# Why do halogen substituents make molecules more lipophilic?

According to my medicinal chemistry text book, halogens increase a drug’s lipophilicty. This makes no sense to me for two reasons:

1. Halogens are all quite electronegative and they will form a $\ce{C-X}$ bond with a large dipole moment. Large dipole moments are associated with increased polarity.

2. As opposed to a methyl group, halogens are capable of donating their lone pair electrons to participate in hydrogen bonding. We see this when looking at water solubility of methyl-halogen compounds:

$$\begin{array}{cc}\hline \text{Compound} & \text{Solubility} / (\mathrm{g \ L^{-1}})\\ \hline \ce{CH4} & 0.055\\ \ce{CH3CH3} & 0.0568\\ \ce{CH3F} & 2.295\\ \ce{CH3Cl} & 5.325\\ \ce{CH3Br} & 15.22\\ \ce{CH3I} & 14\\ \hline \end{array}$$

Halogens, except for fluorine, aren’t all that electronegative. Chlorine is just slightly less electronegative than nitrogen and bromine is even less again. Iodine almost has carbon’s electronegativity. Thus, there won’t be so much of a dipole moment created by substituting a hydrogen with a halogen. This applies even more if a hydroxy group is replaced with a halogen.

To illustrate this, take dichloromethane. Of all the chlorinated methanes it has the highest permanent dipole moment. Yet it is still almost immiscible with water but fully miscible with hexane. Note by the way that your data somewhat reinforces that: the molecule with the highest dipole moment of all your halocarbons, fluoromethane, has the lowest water solubility.

A second point is that halogens come with free lone pairs. Lone pairs are not polarised per se but can easily be polarised in London interactions. Thus, halogens actually increase molecule’s ability to interact with lipophilic, unpolar media by allowing stronger London forces.

Note also that while nitrogen is capable of forming hydrogen bonds with surrounding aquaeous media, chlorine — in spite of its similar electronegativity — cannot since its lone pairs are much more diffuse. Only fluorine is in theory able to accept hydrogen bonds. In principle, this is another consequence of the larger size of higher-period halogens.

In all this, fluorine is somewhat separate and special. But fluorine is special in itself since perfluorinated molecules are typically insoluble in anything but fluorocarbons. Fluorous extraction is actually a technique used by some green chem labs.

To sum it up, just looking at the simple variable polarity cannot explain the solubility of halocarbons adequately. Their size needs to be considered.

• 1.) Actually Jan, fluoromethanes poor solubility in water has to with its unwillingness to form hydrogen bonds compared with the other halocarbons. In fact, I would bet fluoromethane would also have the lowest solubility in hexane as well (because of dipole moment, cant find data to confirm this though). 2.) Not sure what your point about halogen polarizability is saying. A polar molecule like water should have a much stronger polarizing effect (dipole-dipole bonds) on halogens than nonpolar molecules. London dispersion forces are the weakest of all the intermolecular interactions.
– Nova
Oct 4, 2016 at 0:32
• @Nova 2) polariseablility $\ne$ permanent dipole. Permanent dipoles interact with other permanent dipoles preferentially, and something that is ‘just’ large and polariseable does not interact well with permanent dipoles. Like dissolves like.
– Jan
Oct 4, 2016 at 0:40
• 1) I severely contest the opinion that chloro- and bromomethane accept hydrogen bonds from water.
– Jan
Oct 4, 2016 at 0:41
• Hmm, I thought that due to flourines high electronegativity, it was unwilling to share electron density to form hydrogen bonds. And I thought I- was too disperse to form strong hydrogen bonds. But Br and Cl were in the middle and not too big and not too electronegative.
– Nova
Oct 4, 2016 at 0:49
• @Nova I don’t know about your corner of the world, but in my place in school hydrogen bonds are taught as happening between $\ce{N,O,F}$ and a hydrogen bound to one of those, so … ;)
– Jan
Oct 4, 2016 at 0:52