Rust is the common name given to a specific type of corrosion that exclusively affects iron and its alloys, such as steel. This familiar reddish-brown decay is scientifically known as hydrated iron(III) oxide, a compound formed when the metal reverts to a more chemically stable state. Understanding the causes of rust requires examining the environmental conditions and the complex electrochemical process that drives this transformation.
The Necessary Components
The formation of rust depends upon the presence of three specific materials acting together. The process requires iron or steel, the metal being corroded, and oxygen, typically supplied by the air. Finally, water or sufficient moisture must be present to act as the medium that facilitates the chemical reaction.
If any one of these three necessary components is removed, the rusting process cannot begin or sustain itself. Iron exposed to pure, dry oxygen, or iron submerged in de-oxygenated water will not rust. The simultaneous interaction of the metal, the gas, and the liquid is fundamental to initiating the chemical breakdown.
The Chemical Mechanism of Rust Formation
Rusting is a complex electrochemical reaction, similar to what occurs inside a battery. This process, a type of oxidation-reduction (Redox reaction), begins when the iron surface is exposed to water. The water droplet acts as an electrolyte solution, allowing electrons to move across the metal surface.
The initial step is oxidation, occurring at the anodic site on the iron surface. Iron atoms lose electrons and transform into \(\text{Fe}^{2+}\) ions, dissolving the metal into the water droplet. These liberated electrons travel through the conductive iron to a cathodic site, typically the edge of the water droplet where oxygen is readily available. At the cathode, dissolved oxygen accepts the electrons and reacts with water to form hydroxide ions (\(\text{OH}^-\)).
The \(\text{Fe}^{2+}\) ions generated at the anode travel through the water and encounter the \(\text{OH}^-\) ions produced at the cathode. This reaction initially forms unstable iron(II) hydroxide. The iron(II) hydroxide quickly undergoes further oxidation in the presence of oxygen to form \(\text{Fe}^{3+}\) ions. This final step results in the formation of hydrated iron(III) oxide (\(\text{Fe}_2\text{O}_3\cdot\text{nH}_2\text{O}\)), the familiar reddish-brown substance known as rust. Unlike the thin, protective oxide layers that form on metals like aluminum, rust is flaky and porous. This means it does not adhere tightly to the iron and continually exposes fresh metal underneath to the corrosive environment.
Environmental Factors That Speed Up Corrosion
While oxygen, water, and iron are the required reactants, certain environmental conditions can significantly accelerate the rate at which rust forms. The presence of electrolytes, such as salt, drastically increases the conductivity of the water film on the metal surface. Chloride ions from road salt or seawater allow the electrons and ions involved in the electrochemical reaction to move faster between the anode and cathode.
Temperature also plays a role, as higher temperatures increase the rate of most chemical reactions, including the steps in the rusting process. Warmer conditions accelerate the corrosion rate by increasing the kinetic energy of the reacting particles. Conversely, high humidity or prolonged moisture keeps the iron surface wet, ensuring a continuous supply of water necessary to complete the circuit.
Another accelerator is acidity, measured by a low pH level. Acidic environments accelerate the cathodic reaction by providing hydrogen ions (\(\text{H}^+\)), which react with the hydroxide ions, driving the overall process forward. Acid rain, for example, contains dissolved sulfur dioxide and nitrogen oxides, creating an acidic electrolyte that makes the corrosion process more aggressive than in neutral water.
How Rust Can Be Prevented
Preventing rust involves interrupting one or more of the three necessary components or the flow of the electrochemical current. A common strategy is barrier protection, which physically excludes oxygen and water from the metal surface. This method includes applying paint, oil, lacquer, or a protective coating of plastic.
Another effective method involves metallurgical changes, such as alloying iron with chromium to create stainless steel. Chromium reacts with oxygen to form a thin, dense, and non-porous layer of chromium oxide that self-heals if scratched. This stable layer acts as a permanent barrier, preventing the iron underneath from interacting with the environment.
Electrochemical protection is a third strategy that involves altering the electrical nature of the iron. Galvanization is a form of this protection, where steel is coated with a layer of zinc. Since zinc is more chemically reactive than iron, it acts as a sacrificial anode, corroding away instead of the steel. This ensures that the zinc atoms lose their electrons first, keeping the iron a protected cathode and halting the process.