Apparent equilibrium constant ionic strength

With an aqueous fluid phase of high ionic strength, the problem of obtaining activity coefficients may be circumvented simply by using apparent equilibrium constants expressed in terms of concentrations. This procedure is recommended for hydro-metallurgical systems in which complexation reactions are important, e.g., in ammonia, chloride, or sulfate solutions. 

Figure 1.9 Plot of the base 10 logarithm of the apparent equilibrium constant for the hydrolysis of ATP to ADP and Pf at 298.15 K and 0.25 M ionic strength (see Problem 1.7).

In this chapter we have seen that acid dissociation constants are needed to calculate the dependence of apparent equilibrium constants on pH. In Chapter 3 we will discuss the calculation of the effects of ionic strength and temperature on acid dissociation constants. The database described later can be used to calculate pKs of reactants at 298.15 K at desired ionic strengths. Because of the importance of pKs of weak acids, Table 1.3 is provided here. More experimental measurements of acid dissociation constants and dissociation constants of complex ions with metal ions are needed because they are essential for the interpretation of experimental equilibrium constants and heats of reactions. A major database of acid dissociation constants and dissociation constants of metal ion complexes is provided by Martell, Smith, and Motekaitis (2001). 

The procedure for calculating standard formation properties of species at zero ionic strength from measurements of apparent equilibrium constants is discussed in the next chapter. The future of the thermodynamics of species in aqueous solutions depends largely on the use of enzyme-catalyzed reactions. The reason that more complicated ions in aqueous solutions were not included in the NBS Tables (1992) is that it is difficult to determine equilibrium constants in systems where a number of reactions occur simultaneously. Since many enzymes catalyze clean-cut reactions, they make it possible to determine apparent equilibrium constants and heats of reaction between very complicated organic reactants that could not have been studied classically. 

K and went on to calculate A,G ° and AfH ° at pH 7 and ionic strength 0.25 M for the corresponding reactants. This made it possible to calculate apparent equilibrium constants for six biochemical reactions at 283.15 and... 

If the apparent equilibrium constant K for an enzyme-catalyzed reaction has been determined at 298.15K and AfG ° values can be calculated at the experimental pH and ionic strength using known functions of pH and ionic strength for all the reactants but one, the AfG ° of that reactant under the experimental conditions can be calculated using equation 4.4-2. So far functions of pH and ionic strength that yield AfG ° are have been published for 131 reactants at 298.15 K (Alberty, 2001f). 

Table 4.11 Apparent Equilibrium Constants for Pyruvate Dehydrogenase, the Citric Acid Cycle, and Net Reactions at 298.15 K and 0.25 M Ionic Strength...

Source From R. A. Alberty, Arch. Biochem. Biophys. 389, 94 109 (2001). Copyright Academic Press. Note The apparent equilibrium constants for these reactions do not depend on ionic strength because the equilibrium constants for the chemical reference reactions and acid dissociation do not depend on ionic strength. 

Now the data entry for acetylcoA can be calculated from the apparent equilibrium constant (10.8) of this reaction at pH 7.12 and ionic strength 0.05 M. 

Since the hydrolysis of ATP evolves heat, Le Chatelier s principle says raising the temperature will cause the reaction to go less far to the right. But at 313 K the transformed Gibbs energy of reaction is more negative. To apply Le Chatelier we have to look at the apparent equilibrium constants. At pH 7 and ionic strength 0, we obtain... 

Note the ionic strength is not mentioned because the reference reactions and the dissociation of NH4 + do not depend on the ionic strength. Note that the apparent equilibrium constants of the first two reactions change by a factor of 10 when the pH is changed 0.10 in the neutral region. 

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