sulfuric acid

sulfuric acid (H₂SO₄), colorless, odorless, oily, and corrosive liquid that is a widely manufactured industrial chemical. A key raw material for making fertilizers and many other chemical products, it is produced in large quantities and used in petroleum refining, metal processing, and chemical manufacturing.

History

The origin of sulfuric acid is uncertain, but references to its preparation appear before the 10th century. In the late 15th century German alchemist Basil Valentine described methods of obtaining the acid, including burning sulfur with saltpeter (potassium nitrate) and distilling it from a mixture of silica and ferric sulfate, then called vitriol—an association that gave rise to the long-used name oil of vitriol.

Until the 18th century, production was small in scale and limited mainly to the preparation of nitric and hydrochloric acids for assaying and treating nonferrous metals. Large-scale manufacture began in 1746, when English physician John Roebuck developed the lead-chamber process, in which gases from burning sulfur were absorbed in water within lead-lined chambers. This method enabled output far greater than that possible with earlier clay or glass vessels. The soda ash production process introduced by the French chemist Nicolas Leblanc in 1790 further spurred demand. Although his first factory failed amid the turmoil of the French Revolution, the process was more widely adopted after 1807. Because the first step of Leblanc’s process required sulfuric acid, which was difficult to transport, alkali producers usually built their own acid plants.

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Further advances in the 19th century improved the lead-chamber system. In 1827 French chemist Joseph-Louis Gay-Lussac designed a lead tower that recovered nitrogen oxides, allowing producers to recycle them instead of constantly adding new saltpeter. The first tower was installed in France in 1837, but the design was not widely used until after English industrial chemist John Glover patented another tower in 1859 that concentrated the acid and recovered even more nitrogen oxides. During the same period producers began turning to pyrite ores (iron disulfide minerals) as a new source of sulfur dioxide. For much of the century most sulfur had come from Sicily, where a monopoly kept prices high. Roasting pyrites provided a cheaper alternative and had the added advantage of yielding iron or copper after the sulfur was removed.

The lead-chamber process was limited to acid of about 78 percent concentration. The need for stronger acid, especially for dye manufacture, led to adoption of the contact process, patented in 1831 by English merchant Peregrine Phillips. In this method sulfur dioxide is catalytically converted to sulfur trioxide and then absorbed in water to yield concentrated acid. The process spread widely in the late 19th century, particularly in Germany, and in the early 20th century platinum was replaced by vanadium catalysts. By the mid-20th century it had largely replaced the lead-chamber process, and it remains the universal method of production.

Molecular and structural characteristics

Sulfuric acid is built around a central sulfur atom bonded to four oxygen atoms in a tetrahedral arrangement (a shape with four corners like a pyramid with a triangular base). Two of these oxygen atoms are linked by double bonds, while the other two atoms are joined through hydroxyl groups (―OH), giving the formula H₂SO₄.

The sulfur atom is in the +6 oxidation state, meaning it has effectively given up six electrons to the surrounding oxygen atoms. This high oxidation state, combined with the sharing of electrons among the oxygen atoms in a pattern known as resonance, lends the molecule both stability and a strong acidic character.

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When sulfuric acid dissociates in water, it releases hydrogen ions in two steps. The first ion is given up completely, forming the hydrogen sulfate ion (HSO₄⁻). The second ion is only partly released, producing sulfate ions (SO₄²⁻). Because it can donate two hydrogen ions per molecule in sequence, sulfuric acid is known as a diprotic acid.

Properties

Physical properties

Pure sulfuric acid has a specific gravity of 1.83 at 25 °C (77 °F), which indicates that it is 1.83 times as dense as water. It freezes at 10.37 °C (50.7 °F). When heated, the pure acid partially decomposes into water and sulfur trioxide. Sulfur trioxide escapes as a vapor until the concentration of the acid falls to 98.3 percent. This mixture of sulfuric acid and water boils at a constant temperature of 338 °C (640 °F) at one atmosphere of pressure. Sulfuric acid is commonly supplied at concentrations of 78, 93, or 98 percent.

Concentrated sulfuric acid is highly hygroscopic: It readily draws moisture from the air. When exposed for long periods, the liquid gradually absorbs water vapor and becomes more dilute, causing its volume to increase. Volcanic activity can result in the production of sulfuric acid, depending on the emissions associated with specific volcanoes, and sulfuric acid aerosols from an eruption can persist in the stratosphere for many years. These aerosols can then reform into sulfur dioxide (SO₂), a constituent of acid rain, though volcanic activity is a relatively minor contributor to acid rainfall.

Uses and applications

Sulfuric acid is a widely produced chemical. Its uses extend across agriculture, energy, metallurgy, and chemical synthesis.

• Fertilizers: Sulfuric acid is used to produce phosphate fertilizers. Sulfuric acid reacts with phosphate rock to yield superphosphate and other fertilizer products essential for modern agriculture.
• Petroleum and chemical processing: In petroleum refining the acid is employed in the alkylation process, which helps improve the octane rating of gasoline. It is also used to remove impurities from fuels and lubricants. In the chemical industry sulfuric acid serves as a raw material for producing detergents, synthetic resins, dyes, explosives, and numerous inorganic compounds.
• Metallurgy: In metallurgy sulfuric acid is used in pickling, a process that removes oxidation and scale from steel and other metals before further processing, such as galvanizing or cold-rolling.
• Energy storage: Dilute sulfuric acid functions as the electrolyte in lead-acid batteries, which are widely used in automobiles and backup power systems.
• Laboratory and specialized uses: Because of its strong acidity and dehydrating power, sulfuric acid is a common laboratory reagent. A related form, fuming sulfuric acid or oleum (sulfur trioxide dissolved in concentrated sulfuric acid), is employed in the production of dyes and organic chemicals that require stronger conditions than the standard acid can provide.