2. Mixed potential theory and its application

3. Multiple electrodes In a corrosion cell (electrolyte, anode, cathode and a metallic path), multiple reactions occur. For instance, when zinc corrodes in a dilute acid, the following reactions occur:

4. The total reaction is Accordingly, the metal constitutes a multi electrode, because at least two different reactions occur

5. on its surface simultaneously, one oxidation and one reduction. The mixed potential theory partly mentioned earlier, is used with advantage to predict the rate of corrosion of metals and alloys in given environment. It was postulated by Wagner and Traud in 1938. It has two basic assumptions: (a) Electrochemical reactions are composed of two or more partial anodic and cathodic reactions. (b) There cannot be any accumulation of charges. This theory was not applied until 1950, when Stern applied it in analysis of corrosion.

6. Consider now an electrode, such as zinc, immersed in HC1. The following could be the possible reactions:

7. Bubbles of hydrogen are observed from the surface of zinc electrode, and formation of bubbles of hydrogen is a cathodic reaction. Hydrogen is reduced and not oxidized. Similarly, zinc is oxidized and not reduced. Hence, only the two reactions (a) and (b) proceed. Under the condition of rest (no outside current), the potential of the electrode cannot be computed by the Nernst equation as it is not reversible. Also, the above electrode would not corrode in the absence of an external current. The potential assumed by the electrode under the above condition is the mixed potential and its value lies between the value of equilibrium potential of hydrogen and zinc. The value of the potential would depend on the metal

8. and the environment. It is to be observed that the corrosion potential (E corr ) is not the equilibrium potential of either of the reactions, but some intermediate potential determined by the two partial anodic and cathodic reactions. Both the reactions (a) and (b) are polarized to a common potential, the mixed potential (E mix ) (Fig. 2.8).

10. -The mixed potential diagram or E vs I diagram allows accurate prediction of corrosion rate to be made in specific environments. -It also allows the prediction of current based on current density higher than the corrosion current (i Corr ) and utilization of advanced technique for corrosion measurement, such as galvanostatic and potentiostatic techniques.

11. Application of mixed potential theory 1- Effect of an oxidizer

12. It is of interest to observe the effect of oxidizing metal ions on the corrosion rate of a metal in acid solutions, such as the effect of Fe 3+ ions on the corrosion of zinc in hydrochloric acid. Log i vs E diagrams can be constructed to predict the effect of environment on the corrosion rate of metals. Examine diagram (Fig. 2.10). When multiple reactions take place, E corr and i corr can be obtained by summing the currents at different potentials at which reactions occur (anodic and cathodic) to obtain total anodic and cathodic polarization curves. For instance, two partial anodic

13. and two partial cathodic reactions are shown in Fig. 2.10. The total rate of oxidation and reduction is shown by broken line. Charge neutrality must be maintained. Consider a metal, M, in HCl to which Fe 3+ ions have been added. The corrosion rate of metal M in the absence of ferric salts is given by i corr (M). The corrosion increases from i corr (M) to i corr, on adding the oxidizer as shown in the diagram. The rate of hydrogen evolution is decreased from i H2 = i corr (M) to i H2 as shown in the figure. This is due to depolarization. The corrosion potential of the metal E Corr (M) is shifted to a more noble value on addition of the oxidizing ions as shown by the E corr arrowed on the figure obtained by the intersection of total anodic and cathodic curves. Thus, there are three consequences of adding an oxidizer:

15. (a) The corrosion rate of the metal is increased. (b) The corrosion potential is shifted to a more noble direction. Fig.(2.10)Effect of oxidizer on a metal in an air solution (c) The rate of hydrogen evolution is decreased. It is to be noted that the effect of Fe 3+ ion is pronounced on the metal M, because of its high exchange current density. If the exchange current tendency is small, there would be little effect on the corrosion rate of metal M. The exchange current density of the oxidizing ions must be higher than the exchange current density of hydrogen on the metal surface, to have a significant effect.

16. 2-Coupling of an active metal to a noble metal Fig. (2.11)

17. the coupling of zinc (E° = — 0.76 V) with platinum (E° = 1.2 V) in a hydrochloric acid solution. Platinum has a very noble potential and it does not oxidize in HCl. In fact its driving force is negative. Zinc dissolves in acid with the liberation of hydrogen. Platinum does not react. On connecting zinc with platinum electrically, it is observed that:

18. (a) The rate of hydrogen evolution on zinc (H on Zn) decreases. (b) The rate of oxidation of zinc,, in the acid solution increases. (c) The rate of hydrogen evolution on platinum surface (H on Pt) increases vigorously. Consider first, the oxidation reaction of zinc Fig.(2.11)galvanic coupling of active to noble metal (Zn 2+ + 2e — Zn). The oxidation reaction polarizes zinc in the noble direction and the hydrogen reduction reaction (2H+ + 2e —> H 2 ) in the active direction. The anodic and cathodic curves converge at a point d. The point of convergence yields i corr for zinc. The potential corresponding to i corr for zinc is E corr « Now consider the oxidation of platinum.

19. Fig.(2.11)galvanic coupling of active to noble metal

20. There is no oxidation of Pt in HC1, hence, no oxidation curve for platinum is shown in Fig. 2.11. However, the reduction of hydrogen on platinum is clearly shown by the cathodic polarization curve for platinum. It is observed that there is only one oxidation process (Zn—> Zn 2+ + 2e) and

21. two reduction processes on platinum and on zinc. The situation is analogous to Case 2.10.1 above. The total cathodic process is the sum of H 2 on Zn and H 2 on Pt, resulting in the dotted line. As observed in Fig. 2.11, the i corr of zinc (uncoupled) increases from 10 -4 A/cm 2 to a very high value at a when coupled with platinum; consequently, on connecting zinc with platinum, the rate of corrosion of zinc is vigorously increased.

22. As shown by the points of intersection, a and c, the rate of hydrogen evolution on zinc is drastically reduced (from d to c), whereas the rate of evolution of hydrogen on the platinum is vigorously increased. The current density at b is significantly higher than at c as shown in Fig. 2.11. Thus the rate of hydrogen evolution is decreased is shown by the reduction in the current from point d to point c. What does it all amount to? It is shown clearly that zinc is a very poor catalyst for reduction of hydrogen as shown by a very small exchange current density io(H 2 on Zn) = 10 -10 A/cm2, whereas platinum

23. is an excellent catalyst for reduction of hydrogen as shown by a very large exchange current density for reduction of hydrogen io(H 2 on Pt) = 10 -3 A/cm 2. To summarize the effect of coupling of zinc to platinum, the following are the main points of interest: (a) The rate of hydrogen evolution is decreased on zinc and increased on platinum. (b) The rate of oxidation of zinc is increased significantly on coupling and zinc dissolves vigorously. Nothing happens to platinum(c) Platinum is an excellent catalyst for reduction of hydrogen and zinc is a poor catalyst.

24. 3-Effect of galvanic coupling (Fig.2.12) It is a common practice to use the potential of a metal in an emf series to predict its corrosion tendency. In case of galvanic couples, the difference of potential between the couples is taken as a measure of corrosion tendency of the couple. In general, the larger the difference between the thermodynamic potential of the metals forming a galvanic couple, the more severe would be the magnitude of galvanic corrosion. This is on the Fig.(2.12)Effect of galvanic coupling of zinc with gold and platinum basis of thermodynamics. Consider coupling of zinc to gold and zinc to platinum. According to the thermodynamic approach,

25. Fig.(2.12)Effect of galvanic coupling of zinc with gold and platinum

26. the difference between the potential of zinc (E Zn = —0.76V) and gold (E Au = 1.50V) is higher than the difference between the potential of the zinc and platinum (E Zn = -0.76V,E Pt = + 1.2V). According to the thermodynamic approach, the Zn-Au couple should corrode faster than Zn-Pt couple. This is certainly not true and contrary to the experimental observation. The purely thermodynamic approach could not, therefore, be a good basis for prediction. Let us observe now the accuracy of prediction made on the basisof kinetics. Examine Fig. 2.12 showing the galvanic coupling of zinc with gold and platinum. The oxidation of zinc polarizes the electrode in the noble (positive) direction and the cathodic reduction of hydrogen in the active (negative) direction. The intersection of the two curves gives i corr for Zn. The exchange current density for hydrogen on

27. zinc, Au and Pt is shown in the diagram. The exchange current density for zinc is shown in the first curve in Fig. 212. The oxidation curves for gold and platinum are not shown in the diagram as the above metals do not oxidize in HCl. The intersection of hydrogen reduction curve for gold with zinc oxidation curve yields the i corr and E corr of zinc coupled to gold (Au) and the intersection of the oxidation curve of zinc with platinum yields the icon and E corr of zinc when coupled to platinum. It is clearly observed that the highest value of i corr is shown by Zn—Pt coupled.

28. The i corr value of Zn-Au couple is clearly very much lower than Zn-Pt couple. The hydrogen reduction reaction rate is the highest on a platinum surface i o (H on Pt) = 10 -3 A/cm 2 followed by gold i o (H on Au) = 10 -6 A/cm 2. The reduction rate of hydrogen is very low on zinc io (H + /H 2 on Zn) = 10 -10 A/cm 2. As observed in Fig. 2.12, the i corr of Zn-Pt couple is higher than that of Zn- Au couple, hence, the corrosion rate of Zn-Pt couple is higher than that of Zn-Au couple. This result in contrary to that obtained on the basis of only thermodynamic potentials but is true, due to the effect of kinetics. Gold is a poor catalyst of hydrogen evolution io (H on Au) = 10 -6 A/cm 2, compared to i o (H on Pt) = 10 -3 A/cm 2.

29. Zinc when coupled to platinum, therefore, corrodes at a much higher rate than when it is coupled to gold. The above predictions are based on kinetics of reactions and are, therefore, more accurate and complete than the predictions based on thermodynamic potentials. The latter can be often misleading and if an accurate prediction is made, such as in case of active metal coupling, it may be more of a coincidence than the rule.

30. Fig.(2.13)Effect of cathode anode area ratio on corrosion of zinc –platinum galvanic couples

31. Effect of oxgen Fig.(2.14) Figures 2.14a and b show the reactions which occur when iron is placed in a deaerated acidand later oxygen is introduced into the solution. E rev (O 2 ) is the reversible potential of oxygen And E rev (H 2 ), the reversible potential of hydrogen, as shown in the above figures. The reversible potential of any metal is represented by E rev (M/M z+ ). In the deaerated condition, the following reactions take place:

32. The oxidation reaction polarizes iron in the anodic direction and the reduction reaction in the cathodic direction. The point of intersection gives i corr. The magnitude of over-voltage from the reversible potential is given by  c for cathodic polarization and by  A for anodic polarization. Hydrogen is liberated and iron is corroded as Fe 2+. Consider now an aerated condition.

33. Oxygen is passed into the acid until it becomes saturated with oxygen. The reduction of oxygen is shown by the reaction, l/2O 2 +2H++2e -> H 2 O. In Fig. 2.14b, two partial reduction reactions are shown, i.e.

34. fig.(2.14) Effect of aeration and deaeration

35. The summation of two partial reactions is shown in the figure. There are two cathodic processes and one anodic process. The exchange current density of oxygen on iron is very low (-10 -10 A/cm 2 ). Even on a metal like platinum the exchange current density of oxygen is very low. The reason is that charge transfer is slowed down considerably due to film formation on the metal surface by oxygen. The metal electrode under oxidizing conditions becomes a less efficient catalyst for reduction. Oxygen reduction on platinum is approximately ten times slower than hydrogen reduction.

36. The intersection of cathodic reduction curve of oxygen with the anodic polarization curve is given by E corr. The intersection of cathodic reduction curve with the oxidation curve for iron is given by the appropriate icorr (Fig- 2.14b). It is observed that upon aeration of the acid solution, rate of reaction (iCorr) increases and the corrosion potential (E corr ) shifts to more noble value. To summarize the results: (a) Fe is oxidized to Fe 2+. (b) Hydrogen is reduced. (c) Oxygen is reduced. (d) The rate of corrosion increases on aeration. (e) The rate of corrosion decreases on deaeration.