الفهرس 
CHAPTER IOnly 14 pages are availabe for public view
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Abstract
Dissociation is the term used to describe the breakdown of a molecule into simpler
parts. This often (almost always in basic chemistry) occurs in water. The equilibrium constant
for such a reaction is called dissociation constant.
Solute-solvent interactions:
Solvents can be classified according to their chemical bonds: molecular liquids
(molecule melts; covalent bonds only), ionic liquids (molten salts; only ionic bonds), and
atomic liquids (low-melting metals like liquid mercury or liquid sodium; metallic bonds) [2].
A proper choice of solvent, based on the knowledge of its chemical reactivity, helps to avoid
undesired reactions between solute and solvent.
The following three aspects are also of importance in solvation: the stoichiometry of
the solvate complexes (normally described by the coordination or solvation number), the
liability of the solvate complexes (usually described by the rate of exchange of the molecules
of the solvent shell with those of the bulk solvent), as well as the fine structure of the solvation
shell (for water often described by the simple model of ion solvation of Frank and Wen) [3].
Coordination and solvation numbers reflect the simple idea that the solvation of ions
or molecules consists of a coordination of solute and solvent molecules. The coordination
number is defined as the number of solvent molecules in the first coordination sphere of an ion
in solution [4]. This first coordination sphere is composed only of solvent molecules in contact
with or in bonding distance of the ion such that no other solvent molecules are interposed
between them and the ion. This kind of solvation is sometimes termed primary or chemical
solvation. The solvation number is defined as the number of solvent molecules per ion which
remain attached to a given ion long enough to experience its translational movements [5, 6].
Parker [7] divided solvents into two groups according to their specific interactions
with anions and cations, namely dipolar aprotic solvents and protic solvents. The distinction
lies principally in the dipolarity of the solvent molecules and their ability to form hydrogen
bonds. A polar aprotic solvent is characterized by a low relative permittivity (εr < 15) a low
dipole moment (μ < 2.5 D), a low ET
N value of (0-0.3), and the inability to act as a hydrogenbond
donor. Such solvents interact only slightly with the solute since only the non-specific
directional, induction, and dispersion forces can operate. To this group belong aliphatic and
aromatic hydrocarbons, their halogen derivatives, tertiary amines, and carbon disulfide. In
contrast, dipolar aprotic solvents possess large relative permittivities (εr > 15), sizeable dipole
moments (μ > 2.5 D), and average ET
N values of 0.3 to 0.5. These solvents do not act as hydrogen-bond donors since their C---H bonds are not sufficiently polarized. Solvents containing proton-donor groups are designated protic solvents or HBD solvents (water,
ammonia, alcohols, carboxylic acids, and primary amides); solvents containing protonacceptor
groups are called HBA solvents [8, 9] (amines, ethers, ketones, and sulfoxides).
Amphiprotic solvents can act both as HBD and as HBA solvents simultaneously. The abbreviations HBD (hydrogen-bond donor) and HBA (hydrogen-bond acceptor) refer to donation and acceptance of the proton, and not to the electron pair involved in hydrogen bonding. Protic solvents contain hydrogen atoms bound to electronegative elements
(F--H,--O--H, --N--H, etc.) and are, therefore, hydrogen-bond donors i.e. HBD solvents. With
the exception of acetic acid (and its homologues), the relative permittivities are usually larger
than 15, and the ET
N values lie between 0.5 and 1.0, indicating that these solvents are strongly polar. To this class of solvents belong water, ammonia, alcohols, carboxylic acids, and primary amides. The solvation energy is considered as the change in Gibbs energy when an ion or molecule is transferred from a vacuum (or the gas phase) into a solvent. The Gibbs energy of solvation, ΔG°Solv, a measure of the solvation ability of a particular solvent, is the result of a
superimposition of four principal components of a different nature [10,
(a) The cavitation energy associated with the hole that the dissolved molecule or ion produces in the solvent; (b) The orientation energy corresponding to the phenomenon of partial orientation of the
dipolar solvent molecules caused by the presence of the solvated molecule or ion.(c) The isotropic interaction energy corresponding to the unspecific intermolecular forces with a long radius of activity (i.e. electrostatic, polarization, and dispersion energy); (d) The anisotropic interaction energy resulting from the specific formation of hydrogen bonds or electron-pair donor/electron-pair acceptor bonds at well localized points in the dissolved molecules. Liquid water [1] consists both of bound ordered regions of a regular lattice, and regions in which the water molecules are hydrogen-bonded in a random array; it is permeated by monomeric water and interspersed with random holes, lattice vacancies, and cages. There are chains and small polymers as well as bound, free, and trapped water molecules. The currently accepted view of the structure of liquid water treats it as a dynamic threedimensional
hydrogen-bonded network, without a significant number of non-bonded water
molecules, that retains several of the structural characteristics of ice (i.e. tetrahedral molecular
packing with each water molecule hydrogen-bonded to four nearest neighbors), although the
strict tetrahedrality is lost. Its dynamic behaviour resembles that of most other liquids, with
short rotational and translational correlation times of the order of 0.1 to 10 ps, indicating high
hydrogen-bond exchange rates [12, 13].
Acetone, the chosen solvent for the present study, can be classified as a dipolar aprotic solvent [14] although its protons are more reactive than those of typical members of this class of solvents. (This is also the case for nitromethane and, to some extent, for the formyl proton of dimethylformamide.) This proton reactivity constitutes the main limitation of acetone as a solvent. Theintermolecular structure of acetone is determined mainly by steric interactions between the methyl groups and, unexpectedly, only to a small extent by dipole/dipole forces, whereas the inner structure of dimethyl sulfoxide is dictated by strong
dipole/dipole interactions. Other limitations are its exceptionally high hygroscopicity, high
volatility and formation of explosive mixtures with air. On the credit side, acetone is available
commercially in a sufficiently pure state for many (but not all) applications.