sulfuric acid

Structural formula
General
Surname	sulfuric acid
other names	Sulfur (VI) acid Dihydrogen sulfate Monothionic acid E  513 Vitriol Oil (outdated) SULFURIC ACID ( INCI )
Molecular formula	H 2 SO 4
Brief description	colorless and odorless, viscous liquid
External identifiers / databases
CAS number 7664-93-9 EC number 231-639-5 ECHA InfoCard 100,028,763 PubChem 1118 Wikidata Q4118
properties
Molar mass	98.08 g mol −1
Physical state	liquid
density	1.84 g cm −3
Melting point	10.94  ° C (100%)
boiling point	290 ° C anhydrous acid approx. 335 ° C 98% acid (azeotrope)
Vapor pressure	<0.1 Pa
pK s value	−3 (H 2 SO 4 ), −6.62 1.9 (HSO 4 - ), 1.99
solubility	completely miscible with water
safety instructions
GHS hazard labeling from  Regulation (EC) No. 1272/2008 (CLP) , expanded if necessary danger H and P phrases H: 290-314 P: 280-301 + 330 + 331-305 + 351 + 338-308 + 310
MAK	DFG / Switzerland: 0.1 mg m −3 (measured as inhalable dust / aerosol fraction)
Toxicological data	2140 mg kg −1 ( LD 50 ,  rat ,  oral , 25% solution)
As far as possible and customary, SI units are used. Unless otherwise noted, the data given apply to standard conditions .

Sulfuric acid is a chemical compound of sulfur with the empirical formula H ₂ SO ₄ . It is a colorless, oily, very viscous and hygroscopic liquid . Sulfuric acid is one of the strongest acids and is very corrosive . This mineral acid forms two series of salts , the hydrogen sulfates and the sulfates , in which one or two protons are replaced by cations compared to the free acid .

Sulfuric acid is one of the most important technical chemicals and is one of the most widely produced basic chemical substances . In 1993 about 135 million tons of sulfuric acid were produced. It is mainly used in fertilizer production and for the production of other mineral acids, such as hydrochloric or phosphoric acid . Aqueous solutions of various concentrations are mostly used.

The anhydride of sulfuric acid is sulfur trioxide (SO ₃ ). The solution of sulfur trioxide in the sulfuric acid beyond the stoichiometric ratio is called fuming sulfuric acid or oleum , since the sulfur trioxide contained easily escapes from the solution and forms fog (“smoke”) from dilute sulfuric acid with the humidity. Related acids are sulfurous acid (H ₂ SO ₃ ), which is derived from sulfur dioxide and thiosulfuric acid (H ₂ S ₂ O ₃ ), in which one oxygen atom is replaced by sulfur.

history

Galley stove for the extraction of vitriol oil

Sulfuric acid has been known for a long time under the outdated name of vitriol oil . The first clues can be found in the texts of the historically controversial alchemist Jābir ibn Hayyān from the 8th century. After that, possible manufacturing processes are also mentioned in the alchemical writings of Albertus Magnus (1200–1280) and Basilius Valentinus (around 1600). These processes describe how vitriol oil can be obtained from naturally occurring sulphates such as chalcanthite or alum . The name vitriol oil is derived from the outdated term vitriol for these minerals. The first source of large amounts of sulfuric acid was iron vitriol . From the 16th century, sulfuric acid was produced in Bohemia , Saxony and the Harz Mountains using the vitriol method. After the first production site in Nordhausen , the product was called Nordhäuser Vitriol. Johann Rudolph Glauber carried out the first scientific investigations with sulfuric acid . He let the acid act on table salt and received the hydrochloric acid and the Glauber's salt named after him, sodium sulfate .

The processes in which sulfates were used, however, were very complex and expensive. In order to obtain larger quantities, a process was developed in the 18th century in which sulfur and saltpetre were burned in glass vessels. Since the glass vessels were very fragile, the reaction was carried out in lead containers for the first time in 1746 by John Roebuck . In 1778 the first chemical factory in Switzerland was founded in Winterthur with the laboratory , which manufactured vitriol oil as the main product. After Nicolas Clément-Désormes and Charles-Bernard Desormes discovered in 1793 that the use of air could significantly reduce the amount of saltpetre, the lead chamber process could be used on an industrial scale. This was particularly important for the Leblanc process for soda production invented by Nicolas Leblanc in 1789 and first used by him in 1791 . The process was improved several times, for example by the development of methods for absorbing the nitrous gases by Joseph Louis Gay-Lussac . Continuous production management could thus be achieved.

The greatest disadvantage of this process was that only a maximum acid concentration of 78% could be achieved and more concentrated solutions and oleum still had to be produced by the laborious distillation of iron vitriol. A simple production of higher concentrated sulfuric acid was only possible after the development of the contact process from 1870 by Rudolph Messel in England.

Occurrence

Chalcanthite

Free sulfuric acid , which is not dissociated into oxonium and sulfate ions , occurs only very rarely in nature. In the atmosphere , it is formed from sulfur dioxide , which is produced when substances containing sulfur are burned or during volcanic eruptions. The sulfur dioxide is hydroxyl - radicals and oxygen to sulfur oxidized. The free sulfuric acid is finally formed with water. Other oxidizing agents that enable the formation of sulfur trioxide are ozone or hydrogen peroxide . In acid rain it then reaches the earth in the form of dilute acid.

A small amount of free sulfuric acid is also found in some volcanic springs called solfataras .

In contrast to the free acid, its salts, especially the sulfates, are much more common in nature. There are many different sulfate minerals. The best known and most important include gypsum (CaSO ₄  · 2 H ₂ O), barite (BaSO ₄ ), chalcanthite (CuSO ₄  · 5 H ₂ O) or Glauber's salt (Na ₂ SO ₄  · 10 H ₂ O).

Outside the Earth, sulfuric acid is found in the upper atmosphere of Venus . This is caused by photochemical reactions between sulfur dioxide and water. Droplets are formed that contain 80–85% sulfuric acid. In deeper layers the acid decomposes due to the high temperatures into sulfur dioxide, oxygen and water, which rise again and can form sulfuric acid.

Extraction and manufacture

The basic material for the production of sulfuric acid is often elemental sulfur , which is obtained in large quantities (2007: 66 million tons) during the desulphurisation of natural gas and crude oil and is processed using the Claus process or broken down using the Frasch process . The sulfur is burned in order to obtain sulfur dioxide as a starting material for the actual representation.

${\displaystyle {\ce {S + O2 -> SO2}}}$

Reaction of sulfur with oxygen

Another source that produces large amounts of sulfur dioxide is the smelting of sulfur-containing ores . Examples of this are the extraction of copper , zinc or lead from the corresponding sulphides . The sulfur dioxide is formed when roasting with atmospheric oxygen.

${\displaystyle {\ce {2ZnS + 3O2 -> 2ZnO + 2SO2}}}$

Reaction when roasting zinc sulfide

In 1999, three million tons of pyrite were roasted in Europe for sulfuric acid production. In Asia, however, the proportion of pyrite is higher.

Rotary kilns of the gypsum acid plant in the chemical combine Bitterfeld

For countries that are poor in raw materials and have neither sulfur nor sulphidic ores, the production of gypsum sulfuric acid using the Müller-Kühne process is an option . Here, sulfur dioxide is extracted from gypsum and coal in a rotary kiln . The energy-intensive process can be made more profitable if cement is produced as a by-product through the addition of sand and clay . In the GDR , the process was carried out on a large scale.

For further production, sulfur trioxide must be obtained from the sulfur dioxide . The direct reaction of sulfur and oxygen to form sulfur trioxide takes place only to a small extent, since the equilibrium in the reaction of sulfur dioxide to sulfur trioxide is on the side of the sulfur trioxide only at low temperatures. At these temperatures, however, the reaction rate is too slow. Therefore, with the aid of suitable catalysts, the reaction must be controlled in such a way that a sufficiently rapid reaction is ensured at temperatures that are not too high.

${\displaystyle {\ce {2SO2 + O2 <=> 2SO3}}}$

Reaction of sulfur dioxide to sulfur trioxide

Catalytic cycle in the oxidation of sulfur dioxide

In the contact process , which is still used exclusively , vanadium pentoxide is used as the oxygen-transferring catalyst. A salt melt is formed from vanadium (V) oxide and alkali metal sulfates added as co-catalysts. The reactive complex with the composition [(VO) ₂ O (SO ₄ ) ₄ ] ⁴⁻ , which acts as the actual catalyst, is formed in this . Oxygen and sulfur dioxide accumulate on these without changing the oxidation number of the vanadium and react to form sulfur trioxide.

The temperature during the reaction must be between 420 and 620 ° C, since at lower temperatures the catalyst becomes inactive due to the formation of vanadium (IV) compounds and it decomposes at higher temperatures. The reaction is carried out in so-called tray contact ovens , in which the catalyst is arranged in a total of four layers (the "trays") one above the other and the gas flowing through between the trays is cooled to the appropriate temperature.

In the so-called double contact process, the sulfur trioxide present is washed out with concentrated sulfuric acid before the last tray. This enables the yield to be increased to at least 99.8% (First General Administrative Regulation for the Federal Immission Control Act, Technical Instructions for Keeping the Air Clean 2002).

After formation of the sulfur trioxide, it is converted to sulfuric acid. To do this, any remaining sulfur dioxide must first be removed with ammonia or sodium thiosulfate . Since the direct reaction of sulfur trioxide with water is too slow, the gas is passed into concentrated sulfuric acid. Disulfuric acid H ₂ S ₂ O ₇ is quickly formed . If this is diluted with water, it breaks down into two molecules of sulfuric acid.

${\displaystyle {\ce {SO3 + H2SO4 -> H2S2O7}}}$

Conversion of sulfur trioxide with sulfuric acid

${\displaystyle {\ce {H2S2O7 + H2O -> 2H2SO4}}}$

Formation of sulfuric acid

This process does not produce pure sulfuric acid, but rather concentrated acid with 98% acid content. In order to produce pure sulfuric acid, the amount of sulfur trioxide must be introduced into the concentrated acid, which corresponds to the amount of substance of the excess water of the concentrated acid.

In the last few decades, sulfuric acid production has risen sharply, especially in China , while production has declined in European countries such as Germany . Since the beginning of 2000, China has been dependent on additional quantities from Europe. The sharp upheavals in the years 1990 and 1991 can be traced back to the dissolution of the Soviet Union and a change in statistics in the United States .

In the large-scale industrial production of sulfuric acid, it is of considerable economic importance that the three individual steps are exothermic (for values ​​see the contact method ). The amount of heat released is used to generate high-pressure steam for electricity generation and for industrial heating purposes.

properties

Physical Properties

Sulfuric acid molecule with bond lengths

Anhydrous sulfuric acid is a viscous (cross-linked by hydrogen bonds), colorless liquid with a high density (1.8269 g / cm ³ ) that solidifies below 10.371 ° C. The melting point is greatly reduced by small amounts of water and is 3.0 ° C for a 98% sulfuric acid, for example. The common light brown color of technical sulfuric acid is due to organic impurities that are carbonized by dehydration. Above the boiling point of anhydrous sulfuric acid of 279.6 ° C, sulfuric acid vapors are formed which contain excess sulfur trioxide, with the water remaining in the boiling sulfuric acid. The anhydrous sulfuric acid is converted to a 98.33% sulfuric acid with a constant boiling point of 338 ° C. At this temperature, the steam also has an acid content of 98.33% and thus corresponds to an azeotropic mixture of water and sulfuric acid. An acid of the same composition and the same boiling point is obtained if dilute acid is distilled. 100% sulfuric acid cannot therefore be obtained by distilling dilute sulfuric acid, but only by dissolving a certain amount of sulfur trioxide in concentrated sulfuric acid. If the temperature is higher than 338 ° C, sulfuric acid decomposes into water and sulfur trioxide ("smoking off the sulfuric acid") and is almost completely dissociated at 450 ° C.

As a solid, sulfuric acid crystallizes in the monoclinic crystal system in space group C 2 / c (space group no. 15) . The lattice parameters are a  = 814  pm , b  = 470 pm, c  = 854 pm and β  = 111 °. The structure is a corrugated layer structure in which each dihydrogen sulfate tetrahedron is connected to four other tetrahedra via hydrogen bonds . In addition to the crystalline pure sulfuric acid, several sulfuric acid hydrates are known. An example is the dihydrate H ₂ SO ₄ · 2 H ₂ O, which also crystallizes monoclinically with the space group C 2 / c (No. 15) . A total of six different hydrates with one, two, three, four, six and eight equivalents of water are known, in which the acid is completely split into oxonium and sulfate ions. The oxonium ions are associated with a different number of water molecules, depending on the hydrate. The melting point of these hydrates decreases as the number of water molecules increases. The monohydrate melts at 8.59 ° C, while the octahydrate melts at -62 ° C.

Strong hydrogen bonds act between the individual molecules, which cause the high viscosity of 24.6 mPa · s at 25 ° C. In comparison, water has a significantly lower viscosity of 0.89 mPas at 25 ° C.

Similar to pure water, pure sulfuric acid conducts electricity to a small extent . The specific conductivity is 1.044 · 10 ⁻² S / cm. The reason for this lies in the low dissociation of the acid by autoprotolysis . Diluted acid, on the other hand, conducts electricity well because of the oxonium ions it contains .

${\displaystyle {\ce {2H2SO4 <=> HSO4- + H3SO4+}}}$

Autoprotolysis reaction

more likely attachment situation

Sulfuric acid tetrahedron

Individual sulfuric acid molecules are present in the gas phase . These have a tetrahedral structure with bond angles of 101.3 ° between the OH groups and 123.3 ° between the oxygen atoms. The bond lengths of the sulfur-oxygen bonds differ with 157.4 pm (to OH groups) and 142.2 pm (to the oxygen atoms). The molecular structure in the solid corresponds to that in the gas phase.

The bonds in the sulfuric acid molecule can be described by various mesomeric boundary structures. For example, the structure in which double bonds are assumed between sulfur and oxygen or in which only single bonds and at the same time a charge separation exist. In theoretical calculations , it has been found that the d orbitals , which are necessary for the description of an OS double bond, contribute very little to binding. Therefore, the real bond situation in the sulfuric acid molecule is most precisely described by the structure in which only single bonds are drawn. The shortened S – O bond can be explained by additional electrostatic interactions between the charged atoms.

Chemical properties

Hägg diagram of sulfuric acid - H ₂ SO ₄ : black; HSO ₄ ^- : violet; SO ₄ ²⁻ : turquoise; H ⁺ : dashed; OH ^- : dotted

As a very strong acid, sulfuric acid easily releases protons. With a pK _s value of −3.0 (this only applies to dilute solutions) or, more precisely, an H ₀ value of −11.9, sulfuric acid is one of the strong acids in the first protolysis stage .

${\displaystyle {\ce {H2SO4 + H2O -> HSO4- + H3O+}}}$

Reaction with water in the first protolysis stage

It is usually not counted among the super acids , but it is chosen as the starting point for the definition of super acids: All acids that are stronger than pure sulfuric acid and can thus protonate are referred to as super acids.

The second protolysis of hydrogen sulfate to sulfate has a pK _s value of 1.9. The hydrogen sulfate ion is therefore only a moderately strong acid.

${\displaystyle {\ce {HSO4- + H2O -> SO4^{2}- + H3O+}}}$

Reaction with water in the second protolysis stage

For this reason, hydrogen sulfate is mostly present in dilute sulfuric acid (concentration approx. 1 mol / l). The H ₂ SO ₄ molecule is almost completely dissociated , while the reaction to the sulfate takes place only to a small extent (about 1.3% at 1 mol / l). Larger amounts of sulphate are only formed at higher dilutions.

Carbonization of paper by concentrated sulfuric acid

Sulfuric acid has a high affinity for water. If acid and water are mixed, various hydrates of the form H ₂ SO ₄ · n H ₂ O ( n = 1–4, 6, 8) are formed with strong heat development . The strong water affinity of sulfuric acid is also expressed in the fact that it is able to split off hydroxyl groups and protons from organic substances . As a result of this withdrawal, carbon remains, the organic matter turns black and charred. This effect occurs mainly with substances that contain many hydroxyl groups. Examples are many carbohydrates such as glucose or polysaccharides . Furthermore, the great affinity for water can be used for condensation reactions. The water is removed from an organic compound without carbonization . An example of this is the synthesis of 2-pyrone .

Another indication of the strong hygroscopicity is that the acid dehydrates itself to a small extent:

${\displaystyle {\ce {2H2SO4 <=> H3O+ + HS2O7-}}}$

Self-drainage of the sulfuric acid

Concentrated sulfuric acid has an oxidizing effect and is able to dissolve more noble metals such as copper , mercury or silver when heated . The sulfuric acid is reduced to sulfur dioxide. In contrast, even pure, base iron is not attacked by passivation of concentrated sulfuric acid.

${\displaystyle {\ce {Cu + 2H2SO4 -> CuSO4 + SO2 + 2H2O}}}$

Dissolve copper in concentrated sulfuric acid

Dilute sulfuric acid, on the other hand, has only a slight oxidative effect, since the reaction to sulfur dioxide is inhibited by the solvent water . Only those metals are oxidized or dissolved which, as base elements , can be oxidized to hydrogen by the reaction of protons .

use

Sulfuric acid production from 1970 to 2003 in various countries

Sulfuric acid is used in very large quantities and in many areas. Alongside that of chlorine , their production volume is a benchmark for industrial development and a country's level of performance.

It is called differently depending on the concentration. Between 10% and 20% it is called dilute sulfuric acid or dilute acid . Battery acid or battery acid has an acid concentration of 33.5%. These acids remain liquid even below 0 ° C.

Sulfuric acid with a content of up to about 70% is called chamber acid , up to 80% gloversic acid . Concentrated sulfuric acid has a content of at least 98.3% ( azeotrope ). Dilute acid occurs in large quantities as a waste product in the titanium oxide or dye production .

Most of it is used in the production of fertilizers . With the help of sulfuric acid, mainly phosphate and ammonium sulphate fertilizers are obtained. The latter is represented by the reaction of half-concentrated sulfuric acid with ammonia .

${\displaystyle {\ce {H2SO4 + 2NH3 -> (NH4)2SO4}}}$

Reaction of sulfuric acid with ammonia

In the production of phosphate fertilizers , sulfuric acid is required to break down the rock phosphate. The reaction produces superphosphate Ca (H ₂ PO ₄ ) ₂ / CaSO ₄

${\displaystyle {\ce {Ca3(PO4)2 + 2H2SO4 -> Ca(H2PO4)2 + 2CaSO4}}}$

Decomposition of apatite to superphosphate by half-concentrated sulfuric acid

Titanium (IV) oxide is produced on a large scale with the help of sulfuric acid

In addition to ammonium sulfate, other sulfates are also produced by reacting appropriate salts with sulfuric acid. One example is aluminum sulphate obtained from aluminum hydroxide , which is used in large quantities in the paper industry and as a flocculant in water purification.

Since numerous ores are soluble in sulfuric acid, it can be used as a digesting agent. Examples are the wet process for producing zinc from zinc oxide and the sulphate process for obtaining the white pigment titanium dioxide . With the help of sulfuric acid, not only oxidic ores, but also those with other anions such as fluoride or phosphate can be digested. The corresponding acids are formed during the reaction. This process is relevant for the production of some technically important acids. Examples are hydrofluoric acid from fluorite , phosphoric acid from apatite and hydrochloric acid from halite . As a battery acid, sulfuric acid is an important component of the lead accumulator used in automobiles as a starter battery . As in the lead-acid battery, dilute sulfuric acid also serves as an electrolyte in electrolytic processes . The advantages over other electrolytes are their high conductivity and, at the same time, their low tendency to reduce .

In organic chemistry , the sulfonic acid group can be inserted through fuming sulfuric acid ( sulfonation ). It is mainly used to produce surfactants for the detergent industry and dyes . Another functional group that can be introduced with the help of sulfuric acid is the nitro group . This is done with the help of so-called nitrating acid , a mixture of sulfuric and nitric acid . This is mainly used to manufacture explosives such as trinitrotoluene or nitroglycerin .

Sulfuric acid is one of the most commonly used chemicals in chemical laboratories . Besides hydrochloric acid and nitric acid, it is a strong acid that is widely used. It is used, among other things, to adjust the pH value , as a catalyst , for example for esterification and for smoking during digestions. The strong dehydrating effect of sulfuric acid is used to dry organic substances and gases in desiccators and washing bottles .

Biological importance

The sulfuric acid formed in the air from sulfur dioxide is, in addition to the nitric acid formed from nitrogen oxides , a component of acid rain . The acid rain can lead to a drop in the pH value , especially in weakly buffered soils and bodies of water . One effect of a lower pH is a change in the solubility of some metal ions. Aluminum , which is harmful to plants, is more soluble in water at a lower pH value. Biologically important ions such as potassium or magnesium can also be washed out more easily. For these reasons, sulfuric acid is considered a possible cause of forest dieback in the 1980s. Thanks to technical measures such as flue gas desulphurization in coal-fired power plants and the introduction of low-sulfur fuels, so little sulfur dioxide is released in Germany that rainwater contains significantly less sulfuric acid.

Sulfuric acid has a toxic effect on fish and other aquatic life due to its acidity . In soft water without a buffer capacity, the mean lethal concentration (the LC ₅₀ value) for fish is 100–330 mg / l, similar to that of other mineral acids.

In the spoil heaps of ore mines and lignite opencast mines, sulfuric acid is produced by a combination of abiotic and microbial oxidation of exposed sulfide-containing minerals. It is washed out by rainwater and collects as acid pit water in residual lakes in which, due to the low pH value and high heavy metal content, hardly any living organisms can be found.

safety instructions

Sulfuric acid is extremely irritating and corrosive to the skin and mucous membranes . It is able to destroy living tissue ( chemical burns ). The mechanisms of action of concentrated and dilute sulfuric acid must be clearly distinguished. With dilute sulfuric acid, the increased proton concentration has a corrosive effect, i. H. the effect is similar to that of other dilute acids. Upon skin contact, the effect is mainly local irritation, depending on the concentration. It is therefore significantly safer than concentrated sulfuric acid. Due to its strong water-attracting effect, this has a charring effect and damages the skin and eyes even in small amounts. Healing, painful wounds are slow to form. Sulfuric acid can also be absorbed via vapors from the air; the MAK value is 0.1 mg / m ³ , the LC ₅₀ value in rats is 510 mg / m ³ inhaled over four hours .

Since the reaction of concentrated sulfuric acid with water generates a lot of heat, it must only be diluted by pouring it into water and not by adding water to the acid. If water is added to sulfuric acid, it can splash and burn bystanders.

proof

Concentrated sulfuric acid is detected by reacting with organic substances. For example, if a piece of wood is immersed in concentrated sulfuric acid, it slowly turns black. It is possible to distinguish between dilute and concentrated sulfuric acid by different reactions. The different reactivity of the two acids with base metals such as zinc or iron is exploited. While dilute acid forms hydrogen at room temperature, the concentrated acid, which contains almost no free oxonium ions, only reacts when heated to form sulfur dioxide and sulfur .

Since the sulfuric acid is dissociated in aqueous solution, it cannot be directly detected in it. Instead, the proton concentration and thus the acidic pH value can be determined using suitable indicators or a pH meter . The sulfate ion can be determined, for example, by precipitation as sparingly soluble barium sulfate .

literature

• AF Holleman , E. Wiberg , N. Wiberg : Textbook of Inorganic Chemistry . 102nd edition. Walter de Gruyter, Berlin 2007, ISBN 978-3-11-017770-1 .
• Norman N. Greenwood, Alan Earnshaw: Chemistry of the Elements . 1st edition, Wiley-VCH, Weinheim, 1988, ISBN 3-527-26169-9 .
• Thomas M. Klapötke , IC Tornieporth-Oetting: Non-metal chemistry . Wiley-VCH, Weinheim, 1994, ISBN 3-527-29052-4
• Entry to sulfuric acid. In: Römpp Online . Georg Thieme Verlag, accessed on May 26, 2014.
• Hermann Müller: Sulfuric Acid and Sulfur Trioxide. In: Ullmann's Encyclopedia of Industrial Chemistry , 4th edition, Wiley-VCH, 1984, doi : 10.1002 / 14356007.a25_635
• Fritz Ullmann: Encyclopedia of Technical Chemistry. Edition 1922
• Directory of systems installed worldwide

Web links

Wiktionary: sulfuric acid  - explanations of meanings, word origins, synonyms, translations

Commons : Sulfuric Acid  - Collection of pictures, videos and audio files

Wikisource: The Specific Gravity of Sulfuric Acid at Different Levels of Dilution  - Sources and Full Texts

Individual evidence

1. ↑ Entry on E 513: Sulphuric acid in the European database for food additives, accessed on June 27, 2020.
2. ↑ Entry on SULFURIC ACID in the CosIng database of the EU Commission, accessed on August 11, 2020.
3. ↑ ^{a b c d e f g h i j k l m n} Entry on sulfuric acid in the GESTIS substance database of the IFA , accessed on January 10, 2017(JavaScript required) .
4. ↑ ^{a b c d} R. Williams: pKa Data. In: chem.wisc.edu (PDF, 78 kB).
5. ^ ^{A b} Organic Chemistry Michigan State University
6. ↑ Entry on Sulphuric acid in the Classification and Labeling Inventory of the European Chemicals Agency (ECHA), accessed on February 1, 2016. Manufacturers or distributors can expand the harmonized classification and labeling .
7. ↑ Swiss Accident Insurance Fund (Suva): Limit values ​​- current MAK and BAT values (search for 7664-93-9 or sulfuric acid ), accessed on November 2, 2015.
8. ↑ ^{a b} data sheet 96% sulfuric acid (PDF) from Carl Roth , accessed on April 23, 2010.
9. ↑ ^{a b} Entry on sulfuric acid. In: Römpp Online . Georg Thieme Verlag, accessed on May 26, 2014.
10. ^ Sulfuric acid in: Meyers Großes Konversations-Lexikon . Leipzig 1905-1909, Volume 18, pp. 160-163.
11. ↑ Claus Priesner: Johann Christian Bernhardt and the vitriolic acid , in: Chemistry in our time , 1982, 16, 5, pp. 149–159; doi: 10.1002 / ciuz.19820160504 .
12. ^ Walter Jansen: In memory: Johann Rudolph Glauber , in: Chemikon , 2004, 11, 3, p. 155; doi: 10.1002 / ckon.200490022 .
13. ↑ Barbara Winter-Werner: The cradle of Swiss chemistry was in Winterthur. Platform Chemistry, Swiss Academy of Sciences. 12/2009, pp. 8-9.
14. ^ William H. Brock: Viewegs Geschichte der Chemie. Vieweg, Braunschweig 1997, ISBN 3-540-67033-5 .
15. ↑ Hermann Müller: Sulfuric Acid and Sulfur Trioxide , in: Ullmann's Encyclopedia of Industrial Chemistry , 4th edition, 1984, doi : 10.1002 / 14356007.a25_635 .
16. ^ ^{A b} N. N. Greenwood, A. Earnshaw: Chemistry of the elements. 1st edition. VCH Verlagsgesellschaft, 1988, ISBN 3-527-26169-9 .
17. ↑ VA Krasnopolsky, VA Parshev: Chemical composition of the atmosphere of Venus , in: Nature , 1981, 292, pp. 610-613; doi: 10.1038 / 292610a0 .
18. ↑ US Geological Survey, Mineral Commodity Summaries: Sulfur, January 2008 (PDF; 85 kB)
19. ^ ER Riegel and JA Kent: Riegel's Handbook of Industrial Chemistry. Springer 2003. ISBN 0-306-47411-5 , p. 503.
20. ↑ Hans Kühne: Development of the gypsum-sulfuric acid process according to Müller-Kühne . In: Chemical Engineer Technology - CIT . tape 21 , no. June 11-12 , 1949, pp. 227-229 , doi : 10.1002 / cite.330211105 ( PDF ).
21. ↑ ^{a b} O.B. Lapina, BS Bal'zhinimaev, S. Boghosian, KM Eriksen, R. Fehrmann: Progress on the mechanistic understanding of SO ₂ oxidation catalysts , in: Catalysis Today , 1999, 51, pp. 469–479, doi: 10.1016 / S0920- 5861 (99) 00034-6 .
22. ↑ Additional needs in China
23. ↑ Industrial Commodity Statistics Yearbook of the United Nations Statistics Division ( Memento of April 7, 2013 in the Internet Archive )
24. ^ ^{A b c d} A. F. Holleman , E. Wiberg , N. Wiberg : Textbook of Inorganic Chemistry . 102nd edition. Walter de Gruyter, Berlin 2007, ISBN 978-3-11-017770-1 .
25. ↑ ^{a b} Robert L. Kuczkowski, RD Suenram, and Frank J. Lovas: Microwave Spectrum, Structure, and Dipole Moment of Sulfuric Acid , in: J. Am. Chem. Soc. , 1981, 103, pp. 2561-2566; doi: 10.1021 / ja00400a013 .
26. ^ AF Holleman: Textbook of Inorganic Chemistry . Walter de Gruyter GmbH & Co KG, 2019, ISBN 978-3-11-083817-6 , pp. 511 ( limited preview in Google Book search).
27. ↑ C. Pascard-Billy: Structure Précise de l'Acide Sulfurique , in: Acta Cryst. , 1965, 18, pp. 827-829; doi: 10.1107 / S0365110X65002037 .
28. ↑ Inger Taesler, Ivar Olovsson: Hydrogen Bond Studies. XXXVII. The Crystal Structure of Sulfuric Acid Dihydrate (H ₃ O ⁺ ) ₂ SO ₄ ²⁻ , in: J. Chem. Phys. , 1969, 51, p. 4213; doi: 10.1063 / 1.1671781 .
29. ↑ ^{a b} Th. M. Klapötke, IC Tornieporth-Oetting: Nichtmetallchemie , Wiley-VCH, Weinheim, 1994, ISBN 3-527-29052-4 .
30. ↑ Thorsten Stefan and Rudolf Janoschek: How relevant are S = O and P = O Double Bonds for the Description of the Acid Molecules H ₂ SO ₃ , H ₂ SO ₄ , and H ₃ PO ₄ , respectively? in: J. Mol. Model. , 2000, 6, pp. 282-288 doi: 10.1007 / PL00010730
31. ↑ Hans Breuer: General and inorganic chemistry. dtv-Atlas Chemie. Vol. 1. dtv, Munich 2000 (9th edition), ISBN 3-423-03217-0 .
32. ↑ Dilute acid. ( Memento from January 21, 2012 in the Internet Archive ) In: Umweltlexikon-online.de
33. ↑ Acid rain. ( Memento from June 11, 2012 in the Internet Archive ) In: Umweltlexikon-online.de , as of May 8, 2001.
34. ^ Soil - inputs from the atmosphere. ( Memento of October 9, 2010 in the Internet Archive ) In: Federal Environment Agency , as of April 2, 2008.
35. ↑ ^{a b} Jander, Blasius: Introduction to the inorganic-chemical internship , 14th edition, S. Hirzel Verlag, Leipzig 1995, ISBN 3-7776-0672-3 .

This article was added to the list of excellent articles on April 9, 2008 in this version .