hydrogen ions cannot oxidise Cu (or alternatively we can say that hydrogen gas can reduce copper ion) under the standard conditions described above. Thus, Cu does not dissolve in HCl. In nitric acid it is oxidised by nitrate ion and not by hydrogen ion. The negative value of the standard electrode potential in the second case indicates that hydrogen ions can oxidise zinc (or zinc can reduce hydrogen ions).
In view of this convention, the half reaction for the Daniell cell in Fig. . can be written as: Left electrode: Zn(s) ® Zn + (aq, M) + e Right electrode: Cu + (aq, M) + e – ® Cu(s) The overall reaction of the cell is the sum of above two reactions and we obtain the equation: Zn(s) + Cu + (aq) ® Zn + (aq) + Cu(s) emf of the cell = E o cell = E o R – E o L = .34V – (– . )V = .
V Sometimes metals like platinum or gold are used as inert electrodes. They do not participate in the reaction but provide their surface for oxidation or reduction reactions and for the conduction of electrons. For example, Pt is used in the following half-cells: Hydrogen electrode: Pt(s)|H (g)| H + (aq) With half-cell reaction: H + (aq)+ e – ® ½ H (g) Bromine electrode: Pt(s)|Br (aq)| Br – (aq) With half-cell reaction: ½ Br (aq) + e – ® Br – (aq) The standard electrode potentials are very important and we can extract a lot of useful information from them. The values of standard electrode potentials for some selected half-cell reduction reactions are given in Table .
. If the standard electrode potential of an electrode is greater than zero then its reduced form is more stable compared to hydrogen gas. Similarly, if the standard electrode potential is negative then hydrogen gas is more stable than the reduced form of the species. It can be seen that the standard electrode potential for fluorine is the highest in the Table indicating that