Equilibrium constant data tables

Appendix 1 includes a review of SI base units as well as tables of thermodynamic data and equilibrium constants. 

The thermal cracking of propane is practiced industrially for the primary purpose of making ethylene and propylene, but other reactions also occur. A scheme worked out by Sundaram Froment (Chem Eng Sci 32 601, 1977) consists of the nine reactions of the table. Equilibrium constants were deduced from thermodynamic data and the other constants by nonlinear regression from the extensive data on this topic in the literature and laboratory. 

The results calculated are in reasonable agreement with the seawater speciation data from Turner et al. (1981) quoted in Table 6.5 and with the qualitative features of the Garrels and Thompson model. Of course, the selection of different suites of stability constants leads to somewhat different speciation pictures. For example, the calculations made by Garrels and Thompson, Dickson et al., and in Tableau 6.6 are based on the assumption that chloride complexes with the major cations are unimportant. This assumption may be wrong and ion pairs with CP may represent nonnegligible fractions of the major cation concentration. Then, of course, a different speciation picture would result however, the extension of these results to trace metals (see the column for seawater in Table 6.5) would require a reinterpretation of the original experimental coordination data with equilibrium constants with the CP ion pairing model. 

It follows from equation [9.86a] that dependence of lnK s vs. 1/e will be rectilinear. Experimental data of equilibrium constants for process [9.85a] are in agreement with the conclusion. Approximation of data by equation [9.65] is presented in Table 9.3. 

The knowledge on the solubility of the crystalline substances formed and the identification of the ionic species are essential to the understanding of hydrothermal synthesis. As an example, the main equilibrium reactions for AlOOH dissolution in aqueous solution are shown in Table 1. For the estimation of metal oxides, the equilibrium constant K for each reaction is required. Equation (1) can be used for the estimation of the equilibrium constants under supercritical conditions. The data for equilibrium constants or Gibbs free energy changes and enthalpy of reaction for various reactions are available at ambient conditions. 

Table 2 presents the values of Kp calculated for several temperatures by using data on equilibrium constants in Eqs. (9)-(12) found in Ref. [18]. 

The physical data index summarizes the quantitative data given for specific compounds in the text, tables and figures in Volumes 1-7. It does not give any actual data but includes references both to the appropriate text page and to the original literature. The structural and spectroscopic methods covered include UV, IR, Raman, microwave, MS, PES, NMR, ORD, CD, X-ray, neutron and electron diffraction, together with such quantities as dipole moment, pX a, rate constant and activation energy, and equilibrium constant. 

The exceptions are formaldehyde, which is nearly completely hydrated in aqueous solution, and aldehydes and ketones with highly electronegative substituents, such as trichloroacetaldehyde and hexafluoroacetone. The data given in Table 8.1 illustrate that the equilibrium constant for hydration decreases with increasing alkyl substitution. 

The coefficients a, b, and c for hydrogenation were obtained from the literature [13] and those for nitrile and hydrogenated nitrile were calculated from a group contribution method reported by Rihani and Doraiswami [14]. All the necessary data are listed in Table 1. The integration constant / and AHq have been calculated by incorporating the values of AG° and AH° at 298 K in Eqs. (3) and (4). The equilibrium constant at atmospheric pressure and various temperature has been calculated according to the relationship ... 

If any equilibrium constants show this linearity, this behavior is most likely to be found among proton transfers of type (118) and type (120). The expressions for log K given in Table 11 show this linearity they represent, within the experimental error, the accurate data obtained by measurements on three proton transfers in aqueous solution. All three are of the type (120). 

The values of the equilibrium constant K listed in Table A are those obtained from data at low pressures, where the gases behave ideally. At higher pressures the mole percent of ammonia observed is generally larger than the calculated value. For example, at 400°C and 300 atm, the observed mole percent of NH3 is 47 the calculated value is only 41. 

The effect of temp on chemical equilibria is conventially determined via the free energy function AG°/RT and the ideal equilibrium constant K. Table 1 gives the free, energy function G°/RT for the important detonation products of CHNO expls. From these data A G°/(RT) can be obtained for different temps for the reactions of interest, and ideal equilibrium constants computed according to ... 

Some data for the equilibrium constant XXno and the rate constant k2 for the aniline 7V-nitrosation step (Scheme 3-27) are presented in Table 3-1. For comparison purposes the table also includes data on diazotization with N203. The rate constant k2 for the nitrosation step (Scheme 3-27) with NOC1 and NOBr is close to the limits given by the diffusion velocity of particles in a solvent. As postulated by... 

From these rate evaluations it is therefore possible to calculate the equilibrium constants Kx and K2 separately. Table 5-1 presents data for pK and pK2 for a series of diazonium ions taken from the paper of Jahelka et al. (1973 a), and pATm values published by Beranek et al. (1973) for the same diazonium ions. Instructive data from Machackova and Sterba (1972 a) and from Littler (1963) are also included in the table. 

Figure 3-3 depicts the analysis of the first data set in Table 3-2 according to Eq. (3-28). The plot displays the variation of [A]r versus time. The nonlinear least-squares fit gives k = 0.407 0.002 s 1. This value and the equilibrium constant give k- = 3.8 X 102 L mol-1 s 1. (The agreement is impressive because the data were simulated from values that are essentially these—see note a to Table 3-2.)... 

We are free to choose either K or Kc to report the equilibrium constant of a reaction. However, it is important to remember that calculations of an equilibrium constant from thermodynamic tables of data (standard Gibbs free energies of formation, for instance) and Eq. 8 give K, not Kc. In some cases, we need to know Kc after we have calculated K from thermodynamic data, and so we need to be able to convert between these two constants. 

Calculation of the second-order rate constant of carbonylation, kg, and the equilibrium constant, K = [t-C4H9CO+]/[t-C4H ][CO] = A c/fcD> requires knowledge of the concentration of CO. The constant a in Henry s law Pco = [CO] was determined to be 5-3 litre mole atm in HF—SbFs (equimolar) and 53 litre mole atm in FHSOs—SbFs (equimolar) at 20°C. From the ratio [t-C4HBCO+]/[t-C4HJ"] at a known CO pressure, values for k and K were obtained. The data are listed in Table 1, which includes the values for the rate and equilibrium constants of two other tertiary alkyl cations, namely the t-pentyl and the t-adamantyl ions (Hogeveen et al., 1970). 

These four equations are perfectly adequate for equilibrium calculations although they are nonsense with respect to mechanism. Table 7.2 has the data needed to calculate the four equilibrium constants at the standard state of 298.15 K and 1 bar. Table 7.1 has the necessary data to correct for temperature. The composition at equilibrium can be found using the reaction coordinate method or the method of false transients. The four chemical equations are not unique since various members of the set can be combined algebraically without reducing the dimensionality, M=4. Various equivalent sets can be derived, but none can even approximate a plausible mechanism since one of the starting materials, oxygen, has been assumed to be absent at equilibrium. Thermodynamics provides the destination but not the route. 

Construct a table of initial concentrations, changes in concentration, and equilibrium concentrations for each species that appears in the equilibrium constant expression. The equilibrium concentrations from the last row of the table are needed to find Kgq. Start by entering the data given in the problem. The initial concentration of benzoic acid is 0.125 M. Pure water contains no benzoate ions and a negligible concentration of hydronium ions. The problem also states the equilibrium concentration of hydronium ions, 0.0028 M. 

Organize the data using a concentration table. Initial concentrations can be found from the data stated in the problem, and because the equilibrium constant is quite large, we take the reaction to completion and then allow back-reaction to equilibrium ... 

Table 3.3. Equilibrium constants for the dissociation of H2, N2 and O2 and the partial pressures of the atoms at different temperatures calculated from fundamental data given in Tab. 3.4.

With a view to determining the equilibrium constant for the isomerisation, the rates of reduction of an equilibrium mixture of cis- and rra/i5-Co(NH3)4(OH2)N3 with Fe have been measured by Haim S . At Fe concentrations above 1.5 X 10 M the reaction with Fe is too rapid for equilibrium to be established between cis and trans isomers, and two rates are observed. For Fe concentrations below 1 X lO M, however, equilibrium between cis and trans forms is maintained and only one rate is observed. Detailed analysis of the rate data yields the individual rate coefficients for the reduction of the trans and cis isomers by Fe (24 l.mole sec and 0.355 l.mole .sec ) as well as the rate coefficient and equilibrium constant for the cw to trans isomerisation (1.42 x 10 sec and 0.22, respectively). All these results apply at perchlorate concentrations of 0.50 M and at 25 °C. Rate coefficients for the reduction of various azidoammine-cobalt(lll) complexes are collected in Table 12. Haim discusses the implications of these results on the basis that all these systems make use of azide bridges. The effect of substitution in Co(III) by a non-bridging ligand is remarkable in terms of reactivity towards Fe . The order of reactivity, trans-Co(NH3)4(OH2)N3 + > rra/is-Co(NH3)4(N3)2" > Co(NH3)sN3 +, is at va-... 

Table 1.2 gives a representative sampling of eqnilibrinm constants for additions to varions types of carbonyl compounds. Notice that there are nnmerons gaps in the table. This means that mnch remains to be done in the stndy of carbonyl addition reactions. In trying to devise schemes for predicting the eqnilibrinm constants for snch reactions, the scarcity of experimental data is a serions handicap. There are many fewer equilibrium constants for additions to imines, and even fewer cases where... 

The data of Table 6.11 create a wealth of information on the free energies of inorganic reactions. Although reported as emfs, these data are readily transformed into free energies by the expression AG° = —n i it. Such free-energy data are of considerable utility in determining equilibrium properties and, in particular, the equilibrium constant for the overall cell reaction. The possibility of the reduction of ferric ion to ferrous ion by zinc as a reduct-ant is considered as an example. The reaction in which one would be interested might be executed in the cell... 

The foregoing two examples have been taken to convey that the data of Table 6.11 can very well be used to determine the equilibrium constant for any reaction which is the overall reaction for a cell assembled with electrodes contained in the electrochemical series table. 

Here X denotes lb-moles of benzene per lb-mole of pure benzene feed and x, denotes lb-moles of diphenyl per lb-mole of pure benzene feed. The parameters k, and k2 are unknown reaction rate constants whereas K, and K2 are known equilibrium constants. The data consist of measurements of Xi and x2 in a flow reactor at eight values of the reciprocal space velocity t. The feed to the reactor was pure benzene. The experimental data are given in Table 6.2 (in Chapter 6). The governing ODEs can also be written as ... 

The following temperature-time-con version data was obtained for the batch experiments in the gas phase for the isomerization of reactant A to product B. The equilibrium constant for the reaction is large over the temperature range concerned in Table 6.17. 

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