View source for Sulfuric acid

{{chembox
| Watchedfields = changed
| verifiedrevid = 265556728
|   Name = Sulfuric acid
|   ImageFile = Sulfuric-acid-2D-dimensions.svg
|   ImageFile1 = Sulfuric-acid-3D-vdW.png
|   IUPACName = Sulfuric acid
|   OtherNames = Oil of vitriol
| Section1 = {{Chembox Identifiers
|   CASNo = 7664-93-9
|    CASNo_Ref = {{cascite}}
|   RTECS = WS5600000
|   EINECS = 231-639-5
|   UNNumber = 1830
  }}
| Section2 = {{Chembox Properties
|   S=1|O=4|H=2
|   Appearance = Clear, colorless, odorless liquid
|   Density = 1.84 g/cm<sup>3</sup>, liquid
|   Solubility = miscible
|   MeltingPtC = 10
|   BoilingPtC = 337
|   Viscosity = 26.7 [[Poise|cP]] (20 °C)
|   pKa = &minus;3
  }}
| Section7 = {{Chembox Hazards
|   ExternalMSDS = [http://www.ilo.org/public/english/protection/safework/cis/products/icsc/dtasht/_icsc03/icsc0362.htm ICSC 0362]
|   EUIndex = 016-020-00-8
|   FlashPt = Non-flammable
|   EUClass = Corrosive ('''C''')
|   NFPA-H = 3
|   NFPA-F = 0
|   NFPA-R = 2
|   NFPA-O = W
|   RPhrases = {{R35}}
|   SPhrases = {{S1/2}} {{S26}} {{S30}} {{S45}}
  }}
| Section8 = {{Chembox Related
|   Function = [[strong acid]]s
|   OtherFunctn = [[Selenic acid]]<br/>[[Hydrochloric acid]]<br/>[[Nitric acid]]
|   OtherCpds = [[Sulfurous acid]]<br/>[[Peroxymonosulfuric acid]]<br/>[[Sulfur trioxide]]<br/>[[Oleum]]
  }}
}}

'''Sulfuric acid''' ('''sulphuric acid''' in British English) is a [[strong acid|strong]] [[mineral acid]] with the molecular formula {{chem|[[hydrogen|H]]|2}}{{chem|link=sulfate|SO|4}}. It is soluble in [[water]] at all concentrations. Sulfuric acid has many applications, and is one of the top products of the [[chemical industry]]. World production in 2001 was 165 million [[tonne]]s, with an approximate value of US$8 billion.  Principal uses include [[lead-acid battery|lead-acid batteries]] for cars and other vehicles, [[ore]] processing, [[fertilizer]] manufacturing, [[Oil refinery|oil refining]], [[wastewater processing]], and [[chemical synthesis]].

==Occurrence==
Pure sulfuric acid is not encountered naturally on Earth, due to its great [[Hygroscopy|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 the main byproduct produced when sulfur-containing fuels such as coal or oil are burned.

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]] (AMD) or acid rock drainage (ARD).  This acidic water is capable of dissolving metals present in sulfide ores, which results in brightly colored, toxic streams.  The oxidation of [[pyrite]] (iron sulfide) by molecular oxygen produces iron(II), or {{chem|Fe|2+}}:

:2 {{chem|FeS|2}} (s) + 7 {{chem|O|2}} + 2 {{chem|H|2|O}} → 2 {{chem|Fe|2+}} (aq) + 4 {{chem|SO|4|2-}} (aq) + 4 {{chem|H|+}}

The {{chem|Fe|2+}} can be further oxidized to {{chem|Fe|3+}}:

:4 {{chem|Fe|2+}} + {{chem|O|2}} + 4 {{chem|H|+}} → 4 {{chem|Fe|3+}} + 2 {{chem|H|2|O}}

The {{chem|Fe|3+}} produced can be precipitated as the [[hydroxide]] or [[hydrous iron oxides|hydrous oxide]]:

:{{chem|Fe|3+}} (aq) + 3 {{chem|H|2|O}} → {{chem|Fe(OH)|3}} (s) + 3 {{chem|H|+}}

The iron(III) ion ("ferric iron") can also oxidize pyrite:

:{{chem|FeS|2}} (s) + 14 {{chem|Fe|3+}} + 8 {{chem|H|2|O}} → 15 {{chem|Fe|2+}} (aq) + 2 {{chem|SO|4|2-}} (aq) + 16 {{chem|H|+}}

When iron(III) oxidation of pyrite occurs, the process can become rapid. [[pH]] values below zero have been measured in ARD produced by this process.

ARD can also produce sulfuric acid at a slower rate, so that the [[acid neutralizing capacity]] (ANC) of the aquifer can neutralize the produced acid. In such cases, the [[total dissolved solids]] (TDS) concentration of the water can be increased from the dissolution of minerals from the acid-neutralization reaction with the minerals.

===Extraterrestrial sulfuric acid===
====Venus====

Sulfuric acid is produced in the upper atmosphere of [[Venus]] by the [[Sun]]'s [[photochemistry|photochemical]] action on [[carbon dioxide]], [[sulfur dioxide]], and [[water]] vapor. [[Ultraviolet]] [[photon]]s of wavelengths less than 169&nbsp;[[nanometre|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 Venusian atmosphere, the result is [[sulfur trioxide]], which can combine with water vapor, another trace component of Venus's atmosphere, to yield sulfuric acid.  In the upper, cooler portions of Venus's atmosphere, sulfuric acid exists as a liquid, and thick sulfuric acid clouds completely obscure the planet's surface when viewed from above. The main cloud layer extends from 45&ndash;70&nbsp;km above the planet's surface, with thinner hazes extending as low as 30&nbsp;km and as high as 90&nbsp;km above the surface.  The permanent Venusian clouds produce a concentrated acid rain, as the clouds in the atmosphere of Earth produce water rain.

The atmosphere exhibits a sulfuric acid cycle. As sulfuric acid rain droplets fall down through the hotter layers of the atmosphere's temperature gradient, they are heated up and release water vapor, becoming more and more concentrated. When they reach temperatures above 300°C, sulfuric acid begins to decompose into sulfur trioxide and water, both in the gas phase. Sulfur trioxide is highly reactive and dissociates into sulfur dioxide and atomic oxygen, which oxidizes traces of carbon monoxide to form carbon dioxide.  Sulfur dioxide and water vapor rise on convection currents from the mid-level atmospheric layers to higher altitudes, where they will be transformed again into sulfuric acid, and the cycle repeats.

====Europa====

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. Sulfuric acid in solution with water causes significant [[freezing-point depression]] of water's [[melting point]], down to 210 K, and this would make more likely the existence of liquid solutions beneath Europa's icy crust.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|Wet sulfuric acid process}}

Sulfuric acid is produced from [[sulfur]], oxygen and water via the conventional [[contact process]] (DCDA) or the [[wet sulfuric acid process]] (WSA).

===Contact process (DCDA)===

In the first step, sulfur is burned to produce sulfur dioxide.
: S (s) + {{chem|O|2}} (g) → {{chem|SO|2}} (g)

This is then oxidized to sulfur trioxide using oxygen in the presence of a [[vanadium(V) oxide]] [[catalyst]].
: 2 {{chem|SO|2}} (g) + {{chem|O|2}} (g) → 2 {{chem|SO|3}} (g)  (in presence of {{chem|V|2|O|5}})

The sulfur trioxide is absorbed into 97-98% {{chem|H|2|SO|4}} to form [[oleum]] ({{chem|H|2|S|2|O|7}}), also known as fuming sulfuric acid.  The oleum is then diluted with water to form concentrated sulfuric acid.

: {{chem|H|2|SO|4}} (l) + {{chem|SO|3}} → {{chem|H|2|S|2|O|7}} (l)

: {{chem|H|2|S|2|O|7}} (l) + {{chem|H|2|O}} (l) → 2 {{chem|H|2|SO|4}} (l)

Note that directly dissolving {{chem|SO|3}} in water is not practical due to the highly [[Exothermic reaction|exothermic]] nature of the [[Chemical reaction|reaction]] between sulfur trioxide and water.  The reaction forms a corrosive aerosol that is very difficult to separate, instead of a liquid.

: {{chem|SO|3}} (g) + {{chem|H|2|O}} (l) → {{chem|H|2|SO|4}} (l)

===Wet sulfuric acid process (WSA)===
In the first step, sulfur is burned to produce sulfur dioxide:
: S(s) + {{chem|O|2}}(g) → {{chem|SO|2}}(g)

or, alternatively, [[hydrogen sulfide]] ({{chem|H|2|S}}) gas is incinerated to {{chem|SO|2}} gas:
: 2 {{chem|H|2|S}} + 3 {{chem|O|2}} → 2 {{chem|H|2|O}} + 2 {{chem|SO|2}} (−518&nbsp;kJ/mol)

This is then oxidized to sulfur trioxide using oxygen with [[vanadium(V) oxide]] as [[catalyst]].
: 2 {{chem|SO|2}} + {{chem|O|2}} → 2 {{chem|SO|3}} (−99&nbsp;kJ/mol)

The sulfur trioxide is hydrated into sulfuric acid {{chem|H|2|SO|4}}:
: {{chem|SO|3}} + {{chem|H|2|O}} → {{chem|H|2|SO|4}}(g)  (−101&nbsp;kJ/mol)

The last step is the condensation of the sulfuric acid to liquid 97&ndash;98% {{chem|H|2|SO|4}}:
: {{chem|H|2|SO|4}}(g) → {{chem|H|2|SO|4}}(l)  (−69&nbsp;kJ/mol)

===Other methods===
Another method is the less well-known metabisulfite method, in which metabisulfite in placed at the bottom of a beaker, and 12.6 molar concentration hydrochloric acid is added. The resulting gas is bubbled through nitric acid, which will release brown/red vapors. The completion of the reaction is indicated by the ceasing of the fumes. This method does not produce an inseparable mist, which is quite convenient.

Sulfuric acid can be produced in the laboratory by burning sulfur in air and dissolving the gas produced in a hydrogen peroxide solution.

: SO<sub>2</sub> + H<sub>2</sub>O<sub>2</sub> → H<sub>2</sub>SO<sub>4</sub>

Another method is to react hydrochloric acid with copper II sulfate:
: 2 HCl + CuSO<sub>4</sub> → H<sub>2</sub>SO<sub>4</sub> + CuCl<sub>2</sub>{{Citation needed|date=September 2010}}

Prior to 1900, most sulfuric acid was manufactured by the [[chamber process]].<ref>Edward M. Jones, "Chamber Process Manufacture of Sulfuric Acid", Industrial and Engineering Chemistry, Nov 1950, Vol 42, No. 11, pp 2208-10.</ref> As late as 1940, up to 50% of sulfuric acid manufactured in the United States was produced by chamber process plants.

==Physical properties==
===Grades of sulfuric acid===
Although nearly 100% sulfuric acid can be made, this loses {{chem|link=sulfur trioxide|SO|3}} at the boiling point to produce 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:<ref name="Columbia">{{citation | contribution = sulfuric  acid | url = http://www.encyclopedia.com/topic/sulfuric_acid.aspx | title = The Columbia Encyclopedia | edition = 6th | year = 2008 | accessdate = 2010-03-16}}.</ref><ref name="EB11">{{citation | contribution = Sulphuric acid | title = [[Encyclopædia Britannica Eleventh Edition|Encyclopædia Britannica]] | edition = 11th | year = 1910–1911 | volume = 26 | pages = 65–69}}.</ref>
{| class="wikitable"
|-
! Mass fraction<br/>H<sub>2</sub>SO<sub>4</sub>
! Density<br/>(kg/L)
! Concentration<br/>(mol/L)
! Common name
|-
| 10% || 1.07 || align=center | ~1 || dilute sulfuric acid
|-
| 29-32% || 1.25–1.28 || align=center | 4.2–5 || battery acid<br/>(used in [[Lead–acid battery|lead–acid batteries]])
|-
| 62–70% || 1.52–1.60 || align=center | 9.6–11.5 || ''chamber acid''<br/>''fertilizer acid''
|-
| 78–80% || 1.70–1.73 || align=center | 13.5–14 || ''tower acid''<br/>''Glover acid''
|-
| 95–98% || 1.83 || align=center | ~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 (<70% to avoid contamination with [[nitrosylsulfuric acid]]) and tower acid being the acid recovered from the bottom of the [[Glover tower]].<ref name="Columbia"/><ref name="EB11"/> They are now obsolete as commercial concentrations of sulfuric acid, although they may be prepared in the laboratory from concentrated sulfuric acid if needed. In particular, "10M" sulfuric acid (the modern equivalent of chamber acid, used in many [[titration]]s) is prepared by slowly adding 98% sulfuric acid to an equal volume of water, with good stirring: the temperature of the mixture can rise to 80&nbsp;°C (176&nbsp;°F) or higher.<ref name="EB11"/>

When high concentrations of {{chem|SO|3}} gas are added to sulfuric acid, {{chem|H|2|S|2|O|7}}, called ''[[pyrosulfuric acid]]'', ''fuming sulfuric acid'' or ''oleum'' or, less commonly, ''Nordhausen acid'', is formed. Concentrations of oleum are either expressed in terms of % {{chem|SO|3}} (called % oleum) or as % {{chem|H|2|SO|4}} (the amount made if {{chem|H|2|O}} were added); common concentrations are 40% oleum (109% {{chem|H|2|SO|4}}) and 65% oleum (114.6% {{chem|H|2|SO|4}}).  Pure {{chem|H|2|S|2|O|7}} is a solid with melting point 36°C.

Pure sulfuric acid is a viscous clear liquid, like oil, and this explains the old name of the acid ('oil of vitriol').

Commercial sulfuric acid is sold in several different purity grades.  Technical grade {{chem|H|2|SO|4}} 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.  [[Analytical chemistry|Analytical]] grades are also available.

===Polarity and conductivity===

[[Anhydrous]] {{chem|H|2|SO|4}} is a very [[Chemical polarity|polar]] liquid, having a [[Relative static permittivity|dielectric constant]] of around 100.  It has a high electrical conductivity, caused by dissociation through [[protonation|protonating]] itself, a process known as [[autoprotolysis]].<ref name = greenwood>{{Greenwood&Earnshaw}}</ref>
: 2 {{chem|H|2|SO|4}} {{eqm}} {{chem|H|3|SO|4|+}} + {{chem|HSO|4|-}}
The [[equilibrium constant]] for the autoprotolysis is<ref name = greenwood/>
:K<sub>ap</sub>(25°C)= [{{chem|H|3|SO|4|+}}][{{chem|HSO|4|-}}] =  {{val|2.7|e=-4}}.

The comparable equilibrium constant for [[Self-ionization of water|water]], K<sub>w</sub> is 10<sup>−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 {{chem|H|3|SO|4|+}} and {{chem|HSO|4|-}} 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% {{chem|H|2|SO|4}} contains the following species at equilibrium (figures shown as millimoles per kilogram of solvent): {{chem|HSO|4|-}} (15.0), {{chem|H|3|SO|4|+}} (11.3), {{chem|H|3|O|+}} (8.0), {{chem|HS|2|O|7|-}} (4.4), {{chem|H|2|S|2|O|7}} (3.6), {{chem|H|2|O}} (0.1).<ref name = greenwood/>

==Chemical properties==
===Reaction with water===
The [[hydration reaction]] of sulfuric acid is highly [[exothermic reaction|exothermic]]. One should always add the acid to the [[Properties of water|water]] rather than the water to the acid. Because the reaction is in an equilibrium that favors the rapid protonation of water, addition of acid to the water ensures that the ''acid'' is the limiting reagent.  This reaction is best thought of as the formation of [[hydronium]] ions:

: {{chem|H|2|SO|4}} + {{chem|H|2|O}} → {{chem|H|3|O|+}} + HSO<sub>4</sub><sup>−</sup>

: HSO<sub>4</sub><sup>−</sup> + {{chem|H|2|O}} → {{chem|H|3|O|+}} + SO<sub>4</sub><sup>2−</sup>

Because the hydration of sulfuric acid is [[thermodynamic]]ally favorable, sulfuric acid is an excellent dehydrating agent, and is used to prepare many dried fruits. 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]] {{chem|(C|6|H|12|O|6|)|''n''}} and concentrated sulfuric acid will give elemental [[carbon]] and water which is absorbed by the sulfuric acid (which becomes slightly diluted):

:({{chem|C|6|H|12|O|6}})''n'' → 6''n'' C + 6''n'' {{chem|H|2|O}}

The effect of this can be seen when concentrated sulfuric acid is spilled on paper; the cellulose reacts to give a [[combustion|burnt]] 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]] in a [[beaker (glassware)|beaker]]; a rigid column of black, porous carbon will quickly emerge. The carbon will smell strongly of [[caramel (aroma)|caramel]]. Although less dramatic, the action of the acid on [[cotton]], even in diluted form, will destroy the fabric. Clothes like [[jeans]] and [[labcoat]]s that accidentally come in contact with the acid will look perfect until they are received, in a barely recognizable state, from laundry.{{Citation needed|date=July 2010}}

===Other reactions===
As an acid, sulfuric acid reacts with most [[base (chemistry)|bases]] to give the corresponding [[sulfate]]. For example, the blue [[copper]] salt [[copper(II) sulfate]], commonly used for [[electroplating]] and as a [[fungicide]], is prepared by the reaction of [[copper(II) oxide]] with sulfuric acid:

: CuO (s) + {{chem|H|2|SO|4}} (aq) → {{chem|CuSO|4}} (aq) + {{chem|H|2|O}} (l)

Sulfuric acid can also be used to displace weaker acids from their salts. Reaction with [[sodium acetate]], for example, displaces [[acetic acid]], {{chem|CH|3|COOH}}, and forms [[sodium bisulfate]]:

:{{chem|H|2|SO|4}} + {{chem|CH|3|COONa}} → {{chem|NaHSO|4}} + {{chem|CH|3|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 {{chem|NO|2|+}}, 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 {{chem|H|2|SO|4}} attacks [[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 below is typical for most of these metals, but the reaction with tin produces sulfur dioxide rather than hydrogen.

: Fe (s) + {{chem|H|2|SO|4}} (aq) →  {{chem|H|2}} (g) + {{chem|FeSO|4}} (aq)

: Sn (s) + 2 {{chem|H|2|SO|4}} (aq) →  {{chem|SnSO|4}} (aq) + 2 {{chem|H|2|O}} (l) + {{chem|SO|2}} (g)

These reactions may be taken as typical: the hot concentrated acid generally acts as an oxidizing agent whereas the dilute acid acts a typical acid. Hence hot concentrated acid reacts with tin, zinc and copper to produce the salt, water and sulfur dioxide, whereas the dilute acid reacts with metals high in the reactivity series (such as Zn) to produce a salt and hydrogen. This is explained more fully in 'A New Certificate Chemistry' by Holderness and Lambert.

Sulfuric acid undergoes [[electrophilic aromatic substitution]] with [[aromatic compound]]s to give the corresponding [[sulfonic acid]]s:<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 = 27 January 2008}}</ref>

:[[Image:BenzeneSulfonation.png|300px]]

==Uses==
[[Image:2000sulphuric acid.PNG|thumb|right|Sulfuric acid production in 2000]]
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:

: {{chem|Ca|5|F(PO|4|)|3}} + 5 {{chem|H|2|SO|4}}  +  10 {{chem|H|2|O}}  → 5 {{chem|CaSO|4|·2 H|2|O}} + HF + 3 {{chem|H|3|PO|4}}

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]] (appliances) 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 ({{chem|SO|2}}) and sulfur trioxide ({{chem|SO|3}}) 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 [[aluminum sulfate]], also known as paper maker's alum. This can react with small amounts of soap on [[paper pulp]] fibers to give gelatinous aluminum [[carboxylate]]s, which help to coagulate the pulp fibers into a hard paper surface. It is also used for making [[aluminum 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:

: {{chem|Al|2|O|3}} + 3 {{chem|H|2|SO|4}} → {{chem|Al|2|(SO|4|)|3}} + 3 {{chem|H|2|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 [[cyclohexanone oxime]] to [[caprolactam]], used for making [[nylon]]. It is used for making [[hydrochloric acid]] from [[salt]] via the [[Mannheim process]]. Much {{chem|H|2|SO|4}} 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 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 {{chem|H|2|SO|4}} → 2 {{chem|SO|2}} + 2 {{chem|H|2|O}} + {{chem|O|2}} || &nbsp;&nbsp;&nbsp; || (830&nbsp;°C)
|-
| {{chem|I|2}} + {{chem|SO|2}} + 2 {{chem|H|2|O}} → 2 HI + {{chem|H|2|SO|4}} || &nbsp;&nbsp;&nbsp; || (120&nbsp;°C)
|-
| 2 HI → {{chem|I|2}} + {{chem|H|2}} || &nbsp;&nbsp;&nbsp; || (320&nbsp;°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 a large scale.

==History==
{{Disputed-section|date=November 2009}}
[[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 study of vitriol in [[Antiquity|ancient times]]. [[Sumerians]] had a list of types of vitriol that they classified according to substance's color. Some of the earliest discussions on the origin and properties of vitriol are 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".<ref name="Karpenko">Vladimir Karpenko, John A. Norris(2001), [http://www.chemicke-listy.cz/docs/full/2002_12_05.pdf Vitriol in the history of Chemistry], [[Charles University in Prague|Charles University]]</ref>

Iranian alchemists like [[Geber]], [[Rhazes]], Muhammad ibn Ibrahim al-Watwat, who included vitriol in their mineral classification lists. [[Avicenna]] focused on its medical uses. Several Indian alchemical works also mention the different varieties of vitriol.<ref name="Karpenko"/>

Sulfuric acid was discovered by medieval European alchemists. They called it "oil of vitriol". There are mentions to it in the works of [[Vincent of Beauvais]] and in the ''Compositum de Compositis'' ascribed to [[Albertus Magnus]]. A passage from [[Pseudo-Geber]]´s ''Summa Perfectionis'' was long considered to be the first recipe for sulphuric acid, but this was a misinterpretation.<ref name="Karpenko"/>

In the 17th century, the German-Dutch chemist [[Johann Glauber]] prepared sulfuric acid by burning sulfur together with [[Potassium nitrate|saltpeter]] ([[potassium nitrate]], {{chem|KNO|3}}), in the presence of steam. As saltpeter decomposes, it oxidizes the sulfur to {{chem|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 ''[[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.

Sulfuric acid created by John Roebuck's process only approached a 35&ndash;40% concentration.{{Citation needed|date=February 2009}} Later refinements to the lead-chamber process by French chemist [[Joseph-Louis Gay-Lussac]] and British chemist John Glover improved the yield to 78%.{{Citation needed|date=February 2009}}  However, the manufacture of some [[dye]]s and other chemical processes require a more concentrated product.{{Citation needed|date=February 2009}} 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]], {{chem|FeS|2}}) was heated in air to yield [[iron (II) sulfate]], {{chem|FeSO|4}}, which was oxidized by further heating in air to form [[iron(III) sulfate]], {{chem|Fe|2|(SO|4|)|3}}, which, when heated to 480&nbsp;°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 char a piece of tissue paper instantly]]
The corrosive properties of sulfuric acid are accentuated by its highly [[exothermic reaction]] with [[water (molecule)|water]]. 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 secondary thermal damage due to the heat liberated by the reaction with water.

The danger is greater with more concentrated preparations of sulfuric acid, but even the normal laboratory "dilute" grade (approximately 1&nbsp;M, 10%) will char paper by dehydration if left in contact for a sufficient time. Therefore, solutions equal to or stronger than 1.5&nbsp;M are labeled "CORROSIVE", while solutions greater than 0.5&nbsp;M but less than 1.5&nbsp;M are labeled "IRRITANT". Fuming sulfuric acid ([[oleum]]) is not recommended for use in schools as it is 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. Washing is continued for at least ten to fifteen minutes to cool the tissue surrounding the acid burn and to prevent secondary damage. Contaminated clothing is 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. The concentrated acid is always added to water and not the other way round, to take advantage of the relatively high [[heat capacity]] of water. Addition of water to concentrated sulfuric acid leads to the dispersal of a sulfuric acid [[particulate|aerosol]] or worse, an [[explosion]]. Preparation of solutions greater than 6&nbsp;M (35%) in concentration is most dangerous, as the heat produced may be sufficient to boil the diluted acid: efficient mechanical stirring and external cooling (such as an ice bath) are essential.

On a laboratory scale, sulfuric acid is advantageously diluted by pouring the concentrated version onto crushed ice. The ice used is sufficiently chemically pure so as not to interfere with the intended use of the diluted acid. Typically, about half the weight in ice is used for the volume of water (in ml) that would be needed for the solution. The same effect as when salt is added to ice occurs here: the ice melts in a endothermic process while dissolving the acid. The amount of heat needed to melt the ice in this process is greater than the amount of heat evolved by dissolving the acid: the solution becomes ice-cold. In contrast to a solution of 60<sup>o</sup>&nbsp;C, which is dangerously hot but which gives no clues on its temperature, a solution of -10<sup>o</sup>&nbsp;C warnes because of its icy cover. After all the ice has melted, the rest of the needed water is added as water to the cold solution.

===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. 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&nbsp;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 US 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>

==References==
{{Reflist}}
{{Refbegin}}
A New Certificate Chemistry by A Holderness and J Lambert, Heinemann 1976.
* 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==
{{Commons category|Sulfuric acid}}
* {{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]

{{Use dmy dates|date=August 2010}}

{{DEFAULTSORT:Sulfuric Acid}}
[[Category:Sulfur oxoacids]]
[[Category:Sulfates]]
[[Category:Mineral acids]]
[[Category:Equilibrium chemistry]]
[[Category:Alchemical substances]]
[[Category:DEA List II chemicals]]
[[Category:Photographic chemicals]]
[[Category:Oxidizing agents]]
[[Category:Oxidizing acids]]

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