1. Standard Electrode Potentials

2. Metals, because they only have a few valence electrons, like to lose electrons. In other words they tend to be easily oxidized. But metals differ in how easily they lose electrons. A list of metals arranged in order of how easily the metal is oxidized is known as an activity series

3. The fact that different substances are oxidized more readily than others is the driving force behind electrochemical cells, and it is this force that forces electrons through the external circuit from the anode (site of oxidation) to the cathode (site of reduction). This force is known as the potential difference or electromotive force (emf or E).

4. This force is known as the potential difference or electromotive force (emf or E). Potential difference is measured in volts (V), and thus is also referred to as the voltage of the cell. Voltage is a measure of the tendency of electrons to flow. The higher the voltage, the greater the tendency for electrons to flow from the anode to the cathode.

5. Tables of Standard Reduction Potentials for Half-Reactions allow us to determine the voltage of electrochemical cells. These tables compare the ability of different half- reactions to compete for electrons (become reduced). Since half-reactions cannot occur on their own, all values in the table are determined by comparing a half-reaction with a hydrogen half-cell:
2H+(aq) + 2e- → H2 (g)    E° = 0.00 V
the degree symbol following the E (E°) indicates standard conditions: temperature = 25°C; pressure = 100 kPa; concentration of aqueous solutions = 1 mol·L-1

6. For example, if copper and hydrogen half-cells are joined together we find that the copper half-cell will gain electrons from the hydrogen half-cell. Thus the copper half-cell is given a positive voltage and given a relative value of V:
Since both half-reactions cannot undergo reduction, we must reverse the equation of the reaction that will undergo oxidation. This will give us an electrochemical cell voltage of V

7. Calculating Voltages of Electrochemical Cells
Before calculating the voltage of a cell you must first determine which half-cell will undergo oxidation and which one will be reduced. Find the copper and zinc half- reactions in the Table of Standard Reduction Potentials. Be careful - there is often more than one half-reaction for an element. For copper locate the Cu|Cu2+half-reaction

9. When creating an electrochemical cell, always reverse the half-reaction that will result in a positive value for E° when the equations are added together.

10. A positive value of E° indicates a spontaneous chemical reaction
Electrochemical cells always involve a spontaneous chemical reaction

11. Let's try one more example of setting up an electrochemical cell.
We want to create an electrochemical cell using aluminum (Al|Al3+) and lead (Pb|Pb2+) half-cells. Our tasks:
Determine the two half-reactions involved, and which reaction will undergo oxidation and which one will be reduced.
Determine the voltage of the cell.
Diagram the set-up of the electrochemical cell, including the following items: 
the two half-cells, including the electrodes and electrolytic solutions
the external circuit, showing the direction of electron flow
the salt bridge with an electrolyte, including movement of ions
label the anode and the cathode
label the positive and negative posts

12. Activity Series
An activity series is often useful is helping to predict whether certain reactions will occur. An activity series list metals and various other elements in order of their reactivity, with the most reactive elements at the top and the least reactive of the series at the bottom. For any two metals, the metal listed higher in the table is the most readily oxidized.