Butadiene
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Names | |||
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Preferred IUPAC name
Buta-1,3-diene[1] | |||
Other names
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Identifiers | |||
3D model (
JSmol ) |
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605258 | |||
ChEBI | |||
ChEMBL | |||
ChemSpider | |||
ECHA InfoCard
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100.003.138 | ||
EC Number |
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25198 | |||
KEGG | |||
PubChem CID
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RTECS number
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UNII | |||
UN number | 1010
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CompTox Dashboard (EPA)
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Properties[4] | |||
C4H6 CH2=CH-CH=CH2 | |||
Molar mass | 54.0916 g/mol | ||
Appearance | Colourless gas or refrigerated liquid | ||
Odor | Mildly aromatic or gasoline-like | ||
Density |
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Melting point | −108.91 °C (−164.04 °F; 164.24 K) | ||
Boiling point | −4.41 °C (24.06 °F; 268.74 K) | ||
1.3 g/L at 5 °C, 735 mg/L at 20 °C | |||
Solubility | |||
Vapor pressure | 2.4 atm (20 °C)[3] | ||
Refractive index (nD)
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1.4292 | ||
Viscosity | 0.25 cP at 0 °C | ||
Hazards | |||
Occupational safety and health (OHS/OSH): | |||
Main hazards
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Flammable, irritative, carcinogen | ||
GHS labelling:[7] | |||
Danger | |||
H220, H280, H340, H350 | |||
P202, P210, P280, P308+P313, P377, P381, P403 | |||
NFPA 704 (fire diamond) | |||
Flash point | −85 °C (−121 °F; 188 K) liquid flash point[3] | ||
414 °C (777 °F; 687 K)[6] | |||
Explosive limits
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2–12% | ||
Lethal dose or concentration (LD, LC): | |||
LD50 (median dose)
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548 mg/kg (rat, oral) | ||
LC50 (median concentration)
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LCLo (lowest published)
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250,000 ppm (rabbit, 30 min)[5] | ||
NIOSH (US health exposure limits): | |||
PEL (Permissible)
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TWA 1 ppm ST 5 ppm[3] | ||
REL (Recommended)
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Potential occupational carcinogen[3] | ||
IDLH (Immediate danger) |
2000 ppm[3] | ||
Safety data sheet (SDS) | ECSC 0017 | ||
Related compounds | |||
Related
Alkenes and dienes |
Isoprene Chloroprene | ||
Related compounds
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Butane | ||
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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1,3-Butadiene (
Although butadiene
The name butadiene can also refer to the
History
In 1863, French chemist E. Caventou isolated butadiene from the
The butadiene industry originated in the years before World War II. Many of the belligerent nations realized that in the event of war, they could be cut off from rubber plantations controlled by the British Empire, and sought to reduce their dependence on natural rubber.[12] In 1929, Eduard Tschunker and Walter Bock, working for IG Farben in Germany, made a copolymer of styrene and butadiene that could be used in automobile tires. Worldwide production quickly ensued, with butadiene being produced from grain alcohol in the Soviet Union and the United States, and from coal-derived acetylene in Germany.
Production
In 2020, 14.2 million tons were estimated to have been produced.[13]
Extraction from C4 hydrocarbons
In the United States, western Europe, and Japan, butadiene is produced as a byproduct of the
Butadiene is typically isolated from the other four-carbon hydrocarbons produced in steam cracking by extractive distillation using a polar aprotic solvent such as acetonitrile, N-methyl-2-pyrrolidone, furfural, or dimethylformamide, from which it is then stripped by distillation.[14]
From dehydrogenation of n-butane
Butadiene can also be produced by the catalytic
Today, butadiene from n-butane is commercially produced using the Houdry Catadiene process, which was developed during World War II. This entails treating butane over
From ethanol
In other parts of the world, including South America, Eastern Europe, China, and India, butadiene is also produced from ethanol. While not competitive with steam cracking for producing large volumes of butadiene, lower capital costs make production from ethanol a viable option for smaller-capacity plants. Two processes were in use.
In the single-step process developed by
This process was the basis for the Soviet Union's synthetic rubber industry during and after World War II, and it remained in limited use in Russia and other parts of eastern Europe until the end of the 1970s. At the same time this type of manufacture was canceled in Brazil. As of 2017, no butadiene was produced industrially from ethanol.
In the other, two-step process, developed by the Russian emigre chemist
This process was one of the three used in the United States to produce "government rubber" during World War II, although it is less economical than the butane or butene routes for the large volumes. Still, three plants with a total capacity of 200,000 tons per year were constructed in the U.S. (Institute, West Virginia, Louisville, Kentucky, and Kobuta, Pennsylvania) with start-ups completed in 1943, the Louisville plant initially created butadiene from acetylene generated by an associated calcium carbide plant. The process remains in use today in China and India.
From butenes
1,3-Butadiene can also be produced by
In the 1960s, a Houston company known as "Petro-Tex" patented a process to produce butadiene from normal butenes by oxidative dehydrogenation using a proprietary catalyst. It is unclear if this technology is practiced commercially.[19]
After World War II, the production from butenes became the major type of production in USSR.
For laboratory use
1,3-Butadiene is inconvenient for laboratory use because it is gas. Laboratory procedures have been optimized for its generation from nongaseous precursors. It can be produced by the retro-
Uses
Most butadiene is used to make synthetic rubbers for the manufacture of tyres and components of many consumer items.
The conversion of butadiene to synthetic rubbers is called polymerization, a process by which small molecules (monomers) are linked to make large ones (polymers). The mere polymerization of butadiene gives polybutadiene, which is a very soft, almost liquid material. The polymerization of butadiene and other monomers gives copolymers, which are more valued. The polymerization of butadiene and styrene and/or acrylonitrile, such as acrylonitrile butadiene styrene (ABS), nitrile-butadiene (NBR), and styrene-butadiene (SBR). These copolymers are tough and/or elastic depending on the ratio of the monomers used in their preparation. SBR is the material most commonly used for the production of automobile tyres. Precursors to still other synthetic rubbers are prepared from butadiene. One is chloroprene.[17]
Smaller amounts of butadiene are used to make adiponitrile, a precursor to some nylons. The conversion of butadiene to adiponitrile entails the addition of hydrogen cyanide to each of the double bonds in butadiene. The process is called hydrocyanation.
Butadiene is used to make the solvent sulfolane.
Butadiene is also useful in the synthesis of
Butadiene is also a precursor to
Structure, conformation, and stability
The most stable conformer of 1,3-butadiene is the s-trans conformation, in which the molecule is planar, with the two pairs of double bonds facing opposite directions. This conformation is most stable because orbital overlap between double bonds is maximized, allowing for maximum conjugation, while steric effects are minimized. Conventionally, the s-trans conformation is considered to have a C2-C3 dihedral angle of 180°. In contrast, the s-cis conformation, in which the dihedral angle is 0°, with the pair of double bonds facing the same direction is approximately 16.5 kJ/mol (3.9 kcal/mol) higher in energy, due to steric hindrance. This geometry is a local energy maximum, so in contrast to the s-trans geometry, it is not a conformer. The gauche geometry, in which the double bonds of the s-cis geometry are twisted to give a dihedral angle of around 38°, is a second conformer that is around 12.0 kJ/mol (2.9 kcal/mol) higher in energy than the s-trans conformer. Overall, there is a barrier of 24.8 kJ/mol (5.9 kcal/mol) for isomerization between the two conformers.[22] This increased rotational barrier and strong overall preference for a near-planar geometry is evidence for a delocalized π system and a small degree of partial double bond character in the C–C single bond, in accord with resonance theory.
Despite the high energy of the s-cis conformation, 1,3-butadiene needs to assume this conformation (or one very similar) before it can participate as the four-electron component in concerted cycloaddition reactions like the Diels-Alder reaction.
Similarly, a combined experimental and computational study has found that the double bond of s-trans-butadiene has a length of 133.8 pm, while that for ethylene has a length of 133.0 pm. This was taken as evidence of a π-bond weakened and lengthened by delocalization, as depicted by the resonance structures shown below.[23]
A qualitative picture of the molecular orbitals of 1,3-butadiene is readily obtained by applying Hückel theory. (The article on Hückel theory gives a derivation for the butadiene orbitals.)
1,3-Butadiene is also thermodynamically stabilized. While a monosubstituted double bond releases about 30.3 kcal/mol of heat upon hydrogenation, 1,3-butadiene releases slightly less (57.1 kcal/mol) than twice this energy (60.6 kcal/mol), expected for two isolated double bonds. That implies a stabilization energy of 3.5 kcal/mol.[24] Similarly, the hydrogenation of the terminal double bond of 1,4-pentadiene releases 30.1 kcal/mol of heat, while hydrogenation of the terminal double bond of conjugated (E)-1,3-pentadiene releases only 26.5 kcal/mol, implying a very similar value of 3.6 kcal/mol for the stabilization energy.[25] The ~3.5 kcal/mol difference in these heats of hydrogenation can be taken to be the resonance energy of a conjugated diene.
Reactions
The industrial uses illustrate the tendency of butadiene to polymerize. Its susceptibility to 1,4-addition reactions is illustrated by its hydrocyanation. Like many dienes, it undergoes Pd-catalyzed reactions that proceed via allyl complexes.[26] It is a partner in Diels–Alder reactions, e.g. with maleic anhydride to give tetrahydrophthalic anhydride.[27]
Like other dienes, butadiene is a ligand for low-valent metal complexes, e.g. the derivatives Fe(butadiene)(CO)3 and Mo(butadiene)3.
Environmental health and safety
Butadiene is of low acute toxicity.
Long-term exposure has been associated with cardiovascular disease. There is a consistent association with leukemia, as well as a significant association with other cancers.[29]
1,3-Butadiene is also a suspected human
1,3-Butadiene is recognized as a highly reactive
Data sheet
Properties | |
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Phase behavior | |
Triple point | 164.2 K (-109.0 °C) ? bar |
Critical point | 425 K (152 °C) 43.2 bar |
Structure | |
Symmetry group | C2h |
Gas properties | |
ΔfH0 | 110.2 kJ/mol |
Cp | 79.5 J/mol·K |
Liquid properties | |
ΔfH0 | 90.5 kJ/mol |
S0 | 199.0 J/mol·K |
Cp | 123.6 J/mol·K |
Liquid density | 0.64 ×103 kg/m3 |
See also
References
- ^
"Front Matter". Nomenclature of Organic Chemistry : IUPAC Recommendations and Preferred Names 2013 (Blue Book). Cambridge: ISBN 978-0-85404-182-4.
- ISBN 978-1-4987-5429-3.
- ^ a b c d e NIOSH Pocket Guide to Chemical Hazards. "#0067". National Institute for Occupational Safety and Health (NIOSH).
- ^ "1,3-Butadiene". NIST Chemistry WebBook.
- ^ a b "1,3-Butadiene". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
- ^ "1,3-Butadiene". INCHEM. International Programme on Chemical Safety (IPCS).
- ^ Record in the GESTIS Substance Database of the Institute for Occupational Safety and Health
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- US EPA. Retrieved 2 September 2014.
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- ^ S2CID 258122761.
- ^ Sun, H.P. Wristers, J.P. (1992). Butadiene. In J.I. Kroschwitz (Ed.), Encyclopedia of Chemical Technology, 4th ed., vol. 4, pp. 663–690. New York: John Wiley & Sons.
- ^ Beychok, M.R. and Brack, W.J., "First Postwar Butadiene Plant", Petroleum Refiner, June 1957.
- ^ ISBN 0-313-24634-3.
- ^ ISBN 978-3527306732.
- ^ a b Kirshenbaum, I. (1978). "Butadiene". In Grayson, M. (ed.). Encyclopedia of Chemical Technology. Vol. 4 (3rd ed.). New York: John Wiley & Sons. pp. 313–337.
- ^ Welch, L. Marshall; Croce, Louis; Christmann, Harold (November 1978). "BUTADIENE VIA OXIDATIVE DEHYDROGENATION". Hydrocarbon Processing. 57 (11): 131–136. Retrieved 1 June 2019 – via ResearchGate.
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- ^ "NPI sheet". Archived from the original on 22 December 2003. Retrieved 10 January 2006.
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- ^ "1,3-Butadiene". Toxic Substances Portal. Agency for Toxic Substances and Disease Registry (ATSDR). Archived from the original on 9 June 2012.
- ^ a b "1,3-Butadiene: Health Effects". Occupational Safety & Health Administration.
- ^ "Disease Clusters Spotlight the Need to Protect People from Toxic Chemicals". NRDC. 10 May 2011.
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- ^ "1,3-Butadiene CAS No. 106-99-0" (PDF). Report on Carcinogens (11th ed.). Archived (PDF) from the original on 8 May 2009.
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- ^ "Environment Agency - 1,3-Butadiene". Archived from the original on 3 February 2011. Retrieved 20 August 2010.
- ^ "1,3-Butadiene". Technology Transfer Network Air Toxics Web Site. EPA. Archived from the original on 11 May 2012.
- ^ "Controlling HRVOC Emissions". Texas Commission on Environmental Quality.
External links
- 1,3-Butadiene – Agency for Toxic Substances and Disease Registry
- 1,3-Butadiene – CDC - NIOSH Pocket Guide to Chemical Hazards
- National Pollutant Inventory – 1,3-Butadiene