Chemical Education Today
Letters Capsaicin Hazard I would like to call readers’ attention to a potential safety issue related to the handling of the compound capsaicin and related capsainoids. My concerns center on the paper by Batchelor and Jones in the February 2000 issue of JCE, p 266, entitled “Determination of the Scoville Heat Value for Hot Sauces and Chilies: An HPLC Experiment”. The paper sets forth an analytical HPLC experiment that uses analytical standard solutions prepared from pure capsaicin. Students handle solutions of the standards prepared at the 1000 ppm level (however, the paper makes no stipulation about whether the students are to prepare the standard solutions or whether they are provided). The Material Safety Data Sheet for capsaicin characterizes it as “highly toxic, toxic by inhalation, in contact with skin and if swallowed”. Another caution is that the compound “causes severe irritation”. The Merck Index lists the compound as “a powerful irritant; initial administration causes intense pain”. These hazards are not overstated. No safety precautions were included in the article about handling this material. I personally would not allow undergraduate students to handle this compound and would ensure that any lab coordinators or graduate students were thoroughly briefed about the relevant safety issues when handling the pure compound. Pure capsaicin can be dangerous if not handled safely and if the authors have not encountered such problems, then it is testimony to their strict adherence to safety procedures and excellent lab technique. However, an accident involving pure capsaicin and students is not worth the risk.
The author replies: Certainly capsaicin in powder form, or in very concentrated solutions, should be handled with extreme care. The preparation of a 1000-ppm stock solution is performed by the instructor, or an experienced teaching assistant, prior to the laboratory period as stated in the manuscript. Once prepared, the solution is similar in strength to the undiluted hot sauces, which most students are comfortable handling, but proper safety procedures should be strictly followed. Raw chili peppers, if employed, should be handled with gloves (as is normal practice in the kitchen). High concentrations of capsaicin can be avoided altogether if one uses the overthe-counter arthritis creme for standard solution preparation, and only hot sauces as samples. This option is strongly recommended for instructors who are not comfortable handling capsaicin in its powdered form. We would never suggest that the students handle the powder. The supplemental material and the laboratory instructions, available on JCE Online, clearly state these cautions in boldface. We regret that the cautions were not included in the abbreviated manuscript. We assumed that the prelab preparation of the solution would not involve the students.
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Paul E. Vorndam
Bradley T. Jones
Department of Chemistry University of Southern Colorado Pueblo, CO 81001-4901
Department of Chemistry Wake Forest University Winston-Salem, NC 27109
Editor’s Note: Hazards We thank Paul Vorndam for calling attention to the hazards associated with capsaicin. HPLC separation of capsaicin is a popular experiment (see J. Chem. Educ. 1999, 76, 240, as well as the paper by Batchelor and Jones). Jones is correct that the Lab Documentation published in JCE Online makes it clear that students do not handle pure capsaicin and provides warnings in boldface type. Because more information is provided in JCE Online than is provided in print, anyone who plans to use an experiment described in a Lab Summary in these pages must
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also consult the supplemental material in JCE Online before developing materials for students or lab technicians. Not to do so will result in an incomplete picture of the experiment. In order to provide JCE readers with more complete information in print, we will begin requiring a section headed “Hazards” in the Lab Summary of all manuscripts received after January 1, 2000. This will call attention to more detailed safety information in the Lab Documentation that appears on the Web.
Journal of Chemical Education • Vol. 77 No. 4 April 2000 • JChemEd.chem.wisc.edu
Chemical Education Today
Coulombic Models in Chemical Bonding The July 1998 issue of this Journal contains two papers, by Gillespie (1) and Smith (2), on topics closely related to those treated previously in this Journal using “Coulombic Models in Chemical Bonding” (3, 4 ). The substance of the article by Gillespie, that the volatility of certain “ionic molecules” is a function of their structure rather than the type of bonds, was treated in greater generality (3), where a simple yet reliable criterion for molecule formation was proposed: once the charge ratio has determined the formula of the compound, the size ratio of the anions to cation determines the structure. A molecule is a neutral assemblage with a negatively charged surface (no exposed cations). As reiterated by Gillespie, and even acknowledged by Pauling (5), electronegativity considerations fail to predict—or even explain—the abrupt changes in structure type that occur with increasing charge and decreasing size of cations going across the periods. Gillespie’s discussion of atomic charges contributes to verifying that gaseous substances such as BF3 and SiF4 should be considered as “ionic”; but in doing so, the author overlooks the opportunity to include by the same description other molecular compounds now described as “covalent”. In ref 3, hydrides were selected to show trends across period II. Gillespie has verified that the further statement, “Fluorides of period II can be considered in a similar manner with similar results”, is indeed correct, even to the inclusion of analogous interionic distances (see similar tables in both papers). Formulating hydrogenic assemblages exclusively as hydride ion has other significant consequences. It helps clarify acid– base chemistry, since only the relative base strengths determine the extent to which the proton can be extracted from the lone pair of the first base. It denies any significant role of “bond polarity” in determining dipole moments of binary hydrides (6 ) and provides quantitative bond energy results (3), including a calculation of the barrier to internal rotation in ethane (4 ). The qualitative explanation of this phenomenon by Smith (2), in terms of the higher energy of antibonding orbitals, contrasts sharply with that in ref 4 and with many MO calculations, where calculations show good agreement with experiment. Literature Cited 1. 2. 3. 4. 5.
Gillespie, R. J. J. Chem. Educ. 1998, 75, 923–925. Smith, D. W. J. Chem. Educ. 1998, 75, 907–909. Sacks, L. J. J. Chem. Educ. 1986, 63, 288. Sacks, L. J. J. Chem. Educ. 1986, 63, 487. Pauling, L.; Pauling, P. Chemistry; Freeman: San Francisco, 1975; p 275; but see also pp 590–595 for the more commonly found explanation based on electronegativities. 6. Sacks, L. J. J. Chem. Educ. 1986, 63, 373–376.
The author replies: Gillespie and I, in this Journal and elsewhere, have consistently advocated simple models of structure and bonding based on electrostatic/Coulombic rather than molecular orbital considerations. Gillespie et al. (1) argue that structural chemistry at first-year university level can be taught without orbitals at all. I have made considerable use of the ionic model in explaining the relative stabilities of oxidation states (2–4) and shown how periodic variations in the atomization enthalpies of metallic elemental substances can be understood semiquantitatively by an extension of the ionic model (5). I have, incidentally, also addressed the question of why BF3 as an ionic solid is unstable with respect to the molecular form (6 ). I have championed the continued use of the electrostatic crystal field model, while recognizing that for some purposes an MO approach is required (7–9). It is well known that Coulombic models can give remarkably good agreement with experimental thermochemical data. For example, a simple Coulombic description of HF in terms of H+ F ᎑ (which is not how Sacks would view it) gives a bond energy of about 530 kJ mol-1, compared with the experimental value of 569 kJ mol᎑1, better agreement than can be obtained from low-level MO calculations. But if MO theory is to be taught at all, students need to be convinced that it can explain— albeit at a qualitative level—such things as molecular shapes, ligand field splittings and rotational energy barriers, even though electrostatic/Coulombic models might be more appropriate at the elementary level and indeed more amenable to quantitative calculations in some cases. You don’t need a Patriot missile to shoot down a Cessna; but unless the Cessna is demonstrably vulnerable to the Patriot, the credibility of the latter against a ballistic missile is undermined. Literature Cited 1. Gillespie, R. J.; Spencer, J. N.; Moog, R. S. J. Chem. Educ. 1996, 73, 622–627. 2. Smith, D. W. J. Chem. Educ. 1986, 63, 228–231. 3. Smith, D. W. Inorganic Substances: a Prelude to the Study of Descriptive Inorganic Chemistry; Cambridge University Press: Cambridge, 1990. 4. Smith, D. W. J. Chem. Educ. 1996, 73, 1099–1102. 5. Smith, D. W. J. Chem. Educ. 1993, 70, 368–372. 6. Smith, D. W. Inorganic Substances: a Prelude to the Study of Descriptive Inorganic Chemistry; Op. cit.; pp 156–157. 7. Smith, D. W. Inorganic Substances: a Prelude to the Study of Descriptive Inorganic Chemistry; Op. cit.; pp 16–22. 8. Smith, D. W. Encyclopaedia of Inorganic Chemistry; Wiley: Chichester, England, 1994; pp 1065–1082. 9. Smith, D. W. J. Chem. Educ. 1996, 73, 504–507. Derek W. Smith
Lawrence J. Sacks Department of Biology, Chemistry and Environmental Science Christopher Newport University Newport News, VA 23606
Department of Chemistry University of Waikato Private Bag 3105 Hamilton, New Zealand
JChemEd.chem.wisc.edu • Vol. 77 No. 4 April 2000 • Journal of Chemical Education
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