Problem 118 - A Complete Resource Book in Chemistry for JEE Main

Question

\(\mathrm{HBr}\) and HI reduce sulphuric acid, \(\mathrm{HCl}\) can reduce \(\mathrm{KMnO}_{4}\) and HF can reduce (a) \(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) (b) \(\mathrm{KMnO}_{4}\) (c) \(\mathrm{H}_{2} \mathrm{SO}_{4}\) (d) none of the above

Step-by-Step Solution

Answer

HF reduces none of the given compounds.

Step 1: Identifying Reducing Agents and Oxidizing Agents

The problem asks us to determine the compounds that can be reduced by the halide acids (HBr, HI, HCl, HF). It's crucial to know which compounds act as reducing agents (e.g., HBr, HI, HCl, HF) and which compounds act as oxidizing agents (e.g., K2Cr2O7, KMnO4, H2SO4).

Step 2: Analyzing Reactivity with Halides

Given that HBr and HI can reduce H2SO4, and HCl can reduce KMnO4, we need to analyze the reducing capability of HF. HF is known for being a weak reducing agent, meaning it has limited ability to donate electrons compared to the others.

Step 3: Interpreting the Potential Reaction

HF's position in the halide series indicates its poor reducing ability due to its high bond dissociation energy. This suggests that HF is unlikely to reduce strong oxidizing agents like K2Cr2O7, KMnO4, or H2SO4.

Step 4: Conclusion on HF's Reducing Capabilities

Since HF does not have sufficient reducing strength to reduce K2Cr2O7, KMnO4, or H2SO4, the only logical conclusion is that HF cannot reduce any of the given compounds.

Key Concepts

Oxidizing Agents

Oxidizing agents are substances that gain electrons in a chemical reaction, causing the other substance to lose electrons. They are essential players in redox reactions because they facilitate the process of oxidation. In the context of the exercise, compounds like potassium dichromate (\(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\)), KMnO4, and sulfuric acid (H2SO4) are identified as oxidizing agents.

• Potassium dichromate (\(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\)) is a powerful oxidizing agent in the acidic medium. It can convert sulfur dioxide to sulfuric acid, showing its strong electron acceptance ability.
• Potassium permanganate (KMnO4) is another strong oxidizing agent, especially in acidic solutions where it can change from manganese (VII) to manganese (II) while gaining electrons.
• Sulfuric acid (H2SO4) acts as an oxidizing agent due to its ability to accept electrons and transform other substances during the reaction.

Oxidizing agents are crucial for understanding redox reactions because they define which substances will be oxidized. In our problem, recognizing oxidizing agents helps identify which halide acids can act as reducing agents.

Reducing Agents

Reducing agents are substances that donate electrons in a chemical reaction, allowing other substances to gain electrons. In this exercise, the halide acids (HBr, HI, HCl, HF) are reducing agents, each with varying abilities to donate electrons.

• Hydrobromic acid (HBr) and hydroiodic acid (HI) are strong reducing agents because they can easily donate electrons. They are capable of reducing even strong oxidizing agents like sulfuric acid (H2SO4).


• Hydrochloric acid (HCl), while not as strong as HBr or HI, still acts as a reducing agent, capable of reducing potassium permanganate (KMnO4).


• Hydrofluoric acid (HF) is considered a weak reducing agent. Due to its strong bond and high bond dissociation energy, it struggles to donate electrons, making it unlikely to reduce strong oxidizing agents in this context.

Understanding redox reactions involves recognizing which substances act as reducing agents, as they are fundamental in determining the flow of electrons during reactions. Each acid's ability to reduce different oxidizing agents offers insight into their position in the reactivity series.

Halide Acids

Halide acids are acids that consist of hydrogen paired with a halogen from the periodic table. The common halide acids are HCl (hydrochloric acid), HBr (hydrobromic acid), HI (hydroiodic acid), and HF (hydrofluoric acid). The reactivity and reducing power of these acids vary significantly, influencing their behavior in chemical reactions.

• HCl is frequently used in laboratory settings as a strong acid and is known for its moderate reducing power. It can reduce certain oxidizing agents such as KMnO4.
• HBr and HI are regarded as stronger acids with formidable reducing capabilities. They can reduce even strong oxidizers like H2SO4, demonstrating a high electron-donating ability.
• HF, on the other hand, exhibits weak acidity and reducing capability. Its high bond strength between hydrogen and fluorine makes it less reactive and less effective at reducing other substances.

The differences among these halide acids are crucial for predicting outcomes in redox reactions. It is their unique electron-donating capacities that determine their effectiveness as reducing agents. This understanding helps chemists predict reaction products and design chemical processes.