MCQ Detailed View

Ostwald’s dilution law explains the relationship between the dissociation constant (Ka or Kb), the degree of dissociation (α), and the concentration (C) of an electrolyte in solution. It is particularly useful in understanding how weak electrolytes ionize in water.

The law states:

K=Cα21−αK = \frac{C \alpha^2}{1 - \alpha}K=1−αCα2

Here, C is the initial concentration of the weak electrolyte, α is the degree of dissociation, and K is the dissociation constant. As the solution is diluted, α increases because molecules dissociate more in dilute solutions.

This law is only applicable to weak electrolytes. The reason is that weak electrolytes, such as acetic acid (CH₃COOH) or ammonium hydroxide (NH₄OH), ionize only partially in water. Their dissociation follows equilibrium principles that can be described by Ostwald’s equation.

For strong electrolytes (like HCl, NaOH, NaCl), the degree of dissociation is nearly 100% even at moderate concentrations. They do not follow Ostwald’s dilution law because their ionization is essentially complete and does not vary significantly with dilution. Instead, strong electrolytes are better explained by Debye–Hückel theory.

For non-electrolytes (like sugar or alcohol), no ionization occurs in water, so the law has no application at all.

Thus, the correct option is weak electrolytes, because Ostwald’s dilution law provides a clear mathematical explanation for their dissociation behavior with dilution.

This principle is very important in physical chemistry, particularly in ionic equilibrium, acid–base chemistry, and buffer solutions. It also helps in experimentally determining the dissociation constants of weak acids and bases.