Problem 47 - 43 Years JEE ADVANCED (1978-2020) + JEE MAIN Chapterwise & Topicwise Solved Papers Chemistry

Question

Account for the following. Limit your answer to two sentences (i) Hydrogen bromide cannot be prepared by action of concentrated sulphuric acid or sodium bromide. (ii) When a blue litmus paper is dipped into a solution of hypochlorous acid, it first turns red and then later gets decolourised. [1979]

Step-by-Step Solution

Answer

(i) Bromine gas forms instead of HBr due to sulfuric acid's oxidizing nature. (ii) Litmus turns red due to acidity, then gets decolorized by bleaching.

Step 1: Understanding Gas Formation Constraints

Hydrogen bromide (HBr) cannot be prepared using concentrated sulfuric acid and sodium bromide because the experiment would result in the formation of bromine gas (Br₂) instead of HBr. Sulfuric acid is a strong oxidizing agent and would oxidize bromide ions ( ext{Br}⁻) to bromine ( ext{Br}_2) under these conditions.

Step 2: Red to Colorless: Reaction with Litmus Paper

When blue litmus paper is dipped into hypochlorous acid (HOCl), the solution is acidic, causing the litmus paper to turn red initially. The hypochlorous acid then has a bleaching effect, where the ClO⁻ ions bleach the red litmus paper, turning it colorless.

Key Concepts

Sulfuric Acid Oxidizing Agent

Sulfuric acid (\(H_2SO_4\)) is known for its strong oxidizing properties, meaning it can readily accept electrons from other substances during chemical reactions. When it interacts with compounds like sodium bromide (\(NaBr\)), it doesn't just stop at breaking bonds; it further changes the substances' chemical nature. This is why sulfuric acid cannot be used to prepare hydrogen bromide directly from sodium bromide.

• Instead of simply converting sodium bromide to hydrogen bromide, sulfuric acid strongly oxidizes the bromide ions (\(Br^−\)).
• This process results in the formation of bromine gas (\(Br_2\)), a brownish-red vapor, which is different from the desired hydrogen bromide gas.
• This reaction showcases the oxidizing power of sulfuric acid, demonstrating its ability to transform bromide ions into elemental bromine.

Understanding the oxidative behavior of sulfuric acid in this context is crucial for predicting reaction products and ensuring safe laboratory practices.

Bromide Ion Oxidation

Bromide ions (\(Br^-\)) are the anionic form of bromine, stable yet capable of undergoing oxidation to become bromine molecules (\(Br_2\)). In the presence of strong oxidizing agents like sulfuric acid, oxidation occurs because the bromide ion can readily donate its electron.

• This electron removal results in the pairing of bromide ions to form bromine gas, a diatomic molecule.
• It highlights a key chemical property of bromide ions: their susceptibility to oxidation when placed in a sufficiently reactive environment.
• Understanding this reaction is fundamental in chemistry when manipulating halide ions and predicting their behavior under oxidizing conditions.

The conversion of bromide ions into elemental bromine is a classic example of oxidation, demonstrating the movement of electrons in chemical reactions.

Hypochlorous Acid Reaction

Hypochlorous acid (\(HOCl\)) is a weak acid known for its bleaching and disinfecting properties. When you dip blue litmus paper into a solution of hypochlorous acid, it undergoes two distinct changes that provide a clear demonstration of the acid's behavior.

• As an acidic solution, it initially turns the blue litmus paper red, indicating the presence of acidity by lowering the pH around the litmus.
• Over time, hypochlorous acid not only lowers the pH but also bleaches the litmus paper due to its oxidative properties.
• Here, the bleaching effect occurs as the hypochlorite ion (\(ClO^-\)), a byproduct of hypochlorous acid, reacts with the dyes in the litmus paper, removing their colors.

This reaction is a useful reminder of the dual properties of hypochlorous acid: both acidic, turning litmus red, and oxidative, leading to bleaching.