5: Investigating Buffers
- Page ID
- 516590
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)- 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 \(\ce{HCl}\).
The Henderson-Hasselbalch equation tells us that pH depends on the logarithm of the ratio of base to acid.
- The "Sweet Spot": When \([\text{Base}] = [\text{Acid}]\), the ratio is 1. Since \(\log(1) = 0\), the pH equals the p\(K_a\). This is the most effective region of the buffer.
- The "danger zone": If the ratio drops below 0.1 or goes above 10, the log term shifts by \(\pm 1\) unit, and the buffer capacity collapses.
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 covers safety measures and materials for an experiment on buffer solutions using sodium acetate and acetic acid. It details the titration procedure of two buffers with sodium hydroxide and hydrochloric acid to monitor pH changes, highlighting the importance of accurate measurements and data recording. Additionally, it includes instructions for proper disposal and sensor cleaning to ensure safety and accuracy in the experiment.
- 5.2: Investigating Buffers - Pre-lab
- This page covers the preparation of buffer solutions using sodium acetate and acetic acid, focusing on calculations for achieving a specific pH with the Henderson-Hasselbalch equation. Students work with two buffer concentrations (0.1 M and 1.0 M) targeting a pH of 4, and are tasked with selecting an acid/base pair for a buffer at pH 6.10. It emphasizes the impact of concentration on neutralization capacity while maintaining the same pH.
- 5.3: Investigating Buffers - Data and Report
- This page provides a detailed experimental procedure for titrating Buffer A with NaOH and HCl, focusing on data collection related to pH changes and titrant volumes. It explains buffer capacity as the ability to maintain pH when acids or bases are added and includes calculations for buffer capacity (\(\beta\)). The page also discusses how different buffer solutions behave during titration and poses questions to enhance practical understanding of buffer systems.