Galvanic Corrosion

Galvanic corrosion is a form of corrosion that can occur when two dissimilar metals contact each other in the presence of an electrolyte. The alloy with the more active corrosion potential, the more active alloy, forms the anode and corrodes. The alloy with the more noble corrosion potential, the more noble alloy, forms the cathode and is protected. Electrons flow between the two metal. The circuit is completed by the environment in which one of the metals, the cathode, promotes a reduction of a species, often oxygen in the electrolyte and other metal, the anode, promotes the oxidation of a species often the metal itself to a dissolved ionic form. The metal promoting the oxidation reaction tends to corrode at a faster rate than it would suffer if it was not in contact with the other metal. The metal promoting the cathodic reaction tends to corrode at a slower rate (if at all) than it would suffer if it was not in contact with the other metal. For example, iron corrodes in water. When iron contacts zinc in water at room temperature, the corrosion process is altered so that iron corrosion is decreased. Zinc corrosion is increased over that which would occur in the absence of the metallic couple.

While galvanic corrosion is well understood from a qualitative standpoint, the complexities of and geometric influences on the process make quantitative predictions difficult if not impossible in practical systems. A number of factors dictated by the environmental and by the two metals impact galvanic corrosion. While they cannot be included in a general quantitative model, their impact on galvanic corrosion can be summarized. Some of the more important factors that have to be considered are as follows.

• Corrosion potentials of the metals forming the couples
• Kinetics of the anodic and cathodic reactions occurring on the two surfaces
• Electrolyte properties (conductivity, pH, ionic species, temperature, etc.)
• Geometric factors (relative areas, positions, shapes, orientation, etc.)
• Surface condition of the two alloys

The effects of these variables are summarized below.

1 Corrosion Potential and Corrosion Kinetics 2 Electrolyte Properties 2.1 Dissolved Oxygen and Fluid Velocity 2.2 Fluid Conductivity 3 Exposed Areas 4 Condition of Surface

Corrosion Potential and Corrosion Kinetics

The standard electrode potentials as determined from free energy data and tabulated in such sources as Chemical Rubber Company Handbook are not relevant to predicting the effect of coupling two dissimilar metals. The corrosion potentials of the two metals as determined independently in the environment by measuring the potential of the alloy relative to a reference electrode are more helpful in that they provide information as to the possibility of galvanic corrosion and the maximum magnitude of the potential driving force. But such information does not predict that galvanic corrosion will occur. The current between the two metals as supported by the anodic and cathodic reactions with the environment provides that latter information.

Shown in the figure below are partial polarization scans that might be generated from the corrosion potential for each metal when alone in the environment. They are somewhat simplified as straight lines in these figures for the purposes of the illustration.

M_R1 and M_R2 are two more noble metals that have the same corrosion potential in the environment. These metals serve individually as the cathodes in the galvanic couple. M_O is the more active metal. It serves as the anode in the galvanic couple. Only the the cathodic portion of the polarization scan for the more noble metals and the anodic portion of the curve for the more active metal are drawn because those portions provide information needed to describe the ultimate current causing galvanic corrosion. Where the cathodic curve from the more noble metal and the anodic curve from the more active metal cross provides an initial, possibly order of magnitude, guess as to the potential of the couple (V₁ or V₂) and the current density (i₁ or i₂) that could be generated between them.

As seen in this figure, simply knowing the two initial corrosion potentials does not provide enough information from which to conclude that the current can cause problems. Two alloys could have very similar corrosion potentials but show significant differences in their polarization scans (current versus voltage relationship). The current density difference between i₁ and i₂ can be significant.

The potential of the couple (V₁ or V₂) and the resulting current density that might flow between the alloys ((i₁ or i₂)are determined by all of the electrochemical processes occurring in the alloy-environment boundary region in the vicinity of the two alloys and how those processes influence the polarization of each member of the couple.

From a theoretical standpoint, one might quantify the relationships if one could write an equation of the form

(where i_corr is the corrosion current density, V is the applied (assumed to be actual) voltage, V_corr is the corrosion potential, and b_a and b_c are the anodic and cathodic Tafel slopes) for each member of the couple. This equation assumes one anodic and one cathodic reaction on each surface. Additional terms would have to be written for each reaction. The more noble alloy would be polarized in the cathodic (or active) direction increasing the influence of the cathodic reaction (subscript "c"). The more active alloy would be polarized in the anodic (or noble) direction increasing the influence of the anodic portion (subscript "a"). The voltage of the couple is determined when the current densities of each equation are equal. In theory, if all parameters could be determined, the galvanic current density could be calculated. Unfortunately, in practice such a calculation can almost never be made. Measurements using a zero resistance ammeter placed between the two metals tends to be the standard method of measurement.

Electrolyte Properties

Galvanic corrosion is a property of the system containing the two alloys, not a property only of the alloys themselves. Measurements are best made with the metals in the environment of interest. Though general relationships cannot be written for all environmental variables, some generalities are possible.

Dissolved Oxygen and Fluid Velocity

Oxygen reduction to water tends to be the primary cathodic reaction when it participates in the corrosion reaction. The reason is that oxygen is present in most practical situations. Oxygen is sparingly soluble in most liquid environments. This characteristic means that the rate at which oxygen can be reduced is very often equal to the rate at which it reaches the surface from the bulk. This characteristic implies that if dissolved oxygen controls the cathodic reaction, the rate of galvanic corrosion would be influenced by the rate at which oxygen reaches the surface. Since that rate increases with fluid flow rate past the surface, such galvanic corrosion would be influenced by the velocity of the fluid past the surface. The following figure shows how polarization of each alloy individually might appear when dissolved oxygen controls the cathodic reaction of the more noble alloy.

The polarization of the active alloy (M forming M⁺ⁿ) is under activation control meaning that its behavior (current) depends only on potential and is independent of fluid motion. As the more noble alloy is polarized, a current is reached such that the rate of the cathodic reaction becomes limited by the rate of oxygen transfer to the surface. At that point, the current becomes independent of potential and the polarization curve becomes a vertical line. The value of this current limit only increases as fluid velocity is increased (A to B to C). Some examples of this effect are the rate of galvanic corrosion of aluminum (anode) when coupled to steel (cathode) and copper (cathode) coupled to zinc (anode).

This effect of velocity is not limited to only the case of oxygen being the cathodic reaction. The effect of velocity can also occur whenever the anodic or cathodic reaction becomes limited by the rate of a reactant or product migrating to or from the surface.

Fluid Conductivity

Fluid conductivity provides a measure of the ability of the fluid environment to support an electric (ionic) current. The higher the conductivity, the greater the ability to support such a current. The current (electron flow) between two dissimilar alloys is shown in this figure:

In the figure, the arrows show the direction of positive current in the solution environment. Alloy A is oxidized liberating metal-containing ions into solution. These ions (or oxygen) are reduced on alloy B. Electrons flow in the opposite direction between the two alloys where they touch. The electron flow between the metals equals the rate of oxidation which also equals the rate of reduction. The conductivity of the solution environment is a major variable controlling ionic flow (current)in the solution . In a low conductivity environment, the current (ionic flow) is depressed relative to that in a high conductivity environment. As depicted in the figure, as the conductivity decreases, the galvanic current and subsequent galvanic corrosion becomes more concentrated closer to the boundary between the metals. The amount of galvanic corrosion tends to be depressed relative to that in a high conductivity fluid. Situations arise in which the conductivity is so low that galvanic corrosion fails to occur, galvanic current cannot be supported even though the corrosion potentials of the metals forming the couple are different.

If the environment solution is a thin layer contacting the two metals as would occur if the environment is the atmosphere and the liquid is formed by condensation, the total current parallel to the surface is somewhat restricted because of the narrowness of the layer. Only a small region on each of the alloys close to the point of contact tends to undergo corrosion. Even though the concentration of metal-containing ionic species could increase as galvanic corrosion proceeds and the conductivity could increase, the effect on current could be small because of the thinness of the electrolyte layer.

Exposed Areas

The rate of galvanic corrosion on a per unit area basis depends on the current density (current divided by corroding area) on each surface. If the corroding areas are identical, then the current density on each surface is the same. The effect of relative exposed area arises when the two exposed areas differ in size.

If the electrolyte contains dissolved oxygen and oxygen reduction on one of the metals controls the rate of reaction, e.g. aluminum alloys (anode) coupled to copper or stainless steel (cathode), then the total galvanic current becomes (approximately) proportional to the area of the cathode. This type of control is very common. The rate of galvanic corrosion is also proportional to the area of the cathode. Since the rate of oxygen reduction on the more noble alloy controls the galvanic current, the total current on the more active metal (anode) is independent of its area. But, the rate of galvanic corrosion of the more active alloy (anode) as determined by the current density (current divided by area) becomes inversely proportional to its area. From a corrosion testing standpoint, these observations mean that if galvanic corrosion is suspected, the testing might include galvanically coupled specimens that have 1:10, 1:1, and 10:1 area ratios.

There is a classic example of the effect of area on galvanic corrosion. If iron and copper are coupled in a reasonably conductive aqueous environment like seawater, copper is the more noble component (cathode) and iron the more active component (anode). If iron plates (large area) are held together by copper rivets (small area) the steel plates show a small amount of corrosion but the rivets hold. If copper plates are held together by steel rivets, the steel rivets will ultimately corrode completely leaving copper plates no longer held together. The point is that if the more active metal has a much larger surface area than the more noble metal, the deleterious effect of galvanic corrosion on the more active member can sometimes be lessened.

Condition of Surface

Alloys in contact with electrolytes for the most part have surfaces that are not bare metal but usually have a complex metal-oxide, metal-hydroxide structure in immediate contact with the electrolyte. This structure can complicate the ability to predict galvanic corrosion. This structure can affect the galvanic couple. This structure can change with time making galvanic corrosion sometimes a function of time.

If the surface structure varies across the surface, the potential difference across the surface can vary. If the effect is great enough, a local galvanic cell can develop between adjacent areas on the same alloy. The result is a local galvanic cell, increasing the corrosion on the more active area. The ultimate result could be Pitting Corrosion and the formation of pits.

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