1. Unit 8: Electrochemistry Applications
Lesson 2:
Standard Reduction Potentials

2. Measuring Voltage
Voltage, or electrical potential, is the work done per unit of charge (made up of electrons) transferred.
We can’t measure the voltage of a lone half-cell (no e- flow!), but we can measure the voltage between two half-cells.
The hydrogen half-cell is arbitrarily defined as the zero point, and all other half-cell reactions are compared to it.

3. Standard Reduction Potential
E° is the standard reduction potential in volts.
Standard conditions are defined to be:
25°C
All gases at 1 atm
All elements in their standard states
1 M solutions

4. Calculating E°cell
When we add together two half-reactions, we can also add together their voltages.
Note that when you reverse a half-reaction, you must reverse the sign of its E° value.
If the potential of the cell (E°cell) is positive, then the reaction will be spontaneous.
If E°cell is negative, the reaction is non-spontaneous.

5. Examples What is the potential of the following cell?
Ni2+ + Fe → Ni + Fe2+
Your turn! Calculate the potential of each cell.
a) Ni + Fe2+ → Ni2+ + Fe
b) 3 Ag+ + Al → 3 Ag + Al3+
*Note: since voltage is work done per unit charge, it does not need to be multiplied when you multiply a half-reaction to cancel electrons.
25 min

6. Brain Break!
Double Shutter

7. The Big Picture You must consider all species present in the cell.
If a reaction occurs in an acidic solution, the reduction of H+ is a possible reaction.
If a reaction occurs in a neutral solution, the reduction of neutral water is a possible reaction.

8. Cells and Equilibrium
As a cell operates, the [reactant ion] decreases and the [product ion] increases. The solutions do not remain at 1 M!
Because of the changes in concentration, the E of the reduction reaction decreases, and the E of the oxidation reaction increases, until the two values are equal but opposite, and Ecell is zero. The cell is now at equilibrium.

9. Practice
Pg. 224 #36
5 min