ELC_Standard Reduction Potentials Lesson

Standard Reduction Potentials

We mentioned on the previous page that something called a voltmeter can be attached to a galvanic cell. A voltmeter is a device that measures voltage. Voltage is the ability to push electrons through the circuit. In physics you learn that voltage is also called potential difference. This is because it comes from the difference in electrical potential energy for each cell. Voltage or potential difference is measured in volts (V). A volt is equal to one joule of energy per coulomb.

The maximum amount of potential that a galvanic cell can create is known as the  cell potential, Ecell. The cell potential is dependent upon several factors including the concentration of ions in each half-cell, the composition of the electrodes, and the temperature. Because each of these factors change the potential, the  standard cell potential, E°cell , is useful. These are the same standard conditions used in thermochemistry, 25°C and 1 M solutions.

Rather than having a list of cell potentials for oxidation and a different one for reduction, we use a list of reduction potentials for all half-reactions. This list is known as standard reduction potentials. It shows the measure of that half-reaction's tendency to acquire electrons, or undergo reduction. When comparing two half-reactions, the reaction with the most positive reduction potential occurs as written, as a reduction. The other half-cell reaction occurs in reverse, as oxidation. The overall cell potential is the difference between the two half-cells' potential.

CELL POTENTIAL E° cell = (standard reduction potential of the substance reduced) - (standard reduction potential of the substance oxidized) Eᵒcell Eᵒreduction - Eᵒoxidation

Because it is impossible to measure the standard reduction potential of a half-cell alone, a reference electrode, a  standard hydrogen electrode, is used. The standard reduction potentials are available in the table shown here. You will also find a similar table in your book.

Half Reaction Chart

Example: If Cl2 and Br2 are added to a solution that contains both CI- and Br ions, what spontaneous reaction will occur? What is the standard cell potential of the overall reaction that occurs? First, look for the two half-reactions on the standard reduction potential table. Cl2(g) + 2e2Cl(aq) E°= 1.36 V Br2(g) + 2e2 Br(aq) E°= 1.07 V Cl₂ has the more positive value of E°, therefore, it will be reduced. The half-reaction for Br₂ will be reversed, and Br2 will be oxidized. The spontaneous reactions in this cell are: Cl2(g) + 2e2Cl(aq) 2 Br (aq) → Br2(g) + 2e The net reaction is obtained by adding these two half reactions together. Cl2(g) + 2 Br(aq) → Br2(g) + 2Cl(aq) Now let's calculate the cell potential using this equation: E cell = E reduction - Eᵒoxidation E cell = 1.36 V-1.07 V E° cell = 0.29 V Positive cell potentials indicate that the reaction is spontaneous as written. If the cell potential is negative, the reaction will be spontaneous in the reverse direction.

There are a couple of important differences in calculating E°cell and calculating standard thermodynamic values like ∆H°, ∆S°, and ∆G °. When calculating E°cell using the equation E°cell = Ereduction- Eoxidation, you do NOT have to change the sign of E° for the oxidation reaction even though that reaction is reversed. The change in sign is already built in to the equation!

If you have to double, triple, etc., the coefficients in the half-reaction to match your reaction, you do NOT change the values for the standard potentials you look up on the chart. This is because reduction potentials are intensive quantities. Their unit, volt, is a Joule per Coulomb. The same number of joules are available for each coulomb of charge regardless of the total number of electrons shown in the equation.

Therefore, reduction potentials are never multiplied by factors before they are subtracted to give the cell potential!

Remember to work on the module practice problems as you complete each section of content.  

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