So, What Does it Indicate?
Introduction
Phenolphthalein is a common indicator you may have used in a previous science course, such as
Chemistry 184. In solutions with a pH of less then 8.3, this compound remains colorless. In solutions
with a pH greater than 8.3, the compound is reddish pink. Why? A change in molecular conformation
or molecular structure underlies color changes for phenolphthalein--and for all other indicators. pH
indicators are themselves weak acids which have complex molecular structures such as the ones shown
below for phenolphthalein. Like any weak acid, the molecule may be protonated (acid form) or
deprotonated (conjugate base form).
In addition to the oxygen and hydrogen atoms explicitly shown, the substances depicted above contain
many carbon atoms, as well as many additional hydrogen atoms. To understand exactly how many
carbon and hydrogen atoms are present, click here to learn about line formulas--the widely used
organic chemistry shorthand. The structures shown above are line formulas, and additional examples
can be seen everywhere from TV commercials to your physician's medical reference books!
You may notice by comparing the two chemical structures shown above that the phenolphthalein
molecule contains four rings in acidic solutions; in a basic solutions it contains only three rings. This
difference, along with accompanying differences in electronic structure, affect how phenolphthalein
absorbs light--so much so that it is pink in basic solutions, but colorless in acidic solutions. There are
many other compounds that have different colors at different pHs (see Table 16.1 in your Chang
Chemistry textbook). Collectively, these are known as acid-base indicators. Despite the widespread
availability of pH electrodes, the expense and fragility of electrodes prohibit their use in certain
modern applications. pH indicators continue to provide a practical means for acid/base testing,
including endpoint determination in titrations, because when the solution pH exceeds a certain
threshold, the entire solution changes color. An example of this is the determination of acetic acid in
vinegar, an experiment routinely performed in Chemistry 184, employing phenolphthalein as the acidbase indicator.
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In lieu of using the entire molecular structure to represent a given acid-base indicator, the formula of a
protonated indicator is often represented as HIn (where In is a generic abbreviation for indicator). This
HIn species can lose a proton to form its conjugate base, In- , as shown in the equation
HIn ֖ H+ + InThe equilibrium constant for this dissociation can be written as
Ka =
[H + ] [In - ]
[HIn]
This expression can be manipulated to produce the equation below, known as the HendersonHasselbalch equation. Consult your Chang Chemistry text section 16.2 to see how!
pH = pKa + log [In-]/[HIn]
or
pH = pKa + log [conjugate base]/[acid]
Thus, in addition to helping us explain the behavior of pH buffers, we can use this equation to predict
the ratio of the HIn and In- forms of an indicator at a given pH.
The specific pH indicator you will use in today's lab is called phenol red. (The chemical structure of
phenol red, phenolsulfonephthalein, can be found by searching www.chemfinder.com.) The HIn
(acidic) form of phenol red is yellow, whereas the In- (basic) form is red. When the indicator is in a
solution of high pH, the major form of the indicator is In- ; thus the human eye sees this solution as the
color of the In- (red). As the pH decreases with the addition of acid, the H+ of the acid reacts with the
In- to produce HIn according to LeChâtelier's principle. As the HIn becomes the dominant species, the
eye sees the yellow color increasing and the red color of the In- decreasing. The color the human eye
detects really depends upon the relative amounts of the two forms, or the [In-]/[HIn] ratio, so the
observed color of the solution may be intermediate to the two colors. Our spectrophotometers are
more sensitive to changes in color than is the human eye; in fact, in this experiment the
spectrophotometer can simultaneously measure both the acid and base forms of the phenol red
indicator since these two species absorb light at different wavelengths.
The goal of this week's lab is to obtain experimental data for a tabular display that shows how these
concentrations change with pH. Determining the concentrations of chemical species in a systematic
way is often called speciation. In this week’s experiment you will examine the speciation of phenol
red, an acid-base indicator that turns from yellow to red as the solution goes from acidic to basic.
Specifically, 0.15 M HCl will be added in small portions to a solution that contains 0.1 M Na2HPO4
and several drops of phenol red. The pH will change only slightly upon each addition of each
successive 1-mL-portion of 0.15 M HCl because the phosphate buffer will moderate the pH change
upon addition of acid; appreciable concentrations of both forms of the indicator will exist during the
course of the experiment. This will allow close examination of the concentrations of the acidic and
basic forms of phenol red and phosphate species.
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Pre-lab
Safety
•
Always, always wear safety glasses in the laboratory. Your TA can remove you from the lab if
you fail to wear safety glasses.
•
Caution! Acids and bases are hazardous if splashed on clothing, exposed skin or in the eyes.
Prolonged exposure of the skin to even fairly dilute solutions of acid and base can cause serious
burns. If acids and bases splash on skin or clothes, remove the affected clothing and flush the
affected areas thoroughly with cold water.
•
Notice on disposal of chemical wastes: All of the solutions used in this laboratory are dilute and can be discarded down the drain after being neutralized. Consult your TA on the proper method of disposal.
(1) During the course of an experiment, 10.0 mL of water is added to 25.0 mL of aqueous 0.120 M
NaCl. What is the concentration of NaCl in the resulting solution?
(2) Name the following species:
(a) H3PO4
(b) H2PO4-
(c) HPO42-
(d) PO43-
(3) The pKa of H2PO4- at 25ºC is approximately 7.2. At pH 6.2, what is the ratio [HPO42-]/[H2PO4-]?
(4) At pH 7.2, what is the ratio [HPO42-]/[H2PO4-]?
(5) Calculate the pH of a solution resulting from the addition of 10.0 mL of 0.10 M NaOH to 50.0 mL
of 0.10 M benzoic acid (Ka = 6.5 x 10-5).
(6) List the chemical species that will exist in the acidified solution your group will generate in Part 2
of this week’s experiment.
(7) Predict the approximate absorbance wavelength range for the acid and the base forms of the phenol
red indicator you will use in this experiment. For help, study the Introduction and Appendix of this lab
write-up.
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Procedure
Part 1 – Acidic and Basic Forms of Phenol Red
Before undertaking the experiment portion of Part 1, work with your group to reach a consensus on
what chemical species would be produced by acidifying a Na2HPO4 solution using HCl. What
chemical species would be present if the Na2HPO4 solution were treated with NaOH instead of HCl?
1. Prepare a 50-fold diluted phenol red solution by combining approximately 1 mL of stock phenol red
solution in a beaker with about 50 mL of distilled water.
2. Flip a coin. If your coin comes up heads, acidify the 50-fold diluted phenol red solution by adding
approximately 3 mL of 0.10 M HCl. Mix. If your coin turns up tails, treat the 50-fold diluted phenol
red solution with about 3 mL of 0.10 M NaOH. Mix.
3. Gently unplug the USB cable that feeds into the LabPro box. This will ensure that the fullspectrum Ocean Optics signal will be displayed in this Part.
4. Open the LoggerPro software. (If the software is already open, close it and then re-open it, in order
to clear any files left over from previous lab groups.)
5. Fill a cuvette three-quarters full with distilled water and place it in the Ocean Optics
spectrophotometer cuvette holder. Designate this as the absorbance “blank.” Calibrate the
spectrophotometer. Hint: Save this distilled water cuvette so that you can quickly re-calibrate later, if
needed.
6. Fill a cuvette three-quarters full with the diluted solution you prepared in Step 2. Click the Collect
button to acquire its visible spectrum. Do not close the LoggerPro software. You will need the
spectrophotometer for Part 2, below.
7. Make a small sketch of the spectrum in your notebook. Compare your spectrum with spectra taken
by other groups in the laboratory. Why are the spectra similar or different? Discuss this with your lab
mates and lab TA. What do you think the spectra would look like at pHs nearer to neutral?
Part 2 – pH/Concentration Dependence Studies
The goal of Part 2 is to determine the concentrations of protonated and deprotonated species in solution
at a range of pHs. For phenol red, this task will be undertaken by following the absorbance spectrum
with the spectrophotometer as a function of pH. The molar concentrations of the phenol red indicator
species can be determined using the absorbances you measure together with the molar absorptivities
that your TA will give you. The speciation (i.e., the concentrations of the components) of the
phosphate buffer can be determined from the relevant pKa (7.2) and the effect of added acid on the
solution pH. Specifically, the Henderson-Hasselbalch equation allows you to determine the
concentrations of H2PO4- and HPO42- at each pH you measure.
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1. Connect the USB cable from the computer to the LabPro box. The LabPro box should bear a pH
electrode. (The Ocean Optics spectrophotometer must remain connected to the computer at all times
via its separate USB cable.)
2. Open the LoggerPro Experiment titled “PhenolRed,” in the “Chemistry with Vernier” folder.
3. Since the Ocean Optics spectrophotometer had several minutes to “warm up” during Part 1, its light
source should be very stable. Re-calibrate it now using the distilled water blank cuvette.
4. Calibrate the pH electrode.
5. To a 25.0 mL sample of 0.1 M Na2HPO4 solution, add about 10 drops of undiluted phenol red
indicator. Swirl gently to homogenize this buffer/indicator mixture.
6. Carefully transfer a portion of your indicator mixture to a cuvette. (Cuvettes should be about threequarters full when used with the Ocean Optics spectrophotometers.) Note: If you spill any of your
mixture or accidentally dilute it, you will be unable to make reliable concentration calculations!
7. You are about to acquire a series of simultaneous measurements of the indicator mixture’s pH and
its absorbance at 440 nm and 560 nm. Readings for pH and absorbance should now be displayed along
with three corresponding graph windows; prepare the system for data collection by clicking the Collect
button.
8. Place the calibrated pH electrode into your mixture and swirl gently until the pH reading stabilizes.
9. Click the "Keep" button. In the resulting dialogue box, enter 0 (for zero mL) to indicate that no
0.15 M HCl has been added to the mixture yet. This action should record discrete data points for pH,
absorbance at 440 nm, and absorbance at 560 nm.
10. Return the contents of the cuvette to the parent mixture. Do not rinse the cuvette.
11. Obtain about 25 mL of 0.15 M HCl solution in a separate beaker. Use a volumetric pipet to
transfer 1.00 mL of this 0.15 M HCl to your mixture. (Be sure to record the “exact” HCl concentration
shown on the bottle.) Swirl the mixture gently.
12. Place a sample of your mixture into the same unrinsed cuvette you used earlier. Why should you
not rinse this cuvette?
13. Place the pH electrode into your indicator mixture, swirl, and click the Keep button when the pH
reading has stabilized.
14. Continue delivering 1.00-mL portions of 0.15 M HCl to the indicator mixture, recording pH and
absorbance data points after each addition using the “Keep” button.
15. When you are sure that the phenol red has been predominately converted to its acid form, click the
“Stop” button.
16. Perform electronic “cut and paste” operations if you wish to save copies of the spreadsheet window
and/or the graph itself for later use. Open the Word or Excel program and save the file(s) using those
file formats so that you can re-open them on other computers. The Logger Pro file format (.cmbl)
cannot be opened on computers outside of the Chemistry lab!
Now work with your lab mates and lab TA to construct an Excel spreadsheet to determine the
concentration of as many species in the solution as possible for each acid addition.
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Appendix
White light is composed of all colors of visible light. If a molecule in a solution absorbs certain
wavelengths of visible light, the solution appears colored. A color wheel, shown below, can be used to
predict the appearance of the solution. The relationship between the color of absorbance and the
corresponding wavelength is illustrated at http://www.uam.es/docencia/quimcursos/Scimedia/chemed/spec/spectros.htm
The human eye detects the light that is transmitted (not absorbed) by the solution. Therefore when a
species absorbs light of a certain color, the color of the solution appears as the complementary color of
the light that was absorbed. The color wheel shows how colors are related, i.e., with pairs of
complementary colors depicted on the opposite sides of the wheel. For example, if a species absorbs
orange light between 590 and 630 nm, the solution containing that species would be seen as blue.
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