What is galvanic corrosion?

What is galvanic corrosion?

Galvanic corrosion is the accelerated corrosion that occurs when two dissimilar metals are in electrical contact in the presence of an electrolyte — usually water containing dissolved salts. The combination forms a galvanic cell, in effect a small battery, in which one metal is forced to corrode faster than it would on its own while the other is protected.

The crucial word is dissimilar. Two pieces of the same metal produce no driving voltage. But as soon as two metals of different nobility are electrically connected, and moisture closes the circuit, a current begins to flow. With the current, metal is carried away from the less-noble surface — it breaks down, often right at the contact point between the metals.

The phenomenon is sometimes called bimetallic corrosion or contact corrosion. It is one of the most common reasons outdoor structures age prematurely, particularly in infrastructure where different metals are joined: railings, poles, fasteners, cable connections and luminaires. The insidious part is that each metal may be perfectly appropriate on its own — it is the combination that creates the problem. A stainless bolt and an aluminium part each work well alone, but together in salt-laden humid air they become a corrosion cell.

How does galvanic corrosion occur?

To understand the mechanism, it is easiest to follow what happens in the cell. Four things occur simultaneously, and they depend on one another:

• Oxidation occurs at the anode. The less-noble metal releases electrons and goes into solution as ions. For aluminium and iron respectively: Al → Al³⁺ + 3e⁻ and Fe → Fe²⁺ + 2e⁻. This is where material is actually lost.
• The electrons travel through the metallic contact from the anode to the cathode. Without a conductive connection between the metals there is no path for the electrons, and no cell.
• Reduction occurs at the cathode. In neutral, aerated water, dissolved oxygen is reduced: O₂ + 2H₂O + 4e⁻ → 4OH⁻. The cathode consumes the electrons the anode releases, but the cathode metal itself is not broken down.
• The ions close the circuit through the electrolyte. Positive and negative ions move in the liquid and complete the circuit. This is where moisture and salt come in.

The current that flows is called the corrosion current, and it is a direct measure of how fast the anode is broken down. What drives the current is the potential difference between the metals — their different tendencies to release electrons. Just like the voltage in a battery: the greater the difference, the more strongly the cell drives its own corrosion.

The three conditions in practice

The whole mechanism can be summed up in three conditions that must be met simultaneously. This is also the key to preventing it: remove any single condition and the cell cannot form.

The galvanic series

Which metal becomes the anode and which the cathode is determined by their position in the galvanic series — a ranking of metals by how noble they are in a given electrolyte, usually seawater. The nobler metal (higher up) becomes the cathode and is protected. The less-noble one (further down) becomes the anode and is sacrificed.

The order above is simplified. The actual position depends on alloy, temperature, oxygen content and passive film. One important detail: stainless steel is noble only as long as its passive oxide film is intact. In an oxygen-starved crevice, where the passive film cannot re-form, the same stainless steel can become considerably less noble (active) and behave very differently. The galvanic series (measured in a real electrolyte) should therefore not be confused with the pure electrochemical (EMF) series of standard potentials, which does not account for passivation.

Which metals corrode first?

It is always the less-noble metal in the pair that is sacrificed. In practice this means aluminium and zinc corrode when coupled to stainless steel or copper, since the latter are nobler. Magnesium is so base that it is sacrificed against almost anything — a property deliberately exploited in sacrificial anodes (more on that below). The point is that a metal is not "corrosion-prone" in itself; it is its position relative to the other metal that decides.

Why salt water and coastal environments accelerate corrosion

Pure water conducts electricity poorly. As soon as it contains dissolved salts — above all chlorides from sea salt and road salt — its conductivity rises dramatically. Better conductivity means a higher corrosion current, which means faster breakdown of the anode. That is the first and most important effect.

The second effect is at least as important: chlorides break down passive films. Metals such as aluminium and stainless steel are normally protected by a thin, self-healing oxide film. Chloride ions penetrate this film locally and initiate pitting — localised attacks that bore down into the material. In a galvanic cell the two effects reinforce one another.

That is why coastal, harbour and marine environments are among the most demanding there are, and road-salted environments are not far behind. The combination of high humidity, salt and temperature swings means the electrolyte is almost always present and always conductive. A metal combination that would survive a dry inland climate can be broken down in a fraction of the time by the coast. It is also why corrosion classes such as C5-M (marine environment) exist — they describe exactly this kind of exposure. See IP66, IK10 and C5-M for what those classes actually test.

The area ratio — small anode, large cathode

One of the most misunderstood factors is geometry. It is not just which metals meet, but how large their surfaces are relative to one another. The corrosion current the cathode drives must be absorbed by the anode. If the anode is small, the entire current is concentrated on a small area — the current density becomes high and the material is eaten away very quickly right there.

The unfavourable situation is therefore a small anode against a large cathode. A classic example is a small aluminium part or aluminium rivet in a large stainless or copper surface: the aluminium is perforated quickly. The opposite geometry — a large anode against a small cathode — spreads the same current over a large area, and the attack becomes slow and superficial.

This gives a simple design rule: if dissimilar metals must be combined, let the fasteners (the small parts) be of the nobler metal. Stainless bolts in an aluminium structure are manageable — the small cathode limits the rate. Aluminium bolts in a stainless structure, on the other hand, are a classic design error, because the small anode is eaten away fast.

Galvanic corrosion in street lighting and infrastructure

Outdoor infrastructure is full of metal combinations, and therefore of galvanic cells. Lighting poles and luminaires are outdoors around the clock, year-round, often in exactly the coastal and roadside environments where salt is present. Some typical risk points:

• Fasteners: stainless bolts and screws in aluminium housings or aluminium poles — the most common combination of all, and a constant source of local corrosion around the fixings.
• Pole to base and access doors: dissimilar metals in pole doors, earthing devices and fixings where rain and road salt collect.
• Cable connections: copper cable connected directly to aluminium (Al/Cu joints) is notorious for galvanic corrosion combined with oxide build-up. It degrades the contact, increases resistance and can lead to heat and outages.
• Harbours, bridges, industry and marine infrastructure: environments with constant salt loading, where every metal transition is a potential cell.

What these have in common is that corrosion rarely occurs in the middle of an even surface, but precisely at the transitions — where two materials, and often a moisture pocket, meet. It is also where a luminaire or pole usually starts to show its age first. More on the underlying mechanisms in Why Luminaire Housings Corrode.

Aluminium vs stainless steel

Stainless steel in the passive state is clearly nobler than aluminium. The aluminium therefore becomes the anode and is sacrificed. The combination is extremely common — stainless fasteners in aluminium structures — and works tolerably as long as the stainless (the cathode) is small relative to the aluminium. In salt environments, however, the attack accelerates around the bolts, and insulation or sealing is needed for long-term durability.

Aluminium vs galvanised steel

Galvanised steel has a surface of zinc, which is less noble than aluminium. Initially the combination is therefore relatively benign: the zinc is sacrificial and corrodes first, protecting both the steel and the aluminium. The risk arises later — once the zinc layer is consumed, only the steel is exposed, and steel is nobler than aluminium. The situation then reverses and the aluminium begins to be sacrificed. So the combination is better than many assume in the short term, but should still be monitored over the long term.

Aluminium vs copper

Copper is considerably nobler than aluminium, and the potential difference is large. That makes the combination one of the most aggressive in practical electrical engineering. Aluminium connected directly to copper — particularly in outdoor electrical connections — corrodes quickly, and the attack is often combined with an insulating oxide that degrades the electrical contact. Direct Al/Cu connections should therefore be avoided or made with special transitions and joint compounds.

How galvanic corrosion is prevented

Because galvanic corrosion requires three conditions simultaneously, breaking just one of them is enough to stop the process. This gives several possible strategies, which are often combined:

Sacrificial anodes deserve a comment of their own, because the principle is so clear. On boats, in water heaters and on marine steel structures, blocks of zinc, aluminium or magnesium are fitted. These deliberately become the most anodic surface in the system and are sacrificed, while the hull, propeller or tank is protected. The anode is replaced at intervals — it is a consumable part. It is galvanic corrosion turned to advantage.

Why electrically insulating materials eliminate galvanic corrosion

The most fundamental solution is to remove the precondition itself. Galvanic corrosion requires metallic electrodes and a conductive path for the electrons. A non-conductive material — engineering plastic or glass-fibre reinforced composite — can neither act as an electrode nor conduct current between other metals. The cell therefore cannot form.

This is an important distinction in principle: electrical insulation, correct material choice and coatings delay and limit galvanic corrosion, but a surface treatment can be damaged and an insulating washer can be forgotten. A non-metallic material eliminates the mechanism — there is no galvanic cell to protect against. (Metallic fasteners in or on a composite can of course still corrode against each other if they are of dissimilar metals; it is the combination of metals that must be avoided, not metal as such.)

Common misconceptions

• "Stainless never rusts." Stainless can suffer pitting and crevice corrosion in chloride environments, and as a noble cathode it can also accelerate corrosion of adjacent less-noble metals.
• "Aluminium doesn't rust, so it doesn't corrode." Aluminium forms no rust (which is iron oxide), but it does corrode — it forms aluminium oxide and suffers pitting, particularly galvanically and in salt.
• "Galvanic corrosion is the same as ordinary rust." No. Rust is the oxidation of iron. Galvanic corrosion is a mechanism that can affect metals that do not rust.
• "Only the quality of the metal matters." The combination and geometry (the area ratio) are at least as important as the quality of the individual metal.
• "Paint solves everything." Only while the layer is unbroken. A small defect on the cathode side can create a very unfavourable small-anode situation and make matters worse than no paint at all.

Summary

Galvanic corrosion is not a material defect but a predictable electrochemical process. It occurs when two dissimilar metals are electrically connected in a common electrolyte, and the less-noble metal is sacrificed. Salt and moisture make it faster, geometry determines where and how hard it strikes, and it can always be prevented by breaking one of the three conditions.

Frequently asked questions

Sources

1. Fontana, M. G. (1986). Corrosion Engineering (3rd ed.). McGraw-Hill. (Galvanic corrosion, the galvanic series, and the significance of area ratio.)
2. ASM International. (2003). ASM Handbook, Volume 13A: Corrosion — Fundamentals, Testing, and Protection.
3. ASTM. (2014). ASTM G82: Standard Guide for Development and Use of a Galvanic Series for Predicting Galvanic Corrosion Performance.
4. ISO. (2020). ISO 8044: Corrosion of metals and alloys — Vocabulary. (Terminology for anode, cathode, electrolyte and galvanic corrosion.)
5. ISO. (2018). ISO 12944-2: Paints and varnishes — Corrosivity categories (C1–C5, CX).
6. van Bommel, W. (2015). Road Lighting: Fundamentals, Technology and Application. Springer. (Environmental requirements and material considerations for road luminaires.)