Editing Borate (section)

=== Aqueous solution ===
In aqueous solution, [[boric acid]] {{chem2|B(OH)3}} can act as a weak [[Brønsted acid]], that is, a [[proton]] donor, with [[acid dissociation constant|p''K''<sub>a</sub> ~&nbsp;9]]. However, it more often acts as a [[Lewis acid]], accepting an [[electron pair]] from a [[hydroxide ion]] produced by the water [[autoprotolysis]]:<ref name=atki2010/>

: {{chem2|B(OH)3}} + 2 {{H2O}} {{Eqm}} {{chem2|[B(OH)4]-}} + {{H3O+}} {{Spaces|21}} (p''K'' = 8.98)<ref name=ingri1962/>

This reaction is very fast, with a characteristic time less than 10 [[microsecond|μs]].<ref name=momi1967/> Polymeric boron oxoanions are formed in aqueous solution of boric acid at [[pH]] 7–10 if the boron concentration is higher than about 0.025&nbsp;mol/L. The best known of these is the [[tetraborate]] ion {{chem2|[B4O7](2-)}}, found in the mineral borax:

: 4 {{chem2|[B(OH)4]-}} + 2 {{H+}} {{eqm}} {{chem2|[B4O5(OH)4](2-)}} + 7 {{H2O}}

Other anions observed in solution are triborate(1−) and pentaborate(1−), in equilibrium with boric acid and tetrahydroxyborate according to the following overall reactions:<ref name=momi1967/>

: 2 {{chem2|B(OH)3}} + {{chem2|[B(OH)4]-}} {{Eqm}} {{chem2|[B3O3(OH)4](-)}} + 3 {{chem2|H2O}} {{Spaces|5}} (fast, p''K'' = −1.92)

: 4 {{chem2|B(OH)3}} + {{chem2|[B(OH)4]-}} {{Eqm}} {{chem2|[B5O6(OH)4](-)}} + 6 {{chem2|H2O}} {{Spaces|5}} (slow, p''K'' = −2.05)

In the [[pH]] range 6.8 to 8.0, any alkali salts of "boric oxide" anions with general formula {{chem2|[B_{''x''}O_{''y''}(OH)_{''z''}]((''q''-)}} where 3''x'' + ''q'' = 2''y'' + ''z'' will eventually equilibrate in solution to a mixture of {{chem2|B(OH)3}}, {{chem2|[B(OH)4](-)}}, {{chem2|[B3O3(OH)4](-)}}, and {{chem2|[B5O6(OH)4](-)}}.<ref name=momi1967/>

Like the complexed borates mentioned above, these ions are more acidic than boric acid. As a result, the pH of a concentrated polyborate solution will increase more than expected when diluted with water.