Sulfuric acid
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225959136
2008-07-16T05:30:09Z
WinterSpw
2881121
/* Other reactions */ {{clarifyme}}
{{Chembox new
| ImageFileL1 = Sulfuric-acid-2D-dimensions.svg
| ImageSizeL1 = 150px
| ImageFileR1 = Sulfuric-acid-Givan-et-al-1999-3D-vdW.png
| ImageSizeR1 = 120px
| IUPACName = sulfuric acid
| OtherNames = Sulphuric acid<br>Hydrogen sulfate</br>Oil of vitriol
| Section1 = {{Chembox Identifiers
| CASNo = 7664-93-9
| RTECS = WS5600000
}}
| Section2 = {{Chembox Properties
| Formula = H<sub>2</sub>SO<sub>4</sub>
| MolarMass = 98.078 g/mol
| Appearance = clear, colorless,<br/>odorless liquid
| Density = 1.84 g cm<sup>−3</sup>, liquid
| Solubility = fully miscible<br />(exothermic)
| MeltingPtC = 10
| BoilingPtC = 290
| Boiling_notes = bp of pure acid. 98% solution boils at 338°C
| Viscosity = 26.7 c[[Poise|P]] at 20°C
}}
| Section7 = {{Chembox Hazards
| FlashPt = Non-flammable
| EUClass = Highly Corrosive ('''C''')
| NFPA-H = 3
| NFPA-F = 0
| NFPA-R = 2
| NFPA-O = COR
| RPhrases = {{R35}}
| SPhrases = {{S1/2}}, {{S26}}, {{S30}}, {{S45}}
}}
| Section8 = {{Chembox Related
| Function = [[strong acid]]s
| OtherFunctn = [[Hydrochloric acid]]<br />[[Nitric acid]]<br />[[Hydroiodic acid]]<br />[[Hydrobromic acid]]
| OtherCpds = [[Hydrogen sulfide]]<br />[[Sulfurous acid]]<br />[[Peroxymonosulfuric acid]]<br />[[Sulfur trioxide]]<br />[[Oleum]]}}
}}
'''Sulfuric acid''' (also spelled sulphuric acid) is a [[strong acid|strong]] [[mineral acid]] with the molecular formula [[hydrogen|H]]<sub>2</sub>[[sulfur|S]][[oxygen|O]]<sub>4</sub>. It is soluble in [[water]] at all concentrations. It was once known as '''oil of vitriol''', coined by the 8th-century [[Alchemy (Islam)|Muslim alchemist]] [[Geber|Jabir ibn Hayyan]] (Geber) after his discovery of the chemical.<ref>Khairallah, Amin A. ''Outline of Arabic Contributions to Medicine'', chapter 10. Beirut, 1946.</ref>
Sulfuric acid has many applications, and is one of the top products of the [[chemical industry]]. World production in 2001 was 165 million [[tonne|metric tons]], with an approximate value of US$8 billion. Principal uses include [[ore]] processing, [[fertilizer]] manufacturing, [[Oil refinery|oil refining]], [[wastewater processing]], and [[chemical synthesis]]. Its ability to produce foul-smelling sulfur compounds has lent the word ''[[wiktionary:vitriol|vitriol]]'' the additional meaning "bitter, abusive language".
Many [[protein]]s are made of [[sulfur]]-containing amino [[acid]]s (such as [[cysteine]] and [[methionine]]) which produce sulfuric acid (or sulfate ion, SO<sub>4</sub><sup>2-</sup> at neutral pH) when [[metabolism|metabolized]] by the body.
==Occurrence==
Pure undiluted sulfuric acid is not encountered on Earth, due to sulfuric acid's great affinity for water. Apart from that, sulfuric acid is a constituent of [[acid rain]], which is formed by atmospheric [[Redox|oxidation]] of [[sulfur dioxide]] in the presence of [[water (molecule)|water]], i.e., oxidation of [[sulfurous acid]]. Sulfur dioxide is produced when sulfur-containing fuels such as coal or oil are burnt.
Sulfuric acid is formed naturally by the oxidation of sulfide minerals, such as iron sulfide. The resulting water can be highly acidic and is called [[Acid mine drainage|Acid Mine Drainage]] (AMD). This acidic water is capable of dissolving metals present in sulfide ores, which results in brightly-coloured, toxic streams. The oxidation of iron sulfide [[pyrite]] (FeS<sub>2</sub>) by molecular oxygen produces iron(II), or Fe<sup>2+</sup>:
:2 FeS<sub>2</sub> + 7 O<sub>2</sub> + 2 H<sub>2</sub>O → 2 Fe<sup>2+</sub> + 4 SO<sub>4</sub><sup>2−</sup> + 4 H<sup>+</sup>
The Fe<sup>2+</sup> can be further oxidized to Fe<sup>3+</sup>, according to:
:4 Fe<sup>2+</sup> + O<sub>2</sub> + 4 H<sup>+</sup> → 4 Fe<sup>3+</sup> + 2 H<sub>2</sub>O
and the Fe<sup>3+</sup> produced can be precipitated as the [[hydroxide]] or [[hydrous oxide]]. The equation for the formation of the hydroxide is
:Fe<sup>3+</sup> + 3 H<sub>2</sub>O → Fe(OH)<sub>3</sub> + 3 H<sup>+</sup>
The iron(III) ion ("ferric iron", in casual nomenclature) can also oxidize pyrite. When iron(III) oxidation of pyrite occurs, the process can become rapid. pH values below zero have been measured in AMD produced by this process.
AMD can also produce sulfuric acid at a slower rate, so that the [[acid neutralizing capacity]] (ANC) of the aquifer can neutralise the produced acid. In such cases, the [[total dissolved solids]] (TDS) concentration of the water can be increased form the dissolution of minerals from the acid-neutralisation reaction with the minerals.
===Extraterrestrial sulfuric acid===
Sulfuric acid is produced in the upper atmosphere of [[Venus]] by the sun's [[photochemistry|photochemical]] action on [[carbon dioxide]], [[sulfur dioxide]], and [[water]] vapour. [[Ultraviolet]] [[photon]]s of wavelengths less than 169 [[nanometer|nm]] can [[photodissociation|photodissociate]] [[carbon dioxide]] into [[carbon monoxide]] and atomic [[oxygen]]. Atomic oxygen is highly reactive. When it reacts with sulfur dioxide, a trace component of the Venerian atmosphere, the result is [[sulfur trioxide]], which can combine with water vapor, another trace component of Venus's atmosphere, to yield sulfuric acid.
:CO<sub>2</sub> → CO + O
:SO<sub>2</sub> + O → SO<sub>3</sub>
:SO<sub>3</sub> + H<sub>2</sub>O → H<sub>2</sub>SO<sub>4</sub>
In the upper, cooler portions of Venus's atmosphere, sulfuric acid exists as a liquid, and thick sulfuric acid [[cloud]]s completely obscure the planet's surface when viewed from above. The main cloud layer extends from 45–70 [[kilometer|km]] above the planet's surface, with thinner hazes extending as low as 30 and as high as 90 km above the surface.
Infrared spectra from [[NASA]]'s [[Galileo (spacecraft)|Galileo mission]] show distinct absorptions on [[Jupiter]]'s moon [[Europa (moon)|Europa]] that have been attributed to one or more sulfuric acid hydrates. The interpretation of the spectra is somewhat controversial. Some planetary scientists prefer to assign the spectral features to the sulfate ion, perhaps as part of one or more minerals on Europa's surface.<ref>T.M. Orlando, T.B. McCord, G.A Grieves, Icarus 177 (2005) 528–533
</ref>
==Manufacture==
{{main|Contact process}}
Sulfuric acid is produced from [[sulfur]], [[oxygen]] and [[water]] via the [[contact process]].
In the first step, [[sulfur]] is burned to produce [[sulfur dioxide]].
:(1) S([[Solid|s]]) + O<sub>2</sub>([[Gas|g]]) → SO<sub>2</sub>(g)
This is then oxidised to [[sulfur trioxide]] using [[oxygen]] in the presence of a [[vanadium(V) oxide]] [[catalyst]].
:(2) 2 SO<sub>2</sub> + O<sub>2</sub>(g) → 2 SO<sub>3</sub>(g) (in presence of [[Vanadium pentoxide|V<sub>2</sub>O<sub>5</sub>]])
Finally the [[sulfur trioxide]] is treated with water (usually as 97-98% H<sub>2</sub>SO<sub>4</sub> containing 2-3% water) to produce 98-99% sulfuric acid.
:(3) SO<sub>3</sub>(g) + H<sub>2</sub>O([[liquid|l]]) → H<sub>2</sub>SO<sub>4</sub>(l)
Note that directly dissolving SO<sub>3</sub> in water is not practical due to the highly [[Exothermic reaction|exothermic]] nature of the [[Chemical reaction|reaction]], forming a corrosive mist instead of a liquid.
Alternatively, SO<sub>3</sub> can be absorbed into H<sub>2</sub>SO<sub>4</sub> to produce [[oleum]] (H<sub>2</sub>S<sub>2</sub>O<sub>7</sub>), which may then be mixed with water to form sulfuric acid.
:(3) H<sub>2</sub>SO<sub>4</sub>(l) + SO<sub>3</sub> → H<sub>2</sub>S<sub>2</sub>O<sub>7</sub>(l)
Oleum is reacted with water to form concentrated H<sub>2</sub>SO<sub>4</sub>.
:(4) H<sub>2</sub>S<sub>2</sub>O<sub>7</sub>(l) + H<sub>2</sub>O(l) → 2 H<sub>2</sub>SO<sub>4</sub>(l)
==Physical properties==
===Forms of sulfuric acid===
Although nearly 100% sulfuric acid can be made, this loses [[sulfur trioxide|SO<sub>3</sub>]] at the boiling point to produce 98.3% acid. The 98% grade (18M) is more stable in storage, and is the usual form of what is described as ''concentrated sulfuric acid''. Other [[concentration]]s are used for different purposes. Some common concentrations are
* 10%, ''dilute sulfuric acid'' for laboratory use,
* 33.5%, ''battery acid'' (used in [[lead-acid battery|lead-acid batteries]]),
* 62.18%, ''chamber'' or ''fertilizer acid'',
* 77.67%, ''tower'' or ''Glover acid'',
* 98%, ''concentrated acid''.
Different purities are also available. Technical grade H<sub>2</sub>SO<sub>4</sub> is impure and often colored, but is suitable for making fertilizer. Pure grades such as [[United States Pharmacopoeia]] (USP) grade are used for making [[pharmaceutical]]s and [[dye]]stuffs.
When high concentrations of SO<sub>3''([[gas|g]])</sub>'' are added to sulfuric acid, H<sub>2</sub>S<sub>2</sub>O<sub>7</sub>, called ''[[pyrosulfuric acid]]'', ''fuming sulfuric acid'' or ''[[oleum]]'' or, less commonly, ''Nordhausen acid'', is formed. Concentrations of oleum are either expressed in terms of% SO<sub>3</sub> (called% oleum) or as% H<sub>2</sub>SO<sub>4</sub> (the amount made if H<sub>2</sub>O were added); common concentrations are 40% oleum (109% H<sub>2</sub>SO<sub>4</sub>) and 65% oleum (114.6% H<sub>2</sub>SO<sub>4</sub>). Pure H<sub>2</sub>S<sub>2</sub>O<sub>7</sub> is a solid with melting point 36°C.
===Polarity and conductivity===
[[Anhydrous]] H<sub>2</sub>SO<sub>4</sub> is a very [[Chemical polarity|polar]] liquid, having a [[dielectric constant]] of around 100. It has a high electrical conductivity, caused by dissociation through [[protonation|protonating]] itself, a process known as [[autoprotolysis]], or [[autoionization]].<ref name = greenwood>{{Greenwood&Earnshaw}}</ref>
: 2 H<sub>2</sub>SO<sub>4</sub> {{unicode|⇌}} H<sub>3</sub>SO<sub>4</sub><sup>+</sup> + HSO<sub>4</sub><sup>−</sup>
The [[equilibrium constant]] for the autoprotolysis is<ref name = greenwood/>
:K<sub>ap</sub>(25°C)= [H<sub>3</sub>SO<sub>4</sub><sup>+</sup>][HSO<sub>4</sub><sup>−</sup>] = 2.7 × 10<sup><nowiki>−</nowiki>4</sup>.
The comparable equilibrium constant for [[Self-ionization of water|water]], K<sub>w</sub> is 10<sup><nowiki>−</nowiki>14</sup>, a factor of 10<sup>10</sup> (10 billion) smaller.
In spite of the viscosity of the acid, the effective [[Molar conductivity|conductivities]] of the H<sub>3</sub>SO<sub>4</sub><sup>+</sup> and HSO<sub>4</sub><sup>−</sup> ions are high due to an intra-molecular proton-switch mechanism (analogous to the [[Grotthuss mechanism]] in water), making sulfuric acid a good conductor. It is also an excellent solvent for many reactions.
The [[chemical equilibrium|equilibrium]] is actually more complex than shown above; 100% H<sub>2</sub>SO<sub>4</sub> contains the following species at equilibrium (figures shown as millimol per kg solvent): HSO<sub>4</sub><sup>−</sup> (15.0), H<sub>3</sub>SO<sub>4</sub><sup>+</sup> (11.3), H<sub>3</sub>O<sup>+</sup> (8.0), HS<sub>2</sub>O<sub>7</sub><sup>−</sup> (4.4), [[Oleum|H<sub>2</sub>S<sub>2</sub>O<sub>7</sub>]] (3.6), H<sub>2</sub>O (0.1).<ref name = greenwood/>
==Chemical properties==
===Reaction with water===
The [[hydration reaction]] of sulfuric acid is highly [[exothermic reaction|exothermic]]. If [[water]] is added to the concentrated sulfuric acid, it can react, boil and spit dangerously. One should always add the acid to the water rather than the water to the acid. The necessity for this safety precaution is due to the relative densities of these two liquids. Water is less [[density|dense]] than sulfuric acid, meaning water will tend to float on top of this acid. The reaction is best thought of as forming [[hydronium]] ions, by
: H<sub>2</sub>SO<sub>4</sub> + H<sub>2</sub>O → H<sub>3</sub>O<sup>+</sup> + HSO<sub>4</sub><sup><nowiki>−</nowiki></sup>,
and then
: HSO<sub>4</sub><sup><nowiki>−</nowiki></sup> + H<sub>2</sub>O → H<sub>3</sub>O<sup>+</sup> + SO<sub>4</sub><sup>2<nowiki>−</nowiki></sup>.
Because the hydration of sulfuric acid is [[thermodynamic]]ally favorable, sulfuric acid is an excellent dehydrating agent. The affinity of sulfuric acid for [[water (molecule)|water]] is sufficiently strong that it will remove [[hydrogen]] and [[oxygen]] atoms from other compounds; for example, mixing [[starch]] (C<sub>6</sub>H<sub>12</sub>O<sub>6</sub>)<sub>n</sub> and concentrated sulfuric acid will give elemental [[carbon]] and water which is absorbed by the sulfuric acid (which becomes slightly diluted): (C<sub>6</sub>H<sub>12</sub>O<sub>6</sub>)<sub>n</sub> → 6C + 6H<sub>2</sub>O. The effect of this can be seen when concentrated sulfuric acid is spilled on paper; the cellulose reacts to give a [[combustion|burned]] appearance, the [[carbon]] appears much as soot would in a fire. A more dramatic reaction occurs when sulfuric acid is added to a tablespoon of white [[sugar]]; a rigid column of black, porous carbon will quickly emerge. The carbon will smell strongly of [[caramel]].
===Other reactions===
As an acid, sulfuric acid reacts with most [[base (chemistry)|bases]] to give the corresponding [[sulfate]]. For example, [[copper(II) sulfate]]. This blue salt of [[copper]], commonly used for [[electroplating]] and as a [[fungicide]], is prepared by the reaction of [[copper(II) oxide]] (CuO) with sulfuric acid:
:CuO + H<sub>2</sub>SO<sub>4</sub> → CuSO<sub>4</sub>+ H<sub>2</sub>O
Sulfuric acid can also be used to displace weaker acids from their salts. Reaction with [[sodium acetate]], for example, displaces [[acetic acid]]:
:H<sub>2</sub>SO<sub>4</sub> + CH<sub>3</sub>COONa → NaHSO<sub>4</sub> + CH<sub>3</sub>COOH
Similarly, reacting sulfuric acid with [[potassium nitrate]] can be used to produce [[nitric acid]] and a precipitate of [[potassium bisulfate]]. When combined with [[nitric acid]], sulfuric acid acts both as an acid and a dehydrating agent, forming the [[nitronium]] ion NO<sub>2</sub><sup>+</sup>, which is important in [[nitration]] reactions involving [[electrophilic aromatic substitution]]. This type of reaction, where protonation occurs on an [[oxygen]] atom, is important in many [[organic chemistry]] reactions, such as [[Fischer esterification]] and [[dehydration of alcohols]].
Sulfuric acid reacts with most metals via a single displacement reaction to produce [[hydrogen]] gas and the metal sulfate. Dilute H<sub>2</sub>SO<sub>4</sub> attacks{{clarifyme}} [[iron]], [[aluminium]], [[zinc]], [[manganese]], [[magnesium]] and [[nickel]], but reactions with [[tin]] and [[copper]] require the acid to be hot and concentrated. [[Lead]] and [[tungsten]], however, are resistant to sulfuric acid. The reaction with iron (shown) is typical for most of these metals, but the reaction with tin is unusual in that it produces [[sulfur dioxide]] rather than hydrogen.
:Fe(s) + H<sub>2</sub>SO<sub>4</sub>(aq) → H<sub>2</sub>(g) + FeSO<sub>4</sub>(aq)
:Sn(s) + 2 H<sub>2</sub>SO<sub>4</sub>(aq) → SnSO<sub>4</sub>(aq) + 2 H<sub>2</sub>O(l) + SO<sub>2</sub>(g)
Sulfuric acid undergoes [[electrophilic aromatic substitution]] with [[aromatic compound]]s to give the corresponding [[sulfonic acids]]:<ref>{{cite web | url = http://www.chem.ucalgary.ca/courses/351/Carey/Ch12/ch12-4.html| title = Reactions of Arenes. Electrophilic Aromatic Substitution | author = F. A. Carey | | work = On-Line Learning Center for Organic Chemistry | publisher = [[University of Calgary]] | accessdate = 2008-01-27}}</ref>
:[[Image:BenzeneSulfonation.png|300px]]
==Uses==
[[Image:2000sulphuric acid.PNG|thumb|right|Sulfuric acid production in 2000]]
[[Image:Diagram Production sulfuric acid en.svg|thumb|Production trend in some countries]]
Sulfuric acid is a very important [[commodity chemical]], and indeed, a nation's sulfuric acid production is a good indicator of its industrial strength.<ref>Chenier, Philip J. ''Survey of Industrial Chemistry'', pp 45-57. John Wiley & Sons, New York, 1987. ISBN.</ref> The major use (60% of total production worldwide) for sulfuric acid is in the "wet method" for the production of [[phosphoric acid]], used for manufacture of [[phosphate]] [[fertilizer]]s as well as [[trisodium phosphate]] for detergents. In this method, phosphate rock is used, and more than 100 million tonnes are processed annually. This raw material is shown below as [[fluorapatite]], though the exact composition may vary. This is treated with 93% sulfuric acid to produce [[calcium sulfate]], [[hydrogen fluoride]] (HF) and [[phosphoric acid]]. The HF is removed as [[hydrofluoric acid]]. The overall process can be represented as:
:[[fluorapatite|Ca<sub>5</sub>F(PO<sub>4</sub>)<sub>3</sub>]] + 5 H<sub>2</sub>SO<sub>4</sub> + 10 [[water (molecule)|H<sub>2</sub>O]] → 5 [[calcium sulfate|CaSO<sub>4</sub>]]•2 H<sub>2</sub>O + [[hydrogen fluoride|HF]] + 3 [[phosphoric acid|H<sub>3</sub>PO<sub>4</sub>]].
Sulfuric acid is used in large quantities by the [[iron]] and [[steel]]making industry to remove oxidation, [[rust]] and scale from rolled sheet and billets prior to sale to the [[automobile]] and white-goods industry. Used acid is often recycled using a Spent Acid Regeneration (SAR) plant. These plants combust spent acid with natural gas, refinery gas, fuel oil or other fuel sources. This combustion process produces gaseous sulfur dioxide (SO<sub>2</sub>) and sulfur trioxide (SO<sub>3</sub>) which are then used to manufacture "new" sulfuric acid. SAR plants are common additions to metal smelting plants, oil refineries, and other industries where sulfuric acid is consumed in bulk, as operating a SAR plant is much cheaper than the recurring costs of spent acid disposal and new acid purchases.
[[Ammonium sulfate]], an important nitrogen fertilizer, is most commonly produced as a byproduct from coking plants supplying the iron and steel making plants. Reacting the [[ammonia]] produced in the thermal decomposition of [[coal]] with waste sulfuric acid allows the ammonia to be crystallized out as a salt (often brown because of iron contamination) and sold into the agro-chemicals industry.
Another important use for sulfuric acid is for the manufacture of [[aluminium sulfate]], also known as paper maker's alum. This can react with small amounts of soap on [[paper pulp]] fibers to give gelatinous aluminium [[carboxylate]]s, which help to coagulate the pulp fibers into a hard paper surface. It is also used for making [[aluminium hydroxide]], which is used at [[water treatment]] plants to [[filter (water)|filter]] out impurities, as well as to improve the taste of the [[water]]. Aluminum sulfate is made by reacting [[bauxite]] with sulfuric acid:
:[[aluminium oxide|Al<sub>2</sub>O<sub>3</sub>]] + 3 H<sub>2</sub>SO<sub>4</sub> → [[aluminium sulfate|Al<sub>2</sub>(SO<sub>4</sub>)<sub>3</sub>]] + 3 [[water (molecule)|H<sub>2</sub>O]].
Sulfuric acid is used for a variety of other purposes in the chemical industry. For example, it is the usual acid catalyst for the conversion of [[cyclohexanoneoxime]] to [[caprolactam]], used for making [[nylon]]. It is used for making [[hydrochloric acid]] from [[salt]] via the [[Mannheim process]]. Much H<sub>2</sub>SO<sub>4</sub> is used in [[petroleum]] refining, for example as a catalyst for the reaction of [[isobutane]] with [[isobutylene]] to give [[isooctane]], a compound that raises the [[octane rating]] of [[gasoline]] (petrol). Sulfuric acid is also important in the manufacture of [[dye]]stuffs, pigments (such as [[titanium dioxide]]), solutions, and is the "acid" in lead-acid (car) batteries.
Sulfuric acid is also used as a general dehydrating agent in its concentrated form (''see [[sulfuric acid#Reaction with water|Reaction with water]]'').
===Sulfur-iodine cycle===
The [[sulfur-iodine cycle]] is a series of thermo-chemical processes used to obtain [[hydrogen]]. It consists of three chemical reactions whose net reactant is [[water]] and whose net products are hydrogen and [[oxygen]].
:{|
|-
| 2 {{hydrogen|2}}{{sulfur}}{{oxygen|4}} → 2 {{sulfur}}{{oxygen|2}} + 2 {{hydrogen|2}}{{oxygen}} + {{oxygen|2}} || || (830°C)
|-
| {{iodine|2}} + {{sulfur}}{{oxygen|2}} + 2 {{hydrogen|2}}{{oxygen}} → 2 {{hydrogen}}{{iodine}} + {{hydrogen|2}}{{sulfur}}{{oxygen|4}} || || (120°C)
|-
| 2 {{hydrogen}}{{iodine}} → {{iodine|2}} + {{hydrogen|2}} || || (320°C)
|}
The sulfur and [[iodine]] compounds are recovered and reused, hence the consideration of the process as a cycle. This process is [[endothermic]] and must occur at high temperatures, so energy in the form of heat has to be supplied.
The sulfur-iodine cycle has been proposed as a way to supply hydrogen for a [[hydrogen economy|hydrogen-based economy]]. It does not require [[hydrocarbons]] like current methods of [[steam reforming]].
The sulfur-iodine cycle is currently being researched as a feasible method of obtaining hydrogen, but the concentrated, corrosive acid at high temperatures poses currently insurmountable safety hazards if the process were built on large-scale.
==History==
[[Image:Dalton's-sulphuric-acid.jpg|right|thumb|230px|[[John Dalton]]'s 1808 sulfuric acid molecule shows a central sulfur atom bonded to three oxygen atoms.]]
The discovery of sulfuric acid is credited to the 8th century [[Alchemy (Islam)|Arabian chemist and alchemist]], [[Geber|Jabir ibn Hayyan]] (Geber). The acid was later studied by 9th century [[Islamic medicine|Persian physician]] and alchemist [[Al-Razi|Ibn Zakariya al-Razi]] (Rhazes), who obtained the substance by [[dry distillation]] of minerals including [[iron(II) sulfate]] heptahydrate, FeSO<sub>4</sub> • 7H<sub>2</sub>O, and [[copper(II) sulfate]] pentahydrate, CuSO<sub>4</sub> • 5H<sub>2</sub>O. When heated, these compounds decompose to [[iron(II) oxide]] and [[copper(II) oxide]], respectively, giving off [[water (molecule)|water]] and [[sulfur trioxide]], which combine to produce a dilute solution of sulfuric acid. This method was popularized in Europe through translations of Arabic and Persian treatises, as well as books by European alchemists, such as the 13th-century German [[Albertus Magnus]].
Sulfuric acid was known to medieval European alchemists as ''oil of vitriol'', ''spirit of vitriol'', or simply ''vitriol'', among other names. The word vitriol derives from the Latin vitreus, 'glass', referring to the glassy appearance of the sulfate salts, which also carried the name vitriol. Salts called by this name included [[copper(II) sulfate]] (blue vitriol, or rarely [[Rome|Roman]] vitriol), [[zinc sulfate]] (white vitriol), [[iron(II) sulfate]] (green vitriol), [[iron(III) sulfate]] (vitriol of Mars), and [[cobalt(II) sulfate]] (red vitriol).
Vitriol was widely considered the most important [[alchemy|alchemical]] substance, intended to be used as a [[philosopher's stone]]. Highly purified vitriol was used as a medium for reacting other substances. This was largely because the acid does not react with [[gold]], production of which was often the final goal of alchemical processes. The importance of vitriol to alchemy is highlighted in the alchemical motto, ''Visita Interiora Terrae Rectificando Invenies Occultum Lapidem'' which is a [[backronym]] meaning ('Visit the interior of the earth and rectifying (i.e. purifying) you will find the hidden/secret stone'), found in ''[[L'Azoth des Philosophes]]'' by the 15th Century alchemist [[Basilius Valentinus]], .
In the 17th century, the German-Dutch chemist [[Johann Glauber]] prepared sulfuric acid by burning [[sulfur]] together with [[Potassium nitrate|saltpeter]] ([[potassium nitrate]], KNO<sub>3</sub>), in the presence of steam. As saltpeter decomposes, it oxidizes the sulfur to SO<sub>3</sub>, 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 ''[[lead chamber process]]'' allowed the effective industrialization of sulfuric acid production. After several refinements, this method remained the standard for sulfuric acid production for almost two centuries.<ref>{{cite book | author = Derry, Thomas Kingston | coauthors = Williams, Trevor I. | year = 1993 | title = A Short History of Technology: From the Earliest Times to A.D. 1900 | location = New York | publisher = Dover }}</ref><ref>{{cite web | title = Sulfuric Acid: Pumping Up the Volume | author = Kiefer, David M. | year = 2001 | publisher = American Chemical Society | url = http://pubs.acs.org/subscribe/journals/tcaw/10/i09/html/09chemch.html | access-date = 2008-04-21}}</ref>
Sulfuric acid created by John Roebuck's process only approached a 35–40% concentration. Later refinements to the lead-chamber process by French chemist [[Joseph-Louis Gay-Lussac]] and British chemist [[John Glover]] improved the yield to 78%. However, the manufacture of some [[dye]]s and other chemical processes require a more concentrated product. Throughout the 18th century, this could only be made by [[dry distillation|dry distilling]] minerals in a technique similar to the original [[alchemy|alchemical]] processes. [[Pyrite]] ([[iron disulfide]], FeS<sub>2</sub>) was heated in air to yield [[iron (II) sulfate]], FeSO<sub>4</sub>, which was oxidized by further heating in air to form [[iron(III) sulfate]], Fe<sub>2</sub>(SO<sub>4</sub>)<sub>3</sub>, 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.
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 world's sulfuric acid is produced using this method.
==Safety==
===Laboratory hazards===
[[Image:Sulfuric acid burning tissue paper.jpg|thumb|Drops of 98 % sulfuric acid burn a piece of tissue paper instantly]]
The corrosive properties of sulfuric acid are accentuated by its highly [[exothermic reaction]] with [[water (molecule)|water]]. Hence burns from sulfuric acid are potentially more serious than those of comparable strong acids (e.g. [[hydrochloric acid]]), as there is additional tissue damage due to dehydration and particularly due to the heat liberated by the reaction with water; i.e. secondary thermal damage. The danger is obviously greater with more concentrated preparations of sulfuric acid, but it should be remembered that even the normal laboratory "dilute" grade (approx. 1 M, 10%) will char paper by dehydration if left in contact for a sufficient amount of time. Solutions equal to or stronger than 1.5 M should be labeled CORROSIVE, while solutions greater than 0.5 M but less than 1.5 M should be labeled IRRITANT. Fuming sulfuric acid (oleum) is not recommended for use in schools due to it being quite hazardous. The standard first aid treatment for acid spills on the skin is, as for other [[corrosion|corrosive]] agents, irrigation with large quantities of water. Contrary to previous advice, one should NEVER attempt to neutralise the acid: neutralisation is exothermic, and the reaction will increase the severity of the burn. Whilst small amounts of water added to acid do produce considerable heat, a constant stream will quickly remove both the heat and the acid. All standard first aid and health and safety bodies emphasise that plain water is the ONLY medium recommended for immediate treatment of corrosive substance contamination. Advice recommending the use of calcium hydroxide, sodium hydroxide and so on is particularly dangerous: not only is there a strong exothermic reaction with these substances, both are also corrosive in their own right, and will substantially increase both the severity and the area of the burn.
''Washing should be continued for at least ten to fifteen minutes in order to cool the tissue surrounding the acid burn and to prevent secondary damage. Contaminated clothing must be removed immediately and the underlying skin washed thoroughly.''
Preparation of the diluted acid can also be dangerous due to the heat released in the dilution process. It is essential that the concentrated acid is added to water and not the other way round, to take advantage of the relatively high [[specific heat capacity|heat capacity]] of water. Addition of water to concentrated sulfuric acid leads at best to the dispersal of a sulfuric acid [[particulate|aerosol]], at worst to an [[explosion]]. Preparation of solutions greater than 6 M (35%) in concentration is the most dangerous, as the heat produced can be sufficient to boil the diluted acid: efficient mechanical stirring and external cooling (e.g. an ice bath) are essential.
===Industrial hazards===
Although sulfuric acid is non-flammable, contact with metals in the event of a spillage can lead to the liberation of [[hydrogen]] gas. The dispersal of acid aerosols and gaseous [[sulfur dioxide]] is an additional hazard of fires involving sulfuric acid.
Sulfuric acid is not considered toxic besides its obvious corrosive hazard, and the main occupational risks are skin contact leading to burns (see above) and the inhalation of aerosols.<ref>International Labour Organization. (February 2000). Sulfuric Acid. ''International Labour Organization.'' Retrieved March 25, 2008. From http://www.ilo.org/public/english/protection/safework/cis/products/icsc/dtasht/_icsc03/icsc0362.htm.</ref> Exposure to aerosols at high concentrations leads to immediate and severe irritation of the eyes, respiratory tract and mucous membranes: this ceases rapidly after exposure, although there is a risk of subsequent [[pulmonary edema]] if tissue damage has been more severe. At lower concentrations, the most commonly reported symptom of chronic exposure to sulfuric acid aerosols is erosion of the teeth, found in virtually all studies: indications of possible chronic damage to the [[respiratory tract]] are inconclusive as of 1997. In the [[United States]], the [[permissible exposure limit]] (PEL) for sulfuric acid is fixed at 1 mg/m³: limits in other countries are similar. Interestingly there have been reports of sulfuric acid ingestion leading to [[vitamin B12 deficiency]] with subacute combined degeneration. The spinal cord is most often affected in such cases, but the optic nerves may show [[demyelination]], loss of [[axon]]s and [[gliosis]].
==Legal restrictions==
International commerce of sulfuric acid is controlled under the [[United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances|United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances, 1988]], which lists sulfuric acid under Table II of the convention as a chemical frequently used in the illicit manufacture of narcotic drugs or psychotropic substances.<ref name=incb>[http://www.incb.org/pdf/e/list/red.pdf Annex to Form D ("Red List")], 11th Edition, January 2007 (pg. 4). [[International Narcotics Control Board]]. [[Vienna, Austria]]; 2007.</ref>
In the [[United States of America]], sulfuric acid is included in [[DEA list of chemicals#List II chemicals|List II]] of the [[DEA list of chemicals|list of essential or precursor chemicals]] established pursuant to the [[Chemical Diversion and Trafficking Act]]. Accordingly, transactions of sulfuric acid—such as sales, transfers, exports from and imports to the United States—are subject to regulation and monitoring by the [[Drug Enforcement Administration]].<ref name="66 CFR 52670">[http://frwebgate.access.gpo.gov/cgi-bin/getdoc.cgi?dbname=2001_register&docid=01-26013-filed 66 FR 52670—52675.] 17 October 2001.</ref><ref name="21 CFR 1309">[http://www.access.gpo.gov/nara/cfr/waisidx_06/21cfr1309_06.html 21 CFR 1309]</ref><ref name="CSA">[http://www.usdoj.gov/dea/pubs/csa.html 21 USC, Chapter 13 (Controlled Substances Act)]</ref>
==In fiction==
In several films, cartoons and TV shows, especially [[Science-Fiction]] shows and films, sulfuric acid is sometimes depicted as a bubbling green steaming liquid, sometimes capable of dissolving almost anything in an instant. This is purely for visual appeal, since boiling green acid is more dangerous-looking than the actual clear and syrupy form of sulfuric acid. The use of sulfuric acid as a weapon in crimes of assault, known as "[[Vitriolage|vitriol throwing]]", has at times been sufficiently common (if sensational) to make its way into novels and short stories. Examples include [[The Adventure of the Illustrious Client]] by [[Arthur Conan Doyle]], [[The Love of Long Ago]] by [[Guy de Maupassant]], ''[[Nineteen Eighty-Four]]'' by [[George Orwell]] and ''[[Brighton Rock (novel)|Brighton Rock]]'' by [[Graham Greene]]. The novel ''[[Veronika Decides to Die]]'' by [[Paulo Coelho]] talks of a girl who has attempted to commit suicide and ends up with vitriol poisoning. The graphic novel, [[Shake Girl]], is dedicated to over 100 Cambodian victims who have suffered from burns caused by sulfuric acid. In the movie ''[[Untraceable]]'', Griffin Dowd (played by [[Colin Hanks]]) is killed when his kidnapper adds sulfuric acid to a tank of water in which Dowd is submerged. In the popular [[Batman]] comic book series (and some theatrical films), the character of Harvey Dent is scared horribly when sulfuric acid is thrown on the left side of his face. This event triggers his turn into the villainous [[Two-Face]].
==References==
{{reflist}}
{{refbegin}}
* Institut National de Recherche et de Sécurité. (1997). "Acide sulfurique". ''Fiche toxicologique n°30'', Paris: INRS, 5 pp.
* ''Handbook of Chemistry and Physics'', 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
* Agamanolis DP. Metabolic and toxic disorders. In: Prayson R, editor. Neuropathology: a volume in the foundations in diagnostic pathology series. Philadelphia: Elsevier/Churchill Livingstone, 2005; 413-315.
{{refend}}
==External links==
* {{ICSC|0362|03}}
* [http://www.cdc.gov/niosh/npg/npgd0577.html NIOSH Pocket Guide to Chemical Hazards]
* [http://ptcl.chem.ox.ac.uk/MSDS/SU/sulfuric_acid_concentrated.html External Material Safety Data Sheet]
* [http://www2.iq.usp.br/docente/gutz/Curtipot_.html Sulfuric acid analysis - titration freeware]
* [http://twt.mpei.ac.ru/MCS/Worksheets/Chem/Sulfuric-acid-density-pH.xmcd Sulfuric Acid density and pH-value at t=20°C]
[[Category:Sulfur oxoacids]]
[[Category:Sulfates]]
[[Category:Inorganic compounds]]
[[Category:Acids]]
[[Category:Alchemical substances]]
[[Category:DEA List II chemicals]]
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