In the Laboratory
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Acid–Base Indicators: A New Look at an Old Topic Ara S. Kooser, Judith L. Jenkins, and Lawrence E. Welch* Department of Chemistry, Knox College, Galesburg, IL 61401; *
[email protected] Acid–base chemistry is a vital element of the chemistry curriculum. Originally a subject addressed in upper-class offerings, acid–base chemistry has been a common feature of the general chemistry curriculum for a number of years. Despite the continual evolution and refinement of the material in general chemistry, acid–base reactions appear to be a cornerstone of the course for the foreseeable future. This also tends to be a topic that lends itself well to exploration in the laboratory, particularly as an element of a quantitative analysis procedure. The general chemistry laboratory manual used at this institution in the early 1950s (1) did not have a real acid– base experiment. By the early 1960s, our general chemistry laboratory manual (2) did feature a quantitative acid–base experiment, in which a NaOH solution was standardized and then used to titrate an unknown solution of strong acid with a phenolphthalein indicator. Forty years later, all ten laboratory manuals found in the Knox Chemistry Department’s office (3–12) featured the aforementioned NaOH standardization; nine of them described the classic titration with phenolphthalein as the indicator, while one manual (3) added a twist by utilizing a pH meter as an indicator. Acids used were either HCl, KHP, or oxalic acid. In all but one of the manuals (11), the standardized NaOH solution was used in a followup experiment to titrate an unknown acid (often vinegar), all but one (3) using phenolphthalein as indicator. Most of the manuals followed the introductory acid–base experiment with more sophisticated activities within the same subject area, such as a Ka determination (5, 6, 11), activities with buffer systems (3, 4, 6, 10, 11), or some utilization of a pH meter and/or use of indicator groups to measure pH (3, 5–7, 9, 10, 12). While some variety in acid–base experiments clearly does exist, almost every manual introduces acid–base work with essentially the same laboratory exercise. Even worse, that same experiment is no different from what was done in general chemistry 40 years ago, and today commonly duplicates an experience that many of our students have had in the high school chemistry laboratory. Our goal was to develop a fresh experiment to introduce acid–base chemistry in the first-year curriculum, one that embodies some of the important lecture concepts (volumetric analysis in particular), but deviates from the tired repertoire of NaOH standardization experiments. The chemical education literature contains a number of citations of acid– base titration procedures that incorporate alternative nonvisual endpoint determination, often in conjunction with computer interfacing. The most common indicator systems found include potentiometric determination via a pH meter and glass electrode (13, 14 ) and the use of colored dye indicators coupled with colorimetric sensing (15, 16 ). The focus of our work has been the development of an experiment that forces the students to choose the best indicator from a set, using a conductivity probe as a tool to allow them to make an informed choice. Although conductivity has a history of use as an indicator (17) and as a method to probe the physical 1504
properties of a chemical system (18–20), it has seen only minimal use as an indicator in the chemical education literature (21) and none at the introductory level. Materials and Methods All indicator solutions were made from the recipes given in the Acid Base Indicators section of the CRC Handbook (22). Further information regarding the indicators is given in Table 1. Note that alizarin is not the same as the more commonly used alizarin yellow. Other indicators tested and rejected included phenol red, calmagite, universal indicator, litmus, bromcresol purple, methyl red, bromcresol green. The rejections were for a variety of reasons, but typically for lacking a sharp transition, having too many transitions within the titration, or inhabiting the “gray area” between unacceptable and acceptable performance. A diagram of the experimental apparatus is given in Figure 1. The computers were a mixture of 386 and 486
Table 1. Indicator Information Indicator (Abbreviation)
Solvent
Alizarin (Al)
Methanol
pH Transition Color Changea Range (22–24) (acid → base) 5.6 – 7.2 11.0–12.4
Y→R R→V
Bromophenol blue (BpB)
Water
3.0–4.6
Y→B
Bromothymol blue (BtB)
Water
6.0–7.6
Y→B
Congo red (CoR)
Water
3.0–5.0
B→O
m-Cresol purple (mCP)
Water
1.2–2.8 7.6–9.2
R→Y Y→V
Cresol red (CrR)
Water
0.2–1.8 7.2–8.8
R→Y Y→V
Methyl orange (MO)
Water
3.1–4.4
R→O
p-Nitrophenol (pNp)
Water
5.6–7.6
C→Y
Phenolphthalein (Pp)
Ethanol/water
8.3–10.0
C→R
Thymol blue (TB)
Water
1.2–2.8 8.0–9.6
R→Y Y→B
Thymolphthalein (Tp)
Ethanol/water
8.3–10.5
C→B
aY
= yellow, O = orange, B = blue, V = violet, R = red, C = colorless.
Syringe Pump Conductivity Probe
Interface Box
Computer
Stir Plate
Figure 1. Conductivity titration setup.
Journal of Chemical Education • Vol. 78 No. 11 November 2001 • JChemEd.chem.wisc.edu
In the Laboratory
machines. All used the “MPLI Package” from Vernier (Portland, OR), which included software (Windows-based, a DOSbased version is also available), a 12-bit A/D board installed within the computer, and an external interface box for connection of probes. All data were collected with the AIB-Board Timer disabled within the MPLI software via a software command. The conductivity probes were model CON-DIN from Vernier. The probe range was set to 0–2000 µS, and each probe was calibrated prior to use with air and a NaCl conductivity standard supplied with the probe. A wider conductivity range was necessary for some of the development experiments using more concentrated titrants, but the 0–2000 range was always used for the fully developed procedure. The syringe pumps were the Model A type from Razel Scientific Instruments (Stamford, CT). When equipped with a standard 50-mL plastic syringe, these pumps delivered a steady stream at 2.55 mL/min. Procedure The experiment can be performed by individuals, but groups are typically dictated depending on enrollment and the number of laboratory stations available (we use pairs). Each group is given a set of four indicators to be evaluated and is instructed to use only two drops per trial. Each indicator is used in turn for the titration of a 25.00-mL portion of pre-standardized 4.00 × 10᎑3 M HCl with a NaOH titrant of the same molarity. The titrant is delivered at a constant rate from a syringe pump, and the conductivity of the solution is plotted as a function of time after activation of the syringe pump. The students monitor the indicator color, recording the time when the change is observed for each. After the trial, a hard copy of the graph is printed, and the point where the indicator color change took place is marked on the graph. After all four indicators have been tested, the students are challenged to select the best of the set for use in Part II of the experiment. Part II entails using the same NaOH standard and a buret to titrate a 25.00-mL portion of an unknown HCl solution, without the conductivity probe. The endpoint is determined solely by observing the color change of the indicator selected from the original set of four, stopping the addition at the endpoint, and reading the volume delivered from the buret. Replicate trials are run, and the students are asked to report the concentration of the unknown acid. Obviously, if a correct indicator was not chosen in Part I, the quantitative results for Part II will be erroneous.
We did not feel there was undue risk in letting the students use these indicators, but the preparation of the indicator solutions was of more serious concern and was performed by qualified personnel only. Only the p-nitrophenol was gauged toxic enough to collect titration waste for disposal; all titration products from trials with other indicators were flushed down the laboratory drains. Results and Discussion A typical conductivity plot for the titration is shown in Figure 2A; Figure 2B presents a zoom view of the minimum conductance region showing where the different endpoints are observed. Alizarin, bromothymol blue, and p-nitrophenol show acceptable performance correlating with the conductivity minimum. Each set of four indicators dispensed to the student groups contained one of these three, along with three indicators selected from the other eight choices. Although cresol red, phenolphthalein, and m-cresol purple were near the conductivity minimum, experience has shown that a careful student experimentalist can recognize these as being slightly too late.
Hazards The acids and bases used in this experiment are caustic and will require proper eye protection and clothing. However, the low concentrations in use minimize both the safety risk, and waste disposal concerns. Of the indicators, methyl orange is categorized as toxic by ingestion and p-nitrophenol is considered highly toxic by ingestion, inhalation, and skin absorption. This toxicity is mitigated somewhat by the amount of dilution that takes place. The methyl orange solution is made up at 0.01% concentration, the p-nitrophenol at 0.1%. Only two drops are used in each of the titrations, which end up diluted to a total volume of around 50 mL.
Figure 2. A: Typical conductivity data. Titrant is 4.00 × 10᎑3 M NaOH, delivered at 2.55 mL/min. Sample is 25.00 mL of 4.00 × 10᎑3 M HCl. B: Observed endpoints of indicators. Titrant is 4.00 × 10᎑3 M NaOH, delivered at 2.55 mL/min. Sample is 25.00 mL of 4.00 × 10᎑3 M HCl. Abbreviations: BpB, bromophenol blue; CoR, Congo red; MO, methyl orange; pNp, p-nitrophenol; BtB, bromothymol blue; Al, alizarin; CrR, cresol red; Pp = phenolphthalein; mCP, m-cresol purple; TB, thymol blue; Tp, thymolphthalein.
JChemEd.chem.wisc.edu • Vol. 78 No. 11 November 2001 • Journal of Chemical Education
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In the Laboratory
This experiment can be run with a buret replacing the syringe pump for solvent delivery, and in fact was done that way during its initial trials. The buret was filled, set up to drip at a constant rate at the zero time datum, and then just left alone as data were collected. Although we could not expect a constant flow rate, the performance was reproducible enough to function properly. At the end of each trial, a hard copy of the conductivity-versus-time plot was made and the point at which the endpoint was observed was marked onto the plot. The biggest source of irritation with this method was that the flow rate from the buret varied from one trial to the next. The shape of the plots remained the same and it was still easy to note whether an endpoint correlated with the conductivity minimum, but the time axis varied owing to the flow rate variation, and some of the students were confused by this. Being able to reproduce the flow rate with the syringe pump caused less confusion and produced more precise output. Another experimental issue included proper placement of the conductivity probe. The probe needed to be placed well under the surface of the solution, or else it would encounter waves and agitation that caused a highly variable conductivity reading. It also proved beneficial to locate the conductivity probe as far away from the stir bar as possible and to limit the rate of stirring so a significant vortex was not produced, as the same noisy behavior would result. Finally, color blindness is always a vexation when this type of an experiment is undertaken, although if one member of the group was free from these problems, that would suffice. The occurrence of some form of color blindness is 10% for males and 1% for females (25). So if work is done in pairs, even a male duo would have only a 1% likelihood of both members being affected. Acknowledgments We would like to thank Jan Lentz, Frank McAndrew, and Harry Neumiller Jr. for their assistance during the preparation of this manuscript. In memory of Professor Robert G. Kooser, who always liked colorful experiments. W
Supplemental Material
Additional information and a laboratory handout are available in this issue of JCE Online. Literature Cited 1. Department of Chemistry, Knox College, Galesburg, IL. A Laboratory Manual for General Chemistry; John S. Swift Co.: St. Louis, MO, 1949.
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2. Sienko, M. J.; Plane, R. A. Experimental Chemistry; McGrawHill: New York, 1958. 3. Kildahl, N.; Varco-Shea, T. Explorations in Chemistry; Wiley: New York, 1996. 4. Wink, D. J.; Gislason, S. F.; Kuehn, J. E. Working with Chemistry: A Laboratory Inquiry Program; Freeman: New York, 2000. 5. Nelson, J. H.; Kemp, K. C. Laboratory Experiments, 8th ed.; Prentice Hall: Upper Saddle River, NJ, 2000. 6. Weiss, G. S.; Greco, T. G.; Rickard, L. H. Experiments in General Chemistry, 6th ed.; Macmillan: New York, 1993. 7. Bell, J. Chemical Explorations; D. C. Heath: Lexington, MA, 1993. 8. Hein, M.; Best, L. R.; Miner, R. L. Foundations of Chemistry in the Laboratory, 7th ed.; Brooks/Cole: Pacific Grove, CA 1990. 9. Beran, J. A. Chemistry in the Laboratory: A Study of Chemical and Physical Change, 2nd ed.; Wiley: New York, 1996. 10. Murov, S. L. Experiments in General Chemistry, 3rd ed.; Brooks/ Cole: Pacific Grove, CA, 1999. 11. Szafran, Z.; Pike, R. M.; Foster, J. C. Microscale General Chemistry Laboratory with Selected Macroscale Experiments; Wiley: New York, 1993. 12. University of Iowa General Chemistry Faculty and Staff. Experiments in Chemistry, 1996 revision; Stipes: Champaign, IL, 1996. 13. Amend, J. R.; Tucker, K. A.; Furstenau, R. P. J. Chem. Educ. 1991, 68, 857–860. 14. Flowers, P. A. J. Chem. Educ. 1997, 74, 846–847. 15. Mehta, M. A.; Dallinger, R. F. J. Chem. Educ. 1987, 64, 1019– 1020. 16. Patterson, G. S. J. Chem. Educ. 1999, 76, 395-398. 17. Kolthoff, I. M.; Sandell, E. B.; Meehan, E. J.; Bruckenstein, S. Quantitative Chemical Analysis, 4th ed.; Macmillan: Toronto, 1969. 18. El Seoud, O. A.; Bazito, R. C.; Sumodjo, P. T. J. Chem. Educ. 1997, 74, 562–565. 19. Dominguez, A.; Fernandez, A.; Gonzalez, N.; Iglesias, E.; Montenegro, L. J. Chem. Educ. 1997, 74, 1227–1231. 20. Wright, M. R.; Patterson, I. L. J.; Harris, K. D. M. J. Chem. Educ. 1998, 75, 352–357. 21. Rosenthal, L. C.; Nathan, L. C. J. Chem. Educ. 1981, 58, 656– 658. 22. Handbook of Chemistry and Physics, 70th Ed.; Weast, R. C., Ed.; CRC Press: Boca Raton, FL, 1979; pp D150–D151. 23. Harvey, D. Modern Analytical Chemistry; McGraw-Hill: Boston, 2000; p 289. 24. Harris, D. C. Quantitative Chemical Analysis, 4th ed.; Freeman: New York, 1995; p 296. 25. Morris, C. G. Psychology: an Introduction, 9th ed.; Prentice Hall: Upper Saddle River, NJ, 1996; p 97.
Journal of Chemical Education • Vol. 78 No. 11 November 2001 • JChemEd.chem.wisc.edu
