Sulphuric Acid
Sulphuric Acid, H2SO4, is a corrosive, oily, colourless liquid, with a relative density of 1.85. It melts at 10.36° C (50.6° F), boils at 340° C (644° F), and is soluble in all proportions in water. When sulphuric acid is mixed with water, considerable heat is released. Unless the mixture is well stirred, the added water may be heated beyond its boiling point and the sudden formation of steam may blow the acid out of its container . The concentrated acid destroys skin and flesh, and can cause blindness if it gets into the eyes. The best treatment is to flush away the acid with large amounts of water. Despite the dangers created by careless handling, sulphuric acid has been commercially important for many years. The early alchemists prepared it in large quantities by heating naturally occurring sulphates to a high temperature and dissolving in water the sulphur trioxide thus formed. About the 15th century a method was developed for obtaining the acid by distilling hydrated ferrous sulphate, or iron vitriol, with sand. In 1740 the acid was produced successfully on a commercial scale by burning sulphur and potassium nitrate in a ladle suspended in a large glass globe partially filled with water.
Sulphuric acid is a strong acid, that is, in aqueous solution it is largely changed to hydrogen ions (H+) and sulphate ions (SO42-). Each molecule gives two H+ ions, thus sulphuric acid is dibasic. Dilute solutions of sulphuric acid show all the behaviour characteristics of acids. They taste sour, conduct electricity, neutralize alkalis, and corrode active metals with formation of hydrogen gas. From sulphuric acid one can prepare both normal salts containing the sulphate group, SO4, and acid salts containing the hydrogen sulphate group, HSO4.
Concentrated sulphuric acid, formerly called oil of vitriol, is a valuable desiccating ...
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Sulphuric acid is a strong acid, that is, in aqueous solution it is largely changed to hydrogen ions (H+) and sulphate ions (SO42-). Each molecule gives two H+ ions, thus sulphuric acid is dibasic. Dilute solutions of sulphuric acid show all the behaviour characteristics of acids. They taste sour, conduct electricity, neutralize alkalis, and corrode active metals with formation of hydrogen gas. From sulphuric acid one can prepare both normal salts containing the sulphate group, SO4, and acid salts containing the hydrogen sulphate group, HSO4.
Concentrated sulphuric acid, formerly called oil of vitriol, is a valuable desiccating agent. It acts so vigorously in this respect that it removes water from, and therefore chars, wood, cotton, sugar, and paper. It is used in the manufacture of ether, nitroglycerine, and dyes for its property as a desiccant. When concentrated sulphuric acid is heated, it behaves as an oxidizing agent, capable, for example, of dissolving such relatively unreactive metals as copper, mercury, and lead to produce metal sulphate, sulphur dioxide, and water.
During the 19th century, the German chemist Baron Justus von Liebig discovered that sulphuric acid, when added to the soil, increased the amount of soil phosphorus available to plants. This discovery gave rise to an increase in the commercial production of sulphuric acid and led to improved methods of manufacture.
Two processes for the production of sulphuric acid are in use today. In their initial steps, both require the use of sulphur dioxide, which is produced by burning iron pyrites, FeS2, or sulphur, in air. The first of these methods, the lead-chamber process, employs as reaction vessels large lead-sheathed brick towers. In these towers, sulphur-dioxide gas, air, steam, and oxides of nitrogen react to yield sulphuric acid as fine droplets that fall to the bottom of the chamber. Almost all the nitrogen oxides are recovered from the outflowing gas and are brought back to the chamber to be used again. Sulphuric acid produced in this way, and labelled acid, is only about 62 to 70 per cent H2SO4. The rest is water. About 20 per cent of all sulphuric acid is now made by the lead-chamber process, but that percentage is diminishing.
The second method of manufacturing sulphuric acid, the contact process, which came into commercial use about 1900, depends on oxidation of sulphur dioxide to sulphur trioxide, SO3, under the accelerating influence of a catalyst. Finely divided platinum, the most effective catalyst, has two disadvantages: it is very expensive, and it is vitiated by certain impurities in ordinary sulphur dioxide that reduce its activity. Many sulphuric-acid producers use two catalysts in tandem; first, a more rugged but less effective one like iron oxide or vanadium oxide to bring about the bulk reaction; then, a smaller amount of platinum to finish the job. At 400° C (752° F), the conversion of sulphur dioxide to trioxide is nearly complete. The trioxide is dissolved in concentrated sulphuric acid, and at the same time a regulated influx of water maintains the concentration at a selected level, usually about 95 per cent. By reducing the flow of water, a product with more SO3 than shown in the formula H2SO4 may be made. This product, called fuming sulphuric acid, or oleum, or Nordhausen acid, is needed in some organic chemical reactions.
The uses of sulphuric acid are so varied that the volume of its production provides an approximate index of general industrial activity. For example, in the early 1970s, annual United States production of sulphuric acid exceeded 29 million tons, a figure corresponding to a daily production of 1/3 kg (3/4 lb) per person throughout the year. The largest single use of sulphuric acid is for making fertilizers, both superphosphate and ammonium sulphate. It is also used in making organic products, refining petroleum, making paints and pigments, processing metals, and making rayon. One of the few consumer products containing sulphuric acid as such is the lead storage battery, found in cars.
Tim Rogers 11F
