The Corrosion Cell

For corrosion to occur, the presence of all the components of an electrochemical cell is required. These components are an anode, a 

Electric Conference, University of California, Davis, February 1964. 

That mystery often surrounds the process of corrosion is probably because of the hard-to-recognize forms that the electrochemical cell takes. Persons accustomed to the laboratory will visualize an electrochemical cell as a beaker containing electrolyte in which to pieces of metal are immersed and joined externally with a wire. It is difficult to make the translation between this situation and that of a water pipe running through alternate marshy and sandy patches of soil, yet both are electrochemical cells—and both will be subject to the reactions that go to make up corrosion. 

Corrosion cells can be far more subtle. For example, electrical code states that a neutral wire must be bonded to a water pipe. If it is connected to a galvanized pipe in one residence and a copper pipe in a neighboring dwelling, we have the makings of a corrosion cell (Fig. 7-10). The galvanized pipe is the anode the copper pipe is the cathode. The common neutral wire is the external circuit. The moist soil surrounding the pipes is the electrolyte. 

Corrosion cells can be produced by the interaction of small, local, adjacent anodes and cathodes on the same piece of metal. These so-called local-action cells form because the surface of a piece of metal is not uniform. Small variations in composition, local environment, orientation of the grain structure, and differences in the amount of stress and surface imperfections all may contribute to the creation of tiny areas of 

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Precipita.tingInhibitors. As discussed earlier, the localized pH at the cathode of the corrosion cell is elevated due to the generation of hydroxide ions. Precipitating inhibitors form complexes that are insoluble at this high pH (1—2 pH units above bulk water), but whose deposition can be controlled at the bulk water pH (typically 7—9 pH). A good example is zinc, which can precipitate as hydroxide, carbonate, or phosphate. Calcium carbonate and calcium orthophosphate are also precipitating inhibitors. Orthophosphate thus exhibits a dual mechanism, acting as both an anodic passivator and a cathodic precipitator. 

An electrochemical reaction is said to be polarized or retarded when it is limited by various physical and chemical factors. In other words, the reduction in potential difference in volts due to net current flow between the two electrodes of the corrosion cell is termed polarization. Thus, the corrosion cell is in a state of nonequilibrium due to this polarization. Figure 4-415 is a schematic illustration of a Daniel cell. The potential difference (emf) between zinc and copper electrodes is about one volt. Upon allowing current to flow through the external resistance, the potential difference falls below one volt. As the current is increased, the voltage continues to drop and upon completely short circuiting (R = 0, therefore maximum flow of current) the potential difference falls toward about zero. This phenomenon can be plotted as a polarization diagram shown in Figure 4-416. 

For the corrosion process to proceed, the corrosion cell must contain an anode, a cathode, an electrolyte and an electronic conductor. When a properly prepared and conditioned mud is used, it causes preferential oil wetting on the metal. As the metal is completely enveloped and wet by an oil environment that is electrically nonconductive, corrosion does not occur. This is because the electric circuit of the corrosion cell is interrupted by the absence of an electrolyte. Excess calcium hydroxide [Ca(OH)j] is added as it reacts with hydrogen sulfide and carbon dioxide if they are present. The protective layer of oil film on the metal is not readily removed by the oil-wet solids as the fluid circulates through the hole. 

Soluble corrosion products may increase corrosion rates in two ways. Firstly, they may increase the conductivity of the electrolyte solution and thereby decrease internal resistance of the corrosion cells. Secondly, they may act hygroscopically to form solutions at humidities at and above that in equilibrium with the saturated solution (Table 2.7). The fogging of nickel in SO2-containing atmospheres, due to the formation of hygroscopic nickel sulphate, exemplifies this type of behaviour. However, whether the corrosion products are soluble or insoluble, protective or non-protective, the... 

It has been shown that pure distilled water is least corrosive when fully aerated and that some inhibitors function better in the presence of oxygen In these cases oxygen acts as a passivator of the anodic areas of the corrosion cells. These facts do not, however, modify the foregoing statements on the significance of oxygen in waters as used in practice. 

The standard electrode trotential, Ep, 2+ Pb = —Q.126V . shows that lead is thermodynamically unstable in acid solutions but stable in neutral. solutions. The exchange current for the hydrogen evolution reaction on lead is very small (-10 - 10"" Acm ), but control of corrosion is usually due to mechanical passivation of the local anodes of the corrosion cells as the majority of lead salts are insoluble and frequently form protective films or coatings. 

It is at the anode that oxidation takes place, with the anodic metal suffering a loss of negatively charged electrons. The resulting positively charged metal ions dissolve in the water electrolyte and metal wastage occurs. In the corrosion cell, the metal or metal area having the lowest electrical potential becomes the anode. 

Thus, a useful method of corrosion control is to slow down the driving force of the corrosion cell by reducing the difference in potential at the cathode (cathodic polarization). In practice, this can be achieved by maintaining reducing conditions at the boiler surface (typically by dosing an oxygen scavenger), which increases the rate of cathodic polarization and drives down the corrosion rate. 

V, and Tpe= -0.44 V. If the zinc and the iron electrodes connected by an electrical conductor are dipped into standard electrolytes in this cell, then iron would serve as the cathode (Fe2+ + 2 e —> Fe) and zinc as the anode (Zn —> Zn2+ + 2 e ). The result would be a tendency for zinc to dissolve in the electrolyte, and this process is known as the galvanic corrosion of a less noble metal (zinc) in comparison with the more noble metal iron in this system. The reversible emf of the corrosion cell would be... 

Figure 8.18 illustrates the polarization of the corrosion couple between galena and other minerals. Galena has low static potential in the three sulphide minerals. When it contacts with other media, they form the corrosion cell. Galena appears anode and cathode polarization occurs. When galena contacts with Fe medium, which has lower potential than galena, the result is contrary (see Fig. 8.9). 

The reaction produces hydroxyl ions which react directly with the Fe ions to produce an oxide precipitate. The combined anodic and cathodic reactions form the corrosion cell, the electrochemical potential of which lies between the single potential of the two half reactions. This mixed potential is termed the corrosion potential, corr> and for corrosion to proceed beyond the equilibrium state, the corrosion potential must be more positive than the equilibrium single potential of iron. For iron in water at pH 7 and with [Fe j = 10" M, for example, the potential of the anodic reaction is. 

Anodic inhibitors should not, however, be used on galvanized iron, in which protection of the iron depends upon the zinc coating being anodic an anodic inhibitor would coat the zinc, reverse the polarity of the corrosion cell, and cause the iron to corrode merrily away. 

Apart from this important feature of a corroding system, there is another characteristic that arises from the short-circuit condition of the corrosion cell and of the equivalent cell. (It will be recalled that the electron sources and sinks in the corroding metal are internally short-circuited the two electrodes in the equivalent cell are externally short-circuited.) The total potential difference V across the equivalent cell is zero. But this cell potential is composed of the absolute potential differences across the interfaces at the two electrodes and the potential drop IR in the electrolyte ... 

Franke and Winnick [105], using a K2S2O7 based electrolyte dispersed within the interstices of an inert K2Mg2(S04)3 matrix, were able to achieve removal efficiencies of SO2 greater than 99% at current efficiencies near 100%. A porous electrode constructed of a perovskite-type compound, Lao 8Sro.2Co03, was found to be conductive and stable in the corrosive cell environment. 

The steel hulls of ships are constantly in contact with saltwater, so the prevention of corrosion is vital. Although the hull may be painted, another method is used to minimize corrosion. Blocks of metals, such as magnesium, aluminum, or titanium, that oxidize more easily than iron are placed in contact with the steel hull. These blocks rather than the iron in the hull become the anode of the corrosion cell. As a result, these blocks, called sacrificial anodes, are corroded while the iron in the hull is spared. Of course, the sacrificial anodes must be replaced before they corrode away completely, leaving the ship s hull unprotected. A similar technique is used to protect iron pipes that are run underground. Magnesium bars are attached to the pipe by wires, and these bars corrode instead of the pipe, as shown in Figure 21-15b. 

Steel socket welds is a good example of rapid local corrosion. The potential difference between the anodic and the cathodic states drives the corrosion cells (this is an example of galvanic corrosion). Corrosion due to adjacent active-passive sites can be particnlarly rapid if the corrosion cell has an unfavorable anode/cathode area ratio. 

The description of corrosion kinetics in electrochemical terms is based on the use of potential-current diagrams and a consideration of polarization effects. The equilibrium or reversible potentials Involved in the construction of equilibrium diagrams assume that there is no net transfer of charge (the anodic and cathodic currents are approximately zero). When the current flow is not zero, the anodic and cathodic potentials of the corrosion cell differ from their equilibrium values the anodic potential becomes, more positive, and the cathodic potential becomes more negative. The voltage difference, or polarization, can be due to cell resistance (resistance polarization) to the depletion of a reactant or the build-up of a product at an electrode surface (concentration polarization) or to a slow step in an electrode reaction (activation polarization). 

The rate of corrosion (loss of metal) is proportional to the amount of current that flows in the corrosion cell. From Ohm s law we know that, for direct current,... 

Insulation. It is often necessary to connect two types of metal, for example, galvanized water Service pipes to a hot water heater from which copper hot water lines emerge. The corrosion cell that would develop in this case (Fig. 7-9) can be eliminated by inserting a dielectric coupling between the galvanized pipe and the copper pipe which effectively breaks the external circuit and eliminates the corrosion cell. However, this is not common practice because brass fixtures are used in nearly all systems instead of galvanized material. 

Electronic amplifiers have a small bias current across the input. In electrochemical corrosion measurements it is a requirement that this bias current is very small, such as less than 1 jjlA to avoid polarising the electrodes of the corrosion cell. 

The time between placing the specimen in the corrosion cell and the commencement of the measurement itself must also be counted as part of the surface preparation procedure. The value of the electrochemical parameter sought may vary considerably if insufficient attention is paid to this point [6-8]. It is preferable to consider this period as part of the measurement itself, reducing the chances of random variations. 

If one mole of tin is consumed in the corrosion cell Sn/Sn //Cu /Cu, calculate the change of the Gibbs free-energy, AG. 

Determine whether tin is stable in 10 M Sn acid solution of pH = 2. Estimate (a) the Gibbs free-energy change and (b) the cell potential for the corrosion cell. The activity coefficients are assumed to be 1. The hydrogen pressure is 1 atm. 

Estimate the theoretical tendency for zinc to corrode (em ) when immersed in 10 to 10 M ZnCla solution at pH of 3. The corrosion cell is described as ... 

Calculate the driving emffor the corrosion cell and write the spontaneous reaction for the following cell ... 

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