Question

A). A buffer solution is made that is 0.337 M in H₂CO₃ and 0.337 M in KHCO₃.

(1) If K_a for H₂CO₃ is 4.20×10^-7, what is the pH of the buffer solution?  

(2) Write the net ionic equation for the reaction that occurs when 0.075 mol NaOH is added to 1.00 L of the buffer solution

B). A buffer solution is made that is  0.317 M in  H₂C₂O₄ and  0.317 M in  KHC₂O₄.  

(1) If K_a for  H₂C₂O₄ is  5.90×10^-2, what is the pH of the buffer solution?  

(2) Write the  net ionic equation for the reaction that occurs when  0.071 mol  HBr is added to  1.00 L of the buffer solution.  

Use H₃O⁺ instead of H⁺.

  

Answer 1

Pka = -logKa

         = -log4.2*10^-7

         = 6.3767

PH = Pka + log[KHCO3]/[H2CO3]

         = 6.3767 + log0.337/0.337

         = 6.3767

2. naOH -------> Na⁺ + OH^-

       no of moles of H2C2O4 = 0.317

      no of moles of KHC2O4   = 0.317

      Ka = 5.9*10^-2

     PKa = -logka

             = -log0.059   = 1.229

     PH   = Pka + log[KHC2O4]/[H2C2O4]

    PH   = 1.229 + log0.317/0.317

    PH = 1.229

HBr +H2O ------> H3O⁺ + Br^-

0.071 M                0.071M

            = 1.229 + log