Chemical Education Today
Letters Boiling Point and Molecular Weight Thomas T. Earles (J. Chem. Educ. 1995, 72, 727) shows nicely that dispersion forces can be stronger than hydrogen bonds. Twice in this short article, however, boiling points are related to molecular weights. For many readers, especially unsophisticated students, this implies a causal connection where there is essentially none, as shown long ago (1, 2). Molecular polarizabilities, however, do have just such a connection with boiling points (3) and they can be estimated beginning in the same way as with molecular masses, that is, by adding atomic values, then including some corrections (3). And then no quantitative function relation is under discussion anyway, as in the article mentioned, molecular volumes are just as easy to talk about and much more appropriate theoretically than are molecular masses. Earles is far from alone in this, so let’s try, once and for all, to correct the record. Literature Cited 1. Bradley, D. C. Nature 1954, 174, 323. 2. Rich, R. L. Periodic Correlations; Benjamin-Cummings: Menlo Park, CA, 1965; p 69. 3. Rich, R. L. J. Chem. Educ. 1995, 72, 9 and refs therein.
Ronald L. Rich 112 S. Spring St. Bluffton, OH 45817
Osmotic Pressure and Electrochemical Potential The key sentence of a recent paper (J. Chem. Educ. 1995, 72, 713–714) is the first one: “There is a close correspondence between osmotic pressure and electrochemical potential that might help students with better understanding of both.” The recognition of the analogy between osmotic and electrochemical phenomena dates back to W. Nernst, who used the theory of van’t Hoff to explain the dissolution of metals and who has elaborated the theory of galvanic cells and electrodes on the basis of osmotic equilibria (W. Nernst: Die elektromotorische Wirksamkeit der Ionen. Z. Phys. Chem. 1889, 4, 129). Nernst derived his famous equation in the form as follows:
π = RT nε ln (P/ p) where π is the potential difference between a metal and an electrolyte, P is the “electrolytic solution tension” of the metal in water, p is the osmotic tension of metal ions, and nε is the charge that is carried by the ion. My main concern is, however, not the rediscovery of this illusive connection. The problem is that the model is essentially false, as was the reasoning of Nernst in 1889. Unfortunately, it has caused lots of troubles in the course of the history of 20th century electrochemistry. Because within the framework of this theory the concept of ideally polarized electrodes or the behavior of the redox electrodes—where no ions of the metal are in the solution—cannot be explained, Nernst and his followers invented new ideas for the explanation of experimental evidence concerning these systems that certainly hindered the development of a true physical model. For in-
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stance, because it was observed that there is a finite potential difference between an inert metal (e.g., Pt, Hg) and a solution not containing its ions (according to the Nernst model π should be infinite in this case), Nernst concluded that there are not ideally polarized electrodes; that is, a dissolution of the metal to some extent should occur. They also refused the concept of the redox electrodes, but assumed a nonexistent equilibrium (e.g., between Fe3+/Fe2+ and H +/1/2 H 2 couples). Problems also arose with electrodes of second order or with solutions containing complexing ligands. Despite the fact that the calculation using Nernst equation gave very small concentrations, e.g. 10–21 mol dm–3 for Ag+ in solution containing CN– ions, they still considered it as a physical reality and potential-determining quantity. It took decades until the new generation of electrochemists overcame this problem and established modern electrochemistry including electrode kinetics. Therefore, warming up this analogy is highly undesirable. It would not help the student with a better understanding; on the contrary we would mislead them and we would reproduce new cold-fusion cases where the extreme high pressure was assumed on the basis of an argument described in this paper. György Inzelt Department of Physical Chemistry Eötvös University Budapest 112, P. O. Box 32 H-1518, Hungary Bausch replies: It is beyond question that the kinetics of chemical processes at electrodes cannot be explained by using the concept of osmosis. I am grateful to the author of the Comment for stressing this again. But I am sure it is admissible and useful to compare the properties of osmotic and electrochemical equilibrium, and that is what I do in my paper. True, I have considered exchange currents and even allowed small deviations from equilibrium. However, the influence of activation barriers has been explicitly excluded so that the electrodic processes as far as being compared to osmotic processes remain reversible in the sense of thermodynamics. This is a serious limitation, and I readily admit that the parallel between osmotic and electrochemical potential does not exist for many important types of electrodes. But in fact all applications of (classical) thermodynamics to electrochemistry are governed by just this limitation. For example, the important relation ∆G = -nFE applies only if a reaction is proceeding in continuous electrochemical equilibrium. The condition is unrealistic, yet every student has to deal with this relation which so beautifully ties up thermodynamics and a part of electrochemistry. My (much more modest) aim is to tie up the concepts of osmotic and electrochemical potential under the conditions of continuous osmotic and electrochemical equilibrium. If these conditions are seen, new Cold Fusion cases will not follow from my paper. Rainer Bausch FB 19 der Universität Kassel D-34109 Kassel, Germany
Journal of Chemical Education • Vol. 73 No. 12 December 1996
Chemical Education Today
Root Mean Square In an otherwise excellent article, “Kinetic Theory of Gases” (J. Chem. Educ. 1995, 72, 715), a slight error was made. The root mean square speed of He is 1368 m/s at 27 °C, not 926 m/s as stated in the article.
v(rms) =
3RT = M
3 8.314 J/mol K 300 K 0.004 kg/mol
= 1368 m/s Consistent with my calculations, Chemistry, by Stanley R. Radel and Marjorie H. Navidi, West Publishing Company, 1990, page 193. Table 5.5 shows v(rms) for helium as 1360 at 25 °C. Ronald DeLorenzo Department of Chemistry Middle Georgia College 1100 Second Street, S. E. Cochran, GA 31014-1599
the mass spectrum has diagnostically useful fragments (m/e 105 for aldehydes, m/e 119 for 2-alkanones, etc.) and generally a useful M+ peak. In addition, the effect of two sulfurs on fragmentation in place of the carbonyl group may be discussed, as may the effect of chain branching on the mass spectrum. The effect on the MS of isotopes of sulfur, the effect of BF3 on carbonyl reactivity, etc., could also be discussed. Extensions could include ketones with cyclic or unsaturated substituents. As in the method of Rowland, the use of an effective fume cupboard is mandatory. The dithiol is quite malodorous and the use of minimal quantities is recommended. (For the same reason, neither organic nor aqueous waste should be discarded down the sink!) The BF3etherate also needs to be handled cautiously (plastic or rubber gloves, safety glasses, and other normal safety precautions would be adequate), but should not be beyond the capabilities of most students. (Alternatives to BF3-etherate could also be used.) T. G. Harvey CSIRO Division of Materials Science and Technology Rosebank MDC Private Bag 33 Clayton, VIC 3169, Australia Rowland replies:
Acyclic Saturated Ketones I noted with interest an experiment in the August 1995 issue in “the Microscale Laboratory” (“A Microscale GC-MS Experiment: Identification of an Acyclic Saturated Ketone” by A.T. Rowland) which used a deuterium exchange of the alpha protons to assist in the identification of the ketones in question. The variety of principles of mass spectrometry that are illustrated are excellent, but I thought that one aspect which could be usefully included would be another form of derivatization. Specifically, I refer to 1,3-dithiolanes, which may be prepared from acyclic ketones or aldehydes (see T. G. Harvey and T. W. Matheson, J. Chromatog. 1984, 298, 273–277). The preparation procedure is simple and may be conveniently carried out using Pasteur pipets and 2-mL glass vials. Dissolve 1 drop of the ketone in 5–10 drops of ether. Add 1 drop of 1,2-ethanedithiol (which should provide a useful excess) and then two drops of boron trifluoride etherate as a catalyst. (I originally allowed an hour for the reaction to go to completion, but I believe that it is all over in a matter of minutes—a fact that could be easily verified by GC as a classroom experiment.) Gently add a saturated solution of sodium bicarbonate dropwise until the fizzing stops, then extract with 1–2 mL of ether. The ether extract containing the dithiolane product may be separated with a Pasteur pipet and dried according to Rowland’s method, or chromatographed directly. This method has a number of advantages as a classroom preparation, and several useful principles may be introduced as a result. Firstly, it handily illustrates the use of derivatization in both GC and MS (and in synthetic organic chemistry for that matter). It is easily carried out on a small scale using a minimum of reagents. The preparation is simple, rapid, quantitative, and selective. The 1,3-dithiolane product is easily separated by GC-MS and
Tim Harvey’s letter presents an interesting alternate to the method I describe [J. Chem. Educ. 1995, 72, A160–A162] for the identification of acyclic saturated ketones. As I indicated in my paper, the deuterium exchange experiment, in addition to being a help in the identification of the ketone, provides a nice illustration of the acidity of hydrogens, a concept often discussed but rarely examined in the laboratory in the undergraduate level. The preparation of dithiolane derivatives of ketones and aldehydes [Harvey, T.G.; Matheson, T. W. J. Chromatog. 1984, 298, 273–277] represents another method of determining the number of carbons in the R groups substituted on the carbonyl carbon. Most of the examples cited by Harvey and Matheson are contiguous chain aldehydes and ketones in which the longer substituent is lost as R? to give the base peak. In two derivatives containing branching on an α carbon, the smaller, branched radical R? is lost in preference to a long normal chain. These observations do aid in the placement of the carbonyl group, as claimed by Harvey and Matheson. I favor the α hydrogen replacement I reported, however, in order to emphasize the special character of these hydrogens. Also, I avoid ethanedithiol and Oliver Stone movies at every opportunity. Students convey enough wild stories to their friends about the odor in the chemistry building without giving substance to their tales. But odor aside, the suggestion by Harvey is welcomed as another means of encouraging experiments that involve analysis of spectrometric data. Alex T. Rowland Gettysburg College Gettysburg, PA 17325
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Chemical Education Today
Letters Infrared Spectroscopy and Bond Strengths A recent article in this Journal (1) emphasizes the relationship between infrared frequencies, bond lengths, and bond strengths. Nevertheless, it contains two significant errors that can be used to illustrate the difficulties in using infrared data to get information about bond strengths. The first problem found in the interpretation of infrared spectra is band assignment. Indeed, the authors assign to S–S stretching vibrations one band at 1242 cm–1 in the infrared spectrum of K2S2O6 and one band at 1635 cm –1 that they found in the spectrum of (hydrated?) Na2S2O3 (1). However, both bands appear at a frequency too high to be assigned to a S–S stretching vibration, because for S=SF2 and gaseous S2 molecules, which contain a double S=S bond, υS=S appears at 718 cm –1 (2). In fact, for Na2S 2O3 υS–S has been reported at 446 cm, while for K2S2O6?2H2O one polarized Raman band at 293 cm–1 has been assigned to υS–S (the dithionate anion is centrosymmetric and S–S is not active in the infrared spectrum (2, 3). Actually, the band at 1242 cm–1 in the infrared spectrum of K2S2O6 should be assigned to an S–O stretching vibration (2, 3). On the other hand, the band at 1635 cm–1 that the authors assign to υS–S in Na2S2O3 (1) is at a frequency too high to be assigned even to S–O, because these bands appear at 1123 and 995 cm–1 in the vibrational spectra of Na2S2O3 (2). A likely explanation is that the authors isolated hydrated Na2S2O3, and that the band at 1635 cm–1 corresponds to a deformation of the water molecules. Once the bands have been properly assigned, another difficulty in correlating band positions with bond strength comes from coupling. Indeed, bond strengths are related to the force constants, but band positions may not be directly related to the force constants if the vibration of interest is strongly coupled to other vibrations. A striking example of this phenomenon is provided by the infrared spectra of ionic cyanate and thiocyanate salts. The principal resonance structures of the cyanate and thiocyanate ions are show in Figure 1. Bearing in mind that the C–N stretching frequencies are 2168 cm{1 for OCN– and 2049 cm–1 for SCN–, the authors concluded that the C–N bond is stronger, and hence the participation of resonance structures I is higher, for cyanate than for thiocyanate ions (1). Nevertheless, all the available data indicate that the opposite situation is found. Indeed, the NC force constant is higher in NCS– than in NCO– (4, 5). This fact suggest that the weight of resonance structure I is higher for thiocyanate than for cyanate, and in agreement with this idea the negative charge on the N atom is higher for OCN– than for SCN– (4,5). Recent calculations indicate that the atomic charges in NCO– are –0.64 on N and –0.38 on O, while in NCS– they are –0.46 on N and –0.54 on (6). Furthermore, the CN bond order is higher and the EC (E = O,S) bond order is lower in SCN– than in OCN – (4, 5). Therefore, it is clear that resonance structure I is more important in SCN– than in OCN–, and that the CN bond is stronger in SCN–. In agreement with this conclusion, while the C–N distance in ionic thiocyanate salts is close to 1.17 Å (7), the calculated C–N distance in NCO– is 1.19 Å from both ab
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initio studies (8) and matrix infrared spectra of condensed phase NCO– (9). In order to understand why υCN appears at a higher frequency in NCO– than in NCS –, although the CN bond is stronger in the latter ion, we must take into account the υCN and υCE (E=O,S) are not pure vibrations, but they are coupled to each other giving rise to an antisymmetric NCE stretching (mainly υCN) and a symmetric NCE stretching (mainly υCE) (4). Bearing in mind that υCO appears at a higher frequency than υCS, its frequency is closer to that corresponding to CN, so that coupling is more important in NCO– than in NCS–, thus raising the frequency of the antisymmetric NCO stretching vibration υ(CN). Indeed, recent calculations of the atomic displacements during the antisymmetric stretch show that while this mode is significantly delocalized over the molecule in NCO–, it is essentially a CN stretching mode in NCS– (6). The case of NCO – and NCS– provides a very good example, for teaching purposes, of the misconclusions that can be obtained in correlating vibrational frequency with bond strength. Indeed, even the widely accepted idea that shortening of a bond reflects strengthening has been recently challenged (10). According to the accepted rules for drawing Lewis structures, a structure is considered stable if negative formal charges are assigned to atoms with higher electronegativity (11). Therefore, it may be surprising that resonance structure I below has more weight for SCN–, while resonance structure II has more weight for OCN–, although the negative formal charge is not placed on the more electronegative atoms: E
C
E
N
I
C
N
II E = O, S
The higher weight of resonance structure II for OCN– than for SCN– may be related to the higher tendency of second row elements to use p orbitals in p bonding as compared to third and subsequent row elements (12). Literature Cited 1. Wiskamp, V.; Fichtner, W.; Kramb, V.; Nintschew, A.; Schnieder, J. S. J. Chem. Educ. 1995, 72, 952. 2. Weidlein, J; Muller, U.; Dehnicke, K. Schwingungsfrequenzen I; Georg Thieme: Stuttgart, 1981; pp 93, 94, 147, 148. 3. Palmer, W. G. J. Chem. Soc. 1961, 1552. 4. Norburn, A. H. Adv. Inorg. Chem. Radiochem. 1975, 17, 231. 5. Golub, A. M.; Kohler, H.; Skopenko, V. V. Chemistry of Pseudo-Halides; Elsevier: Amsterdam, 1986; Chapter 1. 6. Li, M. ; Owrutsky, J.; Sarisky, M; Culver, J. P.; Yodh, A; Hochstrasser, R. M. J. Chem. Phys. 1993, 98, 5499. 7. Wells, A. F. Structural Inorganic Chemistry, 5th ed.; Clarendon: Oxford, 1986; p 935. 8. Cai, Z. L. Chem. Phys. 1993, 170, 33. 9. Smith, D. F, Jr.; Overend, J.; Decius, J. C.; Gordon, D. J. J. Chem. Phys. 1973, 58, 1636. 10. Ernst, R. D.; Freeman, J. W.; Stahl, L; Wilson, D. R; Arif, A. M.; Nuber, B; Ziegler, M. L. J. Am. Chem. Soc. 1995, 117, 5075. 11. Ahmad, W. Y.; Omar, S. J. Chem. Educ. 1992, 69, 791. 12. Huheey, J. E.; Keiter, E. A.; Keiter, R. L. Inorganic Chemistry, 4th ed.; Harper Collins: New York, 1993; p 861.
David Tudela Departamento de Quimica Inorganica Universidad Autonoma de Madrid 28049–Madrid, Spain
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