Figure 4 A Galvanic Cell Showing Absorbed Hydrogen Atoms on a Cathode

Figure 3 A Galvanic Cell

Corrosion Theory Summary

Chemistry Volume 1 of 2
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CORROSION THEORY DOE-HDBK-1015/1-93 Corrosion CH-02 Rev. 0 Page 8 Figure 4 A Galvanic Cell Showing Absorbed Hydrogen Atoms on a Cathode Now consider a galvanic cell with zinc and platinum electrodes, such as that shown in Figure 4. The half-reactions in the cell are as follows. (2-4) Again, as the cell operates, the cell potential drops. The decrease is partially due to the increase in Zn⁺² concentration and the decrease in H O 3 + concentration, but another type of polarization also occurs in this cell. This second type is associated with the reduction half-reaction. The hydrogen atoms formed by the reaction of Equation (2-4) absorb on the surface of the metal and remain there until removed by one of two processes: combination of two hydrogen atoms to form molecular hydrogen, which is then released as a gas or reaction with dissolved oxygen to form water. In the absence of oxygen (deaerated solutions), the first process applies. (2-6) Combining Equation (2-6) with Equation (2-4), the net reduction half-reaction is obtained. (2-6) (2-7) Until the absorbed hydrogen atoms are removed from the metal surface, they effectively block the sites at which the reaction of Equation (2-4) can occur. At low temperatures the reaction of Equation (2-6) is slow relative to the reaction of Equation (2-4) because, although the reaction is energetically favored, the combination of two hydrogen atoms requires a large activation energy. Equation (2-6) shows the rate-controlling step of the net reduction half-reaction. Because the oxidation half-reaction can occur no faster than the reduction half-reaction, the rate of the overall oxidation-reduction reaction is controlled by the reaction of Equation (2-6).

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