Corrosion And Rusting (2147 words)
Corrosion and RustingIntroduction
Some people may be annoyed by their car “wearing out”. Kids may have trouble with rust forming on their bicycles. One may think how to prevent rusting, but do one knows what is happening when a metal corrode?
“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.
The economic importance of corrosion and corrosion protection can be shown by the following example: “It is estimated that roughly 3% of the annual production of steel is lost by corrosion. In 1974, 140 millions tons of steel were produced in the United States at a cost of approximately $400 per tons. So this gives a monetary loss of about 1.7 billion dollars.” It is clearly of the utmost importance to reduce as far as possible the financial loss by corrosion, which not only affects steel but to the extent all other building metal as well.
It is obvious that corrosion and rusting affect significantly the life of the society, so it is worthy to investigate this topic. In this essay, the cause of the corrosion and rusting and consequently the protection of the corrosion will be explored.
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.
Electrolytes such as acids, alkalis, salt solutions, or even milder media (e.g., rainwater, river water, groundwater, or tap water) also need to be present. “Metallic phases with different electrochemical potentials exhibit electric potential differences. Potential differences may also arise because of impurities, internal stresses, corrosion products, damaged protective coatings, etc. They also occur when different metals are used. The larger the potential difference, the faster the rate of corrosion.”
The electrochemical EMF series (Table 1) gives the electrochemical potential of metals under normal conditions with respect to hydrogen (hydrogen is 0). The farther two metals in electrochemical series are apart, the larger the potential difference between them. A metal is said to be less noble than those which stand to its right in the electrochemical series. In the case of electrochemical corrosion it is always the less noble metal which is removed.
Table 1. Electrochemical Potential Series, Volts.
K Ca Mg Al Zn Cr Fe Ni Sn Pb H Cu Ag Au
-2.92 -2.84 -2.38 -1.66 -0.76 -0.71 -0.44 -0.24 -0.14 -0.13 0.00 0.34 0.80 1.42
not noble —————————————————————–> noble
Likelihood of passing into solution decreases from left to right.
The potential difference does not, however, always fully correspond with the corrosion phenomena experienced in practice. The reason is that oxide and other metal compounds have differing electrochemical potentials.
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.
Figure 1. Chemical corrosion as shown by acid attack..
The chemical reaction of forming the hydroxide or oxide layers is the cause of producing the “rust”.
Rusting
Rusting refers to the corrosion of iron. As irons is the most widely used metal in the world, rusting is the most common type of corrosion.
The Chemistry of Rusting
Rusting of the iron is due to the electrochemical reaction. It requires the presence of both air (i.e. the oxygen) and water. Like the electrochemical cell, electrons given up from anode of an iron atom flow (through metal) to cathode. Consider an iron sheet exposed to open air:
Figure 2. The electrochemical process of rusting on a flat iece of iron
The metal gets wet due to moisture in air. The thin water film dissolves oxygen from air.
In the initial stage of rusting, some iron atoms lose electrons to become Fe2+ (aq) ions:
Fe (s) ? Fe2+ + 2e- (oxidation at anodic area)
The electrons are accepted by the dissolved oxygen and water to form OH- (aq) ion:
O2 (g) + 2H2O (l) + 4e- ? 4OH- (aq) (reduction at cathodic area)
Thus a simple electrochemical cell is set up. One part of the iron piece acts as anode; another part acts as cathode. The cathodic area is usually the region around the outer edges of the water film, where the concentration of dissolved oxygen is higher.
The Fe2+ (aq) and OH- ions in the water film react to form iron (II) hydroxide.
Fe2+ (aq) + 2OH- ? Fe(OH)2 (s)
The precipitate is rapidly oxidized by dissolved oxygen to form iron (III) hydroxide.
4Fe(OH)2 (s) + O2 (aq) + 2H2O (l) ? 4Fe(OH)3 (s)
On standing, this changes to rust, a reddish brown solid. Rust is in fact hydrated iron (III) oxide with variable composition (Fe2O3 ?nH2O).
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. For example, the iron objects by seashore, “where seawater has a high salt content, amounting to about 3.6% in the Atlantic and Pacific Oceans.” The thin water film on iron surface contains dissolved sodium chloride from sea spray. This greatly increases conductivity of the solution, due to a higher concentration of ions. As a result, a special “seawater rust” is formed which is actually very strongly corrosion inducing. One would expect corrosion rates of about 0.1mm/year.
Presence of soluble salts other than sodium chloride may also assist rusting.
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. However, the relative air humidity in enclosed spaces often differs from that existing in the open air; in winter, in heated room it is lower, while in summer it can be higher in cool cellars. On the whole the danger of corrosion in inside rooms is less than in the open air.
“Many water have lime and carbonic acid in equilibrium. This is called equilibrium water, where there is sufficient carbon dioxide in solution to stabilize the carbonate. The equilibrium can be expressed as follow:
CaCO3 + H2CO3 Ca(HCO3)2
Provided the minimum hardness is about 2.2 milliequivalents/liter, these water form layers of mixed lime and rust that safeguard the steel piping against further corrosion.” If the water contains an excess of carbonic acid, which prevents the formation of protective layers, there is a danger of corrosion of unprotected steel in the presence of oxygen.
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 spent every year to protect iron objects and replace rusted articles.
Several methods can be used to prevent rusting or to slow it down.
Applying a Protective Layer
Both air and water are necessary for rusting to occur. Any method which can keep out one or both of them from iron will prevent rusting. The most obvious way is to apply a protective layer.
1) Coating with paint, oil or grease
Objects that are unlikely to be scratched can be coated with paint (or lacquer, or enamel). For example, bridges, ships and car bodies are painted.
Moving parts of a machine are protected by applying oil or grease.
2) Coating with another metal
Iron can be coated with a thin layer of another metal which is resistant to corrosion. Galvanized iron is iron coated with zinc. Some roofs, buckets and dustbins are made from galvanized iron. Tin -plate is iron coated with tin.
3) Using Alloys Of Iron
Steel is produced form iron by carefully controlling the amount of carbon present (0 – 1%). To fight against corrosion, steel can be alloyed with other metals such as chromium and nickel to produce stainless steels.
Cathodic Protection
Rusting is a redox reaction in which iron loses electrons. If iron is connected to a more reactive metal, the other metal will lose electrons in preference, preventing the formation of Fe2+ (aq) ions.
Galvanizing (zinc-plating) provides a good example of cathodic protection. “When the zinc coating is undamaged, the iron is covered up and is protected from rusting. In case the coating is partly damaged, the iron, though exposed, is still protected.” Zinc, being more reactive than iron, will form zinc ions:
At zinc (anode) Zn (s) ? Zn2+ (aq) + 2e-
The electrons flow from zinc to iron, where the following half-reaction take place:
At iron (cathode) O2 (g) + 2H2O (l) + 4e- ? 4 OH- (aq)
Since iron is forced to accept the electrons, it cannot corrode, as corrosion involves giving off electrons. This method of preventing corrosion is called cathodic protection.
Cathodic protection is also used to prevent corrosion in underground iron pipelines. Bags containing magnesium turnings are connected to the buried pipelines at intervals. (Figure 3.) The magnesium corrodes instead of the iron. The magnesium should therefore be replaced from time to time.
Figure 3. Protecting underground iron pipelines from corrosion by cathodic protection
Most ships are made of iron. To prevent rusting, zinc blocks are attached to the hull of a ship. Zinc will corrode in preference to iron.
Electrical Protection
Sometimes rusting can be prevented by using electricity. For example, the negative terminal of the car battery is always connected to the car body. This supplies electrons to the iron body, preventing it from losing electrons. In some piers, the steel structures are protected electrically by connecting them to the negative terminal of a d.c. source.
Conclusion
Corrosion is the gradual deterioration of a metal due to reaction with air, water or other substances in the surroundings. When these substances present, the metal can corrode through the process of electrochemical and chemical corrosion reaction. The most common type of corrosion is rusting which costs several billion dollors a year. To prevent such natural, spontaneous processes, method of protections and constant maintenance work are necessary; only in these ways can a steel structure be adequately protected against corrosion and its value maintained.
