Experimental K values for

Using the equation (3) the contributions of individual groups to retention Aj = lnk can be calculated from the experimental k values for some steroids. The calculated aj values are presented in Table 2. 

Table V. Deviations between Predicted and Experimental K Values for a Multicomponent System...

Figure 8. Comparisons between predicted and experimental K values for...

Figure 3.19 Comparison of predicted with experimental K-values for Roland crude oil at 200 F (from Katz and f irooza-badi, 1978),...

Figure 5. Predicted and experimental K-values for the methane-hydrogen sulfide system at 40°F...

This question was addressed by use of classical trajectory techniques with an ion-quadrupole plus anisotropic polarizability potential to determine the collision rate constant (k ). Over one million trajectories with initial conditions covering a range of translational temperature, neutral rotor state, and isotopic composition were calculated. The results for the thermally average 300 K values for are listed in the last column of Table 3 and indicate that reaction (11) for H2/H2, D2/D2, and HD /HD proceeds at essentially the classical collision rate, whereas the reported experimental rates for H2/D2 and D2/H2 reactions seem to be in error as they are significantly larger than k. This conclusion raises two questions (1) If the symmetry restrictions outlined in Table 2 apply, how are they essentially completely overcome at 300 K (2) Do conditions exist where the restriction would give rise to observable kinetic effects ... 

The logarithm of the n-octanol/water partition coefficient (log Kow is a useful preliminary indicator of the bioconcentration potential of a compound. The calculated log K values for 1,3-DNB and 1,3,5-TNB are 1.52 and 1.18 (Deneer et al. 1987), respectively, suggesting a low potential for bioaccumulation. An experimental bioconcentration factor (BCF) of 1,3-DNB for the guppy, Poecilia reticulata, was reported to be 74.13 (Deneer et al. 1987). This BCF indicates that bioaccumulation in aquatic organisms is not an important fate process. BCF data were not located for 1,3,5-TNB. 

The potential for 2,3-benzofuran to be bioconcentrated by aquatic organisms is likely to be moderate. A bioconcentration factor (BCE) is the ratio of the concentration of a chemical in the tissues of aquatic animals to the concentration of the chemical in the water in which they live. No experimentally measured value for the BCF of 2,3-benzofuran was located, but the octanol-water partition coefficient (K ) of 2,3-benzofuran has been measured as 468 (Leo et al. 1971). The empirical regressions of Neeley et al. (1974) relate the values of and BCF for other compounds, and can be used to estimate that the BCF of 2,3-benzofuran is approximately 40. If this estimate is correct, substantial bioconcentration of 2,3-benzofuran by aquatic organisms would not be expected. 

The very low value of Ashmore and Burnett is difficult to explain. It is easy to demonstrate that the discrepancy is not resolved by assuming the N03 intermediate in nitrogen dioxide decomposition is the pernitrite radical, in contradistinction to the symmetric nitrate radical. Their calculation of k5 depended on an experimentally obtained value for k 5 and an equilibrium constant K5- 5 calculated from thermodynamic properties for N03 measured by Schott and Davidson and Ray and Ogg. These results, obtained in a nitrogen pentoxide system, pertain to the nitrate radical, not the pernitrite radical. Guillory and Johnston176 reported an equilibrium constant based on estimated... 

Bare earth-HEDTA complexes react [456, 457] with an equivalent ambtmt of hydroxide ion to form a complex species [M(HEDTA)(OH)]-. The log k values for the formation of these mono-hydroxo rare earth-HEDTA complexes are presented in Table 36. The difference in log k values for light and heavy rare earths is about 1.5 log units. From La to Sm the increase is less than 0.24 0.09 log units. However, for the heavier rare earths the increase is less than the experimental error (Table 36). 

The Hammett treatment provides a correlation of much experimental data. Tables 26-6 and 26-7 contain 38 substituent constants and 16 reaction constants. This means that we can calculate relative k or K values for 608 individual reactions. To illustrate, let us suppose that we need to estimate the relative rates of Reaction 16 of Table 26-7 for the para-substituents R = OCH3 and R = CF3. According to the p value of 4.92 for this reaction and the cr values of p-OCH3 and p-CF3 in Table 26-6, we may write... 

Equation (359) with m = 0.5 was obtained empirically by M. G. Slin ko from experiments with a nickel catalyst. Starting from this result the general equation (359) was obtained theoretically for reaction (356) with exponent m not necessarily equal to 0.5, but of some value between 0 and 1, depending on the nature of the catalyst. In this form (359) was confirmed for all studied catalysts obtained values of m did not depend much on temperature. The theoretical K values (133) were employed in the calculations after they were checked experimentally. The values of m and absolute (i.e., calculated for unit area) k+ values for the same catalyst obtained in flow and circulation flow systems coincided within the accuracy of kinetic measurements. The table below gives approximated m values for some catalysts. 

FIGURE 5.10 Assessment of stability constant (log K) values of the complexes of crown ethers and their acyclic analogs with K+in methanol.25-77 Linear correlations between predicted and experimental log K values for the test set (a) of crown ethers and for the set of acyclic ligands (b) in Figure 5.6a. As QSPR models were built on the training set containing exclusively crown ethers, it is not surprising that predictions for the test set (a) are more reliable. 

Reactions of type (a) for Na, K, Rb, and Cs can attain equilibrium, and those constants were calculated. Calculations are in good agreement with experimentally determined values. No calculations could be carried out for type (b) reactions. Only for Mg(I03)2 have enthalpy increments above 298 K been measured, but in this case S°(298 K) is not known. There are no experimental studies. For reactions of type (c) only the calculation for Ca could be done since no S°(298 K) values for the periodates of Li, Ba, and Sr are available. 

Good agreement between the CNDO/S semiempirical HAB calculation and the experimental k j for the Ru/Ru-DNA duplex is found. Of course, this comparison requires use of Eq. (4) and a specified value of (0.9 eV) in addition to the measured driving force of 0.7 eV. Combining these data yields a calculated kB1 = 7.1 x 106 s 1 compared to the experimental k j = 1.6 x 106 s 1. Extensive use of the same ruthenium complexes as D/A groups in protein studies means that there is not much uncertainty in X (ca. 0.2 eV). 

In comparison, significant losses of a number of wine components must be expected when in contact with polyolefins because of the large K values for non-polar compounds. Packaging wine in polyolefin-coated containers, which for example is the case for bag-in-box packaging, does not appear to make sense. Quality decreases can occur not only by loss of the aroma but by the alteration of the aroma character due to different K values for different aroma compounds. The uptake (scalping) of non-polar compounds like limonene from fruit juices by polyolefins has been experimentally confirmed (Hirose et al., 1988 Mannheim et al., 1988). 

Several anomalies appear to exist in these data. First, the temperature dependence of rate seems excessive for the last group. Second, the diethyl ketone rates are much higher than those in the ethylene system and led Heller and Gordon to attribute greatly reduced collisional deactivation efficiency to the ketone relative to ethylene. If anything, however, we believe that the efficiency inequality should be reversed. Finally the experimental k values were calculated on the assumption that hydrogen or deuterium gas present in the mixture was completely inefficient as a collisional deactivator. We believe that this assumption is too extreme and that a reasonable lower estimate17 of their collisional efficiency would be 0.20. On this basis, all rate constants would be doubled or tripled. 

A very simple reaction is rotation about a bond. In the compounds in the table, different amounts of energy are needed to rotate about the bonds highlighted in black. See how this activation energy barrier affects the actual rate at which the bond rotates. Approximate values for k have been calculated from the experimentally determined values for the activation energies. The half-life, t1/2> is just the time needed for half of the compound to undergo the reaction. 

Least-squares plots of In [RHgX] vs. t by an IBM 1620 computer gave values for the pseudo-first-order rate constant k. Also, assuming a first-order rate dependence on ozone concentration, k was calculated via experimental saturation values for 03 in CHC13 at 0°C. The results of these calculations are shown in Table III. From Table III a few relative rate sequences can be formulated. These are listed below. 

This value is an estimate from Reference 41, where it was equated with the experimentally measured value for its isomer trispiro[2.0.2.0.2.0]nonane (31) at its melting point of 303 K. (Temperature corrections of the real liquid to a theoretical liquid at 298 K should be small. By comparison, using the enthalpy of sublimation of the solid at 298 K would contain little useful thermochemical information about the compound of interest). 

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