Ionic equilibria, in aqueous solutions

Rossotti, H. (1978). The Study of Ionic Equilibria in Aqueous Solutions. Longmans, New York. A wealth of information on the equilibria of complex formation. 

In treating ionic equilibria in aqueous solution, two activity scales have proved especially useful. The first is the traditional infinite dilution activity scale, which is defined in such a way that the activity coefficient yA = A /[A] approaches unity as the solution approaches pure water. One might refer to this scale as the fresh water scale. 

All the reactions discussed in the previous section could be described as acid/base phenomena, defining acids and bases quite liberally. The importance of ionic equilibria in aqueous solution was recognised in the 1880s by Arrhenius, who proposed that acids were sources of H+(aq) while bases were sources of OH-(aq), and it was soon realised that this definition was closely related to the self-dissociation of water ... 

The general principles of chemical equilibrium apply equally to reactions of neutral molecules and to reactions of ions. Because of the special interest which chemists have in ionic equilibria in aqueous solutions and because of some common methods of treating some of these problems, this chapter and the following are devoted to the specific applications of chemical equilibrium to ionic reactions. As in Chapter 16, concentrations will always be expressed in mol/L, and [X] will represent the numerical value of the concentration of X. 

Most simple ionic equilibria in aqueous solutions tend to be very rapid, their rates often being controlled by diffusion. Further it is probably correct to assume that most equilibria in the dissolved phase are reached rapidly (Horne 1969). Rates of precipitation and even more of dissolution are usually slower (Westall and Stumm 1980). 

In Chapter 1 we discussed the basic principles of chemical equilibrium. We will now apply these principles to ionic equilibria in aqueous solutions. 

There are many t5q>es of equilibria that occur in solution, but for the important analytical conditions of ionic equilibria in aqueous solution, four examples are very important. 

About the same time Beutier and Renon (11) also proposed a similar model for the representation of the equilibria in aqueous solutions of weak electrolytes. The vapor was assumed to be an ideal gas and < >a was set equal to unity. Pitzer s method was used for the estimation of the activity coefficients, but, in contrast to Edwards et al. (j)), two ternary parameters in the activity coefficient expression were employed. These were obtained from data on the two-solute systems It was found that the equilibria in the systems NH3+ H2S+H20, NH3+C02+H20 and NH3+S02+H20 could be represented very well up to high concentrations of the ionic species. However, the model was unreliable at high concentrations of undissociated ammonia. Edwards et al. (1 2) have recently proposed a new expression for the representation of the activity coefficients in the NH3+H20 system, over the complete concentration range from pure water to pure NH3. it appears that this area will assume increasing importance and that one must be able to represent activity coefficients in the region of high concentrations of molecular species as well as in dilute solutions. Cruz and Renon (13) have proposed an expression which combines the equations for electrolytes with the non-random two-liquid (NRTL) model for non-electrolytes in order to represent the complete composition range. In a later publication, Cruz and Renon (J4J, this model was applied to the acetic acid-water system. 

Table 1 Kinetic and thermodynamic parameters determined for various tautomeric equilibria in aqueous solution at 25°C. The symbols for the rate constants k and the equilibrium constants K are explained in the text (first paragraph of section Examples ). Acidity constants are concentration quotients of ionization at ionic strength 7=0.1 m... 

A great many reactions are carried out in a convenient solvent for reactants and products. Dissolved reactants can be rapidly mixed, and the reaction process is easily handled. Water is a specially favored solvent because its polar structure allows a broad range of polar and ionic species to be dissolved. Water itself is partially ionized in solution, liberating and OH ions that can participate in reactions with the dissolved species. This leads to the important subject of acid-base equilibria in aqueous solutions (see Chapter 15), which is based on the equilibrium principles developed in this chapter. We limit the discussion in this subsection to cases in which the solvent does not participate in the reaction. 

FH MacDougall, LE Topel. Ionic equilibria in aqueous and mixed solvent solutions of silver acetate and silver monochloroacetate. J Phys Chem 56 ... 

Micellar effects upon chemical equilibria in aqueous solution were recognized many years ago, and Hartley [12] explained them in terms of the ability of ionic micelles to attract counterions and repel coions. This general explanation was subsequently applied to micellar effects upon chemical reactivity in aqueous solution... 

Perchloric acid is an extremely strong acid in aqueous solution (see Table 7.3). Although [ 104] (Fig. 17.12b) does form complexes with metal cations, the tendency to do so is less than for other common anions. Consequently, NaC104 solution is a standard medium for the investigation of ionic equilibria in aqueous systems, e.g. it is used as a supporting electrolyte in electrochemical experiments (see Box 8.2). Alkali metal perchlorates can be obtained by disproportionation of chlorates (eq. 17.76) under carefully controlled conditions traces of impurities can catalyse decomposition to chloride and O2. Perchlorate salts are potentially explosive and must be handled with particular care. For example, solid NH4CIO4 decomposes at 298 K according to eq. 17.79, and mixtures of ammonium perchlorate and aluminium are standard missile propellants. 

Table 1 Ionic Association Equilibria in Aqueous Solutions of Sulfates Studied by Raman Spectroscopy, A// kJ moP A5 J moF ... 

J. N. Butler—Ionic equilibria A mathematical approach. Addison-Wesley, Reading, MA.1964. (A tme masterpiece. In this book, I have no other ambition than applying Butler s principles of calculations or to analytical chemistry. In my opinion, this book was and is still, purely and simply, the key to understanding and handling equilibria in aqueous solutions without intuitive approximations.)... 

Edwards et al. (1975, 1978A, B) established a molecular-thermodynamic correlation for calculating vapor-liquid equilibria in aqueous solutions containing one or more volatile electrolytes, with special attention to the ternary systems, ammonia-carbon dioxide-water and ammonia-hydrogen sulfide-water. Their 1978 correlation was shown to give results in satisfactory agreement with the limited data then available for temperatures from 0° to 170°C (32° to 338° F) and ionic strengths of about 6 molal (equivalent to total concentrations of the electrolytes between 10 and 20 molal). 

With all the possible equilibria in aqueous systems, most species might be expected to participate in at least one equilibrium. Nonetheless, in many solutions some of the ionic species undergo no significant reactions. These species are classified as spectator ions. 

A more recent examination confirmed the existence of an unusual ionic-strength dependence of the reaction rate , which features a minimum at /r 0.83, but it was noted that the initial kinetics differ from those occurring later in the reaction. Consideration of the equilibria prevailing in aqueous solutions of Cr(VI) produced a simplified rate law... 

Generally, it has been found that the organic acids and bases do exist in aqueous solution as equilibrium mixtures of their respective neutral as well as ionic forms. Thus, these neutral and ionic forms may not have the same identical partition coefficients in a second solvent therefore, the quantity of a substance being extracted solely depends upon the position of the acid-base equilibrium and ultimately upon the pH of the resulting solution. Hence, extraction coefficient (E) may be defined as the ratio of the concentrations of the substance in all its forms in the two respective phases in the presence of equilibria and it can be expressed as follows ... 

The solubility of ionic solids in water covers a wide range of values. Knowing the concentration of ions in aqueous solution is important in medicine and in chemical analysis. In this section, you will continue to study equilibrium. You will examine the solubility equilibria of ionic compounds in water. 

IONIC SPECIATION. Ions interact continually in aqueous solution. Ions are complexed with water molecules. Even when we say that a certain ion is uncomplexed, the fact is that the ion is still complexed, in this case with water molecules. Association constants (also known in the literature as stability or formation constants) allow one to quantitate the extent to which an ion is complexed with any particular substance in solution. They also allow comparisons of the relative affinity of different complex-ing agents for a particular chemical substance. Speciation is a chain of linked binding functions (see Fig. 3). Such diagrams show the relative concentrations of the various complexes in solution, and the reversible equilibria existing between these pools are shown by the arrows. 

Carbonate-rock acid-rain reaction stiochiometry is best examined using a mathematical model of the ionic species present in a dilute carbonate solution. Appropriate models have been recently developed and are used with an electronic digital computer. Although a discussion of the application of mineral-reaction modeling is outside the scope of this paper, several recent publications explain its use in the study of carbonate-mineral reactions in aqueous solution (9-11). A reaction model, such as PHREEQE (11), the model used in this study, incorporates changes in gas-solution and solution-solid equilibria, as well as ionic equilibria in solution, in arriving at a final equilibrium-solution composition (. ... 

The solubility of carbon dioxide in aqueous and non-aqueous solutions depends on its partial pressure (via Henry s law), on temperature (according to its enthalpy of solution) and on acid-base reactions within the solution. In aqueous solutions, the equilibria forming HCO3 and CO3 depend on pH and ionic strength the presence of metal ions which form insoluble carbonates may also be a factor. Some speculation is made about reactions in nonaqueous solutions, and how thermodynamic data may be applied to reduction of CO2 to formic acid, formaldehyde, or methanol by heterogenous catalysis, photoreduction, or electrochemical reduction. 

In the condensed phase, departures from ideality are much more common. A significant departure from ideality results from the effect of the ionic strength of aqueous solutions on the energies of ions. This effect should be considered in any quantitative consideration of ionic equilibria in seawater (Stumm and Morgan, 1981). 

H2O serves as a base in 17-3) and as an acid in 17-4). Note that the bare of 17-1) becomes the hydrated proton or hydronium ion, HsO, of 17-3). In the formulation of equilibrium constants, [H ] and [H30 ] are always equivalent to each other the two forms are used interchangeably in most contexts and will be so used in all ionic equilibrium problems in this book. Although the proton is indeed hydrated in aqueous solutions, the notation H is often used instead of H30 because the reader understands the fact of hydration, because he need not worry about specifying the exact extent of hydration (which exceeds one H2O per proton), and because the specific use of the hydrated formula for the proton might obscure the important fact that all ions in water are extensively hydrated. Note also that the denominator in the Ka expression in 17-3) is identical with that in 17-1) since applications of these equilibria are intended for dilute solutions, H2O is always taken to be in its standard state and therefore need not be represented by a term in the expression. Equation 17-4) avoids describing aqueous ammonia as NH4OH, a species that probably does not exist at ambient temperatures. 

K.W. Herrmann, Non-ionic-cationic micellar properties of dimethyldodecylamine oxide, J. Am. Chem. Soc., 1962, 66, 2895-300 P. Holt and B. Tamami, Relation between pH and viscosity of some poly(alkylvinylpyridine N-oxides) in aqueous solution, Die Makromol. Chem., 1972, 155, 55-60 L. Chmurzynski, A. Liwo and P. Barczynski, A potentiometric study of acid-base equilibria of substituted TV-oxides in nitrobenzene, Anal. Chim. Acta, 1996, 335, 147. 

Hunt, J. P., Metal Ions in Aqueous Solution, Benjamin, 1963 (structures of water and ionic solutions, equilibria involving complex ions, rates and mechanisms and redox reactions). 

The status of Tc(lV) ions in aqueous solution, obtained by dissolution of rc02 aq in diluted HCIO4, was studied by electrophoresis. On the basis of ionic mobilities the existence of the species TcO- and TcO(OH)" at pH 1 and pH 2, respectively, was assumed according to the equilibria... 

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