Cyclopropane

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Cyclopropane[1]
Cyclopropane - displayed formula Cyclopropane - skeletal formula
Cyclopropane-3D-balls.png Cyclopropane-3D-vdW.png
Identifiers
CAS number 75-19-4 YesY
PubChem 6351
ChemSpider 6111 YesY
UNII 99TB643425 YesY
KEGG D03627 YesY
ChEBI CHEBI:30365 YesY
ChEMBL CHEMBL1796999 N
Jmol-3D images Image 1
Properties
Molecular formula C3H6
Molar mass 42.08 g/mol
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
Hazards
MSDS External MSDS
Main hazards Highly flammable
Asphyxiant
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 noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
 N (verify) (what is: YesY/N?)
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. Cyclopropane and propene have the same molecular formula but have different structures, making them structural isomers.

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 In modern anaesthetic practice, it has been superseded by other agents, due to its extreme reactivity under normal conditions: When the gas is mixed with oxygen there is a significant risk of explosion.

History[edit]

Cyclopropane was discovered in 1881 by August Freund, who also proposed the correct structure for the new substance in his first paper.[2] Freund treated 1,3-dibromopropane with sodium, causing an intramolecular Wurtz reaction leading directly to cyclopropane.[3][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]

Anaesthesia[edit]

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, before being largely phased out. Cylinders and flow meters were coloured orange.

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.[11] 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[12],[13] (i.e. 1/6 s density and 5/6 p density) so that the C-C bonds have more p character than normal[14] (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).[15]

An alternative model for describing the bonding in cyclopropane involves Walsh diagrams and aims to do a better job fitting molecular orbital theory in light of spectroscopic evidence and group symmetry arguments. In this model cyclopropane is described as a three-center bonded orbital combination of methylene carbenes.

Synthesis[edit]

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

Cyclopropanation[edit]

Cyclopropane rings are found in numerous biomolecules and pharmaceutical drugs. As such the formation of cyclopropane rings, generally referred to as cyclopropanation, is an active area of chemical research.

Reactions[edit]

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.[16]

Electrophilic addition of HBr to cyclopropane

Safety[edit]

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

See also[edit]

References[edit]

  1. ^ Merck Index, 11th Edition, 2755.
  2. ^ Freund, August (1882) "Über Trimethylen" (On trimethylene), Monatshefte für Chemie … , 3 : 625-635.
  3. ^ a b August Freund (1881). "Über Trimethylen". Journal für Praktische Chemie 26 (1): 625–635. doi:10.1002/prac.18820260125. 
  4. ^ August Freund (1882). "Über Trimethylen". Monatshefte für Chemie 3 (1): 625–635. doi:10.1007/BF01516828. 
  5. ^ a b G. Gustavson (1887). "Ueber eine neue Darstellungsmethode des Trimethylens". J. Prakt. Chem. 36: 300–305. doi:10.1002/prac.18870360127. 
  6. ^ G. H. W. Lucas and V. E. Henderson (1 August 1929). "A New Anesthetic: Cyclopropane : A Preliminary Report". Can Med Assoc J. 21 (2): 173–5. PMC 1710967. PMID 20317448. 
  7. ^ H. B. Hass, E. T. McBee, and 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. doi:10.1136/hrt.12.3.239. PMID 15426685. 
  10. ^ MacDonald, AG (June 1994). "A short history of fires and explosions caused by anaesthetic agents.". British journal of anaesthesia 72 (6): 710–22. doi:10.1093/bja/72.6.710. PMID 8024925. 
  11. ^ Eric V. Anslyn and Dennis A. Dougherty. Modern Physical Organic Chemistry. 2006. pages 850-852
  12. ^ http://isites.harvard.edu/fs/docs/icb.topic93502.files/Lectures_and_Handouts/06-Handouts/deMeijere.pdf
  13. ^ http://isites.harvard.edu/fs/docs/icb.topic1032290.files/lecture%203.pdf
  14. ^ Knipe, edited by A.C. (2007). March's advanced organic chemistry reactions, mechanisms, and structure. (6th ed. ed.). Hoboken, N.J.: Wiley-Interscience. p. 219. ISBN 0470084944. 
  15. ^ 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. 
  16. ^ Advanced organic Chemistry, Reactions, mechanisms and structure 3ed. Jerry March ISBN 0-471-85472-7

External links[edit]