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

80

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

81

(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.

82

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

83

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.

84

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.

85

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

86

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.

87

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

88

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.

93

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),

94

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.

95

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.

97

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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