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Why is the bond between H and F in HF considered polar covalent whereas HCl, HI, and so on are all ionic? The electronegativity difference between them suggests that it too should be ionic, yet all textbooks say that HF is covalent. Is there a good reason why?

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I don't think I've ever seen a reputable source claiming the bonds in HCl, HBr and HI are mainly ionic. The fact that they are all gasses at room temperature clearly suggests otherwise. HF has the most ionic character of all the hydrogen halides, but it too has a rather low boiling point (just below room temperature), which is uncharacteristic of ionic compounds.

It is best to think of all hydrogen halides as covalent polar molecules, with the polarity increasing in the order HI < HBr < HCl < HF, as suggested by the electronegativity differences.

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    $\begingroup$ From the Wikipedia article on hydrogen chloride: "Hydrogen chloride is a diatomic molecule, consisting of a hydrogen atom H and a chlorine atom Cl connected by a covalent single bond." You should find similar sentences in the articles for the other compounds. $\endgroup$ – Nicolau Saker Neto Mar 7 '14 at 2:10
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    $\begingroup$ A few words of support of Nicolau's answer: charge separation can be a bad thing. Noting that H has a proton, there is good reason to leave on the order of one electron in its vicinity. This would imply a covalent-like bonding scenario. $\endgroup$ – Eric Brown Jun 1 '14 at 1:13
  • $\begingroup$ @EricBrown Indeed, that's an interesting fundamental reason why an ionic salt containing discrete non-solvated $\ce{H^+}$ ions is all but impossible; the free hydrogen cation has an immense ability to polarize electron clouds (which I have discussed elsewhere), demanding a high degree of covalent bonding with even the most unwilling species. However, I guess that argument does not strictly preclude an ionic $\ce{HX}$ compound with, for example, discrete $\ce{HX2^{-}}$ and $\ce{H2X^{+}}$ ions, at least qualitatively. $\endgroup$ – Nicolau Saker Neto Jun 1 '14 at 14:33
  • $\begingroup$ @EricBrown See, e.g., bifluoride and fluoronium, for examples of $\ce{HX2-}$ and $\ce{H2X+}$ species (more links to related content therein). $\endgroup$ – hBy2Py Jul 30 '15 at 19:51
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As already stated in the comments, this question includes the FALSE assertion that HCl, HI, and HBr are ionic. No other reply is appropriate. https://en.wikipedia.org/wiki/Hydrogen_halide. Ionic bonds are those which can be nearly completely characterized as electrostatic. That is, by the assumption that one ion is positively electrically charged and the other is negatively charged - or by combinations of various integer charged ions resulting in a net attraction between the ions. The key point is that the electron(s) responsible for bonding are concentrated around the anion(s) and the cation(s) have lost one (or more) of them so can be considered to have a positive charge. This is to be contrasted with covalent bonds in which the electrons are shared between atoms in the bond. This simple definition ignores the amount of "sharing" and the amount of "concentration" which occurs. An ionic bond is "mostly" unshared electrons (substantial charge separation) and a covalent bond is "mostly" shared electrons. In fact, most chemical bonds between dissimilar atoms are somewhere between pure covalent and pure ionic. There is no requirement that the compound be in the solid state. I know of no ionic gasses (in their ground state) near room temperatures (how could you explain the separation of charge?). There are many ionic liquids, and of course ionic solids are common.

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According to dipole moment we can get the information about "percentage ionic character of a bond", so for H-F bond the %age of ionic character comes out to be about 43% ionic and 57% covalent. Therefore this bond is a covalent bond.

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