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Dissolution Process and Classification of Electrolytes

Most chemical reactions occur in solution. The study of such solutions constitutes one of the important branches of physical chemistry. In general, if we analyse the solubility of solutes in various solvents, we find that the polar solutes are more soluble in polar solvents, whereas non-polar solutes are more soluble in nonpolar solvents. For example, sodium chloride (polar solute) is soluble in water (polar solvent), whereas it is insoluble in carbon tetrachloride (nonpolar solvent) and naphthalene (nonpolar solute) is soluble in carbon tetrachloride, whereas it is insoluble in water.

Because of its polar nature* and high dielectric constant * water serves as one of the important solvents for ionic solutes. The high dielectric constant (D = 80) weakens the forces of attraction between the oppositely charged ions of an ionic crystal and its polar nature generates the ion-dipole interaction in which the positive ion is attached to the negative end of the water dipole, whereas negative ion is attached to the positive end of the dipole as shown in Fig. The consequence of this is that the ions are pulled out of the crystal lattice and drift into water and thus a solution is formed. In solution, ions move in the hydrated form. Certain covalent molecules with relatively high dipole moments (e.g. HCl) can also dissolve in water to produce an ionic solution. This is due to the stronger ion-dipole interactions.

Fig : Ion-dipole interactions

A solute when dissolved in water may produce three different types of solution depending upon whether it is a good, poor or bad conductor of electricity. The conduction in the solution is due to the movement of ions. Hence a good conducting solution contains a larger number of ions, a poor one contains lesser number of ions and the bad one will not contain any ions. The ions in the solution are produced by dissolving the solute (also known as electrolyte) in water. Based on the conducting ability, an electrolyte may be classified into three categories as described below.


Conducting ability


Strong electrolyte


NaCl, KCl, KBr, NH4Cl, HCl, H2SO4, NaOH

Weak electrolyte


PbCl2,HgCl2,H2CO3, CH3COOH

Non - electrolyte

non conducting


The classification of compounds in terms of strong and weak electrolytes is based on their behaviour in a particular solvent, namely, water. However, such a classification suffers from a great disadvantage in the sense that a particular electrolyte, though weak in water, might behave as a strong one when dissolved in some other solvent or vice versa. For example, sodium chloride behaves as a strong electrolyte and acetic acid as a weak electrolyte when dissolved in water. However, when acetic acid and sodium chloride are dissolved in ammonia, their conducting abilities are comparable, indicating a strong electrolyte behaviour for acetic acid.

Thus, the above classification depends upon the solvent used. Another classification which is largely based on the characteristics of the solute and not on that of the solvent, is to label them as the true electrolyte and the potential electrolyte. The essential characteristic of a true electrolyte is that even in the pure liquid state it is an ionic conductor. In dissolution process, all that a polar solvent does is to use ion-dipole forces to disengage ions from their lattice sites, solvate them and disperse them into the solution. Examples are NaCl, KCl, etc. The potential electrolyte, however, does not conduct in the pure liquid state, though it provides a conducting solution on dissolution in an ionic solvent. Examples are hydrochloric acid, acetic acid, etc.

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