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Butadiene

"Divinyl" redirects here. For the band, see Divinyls.

1,3-Butadiene (/ˌbjuːtəˈdaɪiːn/ ⓘ)[8] is the organic compound with the formula CH2=CH-CH=CH2. It is a colorless gas that is easily condensed to a liquid. It is important industrially as a precursor to synthetic rubber.[9] The molecule can be viewed as the union of two vinyl groups. It is the simplest conjugated diene.

1,3-Butadiene
Full structural formula of 1,3-butadiene
Full structural formula of 1,3-butadiene
Skeletal formula of 1,3-butadiene
Skeletal formula of 1,3-butadiene
Ball-and-stick model of 1,3-butadiene
Ball-and-stick model of 1,3-butadiene
Space-filling model of 1,3-butadiene
Space-filling model of 1,3-butadiene
Names
Preferred IUPAC name
Buta-1,3-diene[1]
Other names
  • Biethylene
  • Erythrene
  • Divinyl
  • Vinylethylene
  • Bivinyl
  • Butadiene
Identifiers
3D model (JSmol)
605258
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.003.138 Edit this at Wikidata
EC Number
  • 271-039-0
25198
KEGG
RTECS number
  • EI9275000
UNII
UN number 1010
  • InChI=1S/C4H6/c1-3-4-2/h3-4H,1-2H2 checkY
    Key: KAKZBPTYRLMSJV-UHFFFAOYSA-N checkY
  • InChI=1/C4H6/c1-3-4-2/h3-4H,1-2H2
    Key: KAKZBPTYRLMSJV-UHFFFAOYAZ
  • C=CC=C
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
  • 0.6149 g/cm3 at 25 °C, p>1 atm[2]
  • 0.64 g/cm3 at −6 °C, liquid
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]
1.4292
Viscosity 0.25 cP at 0 °C
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
 Flammable, irritative, carcinogen
GHS labelling:[7]
GHS02: Flammable GHS08: Health hazard
Danger
H220, H340, H350
P202, P210, P280, P308+P313, P377, P381, P403
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 4: Will rapidly or completely vaporize at normal atmospheric pressure and temperature, or is readily dispersed in air and will burn readily. Flash point below 23 °C (73 °F). E.g. propaneInstability 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g. white phosphorusSpecial hazards (white): no code
3
4
2
Flash point −85 °C (−121 °F; 188 K) liquid flash point[3]
414 °C (777 °F; 687 K)[6]
Explosive limits 2–12%
Lethal dose or concentration (LD, LC):
548 mg/kg (rat, oral)
  • 115,111 ppm (mouse)
  • 122,000 ppm (mouse, 2 h)
  • 126,667 ppm (rat, 4 h)
  • 130,000 ppm (rat, 4 h)[5]
250,000 ppm (rabbit, 30 min)[5]
NIOSH (US health exposure limits):
PEL (Permissible)
 TWA 1 ppm ST 5 ppm[3]
REL (Recommended)
 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
 Butane
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
checkY verify (what is checkY☒N ?)

Although butadiene breaks down quickly in the atmosphere, it is nevertheless found in ambient air in urban and suburban areas as a consequence of its constant emission from motor vehicles.[10]

The name butadiene can also refer to the isomer, 1,2-butadiene, which is a cumulated diene with structure H2C=C=CH−CH3. This allene has no industrial significance.

History

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In 1863, French chemist E. Caventou isolated butadiene from the pyrolysis of amyl alcohol.[11] This hydrocarbon was identified as butadiene in 1886, after Henry Edward Armstrong isolated it from among the pyrolysis products of petroleum.[12] In 1910, the Russian chemist Sergei Lebedev polymerized butadiene and obtained a material with rubber-like properties. This polymer was, however, found to be too soft to replace natural rubber in many applications, notably automobile tires.

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.[13] 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

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In 2020, 14.2 million tons were estimated to have been produced.[14]

Extraction from C4 hydrocarbons

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In the United States, western Europe, and Japan, butadiene is produced as a byproduct of the steam cracking process used to produce ethylene and other alkenes. When mixed with steam and briefly heated to very high temperatures (often over 900 °C), aliphatic hydrocarbons give up hydrogen to produce a complex mixture of unsaturated hydrocarbons, including butadiene. The quantity of butadiene produced depends on the hydrocarbons used as feed. Light feeds, such as ethane, give primarily ethylene when cracked, but heavier feeds favor the formation of heavier olefins, butadiene, and aromatic hydrocarbons.

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.[15]

From dehydrogenation of n-butane

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Butadiene can also be produced by the catalytic dehydrogenation of normal butane (n-butane). The first such post-war commercial plant, producing 65,000 tons per year of butadiene, began operations in 1957 in Houston, Texas.[16] Prior to that, in the 1940s the Rubber Reserve Company, a part of the United States government, constructed several plants in Borger, Texas, Toledo, Ohio, and El Segundo, California, to produce synthetic rubber for the war effort as part of the United States Synthetic Rubber Program.[17] Total capacity was 68 KMTA (Kilo Metric Tons per Annum).

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 alumina and chromia at high temperatures.[18]

From ethanol

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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 Sergei Lebedev, ethanol is converted to butadiene, hydrogen, and water at 400–450 °C over any of a variety of metal oxide catalysts:[19]

 
2 CH3CH2OH → CH2=CH−CH=CH2 + 2 H2O + H2

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 Ivan Ostromislensky, ethanol is oxidized to acetaldehyde, which reacts with additional ethanol over a tantalum-promoted porous silica catalyst at 325–350 °C to yield butadiene:[19]

 
CH3CH2OH + CH3CHO → CH2=CH−CH=CH2 + 2 H2O

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

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1,3-Butadiene can also be produced by catalytic dehydrogenation of normal butenes. This method was also used by the U.S. Synthetic Rubber Program (USSRP) during World War II. The process was much more economical than the alcohol or n-butane route but competed with aviation gasoline for available butene molecules (butenes were plentiful thanks to catalytic cracking). The USSRP constructed several plants in Baton Rouge and Lake Charles, Louisiana; Houston, Baytown, and Port Neches, Texas; and Torrance, California.[17] Total annual production was 275 KMTA.

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.[20]

After World War II, the production from butenes became the major type of production in USSR.

For laboratory use

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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-Diels-Alder reaction of cyclohexene.[21] Sulfolene is a convenient solid storable source for 1,3-butadiene in the laboratory. It releases the diene and sulfur dioxide upon heating.

Uses

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Most butadiene(75% of the manufactured 1,3-butadiene[9]) 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.[18]

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 cycloalkanes and cycloalkenes, as it reacts with double and triple carbon-carbon bonds through Diels-Alder reactions. The most widely used such reactions involve reactions of butadiene with one or two other molecules of butadiene, i.e., dimerization and trimerization respectively. Via dimerization butadiene is converted to 4-vinylcyclohexene and cyclooctadiene. In fact, vinylcyclohexene is a common impurity that accumulates when butadiene is stored. Via trimerization, butadiene is converted to cyclododecatriene. Some of these processes employ nickel- or titanium-containing catalysts.[22]

Butadiene is also a precursor to 1-octene via palladium catalyzed telomerization with methanol. This reaction produces 1-methoxy-2,7-octadiene as an intermediate.[14]

Structure, conformation, and stability

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Comparison of butadiene (s-trans conformer) and ethylene

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.[23] 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.[24]

 

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.[25] 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.[26] The ~3.5 kcal/mol difference in these heats of hydrogenation can be taken to be the resonance energy of a conjugated diene.

Reactions

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The structure of (butadiene)iron tricarbonyl[27]

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.[28] It is a partner in Diels–Alder reactions, e.g. with maleic anhydride to give tetrahydrophthalic anhydride.[29]

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

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A railway tank car carrying a butadiene and hydrocarbon mixture, displaying hazardous materials information including a diamond-shaped U.S. DOT placard showing a UN number[9]

Butadiene is of low acute toxicity. LC50 is 12.5–11.5 vol% for inhalation by rats and mice.[18]

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.[30]

IARC has designated 1,3-butadiene as a Group 1 carcinogen ('carcinogenic to humans'),[31] and the Agency for Toxic Substances Disease Registry and the US EPA also list the chemical as a carcinogen.[32][33] The American Conference of Governmental Industrial Hygienists (ACGIH) lists the chemical as a suspected carcinogen.[33] The Natural Resource Defense Council (NRDC) lists some disease clusters that are suspected to be associated with this chemical.[34] Some researchers have concluded it is the most potent carcinogen in cigarette smoke, twice as potent as the runner up acrylonitrile[35]

1,3-Butadiene is also a suspected human teratogen.[36][37][38] Prolonged and excessive exposure can affect many areas in the human body; blood, brain, eye, heart, kidney, lung, nose and throat have all been shown to react to the presence of excessive 1,3-butadiene.[39] Animal data suggest that women have a higher sensitivity to possible carcinogenic effects of butadiene over men when exposed to the chemical. This may be due to estrogen receptor impacts. While these data reveal important implications to the risks of human exposure to butadiene, more data are necessary to draw conclusive risk assessments. There is also a lack of human data for the effects of butadiene on reproductive and development shown to occur in mice, but animal studies have shown breathing butadiene during pregnancy can increase the number of birth defects, and humans have the same hormone systems as animals.[40]

1,3-Butadiene is recognized as a highly reactive volatile organic compound (HRVOC) for its potential to readily form ozone, and as such, emissions of the chemical are highly regulated by TCEQ in parts of the Houston-Brazoria-Galveston Ozone Non-Attainment Area.[41]

Data sheet

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Properties
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

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References

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  1. ^ "Front Matter". Nomenclature of Organic Chemistry : IUPAC Recommendations and Preferred Names 2013 (Blue Book). Cambridge: The Royal Society of Chemistry. 2014. p. 374. doi:10.1039/9781849733069-FP001. ISBN 978-0-85404-182-4.
  2. ^ Haynes, William M. (2016). CRC Handbook of Chemistry and Physics (97th ed.). Boca Raton: CRC Press. p. 3-76. ISBN 978-1-4987-5429-3.
  3. ^ a b c d e NIOSH Pocket Guide to Chemical Hazards. "#0067". National Institute for Occupational Safety and Health (NIOSH).
  4. ^ "1,3-Butadiene". NIST Chemistry WebBook.
  5. ^ a b "1,3-Butadiene". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  6. ^ "1,3-Butadiene". INCHEM. International Programme on Chemical Safety (IPCS).
  7. ^ Record in the GESTIS Substance Database of the Institute for Occupational Safety and Health
  8. ^ "BUTADIENE | Meaning & Definition for UK English". Lexico.com. Archived from the original on 20 August 2020. Retrieved 24 August 2022.
  9. ^ a b c PubChem. "1,3-Butadiene". pubchem.ncbi.nlm.nih.gov. Retrieved 8 May 2024.
  10. ^ "1,3-Butadiene". US Environmental Protection Agency US EPA. Retrieved 2 September 2014.
  11. ^ Caventou, E. (1863). "Ueber eine mit dem zweifach-gebromten Brombutylen isomere Verbindung und über die bromhaltigen Derivate des Brombutylens". Justus Liebigs Annalen der Chemie. 127: 93–97. doi:10.1002/jlac.18631270112.
  12. ^ Armstrong, H. E.; Miller, A. K. (1886). "The decomposition and genesis of hydrocarbons at high temperatures. I. The products of the manufacture of gas from petroleum". J. Chem. Soc. 49: 74–93. doi:10.1039/CT8864900074.
  13. ^ Simple Things Won't Save the Earth, J. Robert Hunter
  14. ^ a b Yang, Ji; Wang, Peng; Neumann, Helfried; Jackstell, Ralf; Beller, Matthias (2023). "Industrially applied and relevant transformations of 1,3-butadiene using homogeneous catalysts". Industrial Chemistry & Materials. 1 (2): 155–174. doi:10.1039/D3IM00009E. S2CID 258122761.
  15. ^ 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.
  16. ^ Beychok, M.R. and Brack, W.J., "First Postwar Butadiene Plant", Petroleum Refiner, June 1957.
  17. ^ a b Herbert, Vernon (1985). Synthetic Rubber: A Project That Had to Succeed. Greenwood Press. ISBN 0-313-24634-3.
  18. ^ a b c Grub, J.; Löser, E. (2012). "Butadiene". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a04_431.pub2. ISBN 978-3527306732.
  19. ^ 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.
  20. ^ 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.
  21. ^ Hershberg, E. B.; Ruhoff, John R. (1937). "1,3-Butadiene". Org. Synth. 17: 25. doi:10.15227/orgsyn.017.0025.
  22. ^ 4-Vinylcyclohexene (PDF). IARC. ISBN 9789283212607. Retrieved 19 April 2009.
  23. ^ Feller, David; Craig, Norman C. (26 February 2009). "High Level ab Initio Energies and Structures for the Rotamers of 1,3-Butadiene". The Journal of Physical Chemistry A. 113 (8): 1601–1607. Bibcode:2009JPCA..113.1601F. doi:10.1021/jp8095709. ISSN 1089-5639. PMID 19199679.
  24. ^ Craig, Norman C.; Groner, Peter; McKean, Donald C. (1 June 2006). "Equilibrium Structures for Butadiene and Ethylene: Compelling Evidence for Π-Electron Delocalization in Butadiene". The Journal of Physical Chemistry A. 110 (23): 7461–7469. Bibcode:2006JPCA..110.7461C. doi:10.1021/jp060695b. ISSN 1089-5639. PMID 16759136.
  25. ^ Vollhardt, K. Peter C.; Schore, Neil Eric (2007). Organic chemistry: structure and function (5th ed.). New York: W.H. Freeman. ISBN 978-0716799498. OCLC 61448218.
  26. ^ Carey, Francis A. (2002). Organic chemistry (5th ed.). London: McGraw-Hill. ISBN 978-0071151498. OCLC 49907089.
  27. ^ Reiss, Guido J. (2010). "Redetermination of (η4-s-cis-1,3-butadiene)tricarbonyliron(0)". Acta Crystallographica Section E. 66 (11): m1369. Bibcode:2010AcCrE..66m1369R. doi:10.1107/S1600536810039218. PMC 3009352. PMID 21588810.
  28. ^ Nyström, J. E.; Rein, T.; Bäckvall, J. E. (1989). "1,4-Functionalization of 1,3-Dienes via Palladium-Catalyzed Chloroacetoxylation and Allylic Amination: 1-Acetoxy-4-diethylamino-2-butene and 1-Acetoxy-4-benzylamino-2-butene". Org. Synth. 67: 105. doi:10.15227/orgsyn.067.0105.
  29. ^ Cope, Arthur C.; Herrick, Elbert C. (1950). "cis-Δ4-Tetrahydrophthalic Anhydride". Org. Synth. 50: 93. doi:10.15227/orgsyn.030.0093.
  30. ^ "NPI sheet". Archived from the original on 22 December 2003. Retrieved 10 January 2006.
  31. ^ Grosse, Yann; Baan, Robert; Straif, Kurt; Secretan, Béatrice; El Ghissassi, Fatiha; Bouvard, Véronique; Altieri, Andrea; Cogliano, Vincent (2008). "Carcinogenicity of 1,3-butadiene, ethylene oxide, vinyl chloride, vinyl fluoride, and vinyl bromide". The Lancet Oncology. 8 (8): 679–680. doi:10.1016/S1470-2045(07)70235-8. ISSN 1470-2045. PMID 17726789.
  32. ^ "1,3-Butadiene". Toxic Substances Portal. Agency for Toxic Substances and Disease Registry (ATSDR). Archived from the original on 9 June 2012.
  33. ^ a b "1,3-Butadiene: Health Effects". Occupational Safety & Health Administration.
  34. ^ "Disease Clusters Spotlight the Need to Protect People from Toxic Chemicals". NRDC. 10 May 2011.
  35. ^ Fowles, J.; Dybing, E. (4 September 2003). "Application of toxicological risk assessment principles to the chemical constituents of cigarette smoke". Institute of Environmental Science and Research. 12 (4): 424–430. doi:10.1136/tc.12.4.424. PMC 1747794. PMID 14660781.
  36. ^ Landrigan, P. J. (1990). "Critical assessment of epidemiologic studies on the human carcinogenicity of 1,3-butadiene". Environmental Health Perspectives. 86: 143–147. doi:10.1289/ehp.9086143. PMC 1567758. PMID 2205484.
  37. ^ "1,3-Butadiene CAS No. 106-99-0" (PDF). Report on Carcinogens (11th ed.). Archived (PDF) from the original on 8 May 2009.
  38. ^ Melnick, Ronald L.; Kohn, Michael C. (1995). "Mechanistic data indicate that 1,3-butadiene is a human carcinogen". Carcinogenesis. 16 (2): 157–163. doi:10.1093/carcin/16.2.157. PMID 7859343.
  39. ^ "Environment Agency - 1,3-Butadiene". Archived from the original on 3 February 2011. Retrieved 20 August 2010.
  40. ^ "1,3-Butadiene". Technology Transfer Network Air Toxics Web Site. EPA. Archived from the original on 11 May 2012.
  41. ^ "Controlling HRVOC Emissions". Texas Commission on Environmental Quality.
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