2Fe+2H2O+O2→2Fe2++4OH→2FeOH2 Ferrous(Fet) hydroxide precipitates from solution. However, this compound is unstable in oxygenated solutions and is oxidized to the ferric(Fe)salt 2Fe(OH)2 +H2O+-O2- 2Fe(OH) The final product is the familiar rust The classic example of a replacement reaction, the interaction of zinc with copper sulfate solution, illustrates metal deposition Zn + Cu Zn-+ Cu (Eq18) or, viewed as partial reactions Zn→Zn2++2e The zinc initially becomes plated with copper and eventually, given enough time and reactants, the products are copper sponge and zinc sulfate solution During corrosion, more than one oxidation and one reduction reaction may occur. when an alloy is corroded its component metals go into solution as their respective ions. More importantly, more than one reduction reaction can occur during corrosion. Consider the corrosion of zinc in aerated hydrochloric acid. Two cathodic reactions are possible: the evolution of hydrogen and the reduction of oxygen. Since the rates of oxidation and solutions containing dis c Increasing the total reduction rate increases the rate of zinc solution. Therefore, acid reduction must be equal solved oxygen normally will be more corrosive than air-free acids. Oxygen reduction simply provides another means of electron disposal. The same effect is observed if any oxidizer is present in acid solutions. A frequent impurity in commercial hydrochloric acid is the ferric ion(Fe), present as ferric chloride. Metals corrode much more rapidly in such impure acid because there are two cathodic reactions hydrogen evolution and ferric ion reduction q21) The anodic and cathodic reactions occurring during corrosion are mutually dependent, and it is possible to reduce corrosion by reducing the rates of either reaction. In the above case of impure hydrochloric acid, it can be made less corrosive by removing the ferric ions and consequently reducing the total rate of cathodic reduction. Oxygen reduction is eliminated by preventing air from contacting the aqueous solution or by removing air that has been dissolved. Iron is nearly inert in air-free water or seawater because there is limited athodic reaction possible If the surface of the metal is coated with paint or other nonconducting film, the rates of both anodic and cathodic reactions will be greatly reduced and corrosion will be retarded. A corrosion inhibitor is a substance that, when added in small amounts to a corrosive, reduces its corrosivity. Corrosion inhibitors function by interfering with either the anodic or cathodic reactions, or both. Many of these inhibitors are organic compounds; they function by forming an impervious film on the metal surface or by interfering with either the anodic or cathodic reactions. High-molecular-weight amines retard the hydrogen-evolution reaction and subsequently reduce corrosion rate. Adequate conductivity in both the metal and the electrolyte is required for continuation of the corrosion reaction. Of course, it is not practical to increase the electrical resistance of the metal because the sites of the anodic and cathodic reactions are not known, nor are they predictable. However it is possible to increase the electrical resistance of the electrolyte and thereby reduce corrosion. Very pure water is much less corrosive than impure or natural waters. The low corrosivity of high-purity water is due to its high electrical resistance and few reducible cations Passivity. Essentially, passivity refers to the loss of chemical reactivity experienced by certain metals and alloys under particular environmental conditions. That is, certain metals and alloys become essentially inert and act as if they were noble metals such as platinum and gold. Fortunately, from an engineering standpoint, the metals most susceptible to this kind of behavior are the common engineering and structural materials, including iron metals such as zinc, cadmium, tin, uranium, and thorium have also been observed to exhibit passivity effect r nickel, silicon, chromium, titanium, and alloys containing these metals. Also, under limited conditions, othe2Fe + 2H2O + O2 → 2Fe2+ + 4OH- → 2Fe(OH)2 (Eq 16) Ferrous (Fe2+) hydroxide precipitates from solution. However, this compound is unstable in oxygenated solutions and is oxidized to the ferric (Fe3+) salt: 2Fe(OH)2 + H2O + 1 2 O2 → 2Fe(OH)3 (Eq 17) The final product is the familiar rust. The classic example of a replacement reaction, the interaction of zinc with copper sulfate solution, illustrates metal deposition: Zn + Cu2+ → Zn2+ + Cu (Eq 18) or, viewed as partial reactions: Zn → Zn2+ + 2e (Eq 19) Cu2+ + 2e → Cu (Eq 20) The zinc initially becomes plated with copper and eventually, given enough time and reactants, the products are copper sponge and zinc sulfate solution. During corrosion, more than one oxidation and one reduction reaction may occur. When an alloy is corroded, its component metals go into solution as their respective ions. More importantly, more than one reduction reaction can occur during corrosion. Consider the corrosion of zinc in aerated hydrochloric acid. Two cathodic reactions are possible: the evolution of hydrogen and the reduction of oxygen. Since the rates of oxidation and reduction must be equal, increasing the total reduction rate increases the rate of zinc solution. Therefore, acid solutions containing dissolved oxygen normally will be more corrosive than air-free acids. Oxygen reduction simply provides another means of electron disposal. The same effect is observed if any oxidizer is present in acid solutions. A frequent impurity in commercial hydrochloric acid is the ferric ion (Fe3+), present as ferric chloride. Metals corrode much more rapidly in such impure acid because there are two cathodic reactions, hydrogen evolution and ferric ion reduction: Fe3+ + e → Fe2+ (Eq 21) The anodic and cathodic reactions occurring during corrosion are mutually dependent, and it is possible to reduce corrosion by reducing the rates of either reaction. In the above case of impure hydrochloric acid, it can be made less corrosive by removing the ferric ions and consequently reducing the total rate of cathodic reduction. Oxygen reduction is eliminated by preventing air from contacting the aqueous solution or by removing air that has been dissolved. Iron is nearly inert in air-free water or seawater because there is limited cathodic reaction possible. If the surface of the metal is coated with paint or other nonconducting film, the rates of both anodic and cathodic reactions will be greatly reduced and corrosion will be retarded. A corrosion inhibitor is a substance that, when added in small amounts to a corrosive, reduces its corrosivity. Corrosion inhibitors function by interfering with either the anodic or cathodic reactions, or both. Many of these inhibitors are organic compounds; they function by forming an impervious film on the metal surface or by interfering with either the anodic or cathodic reactions. High-molecular-weight amines retard the hydrogen-evolution reaction and subsequently reduce corrosion rate. Adequate conductivity in both the metal and the electrolyte is required for continuation of the corrosion reaction. Of course, it is not practical to increase the electrical resistance of the metal because the sites of the anodic and cathodic reactions are not known, nor are they predictable. However, it is possible to increase the electrical resistance of the electrolyte and thereby reduce corrosion. Very pure water is much less corrosive than impure or natural waters. The low corrosivity of high-purity water is due to its high electrical resistance and few reducible cations. Passivity. Essentially, passivity refers to the loss of chemical reactivity experienced by certain metals and alloys under particular environmental conditions. That is, certain metals and alloys become essentially inert and act as if they were noble metals such as platinum and gold. Fortunately, from an engineering standpoint, the metals most susceptible to this kind of behavior are the common engineering and structural materials, including iron, nickel, silicon, chromium, titanium, and alloys containing these metals. Also, under limited conditions, other metals such as zinc, cadmium, tin, uranium, and thorium have also been observed to exhibit passivity effects