|16.2 Electrochemical Considerations||●||S-211|
For this reaction to occur spontaneously, �V0must be positive; if it is negative, the spontaneous cell direction is just the reverse of Equation 16.17. When standard half-cells are coupled together, the metal that lies lower in Table 16.1 will experience oxidation (i.e., corrosion), whereas the higher one will be reduced.
INFLUENCE OF CONCENTRATION AND TEMPERATURE ON CELL POTENTIAL
where R is the gas constant, n is the number of electrons participating in either of the half-cell reactions, and F is the Faraday constant, 96,500 C/mol—the magnitude of charge per mole (6.023 � 1023) of electrons. At 25�C (about room temperature),
|�V � (V0 2� V 0 1) � 0.0592 n||
(a) If the cell is a standard one, write the spontaneous overall reaction and calculate the voltage that is generated.
(b) Compute the cell potential at 25�C if the Cd2�and Ni2�concentrations are 0.5 and 10�3M, respectively. Is the spontaneous reaction direction still the same as for the standard cell?
|� Ni � Cd2�||(16.21)|
�V � (V0
�V � V0 Ni� V 0 Cd� � 0.250 V � (�0.403 V) � �0.153 V
The half-cell potentials listed in Table 16.1 are thermodynamic parameters that relate to systems at equilibrium. For example, for the discussions pertaining to Figures 16.2 and 16.3, it was tacitly assumed that there was no current flow through the external circuit. Real corroding systems are not at equilibrium; there will be a flow of electrons from anode to cathode (corresponding to the short-circuiting of the electrochemical cells in Figures 16.2 and 16.3), which means that the half-cell potential parameters (Table 16.1) cannot be applied.