Problem 147 - A Complete Resource Book in Chemistry for JEE Main

Question

For the electrolysis of \(\mathrm{CuSO}_{4}\) solution which is/are correct? (a) Cathode reaction: \(\mathrm{Cu}^{2+}+2 \mathrm{e}^{-} \rightarrow \mathrm{Cu}\) using \(\mathrm{Cu}\) electrode(b) Anode reaction: \(\mathrm{Cu} \rightarrow \mathrm{Cu}^{2}++2 \mathrm{e}^{-}\)using \(\mathrm{Cu}\) electrode (c) Cathode reaction: \(2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \rightarrow \mathrm{H}_{2}\) using Pt electrode (d) Anode reaction: \(\mathrm{Cu} \rightarrow \mathrm{Cu}^{2+}+2 \mathrm{e}^{-}\)using \(\mathrm{Pt}\) electrode

Step-by-Step Solution

Answer

Correct reactions are (a) for cathode and (b) for anode using Cu electrodes.

Step 1: Identify the electrodes used

The problem provides two types of electrodes: Copper (Cu) electrodes for options (a) and (b), and Platinum (Pt) electrodes for options (c) and (d).

Step 2: Determine the cathode reactions

In electrolysis, the cathode is where the reduction occurs. For option (a), with a Cu electrode, the reduction of Cu ions is represented as \( \mathrm{Cu}^{2+}+2 \mathrm{e}^{-} \rightarrow \mathrm{Cu} \). This reaction is correct. For option (c), using a Pt electrode, it suggests \( 2 \mathrm{H}^{+} + 2 \mathrm{e}^{-} \rightarrow \mathrm{H}_2 \), which is possible if the solution allows for hydrogen reduction, but this is less likely in a \( \mathrm{CuSO}_4 \) system without significant acid presence.

Step 3: Determine the anode reactions

The anode is where oxidation takes place. In option (b), using a Cu electrode, \( \mathrm{Cu} \rightarrow \mathrm{Cu}^{2+} + 2 \mathrm{e}^{-} \), is correct as copper is oxidized. In option (d), using a Pt electrode, the oxidation reaction \( \mathrm{Cu} \rightarrow \mathrm{Cu}^{2+} + 2 \mathrm{e}^{-} \) wouldn't happen as Pt is inert and doesn't participate in the electrode reaction; thus, it incorrectly suggests copper evolving into ions at a non-copper electrode.

Step 4: Evaluate correctness of electrode role in each option

With Cu electrodes, option (a) at the cathode and option (b) at the anode correctly reflect the redox processes that can occur. Option (c) assumes hydrogen evolution, which is not typical in a normal \( \mathrm{CuSO}_4 \) electrolysis unless with significant acidity. Option (d) is incorrect since Pt doesn't provide copper for the anode reaction.

Key Concepts

Cathode Reaction

In an electrolysis process, the cathode is where reduction occurs. During the electrolysis of \( \mathrm{CuSO}_4 \) using a copper electrode, the cathode reaction involves the following process: \( \mathrm{Cu}^{2+} + 2 \mathrm{e}^{-} \rightarrow \mathrm{Cu} \). Here, copper ions in the solution gain electrons and are reduced to solid copper, which gets deposited on the cathode.

Using a platinum electrode, particularly in non-copper systems, the reduction of hydrogen ions can occur, represented as \( 2 \mathrm{H}^{+} + 2 \mathrm{e}^{-} \rightarrow \mathrm{H}_2 \). However, in a standard copper sulfate solution, hydrogen evolution is unlikely unless significant acidity is present to facilitate this reaction.

Anode Reaction

At the anode, oxidation takes place. In a \( \mathrm{CuSO}_4 \) electrolysis using copper electrodes, the anode reaction is \( \mathrm{Cu} \rightarrow \mathrm{Cu}^{2+} + 2 \mathrm{e}^{-} \).
This means solid copper from the anode dissolves into the solution, forming copper ions. This process balances the reduction happening at the cathode.When using alternative electrodes like platinum, the assumption in some reactions might incorrectly treat platinum as actively participating in the reaction. Platinum is an inert electrode, meaning it doesn't react with the solution and only serves to transfer electrons.

Copper Electrodes

Copper electrodes are active electrodes that participate in the electrolysis process. In the setup for \( \mathrm{CuSO}_4 \) electrolysis, copper electrodes are crucial for effective redox reactions.

When copper is used as the anode, it releases copper ions into the solution by losing electrons (oxidation). Simultaneously, at the cathode, these copper ions gain electrons and deposit as metallic copper. This cycling of copper ions helps maintain a balance between the amount of metal deposited and dissolved.

• Anode Reaction: Produces \( \mathrm{Cu}^{2+} \) ions.
• Cathode Reaction: Deposits \( \mathrm{Cu} \) from solution.

This interaction exemplifies how copper electrodes ensure efficient and continuous electrolytic processes in solutions like \( \mathrm{CuSO}_4 \).

Platinum Electrodes

Platinum electrodes are different from copper electrodes because they are inert. This means they do not participate chemically in the reactions during electrolysis. Instead, they simply provide a surface for the reaction to occur and facilitate electron transfer.

While platinum is an excellent conductor and highly resistant to corrosion, it doesn't dissolve or deposit substances during the electrolysis of \( \mathrm{CuSO}_4 \). If hydrogen evolution were to occur, for example in highly acidic solutions, platinum can act as an efficient surface.

• Used for scenarios where the electrode should not dissolve.
• Ideal for facilitating reactions without directly participating.

This property makes platinum suitable for many industrial electrochemical processes where electrodes need to remain stable and unaffected by the chemical reactions.

Redox Processes

Redox processes are the fundamental chemical reactions taking place during electrolysis. Redox stands for reduction-oxidation. By definition:

• Reduction: Gain of electrons, as observed at the cathode during electrolysis.
• Oxidation: Loss of electrons, as seen at the anode in the electrochemical setup.

In \( \mathrm{CuSO}_4 \) electrolysis, the copper electrodes facilitate these redox processes efficiently. Copper ions in the solution gain electrons and are reduced to metallic copper at the cathode. Meanwhile, at the anode, copper metal loses electrons, oxidizing to copper ions.

This cycling of electrons between the cathode and anode is essential for maintaining a continuous flow of current through the electrolytic cell and is the essence of all redox-driven electrochemical reactions.