Domanda del libro di testo
Use tabulated electrode potentials to calculate 𝛥𝐺rxn° for each reaction at 25 °C. a. MnO2(s) + 4 H+(aq) + Cu(s)¡Mn2+(aq) + 2 H2O(l) + Cu2+(aq)
Ch.20 - Electrochemistry
Tro6th EditionChemistry: A Molecular ApproachISBN: 9780137832217
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TOC.tableOfContents
- Ch.1 - Matter, Measurement & Problem Solving118
- Ch.2 - Atoms & Elements106
- Ch.3 - Molecules and Compounds131
- Ch.4 - Chemical Reactions and Chemical Quantities56
- Ch.5 - Introduction to Solutions and Aqueous Solutions82
- Ch.6 - Gases114
- Ch.7 - Thermochemistry96
- Ch.8 - The Quantum-Mechanical Model of the Atom60
- Ch.9 - Periodic Properties of the Elements90
- Ch.10 - Chemical Bonding I: The Lewis Model89
- Ch.11 - Chemical Bonding II: Molecular Shapes, VSEPR & MO Theory85
- Ch.12 - Liquids, Solids & Intermolecular Forces51
- Ch.13 - Solids & Modern Materials47
- Ch.14 - Solutions90
- Ch.15 - Chemical Kinetics115
- Ch.16 - Chemical Equilibrium68
- Ch.17 - Acids and Bases136
- Ch.18 - Aqueous Ionic Equilibrium153
- Ch.19 - Free Energy & Thermodynamics93
- Ch.20 - Electrochemistry106
- Ch.21 - Radioactivity & Nuclear Chemistry74
- Ch.22 - Organic Chemistry140
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Capitolo 20, Problema 69c
Use tabulated electrode potentials to calculate 𝛥𝐺rxn° for each reaction at 25 °C. c. O2( g) + 2 H2O(l) + 2 Cu(s)¡4 OH-(aq) + 2 Cu2+(aq)
Guida verificata passo dopo passo1
Identify the half-reactions involved in the given redox reaction. The reaction involves the reduction of O2 to OH- and the oxidation of Cu to Cu2+.
Write the half-reactions: \( \text{O}_2(g) + 4e^- + 2\text{H}_2\text{O}(l) \rightarrow 4\text{OH}^-(aq) \) and \( 2\text{Cu}(s) \rightarrow 2\text{Cu}^{2+}(aq) + 4e^- \).
Look up the standard electrode potentials (E°) for each half-reaction from a table of standard electrode potentials. For the reduction of O2, \( E^\circ = +0.40 \text{ V} \) and for the oxidation of Cu, \( E^\circ = -0.34 \text{ V} \).
Calculate the standard cell potential (E°cell) using the formula: \( E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}} \). Substitute the values from the previous step.
Use the relationship between the standard cell potential and the standard Gibbs free energy change: \( \Delta G^\circ_{\text{rxn}} = -nFE^\circ_{\text{cell}} \), where \( n \) is the number of moles of electrons transferred (4 in this case) and \( F \) is the Faraday constant (approximately 96485 C/mol). Substitute the values to find \( \Delta G^\circ_{\text{rxn}} \).
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Electrode Potentials
Electrode potentials, measured in volts, indicate the tendency of a chemical species to be reduced or oxidized. Standard electrode potentials (E°) are measured under standard conditions and are crucial for predicting the direction of redox reactions. The more positive the electrode potential, the greater the species' ability to gain electrons, thus driving the reaction forward.
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Percorso guidato
01:27
Standard Cell Potential
Gibbs Free Energy (ΔG)
Gibbs Free Energy (ΔG) is a thermodynamic quantity that indicates the spontaneity of a reaction at constant temperature and pressure. A negative ΔG value signifies that a reaction is spontaneous, while a positive value indicates non-spontaneity. The relationship between ΔG and electrode potentials is given by the equation ΔG° = -nFE°, where n is the number of moles of electrons transferred, F is Faraday's constant, and E° is the standard cell potential.
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Percorso guidato
01:51Gibbs Free Energy of Reactions
Nernst Equation
The Nernst equation relates the cell potential of an electrochemical reaction to the concentrations of the reactants and products. It allows for the calculation of the cell potential under non-standard conditions, which can affect the Gibbs Free Energy. The equation is expressed as E = E° - (RT/nF)ln(Q), where Q is the reaction quotient, R is the gas constant, and T is the temperature in Kelvin.
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01:17The Nernst Equation
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