Question

1.Write the equilibrium constant expression (Ka) for the generic weak acid HA.

HA(aq)⇌H+(aq) + A−(aq)

2.Write the Henderson-Hasselbalch equation.

3.Given the Henderson-Hasselbalch equation, under what conditions does the pH= pKa?

4.Sketch a pH versus volume of base curve (a titration curve) for the titration of a weak acid with a strong base. On this sketch indicate

the equivalence point and the point at which the conditions described in #3 are met.

5.When using a buret, do your results depend on the initial volume in the buret? (Or, does the initial volume have to be exactly 0.00 mL?)

6.The Ka for acetic acid is 1.8 x 10-5. What is the pKa of this weak acid?

Answer 1

1. The equilibrium constant (Ka) = [H⁺][A^-]/[HA]

2. pH = pKa + Log([base]/[acid])

i.e. pH = pKa + Log([A^-]/[HA])

i.e. pH = pKa + Log(n_A^-/n_HA)

3. If the no. of moles of acid = no. of moles of base, i.e. at the equivalence point, pH = pKa

Explanation: At the equivalence point, n_A^- = n_HA, i.e. n_A^-/n_HA = 1

i.e. pH = pKa + Log(n_A^-/n_HA)

i.e. pH = pKa + Log(1)

i.e. pH = pKa + 0

i.e. pH = pKa

4. As the volume of base increases, then the pH of weak acid increases.

Reason: As the base is added, it reacts with the weak acid. So, some weak acid (same as the added base) gets consumed. So, [H⁺] decreases, as a result, pH increases because pH = -Log[H⁺]

6. pKa = -Log(Ka)

= -Log(1.8*10^-5)

i.e. pKa = 4.74