Journal of Molecular Structure 1050 (2013) 43–52

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Journal of Molecular Structure

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Spectral studies and thermal analysis of new vanadium complexes ofethanolamine and related compounds

0022-2860/$ - see front matter � 2013 Elsevier B.V. All rights reserved.http://dx.doi.org/10.1016/j.molstruc.2013.07.020

⇑ Corresponding author.E-mail address: hmaahmed@yahoo.co.uk (H.M. Ahmed).

Mamdouh S. Masoud a, Alaa E. Ali b, Hytham M. Ahmed c,⇑, Essam A. Mohamed b

a Chemistry Department, Faculty of Science, Alexandria University, Alexandria, Egyptb Chemistry Department, Faculty of Science, Damanhour University, Damanhour, Egyptc Pharmaceutical Analysis Department, Faculty of Pharmacy, Damanhour University, Damanhour, Egypt

h i g h l i g h t s

� Newly synthesized complexes of VIII,VIV and VO2+ with ethanolamineswere described.� The complexes have been character-

ized by elemental analyses, magneticmoment measurements, IR and UV–Vis spectroscopy.� Absorption spectra in seven different

solvents were recorded.� The solvatochromism was examined

and discussed.� Some complexes were studied by

thermal analysis using DTA and TGtechniques.

g r a p h i c a l a b s t r a c t

a r t i c l e i n f o

Article history:Received 31 May 2013Received in revised form 13 July 2013Accepted 15 July 2013Available online 17 July 2013

Keywords:VanadiumEthanolamineTriethanolamineComplexesThermal analysisSolvatochromism

a b s t r a c t

The electronic absorption spectral behaviors of newly synthesized complexes of VIII, VIV and VO2+ withEthanolamine, Diethanolamine and Triethanolamine were described. The complexes have been charac-terized by elemental analyses, magnetic moment measurements, IR and UV–Vis spectroscopy. Absorptionspectra in seven different solvents were recorded. The solvatochromism was examined and discussed.Dipolar interactions between the solvent and the complexes were used to correlate the observed spectralshifts to solvent polarity. Some of the obtained complexes were studied by thermal analysis using DTAand TG techniques.

� 2013 Elsevier B.V. All rights reserved.

1. Introduction

Vanadium is an early first-row transition metal and forms col-ored compounds in its many different oxidation states dependingon the electronic and steric nature of the coordinating ligands. Inhigher oxidation states, vanadium is very oxophilic, but at low oxi-dation states, p-donating ligands such as dinitrogen and carbon

monoxide are preferred. Vanadium is a micronutrient elementwhich has been shown to produce several important biological ef-fects in living organisms [1–4]. Its deficiency causes growth retar-dation and skeletal deformations in animals. The catalytic andmaterial properties of vanadium compounds and their effects inbiological systems have long provided the impetus to studies ofvanadium science [1,5–10]. Also, the effect of vanadium on livingorganisms was studied [11–13] and vanadium complexes andpolyoxovanadate recently proved to be used as anticancer [14–19]. The bifunctional nature of the examined compounds enables

44 M.S. Masoud et al. / Journal of Molecular Structure 1050 (2013) 43–52

them to serve a variety of commercial applications such as corro-sion inhibitors, surfactants, gas purification, and herbicides andin the manufacture of cosmetics and pharmaceuticals [20]. Theycan also act as multidentate ligands through amino, hydroxyl,and deprotonated hydroxyl groups. Metal with the tested ligandscomplexes have been the subject of many studies arising mainlyfrom their technical applications [21–23]. As a part of our ongoingresearch on the synthesis, spectral, thermal and structural analysisof Ethanolamine (MEA), Diethanolamine (DEA) and Triethanola-mine (TEA) complexes with most classes of metals [24–33]. Thispaper explores the physicochemical properties of complexes de-rived from vanadium of different oxidation states with MEA, DEAand TEA.

2. Materials and methods

All chemicals were of analytical grade and used without furtherpurification. The solutions were prepared in inert atmosphere(bubbling nitrogen) using double distilled water. The VCl3 solu-tions were prepared immediately before each addition. The molarpercentage of VO2+ in each solution was adjusted by direct titrationwith standard KMnO4 solution to be less than 1%. The chemicalstructure, abbreviation and symbol for the three tested compoundsare shown in Fig. 1.

2.1. Preparation of the complexes

2.1.1. Synthesis of VIII complexesAnhydrous VCl3 (6.4 mmol) was dissolved in 10 ml hot N,N-

dimethylformamide (DMF) under a stream of nitrogen, then(12.8 mmol) of the tested ligand was added to the solution. Themixture was kept at 70 �C for 0.5 h with stirring. Dark greenishblue microcrystals of the produced complexes were observed.The complexes were filtered off, washed with dry diethyl etherand finally dried in a desiccator over anhydrous CaCl2. The com-plexes were insoluble in organic solvents but soluble in hot water.Therefore the stock solutions were prepared in hot water and then0.5 ml of the obtained solution was mixed with the used organicsolvent.

2.1.2. Synthesis of VIV complexesRather than using VCl4 which is a convenient starting material

due to its volatility and oxidative nature, it is preferred to useVCl3. The complexes were prepared by addition of VCl3

(0.01 mol) in 25 ml ethanol to a hot ethanol solution of the desiredinvestigated ligand (0.02 mol). The reaction mixtures were re-fluxed with stirring for 2 h. Hot solutions were cooled to room tem-perature and green microcrystalline complexes were formed. The

ethanolamine, diethanolamine, triethanolami

OHH2N

MEA

Name:Abbreviation:

Symbol:

Monoethanolamine

HL

HN

DEA

Diethanola

HL'

HO

Fig. 1. The chemical structure, abbreviation a

complexes were filtered, washed with cold ethanol and finallydried in a vacuum desiccator over anhydrous CaCl2.

2.1.3. Synthesis of VO2+ complexesThe VO2+ complexes were prepared by two techniques:By refluxing solutions of 0.01 mol of VOCl2 in 30 ml of 50%

methanol–water with 0.01 or 0.02 mol of tested compounds onhot plate equipped with a magnetic stirrer. The complexes precip-itated from the boiling solutions and then filtered off under suc-tion. The obtained crystals were washed well with 50%methanol–water solution and then dried in a vacuum desiccatorover anhydrous CaCl2.

VOCl2 þHL! VOLClþHCl VOCl2 þ 2HL! VOL2 þ 2HCl

By dropwise addition of 0.1 mol VOSO4 solution in 30% metha-nol–water solution to 0.01 or 0.02 mol of the desired ligand solu-tion in methanol–water solution. The complexes were separatedimmediately, washed, filtered then dried in a vacuum desiccatorover anhydrous CaCl2.

VOSO4 þ 2HL! VOL2 þH2SO4

2.2. Elemental analysis

2.2.1. Metal ion contentThe complexes were digested and decomposed with aquaregia.

The metal ion contents were determined by the usual complex-metric titration procedures [34] and atomic absorption technique.

2.2.2. Halogen and sulphate contentsThe halogen content was determined by burning 30 mg of the

sample in an oxygen flask in the presence of KOH–H2O2 mixtureand titrating with a standard Hg(NO3)2 solution using diphenylcarbazone as an indicator [35]. The sulphate content was deter-mined gravimetrically as BaSO4 [35] and found to be zero in allcomplexes.

2.3. Physical measurements

2.3.1. Electronic absorption spectraUltraviolet and visible spectra were recorded using Perkin El-

mer spectrophotometer model Lambda 4B covering the wave-length range 190–900 nm. The complexes were measured inNujol mull, following the method described by Lee et al. [36].

2.3.2. Infrared spectraThe IR spectra were recorded using Perkin Elmer spectropho-

tometer model 1430 covering the frequency range of 200–

ne abbreviated as MEA, DEA and TEA

OH

mine

N

OH

OH

HO

TEA

Triethanolamine

HL''

nd symbol for investigated compounds.

M.S. Masoud et al. / Journal of Molecular Structure 1050 (2013) 43–52 45

4000 cm�1. The obtained data similar to the previously publishedIR data [37].

2.3.2.1. Magnetic susceptibility measurements. Molar magnetic sus-ceptibilities, corrected for diamagnetic using Pascal’s constants,were determined at 298 K using Faraday’s method. The apparatuswas calibrated with Hg[Co(SCN)4] [38].

2.3.3. Thermal analysisThermogravimetric and differential thermal analysis performed

on a Du Pont 9900 computerized thermal analyzer. The heatingrate was 10�/min. A sample of 60 mg was placed in a platinum cru-cible. Dry nitrogen was flowed over the sample at a rate of 10 cc/min and a chamber cooling water flow rate was 10 h�1. The speedwas 5 mm/min.

2.4. Solvent effect on the electronic absorption spectra

Electronic absorption spectral measurements were done forMEA complexes by dissolving the investigated complex in theproper spectrograde solvent then the spectra were recorded usingPerkin Elmer spectrophotometer model Lambda 4B. The solvato-chromic behavior for the complexes was studied in seven differentsolvents covering a wide range of solvent parameters, namely eth-anol, methanol, DMF, dimethyl sulfoxide (DMSO), 1,4-dioxane,chloroform and water.

3. Results and discussions

Elemental analyses for the suggested molecular formulae andphysical properties of the isolated pure complexes are given in Ta-ble 1. The obtained metal chelates were stable in the solid stateand in solution under atmospheric conditions.

Table 1The molecular formulae, elemental analysis and physical properties for the isolated comp

Complex Analytical data

V C H

VIII complexes V(L)2Cl 24.66 23.26 5.86(24.65) (23.24) (5.83)

V(HL0)2Cl 17.41 32.83 6.20(17.39) (32.81) (6.17)

V(HL’’)Cl 21.94 30.85 5.61(21.93) (30.84) (5.60)

VIV complexes V(L)2Cl2 21.05 19.85 5.00(21.03) (19.85) (4.97)

V(HL0)2Cl2 15.53 29.29 5.53(15.53) (29.28) (5.52)

V(HL00)2Cl2 12.24 34.63 6.30(12.23) (34.63) (6.31)

VO2+ complexes VO(L)Cl 32.15 15.16 1.27(32.14) (15.16) (1.27)

VO(HL0)Cl (24.67) (23.27) (4.87)24.67 23.26 4.88

VO(H2L00)Cl 20.33 28.76 5.63(20.31) (28.75) (5.61)

VOL2 27.23 25.68 6.46(27.22) (25.69) (6.45)

VO(HL)2 18.51 34.92 7.33(18.52) (34.91) (7.31)

VO(H2L00)2 14.02 39.67 7.77(14.01) (39.65) (7.75)

3.1. Structure of VIII complexes

VCl3 has the very common BI3 structure, a motif that featureshexagonally closest-packed chloride framework with vanadiumions occupying the octahedral holes. It formed adduct with DMFand this adduct is used as starting material for synthesis of thecomplexes with the studied compounds [HZ] according to the fol-lowing reaction:

VCl3 � DMFþ 2HZ! VZ2Clþ 2HClþ DMFðwhere Z ¼ HL;HL0;HL00Þ

It was reported that the IR spectra of the aminoalcohol ligandsgave characteristic bands, where tOH, tNH2 ; tCH2 , dOH, dCH2 ; tCN, sCH2 ;

dCAC, cOH are assigned [39,40]. The low intensity of the waggingvibration band of NH2 group supports the existence of an intramo-lecular hydrogen bond in MEA ligand and the broadness of allvibrations of the OH groups is mainly due to strong intermolecularhydrogen bonds in the three ligands. The IR spectra of the com-plexes indicated that the coordination occurs through both nitro-gen and oxygen sites, because of the shifts of OH and NH2 (orNH) bands beside the appearance of tVAO and tVAN bands at 480and 373 cm�1, respectively. The Nujol mull electronic absorptionspectra of the complexes gave two main bands at �450 and�600 nm with low intensities indicating that the complexes arenon l-oxo dinuclear complexes. The band at 450 nm points tothe 3T1g (F) ? 3T2g (F) transition while the band at 600 nm belongsto d–d transition. The analytical data indicated the metal:ligand ra-tios of 1:2 to both MEA and DEA complexes and 1:1 to TEA com-plex. The magnetic properties of MEA, DEA and TEA complexeswere characteristic for d2 electronic structure, l = 2.52, 2.53 and2.45 BM, respectively. The data indicate the octahedral structureof the complexes with bridged chloride ligands, where the com-plexes of MEA and DEA ligands were polymeric nature and thatof TEA was diametric, Fig. 2a.

lexes.

Color mp. (�C)

N Cl

13.56 17.16 Grey green 235(13.55) (17.15)

9.57 12.12 Grey green 239(9.54) (12.09)

6.00 15.18 Grey green 232(5.97) (15.17)

11.58 29.30 Greenish grey 240(11.58) (29.28)

8.54 21.61 Green 245(8.53) (21.59)

6.73 17.04 Brownish green 253(6.73) (17.01)

8.84 22.38 Gray green >300(8.84) (22.39)

(6.77) (17.16) Gray green >3006.78 17.17

5.59 14.15 Gray green >300(5.59) (14.13)

14.97 – Brownish green >300(14.96)

10.18 – Brownish green >300(10.18)

7.71 – Brownish green >300(7.70)

VIII- MEA and VIII- DEA complexes VIII – TEA complex

Cl

VN

O

O

N

Cl

VN

O

O

NCl

O

VN

O

O

Cl

VCl

O

O

N

O

H

H

Cl

VN

O

O

NCl

(a)

(b)Fig. 2. The general formula of (a) VIII– and (b) VIV-ligand complexes.

46 M.S. Masoud et al. / Journal of Molecular Structure 1050 (2013) 43–52

3.2. Structure of VIV complexes

Interaction of VCl3 with examined ligands in presence of oxygenleaded to formation of 1:2 (metal:ligand) complexes with sponta-neous oxidation of VIII to VIV as represented by the followingequation:

VCl3 þ 2HZ ½O� ! VZ2Cl2 ðwhere Z ¼ HL; HL0; HL00Þ

The IR spectra proved that the coordination bonds between li-gand and vanadium were through both oxygen and nitrogen lead-ing to a five-membered ring chelate structure, where tVAN bandswere recorded at 489, 483 and 479 cm�1, and tVAO bands were re-corded at 375, 373, and 370 cm�1 MEA, DEA and TEA complexes,respectively. The Nujol mull electronic spectra of the complexesgave three bands around 360, 560 and 750 nm. The first peak as-signed to the spin forbidden charge transfer transition, the othertwo peaks assigned to 2B2g ?

2A1g and 2B2g ?2B1g of the octahe-

dral geometry with Jahn–Teller effect where the 2T2g ground statewas splitted into 2Eg and 2B2g levels. The leff values of MEA, DEAand TEA complexes were 1.72, 1.71 and 1.72 BM, respectively. Theywere assigned to the d1-electronic structure of vanadium (IV). Thestructures of the complexes may be represented by the general for-mula as in Fig. 2b.

3.3. Structure of VO2+ complexes

When VOSO4 is used, only the 1:2 (metal:ligand) complexeswere produced where the sulphate group was not involved in thecomplexation. All the prepared complexes were found to be amor-phous with very high melting points, >300 �C. They exhibit poorsolubility in polar and non-polar solvents indicating that thesecomplexes are not dissociated in these solvents. The room temper-ature magnetic moments of the complexes ranged from 1.46 to1.69 B.M. These values are lower than the expected spin-only valueof d1-configuration of VO2+ indicating variable V@O� � �V@O interac-tion. The Nujol mull electronic spectral data for the complexes gavetwo bands at 550 and 800 nm for the 1:1 complexes and threebands at 550, 750 and 900 nm for the 1:2 complexes. The bands as-signed C4V symmetry of the studied complexes. The bands ap-peared at 550 and 800 nm were assigned as dxy ? dx2–y2 anddxy ? dxz transitions, respectively. These data indicated the squarepyramidal geometry. The three bands at 550, 750 and 900 nm for

1:2 complexes pointed to dxy ? dx2-y2, dxy ? dyz and dxy ? dz2 tran-sitions, respectively indicating the octahedral geometry of thesecomplexes with small distortion. The IR spectra for the complexesproved that the complexes of DEA and TEA have stretching vibra-tion of the hydroxyl groups at 3380–3500 cm�1. The tV@O variesbetween 850 and 998 cm�1. The bands due to tVAO and tVACl weredetected at 380 and 415 cm�1, respectively. The tNAH of MEA com-plexes was detected at 3360 and 3500 cm�1 for the 1:1 and 1:2complexes, respectively. Thus the structure of the VO2+ complexesmay be represented as in Fig. 3.

3.3.1. The thermal behavior of the VO2+ complexesThe compositions of VO(L)Cl and VO(L)2 complexes were con-

firmed by DTA and TG. Results of the thermal analysis performedin nitrogen are listed in Table 2.

Decomposition processes started in the 103–522 �C range. Thethermal pattern of the decomposition steps of the two com-plexes showed that MAligand bonds dissociated after removalof small molecules. At the beginning of the DTA-curves of thecomplexes there is a clearly manifested endothermic effect(above 100 �C), which is due to the hygroscopic moisture re-leased. The amount of this mass loss was also determined byKarl Fisher analysis, and were correlated with the intensity ofthese endothermic effects and with the respective decreases inthe mass. On heating the complexes the decomposition step cor-responds to the loss of molecules of the ligands, which is inagreement with the compositions in Table 2. The exothermal ef-fect (272–523 �C) dominates in the curves of the complexes,resulting from the decomposition of the organic matter. A fur-ther mass loss recorded up to higher temperatures indicatedthe formation of thermally stable oxide. The reaction order ofthe thermal reactions was calculated using Kissinger approach[41]. The values of the collision factor, Z were obtained, Table 2,based on Horowitz–Metzger equation [42] by the use of the fol-lowing equation:

Z ¼ DERTm

b expðDE=RT2mÞ

The entropies of activation DS– were obtained from the follow-ing equation:

Z ¼ KTm

hexpðDS–=RÞ

where R is the molar gas constant, b is the rate of heating (K s�1), Kis the Boltzman constant and h is Planck’s constant. The activationenergy, Ea of the thermal decomposition step was determined fromthe slope of the straight line of the lnDTtS 1

T relationship. The slopewas of Arrhenius type and equals to � Ea

R . The change of entropy,DS– values for the two complexes were given in Table 2. The nega-tive values indicated that the activated complexes have more or-dered structures than the reactants [43]. The lower values of thecollision factor, Z indicated the slow nature of the reactions [44].Also, it seems that DS– values are nearly of the same magnitudesuggesting a uniform decomposition pattern. The thermal decom-position of the square pyramidal complex, VO(L)Cl showed fourpeaks at 103.4, 272.2, 418.4 and 522.2 �C with activation energies101.5, 546.1, 344.4 and 196.8 kcal mol�1, respectively, and ordersof reactions of 0.92, 0.95, 0.98 and 0.96, respectively. The first peakwas due to the dehydration reaction of probably adsorbed water. Itsdecomposition began at 103.4 �C and proceeded in three stagesinvolving a loss in mass. The stages probably consisted of a numberof unidentifiable processes. The dehydration of the vanadium com-plex was suggested in stage 1. The processes occurring in the endo-thermic first stage of the decomposition partly coincide with thoseof the second stage, so the corresponding temperature ranges weredifficult to determine precisely. The total thermal effect of the

1:1-complexes 1:2-complexes

NV

Cl

O

O

NV

Cl

O

O

NV

Cl

O

ON

VCl

O

O

ClV

O

NO

NV

Cl

O

O

NV

O

N

O

O

O

O

N

OV

N

NV

N

O

O

O

O

O

O

NV

N

Fig. 3. The probable structures of VO2+–ligands complexes.

Table 2Kinetic parameters of the dissociation steps.

Complex Type Tm (�C) n b Z Ea (kcal mol�1) DS– (cal mol�1) Assignment

VO(L)Cl Endo 103.4 0.92 1.83 47.09 101.5 �0.114 DehydrationExo 272.2 0.95 1.89 273.36 546.1 �0.099 Dissociation of ligand and formation of VOCl2 and VO2

Exo 418.4 0.98 1.95 150.01 344.4 �0.103Exo 522.2 0.96 1.92 96.81 196.8 �0.117 Dissociation to form V2O5 as final product

VO(L)2 Endo 153.3 0.61 1.08 8.27 30.97 �0.233 Dehydration, association with dissociation.Exo 289.2 2.60 3.23 85.90 100.6 �0.211 Dissociation steps, formation of V2O5 as final productExo 515.2 1.39 2.56 153.5 259.5 �0.278

Table 3Physical properties of used solvents.

Solvent D nD e RD l

Chloroform 1.489 1.446 4.81 21 1.15DMF 0.945 1.431 36.7 19.9 3.86DMSO 1.096 1.478 46.7 20.1 3.90Dioxane 1.034 1.422 2.25 21.6 0.45Ethanol 0.789 1.361 24.5 12.8 1.69Methanol 0.791 1.328 32.7 8.2 1.70Water 0.998 1.333 80.1 3.7 1.82

M.S. Masoud et al. / Journal of Molecular Structure 1050 (2013) 43–52 47

second stage is exothermic. No stable intermediate is formed be-tween the first and second stages as the second stage starts at272.2 �C. Above 350 �C, a sharp decrease in mass is noted and thesubsequent slow mass loss was observed until 418.4 �C. Last stepin the thermal decomposition involves the decomposition of inter-mediates and oxidation of metal by atmospheric oxygen. The laststage was started around 522.2 �C suggested the formation ofV2O5 residue. Above 600 �C the mass of the sample did not change.The main gaseous products of the decomposition of the complexwere CO2, small amounts of H2O and trace amounts of HCl. Thefollowing scheme summarizes the decomposition steps:

VOðMEAÞCL!dehydration of adsorbed water�������������������!

—H2O—HCldissociation of the ligand

VOCl2

!VO2!V2O5

In the thermal curves of the octahedral complex, VO(L)2 twomain stages of decomposition can be distinguished, connectedwith continuous mass loss and strong exothermic effects. ItsDTA curve exhibits three peaks. One peak was endothermicand the others were exothermic at 153.2, 289.2 and 515.2 �Cwith activation energies of 30.97, 100.6 and 259.5 kcal mol�1,respectively and reaction orders of 0.61, 2.60 and 1.39, respec-tively, Table 2. The analysis of the mass spectra of the gaseousproducts of this complex-decomposition confirmed that at153.3 �C water molecules were released in the first stage ofdecomposition in an endothermic step and a stable solid inter-mediate was formed. In the temperature range 200–290 �C, theanhydrous salt from the first stage decomposes rapidly (the sec-ond stage), and then more slowly between 290 and 515 �C (thethird stage). The processes which take place in the second andthird stages are exothermic. The presence of V2O5 was confirmedin the sinter obtained at about 515 �C. VO(L)2 have the samemechanism as VO(L)Cl but with higher energies as shown inthe following scheme:

VOðMEAÞ2!dehydration of adsorbed water !dissociation

ðbroad bandÞVO2!V2O5

3.4. Solvent effect on the electronic absorption spectra of somecomplexes

One of the decisive experimental techniques for investigatingsolute–solvent interactions was based on the dependence of theabsorption spectra on the nature of the solvent [45–47]. The elec-tronic absorption spectra of some vanadium complexes with etha-nolamines are studied in seven different solvents. The absorptionspectra were recorded in these solvents of different physicochem-ical properties, mainly dielectric constant, dipole moment andrefractive index. The physical parameters for these solvents at20 �C are collected in Table 3, where D is density, nD is refractiveindex, e is dielectric constant, RD is molar refraction and l is dipolemoment in Debye. Generally, the solvent effect on the absorptionbands of these complexes involves displacements and does not in-volve a fundamental change in the general form of the spectrum.

The electronic absorption spectra of the investigated com-plexes are collected in Table 4. The electronic absorptionspectrum of the complex, V(L)2Cl showed the appearance ofthree distinctive bands, two of them in the ranges 443–470 and591–621 nm in all solvents except MeOH and EtOH, where a thirdband appeared at 363 nm. The observed d–d transition bandscharacteristic of this octahedral complex appeared in therange 443–470 nm were assigned as 3T1g(F) ? 3T1g(P) whilethose appeared in the range 591–621 nm were assigned as3T1g(F) ? 3T2g(F) [48]. The third spin-allowed transition appearedonly in alcohols was assigned as 3T1g(F) ? 3A2g(P). The absorption

Table 4The observed kmax and 107/kmax values of the investigated complexes.

Solvent V(L)2Cl VO(L)Cl V(L)2Cl2

k1 k2 k3 k1 k2 k1 k2 k3

CHCl3 463 – 614 463 620 370 512 759DMSO 443 – 603 442 609 362 490 751DMF 470 – 621 470 619 380 518 760EtOH 446 463 594 448 650 357 498 744MeOH 449 464 598 447 597 361 603 740Dioxane 464 – 591 460 591 391 520 769H2O 450 – 598 449 600 360 500 750

48 M.S. Masoud et al. / Journal of Molecular Structure 1050 (2013) 43–52

band corresponded to the highest transition had appeared atabout 463 nm in solvents with the lower values of e and l, suchas 1,4-dioxane and chloroform. This peak has negative solvato-chromism (hypsochromic shifts) in polar solvents with highervalues of e and l except in DMF, whereas a positive solvatochro-mism (bathochromic shift) was observed. DMF is polar aproticsolvent and is known to be strong hydrogen-bond acceptor(HBA), with the highest HBA basicity among most commonorganic solvents [49,50]. Therefore, the positive solvatochromismmay be attributed to the effect of DMF basicity as well as itsH-bond accepting character.

This indicated that a combination of several solvent character-istics such as polarity, basicity and H-bond-accepting ability isresponsible for the spectral characteristics of the obtained com-plex. The absorption band corresponded to the lowest transitionhad appeared at 591 nm in dioxane which has the lowest dielectricconstant and dipole moment. This peak was bathochromic shiftedin all other solvents with higher values of e and l. This behaviormay be explained as follows, this complex in the ground stateand in the excited state indicated different polarities. Complexeswith non-polarized ground state were more strongly polarized inprotic solvents, because the high-energy polar structure of theexcitation state was stabilized. The excited state was loweredand the ground state is hardly affected leading to a bathochromicshift. Also, as expected the highest red shift was observed inDMF. This confirmed the responsibility of the solvent basicity forthe spectral characteristics of this complex. The spectra of the oxo-vanadium complex, VO(L)Cl in different solvents exhibited twowell-resolved bands of moderate intensity at the wavelengthranges of 442–470 and 591–650 nm. These bands are assigned tothe ligand field transitions, dxy ? dz2 and dxy ? dx2-y2 correspond-ing to b2 ? a�1 and b2 ? b�1 transitions, respectively [51–53]. Theshorter wavelength band have appeared at 460 nm in the lowestpolarity solvent, 1,4-dioxane. This peak undergoes hypsochromicshifts in all more polar solvents with higher values of e and l,whereas a bathochromic shift was observed in chloroform withlow polarity and dielectric parameters. The hypsochromic shiftdue to increasing solvent polarity means that the excited state isless polar than the ground state in this electronic transition. Thelonger wavelength absorption band had appeared at 591 nm in1,4-dioxane. This peak was bathochromic shifted in all other sol-vents with higher values of polarity parameters. Therefore thepolarities of the ground and excited state in this electronic transi-tion were different and the excited state was more polar than theground state. The spectral results for the six-coordinate V(L)2Cl2

complex in solution (Table 4) were characteristic of an octahedralgeometry around the central vanadium (IV) ion. The spectrum ofthis complex showed three bands in the ranges 360–391, 490–603 and 740–769 nm. The first peak assigned to the spin forbiddencharge transfer transition, the second peak assigned to 2B2g ?

2A1g

and the third peak assigned to 2B2g ?2B1g [48]. The shorter wave-

length band appeared at 391 nm in 1,4-dioxane. This peak hypso-chromic shifted in all other more polar solvents indicating that theexcited state was less polar than the ground state. The second band

assigned as 2B2g ?2A1g appeared at 520 nm in 1,4-dioxane. This

peak hypsochromic shifted in all other more polar solvents exceptin methanol, where a bathochromic shift was observed. Finally, theband assigned as 2B2g ?

2B1g appeared at 769 nm in 1,4-dioxane.This peak hypsochromic shifted in all other more polar solventsand had the same solvatochromic behavior as the first electronictransition.

The electronic absorption spectra of the formed complexes hadbeen analyzed in terms of kmax by multiple linear regression tech-nique using the following equation [54].

Y ¼ ao þ a1x1 þ a2x2 þ a3x3 þ . . .þ anxn

where Y is the observed peak location on an absorption band in agiven solvent. ao is the regression intercept, a1, a2, a3, an are coeffi-cients could be determined by multiple regression technique and x1,x2, x3, xn are the various empirical solvent polarity parameters andcould be expressed as

E ¼ 2:859� 10�3t�

K ¼ ðD� 1Þ=ð2Dþ 1Þ

M ¼ ðn2 � 1Þ=ð2n2 þ 1Þ

J ¼ ðD� 1Þ=ðDþ 1Þ

H ¼ ðn2 � 1Þ=ðn2 þ 2Þ

N ¼ J � H

E is related to t�, which is the wave number of the absorptionmaximum in a given solvent [55]. The function was sensitive toboth solvent–solute hydrogen bonding and dipolar interactions.The dielectric function K represents the dipolar interaction and isrelated to the dielectric constant (D) of the solvent. The functionsJ and H have been introduced [56] to account for the non-specificsolute–solvent interactions such as dispersion and dipolar effects.These are related to the dielectric constant (D) and refractive index(n) of the solvents.

The functions M and N had been introduced [56,57] to accountfor the solute permanent dipole-solvent induced dipole and solutepermanent dipole-solvent permanent dipole interactions,respectively.

The regression intercept ao has been considered as peak posi-tion in the gas spectra [58–66]. The intercept ao and the coeffi-cients a1, a2, a3, a4 have been calculated by multiple regressionanalysis using a micro static program and the calculations havebeen carried out on a PC computer. The results of calculations forthe complexes investigated are collected in Tables 5–11. Matthewscorrelation coefficient (MCC) and the probability of variation (P),were deduced from the micro statistic program. This had been con-sidered as a measure of the goodness of the fit as MCC values werenear to unity, alternatively, the lower value of the significantparameter (P), which means the correlation is good. Based onone-parameter equation, the parameter (M) plays the importantrole for determining the spectral shifts for V(L)Cl at k1. The relativehigh value of MCC and the lower value of P for M points to that thesolute permanent dipole-solvent induced dipole interactions areplaying the important role to explain the spectral shifts observed.However, the parameter (K) was important for the determinationof the spectral shifts for VO(L)Cl at k2 and V(L)2Cl2 at k1 .The rela-tive high value of MCC for K for these compounds pointed to thatthe dielectric constant is effective to explain the spectral shiftsrather than the electronic character of the substituents. ForV(L)2Cl2 at k2 and k3, the parameter (N) controlled the spectralshifts, the relative high value of MCC for (N) for these complexes

Table 5Regression analysis for V(L)2Cl in different solvents at k1.

Parameters ao a1 a2 a3 a4 MCC P

E 20483.21 31.497 – – – 0.600 0.155M 23338.13 �7001.17 – – – 0.275 0.550N 21439.56 1012.78 – – – 0.551 0.200K 20894.26 2568.96 – – – 0.512 0.234E,M 16304.10 57.09 15330.15 – – 0.697 0.264E,N 20683.24 23.19 363.49 – – 0.611 0.392E,K 20431.75 24.81 871.17 – – 0.612 0.391M,N 21650.43 �1003.32 980.70 – – 0.552 0.483M,K 21539.85 �2926.34 2378.17 – – 0.529 0.518N,K 22634.99 2936.50 �5254.67 – – 0.575 0.447E,M,N 15204.18 71.56 18562.93 �397.32 – 0.705 0.504E,M,K 15203.15 73.11 19618.99 �1154.84 – 0.708 0.498E,N,K 20373.81 25.23 �84.30 1066.57 – 0.612 0.658M,N,K 22503.65 5178.14 4642.04 �9461.08 – 0.589 0.692E,M,N,K 16056.06 67.49 1763.43 �5380.97 0.714 0.740

Table 6Regression analysis for V(L)2Cl in different solvents at k2.

Parameters ao a1 a2 a3 a4 MCC P

E 16182.35 8.66 – – – 0.278 0.546M 17887.89 �6697.39 – – – 0.444 0.318N 16684.84 �163.63 – – – 0.150 0.748K 16929.99 �783.96 – – – 0.266 0.564E,M 18781.44 �7.25 �9534.21 – – 0.465 0.615E,N 15561.72 34.43 �1727.81 – – 0.680 0.289E,K 16353.64 30.9179 �2899.58 – – 0.734 0.218M,N 18696.11 �9569.67 �469.64 – – 0.588 0.427M,K 18926.93 �9051.817 �1374.11 – – 0.625 0.371N,K 19900.61 5011.26 �14135.3 – – 0.768 0.168E,M,N 15720.16 33.03 �563.82 �1105.81 – 0.680 0.548E,M,K 15820.10 35.84 2001.97 �310,632 0.736 0.447E,N,K 18641.32 14.05 3328.91 �10614.9 – 0.793 0.338M,N,K 19918.68 �712.34 4776.63 �13556.6 – 0.769 0.384E,M,N,K 17637.86 23.877 4872.51 3758.33 �12113.3 0.806 0.578

Table 7Regression analysis for VO(L)Cl in different solvents at k1.

Parameters ao a1 a2 a3 a4 MCC P

E 20631.51 29.25 – – – 0.569 0.183M 23329.59 �6743.86 – – – 0.271 0.557N 21571.31 845.30 – – – 0.470 0.288K 21153.98 2055.38 – – – 0.423 0.344E,M 16939.98 51.85 13541.55 – – 0.653 0.329E,N 20671.66 27.58 72.96 – – 0.569 0.457E,K 20624.83 28.38 113.00 – – 0.569 0.457M,N 21982.57 �1956.78 782.73 – – 0.475 0.600M,K 21954.22 �3627.34 1818.88 – – 0.445 0.643N,K 23389.34 3770.91 �7991.36 – – 0.536 0.508E,M,N 14937.84 78.21 19426.04 �723.22 – 0.682 0.545E,M,K 14952.24 80.78 21285.00 �2085.06 – 0.693 0.525E,N,K 21276.55 23.58242 948.3548 �2084.96 – 0.571 0.716M,N,K 23212.89 6956.66 6062.25 �13642.5 – 0.563 0.728E,M,N,K 16416.78 71.14 23597.76 3028.04 �9341.88 0.711 0.744

M.S. Masoud et al. / Journal of Molecular Structure 1050 (2013) 43–52 49

pointed to that solute permanent dipole-solvent permanent dipoleinteractions are playing the important role to explain the observedspectral shifts. Based on the two parameters equation, the combi-nation (E, M) gave higher correlation shown by its MCC and P forV(L)2Cl and VO(L)Cl at k1. So, the solvent–solute hydrogen bondingcombined with the solute permanent dipole-solvent induced di-pole interactions are major factors for determining the spectralshifts in this case. The combination (N, K) gave higher values of cor-relation for V (L)2Cl and VO(L)Cl at k2, as evident by its MCC. It isconcluded that the studied properties of the solvent are effectiveparameters to explain the spectral shifts for the organic compound.

The effects of the solute permanent dipole solvent permanent di-pole interactions had the major effect in this correlation. The com-bination (M, K) gave higher values of correlation for V(L)2Cl2 at k2

and k3 as evident by its MCC and So, the solute permanent dipolesolvent induced dipole combined with dipolar interaction, whichwas related to the dielectric constant of the solvent. For threeparameters equation. The combination (E, N, K) gave higher valueof correlations for V(L)2Cl at k2 and V(L)2Cl2 at k1 shown by itsMCC So, the effect of interamolecular hydrogen bonding combinedwith solute permanent dipole-solvent permanent dipole interac-tions and the polarity of the solvent given in terms of its dielectric

Table 8Regression analysis for VO(L)Cl in different solvents at k2.

Parameters ao a1 a2 a3 a4 MCC P

E 16369.02 �0.468 – – – 0.009 0.985M 16607.96 �1354.89 – – – 0.052 0.912N 16637.04 �536.32 – – – 0.283 0.539K 17108.96 �1790.641 – – – 0.350 0.441E,M 17577.26 �7.86 �4432.18 – – 0.1 0.980E,N 15569.60 32.73 �1452.70 – – 0.460 0.621E,K 16587.81 27.95 �3703.64 – – 0.499 0.564M,N 17848.00 �5761.80 �720.57 – – 0.345 0.776M,K 18207.44 �4979.23 �2115.27 – – 0.393 0.715N,K 19979.69 4842.75 �14693.0 – – 0.532 0.514E,M,N 12298.07 61.61 11083.88 –1906.98 – 0.503 0.801E,M,K 12557.76 65.18 15121.76 �5265.24 0.566 0.724E,N,K 19199.95 8.70 3801.06 �12513.2 – 0.537 0.761M,N,K 19772.24 8178.98 7536.68 �21337.1 – 0.566 0.724E,M,N,K 15219.96 47.65 19325.79 5504.25 �18456.4 0.630 0.843

Table 9Regression analysis for V(L)2Cl2 in different solvents at k1.

Parameters ao a1 a2 a3 a4 MCC P

E 23750.54 71.10 – – – 0.756 0.049M 30806.60 �18973.2 – – – 0.417 0.352N 25786.96 2512.83 – – – 0.763 0.046K 24237.66 6835.13 – – – 0.770 0.043E,M 16804.62 113.64 25479.58 – – 0.825 0.102E,N 24548.07 37.98 1449.27 – – 0.801 0.128E,K 23505.52 39.27 4147.61 – – 0.822 0.105M,N 26728.84 �4481.53 2369.53 – – 0.769 0.168M,K 26041.18 �8175.04 6302.14 – – 0.788 0.143N,K 24490.82 427.06 5697.31 – – 0.770 0.166E,M,N 18486.62 91.50 20536.04 607.58 – 0.830 0.260E,M,K 19149.54 79.51 16344.73 2459.71 – 0.838 0.249E,N,K 16046.09 94.25 �10854.6 29304.97 – 0.884 0.162M,N,K 24934.27 �17483.00 �5331.33 19899.41 – 0.805 0.314E,M,N,K 14267.68 111.66 8635.48 �10093.6 26649.34 0.888 0.379

Table 10Regression analysis for V(L)2Cl2 in different solvents at k2.

Parameters ao a1 a2 a3 a4 MCC P

E 12595.69 14.26 – – – 0.774 0.041M 14289.62 �5249.66 – – – 0.588 0.165N 13001.97 508.13 – – – 0.788 0.035K 12708.74 1335.25 – – – 0.767 0.044E,M 12336.04 15.85 952.45 – – 0.776 0.158E,N 12762.08 7.35 302.38 – – 0.823 0.104E,K 12551.01 8.46 756.25 – – 0.830 0.097M,N 13561.31 �2661.37 423.21 – – 0.832 0.095M,K 13444.33 �3334.28 1117.87 – – 0.844 0.083N,K 13177.79 791.26 �772.87 – – 0.791 0.14E,M,N 13429.82 1.46 �2262.26 395.10 – 0.832 0.261E,M,K 13353.41 1.04 �3010.81 1067.18 – 0.844 0.238E,N,K 12164.41 11.31 �562.55 2060.07 – 0.834 0.257M,N,K 13294.35 �4595.42 �722.34 2960.17 – 0.852 0.223E,M,N,K 12924.09 3.87 �3688.78 �887.65 3194.47 0.854 0.467

50 M.S. Masoud et al. / Journal of Molecular Structure 1050 (2013) 43–52

constant were effective parameters to explain the spectral shifts.The combination (M, N, K) gave higher values of correlations forV(L)2Cl2 at k2 and k3. So, the effect of solute permanent dipole-solvent induced dipole combined with solute permanent dipole-solvent permanent dipole interactions and the polarity of thesolvent given in terms of its dielectric constant are major ratherthan the electronic character of the substituent. So, the additionof parameter M to N and K leads to better results. Generally, it isconcluded that the addition of the fourth solvent parameter tothe three parameter equations always improve the correlationwith the solvent induced spectral shifts.

4. Kamlet–Taft equation

The overall solvent effects were described by the use of Kamlet–Taft equation, which has successfully been applied to differentiatebetween the influence of nonspecific chemical interactions, includ-ing electrostatic effects (dipolarity/polarizability), and specificinteractions, that is hydrogen bonding, which are related to themolecular structure of a compound [50,67–69]. Kamlet–Taft equa-tion can be written as follows:

mmax ¼ mo þ sp� þ bbþ aa ð1Þ

Table 11Regression analysis for V(L)2Cl2 in different solvents at k3.

Parameters ao a1 a2 a3 a4 MCC P

E 12595.68 14.26 – – – 0.774 0.041M 14289.63 �5249.71 – – – 0.588 0.165N 13001.97 508.31 – – – 0.788 0.035K 12708.74 1335.26 – – – 0.767 0.044E,M 12336.06 15.85 952.38 – – 0.776 0.158E,N 12762.08 7.35 302.38 – – 0.823 0.104E,K 12551.01 8.46 756.25 – – 0.830 0.097M,N 13561.32 �2661.42 423.21 – – 0.832 0.095M,K 13444.34 �3334.34 1117.87 – – 0.844 0.083N,K 13177.807 791.27 �772.91 – – 0.791 0.140E,M,N 13429.85 1.46 �2262.37 395.10 – 0.832 0.261E,M,K 13353.44 1.049 �3010.93 1067.19 – 0.844 0.238E,N,K 12164.43 11.31 �562.53 2060.03 – 0.834 0.257M,N,K 13294.36 �4595.47 �722.34 2960.17 – 0.852 0.223E,M,N,K 12924.12 3.87 �3688.89 �887.64 3194.46 0.854 0.467

Table 12Solvent independent correlation coefficient a, b, s of the Kamlet Taft parameters.

Complex mo a b s MCC

V(L)2Cl 21210.40 417.14 704.47 276.68 0.701VO(L)Cl 21345.15 423.34 574.24 234.94 0.644V(L)2Cl2 25626.22 867.18 1459.11 458.39 0.795

Table 13Percentage contribution of the solvatochromic parameter.

Complex Pp� (%) Pb (%) Pa (%)

V(L)2Cl 19.78 50.38 29.83VO(L)Cl 19.06 46.59 34.34V(L)2Cl2 16.46 52.39 31.14

Fig. 4. Percentage contribution to the solvatochromic effect (1) V(L)2Cl, (2) VO(L)Cl,and (3) V(L)2Cl2.

M.S. Masoud et al. / Journal of Molecular Structure 1050 (2013) 43–52 51

where mmax is the wavenumber (cm�1) in the maximum absorptionband of the investigated compounds in pure solvents, mo the regres-sion intercept corresponds to the gaseous of the spectrally activecompounds, p* is a measure of the solvent dipolarity/polarizability,b is the scale of the solvent hydrogen-bond acceptor (HBA) basici-ties, a is the scale of the solvent hydrogen-bond donor (HBD) acid-ities and mo, a, b, s are solvent independent constants, theirmagnitudes and sign provide measures of the influence of thecorresponding solute–solvent interactions on the wavenumber inthe maximum of electronic absorption band, which have beendetermined by multiple regression analysis, using SPSS statistics

program. The results of the multiple regressions are presented inTable 12. The percentage contribution of (a, b, p*) for the investi-gated complexes are shown in Table 13 and Fig. 4.

5. Conclusions

Different oxidation states of vanadium formed stable complexeswith ethanolamines and that was proved by different spectro-scopic and analytical techniques. The obtained complexes can beused as color additives in different manufacturers such as ceramicsmanufacturing industry because of their thermal stability whichwas proved by different thermal analytical methods.

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