Sample Course work Paper on Corrosion Chemistry Video

Corrosion Chemistry Video


Corrosion is a chemical process through which metals disintegrate. Rust, for example, is a product of corrosion. It occurs when iron reacts with water and oxygen. Rust is a hydrated iron oxide that is formed through a series of reactions. First, iron loses electrons to form iron II ions. Second, water reacts with oxygen to form hydroxide ions. Third, iron II ions combine with hydroxide ions to form iron II hydroxide. The resultant Iron II hydroxide further reacts with water and oxygen to form hydrated iron II oxide—rust. The number of water molecules in rust determines its color. All metals corrode, not just iron. The rate of corrosion depends on the position of the metal in the reactivity series. Oxygen and water are the conditions required for steel or iron to rust. When three nails are each put in different test tubes—the first with water and a layer of oil on top of the water layer, the second with water and the third with air, the nail in the second test tube rusts more while the nail in the third test tube does not rust. When iron is exposed to air for a long period, rusting can occur because air contains water vapor. The presence of salts in water accelerates corrosion.

Corrosion is an oxidation-reduction reaction (redox) reaction. The metal that undergoes corrosion becomes oxidized (loses electrons) while the substance that takes up electrons from the metal becomes reduced. Corrosion thus involves electron transfer. The substance that loses electrons is the reductant while the one that gains electrons is the oxidant. In a galvanic cell, the electrons flow from the more reactive metal (anode) to the less reactive metal (cathode). Therefore, the more reactive metal corrodes. Corrosion destroys metal structures, especially those made of iron and steel because the conditions necessary for these metals to corrode (water and air) are naturally available. Therefore, corrosion minimization activities, such as galvanization are necessary.


Overall, the video makes the topic appear much simpler and more focused than it would have been in a lecture. I could stay engaged throughout the presentation because the narrator displayed many real-life scenarios relevant to the topic while discussing the chemistry behind them. The narrator was also interesting, especially looking like a layperson taking you on an educative tour rather than a professor or a scientist. Initially, I had thought corrosion and rusting meant the same thing but I learned that corrosion is a wider term referring to all chemical reactions in which a substance (mostly metals) disintegrates. Rusting, on the other hand, is the corrosion of iron to form iron III oxide.

However, the narrator in the video made a mistake at 03:40 by displaying an equation for the formation of hydroxide ions from oxygen and water while talking about the formation of iron II ions from iron. The correct equation should have been Fe→Fe2+ + 2e. In addition, the narrator introduced the topic as corrosion but spent more than half of the time talking about iron. Although I now understand corrosion as much broader than rusting, the narrator’s extended discussion of how iron rusts suggests that corrosion and rusting are equivalent in meaning. The question I still have is, must oxygen and water be present for corrosion (not just rusting) to occur? It is clear that water is needed for ions to form and flow in galvanic cells, but the role of oxygen in these reactions is not clear. Furthermore, the narrator uses half-reactions extensively and I had to refresh my knowledge of electrochemistry to catch up. Perhaps the video should state the required prior knowledge right at the beginning so that the viewer can be fully prepared.


Corrosion has a wide range of applications in various fields, especially engineering. For example, the knowledge of redox reactions in a galvanic cell is used to protect underground metallic structures from corroding, a principle called cathodic protection (Bardal, 2004, p. 282). Underground iron structures corrode because of the presence of moisture and air in the soil. To protect these structures, a galvanic cell in which iron is connected to another more reactive metal, such as zinc is created. Zinc, being more reactive than iron, loses electrons (oxidation). The electrons from the zinc electrode are transferred to the iron thus preventing iron from corroding. Zinc corrodes instead and becomes the sacrificial metal in the reaction.

Corrosion is also applied in the manufacturing of tin-coated steel cans for packaging food and drinks. Since tin resists corrosion under natural conditions, it is more suitable for making cans for preserving food. However, the pure tin would be too expensive and less economical to use. Therefore, manufacturers of can use steel because it is readily available. Since steel contains iron, it is highly prone to corrosion. The solution is to coat steel cans with tin. To do this, an electrolytic cell is set up with tin and steel as the anode and cathode respectively. Tin corrodes while steel is covered with a layer of tin. The same principle is used in the making of galvanized iron sheets and gold- or silver-coated ornaments. Furthermore, the redox reactions involved in corrosion are applied in voltaic cells to generate electricity.

Bardal, E. (2004). Corrosion and protection. London: Springer.

Corrosion Chemistry Video. Retrieved April 21, 2014 from