Autodissociation constant

Table 1 Hsts many of acetamide s important physical properties. Acetamide, CH2CONH2, dissolves easily ia water, exhibiting amphoteric behavior. It is slow to hydroly2e unless an acid or base is present. The autodissociation constant is about 3.2 x 10 at 94°C. It combines with acids, eg, HBr, HCl, HNO, to form soHd complexes. The chemistry of metal salts ia acetamide melts has been researched with a view to developing electroplating methods. The hterature of acetamide melts and complexes, their electrochemistry and spectroscopy, has been critically reviewed (9).

Table 6-2 gives some properties of common solvents. The pK,on is the autodissociation constant for the pure solvent, indicating that, among these acids, sulfuric acid dissociates much more readily than any of the others, and that acetonitrile is least likely to autodissociate. The boiling points are given to provide an estimate of the conditions under which each solvent might be used. 

Figure 3.2b shows the variation of ionic product or autodissociation constant of water, K, with temperature, where = [H ][OH Values of log initially increase with temperature from -13.9 at 25 °C up to a maximum value of -11.01 at 250 °C and then steadily decrease until the critical point, after which there is a substantial fall in to below -22.4. The increase in thermal energy upon heating results in two competing effects i) increased heterolytic scission of H2O into H and OH thereby increasing and ii) the reduction in dielectric constant as a result of the disruption of... 

Data taken from ref 28 and 29 (b) The variation in the autodissociation constant of water K with temperature. Data taken from ref. 28 and 31. 

In the low-voltage region of the j-V dependence (region 1, Fig. 2) the current flow is due to the ions present in solution, and the Fermi level of the electrode is expected to match the ionic states in solution (Fig. 5). The low magnitude of current (Figs. 6, 7) is explainable when the autodissociation constant is considered. The situation at the electrode-solvent interface is schematically shown in Fig. 8. The j-V dependence in this potential region is supposed... 

Despite the favorable properties of acetonitrile as a solvent, its use for equilibrium acidity measurements has its definite limitations. The pK range that is tolerable is limited at the high end by onset of solvent deprotonation, and at the low end by substrate autodissociation, as has been implicated for HCo(CO)4 [14a] and TpCr(CO)3H [22b]. These limitations can be overcome by the choice of a less polar solvent, e.g. 1,2-dichloroethane (DCE), dichloromethane, or THE. To make reliable, quantitative comparisons of thermodynamic data obtained in different solvents, it is necessary to link the acidity scales and electrode potential references in the different solvents. This has all too often proven to be a far from trivial task. Although, in principle, 1 1 relationship between the acidity scales in different solvents never exists, pK differences between closely related compounds are often almost constant when compared in different solvents. This is because their solvation properties are similar, because of similarities in size and charge distribution. In less... 

In addition to acid and base catalysis of amide hydrolysis, there is evidence that neutral water can also react with amides. A detailed investigation of the hydrolysis of formamide in aqueous solution was reported by Slebocka-Tilk et al. The value of the observed rate constant (Kbs) was found to depend on temperature because of the variation of the autodissociation equilibrium constant of water, Ky, with temperature. Log fcobs for the reaction at 56°C was a minimum at pH 6.1 and then increased in either more acidic or more basic environments. By measuring the rate constants for the acid- and base-catalyzed reactions, the investigators determined the rate constant for reaction of formamide with neutral water. Thus, the overall observed rate constant for the reaction at 56°C was... 

The self-dissociation of water and the proton/hydroxide mobilities in water Water is amphoteric at 298.15 K and 1 atm, the ionization constant (pKw,) of water is 14.004. The proton and hydroxide concentrations are so small that the water activity is almost unity. The standard enthalpy of self-dissociation is 55.81 kj mol the heat capacity -215 J and the standard volume of self-dissociation is approximately-20 cm mol T Ultrafast mid-infrared spectroscopic measurements have suggested that the first step of the autodissociation of water proceeds through an excited vibrational state of the OH bond, probably with v = 2. 

As indicated above, OH is much stronger base than water itself, and hence, there are a lot more H O than H3O and Off- in pure water, as these two, i.e., H+ (of H3O ) and OH", tend to bind to form H O back. The extent of autodissociation is governed by the so-called mass action law [H30][OH"]=K, and is called an eqnihbrium constant and known to be a very small number, 10" at room temperature. As you recall, [H3O ] represents the molar concentration (mol/L) of the chemical species H3O. More rigorously speaking, [A] in an equilibrium constant expression represents activity which is a number and its magnitude related to the molar concentration. This relationship says that the product of the concentrations of hydronium ion and hydroxide ion in water is constant. So, if you add an acid (which gives off H ) to water, you have increased the hydronium ion concentration, and accordingly the hydroxide ion in the solution will be reduced. 

It is assumed that the concentration of ions present in a liquid is extremely low and that the ions present are formed exclusively in the autodissociation process. All ions and neutrals other than those originating from liquid solvent are removed by purification as impurities. Under such circumstances, the system can be considered as an ideally dilute solution i.e., the solvent mole fractions is 1. Hence, the system, the electrode and the solvent, is assumed to obey Henry s law. The surface concentration of specifically adsorbed anions can be estimated on this basis from the Henry isotherm. Assuming that the specific adsorption equilibrium constant for OH ions in pure water can range from 0.1 to 100 dm /mol, one can obtain the surface concentration of adsorbed anions in the range lO -lO mol/cm i.e., the ratio of the adsorbed anions to the metal atoms of the electrode surface is 10 -10 . Having this in mind and remembering that an amount of possible solvated cations in the bulk of solution is very low, it can hardly be believed that the Helmholtz compact layer is formed in pure liquid. Thus, the electrode-liquid interface seems to... 

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