Magnetic Resonance Image Segmentation Engineering Project Report for B.Sc. Degree
KOLAWOLE OSOKOYA B.Sc Project
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Transcript of KOLAWOLE OSOKOYA B.Sc Project
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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.
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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
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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
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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
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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
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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
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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
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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
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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
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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
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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
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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
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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
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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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