Electrochemical+and+Thermodynamic+Background

=Electrochemical and Thermodynamic Background =  The phenomenon of corrosion can be explained through the basic principles of electrochemistry. It occurs when a metal exhibits an electrochemical reaction with its environment [1]. Corrosion occurs in what can be modeled as a galvanic cell, consisting of an anode and a cathode [1]. At the anode, a metal, such as iron, is oxidized, losing electrons to form a cation [2].

Fe (s) <> Fe 2+ (aq) + 2e - (E 0 =0.44V)

Iron is one constituent of steel, a major material used in breweries, but other components of steel may also undergo corrosion; their associated oxidation reactions are tabulated in the following table:

//Table 1. Oxidation Half-Reactions of Metals Relevant for the Brewery Industr y// The reduction half-reaction, which occurs at the cathode, can be depicted as follows:
 * **Metal** || **Oxidation Half-reaction and Standard Reaction Potential (V)** ||
 * Chromium || Cr (s) <> Cr 3+ (aq) + 3e - (E 0 =0.44V) ||
 * Iron || Fe (s) <> Fe 2+ (aq) + 2e - (E 0 =0.44V) ||
 * Copper || Cu (s) <> Cu 2+ (aq) + 2e - (E 0 =-0.34V) ||
 * Aluminium || Al (s) <> Al 3+ (aq) + 3e - (E 0 = 1.66V) ||

O 2 + 2H 2 O + 4e - <> 4OH - (E 0 =0.40 V)

This explains why oxygen has to be present for corrosion to occur. However, there are other driving forces that cause corrosion to occur; these are explained in other sections, which can be accessed here. Note that the overall reaction potential, for the oxidation of iron, is 0.44V, which denotes a spontaneous reaction. Overall, the resulting reaction can be described by the following schematic:



The metals corrode because they only temporarily exist in metallic form. Therefore, when exposed to a corroding environment, they tend to revert back to their natural, ionized, form, which can be achieved through the process of oxidation [2]. During this process, energy needed to make the metallic form is emitted to the environment [2].

From a thermodynamic perspective, the corrosion reaction only occurs when the Gibbs’ free energy is negative (the reaction is spontaneous). In the corrosion of copper, an important metal used in breweries, copper will corrode and tarnish in a short period of time because the reaction is so spontaneous [1]. Notice how all the reactions in Table 1 are spontaneous since the overall reaction potential is greater than zero.

The oxidation of metals, including copper, follows a similar sequence, which can be summarized by the following steps [1, p.189]:

1. Oxygen adsorbs onto the metal surface. 2. The oxygen nuclei are formed. 3. A continuous oxide film grows.

For copper, in particular, oxidation occurs at low temperatures. Copper dissolves in oxygen and ruptures the grain boundaries when heated in the presence of hydrogen [1]


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**References**
[1] H.H. Uhlig and R.W. Revie, //Corrosion and corrosion control,// 3rd ed. USA: John Wiley and Sons, 1985, pp.1-201.

[2] P. Roberge, //Corrosion Engineering: Principles and practice.// USA: McHGraw-Hill Companies, Inc, 2008.

[3] "National Research Council Canada," //nrc-cnrc.gc.ab.// [Online]. Available: []. [Accessed: November 30, 2009].