INFLUENCE OF POLARIZATION ON CORROSION RATE

5.7 INFLUENCE OF POLARIZATION ON CORROSION RATE

Both resistance of the electrolyte and polarization of the electrodes limit the magnitude of current produced by a galvanic cell. For local - action cells on the surface of a metal, electrodes are in close proximity to each other; consequently, resistance of the electrolyte is usually a secondary factor compared to the more important factor of polarization. When polarization occurs mostly at the anodes, the corrosion reaction is said to be anodically controlled (see Fig. 5.7 ). Under anodic control, the corrosion potential is close to the thermodynamic potential of the cathode. A practical example is impure lead immersed in sulfuric acid, where a lead sulfate fi lm covers the anodic areas and exposes cathodic impurities, such as copper. Other examples are magnesium exposed to natural waters and iron immersed in a chromate solution.

When polarization occurs mostly at the cathode, the corrosion rate is said to

be cathodically controlled . The corrosion potential is then near the thermody- namic anode potential. Examples are zinc corroding in sulfuric acid and iron exposed to natural waters.

Resistance control occurs when the electrolyte resistance is so high that the resultant current is not suffi cient to appreciably polarize anodes or cathodes. An example occurs with a porous insulating coating covering a metal surface. The corrosion current is then controlled by the IR drop through the electrolyte in pores of the coating.

It is common for polarization to occur in some degree at both anodes and cathodes. This situation is described as mixed control . The extent of polarization depends not only on the nature of the metal and electrolyte, but also on the actual exposed area of the electrode. If the anodic area of a corroding metal is very small, caused, for example, by porous surface fi lms, there may be considerable anodic polarization accompanying corrosion, even

INFLUENCE OF POL ARIZATION ON CORROSION R ATE

Figure 5.7. Types of corrosion control.

Figure 5.8. Polarization diagram for corroding metal when anode area equals one - half of cathode area.

though measurements show that unit area of the bare anode polarizes only slightly at a given current density. Consequently, the anode – cathode area ratio is also an important factor in determining the corrosion rate. If current density is plotted instead of total corrosion current, as, for example, when the anode area is half the cathode area, corresponding polarization curves are shown in Fig. 5.8 .

70 KINETICS: POL ARIZATION AND CORROSION R ATES

Figure 5.9. Polarization diagram for zinc amalgam in deaerated HCl.

An important experiment that illustrates the electrochemical mechanism of corrosion was performed by Wagner and Traud [1] . They measured the corrosion rate of a dilute zinc amalgam in an acid calcium chloride mixture and the cathodic polarization of mercury in the same electrolyte. The current density equivalent to the corrosion rate was found to correspond to the current density necessary to polarize mercury to the corrosion potentiual of the zinc amalgam (Fig. 5.9 ). In other words, mercury atoms of the amalgam composing the majority of the surface apparently act as cathodes, or hydrogen electrodes, * and zinc atoms act as anodes of corrosion cells. The amalgam polarizes anodically very little, and the corrosion reaction is controlled almost entirely by the rate of hydrogen evolution at cathodic areas. Consequently, it is the high hydrogen overpotential of mercury that limits the corrosion rate of amalgams in nonoxidizing acids. A piece of plati- num coupled to the amalgam considerably increases the rate of corrosion because hydrogen is more readily evolved from a low - overpotential cathode at the oper- ating emf of the zinc – hydrogen electrode cell.

The very low corrosion rate and the absence of appreciable anodic polariza- tion explain why amalgams in corresponding metal salt solutions exhibit corro- sion potentials closely approaching the reversible potential of the alloyed

component. For example, the corrosion potential of cadmium amalgam in CdSO 4 solution is closer to the thermodynamic value for Cd → Cd 2+

+ 2e − than is observed for pure cadmium in the same solution. The steady - state corrosion rate of pure cadmium is appreciably higher than that of cadmium amalgam, leading

* Mercury in aqueous solutions acts fi rst as a mercury electrode, but, when cathodically polarized, all mercury ions in solution are deposited before H + is discharged. Any conducting surface on which H + ions are discharged acts as a polarized hydrogen electrode and can be so considered in evaluating a corrosion cell.

C ALCUL ATION OF CORROSION R ATES FROM POL ARIZATION DATA

to greater deviation of the measured corrosion potential from the corresponding thermodynamic value. In general, the steady - state potential of any metal more active than hydrogen (e.g., iron, nickel, zinc, cadmium) in an aqueous solution of its ions deviates from the true thermodynamic value by an amount that depends on the prevailing corrosion rate accompanied by H + discharge. The measured value tends to be more noble than the true value. This situation also holds for the steady - state potentials of more noble metals (e.g., copper, mercury) which undergo corrosion in the presence of dissolved oxygen.