Abstract
This chapter addresses the thermodynamic aspects of corrosion, starting from the concept of free energy: indeed, corrosion can take place only if the free energy variation associated with the reaction is negative, i.e., the reaction is thermodynamically favoured. This translates in terms of variation of potential, outlined as driving voltage, or electromotive force, for the reaction. Standard potentials and equilibrium potentials of anodic and cathodic reactions are defined, together with conditions for corrosion and for immunity. Reference electrodes are presented, which allow to measure the potential as difference between a given electrode and a well defined reference electrode that has the property of maintaining its potential constant. Finally, electrochemical cells, as Daniell or concentration cells, are introduced.
There is nothing more practical than a good theory.
W. Nernst
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Notes
- 1.
Walther Herman Nernst (1864–1941) was a German chemist. He received the Nobel Prize for Chemistry in 1920.
- 2.
A platinum wire is subjected to anodic and cathodic cycles, which form a deposit of black platinum powder on the wire surface to increase the effective surface. The platinum so treated is called platinised platinum.
- 3.
The combination of the equation of the hydrogen evolution reaction (H2O + e− = ½H2 + OH-, reaction a) with the water dissociation reaction (H+ + OH− = H2O, reaction c) gives H+ + e− = ½H2 (reaction b). Therefore, ∆Gb = ∆Ga + ∆Gc. At equilibrium conditions ∆Gc = 0, therefore ∆Ga = ∆Gb, as well as the associated potentials with respect to the same reference electrode . From an energy viewpoint, the processes (a) and (b) are equivalent. The potentials of the two processes are mutually correlated through the ionic dissociation constant of water (at 25 °C Kw = aH+ · aOH = 10−14). For example, at pH 14, in conditions where both \(a_{{{\text{H}}_{2} {\text{O}}}} = 1\) and PH2 = 1 bar, depending on which process a) or b) reference is made, the following is obtained: \({E}_{{ {\text{b}}}}^{{}} = {E}_{{ {\text{b}}}}^{ 0} + \frac{{ 2. 3 {\text{R}T}}}{\text{F}}{\text{log }a}_{{{\text{H}}^{ + } }} = 0+ 0. 0 5 9\cdot (- 1 4 )= - 0. 8 2 8\,{\text{V SHE}}\)
- 4.
With similar considerations discussed in note 3, it can easily be proved that the two oxygen reduction reactions (O2 + 4H+ + 4e− = 2H2O and O2 + 2H2O + 4e− = 4OH−) are equivalent
- 5.
Equilibrium conditions are achieved only if a membrane selectively permeable to anions separates the two solutions. In the most common case in which the two solutions are in contact with each other, the potential of the cell is Eeq,I–II =  t a · 0.059/2 log ([Pb I]/[Pb II]), where ta is the transport number of anions.
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Pedeferri (Deceased), P. (2018). Thermodynamics of Aqueous Corrosion. In: Corrosion Science and Engineering. Engineering Materials. Springer, Cham. https://doi.org/10.1007/978-3-319-97625-9_3
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DOI: https://doi.org/10.1007/978-3-319-97625-9_3
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