Question

Assume that 22.00 mL of NaOH (0.100 M) was consumed to titrate a 10.00 mL H2CO3 solution (fully converted to Na2CO3). pKa1= 6.351 and pKa2= 10.329

(a) what is the concentration of the initial H2CO3 solution?

(b) what is the pH of the H2CO3 solution before titration?

(c) which is the principal (dominating) species when 5.00 mL of NaOH has been added?

(d) what is the pH of the solution when 15.00 mL of NaOH has been added?

(e) what are the two most abundant species at pH=10.329?

Answer 1

(a) Millimole of NaOH = 2× millimole of H₂CO₃

22.00ml×0.100M = 2 × 10.00ml × M₂

M₂ = 0.110 mol/L. (Answer)

(b)

[H⁺] = √c.Ka1 = √(0.110×10^-6.351) = 0.0002214

pH = -log (0.0002214) = 3.655 (Answer)

(c)

Millimole of NaOH added = 5.00ml × 0.100M = 0.500mmol

Millimole of H₂CO₃ = 0.110M × 10.00ml = 1.10mmol

NaOH + H₂CO_{3 $\rightarrow$} NaHCO₃ + H₂O

pH = pKa1 + log[salt]/[acid] = 6.351 + log(0.500/0.600)

= 6.272 (Answer)

(d) addition of 11.00ml is the half equivalence point i.e all carbonic acid is converted to NaHCO₃ .

Millimole of Na₂CO₃ formed = 4.00× 0.100 = 0.400mmol

Millimole of NaHCO₃ left = 1.1 - 0.400 = 0.700mmol

pH = 10.329 + log(0.400/0.700) = 10.086 (Answer)

(e)

At pH = pKa2 = 10.329

[CO₃^2-] = [HCO₃^-]

Major two species are carbonate , and bicarbonate ions .