Question

The values of standard oxidation potentials of following reactions are given below: \(\mathrm{Zn} \longrightarrow \mathrm{Zn}^{2+}+2 \mathrm{e}^{-} ; E^{\circ}=0.762 \mathrm{~V}\) \(\mathrm{Fe} \longrightarrow \mathrm{Fe}^{2+}+2 \mathrm{e}^{-} ; E^{\circ}=0.440 \mathrm{~V}\) \(\mathrm{Cu} \longrightarrow \mathrm{Cu}^{2+}+2 \mathrm{e}-E^{\circ}=-0.345 \mathrm{~V}\) \(\mathrm{Ag} \longrightarrow \mathrm{Ag}^{+}+2 \mathrm{e}^{-} ; E^{\circ}=-0.800 \mathrm{~V}\) Which of the following is most easily reduced? (a) \(\mathrm{Fe}^{2+}\) (b) \(\mathrm{Ag}^{+}\) (c) \(\mathrm{Zn}^{2+}\) (d) \(\mathrm{Cu}^{2+}\)

Short answer

(b) \( \mathrm{Ag^{+}} \) is most easily reduced.

Step-by-step solution

Understanding Reduction Potentials

Reduction potentials indicate the tendency of a species to gain electrons and thus become reduced. A more positive reduction potential suggests a greater likelihood of the species being reduced.

Converting Oxidation to Reduction Potentials

The given reactions are written as oxidations, so we need to reverse them to represent reductions and change the signs of their potentials: 1. For Zn, \( \mathrm{Zn^{2+} + 2e^{-} \rightarrow Zn} \); \( E^{\circ} = -0.762 \, \mathrm{V} \)2. For Fe, \( \mathrm{Fe^{2+} + 2e^{-} \rightarrow Fe} \); \( E^{\circ} = -0.440 \, \mathrm{V} \)3. For Cu, \( \mathrm{Cu^{2+} + 2e^{-} \rightarrow Cu} \); \( E^{\circ} = 0.345 \, \mathrm{V} \)4. For Ag, \( \mathrm{Ag^{+} + e^{-} \rightarrow Ag} \); \( E^{\circ} = 0.800 \, \mathrm{V} \)

Comparing Reduction Potentials

Now compare the reduction potentials calculated in Step 2:- Zn: \(-0.762 \, \mathrm{V}\)- Fe: \(-0.440 \, \mathrm{V}\)- Cu: \(0.345 \, \mathrm{V}\)- Ag: \(0.800 \, \mathrm{V}\)The most positive value indicates the species that is most easily reduced.

Determining the Most Easily Reduced Species

From the potentials above, \( \mathrm{Ag^{+}} \) with a reduction potential of \( 0.800 \, \mathrm{V} \) has the highest value. Therefore, \( \mathrm{Ag^{+}} \) is the species most easily reduced.

Key concepts

Reduction Potentials

Reduction potentials are a crucial concept in understanding electrochemical reactions. They measure the tendency of a chemical species to acquire electrons and subsequently be reduced. This is represented by the standard reduction potential, denoted as \( E^{ ext{}} \), which is measured in volts (V).

The more positive the reduction potential, the greater the species' tendency to gain electrons. Therefore, when comparing different reduction potentials, the species with the highest (most positive) value will be the easiest to reduce.

This tendency to be reduced is directly opposed to oxidation potential, which measures the tendency to lose electrons. In practice, if we look at a standard electrochemical table, we can compare various ions and elements by their reduction potentials to predict metal and ion behaviour in redox reactions.

In applications such as galvanic cells, the substance with a higher reduction potential acts as the oxidizing agent. Understanding and utilizing reduction potentials allows chemists to predict and manipulate chemical reactions effectively.

Oxidation Potentials

Oxidation potentials refer to the propensity of a chemical species to lose electrons, hence undergoing oxidation. The oxidation potential is essentially the negative of the reduction potential. To determine the oxidation potential from the reduction potential, simply reverse the reaction and change the sign of the potential.

For example, if the process \( \mathrm{Cu^{2+} + 2e^{-} \rightarrow Cu} \) has a reduction potential of \( 0.345 \, ext{V} \), its oxidation potential would be \( -0.345 \, ext{V} \) for the reaction \( \mathrm{Cu \rightarrow Cu^{2+} + 2e^{-}} \).

Oxidation potentials are important when analyzing galvanic or voltaic cells where the oxidation reaction occurs at the anode. In such cells, metals with lower (more negative) reduction potentials are more likely to be oxidized, meaning they lose electrons more readily.

Understanding both oxidation and reduction potentials is key in predicting reaction trends and behaviors in electrochemical cells. This knowledge can also be used to prevent corrosion and in various industrial processes that require precise control over redox reactions.

Redox Reactions

Redox reactions, short for reduction-oxidation reactions, are chemical processes where the transfer of electrons occurs, involving both reduction and oxidation. One species loses electrons (oxidation) while another gains electrons (reduction).

In these reactions, oxidation and reduction always occur simultaneously -- you cannot have one without the other. The actual movement of electrons from the reducing agent to the oxidizing agent is what drives the reaction forward. This simultaneous process is also depicted in cell diagrams and standard electrode potentials tables.

For instance, in a battery, redox reactions are harnessed to produce electrical energy. Inside the battery, one material is oxidized (giving away electrons) while another is reduced (accepting electrons), driving the electric current through an external circuit.

Understanding the principles of redox reactions such as identifying the oxidation and reduction reactions and calculating net cell potential, helps in various fields such as environmental science, biology, chemical engineering, and beyond. By mastering these concepts, you can predict the feasibility of reactions and design systems that capture, use, or even store energy efficiently.