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Cyclopropane - displayed formula
Cyclopropane - skeletal formula
Preferred IUPAC name
3D model (JSmol)
ECHA InfoCard 100.000.771
Molar mass 42.08 g/mol
Appearance Colorless gas
Odor Sweet smelling
Density 1.879 g/L (1 atm, 0 °C)
Melting point −128 °C (−198 °F; 145 K)
Boiling point −33 °C (−27 °F; 240 K)
Acidity (pKa) ~46
-39.9·10−6 cm3/mol
Main hazards Highly flammable
Safety data sheet External MSDS
NFPA 704
Flammability code 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., propane Health code 1: Exposure would cause irritation but only minor residual injury. E.g., turpentine Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YesYN ?)
Infobox references

Cyclopropane is a cycloalkane molecule with the molecular formula C3H6, consisting of three carbon atoms linked to each other to form a ring, with each carbon atom bearing two hydrogen atoms resulting in D3h molecular symmetry. The small size of the ring creates substantial ring strain in the structure.

Cyclopropane is an anaesthetic when inhaled. In modern anaesthetic practice, it has been superseded by other agents, due to its extreme reactivity under normal conditions: when the gas is mixed with oxygen, there is a significant risk of explosion.


Cyclopropane was discovered in 1881 by August Freund, who also proposed the correct structure for the new substance in his first paper.[3] Freund treated 1,3-dibromopropane with sodium, causing an intramolecular Wurtz reaction leading directly to cyclopropane.[4] The yield of the reaction was improved by Gustavson in 1887 with the use of zinc instead of sodium.[5] Cyclopropane had no commercial application until Henderson and Lucas discovered its anaesthetic properties in 1929;[6] industrial production had begun by 1936.[7]


Cyclopropane was introduced into clinical use by the American anaesthetist Ralph Waters who used a closed system with carbon dioxide absorption to conserve this then-costly agent. Cyclopropane is a relatively potent, non-irritating and sweet smelling agent with a minimum alveolar concentration of 17.5%[8] and a blood/gas partition coefficient of 0.55. This meant induction of anaesthesia by inhalation of cyclopropane and oxygen was rapid and not unpleasant. However at the conclusion of prolonged anaesthesia patients could suffer a sudden decrease in blood pressure, potentially leading to cardiac dysrhythmia; a reaction known as "cyclopropane shock".[9] For this reason, as well as its high cost and its explosive nature,[10] it was latterly used only for the induction of anaesthesia, and has not been available for clinical use since the mid 1980s. Cylinders and flow meters were coloured orange.


Cyclopropane is inactive at the GABAA and glycine receptors, and instead acts as an NMDA receptor antagonist.[11][12] It also inhibits the AMPA receptor and nicotinic acetylcholine receptors, and activates certain K2P channels.[11][12][13]

Structure and bonding[edit]

Orbital overlap in the bent bonding model of cyclopropane

The triangular structure of cyclopropane requires the bond angles between carbon-carbon bonds to be 60°. This is far less than the thermodynamically most stable angle of 109.5° (for bonds between atoms with sp3 hybridised orbitals) and leads to significant ring strain. The molecule also has torsional strain due to the eclipsed conformation of its hydrogen atoms. As such, the bonds between the carbon atoms are considerably weaker than in a typical alkane, resulting in much higher reactivity.

Bonding between the carbon centres is generally described in terms of bent bonds.[14] In this model the carbon-carbon bonds are bent outwards so that the inter-orbital angle is 104°. This reduces the level of bond strain and is achieved by distorting the sp3 hybridisation of carbon atoms to technically sp5 hybridisation[15],[16] (i.e. 16 s density and 56 p density) so that the C-C bonds have more π character than normal[17] (at the same time the carbon-to-hydrogen bonds gain more s-character). One unusual consequence of bent bonding is that while the C-C bonds in cyclopropane are weaker than normal, the carbon atoms are also closer together than in a regular alkane bond: 151 pm versus 153 pm (average alkene bond: 146 pm).[18]

Stability due to cyclic delocalization of the six electrons of cyclopropane's three CC σ bonds was given by Michael J. S. Dewar as an explanation of the only slightly greater strain of cyclopropane ("only" 27.6 kcal/mol) as compared to cyclobutane (26.2 kcal/mol) with cyclohexane as reference with Estr=0 kcal/mol.[19] This stabilization is referred to as σ aromaticity,[20][21] in constrast to the usual π aromaticity, that, for example, is a highly stabilizing effect in benzene. Other studies do not support the role of σ-aromaticity in cyclopropane and the existence of an induced ring current; such studies provide an alternative explanation for the energetic stabilization and abnormal magnetic behaviour of cyclopropane.[22]


Cyclopropane was first produced via a Wurtz coupling, in which 1,3-dibromopropane was cyclised using sodium.[3] The yield of this reaction can be improved by exchanging the metal for zinc.[5]

BrCH2CH2CH2Br + 2 Na → (CH2)3 + 2 NaBr


Cyclopropane rings are found in numerous biomolecules (e.g., pyrethrins, a group of natural insecticides) and pharmaceutical drugs. As such the formation of cyclopropane rings, generally referred to as cyclopropanation, is an active area of chemical research.


Owing to the increased π-character of its C-C bonds, cyclopropane can react like an alkene in certain cases. For instance it undergoes hydrohalogenation with mineral acids to give linear alkyl halides. Substituted cyclopropanes also react, following Markovnikov's rule.[23] Substituted cyclopropanes can oxidatively add to transition metals, in a process referred to as C–C activation.

Electrophilic addition of HBr to cyclopropane


Cyclopropane is highly flammable. However, despite its strain energy it is not substantially more explosive than other alkanes.

See also[edit]


  1. ^ Merck Index, 11th Edition, 2755.
  2. ^ Nomenclature of Organic Chemistry : IUPAC Recommendations and Preferred Names 2013 (Blue Book). Cambridge: The Royal Society of Chemistry. 2014. p. 137. ISBN 978-0-85404-182-4. doi:10.1039/9781849733069-FP001. 
  3. ^ a b August Freund (1881). "Über Trimethylen" [On trimethylene]. Journal für Praktische Chemie. 26 (1): 367–377. doi:10.1002/prac.18820260125. 
  4. ^ August Freund (1882). "Über Trimethylen" [On trimethylene]. Monatshefte für Chemie …. 3 (1): 625–635. doi:10.1007/BF01516828. 
  5. ^ a b G. Gustavson (1887). "Ueber eine neue Darstellungsmethode des Trimethylens" [On a new method of preparing trimethylene]. Journal für Praktische Chemie. 36: 300–305. doi:10.1002/prac.18870360127. 
  6. ^ G. H. W. Lucas; V. E. Henderson (1 August 1929). "A New Anesthetic: Cyclopropane : A Preliminary Report". Can Med Assoc J. 21 (2): 173–5. PMC 1710967Freely accessible. PMID 20317448. 
  7. ^ H. B. Hass; E. T. McBee; G. E. Hinds (1936). "Synthesis of Cyclopropane". Industrial & Engineering Chemistry. 28 (10): 1178–81. doi:10.1021/ie50322a013. 
  8. ^ Eger, Edmond I.; Brandstater, Bernard; Saidman, Lawrence J.; Regan, Michael J.; Severinghaus, John W.; Munson, Edwin S. (1965). "Equipotent Alveolar Concentrations of Methoxyflurane, Halothane, Diethyl Ether, Fluroxene, Cyclopropane, Xenon and Nitrous Oxide in the Dog". Anesthesiology. 26 (6): 771–777. doi:10.1097/00000542-196511000-00012. 
  9. ^ JOHNSTONE, M; Alberts, JR (July 1950). "Cyclopropane anesthesia and ventricular arrhythmias.". British heart journal. 12 (3): 239–44. PMC 479392Freely accessible. PMID 15426685. doi:10.1136/hrt.12.3.239. 
  10. ^ MacDonald, AG (June 1994). "A short history of fires and explosions caused by anaesthetic agents.". British journal of anaesthesia. 72 (6): 710–22. PMID 8024925. doi:10.1093/bja/72.6.710. 
  11. ^ a b Hugh C. Hemmings; Philip M. Hopkins (2006). Foundations of Anesthesia: Basic Sciences for Clinical Practice. Elsevier Health Sciences. pp. 292–. ISBN 0-323-03707-0. 
  12. ^ a b Hemmings, Hugh C. (2009). "Molecular Targets of General Anesthetics in the Nervous System". Suppressing the Mind: 11–31. doi:10.1007/978-1-60761-462-3_2. 
  13. ^ Hara K, Eger EI, Laster MJ, Harris RA (December 2002). "Nonhalogenated alkanes cyclopropane and butane affect neurotransmitter-gated ion channel and G-protein-coupled receptors: differential actions on GABAA and glycine receptors". Anesthesiology. 97 (6): 1512–20. PMID 12459679. doi:10.1097/00000542-200212000-00025. 
  14. ^ Eric V. Anslyn and Dennis A. Dougherty. Modern Physical Organic Chemistry. 2006. pages 850-852
  15. ^
  16. ^
  17. ^ Knipe, edited by A.C. (2007). March's advanced organic chemistry reactions, mechanisms, and structure. (6th ed.). Hoboken, N.J.: Wiley-Interscience. p. 219. ISBN 0470084944. 
  18. ^ Allen, Frank H.; Kennard, Olga; Watson, David G.; Brammer, Lee; Orpen, A. Guy; Taylor, Robin (1987). "Tables of bond lengths determined by X-ray and neutron diffraction. Part 1. Bond lengths in organic compounds". Journal of the Chemical Society, Perkin Transactions 2 (12): S1–S19. doi:10.1039/P298700000S1. 
  19. ^ S. W. Benson, Thermochemical Kinetics, S. 273, J. Wiley & Sons, New York, London, Sydney, Toronto 1976
  20. ^ Dewar, M. J. (1984). "Chemical Implications of σ Conjugation". J. Am. Chem. Soc. 106: 669–682. doi:10.1021/ja00315a036. 
  21. ^ Cremer, D. (1988). "Pros and Cons of σ-Aromaticity". Tetrahedron. 44 (2): 7427–7454. doi:10.1016/s0040-4020(01)86238-4. 
  22. ^ Wu, Wei; Ma, Ben; Wu, Judy I-Chia; von Ragué, Schleyer; Mo, Yirong (2009). "Is Cyclopropane Really the σ-Aromatic Paradigm?". Chemistry: A European Journal. 15 (38): 9730–9736. doi:10.1002/chem.200900586. 
  23. ^ Advanced organic Chemistry, Reactions, mechanisms and structure 3ed. Jerry March ISBN 0-471-85472-7

External links[edit]