Corrosion and Its Prevention

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  • Topic: Corrosion, Metal, Iron
  • Pages : 5 (1669 words )
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  • Published : July 29, 2010
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Corrosion is defined as the involuntary destruction of substances such as metals and mineral building material by surrounding media, which are usually liquid (i.e. corrosive agents)." Most metals corrode. During corrosion, they change into metallic ions. In some cases, the product of corrosion itself forms a protective coating. "For example, aluminium forms a thin protective oxide layer which is impervious to air and water. In other cases (e.g. iron), however, the coating either flakes off or is pervious to both air and water. So the whole piece of metal can corrode right through."  The most common forms of metallic corrosion are caused by electrochemical reactions, wherein two metallic phases (e.g., iron oxide and iron) react in the presence of electrolytic solution. Another mechanism of metallic corrosion is caused by chemical reaction, which explains how the protective layer of the metal is formed.  Rusting is the corrosion of iron which is the most widely used structural metal. Most of it is used in making steel. The wide range of products made from steel includes all types of vehicles, machinery, pipelines, bridges, and reinforcing rods and girders for construction purposes. Therefore, rusting causes enormous economic problem and is the reason why extensive measures of corrosion protection have had to be developed.

Electrochemical corrosion reactions
This type of corrosion takes place when two metallic phases with different electrochemical potentials are connected to each other by means of an electric conductor. 

Chemical corrosion reactions
Metals have a tendency to combine with oxygen to form oxides and this is one of the chemical reactions. This tendency is the stronger the less noble the metal. The layers of oxide on the metal surface which are formed even in dry air may be insoluble and stable against an aqueous medium in contact with them. Therefore, if the oxide layers are dense and adhere well to the metal, they prevent further attack and act as a corrosion prevention layer. An example of this is aluminum oxide. However, iron differs in that, although it does form a surface oxide layer, this layer is loose and enables oxidation to proceed into the depth of the metal.  Chemical corrosion also takes place by the action of acids and alkalis on metals. Hydrochloric acid, for example, reacts with iron, and sodium hydroxide with aluminum (Figure 1). If soluble reaction products are formed, the reaction only ends when either the aggressive medium, or the metal are used up; if salts are formed which are sparingly soluble they can form protective layers.  Factors that speed up rusting

1) Presence of electrolytes
Acid solutions make rusting go faster. In industrial areas where air is seriously polluted, there are high concentrations of carbon dioxide, sulfur dioxide and nitrogen dioxide. These gases dissolve in rain water to give "acid rain", which makes iron objects rust faster.  Sodium chloride also makes iron rust more quickly.

2) Heat
An increase in temperature always increases rate of chemical reactions, including rusting. 3) Humidity
"Corrosion starts when the relative humidity of the air exceeds around 65%. Many areas has a higher humidity in winter (80-95%) than in summer (60-80%)" . In consequence, iron rusts five times faster in winter as it does in summer. 4) Contact with a less reactive metal

Consider iron and copper plates joined together and put in water containing dissolved oxygen. Iron loses electrons more readily than copper. Hence iron forms the anode and copper the cathode of an electrochemical cell. In this case, iron rusts even more quickly than when there was no copper.  5) Other factors

Other factors that speed up rusting include the presence of sharply pointed regions in the iron piece, or a high concentration of dissolved oxygen in water. 

Protection From Rusting
Iron is such a useful metal yet it rusts. Rusting is a serious problem. A very sum of many is...
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