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¬Preparation And Characterization of Some Uracil Complexes of Nickel (Ii) And Copper (Ii) With Ammonia and 2, 2’-Bipyridine as Secondary Ligands.
1.1 Coordination Compounds And Life Science
Coordination compounds generally known as inorganic complexes are of great practical importance in arts and industries and also have important functional value in nature. A coordination compound contains a central metal atom or ion surrounded by a number of oppositely charged ions or neutral molecules known as ligands. Ligand is an ion or molecule that can have an independent existence. For example, a complex is [Co(NH3)6]3+ in which the Co3+ ion is surrounded by six NH3. The number of groups or ligands surrounding the metal is called the coordination number of the metal atom or ion, and is usually 2, 4, or 6. The group of ligands bonded to the metal taken collectively is said to constitute the metal’s coordination sphere. The resulting conglomeration is often called a complex or if it is charged, a complex ion.
The central metal ion is very often a d-block element (transition metals) or an f-block element (lanthanides and actinides), and thus coordination chemistry is for many scientist a synonym for the chemistry of the transition metals, lanthanides and actinides. However, s-block and p-block elements can also form coordination compounds. The difference between a coordination compound and a organometallic compound is that in an organometallic compound a direct metal-carbon (M-C) bond is present, whereas in a coordination compound there is always a heteroatom (O,N, S, P) located between the metal and the carbon atom or there is no carbon atom present at all. Carbonyl compounds (with CO as the ligand) and cyano compounds (with CN- as the ligand) are borderline cases of compounds with direct metal-carbon bonds. In general, carbonyl complexes are considered as organometallic compounds and cyano complexes as coordination compounds. Due to the presence of a metal ion with unpaired electrons, coordination compounds can have interesting spectroscopic and magnetic properties. Often, coordination compounds are intensively colored.
The structures of coordination compounds depend mostly on the size, charge, and electron configuration of the metal ion and the ligands. They follow the points-on-a-sphere pattern, where orbital overlap (between ligand and metal orbitals) and ligand-ligand repulsions tend to lead to certain regular geometries.
The most observed geometries are listed below, but there are many cases that deviate from a regular geometry [1-2].
• Linear for two-coordination
• Trigonal planar for three-coordination
• Tetrahedral or square planar for four-coordination
• Trigonal bipyramidal or square pyramidal for five-coordination
• Octahedral (orthogonal) or trigonal prismatic for six-coordination
• Pentagonal bipyramidal for seven-coordination
• Square antiprismatic for eight-coordination
• Tri-capped trigonal prismatic for nine coordination.
The phenomena of coordination have tremendous significance in the life sciences. Many of the coordination compounds such as hemoglobin, myoglobin, chlorophyll, cytochromes, metallo-enzymes, metallo-porphyrins-tyrosinase, hemocyanin etc. are in an essential and indispensable way, associated with the chemistry of life. Coordination chemistry also plays an important role in facilitating the expansion of the exciting fields of metallo-therapy and bioinorganic chemistry. Fig 1.1 shows two biologically important coordination compounds.
(a)[2] (b)
Fig. 1.1 Examples of coordination compounds, (a) [Co (NH3)6]3+ and (b) Chlorophyll.
1.2 CHEMISTRY OF NICKEL
Nickel is a silvery, hard, ductile, ferromagnetic, metallic element which resembles iron in its strength and toughness. It occurs in group VIII of the periodic table and is a transition element. Nickel makes up 0.008% of the Earth’s crust [3].
It is used extensively in the chemical and dairy industries for the manufacture of machinery and plant. Nickel-plated steel is also used for the same purposes, as well as for the manufacture of domestic articles such as spoons and forks. Nickel, in the form of a very fine powder called Raney nickel, is of great importance in chemistry [4]. This substance acts as a catalyst in many processes which involve the combination of hydrogen with other substances. Nichrome is an alloy of nickel and chromium which has a great resistance to heat. It is used for the wire which makes up the heating elements of electric fires.
Nickel is present in the body in minute amounts, though its exact role is poorly understood. It is thought to activate certain enzymes and may also play a part in stabilizing chromosomal material in the nuclei of cells. Exposure to nickel may cause dermatitis. Lung cancer has been reported in workers in nickel refineries.
(a) (b)
Fig. 1.2 (a) Pure nickel coins and (b) nickel (II) chloride hexahydrate.
The most common oxidation state of nickel is +2, but compounds of Ni0, Ni+, and Ni3+ are well known, and Ni4+ has been demonstrated. Nickel (II) forms compounds with all common anions, i.e. the sulfide, sulfate, carbonate, hydroxide, carboxylates, and halides. Nickel (III) and nickel (IV) only occurs with fluoride and oxides.
1.3 CHEMISTRY OF COPPER
Copper is a ductile metal with very high thermal and electrical conductivity. Pure copper is soft and malleable; a freshly exposed surface has a reddish-orange color. It is used as a conductor of heat and electricity, a building material, and a constituent of various metal alloys.
Copper is an essential trace element in plants and animals, It is the third most abundant element in the human body, following iron and zinc. It plays an important role in action of a multitude of enzyme that catalyzes a variety of reactions. Copper proteins have diverse roles in biological electron transport and oxygen transportation, processes that exploit the easy interconversion of Cu(I) and Cu(II)[5].
Copper deficiency can produce anemia-like symptoms, neutropenia, bone abnormalities, hypopigmentation, impaired growth, increased incidence of infections, osteoporosis, and abnormalities in glucose and cholesterol metabolism. Conversely, an accumulation of copper in body tissues causes Wilson’s disease [6].
(a) (b)
Fig. 1.3 (a) 99.95% pure copper disc, (b) copper (II) chloride dehydrates.
Copper forms a rich variety of compounds with oxidation states +1 and +2, which are often called cuprous and cupric, respectively. The simplest compounds of copper are binary compounds. The principal ones are the oxides, sulfides and halides.
1.4 PROPERTIES OF URACIL
Uracil is one of the four nucleobases in the nucleic acid of RNA. In RNA, uracil (U) binds to adenine(A) via two hydrogen bonds. In DNA, the uracil nucleobase is replaced by thymine. Uracil is a common and naturally occurring pyrimidine derivative [7].
It is a planar, unsaturated compound that has the ability to absorb light. Uracil undergoes amide-imidic acid tautomeric shifts because any nuclear instability the molecule may have from the lack of formal aromaticity is compensated by the cyclic-amidic stability [8]. The amide tautomer is referred to as the lactam structure, while the imidic acid tautomer is referred to as the lactim structure. These tautomeric forms are predominant at pH 7. The lactam structure is the most common form of uracil.
Uracil readily undergoes regular reactions including oxidation, nitration, and alkylation. Uracil (2, 4-Dihydroxypyrimidine) can be synthesized in the following way [9],
Fig. 1.4 Synthesis process of uracil using urea and ethylacrylate.
Uracil can be used for drug delivery and as a pharmaceutical. When elemental fluorine is reacted with uracil, 5-fluorouracil is produced. 5-Fluorouracil is an anticancer drug (antimetabolite) used to masquerade as uracil during the nucleic acid replication process. Uracil’s use in the body is to help carry out the synthesis of many enzymes necessary for cell function through bonding with riboses and phosphates [7]. Uracil serves as allosteric regulator and coenzyme for reactions in the human body and in plants.
1.5 2, 2?-BIPYRIDINE: ITS CHEMISTRY AND PROPERTIES
2, 2?-Bipyridine (bipy or bpy) is an organic compound with the formula (C10H8N2). Its melting point is 70-73?C and boiling point is 273?C. This colorless solid is an important isomer of the bipyridine family. It is a bidentate chelating ligand, forming complexes with many transition metals. Ruthenium complex and platinum complexes of bipy exhibit intense luminescence, which may have practical applications.
2, 2?-bipyridine is prepared by the dehydrogenation of pyridine using Raney nickel [10].
Although uncoordinated bipyridine is often drawn with its nitrogen atoms in cis conformation, the lowest energy conformation both in solid state and in solution is in fact coplanar, with nitrogen atoms in trans position. Only in acidic solution bipyridine adopts a cis conformation [11].
Bipyridine complexes absorb intensely in the visible part of the spectrum. The electronic transitions are attributed to metal-to-ligand charge transfer (MLCT). In the tris(bipy) complexes three bipyridine molecules coordinate to a metal ion, written as [M(bipy)3]n+ (M = metal ion; Cr, Fe, Co, Ru, Rh and so on; bipy = 2,2?-bipyridine). These complexes have six-coordinated, octahedral structures and two enantiomers as follows:
?-isomer ?-isomer
These and other homoleptic tris-2,2′-bipy complexes of many transition metals are electro active. Under strongly reducing conditions, most tris(bipy) complexes can be reduced to neutral derivatives containing bipy ligands. Examples include M (bipy) 3, where M = Al, Cr, Si. In the following figures several bipyridine complexes with different metals are shown:
1.6 AMMONIA AS A LIGAND
Ammonia or azane is a compound of nitrogen and hydrogen with the formula NH3. It is a colourless gas with a characteristic pungent smell. Ammonia contributes significantly to the nutritional needs of terrestrial organisms by serving as a precursor to food and fertilizers. Ammonia is also a building-block for the synthesis of many pharmaceuticals and is used in many commercial cleaning products. Although in wide use, ammonia is both caustic and hazardous. Ammonia, as used commercially, is often called anhydrous ammonia. It is miscible with water. Ammonia in an aqueous solution can be expelled by boiling. The aqueous solution of ammonia is basic. The maximum concentration of ammonia in water (a saturated solution) has a density of 0.880 g/cm3 [12]. Some properties of NH3 are given below:
Table 1.1 some physical properties of ammonia [13].
Boiling point -33.38?C
Freezing point -77.7?C
Density 0.725 g cm-3 (-70?C)
Permittivity (dielectric constant) 26.7?o (-60?C)
Specific conductance 1× 10-11 ?-1 cm-1
Viscosity 0.254 g cm-1 s-1 (-33?C)
Ion product constant 5.1×10-27 mol2 L-2
Ammonia is the most widely used non-aqueous solvent. It has low dielectric constant result in a generally decreased ability to dissolve ionic compounds, especially those containing high charged ions, e.g., carbonates, sulfates, and phosphates are practically insoluble.
Ammonia can act as a ligand in transition metal complexes. It is a pure ?-donor, in the middle of the spectrochemical series, and shows intermediate hard-soft behaviour. Ammonia is named ammine in the nomenclature of coordination compounds. Some notable ammine complexes include tetraamminediaquacopper (II) ([Cu(NH3)4(H2O)2]2+), a dark blue complex formed by adding ammonia to solution of copper(II) salts. It is known as Schweizer’s reagent. Diamminesilver (I) ([Ag(NH3)2]+) is the active species in Tollens’ reagent.
(a) (b)
Fig. 1.5 Ball-and-stick model of the (a) tetraamminediaquacopper (II) cation, [Cu(NH3)4(H2O)2]2+ and (b) diamminesilver(I) cation, [Ag(NH3)2]+.\
Ammine complexes of chromium (III) were known in the late 19th century, and formed the basis of Alfred Werner’s revolutionary theory on the structure of coordination compounds. Werner noted only two isomers (fac- and mer-) of the complex [CrCl3 (NH3)3] could be formed, and concluded the ligands must be arranged around the metal ion at the vertices of an octahedron. This proposal has since been confirmed by X-ray crystallography.
1.7 METAL COMPLEXES OF URACIL or URACIL DERIVATIVES: A BREIF REVIEW
Gupta and Srivastava[14] on 1983, studied on mixed ligand complexes of Cu(II), Ni(II), Co(II) and Zn(II) formed with glycine and uracil or 2-thiouracil have been synthesized and characterized by elemental analysis, conductance, spectral (IR and electronic spectra) and magneto chemical measurements. Results show that glycine is bidentate in all cases; uracil behaves as a bidentate ligand in Cu (II) complex, coordinating through its one carbonyl oxygen and nitrogen, whereas in other cases it is only monodentate, coordinating only through nitrogen. With thiouracil, coordination occurs from carbonyl oxygen and one nitrogen in Cu (II) and Ni(II) complexes, but in the Co(II) complex coordination occurs from thionyl sulphur and nitrogen. In the Zn(II) complex it shows tridentate behaviour, coordinating through oxygen, sulphur and one nitrogen. Mixed Cu(II), Co(II) and Zn(II) complexes of uracil and of Ni(II) and Zn(II) with thiouracil are octahedral, whereas the mixed Ni(II) complex with uracil shows distorted tetrahedral geometry, and the mixed Co(II)-thiouracil complex is square planar.
R. Ghose[15], on 1989 showed that, mixed complexes of Mn(II), Co(II), Ni(II), Cu(II), Zn(II) and Cd(II) ions with 6-aminopurine (adenine, ADN) and uracil (URL) were prepared from aqueous ethanol solution at pH about 7. The new complexes have the formulae M2(ADNURL)(OH)2•H2O for M = Mn(II), Zn(II), Cd(II) and M2(ADNURL)(OH)4•3H2O for M = Co(II), Ni(II), Cu(II). The mixed complexes were characterized by elemental, infrared, electronic, ESR spectral and magnetic measurements. ADN behaves as a bidentate ligand. The probable binding sites of ADN are N3 and N7 nitrogens and of URL is C2=O. Polymeric structures have been suggested with ADN and OH as the bridging ligands.
Maistralis et al.[16] established in 1999 that complexes of orotic acid (6-uracilic acid) with transition metals (Cu2+, Mn2+, VO2+, Zn2+, Hg2+, Cd2+, Fe3+, Cr3+, and Ag+) have been prepared and characterized by elemental, conductivity, magnetic measurements, IR, NMR and diffuse reflectance spectra. The ligand, in its monoanion
Orotic acid (6-uracilic acid)
Form, coordinates through the carboxylic group to the metal. The Co2+ complex was also isolated under alkaline conditions and studied.
In the year 2000, Sarkar and Mandal [17], worked on mixed ligand oxo-peroxo complexes of vanadium (V), M[VO(O2)L2].nH20 where M= K or NH4; HL= uracil or cytosine and n=1 or 2, have been isolated from aqueous methanolic medium. The complexes were characterized by elemental analysis, conductance, TGA, UV-Visible, IR and NMR spectral studies. Both the peroxide and the other ligands acts as chelates coordinating through their oxygen at C(2) and nitrogen at N(3) and the presence of monomeric oxoperoxovanadium(V) species, have been established by IR, 1H and 51V NMR studies. The complexes appeared to possess pentagonal bipyramidal geometry.
In 2010 Srivastava and Gupta[18] were researched on mixed ligand coordination compound of Palladium having square planner stereochemistry, around the metal ion with the general formula [PdL2Cl2] where L=5methyluracil have been isolated in the solid state by the interaction of with the aforesaid ligands. The synthesized coordination compound has been characterized by elemental analysis, electrical conductance, magnetic measurements, molecular weight determination, electron spin resonance, infra red spectral measurements and NMR studies. They observed that the synthesized compound is light yellow in colour and is non hygroscopic. Also it was diamagnetic in nature. This complex has an anti-tumour activity.
Koz et al.[19], on 2010 was worked on the synthesis, spectroscopic and biological activity studies of Ni(II), Cu(II) and Co(II) complexes of Schiff base ligands derived from 5-aminouracil, 2-hydroxyl-1-naphtaldehyde, 2,4-dihydroxybenzaldehyde and salicylaldehyde. In all cases, the complexes appear to be monomeric.
The ligands coordinate in bidentate fashion to Ni (II) and Co(II) but in a tridentate fashion to Cu(II) by coordinating to the carbonyl oxygen atom in the 4th position of uracil ring. The biological activities of the Schiff bases and metal complexes have been tested in vitro against a number of bacteria and a fungus. Ni(II) complexes derived from the saliciylaldehyde. Schiff base ligand showed good antimicrobial activity whereas a Co(II) complex derived from the same ligand showed good anticandidal activity.
Fig. 1.6 some metal complexes formed by uracil and uracil derivatives.
1.8 OBJECTIVE OF THE PRESENT WORK
Coordination compounds are a very important class of chemicals, because some of them play an essential role in the biochemical processes of living beings. For instance, chlorophyll, hemoglobin and vitamin B12 are coordination compounds. Many enzymes contain a metal ion, and as such they can be considered as coordination compounds.
By studying simple coordination compounds, one can gain insight in the mechanism of complex biochemical processes based on the use of a metal ion inside the cell. Many dyes and pigment, for instance the blue color of writing ink, are metal complexes. Coordination compounds are of importance for medical diagnosis and therapy: contrast agents for magnetic resonance imaging (MRI), the active compounds in chemotherapy and in photodynamic therapy for the treatment of cancer contain a metal ion as an essential component. Metal complexes are being studied as potential new drugs (metallopharmaceutics). A number of catalysts used in the chemical industry make use of coordination compounds.
Thus the study of simple coordination compounds with various ligands at different reaction medium is important. The present research work is associated with the following areas of interest:
(i) Synthesis of metal-uracil complexes with ammonia, 2, 2’-bipyridine as secondary ligands.
(ii) Studies of their various physical properties such as melting point, solubility etc.
(iii) Metal analysis of the complexes to determine their chemical compositions of the product.
(iv) Characterization of the complexes with the help of IR as an aid to elucidate the micro-structural features of bonding.
(vi) UV-visible spectral analysis of the complexes to investigate about the number and kind of ligands as well as the geometry of the complex.
(v) Determination of the magnetic properties of the complexes to obtain information about the oxidation state of metal ion, the number of unpaired electrons and the stereochemistry of the molecule.
(vii) Thermal analysis of the complexes to obtain the idea about the coordinated and crystalline water and the geometry of the complexes.
2. EXPERIMENTS AND RESULTS
2.1 MATERIALS, METHODS AND EQUIPMENTS
The materials, methods and the equipments used to carry out the experimental work of this project are reported in this section.
2.1.1 CHEMICALS
Uracil was produced from LOBA CHEMIE, INDIA. Ammonia and 2,2?-bipyridine are all analar grade and was collected from s.d.fine-Chemlimited, Mumbai, India. The copper chloride dihydrate and nickel chloride hexahydrate were purchased from Qualikems, New Delhi, India and JHD™, China respectively.
2.1.2 MELTIING POINT
Melting points of the compounds were recorded in a paraffin oil bath equipped with a thermometer of capacity of recording the temperature up to 300?C.
2.1.3 SOLUBILITY
The solubility of the compounds was determined qualitatively using various solvents in the usual manner. The solvents used were water, methanol, ethanol, acetone, n-hexane and dilute hydrochloric acid in both cold and hot condition.
2.1.4 ELEMENTAL ANALYSIS
The copper content of the compounds was determined complexometrically using Na2EDTA solution as the titrating agent. A stock solution was prepared by treating a known amount of complex under study with dilute HNO3. The solution was then evaporated to dryness (at least two hours), cooled at room temperature and diluted with distilled water in a volumetric flask.
A known volume of stock solution was taken, buffered with concentrated ammonia to the desired pH (pH = 10) and titrated directly with the standard Na2EDTA solution in the presence of Fast Sulphon Black F indicator. The end point was determined by observing the change of colour of the indicator from wine red to green. Calculation of metal content has been done from the equivalence relationship. The following reactions may be occur in this process:[20]
Cu2+ + H2In3- ? CuH2In-
Indicator Wine red
CuH2In- + HY3- ? CuY2- + H2In3-
EDTA Green
Determination of copper was also carried out iodometrically using the same stock solution. In this method a known volume of stock solution was taken and a few drops of NaHCO3 solution were added to make it basic. The concentrated CH3COOH was added drop by drop to make if acidic again. To it 10 mL of 10% KI solution was added. The resulting solution was kept at dark for 5 minutes while iodine was liberated. This solution was then titrated with standard sodium thiosulphate solution (it was standardized using standard K2Cr2O7 solution) using starch as an indicator. The final color of the solution is colourless. The following reactions may occur here [20]:
Cu2+ + 4I- ? Cu2I2(s) + I2
I- + I2 ? I3-
2S2O32- + I2 ? S4O62- + 2I-
The nickel content of the compound was also determined complexometrically using Na2EDTA solution as the titrating agent. A stock solution was prepared by digestion a known amount of complex under study with dilute HNO3. The solution was then evaporated to dryness, cooled at room temperature and diluted with distilled water in 100 mL volumetric flask.
A known amount of stock nickel solution was taken and about 30 mg murexide indicator and 10 mL of ammonium chloride solution was added to it. Then concentrated ammonia solution was added dropwise until the pH was ~7 as shown by the yellow colour of the solution. Then this solution was titrated with standard EDTA solution until the end point was approached. Finally more 10 mL concentrated ammonia solution was added to the titrating mixture and continued the titration until the colour of the solution changed from orange-yellow to violet.
Calculation of nickel content has been done from the equivalence relationship. The following reactions may be occurred [20]:
Ni2+ + H2In3- ? NiH2In- + H+
Indicator Orange-yellow
NiH2In- + HY3- ? NiY2- + H3In2-
EDTA Violet
The chloride content of the sample was determined gravimetrically using silver nitrate. A known amount of complex under study was dissolved in dilute HNO3. Then a solution silver nitrate solution was added drop wise with constant stirring. The suspension of silver chloride formed was heated with constant stirring until the precipitate coagulated and the supernatant liquid was clear. It was then allowed to stand in dark for 1-2 hours before filtration. The precipitate was filtered in a previously weighed sintered glass crucible. The crucible was heated in an oven at 130-150?C for two hours and cooled in a desiccators and weighed. The difference between this weight and the weight of empty sintered glass filtering crucible gave the weight of AgCl precipitate. Calculation of chloride content has been done from the equivalence relationship. The reaction is following [20]:
Cl- + Ag+ ? AgCl(s)
2.1.5 INFRARED SPECTRA
The infrared spectra (IR) of the complexes were recorded on a Shimdzu (Japan) Infrared Spectrometer of model IR-470 in the range of 4000-400 cm-1 using KBr pallets.
2.1.6 ULTRAVIOLET-VISIBLE SPECTRA
The UV-visible spectra (electronic spectra) of the complexes were recorded using a UV-visible recording spectrometer, Model UV- 1800, Shimadzu (Japan) in the wavelength range 200-1100 nm using water or ethanol as solvents.
2.1.7 THERMAL ANALSIS
In thermogravimetric analysis (TGA) the mass of a sample in a controlled atmosphere (generally in nitrogen or any other inert atmosphere) is recorded continuously as a function of time as the temperature of the sample is increased (usually linear with time). A plot of mass or mass percent as a function of time is called a thermogram or a thermal decomposition curve. The instruments for thermogravimetry consists of : (1) a sensitive analytical balance, (2) a furnace, (3) a purge gas system for providing an inert atmosphere, and (4) a microprocessor or a computer for instrument control and data acquisition and display. Thermogravimetric method is largely used for decomposition and oxidation reactions and also use in vaporization, sublimation and desorption [21].
Differential scanning calorimetry (DSC) is also a thermal technique in which differences in heat flow into a substance and a reference are measured as a function of sample temperature while the TGA is subjected to a controlled-temperature program. This process is mainly used in pharmaceutical industry for testing purity of drug samples. It is also used for studying thermal behaviour of inorganic compounds [21].
2.2 SYNTHESIS AND CHEMICAL COMPOSITION OF URACIL COMPOUNDS
2.2.1 Triammine(uracilato)nickel(II), [Ni(C4H2N2O2)(NH3)3]
In two separate beakers 1 mmol of uracil (C4H4N2O2) solution (0.1121 g in 15-20 mL hot water) and nickel chloride (NiCl2.6H2O) solution (0.2376 g in 5 mL water) were prepared. Both the solutions were mixed. 5-6 mL of concentrated ammonia solution was added to it and heated for 2 hours in water bath until the light green precipitate was started to form. It was kept at room temperature for 2-3 hours for complete precipitation. The light green product was then filtered, washed with water for several times and dried in air and in desiccator over silica gel.
NiCl2.6H2O + C4H4N2O2 + NH3 ? [Ni(C4H2N2O2)(NH3)3]
Uracil
Yield = 0.0567 g.
2.2.2 Diammine-bis(uracilato)copper(II).dihydrate,
[Cu (C4H3N2O2)2(NH3)2].2H2O
In a beaker 1 mmol (0.1705 g) copper chloride (CuCl2.2H2O) was dissolved in 2 mL water and in another beaker 1 mmol (0.1121 g) uracil was dissolved in 15-20 mL hot water. Both the solutions were mixed, and 3-4 mL of concentrated ammonia solution was added to it. The solution was filtered and heated for 1.5-2 hours until the greenish-blue precipitated was produced. The solution was kept at room temperature for 24 hours for complete precipitation. The product was then filtered, washed with water and dried in air first and then over silica gel.
CuCl2.2H2O + C4H4N2O2 + NH3 ? [Cu(C4H3N2O2)2(NH3)2].2H2O
Uracil
Yield = 0.0364 g
2.2.3 Bis-bipyridine-bis(uracil)copper(II) chloride.tetrahyradate,
[Cu(C10H8N2)2(C4H4N2O2)2]Cl2.4H2O
In a beaker 1 mmol of uracil (0.1121 g) was dissolved in 15-20 mL of hot water and in another beaker 1 mmol of copper(II) chloride (CuCl2.2H2O) in 5 mL water was dissolved. The two solutions were mixed thoroughly. 2,2?-bipyridine (0.1565 g) was dissolved in a 20 mL of ethanol and water (50:50 by volume) with vigorous shaking, and 5-6 mL of it was added to the mixture. The resulting solution was filtered for removing the impurities. The clear solution was heated for several hours and while the volume of the solution was 5-10 mL, it was kept for crystallization at room temperature. After 4-6 days, deep blue crystalline product was obtained, separated by filtration, washed with water and dried at air first and then over silica gel.
heat
CuCl2.2H2O + C4H4N2O2 + C10H8N2 ? [Cu(C10H8N2 )2(C4H4N2O2)2]Cl2.4H2O
Uracil 2,2?-bipyridine
Yield = 0.0749 g.
2.3 CHARACTERIZATION AND PROPERTIES
2.3.1 MELTING POINT
Melting occurs when a temperature is reached at which the thermal energy of the particles is high enough to overcome the intra crystalline forces that hold them in position. The melting is a unique property of a substance and may provide an approximate idea about the nature of compound- basically ionic or nonionic. It is also a valuable criterion of purity. The sharpness of melting points usually depends on the purity of the compounds. The melting points of studied compounds are given in table 2.1.
Table 2.1 The melting point of the studied complexes.
Compounds Melting point, oC
[Ni(C4H2N2O2)(NH3)3]
[Cu(C4H3N2O2)2(NH3)2].2H2O
[Cu(C10H8N2)2(C4H4N2O2)2]Cl2.4H2O < 250 (d) 188-190 (d) 154-157 (d) It is evident from Table 2.1 that all three compounds have incongruent melting points, i.e., they are unstable and decompose at temperature before melting. 2.3.2 SOLUBILITY Solubility of a substance means its homogeneous mixture in a solvent consisting of a single phase. Polar substances are soluble in polar solvent and non-polar substance soluble in nonpolar solvent. Thus the solubility of a substance in a particular solvent provides information about the polar character of the substance. The solubility of the studied complexes is given in Table 2.2. Table 2.2 demonstrates that the compound C is soluble in water, acetone and conc. HCl (all in hot state) whereas the compound B is soluble in acetone, n-hexane and conc. HCl and compound A is soluble in only n-hexane and conc. HCl. From the solubility test we can say that all the complexes are mainly basic in nature. Table 2.2 the solubility of studied complexes. Solvent Condition Solubility Compound A Compound B Compound C Water Room temperature Insoluble Insoluble Sparingly soluble Warm Insoluble Insoluble Soluble (blue solution) Methanol Room temperature Insoluble Insoluble Insoluble Warm Insoluble Insoluble Insoluble Ethanol Room temperature Insoluble Insoluble Insoluble Warm Insoluble Insoluble Sparingly soluble Acetone Room temperature Insoluble Sparingly soluble Sparingly soluble Warm Sparingly soluble Soluble (blue solution) Soluble (deep blue solution) n-Hexane Room temperature Sparingly soluble Sparingly soluble Insoluble Warm Soluble Soluble Insoluble Conc. HCl Room temperature Sparingly soluble Sparingly soluble Sparingly soluble Warm Soluble (light green solution) Soluble (blue solution) Soluble (deep blue solution) 2.3.3 ELEMENTAL ANALYSIS Elemental analysis is almost a routine and the first step in characterizing any chemical compound. Determination of different elements presents in any unknown sample yields from which the empirical formula can easily deduced. The copper, nickel and chloride contents of the complexes under study were determined and the results are summarized in the following Table 2.3. Table 2.3 Data for the copper, nickel and chloride contents of the complexes. Compound Metal Chloride Methods of Analysis Average (found), % Calcd., % Methods of Analysis Complexo-metric, % Iodometric, % Gravimetric, % Found, % Calcd, % A 27.85 27.85 26.50 – – B 15.22 14.90 15.06 17.85 C 8.23 7.65 7.94 8.52 11.03 9.50 2.3.4 INFRARED SPECTRA Infrared spectra of complexes do not offer a conclusive result but is complementary to other techniques used for structural characterization. The important observation is that the IR spectrum of a complex molecule consists of a number of characteristic group frequencies which are highly useful in identifying the functional groups present in the complex. This is achieved by comparing the number, position and intensities of various absorption bands observed with those reported in group frequency charts. The abbreviations and symbols are used in Table 2.4-2.6 to present the relative intensity of various absorption frequencies and their associated mode of vibration is the following: v = very ? = stretching s = strong ? = bending m = medium ?’ = deformation w = weak sym = symmetric b = broad asym = asymmetric sh = shoulder arom = aromatic The IR bands assignments have been done or the basis of literature survey[22-24]. 2.3.4.1 IR SPECTRUM OF [Ni(C4H2N2O2)(NH3)3] Infrared spectrum of the above compound is shown in Fig. 2.1 and the peaks with tentative assignments are listed in Table-2.4. Table 2.4 Assignment of IR bands of [Ni(C4H2N2O2)(NH3)3]. Bands, cm-1 Relative Intensity Assignment 3432.39 s,b ?(N-H) of amide combine with ?(N-H) of ammonia 1635.66 s ?(C=O) in amide 1462.07 w ?(C=C)arom 1383.95 m ?s(NH3) ammonia symmetric deformation 1292.33 w ?(C-N)arom 1219.03 vw 1054.12 vw ?(C-H)arom (in plane) 832.3 w ?(C-H)arom (out of plane) 655.81 s (N-H) (rocking) Fig. 2.1 Infra-red spectrum of [Ni(C4H2N2O2)(NH3)3]. The presence of (N-H) bond in both uracil and ammonia gives a very strong and broad absorption band centered at 3432 cm-1. The broadness of this peak arises due to the overlapping of (N-H) stretching vibrations of both. The peak with very strong intensity at 1635 cm-1 is assigned to ?(>C=O) vibrations of uracilato ion. However, the characteristic C=O vibration of secondary amide is usually appear at 1660 cm-1. This difference can be explained by the fact that the uracil molecule is deprotonated and then take part in bonding with the Ni(II) ion through N(1) which reduces the electron density on the neighboring C=O group. Compare to the free uracil molecule the pyrimidine ring vibrations of the complex are shifted significantly to lower frequency region. This supports the formation of the uracilato ion during complexation reaction. Peaks at 1383 cm-1 and 655 cm-1 are due to the symmetric deformation and rocking vibrations of ammonia respectively. Peaks in the region of 1292-1219 cm-1 can be assigned to the ?(C-N) merged with ?(C=O) (in plane) mode. The in plane and out of plane (=C-H) vibrations of uracilato ion appear at 1054 cm-1 and 832 cm-1.
2.3.4.2 IR SPECTRUM OF [Cu(C4H3N2O2)2(NH3)2].2H2O
Fig. 2.2 shows the infrared spectrum of above compound. The frequencies are listed in Table 2.5 with relative intensities and band assignments.
This compound exhibits several peaks at 3600-3400 cm-1. This may be due to the N-H stretching vibration uracilato ion coupled with O-H vibration of lattice water. This characteristic N-H stretching vibration of ammonia molecule observes at 3358 cm-1. The ?(>C=O) frequency of uracil at 1690 is shifted to 1602 cm-1 in the complex. This shifting of C=O vibration implies that the N(1)-H group of uracil molecule is deprotonated to N(1)- before taking part in coordination to the Cu(II) ion. A moderate peak of ?(N-H) of uracilato ion is found at 1555 cm-1.
The characteristic peak of vibrations of uracilato ion appears at 1461 cm-1.
Table 2.5 Assignment of IR bands of [Cu(C4H3N2O2)2(NH3)2].2H2O.
Bands, cm-1 Relative Intensity Assignment
3754.5 vw
3516.29 vs ?(N-H) of 2? amide (CONH-)
3436.24 vs
3358.13 w ?(N-H) of NH3
1602.87 s ?(C=O) of amide
1555.62 ms ?(N-H) of amide
1461.10 s ?(C=C)arom
1376.23 w ?s(NH3) ammonia symmetric deformation
1290.4 ms ?(C-N)arom with ?(C=O) (in plane)
1217.1 m
1068.58 vw ?(C-H)arom (in plane)
856.41 vw
816.87 ms ?(C-H)arom (out of plane)
790.83 ms
703.07 s (N-H) (rocking)
507.29 w ?(M-N)
Fig. 2.2 Infra-red spectrum of [Cu(C4H3N2O2)2(NH3)2].2H2O.
Peaks at 1376 cm-1 and 703 cm-1 are due to the symmetric deformation and rocking vibration of ammonia respectively. Bands in the region of 1290-1217 cm-1 can be assigned to the ?(C-N) merged with ?(C=O) (in plane) mode. Again peaks at 1068 cm-1 and 816 cm-1 are obtained due to in plane and out of plane (=C-H) vibrations of uracilato ion. A peak at 507 cm-1 indicates the ?(M-N) vibrations.
2.3.4.3 IR SPECTRUM OF [Cu(C10H8N2)2(C4H4N2O2)2]Cl2.4H2O
Fig. 2.3 shows the infrared spectrum of the above compound. The frequencies are listed in Table 2.6 with relative intensities and band assignments.
Table 2.6 Assignment of IR bands of [Cu(C10H8N2)2(C4H4N2O2)2]Cl2.4H2O.
Bands, cm-1 Relative Intensity Assignment
3360.05 s ?(N-H) of amide
3029.26 w ?(C-H)arom
1648.2 w ?(C=O) of amide
1607.7 s ?(C=C)arom
1489.07 vw
1433.13 s ?(C=C)arom
1315.47 m ?(C-N)arom
1242.18 mw ?(C-H)arom (in plane)
1163.1 m
1107.16 w
1009.75 s ?(C-H)arom (in plane)
907.52 m ?(C-H)arom (ring breathing mode)
771.54 m ?(M-Cl)
726.21 vw ?(C-H) (out of plane)
619.16 s In-plane ring deformation of bipyridine/ ?(M-O)
486.07 s ?(M-N)
423.38 vw Out-of-plane ring deformation of bipyridine
Fig. 2.3 Infra-red spectrum of [Cu(C10H8N2)2(C4H4N2O2)2]Cl2.4H2O.
This compound exhibits a strong and sharp peak at 3360 cm-1. This may due to N-H streching vibration of amide group of uracil present in the complex. Another peak at 3029 cm-1 indicates the aromatic (C-H) stretching of uracil and bipyridine molecules. A moderate peak at 1648 cm-1 shows the presence of (>C=O) ligated to Cu(II) in the complex. The characteristic aromatic ring stretchings found at 1607 cm-1, 1489 cm-1 and 1433 cm-1. A strong band at 1300 cm-1 can be assigned to the ?(C-N) merged with ?(C=O) (in plane). The C-H (in plane) ring deformation peaks are found at 1242 cm-1, 1163 cm-1, 1107 cm-1 and 1009 cm-1. Again peaks at 907 cm-1 and 726 cm-1 are obtained due to the ring breathing vibrations and out of plane (=C-H) vibrations of both uracil and bipyridine. Peaks at 619 cm-1 and 423 cm-1 are attributed to in-plane ring deformation and out-of-plane ring deformation of 2,2′-bipyridine respectively.
A moderately strong peak at 776 cm-1 assigned for ?(M-Cl) indicates the presence of chlorine in the complex. The ?(M-N) vibrates at 486 cm-1.
2.3.5 ULTRAVIOLET-VISIBLE SPECTRA [25-26]
Ultraviolet-visible spectra are associated with a process of electronic transition from relatively lower energy orbitals in the ground state to higher energy state orbitals in the excited state. The energy of the electronic transition is quantized and is highly dependent on the electronic structure of the molecule. The wavelength at which an absorption maxima is found, depends upon the magnitude of the energy involved for a specific electronic transition. The wavelength of absorption maximum is commonly expressed in millimicron (m?) or nanometer (nm). 1 m? = 1 nm = 10-9 m. The three possible types of electronic transitions are associated with the absorption spectrum of a complex compound are
(i) transition occur within the ligand
(ii) transitions occur within the metal ion (d-d transitions) and
(iii) the charge transfer transitions
Most ligands are organic molecules, which contain the ?, ? and the non-bonding n-electrons. Thus the possible transitions occur within the ligands are ???*, n??*, ???* and n??*, can be observed in the UV and visible region. When a transition metal ion binds to a ligand transition metal complexes are formed. These complexes are coloured, that is they possess the property of absorbing certain wavelengths in the visible region of the spectrum causes promotion of d-electrons from a lower to a higher energy level. These transitions are frequently called d-d transitions and can be explained mainly on the basis of crystal field splitting of the d-orbitals by the ligand field. If the metal is easily oxidizable and the ligand is readily reducible, or vice-versa, then the charge transfer transition may occur. In case of ligands having relatively low lying empty orbitals and unsaturated ligands with empty ? anti-bonding orbitals are present. Charge transfer transitions are intense.
Therefore, the particular spectrum exhibited by complex compound is dependent upon the number and kind of ligands, the energy of d-orbitals of the metal ion; there degeneracy and the number of electrons distributed in them, these features in turn are controlled by the oxidation state of the metal and the geometry of the complexes.
The spectrum of studied [Cu (C10H8N2)2(C4H4N2O2)2] Cl2.4H2O complex is shown in Fig. 2.4 and the ?max for various transitions are summarized in Table 2.7.
Table 2.7 Electronic spectral data of [Cu(C10H8N2)2(C4H4N2O2)2]Cl2.4H2O complex.
Compound ?max, nm
[Cu(C10H8N2)2(C4H4N2O2)2]Cl2.4H2O 320, 604
Ligands, uracil and 2, 2’-bipyridine show characteristics high intensity absorption peaks in the UV-region originated from both ???* and n??* transitions. The bands at shorter wavelength results from the electronic transition in the ethylene bonds of the aromatic ring whereas the bands at longer wavelengths are characteristics of the spectra of heteroatomic molecules called benzoid band. These absorptions are found as a single band at 320 nm in the spectrum of the complexes.
The compound is soluble in hot water giving an intense blue solution. This solution shows a band at 320 nm in the UV-region and a broad band at 604 nm in the visible region. The latter band associated with three nearly superimposed bands (dxz, dyz ? dx2-y2; dz2 ? dx2-y2; dxy? dx2-y2) is a characteristic Jahn-Teller distortion phenomenon that we usually observed in case of the simple octahedral complex. In the octahedral complexes of copper give rise to one band in the visible region near 16000 cm-1 (= 625 nm) which often be resolved into at least three components.
These three components may be assigned as transitions from the dxy, dz2 and dxz, dyz pair to the ? anti-bonding and half filled dx2-y2 level. The relative order of these transitions will depend upon the extent of axial metal-ligand interaction. (a)
(b)
Fig. 2.4 (a) UV and (b) Visible spectrum of [Cu(C10H8N2)2(C4H4N2O2)2]Cl2.4H2O.
The ?max obtained from the visible spectrum yields the value of ?o, the crystal field splitting of the octahedral copper complex. From calculation, the ?o value is 198.76 kJ mol-1. It implies that for d-d transition, about 198.76 kJ mol-1 energy is required. To calculate ?o, the equation is
10Dq = ?o = NE
where E = hc/?
and h = 6.626×10-34 Js, c = 3×108 ms-1,
? = (1/604×10-9 m) = 1.66×106 m-1,
N = 6.023×1023 mol-1
So, E = 6.626×10-34 Js×3×108 m s-1 ×
1.66×106 m-1
= 3.30×10-19 J
Now, ?o = NE = 6.023×1023 mol-1 ×
3.30×10-19 J
?o = 198.759 kJ mol-1
2.3.6 THERMAL ANALYSIS
Thermal stability of a compound is very much dependent on the bond strength of various constituent parts. Thermal decomposition studies, therefore, provide important information with respect to the nature of bonding of various ligands attached to the metal ion especially when the complex is non-volatile. Taken through a wide range of temperatures, a substance may undergo physical and chemical changes; react with the ambient atmosphere or water of crystallization and other fragments. All the changes are accompanied by the absorption or release of energy in the form of heat that giving endothermic or exothermic reaction. The portion removed at lower temperature may either be purely lattice component or it may weakly coordinated to the central metal ion as a monodentate molecule or may be an adduct. The weight, particle size and the mode of preparation of a sample, all govern the thermogravimetric results. In practice, a small weight with a small particle size is portable for TGA and DSC.
2.3.6.1 TGA of [Ni(C4H2N2O2)(NH3)3]
Fig. 2.5.1 represents TGA results of [Ni(C4H2N2O2)(NH3)3] compound. The curve shows that the compound starts loosing weight at relatively low temperature i.e. at 40?C. The percentage of weight loss within the temperature range 50-250?C is about 25%. This may be due to the loss of all ammonia molecules present in the complex (calculated loss 23.74%). After 250?C the weight loss process is continuous and at 325?C the total weight loss becomes 38%. At 650?C the compound losses all of its ligands and converted to black NiO (observed weight 28.68%, calculated weight 26.5%).
Fig. 2.5.1 TGA graph of [Ni(C4H2N2O2)(NH3)3].
2.3.6.2 DSC of [Ni(C4H2N2O2)(NH3)3]
Fig. 2.5.2 represents the DSC curve of [Ni(C4H2N2O2)(NH3)3] The DSC curve was obtained by heating the complex in a flowing stream of nitrogen and the ordinate of this plot is energy input in milliwatts (mW). The two minima at 76.5?C and at 300?C are observed represents the exothermic reactions with the release of -5.07mW and -8.88mW heat energy respectively.
Fig. 2.5.2 DSC curve of [Ni(C4H2N2O2)(NH3)3].
2.3.6.3 TGA of [Cu(C4H3N2O2)2(NH3)2].2H2O
Fig. 2.6.1 represents TGA results of [Cu(C4H3N2O2)2(NH3)2].2H2O compound. The curve shows that the compound starts losing weight at relatively low temperature i.e., at 50?C. About 17% of weight loss is observed within the temperature range 50-250?C. This may be due to the loss of all ammonia and water molecules present in the complex (calculated weight loss 17.43%). After 250?C the weight loss process is continuous and at 300?C the weight loss becomes 30.5%. Up to the temperature reached to 650?C the residue attains a constant weight. The residue is black and identified as CuO (observed 17.50%, calculated 16.25%).
Fig. 2.6.1 TGA graph of [Cu(C4H3N2O2)2(NH3)2].2H2O.
2.3.6.4 DSC of [Cu(C4H3N2O2)2(NH3)2].2H2O
Fig. 2.6.2 represents the DSC curve of [Cu(C4H3N2O2)2(NH3)2].2H2O. The DSC curve was obtained by heating the complex in a flowing stream of nitrogen and the ordinate of this plot is energy input in milliwatts (mW). A maxima at 273?C indicates that an endothermic decomposition reaction is occurred here and the heat required, is 106.61 J/g.
Fig 2.6.2 DSC curve of [Cu(C4H3N2O2)2(NH3)2].2H2O.
2.3.6.5 TGA of [Cu(C10H8N2)2(C4H4N2O2)2]Cl2.4H2O
Fig. 2.7.1 represents the TGA result of the above compound. The curve shows that the compound is thermally stable up to 80?C. After that the process of decomposition of the complex commences and up to 150?C the percentage of weight loss is registered about 10% due to the loss of four water molecules from the complex. Above this temperature the decomposition process is rapid and within the region 157-250?C about 28% of weight loss was observed. At 650?C the compound losses all of its ligands and the residue attains a constant weight. The residue is black and identified as CuO (observed 9.94%, calculated 8.52%).
Fig. 2.7.1 TGA graph of [Cu(C10H8N2)2(C4H4N2O2)2]Cl2.4H2O.
2.3.6.6 DSC of [Cu(C10H8N2)2(C4H4N2O2)2]Cl2.4H2O
Fig. 2.7.2 represents the DSC curve of [Cu(C10H8N2)2(C4H4N2O2)2]Cl2.4H2O. The DSC curve was obtained by heating the complex in a flowing stream of nitrogen and the ordinate of this plot is energy input in milliwatts (mW).
The two minima at about 150?C and at 167?C are observed which represents the exothermic reactions with the release of -3.96mW and -11.84mW heat energy respectively. A maximum at 262?C indicates that an endothermic decomposition reaction is also occurred here and the heat required, is 20.30mW.
Fig 2.7.2 DSC curve of [Cu(C10H8N2)2(C4H4N2O2)2]Cl2.4H2O.
3. Discussion
Reactions of MCl2.xH2O [M = Ni(II) and Cu(II)] with uracil in presence of ammonia or bipyridine in aqueous medium yield corresponding metal complexes having either M(II) ? N or M(II) ? O coordination bonds. The properties of the compounds and the results of various physico-chemical analyses (summarized below) suggested that their chemical and structural formulae are:
Triammine(uracilato)nickel(II), [Ni(C4H2N2O2)(NH3)3] A
Table 3.1 Properties of [Ni(C4H2N2O2)(NH3)3].
Physical Properties
State Solid, powder form
Colour Light Green
Melting Point >250?C (d)
Solubility Soluble in n-hexane and conc. HCl
Ni has d8 configuration and the ligands, ammonia and uracil used are either moderately strong or weak. Under this circumstance, the knowledge of coordination chemistry along with the results of elemental analysis, IR, TGA and DSC studies of the above compound suggested that the nickel complex has a tetrahedral geometry having the following structural formula:
[Ni(C4H2N2O2)(NH3)3]
Diammine-bis(uracilato)copper(II).dihydrate, [Cu(C4H3N2O2)2(NH3)2].2H2O B
Table 3.2 Properties of [Cu(C4H3N2O2)2(NH3)2].2H2O.
Physical Properties
State Solid, powder form
Colour Bluish-green
Melting Point 188-190?C (d)
Solubility Soluble in acetone, n-hexane and conc. HCl
Cu(II) has d9 electronic configuration and the ligands used in synthesis of the above compound are ammonia and uracil. Ammonia is a moderately strong ligand while uracil is a weak ligand according to the spectrochemical series. Knowledge of coordination chemistry, the results of IR, TGA and DSC studies and from the elemental analysis of copper yields the complex has a tetrahedral geometry.
[Cu(C4H3N2O2)2(NH3)2].2H2O
Bis-bipyridine-bis(uracil)copper(II) chloride.tetrahyradate,
[Cu(C10H8N2)2(C4H4N2O2)2]Cl2.4H2O C
Table 3.3 Properties of [Cu(C10H8N2)2(C4H4N2O2)2]Cl2.4H2O.
Physical Properties
State Solid, crystalline form
Colour Deep Blue (or royal blue)
Melting Point 154-157?C (d)
Solubility Soluble in water, acetone and conc. HCl
From the data of elemental analysis, IR, UV-Visible, TGA and DSC studies the following structure of the copper complex is suggested. It is octahedral with two bipyridine molecules are in the equatorial and two uracil molecules occupy the two axial positions around the central Cu (II) ion, shown below:
[Cu(C10H8N2)2(C4H4N2O2)2]Cl2.4H2O
4. REFERENCES
1 S. Z. Haider, “Advanced Inorganic Chemistry”, Dacca, 1st ed., 1975.
2 Shriver & Atkins’, “Inorganic Chemistry”, 5th ed., W. H. Freeman and Company, New York, 2010.
3 “The Encyclopedia of Science” Inorganic Chemistry, Nickel.
4 “LENNTECH” Elements, Nickel.
5 S. J. Lippard and J. M. Berg, “Principles of bioinorganic chemistry”, University Science Books; Mill Valley, 1994.
6 M. Bonham, “British Journal of Nutrition”, 87(5), 393–403, 2002.
7 Garrett, H. Reginald, and Charles M. Grisham, “Principals of Biochemistry with a Human Focus”. United States, Brooks/Cole Thomson Learning, 1997.
8 D.J Brown, “Heterocyclic Compounds: Thy Pyrimidines”. Vol 52. New York: Inter science, 1994.
9 I. L. Finar, “Organic Chemistry” Vol 2, ELBS, Longman, 5th ed., p. 631, 2007.
10 W. H. F. Sasse, “2,2’-Bipyridine”, Org. Synth., Vol. 5, p. 102, 1973.
11 A Göller, U.-W. Grummt, “Chem. Phys. Lett.”, 321, 399-405, 2000.
12 G. L. Miessler and D. A. Tarr “Inorganic Chemistry” Pearson Educational International, 3rd ed., 1990.
13 J. E. Huheey, “Inorganic Chemistry” Harper Collins, New York, 4th ed., 1993.
14 M. Gupta and M. N. Srivastava, “Polyhedron”, 4, 475-479, 1984.
15 R. Ghose, “Inorganica Chimica Acta”, 156, 303-306, 1989.
16 G. Maistralis, A. Koutsodimou and N. Katsaros, “Transition Metal Chemistry”, 25, 166-173, 2000.
17 A. R. Sarkar and S. Mandal, “Metal-Based Drugs”, 7(3), 157-164, 2000.
18 A. Srivastava and D. C. Gupta, “Int. J. Appl. Eng. Res.”, 1(2), 222-227, 2010.
19 G. Koz, H. Kaya, D. Astley, D. Yasa, S. T. Astley, “G. U. J. Sci”., 24(3), 407-413, 2011.
20 A. I. Vogel, “Textbook of Quantitative Inorganic Analysis”, ELBS ed., 4th ed., pp. 321, 322. 367 and 433, 1978.
21 D. A. Skoog, F. J. Holler, and T. A. Nieman, “Principles of Instrumental Analysis”, 5th ed., pp.798, 800 and 805, 1997.
22 K. Nakamato, “Infrared and Raman Spectra of Inorganic and Coordination Compounds”, New York, 4th ed., 1986.
23 C. N. R. Rao, “Chemical Applications of Infrared Spectroscopy”, Academic Press, New Delhi, 1st ed., 1963.
24 D. L. Pavia, G. M. Lampman and G. S. Kriz, “Introduction to Spectroscopy”, 3rd ed., 2001.
24 A. B. P. Lever, “Inorganic Electronic Spectroscopy”, Elsevier Publishing Company, Amsterdam, 1st ed., pp. 384-386, 1968.
26 “The Royal Society of Chemistry”, Modern Chemical Techniques, Ultraviolet/Visible Spectroscopy, p. 92.
27 Wikipedia, the free encyclopedia (www.wikipedia.org).