Buffers and Titration of Acids and Bases | General Chemistry 3

General conditions to apply the Henderson-Hasselbalch equation:

• K_a should be in the range of about 10^-4 M to 10^-11 M
• The ratio $\frac{{\left[\mathrm{base}\right]}_0}{{\left[\mathrm{acid}\right]}_0}$ must be between 0.1 and 10
• The values of [base]₀ and [acid]₀ should be in the range of 10^-3 M to 1 M

A buffer solution is an aqueous solution consisting of a mixture of a weak acid and its conjugate base. Its pH changes very little when a small amount of strong acid or base is added to it and does not change when the solution is diluted
 

CH₃COOH is a weak acid and CH₃COO^- is its conjugate base
⇒ a mixture of a solution of CH₃COOH and a solution of CH₃COO^- results in a buffer solution
Adding few mL of a strong acid (for example HCl) to this mixture will change the pH of the solution by around 0.1 pH unit

pH indicators: substances which change color with pH ⇒ they are used to signal the end point of titrations
The pH region in which an indicator changes color is equal to the pK_ai for this indicator
 

Bromthymol blue (pK_ai = 7.3) changes from yellow to blue around a pH = 7.3

If the solution is yellow ⇒ acids are in excess
If the solution is blue ⇒ bases are in excess
When the solution changes color ⇒ equivalence point

At the equivalence point: stoichiometrically equivalent amounts of acid and base have reacted

Titration of a strong acid with a strong base (or vice versa): pH changes abruptly at the equivalence point
Titration curve for the titration of a 0.100M HCl (aq) solution with a 0.100M NaOH (aq) solution:
 

Titration of a weak acid with a strong base (or vice versa): pH = pKa at the Midpoint
Titration curve for the titration of a 0.100M CH₃COOH (aq) solution with a 0.100M NaOH (aq) solution: