Acid-Base Equilibrium in Water

Acids and bases were defined and described by early chemists, including Boyle, Lavoisier, Davy, Berzelius, Liebig, and Arrhenius. At the present time, depending on objectives, one of two definitions of acids and bases is likely to be accepted. These two definitions, by Bronsted and Lowry and by Lewis, were proposed about the same time. According to the Bronsted definition acids are substances having a tendency to lose a proton, and bases are those having a tendency to accept a proton. Thus, for an acid HA the acid-base half-reaction is 

It is apparent from the above definition that a substance cannot act as an acid unless a base is present to accept the protons. Thus, acids will undergo complete or partial ionization in basic solvents such as water, liquid ammonia, or ethanol, depending on the basicity of the solvent and the strength of the acid. But in neutral or inert solvents, ionization is insignificant. However, ionization in the solvent is not a prerequisite for an acid-base reaction, as in the last example in the table, where picric acid reacts with aniline. 

Also in 1923, G. N. Lewis introduced the electronic theory of acids and bases. In the Lewis theory, an acid is a substance that can accept an electron pair and a base is a substance that can donate an electron pair. The latter frequently contains an oxygen or a nitrogen as the electron donor. Thus, nbnhydrogen-containing substances are included as acids. Examples of acid-base reactions in the Lewis theory are as follows  

The Lewis theory assumes a donation (sharing) of electrons from a base to an acid. 

In the second example, aluminum chloride is an acid and ether is a base. 

We see from the above that when An acid or base is dissolved in water, it will. dissociate, or ionize, the amount of ionization being dependent on the strength of the acid. A strong electrolyte is completely dissociated, while a weak electrolyte is partially dissociated. Table 7.2 lists some common electrolytes, some strong and some weak. Other weak acids and bases are listed in Appendix C. 

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Methods of studying acid-base equilibria in water... 

ACID-BASE EQUILIBRIA IN WATER OSTWALDS DILUTION LAW... 

For a complete description of acid-base equilibria in water, one must include the acid-base behavior of the solvent itself ... 

Monoprotic Acid—Base Equilibria in Water Solutions... 

The positions of acid-base equilibria in water are not only controlled by the relative acid-base strengths, but also by the pH. The pH of a solution is given by Eq. 5.11. It tells... 

Due to the biological significance of some azoles (pyrrole, indole, imidazole, benzimidazole) and the consequences of acid-base equilibria in their functions, a continuous interest in the behavior in water is to be expected. To quote a significant approach, imidazole is being used to determine the intra- and extracellular acidity by H-NMR (82MI4 86UP13). 

Among amphiprotic solvents of high permittivities, there are water-like neutral solvents (e.g. methanol and ethanol), more acidic protogenic solvents (e.g. formic acid), and more basic protophilic solvents (e.g. 2-aminoethanol). There are also amphiprotic mixed solvents, such as mixtures of water and alcohols and water and 1,4-dioxane. The acid-base equilibria in amphiprotic solvents of high permittivity can be treated by methods similar to those in aqueous solutions. If the solvent is expressed by SH, the acid HA or BH+ will dissociate as follows ... 

Sections 3.3.1 and 4.2.1 dealt with Bronsted acid/base equilibria in which the solvent itself is involved in the chemical reaction as either an acid or a base. This Section describes some examples of solvent effects on proton-transfer (PT) reactions in which the solvent does not intervene directly as a reaction partner. New interest in the investigation of such acid/base equilibria in non-aqueous solvents has been generated by the pioneering work of Barrow et al. [164]. He studied the acid/base reactions between carboxylic acids and amines in tetra- and trichloromethane. A more recent compilation of Bronsted acid/base equilibrium constants, determined in up to twelve dipolar aprotic solvents, demonstrates the appreciable solvent influence on acid ionization constants [264]. For example, the p.Ka value of benzoic acid varies from 4.2 in water, 11.0 in dimethyl sulfoxide, 12.3 in A,A-dimethylformamide, up to 20.7 in acetonitrile, that is by about 16 powers of ten [264]. 

In Chapter 8, we learned the Arrhenius definition of acids and bases—that an acid is a snbstance that can increase the concentration of ions in water and a base is a snbstance that can increase the concentration of OH ions in water. In Chapter 18, we learned about equilibrium systems. This chapter extends both of these concepts in discussing acid-base equilibria in aqueous solutions, which are extremely important to biological as well as chemical processes. 

The ionization constant should be a function of the intrinsic heterolytic ability (e.g., intrinsic acidity if the solute is an acid HX) and the ionizing power of the solvents, whereas the dissociation constant should be primarily determined by the dissociating power of the solvent. Therefore, Ka is expected to be under the control of e, the dielectric constant. As a consequence, ion pairs are not detectable in high-e solvents like water, which is why the terms ionization constant and dissociation constant are often used interchangeably. In 1ow-e solvents, however, dissociation constants are very small and ion pairs (and higher aggregates) become important species. For example, in ethylene chloride (e = 10.23), the dissociation constants of substituted phenyltrimethylanunonium perchlorate salts are of the order 10" . Overall dissociation constants, expressed as pXrx = log Xrx, for some substances in acetic acid (e = 6.19) are perchloric acid, 4.87 sulfuric acid, 7.24 sodium acetate, 6.68 sodium perchlorate, 5.48. Acid-base equilibria in acetic acid have been carefully studied because of the analytical importance of this solvent in titrimetiy. 

EQUILIBRIA IN A SINGLE ACID-BASE SYSTEM IN WATER ... 

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. 

The lower the value of this constant, the larger the deferences in acidity indices (pH) between the standard solutions of strong acids and bases, that results in a wider acid-base range for the solvent. This refers not only to the acid-base equilibria in aqueous solutions but also applies to any donor-acceptor interaction in molecular solvents which are prone to heterolytic dissociation with the formation of acidic and basic particles, as provided by an appropriate definition of acids and bases. It follows from equations (1.1.3) and (1.1.4) that the Arrhenius definition can only be used for the description of acid-base interactions in aqueous solutions, since the reaction between the acid of solvent and the base of solvent can result in the formation only of the solvent molecules. In the case considered, this solvent is water. 

The equilibrium constant, K, is defined to include the concentration of water, which does not change within the limits of error for weak acid-base equilibria in dilute solutions. Since the constant concentration of water molecules does not affect the point of equilibrium, water is omitted from the expression. 

Acid-based equilibria in the aquifer, such as pH buffering and speciation of TIC (total inorganic carbon), affect additionally the mobility of Fe and Mn with respect to saturation of rhodochrosite and siderite (Reaction (1C) and (2C)). Similar to the model STEADYSEDl (van Cappellen and Wang, 1995) TIC is considered as an excess component as indicated by the hydrogeochemical analyses. The contribution of TIC as reaction product from DOC degradation may be of minor importance. The infiltration water of the Oder shows only supersatrrration of calcite (but not of rhodochrosite and siderite) whereas the aquifer apart from the Oder is generally supersaturated with respect to rhodochrosite and siderite. 

This chapter is primarily about acid-base equilibria in anhydrous organic solvents. Equilibria in water-organic solvent mixtures, on which much effort has been expended for both practical and theoretical reasons, have been the subject of recent authoritative reviews. ... 

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