Chapter 3
What is Corrosion?
Written and reviewed by: Michel Comtois, M. Sc. Chemistry and Anne-Marie Comtois, ChE — August 1, 2026
Corrosion, by definition, is the chemical or electrochemical reaction occurring between a material and its environment. The reaction can deteriorate the material, its structure, and its properties, which can lead to various complications depending on the affected area.
DEFINITION OF CORROSION
More Than Just Rust
To some extent, everyone knows what corrosion is. It is a common phenomenon that can be observed in our daily lives — just think of the rust that can be found on a car. Knowing that corrosion is a chemical or electrochemical reaction involving electron transfer is far less intuitive.
In automotive applications, ferrous alloys can react with oxygen and humidity in the air to form hydrated iron oxide, commonly known as rust. Although seemingly simple, corrosion is actually more complex, as it affects a large variety of surfaces and can involve many chemical reactions. Corrosion affects mostly metallic parts. However, ceramics and polymers may also degrade in the presence of environmental factors such as acid rain or sulfates sometimes present in groundwater due to pollution. The combination of these factors can destroy materials considered very stable and thus difficult to degrade over time. Despite being part of our day-to-day life, from a scientific standpoint, corrosion is a world in itself.1 2
REAL-WORLD EXAMPLES
Well-Known Examples of Corrosion
The Statue of Liberty
You are probably familiar with the Statue of Liberty, which was gifted to the United States by France. The monument was originally built in Paris in the 1880s and designed by Frédéric Auguste Bartholdi, with the internal iron framework engineered by Gustave Eiffel. Made of copper, the statue was brown at first, but turned green over time. What happened? The answer is simple: corrosion or, more precisely, oxidation.
In this case, the oxidation of copper protected the rest of the statue by coating it in a green layer, also known as patina, composed primarily of copper carbonate and related compounds. This patina acts as a barrier, significantly reducing further corrosion of the underlying metal.
In other cases, corrosion can lead to severe consequences. The Statue of Liberty, for example, required significant restoration in the late 1900s due to severe corrosion of its iron framework. When the copper skin came into contact with the iron, it produced an electrical current similar to a battery, resulting in what is known as galvanic corrosion.
Iron–Copper Galvanic Corrosion Reactions
Simply put, an improper combination of dissimilar metals can lead to an accelerated corrosion process.
Other Famous Corrosion Stories
Pipeline Failure — Santa Barbara, 2015
Plains’ Line 901 discharged 105,000 gallons of crude oil along the coastline due to a 45% reduction in pipeline thickness. The presence of severe corrosion led to crack propagation. Result: 2,934 barrels spilled, USD $24M in penalties, USD $32M in natural resource damages, USD $4M for Coast Guard cleanup.
Gas Explosion — Guadalajara, 1992
Corrosion of water and gas pipes caused leaks into the sewer system, which subsequently led to explosions. Result: 8 km of streets destroyed, 252 deaths, 1,500 injuries, 15,000 people left homeless.
Toxic Gas — Bhopal, 1984
Toxic gas was released into the atmosphere from the UCIL pesticide plant, where corrosion was cited as a contributing factor. Official death toll: 2,259. Estimated death toll: 8,000.
On a general level, statistical data indicate that 42% of engineering component failures are due to corrosion in one of its many forms 4. Investigations following several major incidents have concluded that corrosion was the main culprit.
THE SCIENCE
Electrochemical Cells — Electric Circuit
The corrosion process is generally classified into two categories: wet corrosion (corrosion requiring an electrolyte) and dry corrosion (corrosion associated with elevated temperatures or sulfidation). Wet corrosion is the most common type and is the primary focus of this guide.
Dry corrosion, or oxidation, typically occurs at elevated temperatures when oxygen in the air reacts with the metal, forming an oxide layer. Since most engineering metals oxidize slowly at ambient temperature, dry corrosion is less common than wet corrosion in typical service environments.
Wet corrosion is a slow, measurable process that follows Faraday’s law. For this type of corrosion to occur, the following elements are required: an electrolyte, an anode, and a cathode.
The combination of these components forms a system called an electrochemical cell. Corrosion is essentially an electric circuit with a flow of current between two electrodes, commonly called an anode and a cathode. These electrodes may come from different materials or simply from dissimilar areas on the same surface.
In this system, these two elements are in metallic contact and are immersed in an electrolyte. This electrolyte is usually a liquid through which ionic conduction occurs, or that separates into ions when electricity passes through it. It is a phase through which the movement of ions carries a charge. This conductive substance is required in order for the corrosion circuit to be completed. A common example of an electrolyte would be salt water. This is a key element in corrosion protection. Indeed, if this conductive substance can be prevented from contacting the metal structure we want to protect, for example, with a protective coating, then the electron exchange cannot proceed, and corrosion is significantly reduced.
Oxidoreduction Reactions
The corrosion of a material involves a pair of reactions, which are called oxidoreduction (oxidation-reduction) reactions, that happen simultaneously. The first half-reaction is called the oxidation reaction. It is defined by the loss of electrons at the anodic electrode. At the same time, a reduction reaction, also called a cathodic reaction, occurs.
The reduction reaction refers to a gain of electrons, which will attach to the other material in the electrochemical cell called the cathode. In other words, there is a current in the electrochemical cell that induces a flow of electrons away from the anode and through the metallic path toward the cathode. Ions move through the electrolyte solution to maintain charge balance. These oxidation-reduction reactions induce the corrosion of the anode. The cathode is not consumed during corrosion.
Types of Electrochemical Cells
Electrochemical cells are generally classified into two main types: galvanic cells and electrolytic cells5.
Also called a Voltaic cell or a Daniell cell, this type of electrochemical cell derives its energy from spontaneous redox reactions. Galvanic cells are generally used as sources of DC electrical power (e.g., a disposable battery). It may contain only one electrolyte separated by a semi-porous membrane or two separate half-cells connected by a salt bridge. The anode is the negative part in this type of cell.
This type of electrochemical cell allows the application of a reverse voltage to a galvanic cell. An external source is used to carry out a naturally non-spontaneous reaction. They are generally encountered during the charging phase of rechargeable batteries (e.g., charging your smartphone). In this case, the cathode will become the negative part of the cell.
The figure on the right illustrates a galvanic cell, specifically, a Daniell cell, and shows the direction of electron flow associated with the two half-reactions described above.
Galvanic cells are mostly discussed throughout this guide as they are the most common form of corrosion found in the environment.
Summary Comparison of Electrolytic and Galvanic Cells7
| Electrolytic (Nickel-Iron Cell) |
Galvanic (Zinc-Cerium Cell) |
|
|---|---|---|
| Oxidation | ||
| Occurs at | Anode | Anode |
| Sign of Terminal | Positive (+) | Negative (-) |
| Half-Reaction | Ni(OH)2(s) + OH− ⇌ NiOOH(s) + H2O(l) + e− |
Zn(s) → Zn2+(aq) + 2e− |
| EMF | E0 = −0.52 V | E0 = +0.76 V |
| Reduction | ||
| Occurs at | Cathode | Cathode |
| Sign of Terminal | Negative | Positive |
| Half-Reaction | Fe(OH)2(s) + 2e− ⇌ Fe(s) + 2OH− |
Ce4+(aq) + e− ⇌ Ce3+(aq) |
| EMF | E0 = −0.89 V | E0 = +1.44 V |
| Electrolytic Cell | ||
| EMF Signage | Negative (-) | Positive (+) |
| EMF Value | E0 = -1.41 V | E0 = +2.20 V |
| Electron Flow | From anode to cathode | From anode to cathode |
| Gibbs Free Energy (ΔG) | Positive (+) | Negative (-) |
| Spontaneous | No | Yes |
| Characteristics | Aqueous or molten electrolyte | Salt bridge with inert electrolyte (i.e. K2SO4) |
Reduction Potential
How can one know which metal combination to use or not to use, depending on the goal? Is it possible to deduce which of the materials involved in the corrosion system will corrode? Yes, each reduction reaction of an element is rated on a scale by its reduction potential, which refers to its ability to accept electrons. It is measured relative to the standard reference electrode (hydrogen), which is assigned a potential of 0.00 V under standard conditions.
The figure below illustrates this reduction potential, with negative values being the best electron donor (anode), also called the least noble materials in reference to the oxidation half-reaction, i.e., the more likely they are to oxidize. As another reference, Lithium (as in lithium-ion batteries) has a potential of -3.045V. Positive values indicate a good electron receiver (cathode) or noble materials for the reduction half-reaction. These indicators are also useful for evaluating the potential of an electrochemical cell, which will be defined by the added potentials of both its constituent electrodes.
Material Nobility and Potentials in Seawater8
NINE FAILURE MODES
Types of Corrosion
There are many types of corrosion that can occur on materials, as highlighted by the non-exhaustive list shown below9, which can induce damage to various parts of a structure or material. Sometimes, more than one corrosion type occurs at once, which can potentially compound the deleterious effects.
Galvanic
- Symptoms
- Powder-like white or grey deposit.
- Cause
- Two dissimilar metals in contact with each other in the presence of an electrolyte.
- Prevention
- Detail design, protective treatment, special assembly techniques (sealing, electrical insulation of metal).
Exfoliation
- Symptoms
- Flaking and loss of metal thickness.
- Cause
- Swelling and flaking at grain ends exposed by machining.
- Prevention
- Pre-heat treatment, material selection.
Pitting
- Symptoms
- Holes in metal surface.
- Cause
- Halogen ions present in attacking electrolyte (corrosive agent), destroying surface treatment.
- Prevention
- Protective treatment.
Filiform
- Symptoms
- Paint bulging and longitudinal tracking.
- Cause
- Paint damage.
- Prevention
- Corrosion-resistant primer, restoration of paint system.
Crevice
- Symptoms
- Severe local corrosion along adjoining surfaces.
- Cause
- Penetration of oxygen and corrosive agent into a joint (due to flexing).
- Prevention
- Efficient sealing of adjoining surfaces from corrosive substances.
Intergranular
- Symptoms
- Normally only perceived by cracking.
- Cause
- Chemical action along grain boundaries within the material; difference of electrical potential between grain and grain boundaries.
- Prevention
- Material selection, protective treatment.
Fretting
- Symptoms
- Destruction of natural protective film resulting from slight relative movement between mating surfaces and loss of metal from surfaces followed by oxidation.
- Cause
- Abrasion of metal under load in humid environmental conditions.
- Prevention
- Detail design, protective treatment, material selection.
Stress
- Symptoms
- Normally only perceived by cracking with fast crack propagation leaving bare metal subject to corrosion.
- Cause
- Residual stress from manufacturing process, or stress concentrations due to design features in a corrosive environment.
- Prevention
- Material selection, handling care, detail design, assembly techniques, background surface protection.
Microbiological
- Symptoms
- Local surface attack or formation of deposit such as fungi.
- Cause
- Growth of micro-organisms in moisture traps.
- Prevention
- Detail design, protective treatment, assembly techniques, use of inhibitors and primers.
One of the most common and destructive forms of corrosion is galvanic corrosion. This type of corrosion is defined as the accelerated corrosion of a metal due to electrical contact with a more noble metal or a non-metallic contactor in a corrosion electrolyte.
For galvanic corrosion to occur, the conditions of an electrochemical cell must be met. The materials must be electrochemically dissimilar, meaning that their reduction potentials are sufficiently different. As mentioned above, the most noble material, or the one with the highest potential, will act as the cathode in the corrosion cell, whereas the other will serve as the anode.
The anode will then corrode at an accelerated rate, while the other will corrode more slowly (or even stop). Galvanic corrosion is more severe at the junction between two dissimilar materials. It is a localized form of corrosion.
One might think that the further apart the materials are from each other in the previous scale illustration (see Material Nobility and Potentials in Seawater), the faster the corrosion rate will be. The galvanic series is indeed extremely helpful in predicting what happens when two metals are placed in electrical contact in a certain environment. However, it is not exactly accurate to say that the farther apart two metals are in the galvanic series, the greater the corrosion rate will be.
To predict the rate of galvanic corrosion, an Evans Diagram (see example below) of both metals involved must be considered. This diagram describes the four main reactions that can occur during the corrosion process, as electrons can flow freely through the metallic connection. The charge balance must be considered differently from that of individual metal corrosion, as shown in the galvanic series.
Below is an example of corrosion involving zinc and iron:
In this example, what is called the mixed potential of the couple is shown as a dotted red line, indicating the corrosion rate (Icorr) of each material. As the corrosion rate of zinc increases, the corrosion rate of iron decreases.
Of course, galvanic corrosion, as well as any other type of corrosion, depends as much on the materials used as on the environment to which they are subjected. Zinc, for example, is usually anodic to steel, although it can become cathodic when placed in fresh water between 60 °C and 90 °C.
Evans Diagram for Iron and Zinc in Acid10
CORROSION CAN BE ANYWHERE
Everyday Examples
Tomato Sauce
Let’s say you are home and want to prepare a good tomato sauce to go with your spaghetti. You take an aluminum pan and start following the recipe. After all your efforts, the sauce looks weird and tastes bitter. The aluminum pan had a chemical reaction due to the acidity of the tomatoes, forming an aluminum salt and some dihydrogen gas, which changed the taste and the look of your sauce. This type of corrosion is called uniform corrosion and occurs, in this case, in the presence of an acidic environment. This might surprise you as aluminum is well-known for being relatively resistant to corrosion, proving that understanding the phenomena of corrosion allows us to avoid the problems they generate as much as possible.11
Contrary to the general belief, metallic aluminum is much more reactive than iron, therefore less resistant to corrosion. What happens is that aluminum oxidizes quickly to form an aluminum oxide layer, which prevents further corrosion under neutral conditions. If you expose the aluminum to an acidic solution, e.g. tomato sauce, the aluminum oxide layer dissolves, further exposing the bare reactive aluminum metal. To make a long story short, don’t use an aluminum pan to make your tomato sauce.
What About Aluminum and Paint Stripper?
An issue that we came across more than once was an aluminum can into which a worker had poured methylene chloride to rinse paintbrushes. Methylene chloride, over the weekend, attacked the aluminum oxide layer, and the reactive aluminum reacted with it to produce phosgene and hydrochloric acid, leaving all metal surfaces in the immediate area with a corrosion film.
PROTECTION STRATEGIES
Preventing Corrosion
Large engineered systems employing many types of metal in their construction are susceptible to galvanic corrosion if care is not exercised during the design phase. Choosing metals that are as close together as practicable on the galvanic series helps reduce the risk of galvanic corrosion. The infrastructure of the Statue of Liberty, for example, had to be rebuilt using a duplex stainless steel structural frame, as copper alloys and stainless steels are quite close in their nobility. A small anode/cathode area ratio is also highly undesirable since the current will be concentrated onto a small anodic area, leading to rapid thickness loss. Preventing moisture/water from contacting the surface using an appropriate coating will also greatly reduce the corrosion risks.
Corrosion Protection Methods
Some metals or alloys react with their environment to produce a layer of oxides (e.g. PbO, Mg(OH)₂) which offers considerable protection to the pure metal or alloy underneath. Anodized aluminum is probably the most common natural protective film. This occurs usually in normal atmospheric conditions, though materials can be sensitive to particular environmental factors such as acidic or alkaline components.
Protective coatings can be applied to metals or alloys to enhance the thickness and quality of the natural oxide layer. Processes such as anodizing in acid electrolytes could multiply the level of protection depending on the material. Sometimes, a simple paint job can also go a long way in protecting products against corrosion.
Chemicals called inhibitors can be added to the product’s environment (fluid phase) in order to control corrosion. They form barriers to isolate the metal surface (solid phase) from the environment or control the anodic and cathodic corrosion reactions. Inhibitors include oxygen scavengers, filming inhibitors, biocides, and passivation agents. A major aspect of corrosion inhibition concerns the transport of the inhibitor to the interface, as the particles must reach the solid surface to be effective. This may help extend the useful life of a structure if the chemical and the corroding system are compatible as well as chemically, thermally, and, in some cases, physically stable. An example of this is the use of calcium nitrite as an admixture material in concrete, which will be exposed to high chloride environments.
Many industries use sacrificial anodes to slow down galvanic corrosion on their product. This type of protection is also called “cathodic protection”. The anode has to be changed after it has been completely consumed, although the usual standard is once only 10% of the anode is left. A sacrificial anode is a metal or alloy purposely put in electrical contact with a material that needs protection against galvanic corrosion and which will be placed in a cathodic position. That way, with time, the material protected will stay healthy and strong as the anode deteriorates. It goes without saying that the material composing the anode must have a more negative electrochemical potential than the other metal it will be used to protect. This means of prevention is also known as a sacrificial anode.
Zinc Protection — Real World Application
Zinc protection is a very common type of sacrificial anode often used on ships in order to protect the hull, the ballast tanks, and the heat exchangers, as well as the storage chests close to the bottom of the vessel, from corrosion. Pieces of zinc are bolted to the shaft of the boats and are more reactive, thus protecting the hull, which is the body of the ship that is exposed to water. Even non-metal boats use zinc anodes, which are connected by copper wires throughout the boat to protect underwater metal parts.
Zinc coatings are also often applied to provide barrier protection. The coating starts to corrode sacrificially first, therefore protecting the material underneath. This type of protection is usually used on smaller parts, such as screws or light switch plates, that will be exposed to corrosive environments.13
With a clear picture of what corrosion is, how it works at the electrochemical level, and what forms it takes in real materials and environments, the next step is practical: how do you actually test for it? Chapter 4 walks through the most widely used corrosion testing methods, how to choose between them, and how to build a test plan that generates results you can act on.
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- American Galvanizers Association. (n.d.). Zinc plating. In GALVANIZE IT! Retrieved January 29, 2026.