CHAPTER 1 INTRODUCTION 1.1 GENERAL: The science of solution chemistry is very complex. It requires for its clarification, the help of many branches such as mechanics, electrostatics, and hydrodynamics. In 1887 an organized theoretical and experimental investigation of conducting solution was begun, van't Hoff[l], predicted that solutions which readily conduct electric current possess freezing points, boiling points, osmotic pressures and vapor pressures, characteristics of a special class of a system, and the simultaneous and even more important discovery of Arrhenius[2] that such systems contain electrically charged particles or ions. The powerful methods of thermodynamics were for the first time applied by van't Hoff to solutions in a systematic manner. His treatment, however, lacked the generality, which might have been achieved at that time if the system of thermodynamics developed by by Gibbs [3], ten years earlier had been employed. Gibbs great treatise provided all the essential basic principles required for the thermodynamics of solutions. The most important contribution of thermodynamics has been to reduce all measurements of system in equilibrium to the determination of a single thermodynamic function. Traditional theories of solutions well exclusively on two aspects of liquid mixtures. One of these is the entropy associated with dispersion of the two molecular species, or of their constituent elements in the case of complex molecules among one another. A lattice model often serves as the device for estimating this "combinational" entropy. The other aspect relates to the interactions between neighboring molecules and, in particular, to the difference in the interactions between unlike and like neighbor pairs. Treatment of

l the properties of the liquid mixtures has progressed little beyond the level of interpretation possible within the framework supported by these two considerations alone [4,5]. The equilibrium properties of a liquid are strongly dependent on what may be loosely called its local structure, often expressed in terms such as packing density, free volume or more exactly in terms of the radial distribution function. In as much as this local structure depends on the forces between molecules and on the forms and volumes of the molecules, in general, it will change with the composition. This change in turn will be reflected in the thermodynamic properties of the mixture. Contributions of this nature have either been ignored altogether, or correction to a state of null volume change on mixing has been adopted as a means of compensating for the effects referred to [5]. This device alters the various properties of the mixtures altered by the changes they would sustain if the volumes were adjusted to the value linearly interpolated between the volumes of the pure components. It will be apparent, however, that adjustments of one thermodynamic quantity (e.g. volume) in this manner will not, in general, effect a simultaneous correction of others (e.g. the free energy or that part of it relating to the local structure) to their linearly interpolated values. The choice of volume as the property to be "conserved" is arbitrary, and there is no assurance that nullifying volume change obviates consideration of other characteristic properties of the liquid. Interest in the thermodynamic properties of binary liquid mixtures extended over most of the 19th century. Binary liquid mixtures are often used as solvents for chemical equilibria and for media in which to carry out chemical reactions. Binary liquid mixtures can be broadly classified into simple mixtures and

2 associated mixtures. Simple mixtures are defined as those in which the nominal components are identical with the actual species in the pure liquids as well as in the mixtures. The associated mixtures are defined as those in which there is an evidence for the association of at least one of the components either with each other (self association) or with those of another component. It is possible to approach the description of such mixtures in terms of the properties of the components in a formal thermodynamic manner or with respect to the molecular interactions, such as dipole-interactions, hydrogen bonding, and coordinative bond formation etc. 1.2 DENSITY: The density of a material is defined as its mass per unit volume. The symbol of density is p (the Greek letter rho). In some countries (for instance, in the United States), density is also defined as its weight per unit volume [6]. Mathematically: Density = Mass / Volume m p"v Where, p (rho) is the density, m is the mass, V is the volume. Different materials usually have different densities, so density is an important concept regarding buoyancy, metal purity and packaging. In some cases density is expressed as the dimensionless quantities specific gravity (SG) or relative density (RD), in which case it is expressed in multiples of the density of some other standard material, usually water or air/gas. For a homogeneous object, the mass divided by the volume gives the density. The mass is normally measured with an appropriate scale or balance; the volume may be measured directly (from the geometry of the object) or by the

3 displacement of a fluid. Hydrostatic weighing is a method that combines these two[7]. The density of a solution is the sum of mass (massic) concentrations of the components of that solution. Mass (massic) concentration of a given component p; in a solution can be called partial density of that component. The SI unit for density is: • kilograms per cubic metre (kg/m3) Densities using the following metric units all have exactly the same numerical value, one thousandth of the value in (kg/m3). Liquid water has a density of about 1 kg/dm3, making any of these SI units numerically convenient to use as most solids and liquids have densities between 0.1 and 20 kg/dm3[8]. • kilograms per cubic decimeter (kg/dm3) • grams per cubic centimeter (g/cc, gm/cc or g/cm3) • mega grams per cubic meter (Mg/m3) Liters and metric tons are not part of the SI, but are acceptable for use with it. Since 1 L = 1 dm3, we also have these of the same size: • kilograms per liter (kg/L) • grams per milliliter (g/mL) • metric tons per cubic meter (t/m3) 1.2.1 Changes of density: In general, density can be changed by changing either the pressure or the temperature. Increasing the pressure will always increase the density of a material. Increasing the temperature generally decreases the density, but there are notable exceptions to this generalization. For example, the density of water increases between its melting point at 0 °C and 4 °C; similar behaviour is observed in silicon at low temperatures.

4 The effect of pressure and temperature on the densities of liquids and solids is small. The compressibility for a typical liquid or solid is 10~6 bar-1 (1 bar=0.1 MPa) and a typical thermal expansivity is 1CT5 K"1. In contrast, the density of gases is strongly affected by pressure. 1.3 EXCESS VOLUMES OF BINARY LIQUID MIXTURES: "Of the various thermodynamic functions for the mixing process, the volume change on mixing at constant pressure is one of the most interesting, yet certainly still one of the least understood." Hildebrand and Scott[5] wrote this in 1962 at a time when the subject of excess volumes of binary liquid mixtures was entering a phase of renewed interest and development. This renewed interest is closely tied to the advances in the theories of solutions which were made in the 1960's. For a long time the precision attainable on measuring volume changes on mixing was several orders of magnitude better than theory could calculate. The first breakthrough in the corresponding states theories was in predicting accurately the sign of the volume change on mixing. Current theories of solution are much more accurate with respect to sign and magnitude, and this has spawned a vigorous expansion in techniques and measurement. Volume changes on mixing have been determined mainly by two principle methods, a) by directly mixing the liquids and observing volume changes in dilatometer and b) indirectly by measuring the density of liquid mixtures. The volume change on mixing at constant pressure for the binary liquid mixtures is one of the most interesting thermodynamic functions of mixing. This property remains to be of interest because the experimental procedures are relatively easy to perform with great precision and

5 because the volume change on mixing is a sensitive indicator to the accuracy of theories of solution. 1.4 VISCOSITY : Viscosity is a measure of the resistance of a fluid which is being deformed by either shear stress or tensile stress. In everyday terms (and for fluids only), viscosity is "thickness". Thus, water is "thin", having a lower viscosity, while honey is "thick", having a higher viscosity. Put simply, the less viscous the fluid is, the greater its ease of movement (fluidity)[9]. Viscosity describes a fluid's internal resistance to flow and may be thought of as a measure of fluid friction. All real fluids have some resistance to stress and therefore are viscous, but a fluid which has no resistance to shear stress is known as an ideal fluid or inviscid fluid. In general, in any flow, layers move at different velocities and the fluid's viscosity arises from the shear stress between the layers that ultimately opposes any applied force. 1.4.1 Viscosity coefficients: Viscosity coefficients can be defined in two ways: • Dynamic viscosity, also absolute viscosity, the more usual one; • Kinematic viscosity is the dynamic viscosity divided by the density. Viscosity is a tensorial quantity that can be decomposed in different ways into two independent components. The most usual decomposition yields the following viscosity coefficients: • Shear viscosity, the most important one, often referred to as simply viscosity, describing the reaction to applied shear stress; simply put, it is the ratio between the pressure exerted on the surface of a fluid, in the lateral or horizontal direction, to the

6 change in velocity of the fluid as you move down in the fluid (this is what is referred to as a velocity gradient). • Volume viscosity or bulk viscosity, describes the reaction to compression, essential for acoustics in fluids, see Stokes' law (sound attenuation). Alternatively, • Extensional viscosity, a linear combination of shear and bulk viscosity, describes the reaction to elongation, widely used for characterizing polymers. 1.4.2 Viscosity measurement: Viscosity is measured with various types of rheometers. Close temperature control of the fluid is essential to accurate measurements, particularly in materials like lubricants, whose viscosity can double with a change of only 5 °C. For some fluids, it is a constant over a wide range of shear rates. These are Newtonian fluids. The fluids without a constant viscosity are called non-Newtonian fluids. Their viscosity cannot be described by a single number. Non-Newtonian fluids exhibit a variety of different correlations between shear stress and shear rate. One of the most common instruments for measuring kinematic viscosity is the glass capillary viscometer. 1.4.3 Units: Dynamic viscosity: The usual symbol for dynamic viscosity used by mechanical and chemical engineers — as well as fluid dynamicists — is the Greek letter mu (//)[10-13]. The symbol // is also used by chemists, physicists, and the IUPAC[13]. The SI physical unit of dynamic viscosity is the pascal-second (Pa-s), (equivalent to N-s/m , or kg/ms). If a fluid with a viscosity of one Pas is placed between two plates, and one plate is pushed

7 sideways with a shear stress of one pascal, it moves a distance equal to the thickness of the layer between the plates in one second. The cgs physical unit for dynamic viscosity is the poise(P)[14], named after Jean Louis Marie Poiseuille. It is more commonly expressed, particularly in ASTM standards, as centipoise (cP). Water at 20 °C has a viscosity of 1.0020 cP or 0.001002 kilogram/meter second. 1 P= 1 g-cnfV. The relation to the SI unit is lP = 0.1Pa-s, 1 cP=l mPas = 0.001 Pas. The cgs physical unit for kinematic viscosity is the stokes (St), named after George Gabriel Stokes. It is sometimes expressed in terms of centistokes (cSt or ctsk). In U.S. usage, stoke is sometimes used as the singular form. 1 St=lcmV = 10"4mV. 1 cSt= 1 mmV = 10"6mV. Water at 20 °C has a kinematic viscosity of about 1 cSt. The kinematic viscosity is sometimes referred to as diffusivity of momentum, because it has the same unit as and is comparable to diffusivity of heat and diffusivity of mass. It is therefore used in dimensionless numbers which compare the ratio of the diffusivities. /. 4.4 Viscosity of binary liquid mixtures: The prediction of the viscosity of liquid mixtures is a goal of long standing, with both theoretical and practical importance. A truly fundamental theory would predict the viscosity, along with other thermodynamic and transport properties, from knowledge of the intermolecular forces and radial distribution function alone. Such a program has an appreciable success in application to pure, simple

8 liquids such as the liquefied rare gases [15]. For solutions, however, although the general theory has been formulated[16], it has not yet been reduced successfully to numerical results. One is thus forced to approximate approaches of which two general types may be distinguished. The first is that of continuum hydrodynamics whose application to molecular problem is identified with the names of Einstein and Stokes. This approach, in which the discrete molecular nature of the solvent is neglected, has been remarkably successful in explaining the viscosity of dilute solutions of high polymers. Its application to solutions in which both components are of comparable size is less appropriate. The second general approach is to correlate the viscosity of the mixtures with properties of the pure components and with thermodynamic parameters characteristics of the interactions between components. In simplest form, the interactions may be neglected and simple additive relation assumed. One such relation commonly used is,

In r| =Xjlnr|j + x2lnr|2 —(1.1)

where r| is the viscosity of the solution, Xi, x2, r\{ and r|2 are the mole fractions and viscosities of the two components, respectively, in a binary mixture. However, it will be noted that this equation and other similar ones, predict monotonic increase or decrease of viscosity with composition. Such monotonicity is often not observed in practice. It is, therefore, generally necessary to take explicit account of interactions. There are two major semi-empirical theories of liquid viscosity. The first is the absolute reaction rate theory of Eyring et al.[17] This relates the viscosity to the free energy needed for a molecule to

9 overcome the attractive force field of its neighbour, so that it can jump (flow) to a new equilibrium position. Thus, the deviation of the mixture viscosity from equation 1.1 should be related to the free energy, more precisely, the excess free energy of mixing[18,19]. The second semi-empirical theory is the free volume theory[20-22], which relates the viscosity to the probability of occurrence of an empty neighbouring site into which a molecule can jump. This probability is exponentially related to the free volume of the liquid, so deviations from equation 1.1 can be attributed to variations in the free volume of the solution. Macedo and Litovitz[23] have argued that neither of these theories are entirely adequate in its own right and that better results for pure liquids are obtained by combining them. Thus, the probability for viscous flow is taken as the product of the probabilities of acquiring sufficient activation energy and of the occurrence of an empty site. It has been shown[24], that the assumptions of these two probabilities are independent, and may be multiplied, is incorrect for one-dimensional liquids. While this conclusion also presumably holds in three dimensions, it will be of interest to investigate the consequences of the assumptions of independent probabilities for viscosities of liquid mixtures in the treatment that follows. The free energy and free volume data, required in these viscosity theories may occasionally be available directly from experiment. This is often not the case, however, particularly for the free volume whose operational definition is in any event somewhat ambiguous [23]. They have, therefore, preferred to obtain the thermodynamic parameters from a statistical thermodynamic theory of liquid mixtures proposed by Flory et al.[25-28]. This theory, which

10 bears some similarities to the corresponding states theory of Prigogine et al.[29], treats the properties of mixtures in terms of reduced properties of the pure components and a single interaction parameter. It thus, has a minimum of adjustable parameters. The theory is applicable to mixtures of molecules of different sizes and lays particular stress on equation of state contributions to the thermodynamic excess functions, which are of central importance in determining the free volume of the solution. Use of theoretical expressions for the thermodynamic properties of mixtures has the further advantage that an equation for the intrinsic viscosity or the rate of change of viscosity with concentration at infinite dilution of one component can be explicitly formulated by expanding the thermodynamic functions in a power series in concentration and keeping only the first non-trivial terms. Several empirical and semi-empirical relations have been used to represent the dependence of viscosity on concentration of components in binary liquid mixtures[30]. 1.5 INTERMOLECULAR FORCES: Forces binding atoms in a molecule are due to chemical bonding. The energy required to break a bond is called the bond- energy. The forces holding molecules together are generally called intermolecular forces. The energy required to break molecules apart is much smaller than a typical bond-energy, but intermolecular forces play important roles in determining the properties of a substances. Intermolecular forces are particularly important in terms how molecules interact and form biological organisms or even life. This link gives an excellent introduction to the interactions between molecules.

11 As far as the intermolecular forces are concerned, there is little difference between the simple fluids (i.e. liquefied inert gases), which are monoatomic, and spherically symmetrical non-polar molecular fluids, such as methane, nitrogen or carbon dioxide. These liquids differ from other non-polar liquids with non-spherical molecules or large non-spherical molecules which are highly polarizable, where the intermolecular forces are not central. Because of the short range of the forces, interaction occurs mainly between adjacent segments of neighbouring molecules, rather than between the molecules as entities. Strong intermolecular forces that lead to self-association of liquids may be of several kinds. Dipole-dipole interactions may lead to pair-wise association, but rarely to higher aggregates. Hydrogen bonding may lead to dimers primarily, to chain like oligomers. The properties that will be bulk thermodynamics and mechanical properties of liquids, and the static optical and electrical properties, as well as the relevant molecular properties (size, shape, charge, polaizability, dipole moment). The structural features that can be discussed are mainly the distances to the nearest- and next-nearest- neighbours and the appropriate coordination numbers. These properties and features depend on the condition under which the liquid is studied: the temperature and the pressure, which, in turn, determine the density. All matter is held together by force. The force between atoms within a molecule is a chemical or intramolecular force. The force between molecules is a physical or intermolecular force. The intermolecular forces (forces between molecules) are weaker than intramolecular forces i.e. the chemical bonds within an individual molecule. The properties of matter result from the

12 properties of the individual molecule (resulting from chemical bonding) and how the molecules act collectively (resulting from intermolecular forces). Intermolecular Forces are longest-ranged (act strongly over a large distance) when they are electrostatic. An uncharged molecule can still have an electric dipole moment. Electric dipoles arise from opposite but equal charges separated by a distance. Molecules that possess a dipole moment are called polar molecules. The force may be understood by decomposing each of the dipole into two equal but opposite charges and adding up the resulting charge-charge forces. The forces can be attractive or repulsive depending on whether like or unlike charges are closer together. On average, dipoles in a liquid orient themselves to form attractive interactions with their neighbours, but thermal motion makes some instantaneous configurations exist fleetingly that are, in fact, repulsive. /. 5. / Classifying Intermolecular Forces: In general, intermolecular forces can be divided into several categories. The prominent types are: 1.5.1.1 Strong ionic attraction: Recall lattice energy and its relations to properties of solid. The more ionic, the higher the lattice energy. Examine the following list and see if you can explain the observed values by way of ionic attraction. 1.5.1.2 Intermediate dipole-dipole forces: Substances whose molecules have dipole moment have higher melting point or boiling point than those of similar molecular mass, but their molecules have no dipole moment.

13 1.5.1.3 Weak London dispersion forces or van der Waal's force: These forces always operate in any substance. The force arisen from induced dipole and the interaction is weaker than the dipole-dipole interaction. In general, the heavier the molecule, the stronger the van der Waal's force of interaction. For example, the boiling points of inert gases increase as their atomic masses increases due to stronger Landon dispersion interactions. 1.5.1.4 Hydrogen bond:

Certain substances such as H20, HF, NH3 form hydrogen bonds, and the formation of which affects properties (mp, bp, solubility) of substance. Other compounds containing

OH and NH2 groups also form hydrogen bonds. Molecules of many organic compounds such as alcohols, acids, amines, and aminoacids contain these groups, and thus hydrogen bonding plays a important role in biological science. 1.5.1.5 Covalent bonding: Covalent is really intramolecular force rather than intermolecular force. It is mentioned here, because some solids are formed due to covalent bonding. For example, in diamond, silicon, quartz etc., the all atoms in the entire crystal are linked together by covalent bonding. These solids are hard, brittle, and have high melting points. Covalent bonding holds atoms tighter than ionic attraction. 1.5.1.6 Metallic bonding: Forces between atom in metallic solids belong to another category. Valence electrons in metals are rampant. They are not restricted to certain atoms or bonds. Rather they run

14 freely in the entire solid, providing good conductivity for heat and electric energy. These behaviour of electrons give special properties such as ductility and mechanical strength to metals. Intermolecular forces also play important roles in solutions, which is given in Hydration, solvation in water. 1.6 HOMO/LUMO: HOMO and LUMO are acronyms for highest occupied molecular orbital and lowest unoccupied molecular orbital, respectively. The difference of the energies of the HOMO and LUMO, termed the band gap, can sometimes serve as a measure of the excitability of the molecule: the smaller the energy, the more easily it will be excited. The HOMO level is to organic semiconductors and quantum dots what the valence band is to inorganic semiconductors. The same analogy exists between the LUMO level and the conduction band. The energy difference between the HOMO and LUMO level is regarded as band gap energy. When the molecule forms a dimer or an aggregate, the proximity of the orbitals of the different molecules induces a splitting of the HOMO and LUMO energy levels. This splitting produces vibrational sublevels which each have their own energy, slightly different from one another. There are as many vibrational sublevels as there are molecules that interact together. When there are enough molecules influencing each other (e.g. in an aggregate), the number of sublevels are large enough to be perceived as a continuum rather than discrete levels. We no longer consider energy levels, but energy bands [31].

15 1.6.1 Identifying the HOMO and the LUMO: To identify the HOMO and the LUMO of a given molecule, one can 1. find out all the molecular orbitals and fill them with the available electrons, or 2. use a generic ordering of orbitals, and use valence bonding for a-type and lone pair orbitals, and molecular orbitals for n- systems as an approximation. The generic ordering of molecular orbitals (from highest energy to lowest energy): • a* - almost never occupied in the ground state • 7i* - very rarely occupied in the ground state • n - nonbonding (ie. lone pairs) • 7i - always occupied in compounds with multiple bondT^PT*" Bond) • a - at least one occupied in all molecules (Sigma Bond) • a - an empty orbital The kinds of electrons present are identified through observation of the Lewis structure. Although this generic system will identify the correct HOMO and LUMO most of the time, there are exceptions to this classification. To identify any resonances in the system, the entire TT system must be considered. 1.6.2 SOMO: A SOMO is a singly occupied molecular orbital and defined as half the HOMO of a free radical[32]. It shows the difference of energies between orbitals. 1.7 ETHERS: Ether is a class of organic compounds that contain an ether group — an oxygen atom connected to two alkyl or aryl groups — of general formula R-0-R[33]. A typical example is the solvent and

16 anesthetic diethyl ether, commonly referred to simply as "ether" (CH3-CH2-O-CH2-CH3). Ethers are common in organic chemistry and pervasive in biochemistry, as they are common linkages in carbohydrates and lignin. Ethers feature C-O-C linkage defined by a bond angle of about 120° and C-0 distances of about 1.5 A. The barrier to rotation about the C-0 bonds is low. The bonding of oxygen in ethers, alcohols, and water is similar. In the language of valence bond theory, the hybridization at oxygen is sp3. Oxygen is more electronegative than carbon, thus the hydrogens alpha to ethers are more acidic than in simple hydrocarbons. They are far less acidic than hydrogens alpha to ketones, however. Ether molecules cannot form hydrogen bonds amongst each other, resulting in a relatively low boiling point compared to that of the analogous alcohols. The difference, however, in the boiling points of the ethers and their isometric alcohols become smaller as the carbon chains become longer, as the van der waals interactions of the extended carbon chain dominate over the presence of hydrogen bonding. Ethers are slightly polar, as the C-O-C bond angle in the functional group is about 110 degrees, and the C-0 dipoles do not cancel out. Ethers are more polar than alkenes but not as polar as alcohols, esters, or amides of comparable structure. However, the presence of two lone pairs of electrons on the oxygen atoms makes hydrogen bonding with water molecules possible. Cyclic ethers such as tetrahydrofuran and 1,4-dioxane are miscible in water because of the more exposed oxygen atom for hydrogen bonding as compared to aliphatic ethers[34]. Ethers in general are of low chemical reactivity, but they are more reactive than alkanes. Although ethers resist hydrolysis, they are cleaved by mineral acids such as hydrobromic acid and hydroiodic acid.

17 Hydrogen chloride cleaves ethers only slowly. Some ethers rapidly cleave with boron tribromide to give the alkyl bromide[35]. Depending on the substituents, some ethers can be cleaved with a variety of reagents, e.g. strong base. 1.7.1 Uses of Ether: The following are the general uses of ether, • Diethyl ether is used in anaesthia. • In Grignard reaction, it is used as a solvent. • Used as a solvent for resins, oils, fats, and gums. • Used for refrigerator purposes. 1.7.2 2-Methoxyethanol: 2-Methoxyethanol, or methyl cellosolve, is an organic compound that is used mainly as a solvent. It is a clear, colorless liquid with an ether-like odor. It is in a class of solvents known as glycol ethers which are notable for their ability to dissolve a variety of different types of chemical compounds and for their miscibility with water and other solvents. It can be formed by the nucleophilic attack of methanol on protonated oxirane followed by proton transfer. 2-Methoxyethanol is used as a solvent for many different purposes such as varnishes, dyes, and resins. It is also used as an additive in airplane deicing solutions [3 6]. 2-Methoxyethanol is toxic to the bone marrow and testicles. /. 7.3 Cyclohexane: Cyclohexane is a cycloalkane with the molecular formula

C6Hi2. Cyclohexane is used as a nonpolar solvent for the chemical industry, and also as a raw material for the industrial production of adipic acid and caprolactam, both of which are intermediates used in the production of nylon. On an industrial scale, cyclohexane is

18 produced by reacting benzene with hydrogen. Due to its unique chemical and conformational properties, cyclohexane is also used in labs in analysis and as a standard. Pure cyclohexane in itself is rather unreactive, being a non- polar, hydrophobic hydrocarbon. It can react with very strong acids such as the superacid system HF + SbF5 which will cause forced protonation and "hydrocarbon cracking". Substituted cyclohexanes, however, may be reactive under a variety of conditions, many of which are important to organic chemistry. Cyclohexane is highly flammable[37]. Cyclohexane is also used for calibration of Differential scanning calorimetry (DSC) instruments, because of a convenient crystal-crystal transition at -87.1 °C[38]. 1.7.4 Ethanolamine: Ethanolamine, also called 2-aminoethanol or monoethanol- amine (often abbreviated as EA or MEA), is an organic chemical compound that is both a primary amine (due to an amino group in its molecule) and a primary alcohol (due to a hydroxyl group). Like other amines, monoethanolamine acts as a weak base. Ethanolamine is a toxic, flammable, corrosive, colorless, viscous liquid with an odor similar to that of ammonia. Ethanolamine is commonly called monoethanolamine or EA in order to be distinguished from diethanolamine and triethanolamine. Ethanolamine is the second- most-abundant head group for phospholipids, substances found in biological membranes[39]. EA is used in aqueous solutions for scrubbing certain acidic gases. It is used as feedstock in the production of detergents, emulsifiers, polishes, pharmaceuticals, corrosion inhibitors, chemical intermediates[40,41]. In pharmaceutical formulations, EA is primarily

19 used for buffering or preparation of emulsions. Ethanolamine is often used for alkalinization of water in steam cycles of power plants, including nuclear power plants with pressurized water reactors. 1.7.5 1, 4-Dioxane: 1, 4-Dioxane, often called dioxane because the other isomers of dioxane are rare, is a colorless heterocyclic organic compound, which is a liquid at room temperature and pressure. It is classified as an ether. This colorlesss liquid is mainly used as a stabilizer for the solvent trichloroethane. It is an occasionally used solvent for a variety of practical applications as well as in the laboratory. It has a faint sweet odor similar to that of diethyl ether. Dioxane is an occassionally controversial byproduct of the ethoxylation process in the production of materials used in cosmetics, notably sodium myreth sulfate and sodium laureth sulfate [42]. Dioxane is primarily used as a stabilizer for 1,1,1-trichloro­ ethane for storage and transport in aluminium containers [43]. Apart from its use as a stabilizer, dioxane is used in a variety of applications as a solvent, e.g. in inks and adhesives. Dioxane is relatively nonpolar but has superior dissolving power relative to diethyl ether. Diethyl ether is rather insoluble in water, whereas dioxane is miscible and in fact is hygroscopic. It is a versatile polar aprotic solvent. The oxygen atom is Lewis basic, so it is able to solvate many inorganic compounds. Because of its lower toxicity, it is sometimes substituted with tetrahydrofuran (THF) in some processes. However, it has a higher boiling point, which is important when reactions are to be conducted at a higher temperature. Dioxane is a relatively nontoxic substance with an LD50 of 5170 mg/kg[43]. This compound is irritating to the eyes and respiratory tract. It is suspected of causing damage to the central

20 nervous system, liver and kidneys[44]. Accidental worker exposure to 1,4-dioxane has resulted in several deaths[45]. Dioxane is classified by the IARC as a Group 2B carcinogen: possibly carcinogenic to humans because it is a known carcinogen in animals[46]. The U.S. Environmental Protection Agency classifies dioxane as a probable human carcinogen[47,48]. Dioxane can contaminate cosmetics and personal care products such as deodorants, shampoos, toothpastes and mouthwashes[49-51]. The ethoxylation process makes the cleansing agents, such as sodium lauryl sulfate, less abrasive and offers enhanced foaming characteristics [52- 5 5 ]. 1.7.6 Tetrahydrofuran: Tetrahydrofuran (THF) is a colorless, water-miscible organic liquid with low viscosity at standard temperature and pressure. This heterocyclic compound has the chemical formula (CH2)40. As one of the most polar ethers with a wide liquid range, it is a useful solvent. Its main use, however, is as a precursor to polymers. THF has an odor similar to its chemical cousin, diethyl ether, but is a much less potent anesthetic than diethyl ether. THF can be polymerized by strong acids to give a linear polymer called poly glycol[56]. The other main application of THF is as an industrial solvent for PVC and in varnishes[57]. It is an aprotic solvent with a dielectric constant of 7.6. It is a moderately polar solvent and can dissolve a wide range of nonpolar and polar chemical compounds[58]. THF is water-miscible, and can form solid clathrate hydrate structures with water at low temperatures[59]. THF can be used in hydroboration reactions to synthesize primary alcohols, and as a solvent for organometallic compounds

21 such as organolithium and Grignard reagents[60]. Although similar to ether, THF is a stronger base [61]. THF is considered a relatively nontoxic solvent, it penetrates the skin causing rapid dehydration. THF readily dissolves latex and is typically handled with nitrile or neoprene rubber gloves. It is highly flammable. The greatest danger posed by THF follows from its tendency to form highly-explosive peroxides on storage in air. To minimize this problem, commercial samples of THF are often inhibited with BHT. THF should not be distilled to dryness, because the explosive peroxides concentrate in the residue. 7.7.7 Diethyl ether: Diethyl ether, also known as ether, is the organic compound with the formula (C2H5)20. It is a colorless and highly flammable liquid with a low boiling point and a characteristic odor. It is the most common member of a class of chemical compounds known generically as ethers. It is a common solvent and was once used as a general anesthetic. Ether is sparingly soluble in water (6.9 g/100 mL).It is particularly important as a solvent in the production of cellulose plastics such as cellulose acetate [62]. Diethyl ether has a high cetane number of 85 - 96 and is used as a starting fluid for diesel and gasoline engines[63], because of its high volatility and low autoignition temperature. For the same reason it is also used as a component of the fuel mixture for carbureted compression ignition model engines. Diethyl ether is a common laboratory solvent. It has limited solubility in water, thus it is commonly used for liquid-liquid extraction. Being less dense than water, the ether layer is usually on top. Diethyl ether is a common solvent for the Grignard reaction, and for many other reactions involving organometallic reagents [64]. A

22 cough medicine called Hoffmann's Drops was marketed at the time as one of these drugs, and contained both ether and alcohol in its capsules[65]. Ether tends to be difficult to consume alone, and thus was often mixed with drugs like ethanol for recreational use. Ether may also be used as an inhalant. In the 19th century and early 20th century ether drinking was popular among Polish peasants [66]. It is traditional and still relatively popular recreational drug among Lemkos [67]. Water and peroxides can be removed by either distillation from sodium and benzophenone, or by passing through a column of activated alumina[68]. /. 7.8 Diphenyl ether: Diphenyl ether is the organic compound with the formula

0(C6H5)2. The molecule is subject to reactions typical of other phenyl rings, including hydroxylation, nitration, halogenation, sulfonation, and Friedel-Crafts alkylation or acylation. The main application of diphenyl ether is as a eutectic mixture with biphenyl, used as a heat transfer medium. Such a mixture is well-suited for heat transfer applications because of the relatively large temperature range of its liquid state.Diphenyl ether is a starting material in the production of phenoxathiin via the Ferrario reaction [69]. Phenoxathiin is used in polyamide and polyimide production [70]. Because of its odor reminiscent of scented geranium, as well as its stability and low price, diphenyl ether is used widely in soap perfumes[71]. Diphenyl ether is also used as a processing aid in the production of polyesters[72,73]. 1.7.9 Diisopropyl ether: Diisopropyl ether is secondary ether that is used as a solvent. It is a colorless liquid that is slightly soluble in water, but miscible with most organic solvents. It is also used as an oxygenate gasoline

23 additive. Diisopropyl ether is sometimes represented by the abbreviation "DIPE". Diisopropyl ether tends to form explosive peroxides upon standing in air for long periods (years). This reaction proceeds more easily than for ethyl ether, due to the secondary carbon next to the oxygen atom, which makes storage of diisopropyl ether more dangerous. The stored solvent should therefore be tested for the presence of peroxides more often (recommended once every 3 months for diisopropyl ether vs. once every 12 months for ethyl ether)[74]. For safety reasons, methyl tert-butyl ether is often used as an alternative solvent. 1.7.10 Methyl tert-butyl ether: Methyl tert-butyl ether, also known as methyl tertiary butyl ether and TBME, is a chemical compound with molecular formula C5H12O. TBME is a volatile, flammable and colorless liquid that is immiscible with water. TBME has a minty odor vaguely reminiscent of diethyl ether, leading to unpleasant taste and odor in water. TBME is a gasoline additive, used as an oxygenate and to raise the octane number, although its use has declined in the United States in response to environmental and health concerns. It has been found to easily pollute large quantities of groundwater when gasoline with TBME is spilled or leaked at gas stations. TBME is also used in organic chemistry as a relatively inexpensive solvent with properties comparable to diethyl ether but with a higher boiling point and lower solubility in water. It is also used medically to dissolve gallstones. TBME is almost exclusively used as a fuel component in engine gasoline. It is one of a group of chemicals commonly known as oxygenates because they raise the oxygen content of gasoline. Due

24 to its higher solubility in water TBME moves more quickly than other fuel components [75,76]. TBME gives water an unpleasant taste at very low concentrations, and thus can render large quantities of groundwater non-potable[77]. The United States Environmental Protection Agency (EPA) has concluded that available data are not adequate to quantify health risks of TBME at low exposure levels in drinking water, but that the data support the conclusion that TBME is a potential human carcinogen at high doses [78-80]. 1.8 SPECTROSCOPY : Spectroscopy is the study of the interaction between radiation (electromagnetic radiation, or light, as well as particle radiation) and matter. Spectrometry is the measurement of these interactions; a machine which performs such measurements is a spectrometer or spectrograph. A plot of the interaction is referred to as a spectrogram, or, informally, a spectrum. The spectroscopic properties of atoms and molecules result from interaction with electromagnetic waves. Atoms of molecules in their normal state are in the most stable electronic, vibrational and rotational state. In the stable state the atom or molecule possesses a particular minimum energy E0. By interaction with radiations, atoms or molecules may undergo transitions into energy enriched, excited states. These then posses a higher energy written as E*. The energy

AE = E*- E0 is needed to bring about this transitions from a stable ground state to excited state. This transition is known as excitation and so the energy needed is called excitation energy. This energy is supplied by one quantum of radiation and is absorbed by an atom or molecule. When energy associated with one quantum is smaller than excitation energy (AE) no excitation takes place. Two or more

25 photons with a lowest energy than that required for excitation can not combine their energies to produce the AE. Photon with the energy hv striking atoms or molecules which required the energy AE = hv to get excited are absorbed by these atoms or molecules. Thus the intensity of radiations decreases after passing through it. The measurement of radiation intensity during irradiation of a substance as a function of wavelengths, wave number or radiation energy results in the absorption spectrum of that substance. In an atom, electrons occupy only particular energy levels known as atomic orbital. Thus atoms can be excited only those photons which supply the energy difference between these energy levels. Absorption can occur only at particular energies or wavelengths. Therefore absorption spectra are line spectra. For spectroscopic analysis of organic molecules ultraviolet, infrared, microwave and NMR spectroscopy are found to useful. They provide rapid and exact information with instruments which can be operated easily. The various branches of spectroscopy involve measurement of two important experimental parameters, namely, the energy of the radiation absorbed or emitted by the system and the intensity of spectral lines. Most of the spectroscopic studies are done under two important heads i.e. a) atomic spectroscopy and b) molecular spectroscopy. (a) Atomic Spectroscopy: It is concerned with the interaction of electromagnetic radiations with atoms which are most commonly in their lowest energy state known as ground state. In this case transitions of electron occur from one electronic energy level to another.

26 (b) Molecular Spectroscopy: It is concerned with the interaction of electromagnetic radiations with molecules. In this case transitions occur between rotational and vibrational energy level in addition to electronic transitions. The spectra obtained are therefore quite complicated in this case than those of atomic spectroscopy. Molecular spectra extend from visible to microwave region and it can provide information about molecular rotations and vibrations to reveal a great deal about molecular structure. It is possible to get a lot of information about molecular structure from atomic and molecular spectra. The introduction of quantum mechanics made possible the quantitative treatment of spectroscopic. 1.8.1 Infrared spectroscopy: Infrared spectroscopy (IR spectroscopy) is the subset of spectroscopy that deals with the infrared region of the electromagnetic spectrum. It covers a range of techniques, the most common being a form of absorption spectroscopy. As with all spectroscopic techniques, it can be used to identify compounds and investigate sample composition. A common laboratory instrument that uses this technique is an infrared spectrophotometer. The infrared portion of the electromagnetic spectrum is usually divided into three regions; the near-, mid- and far- infrared, named for their relation to the visible spectrum. The far-infrared, approximately 400-10 cm-1 (1000-30 um), lying adjacent to the microwave region, has low energy and may be used for rotational spectroscopy. The mid-infrared, approximately 4000-400 cm-1 (30-2.5 um) may be used to study the fundamental vibrations and associated rotational- vibrational structure. The higher energy near-IR, approximately 14000-4000 cm-1 (2.5-0.8 um) can excite overtone or harmonic

27 vibrations. The names and classifications of these subregions are merely conventions. They are neither strict divisions nor based on exact molecular or electromagnetic properties[81]. Infrared spectroscopy exploits the fact that molecules absorb specific frequencies that are characteristic of their structure. These absorptions are resonant frequencies, i.e. the frequency of the absorbed radiation matches the frequency of the bond or group that vibrates. The energies are determined by the shape of the molecular potential energy surfaces, the masses of the atoms, and the associated vibronic coupling. In order for a vibrational mode in a molecule to be "IR active," it must be associated with changes in the permanent dipole. A molecule can vibrate in many ways, and each way is called a vibrational mode. Linear molecules have 3N-5 degrees of vibrational modes whereas nonlinear molecules have 3N-6 degrees of vibrational modes (also called vibrational degrees of freedom). As an example

H20, a non-linear molecule, will have 3><3-6 = 3 degrees of vibrational freedom, or modes. The simplest and most important IR bands arise from the "normal modes," the simplest distortions of the molecule. In some cases, "overtone bands" are observed. These bands arise from the absorption of a photon that leads to a doubly excited vibrational state. Such bands appear at approximately twice the energy of the normal mode. Some vibrations, so-called 'combination modes," involve more than one normal mode. The phenomenon of Fermi resonance can arise when two modes are similar in energy, Fermi resonance results in an unexpected shift in energy and intensity of the bands.

28 Absorptions bands : r--r C ° C "C C N C N C C 4000 N H O H 3200 2800 2300 2100 1800 1500 Fingerprint l^t-H Attachedto hdwjj| M C-H | Triples frapnj| Oingtea J

! 2.380 I 1460.1380 ' CITI"1 4000 3000 COa 2000 nujcl 1000 Wave numbers listed in cm-1 Fig. 1.1: Absorption bands Uses and applications: Infrared spectroscopy is widely used in both research and industry as a simple and reliable technique for measurement, quality control and dynamic measurement. It is also used in forensic analysis in both criminal and civil cases, enabling identification of polymer degradation. Infrared spectroscopy has been highly successful for applications in both organic and inorganic chemistry. Infrared spectroscopy has also been successfully utilized in the field of semiconductor microelectronics [62,82]. For example, infrared spectroscopy can be applied to semiconductors like silicon, gallium arsenide, gallium nitride, zinc selenide, amorphous silicon, silicon nitride, etc. An IR is widely used for: 1) identification and determination of structure 2) determining the purity & quantitative analysis 3) following the course of reaction 4) hydrogen bonding 5) molecular geometry & conformational analysis 6) chemistry of organic polymers and 7) reactions of reactive species like free radicals and ions. The stretching frequency of chemical bond depend upon i) bond strength ii) reduced masses of the atoms forming bond. Any

29 factor which will increase bond strength will increase stretching frequency of the bond, and if the mass of atoms forming the chemical bond is increased, the reduced mass will increase and stretching frequency will decrease 1.8.1.1 Factors affecting on IR Frequencies : Inductive effect: The inductive effect in Chemistry and Physics is an experimentally observable effect of the transmission of charge through a chain of atoms in a molecule by electrostatic induction [83]. The net polar effect exerted by a substituent is a combination of this inductive effect and the mesomeric effect. The electron cloud in a a-bond between two unlike atoms is not uniform and is slightly displaced towards the more electronegative of the two atoms. This causes a permanent state of bond polarization, where the more electronegative atom has a slight negative charge(8-) and the other atom has a slight positive charge(5+). If the electronegative atom is then joined to a chain of atoms, usually carbon, the positive charge is relayed to the other atoms in the chain. This is the electron-withdrawing inductive effect, also known as the -I effect. Some groups, such as the alkyl group are less electron- withdrawing than hydrogen and are therefore considered as electron- releasing. This is electron releasing character is indicated by the +1 effect. As the induced change in polarity is less than the original polarity, the inductive effect rapidly dies out, and is significant only over a short distance. The inductive effect is permanent but feeble, as it involves the shift of strongly held a-bond electrons, and other stronger factors may overshadow this effect. The inductive effect may

30 be caused by some molecules also. Relative inductive effects have been experimentally measured with reference to hydrogen. Inductive Effect can also be used to determine whether a molecule is stable or unstable depending on the charge present on the atom under consideration and the type of groups bonded to it. For example, if an atom has a positive charge and is attached to a -I group its charge becomes 'amplified' and the molecule becomes more unstable than if I-effect was not taken into consideration. Similarly, if an atom has a negative charge and is attached to a +1 group its charge becomes 'amplified' and the molecule becomes more unstable than if I-effect was not taken into consideration. But, contrary to the above two cases, if an atom has a negative charge and is attached to a -I group its charge becomes 'de-amplified' and the molecule becomes more stable than if I-effect was not taken into consideration. Similarly, if an atom has a positive charge and is attached to a +1 group its charge becomes 'de-amplified' and the molecule becomes more stable than if I-effect was not taken into consideration. The explanation for the above is given by the fact that more charge on an atom decreases stability and less charge on an atom increases stability. Polarity induced in a covalent bond due to the difference in electronegativities of the bonded atoms is called Inductive Effect. 1.8.1.2 Resonance: In chemistry, resonance or mesomerism[84], is a way of describing delocalized electrons within certain molecules or polyatomic ions where the bonding cannot be expressed by one single Lewis formula. A molecule or ion with such delocalized electrons is represented by several contributing structures[85].

31 Each contributing structure can be represented by a Lewis structure, with only an integer number of covalent bonds between each pair of atoms within the structure[86]. These individual contributors cannot be observed in the actual resonance-stabilized molecule; resonance is not a rapidly-interconverting set of contributors. Several Lewis structures are used collectively to describe the actual molecular structure. The actual structure is an approximate intermediate between the canonical forms, but its overall energy is lower than each of the contributors. This intermediate form between different contributing structures is called a resonance hybrid [87]. Contributing structures differ only in the position of electrons, not in the position of nuclei. Resonance is a key component of valence bond theory. Electron derealization lowers the potential energy of the substance and thus makes it more stable than any of the contributing structures. The difference between the potential energy of the actual structure and that of the contributing structure with the lowest potential energy is called the resonance energy [88,89] or derealization energy. Resonance is distinguished from tautomerism and conformational isomerism, which involve formation of isomers. One contributing structure may resemble the actual molecule more than another (in the sense of energy and stability). Structures with a low value of potential energy are more stable than those with high values and resemble the actual structure more. The most stable contributing structures are called major contributors. The greater the number of contributing structures, the more stable the molecule. This is because the more states at lower energy are available to the electrons in a particular molecule, the more stable the electrons are.

32 Also the more volume electrons can occupy at lower energy the more stable the molecule is [90]. Equivalent contributors contribute equally to the actual structure; those with low potential energy (the major contributors) contribute more to the resonance hybrid than the less stable minor contributors. Especially when there is more than one major contributor, the resonance stabilization is high. High values of resonance energy are found in aromatic molecules. 1.8.1.3 Steric effects: The steric effect of tri-(tert-butyl)amine makes electrophilic reactions, like forming the tetraalkylammonium cation, difficult. It is difficult for electrophiles to get close enough to allow attack by the lone pair of the nitrogen. Steric effects arise from the fact that each atom within a molecule occupies a certain amount of space. If atoms are brought too close together, there is an associated cost in energy due to overlapping electron clouds (Pauli or Born repulsion), and this may affect the molecule's preferred shape (conformation) and reactivity. Steric hindrance or steric resistance occurs when the size of groups within a molecule prevents chemical reactions that are observed in related smaller molecules. Although steric hindrance is sometimes a problem, it can also be a very useful tool, and is often exploited by chemists to change the reactivity pattern of a molecule by stopping unwanted side-reactions (steric protection). Steric hindrance between adjacent groups can also restrict torsional bond angles. However, hyperconjugation has been suggested as an explanation for the preference of the staggered conformation of ethane because the steric hindrance of the small hydrogen atom is far too small[91-93].

33 Steric attraction occurs when molecules have shapes or geometries that are optimized for interaction with one another. In these cases molecules will react with each other most often in specific arrangements. Steric effect can even induce a mechanism switch in the catalytic reaction[94]. A special computational procedure was developed to separate electronic and steric effects of an arbitrary group in the molecule and to reveal their influence on structure and reactivity [95]. The conjugations with aromatic ring results in the delocalisation of n electrons which reduces the double bond character or strength of double bond and therefore drop in IR frequency is observed. 1.8.2 Ultraviolet-visible spectroscopy: Ultraviolet-visible spectroscopy or ultraviolet-visible spectrophotometry (UV-Vis or UV/Vis) refers to absorption spectroscopy in the ultraviolet-visible spectral region. This means it uses light in the visible and adjacent (near-UV and near-infrared (NIR)) ranges. The absorption in the visible range directly affects the perceived color of the chemicals involved. In this region of the electromagnetic spectrum, molecules undergo electronic transitions. This technique is complementary to fluorescence spectroscopy, in that fluorescence deals with transitions from the excited state to the ground state, while absorption measures transitions from the ground state to the excited state[96]. UV/Vis spectroscopy is routinely used in the quantitative determination of solutions of transition metal ions and highly conjugated organic compounds. • Solutions of transition metal ions can be colored (i.e., absorb visible light) because d electrons within the metal atoms can be

34 excited from one electronic state to another. The colour of metal ion solutions is strongly affected by the presence of other species, such as certain anions or ligands. For instance, the colour of a dilute solution of copper sulfate is a very light blue; adding ammonia intensifies the colour and changes the

wavelength of maximum absorption (kmax). • Organic compounds, especially those with a high degree of conjugation, also absorb light in the UV or visible regions of the electromagnetic spectrum. The solvents for these determinations are often water for water soluble compounds, or ethanol for organic-soluble compounds. (Organic solvents may have significant UV absorption; not all solvents are suitable for use in UV spectroscopy. Ethanol absorbs very weakly at most wavelengths.) Solvent polarity and pH can affect the absorption spectrum of an organic compound. Tyrosine, for example, increases in absorption maxima and molar extinction coefficient when pH increases from 6 to 13 or when solvent polarity decreases. • While charge transfer complexes also give rise to colours, the colours are often too intense to be used for quantitative measurement[97]. The wavelengths of absorption peaks can be correlated with the types of bonds in a given molecule and are valuable in determining the functional groups within a molecule. The most widely applicable cuvettes are made of high quality fused silica or quartz glass because these are transparent throughout the UV, visible and near infrared regions. Glass and plastic cuvettes are also common, although glass and most plastics absorb in the UV, which limits their usefulness to visible wavelengths[98].

35 1.8.2.1 Factors affecting on UV: Bathochromic shift: Bathochromic shift is a change of spectral band position in the absorption, reflectance, transmittance, or emission spectrum of a molecule to a longer wavelength (lower frequency). Because the red color in the visible spectrum has a higher wavelength than most other colors, this effect is also commonly called a red shift, although this usage is considered informal and has no relation to Doppler shift or other wavelength-independent meanings of red shift. This can occur because of a change in environmental conditions: for example, a change in solvent polarity will result in solvatochromism. A series of structurally related molecules in a substitution series can also show a bathochromic shift. Bathochromic shift is a phenomenon seen in molecular spectra, not atomic spectra; it is thus more common to speak of the movement of the- peaks in the spectrum rather than lines. Bathochromic shift is typically demonstrated using a spectrophotometer, colorimeter, or spectroradiometer. Hypsochromic shift: Hypsochromic shift is a change of spectral band position in the absorption, reflectance, transmittance, or emission spectrum of a molecule to a shorter wavelength (higher frequency). Because the blue color in the visible spectrum has a lower wavelength than most other colors, this effect is also commonly called a blue shift. This can occur because of a change in environmental conditions: for example, a change in solvent polarity will result in solvatochromism. A series of structurally related molecules in a substitution series can also show a hypsochromic shift. Hypsochromic shift is a phenomenon seen in molecular spectra, not atomic spectra -

36 it is thus more common to speak of the movement of the peaks in the spectrum rather than lines. 1.8.3 NMR spectroscopy: Nuclear magnetic resonance spectroscopy, most commonly known as NMR spectroscopy, is the name given to a technique which exploits the magnetic properties of certain nuclei. For details regarding this phenomenon and its origins, refer to the nuclear magnetic resonance article. The most important applications for the organic chemist are proton NMR and carbon-13 NMR spectroscopy. In principle, NMR is applicable to any nucleus possessing spin. Many types of information can be obtained from an NMR spectrum. Much like using infrared spectroscopy (IR) to identify functional groups, analysis of a NMR spectrum provides information on the number and type of chemical entities in a molecule. However, NMR provides much more information than IR. The impact of NMR spectroscopy on the natural sciences has been substantial. It can, among other things, be used to study mixtures of analytes, to understand dynamic effects such as change in temperature and reaction mechanisms, and is an invaluable tool in understanding protein and nucleic acid structure and function. It can be applied to a wide variety of samples, both in the solution and the solid state. The basic arrangement of an NMR spectrometer is shown to the Fig. 1.2. The sample is positioned in the magnetic field and excited via pulsations in the radio frequency input circuit. The realigned magnetic fields induce a radio signal in the output circuit which is used to generate the output signal. Fourier analysis of the complex output produces the actual spectrum. The pulse is repeated

37 as many times as necessary to allow the signals to be identified from the background noise.

Sample tube

radio frequency •I.I * radio frequency output input

spectrum. --J Fig. 1.2: Basic arrangement of NMR spectrometer The NMR sample is prepared in a thin-walled glass tube - an NMR tube. When placed in a magnetic field, NMR active nuclei (such as H or C) absorb at a frequency characteristic of the isotope. The resonant frequency, energy of the absorption and the intensity of the signal are proportional to the strength of the magnetic field. 1.8.3.1 Factors affecting on NMR: Chemical Shift: The chemical shift in absolute terms is defined by the frequency of the resonance expressed with reference to a standard compound which is defined to be at 0 ppm. The scale is made more manageable by expressing it in parts per million (ppm) and is independent of the spectrometer frequency.

.,..»* frequency of signal - frequency of reference , Chemical shift, 8 - - - x 10 spe:trometer frequency

38 It is often convenient to describe the relative positions of the resonances in an NMR spectrum. For example, a peak at a chemical shift, 8, of 10 ppm is said to be downfield or deshielded with respect to a peak at 5 ppm, or if you prefer, the peak at 5 ppm is upfield or shielded with respect to the peak at 10 ppm.

^ DownfiriTl | Upfield "%

—| i | i | i |—i—| 1—| 1—| 1 1 1—| 1—| 1—(- 10 38763432 10 ppm Fig. 1.3: Chemical shifts Typically for a field strength of .4.7 T the resonance frequency of a proton will occur around 200MHz and for a carbon, around 50.4MHz. The reference compound is the same for both, tetramethysilane (Si(CH3)4). Depending on the local chemical environment, different protons in a molecule resonate at slightly different frequencies. Since both this frequency shift and the fundamental resonant frequency are directly proportional to the strength of the magnetic field, the shift is converted into a field-independent dimensionless value known as the chemical shift. The chemical shift is reported as a relative measure from some reference resonance frequency. (For the nuclei 'H, 13C, and 29Si, TMS (tetramethylsilane) is commonly used as a reference.) This difference between the frequency of the signal and the frequency of the reference is divided by frequency of the reference signal to give the chemical shift. The frequency shifts are extremely small in comparison to the fundamental NMR frequency. A typical frequency shift might be 100 Hz, compared to a fundamental NMR frequency of

39 100 MHz, so the chemical shift is generally expressed in parts per million (ppm)[99]. To be able to detect such small frequency differences it is necessary, that the external magnetic field varies much less throughout the sample volume. High resolution NMR spectrometers use shims to adjust the homogeneity of the magnetic field to parts per billion (ppb) in a volume of a few cubic centimeters. The shape and size of peaks are indicators of chemical structure too. In the example above—the proton spectrum of ethanol—the CH3 peak would be three times as large as the OH. Similarly the CH2 peak would be twice the size of the OH peak but only 2/3 the size of the CH3 peak. The integral of the NMR signal is very difficult to interpret in more complicated NMR experiments [100]. J-coupling: Some of the most useful information for structure determination in a one-dimensional NMR spectrum comes from J- coupling or scalar coupling between NMR active nuclei. This coupling arises from the interaction of different spin states through the chemical bonds of a molecule and results in the splitting of NMR signals. These splitting patterns can be complex or simple and, likewise, can be straightforwardly interpretable or deceptive. This coupling provides detailed insight into the connectivity of atoms in a molecule. Coupling combined with the chemical shift (and the integration for protons) tells us not only about the chemical environment of the nuclei, but also the number of neighboring NMR active nuclei within the molecule. In more complex spectra with multiple peaks at similar chemical shifts or in spectra of nuclei other than hydrogen, coupling is often the only way to distinguish different nuclei[101-1031. As is

40 the case for NMR the chemical shift reflects the electron density at the atomic nucleus [104]. Hydrogen Bonding: Protons that are involved in hydrogen bonding (this usually means -OH or -NH) are typically observed over a large range of chemical shift values. The more hydrogen bonding there is, the more the proton is deshielded and the higher its chemical shift will be. However, since the amount of hydrogen bonding is susceptible to factors such as solvation, acidity, concentration and temperature, it can often be difficult to predict. 1.9 DENSITY FUNCTIONAL THEORY: Density functional theory (DFT) is a quantum mechanical theory used in physics and chemistry to investigate the electronic structure (principally the ground state) of many-body systems, in particular atoms, molecules, and the condensed phases. With this theory, the properties of a many-electron system can be determined by using functionals, i.e. functions of another function, which in this case is the spatially dependent electron density. Hence the name density functional theory comes from the use of functionals of the electron density. DFT is among the most popular and versatile methods available in condensed-matter physics, computational physics, and computational chemistry. DFT has been very popular for calculations in solid state physics since the 1970s. In many cases the results of DFT calculations for solid-state systems agreed quite satisfactorily with experimental data. Also, the computational costs were relatively low when compared to traditional ways which were based on the complicated many-electron wave function, such as Hartree-Fock theory and its descendants. However, DFT was not considered

41 accurate enough for calculations in quantum chemistry until the 1990s, when the approximations used in the theory were greatly refined to better model the exchange and correlation interactions. DFT is now a leading method for electronic structure calculations in chemistry and solid-state physics. Despite recent improvements, there are still difficulties in using density functional theory to properly describe intermolecular interactions, especially van der Waals forces (dispersion); charge transfer excitations; transition states, global potential energy surfaces and some other strongly correlated systems; and in calculations of the band gap in semiconductors. Its incomplete treatment of dispersion can adversely affect the accuracy of DFT (at least when used alone and uncorrected) in the treatment of systems which are dominated by dispersion (e.g. interacting noble gas atoms) or where dispersion competes significantly with other effects (e.g. in biomolecules). The development of new DFT methods designed to overcome this problem, by alterations to the functional or by the inclusion of additive terms, is a current research topic. Although density functional theory has its conceptual roots in the Thomas-Fermi model, DFT was put on a firm theoretical footing by the two Hohenberg-Kohn theorems (H-K)[105]. The original H-K theorems held only for non-degenerate ground states in the absence of a magnetic field, although they have since been generalized to encompass these[106,107]. Recently, another foundation to construct the DFT without the Hohenberg-Kohn theorems is getting popular, that is, as a Legendre transformation from external potential to electron density [108].

42 1.9.1 Applications: In practice, Kohn-Sham theory can be applied in several distinct ways depending on what is being investigated. In solid state calculations, the local density approximations are still commonly used along with plane wave basis sets, as an electron gas approach is more appropriate for electrons delocalised through an infinite solid. In molecular calculations, however, more sophisticated functionals are needed, and a huge variety of exchange-correlation functionals have been developed for chemical applications. Some of these are inconsistent with the uniform electron gas approximation; however, they must reduce to LDA in the electron gas limit. Among physicists, probably the most widely used functional is the revised Perdew- Burke-Ernzerhof exchange model; however, this is not sufficiently calorimetrically accurate for gas-phase molecular calculations. In the chemistry community, one popular functional is known as BLYP (from the name Becke for the exchange part and Lee, Yang and Parr for the correlation part). Even more widely used is B3LYP which is a hybrid functional in which the exchange energy, in this case from Becke's exchange functional, is combined with the exact energy from Hartree-Fock theory. Along with the component exchange and correlation functionals, three parameters define the hybrid functional, specifying how much of the exact exchange is mixed in. The adjustable parameters in hybrid functionals are generally fitted to a 'training set' of molecules. Unfortunately, although the results obtained with these functionals are usually sufficiently accurate for most applications, there is no systematic way of improving them. Hence in the current DFT approach it is not possible to estimate the error of the calculations without comparing them to other methods or experiments.

43 For molecular applications, in particular for hybrid functional, Kohn-Sham DFT methods are usually implemented just like Hartree- Fock itself. 1.10 PRESENT WORK: Accurate knowledge of thermodynamic mixing properties of binary and ternary mixtures has great relevance in theoretical and applied areas of research. These data, are needed for design processes in chemical, petrochemical and pharmaceutical industries. Usually, for non ideal mixtures, direct experimental measurements are performed over the entire composition range. Many times predictive methods for the excess quantities would be more useful than the direct experimental measurements especially when quick estimates are needed. Most empirical approaches for calculating the excess properties attempt to explain solution non idealities in terms of specific and nonspecific intermolecular interactions. These specific and nonspecific intermolecular interactions are sensitive to composition, temperature and presence of different functional groups in the mixtures under investigation. Group contribution methods are powerful tools for the prediction of thermodynamic and transport properties of liquid mixtures, in particular the excess molar volume and deviation function in viscosity and in ultrasonic velocities. The knowledge of the group contribution parameters of the solvent is useful in formulating solvent mixtures for specific applications. IR is the powerful tool in studying inter and intramolecular associations from the position of the O-H band, band width, and intensity of the first overtone band. Electron donating groups lowers IR absorption frequency while electron withdrawing groups increase

44 absorption. Conjugation increases electron density hence IR absorption frequency decreases with increase in conjugation. On absorption of energy of a molecule in the UV region, changes are produced in the electronic energy of the molecule due to transitions of valence electrons from an occupied molecule. These transitions consists of the excitation of an electron from an occupied molecular orbital e.g., a non bonding a or bonding Tt-orbital to the next higher energy orbital i.e., an antibonding 7t* or o*, orbital. Thus, the promotion of an electron, e.g., from a Tt-bonding orbital to an anti- bonding (71*) orbital is designated: n —• TT*. It is clear that n —> n* transition requires less energy compared to n —> it* or a —> o* transition. Since the first application of NMR to structural problems in organic chemistry in 1953, the growth of this technique[109-115] in the seventeenth century has been phenomenal. Within a short period, this technique has thrown light on many difficult organic problems and has continued to solve many intricate problems which remained unresolved when only conventional methods were available to organic chemist. Chemical shifts and coupling constants depend on their electronic environment. This is a function of the type of functional groups present and the number and type of intervening bonds in the molecule. NMR technique is useful for: i) geometric isomerism and conformational analysis; ii) time averaging and its application to rate process i.e. proton exchange, kinetic and thermodynamic study of an equilibrium involving rotation about bond axes, hydride ion shift, spin decoupling and spin ticking by double resonance, and hydrogen bond); iii) determination of aromaticity; iv) determination of

45 molecular weight; and v) study of nuclei other than those of hydrogen. In order to investigate the presence of specific and non-specific interactions between constituent molecules of binary mixtures of 2-methoxyethanol, ethanolamine and cyclohexane with some cyclic and acyclic ethers (i.e. diisopropyl ether, diethyl ether, 1,4-dioxan, tetrahydrofuran, diphenyl ether, t-butyl methyl ether), measurements of densities, viscosities, NMR, ultraviolet and infrared radiations as a function of mixture composition (mole fraction). From the measured values of density and viscosity of the binary mixtures of 2-methoxyethanol, ethanolamine and cyclohexane with some ethers, excess molar volume and deviation in viscosity have been calculated. All the theoretical calculations on 2-methoxyethanol and their mixtures with ethers were performed by the use of G03W series of programs. Geometries of optimizations for all of the investigated molecules in this work were carried out using the DFT/B3LYP method with a medium size 6-311++G(d, p) basis set.

46 REFERENCES 1. J H van't Hoff, Z Phys Chem, 1 (1887) 481. 2. S Arrhenius, Z Phys Chem, 1 (1887) 631. 3. J W Gibbs ,Trans Conn Acad Sci, (1873) 309, 382. 4. J H Hildebrand and R L Scott, "Solubility of Non- Electrolytes" 3rd, ed., Reinhold Publ Co. New York, 1964. 5. J H Hildebrand and R L Scott "Regular Solutions," Prentice Hall, Englewood Cliffs, New Jersey, 1962. 6. The first Eureka moment, Science 305: 1219, August 2004. 7. Test Methods for Density and Specific Gravity (Relative Density) of Plastics by Displacement. ASTM Standard D792- 00. Vol 81.01. American Society for Testing and Materials. West Conshohocken. PA. 8. Nuclear Size and Density, Hyper Physics, Georgia State University. Accessed on line June 26, 2009. 9. Symon, Keith (1971). Mechanics (Third ed.). Addison-Wesley. ISBN 0-201-07392-7. 10. ASHRAE handbook, 1989 edition . 11. Streeter and Wylie Fluid Mechanics, McGraw-Hill, 1981. 12. Holman Heat Transfer, McGraw-Hill, 2002 . 13. Incropera & DeWitt, Fundamentals of Heat and Mass Transfer, Wiley, 1996. 14. IUPAC definition of the Poise. 15. Liquids," Interscience", New York , 1965. 16. a. R J Bearman and J G Kirkwood, J Chem Phys, 28 (1958) 136. b. R J Bearman, J Phys Chem, 65(1961)1961.

47 17. S Glasstone, K J Laidler and H Eyring, "The Theory of Rate Processes", McGraw Hill, New York, 1941, Chapter 9. 18. W E Roseveare, R E Powell and H Eyring, J Appl Phys, 12 (1941) 669. 19. M R V Krishna and G S Laddha, Indust Eng Chem, 7 (1968) 324. 20. A K Doolittle, J Appl Phys, 22 (1951) 1471, 23 (1952) 236. 21. M L Williams, R F Landel and J D Ferry, JAm Chem Soc, 77 (1955)3701. 22. M H Cohen and D Turnbull, J Chem Phys, 31 (1959) 1164. 23. P B Macedo and T A Litovitz, J Chem Phys, 42 (1965) 245. 24. V A Bloomfield, J Chem Phys, 52 (1970) 2781. 25. P J Flory, R A Orwoll and A Vrij, J Am Chem Soc, 86 (1964) 3507. 26. P J Flory, R A Orwoll and A Vrij, J Am Chem Soc, 86 (1964) 3515. 27. P J Flory and A Abe, J Am Chem Soc, 86 (1964) 3563. 28. J Flory P, J Am Chem Soc, 87 (1965) 1833. 29.1 Prigogine, V Mathort and A Bellemans, "The Molecular Theory of Solutions," North-Holland Publishing Co. Amsterdam, 1957. 30. Elert and Glenn. "Viscosity". The Physics Hypertextbook. http://hypertextbook.com/physics/matter/viscosity/. 31. M Pope and C E Swenberg, "Electronic Processes in Organic Crystals and Polymers", 2nd ed., Oxford Science Publications, Oxford University Press, New York, 1999 . 32. IUPAC Gold Book http://www.iupac.org/goldbook/S05765.pdf Retrieved from http://en.wikipedia.org/wiki/HOMO/LUMO.

48 33. International Union of Pure and Applied Chemistry, "ethers". Compendium of Chemical Terminology Internet edition. 34. W Heitmann, G Strehlke and D Mayer "Ethers, Aliphatic" in Ullmann's Encyclopedia of Industrial Chemistry" Wiley-VCH, Weinheim, 2002. 35. J F W McOmie and D E West (1973), 3,3'-ihydroxylbiphenyl Org. Synth., http://www.orgsyn.org/orgsyn/orgsvn/ prepContent.asp ?prep=C V5P0412; Coll. Vol. 5, 412. 36. Occupational exposure guidelines Retrieved from http://en.wikipedia.org/wiki/2-Methoxyethanol 37. The curiously intertwined histories of benzene and cyclohexane E W Warnhoff, J. Chem. Ed. 1( 1996) 494. 38. D M Price, "Temperature Calibration of Differential Scanning Calirimeters", Journal of Thermal Analysis, 45 (1995) 1285. 39. "Ethanolamine MSDS", Acros Organics. http://www.paclp.com/content/documents/MSDS/Ethanolamin e.pdf. 40. K Weissermel, H-J Arpe and C R Lindley, Stephen Hawkins (2003). "Chap. 7. Oxidation Products of Ethylene". Industrial Organic Chemistry. Wiley-VCH. pp. 159, ISBN 3527305785. 41."Ethanolamine". Occupational Safety & Health Administration. http://www.osha.gov/SLTC/healthguidelines/ethanolamine/rec ognition.html. 42. R E Black, F J Hurley and D C Havery (2001). "Occurrence of 1,4-dioxane in cosmetic raw materials and finished cosmetic products". Journal of AOAC International 84 (3), 666. 43. K S Surprenant, "Dioxane" in Ullmann's Encyclopedia of Industrial Chemistry, 2002, Wiley-VCH, Weinheim.

49 44. "International Chemical Safety Card", National Institute for Occupational Safety and Health. http://www.cdc.gov/niosh/ipcsneng/neng0041 .html. 45. "OPPT Chemical Fact Sheets 1,4-Dioxane (CAS No. 123-91- 1)". United States Environmental Protection Agency. http://www.epa.gov/opptintr/chemfact/dioxa-fs.txt. 46. "Eleventh Report on Carcinogens". United States Department of Health and Human Services' National Toxicology Program. http://ntp-server.niehs.nih.gov/ntp/roc/eleventh/profiles/ s080diox.pdf. 47. 1,4-Dioxane (1,4-Diethyleneoxide). Hazard Summary U S Environmental Protection Agency, Created in April 1992; Revised in January 2000. 48. "Proposition 65 List of Chemicals" (PDF), Office of Environmental Health Hazard Assessment, 1,4-Dioxane cancer 123-91-1 January 1, 1988. 49. "Tenth Report on Carcinogens", U S Department of Health and Human Services, Public Health Service, National Toxicology Program, December 2002. http://ehp.niehs.nih.gov/roc/tenth/profiles/s080diox.pdf 50. "California Files Prop 65 Lawsuit Against Whole Foods, Avalon", Bloomberg. http://www.bloomberg.com/apps/news?pid=conewsstory&refer =conews&tkr=WFMI:US&sid=aDC5vdNMQCiM. 51. "Chemical Encyclopedia: 1,4-dioxane", Healthy Child Healthy World. http://healthychild.org/issues/chemical-pop/L4- dioxane/. Retrieved 2009-12-14. 52. http://www.caslab.com/l-4-Dioxane-Testing-Soaps- Shampoos/

50 53. Xinhua (March 21, 2009). "Watchdog issues inspection results on Johnson & Johnson". China Daily. http://www.chinadaily.com.cn/china/2009- 03/21/content 7603271.htm. 54. FDA/CFSAN-Cosmetics Handbook Part 3, " Cosmetic Product-Related Regulatory Requirements and Health Hazard Issues", Prohibited Ingredients and other Hazardous Substances: 9. Dioxane. 55. http://en.wikipedia.org/wiki/K4-Dioxane 56. "Polyethers, Tetrahydrofuran and Oxetane Polymers", G Pruckmayr, P Dreyfuss and M P Dreyfuss Encyclopedia of Chemical Technology, John Wiley & Sons, Inc. 1996. 57. H Muller, "Tetrahydrofuran" in Ullmann's Encyclopedia of Industrial Chemistry 2002, Wiley-VCH, Weinheim. 58. "Chemical Reactivity", Cem.msu.edu. http://www.cem.msu.edU/-reusch/VirtualText/enrgtop.htm#top 4. Retrieved 2010-02-15. 59. "FileAve.com". Gashydrate.fileave.com. http://gashydrate.fileave.com/NMR- MRl%20study%20of%20clathrate%20hvdrate%20mcchanisms .pdf. Retrieved 2010-02-15. 60. C Elschenbroich and A Salzer, "Organometallics : A Concise Introduction" (2nd Ed) (1992) Wiley-VCH, Weinheim. ISBN 3-527-28165-7. 61. BLLuchtand D B Collum, "Lithium Hexamethyldisilazide: A View of Lithium Ion Solvation through a Glass-Bottom Boat", Accounts of Chemical Research, 32 (1999) 1035. 62. "Ethers", L Karas and W J Piel, Kirk-Othmer Encyclopedia of Chemical Technology, John Wiley & Sons, Inc. 2004.

51 63. "Extra Strength Starting Fluid: How it Works", Valvovine. http.V/www.valvoline.com/pages/products/product detail.asp?p roduct=3 8§ion=402. 64. Microsoft Word - RedListE2007.doc. 65. Erowid Ether Vaults : Hoffmann's Drops. 66. Zandberg and Adrian (2010). ""Villages... Reek of Ether Vapours": Ether Drinking in Silesia before 1939", Medical History 54 (3), 387. http://www.ncbi.nlm.nih.gov/pmc/articles/PMC2890321. 67. Kaszycki and Nestor (2006-08-30), "Lemkowska Watra w Zdyni 2006 - pilnowanie ognia pami?ci" (in Polish), Histmag.org - historia od podszewki, Krakow, Poland. http://histmag.org/7idK349. Retrieved 2009-11-25. 68. W L F Armarego and C L L Chai (2003). Purification of laboratory chemicals, Butterworth-Heinemann. ISBN 978- 0750675710. 69. C M Suter and C E Maxwell (1943), "Phenoxthin", Org Synth, http://www.orgsyn.orR/orgsvn/orgsyn/ prepContent.asp?prep=cv2p0485; Coll Vol 2, 485 . 70. M Ueoda, T Aizawa and Y Imai (1977), "Preparation and properties of poly amides and polyimides containing phenoxathiin units", Journal of Polymer Science: Polymer Chemistry Edition 15 (11) 2739. 71. H E Ungnade and E F Orwoll (1955), "2-Methoxy Diphenyl Ether", Org Synth, http://www.orgsvn.org/orgsyn/orgsvn/ prepContent.asp?prep=:cv3p0566; Coll Vol 3, 566. 72. http://eur-lex.europa.eu/LexUriServ/site/en/oj/2003/l_042/ 1 04220030215en00450046.pdf.

52 73. B J Sutker (2005), "Flame Retardants". Ullmann's Encyclopedia of Industrial Chemistry (Weinheim: Wiley- VCH). 74.http://www.ccohs.ca/oshanswers/chemicals/organic/organic_pe roxide.html 75. CRC Handbook of Chemistry and Physics, 90th edition. 76. http://www.epa.gov/otaq/rfg regs.htm#usage 77. San Francisco Bay Area Regional Water Quality Control Board Integrated Basin Management Plan (2004). 78. A Fischer, C Oehm, M Selle and P Werner (2005), "Biotic and abiotic transformations of methyl tertiary butyl ether (MTBE)". Environ Sci Pollut Res Int 12 (6): 381. 79. http://monographs.iarc.fr/ENG/Monographs/vol73/ volume73.pdf 80. http://www.epa.gOv/mtbe/faq.htm#concerns 81. L M Harwood and C J Moody (1989), Experimental organic chemistry: Principles and Practice (Illustrated ed.), Wiley- Blackwell. p. 292, ISBN 0632020172. 82. WSLau(1999), Infrared characterization for microelectronics, World Scientific. ISBN 9810223528. http://books.google.com/?id=rotNlJDFJWsC&printsec=frontco ver 83. International Union of Pure and Applied Chemistry, "inductive effect", Compendium of Chemical Terminology Internet edition. Retrieved from http://en.wikipedia.org/wiki/Inductive_effect 84. IUPAC Gold Book mesomerism PDF. 85. IUPAC Gold Book resonance PDF. 86. IUPAC Gold Book contributing structure PDF.

53 87. L Pauling, "The Nature of the chemical bond - An Introduction to Modern Structural Chemistry ", Cornell University Press, Third Edition 1960, The Concept of Resonance, pp.10. 88. IUPAC Gold Book resonance energy PDF. 89. L Pauling, "Resonance", Manuscript for publication in Encyclopedia Britannica, p. 13; July 29, 1946. 90. R Morrison and R Boyd (1989), "Chapter 10", Organic Chemistry (Fifth Edition ed.), Prentice Hall of India, pp. 372, ISBN 0-87692-560-3. 91. Hyperconjugation not steric repulsion leads to the staggered structure of ethane, V Pophristic and L Goodman , 565. 92. Chemistry: A new twist on molecular shape, Frank Weinhold Nature 411,539. 93. Gait and Michael, Oligonucleotide synthesis: a practical approach, Oxford: IRL Press, ISBN 0904147746. 94. K Morokuma et al., "Can Steric Effects Induce the Mechanism Switch in the Rhodium-Catalyzed Imine Boration Reaction? A Density Functional and ONIOM Study" , Organometallics 2005, 24, 1938. 95. V P Ananikov, D G Musaev and K Morokuma, Eur. J. Inorg. Chem.( 2007) 5390. 96. Skoog, Principles of Instrumental Analysis, Sixth ed. Thomson Brooks/Cole. 2007, 169. 97. "Beer's Law - Alisdair Boraston". http://www.brewingtechniques.com/brewingtechniques/beersla w/boraston.html. Retrieved 2009-02-06. 98. Skoog, Principles of Instrumental Analysis. Sixth ed. Thomson Brooks/Cole. 2007, 351.

54 99. J Keeler, "Chapter 2: NMR and energy levels" (reprinted at University of Cambridge), Understanding NMR Spectroscopy, University of California, Irvine. http://www-keeler.ch.cam.ac.uk/lectures/Irvine/chapter2.pdf. 100. G E Martin and A S Zekter, Two-Dimensional NMR Methods for Establishing Molecular Connectivity, VCH Publishers, Inc: New York, 1988, p.59. 101. Spectrometric Identification of organic Compounds Silverstein, Bassler, Morrill Fourth Ed, ISBN 047109700. 102. Organic Spectroscopy, William Kemp Third Ed, ISBN 0333417674. 103. Basic !H - 13C-NMR spectroscopy, MetinBalei, ISBN 0444518118. 104. "A Short History of Three Chemical Shifts", Journal of Chemical Education, Shin-ichi Nagaoka, Vol. 84 No. 5 May 2007,801. 105. H Pierre and W Kohn (1964), "Inhomogeneous electron gas", Physical Review 136 (3B): B864. 106. L Mel (1979), "Universal variational functional of electron densities, first-orderdensit y matrices, and natural spin-orbitals and solution of the v-representability problem", Proceedings of the National Academy of Sciences (United States National Academy of Sciences) 76 (12) 6062. 107. G Vignaleand MRasolt (1987), "Density-functional theory in strong magnetic fields", Physical Review Letters (American Physical Society) 59 (20): 2360. 108. Density Functional Theory, an introduction, Rev. Mod. Phys. 78(2006) 865.

55 109. L R Subramanian and G S Krishna Rao, Tetrahedron Letters, (1967)3693. 110. G A Neville and I C Nigam, Can J Chem, 47 (1969) 2901. 111. S Meiboon and L C Snyder, J Am Chem Soc, 89 (1967) 1083. 112. D F Leyden and W R Morgan, J Che Educ, 46 (1969) 169. 113. M Sauders, P von R Schleyer and G Olah, J Am Chem Soc, 86 (1964) 5680. 114. W McFarlane, Chem Brit, 5 (1969) 142. 115. S Barcza, J Org Chem, 28 (1963) 1964.

56