Hydrogen bond bond dissociation enthalpies

The enthalpy of dissociation of hydrogen bonds, A//nBomi, is a measure of their strength. Explain the trend seen in the data for the following pure substances, which were measured in the gas phase ... 

Table I includes the relative bond dissociation enthalpies obtained for some group 14 hydrides by photoacoustic calorimetry,7 10 The data demonstrate that, for the trialkyl-substituted series, the bond strengths decrease by 6.5 and 16.5 kcal/mol on going from silane to germane and to stannane, respectively. The silicon-hydrogen bonds can be dramatically weakened by successive substitution of the Me3Si group at the Si-H functionality. A substantial decrease in the bond strength is also observed by replacing alkyl with methylthio groups.

In Table 2.4, we have collected background information for discussion in the following chapters. Recommended C—H bond dissociation enthalpies of selected organic compounds are reported in the first two columns, followed by a variety of heteroatom-hydrogen bond strengths including N—H, O—H, S—H, Ge—H, and Sn—H bonds. 

To be effective as autoxidation inhibitors radical scavengers must react quickly with peroxyl or alkyl radicals and lead thereby to the formation of unreactive products. Phenols substituted with electron-donating substituents have relatively low O-H bond dissociation enthalpies (Table 3.1 even lower than arene-bound isopropyl groups [68]), and yield, on hydrogen abstraction, stable phenoxyl radicals which no longer sustain the radical chain reaction. The phenols should not be too electron-rich, however, because this could lead to excessive air-sensitivity of the phenol, i.e. to rapid oxidation of the phenol via SET to oxygen (see next section). Scheme 3.17 shows a selection of radical scavengers which have proved suitable for inhibition of autoxidation processes (and radical-mediated polymerization). 

In the chlorination of propane, the secondary hydrogen atom is abstracted more often because the secondary radical and the transition state leading to it are lower in energy than the primary radical and its transition state. Using the bond-dissociation enthalpies in Table 4-2 (page 143), we can calculate AH° for each of the possible... 

The energy differences between chlorination and bromination result from the difference in the bond-dissociation enthalpies of H—Cl (431 kJ) and H—Br (368 kJ). The HBr bond is weaker, and abstraction of a hydrogen atom by Br- is endothermic. This endothermic step explains why bromination is much slower than chlorination, but it still does not explain the enhanced selectivity observed with bromination. 

Peroxides are often added to free-radical reactions as initiators because the oxygen-oxygen bond cleaves homolytically rather easily. For example, the bond-dissociation enthalpy of the O —O bond in hydrogen peroxide (H —O —O —H) is only 213 kJ/mol (51 kcal/mol). Give a mechanism for the hydrogen peroxide-initiated reaction of cyclopentane with chlorine. The BDE for HO — Cl is 210 kJ/mol (50 kcal/mol). 

Stability of Allylic Radicals Why is it that (in the first propagation step) a bromine radical abstracts only an allylic hydrogen atom, and not one from another secondary site Abstraction of allylic hydrogens is preferred because the allylic free radical is resonance-stabilized. The bond-dissociation enthalpies required to generate several free radicals are compared below. Notice that the allyl radical (a primary free radical) is actually 13 kJ/mol (3 kcal/mol) more stable than the tertiary butyl radical. 

The bond dissociation enthalpy for normal hydrogen bonds is ca. 13... 42 kJ/mol (3. .. 10 kcal/mol) For comparison, covalent single bonds have dissociation enthalpies of 210... 420 kJ/mol (50... 100 kcal/mol). Thus, hydrogen bonds are approx, ten times weaker than covalent single bonds, but also approx, ten times stronger than the non-... 

The bond dissociation enthalpies A// of the carbon-hydrogen bond in a series of environments is as follows ... 

The use of isotopes has found surprisingly few applications in the determination of mechanisms of dehydrations. Below the temperature of appreciable water evolution, DjO/HjO exchange could, in principal be used to measure, in partially decomposed salt, e quantities of water held at surface sites of different reactivities, i.e. water physically adsorbed, chemisorbed, within amorphous phases and at the reaction interface. Similarly, kinetic isotope effects in dehydration reactions have not been extensively investigated. If corresponds to the enthalpy of dissociation, the substitution of a deuterium bond for a hydrogen bond could result in a change of the magnitude of E. ... 

It is well known that substituents have a profound effect on the hydrogen atom donating ability of phenols. Indeed, only those phenols bearing electron donating substituents, particularly at the ortho and/or para positions, are active as antioxidants. In general, this is as expected since such groups are expected to lower the phenolic O—H bond dissociation enthalpy and increase the reaction rates with peroxyl radicals. 

Table 9.4 shows typical values of bond dissociation enthalpies of some hydrogen bonds. The data in the table have been obtained from calculations on isolated species. These enthalpy values are therefore only approximate when applied to hydrogen bonds between molecules in a solid state lattice enthalpy values for these interactions cannot be measured directly. An example of how the strengths of hydrogen bonds can be obtained experimentally comes from the dissociation of a carboxylic acid dimer in the vapour state (equation 9.25). 

Chapter 11 deals with free radicals and their reactions. Fundamental structural concepts such as substituent effects on bond dissociation enthalpies (BDE) and radical stability are key to understanding the mechanisms of radical reactions. The patterns of stability and reactivity are illustrated by discussion of some of the absolute rate data that are available for free radical reactions. The reaction types that are discussed include halogenation and oxygenation, as well as addition reactions of hydrogen halides, carbon radicals, and thiols. Group transfer reactions, rearrangements, and fragmentations are also discussed. 

So far, our laboratories have primarily focused on the development, testing, and implementation of ccCA, but early success has been obtained in utilizing ccCA to solve chemical problems. For instance, a prototype version of ccCA was used by our group in collaboration with the experimental group of Professor T. Brent Gunnoe (then of North Carolina State University) to compute bond-dissociation enthalpies (BDEs) of ethylene, formaldehyde, methylene imine, carbodiimide, isocyanide, and their hydrogenated counterparts in order to probe useful correlations between the free BDEs of these model substrates and the proclivity of their insertion into the Ru(II)-phenyl bond of a catalyst-active species [108]. The BDEs of these small model systems provided a useful diagnostic to explain the thermochemical preferences of certain types of bond insertions. 

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