Rust is the common term for the reddish-brown decay that appears on iron and its alloys, such as steel. This process is a chemical reaction known as oxidation, where the metal slowly combines with oxygen to form a new compound called iron oxide. Iron oxide is brittle and flaky, and understanding this common form of corrosion requires a look at the ingredients and conditions that drive this natural chemical process.
The Essential Ingredients for Rust
Rusting is an electrochemical process, which means it involves the movement of electrons. For this reaction to occur, three components must be present: iron, oxygen, and water. The iron in the nail acts as the anode, the site where iron atoms lose electrons (oxidation), turning the metal into positively charged iron ions.
The electrons released by the iron atoms travel through the metal to the cathode site. Here, oxygen reacts with water and the incoming electrons (reduction) to form hydroxide ions. Water is necessary because it acts as an electrolyte, a medium that facilitates the movement of these ions and completes the electrical circuit.
The iron ions and hydroxide ions then meet within the water droplet on the surface of the nail. They combine and react further with oxygen to produce hydrated iron(III) oxide, which is the chemical composition of the flaky, reddish-brown substance we recognize as rust. Unlike the protective oxide layer that forms on aluminum, iron oxide is porous and non-adherent, meaning it continuously flakes off to expose fresh metal to the environment, allowing the corrosive cycle to continue.
Environmental Factors That Speed Up Corrosion
While iron, oxygen, and water are necessary for rust to form, certain environmental conditions accelerate the electrochemical reaction. The presence of dissolved salts significantly speeds up corrosion because they increase the electrical conductivity of the water. Saltwater acts as a superior electrolyte, allowing electrons to transfer more quickly between the anode and cathode sites on the nail’s surface.
The acidity of the environment, measured by pH, also affects the speed of decay. Extremely acidic conditions, such as those found in acid rain or near certain industrial chemicals, accelerate the rust rate. In environments with a pH below 4.0, the initial protective layer of iron oxide is soluble, so it dissolves as quickly as it forms, continuously exposing the fresh metal underneath to corrosive elements.
Temperature and humidity influence the speed of the reaction. Chemical reactions occur faster at higher temperatures, meaning warm, moist environments promote quicker rusting. High relative humidity, typically above 50%, provides the necessary concentration of water on the metal surface for the electrochemical process to sustain itself.
Practical Ways to Stop Nails From Rusting
Preventing rust requires interrupting the chemical process by eliminating one or more of the three essential ingredients. The simplest method is to apply a protective coating, such as paint, oil, or lacquer, which creates a physical barrier between the iron and the surrounding oxygen and moisture. However, this method is only effective as long as the coating remains intact.
A more robust method is galvanization, which involves coating the iron nail with a layer of zinc. Zinc acts as a physical barrier to the elements. If the zinc coating is scratched and the iron is exposed, the zinc still protects the iron through sacrificial protection.
Since zinc is more chemically reactive than iron, it corrodes preferentially, sacrificing itself to protect the underlying nail. The zinc acts as the anode, continuously giving up electrons to the iron. This forces the iron to remain the cathode, suppressing its tendency to oxidize into rust. Another approach is to use stainless steel nails, an iron alloy containing at least 10.5% chromium. The chromium rapidly reacts with oxygen to form a thin, tightly bonded layer of chromium oxide on the surface, known as a passive film. This film acts as a self-healing protective shield that prevents oxygen and water from reaching the iron, effectively halting the corrosion process.