Zinc corrosion reaction

These two reactions then add up to the overall zinc corrosion reaction (eq. 11). 

Table 2.16 Zinc Corrosion Reactions in Different Atmospheres (Schematic)...

Since metals have very high conductivities, metal corrosion is usually electrochemical in nature. The tenn electrochemical is meant to imply the presence of an electrode process, i.e. a reaction in which free electrons participate. For metals, electrochemical corrosion can occur by loss of metal atoms tluough anodic dissolution, one of the fiindamental corrosion reactions. As an example, consider a piece of zinc, hereafter referred to as an electrode, inunersed in water. Zinc tends to dissolve in water, setting up a concentration of Zn ions very near the electrode... 

The standard electrode potentials , or the standard chemical potentials /X , may be used to calculate the free energy decrease —AG and the equilibrium constant /T of a corrosion reaction (see Appendix 20.2). Any corrosion reaction in aqueous solution must involve oxidation of the metal and reduction of a species in solution (an electron acceptor) with consequent electron transfer between the two reactants. Thus the corrosion of zinc ( In +zzn = —0-76 V) in a reducing acid of pH = 4 (a = 10 ) may be represented by the reaction ... 

The rate (or kinetics) and form of a corrosion reaction will be affected by a variety of factors associated with the metal and the metal surface (which can range from a planar outer surface to the surface within pits or fine cracks), and the environment. Thus heterogeneities in a metal (see Section 1.3) may have a marked effect on the kinetics of a reaction without affecting the thermodynamics of the system there is no reason to believe that a perfect single crystal of pure zinc completely free from lattic defects (a hypothetical concept) would not corrode when immersed in hydrochloric acid, but it would probably corrode at a significantly slower rate than polycrystalline pure zinc, although there is no thermodynamic difference between these two forms of zinc. Furthermore, although heavy metal impurities in zinc will affect the rate of reaction they cannot alter the final position of equilibrium. 

During the operation of the cell (or during the direct interaction of zinc metal and cupric ions in a beaker) the zinc is oxidised to Zn and corrodes, and the Daniell cell has been widely used to illustrate the electrochemical mechanism of corrosion. This analogy between the Daniell cell and a corrosion cell is perhaps unfortunate, since it tends to create the impression that corrosion occurs only when two dissimilar metals are placed in contact and that the electrodes are always physically separable. Furthermore, although reduction of Cu (aq.) does occur in certain corrosion reactions it is of less importance than reduction of HjO ions or dissolved oxygen. 

For these reasons a somewhat different approach will be adopted here, and an attempt will be made to show how a corrosion reaction may be represented by a well-defined reversible electrochemical cell, although again there are a number of difficulties. Consider the corrosion of metallic zinc in a reducing acid... 

Zembura has made specific use of the rotating disc for investigation of the effect of flow on corrosion reactions. This work has shown that it is possible to determine the type of control (activation or concentration polarisation) of zinc dissolving in 0.1 N Na2S04 (de-aerated), which followed closely the predicted increase in hydrogen ion reduction as the flow rate increased, and proved that in this example... 

In atmospheric exposure to industrial environments its corrosion rate is only about one-third that of zinc and the corrosion reaction is stifled by the tenacious oxide which is produced nevertheless it can frequently function as an anodic coating both for steel and for the less corrosion-resistant aluminium alloys. 

These considerations show the essentially thermodynamic nature of and it follows that only those metals that form reversible -i-ze = A/systems, and that are immersed in solutions containing their cations, take up potentials that conform to the thermodynamic Nernst equation. It is evident, therefore, that the e.m.f. series of metals has little relevance in relation to the actual potential of a metal in a practical environment, and although metals such as silver, mercury, copper, tin, cadmium, zinc, etc. when immersed in solutions of their cations do form reversible systems, they are unlikely to be in contact with environments containing unit activities of their cations. Furthermore, although silver when immersed in a solution of Ag ions will take up the reversible potential of the Ag /Ag equilibrium, similar considerations do not apply to the NaVNa equilibrium since in this case the sodium will react with the water with the evolution of hydrogen gas, i.e. two exchange processes will occur, resulting in an extreme case of a corrosion reaction. 

The corrosion reactions may be slowed down by using zinc alloys (with lead and cadmium, also improving the mechanical properties of zinc to simplify the production process) instead of the pure metal, or by amalgamating the inner surface of the can by adding a small amount of a mercury compound to the electrolyte. 

The life of a dry-cell battery is relatively short. Oxidation causes the zinc cup to deteriorate, and eventually the contents leak out. Even while the battery is not operating, the zinc corrodes as it reacts with ammonium ions. This zinc corrosion can be inhibited by storing the battery in a refrigerator. As discussed in Chapter 9, chemical reactions slow down with decreasing temperature. Chilling a battery therefore slows down the rate at which the zinc corrodes, which increases the life of the battery. 

Although not in common use today, thin metallic rods of high purity lead (Pb) and, to a lesser extent, zinc (Zn) have been used as references in concrete. They are simple and easy to make, but their potentials are produced by corrosion reactions (mixed potentials) as opposed to known reversible... 

Fig. 12.6. The hydrogen-evolution reaction is the electronation reaction that occurs in zinc corrosion in acid solution.

In order to suppress or inhibit zinc corrosion and hydrogen evolution at zinc anode, a number of additives are selected and used for the anode. Because of its high over potential to the hydrogen evolution reaction, mercury is an effective gassing suppresser and used to be widely employed in alkaline Zn/Mn02... 

Estimating the deposition velocities of gaseous species is considerably more complex than estimating those for substances in particles, in part due to the uncertainties in the sticking and reaction probabilities. Such estimates have not been made but the potential effects of some of the typical gases can be surmised from available data on surface accumulation rates, e.g. sulfate accumulation on indoor zinc and aluminum surfaces is predominantly a result of particulate sulfate deposition rather than a corrosion reaction involving sulfur dioxide (0. 

Galvanized iron is steel sheet that has been coated with zinc tin cans are made of steel sheet coated with tin. Discuss the functions of these coatings and the electrochemistry of the corrosion reactions that occur if an electrolyte contacts the scratched surface of a galvanized iron sheet or a tin can. 

Anode Corrosion Reaction. Zinc might at first appear to be an unusual choice for battery anode material, because the metal is thermodynamically unstable in contact with water... 

In a battery, the anode and cathode reactions occur in different compartments, kept apart by a separator that allows only ionic, not electronic conduction. The only way for the cell reactions to occur is to mn the electrons through an external circuit so that electrons travel from the anode to the cathode. But in the corrosion reaction the anode and cathode reactions, equations 8 and 12 respectively, occur at different locations within the anode. Because the anode is a single, electrically conductive mass, the electrons produced in the anode reaction travel easily to the site of the cathode reaction and the zinc acts like a battery where the positive and negative terminals are shorted together. 

The rate at which the corrosion of the zinc proceeds depends on the rates of the two half reactions (eqs. 8 and 12). Equation 8, a necessary part of the desired battery reaction, fortunately represents a reaction that proceeds rather rapidly, whereas the reaction represented by equation 12 is slow. Ie, the generation of hydrogen on pure zinc is a sluggish reaction and thus limits the overall corrosion reaction rate. 

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