The carbonâfluorine 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 BâF single bond, SiâF single bond, and HâF 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.
The high electronegativity of fluorine (4.0 for fluorine vs. 2.5 for carbon) gives the carbonâfluorine 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 (CÎŽ+—FÎŽ−). The partial charges on the fluorine and carbon are attractive, contributing to the unusual bond strength of the carbonâfluorine bond. The bond is labeled as "the strongest in organic chemistry," [1] because fluorine forms the strongest single bond to carbon. Carbonâfluorine bonds can have a bond dissociation energy (BDE) of up to 130 kcal/mol. [2] The BDE (strength of the bond) of CâF is higher than other carbonâ halogen and carbonâ hydrogen bonds. For example, the BDEs of the CâX bond within a CH3âX molecule is 115, 104.9, 83.7, 72.1, and 57.6 kcal/mol for X = fluorine, hydrogen, chlorine, bromine, and iodine, respectively. [3]
The carbonâfluorine bond length is typically about 1.35 Ă„ngström (1.39 Ă in fluoromethane). [1] It is shorter than any other carbonâhalogen bond, and shorter than single carbonâ nitrogen and carbonâ 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 carbonâfluorine bond length varies by several hundredths of an Ă„ngstrom 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).
Bond | Mean bond length (Ă ) [4] |
---|---|
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 fluorineâfluorine 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]
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. [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.
Compound | C-F bond length (Ă ) | BDE (kcal/mol) | qC | qF |
---|---|---|---|---|
CH3F | 1.385 | 109.9 ± 1 | 0.01 | â0.23 |
CH2F2 | 1.357 | 119.5 | 0.40 | â0.23 |
CHF3 | 1.332 | 127.5 | 0.56 | â0.21 |
CF4 | 1.319 | 130.5 ± 3 | 0.72 | â0.18 |
When two fluorine atoms are in vicinal (i.e., adjacent) carbons, as in 1,2-difluoroethane (H2FCCFH2), the gauche conformer is more stable than the anti conformerâthis is the opposite of what would normally be expected and to what is observed for most 1,2-disubstituted ethanes; this phenomenon is known as the gauche effect. [9] In 1,2-difluoroethane, the gauche conformation is more stable than the anti conformation by 2.4 to 3.4 kJ/mole in the gas phase. This effect is not unique to the halogen fluorine, however; the gauche effect is also observed for 1,2-dimethoxyethane. A related effect is the alkene cis effect. For instance, the cis isomer of 1,2-difluoroethylene is more stable than the trans isomer. [10]
There are two main explanations for the gauche effect: hyperconjugation and bent bonds. In the hyperconjugation model, the donation of electron density from the carbonâhydrogen Ï bonding orbital to the carbonâfluorine Ï* antibonding orbital is considered the source of stabilization in the gauche isomer. Due to the greater electronegativity of fluorine, the carbonâhydrogen Ï orbital is a better electron donor than the carbonâfluorine Ï orbital, while the carbonâfluorine Ï* orbital is a better electron acceptor than the carbonâhydrogen Ï* 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 carbonâfluorine 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 carbonâcarbon 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]
The carbonâfluorine bond stretching appears in the infrared spectrum between 1000 and 1360 cmâ1. 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 cmâ1; with more than one fluorine atoms, the band splits into two bands, one for the symmetric mode and one for the asymmetric. [13] The carbonâfluorine bands are so strong that they may obscure any carbonâhydrogen 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]
Type of Compound | Chemical Shift Range (ppm) Relative to neat CFCl3 |
---|---|
FâC=O | â70 to â20 |
CF3 | +40 to +80 |
CF2 | +80 to +140 |
CF | +140 to +250 |
ArF | +80 to +170 |
Breaking CâF bonds is of interest as a way to decompose and destroy organofluorine " forever chemicals" such as PFOA and perfluorinated compounds (PFCs). Candidate methods include catalysts, such as platinum atoms; [16] photocatalysts; UV, iodide, and sulfite, [17] radicals; etc.
The carbonâfluorine 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 BâF single bond, SiâF single bond, and HâF 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.
The high electronegativity of fluorine (4.0 for fluorine vs. 2.5 for carbon) gives the carbonâfluorine 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 (CÎŽ+—FÎŽ−). The partial charges on the fluorine and carbon are attractive, contributing to the unusual bond strength of the carbonâfluorine bond. The bond is labeled as "the strongest in organic chemistry," [1] because fluorine forms the strongest single bond to carbon. Carbonâfluorine bonds can have a bond dissociation energy (BDE) of up to 130 kcal/mol. [2] The BDE (strength of the bond) of CâF is higher than other carbonâ halogen and carbonâ hydrogen bonds. For example, the BDEs of the CâX bond within a CH3âX molecule is 115, 104.9, 83.7, 72.1, and 57.6 kcal/mol for X = fluorine, hydrogen, chlorine, bromine, and iodine, respectively. [3]
The carbonâfluorine bond length is typically about 1.35 Ă„ngström (1.39 Ă in fluoromethane). [1] It is shorter than any other carbonâhalogen bond, and shorter than single carbonâ nitrogen and carbonâ 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 carbonâfluorine bond length varies by several hundredths of an Ă„ngstrom 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).
Bond | Mean bond length (Ă ) [4] |
---|---|
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 fluorineâfluorine 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]
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. [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.
Compound | C-F bond length (Ă ) | BDE (kcal/mol) | qC | qF |
---|---|---|---|---|
CH3F | 1.385 | 109.9 ± 1 | 0.01 | â0.23 |
CH2F2 | 1.357 | 119.5 | 0.40 | â0.23 |
CHF3 | 1.332 | 127.5 | 0.56 | â0.21 |
CF4 | 1.319 | 130.5 ± 3 | 0.72 | â0.18 |
When two fluorine atoms are in vicinal (i.e., adjacent) carbons, as in 1,2-difluoroethane (H2FCCFH2), the gauche conformer is more stable than the anti conformerâthis is the opposite of what would normally be expected and to what is observed for most 1,2-disubstituted ethanes; this phenomenon is known as the gauche effect. [9] In 1,2-difluoroethane, the gauche conformation is more stable than the anti conformation by 2.4 to 3.4 kJ/mole in the gas phase. This effect is not unique to the halogen fluorine, however; the gauche effect is also observed for 1,2-dimethoxyethane. A related effect is the alkene cis effect. For instance, the cis isomer of 1,2-difluoroethylene is more stable than the trans isomer. [10]
There are two main explanations for the gauche effect: hyperconjugation and bent bonds. In the hyperconjugation model, the donation of electron density from the carbonâhydrogen Ï bonding orbital to the carbonâfluorine Ï* antibonding orbital is considered the source of stabilization in the gauche isomer. Due to the greater electronegativity of fluorine, the carbonâhydrogen Ï orbital is a better electron donor than the carbonâfluorine Ï orbital, while the carbonâfluorine Ï* orbital is a better electron acceptor than the carbonâhydrogen Ï* 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 carbonâfluorine 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 carbonâcarbon 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]
The carbonâfluorine bond stretching appears in the infrared spectrum between 1000 and 1360 cmâ1. 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 cmâ1; with more than one fluorine atoms, the band splits into two bands, one for the symmetric mode and one for the asymmetric. [13] The carbonâfluorine bands are so strong that they may obscure any carbonâhydrogen 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]
Type of Compound | Chemical Shift Range (ppm) Relative to neat CFCl3 |
---|---|
FâC=O | â70 to â20 |
CF3 | +40 to +80 |
CF2 | +80 to +140 |
CF | +140 to +250 |
ArF | +80 to +170 |
Breaking CâF bonds is of interest as a way to decompose and destroy organofluorine " forever chemicals" such as PFOA and perfluorinated compounds (PFCs). Candidate methods include catalysts, such as platinum atoms; [16] photocatalysts; UV, iodide, and sulfite, [17] radicals; etc.