Redox reactions are those in which electrons are transferred from one atom to another, and they are at the heart of many chemical and biological processes. From rusting iron to forest fires to the beating of a human heart, redox reactions are the hero of many chemical and biological processes.
Rust is the common name for a very common compound, iron oxide. Iron oxide, the chemical Fe2O3, is common because iron combines very readily with oxygen — so readily, in fact, that pure iron is only rarely found in nature. Iron rusting is an example of corrosion — an electrochemical process involving an anode (a piece of metal that readily gives up electrons), an electrolyte (a liquid that helps electrons move) and a cathode (a piece of metal that readily accepts electrons). When a piece of metal corrodes, the electrolyte helps provide oxygen to the anode. As oxygen combines with the metal, electrons are liberated. When they flow through the electrolyte to the cathode, the metal of the anode disappears, swept away by the electrical flow or converted into metal cations in a form such as rust. For iron to become iron oxide, three things are required: iron, water and oxygen. Here’s what happens when the three get together:
When a drop of water hits an iron object, two processes begin to happen almost immediately. First, the water, a good electrolyte, combines with carbon dioxide in the air to form a weak carbonic acid, an even better electrolyte. As the acid is formed and the iron dissolved, some of the water will begin to break down into its component pieces — hydrogen and oxygen. The free oxygen and dissolved iron bond into iron oxide, in the process freeing electrons. The electrons liberated from the anode portion of the iron flow to the cathode, which may be a piece of a metal less electrically reactive than iron, or another point on the piece of iron itself.
Protection for corrosion can be done in many types; they are galvanization, cathodic protection and painting. Galvanizing is the process where the iron material is coated with zinc. Since zinc is a better reducing agent than iron, it readily oxidizes instead of iron, thereby protecting the iron material. This method very effective because, even after a part of the zinc coating is scratched off, the iron is still protected. Cathodic protection in its simplest form is achieved by attaching a sacrificial anode thus making the iron or steel the cathode in the cell formed. The sacrificial anode must be made from something with a more negative electrode potential than the iron or steel, commonly zinc, aluminium or magnesium. Rust formation can also be controlled with coatings, such as paint, that isolate the iron from the environment.
Teaching in Australiaemphasizes a lot on redox chemistry in both years 11 and 12 stages. It provides a lot of different examples and varieties which makes the subject more interesting to study.