Sulfuric Acid

I

Introduction

Sulfuric Acid, H₂SO₄, corrosive, oily, colorless liquid, with a specific gravity 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 sulfuric 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 (see Acids and Bases). 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, sulfuric acid has been commercially important for many years. The early alchemists prepared it in large quantities by heating naturally occurring sulfates to a high temperature and dissolving in water the sulfur trioxide thus formed. About the 15th century a method was developed for obtaining the acid by distilling hydrated ferrous sulfate, or iron vitriol, with sand. In 1740 the acid was produced successfully on a commercial scale by burning sulfur and potassium nitrate in a ladle suspended in a large glass globe partially filled with water.

II

Properties

Sulfuric acid is a strong acid, that is, in aqueous solution it is largely changed to hydrogen ions (H⁺) and sulfate ions (SO₄^2-). Each molecule gives two H⁺ ions, thus sulfuric acid is dibasic. Dilute solutions of sulfuric acid show all the behavior characteristics of acids. They taste sour, conduct electricity, neutralize alkalies, and corrode active metals with formation of hydrogen gas. From sulfuric acid one can prepare both normal salts (see Salt) containing the sulfate group, SO₄, and acid salts containing the hydrogen sulfate group, HSO₄.

Concentrated sulfuric 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 sulfuric acid is heated, it behaves also as an oxidizing agent, capable, for example, of dissolving such relatively unreactive metals as copper, mercury, and lead to produce metal sulfate, sulfur dioxide, and water.

During the 19th century, the German chemist Baron Justus von Liebig discovered that sulfuric 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 sulfuric acid and led to improved methods of manufacture.

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III

Manufacture

Two processes for the production of sulfuric acid are in use today. In their initial steps, both require the use of sulfur dioxide, which is produced by burning iron pyrites, FeS₂, or sulfur, in air. The first of these methods, the lead-chamber process, employs as reaction vessels large lead-sheathed brick towers. In these towers, sulfur-dioxide gas, air, steam, and oxides of nitrogen react to yield sulfuric 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. Sulfuric acid produced in this way, and labeled acid, is only about 62 to 70 percent H₂SO₄. The rest is water. About 20 percent of all sulfuric acid is now made by the lead-chamber process, but that percentage is diminishing.

The second method of manufacturing sulfuric acid, the contact process, which came into commercial use about 1900, depends on oxidation of sulfur dioxide to sulfur trioxide, SO₃, under the accelerating influence of a catalyst (see Catalysis). Finely divided platinum, the most effective catalyst, has two disadvantages: It is very expensive, and it is vitiated by certain impurities in ordinary sulfur dioxide that reduce its activity. Many sulfuric-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 sulfur dioxide to trioxide is nearly complete. The trioxide is dissolved in concentrated sulfuric acid, and at the same time a regulated influx of water maintains the concentration at a selected level usually about 95 percent. By reducing the flow of water, a product with more SO₃ than shown in the formula H₂SO₄ may be made. This product, called fuming sulfuric acid, or oleum, or Nordhausen acid, is needed in some organic chemical reactions.

IV

Production

The uses of sulfuric acid are so varied that the volume of its production provides an approximate index of general industrial activity. American production of sulfuric acid exceeded 29 million tons annually in the early 1970s, a figure corresponding to a daily production of 1/3 kg (3/4 lb) per person throughout the year. The largest single use of sulfuric acid is for making fertilizers, both superphosphate and ammonium sulfate. 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 sulfuric acid as such is the lead storage battery, or car energizer.