Corrosion

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Submitted By EmersonBryan
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The truth that many people never understand, until it is too late, is that the more you try to avoid suffering the more you suffer because smaller things begin to torture you in proportion to your fear of suffering. --Thomas Merton

The net reaction for this first simple step is therefore: 2(Fe O2 + 2 H2O + 4 e2 Fe + O2 + 2 H2O Fe2+ + 2 e-) 4 OH2 Fe(OH)2

Iron(II) hydroxide is insoluble but its green color is almost never observed because it is ordinarily further oxidized by the oxygen: 2 Fe(OH)2 + ½ O2 + H2O → 2 Fe(OH)3 The final product (when dry) has the reddish-brown flaky character we associate with rust. Although the reaction that produces Fe(OH)2 is technically an equilibrium process (all electrochemical processes are) the value of Kc is very large (>1099 at 298 K) and left unchecked it will go to completion. But the rate is relatively slow under normal atmospheric conditions and so it is still possible to manipulate the equilibrium somewhat by changing appropriate factors. The rates of corrosion reactions--and presumably their mechanisms--vary widely. Factors which influence the progress of the net reaction in the first step of the oxidation of iron may have an effect on the overall rate. The nature of the oxide product is also very important in affecting the extent of the corrosion. For example, aluminum is a very active metal, but its oxide, Al2O3, is very dense and forms a thin protective layer on the metal which discourages further corrosion. In contrast, iron rust (hydrated forms of Fe2O3 such as reddish-brown Fe(OH)3) is typically flaky and easily crumbles off to continually expose fresh metal for reaction. Although the mechanism for corrosion is not always well understood, it is clear how to prevent it. The surface of the metal must be protected from contact with oxygen. Paints, oils and other coatings are often used for this purpose. But it is…...

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