Sulfuric acid (alternative spelling sulphuric acid) is a highly corrosive strong mineral acid with the molecular formula H2SO4. It is a pungent-ethereal, colorless to slightly yellow viscous liquid which is soluble in water at all concentrations.[4] Sometimes, it is dyed dark brown during production to alert people to its hazards.[5] The historical name of this acid is oil of vitriol.[6] Sulfuric acid is a diprotic acid and shows different properties depending upon its concentration. Its corrosiveness on other materials, like metals, living tissues (e.g. skin and flesh) or even stones, can be mainly ascribed to its strong acidic nature and, if concentrated, strong dehydrating and oxidizing property. Sulfuric acid at a high concentration can cause very serious damage upon contact, as it not only causes chemical burns via hydrolysis, but also secondary thermal burns via dehydration. It burns the cornea and can lead to permanent blindness if splashed onto eyes. Accordingly, safety precautions should be strictly observed when handling it. Moreover, it is hygroscopic, readily absorbing water vapour from the air.[4] Sulfuric acid has a wide range of applications including domestic acidic drain cleaner,[7] electrolyte in lead-acid batteries and various cleaning agents. It is also a central substance in the chemical industry. Principal uses include mineral processing, fertilizer manufacturing, oil refining, wastewater processing, and chemical synthesis. It is widely produced with different methods, such as contact process, wet sulfuric acid process and some other methods. Contents [hide] 1 History 2 Physical properties 2.1 Grades of sulfuric acid 2.2 Polarity and conductivity 3 Chemical properties 3.1 Reaction with water and dehydrating property 3.2 Acid-base properties 3.3 Reactions with metals and strong oxidizing property 3.4 Reactions with non-metals 3.5 Reaction with sodium chloride 3.6 Electrophilic aromatic substitution 4 Occurrence 4.1 Extraterrestrial sulfuric acid 4.1.1 Venus 4.1.2 Europa 5 Manufacture 5.1 Contact process 5.2 Wet sulfuric acid process 5.3 Other methods 6 Uses 6.1 Industrial production of chemicals 6.2 Sulfur-iodine cycle 6.3 Industrial cleaning agent 6.4 Catalyst 6.5 Electrolyte 6.6 Domestic uses 6.7 Health 7 Safety 7.1 Laboratory hazards 7.2 Industrial hazards 8 Legal restrictions 9 See also 10 References 11 External links History[edit] John Daltons 1808 sulfuric acid molecule shows a central sulfur atom bonded to three oxygen atoms, or sulfur trioxide, the anhydride of sulfuric acid. The study of vitriol began in ancient times. Sumerians had a list of types of vitriol that they classified according to the substances color. Some of the earliest discussions on the origin and properties of vitriol is in the works of the Greek physician Dioscorides (first century AD) and the Roman naturalist Pliny the Elder (23–79 AD). Galen also discussed its medical use. Metallurgical uses for vitriolic substances were recorded in the Hellenistic alchemical works of Zosimos of Panopolis, in the treatise Phisica et Mystica, and the Leyden papyrus X.[8] Islamic alchemists Jābir ibn Hayyān (c. 721 – c. 815 AD), Razi (865 – 925 AD), and Jamal Din al-Watwat (d. 1318, wrote the book Mabāhij al-fikar wa-manāhij al-ibar), included vitriol in their mineral classification lists. Ibn Sina focused on its medical uses and different varieties of vitriol.[8] Sulfuric acid was called oil of vitriol by medieval European alchemists because it was prepared by roasting green vitriol (iron (II) sulfate) in an iron retort. There are references to it in the works of Vincent of Beauvais and in the Compositum de Compositis ascribed to Saint Albertus Magnus. A passage from Pseudo-Geber´s Summa Perfectionis was long considered to be the first recipe for sulfuric acid, but this was a misinterpretation.[8] In the seventeenth century, the German-Dutch chemist Johann Glauber prepared sulfuric acid by burning sulfur together with saltpeter (potassium nitrate, KNO 3), in the presence of steam. As saltpeter decomposes, it oxidizes the sulfur to SO 3, 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 adapted this method to produce sulfuric acid in lead-lined chambers, which were stronger, less expensive, and could be made larger than the previously used glass containers. This process allowed the effective industrialization of sulfuric acid production. After several refinements, this method, called the lead chamber process or chamber process, remained the standard for sulfuric acid production for almost two centuries.[1] Sulfuric acid created by John Roebucks process approached a 65% concentration. Later refinements to the lead chamber process by French chemist Joseph Louis Gay-Lussac and British chemist John Glover improved concentration to 78%. However, the manufacture of some dyes and other chemical processes require a more concentrated product. 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, FeS 2) was heated in air to yield iron(II) sulfate, FeSO 4, which was oxidized by further heating in air to form iron(III) sulfate, Fe2(SO4)3, which, when heated to 480 °C, decomposed to iron(III) oxide and sulfur trioxide, which could be passed through water to yield sulfuric acid in any concentration. However, the expense of this process prevented the large-scale use of concentrated sulfuric acid.[1] In 1831, British vinegar merchant Peregrine Phillips patented the contact process, which was a far more economical process for producing sulfur trioxide and concentrated sulfuric acid. Today, nearly all of the worlds sulfuric acid is produced using this method.[1] Physical properties[edit] Grades of sulfuric acid[edit] Although nearly 99% sulfuric acid can be made, the subsequent loss of SO 3 at the boiling point brings the concentration to 98.3% acid. The 98% grade is more stable in storage, and is the usual form of what is described as concentrated sulfuric acid. Other concentrations are used for different purposes. Some common concentrations are:[9][10] Mass fraction H2SO4 Density (kg/L) Concentration (mol/L) Common name 10% 1.07 ~1 dilute sulfuric acid 29–32% 1.25–1.28 4.2–5 battery acid (used in lead–acid batteries) 62–70% 1.52–1.60 9.6–11.5 chamber acid fertilizer acid 78–80% 1.70–1.73 13.5–14 tower acid Glover acid 98% 1.83 ~18 concentrated sulfuric acid Chamber acid and tower acid were the two concentrations of sulfuric acid produced by the lead chamber process, chamber acid being the acid produced in lead chamber itself (
Posted on: Sun, 26 Oct 2014 04:50:33 +0000
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