.. 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.