Problem 61 - A Complete Resource Book in Chemistry for JEE Main

Question

The pH \(0.1 \mathrm{M}\) solution the following salts increases in the order (a) \(\mathrm{NaCl}<\mathrm{NH}_{4} \mathrm{Cl}<\mathrm{NaCN}<\mathrm{HCl}\) (b) \(\mathrm{HCl}<\mathrm{NH}_{4} \mathrm{Cl}<\mathrm{NaCl}<\mathrm{NaCN}\) (c) \(\mathrm{NaCN}<\mathrm{NH}_{4} \mathrm{Cl}<\mathrm{NaCl}<\mathrm{HCl}\) (d) \(\mathrm{HCl}<\mathrm{NaCl}<\mathrm{NaCN}<\mathrm{NH}_{4} \mathrm{Cl}\)

Step-by-Step Solution

Answer

(b) HCl < NH₄Cl < NaCl < NaCN

Step 1: Understand pH and Acid-Base Properties

The pH of a solution is a measure of how acidic or basic it is. Acids have a pH less than 7, bases have a pH greater than 7, and neutral solutions have a pH of around 7. To solve this exercise, we need to understand the acidic or basic nature of the ions produced by the salts in solution.

Step 2: Analyze Each Salt Solution

Each salt dissociates into its constituent ions in solution. 1. **HCl** is a strong acid, which means it dissociates completely in water to form H⁺ ions, yielding a low pH. 2. **NH₄Cl** dissociates to NH₄⁺ and Cl⁻. NH₄⁺ acts as a weak acid because it can donate an H⁺ to water. 3. **NaCl** dissociates into Na⁺ and Cl⁻, neither of which affects the pH significantly, so the solution is nearly neutral. 4. **NaCN** dissociates into Na⁺ and CN⁻. CN⁻ acts as a base, accepting protons from water, thus increasing the pH.

Step 3: Order the Salts by Increasing pH

Based on the acidic or basic nature of the ions: 1. **HCl** is the most acidic with the lowest pH. 2. **NH₄Cl**, being slightly acidic, follows. 3. **NaCl** is approximately neutral, so it comes next. 4. **NaCN** is the most basic, increasing the pH the most.

Key Concepts

Acidic and Basic Nature of Ions

When a salt dissolves in water, it splits up into its ions. These ions can influence the acidity or basicity of the solution. Some ions can donate or accept protons, changing the pH of the solution. For instance, the ammonium ion \( \text{NH}_4^+ \) from \( \text{NH}_4\text{Cl} \) can donate protons, making the solution slightly acidic. In contrast, cyanide ions \( \text{CN}^- \) from \( \text{NaCN} \) are proton acceptors, which makes the solution basic by raising the pH. Understanding if an ion is acidic or basic helps predict how it alters the pH.

Strong Acid Dissociation

Strong acids, like \( \text{HCl} \), dissociate completely in water. This means each molecule of \( \text{HCl} \) will release a proton \( \text{H}^+ \) into the solution. As a result, the solution becomes very acidic because of the high concentration of \( \text{H}^+ \) ions. Even in a dilute solution, the complete dissociation ensures the pH remains low, often below 4. The presence of more \( \text{H}^+ \) ions correlates with a more acidic environment and lower pH. It's crucial in determining the order of acidity among different salt solutions.

Weak Acid Behavior

In contrast to strong acids, weak acids do not fully dissociate in water. For example, the \( \text{NH}_4^+ \) ion from \( \text{NH}_4\text{Cl} \) only partially donates protons to the solution. Because it is in equilibrium, most of the \( \text{NH}_4^+ \) remains in its non-ionized form, contributing to a less acidic solution compared to strong acids like \( \text{HCl} \). This limited proton donation means a higher pH than solutions of strong acids, though still below neutral. Understanding the limited dissociation is key in realizing why salts like \( \text{NH}_4\text{Cl} \) result in mildly acidic solutions.

Neutral Salt Solutions

Salts like \( \text{NaCl} \) are classified as neutral salts because, upon dissociation, neither of their constituent ions affects the pH significantly. Sodium ions \( \text{Na}^+ \) and chloride ions \( \text{Cl}^- \) do not react with water to change the concentration of \( \text{H}^+ \) or \( \text{OH}^- \) in the solution. Thus, the pH remains close to neutral, around 7. This inert behavior towards altering water's acidity or basicity sets neutral salts apart from those whose ions can donate or accept protons. Recognizing this neutrality helps clearly place \( \text{NaCl} \) in order within pH trends.