two chemists measured heats of pro tonation (complexation) of various ox ygen bases, both in magic acid mix tures and in 30% sulfuric acid in acetic acid. In both systems the heats of protonation vary over only 2 kcal. per mole for oxygen bases as diverse as water, dioxane, and simple alcohols. Since the heats of protonation of the oxygen bases were measured with ref erence to a carbon tetrachloride solu tion, the heat-of-protonation values in clude the heats of transfer, (AH) t r a n s , for the hypothetical process of trans ferring the carbonium ion in solution from carbon tetrachloride at 25° C. to acid at - 6 0 ° C. Thus, the heat-ofprotonation results also suggest that heats of transfer are not significantly affected by variations in the structure of the solute. Proof that this is true for carbonium ion precursors was already in hand. Dr. Arnett and his coworkers had measured the partial molal heats of solution of groups of related carbonium-ion precursors into various organic solvents. They found that the heats of solution vary by much less than 1 kcal. per mole within each group of compounds. Since there is no signif icant difference within each group in the heats of solution for various sol vents, the heat of solution in super acid, ( Δ Η 8 ) _ 6 0 Η Α , should be unaf fected by structural variations for re lated carbonium ion precursors. Thus, variations in structure will have the same effect on the measured heat, (AH) _ 6 0 r e a c t , as on the heat of forma tion, ( Δ Η ) - 6 ° 0 + . Relating the measured heats to ΔΗ Ο1)Θ requires that the initial state of the precursor (pure compound at
Measured Heat of reaction
Partial molal heat of solution of RX in HA
Heat of complexation
Heat of formation of carbonium ion from complex precursor
Calculated from hypothetical transfer process from standard state (RX in CCU at 25° C.)
Found, by independent studies, to be unaffected by structure of R
HSAB rule predicts reaction direction
EXPERIMENTALIST. Dr. John W. Larsen of University of Pittsburgh attends to low-temperature calorimeter
- 6 0 ° C.) be related to the stand ard state (dilute carbon tetrachloride solution at 25° C ) . Dr. Arnett has shown that, if corrections for phase changes are made, the partial molal heats of solution in carbon tetrachlo ride, ( A H s ) c c l 4 , of typical carbonium ion precursors are insensitive to tem perature over a range of 50° C. Thus, AH obs is equivalent to the dif ference between (AH)- 6 0 r e a c t , which is measured, and ( A H s ) c c l 4 , which is measured at 25° C. and corrected for phase changes at 25° C. and —60° C. Conclusion: Variations in the reported heats of reaction, AH obs , within a series of related compounds reflect primarily the influence of struc ture on ( Δ Η ) _ 6 0 0 + , the enthalpy of carbonium-ion formation. There are no recognized standards for calorimetry in low-temperature su peracids, so Dr. Arnett and Dr. Larsen checked their procedure several ways. They measured the heat of reaction of mesitylene at —60° C. in both the supercooled liquid state and the solid state. The difference between the values was 3.0 ± 1 . 2 kcal. per mole, equal within experimental error to the heat of fusion of mesitylene. Another check makes use of the fact that both cyclohexene and methylcyclopentene go cleanly to the methylcyclopentyl cation in superacid at —60° C. Any difference in the heats of carbonium ion formation from these two precursors would be due to a dif ference in the heats of formation of the precursors. The observed differ ence in their heats of ion formation, 1.5 ± 0.9 kcal. per mole, agrees closely with the difference in their precursors' heats of formation, 1.4 kcal. per mole. This technique is a general method for determining the differences between ground-state en ergies of any two initial states which yield a common carbonium ion. The reaction need not be clean nor the products identified.
Dr. Ralph Pearson of Northwestern University has formulated a rule based on the principle of hard and soft acids and bases (HSAB) which correctly predicts the sign of a large number of heats of reaction [Chem. Commun., 2, 65 (1968)]. Qualita tively, Dr. Pearson's rule states that a reaction will proceed in the direc tion in which the hardest Lewis acid coordinates with the hardest Lewis base. Hardness and softness of Lewis acids and bases are qualitative prop erties, related to their polarizabilities, electronegativities, and comparative ease of oxidation and reduction. Dr. Pearson shows that, for a re action of the type AB + CD -> AD + CB (where A and C are the more metallic species), his rule predicts the correct sign for the heat of reaction in cases when Pauling's bond energy equation is not applicable. According to the Pauling equation, the heat of reaction is: ΔΗ
=
46
(XC-XA) ( X B - X D )
kcal.
where χ Α , χ Β , and the like are the electronegativities of the bonded atoms. The equation works best for simple reactions where electronega tivity differences are not very great. That is, only reactions involving changes in covalent bonds with small partial ionic character are covered by Pauling's equation. However, some scientists have tried to extend the equation to include reactions such as: LiF(g) + CsI(g)->LiI(g) + CsF( g ) , or CH 3 F(g) + CF 3 H(g) -> CH 4 (g) + C F 4 ( g ) . Trying to apply the bond energy equation to such reactions just won't work, Dr. Pauling says. In the first case, the ionic character of the bonds involved is too great; in the second case, strict application of the bond energy equation results in no heat of reaction at all. Dr. Pauling disclaims any responsibility for other people's misuse of the bond energy equation. Indeed, misuse of the bond energy equation predicts that the first reac tion is exothermic and the second reaction endothermic, Dr. Pearson points out. Calculation of the heats of reaction from Dr. R. C. Feber's compilation of heats of formation (available as Los Alamos Report No. LA-3164) reveals the reverse to be true—the first reaction is endothermic, consuming 17 kcal., while the second reaction is exothermic, releasing 19 kcal. In the first example, a back-ofthe-envelope electrostatic computation shows that the dominant energy term is the attraction between the small lithium and fluorine ions, Dr. Pearson says. FEB. 26, 1968 C&EN
37
The HSAB principle correctly pre dicts endothermicity in the lithium fluoride plus cesium iodide reaction. The hardest Lewis acid is Li+, and the hardest base is F~. Since they are already coordinated in the reactants, the reaction will be endothermic as written. In the reaction of methyl fluoride with trifluoromethane, the hardest acid is CF 3 + and the hardest base is F~. They are coordinated in the products, and the reaction as written will produce heat.
4th Edition
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RESEARCH IN BRIEF Partial Contents • 524 Pages, hard cover, 91 Charts & Tables • Specifications, Physical Constants, Chemical Properties, Physical Data on 100 gases • Recommended Controls and Regulators; Dimensional Drawings • Safety Procedures; Leak Detection • Cylinder Handling; Storage
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Lamia Circulators 38 C&EN FEB. 26, 1968
Nitrogen triiodide—easy to prepare from iodine and ammonia—is highly explosive when dry. It detonates at the slightest touch. This hazard has long been known and is mentioned in classic texts such as "Reference Book of Inorganic Chemistry" by Latimer and Hildebrand. Recent information coming to C&EN from Drew Univer sity student John Terebey indicates, however, that the hazard may not be fully appreciated. Since nitrogen tri iodide seems to enjoy considerable popularity among those looking for spectacular pseudochemistry demon strations, an example of its potential danger may serve as a useful warning. At an ACS student affiliate chapter in duction meeting, a demonstrator tried quadrupling a nitrogen triiodide prep aration recipe, making 20 grams and spreading it on paper whose edges were quite dry. He touched the mess and everyone present was lucky to escape uninjured from the resulting ex plosion. Haphazard demonstrations and indiscriminate scaling up of reac tion recipes should be avoided. New specifics on another old hazard—
mixtures of active metals and halogenated hydrocarbons—have been offered by H. D. Moshenrose and H. L. Tracy of North American Rockwell Corp., Downey, Calif. NARC con ducted impact-sensitivity tests on mix tures of finely divided barium and a number of halogenated hydrocarbons containing fluorine as well as chlorine. Among those compounds that deto nated in contact with granular barium were monofluorotrichloromethane, trichlorotrifluoroethane, carbon tetrachlo ride, trichloroethylene, and tetrachloroethylene. The secretary of the ACS Committee on Chemical Safety, How ard H. Fawcett, notes that halogen ated hydrocarbons are not chemically inert, and anyone using these materials would be wise to check their reactivity thoroughly. He cites the high-tem perature reaction of carbon tetrachlo ride to give phosgene as an example of halohydrocarbons being dangerous even in the absence of active metals.
