الفهرس 
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Abstract
Silver is used in the medical practice to avoid infections in wound dressings such as burns and chronic wounds where Ag (I) compounds have known antimicrobial activity 1-18. In addition, Silver was used as an antimicrobial agent for a long time before the discovery of microorganisms, and in ’’ modern ’’ medicine, prior to the introduction of efficient antibiotics. Silver in various forms was used in a number of clinical situations. However, since World War II, its use has declined; for example, the compulsory dropping of AgNO3 solutions in the eyes of newly born babies was discontinued in Sweden 19.
Since the drug silver sulfadiazine [Ag ((4-aminophenyl)sulfonyl)(pyrimidin-2-yl) azanide]n, was used as antibiotic ointment 20, there was a great interest to study chemistry of such silver complexes. The behavior of silver (I) complexes in solution is not yet elucidated. from this point of view, we use the conductance technique to study the solution chemistry of these compounds. Measurements of the conductivity of solutions of electrolytes have been made for nearly a century 21, and from the earliest times have been one of the most accurate physical measurements made on solutions Moreover, interest has never been restricted to aqueous solutions and even in 1888, and the conductivity of non-aqueous solutions 22 and mixed electrolytes 23 was reported. However, throughout this long history most emphasis has been on the conductance of signal symmetrical electrolytes for the simple reason that the interpretation of the conductance of a solution containing only two types of ion of equal and opposite charge has been relatively straightforward.
A variety of theoretical approaches has been developed to account successfully for the conductance of 1:1 electrolytes in water up to a concentration of about 0.1 mol dm-3 but to rather lower concentrations for higher charged electrolytes and/or solvents with a lower dielectric constant. These have been extensively reviewed elsewhere 24, 25 and will not be considered further here.
1.2. Ion association :
Random thermal motion of ions in solution of strong electrolytes constantly bring anions and cations into proximity. Electrostatic attraction then alters the ionic motions to bring the ions closer together, occasionally to contact or near contact distances. Such pairs remain together until sufficiently large thermal fluctuations again send the ions apart. The tendency of the ions to associate in this way depends upon the balance the interionic coulomb forces and thermal energy - thus upon the ion valances, the dielectric constant and the temperature. It depends also on the salvation energy of the ions. Since the heats of solutions are small, dissolution of soluble salt as free ions requires sufficient salvation energy almost completely offset the lattice energy.
Dissolution as ion-pairs requires salvation energy only approximately equivalent to the sublimation energy. Thus, poor solvation promotes ion association.
1.2.1. Solvent effect on ion association:
The effect of a solvent on the equilibrium constant of ionic association (Ka) is of major interest for the theory of electrolytic solutions 26. Many investigations have established a relation between Ka and the micro- and macro properties of the electrolytic system, starting from the Bronsted equation 27 to the most general Izmallov equation. The crystallographic or effective ionic radii, ion-ion and ion-dipole interaction energies belong to the former, the latter are crystalline lattice energy and solvent permittivity ().
For symmetrical electrolytes, where cospheres of a pair of ions Mn+ and Xn- overlap, the charge equal to zero and thus taken no part in charge transport. On the other hand, in case of asymmetrical electrolytes the formed ion-pair has charge and contribute in conducting process but less than that of the free ion.
Bjerrum 28 that defined, for dilute solution, the probability of an ion of charge (z-eo) in a dr thickness spherical shell of radius r around a reference ion of charge (z+eo) by the equation:
Pr = (4 π nio) eλ/r r2 dr ………………………………. (1-1)
where
= z+z-eo2 / DkT
Bjerrum 28 plotted, for dilute solution, the probability of finding an oppositely charged ion at a given distance from a central ion. This distribution curve, Figure (1), shows that pr goes through a minimum for a particular value of r, which was named the critical distance and Bjerrum put it equals to the distance q. This minimum occurs at q = /2= (zizjeo2/2DkT) =3.57 Ao for (1:1) electrolyte at 25oC in water.