1

ABSTRACT

The Mn(II), Co(II), Ni(II), Cu(II), Zn(II) and Fe(II) complexes of ligand Thiosalicylicacid were

synthesized and characterized by electronic and infrared spectroscopy, percentage metal,

magnetic susceptibility and melting point measurements.

The ligands were bidentate, co-ordinating through the carboxylic oxygen of thiosalicylic acid

and through the deprotonated hydroxide.

2

TABLE OF CONTENTS

TITLE PAGE………………………………………………………………………………………i

CERTIFICATION………………………………………………………………………………..ii

DEDICATION……………………………………………………………………………………iii

ACKNOWLEDGEMENT…………………………………………………………...…………...iv

ABSTRACT .................................................................................................................................... 1

TABLE OF CONTENTS…..…………………………………………………………………...…2

LIST OF TABLES……..….……………………………………………………………………....4

LIST OF FIGURES……………………………………………………………………………….5

SYMBOLS AND ABBREVIATION……………………………………………………………..6

LIST OF APPENDICES………….……………………………………………………………….7

CHAPTER ONE:INTRODUCTION

1.1. COORDINATION COMPOUNDS ..................................................................................... 8

1.1.1. The Nature Of Ligands ..................................................................................................... 9

1.1.2 Classification of ligands ......................................................................................................... 9

1.1.3. Coordination Number ......................................................................................................... 10

1.1.3.1. Factors Affecting Coordination Number ....................................................................... 100

1.2. THIOSALICYLIC ACID .................................................................................................. 11

1.2.1. Preparation of Thiosalicylic Acid ....................................................................................... 11

1.3. Structural review of some metal(II) complexes ................................................................. 13

1.4. Electronic spectroscopy and magnetic moments of some metal(II) complexes ................... 16

1.5. Infrared Spectroscopy of Some Metal(II) Complexes ....................................................... 20

1.6. AIM OF THE PROJECT ................................................................................................... 22

CHAPTER TWO:THEORETICAL BACKGROUND

2.1. ELECTRONIC SPECTROSCOPY ..................................................................................... 233

2.1.1. ELECTROMAGNETIC RADIATION ............................................................................ 233

2.1.2. INFRARED SPECTROSCOPY ....................................................................................... 244

2.1.3. WAVE NUMBERS .......................................................................................................... 244

3

2.1.4. MOLECULAR VIBRATIONS ........................................................................................ 244

2.1.5. ABSORPTION REQUIREMENTS ................................................................................. 255

2.1.6. ABSORPTION RANGES ................................................................................................ 256

2.1.7. ELECTRON EXCITATION ............................................................................................ 266

2.1.8. TERMINOLOGIES OF ELECTRONIC SPECTROSCOPY ........................................... 277

CHAPTER THREE: METHODOLOGY

3.1. REAGENTS AND SOLVENTS ......................................................................................... 288

3.2. SYNTHETIC METHODS ..................................................................................................... 28

3.2.1. PREPARATION OF THE METAL(II)COMPLEXES ...................................................... 28

3.3. PHYSICAL MEASUREMENT............................................................................................. 28

3.3.1. MELTING POINT AND DECOMPOSITION TEMPERATURE .................................... 28

3.3.2. INFRARED SPECTRA .................................................................................................. 2929

3.3.3. ELECTRONIC SPECTRA ................................................................................................. 29

3.3.4. SOLUBILITY ..................................................................................................................... 29

3.3.5. MAGNETIC MOMENT ..................................................................................................... 29

3.4. METAL ANALYSIS ............................................................................................................. 29

CHAPTER FOUR:RESULTS AND DISCUSSIONS

4.1. GENERAL PROPERTIES OF THE COMPLEX ............................................................... 344

4.1.1. COLOUR .......................................................................................................................... 344

4.1.2. MELTING POINT/ DECOMPOSITION TEMPERATURE ........................................... 344

4.1.3. PERCENTAGE METAL DETERMINATION ................................................................ 344

4.1.4. SOLUBILITY ................................................................................................................... 344

4.1.5. MAGNETIC MEASUREMENT ...................................................................................... 344

4.1.6. ELECTRONIC SPECTRA OF THE LIGANDS……………………………………..…..35

4.1.7. INFRARED SPECTRA OF THE LIGANDS AND COMPLEXES .................................. 36

CONCLUSION ........................................................................................................................... 411

RECOMMENDATION .............................................................................................................. 424

REFERENCES………….……………………………………………………………………….58

4

LIST OF TABLES

Table 4.1: Analytical data of the ligand and complexes………..………………………………..36

Table 4.2: Solubility data of the complexes………………….………………………………..…38

Table 4.3: Electronic spectra…………………………………….……………………………….39

Table 4.4 Infrared spectra of the ligand and metal complexes…………………………..………40

5

LIST OF FIGURES

Figure 1.1: Three-step reactive process of producing thiosalicylic acid………..……………….13

Figure 1.2:Proposed structure of the metal complexes….………………………………………45

Figure 1.3:Proposed structure of the copper complex….……………………………………….46

6

SYMBOLS AND ABBREVIATION

M: Metal

L =Ligand

Mn: Manganese

Fe: Iron

Co: Cobalt

Ni: Nickel

Cu: Copper

Zn: Zinc

TSA: Thiosalicylic acid

B.M: Bohr Magneton

IR: Infrared

UV: Ultraviolet

NMR: Nuclear Magnetic Resonance

DMSO: Dimethylsulfoxide

ETOH: Ethanol

EDTA: Ethylenediaminetetraacetic acid

MeOH: Methanol

CH2Cl2: Dichloromethane

(C2H5)2O: Diethyl ether

CT = Charge Transfer

7

LIST OF APPENDICES

The Infrared Spectrum of Thiosalicylic acid…………………………………………………….45

The Infrared Spectrum of [Mn(TSA)2].3H2O…………………………………………..……….46

The Infrared Spectrum of [Fe(TSA)2].3H2O…………………………………………………….47

The Infrared Spectrum of [Co(TSA)2].3H2O…………………………………………………….48

The Infrared Spectrum of [Ni(TSA)2].3H2O…………………………………………………….49

The Infrared Spectrum of [Cu(TSA)ClH2O]……………………………………………………50

The Infrared Spectrum of [Zn(TSA)2].3H2O…………………………………………………….51

The UV Spectrum of [Mn(TSA)2].3H2O………………………………………………………...52

The UV Spectrum of [Fe(TSA)2].3H2O………………………………………………………….53

The UV Spectrum of [Co(TSA)2].3H2O…………………………………………………………54

The UV Spectrum of [Ni(TSA)2].3H2O………………………………………………………….55

The UV Spectrum of [Cu(TSA)ClH2O]…………………………………………………………56

The UV Spectrum of [Zn(TSA)2].3H2O…………………………………………………………57

8

CHAPTER ONE

INTRODUCTION

1.1. COORDINATION COMPOUNDS

Coordination refers to the "coordinate covalent bonds" (dipolar bonds) between the ligands and

the central atom. Originally, a complex implied a reversible association of molecules, atoms, or

ions through such weak chemical bonds. As applied to coordination chemistry, this meaning has

evolved. Some metal complexes are formed virtually irreversibly and many are bound together

by bonds that are quite strong. (Cotton and Wilkinson, 1998)

Coordination compounds are also known as coordination complexes, complex compounds,

orsimplycomplexes. A coordination complex consists of a central atom or ion, which is usually

metallic and is called the coordination center, and a surrounding array of bound molecules or

ions, that are in turn known as ligands or complexing agents. Many metal-containing

compounds, especially those of transition metals, are coordination complexes.(Greenwood and

Norman, 1997)

The essential feature of coordination compounds is that coordinate bonds form between electron

pair donors, known as the ligands, and electron pair acceptors, the metal atoms or ions. The

number of electron pairs donated to the metal is known as its coordination number. Although

many complexes exist in which the coordination numbers are 3, 5, 7, or 8, the majority of

complexes exhibit coordination numbers of 2, 4, or 6. (James E. House, 2008)

1.1.1. The Nature Of Ligands

Ligands also called donor molecules (since they donate a pair of electron to form a coordinate

bond) include simple ligands such as water, ammonia, chloride ion, etc. They all have active lone

pairs of electrons in the outer energy level. These are used to form co-ordinate bonds with the

metal ion. There are also organic ligands such as alkenes and benzene whose pi bonds can

coordinate to empty metal orbitals. Ligands can be anions, cations, or neutral molecules. All

9

ligands are lone pair donors. In other words, all ligands function as Lewis bases. (Jim Clark,

2003).

1.1.2 Classification of ligands

Ligands can be classified based on different criteria such as; number of donor atoms, nature, size,

and so on. Based on number of donor atoms, we have;

1. Monodentate Ligands

A monodentate ligand has only one donor atom used to bond to the central metal atom or ion.i.e

it has only one lone pair of electron that it can use to bond to the metal – any other lone pairs are

pointing in the wrong direction. (Jim Clark, 2003). The term "monodentate" can be translated as

"one tooth," referring to the ligand binding to the center through only one atom. Some examples

of monodentate ligands are: chloride ions (referred to as chloro when it is a ligand), water

(referred to as aqua when it is a ligand), hydroxide ions (referred to as hydroxo when it is a

ligand), and ammonia (referred to as ammine when it is a ligand).

2. Bidentate Ligands

Bidentate ligands have two donor atoms which allow them to bind to a central metal atom or

ion at two points. Common examples of bidentate ligands are 1,2-diaminoethane (oldname:

ethylenediamine – often given the abbreviation „en‟), and the ethanedioate (old name: oxalate ion

„ox‟).

3. Tridentate ligands

These are ligands with three donor atoms. They have three possible sites of attachment. Example

is diethylenetriamine. They make use of their three donor sites in bonding to the central metal

ion.

4. Polydentate Ligands

Polydentate ligands range in the number of atoms used to bond to a central metal atom or ion.

They usually form chelates with the central metal ion which is ring-like and formed through a

10

process called chelation. EDTA, a hexadentate ligand, is an example of a polydentate ligand that

has six donor atoms with electron pairs that can be used to bond to a central metal atom or

ion.(Jim Clark, 2003).

1.1.3. Coordination Number

Coordination number is the number of donor atoms attached to the metal ion and this can vary

from 2 to as many as 16. Higher coordination numbers are rare but are found in lanthanides and

actinides. The most stable are coordination number 2 and 6. Coordination number 4 and 6 are the

most common. In simple terms, the coordination number of a complex is influenced by the

relative sizes of the metal ion and the ligands and by electronic factors, such as charge which is

dependent on the electronic configuration of the metal ion.

In describing complexes, the ligands directly attached to the metal (usually as Lewis bases,

donating electrons to the metal), are counted to determine the coordination number of the

complex. Ions that are directly coordinated to the metal are written within the brackets of the

formula, and are referred to as inner sphere. Ions that are serving as counter ions in order to

produce a neutral salt and are not coordinated to the metal are called outer sphere,andare written

outside brackets in the formula. For example:

[Pt(NH3)6)]Cl4: coordination number=6, and chloride is outer sphere

[Pt(NH3)2Cl4] : coordination number =6, and chloride is inner sphere

1.1.3.1. Factors Affecting Coordination Number

The following factors influence or determine the most suitable coordination number for a given

metal and ligand.

1. Size and steric effect: the larger the central atom, the more ligands it can accommodate and

the smaller the ligand size, the more ligands around the central atom. Large central atoms

are common in transition metals, lanthanides and actinides. For example; Molybdenum

(Mo) and Tungsten (W) have coordination numbers up to 8 in cyano complexes.

2. Oxidation state of the central metal: the higher the oxidation state, the more ligands that

can be accommodated. For example; M4+accommodates more ligands than M3+

. The higher

11

the charge on the ligand, the more deficient in electron and the higher the tendency to

accommodate more ligands. (Omoregie 2013)

3. Availability of empty d-orbitals: as we go down from Scandium to Zinc, more orbitals

become available. The more the orbitals available, the more the chances of having more

coordination number. As the shell increases, orbital energy level increases, therefore, the

metal becomes larger and can accommodate more ligands. (Odiaka 2014)

1.2. THIOSALICYLIC ACID

Thiosalicylic acid could be o-thiosalicylic acid also known as 2-Mercaptobenzoic acid with

systematic name, 2-Sulfanylbenzoic acid. Thiosalicylic acid is an organosulfur compound

containing carboxyl and sulfhydryl functional groups. Its molecular formula is C6H4(SH)(CO2H).

It is a yellow solid that is slightly soluble in water, ethanol and diethyl ether, and alkanes, but

more soluble in DMSO. (Lide et al, 2009)

1.2.1. Preparation of Thiosalicylic Acid

Thiosalicylic acid is prepared fromanthranilic acid with IUPAC name 2-aminobenzoic acid via

diazotization.The present invention relates to a method for causing sodium sulfide or a mixture

of sodium sulfide and sulfur to react with diazonium salt formed by diazotizing anthranilic acid,

wherein the Na/S atomic ratio as calculated on the basis of the employed sodium sulfide and

sulfur is adjusted to within the range of 1.33 to 2.0 during the reaction. Thus is made practicable

the manufacture of thiosalicylic acid in a high yield without going through the individual route of

isolating and reducing dithiosalicylic acid.

The present invention relates to a method for efficiently producing thiosalicylic acid which is a

useful intermediate in the manufacture of medicines, pesticides, dyes, etc. More particularly, the

present invention relates to a process for manufacturing thiosalicylic acid from anthranilic acid in

three steps of reactive process.

The process for manufacturing thiosalicylic acid from anthranilic acid as the raw material has

been historically known. This process generally necessitates the following three-step reactive

process. That is to say, firstly anthranilic acid is converted to diazonium salt by reaction with

12

sodium nitrite in hydrochloric acid; secondly, the resulting salt is converted to dithiosalicylic

acid by reaction with an equimolar mixture of sodium sulfide and sulfur; and thirdly and lastly

the resulting dithiosalicylic acid is converted to thiosalicylic acid by reduction by zinc in acetic

acid solvent (Org. Synth. Coll., vol. 2, p 580(1943)). This three-step reactive process is

illustrated in the following diagrams.

Fig 1.1: Three-step reactive process of producing thiosalicylic acid

As for the third step in the above diagram which constitutes the process to reduce dithiosalicylic

acid, there have been proposed some number of various processes, such as the method based on

reduction by a metal selected from among zinc, aluminum, and tin in aqueous solution of an

alkali metal hydroxide,the method based on reduction by a metal like zinc, iron, etc., which

generates hydrogen by reaction with acid, and a hydrogen halide in a lower aliphatic alcohol

13

solvent, and the method based on reduction by hydrogen in the presence of a Raney nickel

catalyst in aqueous solution of an alkali.

Notwithstanding their time-honored positions, those manufacturing methods can be hardly said

satisfactory from the contemporary industrial standpoint, for they disclose certain problematic

aspects. Above all, in those processes for synthesizing thiosalicylic acid from anthranilic acid

some sort of reductive step has to be followed through after the isolation of dithiosalicylic acid.

Obviously, the isolation of dithiosalicylic acid constitutes an indispensable element of the whole

process. Such need to run a multi-step operation no doubt makes those processes less

economical, and, moreover, the reductive step adds to disadvantages due to inevitable

requirements for disposal of waste liquids, waste water, etc. In particular, so long as zinc or

another metal is employed for the reduction, the waste disposal problem involving metal

containing sulfur compounds remains to be resolved.

1.3. Structural review of some metal(II) complexes

The complex forming ability of TSA with many metal ions has been the subject of several

investigations. However, no detailed studies on the complexation equilibria of metal(II) with

TSA has been reported. A fewmetal(II) complexes of thiosalicylic acid have been reported.

A.N Kumar (2009)reported that the complexation equilibria of Cu(II) with thiosalicylic acid

(TSA) have been studied spectrophotometrically in aqueous ethanol (17.08 mol% ethanol)

atI=0.1M (NaClO4) and 25±0.1°C. Analysis of the absorbancevs. pH graphs afforded the

equilibria in solution and the stability constants of the complexes formed. A simple, rapid, and

sensitive method for the spectrophotometric determination of trace amounts of copper is

proposed. The effect of interference of a large number of foreign ions was studied. The method

has been applied successfully to the analysis of some synthetic mixtures and non-ferrous alloys

containing copper.

In this present work, fundamental studies on the complexation equilibria of TSA with copper(II)

in an ethanol-water mixture containing 17.08mol % ethanol at an ionic strength of I = 0.1M

(NaClO4) are presented. The aim of this investigation is to study the equilibria that exist in

solution and to determine the basic characteristics of the complexes formed. The acid

14

dissociation constant of TSA in water-ethanol medium containing 17.08mol %ethanol have been

determined previously. The two acid dissociation constants pKa1 and pKa2 corresponding to the

deprotonation of the carboxyl and sulfydryl groups of TSA are 4.65 and 9.40 respectively.

Humphrey SM (2001)et al reported that the hydrothermal reaction of thiosalicylic acid,

(C6H4(CO2H)(SH)-1,2) with manganese(III) acetate leads to formation of the coordination solid

[Mn5((C6H4(CO2)(S)-1,2)2)4(μ3-OH)2] (1) via a redox reaction, where resulting manganese(II)

centres are coordinated by oxygen donor atoms and S–S disulfide bridge formation is

simultaneously observed. Reaction of the same ligand under similar conditions with zinc(II)

chloride yields the layered coordination solid [Zn(C6H4(CO2)(S)-1,2)] (2). Hydrothermal

treatment of manganese(III) acetate with 2-mercaptonicotinic acid, (NC5H3(SH)(CO2H)-2,3) was

found to produce the 1-dimensional chain structure [Mn2((NC5H3(S)(CO2)-2,3)2)2(OH2)4]·4H2O

(3) which also exhibits disulfide bridge formation and oxygen-only metal interactions.

Compound 3 has been studied by thermogravimetric analysis and indicates sequential loss of

lattice and coordinated water, prior to more comprehensive ligand fragmentation at elevated

temperatures. The magnetic behaviour of 1 and 3 has been investigated and both exhibit

antiferromagnetic interactions. The magnetic behaviour of 1 has been modelled as two corner-

sharing isosceles triangles whilst 3has been modelled as a 1-dimensional chain.

Tewari PK et al (2000) reportedThiosalicylic acid (TSA)-modified Amberlite XAD-2 (AXAD-2)

was synthesized by coupling TSA with the support matrix AXAD-2 through an azo spacer. The

resulting chelating resin was characterized by elemental analyses, thermogravimetric analysis

(TGA) and infrared spectra. The newly designed resin quantitatively sorbs CdII, Co

II, Cu

II, Fe

III,

NiII and Zn

II at pH 3.5–7.0 when the flow rate is maintained between 2 and 4 ml min

−1. The HCl

or HNO3 (2 mol l−1

) instantaneously elutes all the metal ions. The sorption capacity is 197.5,

106.9, 214.0, 66.2, 309.9 and 47.4 μmol g−1

of the resin for cadmium, cobalt, copper, iron, nickel

and zinc, respectively, whereas their preconcentration factor is between 180–400. The

breakthrough volume of HCl or HNO3 for elution of these metal ions was found to be 4–8 ml.

The limit of detection (LOD) for CdII, Co

II, Cu

II, Fe

III, Ni

II and Zn

II was 0.48, 0.20, 4.05, 0.98,

1.28 and 3.94 μg l−1

, respectively, and the limit of quantification (LOQ) was found to be 0.51,

0.29, 4.49, 1.43, 1.58 and 4.46 μg l−1

, respectively. The loading half time, t1/2, for the cations was

found to be less than 2.0 min, except for nickel for which the value was 13.1 min. The

15

determination of each of these six cations is possible in the presence of other five, if their

concentration is up to 4 times. All six metals were determined in river water (RSD ≈ 0.7–7.7%)

and tap water samples (RSD ≈ 0.3–5.7%). The estimation of Co was made in the samples of

multivitamin tablets (RSD <2.3%). The results agree with those quoted by manufacturers.

Najlaa S. Al-Radad (2007) reported that thenew types of binary and ternary CuII complexes have

been synthesized and characterized by elemental analyses, molar conductivities, spectral (I.R.,

U.V.-vis., E.S.R., mass), magnetic, and thermal analyses measurements. The binary CuII

complexes are synthesized by the reaction of thiosalicylic acid (TSA), 2,2′-dipyridyl (dipy),

1,10-phenanthroline (phen), and/or tricine. {N-{tris-(hydroxymethyl) methyl} glycine}(tric) with

CuCl2·2H2O in aqueous ethanol solution (50%). All the above compounds behave as neutral

bidentate ligands coordinating via N, N; O, S; N, N and O, O, respectively. The results of

thermal analyses, elemental analyses, as well as the weight loss methods suggest that the binary

CuII complexes are dimeric in nature and contain at least four H2O molecules inside the

coordination sphere. All the ternary CuII complexes were prepared by adding dl-serine (ser) to

the abovementioned binary CuII complexes in which dl-serine behaves as a bidentate ligand and

substitutes two water molecules from the coordination sphere of the binary complexes to

complete the octahedral structure around the CuII ion together with the cleavage of the dimmer

structure forming two monomers. The isolated solid complexes (binary and ternary) are blue in

color, stable in air, and easily soluble in polar solvents (H2O, EtOH), indicating the electrolytic

nature of these complexes, except the non-electrolyte bluish-grey CuII complex with the general

formula, [Cu2(TSA)(ser)Cl4·(H2O)2]3H2O, which is insoluble in polar solvents. The geometries

of the isolated solid CuII complexes are elucidated from the results of the molar conductivities,

spectral (I.R., electronic, E.S.R., and mass), thermal (T.G.A., D.T.A.), and magnetic

measurements. The room temperature solid state E.S.R. spectra of the binary complexes indicate

the existence of dimeric structures around the CuII ions, while the ternary complexes show

monomer form, except the bluish-grey complex. Also, the E.S.R. spectra of the CuII

complexes

suggest that the copper site has a d ground state.

Israa A. Saeed(2009) reported thata series of new complexes of the type [M(L)Cl2], [M(L)2Cl2]

and [M(L)2]Cl2, where L= L1 or L2 , L1= 2,2`-thiosalyciylic acid disulfide, L2=dibenzyl

disulfide, M=Co(II), Ni(II) and Cu(II), were prepared and characterized by molar conductance,

16

IR, UV/Vis spectral studies, magnetic measurements and metal content analysis. Magnetic

moment and electronic spectra indicate that the some of the complexes show a tetrahedral

geometry and, the others show an octahedral geometry. He reported that Direct reaction of the

ligands 2,2`-disulfide of salicylic acid or dibenzyl disulfide with Co+2, Ni+2 and Cu+2 in

ethanol using 1:1 or 1:2 metal to ligand molar ratio afford complexes of the type [M(L1)Cl2]

,[M(L1)2]Cl2 and [M(L2)2Cl2]. Ligands were coordinate through sulphur atom (according to

Hoiduc and Goh)These reactions indicated the involvement of simple disulfide or polyfunctional

disulfide with metal ions. The prepared complexes are colored solids, stable in air at room

temperature. Molar conductances of the complexes in DMSO are within the range 20-32 ohm-

1.cm

2.mol

-1.

A.N Kumar (2009) reported on the polarographic measurements of the complex formed between

cadmium and thiosalicylic acid. Polarographic measurements on the system Cd-thiosalicylic acid

reveal the formation of a complex of ratio 1:2 metal to ligand (ammonia buffer pH 9.1, µ= 0.2M,

50% alcoholic medium). Potentiometric titrations confirmed this stoichiometry, stepwise

stability constants; logk1 and logk2 being 7.85 and 8.10 respectively. The complex has also been

isolated and studied in solid state. Studies were done in 50% alcoholic medium and ammonia

buffer (pH =9.1). cadmium(II) gives a well-defined single reversible wave.

1.4. Electronic spectroscopy and magnetic moments of some metal(II)

complexes

The tentative assignments of the absorption bands from the electronic spectra of the complexes

have been reported. The π-π* transition in the spectrum of the ligands was observed at 34965-

38461cm-1

, and a bond at 33333-32258 cm-1

which associated with ligand to metal charge

transfer transition. The Co(II) complex showed a band observed in the visible region 14286cm-1,

general considered to correspond to a transition 4A2(F)→

4T1g(P) , this band suggest tetrahedral

geometry while cobalt(II)complex showed a band at 10204 cm-1

which could be assigned to spin-

allowed transition 4T1g(F)→

4T2g(υ1). The position of these bands indicated an octahedral

geometry around Co(II) ion. The nickel complex shows three bands at 10183, 16447and 19011

cm-1

corresponding to the transitions 3A2g(F)→

3T2g(F),

3A2g(F)→

3T1g(F) and

3A2g(F)→

3T2g(P)

respectively , as expected for octahedral Ni2+

ion. The copper complex showed the presence of

17

one band at 12406 cm-1

which was assigned as2T2→

2E which consistent with distorted

tetrahedral geometry.

MadhavanSivasankaran Nair et al (2012) reported that the electronic spectrum of free Schiff base

ligand showed a broad band at 348 nm, which is assigned to π–π∗ transition of the C=N

chromophore. On complexation this band was shifted to lower wavelength region suggesting the

coordination of azomethine nitrogen to the central metal ion.The electronic spectrum of

tetrahedral Co(II) complexes is reported to have only one absorption band in the visible region

due to 4A2(F) →

4T1(P) transition. The spectrum of the present Co(II) complex had only one

band in the visible region at 693 nm, which indicates tetrahedral geometry for the complex. The

electronic spectrum of the Ni(II) complex showed an intense absorption band at 604 nm, which

is due to the 3T1(F) →

3T1(P) transition indicating tetrahedral geometry (Lever, 1984). The

electronic spectrum of Cu(II) complex showed a broad band centered at 616 nm due to

2B1g →

2A1g transition corresponding to square planar geometry. Generally, Zn(II) complexes do

not exhibit any d–d electronic transition due to its completely filled d10

electronic configuration,

however, often exhibit charge transfer spectra. The Zn(II) complex shows an absorption band at

414 nm attributed to the L → M charge transfer transition, which is compatible with this

complex having a tetrahedral geometry (Temel et al., 2002).

Wahab A. Osunniran et al (2004) reported thatthe electronic spectra of Cu(II)complex in

chloroform showed two bands 396(25,253) and 554(18,051). These were assignedto 2Bg→

2Eg

and 2B1g→ 2A1g respectively. The spectra pattern suggested a squareplanar geometry around the

copper(II) ion. The solution spectra data of Co(II) complex revealed three bands: 423(23,641),

470(21,277) and 632(15,823). 470nm and 632nm wereassigned to 4T1g→

4T1g(P) and

4T1g→

4A2g respectively. The band at 423(23,641)is considered as shoulder band and was also assigned

to 4T1g→

4T2g. The spectra of Ni(II) complex consists of two bands in the range between 558nm

and429nm(17,921 and 23,310)cm-1

and were assigned 3T1(F)→

2A2(F) and

3T1(F) →

3T1(P)

respectively. This is a probable indicative of four coordinate squareplanar geometry. The

brownish yellow colour of the complexes, its diamagnetism and theposition of electronic

absorption bands of medium intensity are characteristics of low-spin squareplanar Ni(II)

complexes.Three bands were observed in the spectra of oxovanadium(IV): 556(17,668),

461(21,692)and 400(25,000). These are classified as band II (17,668cm-1) and band III (21,692,

18

25,000)cm-1

and were assigned (b2 – b1*) and (b2 – a1*) transitions respectively. The five-

coordinated Schiffbase complex of VO2+

may have the usual tetragonal pyramidal structure. In

all the complexes Bis(2-hydroxyl-4-methoxyacetophenone)ethylenediimine acts as atetradentate

ligand and binds through the oxygen of the phenoxyl group and nitrogen of C=Nazomethine to

give square planar and tetrahedral geometries while VO(IV) gave tetragonalpyramidal geometry.

Abeer A. Alhadi et al (2011) reported that a new hydrazideShiff base ligand GHL1 (5-bromo-2-

hydroxybezylidene)-3,4,5-trihydroxybenzohydrazide) was prepared by refluxing of

trihydroxybenzhydrazide with an ethanolic of 5-bromo-2-hydroxybenzaldehyde. The ligand

reacted with Ni(II), Cu(II), Zn(II) and Cd(II) (acetate salts). All the complexeswere characterized

by elemental analysis, molar conductivity, TGA, UV-Vis and FT-IR spectral studies. All

thecomplexes have octahedral geometry except for Ni(II) complex which had tetrahedral

geometry. The electronic spectrum of the ligand GHL1 showed two bands at 307 and 340 nm

due to then __> π* transition of the chromophore (-C=N-NH-CO). In the spectra of the complexes,

thesebands were shifted to the lower frequencies which indicated that the imino- nitrogen atom

and theoxygen atom were involved in coordination with the metal ions.

R. H. Holm (2007) reported that the electronic spectrum depends on the energy of metal d

orbital, theirdegeneracy and the number of electrons distributed. These features are in

turncontrolled by the oxidation state of the metal, number and kind of the ligandand the

geometry of the Complexes of Cu(II), Ni(II), Co(II), Mn(II),Zn (II) with furoin-2-

aminothiophenol(FATP),a potential tridendate Schiff baseligand which has been synthesized for

the first time.. The electronic spectral data obtainedwere found to agree with conclusions arrived

from magnetic susceptibilitymeasurements.The expected octahedral transitions of Co(II) are 4T1g

(F) → 4T2g (F),

4T1g (F) →

4A2g (F) and

4T1g (F) →

4T1g(P). The middle band was due to the

transition of two electron which is forbidden and gave a weak band and4A2g (F) and

4T1g(P) are

very close in octahedral geometry. Due to thesefactors detection of middle band is very difficult.

The electronic spectrum ofCo(II) gives two peaks at 1075nm and 442nm due to 4T1g (F) →

4T2g

(F) and4T1g (F) →

4T1g(P) transitions corresponding to octahedral geometry. Thepurple colour

of Co(II) complex is also suggestive of octahedral geometry.Ni(II) complex of FATP exhibit two

19

d-d transitions in the electronicspectra at about 544nm and 978nm due to 3A2g(F) →

3T1g(F) and

3A2g(F) →

3T2g(F) transitions of octahedral geometry. The distorted octahedralgeometry for

Cu(II) complex is indicated by a peak at 666nm. The Zn(II)complexes do not show any

characteristic d-d transition band due to its d10

configuration.

The electronic spectrum of Ni(II) complex displayed two bands in the visible regionobserved at

422 and 626 nm which are assigned to the electronic transitions 3T1(F)→

3T1(P) (ν3)and

3T1(F)→

3T2(F) (ν1), respectively. The band (ν2) is attributed to the transition

3T1

(F)→3A2(F)which corresponds to the charge transfer (C.T.) at 385 nm. The calculated value of

theligand field parameter 10Dq is 19967 cm-1

for (ν1). Thus, the interelectronic repulsion

parameterB was calculated and found to be 116 cm-1

for Ni(II) complex, this value is less than

the freeNi2+

ion value of 1040 cm-1

which was due to overlapping and delocalization of electrons

overthe molecular orbital that encompasses both the metal and ligands. Moreover, the

nephelauxeticratio = B/Bo = 0.11 indicates appreciable covalent character in this complex.

So,the magnetic moment value is 3.4 B.M., which demonstrates that the Ni(II) complex

isparamagnetic and has a high spin tetrahedral configuration with 3T1(F) ground state.

The Co(II) complex has a magnetic moment value of 4.50 BM, which is in agreement with the

reported value for tetrahedral (Kettle, 1969, Cotton and Wilkinson, 1998, Day and Selbin,

1969 and Banerjea, 1998) Co(II) complex. Generally, square planar Ni(II) complexes are

diamagnetic while tetrahedral (Kettle, 1969 and Cotton and Wilkinson, 1998, Banerjea, 1998)

complexes have moments in the range of 3.2–4.1 BM. The Ni(II) complex reported herein has a

room temperature magnetic moment value of 3.33 BM, which is within the normal range

observed for tetrahedral Ni(II) complex. The magnetic moment value of the Cu(II) complex was

observed to be 1.91 BM, which indicates that the complex is monomeric and paramagnetic

(Kettle, 1969, Cotton and Wilkinson, 1998, Day and Selbin, 1969 and Banerjea, 1998). From the

results obtained from elemental analysis, conductance, infrared, electronic and magnetic moment

studies, the proposed geometry of the complexes were assigned. The proposed structure of Schiff

base metal complexes were tetrahedral geometry for Co(II), Ni(II) and Zn(II) complexes and

square planar geometry for Cu(II) complex.

20

Israa A. Saeed(2009) reported from his analysis that the magnetic moment of cobalt complex

was 4.11 B.M corresponding to a high spin tetrahedral Co(II) complexes. Whereas the Mn(II)

complex has a magnetic moment of 4.78 B.M which corresponds to an octahedral geometry. The

magnetic moment of nickel complex was 3.98 B.M which suggests the presence of two unpaired

electrons, corresponding to a tetrahedral geometry. The high value result is probably from an

orbital contribution. The nickel complex has a magnetic moment of 3.10 B.M which corresponds

to an octahedral geometry. The magnetic moment value of copper complex was 1.95 B.M which

is in agreement with distorted tetrahedral geometry.

1.5. Infrared Spectroscopy of Some Metal(II) Complexes

The infrared spectroscopic results provide support for the molecular constitution of these

complexes. The assignments are made on the basis of comparison with the spectra of similar type

of compounds.

From Israa A. Saeed‟s analysis (2009), IR spectra were recorded in the range 4000-250 cm-1

range using CsI discs. The IR spectra of the ligand recorded using CsI, showed medium bands at

467 and 493,1700-1710,3400 cm-1

assigned to υ(S-S), υ(CO) and υ(OH) respectively. The υ(S-S)

stretching vibration has been used as a probe for studying the disulfide group. The change in υ(S-

S) upon coordination vary from 15-35cm-1

. Much larger changes in υ(S-S) were reported by Seff

and et al for complexes of amino alkyldisulfide and pridyl alkyl disulfide The υ(OH) band

remained almost unchanged upon coordination with the metal ions which indicates that this

group is not involved in the coordination. The frequency of υ(CO) band decreased upon

complexation with the metal ions so as the υ(S-S) frequency were shifted to lower position by

15-30cm-1

which indicates the it is shared in coordination. Further support for the formation of

new complexes were provided by the a appearance of a new bands with in the 340-360 and 551-

569 cm-1

range characteristic for υ(M-S) and υ(M-O) respectively. Furthermore the IR spectra of

the complexes show a new band with the range 290-310 cm-1

which may be due υ(M-Cl)

R. H. Holm (2007) reported on representative IR spectrum of the ligand FATP furoin-2-

aminothiophenol and its Cu(II) complex. The selected infrared absorption frequencies of the

ligand and complexes. On complex formation most of the bands in the IR spectrum of the ligand

21

FATP undergo frequency shift and in many cases intensity changes. A strong intense band

approximately at 1676cm-1

in the spectrum of the ligand may be assigned to νC=N stretch. This

band shows a downward shift by about 25-35cm-1

in the spectra of all the metal complexes,

indicating the participation of the azomethine nitrogen in coordination with metals. The

depression in stretching frequency may tentatively attributed to a lowering of the C=N bond

order as a result of the M-N bond formation in the complexes. The shifted band in many cases is

coincident with the C=C band, which then shows greater intensity or broadening. Further

evidence for bonding by nitrogen and oxygen atoms is provided by far IR spectra of complexes.

Due to interference of skeletal vibrations of ligands with M-N and M-O vibrations, definite

assignments of bands are difficult. Therefore only tentative assignments are made on the basis of

information available in literature. Spectra of all complexes showed bands at 586-579cm-1

and

483-478cm-1

which may be assigned to the νM-N and νM-O stretching vibrations .It was

observed that the symmetric vibrations of C-S, which appeared as a band near 701cm-1

in the

ligand spectrum, has been shifted to lower frequencies after complexation. Similarly a weak

band of S-H, which appeared at 2650cm-1

in the case of ligand, has been disappeared in the

spectrum of all complexes. This suggests that the –SH group is involved in coordination. A

broad band at 3450-3400cm-1

in the spectra of several complexes is attributed to the hydroxyl

stretching mode of water molecule. In addition, a medium band approximately at 870-950cm-1

suggests that water molecules are coordinated

Wahab A. Osunniran et al (2004) reported thatthe important bands in the IR spectra of the Schiff

base ligand [Bis(2-hydroxy-4-methoxyacetophenone)ethylenediimine] as well its complexes. In

the IR spectra of the ligand [(2H-4-MA)2en], a weak broad band at 3000 – 2875cm-1

whichhas

been assigned to the intramolecularly hydrogen bonded phenolic-OH in the spectra of theligand

is not observed in the IR spectra of these complexes. There is no corresponding band inthe metal

complexes. A broad band around 3400 – 3500cm-1

in all the complexes can beassigned to the O-

H vibrations of lattice water molecules. The free ligand also exhibited strong band at 1584cm-1

which undergone a slighthypsochromic shift in the complexes and this was assigned to V(C=N)

stretching. This bandwas shifted to higher frequency by 17-23cm-1

as compared to the ligand.

This suggest theinvolvement of azomethine group(C=N) in the coordination with the metal ions

and bondedthrough the Nitrogen atom. The band due to phenolic C-O stretching vibration that

22

appearedin 1336cm-1

in the Schiff base has shifted towards higher frequency by Ca 14-51cm-1

in

the complexes. The positive shift of the band suggests the coordination of the phenolate anions

withthe metal ions via deprotonation. The shift equally confirms the participation of oxygen in

theC-O-M bond. The ring skeletal vibrations (C=C) were consistent in all derivatives

andunaffected by complexation. In the low frequency region, the new absorption bands observed

inthe complexes in the region 476-491cm-1

and 519-521cm-1

were attributed to V(M-phenolic-

O)and (M-N) respectively. The M-O and M-N stretching vibration provide direct evidence forthe

complexation. A very sharp peak at 978cm-1

suggests the presence of V = O bond in

VO(IV)complex. The value is in the range observed for monomeric VO2+

complexes. All the IR

datasuggest that the metal is bonded to Schiff base through the phenolic oxygen and the imino-

Nitrogen.

1.6. AIM OF THE PROJECT

A detailed search through literature has reviewed that very little work has been carried out on

metal complexes derivedfrom thiosalicylic acid andCu(II), Zn(II), Fe(II), Mn(II), Co(II) and

Ni(II) salts

Consequently, the aims and objectives of this project are;

a) Synthesis and isolation of Cu(II), Zn(II), Fe(II), Mn(II), Co(II) and Ni(II) complexes of

thiosalicylic acid

b) The various metal complexes will be characterized by electronic spectroscopy, magnetic

moments, infrared (IR) spectroscopy, metal analysis and melting point measurements.

c) Geometry will be proposed for the metal complexes

23

CHAPTER TWO

THEORETICAL BACKGROUND

2.1. ELECTRONIC SPECTROSCOPY

Spectroscopy is the study of methods of producing and analyzing spectra using spectroscopes,

spectrometers, spectrographs and spectrophotometers. It is the study of interaction of

electromagnetic spectrum. The spectrum is subdivided into regions: radio wave, microwaves,

infrared (IR) radiation, visible light, ultraviolet (UV) radiation, x-rays and gamma rays. The

interpretation of the spectra so produced can be used for chemical analysis, examining molecular

and atomic energy levels and molecular structures and for determining compositions.

2.1.1. ELECTROMAGNETIC RADIATION

Electromagnetic radiation is a fundamental phenomenon of electromagnetism, behaving as

waves and also as particles called photons which travel through space carrying radiant energy. In

physics, all EMR is often referred to broadly as "light," whereas in other colloquial uses, "light"

is reserved for visible light, which is only a very small section of the spectrum of EMR. In

chemistry, the term "light" refers also to those parts of the electromagnetic spectrum that are next

to the visible spectrum, such as ultraviolet and infrared "light." The wavelength λ is the distance

between crests of the wave. The frequency, v, is the number of crests that pass a given point in

one second. Frequency is expressed in Hertz which has unit of 1/sec. frequency times

wavelength equals a constant, the speed of light (vλ = constant). Frequency is inversely

proportional to wavelength, according to the equation: v= fλ where v is the speed of the wave ,f

is the frequency and λ is the wavelength. As waves cross boundaries between different media,

their speeds change but their frequencies remain constant. The energy of the photons is related to

the frequency of radiation by the equation E = hv = h(c λ), where h is Planck's constant, λ is the

wavelength and c is the speed of light. This is sometimes known as the Planck–Einstein

equation. The energy of the photons is thus directly proportional to the frequency of wave and

inversely proportional to wavelength. Likewise, the momentum p of a photon is also proportional

to its frequency and inversely proportional to its wavelength; p= E/c =hf/c = h/ λ(Monk, 2004)

24

2.1.2. INFRARED SPECTROSCOPY

Infrared spectroscopy (IR spectroscopy) is the spectroscopy that deals with the infrared region of

the electromagnetic spectrum, which is light with a longer wavelength and lower frequency than

visible light. For a given sample which may be solid, liquid, or gaseous, the method or technique

of infrared spectroscopy uses an instrument called an infrared spectrometer (or

spectrophotometer) to produce an infrared spectrum. A basic IR spectrum is essentially a graph

of infrared light absorbance (or transmittance) on the vertical axis vs. frequency or wavelength

on the horizontal axis. Typical units of frequency used in IR spectra are reciprocal centimeters

(sometimes called wave numbers), abbreviated as cm−1

. Units of IR wavelength are commonly

given in microns, abbreviated as μm, which are related to wave numbers in a reciprocal way. The

bonds between atoms are usually given as specific lengths, implying rigid bonds between atoms.

The bond vibrates with a frequency that is characteristic of that specific bond. A C-H bond, C-C

single bond and C=C double bond have different vibrational frequencies of range 2850- 3000,

680-700 and 1630-1680 respectively(Paula et al, 2009).

2.1.3. WAVE NUMBERS

In spectroscopy, the wavenumber of electromagnetic radiation is defined as ṽ = 1/λ expressed in

reciprocal centimeters (cm−1

) where λ is the wavelength of the radiation. The wave number range

of IR spectra is from 400 to 4000cm−1

. Larger values of wave numbers represent higher energies

and higher frequencies of vibration. However spectroscopic data are being tabulated in terms of

wave number rather than frequency or energy, since spectroscopic instruments are typically

calibrated in terms of wavelength, independent of the value for the speed of light or Planck's

constant.

2.1.4. MOLECULAR VIBRATIONS

A covalent bond between two atoms can be envisaged as a spring holding them together.If the

bond is compressed, there is a restoring force which pushes the atoms apart, back to the

equilibrium bond length. Molecules undergo two main types of vibrations; stretching and

bending. More energy is required for a stretching vibration than for a bending vibration. When a

25

stretching vibration absorbs IR radiation, the amplitude of the vibration changes but the

frequency of the vibration does not change. If a bond is stretched, there is a restoring force that

forces the atoms back closer together, again restoring the equilibrium bond length. A bending

vibration involves at least three atoms. As an analogy, hold your arms straight out from your

shoulders, parallel to the ground. This is analogous to a symmetric bending (scissoring)

vibration. The frequency of the stretching vibration depends on two factors:

(1) The mass of the atoms

(2) The stiffness of the bond

Heavier atoms vibrate more slowly than lighter ones, so a C-D bond will vibrate at a lower

frequency than a C-H bond. Stronger bonds are stiffer than weaker bonds, and therefore require

more force to stretch or compress them.Thus, stronger bonds generally vibrate faster than

weaker bonds. So O-H bonds which are stronger than C-H bonds vibrate at higher frequencies.

2.1.5. ABSORPTION REQUIREMENTS

IR radiation does not have enough energy to induce electronic transitions as seen with UV.

Absorption of IR is restricted to compounds with small energy differences in the possible

vibrational and rotational states.

For a molecule to absorb IR, the vibrations or rotations within a molecule must cause a net

change in the dipole moment of the molecule. The alternating electrical field of the radiation

(remember that electromagnetic radiation consists of an oscillating electrical field and an

oscillating magnetic field, perpendicular to each other) interacts with fluctuations in the dipole

moment of the molecule. If the frequency of the radiation matches the vibrational frequency of

the molecule then radiation will be absorbed, causing a change in the amplitude of molecular

vibration.

2.1.6. ABSORPTION RANGES

The term "infrared" covers the range of the electromagnetic spectrum between 0.78 and 1000

mm. In the context of infrared spectroscopy, wavelength is measured in wavenumbers, which

26

have the units cm-1

i.e. wavenumber = 1 / wavelength in centimeters. It is useful to divide the

infrared region into three sections; near, mid and far infrared;

Region

Wavelength range (mm)

Wavenumber range (cm-1

)

Near 0.78 - 2.5 12800 - 4000

Middle 2.5 - 50 4000 - 200

Far 50 -1000 200 - 10

The most useful I.R. region lies between 4000 –670cm-1

. Heavier atoms vibrate at lower wave

numbers or frequencies. The positions of the absorption bands for specific functional groups do

not change much in different compounds and are usually found within a given range. For

example, the carbonyl absorptions in aldehydes, ketones, acids and esters are usually found

between 1780 to 1650cm-1

.

2.1.7. ELECTRON EXCITATION

To understand why some compounds are colored and others are not, and to determine the

relationship of conjugation to color, we must make accurate measurements of light absorption at

different wavelengths in and near the visible part of the spectrum. The visible region of the

spectrum comprises photon energies of 36 to 72 kcal/mole, and the near ultraviolet region, out to

200 nm, extends this energy range to 143 kcal/mole. Ultraviolet radiation having wavelengths

less than 200 nm is difficult to handle, and is seldom used as a routine tool for structural analysis.

The energies noted above are sufficient to promote or excite a molecular electron to a higher

energy orbital. When a conjugated molecule absorbs UV radiation, an electron is promoted from

a ground state, bonding orbital into an unfilled higher energy orbital level. Electrons promoted

are usually π (pi) electrons. The UV regions affected by coordination of metal ions include;

n __

> σ*: this transition occurs when atoms in molecule have lone pair of electrons not

involved in the internal bonding. Since the σ bonding level is lower than the non-bonding

level (containing lone pair of electrons). This transition gives the lowest energy

transition. It is found in molecules like alcohols, water, amines, etc.

n __

> π*: this transition occurs in unsaturated heteroatoms that are involved in π-bonding

and also non-bonding electron pair e.g. aldehyde, ketones, esters, etc.

27

π __

> π*: the difference between π and π* energy levels decreases as the extent of

conjugation increases. It usually occur in the UV and visible regions of the spectrum

(10,000- 50,000cm-1

). This transition occur in molecules with double or triple bonds but

no non-bonding electron

2.1.8. TERMINOLOGIES OF ELECTRONIC SPECTROSCOPY

Chromophore: this is the functional group in a molecule that is responsible for particular

absorption. The chromophore is a region in the molecule where the energy difference between

two different molecular orbitals falls within the range of the visible spectrum. Visible light that

hits the chromophore can thus be absorbed by exciting an electron from its ground state into an

excited state.

Auxochrome: this is a group of atoms attached to a chromophore which modifies the ability of

that chromophore to absorb light. An auxochrome is a functional group of atoms with non-

bonded electrons which when attached to a chromophore, alters both the wavelength and the

intensity of absorption.

Hyperchromicity: this involves an increase in the intensity of absorption of light

Hypochromicity: this involves a decrease in the intensity of absorption of light

Hypsochromic shift: this is a change in band position to a shorter wavelength (higher

frequency) which can be as a result of removal of auxochrome or as a result of substitution in a

moleculeor as a result of a change inthe physical conditions.

Bathochromic shift: this is a change in band position to a longer wavelength (lower frequency)

which may be due to the presence of auxochrome or as a result of substitution in a moleculeor a

change in the conditions. (John Daintith, 2008)

28

CHAPTER THREE

METHODOLOGY

3.1. REAGENTS AND SOLVENTS

Reagent grade thiosalicylic acid, manganese(II)chloride tetrahydrate, iron(II)

tetraoxosulphate(VI) heptahydrate, cobalt(II) chloride hexahydrate, nickel(II)chloride

hexahydrate, copper(II)chloride dihydrate, zinc(II)tetraoxosulphate(VI) heptahydrate, ethanol,

methanol, dimethyl sulfoxide, water, diethyl ether, dichloromethane, perchloric acid/ nitric acid,

ethylenediamminetetraacetic acid, ammonia/ ammonium chloride, triethlamine, murezide and

solochrome black T were obtained from Adrich and BDH chemicals and were used as received.

Ethanol, methanol and DMSO were purified by distillation.

3.2. SYNTHETIC METHODS

3.2.1. PREPARATION OF THE METAL(II)COMPLEXES

A solution of metal(II) salts (M= Mn, Co, Fe, Ni, Cu, Zn) in 10ml of methanol was added to a

string solution of the ligand (i.e. 0.514g of MnCl2.4H2O salt was added to 0.8g of thiosalicylic

acid for the preparation of the Mn(II)complex) in 10ml of methanol at room temperature

followed by gradual addition of six dropsof buffer; triethylamine to raise to pH of 8 from pH of

4. The resulting homogenous solution was refluxed for 3 hours during which precipitation

occurred. The resulting precipitate formed was filtered under gravity, stored and dried in

desiccator over activated silica gel. The same method was used for the preparation of Co(II),

Fe(II), Ni(II), Cu(II), Zn(II) complexes from their chloride and sulphate salts respectively

3.3. PHYSICAL MEASUREMENT

3.3.1. MELTING POINT AND DECOMPOSITION TEMPERATURE

The melting points/ decomposition temperatures of the ligand and its complexes were

determined using Gallenkamp Melting Point Apparatus. The results are presented in Table 4.1.

29

3.3.2. INFRARED SPECTRA

The infrared spectra of the ligand and its complexes were recorded as KBr discs on a Perkin-

Elmer FTIR spectrum in the range 4000-400cm-1

at the department of Chemistry, University of

Ibadan. The results are shown in Table 4.4

3.3.3. ELECTRONIC SPECTRA

The UV-Visible spectra of the complexes were recorded using UV-Visible beam PC scanning

spectrophotometer UVD-2960 machine at room temperature in the range of 400 to 900nm at

department of Chemistry, University of Ibadan. The results are shown in Table 4.3.

3.3.4. SOLUBILITY

The solubility of the metal complexes were determined in the following solvents; DMSO, water,

ethanol, dichloromethane, methanol and diethyl ether. The results are shown in Table 4.2

3.3.5. MAGNETIC MOMENT

The room temperature magnetic susceptibilities of the metal complexes were determined using

Sherwood susceptibility balance MSB mark 1 at the Department of Chemistry, University of

Ibadan. The diamagnetic corrections were calculated using Pascal Constants. The results are

recorded in Table 4.1.

3.4. METAL ANALYSIS

A. Preparation of EDTA

3.362g of EDTA was weighed and dissolved in 100ml volumetric flask. This was then made up

to the mark with distilled water.

B. Preparation of 0.005M of ZnSO4.7H20 Solution

0.1438g (0.005mol) of analarzinc(II)sulphate heptahydrate was weighed and transferred into

100ml volumetric flask. It was then dissolved with distilled water and made up to the mark.

C. Standardization of EDTA

25mL of the dissolved zinc(II)sulphate heptahydrate was pipetted into a conical flask of which 2

drops of solochrome black T was added. On addition of the indicator, the colourless solution of

30

the zinc(II)sulphate heptahydrate changed to light purple. The resulting solution was then titrated

against the EDTA solution until a blue end point was observed. The titration was repeated in

triplicates in order to ensure accuracy.

Equation of the reaction:

Zn2+

+ EDTA4-

→ [Zn(EDTA)]2-

1 mol of EDTA4-

= 1mole of Zn2+

i.e. nA= nB

Molarity of ZnSO4 solution (MB) = 0.005M

Volume of ZnSO4 (VB) = 25cm3

Titration result table

Volume (cm3) 1

st titre (cm

3) 2

nd titre (cm

3)

Final burette reading 13.40 26.60

Initial burette reading 0.00 13.50

Titre volume 13.40 13.10

Average volume of EDTA used = (13.40 + 13.10) cm3

2

Volume of EDTA used (VA) = 13.25cm3

To calculate the molarity of EDTA (MA)

Using;

MA × VA = MB × VB

31

MA =MB × VB

VA

= 0.005 × 25

13.25

Therefore, MA = 9.4 × 10-3

M

D. PERCENTAGE METAL DETERMINATION

i. Digestion process

0.01g of each metal complex was weighed into digestion tubes and 5 drops of perchloric acid /

nitric acid (1:1 of HClO4/ HNO3) was added, this was placed on a hot plate and heated to almost

dryness. After the acid treatment, few drops of distilled water was added and heated until the

solution was almost dry. The digested sample obtained were then washed into 100mL standard

flask and made up to the mark with distilled water.

ii. Titrimetric analysis

25mL/10mL of digested samples were pipetted into a conical flask. Few drops of ammonia/

ammonium chloride (NH3/NH4Cl) were added as a buffer and a pinch of murexide (for Mn, Fe,

Co, Ni, Cu metal complexes) or solochrome black T (for Zn metal complex only) was added as

indicator. The standardized EDTA solution was then used to titrate the orange colour Cu(II)

metal complex to a pink colour end-point. This procedure was used for other metal complexes.

The equation of reaction;

M2+

+ EDTA4-

→ [M(EDTA)]2-

The calculation for metal analysis of Cu complex is shown below:

Metal complex = CuTSA

Weight = 0.01g

32

Titration table for [Cu(TSA)ClH20]

Volume (cm3) Rough (cm

3) 1

st titre (cm

3) 2

nd titre (cm

3) 3

rd titre (cm

3)

Final burette

reading

20.40 24.80 26.50 25.60

Initial burette

reading

19.40 23.90 25.60 24.60

Titre volume 1.00 0.90 0.9 1.00

Average volume of EDTA used = (0.90 + 1.00) cm3

2

= 0.95cm3

Number of moles of EDTA used = Volume × Molarity

1000

0.95 × 0.0094

1000

8.93 × 10-6

moles

% of Cu = 8.93 × 10-6

× 63.55 × 4 × 100

0.01

= 23.70%

The experimental value for the complex is 23.70%

The theoretical value

63.55 × 100%

(63.55 + 154.19 + 35.4527 + 18)

=23.43%

33

This procedure of metal analysis was used in calculating the percentage metal for other

complexes. The experimental values obtained were very close in comparison with their

theoretical values and presented in Table 4.1

34

CHAPTER FOUR

RESULTS AND DISCUSSIONS

4.1. GENERAL PROPERTIES OF THE COMPLEX

4.1.1. COLOUR

The complexes synthesized were of different colours due to d-d transition. The results are

presented in Table 4.1.

4.1.2. MELTING POINT/ DECOMPOSITION TEMPERATURE

The ligand and the metal complexes were found to melt/decompose within the range of 162-167

and 200- 265oC respectively thereby showing coordination. The results are presented in Table

4.1.

4.1.3. PERCENTAGE METAL DETERMINATION

All the metal(II) complexes had their percentage metal very close to the theoretical values. Thus

confirming the formulated masses result as shown in Table 4.1.

4.1.4. SOLUBILITY

All the metal complexes were soluble in DMSO while they were slightly soluble in diethyl ether

and water and insoluble in dichloromethane, ethanol and methanol which suggested their likely

polymeric nature. The results are presented in Table 4.1.

4.1.5. MAGNETIC MEASUREMENT

All the complexes were paramagnetic with the exception of Zn(II) complex which was

diamagnetic as expected for d10

configuration.

35

The Fe(II) complex had a moment of 3.80BM which showed equilibrium between tetrahedral

and square planar complex such that the expected magnetic moment for tetrahedral Fe(II)

complex is 4.9BM and the expected magnetic moment for square planar Fe(II) complex is 2.82

BM.

The Co(II) complex had a moment of 1.66BM which was slightly lower than the expected value

of 1.73 BM due to antiferromagnetism and indicative of square planar geometry

The Ni(II) complex had a momentof 2.10BM instead of the expected 2.8 BM due to

antiferromagnetism.The complex was tetrahedral in geometry

The Mn(II) complex had a momentof 1.51BM which was lower than the expected value of 1.73

BM showing antiferromagnetism and corroborative of square planar geometry.

The Cu(II) complexhad a momentof 0.38BM thereby exhibiting antiferromagnetism since the

value is lower than the expected value of 1.73BM.

The Zn(II) complex was expectedly diamagnetic with tetrahedral geometry.The results are

presented in Table 4.4.

4.1.6. ELECTRONIC SPECTRA OF THE LIGANDS

The Mn(II) complex showed a lone band at 18.83kK which was assigned to 2A1→

2B2 transition

of a square planar geometry

The Fe(II) complex exhibited equilibrium between tetrahedral geometry and square planar

geometry with bands at 15.15kK and 13.14kK and were assigned to 5E →

5T2 and

1A1→

1B2

indicative of tetrahedral geometry and square planar geometry respectively.

The Co(II) complex exhibited a band at 13.66kK which indicated a square planar geometry and

was assigned to 2A1→

2B2

36

The Ni(II) complex exhibited 2 bands at 14.85kK and 13.16kK which were suggestive of a

tetrahedral geometry assigned to 3A2→

3T1(P) and

3A2→

3T1 (F) transitions respectively

The Cu(II) complex exhibited a band at 14.85kK which was assigned to 2E →

2T2 which is

indicative of tetrahedral geometry

The Zn(II) complex exhibited a lone band at 13.14kK which was attributed to metal → ligand

charge transfer since no d-d transition was expected,the complex assumed a tetrahedral

geometry.

4.1.7. INFRARED SPECTRA OF THE LIGANDS AND COMPLEXES

The broad band at 3063 cm-1

in thiosalicylic acid was assigned as ν(O-H) band (Table 4.4). This

band remained in the metal complexes but was shifted to 3646-3428cm-1

,indicating

coordination through oxygen atom of the hydroxyl group

The medium and strong band of ν(C=O) in thiosalicylic acid at 1681cm-1

and 1586cm-1

respectively shifted to 1682- 1562cm-1

in the metal complexes, confirming the coordination of the

carbonyl oxygen atom to the metal ions.

Furthermore, the new bands in the range 495-400cm-1

and 397-354cm-1

of which were absent in

thiosalicylic acid, were assigned to v(M-O)/ v(M-S) and v(M-Cl) respectively.This corroborated

coordination of the metals to the ligand.

37

Table 4.1 Analytical data of the ligand and complexes

Compound

(Empirical

formular)

Formular

weight

Colour Melting

point/decom

position

temperature

(oC)

Magnetic

moment

(BM)

% Metal

Theoretical

(Experimental)

% Yield

TSA 154.19 Yellow 162-169

[Mn(TSA)2].3H2O 417.32 Dark

brown

230* 1.51 13.16

13.08

18.43

[Fe(TSA)2].3H2O 418.23 Light

yellow

206* 3.80 13.35

13.12

91.96

[Co(TSA)2].3H2O 421.31 Dark

purple

262* 1.66 13.98

13.84

18.26

[Ni(TSA)2].3H2O 421.08 brown 202* 2.10 13.94

13.79

18.27

[Cu(TSA)ClH2O] 271.24 Greenish

brown

236* 0.38 23.43

23.70

98.30

[Zn(TSA)2].3H2O 427.75 White 240* 0 15.28

15.36

71.93

* = Decomposition temperature

TSA = Thiosalicylic acid

BM = Bohr Magneton

38

Table 4.2 Solubility data of the complexes

Complex H2O DMSO ETOH MeOH CH2CL2 (C2H5)2O

[Mn(TSA)2].3H2O SS (SH) S I I I SS (SH)

[Fe(TSA)2].3H2O SS (SH) S I I I SS (SH)

[Co(TSA)2].3H2O SS (SH) S I I I SS (SH)

[Ni(TSA)2].3H2O SS (SH) S I I I SS (SH)

[Cu(TSA)ClH2O] SS (SH) S I I I SS (SH)

[Zn(TSA)2].3H2O SS (SH) S I I I SS (SH)

S= Soluble

SS= Slightly Soluble

SH= Soluble when heated

I=Insoluble

39

Table 4.3. Electronic spectra

Compounds Absorption region (kK) Transition Geometry

[Mn(TSA)2].3H2O 18.83 2A1→

2B2 Square Planar

[Fe(TSA)2].3H2O 15.15

13.14

5E →

5T2

1A1→

1B2

Tetrahedral

Square Planar

[Co(TSA)2].3H2O 13.66 2A1→

2B2 Square Planar

[Ni(TSA)2].3H2O 14.85

13.16

3A2→

3T1(P)

3A2→

3T1 (F)

Tetrahedral

[Cu(TSA)ClH2O] 14.85

2E →

2T2 Tetrahedral

[Zn(TSA)2].3H2O 13.14 M → LCT Tetrahedral

M = Metal

L =Ligand

CT = Charge Transfer

TSA = Thiosalicylic acid

40

Table 4.4 Infrared spectra of the ligand and metal complexes

Compounds v(O-H) v(C=O) v(C-O) v(C-S) v(M-O)/

v(M-S)

v(M-Cl)

TSA 3063(b) 1681(m)

1586(s)

1267(m) 1039(s)

1151(s)

[Mn(TSA)2].3H2O 3646(m) 1682(m)

1562(s) 1260(s) 1039(s)

1151(s)

400(s) 355(s)

365(s)

373(s)

378(s)

[Fe(TSA)2].3H2O 3565(b) 1682(m)

1561(s) 1270(m) 1038(s)

1151(s)

490(s) 355(s)

369(s)

362(s)

376(s)

[Co(TSA)2].3H2O 3625(b) 1651(m)

1562(m) 1260(m) 1038(s)

1151(s)

402(s) 354(s)

364(s)

377(s)

[Ni(TSA)2].3H2O 3428(b) 1682(m)

1592(s) 1261(s) 1038(s)

1151(m)

413(s) 373(s)

361(s)

385(s)

397(s)

[Cu(TSA)ClH2O] 3444(b) 1562(m)

1587(m) 1260(m) 1038(w)

1151(m)

495(s) 352(s)

368(s)

376(s)

[Zn(TSA)2].3H2O 3478(b) 1634(m)

1568(m) 1270(m) 1058(m)

1122(m)

401(s) 365(s)

355(s)

376(s)

s= stong

w = weak

b= broad

m= medium

41

CONCLUSION

The metal(II) complexes were all found to have a probable 4- coordinate tetrahedral or square

planar geometry. This assignment of geometry was corroborated by infrared, room temperature

magnetic moment and electronic spectra measurement. The ligand coordinated to the metals

through the carbonyl group and hydroxyl group oxygen atoms.The proposed structures of the

complexes are therefore shown below

42

M

OO

O

CC

HSSH

3H20

O

M

OO

O

CC

HSSH

O

M = Mn, Fe, Co, Ni and Zn

Fig 1.2: Proposed structure of the metal complexes

43

Cu

O

OH2

C

SH

Cu

ClO

C

SH

n

Cl

OH2

O

O

Fig 1.3: Proposed structure of the coppercomplex

44

RECOMMENDATION

Further work should be carried out on the following areas;

1. X-ray crystallography measurement should be carried outto confirm the structures.

2. The probable antimicrobial activity should be investigated.

3. In the eventual case of good antimicrobial activity, the toxic nature of the compounds

should be probed in order to find a safe way of usage to fight microorganism at the same

time not being detrimental to human health.

4. The use of these compounds as probable industrial catalysts should also be investigated.

45

The Infrared Spectrum of Thiosalicylic acid

46

The Infrared Spectrum of [Mn(TSA)2].3H2O

47

The Infrared Spectrum of [Fe(TSA)2].3H2O

48

The Infrared Spectrum of [Co(TSA)2].3H2O

49

The Infrared Spectrum of [Ni(TSA)2].3H2O

50

The Infrared Spectrum of [Cu(TSA)ClH2O]

51

The Infrared Spectrum of [Zn(TSA)2].3H2O

52

The UV Spectrum of [Mn(TSA)2].3H2O

53

The UV Spectrum of [Fe(TSA)2].3H2O

54

The UV Spectrum of [Co(TSA)2].3H2O

55

The UV Spectrum of [Ni(TSA)2].3H2O

56

The UV Spectrum of [Cu(TSA)ClH2O]

57

The UV Spectrum of [Zn(TSA)2].3H2O

58

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