Equilibrium: Indicators

• To gain experience in the use of the electronic absorption spectrometer
• To understand indicator color changes and their relationship to the equilibrium constant for reaction of the indicator with water
• To determine the equilibrium constant, K_In, for a particular indicator
• To use the indicator to signal the equivalence point in an acid-base titration

An acid-base indicator is a weak acid that reacts with water, like any weak acid, according to the following general reaction:

The usefulness of an indicator lies in the fact that one (or both) of the forms, HIn or In^-, is intensely colored. The appearance, disappearance, or change of indicator color in a solution during a titration of an acid with a base may be used as a signal that a certain pH has been reached, and that it is time to stop the titration. An indicator is chosen for a titration so that its color change coincides with the equivalence point of the titration--the point at w hich the amount of added base (acid) is stoichiometrically equivalent to the amount of acid (base) originally present in the solution. The equivalence point in the titration of an acid with a base occurs at pH 7 or higher; when a base is titrated with an acid, the equivalence point has pH 7 or lower. Generally, an indicator changes color over a pH range of 2 units, centered about the value of pK_In (-log K_In). For example, the indicator thymolphthalein, HTP, has a pK_In of 10.0. Its acid form, HTP, is colorless, and its base form, TP^-, is blue. A change from colorless to blue occurs over the pH range 9 to 11, and is useful in titrations of acids with base. The indicator alizarin red S, HARS, has a pK _In of 5.4. Its acid form is red and its base form, ARS^- is yellow. It undergoes a change from yellow to red over the pH range 6.4 to 4.4, and is useful in titrations of bases with acid. Clearly, thymolphthalein would not be useful as an indicator for a titration where the pH at the equivalence point is 6, because at this pH, it would not have begun to change color. Similarly, alizarin red S would not be useful as an indicator for a titration where the pH at the equivalence point is 8 , because it would undergo a color change long before this pH was achieved. The point in a titration where the indicator color change occurs is called the endpoint of the titration. It is essential in choosing an indicator for a titration that the endpoint coincide with the equivalence point.

In this experiment, your section will be assigned one of the indicators, methyl red (HMR) or phenolphthalein (HPP). Use electronic absorption spectroscopy to discover the answers to the following questions:

1. Write the equilibrium constant expression for an indicator, HIn, so as to show directly what factors influence the concentration ratio, [In^-]/[HIn].
2. What is the value of this concentration ratio when pH = pK_In-1; when pH =pK_In; when pH = pK_In+1?
3. Write a paragraph explaining why an indicator undergoes a color change over the pH range pK_In ± 1.
4. Suppose we examine the electronic absorption spectrum of an indicator over a pH range including pK_In. How can we use these spectra to determine the value of pK_In?
5. What is pK_In for your assigned indicator?

• electronic absorption spectrometer
• 1-cm sample cells
• 10-mL volumetric flasks
• 1-mL measuring pipets
• 2-mL measuring pipets
• syringe pumps
• Pasteur pipets/bulb
• 50-mL burets
• stock solution of phenolphthalein
• stock solution of methyl red
• glycine buffers covering the pH range »7-11
• citrate buffers covering the pH range »2.75-6.75
• Vitamin C tablets
• standard NaOH, »0.1 M
• household ammonia cleaner
• standard HCl, »0.1 M

Safety glasses must be worn at all times in the laboratory. You will work with solutions of acids and bases. Avoid ingestion and contact of these solutions with the skin. If you spill an acid or base solution on your skin , flush the area immediately with plenty of water.

Record all data in your lab notebook. Your instructor will tell you which indicator to investigate, and will assign your group 4 pH values to investigate. Obtain an equipment box and 4 buffers. Obtain no more than 3 mL in dicator solution in the 5-mL graduated cylinder in your equipment box. Pipet 0.50 mL of indicator stock solution into each of your four 10-mL volumetric flasks. Fill the first volumetric flask to the mark with one of the buffer solutions, and label the fl ask with the pH. Quickly obtain the electronic absorption spectrum of the solution in the wavelength range 700-300 nm. Read and record the absorbance at 520 nm for methyl red or 552 nm for phenolphthalein. Fill the second volumetric flask to the mark with the second buffer solution, label, and obtain the spectrum and required absorbance value. Repeat for the third and fourth buffer solutions. When you are finished collecting data, process it as indicated by the instructor. While waiting for the midlab dis cussion, work on the Questions. Put answers in your notebook.

Clean-up. When you have finished all of your work:

• Clean all glassware with Alconox and water and shake dry.
• Rinse Pasteur pipets by inserting the nozzle of your wash bottle in the top end and squirting water through. Aspirate dry.
• Return all components to the labkit.
• Return all borrowed equipment to the instructor before leaving lab.

Pour used solutions into the sink, and flush down the drain with plenty of water.

Obtain 60 mL standard NaOH (»0.1 M) and two vitamin C tablets. Rinse a 50-mL buret with two 5-mL portions of standard NaOH solution, then fill the buret with NaOH solution. Using a mortar and pe stle, thoroughly crush and grind the vitamin C tablets. Measure out about 0.25 g of powder into each of three 250-mL Erlenmeyer flasks. The masses should be known to 4 decimal places. Add about 75 mL distilled water to the first sample. Stopper the other two flasks. Swirl the first flask for 3-5 minutes to achieve maximum dissolution of the powder (the binders and fillers used in the tablet will not dissolve). Add 3 drops of phenolphthalein to the flask and titrate with standard NaOH from the buret until the solution has a persistent faint pink color (the solid in the flask makes the pink color difficult to see).

Titrate the other two samples in the same manner.

Calculate the amount of ascorbic acid in each sample, the weight percent of ascorbic acid in the sample, and the average weight percent of ascorbic acid in a vitamin C tablet. Estimate your precision. The formula and structure of ascorb ic acid are given below.

Ascorbic acid formula: HC₆H₇O₆ structure:

Obtain 7 mL of household ammonia solution. Obtain 60 mL standard HCl solution (»0.1 M). Rinse a 50-mL buret with two 5-mL portions of standard HCl solution then fill the buret with standard HCl solution. Pipet 2.00 mL of ammonia cleaner into each of three 250-mL Erlenmeyer flasks, and add about 25 mL distilled water. Stopper two flasks. To the other flask, add 5 drops methyl red solution, and titrate with standard HCl (» 0.1M) from the buret until the solution has a persistent pink color.

Titrate the other samples in the same manner.

Calculate the concentration of ammonia in each sample in moles per liter of cleaner, and the average concentration of the three samples. Estimate your precision.

1. This experiment.
2. The appropriate sections of your textbook.

1. An indicator, HIn, has K_In = 4.2 x 10^-8. Over what pH range will it change color?

2. The following absorbance data are obtained as a function of pH for an unknown indicator. From the data, determine the value of K_In for the indicator.