Sulfuric acid: Difference between revisions
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'''Sulfuric acid''' ([[British English]]: '''sulphuric acid'''), [[hydrogen|H]]<sub>2</sub>[[sulfur|S]][[oxygen|O]]<sub>4</sub>, is a strong [[mineral acid]]. It is soluble in [[water]] at all concentrations. The old name for sulfuric acid was '''Zayt al-Zaj''', or '''oil of [[vitriol]]''', coined by the 8th-century |
'''Sulfuric acid''' ([[British English]]: '''sulphuric acid'''), [[hydrogen|H]]<sub>2</sub>[[sulfur|S]][[oxygen|O]]<sub>4</sub>, is a strong [[mineral acid]]. It is soluble in [[water]] at all concentrations. The old name for sulfuric acid was '''Zayt al-Zaj''', or '''oil of [[vitriol]]''', coined by the 8th-century [[alchemy|alchemist]] [[Geber|Jabir ibn Hayyan]], the chemical's probable discoverer{{Ref|4}}. Sulfuric acid has many applications, and is produced in greater amounts than any other chemical besides [[water]]. World production in 2001 was 165 million [[tonne]]s, with an approximate value of $8 billion. Principal uses include [[fertilizer]] manufacturing, ore processing, chemical synthesis, [[wastewater]] processing, and oil refining. |
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== Physical properties == |
== Physical properties == |
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== History of sulfuric acid == |
== History of sulfuric acid == |
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The discovery of sulfuric acid is credited to the 8th centuary [[alchemy|alchemist]] [[Geber|Jabir ibn Hayyan]]. It was studied later by the 9th century |
The discovery of sulfuric acid is credited to the 8th centuary [[alchemy|alchemist]] [[Geber|Jabir ibn Hayyan]]. It was studied later by the 9th century physician and [[alchemy|alchemist]] [[Abu Bakr Mohammad Ibn Zakariya al-Razi|Ibn Zakariya al-Razi]] (Rhases), who obtained the substance by [[Dry distillation | dry distillation]] of minerals including [[iron(II) sulfate]] heptahydrate, FeSO<sub>4</sub> • 7H<sub>2</sub>O, called green [[vitriol]], and [[copper(II) sulfate]] pentahydrate, CuSO<sub>4</sub> • 5H<sub>2</sub>O, called blue vitriol. 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 treatises and books by European alchemists, such as the 13th-century German [[Albertus Magnus]]. For this reason, sulfuric acid was known to medieval European alchemists as ''oil of vitriol'' and ''spirit of vitriol'', among other names. |
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In the 17th century, the German-Dutch chemist [[Johann Glauber]] prepared sulfuric acid by burning [[sulfur]] together with [[saltpeter]] ([[potassium nitrate]], KNO<sub>3</sub>), in the presence of steam. As the 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 the 17th century, the German-Dutch chemist [[Johann Glauber]] prepared sulfuric acid by burning [[sulfur]] together with [[saltpeter]] ([[potassium nitrate]], KNO<sub>3</sub>), in the presence of steam. As the 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. |
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Revision as of 11:41, 4 April 2006
Sulfuric acid (British English: sulphuric acid), H2SO4, is a strong mineral acid. It is soluble in water at all concentrations. The old name for sulfuric acid was Zayt al-Zaj, or oil of vitriol, coined by the 8th-century alchemist Jabir ibn Hayyan, the chemical's probable discoverer[1]. Sulfuric acid has many applications, and is produced in greater amounts than any other chemical besides water. World production in 2001 was 165 million tonnes, with an approximate value of $8 billion. Principal uses include fertilizer manufacturing, ore processing, chemical synthesis, wastewater processing, and oil refining.
Physical properties
Forms of sulfuric acid
Although 100% sulfuric acid can be made, this loses SO3 at the boiling point to produce 98.3% acid. The 98% grade is also more stable for storage, making it the usual form for "concentrated" sulfuric acid. Other concentrations of sulfuric acid are used for different purposes. Some common concentrations are:
- 10%, dilute sulfuric acid for laboratory use
- 33.5%, battery acid (used in lead-acid batteries)
- 62.18%, chamber or fertilizer acid
- 77.67%, tower or Glover acid
- 98%, concentrated
Different purities are also available. Technical grade H2SO4 is impure and often colored, but it is suitable for making fertiliser. Pure grades such as US Pharmacopoeia (USP) grade are used for making pharmaceuticals and dyestuffs.
When high concentrations of SO3(g) are added to sulfuric acid, H2S2O7 forms. This is called fuming sulfuric acid or oleum or, less commonly, Nordhausen acid. Concentrations of oleum are either expressed in terms of % SO3 (called % oleum) or as "% H2SO4 (the amount made if H2O were added); common concentrations are 40% oleum (109% H2SO4) and 65% oleum (114.6% H2SO4). Pure H2S2O7 is in fact a solid, melting point 36 °C.
Polarity and conductivity
Anhydrous H2SO4 is a very polar liquid, with a dielectric constant of around 100. This is due to the fact that it can dissociate by protonating itself, a process known as autoprotolysis,[2] which occurs to a high degree, more than 10 billion times the level seen in water:
- 2 H2SO4 ⇌ H3SO4+ + HSO4−
This allows protons to be highly mobile in H2SO4. It also makes sulfuric acid an excellent solvent for many reactions. In fact, the equilibrium is more complex than shown above. 100% H2SO4 contains the following species at equilibrium (figures shown as mmol per kg solvent): HSO4− (15.0), H3SO4+ (11.3), H3O+ (8.0), HS2O7− (4.4), H2S2O7 (3.6), H2O (0.1).
Chemical properties
Reaction with water
The hydration reaction of sulfuric acid is highly exothermic. If water is added to concentrated sulfuric acid, it can boil and spit dangerously. One should always add the acid to the water rather than the water to the acid. This can be remembered through mnemonics such as "Do as you oughta: add acid to water", "A.A.: Add Acid", or "Drop acid, not water." 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 float above the acid. The reaction is best thought of as forming hydronium ions, as such so:
- H2SO4 + H2O → H3O+ + HSO4-
And then:
- HSO4- + H2O → H3O+ + SO42-
Because the hydration of sulfuric acid is thermodynamically favorable (ΔH = -880 kJ/mol), sulfuric acid is an excellent dehydrating agent, and is used to prepare many dried fruits. The affinity of sulfuric acid for water is sufficiently strong that it will take hydrogen and oxygen atoms out of other compounds; for example, mixing starch (C6H12O6)n and concentrated sulfuric acid will give elemental carbon and water which is absorbed by the sulfuric acid (which becomes slightly diluted): (C6H12O6)n → 6C + 6H2O. The effect of this can be seen when concentrated sulphuric acid spilled on paper; the starch reacts to give a burned appearance, the carbon appears as soot would in a fire. A more dramatic illustration ocurs when sulfuric acid is added to a tablespoon of white sugar in a cup when a tall rigid column of black porous carbon smelling strongly of caramel emerges from the cup.
Other reactions of sulfuric acid
As an acid, sulfuric acid reacts with most bases to give the corresponding sulfate. For example, copper(II) sulfate, the familiar blue salt of copper used for electroplating and as a fungicide, is prepared by the reaction of copper(II) oxide with sulfuric acid:
Sulfuric acid can be used to displace weaker acids from their salts, for example sodium acetate gives acetic acid:
H2SO4 + CH3COONa → NaHSO4 + CH3COOH
Likewise the reaction of sulfuric acid with potassium nitrate can be used to produce nitric acid, along with a precipitate of potassium bisulfate. With nitric acid itself, sulfuric acid acts as both an acid and a dehydrating agent, forming the nitronium ion NO2+, 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 reactions in organic chemistry, such as Fischer esterification and dehydration of alcohols.
Sulfuric acid reacts with most metals in a single displacement reaction to produce hydrogen gas and the metal sulfate. Dilute H2SO4 attacks iron, aluminium, zinc, manganese and nickel, but tin and copper require hot concentrated acid. Lead and tungsten are, however, 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.
Environmental aspects
Sulfuric acid is a constituent of acid rain, being formed by atmospheric oxidation of sulfur dioxide in the presence of water. Sulfur dioxide is the main product when sulfur-containing fuels such as coal or oil are burned together.
Extraterrestial sulfuric acid
Sulfuric acid is produced in the upper atmosphere of Venus by the sun's photochemical action on carbon dioxide, sulfur dioxide, and water vapor. Ultraviolet photons of wavelengths less than 169 nm can photodissociate carbon dioxide into carbon monoxide and atomic oxygen. Atomic oxygen is highly reactive; when it reacts with with sulfur dioxide, a trace component of the Venusian atomsphere, the result is sulfur trioxide, which can combine with water vapor, another trace component of Venus' atomsphere, to yield sulfuric acid.
In the upper, cooler portions of Venus' atmosphere, sulfuric acid can exist as a liquid, and thick sulfuric acid clouds completely obscure the planet's surface from above. The main cloud layer extends from 45–70 km above the planet's surface, with thinner hazes extending as low as 30 and as high as 90 km above the surface.
History of sulfuric acid
The discovery of sulfuric acid is credited to the 8th centuary alchemist Jabir ibn Hayyan. It was studied later by the 9th century physician and alchemist Ibn Zakariya al-Razi (Rhases), who obtained the substance by dry distillation of minerals including iron(II) sulfate heptahydrate, FeSO4 • 7H2O, called green vitriol, and copper(II) sulfate pentahydrate, CuSO4 • 5H2O, called blue vitriol. When heated, these compounds decompose to iron(II) oxide and copper(II) oxide, respectively, giving off water and sulfur trioxide, which combine to produce a dilute solution of sulfuric acid. This method was popularized in Europe through translations of Arabic treatises and books by European alchemists, such as the 13th-century German Albertus Magnus. For this reason, sulfuric acid was known to medieval European alchemists as oil of vitriol and spirit of vitriol, among other names.
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. Later refinements in the lead-chamber process by the French chemist Joseph-Louis Gay-Lussac and the British chemist John Glover 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 iron (II) sulfate, FeSO4, 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. The expense of this process prevented the large-scale use of concentrated sulfuric acid.
In 1831, the British vinegar merchant Peregrine Phillips patented a far more economical process for producing sulfur trioxide and concentrated sulfuric acid, now known as the contact process. Essentially all of the world's supply of sulfuric acid is now produced by this method.
Manufacture
- Main Article: 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) Template:Sulfur(s) + O2(g) → SO2(g)
This is oxidised to sulfur trioxide using oxygen in the presence of a vanadium(V) oxide catalyst.
Finally the sulfur trioxide is treated with water (usually as 97-98% H2SO4 containing 2-3% water) to produce 98-99% sulfuric acid.
Note that directly dissolving SO3 in water is impractical due to the highly exothermic nature of the reaction. Mists are formed instead of a liquid. Alternatively, the SO3 is absorbed into H2SO4 to produce oleum (H2S2O7), which is then diluted to form sulfuric acid.
- (3) H2SO4(l) + SO3 → H2S2O7(l)
Oleum is reacted with water to form concentrated H2SO4.
- (4) H2S2O7(l) + H2O(l) → 2 H2SO4(l)
In 1903, US production of sulfuric acid amounted to 346754323.237748 million tonnes. World production in 242 was 4367647264761 million tonnes.
Uses
Sulfuric acid is a very important commodity chemical, and indeed a nation's sulfuric acid production is a good indicator of its industrial strength.[3] The major use (60% of total worldwide) for sulfuric acid is in the "wet method" for the production of phosphoric acid, used for manufacture of phosphate fertilisers as well as trisodium phosphate for detergents. In this method phosphate rock is used, and more than 100 million tonnes is 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:
- Ca5F(PO4)3 + 5 H2SO4 + 10 H2O → 5 CaSO4·2 H2O + HF + 3 H3PO4
Sulfate fertilisers such as ammonium sulfate are manufactured using sulfuric acid, although in smaller quantities than phosphates.
Another important use for sulfuric acid is for the manufacture of aluminium sulfate, also known as papermaker's alum. This can react with small amounts of soap on paper pulp fibres to give gelatinous aluminium carboxylates, which help to coagulate the pulp fibres into a hard paper surface. It is also used for making aluminium hydroxide, which is used at water treatment plants to filter out impurities, as well as to improve the taste of the water. Aluminium sulfate is made by reacting bauxite with sulfuric acid:
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 H2SO4 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 dyestuffs.
A mixture of sulfuric acid and water is used as the electrolyte in various types of lead-acid battery where it undergoes a reversible reaction where lead and lead dioxide are converted to lead(II) sulfate. Sulfuric acid is also the principal ingredient in some drain cleaners, used to clear blockages consisting of paper, rags, and other materials not easily dissolved by caustic solutions.
Sulfuric acid is also used as a general dehydrating agent in its concentrated form. See Reaction with water.
Safety
Laboratory hazards
The corrosive properties of sulfuric acid are accentuated by its highly exothermic reaction with 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 length of time. The standard first aid treatment for acid spills on the skin is, as for other corrosive agents, irrigation with large quantities of water: in the case of sulfuric acid it is important that the acid should be removed before washing, as a further heat burn could result from the exothermic dilution of the acid. Washing should be continued for a sufficient length of time—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 heat capacity of water. Addition of water to concentrated sulfuric acid leads at best to the dispersal of a sulfuric acid 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. Water should not be used as the extinguishing agent because of the risk of further dispersal of aerosols: carbon dioxide is preferred where possible.
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 mg/m3: limits in other countries are similar.
Comic rhyme
Sulfuric acid is one of the few compounds whose chemical formula is well known by the general public because of many comic rhymes, such as this one popular in the UK:
- Johnny was a chemist's son, but Johnny is no more.
- What Johnny thought was H2O was H2SO4.
In the U.S., a common variant is:
- Little Johnny took a drink, but he shall drink no more.
- For what he thought was H2O was H2SO4.
Source
Institut National de Recherche et de Sécurité. (1997). "Acide sulfurique". Fiche toxicologique n°30, Paris: INRS, 5 pp.
References
- ^ Chenier, Philip J. Survey of Industrial Chemistry, pp 45-57. John Wiley & Sons, New York, 1987. ISBN 0471010774.
- ^ Greenwood, N.N. and A. Earnshaw. Chemistry of the Elements, pp 837-845. Pergamon Press, Oxford, UK, 1984. ISBN 0080220576.
- ^ Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
- ^ Khairallah, Amin A. Outline of Arabic Contributions to Medicine, chapter 10. Beirut, 1946.