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Ch.20 - Electrochemistry
All textbooksTro 5th EditionCh.20 - ElectrochemistryProblem 45
Chapter 20, Problem 45

Calculate the standard cell potential for each of the electro- chemical cells in Problem 43.

Verified step by step guidance
1
First, identify the half-reactions involved in the electrochemical cell. The half-reactions are usually given in the problem or can be determined from the cell diagram.
Next, look up the standard reduction potentials for each half-reaction in a table of standard reduction potentials. These tables can be found in your textbook or online.
Remember that the standard cell potential is calculated by subtracting the standard reduction potential of the reaction at the anode (where oxidation occurs) from the standard reduction potential of the reaction at the cathode (where reduction occurs). This can be represented by the equation: E°cell = E°cathode - E°anode.
Substitute the standard reduction potentials you found in step 2 into the equation from step 3. Be careful with the signs of the potentials, as they can be positive or negative.
Finally, calculate the standard cell potential. This will give you the maximum potential difference, or voltage, that the cell can produce under standard conditions. Remember that a positive cell potential indicates a spontaneous reaction, while a negative cell potential indicates a non-spontaneous reaction.

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Key Concepts

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

Standard Cell Potential

The standard cell potential (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 is calculated using the standard reduction potentials of the half-reactions involved in the cell. A positive E° indicates a spontaneous reaction, while a negative E° suggests non-spontaneity.
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Electrochemical Cells

Electrochemical cells consist of two half-cells, each containing an electrode and an electrolyte. The oxidation reaction occurs at the anode, while the reduction reaction takes place at the cathode. The flow of electrons from the anode to the cathode through an external circuit generates electrical energy, which can be harnessed for work.
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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 and is expressed as E = E° - (RT/nF)ln(Q), where R is the gas constant, T is the temperature in Kelvin, n is the number of moles of electrons transferred, F is Faraday's constant, and Q is the reaction quotient.
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Related Practice
Textbook Question

Balance each redox reaction occurring in basic aqueous solution. b. Ag(s) + CN–(aq) + O2(g) → Ag(CN)2–(aq)

Textbook Question

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

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. Ni2+(aq) + Mg(s) → Ni(s) + Mg2+(aq)

Textbook Question

Consider the voltaic cell:

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

Textbook Question

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