Question

How do we assess the strength of an acid?

Answer 1

How else but by experiment.........?

Explanation:

Of course we know that `HBr` is a POTENT `"Bronsted acid"`. And all `"Bronsted acids"` are by definition potent `"Lewis acids"`, electron pair acceptors.

Solutions of hydrogen bromide in water are stoichiometric in `"hydronium ion"`:

`HBr(g) +H_2O(l) rarr H_3O^(+) + Br^(-)`.

With the notable exception of `HF`, ALL of the hydrogen halides, `HX`, are strong `"Bronsted acids"`, i.e. the defining equilibrium lies strongly to the RIGHT:

`HX(aq) + H_2O(l) rightleftharpoons H_3O^+ + X^(-)`

Because the fluoride ion is smaller, and more charge dense, and the `H-F` bond is strong, `HF` is a relatively weak Bronsted acid, and fluoride salts give rise to (slightly) basic solutions:

`F^(-) + H_2O(l) rightleftharpoonsHF(aq) + HO^-`

This is well-known to be an entropy NOT an enthalpy phenomenon.