These types of titrations revolve around the two Ka values of the diprotic acid.
Concept #1: With the presence of 2 equivalence points we must now determine 2 equivalence volumes.
Concept #2: Before the first equivalence point has been reached we have the formation of a buffer region.
Concept #3: At the first equivalence point we now have an excess of the intermediate form for the diprotic acid.
Concept #4: Before the second equivalence point has been reached we have the formation of another buffer region.
Concept #5: At the second equivalence point we now have an excess of the basic form for the diprotic acid.
Concept #6: After the second equivalence point there is an excess of strong base remaining.
Example #1: Calculate the pH of 100 mL of a 0.25 M H2CO3 when 70.0 mL of 0.25 M NaOH are added. Ka1 = 4.3 x 10-7 and Ka2 = 5.6 x 10-11.
Example #2: Calculate the pH of 75.0 mL of a 0.10 M of phosphorous acid, H3PO3, when 80.0 mL of 0.15 M NaOH are added. Ka1 = 5.0 x 10-2, Ka2 = 2.0 x 10-7.
Example #3: Find the pH when 100.0 mL of a 0.1 M dibasic compound B (pKb1 = 4.00; pKb2 = 8.00) was titrated with 11 mL of a 1.00 M HCl.
Example #4: Carbonic acid, H2CO3, is a diprotic acid that dissociates, losing its two protons, to create bicarbonate, HCO3–, and carbonate, CO32–, according to the following reactions given below:
H2CO3 (aq) ⇌ HCO3– (aq) + H+ (aq) ⇌ CO32– (aq) + H+ (aq)
As a diprotic acid system, it has two dissociation constants that pKa1 = 6.30 and pKa2 = 10.30 for the two steps. In the reaction you titrate 50.0 mL solution of 0.50 M H2CO3 with 1.00 M solution of NaOH. What would be the expected pH after the addition of 35 mL of the NaOH titrant?