| 词条 | Carbon–fluorine bond | |||||||||||||||||||||||||||||||||||||||||||||||||||||||||
| 释义 |
The carbon–fluorine bond is a polar covalent bond between carbon and fluorine that is a component of all organofluorine compounds. It is the fourth strongest single bond in organic chemistry—behind the B-F single bond, Si-F single bond and the H-F single bond, and relatively short—due to its partial ionic character. The bond also strengthens and shortens as more fluorines are added to the same carbon on a chemical compound. As such, fluoroalkanes like tetrafluoromethane (carbon tetrafluoride) are some of the most unreactive organic compounds. Electronegativity and bond strengthThe high electronegativity of fluorine (4.0 for F vs. 2.5 for carbon) gives the carbon–fluorine bond a significant polarity/dipole moment. The electron density is concentrated around the fluorine, leaving the carbon relatively electron poor. This introduces ionic character to the bond through partial charges (Cδ+—Fδ−). The partial charges on the fluorine and carbon are attractive, contributing to the unusual bond strength of the carbon–fluorine bond. The bond is labeled as "the strongest in organic chemistry,"[1] because fluorine forms the strongest single bond to carbon. Carbon–fluorine bonds can have a bond dissociation energy (BDE) of up to 544 kJ/mol.[2] The BDE (strength of the bond) is higher than other carbon–halogen and carbon–hydrogen bonds. For example, the molecule represented by CH3X has a BDE of 115 kcal/mol for carbon–fluorine while values of 104.9, 83.7, 72.1, and 57.6 kcal/mol represent carbon–X bonds to hydrogen, chlorine, bromine, and iodine, respectively.[3] Bond lengthThe carbon–fluorine bond length is typically about 1.35 ångström (1.39 Å in fluoromethane).[1] It is shorter than any other carbon–halogen bond, and shorter than single carbon–nitrogen and carbon–oxygen bonds, despite fluorine having a larger atomic mass. The short length of the bond can also be attributed to the ionic character/electrostatic attractions between the partial charges on carbon and fluorine. The carbon–fluorine bond length varies by several hundredths of an ångstrom depending on the hybridization of the carbon atom and the presence of other substituents on the carbon or even in atoms farther away. These fluctuations can be used as indication of subtle hybridization changes and stereoelectronic interactions. The table below shows how the average bond length varies in different bonding environments (carbon atoms are sp3-hybridized unless otherwise indicated for sp2 or aromatic carbon).
The variability in bond lengths and the shortening of bonds to fluorine due to their partial ionic character are also observed for bonds between fluorine and other elements, and have been a source of difficulties with the selection of an appropriate value for the covalent radius of fluorine. Linus Pauling originally suggested 64 pm, but that value was eventually replaced by 72 pm, which is half of the fluorine–fluorine bond length. However, 72 pm is too long to be representative of the lengths of the bonds between fluorine and other elements, so values between 54 pm and 60 pm have been suggested by other authors.[5][6][7][8] Bond strength effect of geminal bondsWith increasing number of fluorine atoms on the same (geminal) carbon the other bonds become stronger and shorter. This can be seen by the changes in bond length and strength (BDE) for the fluoromethane series, as shown on the table below; also, the partial charges (qC and qF) on the atoms change within the series.[2] The partial charge on carbon becomes more positive as fluorines are added, increasing the electrostatic interactions, and ionic character, between the fluorines and carbon.
Gauche effectThere are two main explanations for the gauche effect: hyperconjugation and bent bonds. In the hyperconjugation model, the donation of electron density from the carbon–hydrogen σ bonding orbital to the carbon–fluorine σ* antibonding orbital is considered the source of stabilization in the gauche isomer. Due to the greater electronegativity of fluorine, the carbon–hydrogen σ orbital is a better electron donor than the carbon–fluorine σ orbital, while the carbon–fluorine σ* orbital is a better electron acceptor than the carbon–hydrogen σ* orbital. Only the gauche conformation allows good overlap between the better donor and the better acceptor.[11] Key in the bent bond explanation of the gauche effect in difluoroethane is the increased p orbital character of both carbon–fluorine bonds due to the large electronegativity of fluorine. As a result, electron density builds up above and below to the left and right of the central carbon–carbon bond. The resulting reduced orbital overlap can be partially compensated when a gauche conformation is assumed, forming a bent bond. Of these two models, hyperconjugation is generally considered the principal cause behind the gauche effect in difluoroethane.[1][12] SpectroscopyThe carbon–fluorine bond stretching appears in the infrared spectrum between 1000 and 1360 cm−1. The wide range is due to the sensitivity of the stretching frequency to other substituents in the molecule. Monofluorinated compounds have a strong band between 1000 and 1110 cm−1; with more than one fluorine atoms, the band splits into two bands, one for the symmetric mode and one for the asymmetric.[13] The carbon–fluorine bands are so strong that they may obscure any carbon–hydrogen bands that might be present.[14] Organofluorine compounds can also be characterized using NMR spectroscopy, using carbon-13, fluorine-19 (the only natural fluorine isotope), or hydrogen-1 (if present). The chemical shifts in 19F NMR appear over a very wide range, depending on the degree of substitution and functional group. The table below shows the ranges for some of the major classes.[15]
See also
References1. ^1 2 {{cite journal |author=O'Hagan D |title=Understanding organofluorine chemistry. An introduction to the C–F bond |journal=Chem Soc Rev |volume=37 |issue=2 |pages=308–19 |date=February 2008 |pmid=18197347 |doi=10.1039/b711844a |url=}} {{ChemicalBondsToCarbon}}{{DEFAULTSORT:Carbon-fluorine bond}}2. ^1 Lemal DM. [https://www.ncbi.nlm.nih.gov/pubmed/14703372 "Perspective on Fluorocarbon Chemistry"] J Org Chem. 2004, volume 69, p 1–11. {{doi|10.1021/jo0302556}} 3. ^{{cite journal |vauthors=Blanksby SJ, Ellison GB |title=Bond dissociation energies of organic molecules |journal=Acc. Chem. Res. |volume=36 |issue=4 |pages=255–63 |date=April 2003 |pmid=12693923 |doi=10.1021/ar020230d |url=|citeseerx=10.1.1.616.3043 }} 4. ^F. H. Allen, O. Kennard, D. G. Watson, L. Brammer, A. G. Orpen. Tables of bond Lengths determined by X-Ray and Neutron Diffraction. Part 1. Bond Lengths in Organic Compounds. J. Chem. Soc. Perkin Trans. II 1987, S1-S19. 5. ^{{cite journal | author = Gillespie Ronald | author2 = Robinson Edward | year = 1992 | title = Bond Lengths in Covalent Fluorides. A New Value for the Covalent Radius of Fluorine | url = | journal = Inorganic Chemistry | volume = 31 | issue = 10| pages = 1960–1963 | doi=10.1021/ic00036a045}} 6. ^{{cite journal | author = Robinson Edward | author2 = Johnson Samuel | author3 = Tang Ting-Hua | author4 = Gillespie Ronald | year = 1997 | title = Reinterpretation of the Lengths of Bonds to Fluorine in Terms of an Almost Ionic Model | url = | journal = Inorganic Chemistry | volume = 36 | issue = 14| pages = 3022–3030 | doi=10.1021/ic961315b | pmid=11669953}} 7. ^{{cite journal | author = Cordero Beatriz | author2 = Gómez Verónica | author3 = Platero-Prats Ana E | author4 = Revés Marc | author5 = Echeverría Jorge | author6 = Cremades Eduard | author7 = Barragán Flavia | author8 = Alvarez Santiago | year = 2008| title = Covalent radii revisited | url = | journal = Dalton Trans | volume = 2008 | issue = 21| pages = 2832–2838 | doi = 10.1039/b801115j | pmid = 18478144 }} 8. ^{{cite journal | author = Pyykkö P. | author2 = Atsumi M. | year = 2009 | title = Molecular Single-Bond Covalent Radii for Elements 1-118| url = | journal=Chemistry: A European Journal| volume = 15 | issue = 1| pages = 186–197 | doi = 10.1002/chem.200800987 | pmid = 19058281 }} 9. ^Contribution to the Study of the Gauche Effect. The Complete Structure of the Anti Rotamer of 1,2-Difluoroethane Norman C. Craig, Anthony Chen, Ki Hwan Suh, Stefan Klee, Georg C. Mellau, Brenda P. Winnewisser, and Manfred Winnewisser J. Am. Chem. Soc.; 1997; 119(20) pp 4789 - 4790; (Communication) {{DOI|10.1021/ja963819e}} 10. ^The stereochemical consequences of electron delocalization in extended .pi. systems. An interpretation of the cis effect exhibited by 1,2-disubstituted ethylenes and related phenomena Richard C. Bingham J. Am. Chem. Soc.; 1976; 98(2); 535-540 Abstract 11. ^Alabugin, I. V. Stereoelectronic Effects: the Bridge between Structure and Reactivity. John Wiley & Sons Ltd, Chichester, UK, 2016 12. ^Goodman, L.; Gu, H.; Pophristic, V.. Gauche Effect in 1,2-Difluoroethane. Hyperconjugation, Bent Bonds, Steric Repulsion. J. Phys. Chem. A. 2005, 109, 1223-1229. {{doi|10.1021/jp046290d}} 13. ^{{cite book |title=Infrared and Raman characteristic group frequencies: tables and charts |author=George Socrates |author2=Socrates |publisher=John Wiley and Sons |year=2001 |isbn=978-0-470-09307-8 |pages=198}} 14. ^{{cite book |title=Infrared Spectroscopy: Fundamentals and Applications |author=Barbara H. Stuart |publisher=John Wiley and Sons |year=2004 |isbn=978-0-470-85428-0 |pages=82}} 15. ^{{cite web|url=http://nmr.chem.indiana.edu/NMRguide/misc/19Fshifts.html |title=Archived copy |accessdate=2008-11-09 |deadurl=yes |archiveurl=https://web.archive.org/web/20080515131118/http://nmr.chem.indiana.edu/NMRguide/misc/19Fshifts.html |archivedate=2008-05-15 |df= }} 2 : Fluorine|Organic chemistry |
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