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.
