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'''Sulfuric acid''' ([[Sulfur#Spelling and etymology|alternative spelling]] |
'''Sulfuric acid''' ([[Sulfur#Spelling and etymology|alternative spelling]] ) is a [[strong acid|strong]] [[mineral acid]] with the molecular formula {{chem|[[hydrogen|H]]|2}}{{chem|[[sulfur|S]][[oxygen|O]]|4}}. Its historical name is oil of '''vitriol'''. Pure sulfuric acid is a highly corrosive, colorless, viscous liquid. The [[salt (chemistry)|salts]] of sulfuric acid are called [[sulfate]]s. Sulfuric acid is soluble in [[water]] at all concentrations. |
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Sulfuric acid has many applications, and is a central substance in the [[chemical industry]]. 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]]. |
Sulfuric acid has many applications, and is a central substance in the [[chemical industry]]. 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]]. |
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Revision as of 17:34, 23 September 2011
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| Names | |
|---|---|
| IUPAC name
Sulfuric acid
| |
| Other names
Oil of vitriol
| |
| Identifiers | |
3D model (JSmol)
|
|
| ChEBI | |
| ChEMBL | |
| ChemSpider | |
| ECHA InfoCard | 100.028.763 |
| EC Number |
|
| E number | E513 (acidity regulators, ...) |
| KEGG | |
| RTECS number |
|
| UNII | |
| UN number | 1830 |
CompTox Dashboard (EPA)
|
|
| |
| |
| Properties | |
| H 2SO 4 | |
| Molar mass | 98.079 g/mol |
| Appearance | Clear, colorless, odorless liquid |
| Density | 1.84 g/cm3, liquid |
| Melting point | 10 °C (50 °F; 283 K) |
| Boiling point | 337 °C (639 °F; 610 K) |
| miscible | |
| Acidity (pKa) | −3, 1.99 |
| Viscosity | 26.7 cP (20 °C) |
| Hazards | |
| NFPA 704 (fire diamond) | |
| Flash point | Non-flammable |
| Related compounds | |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
| |
Sulfuric acid (alternative spelling ) is a strong mineral acid with the molecular formula H
2SO
4. Its historical name is oil of vitriol. Pure sulfuric acid is a highly corrosive, colorless, viscous liquid. The salts of sulfuric acid are called sulfates. Sulfuric acid is soluble in water at all concentrations.
Sulfuric acid has many applications, and is a central substance in the chemical industry. Principal uses include lead-acid batteries for cars and other vehicles, ore processing, fertilizer manufacturing, oil refining, wastewater processing, and chemical synthesis.
History
The study of vitriol began 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".[1]
Islamic alchemists Jabir Ibn Hayyan (c. 721 – c. 815 AD), Al-Razi (865 – 925 AD), and Jamal Din al-Watwat (d. 1318, wrote the book Mabāhij al-fikar wa-manāhij al-'ibar), included vitriol in their mineral classification lists. Ibn Sina focused on its medical uses and different varieties of vitriol.[1]
Sulfuric acid was called "oil of vitriol" by medieval European alchemists. 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 sulfuric acid, but this was a misinterpretation.[1]
In the 17th century, the German-Dutch chemist Johann Glauber prepared sulfuric acid by burning sulfur together with saltpeter (potassium nitrate, KNO
3), in the presence of steam. As saltpeter decomposes, it oxidizes the sulfur to 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 35–40% 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, FeS
2) was heated in air to yield iron (II) sulfate, FeSO
4, which was oxidized by further heating in air to form iron(III) sulfate, Fe
2(SO
4)
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.
Physical properties
Grades of sulfuric acid
Although nearly 99% sulfuric acid can be made, this loses 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:[2][3]
| 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.[2][3] 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.[3]
When high concentrations of SO
3 gas are added to sulfuric acid, H
2S
2O
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 % SO
3 (called % oleum) or as % H
2SO
4 (the amount made if H
2O were added); common concentrations are 40% oleum (109% H
2SO
4) and 65% oleum (114.6% H
2SO
4). Pure H
2S
2O
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 H
2SO
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 pharmaceuticals and dyestuffs. Analytical grades are also available.
Polarity and conductivity
Anhydrous H
2SO
4 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.[4]
- 2 H
2SO
4 ⇌ H
3SO+
4 + HSO−
4
The equilibrium constant for the autoprotolysis is[4]
- Kap(25°C)= [H
3SO+
4][HSO−
4] = 2.7×10−4.
The comparable equilibrium constant for water, Kw is 10−14, a factor of 1010 (10 billion) smaller.
In spite of the viscosity of the acid, the effective conductivities of the H
3SO+
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% H
2SO
4 contains the following species at equilibrium (figures shown as millimoles per kilogram of solvent): HSO−
4 (15.0), H
3SO+
4 (11.3), H
3O+
(8.0), HS
2O−
7 (4.4), H
2S
2O
7 (3.6), H
2O (0.1).[4]
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:
- H
2SO
4 + H
2O → H
3O+
+ HSO4− K1 = 2.4 x 106 (strong acid)
- HSO4− + H
2O → H
3O+
+ SO42− K2 = 1.0 x 10−2 [5]
HSO4- is the bisulfate anion and SO42- is the sulfate anion. K1 and K2 are the acid dissociation constants. Because the hydration of sulfuric acid is thermodynamically favorable, sulfuric acid is an excellent dehydrating agent. 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 (C
6H
12O
6)
n and concentrated sulfuric acid will give elemental carbon and water which is absorbed by the sulfuric acid (which becomes slightly diluted):
- (C
6H
12O
6)n → 6n C + 6n H
2O
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.[6] The carbon will smell strongly of caramel due to the heat generated. Although less dramatic, the action of the acid on cotton, even in diluted form, will destroy the fabric.
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) + H
2SO
4 (aq) → CuSO
4 (aq) + H
2O (l)
Sulfuric acid can also be used to displace weaker acids from their salts. Reaction with sodium acetate, for example, displaces acetic acid, CH
3COOH, and forms sodium bisulfate:
- H
2SO
4 + CH
3COONa → NaHSO
4 + CH
3COOH
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.
Concentrated sulfuric acid reacts with sodium chloride, and gives hydrogen chloride gas and sodium bisulfate:
- NaCl + H2SO4 → NaHSO4 + HCl
As mentioned above, concentrated sulfuric acid is a powerful dehydrating agent, removing water from sugar and other carbohydrates, to produce carbon, heat, steam, and a more dilute acid containing increased amounts of hydronium and bisulfate ions.
- (CH2O)n + Sulfuric acid → C (graphitic foam) + steam + Sulfuric acid/water mixture
Sulfuric acid reacts with most metals via a single displacement reaction to produce hydrogen gas and the metal sulfate. Dilute H
2SO
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) + H
2SO
4 (aq) → H
2 (g) + FeSO
4 (aq)
- Sn (s) + 2 H
2SO
4 (aq) → SnSO
4 (aq) + 2 H
2O (l) + 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.
Benzene undergoes electrophilic aromatic substitution with sulfuric acid to give the corresponding sulfonic acids:[7]
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 FeS
2 (s) + 7 O
2 + 2 H
2O → 2 Fe2+
(aq) + 4 SO2−
4 (aq) + 4 H+
The Fe2+
can be further oxidized to Fe3+
:
- 4 Fe2+
+ O
2 + 4 H+
→ 4 Fe3+
+ 2 H
2O
The Fe3+
produced can be precipitated as the hydroxide or hydrous oxide:
- Fe3+
(aq) + 3 H
2O → Fe(OH)
3 (s) + 3 H+
The iron(III) ion ("ferric iron") can also oxidize pyrite:
- FeS
2 (s) + 14 Fe3+
+ 8 H
2O → 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 (−63 °C), 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.[8]
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) + O
2 (g) → SO
2 (g)
This is then oxidized to sulfur trioxide using oxygen in the presence of a vanadium(V) oxide catalyst. This reaction is reversible and the formation of the sulfur trioxide is exothermic.
- 2 SO
2 (g) + O
2 (g) ⇌ 2 SO
3 (g) (in presence of V
2O
5)
The sulfur trioxide is absorbed into 97–98% H
2SO
4 to form oleum (H
2S
2O
7), also known as fuming sulfuric acid. The oleum is then diluted with water to form concentrated sulfuric acid.
- H
2SO
4 (l) + SO
3 → H
2S
2O
7 (l)
- H
2S
2O
7 (l) + H
2O (l) → 2 H
2SO
4 (l)
Note that directly dissolving SO
3 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.
- SO
3 (g) + H
2O (l) → H
2SO
4 (l)
In the first step, sulfur is burned to produce sulfur dioxide:
- S(s) + O
2(g) → SO
2(g)
or, alternatively, hydrogen sulfide (H
2S) gas is incinerated to SO
2 gas:
- 2 H
2S + 3 O
2 → 2 H
2O + 2 SO
2 (−518 kJ/mol)
This is then oxidized to sulfur trioxide using oxygen with vanadium(V) oxide as catalyst.
- 2 SO
2 + O
2 → 2 SO
3 (−99 kJ/mol) (reaction is reversible)
The sulfur trioxide is hydrated into sulfuric acid H
2SO
4:
- SO
3 + H
2O → H
2SO
4(g) (−101 kJ/mol)
The last step is the condensation of the sulfuric acid to liquid 97–98% H
2SO
4:
- H
2SO
4(g) → H
2SO
4(l) (−69 kJ/mol)
Other methods
Another method is the less well-known metabisulfite method, in which metabisulfite is 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
Prior to 1900, most sulfuric acid was manufactured by the chamber process.[9] As late as 1940, up to 50% of sulfuric acid manufactured in the United States was produced by chamber process plants.
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.[10] World production in 2001 was 165 million tons, with an approximate value of US$8 billion.[citation needed] 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.[citation needed] 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:
- Ca
5F(PO
4)
3 + 5 H
2SO
4 + 10 H
2O → 5 CaSO
4·2 H
2O + HF + 3 H
3PO
4
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[citation needed]. Used acid is often recycled using a Spent Acid Regeneration (SAR) plant. These plants combust spent acid with natural gas, refinery gas, fuel oil or other fuel sources. This combustion process produces gaseous sulfur dioxide (SO
2) and sulfur trioxide (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 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:
- Al
2O
3 + 3 H
2SO
4 → Al
2(SO
4)
3 + 3 H
2O
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 H
2SO
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 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 H
2SO
4 → 2 SO
2 + 2 H
2O + O
2(830 °C) I
2 + SO
2 + 2 H
2O → 2 HI + H
2SO
4(120 °C) 2 HI → I
2 + H
2(320 °C)
The sulfur and iodine compounds are recovered and reused, hence the consideration of the process as a cycle. This process is endothermic and must occur at high temperatures, so energy in the form of heat has to be supplied.
The sulfur-iodine cycle has been proposed as a way to supply hydrogen for a hydrogen-based economy. It does not require hydrocarbons like current methods of steam reforming. But note that all of the available energy in the hydrogen so produced is supplied by the heat used to make it.
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.
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 around, 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 can be diluted by pouring concentrated acid onto crushed ice made from de-ionized water. The ice melts in an 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 so the solution remains cold. After all the ice has melted, further dilution can take place using water.
Pure sulfuric acid may be safely stored in glass vessels or bottles.
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. 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.[11]
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.[12][13][14]
See also
References
- ^ a b c Vladimir Karpenko, John A. Norris (2001), Vitriol in the history of Chemistry, Charles University
- ^ a b "sulfuric acid", The Columbia Encyclopedia (6th ed.), 2008, retrieved 16 March 2010
- ^ a b c "Sulphuric acid", Encyclopædia Britannica, vol. 26 (11th ed.), 1910–1911, pp. 65–69
- ^ a b c Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
- ^ "Ionization Constants of Inorganic Acids". .chemistry.msu.edu. Retrieved 30 May 2011.
- ^ sulphuric acid on sugar cubes chemistry experiment 8. Old Version. YouTube. Retrieved on 2011-07-18.
- ^ F. A. Carey. "Reactions of Arenes. Electrophilic Aromatic Substitution". On-Line Learning Center for Organic Chemistry. University of Calgary. Retrieved 27 January 2008.[dead link]
- ^ Orlando, T. M.; McCord, T. B.; Grieves, G. A. (2005). "The chemical nature of Europa surface material and the relation to a subsurface ocean". Icarus. 177 (2): 528–533. Bibcode:2005Icar..177..528O. doi:10.1016/j.icarus.2005.05.009.
- ^ Jones, Edward M. (1950). "Chamber Process Manufacture of Sulfuric Acid". Industrial and Engineering Chemistry. 42 (11): 2208–2210. doi:10.1021/ie50491a016.
- ^ Chenier, Philip J. (1987). Survey of Industrial Chemistry. New York: John Wiley & Sons. pp. 45–57. ISBN 0471010774.
- ^ Annex to Form D ("Red List"), 11th Edition, January 2007 (p. 4). International Narcotics Control Board. Vienna, Austria; 2007.
- ^ 66 FR 52670—52675. 17 October 2001.
- ^ "21 CFR 1309". Access.gpo.gov. Retrieved 30 May 2011.
- ^ "21 USC, Chapter 13 (Controlled Substances Act)". Usdoj.gov. Retrieved 30 May 2011.
Further reading
- 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.
External links
- International Chemical Safety Card 0362
- NIOSH Pocket Guide to Chemical Hazards
- External Material Safety Data Sheet
- Sulfuric acid analysis – titration freeware