Standard cell potentials find many uses as sources of thermodynamic information of interest. We have already encountered the relation ΔGº_{r} = – νFEº. From this, and the thermodynamic relation (δG/δT)_{P} = – S , we obtain:

we can now use the fundamental relation ΔGº_{r} = ΔHº_{r} – T ΔSº_{r} to write

Standard potentials may also be used to construct an electrochemical series:

Consider a general cell, composed of two redox couples Ox_{1}/Red_{1} and Ox_{2}/Red_{2} , with associated standard potentials Eº_{1} and Eº_{2}. We know that for the following cell, the cell potential may be written as shown:

Red_{1} | Ox_{1} || Ox_{2} | Red_{2} , Eº = Eº_{2} – Eº_{1}

Further, we know that the cell reaction, Red_{1} + Ox_{2} => Ox_{1} + Red_{2} , is spontaneous as written if Eº > 0, i.e. if Eº_{2} > Eº_{1} .

Since in the cell reaction Red_{1} reduces Ox_{2} to Red_{2} , we can conclude that Red_{1} is more reducing than Red_{2} (otherwise the reaction would go in the other direction) , and hence, quite generally, the lower the standard potential of a couple, the more powerfully reducing the reduced species in the couple is.

This knowledge may be used to construct the electrochemical series, a list of the metallic elements in order of their reducing ability.

Quite simply, this series is the metals from metal/metal ion couples listed in order of their standard potentials in aqueous solution; any metal may reduce the ions of a metal with a larger value of Eº.

The list can be extended to include non-metallic elements. (Note that this information is purely thermodynamic, and as always reactions which are thermodynamically possible may not be observed for kinetic reasons.)

#### Partial Electrochemical series in order from least reducing (largest Eº) to most reducing (smallest Eº)

Electrochemical Series |
Redox Couple (Ox/Red) |
Eº/V |

Fluorine |
F |
+ 2.87 |

Gold |
Au |
+ 1.40 |

Chlorine |
Cl |
+ 1.36 |

Platinum |
Pt |
+ 1.20 |

Bromine |
Br |
+ 1.09 |

Silver |
Ag |
+ 0.80 |

Mercury |
Hg |
+ 0.79 |

Iodine |
I |
+ 0.54 |

Copper |
Cu |
+ 0.34 |

Hydrogen |
H |
0 |

Lead |
Pb |
– 0.13 |

Tin |
Sn |
– 0.14 |

Nickel |
Ni |
– 0.23 |

Iron |
Fe |
– 0.44 |

Zinc |
Zn |
– 0.76 |

Chromium |
Cr |
– 0.91 |

Aluminium |
Al |
– 1.66 |

Magnesium |
Mg |
– 2.36 |

Sodium |
Na |
– 2.71 |

Calcium |
Ca |
– 2.87 |

Barium |
Ba |
– 2.91 |

Potassium |
K |
– 2.93 |

Lithium |
Li |
– 3.05 |

(Note that the table can also be viewed as reading from most to least reducing.) As can be seen from the table, typical electrode potentials lie within the range ± 3 V.