Question

Which of the following statement is true? (a) Ostwald's dilution law holds good only for strong electrolytes and fails completely when applied to weak electrolytes. (b) Ostwald's dilution law holds good for both weak and strong electrolytes. (c) Ostwald's dilution law holds good only for weak electrolytes and fails completely when applied to strong electrolytes. (d) Ostwald's dilution law does not good hold good for both weak and strong electrolytes.

Short answer

The correct statement is (c) - Ostwald's dilution law holds good only for weak electrolytes and fails completely when applied to strong electrolytes.

Step-by-step solution

Understand Ostwald's Dilution Law

Ostwald's dilution law is applicable to weak electrolytes. This law states that as a solution of a weak electrolyte is diluted, or made more dilute, the proportion of the electrolyte that exists as ions (the degree of dissociation) increases. In other words, weak electrolytes dissociate more extensively into ions as their concentration decreases.

Analyze Each Statement

Now let's use our understanding to evaluate each of the provided options. We know that Ostwald's dilution law is specifically applicable for weak electrolytes.

Choose the Correct Statement

Option (c) states 'Ostwald's dilution law holds good only for weak electrolytes and fails completely when applied to strong electrolytes.' This option is in agreement with the definition and understanding of Ostwald's dilution law and is thus correct.

Key concepts

Weak Electrolytes

In chemistry, weak electrolytes are substances that only partially ionize in solution. Unlike strong electrolytes, which completely dissociate into ions, weak electrolytes exist in a dynamic equilibrium between the ionized (dissociated) and the non-ionized (undissociated) states when dissolved in a solvent, typically water.

To understand this better, let's take the example of acetic acid in water. Acetic acid (\( CH_3COOH \)) doesn't fully break down into its ions (\( CH_3COO^- \) and H+). Instead, some molecules remain intact in the solution, which is why conductivity measurements of these solutions are lower compared to those of strong electrolytes. The key aspect of weak electrolytes is this partial dissociation, which is heavily dependent on the concentration of the electrolyte—this is where Ostwald's Dilution Law plays a significant role in predicting the behavior of weak electrolytes in dilute solutions.

Degree of Dissociation

The term degree of dissociation, represented by the symbol \( \alpha \), describes the extent to which a substance breaks into ions in a solution. It is defined as the fraction of the total number of solute molecules that dissociate.

For instance, for every 1 mole of a weak electrolyte dissolved, if only 0.1 moles actually dissociate into ions, then \( \alpha = 0.1 \), indicating a 10% dissociation. This concept is crucial for measuring the strength of an electrolyte. The degree of dissociation is affected by dilution; as a solution becomes more dilute, the degree of dissociation tends to increase until it reaches an equilibrium. Ostwald's Dilution Law mathematically relates the degree of dissociation to the dilution of the solution, providing a quantitative measure to understand this relation.

Electrolytic Dissociation

Electrolytic dissociation is the process by which an electrolyte breaks apart into its component ions when dissolved in a solvent. This phenomenon is a reversible reaction for weak electrolytes, meaning the ions can recombine to form the undissociated substance.

In a highly diluted solution, the ions are spread out, reducing their chance of recombining. Consequently, dissociation is encouraged, and therefore, the presence of more free ions is noted. Ostwald's dilution law provides a straightforward relationship between the concentration of the electrolyte and its degree of dissociation, which is especially relevant for weak electrolytes. This law indicates that as the concentration of a weak electrolyte decreases, its degree of dissociation increases, having a vital application in predicting the conductivity and reactivity of dilute solutions of weak electrolytes.