Oxidation, visible as the rusting of iron or the greenish patina on copper, is a fundamental chemical process. This degradation, often called corrosion, is the natural tendency of refined metals to return to a more chemically stable state, similar to the ores from which they were originally extracted. This process is driven by the interaction of the metal with its surrounding environment.
Oxidation Defined as Electron Loss
Metal oxidation is a chemical reaction involving the transfer of electrons, a process known as a redox reaction. Specifically, oxidation is defined as the half-reaction where an atom loses one or more electrons to become a positively charged ion. This electron loss must be coupled with a gain of electrons by another substance in a simultaneous process called reduction. The metal acts as the reducing agent, giving up its electrons, while the environmental component that accepts the electrons is the oxidizing agent. This transfer of electrical charge creates an electrochemical circuit that powers metal degradation.
The Role of Oxygen and Water
For the electron loss required for oxidation to occur efficiently, two environmental components must be present: oxygen and water. Oxygen acts as the primary electron acceptor, or oxidizing agent, in the reduction half-reaction, combining with the electrons released by the metal. Without oxygen, the metal’s electrons have no immediate destination, slowing the entire process significantly.
Water or high moisture levels are necessary because they act as an electrolyte, a medium that allows ions to move freely. The electrolyte facilitates the flow of electrons between the area where the metal is dissolving and the area where oxygen is accepting the electrons. Common impurities dissolved in the water, such as salts or acids, accelerate the degradation process by making the water a much better electrolyte. For example, saltwater allows electrical current to flow more easily, speeding up the rate at which the metal can lose its electrons. Carbon dioxide in the air dissolves in water to form a weak acid, which also enhances the water’s ability to conduct ions and speed up the reaction.
Why Reaction Rates Vary Between Metals
The variation in degradation rates is primarily due to the metal’s inherent chemical reactivity, which dictates how easily it gives up its electrons. Some metals, like gold, are naturally less reactive and highly resistant to oxidation.
A more common difference is seen in a phenomenon called passivation, which dramatically alters the degradation rate. Metals like aluminum, chromium, and the alloys in stainless steel are highly reactive, but upon initial exposure to air, they form an extremely thin, dense, and stable oxide layer on their surface. This layer acts as a physical barrier, preventing further contact between the underlying metal and the environment. In contrast, the oxide formed on iron, commonly known as rust, is porous, flaky, and loosely bound to the metal’s surface. This non-passivating oxide flakes away to expose fresh metal, allowing the oxidation process to continue unchecked until the entire material is consumed.
Methods for Blocking the Process
Preventing metal degradation requires interrupting the conditions necessary for the process to occur. One strategy involves using barrier coatings, such as paint, oil, or plating with a non-reactive metal like tin, which physically separate the metal surface from surrounding oxygen and moisture.
Another approach is to modify the metal itself through alloying, which forces a protective outcome. Stainless steel, for example, contains chromium, which spontaneously forms a dense, self-repairing passive oxide layer that protects the underlying iron. This internal modification ensures the metal is resistant to environmental attack even if the surface is scratched.
A third method is sacrificial protection, most commonly seen in galvanization, where a layer of zinc is applied to steel. Zinc is more reactive than iron, meaning it loses its electrons more easily. If the coating is damaged, the zinc sacrifices its own electrons to protect the underlying iron, ensuring the iron remains intact until the zinc layer is fully consumed.