Strengths of Oxidizing and Reducing Agents

The strengths of oxidizing and reducing agents are indicated by their standard electrode potentials. A sample from the table of standard potentials shows the extremes of the table.

Cathode (Reduction) Half-Reaction	Standard Potential E^° (volts)
Li⁺(aq) + e^- → Li(s)	-3.04
K⁺(aq) + e^- → K(s)	-2.92
Ca²⁺(aq) + 2e^- → Ca(s)	-2.76
Na⁺(aq) + e^- → Na(s)	-2.71
Zn²⁺(aq) + 2e^- → Zn(s)	-0.76
Cu²⁺(aq) + 2e^-→ Cu(s)	0.34
O₃(g) + 2H⁺(aq) + 2e^- → O₂(g) + H₂O(l)	2.07
F₂(g) + 2e^-→ 2F^-(aq)	2.87

The values for the table entries are reduction potentials, so lithium at the top of the list has the most negative number, indicating that it is the strongest reducing agent. The strongest oxidizing agent is fluorine with the largest positive number for standard electrode potential.

Table of Standard Electrode Potentials

Index

Oxidation/
Reduction concepts

Electrochemistry concepts

Reference
Hill & Kolb
Ch 8

Ebbing
Ch 19

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Free Energy and Electrode Potentials

The cell potential of a voltaic cell is a measure of the maximum amount of energy per unit charge which is available to do work when charge is transferred through an external circuit. This maximum work is equal to the change in Gibbs free energy, ΔG, in the reaction. These relationships can be expressed as

Maximum work = ΔG = -nFE^°_cell

where n is the number of electrons transferred per mole and F is the Faraday constant.

Consider the historic Daniell cell in which zinc and copper were used as electrodes. The data from the table of standard electrode potentials is

Cathode (Reduction) Half-Reaction	Standard Potential E^° (volts)
Zn²⁺(aq) + 2e^- → Zn(s)	-0.76
Cu²⁺(aq) + 2e^- → Cu(s)	0.34

The standard cell potential is then E^°_cell = 1.1 volt and 2 electrons are transferred per mole of reactant. The change in free energy is then

ΔG = -nFE^°_cell = -2 x 96,485 coul/mole x 1.10 joule/coul = -212 kJ

This relationship with free energy can be used in the opposite direction as well. From a table of thermodynamic quantities, the free energy changes for the ions under standard conditions are

Zn²⁺(aq), ΔG = -147.21 kJ/mol

Cu²⁺(aq), ΔG = 64.98 kJ/mol

Since the Zn ion is produced and the Cu ion is reduced in the cell process, the net change in free energy is -212 kJ/mol, as we obtained above. Starting from these free energy changes, we could have calculated the cell potential of 1.1 volts by reversing the above calculation.

Table of Standard Electrode Potentials

Index

Oxidation/
Reduction concepts

Electrochemistry concepts

Reference
Hill & Kolb
Ch 8

Ebbing
Ch 19

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Electrode Potentials and Equilibrium Constants

Maximum work = ΔG = -nFE^°_cell

where n is the number of electrons transferred per mole and F is the Faraday constant.

This free energy change can also be related to the equilibrium constant K

ΔG = -RT ln K

Combining these relationships allows us to express the cell potential in terms of the equilibrium constant.

Consider the historic Daniell cell in which zinc and copper were used as electrodes. The data from the table of standard electrode potentials is

Cathode (Reduction) Half-Reaction	Standard Potential E^° (volts)
Zn²⁺(aq) + 2e^- → Zn(s)	-0.76
Cu²⁺(aq) + 2e^- → Cu(s)	0.34

The standard cell potential is then E^°_cell = 1.1 volt and 2 electrons are transferred per mole of reactant. The relationship for the equilibrium constant is then

This extremely high value for the equilibrium constant confirms that the reaction of the Daniell cell is indeed spontaneous and that it will proceed until the reactants are exhausted.

Table of Standard Electrode Potentials

Index

Oxidation/
Reduction concepts

Electrochemistry concepts

Reference
Hill & Kolb
Ch 8

Ebbing
Ch 19

HyperPhysics***** Electricity and Magnetism ***** Chemistry

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