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Ch.19 - Electrochemistry
Tro - Chemistry: A Molecular Approach 4th Edition
Tro4th EditionChemistry: A Molecular ApproachISBN: 9780134112831Not the one you use?Change textbook
All textbooksTro 4th EditionCh.19 - ElectrochemistryProblem 46
Chapter 19, Problem 46

Calculate the standard cell potential for each of the electrochemical cells in Problem 44.

Verified step by step guidance
1
Identify the half-reactions involved in the electrochemical cell from Problem 44.
Write the reduction half-reaction and its standard reduction potential, E° (from a standard reduction potential table).
Write the oxidation half-reaction and its standard reduction potential, E° (from a standard reduction potential table).
Calculate the standard cell potential, E°_cell, using the formula: E°_cell = E°_cathode - E°_anode.
Ensure that the units are consistent and the final E°_cell is expressed in volts (V).

Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Standard Cell Potential

The standard cell potential, denoted as E°, is the measure of the voltage produced by an electrochemical cell under standard conditions (1 M concentration, 1 atm pressure, and 25°C). It indicates the tendency of a chemical reaction to occur spontaneously; a positive E° value suggests a spontaneous reaction, while a negative value indicates non-spontaneity.
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Standard Cell Potential

Nernst Equation

The Nernst equation relates the cell potential to the concentrations of the reactants and products in an electrochemical reaction. It allows for the calculation of the cell potential under non-standard conditions, showing how changes in concentration affect the voltage. The equation is given by E = E° - (RT/nF)ln(Q), where Q is the reaction quotient.
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Electrochemical Cells

Electrochemical cells consist of two half-cells, each containing an electrode and an electrolyte. The oxidation and reduction reactions occur at the anode and cathode, respectively. Understanding the components and functions of these cells is crucial for calculating the standard cell potential, as it involves identifying the correct half-reactions and their standard potentials.
Related Practice
Textbook Question

Balance each redox reaction occurring in basic aqueous solution. c. NO2–(aq) + Al(s) → NH3(g) + AlO2–(aq)

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Textbook Question

Consider the voltaic cell:

d. Indicate the direction of anion and cation flow in the salt bridge

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Textbook Question

Sketch a voltaic cell for each redox reaction. Label the anode and cathode and indicate the half-reaction that occurs at each electrode and the species present in each solution. Also indicate the direction of electron flow.

a. 2 Ag+(aq) + Pb(s) → 2 Ag(s) + Pb2+(aq)

b. 2 ClO2(g) + 2 I–(aq) → 2 ClO2–(aq) + I2(s)

c. O2(g) + 4 H+(aq) + 2 Zn(s) → 2 H2O(l) + 2 Zn2+(aq)

Textbook Question

Use line notation to represent each electrochemical cell in Problem 43.

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Textbook Question

Sketch a voltaic cell for each redox reaction. Label the anode and cathode and indicate the half-reaction that occurs at each electrode and the species present in each solution. Also indicate the direction of electron flow.

b. 2 H+(aq) + Fe(s) → H2(g) + Fe2+(aq)

c. 2 NO3–(aq) + 8 H+(aq) + 3 Cu(s) → 2 NO(g) + 4 H2O(l) + 3 Cu2+(aq)