How does oxygen influence corrosion?

We see it all the time: the reddish-brown flaking along an old metal gate, the dull green patina on a copper sculpture, the subtle pitting on a car part that seems to be suffering from neglect. Corrosion is common and costs industry billions, and yet when it is at work, it is more subtle than obvious.

Many have partnered with oxygen in this degradation process, and while corrosion is affected by many things, it is often catalyzed by one thing, an unseen catalyst that enables the destruction: oxygen.

But how does the life-giving gas, oxygen, become a destructive force for metals? How does it react, and why should this be any different? Let’s look at our relationship with oxygen and how it can both corrode and be a friend, how it develops and changes its form, and identify when it can make a difference to our metallic world.

Table of Contents

What is Corrosion, and Why Does it Happen?

Before going into the role played by oxygen, we will briefly review what corrosion is. Broadly speaking, corrosion is the natural degradation of a material usually a metal due to a chemical or electrochemical reaction with its environment.

In terms of purity, metals are not very stable materials. The extraction of metals from ores (i.e. usually oxides or other compounds) is a very energy intensive process. Corrosion is nature’s way of bringing metals back to a lower energy, stable state usually back to oxides.

The most common type of corrosion and the type of corrosion most metals including iron experiences is purely an electrochemical process.

That is, there is a transfer of electrons, which is fundamentally like a tiny battery which forms on the surface of the metal itself. For an electrochemical reaction to occur there are four fundamental components:

An Anode: The area of the metal which is losing some of its electrons (oxidation is occurring) and is corroding.
A Cathode: The area of the metal where some other substance is gaining some of the electrons (reduction is occurring).
An Electrolyte: A conducting medium (example – water or moisture containing dissolved salts) which allows the ions to move between the anode and the cathode.
A Metallic Path: The metal – which allows electrons to flow from anode to the cathode.

Without all four of these components significant corrosion typically cannot occur.

How Does Oxygen Directly Participate in the Corrosion Process?

Oxygen is a potent oxidizer and will accept electrons readily and is required for the reaction on the cathode or reduction site, especially when moisture or water is present (this is the same as the electrolyte in this electrochemical corrosion reaction).

To illustrate, we will take the corrosion of iron (Fe) through the case of Fe (while it is general corrosion, iron is common).

The area of the iron that is eroding is losing electrons and dissolving into iron (Fe) ions at the anode:

Fe→Fe²⁺+2e⁻

The electrons will travel through the conductor (the iron metal) to the cathode (another area on the iron surface where the reduction reaction would occur, or usually where oxygen is available in a good concentration). The oxygen will accept the electrons in this reduction reaction in the presence of water:

O₂+2H₂O+4e⁻→4OH⁻ (in neutral or alkaline conditions).

The iron (Fe) ion (Fe²⁺) and hydroxide (OH⁻) ions would combine to form iron hydroxide (Fe(OH)₂). The iron hydroxide would oxidize again in the presence of oxygen to produce hydrated iron (III) oxide or rust (Fe₂O₃⋅xH₂O).

Oxygen allows electrons to be removed from the metal surface that permits for the anodic reaction (metal dissolution) to proceed.

If the electrons are not “consumed” by the oxygen at the cathode, they would not travel through the anode to complete the electrochemical circuit, effectively stopping or greatly reducing anticipated corrosion. Oxygen is able to “depolarize” the cathode to permit for corrosion current to flow.

Why Does More Oxygen Often Mean Faster Corrosion?

If oxygen is a reactant in the corrosion process, it only makes sense that more oxygen should mean more corrosion. They said this would be true, and it usually is.

The more dissolved oxygen in the electrolyte, the more electron accepting oxygen molecules are at the cathodic sites, thus accelerating the cathodic reaction, allowing the anodic reaction to keep going faster and producing faster metal degradation.

You might think of it like a fire; the oxygen is the fuel. The more oxygen there is, the more it fuels and burns hotter and faster. In an electrochemical corrosion cell, more oxygen means a higher “burn rate” for your metal.

This continual level of availability of fresh oxygen is why structures that experience oscillating water levels like pipelines at the waterline or storage tanks that are filled and emptied experience severe corrosion. The fresh supply of oxygen continuously at the interface accelerates the attack.

Can Oxygen Concentration Differences Also Cause Corrosion?

Yes, and this is a very sneaky type of corrosion known as differential aeration corrosion or oxygen concentration cell corrosion.

Differential aeration corrosion happens when different parts of the same metal surface are in contact with the same electrolyte, but each metal surface is different because of differences in the amount of oxygen present.

Here’s how it works:

Low oxygen content places (e.g. under a dirt deposit; in a crevice; at the bottom of a water tank) become anodic and corrode.
High oxygen content places (e.g. exposed surfaces; areas with good circulation) become cathodic and are protected.

Why? Because in these low-oxygen regions, oxygen reduction (the cathodic reaction) is inhibited, making the dissolved metal in these areas preferentially donate its electrons to the oxygen-rich regions that can readily carry out the cathodic reaction. This results in localized and often severe corrosion in the low oxygen areas.

Examples entail:

Crevice corrosion: Occurs in a confined space or under a gasket . Oxygen could not circulate in confinement.
Pitting corrosion: Can be initiated where a small particle or deposit creates a local low-oxygen environment at localized areas of the metal surface which leads to a deep and localized hole.
Water line corrosion: Can occur occasionally on tanks or ships; here the metal just below the water line could be anodic (lower oxygen than the splash zone, higher oxygen than the submerged area).

How Do Other Environmental Factors Influence Oxygen’s Role?

Oxygen does not exist in isolation; its role in corrosion involves other factors in the environment, sometimes enhancing the corrosion process and other times reducing it.

Temperature: Generally, increasing temperature increases chemicals interaction rates, which includes corrosion. Generally, as temperature rises the diffusion rate of oxygen to the metal surface increases as does the kinetic energy of the species reacting, which speeds consumption of oxygen in a corrosion process. This is of special concern in systems, such as boilers, where hot water with dissolved oxygen present has a potential to cause catastrophic “oxygen attack.”
pH of solution: Acidity or alkalinity of an electrolyte has a dramatic effect on oxygen’s participatory role.
- In acidic environments (low pH), there are more readily available hydrogen ions (H+) to accept electrons, so oxygen’s role as the primary electron acceptor is weaker but still involving. The overall corrosion rate is higher in acidic environments.
- When the environment is neutral or slightly alkaline (as in most natural waters), oxygen reduction is the primary cathodic reaction and therefore presents the most risk for corrosion.
- In strongly alkaline environments (high pH), the reaction for some metals, such as aluminum, may become passive as a protective oxide is formed even in the presence of oxygen. However, other metals under alkaline conditions may still not produce the protective oxide and corrosion would not be inhibited.
Presence of other ions (especially chlorides): Chloride ions (Cl−) are particularly damaging and damaging because they can destroy passive oxide films that would otherwise protect the metal surface from the actions of oxygen. Also from chlorides, localized corrosion cells can form and oxygen can be a contributor in the process of unified loss, localized corrosion.
Flow rate/ Aeration: As has been mentioned previously, increased flow will introduce more oxygen to the surface of the metal, effectively increasing the rate of corrosion, however, stagnant conditions can also yield local areas from which oxygen can be depleted aiding a corrosion form called differential aeration corrosion.
Surface condition: A rough and/or scratched surface offers more initiation sites for corrosion to begin and a potential for micro-crevice that have differing concentrations of oxygen sets the stage for the accelerated localized corrosion that can occur.

Can Oxygen Ever Be Beneficial in Preventing Corrosion?

Yes! Oxygen is often the enemy, but also plays an important role in something called passivation.

Some metals such as stainless steel, aluminum and titanium react with oxygen to form a very thin, stable, and non-porous oxide layer at the surface.

This oxide layer, which is often only a few nanometers thick, serves as a barrier to corrosion and as such prevents further oxygen and other corrosive agents from reaching the underlying metal, thereby making it “passive” or resistant to corrosion for normal conditions.

When this passive layer is scratched or damaged, if there is still oxygen in the environment, the metal usually can “repassivate” by quickly reforming the oxide layer.

But without oxygen in a crevice or underneath a deposit on a passivating metal, localized breakdown of the passive film can occur, leading to severe pitting or crevice corrosion as mentioned above.

How Do We Control Oxygen’s Influence to Prevent Corrosion?

Understand oxygen’s role in corrosion is necessary to apply corrosion prevention techniques. The following are some common avenues of corrosion prevention:

Deaeration/Oxygen scavengers. In closed systems (i.e., boilers and pipelines), dissolved oxygen in water can be eliminated. The traditional method is to mechanically deaerate water by using heat and vacuum to remove oxygen. Then, later chemicals are used as an oxygen scavengers using sodium sulfate or hydrazine to deal with the little remaining dissolved oxygen after mechanical deaeration.
Protective coatings. The use of paints, polymers and metallic coatings (e.g., galvanizing with zinc) physically block access to the metal surface by providing an impermeable physical barrier through which oxygen and moisture cannot reach.
Cathodic Protection. The protective action involves making the structure number in terms of protection that the metal to be protected is the cathode of an electrochemical cell and the corrosion (anodic) reaction is displaced to another surface. While cathodic protection does not physically remove oxygen, oxygen reduction and consumption will occur on the sacrificial anode or inert electrode, where the reduction of oxygen does not form any corroding deposits or other products on the structure that is being protected from corrosion.
Material selection. Use metals and alloys that can resist corrosion, such as stainless steels that passivate and noble metals (gold and platinum) that have a noble oxidation state and have low reactivity.
Environmental control. The impact that oxygen can have on a cathodic reaction can be reduced depending on the control of temperature, and pH, or, and control the concentrations of aggressive ions such as chlorides close to the metal in the environmental surrounding. For example, a slightly alkaline pH in the water of and boiler will limit “oxygen” attack.

Why is This Knowledge Crucial for Everyone?

From the common spoon you have in your kitchen, to the vast pipeline network essentially delivering energy to the globe, learning how oxygen promotes corrosion is not solely the job for scientists and engineers. Here’s what you, the consumer in public and wannabe scientist or engineer will gain:

Homeowners: You will know what material is technically appropriate for your outdoor fixtures, know what proper drainage to consider around your foundations, and identify the first signs of rust.
Vehicle Owners: You will recognize the importance of rustproofing and your responsibility to regularly cleanaway contaminants from your vehicle to prolong life, particularly in humid areas or coastal environments.
Industry: You will know how to design safe durable infrastructure, reduce maintenance costs, and prevent catastrophic failures in vehicles, bridges, and chemical plants.
Environmentalists: You will learn how corrosion can be the cause of leaks and spills, which damages ecosystems.

The struggle against corrosion is still real, and our very own friend oxygen is often the biggest culprit. While we may not be able to stop corrosion entirely, understanding its capabilities and influences allows consumers to equip themselves with the information and resources to preserve or protect our metallic world for longer periods of time, safety, and efficiencies to come.