Question

Phosphoric acid, H3PO4, is tribasic with pKa values of 2.14, 6.86, and 12.4. The ionic form that predominates at pH 9.0is:

H3PO4+ H2O ⇌H3O++ H2PO4-⇌H3O++ HPO42-⇌H3O++ PO43-

the answer is HPO42–

THANKS!

Answer 1

a)H3PO4+ H2O ⇌H3O++ H2PO4-

Ka1 = [H3O+]. [H2PO4-]/ [H3PO4]

Given pKa1 = 2.14

-log(Ka1) = 2.14

Ka1 = 0.00724

Hence, [H3O+]. [H2PO4-] = 0.00724

=> [H3O+]= square root of (0.00724) = 0.0851

therefore pH = -log(0.0851)

                    =1.07

b)H2PO4-⇌H3O++ HPO42-

Ka2 = [H3O+]. [HPO42-]/ [H2PO4-]

Given pKa2 = 6.86

-log(Ka2) = 6.86

Ka2 = 1.38x10^-7

Hence, [H3O+]. [HPO42-] = 1.38x10^-7

=> [H3O+]= square root of (1.38x10^-7 ) = 3.71x10^-4

Hence, [HPO42-]=3.71x10^-4

therefore pH = -log(3.71x10^-4)

                    =3.43

c)HPO42-⇌H3O++ PO43-

Ka3 = [H3O+]. [PO43-]/ [HPO42-]

Given pKa3 = 12.4

-log(Ka3) = 12.4

Ka3 = 3.98x10^-10

Hence, [H3O+]. [PO43-] = 3.98x10^-12

=> [H3O+]= square root of (3.98x10^-12) = 1.99x10^-6

[PO43-] = 1.99x10^-6

therefore pH = -log(1.99x10^-6)

                    =5.7

Use Handerson's equation to find the predominant ionic form at the given pH.

pH =pKa3 + log[salt/acid]

= 12.4+log[PO43-/HPO42-]

= 12.4+ log[1.99x10^-6/3.71x10^-4]

=10.1

Hence, the predominant ionic form is HPO42-