Galvanic cell rusting

Oxidation of the iron occurs at a location out of contact with the oxygen of the air. The surface of the metal acts as an anode in a tiny galvanic cell, with the metal at the outer edge of the drop serving as the cathode, (b) Further oxidation of Fe2+ results in the formation of Fe3+ ions, (c) Protons are removed from H,0 as oxide ions combine with Fe3+ ions to deposit as rust. These protons are recycled, as indicated by the dotted line. 

Rust is a hydrated iron(III) oxide, Fe203 XH2O. The electrochemical formation of rust occurs in small galvanic cells on the surface of a piece of iron, as shown in Figure 11.28. In each small cell, iron acts as the anode. The cathode is inert, and may be an impurity that exists in the iron or is deposited onto it. For example, the cathode could be a piece of soot that has been deposited onto the iron surface from the air. 

The rusting of iron involves the reaction of iron, oxygen, and water in a naturally occurring galvanic cell on the exposed surface of the metal. There may be many of these small cells on the surface of the same piece of iron. 

Cathodic protection is another method of preventing rusting, as shown in Figure 11.29. As in galvanizing, a more reactive metal is attached to the iron object. This reactive metal acts as a sacrificial anode, and the iron becomes the cathode of a galvanic cell. Unlike galvanizing, the metal used in cathodic protection does not completely cover the iron. Because the sacrificial anode is slowly destroyed by oxidation, it must be replaced periodically. 

If iron is covered with a protective layer of a metal that is less reactive than iron, there can be unfortunate results. A tin can is actually a steel can coated with a thin layer of tin. While the tin layer remains intact, it provides effective protection against rusting. If the tin layer is broken or scratched, however, the iron in the steel corrodes faster in contact with the tin than the iron would on its own. Since tin is less reactive than iron, tin acts as a cathode in each galvanic cell on the surface of the can. Therefore, the tin provides a large area of available cathodes for the small galvanic cells involved in the rusting process. Iron acts as the anode of each cell, which is its normal role when rusting. 

Sometimes, the rusting of iron is promoted accidentally. For example, by connecting an iron pipe to a copper pipe in a plumbing system, an inexperienced plumber could accidentally speed up tbe corrosion of tbe iron pipe. Copper is less reactive than iron. Therefore, copper acts as the cathode and iron as the anode in numerous small galvanic cells at the intersection of the two pipes. 

Zinc has been nscd for ages to coat iron pails and pipes to prevent them from rusting — "galvanized iron. Zinc is also a part of many alloys (German silver and brass) and is important in the making of dry-cell batteries. 

Use and exposure Zinc is available as a silver or bluish-white foil or powder. It is incompatible with amines, cadmium, sulfur, chlorinated solvents, strong acids, and strong bases. The important use of zinc is to coat iron or steel in a process called galvanization to prevent rust. Zinc powder is very flammable. Zinc is another essential micronutrient that is important in immunity and antioxidation. Zinc is an essential mineral that is found in almost every cell function. It stimulates the activity of approximately 100 enzymes, which are substances that promote biochemical reactions in the body. Zinc supports a healthy immune system that the body requires for wound healing. It helps to maintain a sense of taste and smell and is needed for DNA synthesis. Zinc supports normal growth and development during pregnancy, childhood, and adolescence. ... 

Iron or steel is often covered by a thin layer of a second metal to prevent rusting Tin cans consist of steel covered with tin, and galvanized iron is made by coating iron with a layer of zinc. If the protective layer is broken, however, iron will rust more readily in a tin can than in galvanized iron. Explain this observation by comparing the half-cell potentials of iron, tin, and zinc. 

All of the general t5 es of corrosion attack occur in the atmosphere. Since the corroding metal is not bathed in large quantities of electrolyte, most atmospheric corrosion operates in highly localized corrosion cells, sometimes producing patterns difficult to explain as in the example of the rusting galvanized roof shown in Fig. 9.1. 

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