Beer's Law: The Real Hazards (the author replies)

Chemical Education Today

Letters Beer’s Law: The Real Hazards A recent article (Muyskens, M. A.; Sevy, E. T. J. Chem. Educ., 1997, 74, 1138) gave an example of the linearity of the absorbance, A, with the concentration of the absorbing species, up to A = 3.5. However, it gave the wrong impression as to why absorbance measurements in excess of 2 have tended to be treated with some suspicion. If the absorbance exceeds 2, then the light transmitted through the cell is less than 1% of the intensity incident upon the cell. Should any stray light reach the detector, by whatever means or from whatever origin, its intensity may then be a significant fraction of that of the transmitted light. The consequential error becomes the greater the higher the absorbance. For example, if the transmitted light were 1.0% of the incident intensity and the stray light were 0.1% of it, then the perceived absorbance would be log10(100/1.1) = 1.96, rather than the correct value of 2.0. This represents an error of 2%. But if only 0.1% of the incident light were to be transmitted so that the correct absorbance was 3.0, the absorbance would appear to be log10(100/0.2) = 2.70 with the same amount of stray light. There would now be an error of 10%. From the linearity achieved by Muyskens and Sevy up to A = 3.5, one may deduce that in their set-up, the stray light could not have been more than 0.003% of the intensity of the incident beam. This considerable achievement is helped by (i) having an intense light beam and (ii) using a

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monochromatic light source. For both features, a laser source is ideal. However, one cannot expect a commercial scanning spectrophotometer to employ a laser light source. S. R. Logan University of Ulster Coleraine, N. Ireland

The author replies: In the response to our article, S. R. Logan makes the valid point that, in considering deviations from Beer’s law, stray light effects should not be neglected. This can be an important factor in commercial scanning spectrophotometers where most students will encounter the application of Beer’s law. Our article is intended to give an interesting counter-example for deviations from Beer’s law, and it considers the usual causes for deviations including chemical effects and stray light. We did not intend to suggest that chemical effects are necessarily more important than stray light in causing deviations. In his analysis of how stray light affects our experimental measurements, Logan mentions two important laser light properties but neglects to mention several other features of our experiment that contribute to avoiding stray light problems. First, the laser is pulsed which allows us to distinguish between laser light and stray light that is not synchronized with the pulse frequency, such as room light. Second, we need good sensitivity for accurate intensity measurements when the high in-

Journal of Chemical Education • Vol. 75 No. 12 December 1998 • JChemEd.chem.wisc.edu


Letters Letters continued from page 1515

Qual from a Different Viewpoint With regard to the article “Qual from a Different Viewpoint” by Michael Laing (J. Chem. Educ. 1993, 70, 666) I have some concerns about his Comments in the Instruction Sheet for Block 2 which is presented as typical. Because the author offered copies of the complete set of instructions, I recommend people to give it a critical reading (followed by the appropriate changes) before incorporating it in lab. This may avoid imprecisions in the freshman class. In the mentioned Comments, I suggest changing: Comment 1: “The colors of solutions of Cr3+ and Fe3+ will vary depending on the identity and concentration of the anion present” to “The colors of solutions of Cr3+ and Fe3+ might vary depending on the identity and concentration of the species present”. (The term anion may be confusing since the involved anions are colorless.) Comments 2, 3: “Both Al3+ and Cr3+ are clearly amphoteric” to “Both Al(OH)3 and Cr(OH)3 are clearly amphoteric” (writing hydroxides seems to me acceptable at this level and facilitates illustrating amphoterism). Adding another sentence such as “Mn(OH)2 is very air sensitive” should be also appropriate. Comment 8: “The dark red brown compound [Fe(SCN)]2+ is quite characteristic of iron” to “The red complexes [Fe(SCN)n](3-n)+ are quite characteristic of Fe3+”. Comment C: “Of all these metals, only Mn2+ will form a sulfide…” to “Of all these ions,…” (Fe2+ also forms a sulfide in NH3-H2S). Comment 7: “Extract this into dichloromethane: note the change from brown (I2 in water) to purple (I2 in nonpolar organic)”. It is probably a slip by both author and reviewers (it is quite evident that CH2Cl2 is not a nonpolar solvent, and also that a correspondence between the color of I2 solutions and solvent polarity is missing), but necessary to be corrected. Thus, the words “nonpolar organic” must be replaced by, for example, “common chlorinated solvents”. Although in the mentioned context, if we respect “I 2 in water”, “dichloromethane” should be appropriate. Ending the sentence with “The iodide is oxidized to iodine by Fe3+” is convenient. Comment 4: “Mn(OH)2 does not form because [OH–] is too low in aqueous NH3”. This is doubly incorrect. According to textbooks and my personal experience, the characteristic behavior of Mn(II) common salts in solution, when ammonia is added, is to form a white (very pale pink) precipitate which darkens progressively by oxidation. In dilute solutions this precipitate usually manifests itself by muddying the solution. After some minutes on standing, a cloudy gelatinous pale brown (white when conducting the test with rigorous exclusion of air and any other oxidant) precipitate separates. The only time I have failed to find precipitation, except when a complexing agent was present, occurred years ago as a consequence of missing all the Mn(II) soluble salts from our lab. I ordered a sophomore student to prepare a solution of MnCl2 by treating a stirred suspension of MnCO3 pow-

der in some water with “sufficient” 6 M HCl. In this process, and because MnCO3 was not pure (it is usually accompanied by minor amounts of higher oxidation state manganese species), the preparation of a clear solution was not straightforward even after heating. The student, having decided to obtain a clear solution (being aware of the difference between solution and suspension and its importance in analytical tests) at whatever the costs, added a considerable amount of HCl without success (I recommend adding N2H5Cl, by portions of a few mg, to the warm acid mixture until a clear solution is produced). After filtering, the resulting MnCl2 and HCl solution was used for reactivity tests. Observing that no precipitate was formed with aqueous NH3, I asked the student for the detailed procedure of preparing the solution. Thus, I could explain this “anomalous” behavior. It is known that Mn(OH)2 does not form from an ammonium-rich solution (consistently, Mn(OH)2 may be dissolved by aqueous ammonium salts). So when NH 3 was added to the mentioned solution, HCl was primarily transformed into NH4Cl, and we, in fact, were intending to precipitate Mn(OH)2 by adding NH3 to a NH4Cl-rich MnCl2 solution. This was the reason for our failure. This also could explain why the author did not observe precipitation, provided that the parent solution was sufficiently acidic (of course, he uses solutions “…in acid medium, as appropriate”). However, it is quite improbable that a teacher ignores this. I suppose that the author is using a sufficient acid solution to guarantee that Mn(OH)2 does not form, thus generating useful discussions. If this was the case, some more information should have been included in order to avoid readers becoming bewildered. And, of course, comment 4 would still remain inadmissible, requiring a change to “Mn(OH)2 does not form because [OH–] is too low when [NH4+] is too high. This clearly invites discussion. A complementary sentence on the possible formation of Cr(III)-NH3 soluble complexes should be also appropriate in this Comment 4. Francisco J. Arnáiz Departamento de Química Inorgánica Universidad de Burgos 09001, Burgos, Spain

/ About Letters to the Editor Letters to the Editor may be submitted to the editorial office by regular mail (JCE, University of Wisconsin– Madison, Department of Chemistry, 209 N. Brooks, Madison, WI 53715-1116), by fax (608/262-7145), or by email ( [email protected]) . Be sure to include your complete address including email, your daytime phone number, and your signature. Your letter should be brief (400 words or less) and to the point; it may be edited for style, consistency, clarity, or for space considerations.

JChemEd.chem.wisc.edu • Vol. 75 No. 12 December 1998 • Journal of Chemical Education

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cident laser intensity has been reduced to very low levels by absorption. Two features that contribute to this sensitivity are adjustment of the sensitivity of the digital oscilloscope and signal averaging, which improves our signal-to-noise ratio. Mark Muyskens Department of Chemistry and Biochemistry Calvin College Grand Rapids, MI 49546

. Photodimerization of Maleic Anhydride I have introduced the photochemical reaction of maleic anhydride described in “Three Puzzles for Organic Laboratory” by David Todd and Miles Pickering (J. Chem. Educ. 1988, 65, 1100) to our undergraduate organic laboratory. As the authors and I expected, the responses from the students who conducted the experiment are positive and encouraging. The Puzzle 3 can be solved by measuring the melting point of the product. But why is the cis, trans, cis product, rather than other possible isomers, formed in the photodimerization of maleic anhydride? The two authors give no definite and convincing explanation. The stereochemical process, in my opinion, may be described below as scheme 1, instead of scheme 2 (where X = COOMe).

Because of the stereohindrance, the configuration of the photochemical reaction intermediate can only be Ia, rather than Ib, and then the esterification of Ia gives cis, trans, cis product. This explains why the other possible isomers, including the thermodynamically stable trans, trans, trans form is not observed in the experiment. Zhijun Zhang Department of Chemistry Kwansei Gakuin University Nishinomiya 662, Japan

Letters continued on page 1551


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Letters Letters continued from page 1515

Qual from a Different Viewpoint With regard to the article “Qual from a Different Viewpoint” by Michael Laing (J. Chem. Educ. 1993, 70, 666) I have some concerns about his Comments in the Instruction Sheet for Block 2 which is presented as typical. Because the author offered copies of the complete set of instructions, I recommend people to give it a critical reading (followed by the appropriate changes) before incorporating it in lab. This may avoid imprecisions in the freshman class. In the mentioned Comments, I suggest changing: Comment 1: “The colors of solutions of Cr3+ and Fe3+ will vary depending on the identity and concentration of the anion present” to “The colors of solutions of Cr3+ and Fe3+ might vary depending on the identity and concentration of the species present”. (The term anion may be confusing since the involved anions are colorless.) Comments 2, 3: “Both Al3+ and Cr3+ are clearly amphoteric” to “Both Al(OH)3 and Cr(OH)3 are clearly amphoteric” (writing hydroxides seems to me acceptable at this level and facilitates illustrating amphoterism). Adding another sentence such as “Mn(OH)2 is very air sensitive” should be also appropriate. Comment 8: “The dark red brown compound [Fe(SCN)]2+ is quite characteristic of iron” to “The red complexes [Fe(SCN)n](3-n)+ are quite characteristic of Fe3+”. Comment C: “Of all these metals, only Mn2+ will form a sulfide…” to “Of all these ions,…” (Fe2+ also forms a sulfide in NH3-H2S). Comment 7: “Extract this into dichloromethane: note the change from brown (I2 in water) to purple (I2 in nonpolar organic)”. It is probably a slip by both author and reviewers (it is quite evident that CH2Cl2 is not a nonpolar solvent, and also that a correspondence between the color of I2 solutions and solvent polarity is missing), but necessary to be corrected. Thus, the words “nonpolar organic” must be replaced by, for example, “common chlorinated solvents”. Although in the mentioned context, if we respect “I 2 in water”, “dichloromethane” should be appropriate. Ending the sentence with “The iodide is oxidized to iodine by Fe3+” is convenient. Comment 4: “Mn(OH)2 does not form because [OH–] is too low in aqueous NH3”. This is doubly incorrect. According to textbooks and my personal experience, the characteristic behavior of Mn(II) common salts in solution, when ammonia is added, is to form a white (very pale pink) precipitate which darkens progressively by oxidation. In dilute solutions this precipitate usually manifests itself by muddying the solution. After some minutes on standing, a cloudy gelatinous pale brown (white when conducting the test with rigorous exclusion of air and any other oxidant) precipitate separates. The only time I have failed to find precipitation, except when a complexing agent was present, occurred years ago as a consequence of missing all the Mn(II) soluble salts from our lab. I ordered a sophomore student to prepare a solution of MnCl2 by treating a stirred suspension of MnCO3 pow-

der in some water with “sufficient” 6 M HCl. In this process, and because MnCO3 was not pure (it is usually accompanied by minor amounts of higher oxidation state manganese species), the preparation of a clear solution was not straightforward even after heating. The student, having decided to obtain a clear solution (being aware of the difference between solution and suspension and its importance in analytical tests) at whatever the costs, added a considerable amount of HCl without success (I recommend adding N2H5Cl, by portions of a few mg, to the warm acid mixture until a clear solution is produced). After filtering, the resulting MnCl2 and HCl solution was used for reactivity tests. Observing that no precipitate was formed with aqueous NH3, I asked the student for the detailed procedure of preparing the solution. Thus, I could explain this “anomalous” behavior. It is known that Mn(OH)2 does not form from an ammonium-rich solution (consistently, Mn(OH)2 may be dissolved by aqueous ammonium salts). So when NH 3 was added to the mentioned solution, HCl was primarily transformed into NH4Cl, and we, in fact, were intending to precipitate Mn(OH)2 by adding NH3 to a NH4Cl-rich MnCl2 solution. This was the reason for our failure. This also could explain why the author did not observe precipitation, provided that the parent solution was sufficiently acidic (of course, he uses solutions “…in acid medium, as appropriate”). However, it is quite improbable that a teacher ignores this. I suppose that the author is using a sufficient acid solution to guarantee that Mn(OH)2 does not form, thus generating useful discussions. If this was the case, some more information should have been included in order to avoid readers becoming bewildered. And, of course, comment 4 would still remain inadmissible, requiring a change to “Mn(OH)2 does not form because [OH–] is too low when [NH4+] is too high. This clearly invites discussion. A complementary sentence on the possible formation of Cr(III)-NH3 soluble complexes should be also appropriate in this Comment 4. Francisco J. Arnáiz Departamento de Química Inorgánica Universidad de Burgos 09001, Burgos, Spain

/ About Letters to the Editor Letters to the Editor may be submitted to the editorial office by regular mail (JCE, University of Wisconsin– Madison, Department of Chemistry, 209 N. Brooks, Madison, WI 53715-1116), by fax (608/262-7145), or by email ( [email protected]) . Be sure to include your complete address including email, your daytime phone number, and your signature. Your letter should be brief (400 words or less) and to the point; it may be edited for style, consistency, clarity, or for space considerations.


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Beer's Law: The Real Hazards (the author replies)

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