19 Buffers, Indicators, and Solution Standardization

Purpose

To explore the properties of buffers and indicators, and to review experimental techniques associated with acid-base titration

Expected Learning Outcomes

After completing this experiment, students are expected to be able to

  • Calibrate and use a pH meter.
  • Prepare and recall the properties of a buffer solution.
  • Select appropriate indicators for use in a titration.
  • Prepare and standardize a solution of sodium hydroxide.

Textbook Reference

From Tro, Chemistry: Structures and Properties, 2nd Ed:

Introduction

Buffers

Buffers are formed by mixing together similar amounts of weak acids and their conjugate base, such that

[latex]0.1 < \frac{[\mbox{base}]}{[\mbox{acid}]} < 10[/latex]

In practice, most of the time buffers are prepared by adding to a solution of the weak base HCl (or another strong acid) in a dropwise fashion until the desired pH is obtained.
Buffer solutions are able to resist changes to pH when small amounts of acid or base are added to the solution, as illustrated below.  As the concentration of buffer solution increases, the ability of the buffer to resist change in pH increases, as does the buffering capacity (the moles of acid or base that can be added before a significant jump in pH is observed).
pH change when NaOH is added to distilled water and different volumes of water. Distilled water changes pH v quickly, lower concentration buffers change pH slowly until a certain volume then jumps up; higher concentration buffers barely change pH
pH curves of different solutions as NaOH is added. Source: Averill and Eldredge, General Chemistry: Principles, Patterns, and Applications.

The pH of a buffer solution can be calculated using the Henderson-Hasselbalch equation.[1]

\begin{equation}
\mbox{pH} = \mbox{p}K_a + \log_{10} \left(\frac{[\mbox{base}]}{[\mbox{acid}]}\right) \label{414:HH}
\end{equation}

Indicators

Indicators are weak acids that absorb different colors in their acid and conjugate base forms. In environments with different pHs, they have different colors.[2]

Examples

Phenol red changes color from yellow at pH 6 to red at pH 8:

 

The color of phenol red at different pHs (6.0 (yellow) at 0.1 pH increments to 8.0 (red)
Color of phenol red at different pH. Credit: Max Schwalbe/Wikimedia.

As part of this experiment, you will examine the colors of different indicators at different pHs. You may have used indicators before for standardization of acid and base solutions.

Standardization of Acids and Bases

Often, it is vital that the concentration of solutions are critical for human life. These are often determined using various titrimetric and other analytical techniques. In these cases, it is vital that reference solutions be prepared accurately.

Examples

The concentration of chloride ions in water is typically determined by titration with silver, using the precipitation of silver chloride as a standard.

\begin{equation}
\ce{Ag+}(aq) + \ce{Cl-} (aq) \rightleftharpoons \ce{AgCl}(s)
\end{equation}

When this is performed, it is vital that the concentration of silver ions in the analyte (we typically use the water sample as the titrant) is known accurately.

For acid-base experiments, the challenge is even greater. The solid form of many acids and bases are hygroscopic; they will absorb water from the ambient environment. This is true in particular for sodium hydroxide, which will form swollen pellets that may clump together as it absorbs moisture from air:

Pellets of sodium hydroxide clumped together
Solid sodium hydroxide. Credit: Martin Walker/Wikimedia.

For this reason, specific primary standards of acids and bases are used to determine the concentration of other acids and bases using titrimetric methods rather than simply by dissolving the appropriate mass of solution.

In this experiment, you will prepare a solution of approximately 0.1 M sodium hydroxide and standardize this solution using potassium hydrogen phthalate (KC8H5O4, often referred to as “KHP”; molar mass = 204.2 g/mol) as the primary standard, as you did previously in the CHEM-C 125 experiment Analysis of Acids and Bases by Titration.  Please review the discussion about acid-base titration in the lab manual for that class.

Procedures

  • You will perform this experiment in pairs.  However, for the solution preparation and standardization part, if your lab partner is absent you should conduct this part of the experiment individually.
  • Review the directions on using a pH meter in Using Laboratory Equipment before coming to lab.
  • It is critical that
    • All solutions are properly labeled.
    • Volumes are measured accurately using volumetric glassware when preparing solutions. The exact volume for the pH measurement, however, does not matter. You will need to ensure that you have at least about 10 mL of solution to take good measurements such that the minimum liquid level on the pH probe is exceeded.
  • You must label and turn in your solution of sodium hydroxide. In the next lab period, you will use this solution as a secondary standard} to perform acid-base titration experiments.

The Action of Buffers

  1. Turn on your device and open the Vernier Graphical Analysis app. Connect the pH probe following the directions in this manual. Calibrate the pH probe using pH 4, 7 and 10 buffers following the directions here.[3]
  2. Prepare a buffer solution by mixing 15.0 mL of 1.0 M acetic acid (CH3COOH) with 15.0 mL of 1.0 M sodium acetate (CH3COONa) in a graduated cylinder. Mix these solutions thoroughly and measure its pH in a test tube.
  3. Measure out 5 mL of this solution into each of three different test tubes.
  4. Mix into the three different test tubes the following:
    1. 5 mL deionized water
    2. 5 mL of 0.05 M sodium hydroxide
    3. 5 mL of 0.05 M hydrochloric acid.

    Measure the pH of each of these solutions

  5. Repeat steps 3 and 4 using deionized water instead of the buffer solution you made in step 2.

Indicators

This part of the experiment will be done individually; each student will be assigned a different pH buffer.

As you may know different indicator have different colors depending on their pH.  These are often summarized on charts like this:

Different indicators and their colors at different pH values
Indicators and their color ranges. Source: OpenStax Chemistry 2e.

You will be assigned one of the buffers provided (pH 4, 7, 10) to study.  The goal of this experiment is to use two different indicators from this list to identify the pH as best you can.  You will be given the following indicator solutions:

  • alizarin yellow
  • bromothymol blue
  • methyl orange
  • methyl red
  • phenolphthalein
  • phenyl red

Add a small amount of your chosen indicator solutions into separate samples of your assigned buffer solution.  Record the colors of each of these solutions.

This solution need not be prepared accurately concentration-wise. The reason for this is, as explained above, it is not possible to prepare such a sodium hydroxide solution accurately. Doing that would be a waste of time.

Preparation of a Sodium Hydroxide Solution

In this part of the experiment, you will prepare using solid sodium hydroxide approximately 500 mL of 0.1 M standardized sodium hydroxide. To determine the pH, you will prepare a standard potassium hydrogen phthalate (KHP) solution as a primary reference.

  1. Calculate the mass of solid sodium hydroxide required to prepare 500 mL of 0.1 M sodium hydroxide solution.
  2. Prepare 500 mL solution in a beaker, dissolving approximately the amount of sodium hydroxide required into it. Be sure that the solution is properly stirred. Pour this solution into the plastic bottle provided. Label the plastic bottle with the label provided with the following information
    1. Your names/initials.
    2. The identity of the solution (including the approximate concentration).
    3. Today’s date.
    4. The hazard class (look up from an SDS sheet)

Standardization of Sodium Hydroxide

Following the directions in the Analysis of Acids and Bases by Titration experiment, prepare a sodium hydrogen phthalate (KHP) solution and use this to determine (accurately) the concentration of sodium hydroxide solution you prepared.

The bottle of sodium hydroxide you prepared should be placed in your drawer for use next week.

Waste Disposal

All waste from this experiment can be disposed of in the drain with plenty of water.

  1. This is equivalent in assumption to neglecting x in ICE tables. ↵
  2. Universal indicator is a mixture of different indicators designed to have different colors at different pHs. ↵
  3. These may be substituted by any two or three pH values.  Be sure to pay attention to the pH values on bottle. ↵

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IU East Experimental Chemistry Laboratory Manual Copyright © 2022 by Yu Kay Law is licensed under a Creative Commons Attribution-NonCommercial 4.0 International License, except where otherwise noted.

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