Equilibrium Calculations for Salts and Buffers

An Introduction to Aqueous Electrolyte Solutions. By Margaret Robson Wright 2007 John Wiley Sons Ltd ISBN 978-0-470-84293-5 (cloth) ISBN 978-0-470-84294-2 (paper) 

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The closeness of fit may be gauged from the experimental and theoretical rate vs. concentration curves for hydrolysis of p-nitrophenyl carboxylates catalysed by quaternary ammonium surfactant micelles (Figure 3). The shape of the curve is satisfactorily explained for unimolecular, bimolecular, and termolecular reactions. An alternative speculative model is effectively superseded by this work. Romsted s approach has been extended in a set of model calculations relating to salt and buffer effects on ion-binding, acid-dissociation equilibria, reactions of weakly basic nucleophiles, first-order reactions of ionic substrates in micelles, and second-order reactions of ionic nucleophiles with neutral substrates. In like manner the reaction between hydroxide ion and p-nitrophenyl acetate has been quantitatively analysed for unbuffered cetyltrimethylammonium bromide solutions. This permits the derivation of a mieellar rate constant km = 6-5 m s compared to the bulk rate constant of kaq =10.9m s . The equilibrium constant for ion-exchange at the surface of the micelle Xm(Br was estimated as 40 10. The... 

The pH of sea salt aerosol is an important property as many important aqueous phase reactions are pH dependent. For example, oxidation of S(IV) (SO2 + HSOs + SO ) by O3 is only important for pH of more than 6. Sea salt aerosol is buffered with HC03. Uptake of acids from the gas phase leads to acidification of the particles. According to the indirect sea salt aerosol pH determinations by Keene and Savoie (1998, 1999), the pH values for moderately polluted conditions at Bermuda were in the mid-3s to mid-4s. The equilibrium model calculations of Fridlind and Jacobson (2000) estimated marine aerosol pH values of 2-5 for remote conditions during ACE-1. Using a one-dimensional model of the MBL which includes gas phase and aqueous phase chemistry of sulfate and sea salt aerosol particles, von Glasow and Sander (2001) predicted that under the chosen initial conditions the pH of sea salt aerosol decreases from 6 near... 

The acidity or basicity of a solution is frequently an important factor in chemical reactions. The use of buffers of a given pH to maintain the solution pH at a desired level is very important. In addition, fundamental acid-base equihbria are important in understanding acid-base titrations and the effects of acids on chemical species and reactions, for example, the effects of complexation or precipitation. In Chapter 6, we described the fundamental concept of equilibrium constants. In this chapter, we consider in more detail various acid-base equilibrium calculations, including weak acids and bases, hydrolysis, of salts of weak acids and bases, buffers, polyprotic acids and their salts, and physiological buffers. Acid-base theories and the basic pH concept are reviewed first. 

A more comprehensive analysis of the influences on the ozone solubility was made by Sotelo et al., (1989). The Henry s Law constant H was measured in the presence of several salts, i. e. buffer solutions frequently used in ozonation experiments. Based on an ozone mass balance in a stirred tank reactor and employing the two film theory of gas absorption followed by an irreversible chemical reaction (Charpentier, 1981), equations for the Henry s Law constant as a function of temperature, pH and ionic strength, which agreed with the experimental values within 15 % were developed (Table 3-2). In this study, much care was taken to correctly analyse the ozone decomposition due to changes in the pH as well as to achieve the steady state experimental concentration at every temperature in the range considered (0°C [Pg.86]

If the buffered solution is dilute, this is its hydrogen-ion concentration. Because the activities of ions are affected by other ions, however, there is appreciable deviation from the calculated values in salt solutions as concentrated as 0,1 M, This fact accounts for the small discrepancies beuveen the pH values calculated from equilibrium con status and those gi en in the buffer tables. 

To calculate how the pH of a buffer solution changes when small amounts of a strong acid or base are added, we must first use stoichiometric principles to establish how much of one buffer component is consumed and how much of the other component is produced. Then the new concentrations of weak acid (or weak base) and its salt can be used to calculate the pH of the buffer solution. Essentially, this problem is solved in two steps. First, we assume that the neutralization reaction proceeds to completion and determine new stoichiometric concentrations. Then these new stoichiometric concentrations are substituted into the equilibrium constant expression and the expression is solved for [H30 ], which is converted to pH. This method is applied in Example 17-6 and illustrated in Figure 17-6. 

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