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Why does fluorine make compounds volatile?

  1. Jan 19, 2007 #1
    In medicinal chemistry fluorine atoms are often used on scaffolds in order to increase metabolic stability of a particular compound. The problem with working with reactants that contain fluorine is that they are a lot of times very volatile. For example, I am thinking along the lines of coupling a small carboxylic acid that contains a fluorine in its structure with an amine. I can't count the number of times that the yields of my reaction have been very poor because the acid chloride I made (which contained a fluorine) got sucked up into my rotavap because it was very volatile. So why is it that introducing fluorines makes certain types of things more volatile? I would think that with fluorines the higher mw and possible H bonding would make evaporating them harder but a lot of times it is the exact opposite.

    Here is a list of volatile organic compounds I found:
    reactivity: methane; ethane; methylene chloride (dichloromethane); 1,1,1-trichloroethane (methyl chloroform); 1,1,2-trichloro-1,2,2-trifluoroethane (CFC–113); trichlorofluoromethane (CFC–11); dichlorodifluoromethane (CFC–12); chlorodifluoromethane (HCFC–22); trifluoromethane (HFC–23); 1,2-dichloro-1,1,2,2-tetrafluoroethane (CFC–114); chloropentafluoroethane (CFC–115); 1,1,1-trifluoro-2,2-dichloroethane (HCFC–123); 1,1,1,2-tetrafluoroethane (HFC–134a); 1,1-dichloro-1-fluoroethane (HCFC–141b); 1-chloro-1,1-difluoroethane (HCFC–142b); 2-chloro-1,1,1,2-tetrafluoroethane (HCFC–124); pentafluoroethane (HFC–125); 1,1,2,2-tetrafluoroethane (HFC–134); 1,1,1-trifluoroethane (HFC–143a); 1,1-difluoroethane (HFC–152a); parachlorobenzotrifluoride (PCBTF); cyclic, branched, or linear completely methylated siloxanes; acetone; perchloroethylene (tetrachloroethylene); 3,3-dichloro-1,1,1,2,2-pentafluoropropane (HCFC–225ca); 1,3-dichloro-1,1,2,2,3-pentafluoropropane (HCFC–225cb); 1,1,1,2,3,4,4,5,5,5-decafluoropentane (HFC 43–10mee); difluoromethane (HFC–32); ethylfluoride (HFC–161); 1,1,1,3,3,3-hexafluoropropane (HFC–236fa); 1,1,2,2,3-pentafluoropropane (HFC–245ca); 1,1,2,3,3-pentafluoropropane (HFC–245ea); 1,1,1,2,3-pentafluoropropane (HFC–245eb); 1,1,1,3,3-pentafluoropropane (HFC–245fa); 1,1,1,2,3,3-hexafluoropropane (HFC–236ea); 1,1,1,3,3-pentafluorobutane (HFC–365mfc); chlorofluoromethane (HCFC–31); 1- chloro-1-fluoroethane (HCFC–151a); 1,2-dichloro-1,1,2-trifluoroethane (HCFC–123a); 1,1,1,2,2,3,3,4,4-nonafluoro-4-methoxy-butane (C4F9OCH3 or HFE-7100); 2-(difluoromethoxymethyl)-1,1,1,2,3,3,3-heptafluoropropane ((CF3)2CFCF2OCH3); 1-ethoxy-1,1,2,2,3,3,4,4,4-nonafluorobutane (C4F9OC2H5 or HFE-7200); 2-(ethoxydifluoromethyl)-1,1,1,2,3,3,3-heptafluoropropane ((CF3)2CFCF2OC2H5); methyl acetate; 1,1,1,2,2,3,3-heptafluoro-3-methoxy-propane (n-C3F7OCH3 or HFE–7000); 3-ethoxy-1,1,1,2,3,4,4,5,5,6,6,6-dodecafluoro-2-(trifluoromethyl) hexane (HFE–7500); 1,1,1,2,3,3,3-heptafluoropropane (HFC 227ea); methyl formate (HCOOCH3); and perfluorocarbon compounds which fall into these classes:
    (i) Cyclic, branched, or linear, completely fluorinated alkanes;
    (ii) Cyclic, branched, or linear, completely fluorinated ethers with no unsaturations;
    (iii) Cyclic, branched, or linear, completely fluorinated tertiary amines with no unsaturations; and
    (iv) Sulfur containing perfluorocarbons with no unsaturations and with sulfur bonds only to carbo

  2. jcsd
  3. Jan 21, 2007 #2


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    It may have something to do with decreasing the Van Der Wall interactions between organic compounds which would otherwise be more soluble. Also, it may simply be the case that the synthesis and thus the existence for the fluoro-organic compounds are more prevalent, considering the fact that the chlorine based ones are ozone depleting as well as toxic. I'm going to need to do some research on the relative toxicity regarding the Fluorine compounds.
  4. Jan 21, 2007 #3
    Yes also a low atomic radius, which contributes to the Van Der Walls, if you consider the mean free path of a molecule, those with a high electronegativity and small radii have the largest mean free path and lowest viscosities. For molecules, one would consider the bond length also, which is also low for the C-F bond.

    C-F -> 135pm (shortest of all C-X single bonds)
    C=N -> 129pm
    C(triple)N -> 116pm

    EDIT: btw H-Bonding is lower in Flouro compounds because the atomic radii and dipole moments are less, compared with say Oxygen, which has lone pairs and a greater affinity for H-bonding.

    I think your correct with the higher MW being a contributing factor to viscosity (positive), inertial, ive not done the math though, and i guess it must be offset by the other factors.
    Last edited: Jan 21, 2007
  5. Jan 22, 2007 #4


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    You've gotten it backwards. Fluorine substitution is making them less volatile by increasing dipole moments.
    Last edited: Jan 22, 2007
  6. Jan 22, 2007 #5
    Maybe I am misunderstanding something?

    From wiki (take this with a grain of salt):

    "As a result of these unique features of the carbon-fluorine bond, an overarching theme in fluorocarbon chemistry is the contrasting set of physical and chemical properties in comparison to the corresponding hydrocarbons. Case studies follow....

    The molecule hexafluoroacetone ((CF3)2CO), the fluoro-analogue of acetone, has a boiling point of −27 °C compared to +55 °C for acetone itself. This difference illustrates one of the remarkable effects of replacing C-H bonds with C-F bonds. Normally, the replacement of H atoms with heavier halogens results in elevated boiling points due to increased van der Waals interactions between molecules. "

    It seems that fluorination, even though it is expected to increase dipole moments of a compound, causes certain things to increase in volatility (as described above).

    H-bonding maybe lower in flouro compounds, but H-bonding may nonetheless exist in carbonfluoro compounds. I would expect that they would have a higher bp than their C-H counterparts because of some H-bonding/vander waal forces, but this doesn't seem to be the case.
    Last edited: Jan 22, 2007
  7. Jan 22, 2007 #6


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    1. Can I see a link to the wiki?

    2. Draw the structure of hexafluoroacetone, and then the structure of acetone. It's then a trivial observation that the hexafluoro-substitution merely introduces a pair of dipole moments that oppose the existing dipole moment from the C=O, thus it's more than plausible there's a reduction of the net dipole moment producing the expected result.

    3. Of the bazillion compounds listed in the OP, very few of them are fluorinated ketones, aldehydes, esters or ethers. Those that are, are derived from hexane or higher (boiling points above room temperature); it's very plausible that something like what's happening with hexafluoroacetone happens with these too. I'm not about to go about drawing structures for them to look for anomalies.

    4. All the rest (making up the vast majority) of the fluorinated/chlorinated compounds are derived from alkanes (1 from butane, most from propane and ethane, and a half-dozen or so from methane), with the exception of 2 compounds (1 of which is derived from ethylene).

    Boiling Points of the parent compounds (approx):
    methane: -162*C
    ethane: -89*C
    ethylene: -104*C
    propane: -42*C
    butane: 0*C

    Now, I'm not about to look up all the BPs for the substituted compounds, but I expect them to be no less than their parents above. If you find they're not, then there's something interesting here. Else, it's business as usual.
    Last edited: Jan 22, 2007
  8. Jan 22, 2007 #7
    Ahh thats probably what is going on. Reduction in the net dipole of the entire compound which would reduce london forces. Man, I haven't drawn a structure out in years, this should be a good exercise. Thanks.
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