Equilibrium constant recalculation

In the Hg -Se system, and at pH > 7, a significant solubilisation of HgSe(s) was recorded. These data cannot be questioned by the experimental uncertainty discussed above, and clearly indicate the formation of a soluble complex, (HgSe ). This species has therefore been accepted with an equilibrium constant recalculated by the review. 

The equilibrium constant recalculated from the shift values reported is in reasonable agreement with the fractionation factor for a simple model reaction of isotopic exchange equilibrium (38). The fractionation factor, Kf = 1.07 (25°C), for d,-cyclopropane relative to 2-dj-propane [39] (Shiner and Hartshorn, 1972) might serve as a model value for the maximum expected a-deuterium isotope effect. 

These mle-of-thumb generalizations must be treated cautiously. For precise work it is necessary to recalculate values of the equilibrium constants at the new temperature and then replot the diagram. The equilibrium constants have a general form (Chapter 2) ... 

For the equation (CgHg) K+ + CgDg = CgHg + (CgDg) K+, the ratio of isotopomers before and after the electron transfer can be recalculated into the corresponding equilibrium constants. Table 2.1 gives the order of these constants and experimental conditions typical for their determination. 

Making use of the data of Bodenstein and Lindner, as well as of some earlier investigators,50,363 Giauque and Kemp162 recalculated the thermodynamic functions and equilibrium constant for the reaction. Selected values from their results are shown in Table 4-2, along with... 

The concentration of any of these species depends on the total concentration of dissolved aluminum and on the pH, and this makes the system complex from the mathematical point of view and consequently, difficult to solve. To simplify the calculations, mass balances were applied only to a unique aluminum species (the total dissolved aluminum, TDA, instead of the several species considered) and to hydroxyl and protons. For each time step (of the differential equations-solving method), the different aluminum species and the resulting proton and hydroxyl concentration in each zone were recalculated using a pseudoequilibrium approach. To do this, the equilibrium equations (4.64)-(4.71), and the charge (4.72), the aluminum (4.73), and inorganic carbon (IC) balances (4.74) were considered in each zone (anodic, cathodic, and chemical), and a nonlinear iterative procedure (based on an optimization method) was applied to satisfy simultaneously all the equilibrium constants. In these equations (4.64)-(4.74), subindex z stands for the three zones in which the electrochemical reactor is divided (anodic, cathodic, and chemical). 

Apparent equilibrium constants A recalculation by JANAF using... 

A 3rd law analysis of the experimental equilibrium constants tabulated by Cotton and Jenkins (3 ) using current JANAF auxiliary data (4) leads to Dq = 204.8 kcal mol which is 2.6 kcal raol higher than the 202.2 kcal mol derived by Cotton and Jenkins (3 ). Applying this difference to the data of Ryabova and Gurvich ( ) and of Sugden and Schofield (2) as recalculated by... 

Equilibria for the reaction N20g(g) + N0(g) + N02(g) have been studied by Beattie and Bell ( ), Verhoek and Daniels (2), and Abel and Proisl (3). The results of Beattie and Bell are the most extensive, but they depend on the early analysis of Giauque and Kemp (4) for the simultaneous equilbrium N20 (g) 2 N02(g). The JANAF analysis of the tetroxide dissociation has been used to recalculate the data of Beattie and Bell. Non-ideality of N20 and NOg was allowed for by use of the equations of state of Giauque and Kemp, while non-ideality of NO and N20g was removed by extrapolation of the equilibrium constants to zero pressure in a manner similar to that of Beattie and Bell. The data of Verhoek and Daniels and of Abel and Proisl have not been recalculated. 2nd and 3rd... 

The data have been reconsidered by the review and the following observations were made, see also Appendix A. The solid selenium used in the experiments was obtained by reduction of a selenite solution by thiosulphate. The selenium might therefore not be in its standard state. The activity of the specimen is most likely elose enough to the standard state activity, however, since the precipitate was kept at boiling temperature for several hours. A recalculation of the side-reactions with more recent values of the auxiliary equilibrium constants made little difference to the result. The analytical data are not always consistent with the stoichiometry of Reaction (V.24) and the known initial composition of the test solution. The authors also observed this and held oxidation of iodide by initially present oxygen responsible for the discrepancies. Flowever, in some instanees the deviations from the expected concentrations are remarkably large. The deviations do not invalidate the results if equilibrium prevails, which was tested. 

The notation Se2Cl2(sln) represents Se2Cl2(l) saturated with Se(cr). The activity of Se2Cl2(sln) was estimated from determinations of the solubility of Se(cr). The temperature dependence of the equilibrium constant was recalculated by this review to be ... 

Aruga [78ARU] made calorimetric titrations of 0.17 M zinc nitrate solution with 0.17 M tetraethylammonium selenate, see Appendix A. The concentration of ZnSe04(aq) formed was calculated with the equilibrium constant determined in [34BAN] corrected to / = 0.5 M by Davies equation. After experimentally determined corrections for heats of dilution had been applied, the enthalpy change of Reaction (V.80) was calculated to be A //, ((V.80), / = 0.5 M, 298.15 K) = (0.20 0.05) kJ-moP. The recalculation of the equilibrium constant to / = 0.5 M is uncertain. The uncertainty assigned by the author to A,//° has therefore been increased to include the uncertainty in the value of the equilibrium constant. 

Bakeeva, Pashinkin, Bakeev, and Buketov [73BAK/PAS] measured the selenium dioxide pressure over gold selenite in the interval 489 to 599 K by the dew point method. The pressure was calculated from the dew point temperature by the relationship for the saturated vapour pressure in [69SON/NOV]. The data in the deposited VlNITl document (No. 4959-72) have been recalculated with the relationship selected by the review. The enthalpy and entropy changes obtained from the temperature variation of the equilibrium constant are A //° ((V.123), 544 K) = (576.8 13.0) kJ-mol and A,S° ((V.123), 544 K) = (899.4 + 24.0) J-K -mor. The uncertainties are entered here as twice the standard deviations from the least-squares calculation. 

The majority of equilibrium constants (including those given in Tables I-III) have been determined in solutions containing high (usually >1 M) and constant concentration of an inert electrolyte (e.g., alkali perchlorate). In this way the variation of the activity coefficients of the studied species (kept below 0.1 M) is so small that no correction factors have to be applied. The equilibrium constants, however, are strictly valid only in the ionic medium in which they have been determined. In order to avoid the burden of experimental determination of equilibrium constants in each ionic medium encountered, semi-empirical methods have been developed to recalculate the constants from one ionic medium to another. One such method is the specific interaction theory (SIT), developed by Guggenheim il55) and Scatchard (156,157) on the basis... 

Note that As(OH)3 is defined in terms of a different basis species (H2AsOJ), so that equilibrium constants cannot be directly compared between databases. They can of course be recalculated quite easily. In the eq3/eq6 database, we are also supplied with thermodynamic data about the species which is not directly used by the program, but which might be useful. 

The activities of Li2ZrCl6, Na2ZrCl6, and K2ZrCl0 in the systems MCl-MaZrClg were recalculated from available vapor pressure data (186). Previously reported calculations (342) for the sodium and potassium chloride systems are now known to be in error. The calculations indicate positive deviations from ideality that decrease with increasing size of the alkali metal cation. Equilibrium constants for the dissociation... 

Osthols et al. have also reanalysed the experimental data in [1987JOA/BIG2] they point out that the Th(IV) EDTA complex used to evaluate the equilibrium constant for the formation of Th(C03)5 is Th(EDTA)(OH) , not Th(EDTA)(aq), but it is not clear from [19940ST/BRU] how they arrived at the recalculated value of logj this review has therefore reanalysed the data in [1987JOA/BIG2] as described in the Appendix A entry for that paper. 

Check We can always check our answer by using it to recalculate the value of the equilibrium constant Kp =... 

A different approach to a systematic study of the effects of experimental errors has also appeared. A computer program was developed which enabled ready calculation of the equilibrium constant from input experimental data. Beginning with synthetic data (no errors) small errors were deliberately introduced into the input data—for example, amounting to a weighing error of 0.3 mg in 20-500 mg, or an instrumental error of +0.003 absorbance units. Recalculation of the constant revealed that for certain concentration situations the determined equilibrium constant could be extremely sensitive to small experimental errors. The same conclusion was reached when K was determined by a graphical method with the same data. This again emphasizes the need for careful planning of experimental conditions. 

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