUMO HOMO→LUMO
125→131 (2.275 eV) 127→129 (1.332 eV) 128→129 (0.287eV)
455.32 0.2553 HOMO→LUMO+1 128→130 (0.5723eV)
Eu(H4L) 243.15,
351.98,
479.71
0.2034
0.1012
0.0201
HOMO-3→LUMO+2 HOMO→LUMO+1 HOMO→LUMO
125→131 (2.058 eV) 128→130 (1.124 eV) 128→129 (0.215 eV)
435.39,
480.75
0.0691
0.2567
HOMO→LUMO
HOMO→LUMO+1
128→129 (0.1249 eV)
128→130 (0.5946 eV)
Eu(H3L) 245.29,
350.62,
463.19
0.1753
0.0952
0.0163
HOMO-2→LUMO+1 HOMO-2→LUMO HOMO→LUMO
126→130 (1.817 eV) 126→129 (0.876 eV) 128→129 (0.165 eV)
445.13,
480.76
0.0942
0.4034
HOMO-1→LUMO+1
HOMO-2→LUMO+1
127→130 (0.1884 eV)
126→130 (0.8069 eV)
Eu(H2L) 247.42,
345.37,
454.61
0.1322
0.0723
0.0145
HOMO-3→LUMO+1 HOMO→LUMO+2 HOMO→LUMO
125→130 (1.413 eV) 128→131 (0.632 eV) 128→129 (0.139 eV)
446.27,
481.43
0.0994
0.4759
HOMO-1→LUMO+1
HOMO-2→LUMO+2
127→130 (0.715 eV)
126→131 (0.9417 eV)
Eu(HL) 250.92,
341.58,
441.43
0.1034
0.0532
0.0107
HOMO-2→LUMO+1 HOMO→LUMO HOMO-1→LUMO
126→130 (1.251eV) 128→129 (0.416 eV) 127→129 (0.076 eV)
451.42
479.54
0.1724
0.6532
HOMO→LUMO
HOMO-2→LUMO+3
128→129 (0.2448eV)
126→132 (1.3064 eV)
EuL 265.94,
369.86
0.0520
0.0132
HOMO-1→LUMO+2 HOMO-1→LUMO
127→ 131 (0.504 eV) 127→129 (0.327 eV)
455.95
478.37
0.2243
0.7813
HOMO→LUMO+2
HOMO-3→LUMO+3
128→131 (0.4486 eV)
125→132 (1.5626 eV)
EuL(OH) 239.26,
338.37
0.0574
0.0139
HOMO-1→LUMO+1 HOMO-2→LUMO
127→130 (0.623 eV) 126→129 (0.345 eV)
448.52 0.2615 HOMO→LUMO+1 128→130 (0.6453 eV)
Tb(H5L) 245.59,
272.30,
319.87
0.0273
2.5504
0.3106
HOMO→LUMO HOMO-3→LUMO+2 HOMO-1→LUMO
128→129 (0.2321 eV) 125→131 (2.5754 eV) 127→128 (0.3871 eV)
455.32 0.2553 HOMO→LUMO+1 128→130 (0.5723eV)
Tb(H4L) 245.64,
272.91,
320.02
0.0214
2.1613
0.3012
HOMO→LUMO HOMO-3→LUMO+2 HOMO→LUMO+1
128→129 (0.2232 eV) 125→131 (2.0122 eV) 128→130 (0.3564 eV)
435.39,
480.75
0.0691
0.2567
HOMO→LUMO
HOMO→LUMO+1
128→129 (0.1249 eV)
128→130 (0.5946 eV)
Tb(H3L) 246.01,
273.14,
320.78
0.0198
1.8612
0.2815
HOMO→LUMO HOMO-2→LUMO+2 HOMO-2→LUMO
128→129 (0.2104 eV) 126→131 (1.8135 eV) 126→129 (0.3247eV)
445.13,
480.76
0.0942
0.4034
HOMO-1→LUMO+1
HOMO-2→LUMO+1
127→130 (0.1884 eV)
126→130 (0.8069 eV)
Tb(H2L) 248.08,
274.13,
321.05
0.0163
1.5238
0.2461
HOMO→LUMO HOMO-2→LUMO+1 HOMO→LUMO+1
128→129 (0.1932 eV) 126→130 (1.6710 eV) 128→130 (0.2869 eV)
446.27,
481.43
0.0994
0.4759
HOMO-1→LUMO+1
HOMO-2→LUMO+2
127→130 (0.715 eV)
126→131 (0.9417 eV)
Tb(HL) 255.31,
282.23,
330.07
0.0119
1.3107
0.2125
HOMO→LUMO HOMO-2→LUMO+1 HOMO-1→LUMO
129→130 (0.1045 eV) 126→130 (1.1250 eV) 128→130 (0.2173 eV)
451.42
479.54
0.1724
0.6532
HOMO→LUMO
HOMO-2→LUMO+3
128→129 (0.2448eV)
126→132 (1.3064 eV)
TbL 258.83,
285.15
332.23
0.0089
1.0721
0.1824
HOMO-1→LUMO HOMO-2→LUMO+3 HOMO→LUMO+1
127→ 129 (0.0752 eV) 126→ 132 (1.0014 eV) 128→130 (0.1735 eV)
455.95
478.37
0.2243
0.7813
HOMO→LUMO+2
HOMO-3→LUMO+3
128→131 (0.4486 eV)
125→132 (1.5626 eV)
TbL2 260.46,
288.51,
335.50
0.0062
1.0985
0.1472
HOMO→LUMO HOMO-3→LUMO+2 HOMO-1→LUMO
128→129 (0.0526 eV) 125→131 (1.0134 eV) 127→129 (0.1378 eV)
448.52 0.2615 HOMO→LUMO+1 128→130 (0.6453 eV)
TbL(OH) 259.74,
286.12,
0.0051
0.7135
HOMO-1→LUMO+1 HOMO-2→LUMO+1
127→130 (0.0226 eV) 126→130 (0.8429 eV)
410.53 0.2042 HOMO→LUMO 128→129 (0.4084 eV)
To scrutinize the pH-dependent optical properties of above said Ln complexes
of TAME5OX with unique pyridyl –N and –OH structural characteristics, the pH
167
vs. fluorescence titration experiments were carried out in aqueous medium with an
appropriate volume of 0.1M HCl and 0.1M NaOH used for protonation of pyridyl
N-group and deprotonation of the OH-group, respectively (discussed in
complexation section). As shown in Figures 3.37, the electronic absorption spectra
of complexes of TAME5OX with each lanthanide ion exhibit the exemplary
behaviour (red and hypochromic shifts) along with increasing the pH value of the
system. In good contrast, the changes in the blue–green emission maxima located at
425 nm (for La and Er), with excitation wavelength at 345 nm and 385 nm, and at
410 nm and 425 nm (for Eu and Tb), which were excited at 325 nm and 380 nm,
were monitored at pH of aqueous solution of these metal ions with TAME5OX in
1:1 ratio, ranging from 1.93-10.88 pH (Figure 3.38) at room temperature. The
emission intensity at 425 nm, 410 nm and 425 nm for La, Er, Eu and Tb with meager
quantum yield as well as maximum wavelength remains almost unchanged under
acidic conditions within the pH range of 1.93.0. Most interestingly, with an
increase in the pH from 3.0, the evolution of emission peak at about 474 nm, 463
nm, 475 nm and 487 nm, for La, Eu, Tb and Er complexes, respectively, result
along with no change in blue–green emission at 410 nm and 425 nm. The maximum
wavelength of the emission at 463 nm and 475 nm get red shifted upto 490-508 nm
and 510 nm with the enhancement of fluorescence intensity upto 7-11 and 10-15
folds, along with an increase in the pH upto 7.5 and 7.4, respectively (Figure 3.38
bold lines). The intense fluorescence of all complexes gradually diminished at 500
nm upto 5-8 and 7-12 folds after 9.4 to 9.5 and 7.4 to 9.6 pH (Figure 3.38 dashed
lines) get blue shifted towards 465 nm, 460 nm and 470 nm in case of La, Eu (Er)
and Tb, respectively. With further increase in pH above 9.6, little growth of
emission in case of La(TAME5OX) and Eu(TAME5OX) whereas significant
168
enhancement of blue–green emission of Tb(TAME5OX) was observed, which
shows insignificant change under basic conditions (pH 10.0-10.9) in a similar
manner as observed under acidic conditions. Similarly, considerable growth of
emission at 438 nm of Er(TAME5OX) was observed upto 9-folds, which
intermittently get quenched upon further increase in pH and also display
insignificant change under basic conditions. It is worth noting that more significant
changes under physiological pH and less significant changes under acidic and basic
medium confirm that protonation of N atom lowers the LUMO energy and
decreases the vertical S1→S0 transition energy whereas deprotonation of the
hydroxyl functions destabilize the HOMO, thus closing the gap and giving rise to
smaller transition; while in neutral form the HOMO-LUMO gap increases which
corresponds significant red shifts along with quenching of intense fluorescence.
Thus, among all the different species formed during the pH range of 1.910.8,
[LnL] was found to be more fluorescent which was formed under physiological pH
in case of each Ln complex while the protonated and hydroxo complexes were
found to be less fluorescent.
These results reveal the typical characteristics of a pH-fluorescent probe
through the photoinduced proton transfer (PPT) enhancement process and
photoinduced electron transfer (PET) quenching processes [68] between receptor
and fluorescent signaling group for this tripodal sensor. In other words, these two
processes between the TAME5OX receptor and metal ions are responsible for the
particular fluorescent properties along with the pH range of 1.9 to 10.8. The unique
pH-dependent fluorescent property of these complexes indicates the potential
application, as pH indicators under both acidic and basic conditions and will act as
novel OFF-ON-OFF type of fluorescent pH sensors.
169
3.2.3.2.2.1 pH-Dependent PET and PPT Mechanisms
For the intention of understanding the remarkable pH-dependent fluorescence
behaviour of these Ln complexes of TAME5OX, the PET quenching and PPT
enhancement mechanisms may be followed as discussed in the ligand, TAME5OX
in section 3.2.1. The two opposite photoinduced electron transfer processes are
responsible for the above argued unusual fluorescence characteristics. The Ln
complexes of TAME5OX exhibit similar emission behaviour upon complexation,
where the H+ ions are replaced by Ln ions along with the change in the pH value
under acidic and basic conditions.
With the endeavour of cognizance, the interesting pH-dependant
fluorescence characteristics of these Ln complexes of TAME5OX, due to above-
mentioned two processes (PPT and PET) observed during pH dependant
fluorescence are clearly validated by the theoretical calculations for excited state
sparkle/PM7 optimized structures based on CIS method using ZINDO/s were
carried out for all corresponding species. The calculated emission spectra of these
species are shown in Figure 3.45. In comparison, the computed emission
wavelengths (em) are in excellent agreement not only with experimental data,
particularly with the numerical values of 425 nm, 437 nm and 508 nm (for La and
Er) and 410 nm, 480 nm and 500 nm (for Eu and Tb) for protonated, hydroxo and
neutral complexes, respectively, but also the shape. Computed vertical S1→S0
emission energies (that is, fluorescence) reported in Table 3.9, are all red-shifted
with respect to the corresponding absorption values. These observations are in
line with the very small structural changes computed for each form when
going from the ground to the excited state. Indeed, the largest structural evolutions
are related to a modification of the hydrogen bond network at the excited state due
170
to the photoacidity or photobasicity of the different functional groups. In particular,
the NH+/N pair seems to act as a photobase, the nitrogen atoms interacting stronger
with water‟s hydrogens at the excited state than at the ground state. On the contrary,
the OH function acts as a photoacid, its hydrogen interacting stronger with water‟s
oxygen at the excited state than at the ground state. These results can justify the
fluorescence behaviour observed for these complexes in solution and appear to
represent the novel pH fluorescent sensors, with TAME5OX as receptor, in
particular with OFF-ON-OFF type nature, which proved novel and versatile
fluorescent sensors with potential applications in chemical and biological fields
La[TAME5OX] Er[TAME5OX]
Eu[TAME5OX] Tb[TAME5OX]
Figure 3.45 Emission spectra of complexes computed at the sparkle/PM7 optimized excited
state geometries. In CIS-ZINDO/S computation of the spectra, the coordinating centre has
been assumed to be a 3+ point charge. For La[TAME5OX], Er[TAME5OX],
Eu[TAME5OX] and Tb[TAME5OX] species, respectively.
171
3.3 Conclusions
The new C3 symmetric tripod, TAME5OX, derived from three bidentate 8-
hydroxyquinoline subunits with nine donor atoms (three O and six N) has been
successively developed for transition metal, main group and lanthanide ions. The
molecular edifice is expected to form less stable metal coordinate compounds than
do their 2, 3 and 7- substituted analogs, but surprisingly, it showed high complexing
ability for ferric ion as well as other targeted metal ions over a large range of pH, of
the same order of magnitude as the tris-quinolinate ferric TRENOX. The ligand
exhibits high pFe value of 31.16, high selectivity towards iron than O–Trenox and
only 1.34 pFe log unit less than oxinobactin. The tripod also leads to soluble and
thermodynamically stable Ln3+
complexes in water with high stability constants,
featuring resistance toward hydrolysis in physiological pH with pLn value of 14.6,
16.14, 19.8 and 19.48 for La, Eu, Tb and Er ions, respectively. Among major
species present at physiological pH the species LaL, EuL, TbL and ErL with
tremendously enhanced stability, show interesting photophysical properties in
ultraviolet and visible range. The three major protonated species of the 1:1 podate
present between acidic to physiological pH are: [M(H3L)], [M(H2L)] and [(MHL)]
for Fe and Cr, while a single major species [Al(H3L)] was obtained for Al complex.
The neutral complex [ML] is present at basic pH with each metal ion. This
remarkable greater stability of these complexes is incipient due to the chelation and
anchoring effects of TAME5OX arose from the topology and pre-organisation of
the ligand. This is supported by the correlation between the experimental log K and
calculated ∆Gº.
The results of photoluminescence titrations of TAME5OX with the trivalent
metal ions Al3+
and Cr3+
other than Fe3+
led different results. In contrast to Al3+
, for
172
Fe3+
and Cr3+
there was a decrease in the fluorescence emission, which might be
due to the coordination of the oxygen atoms and the uncoordinated nitrogen donor
atoms present in the ligand, activating the quenching through the PET phenomena
by energy or electron transfer. For Al3+
, reverse effect of PPT was observed which
greatly enhanced the fluorescence of the ligand. This observed change is likely to
be particularly important for the potential application of TAME5OX as a
luminescent sensor for the detection and remediation of Al3+
in surface waters,
biological fluids and these complexes may also have potential in organic light
emitting devices.
The results obtained from time dependent density functional theory (TD–DFT)
show good agreement with the available experimental determined properties. From
a mechanistic point of view, the absence of fluorescence of TAME5OX in acidic as
well as basic experimentally observed pH are and claimed to be related to the
excited-state intramolecular proton transfer (ESIPT), from 8-OH to the quinoline-N
atom are proven and validated. NBOs and frontier molecular orbital analysis for
ground as well as excited states are in good agreement with the observed trends in
different species of M(TAME5OX) for Fe, Al and Cr. Similarly, computed
absorption and emission spectra of all the pH dependent species of
Ln(TAME5OX), studied by means of interaction with single excitations (CIS)
based on ZINDO/S methodology, were found in concurrent with the
experimentally claimed data. Some insights are shared on the combined effects
of 8-hydroxyquinoline binding units of TAME5OX and Ln ions on the
geometry, electronic structure and optical properties of complexes. Our findings
also suggest that experimental spectroscopic data are better reproduced in
ZINDO calculations when the lanthanide is represented by a point charge; these
173
methods also enabled the calculation of luminescent properties. The results of
these methods not only reproduced well the experimental values but also helped
in explaining the low values of quantum efficiency observed for these
complexes. The excellent photophysical properties of these species make them very
good candidates with potential applications in chemical and biological fields and
procedure followed in this investigation could provide valuable information in the
ligand design to improve this pathway.
3.4 Experimental
3.4.1 Materials and Methods
All manipulations were carried out under an atmosphere of nitrogen using standard
cannula technique. The reagents were purchased from SigmaAldrich at the highest
commercial quality, Alfa Aesar or Acros Organics, and used without further
purification unless otherwise stated. Solvents were distilled freshly, refluxed over
appropriate drying agents following standard procedures and kept under argon
atmosphere. Melting points were determined on conventional melting point
apparatus and are uncorrected. Reactions were monitored by thin layer
chromatography (TLC), performed on silica gel G plates using hexane and
ethyleacetate (7:3), chloroform and methanol (4:1) as eluents; developed TLC
plates were visualized by iodine vapours. Elemental analysis (CHN) was performed
on the elemental analyzer Euro Vector S.P.A81328 (Euro EA). Infrared spectra
were recorded using an Agilent Technologies Cary 630 FTIR spectrometer
equipped with a universal attenuated total reflection (ATR) sampler and samples
were directly used for obtaining IR spectrum. 1H and
13C NMR spectra were
recorded in DMSO–d6 and D2O, using a Joel 300 (MHz) ECX spectrometer.
Chemical shifts are reported in values related to an internal reference of
tetramethylsilane (TMS). Mass analysis was performed in a TOF MS ES+
174
instrument. For solubility reason the ligand was converted to its hydrochloride salt
and potentiometric and spectrophotometric studies were carried out in water
purified by a Millipore MilliQ water purification system. Ionic strength was
adjusted with 0.1M KCl. All the stock solutions were prepared by weighing
appropriate amounts by using GR202 electronic balance (precision 0.01 mg) in
Millipore grade deionized water. The exact concentration of KOH (0.1 M) was
determined potentiometrically using 0.1 M solutions of succinic acid, potassium
hydrogen phthalate (KHP) and oxalic acid as primary standards and then the exact
strength of HCl (0.1N) was determined by same method using standardized KOH.
The solutions were prepared with Millipore grade deionized water immediately
before use, which was deoxygenated and flushed continuously with Ar (U grade) to
exclude CO2 and O2. All measurements were carried out at 25.0 ± 0.2°C maintained
with the help of Julabo F25 thermostat.
3.4.2 Synthesis of TAME5OX
The synthesis of ligand was carried out by chloromethylation of the oxine 2
followed by condensation with tripodal 1,1,1 tris (aminomethyl) ethane) 1c as
depicted in Scheme 1. The synthesis of 5-chloromethyl-8-hydroxyquinoline
hydrochloride 2a was carried out with higher yields by the reported procedure [69].
The synthesis of 1,1,1 Tris (aminomethyl)ethane (TAME) 1c was carried out under
multistep reactions by the reported procedure [33]. The detailes of procedure for
synthesis of 1c and 2a are presented below.
3.4.2.1 Synthesis of 1,1,1 Tris (aminomethyl)ethane (TAME)
The synthesis of 1,1,1 Tris (aminomethyl)ethane (TAME) 1c was carried out under
multistep reactions as detailed below:
3.4.2.1.1 tris[(4-tolylsulfonyl)methyl]ethane (1a)
175
Tris(hydroxymethyl)ethane (10 g, 0.083 mol) was dissolved in THF (45 mL)
and added to 70 ml of aqueous solution of NaOH (16.66 g, 0.41 mol). The solution
was cooled in an icewater bath and THF solution of tosyl chloride (47.66 g, 0.25
mol) was added dropwise with stirring for about 2 hours. During addition, the
solution became pale yellow while an offwhite slurry appeared progressively; the
mixture was stirred for 4 hrs further at 0°C. The reaction mixture was poured into
ice water (250 mL) and extracted with toluene (3×100 mL). The combined organic
phases were washed with water (2×100 mL), dried over anhydrous MgSO4, and
concentrated on rotary evaporator. The thick jelly residue was poured into methanol
(500 mL) with stirring. The resulting suspension was filtered on a Buchner funnel
and the precipitate was washed with water followed by methanol and diethyl ether,
and dried under vacuum. The pure product was isolated as a white solid after
recrystallization from npropanol. Yield: 87%, mp: 105ºC. FTIR (KBr): 3023,
2954, 1889, 1686, 1605, 1478, 1458, 1367, 1302, 1290, 1153, 1099, 1012, 991,
941, 882, 805, 797, 701, 677, 532, 498 cm-1
.
3.4.2.1.2 tris(azidomethyl)ethane (1b)
Solution of tris[(4-tolylsulfonyl)methyl]ethane (1a) (25.0 g, 42.76 mmol) and
sodium azide (29.58 g, 357.17 mmol) in dry DMSO (250 mL) was stirred under N2
atmosphere at 125ºC for 8 hrs. The reaction mixture was poured into cold water (1.5
L), after cooling to room temperature. The aqueous solution was extracted with
ethyl acetate (3×250 mL). The organic phases were washed with water (2×100 mL),
dried over anhydrous NaSO4 and reduced on a rotary evaporator. The residual oil
shows strong absorption band at 2123 cm-1
in the infrared spectrum which confirms
the presence of an azide and the product was used readily for the next step.
3.4.2.1.3 tris(aminomethyl)ethane (1c)
176
tris(azidomethyl)ethane (1b) (21g, 107.69 mmol) in freshly distilled THF(60
mL) was added to a vigrously stirred mixture of LiAlH4 (18.63g 374.39mmol) in
dry THF(300 mL) over a period of 3h under dry nitrogen. After complete addition,
the mixture was refluxed for 12h, cooled, followed by addition of water (22 mL),
then by NaOH (34mL of 15% aqueous solution) and again with water (50mL). The
resulting slurry was stirred for 25 min and filtered through sintered glass G4 funnel
under vaccum. The white residue was further refluxed over night with THF to
extract remaining product. The solvent was removed from all combined filtrates to
give pale yellow oil. The crude product was taken up in super dry EtOH (300mL)
and concentrated HCl (12mL) was added dropwise to precipitate the compound as
the trihydrochloride salt. The pure product was reprecipitated from an ethanol
solution to yield a fine white powder. Yield: 9.1g, 77.77mmol (72.22%),
decomposes at ≥ 250°C.
1HNMR (D2O) (δ): 3.24(s, 6H, 3CH2) 1.20 (s, 3H ,CH3).
3.4.2.2 8-hydroxyquinoline-5-methylchloride hydrochloride (2a)
Dry hydrogen chloride gas was passed through a solution of 8–hydroxyquinoline 2
(14.51g, 0.1 mol) and formaldehyde (20 mL, 40%) in 37% hydrochloric acid (50
mL) for 5 hours at 60°C. After filtration, the yellow solid product was washed
several times with acetone and dried in vaccuo to afford 2a (14.01g) in 96.5% yield.
m.p. 281ºC;
1HNMR (300 MHz, D2O, 25°C, TMS): = 8.81 (d,
3JH,H = 4.1 Hz, 1H,
Ar–H), 8.51 (d, 3JH,H = 8.5 Hz, 1H, Ar–H), 7.59 (dd,
3J H,H = 8.2 Hz,
3J H,H = 4.1 Hz,
1H, Ar–H), 7.40 (d, 3J H,H = 7.8 Hz, 1H, Ar–H), 7.01 (d,
3J H,H = 7.8 Hz, 1H, Ar–H),
4.82 (s, 2H, ArCH2Cl) ppm. IR (ATR) (cm-1
) = 3032, 2731, 1564, 1510, 1439,
1330, 1006. Elemental Analysis calcd (%) for C10H9NOCl2: 52.20 C, 3.94 H, 6.09
N; found: 52.29 C, 3.89 H, 6.13 N; MS (ES+), m/Z (I,%) = 320.07[M+1]
+(90).
3.4.2.2.1 5,5'-(2-(((8-hydroxyquinolin-5-yl)methylamino)methyl)-2-methyl
propane -1,3-diyl) bis(azanediyl)bis(methylene)diquinolin-8-ol (TAME5OX)(3).
177
To a suspension of 2a (138.28 mg, 0.6mmol) was added 1,1,1 tris (aminomethyl)
ethane 1c (23.4 mg, 0.2 mmol) in an acetonewater mixture and K2CO3 (82.92 mg,
0.6 mmol) was used as an acid acceptor. The mixture was refluxed with stirring on
magnetic stirrer for 3 hrs. The suspension gets dissolved upon reflux and the colour
changed from yellow to brown with the progress of reaction. The resulting mixture
which contained brownish green precipitate after cooling at room temperature, was
poured into ice cold (50 mL) water and then filtered. The solid product was
collected, washed several times with diethylether and dried in vaccuo to afford
TAME5OX 3 as a off white colour powder (0.085mg, 0.17mmol, 77%). m.p.
160°C; 1HNMR (DMSO d6, 25°C, TMS): = 1.18(s, 3H, CH3 ); 2.05(s, 6H, CH2);
2.4 (s, 6H, CH2); 4.61 (s, 3H, NH); 8.42(d, 3J H,H = 6.8 Hz, 3H, Ar–H); 8.84(d,
3J H,H
= 7.8 Hz, 3H, Ar–H); 6.91–6.94(t, 3J H,H = 7.6 Hz, 3H, Ar–H); 7.49-7.55(dd,
3J H,H =
8.0 Hz, 6H, Ar–H); 9.74(s, 1H, OH) ppm. 13
CNMR (DMSO d6): δ = 37.01 (CH3),
57.45 (CH2), 62.66 (CH2), 111.44, 117.88, 121.93, 127.57; 135.97, 138.34, 148.87,
152.51 (C=N), 153.58 (C–O). IR (ATR) (cm-1
) = 3470, 1594, 1454, 1396, 1239,
1172, 1080. Elemental analysis Calcd (%) for C35H36N6O3: 71.41 C; 6.16 H; 14.28
N; found: 71.66 C; 6.21 H; 14.35.N. MS (ES+), m/Z (I,%) 598.2[M+3]
+(20).
3.4.3 Synthesis of metal complexes of TAME5OX
The synthesis of metal complexes was carried out maintaing ligand-to-metal molar
ratio 1:1. Appropriate amount of FeCl3, CrCl3 or Al(NO3)3 salt (0.1mmol) in dry
ethanol (5mL) was added dropwise under N2 atmosphere to the stirred solution of
the TAME5OX (63.11 mg, 0.1 mmol) in 15 mL ethanol. The reaction mixture
turned cloudy with the addition of the metal. Two drops of triethylamine was then
added to complete the precipitation. After stirring at room temperature for
overnight, the total volume was reduced to 5mL on rotary evaporator. The product
178
was filtered off, washed with cold absolute ethanol followed by dry diethyl ether,
and dried in vaccuo.
Fe(TAME5OX): Colour: black, yield: 60.3 mg, 0.079 mmol, 85.9%, m.p > 300°C.
IR (ATR) (cm-1
) =2932, 1612, 1432, 1345, 1136, 1085, 991, 603, 461. Elemental
analysis Calcd (%) for C35N6O3H30Fe: 65.23 C, 4.69 H, 6.52 N; found 64.95 C, 5.01
H, 6.71 N. MS (ES+), m/Z (I,%) 653.0[M+1]
+(10).
Al(TAME5OX): Colour: green, yield: 59.7 mg, 0.078 mmol, 79.9%, m.p > 277°C).
IR (ATR) (cm-1
) = 2930, 1668, 1543, 1456, 1392, 1223, 1082, 965, 638, 479, 417.
Elemental analysis Calcd (%) for C35N6O3H30Al: 68.29 C, 4.91 H, 6.83 N; found
68.03 C, 4.83 H, 6.92 N. MS (ES+), m/Z (I,%) 624.0[M+2]
+(20).
Cr (TAME5OX): Colour: black, yield: 52.3 mg, 0.065 mmol, 76.4%, m.p >
260°C). IR (ATR) (cm-1
) = 2927, 1652, 1510, 1437, 1386, 1209, 1103, 603, 461,
403. Elemental analysis Calcd (%) for C35N6O3H30Cr: 65.62 C, 4.72 H, 6.56 N;
found 65.09 C, 5.07 H, 6.64 N. MS (ES+), m/Z (I,%) 649.0[M+1]
+(15).
3.4.4 Solution Thermodynamics: Potentiometric, Spectrophotometric and
Fluorescence measurements
3.4.4.1 Potentiometric Measurements
Potentiometric titrations were carried out for the determination of the protonation of
the ligand and formation constants of the metal complexes; the double wall glass
jacketed titration cell connected to a constant temperature circulatory bath was used
to maintain temperature at 25 ± 1°C. The pH measurements were performed using
Thermo Scientific ORION STAR–A211 pH meter equipped with combined Ross
Ultra pH/ATC glass electrode and the observed pH was measured as -log [H+]. The
titrations were employed with borosilicate glass burette (1 ml, least count 0.01 ml).
A standard KOH (0.095M) solution was used to titrate a standard hydrochloric acid
(0.1045 M) solution and the pH-meter readings were converted to hydrogen ion
179
concentration by calculated hydrogen ion concentrations (pKw = 13.77). The
hydrochloride salt of ligand (5 ml, 1 × 10-4
M) and mixtures of 1:1 L: M+3
(1×10-4
M) (M= Fe3+
, Al3+
, Cr3+
, La3+
, Eu3+
, Tb3+
and Er3+
) were titrated with standardized
0.095 M KOH for determination of protonation and formation constants,
respectively. Final concentration of ligand (1 × 10-5
M) and metals (1 × 10-5
M) were
maintained for the different titrations. The solution was acidified to a pH of 1.98
with standardized HCl (5 ml, 0.1045 M) and the ionic strength was fixed at 0.1 M
with KCl were monitored as a function of pH over the range of 1.98-11.5. After
each addition of KOH and adjustment of pH, it was ensured that complete
equilibrium was attained (ca. 5 minutes) before measurement of pH. The titration
data (pH range 2.1−11.0, 158 points) were refined by the nonlinear least-squares
refinement program HYPERQUAD to determine equilibrium constants (protonation
and complexation), while the distribution of species were plotted with the program
HYSS 2009 [20].
3.4.4.2 Spectrophotometric Measurements
Electronic absorption spectra were recorded on an Agilent-8453 Diode Array
Spectrophotometer using 1.0 cm path length Hellma quartz cell and externally
connected to a HEWLETT PACKARD Vectra P3327G computer. All
spectrophometric titrations were performed in a thermostated (25.0 ± 0.1°C) glass-
jacked vessel. Spectrophotometric titrations of a 50mL solution of stoichiometric
1:1 quantities of TAME5OX (5×10-5
M) and M (M = Fe+3
, Al
+3 and Cr
+3) in
aqueous medium and acidified with approximately 5mL (0.1M) HCl, ionic strength
maintained at 0.1 M with KCl were monitored as a function of pH over the range
1.6-10.8 by titrating with freshly prepared standardized (0.1M) KOH solution.
Similarly, Spectrophotometric titrations were carried out by varying the pH in
180
aqueous medium, in a typical experiment, for monitoring the UV-visible spectra in
the titration of stoichiometric 1:0 and 1:1 quantities of TAME5OX with metals
Ln(NO3).nH2O{5 ml (5 × 10-5
M)} (Ln = La, Eu, Tb and Er; n= 5)., acidified with
approximately 5 ml (0.1045M) HCl, ionic strength maintained at 0.1 M with KCl
were monitored as a function of pH over the range 1.7-12.5 by titrating with
freshly prepared standardized (0.095 M) NaOH solution. After each addition of
standardized base and adjustment of pH, it was ensured that complete equilibrium
was attained before measurement of the pH and recording of the absorption
spectrum. The equilibrium and the formation constants from the spectral data were
calculated by using a non-linear least-square fitting program, HYPSPEC.
3.4.4.3 Fluorescence Measurements
Fluorescence measurements were carried out on a Perkin Elmer LS-55
luminescence spectrophotometer equipped with quartz cuvettes with a path length
of 10 mm at 25.0 ± 0.1°C. Fluorescence data were recorded from two sets of
solutions. The first set of titrations was recorded from same sets of solutions as used
for absorption spectra. The second set for the formation of the luminescent metal
complexes of TAME5OX was ascertained by luminescence titrations of a 1.5 mL
aqueous solution of the TAME5OX (1×10-5
M) with (1×10-5
M) Fe+3
, Cr+3
, or (1×10-
6M) Al
+3 (0 to 1 or above 1 equivalent) at physiological pH with micropipette in the
cell. But for lanthanide complexes only first method was used, as no significant
variations were observed by incremental titration of metal ions. Fluorescence
spectra were registered with excitation at 385 nm (for ligand and Fe complex), 370
nm and 345 nm (for Al and Cr), 335 nm and 390 nm (for Eu and Tb) and at 332 nm
and 393 nm (for La and Er) complexes, respectively, and all excitation and emission
slit widths at 5.0 nm unless otherwise indicated. In order to allow correlation of
181
emission intensities, adjustment for instrumental response, inner filter effects, and
phototube sensitivity were implemented. The linearity of the fluorescence emission
vs. concentration was checked in the concentration range used (10-5
–10-6
M). A
correction for wavelength response of the system was performed when necessary.
For the formation of the luminescent metal complexes of TAME5OX, titrations
were ascertained as after each addition of base, the pH of the solution was measured
using combined Ross Ultra pH/ATC glass electrode and emission spectra was
recorded. The formation constants from the spectral data were calculated by using a
non-linear least-square fitting program, HYPSPEC [20]. Quantum yields were
determined using the relationship:
∅ = ∅𝑟𝑒𝑓
𝐼 𝐴
𝐼𝑟𝑒𝑓𝐴𝑟𝑒𝑓
𝜂
𝜂𝑟𝑒𝑓
2
where Φ is the radiative quantum yield of the sample, Φref is the known quantum
yield of quinine sulphate in 1M aqueous H2SO4 (=0.546), A is the absorbance at the
excitation wavelength, I is the integrated emission, and η is the refractive index of
the solvent, which is assumed to be the same for the solutions of sample and
reference.
3.4.4.4 Theoretical Calculations
All computations were carried out on a Pentium IV 3.2 GHz machine with a Linux
operating system. Starting structures for optimization were manually drawn using
ChemDraw or ISIS Draw. The initial geometry optimization of the TAME5OX, its
different protonated and deprotonated species obtained in solution and its above
mentioned metal complexes leading to minimum strain energy were achieved
through molecular mechanics calculation using MM force field by Gabedit
version
2.4.6 [27]. The periodically search to global minimum energy conformer of the
182
ligand and its metal complex were achieved using molecular dynamics simulation
upto 2000K followed by the geometrical optimization calculations. A bath
relaxation time of 0.1ps and a step size of 0.001ps were used for dynamic
simulation. Then, the minimized structures were further re–optimized through
semi–empirical method by applying PM6 self consistent fields (SCF) method, at the
Restricted Hartree-Fock (RHF) level using MOPAC2012 [70]. The geometry
optimizations were obtained by the application of the Polak–Ribiere algorithm with
convergence limit of 0.0001 kcal/mol and RMS gradient of 0.001kcal/mol.
3.4.4.4.1 DFT Calculations
The PM6 optimized molecular structures were then used for DFT level calculations.
All the calculations including geometry optimization, harmonic vibrational
frequency analysis, NBO analysis and ground state electronic transitions were
performed by employing the B3LYP (Becke three parameter Lee-Yang-Parr) [71]
exchange-correlation functional with the all-electron 6-31G* basis set except for
Cr(TAME5OX) where LACVP* was used. All possible protonated and
deprotonated states of ligand (TAME5OX) described experimentally [(LH9)+6
,
(LH8)+5
, (LH7)+4
, (LH6)+3
, (LH5)+2
, (LH4)+1
, (LH3), (LH2)-1
, (LH)-2
and (L)-3
] were
also optimized at DFT level. Similarly, protonated and neutral metal complexes of
ligand were also optimized in the gas phase at DFT level. To each optimized
molecule, the Hessian was calculated for the resultant stationary points and all were
characterized as true minima (i.e., no imaginary frequencies). NBO analysis has
been performed to examine the bonds forming molecule using NBO5. The
calculated vibrational wavenumbers and chemical shifts were compared with the
experimental data of M(TAME5OX). The cartesian (x, y, z) coordinates of
minimized structures were evaluated. Dispersion corrected DFT calculations were
carried out for all the molecules using B3LYP-D3 functional [72] with zero-
183
damping, using coordinates from the file (LHn)n-3
. xyz [n= 0-9]. 1H and
13C NMR
chemical shifts were calculated from the optimized molecule, by the
GuageIncluding Atomic Orbital (GIAO). All relative chemical shifts () are given
with respect to the absolute shielding values (ζ) obtained at the same computational
level ( = (ζref-ζ). The calculated vibrational wavenumbers and chemical shifts
were compared with the experimental data of TAME5OX. To obtain the electronic
transition energies (which include some account of electron correlation) in ground
state for various acid-base species of TAME5OX and M(TAME5OX), Time-
Dependent Density-Functional Theory (TD-DFT) with same functional and basis
sets as for geometry optimizations, were used.
The excited states calculations were computed at the TD-DFT level, the
camB3LYP density functional [73] with the all electron Pople double-ζ basis set
including one diffusion functions on heavier atoms 6-31G*. The lowlying excited-
state structure (S1) was also optimized using the cam-B3LYP. This functional
makes use of the Coulomb attenuation method and has been specifically
parameterized with electron excitation processes in mind. To obtain properties of
the first excited state (S1) such as structure, free energy, absorption and emission
energies, three excited states were always included in the calculations. Emission
wavelengths were evaluated on the TD-DFT (cam-B3LYP) optimized structures of
the excited state. All of the computational calculations have been performed using
the Gaussian 09 [40] being used for spectra convolution and molecular orbital plots.
3.4.4.4.2 Semiempirical Sparkle Model Calculations
The MM optimized molecular structures were used for sparkle mode calculations
for lanthanides. Parametric method number 7, PM7, is one of the latest in a series of
semiempirical methods which encompass MNDO, AM1, PM3 and RM1 was used.
184
3.4.4.4.2.1 Co-ordination Scan: The coordination scan technique was used to
determine the coordination number of Ln(III) in the various complexes studied. The
technique utilized in these studies is the calculation of the complex strain energy as
a function of varying coordination numbers by covalently binding the appropriate
number of waters to the lanthanide metal ion [74]. The preferred coordination
number 9 was observed for Ln[TAME5OX] complexes calculated by the
coordination scan technique with the stepwise addition of three or more water
molecules and calculating strain energy of the molecules using SYBYL. It
calculates water as having strain energy of 0.00 kcal/mol; thus the waters added to
the complex add no energy other than steric interactions with the ligand. This is
accomplished by modeling the complex with a generic metal while varying its ionic
radius. The resultant curves give a minimum energy which corresponds to the best
fitting metal ion radius. The Ln(III) ionic radius was effectively varied by
systematically altering the Ln-N and Ln-O equilibrium bond lengths in the
following manner. The Ln-N bond length was assigned through the following
relationship: equilibrium bond length = Ln ionic radius + 1.7 Å. Similarly the Ln-O
equilibrium bond lengths were assigned using: equilibrium bond length = Ln ionic
radius + 1.4 Å. The force constant for the bond stretching was kept at a constant
value of 500 kcal mol-1
Å-1
. The complex was then minimized and the energy of the
complex found. The ionic radius was then increased by 0.1 Å, the equilibrium bond
length modified to the new value and the complex again minimized. This process is
continued until the Ln ionic radius has reached 1.5 Å; this range of 0.3-1.5 Å is
sufficiently large that it covers the radii of all the possible coordination states [74].
This procedure is then repeated for the complex with one water, two waters, three
waters and so on until all the possible coordination states have been examined. Plots
185
of the complex energy versus metal ionic radius were then generated for each
coordination number.
3.4.4.4.2.2 The Ground State Geometries Calculation: Full geometry optimizations of
all the molecular systems were carried out without symmetry constraint. Frequency
calculations and dispersion corrections were performed, and the minima on the
potential-energy surfaces of the reported structures were characterized by the
absence of negative eigen values in the Hessian matrix. The ground state geometries
of [EuTAME5OX(H2O)3], [TbTAME5OX (H2O)3], [LaTAME5OX (H2O)3] and
[ErTAME5OX (H2O)3] were calculated with the sparkle/PM7 [64] model
implemented in the MOPAC2012 software [70] package. The key words used were
SPARKLE, PM7-DH2, PRECISE, FORCE, BFGS, GNORM = 0.25, SCFCRT =
1.D-10 (to increase the SCF convergence criterion) and XYZ (for Cartesian
coordinates).
3.4.4.4.2.3 Excited States Energies and Absorption Spectra: Absorption spectra of
lanthanide complexes was calculated from sparkle model optimized geometries,
followed by Orca calculations, in which the lanthanide ion is replaced by a +3e
point charge and the "spin multiplicity" is set to 1 (singlet). The singlet and triplet
excited states of all the calculated ground state geometries have been predicted
using the configuration interaction with single excitations (CIS) based on the
Zerner's intermediate neglect of differential overlap/spectroscopic (ZINDO/S)
methodology [75], using a point charge of +3e to represent the trivalent lanthanide
ion. The 28 nroots were chosen for excited state calculation with maxdim set value
to at least equal to five times the value of nroots. The CIS space was gradually
increased until there were no further meaningful changes in the calculated triplet
energies and absorption spectra. A Lorentzian line shape was fitted to the calculated
186
singlet transitions, together with the relative intensities obtained from oscillator
strengths and all the simulated spectra had a half-height band width of 20 – 30 nm,
which properly allows the comparison between the calculated and the observed
absorption electronic spectra.
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