5: Investigating Buffers
- Page ID
- 516590
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)PURPOSE
- To use the Henderson-Hasselbalch equation to determine the amounts of acetic acid and sodium acetate needed to prepare two acidic buffer solutions (Buffer A and Buffer B), and then prepare these solutions.
- To determine the buffer capacities of the prepared Buffer A and Buffer B by adding solutions of \(\ce{NaOH}\) (strong base) and \(\ce{HCl}\) (strong acid).
INTRODUCTION
A buffer is a mixture of a weak acid and its conjugate base, or a weak base and its conjugate acid. A buffer’s function is to absorb acids (\(\ce{H+}\) or \(\ce{H3O+}\) ions) or bases (\(\ce{OH-}\) ions) so that the system's pH changes very little.
Buffers are critical in many systems. Blood plasma, a natural example in humans, serves as a bicarbonate buffer, maintaining the pH of blood between 7.2 and 7.6.
By design, a buffer is a system at equilibrium. For example, a buffer can be prepared with nitrous acid, \(\ce{HNO2}\). The weak acid establishes an aqueous equilibrium as shown below.
\[\ce{HNO2} (aq) + \ce{H2O} (l) \rightleftharpoons \ce{H3O+} (aq) + \ce{NO2-} (aq)\]
The equilibrium constant expression is shown below.
\[ K_a = \frac{\left[ \ce{H3O+} \right] \left[ \ce{NO2-} \right]}{\left[ \ce{HNO2} \right]} \]
To prepare a buffer system with nitrous acid, a salt containing the conjugate base is added, such as sodium nitrite (\(\ce{NaNO2}\)). The resulting system is a mixture of \(\ce{HNO2}\) and \(\ce{NO2-}\) ions. The nitrous acid will neutralize additional hydroxide ions, and the nitrite ion will neutralize additional hydronium ions.
A variation of the equilibrium expression above, called the Henderson-Hasselbalch equation, is the best reference in preparing a buffer solution. For our nitrous acid-nitrite buffer example, the Henderson-Hasselbalch equation is shown below.
\[ \text{pH} = \text{p}K_a + \log \left( \frac{\left[ \ce{NO2-} \right]}{\left[ \ce{HNO2} \right)} \right) \]
The pH range in which a buffer solution is effective is generally considered to be ±1 of the pKa.
In this experiment, you will use the Henderson-Hasselbalch equation to determine the amount of acetic acid and sodium acetate needed to prepare two acidic buffer solutions. You will then prepare the buffers and test their buffer capacities by adding solutions of \(\ce{NaOH}\) and \(\ce{HCl}\).
Henderson-Hasselbalch Equation:
\[ \text{pH} = \text{p}K_a + \log \left( \frac{\left[ \ce{A-} \right]}{\left[ \ce{HA} \right)} \right) \]
- 5.1: Investigating Buffers - Experiment
- This page details safety measures for handling caustic substances such as sodium hydroxide and hydrochloric acid. It includes a materials list and a step-by-step procedure for preparing and testing two buffer solutions. The procedure involves titrating the buffers to modify pH levels while maintaining accurate measurements and thorough documentation of results. Adequate caution in handling materials is emphasized throughout the process.
- 5.2: Investigating Buffers - Pre-lab
- This page provides guidance on using the Henderson-Hasselbalch equation to determine the mass of sodium acetate required for buffer solutions at a specific pH using acetic acid. It includes calculations for buffers made from 50.0 mL of both 0.1 M and 1.0 M acetic acid aiming for a pH of 4. Furthermore, it encourages readers to find an appropriate acid/conjugate base pair for creating a buffer at pH 6.10, fostering exploration of weak acid options.
- 5.3: Investigating Buffers - Data and Report
- This page outlines an experimental setup for titrating two buffer solutions with NaOH and HCl. It includes instructions for recording pH changes, using data analysis software to plot results, and analyzing buffer behavior through tables. Key tasks include calculating the titrant volume needed for a 2-unit pH change and identifying titration endpoints. Post-lab questions focus on buffer reactions, capacity, and predictions under varying conditions.