A 1.49 L buffer solution consists of 0.186 M butanoic acid and 0.258 M sodium butanoate. Calculate the pH of the solution following the addition of 0.062 moles of NaOH. Assume that any contribution of the NaOH to the volume of the solution is negligible. The Ka of butanoic acid is 1.52 × 10-5.
mol of NaOH added = 0.062 mol
C3H8COOH will react with OH- to form C3H8COO-
Before Reaction:
mol of C3H8COO- = 0.258 M *1.49 L
mol of C3H8COO- = 0.3844 mol
mol of C3H8COOH = 0.186 M *1.49 L
mol of C3H8COOH = 0.2771 mol
after reaction,
mol of C3H8COO- = mol present initially + mol added
mol of C3H8COO- = (0.3844 + 0.062) mol
mol of C3H8COO- = 0.4464 mol
mol of C3H8COOH = mol present initially - mol added
mol of C3H8COOH = (0.2771 - 0.062) mol
mol of C3H8COOH = 0.2151 mol
Ka = 1.52*10^-5
pKa = - log (Ka)
= - log(1.52*10^-5)
= 4.8182
since volume is both in numerator and denominator, we can use mol instead of concentration
we have below equation to be used:
This is Henderson–Hasselbalch equation
pH = pKa + log {[conjugate base]/[acid]}
= 4.8182+ log {0.4464/0.2151}
= 5.14
Answer: 5.14
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