The Blue Bottle Revisited - ACS Publications

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Letters The Blue Bottle Revisited The article “The Blue Bottle Revisited” by Walter R. Vandaveer IV and Mel Mosher (J. Chem. Educ. 1997, 74, 402) builds on the classic demonstration in which an alkaline glucose solution containing methylene blue is repeatedly shaken, causing it to turn from clear to blue. Their suggestion to add different dyes to this “blue bottle” provides colorful variations. However, the footnote referring to the fact that the intensity of the pink color resulting from the use of phenolphthalein indicator “with each cycle…becomes lighter until it is pale pink in the reduced form” is misleading. As noted in this Journal (1), the familiar pink form of phenolphthalein (and other similar dyes including malachite green, crystal violet, and bromophenol blue) actually reacts with hydroxyl ion which converts the resonant colored quinoid form and produces the nonresonant colorless carbinol form. This fading will occur independent of the shaking “cycles” which dissolve oxygen in the solution. At room temperature, the half-life for the alkaline fading reaction of phenolphthalein in 0.3 N hydroxide solution is about 2 minutes. Since the reaction is actually reversible (2), the final solution remains slightly pink. If a pinkto-purple demonstration is desired, a dye that doesn’t react with base (perhaps rhodamine B) might be better. Or, even better, add resazurin to the “blue bottle” mixture. Several sources (3, 4) refer to using resazurin as a variation to using methylene blue. This dye produces a brilliant fluorescent pink in its oxidized form with a slight swirling of the solution; i.e. a low concentration of dissolved oxygen is sufficient to produce this color. In contrast, much more vigorous shaking (a higher dissolved oxygen concentration) is needed to oxidize the methylene blue to its blue form. Also, it is interesting that the reduction of the methylene blue as it oxidizes the glucose is faster than the corresponding reduction of the resazurin (4). The net effect is quite amazing: an initially colorless solution is swirled and becomes pink; it is then more vigorously shaken and becomes purple (blue + pink); on sitting, the process reverses as the purple fades away, leaving a pink solution which then fades back to colorless! And this whole cycle can be repeated dozens of times with the same solution. While both of these demonstrations are well known separately, the suggestion to combine them (5) is less commonly seen. Yet this provides an excellent example of how different chemical reactions occur at different rates— in both the oxidation step which produces the color and the reduction step which causes the fading. Literature Cited 1. Nicholson, L. J. Chem. Educ. 1989, 66, 725. 2. Hile, L. J. Chem. Educ. 1991, 68, 446. 3. Cook, A. G.; Tolliver, R. M.; Williams, J. E. J. Chem. Educ. 1994, 71, 160.

4. Tested Demonstrations in Chemistry, Vol. 1; Gilbert G. L., et al., Eds.; ACS Publishing: Washington, DC, 1994; p G-3. 5. Shakhashiri, B. Chemical Demonstrations: A Handbook for Teachers of Chemistry, Vol. 2; University of Wisconsin Press: Madison, 1985; p 143. Lloyd Hile Chemical Engineering Department California State University Long Beach, CA 90840

The authors reply: The comments made by Hile about a footnote in our demonstration, “The Blue Bottle Revisited”, (1) suggest that the slow fading of the red color of phenolphthalein is due to the reaction of the basic form of phenolphthalein with excess hydroxide ion (2) to form a tricharged, colorless form of phenolphthalein (see ref. 2 for structures). We believe that in the demonstration presented the formation of the tricharged colorless form is not the only reason why the red color of the reduced form of the methylene blue/phenolphthalein slowly fades from an intense red to a pale pink after several oxidation-reduction cycles. Phenolphthalein in the presence of strong reducing agents is known to be reduced to leucophenolphthalein (also called phenolphthalin) (3) which is the basis of the classic Kastle–Meyer reagent for the detection of blood. The slow fading of the red to pink color of the phenolphthalein is most likely due to a combination of the glucose/hydroxide reduction to leucophenolphthalein and the reaction of the basic form of phenolphthalein with hydroxide to the tricharged form. Both of these products are colorless. As support for this idea, intense shaking of the pink colored reduced solution of methylene blue/phenolphthalein will regenerate the solution into a purple oxidized solution that will produce a more intense pink-red solution in the next cycle of the reduced solution. This is consistent with the air oxidation of leucophenolphthalein back to phenolphthalein. If this is indeed the case, this demonstration becomes similar to the two cited by Hile (4, 5) and one using methylene blue/indigo carmine (6). Literature Cited 1. Vandaveer, W. R. IV; Mosher, M. J. Chem. Educ. 1997, 74, 402. 2. Nicholson, L. J. Chem. Educ. 1989, 66, 725. 3. The Merck Index, 12th ed.; Budavari, S. et al., Eds.; Merck Research Laboratories: Whitehorse Station, NJ, 1996; entry 7394. 4. Cook, A. G.; Tolliver, R. M.; Williams, J. E. J. Chem. Educ. 1994, 71, 160. 5. Tested Demonstrations in Chemistry, Vol. 1; Gilbert G. L., et al., Eds.; ACS Publishing: Washington, DC, 1994; p G-3. 6. Lang, C. M.; Showalter, D. L.; Shulfer, G. J. Yes Virginia…Learning Chemistry Can Be Fun; University of Wisconsin: Stevens Point, 1989; p 18–19. Walter R. Vandaveer, IV, and Melvyn M. Mosher Department of Chemistry Missouri Southern State University Joplin, MO 64801

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Letters Chemistry in the Public Domain

Tetrahedral Bond Angle

I enjoyed the article by Sidney Toby in the November 1997 issue of this Journal (J. Chem. Educ. 1997, 74, 1285– 1287). I must take exception to the author’s assertion that ammonia–air mixtures are nonexplosive. During the period between my two teaching careers, I spent 17 years as a fire and explosion investigator, and several times found that simple answers are not always right. In this letter I will refer only to the literature that is on most chemistry teacher’s desks, not the fire and explosion profession’s literature. The Merck Index, 11th edition in article 510 states: “Mixtures of ammonia in air will explode when ignited under favorable conditions but ammonia is generally considered to be nonflammable.” Lange’s handbook, 11th edition in Table 11-10 lists the lower explosive limit for ammonia in air at 4.5%. There are many other references that support the fact that ammonia–air mixtures are explosive, but difficult to ignite. These two are available and familiar to most chemists. Chlorinated solvents are also a class of compounds that, while generally considered nonflammable, will also explode under favorable conditions. I once investigated a case where a chemist had assured an appliance manufacturer that methylene chloride was nonflammable and therefore its vapors were nonexplosive. A large (multistory) vapor degreaser exploded as a result of this ill-considered advice. I have experimentally exploded mixtures of both ammonia in air and methylene chloride in air.

I have been reading the Journal of Chemical Education and even contributed one paper four years ago (Ferreira, R. J. Chem. Educ. 1993, 70, 483). So I was pleased to learn that there are still people interested in how to calculate the tetrahedral bond angle (Glaister, P. J. Chem. Educ. 1997, 74, 1086). Fifty years ago last May, I published my first note in your magazine on this very problem (Ferreira, R. J. Chem. Educ. 1947, 24, 246). The most meaningful way to find the value of θ was shown to me by, of all people, Dick Feynman. Methane has no permanent electric dipole moment, but each C–H bond has one. Hence, any bond moment must be equal (but with opposite sign) to three other bond moments which make an angle θ, the cosine of which must be –1/3. Hence, θ = 109o28′.

John L. Odom Chattanooga School for the Arts and Sciences Chattanooga, TN 37403

The author replies: I thank John Odom for his comments. Before I wrote the paper, I put a lecture bottle of ammonia in a fume hood, opened the valve slightly and, somewhat timorously, put a lighted match in the ammonia stream several times. The flame flickered but whenever the match was removed, the flame went out. I was unable to make ammonia burn in the absence of the flame and I therefore support the U. S. Department of Transportation’s classification of it as a “nonflammable gas” under ordinary conditions. Can an ammonia–air mixture explode? Handbooks vary in their assessment but Bretherick’s Handbook (4th Ed., Butterworth) quotes 15.8% in air as a lower explosion or flammability limit. This is an order of magnitude higher than the value quoted in the article; nevertheless my statement that ammonia does not form potentially explosive fumes in air under ordinary conditions is probably wrong. Sidney Toby Department of Chemistry Rutgers University Piscataway, NJ 08855-0939

Ricardo Ferreira Department of Physics University of Calfornia, San Diego La Jolla, CA 92037-0345

Formation and Dimerization of NO2 I have strong misgivings about the article, “Formation and Dimerization of NO2” (J. Chem. Educ. 1997, 74, 1340– 1342). Aside from the unfortunate implication in the calculations that individual gases in a mixture occupy separate volumes, there is an overriding disadvantage to the experiment described: The results are terrible! A student who has any intuitive feel for quantitative relationships will be appalled by the scatter in the values obtained for the equilibrium “constant” and its distance from the accepted value. What appalls me, as one who spent 30 years trying to impart to students the faintest glimmer of the notion of “significant figures”, is the statement by the authors that 2.7±1.0 “compares reasonably well” with 8.6! More important, the beginning scientist may not know whether this is the result of poor data, or an innate characteristic of the scientific method. The implication that these are the kinds of data upon which the law of combining volumes is based will mislead the typical student, and confuse the circumspect one. The authors begin by observing that “experiments demonstrating…law of combining volumes are virtually absent” from general chemistry lab manuals. What their article illustrates most clearly is the reason for this observation: Successful performance of such an experiment is technically very demanding, and beyond the abilities of the typical general chemistry class. What would be the effect on a neophyte piano student of expecting him or her to play “The Minute Waltz” before ever having practiced scales? Discouragement and disillusion, surely, and perhaps the urge to try a different instrument! Edwin F. Meyer Emeritus Professor of Chemistry, DePaul University 1022 Dobson St. Evanston, IL 60202


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Letters Let’s Dot Our I’s and Cross Our T’s I read with great interest the letters “Redox Challenges” by David M. Hart and by Noel S. Zaugg (J. Chem. Educ. 1996, 73, A226), the article by R. Stout that they both had cited, and the author’s reply to the polemics. I was intrigued also because I knew very well that I had read about this “Chemical Monster” many years ago in this Journal ! Unfortunately I didn’t remember even the year of publication; Chem. Abstr. also didn’t mention that particular article. Nevertheless after two days of searching I found what I was looking for in the November issue, 1943, on page 570. Here is this very short unsigned article: “Try This One! L. S. Foster, until recently one of our associate editors, sent us the following equation to balance. The editor bears witness that it can be done—although it took $1.27 worth of his time (at current rates) to do it. We’d like to know who can do it in the shortest time. Ready—go!” And below was the famous equation [Cr(N2H4CO)6]4[Cr(CN)6]3 + MnO 4– + H+ → Cr2O7= + CO2+ NO 3– + Mn ++ + H2O

So we can see again that “There is no new thing under the sun” (Ecclesiastes, 1:9)! But I wonder whether anybody knows the salary of the editor in the midst of 1940s to estimate the time necessary to balance the equation in minutes instead of hard currency? Equally marvelous was the first reply to that challenge in the “Out of the Editor’s Basket” column. Norris Rakestraw wrote on page 30 in the January issue, 1944: “We received a number of solutions to the equation which appeared in the November issue (page 570). All three of the common methods were represented: the “valence change method,” the “ion-electron method,” and the “algebraic method.” The best time was reported by Miss Dorothy Elkind, of Villa Park, Illinois, who used the algebraic method and who says, “By the method outlined, it does not take more than 15 minutes to balance the equation.” We assume that this was her time, although she doesn’t actually say so. If not, the speed record goes to another contestant who reports “less than 16 minutes”—but whose solution didn’t balance!... Various interesting comments were offered, including the suggestion of assigning arbitrary zero valence numbers to all the elements in the complex, which of course leads to a correct solution if consistently followed...” But my own opinion is that all these puzzles are nothing more than mathematical and, partly, chemical “games” because it is highly improbable that such “monsters” are stoichiometric equations! Directly connected with this question is the article by William C. McGavock “Nonstoichiometric Equations” (J. Chem. Educ. 1945, 22, 269) and many others published in J. Chem. Educ. from the first issues (see, for example, 1926, 3, 425–431, 1430–1431; 1927, 4, 1021–1030, 1158–1167; 1929, 6, 479–485, 527–535, 1136–1138; 1930, 7, 1180, 1688–1693). In the end I would like to quote Leo Lehrman from The College of the City of New York. He

wrote 66 years ago: “To make such practice a mere mathematical exercise…is to put us back many years” (J. Chem. Educ. 1932, 9, 944–945). Ilya A. Leenson Department of Chemistry Moscow State University 119899 Russia

Formation and Dimerization of NO2 I have strong misgivings about the article, “Formation and Dimerization of NO2” (J. Chem. Educ. 1997, 74, 1340– 1342). Aside from the unfortunate implication in the calculations that individual gases in a mixture occupy separate volumes, there is an overriding disadvantage to the experiment described: The results are terrible! A student who has any intuitive feel for quantitative relationships will be appalled by the scatter in the values obtained for the equilibrium “constant” and its distance from the accepted value. What appalls me, as one who spent 30 years trying to impart to students the faintest glimmer of the notion of “significant figures”, is the statement by the authors that 2.7±1.0 “compares reasonably well” with 8.6! More important, the beginning scientist may not know whether this is the result of poor data, or an innate characteristic of the scientific method. The implication that these are the kinds of data upon which the law of combining volumes is based will mislead the typical student, and confuse the circumspect one. The authors begin by observing that “experiments demonstrating…law of combining volumes are virtually absent” from general chemistry lab manuals. What their article illustrates most clearly is the reason for this observation: Successful performance of such an experiment is technically very demanding, and beyond the abilities of the typical general chemistry class. What would be the effect on a neophyte piano student of expecting him or her to play “The Minute Waltz” before ever having practiced scales? Discouragement and disillusion, surely, and perhaps the urge to try a different instrument! Edwin F. Meyer Emeritus Professor of Chemistry, DePaul University 1022 Dobson St. Evanston, IL 60202

Presidential Education Policy Aside from the fact that the editor of the Association Reports: 2YC3 is a member of the faculty of San Jacinto Community College, I cannot find a single additional reason for publishing President Clinton’s educational/political remarks which he presented at San Jacinto Community College (J. Chem. Educ. 1997, 74, 1392–1393). What’s next, a similarly vapid political reply from Gingrich outlining his contributions to balancing the budget and his desire to have “an educational system that works for all”? Harold T. McKone Department of Chemistry Saint Joseph College West Hartford, CT 06117

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Journal of Chemical Education • Vol. 75 No. 9 September 1998 • JChemEd.chem.wisc.edu


Oxygen vs Dioxygen: Diatomic/Monatomic Usage Are you as tired as I am of having freshman chemistry students who divide by 32 when determining the empirical formula of a compound containing oxygen or by 35.5 when finding the number of moles of gaseous elemental chlorine in a stoichiometry problem? How about students who persistently write weak acids as ions when writing gross and net ionic equations? I am using two techniques that have all but eliminated these errors in my classes and hope, by this letter, to influence you to do the same. Here is what I do, thanks to the example of Bill Plachy in the Department of Chemistry at San Francisco State University. I consistently, in writing and speaking, refer to hydrogen gas as “dihydrogen”, chlorine gas as “dichlorine”, nitrogen gas as “dinitrogen”, etc. I use “hydrogen” only when referring to the monatomic species. In the case of weak acids, the subscript (aq) is misinterpreted by many students to mean that the substance ionizes completely, as it does when used for salts. Of course (aq) means only that the substance is soluble in water. Strong acids and bases can be labeled (aq) without confusion because they too ionize completely. Why not label weak acids, instead of (aq), as (wa), weak acid? Weak bases could be labeled (wb). If you agree with these ideas, please feel free to adopt them in your class presentations, written handouts, and the next edition of the text you have written. If enough people start to adopt this system it will become commonly used, and we will have blotted out these two errors once and for all. Jared B. Sharon Department of Chemistry Canada College Redwood City, California 94061

Ionization or Dissociation? I applaud the article by Emeric Schultz in the July, 1997 issue of the Journal titled “Ionization or Dissociation?” Finally someone has raised the issue of the confusion surrounding these two terms in elementary texts and begun the dialog to correct it. The article is logical and consistent as far as it goes, and I concur with its final recommendation. However, a set of definitions for these two terms that could be used by educators and authors working at the introductory level is needed. As a starting point, and after consultation and recommendations from Schultz, I propose the following definitions:

In general Ionization: a chemical change involving the transformation of an uncharged species into an ion (or ions) or other charged entities. Dissociation: a change involving the separation of a chemical entity into its component parts without creation of new ions or other charged entities.

In aqueous solution Ionization: a chemical change involving the transformation of water into H3O+ or OH– or both. Dissociation: a change involving the separation of a chemical entity into its component parts in water without creation of H3O+ or OH–. Hopefully others will add to and enhance these definitions so that we can communicate these concepts as clearly as possible. David L. Adams Babson College Babson Park, MA 02157-0310

Cheating Probabilities The note by Rizzuto and Walters on “Cheating Probabilities on Multiple Choice Tests” (J. Chem. Educ. 1997, 74, 1185) implies that a significant measure of the probability that two examination papers were not created independently can be obtained using the binomial probability distribution. They do point out that the probability of correlated answers is most probably higher than predicted by the formula, and suggest a value of .001 may be appropriate for considering the possibility of cheating. It can easily be shown that this limit may be completely inappropriate. For example, consider an examination with 50 questions. If two students each get an 80% on the examination, the minimum number of questions that will have the same answer is 30, (each misses 10 questions but a different 10 questions). The probability of this occurring by chance, using Table 1 is 5.8 ⫻ 10–10. This would then suggest that it is very unlikely that two students would have exam grades around 80% without cheating! In summary, the reason that “it is not well known that the probability of a student having the same answers as another student can be easily calculated using the binomial distribution” is that it is not true. Charles D. Jonah Argonne National Laboratory Chemistry Division Argonne, IL 60439

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The Blue Bottle Revisited - ACS Publications

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