Problem 113 - A Complete Resource Book in Chemistry for JEE Main

Question

\([\mathrm{X}]+\mathrm{H}_{2} \mathrm{SO}_{4} \rightarrow[\mathrm{Y}]\), a colourless gas with irritating smell. \([\mathrm{Y}]+\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}+\mathrm{H}_{2} \mathrm{SO}_{4} \rightarrow\) green solution. \([\mathrm{X}]\) and \([\mathrm{Y}]\) are (a) \(\mathrm{SO}_{3}^{-2}, \mathrm{SO}_{2}\) (b) \(\mathrm{Cl}^{-}, \mathrm{HCl}\) (c) \(\mathrm{S}^{-2}, \mathrm{H}_{2} \mathrm{~S}\) (d) \(\mathrm{CO}_{3}^{-2}, \mathrm{CO}_{2}\)

Step-by-Step Solution

Answer

(a) SO3^{-2}, SO2

Step 1: Understanding the Reaction of X with H2SO4

The reaction of a substance X with sulfuric acid, H2SO4, produces a colorless gas Y. This gas is described as having an irritating smell. We need to identify X and Y based on these clues.

Step 2: Identify Y Based on Characteristics

The gas Y is described as colorless and having an irritating smell. One common gas that fits this description is sulfur dioxide, SO2.

Step 3: Determining X From Y's Formation

Since Y is SO2, identify what X can release SO2 when reacting with sulfuric acid. A common reaction yielding SO2 is when sulfite ions (SO3^{2-}) react with acid.

Step 4: Reaction of Y with K2Cr2O7 and H2SO4

The second reaction is described as Y reacting with K2Cr2O7 and H2SO4 to form a green solution. SO2 can reduce K2Cr2O7 (potassium dichromate) in acidic conditions to form Cr^3+ ions, which are green.

Step 5: Match X and Y to Given Options

Based on the reactions, X must be SO3^{2-}, which produces SO2 when reacted with H2SO4. The corresponding options are then to be matched with the descriptions and outcomes provided, leading to identification.

Key Concepts

Sulfur Dioxide

Sulfur dioxide (SO₂) is a fascinating compound with notable properties and reactions. It is a colorless gas known for its pungent, irritating smell, often described akin to burnt matches. This characteristic odor, along with its gaseous state, makes sulfur dioxide a distinctive compound.

• Sulfur dioxide is primarily recognized for its role in the atmosphere as it contributes to air pollution and is a precursor to acid rain.
• When released into the atmosphere, it reacts with water vapor to form sulfurous acid (H₂SO₃), which can further oxidize to sulfuric acid (H₂SO₄), posing environmental hazards.
• In industrial settings, sulfur dioxide is widely used as a preservative and for bleaching.

An interesting aspect of sulfur dioxide is its reducing property, which makes it a key player in various chemical reactions, such as the reduction of potassium dichromate, turning the solution from orange to green.

Potassium Dichromate Reduction

Potassium dichromate (\( ext{K}_2 ext{Cr}_2 ext{O}_7\)) serves as a potent oxidizing agent, widely used across laboratories and in various industrial applications due to its effective reaction capabilities. When it comes in contact with reducing agents like sulfur dioxide in acidic conditions, a fascinating transformation occurs:

• The bright orange color of potassium dichromate is attributed to the presence of the chromium (VI) ion (CrO₄²⁻). When this is reduced, the chromium (VI) changes to chromium (III) (Cr³⁺), which imparts a green color to the solution.
• The reduction process involves the transfer of electrons from the reducing agent, sulfur dioxide, to the oxidizing agent, leading to a change in oxidation state and the characteristic color shift.
• This reaction is not only visually striking but also a classic demonstration of redox reactions in chemistry.

This process is important in analytical chemistry, as the color change serves as an indicator for the presence of specific reducing agents.

Sulfite Ion Reaction

Sulfite ions (\( ext{SO}_3^{2-}\)) are crucial in various chemical reactions, particularly in conjunction with acids like sulfuric acid, where they undergo a transformation leading to the release of a distinct gas.

• When sulfite ions react with sulfuric acid (H₂SO₄), they undergo an acid-base reaction where sulfur dioxide (\( ext{SO}_2\)), water, and sulfate ions (\( ext{SO}_4^{2-}\)) are the end products.
• This reaction is a practical example of how to liberate sulfur dioxide gas from non-gaseous compounds.
• The ability of sulfites to release SO₂ upon acidification is utilized in processes where controlled release of SO₂ is necessary, such as in some winemaking procedures where it acts as an antimicrobial and antioxidant agent.

Understanding this reaction not only helps illuminate core concepts in inorganic chemistry but also highlights practical applications of these reactions in real-world industrial processes.