sulfuric acid

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sulfuric acid,

chemical compound, H2SO4, colorless, odorless, extremely corrosive, oily liquid. It is sometimes called oil of vitriol.

Concentrated Sulfuric Acid

When heated, the pure 100% acid loses sulfur trioxide gas, SO3, until a constant-boiling solution, or azeotrope, containing about 98.5% H2SO4 is formed at 337°C;. Concentrated sulfuric acid is a weak acid (see acids and basesacids and bases,
two related classes of chemicals; the members of each class have a number of common properties when dissolved in a solvent, usually water. Properties
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) and a poor electrolyteelectrolyte
, electrical conductor in which current is carried by ions rather than by free electrons (as in a metal). Electrolytes include water solutions of acids, bases, or salts; certain pure liquids; and molten salts.
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 because relatively little of it is dissociated into ions at room temperature. When cold it does not react readily with such common metals as iron or copper. When hot it is an oxidizing agent, the sulfur in it being reduced; sulfur dioxide gas may be released. Hot concentrated sulfuric acid reacts with most metals and with several nonmetals, e.g., sulfur and carbon. Because the concentrated acid has a fairly high boiling point, it can be used to release more volatile acids from their salts, e.g., when sodium chloride (NaCl), or common salt, is heated with concentrated sulfuric acid, hydrogen chloride gas, HCl, is evolved.

Concentrated sulfuric acid has a very strong affinity for water. It is sometimes used as a drying agent and can be used to dehydrate (chemically remove water from) many compounds, e.g., carbohydrates. It reacts with the sugar sucrose, C12H22O11, removing eleven molecules of water, H2O, from each molecule of sucrose and leaving a brittle spongy black mass of carbon and diluted sulfuric acid. The acid reacts similarly with skin, cellulose, and other plant and animal matter.

When the concentrated acid mixes with water, large amounts of heat are released; enough heat can be released at once to boil the water and spatter the acid. To dilute the acid, the acid should be added slowly to cold water with constant stirring to limit the buildup of heat. Sulfuric acid reacts with water to form hydrates with distinct properties.

Dilute Sulfuric Acid

Dilute sulfuric acid is a strong acid and a good electrolyte; it is highly ionized, much of the heat released in dilution coming from hydration of the hydrogen ionsion,
atom or group of atoms having a net electric charge. Positive and Negative Electric Charges

A neutral atom or group of atoms becomes an ion by gaining or losing one or more electrons or protons.
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. The dilute acid has most of the properties of common strong acids. It turns blue litmus red. It reacts with many metals (e.g., with zinc), releasing hydrogen gas, H2, and forming the sulfatesulfate,
chemical compound containing the sulfate (SO4) radical. Sulfates are salts or esters of sulfuric acid, H2SO4, formed by replacing one or both of the hydrogens with a metal (e.g., sodium) or a radical (e.g., ammonium or ethyl).
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 of the metal. It reacts with most hydroxides and oxides, with some carbonates and sulfides, and with some salts. Since it is dibasic (i.e., it has two replaceable hydrogen atoms in each molecule), it forms both normal sulfates (with both hydrogens replaced, e.g., sodium sulfate, Na2SO4) and acid sulfates, also called bisulfates or hydrogen sulfates (with only one hydrogen replaced, e.g., sodium bisulfate, NaHSO4).

Production of Sulfuric Acid

There are two major processes (lead chamber and contact) for production of sulfuric acid, and it is available commercially in a number of grades and concentrations. The lead chamber process, the older of the two processes, is used to produce much of the acid used to make fertilizers; it produces a relatively dilute acid (62%–78% H2SO4). The contact process produces a purer, more concentrated acid but requires purer raw materials and the use of expensive catalysts. In both processes sulfur dioxidesulfur dioxide,
chemical compound, SO2, a colorless gas with a pungent, suffocating odor. It is readily soluble in cold water, sparingly soluble in hot water, and soluble in alcohol, acetic acid, and sulfuric acid.
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 is oxidized and dissolved in water. The sulfur dioxide is obtained by burning sulfur, by burning pyrites (iron sulfides), by roasting nonferrous sulfide ores preparatory to smelting, or by burning hydrogen sulfide gas. Some sulfuric acid is also made from ferrous sulfate waste solutions from pickling iron and steel and from waste acid sludge from oil refineries.

Lean Chamber Process

In the lead chamber process hot sulfur dioxide gas enters the bottom of a reactor called a Glover tower where it is washed with nitrous vitriol (sulfuric acid with nitric oxide, NO, and nitrogen dioxide, NO2, dissolved in it) and mixed with nitric oxide and nitrogen dioxide gases; some of the sulfur dioxide is oxidized to sulfur trioxide and dissolved in the acid wash to form tower acid or Glover acid (about 78% H2SO4). From the Glover tower a mixture of gases (including sulfur dioxide and trioxide, nitrogen oxides, nitrogen, oxygen, and steam) is transferred to a lead-lined chamber where it is reacted with more water. The chamber may be a large, boxlike room or an enclosure in the form of a truncated cone. Sulfuric acid is formed by a complex series of reactions; it condenses on the walls and collects on the floor of the chamber. There may be from three to twelve chambers in a series; the gases pass through each in succession. The acid produced in the chambers, often called chamber acid or fertilizer acid, contains 62% to 68% H2SO4. After the gases have passed through the chambers they are passed into a reactor called the Gay-Lussac tower where they are washed with cooled concentrated acid (from the Glover tower); the nitrogen oxides and unreacted sulfur dioxide dissolve in the acid to form the nitrous vitriol used in the Glover tower. Remaining waste gases are usually discharged into the atmosphere.

Contact Process

In the contact process, purified sulfur dioxide and air are mixed, heated to about 450°C;, and passed over a catalyst; the sulfur dioxide is oxidized to sulfur trioxide. The catalyst is usually platinum on a silica or asbestos carrier or vanadium pentoxide on a silica carrier. The sulfur trioxide is cooled and passed through two towers. In the first tower it is washed with oleum (fuming sulfuric acid, 100% sulfuric acid with sulfur trioxide dissolved in it). In the second tower it is washed with 97% sulfuric acid; 98% sulfuric acid is usually produced in this tower. Waste gases are usually discharged into the atmosphere. Acid of any desired concentration may be produced by mixing or diluting the products of this process.

Uses of Sulfuric Acid

Sulfuric acid is one of the most important industrial chemicals. More of it is made each year than is made of any other manufactured chemical; more than 40 million tons of it were produced in the United States in 1990. It has widely varied uses and plays some part in the production of nearly all manufactured goods. The major use of sulfuric acid is in the production of fertilizers, e.g., superphosphate of lime and ammonium sulfate. It is widely used in the manufacture of chemicals, e.g., in making hydrochloric acid, nitric acid, sulfate salts, synthetic detergents, dyes and pigments, explosives, and drugs. It is used in petroleum refining to wash impurities out of gasoline and other refinery products. Sulfuric acid is used in processing metals, e.g., in pickling (cleaning) iron and steel before plating them with tin or zinc. Rayon is made with sulfuric acid. It serves as the electrolyte in the lead-acid storage battery commonly used in motor vehicles (acid for this use, containing about 33% H2SO4 and with specific gravity about 1.25, is often called battery acid).

History of Sulfuric Acid

Although sulfuric acid is now one of the most widely used chemicals, it was probably little known before the 16th cent. It was prepared by Johann Van Helmont (c.1600) by destructive distillation of green vitriol (ferrous sulfate) and by burning sulfur. The first major industrial demand for sulfuric acid was the Leblanc process for making sodium carbonate (developed c.1790). Sulfuric acid was produced at Nordhausen from green vitriol but was expensive. A process for its synthesis by burning sulfur with saltpeter (potassium nitrate) was first used by Johann Glauber in the 17th cent. and developed commercially by Joshua Ward in England c.1740. It was soon superseded by the lead chamber process, invented by John Roebuck in 1746 and since improved by many others. The contact process was originally developed c.1830 by Peregrine Phillips in England; it was little used until a need for concentrated acid arose, particularly for the manufacture of synthetic organic dyes.

Sulfuric Acid

 

H2 SO4, a strong dibasic acid corresponding to the highest oxidation state of sulfur (+ 6). Under usual conditions, sulfuric acid is a heavy, oily, colorless, and odorless liquid. In industry, mixtures of sulfuric acid both with water and sulfur trioxide are also called sulfuric acid. If the SO3: H2 O molecular ratio is less than 1, the mixture is an aqueous solution of sulfuric acid; if it is more than 1, the mixture is a solution of SO3 in sulfuric acid.

Physical and chemical properties. When existing in a concentration of 100 percent, sulfuric acid (monohydrate, SO3 · H2 O) crystallizes at 10.45°C, boils at 296.2°C, and has a density of 1.9203 g/cm3 and a heat capacity of 1.62joules/g•°K. H2 SO4 is miscible with water and SO3 in all proportions, forming the compounds

H2 SO4·4H2 O (melting point [mp] - 28.36°C)

H2 SO4·3H2 O (mp -36.31 °C)

H2 SO4·2H2 O (mp -39.60°C)

H2 SO4·H2 O (mp 8.48°C)

The compounds H2 SO4·SO3 (H2 S2 O7, disulfuric, or pyrosulfuric acid, mp 35.15°C) and H2 SO4·2SO3 (H2 S3 O10, trisulfuric acid, mp 1.20°C) are also formed. Only water vapor is given off into the vapor phase upon heating and boiling aqueous solutions of sulfuric acid containing up to 70 percent H2 S04. Vapors of sulfuric acid are formed above more concentrated solutions. A solution of 98.3 percent H2 SO4 (azeotropic mixture) upon boiling (336.5°C) is completely distilled. Sulfuric acid containing more than 98.3 percent H2 SO4 releases vapors of SO3 upon heating.

Concentrated sulfuric acid is a strong oxidizing agent. It oxidizes HI and HBr to the free halogens, and upon heating, it will oxidize all metals except Au and the platinum metals (with the exception of Pd). In the cold, concentrated sulfuric acid passi-vates many metals, including Pb, Cr, Ni, steel, and pig iron. Dilute sulfuric acid reacts with all the metals (except Pb) that are above hydrogen in the electromotive force series, for example,

Zn + H2 SO4 = ZnSO4 + H2

As a strong acid, sulfuric acid displaces weaker acids from their salts, for example, boric acid from borax:

Na2 B4 O7 + H2 SO4 + 5H2 O = Na2 SO4 + 4H3 BO3

Upon heating, it displaces more volatile acids, for example,

NaNO3 + H2 SO4 = NaHSO4 + HNO3

Sulfuric acid removes water that is chemically bound to organic compounds containing OH, or hydroxyl, groups. The dehydration of ethyl alcohol in the presence of concentrated sulfuric acid results in the formation of ethylene or diethyl ether. The charring of sugar, cellulose, starch, and other carbohydrates upon contact with sulfuric acid also derives from the dehydration of these substances. As a dibasic acid, sulfuric acid forms two types of salts: sulfates and bisulfates.

Production. The first descriptions of oil of vitriol, that is, concentrated sulfuric acid, were given by the Italian scientist V. Biringuccio in 1540 and the German alchemist whose works were published under the name of Basilius Valentinus in the late 16th and early 17th centuries. By 1690, the French chemists N. Lemery and N. Lefebvre had laid the basis for the first industrial method of obtaining sulfuric acid, a method applied in England in 1740. According to this method, a mixture of sulfur and saltpeter was burned in a ladle suspended in a glass jar containing a certain amount of water. The SO3 generated reacted with the water to form sulfuric acid. In 1746, J. Roebuck in Birmingham replaced the glass jars with chambers made of sheet lead, thus laying the basis for the chamber process for the production of sulfuric acid. Continuous improvement in the process for the production of sulfuric acid in Great Britain and France resulted in 1908 in the first tower system. In the USSR, the first tower installation went into operation in 1926 at the Po-levskoi Metallurgical Plant in the Urals.

The raw material for the production of sulfuric acid can be sulfur, pyrite (FeS2), or exhaust gases containing SO2 from furnaces for the oxidative roasting of the sulfide ores of Cu, Pb, Zn, and other metals. In the USSR, most sulfuric acid is obtained from pyrite. Here, the FeS2 is burned in furnaces in the state of a fluidized bed, a state achieved by blowing a rapid stream of air through a layer of finely ground pyrite. The gaseous mixture obtained contains SO2, O2, N2, impurities of SO3, and vapors of H2 O, As2 O3, and SiO2 and holds considerable cinder dust, which is removed from the gas in electrostatic precipitators.

Sulfuric acid is produced from SO2 by the nitrous (tower) and contact methods. The conversion of SO2 into sulfuric acid by the nitrous method is carried out in cylindrical production towers 15 m and more in height and filled with a packing of ceramic rings. Nitrous vitriol—a mixture of dilute sulfuric acid and ni-trosylsulfuric acid (NOOSO3 H) is sprayed from above into a rising stream of gases. The nitrosylsulfuric acid is produced by the reaction

N2 O3 + 2H2 SO4 = 2NOOSO3 H + H2 O

The oxidation of SO2 by nitrogen oxides occurs in solution after the absorption of SO2 by nitrous vitriol. The mixture is hydro-lyzed by water:

NOOSO3 H + H2 O = H2 SO4 + HNO2

The sulfur dioxide gas entering the tower reacts with water to form sulfurous acid:

SO2 + H2 O = H2 SO3

The reaction of HNO2 with H2 SO3 results in the formation of sulfuric acid:

2HNO2 + H2 SO3 = H2 SO4 + 2NO + H2 O

The NO produced is converted into N2 O3, or, more precisely, a mixture of NO and NO2, in the oxidizing tower. The gases are then introduced into absorption towers, where they encounter sulfuric acid supplied from the top. It is here that nitrous vitriol is obtained; the mixture is then transferred to the production towers. Thus, there is continuous production and a circulation of nitrogen oxides. The inevitable losses of nitrogen oxides with the exhaust gases are balanced by the addition of HNO3.

The sulfuric acid produced by the nitrous method is of insufficient concentration and contains harmful impurities, for example. As. Production is accompanied by the release of nitrogen oxides into the atmosphere (“foxtails,” named for the color of NO2).

The principle of the contact process for the production of sulfuric acid was discovered in 1831 by P. Phillips in Great Britain. The first catalyst was platinum. In the late 19th century and early 20th, it was found that the oxidation of SO2 to SO3 could be accelerated by vanadium pentoxide (V2 O5). The studies of the Soviet scientists A. E. Adadurov, G. K. Boreskov, and F. N. Iushkevich were important in determining the action and selection of vanadium catalysts. Modern sulfuric acid plants are constructed for the operation of the contact process. Vanadium oxides with additives of SiO2, AI2 O3, K2 O, CaO, and BaO in various proportions are used as catalyst bases. All vanadium contact materials are reactive only at temperatures above ~420°C. In the contact apparatus, the gas usually passes through four or five tiers of contact material. In the production of sulfuric acid by the contact process, the gas undergoing oxidation is first subjected to removal of impurities that might poison the catalyst. As, Se, and traces of dust are removed in scrubbing towers in which there is a countercurrent trickling of sulfuric acid. The sulfuric acid mist formed from the SO3 and H2 O present in the gaseous mixture is eliminated in wet electrostatic precipitators. Vapors of H2 O are absorbed by concentrated sulfuric acid in drying towers. The mixture of SO2 and air then passes through the catalyst (contact material) and undergoes oxidation to SO3:

The sulfur trioxide is then absorbed by the water in dilute H2 SO4:

SO3 + H2 O = H2 SO4

Depending on the amount of water introduced into the process, either oleum or a solution of sulfuric acid in water is obtained.

In 1973 the production of sulfuric acid (in the monohydrate) was (in millions of tons): 14.9 in the USSR, 28.7 in the United States, 7.1 in Japan, 5.5 in the Federal Republic of Germany, 4.4 in France, 3.9 in Great Britain, 3.0 in Italy, 2.9 in Poland, 1.2 in Czechoslovakia, 1.1 in the German Democratic Republic, and 0.9 in Yugoslavia.

USE. Sulfuric acid is one of the most important products of the heavy chemical industry. The available grades include chamber acid (not less than 75 percent H2 SO4), oil of vitriol (not less than 92.5 percent), and oleum, or fuming sulfuric acid (a solution of 18.5–20 percent SO3 in H2 SO4), as well as especially pure battery acid (92–94 percent; when diluted by water to 26–31 percent, it serves as the electrolyte in lead batteries). In addition, reagent-grade sulfuric acid (92–94 percent) is produced by the contact process in quartz or platinum apparatus. The strength of sulfuric acid is determined by the density, which is measured with a hydrometer. Most of the chamber acid is used in the production of mineral fertilizers. Sulfuric acid is used in the production of, for example, phosphoric, hydrochloric, boric, and hydrofluoric acids because of its ability to displace these acids from their salts. Concentrated sulfuric acid is used in separating organosulfur compounds and unsaturated organic compounds from petroleum products. Dilute sulfuric acid is used for the removal of scale from wire and sheets before plating with tin or zinc and for the pickling of metal surfaces before plating with chromium, nickel, or copper. It is used in metallurgy for the decomposition of complex ores, in particular, those of uranium. In organic synthesis, concentrated sulfuric acid is a necessary component of nitrating mixtures and a sulfonating agent in the production of many dyes and pharmaceuticals. Owing to its high hygroscopicity, sulfuric acid is used in drying gases and in concentrating nitric acid.

Safety measures. Poisonous gases—SO2 and NO2—as well as vapors of SO3 and H2 SO4, present a danger in the production of sulfuric acid. Proper ventilation and hermetically sealed production apparatus are therefore mandatory. Since sulfuric acid causes serious burns of the skin, handling requires extreme care and protective devices (goggles, rubber gloves, aprons, boots). When diluting sulfuric acid, the acid must be poured into water in a thin stream while stirring. Pouring the water into the acid leads to spattering because of the evolution of a great amount of heat.

REFERENCES

Spravochnik sernokislotchika, 2nd ed. Edited by K. M. Malin. Moscow, 1971.
Malin, K. M., N. L. Arkin, G. K. Boreskov, and M. G. Slin’ko. Tekhnologiia sernoi kisloty. Moscow, 1950.
Boreskov, G. K. Kataliz v proizvodstve sernoi kisloty. Moscow-Leningrad, 1954.
Amelin, A. G., and E. V. Iashke. Proizvodstvo sernoi kisloty. Moscow, 1974.
Luk’ianov, P. M. Kratkaia istoriia khimicheskoi promyshlennosti SSSR. Moscow, 1959.

I. K. MALINA

sulfuric acid

[¦səl¦fyu̇r·ik ′as·əd]
(inorganic chemistry)
H2SO4 A toxic, corrosive, strongly acid, colorless liquid that is miscible with water and dissolves most metals, and melts at 10°C; used in industry in the manufacture of chemicals, fertilizers, and explosives, and in petroleum refining. Also known as dipping acid; oil of vitriol, vitriolic acid.