{"id":343,"date":"2022-11-17T20:37:49","date_gmt":"2022-11-17T20:37:49","guid":{"rendered":"https:\/\/public.ksc.nasa.gov\/corrosion\/?page_id=343"},"modified":"2022-11-17T20:37:50","modified_gmt":"2022-11-17T20:37:50","slug":"corrosion-fundamentals","status":"publish","type":"page","link":"https:\/\/public.ksc.nasa.gov\/corrosion\/corrosion-fundamentals\/","title":{"rendered":"Corrosion Fundamentals"},"content":{"rendered":"\n
<\/span>\"Corrosion
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Corrosion Fundamentals<\/p>\n<\/div><\/div>\n\n\n\n

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Fundamentals\u00a0of Corrosion<\/strong><\/p>\n\n\n\n

Corrosion can be defined as the degradation of a metal due to a reaction with its environment.

Degradation implies deterioration of physical properties of the material. This can be a weakening of the material due to a loss of cross-sectional area, it can be the shattering of a metal due to hydrogen embrittlement, or it can be the cracking of a polymer due to sunlight exposure. Materials can be metals, polymers (plastics, rubbers, etc.), ceramics (concrete, brick, etc.) or composites-mechanical mixtures of two or more materials with different properties. Because metals are the most used type of structural materials most of this web site will be devoted to the corrosion of metals.<\/p>\n\n\n\n

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Why Metals Corrode<\/strong><\/p>\n\n\n\n

Metals corrode because we use them in environments where they are chemically unstable. Only copper and the precious metals (gold, silver, platinum, etc.) are found in nature in their metallic state. All other metals, to include iron-the metal most commonly used-are processed from minerals or ores into metals which are inherently unstable in their environments.<\/p>\n\n\n\n

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This golden statue in Bangkok, Thailand, is made of the only metal which is thermodynamically stable in room temperature air. All other metals are unstable and have a tendency to revert to their more stable mineral forms. Some metals form protective ceramic films (passive films) on their surfaces and these prevent, or slow down, their corrosion process. The woman in the picture below is wearing anodized titanium earrings. The thickness of the titanium oxide on the metal surface refracts the light and causes the rainbow colors on her earrings. Her husband is wearing stainless steel eyeglasses. The passive film that formed on his eyeglasses is only about a dozen atoms thick, but this passive film is so protective that his eyeglasses are protected from corrosion. We can prevent corrosion by using metals that form naturally protective passive films, but these alloys are usually expensive, so we have developed other means of corrosion control.<\/p>\n<\/div>

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Electrochemical Cells<\/strong><\/p>\n\n\n\n

Oxidation and Reduction:<\/strong><\/p>\n\n\n\n

Metals are elements that tend to lose electrons when they are involved in chemical reactions, and nonmetals are those elements that tend to gain electrons.\u00a0
Sometimes these elements form ions, charged elements or groups of elements. Metallic ions, because they are formed from atoms that have lost electrons, are positively charged (the nucleus is unchanged). When an atom or ion loses electrons it is said to have been\u00a0oxidized.\u00a0
A common oxidation reaction in corrosion is the oxidation of neutral iron atoms to positively charged iron ions:<\/p>\n\n\n\n

Fe \u00bb Fe+2 + 2e-<\/p>\n\n\n\n

The electrons lost from a metal must go somewhere, and they usually end up on a nonmetallic atom forming a negatively charged nonmetallic ion. Because the charge of these ions has become smaller (more negative charges) the ion or atom which has gained the electron(s) is said to have been\u00a0reduced.<\/p>\n\n\n\n

4H+ +O2 + 4e- \u00bb 2H2O
or
2H+ +2e- \u00bb H2<\/p>\n\n\n\n

While other reduction reactions are possible, the reduction of oxygen is involved in well over 90% of all corrosion reactions. Thus the amount of oxygen present in an environment, and its ability to absorb electrons, is an important factor in determining the amount of oxidation, or corrosion, of metal that occurs.<\/p>\n\n\n\n

Electrochemical Reactions:<\/strong><\/p>\n\n\n\n

The two metal strips shown below are exposed to the same acid.<\/p>\n\n\n\n

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Both metals undergo similar oxidation reactions:<\/p>\n\n\n\n

Cu \u00bb Cu+2 + 2e-
and
Zn \u00bb Zn+2 + 2e-<\/p>\n\n\n\n

The electrons freed by the oxidation reactions are consumed by reduction reactions.<\/p>\n\n\n\n

On the copper the reduction reaction is: <\/p>\n\n\n\n

4H+ +O2 +4e- \u00bb 2H2O<\/p>\n\n\n\n

The corrosion rate of the copper is limited by the amount of dissolved oxygen in acid.\u00a0<\/p>\n\n\n\n

On the zinc the reduction reaction is: <\/p>\n\n\n\n

2H+ +2e- \u00bb H2<\/p>\n\n\n\n

The hydrogen ions are converted to hydrogen gas molecules and can actually be seen bubbling off from the acid.<\/p>\n\n\n\n

If we now connect the two metal samples with a wire and measure the electricity through the connecting wire, we find that one of the electrodes becomes different in potential than the other and that the corrosion rate of the copper decreases while the corrosion rate of the zinc increases. By connecting the two metals, we have made the copper a\u00a0cathode\u00a0in an electrochemical cell, and the zinc has become an\u00a0anode. The accelerated corrosion of the zinc may be so much that all of the oxidation of the copper stops and it becomes protected from corrosion. We call this method of corrosion control cathodic protection.<\/p>\n\n\n\n

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The reaction at the copper (cathode) becomes:<\/p>\n\n\n\n

2H+ +2e- \u00bb H2<\/p>\n\n\n\n

The voltage of the copper shifts to a point where hydrogen ion reduction can occur at the copper surface. The oxidation (corrosion) of the copper cathode may completely stop due to the electrical connection to the zinc anode. <\/p>\n\n\n\n

The reaction at the zinc (anode) remains the same: <\/p>\n\n\n\n

Zn \u00bb Zn+2 + 2e-<\/p>\n\n\n\n

The reaction rate increases due to the fact that the surface area of the clean (uncorroding) copper surface can now support a reduction reaction at a high rate. <\/p>\n\n\n\n

Thus connecting these two metals virtually stopped the corrosion of the copper and increased the corrosion rate of the zinc. We say that the zinc cathodically protected the copper from corrosion. Cathodic protection is a common means of corrosion control. 

Anodes are those portions of an electrochemical cell that have mostly oxidation reactions. Cathodes are those locations of an electrochemical cell that have mostly reduction reactions. One way to remember which kind of reaction predominates at each kind of electrode is to note that anode comes before cathode in the alphabet just like oxidation comes before reduction. Anodes oxidize; cathodes reduce.<\/p>\n\n\n\n

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