Equilibrium constant aqueous solution, reactions involving

Most reactions involve reactants and products that are dispersed in a solvent. If the amount of solvent is changed, either by diluting or concentrating the solution, the concentrations of ah reactants and products either decrease or increase. The effect of these changes in concentration is not as intuitively obvious as when the concentration of a single reactant or product is changed. As an example, let s consider how dilution affects the equilibrium position for the formation of the aqueous silver-amine complex (reaction 6.28). The equilibrium constant for this reaction is... 

The standard electrode potentials , or the standard chemical potentials /X , may be used to calculate the free energy decrease —AG and the equilibrium constant /T of a corrosion reaction (see Appendix 20.2). Any corrosion reaction in aqueous solution must involve oxidation of the metal and reduction of a species in solution (an electron acceptor) with consequent electron transfer between the two reactants. Thus the corrosion of zinc ( In +zzn = —0-76 V) in a reducing acid of pH = 4 (a = 10 ) may be represented by the reaction ... 

The kinetics of the reactions of phthalic and maleic anhydrides with Z-substituted phenols (Z = H, m-Me, p-Me, m-Cl, p-C 1, and -CN) (Scheme 8) were studied in aqueous solution at pH 8.5. Two kinetic processes well separated in time were observed. The fast process was attributed to the formation of the aryl ester in equilibrium with the anhydride and allowed the determination of the rate of nucleophilic attack of the phenol on the anhydride. From the slow kinetic process, the equilibrium constant for this reaction was determined. The Brpnsted-type plots for the nucleophilic attack of substituted phenols on the anhydrides were linear with slopes /SNuc of 0.45 and 0.56 for phthalic and maleic anhydride, respectively. The results are consistent with a mechanism involving rate-determining nucleophilic attack and also with a concerted mechanism.27... 

We have devised a method of measurement of 6-oxo-PGF which depends on the oximation of the compound in aqueous solution. This approach has the advantages of 1. the reduction of the polarity of 6-oxo-PGF, allowing its easier extraction and 2. a reliable oximation procedure with minimum manipulation of the sample. The process of oximation in aqueous solution seems anomalous since the formation of an oxime involves the elimination of water, however, the equilibrium constant for the reaction... 

In Chapter 7, we learned how to combine thermodynamic properties of different substances to calculate the heat absorbed or released by a reaction. In Chapter 13, we learned how to combine thermodynamic properties of different substances to calculate equilibrium constants for reactions. Many chemical reactions occur in aqueous solution and involve ions. To be able to calculate heats of reaction or equilibrium constants for such reactions, we need values for the thermodynamic properties of the aqueous ions involved. In this section, we discuss the standard thermodynamic properties of ions in solution, particularly with respect to how their values are established and interpreted. 

In a simple liquid-liquid extraction the solute is partitioned between two immiscible phases. In most cases one of the phases is aqueous, and the other phase is an organic solvent such as diethyl ether or chloroform. Because the phases are immiscible, they form two layers, with the denser phase on the bottom. The solute is initially present in one phase, but after extraction it is present in both phases. The efficiency of a liquid-liquid extraction is determined by the equilibrium constant for the solute s partitioning between the two phases. Extraction efficiency is also influenced by any secondary reactions involving the solute. Examples of secondary reactions include acid-base and complexation equilibria. 

Whilst this will be satisfactory when dealing with kinetic data in which reactions involving the solvent will not explicitly appear in the rate equations, it is not appropriate when we consider equilibrium constants. As an exercise, consider the formation of [Ni(en)3] from aqueous solutions of nickel(ii) chloride and en (en = H2NCH2CH2NH2) write the equations with the inclusion and the omission of the water molecules. Can you recognize the driving force for the formation of the chelate in each case ... 

Since the equilibrium constant of reaction 15.19 is 0.93, purple solid iodine is significantly more soluble in aqueous iodide solutions than in pure water. For dilute solutions, however, this can be ignored, as can a similar reaction involving chlorine and chloride in connection with reaction 15.7. Finally, we should remind ourselves that we have assumed that all the activity coefficients are unity (although here again the discrepancy so introduced can be ignored for simplicity). 

Dissolved carbon dioxide is different from species like S03 and NH3 in aqueous solutions in that the hydration reaction is slow enough (r, /2 = 15 seconds at pH 7 and 298 K) so that the rate constants involved can be determined and can be used to calculate the hydrolysis equilibrium constant (Edsall, 1969) at 298.15 K in terms of species for... 

The equilibrium constants involved in the reaction Fe3+ + 3 cat2" Fe(cat)33- were determined as follows. An aqueous solution of Fe3+ (5.5 X 10-3M) and catechol (1.48 X 10-2M), initially made basic with the addition of KOH, was titrated with 1.24M HC1 under an oxygen-free atmosphere at 22° and ionic strength (KC1) 0.16-0.22M (Figure 12). The acid dissociation constants for catechol were determined independently (under similar experimental conditions) to be pKai = 9.38 and... 

It is unfortunate that there has been so little work devoted to quantitative measurements of cation-pseudobase equilibria in methanol and ethanol since these media have several advantages over water for the determination of the relative susceptibilities of heterocyclic cations to pseudobase formation. The enhanced stability of the pseudobase relative to the cation in alcohols compared to water is discussed earlier this phenomenon will permit the quantitative measurement of pseudobase formation in methanol (and especially ethanol) for many heterocyclic cations for which the equilibrium lies too far in favor of the cation in aqueous solution to allow a direct measurement of the equilibrium constant. Furthermore, the deprotonation of hydroxide pseudobases (Section V,B) and the occurrence of subsequent irreversible reactions (Sections V,C and D), which complicate measurements for pKR+ > 14 in aqueous solutions, are not problems in alcohol solutions. Data are now available for the preparation of buffer solutions in methanol over a wide range of acidities.309-312 An appropriate basicity function scale will be required for more basic solutions. The series of -(substituted phenyl)pyridinium cations (163) studied by Kavalek et al.i2 should be suitable for use as indicators in at least some of the basic region. The Hm and Jm basicity functions313 should not be assumed90 to apply to methoxide ion addition to heterocyclic cations because of the differently charged species involved in the indicators used to construct these scales. 

Stopped flow spectrophotometry has also been used to measure the forward and reverse rate constants for the nitrosation of a number of alcohols and carbohydrates in aqueous solution (Aldred and Williams, 1980 Aldred et al., 1982). Both reactions were acid-catalysed and were also catalysed by added chloride and bromide ion. Structural effects were not marked but the sequence MeOH > EtOH > i-PrOH > t-BuOH was established for the forward reaction and was attributed to a steric effect, whilst for the reverse reaction the rate constants were not very dependent on the alcohol structure. The halide ion-catalysed reactions were believed to involve equilibrium concentrations of reactive NOCl and NOBr rate constants for the bimolecular process (39) were determined as 2.1x10 and 2.0 x 10 dm mol s for NOCl and NOBr respectively, for reaction with methanol at 25°C. Thus, as... 

Actually this equilibrium constant was measured for an analogous reaction involving a similar carbon acid (9-phenylfluorene). The value of KT /KD for (83) was calculated from Kr /KH with the Swain—Schaad relation. Using this analysis for the isotope exchange of triphenylmethane, a value of aT = 0.66 0.04 was obtained which is similar to the value deduced from other experiments for a closely related reaction [22]. However, the assumptions involved in the analysis may not be entirely valid and the derived values of a1 which measure the amount of internal return may not always be reliable. For the reaction of triphenylmethane the discrepancy between the measured value obs/ obs = 1-77 0.05 and the value 2.6 0.2 which is calculated for this isotope effect from the measured value of kobs/ obs = 1.34 0.03 using the Swain—Schaad relation is well outside experimental error. For the malononitriles discussed earlier in this section the experimental isotope effects for reaction in aqueous solution (Table 3) are well correlated by the Swain—Schaad relation which probably means that in these cases internal return is not important. 

Much of the work of analytical chemists involves reactions that take place at appreciable concentrations yet the equilibrium constants of fundamental importance are the thermodynamic values obtained by extrapolation to infinite dilution or to zero ionic strength. The purpose in this chapter is to examine some of the basic thermodynamic concepts that apply to solutions estimate the magnitude of the errors introduced by neglecting effects of ionic strength in aqueous solutions consider the extent to which these errors can be minimized by suitable corrections and examine the behavior of nonelectrolytes in solution. 

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