LABORATORY EXPERIMENT pubs.acs.org/jchemeduc
The Equilibrium Constant for Bromothymol Blue: A General Chemistry Laboratory Experiment Using Spectroscopy Elsbeth Klotz, Robert Doyle, Erin Gross,* and Bruce Mattson* Department of Chemistry, Creighton University, Omaha, Nebraska 68178, United States
bS Supporting Information ABSTRACT: A simple, inexpensive, and environmentally friendly undergraduate laboratory experiment is described in which students use visible spectroscopy to determine a numerical value for an equilibrium constant, Kc. The experiment correlates well with the lecture topic of equilibrium even though the subject of the study is an acid-base indicator, bromothymol blue. The experiment gives excellent results and offers an inexpensive, zero-waste alternative to traditional equilibrium experiments, such as the Fe(SCN)2þ equilibrium. KEYWORDS: First-Year Undergraduate/General, Laboratory Instruction, Hands-On Learning / Manipulatives, Acids/Bases, Equilibrium, Green Chemistry, Physical Chemistry, Dyes/Pigments, UV-Vis Spectroscopy
e describe an update to a traditional “determination of a pKa” experiment, typically taught in general chemistry during the acid-base unit. The experiment is presented so that it could be integrated into a general equilibrium unit, prior to acid-base chemistry. It is a simple, inexpensive, and environmentally friendly undergraduate laboratory experiment in which students use visible spectroscopy to determine a numerical value for the equilibrium constant, Kc. It correlates well with the lecture topic of equilibrium even though the subject of the study is an acid-base indicator, bromothymol blue. The experiment gives excellent results and offers a less expensivea and green alternative to traditional equilibrium experiments, such as the Fe(SCN)2þ equilibrium1-5 and I2/I3- equilibrium.6-8 The Fe(SCN)2þ experiment uses 0.5 M nitric acid and generates heavy metal waste. The I2/I3- experiment uses organic solvents or halogenated organic solvents. This equilibrium experiment with bromothymol blue produces aqueous waste at a neutral pH or a pH that can easily be neutralized and uses small quantities of acids and bases. One of the oldest and most familiar quantitative equilibrium experiments used in first-year chemistry laboratory programs involves the formation of the Fe(SCN)2þ complex. Students use spectrophotometry to calculate the formation constant of the complex. Typically, students prepare a Beer-Lambert law plot to determine the molar absorptivity of the complex. Then, under conditions of dynamic equilibrium, they use the molar absorptivity value to determine the equilibrium concentration of the complex. Using an ICE table,9 students determine the
W
Copyright r 2011 American Chemical Society and Division of Chemical Education, Inc.
equilibrium concentration of all relevant species and calculate Kc. Values for Kc vary dramatically with the concentration of reagents, probably because other complexes such as Fe(SCN)2þ form at higher thiocyanate concentrations. Other experiments have been designed to determine the pKa of an acid-base indicator in general chemistry using spectrophotometry.10,11 These experiments typically accompany the acid-base unit during a general chemistry course. They generate little waste as most solutions can be poured down the drain immediately after use or after neutralization. They give accurate results when compared to literature values. Because Ka is described by the general equilibrium constant, Kc, this type of experiment can be modified to accompany the chemical equilibrium unit.
’ EXPERIMENT The experiment described here assumes that students have no background in acid-base equilibria.b The primary focus of this experiment is the determination of an equilibrium constant Kc for a chemical equilibrium: HBB þ H2 O H BB- þ H3 Oþ
ð1Þ
Published: March 04, 2011 637
dx.doi.org/10.1021/ed1007102 | J. Chem. Educ. 2011, 88, 637–639
Journal of Chemical Education Kc ¼
LABORATORY EXPERIMENT
½H3 Oþ ½BB- ½HBB
Scheme 1. Equilibrium of Bromothymol Bluea
ð2Þ
where HBB is the acid form of bromothymol blue and BB- is the basic form (Scheme 1). Students are given background information on the Beer-Lambert law and spectrophotometry and the chemical equation for the acid-base equilibrium of the indicator, bromothymol blue. More detail on the student experimental background and procedure can be found in the Supporting Information. The students derive an expression to calculate the equilibrium constant of the reaction using the absorbance data: ½H3 Oþ A616nm A453nm green A453nm Ablue 616nm green
Kc ¼
a The acidic form is HBB (yellow) and the basic form is BB- (blue). A mixture containing significant quantities of each species would appear green.
yellow
ð3Þ
The absorbance measurements at 616 and 453 nm are the λmax values for the base and acid forms of bromothymol blue, respectively. Students make the four absorbance measurements indicated in the equation along with a pH measurement and calculate the value of Kc. If time permits or with appropriate instrumentation, students can collect entire spectra of the yellow and blue solutions and determine the λmax values. Sample spectra are shown in Figure 1. If there is not time for students to collect the spectra, they could be provided for students to view. In a previous experiment using indicator (HIn) solutions,11 HIn þ H2 O H In- þ H3 Oþ
ð4Þ
students measured the absorbance at only the λmax of the basic form, In-. Students use algebraic or graphical methods to calculate pKa of the indicator using the expression: pKa ¼ pH þ log½ðA - AIn- Þ=ðAHIn - AÞ
Figure 1. Visible absorbance spectra of bromothymol blue at low pH (yellow, solid line) and high pH (blue, dashed line).
ð5Þ
Table 1. Student Absorbance Measurements
where A is the absorbance of the solution containing a certain total concentration of the acid-base mixture, AIn- is the absorbance of the base form at the same concentration, and AHIn is the absorbance of the acid form at the same concentration. In the bromothymol blue experiment, students make absorbance measurements at 616 and 453 nm on each solution (yellow, green, and blue). In Figure 1, it can be seen that the blue form absorbs near the λmax of the yellow form. Therefore, for the green solutions, where both yellow and blue forms are present, the absorbance of the “interfering” form is subtracted. ½H3 Oþ ðA616nm - A616nm Þ ðA453nm Þ green blue ðA453nm - Ablue 453nm Þ ðA616nm Þ green
Kc ¼
yellow
Absorbance
ð6Þ
More details on this procedure can be found in the Supporting Information.
’ RESULTS In a laboratory test of this experiment, 22 pairs of students performed the experiment and calculations successfully out of 23
Yellow Solution
Green Solution
Blue Solution
453 616
0.163 0.008
0.099 0.170
0.030 0.411
pairs submitting results. Results were ultimately reported as pKc values for simpler comparison to the literature value of 7.1, reported at an ionic strength of 0.1.12-15 Using eq 6, the 22 pairs of students obtained an average value for pKc of 7.04 ( 0.03. A typical set of data obtained from one pair of students, who also reported their green solution had a pH = 7.00, is given in Table 1.
yellow
’ HAZARDS Appropriate eye protection and clothing must be worn at all times. Use appropriate caution when handing 1 M HCl(aq) and NaOH(aq) solutions. Both solutions are irritants at the concentrations used. All solutions should be neutralized before disposal. Bromothymol blue may cause irritation to skin, eyes, and respiratory tract and may be harmful if swallowed or inhaled.
Wavelength/nm
’ LEARNING OUTCOMES The student learning outcomes of this experiment are • To reinforce the concept that an equilibrium constant can be calculated from concentration measurements (as learned in lecture). • To use the Beer-Lambert law expressions to relate concentration to absorbance. • To reinforce laboratory techniques of using volumetric glassware and spectrophotometry. • To recognize good laboratory techniques with a comparison of the student-determined Kc value to the literature value. 638
dx.doi.org/10.1021/ed1007102 |J. Chem. Educ. 2011, 88, 637–639
Journal of Chemical Education
LABORATORY EXPERIMENT
’ ASSOCIATED CONTENT
Raton, FL, 2007-2008, which referencesIndicators; Bishop, E., Ed.; Pergamon: Oxford, 1972.
bS
Supporting Information Written directions used by students (including appropriate safety instructions and CAS registry numbers for all chemicals); instructor notes, including a color version of Figure 1 from the manuscript, a photograph of bromothymol blue solutions (1-7) in conditions of high pH to low pH, and absorbance spectra of solutions 1-7; complete information regarding potential hazards to students and instructors. This material is available via the Internet at http://pubs.acs.org.
’ AUTHOR INFORMATION Corresponding Author
*E-mails: (E.G.)
[email protected], (B.M.) brucemattson@ creighton.edu.
’ ACKNOWLEDGMENT The authors wish to acknowledge the Creighton University Department of Chemistry. E.K. acknowledges support from the Clare Boothe Luce Scholarship Program. ’ ADDITIONAL NOTE a Chemical costs ran $0.14/pair of students using 2009 prices from Fisher. Using the same price list, the iron(III) thiocyanate experiment costs approximately the same whereas the I2/I3experiment runs over $2.00/pair. b
The experiment can serve as an introduction to the concepts of pH, acids and bases, indicators, and buffers and when those topics are introduced later in the semester, experiences recalled from this experiment will help with the segue. Students are called upon to convert pH into the hydronium ion concentration.
’ REFERENCES (1) Ramette, R. W. J. Chem. Educ. 1963, 40, 71–72. (2) Cobb, C. L.; Love, G. A. J. Chem. Educ. 1998, 75, 90–92. (3) Postma, J. M.; Roberts, J. L.; Hollenberg, J. L. Chemistry in the Laboratory, 6th ed.; W.H. Freeman: New York, 2004; pp 24-1-24-11. (4) Bauer, R. D.; Birk, J. P.; Sawyer, D. J. Laboratory Inquiry in Chemistry, 2nd ed.; Brooks/Cole: Belmont, CA, 2005; pp 155-159. (5) Beran, J. A. Laboratory Manual for Principles of General Chemistry, 7th ed.; John Wiley and Sons: New York, 2004; pp 293-304. (6) Shoemaker, D. P.; Garland, C. W. Experiments in Physical Chemistry; McGraw-Hill: New York, 1962; pp 180-184. (7) Ackermann, M. N. J. Chem. Educ. 1978, 55, 795. (8) Petrovic, S. C.; Bodner, G. M. J. Chem. Educ. 1991, 68, 509. (9) ICE Table. http://en.wikipedia.org/wiki/ICE_table (accessed Jan 2011). (10) Wink, D. J.; Fetzer Gislason, S.; Kuehn, J. E. Working with Chemistry: A Laboratory Inquiry Program, 2nd ed.; W.H. Freeman: New York, 2005; pp L-1-L-10. (11) Patterson, G. S. J. Chem. Educ. 1999, 76, 395–398. (12) Skoog, D. A.; West, D. M.; Holler, F. J.; Crouch, S. R. Fundamentals of Analytical Chemistry, 8th ed.; Brooks/Cole: Belmont, CA, 2004; pp 372. (13) Patnaik, P. Dean’s Analytical Chemistry Handbook, 2nd ed.; McGraw-Hill: New York, 2004; p 4.30. (14) Dean, J. A., Lange’s Handbook of Chemistry, 14th ed.; McGrawHill: New York, 1992; pp 8-116. (15) A different value for pKa (7.3 at 20 °C) is given in the CRC Handbook of Chemistry and Physics; Lide, D. R., Ed.; CRC Press:: Boca 639
dx.doi.org/10.1021/ed1007102 |J. Chem. Educ. 2011, 88, 637–639
