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At the equivalence point during the titration of a weak acid with a strong base, would the pH be acidic, neutral, or basic?
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17. Acid and Base Equilibrium
pH of Weak Acids
Scelta multipla
An unknown weak acid has an initial concentration of 0.55 M. What is the pH of the solution if the weak acid also has a pKa of 5.79?
A
0.60
B
6.05
C
3.02
D
5.75
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Identify the given information: the initial concentration of the weak acid, \([HA]_0 = 0.55\,M\), and the \(pK_a = 5.79\). We need to find the pH of the solution.
Recall the relationship between \(pK_a\) and the acid dissociation constant \(K_a\):
\(pK_a = -\log K_a\)
From this, calculate \(K_a\) using
\(K_a = 10^{-pK_a} = 10^{-5.79}\).
Set up the expression for the acid dissociation equilibrium:
\(HA \rightleftharpoons H^+ + A^-\)
Let \(x\) be the concentration of $H^+$ ions produced at equilibrium. Then,
\([H^+] = x\),
\([A^-] = x\),
and \([HA] = 0.55 - x\).
Write the expression for the acid dissociation constant \(K_a\) in terms of \(x\):
\(K_a = \frac{[H^+][A^-]}{[HA]} = \frac{x^2}{0.55 - x}\)
Since \(K_a\) is small (weak acid), assume \(x \ll 0.55\) to simplify the denominator to approximately 0.55.
Solve for \(x\) (the \([H^+]\) concentration) using the simplified equation:
\(x^2 = K_a \times 0.55\)
Then calculate pH using
\(pH = -\log x\).
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