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Standard Half-Cell Potentials

The half-cell potential is said to be the standard half-cell potential if the concentration of ions and pressure of gaseous species appearing in the half-cell reaction have values of 1 mol dm-3 and 101.325 kPa (= 1 atm), respectively.

Determination of Standard Potentials

At the very outset, it may be stated that the absolute value for the reduction potential of any single half-reaction cannot be determined experimentally. However, constructing a suitable cell and then determining its emf experimentally can determine the difference between the two reduction potentials. By definition, the emf of a cell is given by
Ecell = ERHC - ELHC

Where ERHC and ELHC are the reduction potentials of the right-hand and left-hand half-cells, respectively.
If ELHC is arbitrarily assigned some value, then the value of ERHC can be determined using the expression
ERHC = Ecell + ELHC

In the study of an electrochemical cell, the hydrogen half-cell has been adopted as the reference half-cell and its standard potential has been assigned the value zero at 298.15 K. By standard potential of hydrogen half-cell, we mean that the hydrogen ion and hydrogen gas involved in half-cell H+|H2|Pt  are present in their standard states of unit concentration (1 mol dm-3) and unit pressure (taken as 1 atm, i.e. 1.01325 bar pressure), respectively (Fig. 10.2).

Thus, if we have a cell in which the left-hand half-cell is the standard hydrogen half-cell and the right-hand half-cell constitutes the half-cell system whose potential, relative to that of standard hydrogen half-cell, is required then
ERHC = Ecell + 0 = Ecell

that is, the reduction potential of the given half-cell is numerically equal to the emf of the cell. Proceeding in this way, the standard potentials of other half-cells relative to the standard hydrogen half-cell can be determined. The sign of the electrode potential is the experimentally measured sign of the cell emf if the standard hydrogen electrode is on the left and the electrode in question is on the right.

To illustrate the procedure, we cite below two typical examples of silver-silver ion and zinc-zinc ion half-cells. If the silver-silver ion half-cell is coupled with the standard hydrogen half-cell (to be kept on the left-hand side), we get a cell
Pt| H2(1.01325 bar)| H+(1M)|| Ag+(1M)|Ag

The emf of the cell is given by

and its value as determined experimentally is found to be 0.7991 V. Hence,

Taking the example of zinc-zinc ion half-cell, we have
Pt| H2(1.01325 bar)| H+(1 mol dm-3)||Zn2+ (1 mol dm-3)|Zn
Its emf as determined experimentally is found to be - 0.763 V.


Note that the cell emf is negative. It implies that the cell as written above will not produce a spontaneous reaction. In fact, while determining emf of the cell, the electrode of zinc-zinc ion half-cell will serve as the negative terminal and Pt electrode of hydrogen-hydrogen ion half-cell as the positive terminal, in order to get a positive potential of 0.763 V. Since the cell has been written in the reverse direction (i.e. Pt electrode as the negative terminal and Zn as the positive terminal), it follows that the emf of the written cell will be - 0.763 V.

From the two examples cited above, it may be concluded that if an electrode of a half-cell with a positive reduction potential is coupled with the standard hydrogen-hydrogen ion half-cell, it will constitute the positive terminal of the cell in order to have a spontaneous cell reaction. Similarly, an electrode of a half-cell with a negative reduction potential will constitute the negative terminal of the cell to get the spontaneous cell reaction. In other words, the nature of the electrode of a half-cell (whether positive or negative), in a cell in which the other half-cell is the standard hydrogen-hydrogen ion half-cell, is determined by the sign of the reduction potential of the given half-cell.

In brief, the standard electrode potential of an electrode has a positive value if this electrode is more positive than the standard hydrogen electrode and a negative sign if it is more negative than the standard hydrogen electrode.

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