3: Acid Dissociation Constant
- Page ID
- 516588
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)- To titrate an unknown weak acid using a pH probe.
- To use the titration curve to determine the pKa (and therefore the Ka) of the acid and determine a likely identity of the acid.
INTRODUCTION
This lab experiment focuses on titrating an unknown weak acid using a pH probe. The primary goal is to determine the acid dissociation constant (Ka) of the unknown weak acid.
The procedure involves titration of the acid using a 0.1 M \(\ce{NaOH}\) solution as the titrant, a pH sensor, and a magnetic stirrer to ensure mixing.
Data collection begins by slowly adding the \(\ce{NaOH}\) solution to a known volume of the unknown weak acid, while the pH is continuously monitored and recorded. The titration continues until the pH is over 11.
The Ka value is calculated using the Henderson-Hasselbalch equation. Specifically, at the half-equivalence point of the titration, the pH is equal to the pKa of the weak acid. From the pKa, the Ka can then be determined. The experiment also involves analyzing the titration graphs to estimate the concentration of the unknown acid from the equivalence point.
Calculating the Ka
For a monoprotic weak acid, \(\ce{HA}\), the Henderson-Hasselbalch equation is
\[ \text{pH}=\text{p}K_{\text{a}}+\log \left( \frac{[\ce{A^-}]}{[\ce{HA}]} \right)\]
where \( [\ce{HA}] \) and \( [\ce{A^-}] \) are the equilibrium molar concentrations of the weak acid and its conjugate base, respectively.
At the half-equivalence point, \( V_{1/2} = \frac{1}{2} V_{\text{eq}} \),
\[ [\ce{HA}] = [\ce{A^-}],\ \ \text{so}\ \ \frac{[\ce{A^-}]}{[\ce{HA}]} = 1,\ \ \text{and}\ \ \log\left( \frac{[\ce{A^-}]}{[\ce{HA}]} \right) = 0 \]
Therefore, at the half-equivalence point,
\[ \text{pH}=\text{p}K_{\text{a}} \]
The \( K_{\text{a}} \) can be calculated using the definition of \( \text{p}K_{\text{a}} \),
\[ \text{p}K_{\text{a}} = -\log{K_{\text{a}}},\ \ \text{so}\ \ K_{\text{a}} = 10^{-\text{p}K_{\text{a}}} \]
You will determine the equivalence point (\(V_{eq}\)) visually from the steep part of your graph.
- Steepness Matters: The sharper the jump in pH, the more precisely you can read \(V_{eq}\).
- Data Density: If you add large volumes of base (e.g., 1.0 mL) near the equivalence point, you will miss the turn. You need more data points (small drops) in the steep region to get a precise reading.
Henderson-Hasselbalch Equation:
\[ \text{pH}=\text{p}K_{\text{a}}+\log \left( \frac{[\ce{A^-}]}{[\ce{HA}]} \right)\]
Definition of pK_a:
\[ \text{p}K_\text{a} = -\log{K_\text{a}}\ \ \ \ \ \text{or}\ \ \ \ \ K_\text{a} = 10^{-\text{p}K_\text{a}} \]
- 3.1: Acid Dissociation Constant - Experiment
- This page details the safety precautions and materials required for titrating a monoprotic weak acid with NaOH. It includes a three-part experimental procedure: setting up the titration, gradually adding NaOH while recording pH changes, and repeating with a second sample. The page stresses the importance of safe chemical waste disposal and proper handling of caustic substances.
- 3.2: Acid Dissociation Constant - Pre-lab
- This page discusses the titration of a monoprotic acid, detailing the acid dissociation reaction and its constant (Ka). It highlights the importance of the half-equivalence point where acid and conjugate base concentrations are equal. A practical titration example with a weak acid and \(\ce{NaOH}\) is provided, stressing the importance of careful volume additions near the equivalence point for accurate endpoint determination, supported by hypothetical titration curves.
- 3.3: Acid Dissociation Constant - Data and Report
- This page details a titration lab activity focused on recording and analyzing data. Students learn to create titration curves, identify key points, estimate pH and acid dissociation constants (Ka), and assess the precision of their techniques through percent difference. They also identify an unknown acid by comparing their experimental Ka to literature values and calculating percent error, enhancing their data analysis and critical thinking skills in chemistry.