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Sulfuric acid
File:Sulfuric-acid-2D-dimensions.svg
File:Sulfuric-acid-3D-vdW.png
IUPAC name Sulfuric acid
Other names Oil of vitriol
style="background: #F8EABA; text-align: center;" colspan="2" | Identifiers
CAS number	7664-93-9 Y
EC number	231-639-5
UN number	1830
RTECS number	WS5600000
style="background: #F8EABA; text-align: center;" colspan="2" | Properties
Molecular formula	H2O4S
Molar mass	98.08 g mol−1
Appearance	Clear, colorless, odorless liquid
Density	1.84 g/cm3, liquid
Melting point	10 °C, 283 K, 50 °F
Boiling point	337 °C, 610 K, 639 °F
Solubility in water	miscible
Acidity (pKa)	−3
Viscosity	26.7 cP (20 °C)
style="background: #F8EABA; text-align: center;" colspan="2" | Hazards
MSDS	ICSC 0362
EU Index	016-020-00-8
EU classification	Corrosive (C)
R-phrases	R35
S-phrases	(S1/2) S26 S30 S45
NFPA 704	75px 0 3 2 W
Flash point	Non-flammable
style="background: #F8EABA; text-align: center;" colspan="2" | Related compounds
Related strong acids	Selenic acid Hydrochloric acid Nitric acid
Related compounds	Sulfurous acid Peroxymonosulfuric acid Sulfur trioxide Oleum
Y (what is this?)  (verify) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Sulfuric acid (sulphuric acid in British English) is a strong mineral acid with the molecular formula H₂SO₄. 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 tonnes, with an approximate value of US$8 billion. Principal uses include lead-acid batteries for cars and other vehicles, ore processing, fertilizer manufacturing, oil refining, wastewater processing, and chemical synthesis.

Occurrence

Pure sulfuric acid is not encountered naturally on Earth, due to its great affinity for water. Apart from that, sulfuric acid is a constituent of acid rain, which is formed by atmospheric oxidation of sulfur dioxide in the presence of 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 Fe²⁺:

2 FeS₂ (s) + 7 O₂ + 2 H₂O → 2 Fe²⁺ (aq) + 4 SO2−4 (aq) + 4 H⁺

The Fe²⁺ can be further oxidized to Fe³⁺:

4 Fe²⁺ + O₂ + 4 H⁺ → 4 Fe³⁺ + 2 H₂O

The Fe³⁺ produced can be precipitated as the hydroxide or hydrous oxide:

Fe³⁺ (aq) + 3 H₂O → Fe(OH)₃ (s) + 3 H⁺

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

FeS₂ (s) + 14 Fe³⁺ + 8 H₂O → 15 Fe²⁺ (aq) + 2 SO2−4 (aq) + 16 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 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 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 45–70 km above the planet's surface, with thinner hazes extending as low as 30 km and as high as 90 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 mission show distinct absorptions on Jupiter's 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.^[1]

Manufacture

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) + O₂ (g) → SO₂ (g)

This is then oxidized to sulfur trioxide using oxygen in the presence of a vanadium(V) oxide catalyst.

2 SO₂ (g) + O₂ (g) → 2 SO₃ (g) (in presence of V₂O₅)

The sulfur trioxide is absorbed into 97-98% H₂SO₄ to form oleum (H₂S₂O₇), also known as fuming sulfuric acid. The oleum is then diluted with water to form concentrated sulfuric acid.

H₂SO₄ (l) + SO₃ → H₂S₂O₇ (l)

H₂S₂O₇ (l) + H₂O (l) → 2 H₂SO₄ (l)

Note that directly dissolving SO₃ in water is not practical due to the highly exothermic nature of the reaction between sulfur trioxide and water. The reaction forms a corrosive aerosol that is very difficult to separate, instead of a liquid.

SO₃ (g) + H₂O (l) → H₂SO₄ (l)

Wet sulfuric acid process (WSA)

In the first step, sulfur is burned to produce sulfur dioxide:

S(s) + O₂(g) → SO₂(g)

or, alternatively, hydrogen sulfide (H₂S) gas is incinerated to SO₂ gas:

2 H₂S + 3 O₂ → 2 H₂O + 2 SO₂ (−518 kJ/mol)

This is then oxidized to sulfur trioxide using oxygen with vanadium(V) oxide as catalyst.

2 SO₂ + O₂ → 2 SO₃ (−99 kJ/mol)

The sulfur trioxide is hydrated into sulfuric acid H₂SO₄:

SO₃ + H₂O → H₂SO₄(g) (−101 kJ/mol)

The last step is the condensation of the sulfuric acid to liquid 97–98% H₂SO₄:

H₂SO₄(g) → H₂SO₄(l) (−69 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.

SO₂ + H₂O₂ → H₂SO₄

Another method is to react hydrochloric acid with copper II sulfate:

2 HCl + CuSO₄ → H₂SO₄ + CuCl₂^{[citation needed]}

Prior to 1900, most sulfuric acid was manufactured by the chamber process.^[2] 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 SO₃ 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:^[3][4]

Mass fraction H2SO4	Density (kg/L)	Concentration (mol/L)	Common name
10%	1.07	~1	dilute sulfuric acid
29-32%	1.25–1.28	4.2–5	battery acid (used in lead–acid batteries)
62–70%	1.52–1.60	9.6–11.5	chamber acid fertilizer acid
78–80%	1.70–1.73	13.5–14	tower acid Glover acid
95–98%	1.83	~18	concentrated sulfuric acid

"Chamber acid" and "tower acid" were the two concentrations of sulfuric acid produced by the lead chamber process, chamber acid being the acid produced in lead chamber itself (<70% to avoid contamination with nitrosylsulfuric acid) and tower acid being the acid recovered from the bottom of the Glover tower.^[3][4] 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 titrations) 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 °C (176 °F) or higher.^[4]

When high concentrations of SO₃ gas are added to sulfuric acid, H₂S₂O₇, called pyrosulfuric acid, fuming sulfuric acid or oleum or, less commonly, Nordhausen acid, is formed. 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 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 H₂SO₄ is impure and often colored, but is suitable for making fertilizer. Pure grades such as United States Pharmacopoeia (USP) grade are used for making pharmaceuticals and dyestuffs. Analytical grades are also available.

Polarity and conductivity

Anhydrous H₂SO₄ is a very polar liquid, having a dielectric constant of around 100. It has a high electrical conductivity, caused by dissociation through protonating itself, a process known as autoprotolysis.^[5]

2 H₂SO₄ 15px H₃SO+4 + HSO−4

The equilibrium constant for the autoprotolysis is^[5]

K_ap(25°C)= [H₃SO+4][HSO−4] = 2.7×10⁻⁴.

The comparable equilibrium constant for water, K_w is 10⁻¹⁴, a factor of 10¹⁰ (10 billion) smaller.

In spite of the viscosity of the acid, the effective conductivities of the H₃SO+4 and 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 equilibrium is actually more complex than shown above; 100% H₂SO₄ contains the following species at equilibrium (figures shown as millimoles per kilogram of solvent): HSO−4 (15.0), H₃SO+4 (11.3), H₃O⁺ (8.0), HS₂O−7 (4.4), H₂S₂O₇ (3.6), H₂O (0.1).^[5]

Chemical properties

Reaction with water

The hydration reaction of sulfuric acid is highly exothermic. One should always add the acid to the 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:

H₂SO₄ + H₂O → H₃O⁺ + HSO₄⁻

HSO₄⁻ + H₂O → H₃O⁺ + SO₄²⁻

Because the hydration of sulfuric acid is thermodynamically favorable, 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 remove hydrogen and oxygen atoms from 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 → 6n C + 6n H₂O

The effect of this can be seen when concentrated sulfuric acid is spilled on paper; the cellulose reacts to give a 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; a rigid column of black, porous carbon will quickly emerge. The carbon will smell strongly of caramel. Although less dramatic, the action of the acid on cotton, even in diluted form, will destroy the fabric. Clothes like jeans and labcoats that accidentally come in contact with the acid will look perfect until they are received, in a barely recognizable state, from laundry.^{[citation needed]}

Other reactions

As an acid, sulfuric acid reacts with most 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) + H₂SO₄ (aq) → CuSO₄ (aq) + H₂O (l)

Sulfuric acid can also be used to displace weaker acids from their salts. Reaction with sodium acetate, for example, displaces acetic acid, CH₃COOH, and forms sodium bisulfate:

H₂SO₄ + CH₃COONa → NaHSO₄ + CH₃COOH

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+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 H₂SO₄ 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) + H₂SO₄ (aq) → H₂ (g) + FeSO₄ (aq)

Sn (s) + 2 H₂SO₄ (aq) → SnSO₄ (aq) + 2 H₂O (l) + SO₂ (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 compounds to give the corresponding sulfonic acids:^[6]

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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.^[7] 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 fertilizers 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:

Ca₅F(PO₄)₃ + 5 H₂SO₄ + 10 H₂O → 5 CaSO₄·2 H₂O + HF + 3 H₃PO₄

Sulfuric acid is used in large quantities by the iron and steelmaking 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 (SO₂) and sulfur trioxide (SO₃) 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 carboxylates, 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 out impurities, as well as to improve the taste of the water. Aluminum sulfate is made by reacting bauxite with sulfuric acid:

Al₂O₃ + 3 H₂SO₄ → Al₂(SO₄)₃ + 3 H₂O

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 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 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 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 H2SO4 → 2 SO2 + 2 H2O + O2		(830 °C)
I2 + SO2 + 2 H2O → 2 HI + H2SO4		(120 °C)
2 HI → I2 + H2		(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-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

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This section's factual accuracy is disputed. (November 2009)

The study of vitriol in 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".^[8]

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.^[8]

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.^[8]

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 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 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 35–40% concentration.^{[citation needed]} 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]} However, the manufacture of some dyes and other chemical processes require a more concentrated product.^{[citation needed]} 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. 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

The corrosive properties of sulfuric acid are accentuated by its highly exothermic reaction with 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 M, 10%) will char paper by dehydration if left in contact for a sufficient time. Therefore, solutions equal to or stronger than 1.5 M are labeled "CORROSIVE", while solutions greater than 0.5 M but less than 1.5 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 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 aerosol or worse, an explosion. Preparation of solutions greater than 6 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^o C, which is dangerously hot but which gives no clues on its temperature, a solution of -10^o 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 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 axons 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, 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.^[9]

In the US sulfuric acid is included in List II of the 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.^[10][11][12]

References

External links

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Wikimedia Commons has media related to Sulfuric acid.

• International Chemical Safety Card 0362
• NIOSH Pocket Guide to Chemical Hazards
• External Material Safety Data Sheet
• Sulfuric acid analysis - titration freeware

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1. ↑ T.M. Orlando, T.B. McCord, G.A Grieves, Icarus 177 (2005) 528–533
2. ↑ Edward M. Jones, "Chamber Process Manufacture of Sulfuric Acid", Industrial and Engineering Chemistry, Nov 1950, Vol 42, No. 11, pp 2208-10.
3. ↑ ^{3.0 3.1} "sulfuric acid", The Columbia Encyclopedia (6th ed.), 2008, retrieved 2010-03-16 .
4. ↑ ^{4.0 4.1 4.2} "Sulphuric acid", Encyclopædia Britannica, 26 (11th ed.), 1910–1911, pp. 65–69 .
5. ↑ ^{5.0 5.1 5.2} Greenwood, Norman N.; Earnshaw, A. (1997), Chemistry of the Elements (2nd ed.), Oxford: Butterworth-Heinemann, ISBN 0080379419 CS1 maint: Multiple names: authors list (link)
6. ↑ F. A. Carey. "Reactions of Arenes. Electrophilic Aromatic Substitution". On-Line Learning Center for Organic Chemistry. University of Calgary. Retrieved 27 January 2008. 
7. ↑ Chenier, Philip J. Survey of Industrial Chemistry, pp 45-57. John Wiley & Sons, New York, 1987. ISBN.
8. ↑ ^{8.0 8.1 8.2} Vladimir Karpenko, John A. Norris(2001), Vitriol in the history of Chemistry, Charles University
9. ↑ Annex to Form D ("Red List"), 11th Edition, January 2007 (pg. 4). International Narcotics Control Board. Vienna, Austria; 2007.
10. ↑ 66 FR 52670—52675. 17 October 2001.
11. ↑ 21 CFR 1309
12. ↑ 21 USC, Chapter 13 (Controlled Substances Act)