Re: HF / HI and acid strength

From: Angelo (patrik56_at_libero.it)
Date: 03/16/05 

Date: 16 Mar 2005 05:45:59 -0800

Wilco Oelen wrote:
> Mark1972 wrote:
> > Hello everybody,
> >
> > I am NOT a chemist, but Iīm interested in the subject, thus
> > reading and reading whenever I find the time to.
> > "Greenhorning" several books of anorganic chemistry, I found myself
> quite alone
> > with the (for you maybe and hopefully ;- simple) question :
> > WHY is HF a WEAKER acid than HI, HCl, or HBr ?. Why does the
> > strength of the acid INCREASE while the electronegativity of
> > the halogen atom DECREASES ?
> >
> > Shouldnīt it be the other way round ?
> >
> > I know, it isnīt, but WHY ? --> I guess I lack some of the basics
> :-(((
> >
> > All help very much appreciated,
> >
> > THXalot
> >
> > Mark
> >
> > PS : by the way : Iīm german, IS the term "acid strength"
> appropriate ?
> > --> the lower the pKa is, the "stronger" the acid. Did I get that
> right ?
> > Or do the english / americans use a different term for this ?
> > THX again
>
> I also am not a chemist. I read the answers of the other contributors
> with great interest.
>
> I would like to make this discussion somewhat wider. What puzzles me
> most is the totally different behavior of fluorine, compared to the
> other halogens.

Indeed: some call it a super-halogen, others argue it should be
dubbed as a sub-halogen :-)

> As noted by other contributors, Ka(HF) < Ka(HCl) < Ka(HBr) < Ka(HI),
> but if we look at the actual behavior, then HCl, HBr and HI are quite
> similar with respect to acidity, they all ionize almost 100%, while
HF
> only ionizes slightly.

Right, that is a manifestation of the so called 'levelling effect'.
Increasing the acid strength, and hence its Ka, at a given
starting concentration (Ca) the percentage of dissociation
increases, but cannot eventually trespass 100 p.c. IOW,
the basicity of water (as a solvent, i.e. in bulk) cannot
differentiate the strength of, say, HCl, HBr, HI, etc,
because such basicity is simply too high.
The reported approximate values for those pKas in water
(from HF to HI) are: 3, -7, -9, -10.
Nonetheless, in suitably less basic solvent the intrinsic
correct ordering (qualitatively speaking, 'cause solvation
can be quite different) of the mentioned Kas can be
experimentally demonstrated. The difficulty lies in the
choice of such solvent, which should address at least
the requirements of 1) high dielectric constant,
2) unreactivity (besides acid-base equilibria) towards
redox, substitution, elimination, polymerization, and
other reactions prompted by a high H(+) activity.

As for the weakness of HF in water, please see other
pertinent posts in this thread.

> Another important difference: CaF2 is very insoluble, the calcium
salt
> of the other halogens is very soluble. A similar behavior is
exhibited
> by the lithium salts.
> For AgF and the other silver halogenides it is exactly the other way
> around.

Solubility of salts is somewhat tricky, in that both enthalpy
and entropy come into play. When delta H is negative, there's
should be no conflict (as in MgCl2, or CaCl2). When delta H is
positive the results are by far less easily predicted. In such cases
the HASAB rules may give some good hints.
As a general rule, the more the ionic radii are mismatched in the
crystal structure, the more solubility is expected (LiClO4 vs KClO4;
LiBF4 vs KBF4, and also LiI vs LiF).

> Why is there such a strong discontinuity in behavior between
compounds
> of fluorine and the other halogens?

Most of the anomalies of Fluorine can be traced back to the small
size of F(-), and a very much greater solvation energy (in water).
HF is completely miscible with H2O in any mole fraction, and the
relevant phase diagram shows, besides other cases, the phase
 [HF.4 H2O], whose m.p. is -100 oC, i.e. well below that of neat HF
(-83.5 oC). OTOH, HCl is partially soluble in H2O; when its mole
fraction goes beyond ca. 0.5 the mixture splits in two phases.
Similar behavior for the other hydrohalides.

> Wilco

Best regards,
Angelo

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