Dissociation of electrolytes in solutions

Feb. 19,1859, Wijk, Sweden - Oct. 2,1927, Stockholm, Sweden). Arrhenius developed the theory of dissociation of electrolytes in solutions that was first formulated in his Ph.D. thesis in 1884 Recherches sur la conductibilit galvanique des dectrolytes (Investigations on the galvanic conductivity of electrolytes). The novelty of this theory was based on the assumption that some molecules can be split into ions in aqueous solutions. The - conductivity of the electrolyte solutions was explained by their ionic composition. In an extension of his ionic theory of electrolytes, Arrhenius proposed definitions for acids and bases as compounds that generate hydrogen ions and hydroxyl ions upon dissociation, respectively (- acid-base theories). For the theory of electrolytes Arrhenius was awarded the Nobel Prize for Chemistry in 1903 [i, ii]. He has popularized the theory of electrolyte dissociation with his textbook on electrochemistry [iv]. Arrhenius worked in the laboratories of -> Boltzmann, L.E., -> Kohlrausch, F.W.G.,- Ostwald, F.W. [v]. See also -> Arrhenius equation. 

The most elaborate treatment of the dissociation of electrolytes in solutions is the one given by Fuoss and Onsager (9, 10). The so-called F.O. equation, applied to I-I associated electrolytes is... 

In electrolyte solutions, the interaction between ions leads to the formation of ion pairs. The van t Hoff factor provides a measure of the exteut of dissociation of electrolytes in solution. 

As has been pointed out, the equations are all integrations of the Gibbs-Duhem relationship. They consequently cannot be applied to systems which when treated in the ordinary fashion apparently do not follow this basic relation, as in the case of dissociation of electrolytes in solution. 

A century ago, van t Hoff s (1 ) pioneering work on the gas-solution analogy was followed by Arrhenius (2 ) theory of partial dissociation of electrolytes in solutions. Later, electrolytes came to be classified as weak or strong with the supposition that the former are partially dissociated whereas the latter are completely dissociated in the given solvent However, with... 

This is the final evidence provided here in support of Arrhenius idea of partial dissociation of electrolytes in solutions. 

Table 1.3 Degree of dissociation of electrolytes in O 1m aqueous solutions...

An imponant characteristic of electrolyte solutions is that they are electrically conductive. A useful measure is the equivalent conductance, Q, the conductance per mole of charge. A strong electrolyte is one that is completely dissociated into ions. In this case the equivalent conductance is high, and decreases only slowly with increasing concentration. A weak electrolyte is only partially dissociated into its constituent ions, and its equivalent conductance is less than that of a strong electrolyte at any concentration but increases rapidly as the concentration decreases. This is because there is more complete dissociation and therefore more ions per mole of electrolyte in solution as the concentration of a weak electrolyte decreases. Sodium chloride, which completely dissociates into sodium and chloride ions,... 

Jacobus Henricus van t Hoff (1852-1911) a Dutch physical and organic chemist (Nobel Prize in Chemistry in 1901). His research work concentrated on chemical kinetics, chemical equilibrium, osmotic pressure, and crystallography. He is one of the founders of the discipline of physical chemistry. He explained the phenomenon of optical activity by assuming that the chemical bonds between carbon atoms and their neighbors were directed towards the corners of a regular tetrahedron, applied the laws of thermodynamics to chemical equilibriums, showed similarities between the behavior of dilute solutions and gases, and worked on the theory of the dissociation of electrolytes. In 1878, he became professor of chemistry at the University of Amsterdam, and in 1896 he became professor at the Prussian Academy of Science at Berlin, where he worked until his death. 

Svante Arrhenius, 1859-1927, was a Swedish chemist who won the 1905 Nobel Prize in chemistry for his theory of dissociation and ionization of electrolytes In solution. 

The above examples assume that the strong base KOH is completely dissociated in solution and that the concentration of OH ions was thus equal to that of the KOH. This assumption is valid for dilute solutions of strong bases or acids but not for weak bases or acids. Since weak electrolytes dissociate only slightly in solution, we must use the dissociation constant to calculate the concentration of [H" ] (or [OH ]) produced by a given molarity of a weak acid (or base) before calculating total [H" ] (or total [OH ]) and subsequendy pH. 

Debye-Huckel theory assumes complete dissociation of electrolytes into solvated ions, and attributes ionic atmosphere formation to long-range physical forces of electrostatic attraction. The theory is adequate for describing the behaviour of strong 1 1 electrolytes in dilute aqueous solution but breaks down at higher concentrations. This is due to a chemical effect, namely that short-range electrostatic attraction occurs... 

Because of dissociation and the resulting increase in the total number of particles in solution, the parameters of the colUgative properties assume higher values. These values are proportional to the total concentration, c, of particles (ions and undissociated molecules) in the solution, which for a binary electrolyte is given by [1 + a(X(, - l)]q. The isotonic coefficient i is the ratio of and the concentration c, that would be observed in the absence of dissociation ... 

A difference between electrolytes in solution and in the molten state is that the latter, in general, do not need solvents to dissociate. When an electrolyte exists in the molten state as the only component present and not as a solution of one electrolyte in another molten electrolyte, all phenomena associated with the ionic concentration during electrolysis, such as concentration polarization, cease to be relevant. 

Arrhenius postulated in 1887 that an appreciable fraction of electrolyte in water dissociates to free ions, which are responsible for the electrical conductance of its aqueous solution. Later Kohlrausch plotted the equivalent conductivities of an electrolyte at a constant temperature against the square root of its concentration he found a slow linear increase of A with increasing dilution for so-called strong electrolytes (salts), but a tangential increase for weak electrolytes (weak acids and bases). Hence the equivalent conductivity of an electrolyte reaches a limiting value at infinite dilution, defined as... 

The secondary salt effect is important when the catalytically active ions are produced by the dissociation of a weak electrolyte. In solutions of weak acids and weak bases, added salts, even if they do not exert a common ion effect, can influence hydrogen and hydroxide ion concentrations through their influence on activity coefficients. 

A solute in a given solvent may remain unionized (nonelectrolyte) or may ionize (electrolyte). For nonelectrolytes, 1 millimole (mmol i.e., one formula weight in mg) represents 1 mOsm. For electrolytes, osmolarity depends on the total number of particles in solution which in turn depends on the degree of dissociation of a solute. For example, 1 mmol of completely dissociated KC1 represents 2 mOsm of total particles (i.e., K+ + CT). Similarly, 1 mmol of CaCl2 represents 3 mOsm of total particles (i.e., Ca++ + CT + CT). 

It is estimated that half of all pharmaceuticals are formulated as salts, to achieve increased stability and bioavailability [13]. Predictive solubility methods are very limited for this area, and the development of new models to address this category is very important. The NRTL-SAC model has recently been extended by C.-C. Chen, and Y. Song to represent such electrolytic solutes, that partly dissociate to ions in solution. This extension has been achieved by the addition of one new segment type into the preceding NRTL-SAC model. NRTL-SAC thus becomes a limiting case of the eNRTL-SAC formulation [27]. 

The dissolved state of the electrolytes in water has long been of great interest. More than a century ago, Mendelejew [12] suggested that sulphuric acid forms hydrates, and Arrhenius [13] put forward the theory of partial dissociation of electrolytes, both in the same Jorrmal. These pioneering ideas have eventually proved to be correct [14] for electrolyte solutions from zero to saturation. 

For a weak electrolyte such as acetic acid. CH3C02H, some of the molecules dissociate into ions in solution ... 

According to modem theory, many strong electrolytes are completely dissociated in dilute solutions. The freezing-point lowering, however, does not indicate complete dissociation. For NaCl, the depression is not quite twice the amount calculated on the basis of the number of moles of NaCl added. In the solution, the ions attract one another to some extent therefore they do not behave as completely independent particles, as they would if they were nonelectrolytes. From the colligative properties, therefore, we can compute only the "apparent degree of dissociation" of a strong electrolyte in solution. 

Solvent effects in electrochemistry are relevant to those solvents that permit at least some ionic dissociation of electrolytes, hence conductivities and electrode reactions. Certain electrolytes, such as tetraalkylammonium salts with large hydrophobic anions, can be dissolved in non-polar solvents, but they are hardly dissociated to ions in the solution. In solvents with relative permittivities (see Table 3.5) s < 10 little ionic dissociation takes place and ions tend to pair to neutral species, whereas in solvents with 8 > 30 little ion pairing occurs, and electrolytes, at least those with univalent cations and anions, are dissociated to a large or full extent. The Bjerrum theory of ion association, that considers the solvent surrounding an ion as a continuum characterized by its relative permittivity, can be invoked for this purpose. It considers ions to be paired and not contributing to conductivity and to effects of charges on thermodynamic properties even when separated by one or several solvent molecules, provided that the mutual electrostatic interaction energy is < 2 kBT. For ions with a diameter of a nm, the parameter b is of prime importance ... 

There are some cases where a reaction, that is, the formation or dissolution of a chemical bond, is involved along with ion exchange phenomena (Helfferich, 1983). Examples of this are acid-base neutralization, dissociation of weak electrolytes in solution or weak ionogenic groups in ion exchangers, complex formation, or combinations of these (Table 5.2). With some of these, very low apparent D in ion exchangers have been noted. 

We will see in Chapter 10, when we deal with the dissociation of electrolytes into ions, that, in general, the chemical potential of a dissociating solute is the chemical potential of its component parts, whereas the activity of the solute is the product of the component parts. 

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