Problem 102 - Chemistry The Central Science

Question

Hydrazine has been employed as a reducing agent for metals. Using standard reduction potentials, predict whether the following metals can be reduced to the metallic state by hydrazine under standard conditions in acidic solution: (a) \(\mathrm{Fe}^{2+},(\mathbf{b}) \mathrm{Sn}^{2+},(\mathbf{c}) \mathrm{Cu}^{2+},(\mathbf{d}) \mathrm{Ag}^{+},(\mathbf{e}) \mathrm{Cr}^{3+},(\mathbf{f}) \mathrm{Co}^{3+}\)

Step-by-Step Solution

Answer

Based on the comparison of reduction potentials, hydrazine can reduce the following metal ions to their metallic state in acidic solution: (a) iron(II), (b) tin(II), (c) copper(II), (e) chromium(III), and (f) cobalt(III). However, hydrazine cannot reduce silver(I) ions to silver under standard conditions in an acidic solution.

Step 1: Find the Standard Reduction Potential of Hydrazine

The standard reduction potential of hydrazine in an acidic solution is given by the following equation: \( N_2H_4 + 4 H^+ + 4 e^- \rightarrow N_2 + 4 H_2O \) For this equation, the standard reduction potential, \(E^0\), is +0.76 V.

Step 2: Determine Standard Reduction Potentials for Metals

The standard reduction potentials for the given metal ions are as follows: (a) \( \mathrm{Fe}^{2+} + 2 e^- \rightarrow \mathrm{Fe} \), \(E^0 = -0.44 \, V\) (b) \( \mathrm{Sn}^{2+} + 2 e^- \rightarrow \mathrm{Sn} \), \(E^0 = -0.14 \, V\) (c) \( \mathrm{Cu}^{2+} + 2 e^- \rightarrow \mathrm{Cu} \), \(E^0 = +0.34 \, V\) (d) \( \mathrm{Ag}^{+} + e^- \rightarrow \mathrm{Ag} \), \(E^0 = +0.80 \, V\) (e) \( \mathrm{Cr}^{3+} + 3 e^- \rightarrow \mathrm{Cr} \), \(E^0 = -0.74 \, V\) (f) \( \mathrm{Co}^{3+} + 3 e^- \rightarrow \mathrm{Co} \), \(E^0 = -0.29 \, V\)

Step 3: Compare the Reduction Potentials

To determine whether hydrazine can reduce the metals, compare the standard reduction potentials of hydrazine and the metal ions. If the standard reduction potential of metal ion is lower than that of hydrazine (i.e., more negative), hydrazine can reduce the metal ion to its metallic state. (a) \(E^0_\mathrm{Hydrazine} > E^0_\mathrm{Fe^{2+}}\), so hydrazine can reduce iron(II) ions to iron. (b) \(E^0_\mathrm{Hydrazine} > E^0_\mathrm{Sn^{2+}}\), so hydrazine can reduce tin(II) ions to tin. (c) \(E^0_\mathrm{Hydrazine} > E^0_\mathrm{Cu^{2+}}\), so hydrazine can reduce copper(II) ions to copper. (d) \(E^0_\mathrm{Hydrazine} < E^0_\mathrm{Ag^{+}}\), so hydrazine cannot reduce silver(I) ions to silver. (e) \(E^0_\mathrm{Hydrazine} > E^0_\mathrm{Cr^{3+}}\), so hydrazine can reduce chromium(III) ions to chromium. (f) \(E^0_\mathrm{Hydrazine} > E^0_\mathrm{Co^{3+}}\), so hydrazine can reduce cobalt(III) ions to cobalt.

Step 4: Summarize the Results

Based on the comparison of reduction potentials, hydrazine can reduce the following metal ions to their metallic state: (a) iron(II), (b) tin(II), (c) copper(II), (e) chromium(III), and (f) cobalt(III). However, hydrazine cannot reduce silver(I) ions to silver under standard conditions in an acidic solution.

Key Concepts

Redox Reactions

Redox reactions, or oxidation-reduction reactions, are chemical processes that involve the transfer of electrons between substances. In these reactions, one substance loses electrons and is oxidized, while another gains electrons and is reduced.
The substance that gets reduced often serves as an oxidizing agent, whereas the one that is oxidized is the reducing agent. Understanding redox reactions is essential in electrochemistry where chemical changes are closely tied to the flow of electric charge.
In the given exercise, hydrazine acts as the reducing agent. This means that it donates electrons to certain metal ions, causing these metals to gain electrons, which results in their reduction to a metallic state. This electron donation is evaluated by comparing standard reduction potentials, which gives us insights into which metals can be effectively reduced by hydrazine.

Metal Reduction

Metal reduction refers to the process of converting metal ions into their metallic, elemental forms through the gain of electrons.
In our exercise, hydrazine serves as a powerful reducing agent that can cause reduction of various metal ions to their pure metals in acidic solutions.
Several factors determine whether a metal can be reduced, including the metal's standard reduction potential, the chemical environment (like pH), and the presence of other reactants.
The comparison of standard reduction potentials is key to predicting metal reduction. Here, if a metal ion has a lower (more negative) standard reduction potential than hydrazine, it indicates that hydrazine can effectively donate electrons to that metal ion. Hydrazine can thus reduce metals like iron, tin, copper, chromium, and cobalt as they have lower potentials than hydrazine, leading to their reduction.

Electrochemistry

Electrochemistry is the branch of chemistry that studies chemical processes which involve the movement of electrons. This field bridges the gap between chemical reactions and electrical energy, enabling the understanding of processes like battery operation, corrosion, and electrolysis.
Central to electrochemistry is the concept of standard reduction potentials. These are measured under standard conditions and are crucial for predicting the direction of redox reactions. A higher reduction potential indicates a greater tendency to gain electrons and be reduced.
In the provided solution, we use standard reduction potentials to predict outcomes of metal reductions. By analyzing these values, we ascertain hydrazine's ability to serve as a reducing agent. The tabulated potentials enable chemists and students alike to determine whether a specific redox reaction will proceed and, consequently, which metals will form from their ions.