Problem 37 - A Complete Resource Book in Chemistry for JEE Main

Question

When a Zn strip is placed in \(\mathrm{CuSO}_{4}\) solution, Cu gets precipitated, because standard oxidation potential of \(\mathrm{Zn}\) is (a) \(<\mathrm{Cu}\) (b) \(>\mathrm{Cu}\) (c) \(<\mathrm{SO}_{4}\) \((\mathrm{d})>\mathrm{SO}_{4}\)

Step-by-Step Solution

Answer

(b) \(> \text{Cu} \)

Step 1: Understanding the Reaction

When you place a zinc strip in a copper sulfate solution, a redox reaction occurs. Zinc metal, Zn, reacts with aqueous copper ions, \( \text{Cu}^{2+} \), to form zinc ions, \( \text{Zn}^{2+} \), and solid copper metal, Cu. This displacement happens because zinc is more reactive than copper.

Step 2: Standard Electrode Potentials

The potential of the zinc electrode and the copper electrode are defined by their standard reduction potentials: \( E^0(\text{Zn}^{2+}/\text{Zn}) = -0.76 \text{ V} \) and \( E^0(\text{Cu}^{2+}/\text{Cu}) = +0.34 \text{ V} \). Here we can see that zinc has a more negative standard reduction potential than copper, meaning it is easier for zinc to lose electrons and get oxidized.

Step 3: Comparing Oxidation Potentials

The oxidation potential is the negative of the reduction potential. Therefore, the oxidation potential of zinc is \( E^0(\text{Zn}/\text{Zn}^{2+}) = +0.76 \text{ V} \) and for copper, \( E^0(\text{Cu}/\text{Cu}^{2+}) = -0.34 \text{ V} \). The oxidation potential of zinc is greater than that of copper, allowing zinc to displace copper from the solution.

Key Concepts

Standard Electrode Potentials

Standard electrode potentials are a measure of the ability of a half-cell to be reduced, which essentially refers to its tendency to gain electrons. These potentials are measured under standard conditions, which typically include a concentration of 1 M, a pressure of 1 atm, and a temperature of 25°C. In electrochemical cells, these standard potentials (denoted by \( E^0 \)) are used to determine the direction of electron flow. For our reaction:

• The standard reduction potential for \( \text{Zn}^{2+}/\text{Zn} \) is \( -0.76 \text{ V} \).
• For \( \text{Cu}^{2+}/\text{Cu} \), it is \( +0.34 \text{ V} \).

The more negative the \( E^0 \) value, the less noble (more likely to donate electrons) the metal is. Zinc, having a more negative value, is more likely to oxidize compared to copper, thereby driving the displacement reaction.

Displacement Reactions

A displacement reaction, also known as a single replacement reaction, involves one element being replaced by another in a compound. These reactions occur based on the reactivity of elements involved. In the case of zinc and copper sulfate, zinc displaces copper from its compound because zinc is more reactive:

• The reactivity series of metals is a helpful guide to predict displacement reactions.
• Zinc lies above copper in the reactivity series, indicating that it can replace copper ions in the solution to form \( \text{ZnSO}_4 \) and solid copper \( \text{Cu} \).

This reaction is also an example of a redox process, where zinc gets oxidized, and copper ions are reduced.

Electrochemistry

Electrochemistry is the branch of chemistry that studies the relationship between electricity and chemical changes. It involves redox reactions, where the movement of electrons is a key activity. In our exercise, the redox reaction between zinc and copper ions is measured electrochemically by their respective electrode potentials. Some key points to remember include:

• Redox reactions involve the transfer of electrons from one substance to another.
• The species that loses electrons is oxidized, while the one gaining electrons is reduced.
• Electrochemical cells, like galvanic cells, allow us to harness the energy from these reactions.

Understanding electrochemistry helps in applications like batteries, chemical synthesis, and corrosion prevention.

Oxidation Potential

Oxidation potential is essentially the tendency of a species to lose electrons — the opposite of reduction potential. It is crucial for determining which metals can effectively act as reducing agents. To switch the standard reduction potential values to oxidation potentials, you simply reverse the sign:

• Oxidation potential for \( \text{Zn/}\text{Zn}^{2+} \) is \( +0.76 \text{ V} \).
• For \( \text{Cu/}\text{Cu}^{2+} \), it is \( -0.34 \text{ V} \).

Zinc, with a higher oxidation potential, is more willing to lose electrons compared to copper, and this underpins its ability to displace copper in the reaction. Understanding oxidation potential alongside reduction potential allows chemists to predict the spontaneity of redox processes, such as displacement reactions.