| Contents for this page | Related topics |  |
|---|---|---|
|
The cell potential The standard hydrogen electrode Additional questions |
Electrolysis The Daniell (Cu-Zn) cell How to balance redox reactions |
 Data Glossary |
| Learning Outcomes | ||
| After studying this section, you will (a) know the meaning of the term "standard electrode potential", (b) understand the relationship between the electromotive force of a cell and the standard electrode potentials of its half cells, (c) be able to use a table of standard electrode potentials in order to calculate the e.m.f. of a cell, and (d) predict whether a given metal will displace another metal or hydrogen from solution. | ||
We saw in a previous section that the redox reaction
can be exploited to create an electrochemical cell which is written conventionally as
The cell potential (emf) of this cell under standard conditions is 1.10 V (volt). Further, the electrons flow in the external circuit from the anode (Zn), where oxidation takes place, to the cathode (Cu), where reduction occurs.
Now consider the two half-reactions which contribute to the cell, written down as reductions (gain of electrons):
Since the cell causes electrons to flow in the external circuit from Zn to Cu, it means that reaction (2) takes place rather than reaction (1). Reaction (2) takes place because the REVERSE of reaction (1) supplies the 2 electrons required.
The emf of the cell (1.10 V) is the POTENTIAL DIFFERENCE between the electrodes of the cell, which is the difference between the abilities of the half-cells to attract electrons.
We could therefore write down the emf of the above cell (under standard conditions) as:
emf of cell = potential of Cu - potential of Zn
or conventionally
(cathode) and (anode) are the STANDARD ELECTRODE POTENTIALS of the cathode and the anode respectively. Since these represent the ability of the electrode to accept electrons, they are in fact REDUCTION POTENTIALS.
(cell) is the emf of the cell, which can be determined by actually constructing the cell and measuring its potential (in volt).
In principle, any redox reaction could be utilised to make an electrochemical cell. For example, one could use the reaction
to construct the cell
and this cell could be shown experimentally to have an emf of 1.24 V at a temperature of 25 ºC.
If one only knew the values of the standard potentials of the half-reactions adding up to the whole redox process, one could easily calculate the emf of any cell using the formula
It is not possible to determine the absolute values of the electrode potentials. We can however determine the 
values of half-reactions relative to a standard electrode whose value is set arbitrarily at 0.00 V.
Consider the half-reaction
and let us use it to construct the electrochemical cell
This cell has an experimentally determined of 0.34 V at 25 ºC. If we assign a standard electrode potential of 0.00 V to the hydrogen half-reaction (which is the anode reaction in this case),
Now let us look at the cell
This cell has = -0.76 V at 25 ºC. (Normally, a negative emf means the anode and cathode have been wrongly identified. As the cell is written, electrons should flow from H2 to Zn, but
in fact it is the other way round!)
However, by convention, the standard hydrogen electrode is always written on the left-hand side of the cell. In other words, we always write such cells as if the standard hydrogen electrode were the anode. Thus ,
Note that a hydrogen half-cell is constructed by bubbling pure H2 gas over a piece of platinum foil connected to the external circuit. This enables contact to be made between the gas and the solution. For this reason, the standard hydrogen half-cell is normally written
A typical cell involving a hydrogen half-cell is shown on the left. The half-cell whose E is measured is the left hand electrode. The hydrogen gas is at a pressure of 101.3 kPa.
In this way one can derive the values for a wide variety of half-reactions. In the examination you will normally be supplied with a table of standard electrode potentials, so do not try to memorise any of them, except the one for the reduction of H2, which is easy to remember, as
it is simply 0.00 by definition!
Note that a metal which has a lower electrode potential will displace from solution a metal which has a higher electrode potential.
See a table of standard electrode potentials.