Enthalpy of interaction

The enthalpy of interaction of the grafted polymer with the filler surface is given by... 

Surface area around a molecule where optimum enthalpy of interactions of acceptor atoms with H-bond donor probe is realized [38, 39]... 

Surface area around a molecule where optimum enthalpy of interactions of donor atoms with H-bond acceptor probe is realized [38, 39[ Surface area around a molecule where optimum free energy of interactions of donor atoms with H-bond acceptor probe is realized [38, 39[ Sum of enthalpy values (kcalmof interactions between the acceptor atoms in a molecule and donor probe on OEASA [38, 39[ Sum of enthalpy values (kcalmof interactions between the donor atoms in a molecule and an acceptor probe on OEDSA [38, 39[... 

In Equation (10), Vc represents a hard-core repulsion that is entropic in nature since it is linearly dependent on temperature in the expression for energy. Repulsion is generally associated with enthalpic interactions and we can consider the effect of an enthalpic interaction. Since Vc is associated with a single Kuhn unit we consider the average enthalpy of interaction per pair-wise interaction and the number of pair-wise interactions per Kuhn unit,... 

In the case of the four-parameter equation, the enthalpies of interaction of a large number of acids and bases were determined calorimetrically in an inert solvent. With these values being known, a value of 1.00 was assigned for EA and CA for the Lewis acid iodine. The experimental enthalpies for the interaction of iodine with several molecular Lewis bases were fitted to the data to determine EB and CB values for the bases. Values were thus established for the four parameters for many acids and bases so that they can be used in Eq. (9.112) to calculate the enthalpies of the interactions. The agreement of the experimental and calculated enthalpies is excellent in most cases. However, the four-parameter approach is used primarily in conjunction with interactions between molecular species, although extensions of the approach to include interactions between charged species have been made. Table 9.7 gives the parameters for several acids and bases. 

Enthalpies of interaction of Ca2+ with several aldopentoses and aldohexoses have been determined calorimetrically (652,653,656-658). Complex formation is generally characterized by the balance between a fairly large favorable enthalpy term and an almost equal but unfavorable entropy term - for ribose both AH and TAS° are — 24 kJ mol-1 a value of +5 cm3 mol 1 for AV° can be understood in terms of electrostriction (658). 

The donicity represents the total enthalpy of interaction including the electrostatic constributions between D and SbCls in high dilution of 1,2-dichloro-ethane 3> 1 ... 

Drago and co-workers Introduced an empirical correlation to calculate the enthalpy of adduct formation of Lewis acids and bases ( 5). In 1971, he and his co-workers expanded the concept to a computer-fitted set of parameters that accurately correlated over 200 enthalpies of adduct formation ( ). These parameters were then used to predict over 1200 enthalpies of interaction. The parameters E and C are loosely Interpreted to relate to the degree of electrostatic and covalent nature of the Interaction between the acids and bases. This model was used to generalize the observations involved in the Pearson hard-soft acid-base model and render it more quantitatively accurate. 

Charge-transfer interactions have been observed between NbFj and MX4 (M = Si, Ge. or Sn. X = alkyl M = C, Si, or Sn, X = Cl) and several organic solvents, and for NbClj and TaClj with aromatic hydrocarbons and fluorocarbons. Tn these latter cases the enthalpies of interaction were estimated as 6.3 kJ mol The new complexes [MF L] (M = Nb or Ta, L = DMSO, EtCN, or CH CICN), [TaF, 2-Mepy)]. and [tap5(4-Mepy)2] have been prepared. Their vibrational spectra have been recorded and. together with those of [MF5(MeCN)] and [MFjLJ (M = Nb or Ta and L - DMSO or py), discussed in terms of MFjL and [MF L4][MFg] structures." ... 

We are at a loss to explain the discrepancy in the BF3 enthalpies of interaction with the sulfur donors. Steric effects may be operative, but this is far from the whole story for the BCI3 interaction is much larger than BF3 with these donors. Furthermore, using the tentative ( 113)3 parameters to estimate those of ( 2115)3 , we calculate an enthalpy from E and of 11.1 k.cal mole- for the BF3-P( 2H6)3 adduct compared to a measured value of 9.5 k.cal mole i. The authors report much difficulty with the sulfur donor system, but their error estimates could not possibly account for the difference between our calculated and the observed result. The behavior of ( 2115)35 compared to ( 2115)3 is clearly inconsistent with the behavior of these two donors toward ( 2H5)sAl where both enthalpies are correctly predicted with our parameters. It may be that the BF3-( 2115)25 system has an even lower equilibrium constant than reported and is completely dissociated over the temperature range studied. (This would require a very different entropy if the — AH predicted by E and were correct.) A slight impurity (reported to be less than 0.1%) or decomposition product could interact appreciably with BF3 and changing pressure contributions from this adduct with temperature could be attributed incorrectly to the sulfur donor adduct. The actual BF3-sulfur donor adduct would then be a very common example of an adduct which cannot be studied by the vapor pressure technique because it is completely dissociated at the temperatures at which one of the components has appreciable vapor pressure. We have examined the reaction of BF3 ( 2Hs) 2O with large excess of ( H2) 4S in dichloroethane solution at 25 ° and have found the equilibrium constant to be too low to be measured calorimetrically. 

The potential existence of these interrelationships has considerable significance both from a practical and theoretical standpoint. With them, it is possible to measure the O—H frequency shifts of any new hydroxyl compound with the three donors in Fig. 8 and locate these frequencies on the three straight constant base lines. The points can be connected to produce the constant-acid line whose slope and intercept can then be determined. The enthalpy of interaction of this acid with all the donors in the phenol correlation can then be predicted from the measured frequency shift for these donors and the constant acid line. By use of the E and C correlation, the internal consistency of the enthalpies predicted from the frequency shifts can be checked because the three enthalpies from the constant base line give three heats for the determination of the Ea and Ca parameters of the new add. 

The enthalpies of interaction of tetrahydro-furan (THF) and 2,5-dimethyltetrahydrofuran (l THF) with 0.03M benzene solutions of poly-(styryl)1ithium (PSLi) and poly(isoprenyl)lithium (PILi) have been measured as a function of R ([THF]/[Li]) at 25°C using high dilution solution calorimetry. At low R values (ca. 0.2) the enthalpy of interaction of THF with PILi (-5.8 kcal/mole) is more exothermic than with PSLi (-4.5 kcal/mole). However, the decrease in enthalpy for Me,THF versus THF is larger for PILi (3.2 kcal/mole) than for PSLi (2.2 kcal/mole) at low R values. The enthalpies decrease rapidly with increasing R values for PILi, but are relatively constant for PSLi. It is suggested that THF interacts with intact dimer for PILi, but for PSLi this base coordinates to form the unassociated, THF-solvated species. 

We have measured the enthalpies of interaction of tetra-hydrofurans with polymeric organolithiums to characterize the specific nature of base-alkyllithium interactions and these results are reported herein. 

Figure 1. Enthalpies of interaction of THF ( ) and 2 J-dimethyltetrahydrofuran (0) with 0.03M solutions of polystyryllithium in benzene at 25°C R - [base]/[Li] where [Li] is the concentration of carbon-bound lithium atoms in the solution.

With a ternary system of type biopolymer/ + biopolymey + solvent, in order to characterize all the different pair interactions, the following heat effects, Q, should be measured in flow mode (Semenova et al., 1991) (i) biopolymer, solution diluted by pure buffer, Q (ii) biopolymey solution diluted by pure buffer, Qp and (iii) mixed (biopolymer, + biopolymey) solution diluted by pure buffer, Qijh. The specific enthalpy of interaction between biopolymer, and biopolymey can then be obtained from... 

Barbour and Petersen (1974) observed strong H-bonding basicity of asphaltenes with phenol, having enthalpy of interactions between 6 and 8 kcal/mole. Phenol-base interactions for ir aromatic bases are typically as low as 0.5 kcal/mole but can reach 9 kcal/mole for nitrogen- or sulfur-containing molecules (Barton et al., 1972 Osawa et al., 1967 Whetsel, 1968). 

Table 6. Enthalpies of Interaction of Bases with Alkyllithiums at low R Values 80-821... 

Table 7. Comparison of the Initial Enthalpies of Interaction of Bidentate versus Monodentate Bases with n-Butyllithium8...

Initial enthalpies for addition of small amounts of bases to dilute solutions (0.2 M) of polymeric organolithiums at low R values ([base]/[Li]) provide direct information on the strength of the base interactions as well as the steric requirements of the bases. Data for initial enthalpies of interaction for a variety of bases with poly(styryl)lithium in benzene are listed in Table 8 88,89>. It is especially significant to note that the basicity order observed for poly(styryl)lithium (TMEDA > diglyme > THE > 2,5-Me2THF > dioxane > TMEDP > Et3P > EtzO = Et3N) is very similar to the order for simple alkyllithiums (see Tables 6 and 7) TMEDA > THF >... 

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