A sample of a certain monoprotic weak acid was dissolved in...

A sample of a certain monoprotic weak acid was dissolved in water and titrated with 0.125 M NaOH , requiring 16.00 mL to reach the equivalence point. During the titration, the pH after adding 2.00 mL of NaOH was 6.912 . Calculate $K_{\mathrm{a}}$ for the weak acid.

Key Concepts

Acid-Base Titration

Acid-base titration is an analytical technique where a solution of known concentration (titrant) is added incrementally to a solution containing the analyte, which in this context is a weak acid. The goal is to determine properties like concentration or dissociation constants by observing the change in pH until the reaction reaches a defined endpoint.

Equivalence Point

The equivalence point in a titration is the stage at which the moles of titrant added exactly neutralize the moles of analyte present. At this point, the reaction has reached stoichiometric balance, and in acid-base titrations, it is typically identified by a sudden change in pH.

Buffer Solution

A buffer solution is formed when a weak acid and its conjugate base exist together in significant amounts in solution. Such mixtures resist drastic changes in pH when small amounts of an acid or base are added, making them especially useful for studying pH-related equilibria in titrations.

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation provides a relation between the pH of a buffer solution, the pKa of the weak acid, and the ratio of the concentrations of the conjugate base to the weak acid. This equation is critical for estimating pH or pKa values in systems where both species are present.

Acid Dissociation Constant (Ka)

The acid dissociation constant, Ka, measures the strength of a weak acid by quantifying the extent to which it dissociates into its ions in water. A small Ka value indicates limited ionization and hence a weaker acid, while a larger Ka signifies a stronger acid tendency to donate protons.

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