There are two definitions of oxidation, one involving oxygen, the other not. (a) State the two definitions of oxidation. (b) Are they compatible? Explain. (c) When we speak of corrosion of metals, what chemical reaction are we usually talking about? Use Fe as an example.

Oxidation can be defined in two different ways: 1. In terms of oxygen: Oxidation is the process where an element or compound reacts with oxygen, leading to the formation of an oxide. 2. In terms of electron transfer: Oxidation is the process of losing electrons by an atom or an ion during a chemical reaction. This is a part of the redox (reduction-oxidation) reaction where another species gains the lost electrons, called the reduction process.

Both of these definitions of oxidation are indeed compatible. In the first definition, the process involves reacting with oxygen, which is an electron acceptor, meaning that it tends to gain electrons during the reaction. When an atom or ion loses electrons to react with oxygen, it is going through the process of oxidation as defined in the second definition. So, in essence, both definitions describe similar processes, involving the transfer of electrons and, in most cases, the reaction with an electron acceptor, which is often oxygen.

The corrosion of metals generally refers to the chemical reaction of metals when they are exposed to the environment, leading to their eventual deterioration and loss of structural strength. A common example of corrosion is the rusting of iron (Fe). The rusting of iron involves oxidation and reduction reactions. The iron reacts with water and oxygen in the presence of moisture and forms hydrated iron(III) oxide, also known as rust. The process can be represented by the following equations: 1. Oxidation half-reaction (anodic reaction): \( \textrm{Fe} \rightarrow \textrm{Fe}^{2+} + 2\, \textrm{e}^{-} \) 2. Reduction half-reaction (cathodic reaction): \( \frac{1}{2}\, \textrm{O}_2 + \textrm{H}_2\textrm{O} + 2\, \textrm{e}^{-} \rightarrow 2\, \textrm{OH}^{-} \) By combining these half-reactions, we can get the overall redox reaction: \( \textrm{2Fe} + \frac{1}{2}\, \textrm{O}_2 + \textrm{H}_2\textrm{O} \rightarrow 2\, \textrm{Fe(OH)}_2 \) Further, the iron(II) hydroxide (Fe(OH)2) formed may also get oxidized by the dissolved oxygen to form iron(III) hydroxide: \( \textrm{4Fe(OH)}_2 + \textrm{O}_2 \rightarrow 4\, \textrm{Fe(OH)}_3 \) Finally, iron(III) hydroxide dehydrates to form hydrated iron(III) oxide, which is rust: \( \textrm{Fe(OH)}_3 \rightarrow \textrm{Fe}_2\textrm{O}_3\cdot \textrm{nH}_2\textrm{O} \) So, in the context of corrosion, the chemical reaction we are usually discussing is the redox reaction involving the transfer of electrons between the metal and its surrounding environment, ultimately leading to the formation of an oxide or hydroxide, as in the case of iron.