Question

Boric acid is written as H3BO3 or B(OH)3...but if it is written as B(OH)3...why is it still an acid and not a base?

Answer 1

In fact most oxyacids have hydroxyl functions...

Explanation:

Consider nitric acid, i.e. `HNO_3`...the way I would write the Lewis structure is `HO-stackrel+N(=O)O^(-)`...and this is more or less a strong acid...that dissociates according to the equation...

`HO-stackrel+N(=O)O^(-)(aq)+H_2O(l)rarrH_3O^+ + ""^(-)O_2stackrel+N(=O)`

Any resonance isomer of the nitrate anion has 3 atoms with formal charge. But the parent is manifestly an hydroxide....

And we could got to a stronger acid...`"sulfuric acid"`:

`(HO)_2S(=O)_2(aq) + 2H_2Orarr2H_3O^+ +""^(-)(O-)_2S(=O)_2`

As written this is another hydroxide..

Of course the sulfate anion is resonance stabilized but I think you get the idea. Are you happy with this....? And of course there is also `"perchloric acid"`...`HOCl(=O)_3`...another strong acid....these ARE ALL HYDROXIDES....but the `H-O` bond is attenuated and weak...and so the proton pops off to torment the solvent.

And so another generalization....the hydroxides of non-metals feature fairly weak bonds between hydrogen and the oxygen...on the other hand, for metal hydroxides, the `H-O` bond is strong.

Answer 2

It can accept electron density into the empty `2p` orbital of boron, OR it can dissociate an `"H"^(+)`. Analogize with `"BH"_3`:

Those `"OH"^-` cannot dissociate on their own. Instead, we have that it accepts an `"OH"^(-)` from water's autoionization as a Lewis acid:

`2"H"_2"O"(l) rightleftharpoons "H"_3"O"^(+)(aq) + cancel("OH"^(-)(aq))`
`ul("B"("OH")_3(aq) + cancel("OH"^(-)(aq)) rightleftharpoons "B"("OH")_4^(-)(aq))`
`"B"("OH")_3(aq) + 2"H"_2"O"(l) \rightleftharpoons "B"("OH")_4^(-)(aq) + "H"_3"O"^(+)(aq)`

`K_(a,L) = 7.3 xx 10^(-10)`

So the equilibrium constant for the addition of `"OH"^(-)` onto `"B"("OH")_3` (step 2) is quite thermodynamically favorable, at `7.3 xx 10^4` at `25^@ "C"`, making it quite the Lewis acid.

Or, since it has `"OH"^(-)` on it, it can also dissociate an `"H"^(+)`, apparently... which is how we normally know it---as a Bronsted-Lowry acid.

`"B"("OH")_3(aq) + "H"_2"O"(l) rightleftharpoons "BO"("OH")_2^(-)(aq) + "H"_3"O"^(+)(aq)`

or

`"H"_3"BO"_3(aq) + "H"_2"O"(l) rightleftharpoons "H"_2"BO"_3^(-)(aq) + "H"_3"O"^(+)(aq)`

`K_(a,BL) = 5.8 xx 10^(-10)`

However, it's not much of a Bronsted-Lowry acid (`K_(a,BL)` is small).