Galvanic corrosion is an electrochemical process that occurs when two different metals are in electrical contact and exposed to a conductive liquid environment. This phenomenon is a form of accelerated degradation where one metal is preferentially consumed while the other remains protected. The process mimics a simple battery, resulting in the spontaneous dissolution of the more chemically active material. The corrosive attack is concentrated at the junction or on the surface of the susceptible metal, often leading to rapid material failure.
The Required Elements
For galvanic corrosion to be initiated, three distinct elements must be present simultaneously, forming what is known as a galvanic cell. The first requirement is the presence of two dissimilar metals or alloys that possess different electrical potentials. Metals are ranked by their tendency to lose electrons in a sequence called the Galvanic Series; the further apart two metals are on this series, the greater the potential difference and the stronger the driving force for corrosion.
The second element is an electrolyte, an electrically conductive liquid. Common examples include seawater, rainwater, moisture in soil, or high humidity in the air. Without this conductive path, the electrochemical circuit cannot be completed, and the corrosive reaction cannot proceed.
The final requirement is a direct electrical connection between the two dissimilar metals. This connection allows electrons to flow freely from one metal to the other, often occurring when they are physically touching, such as a bolt fastening two plates together.
How the Reaction Proceeds
The mechanism of galvanic corrosion is fundamentally an oxidation-reduction reaction driven by the difference in electrochemical potential between the two metals. The metal that is more chemically active, or less noble, becomes the anode, and it is here that the corrosive attack begins. The anodic reaction involves the metal atoms losing electrons, a process called oxidation, which turns the solid metal into positively charged ions that dissolve into the electrolyte.
The electrons released by the dissolving anodic metal then travel through the direct electrical connection to the more noble metal (the cathode). At the cathode, the electrons are consumed by a reduction reaction, often involving the reduction of oxygen or the production of hydrogen gas. This electron flow sustains the corrosion process, protecting the cathodic metal from degradation.
The electrolyte facilitates the movement of ions, which completes the electrical circuit. Positive ions from the anode and negative ions from the electrolyte migrate to maintain electrical neutrality, preventing the buildup of charge that would stop the flow of electrons. This continuous cycle leads to the accelerated consumption and failure of the anodic metal.
Variables Affecting Corrosion Rate
Several variables dictate the speed and severity of the degradation. One significant factor is the relative surface area of the two coupled metals. An unfavorable ratio, where a small anodic area is connected to a large cathodic area, dramatically accelerates the corrosion of the anode. The large cathode collects electrons over a wide area, concentrating the corrosive current onto the anode’s tiny surface and causing rapid material loss.
The conductivity of the electrolyte influences the rate of corrosion. Highly conductive solutions, such as saltwater or industrial runoff, allow for a faster flow of ions and current, speeding up the electrochemical process. Conversely, a weakly conductive electrolyte, like pure distilled water, slows the reaction significantly.
Temperature is another influencing factor, as higher temperatures increase the speed of chemical reactions, including the oxidation and reduction processes. Elevated temperatures also increase the diffusion rate of oxygen through the electrolyte, accelerating the cathodic reaction. The environment’s pH level is also important; a lower, more acidic pH favors the corrosive reaction and increases the overall degradation rate.
Recognizing Galvanic Corrosion
Galvanic corrosion occurs where different metals are joined, particularly in environments exposed to moisture. A frequent example is in household plumbing, where copper pipes connect to steel or galvanized fittings. The steel, being the more active metal (the anode), corrodes rapidly at the junction with the more noble copper.
In construction and marine environments, aluminum structures fastened with stainless steel bolts are a classic scenario for this type of attack. The aluminum corrodes preferentially around the fasteners, which can weaken the structural integrity of the assembly. Similarly, the use of steel fasteners in aluminum window frames can lead to localized failure where rainwater acts as the electrolyte.
The corrosion is localized and concentrated on the anodic metal, usually occurring close to the point of contact with the cathodic metal. Signs of this accelerated attack include pitting (small, deep holes in the surface) or a white, powdery residue on aluminum or zinc alloys. Discoloration or heavy rust staining near the junction point indicates a galvanic reaction is underway.