Textbook Question
Determine whether or not each metal dissolves in 1 M HNO3. For those metals that do dissolve, write a balanced redox reaction showing what happens when the metal dissolves. a. Cu b. Au
Ch.20 - Electrochemistry
Tro5th EditionChemistry: A Molecular ApproachISBN: 9780134874371
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Table of contents
- Ch.1 - Matter, Measurement & Problem Solving117
- Ch.2 - Atoms & Elements106
- Ch.3 - Molecules and Compounds128
- Ch.4 - Chemical Reactions and Chemical Quantities56
- Ch.5 - Introduction to Solutions and Aqueous Solutions84
- Ch.6 - Gases112
- Ch.7 - Thermochemistry96
- Ch.8 - The Quantum-Mechanical Model of the Atom60
- Ch.9 - Periodic Properties of the Elements86
- Ch.10 - Chemical Bonding I: The Lewis Model91
- Ch.11 - Chemical Bonding II: Molecular Shapes, VSEPR & MO Theory85
- Ch.12 - Liquids, Solids & Intermolecular Forces51
- Ch.13 - Solids & Modern Materials48
- Ch.14 - Solutions93
- Ch.15 - Chemical Kinetics114
- Ch.16 - Chemical Equilibrium68
- Ch.17 - Acids and Bases137
- Ch.18 - Aqueous Ionic Equilibrium154
- Ch.19 - Free Energy & Thermodynamics91
- Ch.20 - Electrochemistry105
- Ch.21 - Radioactivity & Nuclear Chemistry74
- Ch.22 - Organic Chemistry141
Chapter 20, Problem 56
Which metal can be oxidized with an Sn2+ solution but not with an Fe2+ solution?
Verified step by step guidance1
Identify the standard reduction potentials (E°) for Sn2+/Sn and Fe2+/Fe from a standard reduction potential table.
Understand that a metal can be oxidized by an ion if the ion has a higher reduction potential than the metal's corresponding ion.
Compare the reduction potential of Sn2+ with the reduction potential of the metal in question to determine if Sn2+ can oxidize the metal.
Similarly, compare the reduction potential of Fe2+ with the reduction potential of the metal in question to determine if Fe2+ can oxidize the metal.
Find a metal whose corresponding ion has a reduction potential between those of Sn2+ and Fe2+, indicating that it can be oxidized by Sn2+ but not by Fe2+.
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Key Concepts
Here are the essential concepts you must grasp in order to answer the question correctly.
Oxidation and Reduction
Oxidation refers to the loss of electrons by a substance, while reduction is the gain of electrons. In redox reactions, one species is oxidized and another is reduced. Understanding these processes is crucial for determining which metals can be oxidized by specific ions, as the ability of a metal to be oxidized depends on its position in the electrochemical series.
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Electrochemical Series
The electrochemical series is a list of metals and their standard electrode potentials, which indicates their tendency to be oxidized or reduced. Metals higher in the series are more easily oxidized than those lower down. This series helps predict which metals can be oxidized by certain ions, such as Sn2+ and Fe2+, based on their relative positions.
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Metal Reactivity
Metal reactivity refers to how readily a metal can lose electrons and undergo oxidation. More reactive metals can displace less reactive metals from their compounds in solution. In the context of the question, identifying a metal that can be oxidized by Sn2+ but not by Fe2+ requires understanding the reactivity of the metals involved and their respective positions in the electrochemical series.
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Related Practice
Textbook Question
Determine whether or not each metal dissolves in 1 M HCl. For those metals that do dissolve, write a balanced redox reaction showing what happens when the metal dissolves. a. Al b. Ag c. Pb
Textbook Question
Which metal could you use to reduce Mn2+ ions but not Mg2+ ions?
Textbook Question
Determine whether or not each redox reaction occurs spontaneously in the forward direction.
a. Ca2+(aq) + Zn(s) → Ca(s) + Zn2+(aq)
b. 2 Ag+(aq) + Ni(s) → 2 Ag(s) + Ni2+(aq)
c. Fe(s) + Mn2+(aq) → Fe2+(aq) + Mn(s)
d. 2 Al(s) + 3 Pb2+(aq) → 2 Al3+(aq) + 3 Pb(s)
Textbook Question
Determine whether or not each metal dissolves in 1 M HIO3. For those metals that do dissolve, write a balanced redox equation for the reaction that occurs. a. Au b. Cr
Textbook Question
Determine whether or not each redox reaction occurs spontaneously in the forward direction.
a. Ni(s) + Zn2+(aq) → Ni2+(aq) + Zn(s)
b. Ni(s) + Pb2+(aq) → Ni2+(aq) + Pb(s)
c. Al(s) + 3 Ag+(aq) → Al3+(aq) + 3 Ag(s)
d. Pb(s) + Mn2+(aq) → Pb2+(aq) + Mn(s)