Question

Strong base is dissolved in 645 mL of 0.400 M weak acid (Ka = 4.91 × 10-5) to make a buffer with a pH of 4.11. Assume that the volume remains constant when the base is added. Calculate the pKa value of the acid and determine the number of moles of acid initially present. When the reaction is complete, what is the concentration ratio of conjugate base to acid? How many moles of strong base were initially added?

Answer 1

Ka of weak acid =4.91*10-5

PKa= -log (4.91*10-5)=4.31

Moles of acid initially present = 0.4*645/1000 moles=0.258 moles

let the moles of NaOH added =x

HA+ NaOH----> ANa+ H2O

ANa----> A- + Na+

Let the acid be named as [HA]

HA-----> H+ +A-

the Henderson-Hasselbach Eq.

pH= pKa+ log [A-/[HA]

4.11 = 4.31+log [A-]/[HA]

-0.2= log [A-]/[HA]

[A-]/[HA] =10^-0.2 =0.631

let x= moles of based added, mole of salt added

x/(0.258-x)= 0.631

x= 0.631*0.258- 0.631x

1.631x= 0.631*0.258

x= 0.631*0.258/1.631=0.0998 moles