Chemical Education Today
Letters Another Hazard The most serious hazard is missing from the list on page 95 of the article Soap from Nutmeg: An Integrated Introductory Organic Chemistry Laboratory Experiment (1). On standing in a partly-filled container, undergoing distillation or evaporation, diethyl ether can develop explosive peroxides. Years ago when I was distilling ether from a few milliliters of solution in a micro apparatus, an explosion occurred that blew pieces of ringstand a number of feet. I escaped serious injury only because I happened to have gone to my desk on the other side of the lab to write something in my notebook. The ether solution had given a negative test for peroxides before distillation was started. Literature Cited
Literature Cited 1. Hazards in the Chemical Laboratory; Bretheerick, L., Ed.; Royal Society of Chemistry: London, 1986; p 292; Roberts, R. M.; Gilbert, J. C.; Rodewald, L. B.; Wingrove, A. S. Modern Experimental Organic Chemistry, 4th ed; Holt Saunders International: Tokyo, 1995; pp 548-549. 2. de Mattos, M. C. S.; Nicodem, D. E. J. Chem. Educ. 2002, 79, 94–95. Marcio C.S. de Mattos and David E. Nicodem Departamento de Química Orgânica Instituto de Química Universidade Federal do Rio de Janeiro Cx. Postal 68545 21945-970 Rio de Janeiro, Brazil
[email protected] or
[email protected] 1. de Mattos, M. C. S.; Nicodem, D. E. J. Chem. Educ. 2002, 79, 94–95. Jean B. Umland 21 South End Avenue, #423 New York, NY 10280-1063
[email protected] The authors reply: Dr. Umland is correct in her preoccupation concerning the risk of explosion when distilling ether. The formation of unstable peroxides by oxidation of diethyl ether is well known and documented in the literature (1). Explosion can occur when the ether is distilled off and the residue, rich in peroxides, is heated to a higher temperature by the heating element. In our experiment, described in the article (2), heating and distillation of the ether is done using previously heated hot water to prevent a fire hazard (see notes for the instructor in the supplemental material). Using this procedure, the ether is heated by warm water and the flask never reaches a high temperature. At the end of the distillation the flask still contains some ether along with a large amount of trimyristin. Under these conditions the risk of explosion is minimized even if the ether used initially contained peroxides. We thank Dr. Umland for her comments and observations.
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Color Changes in Indicator Solutions Silva, Pereira, and Sabadini, in their recent paper in this Journal (1), ask “Can an indicator solution change color upon dilution? Our usual experience in chemistry suggests to us that such behavior is impossible.” They then complete the paradox by saying “However, the answer to the question is affirmative, as shown in this paper.” The resolution of the paradox is to realize that dilution of a buffered solution of an indicator will not appreciably change the pH (no change in color), but dilution of a solution that contains only the indicator will certainly change the pH if the pH is not initially 7 (possibly a change in color). The experiment they describe involves a special case in which there is a change in color upon dilution of an unbuffered indicator solution. This special case is characterized by a pKa value for the indicator, called pKIn, of 5. We can estimate (2) that a 0.0001 molar solution of a weak acid whose pKa is 5 should have a pH of about 2.5 + 2, or 4.5. This sets the pH of the solution on the acidic side of the pKa, where the acidic form of the indicator will predominate. This possibility corresponds quite well to point 6 in Table 1 of the supplemental materials. For this point the reported indicator concentration is 1.092 × 10᎑4 M, the reported pH is 4.13, the ratio of [In᎑] to [HIn] was determined to be 0.123, and the log of this ratio is ᎑ 0.91.
Journal of Chemical Education • Vol. 79 No. 9 September 2002 • JChemEd.chem.wisc.edu
Chemical Education Today
The Henderson–Hasselbalch equation states that pKIn = pH – log[In᎑]/[HIn], and so pKIn must equal 4.13 – (᎑ 0.91), or 5, the value reported in the paper. Dilution of this solution by a factor of 10 should change the pH to 4.5 + 0.5, or 5.0 (2). This possibility corresponds pretty nearly to point 1 in Table 1 of the supplemental materials. For this point the reported indicator concentration is 0.182 × 10᎑4 M, the reported pH is 5.28, and the ratio of [In᎑] to [HIn] was determined to be 1.013. Thus dilution by a factor of 10 changed the ratio of [In᎑] to [HIn] from 0.123 to 1.013; that is, as the solution was diluted by a factor of 10 the fraction of the indicator in the form of the conjugate base went from about 1/9 to about 1/2. Further dilution would make the solution even less acidic, and finally, as the pH approached 7, practically all of the indicator would be present in the basic form; the color change would be complete. In order to see a color change upon dilution of an unbuffered solution of an indicator of the form HIn, the pKa of the indicator should be 2 or more units on the acidic side of 7 so that dilution can take the pH of the solution to the basic side of the pKa, but not so far on the acidic side of 7 that the solubility limit of the indicator will prevent the indicator from setting the pH initially to the acidic side of the pKa. For example, if the pKa is 3, it will take a 0.01 M solution to set the pH to 2.5 (2). “Can an indicator solution change color upon dilution?” A better answer is: Only if the solution is unbuffered and the pKa of the indicator, HIn, lies between about 3 and 5. Literature Cited
The author replies: First, we would like to thank Addison Ault for the useful comments about our recent paper in this Journal (1). The main goal of our paper is the resolution of the paradox that a simple dilution can change the color of an indicator (bromocresol green) solution. Students, trying to solve this paradox, will realize that the color change is due to the shift in the equilibrium HInyellow = H+ + Inblue᎑; that is, the degree of ionization increases upon dilution. Therefore, this unusual color change behavior can be easily understood by a very common equilibrium concept. Certainly, this phenomenon will only be clearly observed in water for indicators whose pKIn is near 5. The approach developed by Ault (2), which is useful to estimate pH from the pK values, can be also used to estimate the pH value of the bromocresol green solution upon dilution. Ault’s comment that no color changes will be observed in a buffered indicator solution is correct, and we can also expect no color changes for other cases involving a diluted aqueous acid or basic solution as a solvent. Literature Cited 1. Silva, C. R.; Pereira, R. B.; Sabadini, E. J. Chem. Educ. 2001, 78, 939–940. 2. Ault, Addison. J. Chem. Educ. 1999, 76, 936–938. E. Sabadini Instituto de Química Universidade Estadual de Campinas Caixa Postal 6154, 13083-970 Campinas, SP, Brazil
1. Silva, C. R.; Pereira, R. B.; Sabadini, E. J. Chem. Educ. 2001, 78, 939–940. 2. Ault, Addison. J. Chem. Educ. 1999, 76, 936–938. Addison Ault Department of Chemistry Cornell College Mt. Vernon, IA 52314
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