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Looking at electronegativity tables, chlorine consistently has a higher electronegativity than carbon. However, when I use simple algorithms (molcalc, acc2), I get a positive partial charge on the chlorine atoms in $\ce{CCl4}$ and $\ce{CHCl3}$. This correlates with chloroform having a lower dipole moment than chloromethane (see this question and this duplicate question).

Is this an artifact of the simple algorithms? Or is this a reflection of different electron distribution around the carbon:chlorine bond in the different chloroalkanes?

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  • $\begingroup$ en.wikipedia.org/wiki/Sigma_hole_interactions ? $\endgroup$
    – Mithoron
    Commented Nov 18, 2023 at 18:17
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    $\begingroup$ I mean this positive partial charge may represent this sigma hole. $\endgroup$
    – Mithoron
    Commented Nov 18, 2023 at 22:59
  • $\begingroup$ Molcalc appears to use a polarizable continuum model to model a solvent (assumed to be water) and a PM3 semi-empirical model for the solute embedded in the solvent. A PCM models the solvent as a dielectric, with counter-charges in the solvent mirroring those of the solute. The use of a PCM model would change the charge distribution in the solute. However I don't see how this would lead to the observed partial charges. One way to check this is to perform a calculation on a bigger molecule? $\endgroup$
    – Buck Thorn ♦
    Commented Nov 19, 2023 at 13:26
  • $\begingroup$ @BuckThorn According to doi.org/10.1021/ed400164n in MolCalc "These properties are computed using the GAMESS program at either the RHF/STO-3G (orbitals and orbital energies) or PM3 level of theory (all other properties) " From the question it's unclear what is meant by a "positive partial charge on the chlorine atoms", but my guess would be from a Mulliken analysis on the output of the SCF - in which case given the miserable quality of the basis set I wouldn't trust a word of it. Population analyses are notoriously basis set dependent, especially for poor quality ones. $\endgroup$
    – Ian Bush
    Commented Nov 19, 2023 at 15:48
  • $\begingroup$ @IanBush So is there a good source of reliable partial charge values for these molecules? From the difference in dipole moments, we know that polarization of the bonds is different in chloroform vs chloromethane, but it is not clear which bond "to blame" for this difference. $\endgroup$
    – Karsten ♦
    Commented Nov 19, 2023 at 17:29

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Calculation issues aside, chlorine may become positively charged through pi back-donation similar to the role of the oxygen in trifluoramine oxide.

In this interaction, formally nonbonding electron pairs on the chlorine atoms overlap antibonding orbitals on adjacent sigma bonds, weakening the sigma bonds but creating pi bonding in return. We may render this as a delocalization of the bonding, as in carbon tetrachloride:

$\ce{Cl–CCl2–Cl <-> [\overset{+}{Cl}=CCl2]Cl^-}$

Because of the polarity of the carbon-chlorine bond, the carbon-chlorine antibonds are a better pi acceptor than carbon-hydrogen ones. So the additional chlorine atoms in more chlorinated molecules strengthen the pi interactions of each one. Since the molecular-orbital model (unlike the valence-bond structures given above) has both carbon and chlorine accepting the charge from this pi back-donation, the net effect is to transfer more electronic charge from chlorine to carbon in chloroform or carbon tetrachloride versus chloromethane.

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