Redirected from Sulfuric acid
| General | |
|---|---|
| Name | Sulfuric acid |
| Chemical formula | H2SO4 |
| Appearance | Colourless liquid |
| Physical | |
| Formula weight | 98.1 amu |
| Melting point | 283 K (10 °C) |
| Boiling point | 610 K (337 °C) |
| Density | 1.8 ×103 kg/m3 |
| Solubility | miscible |
| Thermochemistry | |
| ΔfH0liquid | -814 kJ/mol |
| S0liquid, 1 bar | 19 J/mol·K |
| Safety | |
| Ingestion | Severe and permanent damage may result. |
| Inhalation | Very dangerous, possibly fatal. Long-term effects known. |
| Skin | Causes burns. |
| Eyes | Causes burns. |
| More info | Hazardous Chemical Database (http://ull.chemistry.uakron.edu/erd/chemicals/8/7116.html) |
| SI units were used where possible. Unless otherwise stated, standard conditions were used. | |
Sulphuric acid has many applications, including in many chemical reactions and production processes. It is the most widely used chemical. Principal uses include fertilizer manufacturing, ore processing, chemical synthesis, wastewater processing and oil refining.
In combination with nitric acid it forms the nitronium ion, which is used in the nitration[?] of compounds. The process of nitration is used to manufacture a great many explosives, including trinitrotoluene, nitroglycerine, and guncotton. It is also the acid used in lead-acid batteries, and so is sometimes known as battery acid.
The energy of the hydration reaction with sulphuric acid is highly exothermic, and if water is added to concentrated sulphuric acid it can boil. Always add the acid to the water rather than the water to the acid. Note that part of this problem is due to the relative densities of the two liquids. Water is less dense than sulfuric acid and will tend to layer above the acid, and not mix well, if added to the acid.
Because the hydration of sulfuric acid is thermodynamically favourable, sulphuric acid is an excellent dehydration agent, and is used to prepare many dried fruits.
When in the atmosphere it is part of many chemicals which make up acid rain.
Sulfuric acid was known to medieval alchemists under of variety of names including oil of vitriol and spirit of vitriol. These substances were produced by the dry distillation of minerals including ferrous sulfate heptahydrate, FeSO4 • 7 H2O, called green vitriol, and cupric sulfate pentahydrate, CuSO4 • 5 H2O, called blue vitriol. When heated, these compounds decompose to ferrous and cupric oxides, respectively, giving off water and sulfur trioxide, which combine to produce a dilute solution of sulfuric acid. Preparations like these have been ascribed to alchemists including the 12th-century Arab Abou Bekr al-Rhases and the 13th-century German Albertus Magnus.
In the 17th century, the German-Dutch chemist Johann Glauber[?] prepared sulfuric acid by burning sulfur together with saltpeter (potassium nitrate, KNO3), in the presence of steam. As the saltpeter decomposes, it oxidizes the sulfur to SO3, which combines with water to produce sulfuric acid. In 1736, Joshua Ward, a London pharmacist, used this method to begin the first large-scale production of sulfuric acid.
In 1746 in Birmingham, John Roebuck began producing sulfuric acid this way in lead-lined chambers, which were stronger, less expensive, and could be made larger than the glass containers which had been used previously. This lead-chamber process allowed the effective industrialization of sulfuric acid production, and with several refinements remained the standard method of production for almost two centuries.
John Roebuck's sulfuric acid was only about 35-40% sulfuric acid, and later refinements in the lead-chamber process improved this to 78%. However, the manufacture of some dyes and other chemical processes require a more concentrated product, and throughout the 18th century, this could only be made by dry distilling minerals in a technique similar to the original alchemical processes. Pyrite (iron disulfide, FeS2) was heated in air to yield ferrous sulfate, FeSO4, which was oxidzied by further heating in air to form ferric sulfate, Fe2(SO4)3, which when heated to 480°C decomposed to ferric oxide and sulfur trioxide, which could be passed through water to yield sulfuric acid in any concentration. The expense of this process prevented the large-scale use of concentrated sulfuric acid.
In 1831, the British merchant Peregrine Phillips patented a far more economical process for producing sulfur trioxide and concentrated sulfuric acid. In this process sulfur dioxide, SO2, produced by roasting either sulfur or pyrite in air, is combined with additional air and passed over a platinum catalyst at high temperatures, where it combines with oxygen from the air to produce nearly pure SO3. Even so, there was little demand for highly concentrated sulfuric acid at the time, and the first sulfuric acid plant using this contact process was not built until 1875 in Freiburg, Germany.
The development of the less expensive and less easily contaminated vanadium pentoxide (V2O5) catalyst by BASF in Gemany in 1915, combined with increasing demand for concentrated sulfuric acid by the chemical industry, has led to the gradual replacement of the lead-chamber process by the contact process. In 1930, sulfuric acid produced by the contact process accounted for only 25% of sulfuric acid production, while today nearly all sulfuric acid is manufactured in this way.
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