The carbon–fluorine bond is a polar covalent bond between carbon and fluorine that is a component of all organofluorine compounds. It is one of the strongest single bonds in chemistry (after the B–F single bond, Si–F single bond, and 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 strength
The high electronegativity of fluorine (4.0 for fluorine vs. 2.5 for carbon) gives the carbon–fluorine bond a significant polarity or 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," because fluorine forms the strongest single bond to carbon. Carbon–fluorine bonds can have a bond dissociation energy (BDE) of up to 130 kcal/mol. The BDE (strength of the bond) of C–F is higher than other carbon–halogen and carbon–hydrogen bonds. For example, the BDEs of the C–X bond within a CH3–X molecule is 115, 104.9, 83.7, 72.1, and 57.6 kcal/mol for X = fluorine, hydrogen, chlorine, bromine, and iodine, respectively.
Bond length
The carbon–fluorine bond length is typically about 1.35 ångström (1.39 Å in fluoromethane). It is shorter than any other carbon–halogen bond, and shorter than single carbon–nitrogen and carbon–oxygen bonds. The short length of the bond can also be attributed to the ionic character of the bond (the electrostatic attractions between the partial charges on the carbon and the 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).
{ | class="wikitable" |- ! Bond !! Mean bond length (Å) |- | CCH2F, C2CHF || 1.399 |- | C3CF || 1.428 |- | C2CF2, H2CF2, CCHF2 || 1.349 |- | CCF3 || 1.346 |- | FCNO2 || 1.320 |- | FCCF || 1.371 |- | Csp2F || 1.340 |- | CarF || 1.363 |- | FCarCarF || 1.340 |}
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.
Bond strength effect of geminal bonds
With 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. The partial charge on carbon becomes more positive as fluorines are added, increasing the electrostatic interactions, and ionic character, between the fluorines and carbon.
{ | class="wikitable" |- !Compound !C-F bond length (Å) !BDE (kcal/mol) !qC !qF |- |CH3F |1.385 |109.9 ± 1 |0.01 | −0.23 |- |CH2F2 |1.357 |119.5 |0.40 | −0.23 |- |CHF3 |1.332 |127.5 |0.56 | −0.21 |- |CF4 |1.319 |130.5 ± 3 |0.72 | −0.18 |}