Acid dissociation constant calculating from

The second group of values came from studies where it was assumed that polymerization reactions occurred, such as the formation of H5As206 (aq>, in addition to the deprotonation reaction. For chemical and mathematical reasons, the dissociation constant calculated from a set of measurements becomes smaller as one introduces polymeric anions into the model. The differences of the models chosen, at first appearance, could serve to explain the differences of the equilibrium constants given in the previous table. Unfortunately, the situation, from the perspective of data evaluation, is more complex. In principle, there should be a sufficient dilution of arsenious acid for which one would not expect the formation of a significant proportion of species like HsAsaOe caq) upon addition of base. For such a condition, the equilibrium constant determined assuming that only the monomer exists, should approach that determined for the multi-species model. Britton and Jackson (1934) performed potentiometric titration at two concentrations of arsenious acid (0.0170 and 0.0914 molar) and obtained essentially the same... 

The line-broadening data as a function of pH, typically shown for the W(IV) in Figs. 13 and 14, incorporating the known pKa values (Table II), were fitted in 5 X 5 Kubo-Sack matrices describing the exchange based on the above schemes (6, 57). The experimentally determined chemical shift and linewidth data in the absence of exchange for the aqua oxo, hydroxo oxo, and dioxo species and the pH-dependent species distribution as calculated from the acid dissociation constants for the four systems were all introduced in the different matrices and the spectra were computer simulated. For each set of chosen rate con-... 

The value of / has been determined by NMR measurements on the mixed solvent as 0.69+0.02 and is not to be interpreted as an independent fitting parameter. With / established, can be calculated from the acid dissociation constants in pure H20 and D20, and Equation 11.68 interpreted as a zero parameter theoretical prediction of... 

Fractionation factor acid-dissociation constants in H2O and D2O by Kreevoy (Kreevoy et al., 1977 Kresge and Chiang, 1973). This value is not as low as that for the many other systems which are known to have weaker hydrogen bonds. 

Similar expressions have been obtained for the particular cases of mono-protic acids and bases, diprotic acids and bases, and zwitterions (207, 208), and in each case the data conformed well to Eq. (111). It has also been shown (207) that the acid dissociation constants can be determined by using reversed phase chromatography. The pIK, values of 10 aromatic acids calculated from chromatographic data by employing Eq. (91) were... 

See Eq. (4). Data from Roecker et al. (7), Gilbert et el. (9), and Pipes and Meyer (19). b Standard reduction potentials for M(III)OH2+ + e —> M(II)OH calculated from the standard reduction potentials for the aquapentakis(imine) species and the acid dissociation constants. 

Figure 2.3 Speciation of arsenious acid with pH. Curves calculated from acid dissociation constants listed in Table 2.10.

The acid dissociation constants can be calculated from In P by fitting the plot of P versus [H+] with a power series in [H+],... 

Since the acid dissociation constants are known, the value of Krcf can be calculated from the value of K at a pH in the neutral region in the absence of metal ions by using equation 1.4-3. Values of Kief at zero ionic strength are given in Table 1.2 for six reference reactions. 

Some aqueous models accept only total inorganic carbon rather than titration alkalinity or carbonate alkalinity. For this reason, the sea water analysis of Table III includes total inorganic carbon which was calculated from pH, total alkalinity and salinity using the apparent sea water constants of Mehrbach et al.(82) for the dissociation of carbonic acid and the boric acid dissociation constant of Lyman (83), as expressed by Li et al. (84). 

When HX is a carbon acid the value of the rate coefficient, ) for a thermodynamically favourable proton transfer rarely approaches the diffusion limit. Table 1 shows the results obtained for a few selected carbon acids which are fairly representative of the different classes of carbon acids which will be discussed in detail in Sect. 4. For compounds 1—10, the value of k i is calculated from the measured value of k, and the measured acid dissociation constant and, for 13, k, is the measured rate coefficient and k1 is calculated from the dissociation constant. For 11 and 12, both rate coefficients contribute to the observed rate of reaction since an approach to equilibrium is observed. Individual values are obtained using the measured equilibrium constant. In Table 1, for compounds 1—10 the reverse reaction is between hydronium ion and a carbanion whereas for 11, 12 and 13 protonation of unsaturated carbon to give a carbonium ion is involved. For compounds 1—12 the reverse reaction is thermodynamically favourable and for 13 the forward reaction is the favourable direction. The rate coefficients for these thermodynamically favourable proton transfers vary over a wide range for the different acids. In the ionization of ketones and esters, for which a large number of measurements have been made [38], the observed values of fe, fall mostly within the range 10s—101 0 1 mole-1 sec-1. The rate coefficients observed for recombination of the anions derived from nitroparaffins with hydronium ion are several orders of magnitude below the diffusion limit [38], as are the rates of protonation and deprotonation of substituted azulenes [14]. For disulphones [65], however, the recombination rates of the carbanions with hydronium ion are close to 1010 1 mole-1 sec-1. Thermodynamically favourable deprotonation by water of substituted benzenonium ions with pK values in the range —5 to —9 are slow reactions [27(c)], with rate coefficients between 15 and 150 1 mole-1 sec-1 (see Sect. 4.7). 

The original Hammett substituent constants [Hammett, 1937 Hammett, 1970] measuring the overall electronic effect of the meta- and paro-substituents of benzene derivatives having the functional group in the side chain. They were originally calculated from the variation of the acid dissociation constant of substituted benzoic acids (m-, P-XC6H4COOH) in water at 25°C, with respect to the unsubstituted compound (i.e. benzoic acid) ... 

To determine Ki and K2 for H3PO4 from titration data, careful pH measurements ar e made after 0.5 and 1.5 mol of base are added for each mole of acid. It is then assumed that the hydrogen ion activities computed from these data are identical to the desired dissociation constants. Calculate the relative error incurred by the assumption if the ionic strength is 0.1 at the time of each measurement. 

By way of illustration we present the two following tables which include the thermodynamic dissociation constants of acetic acid and o-nitrobenzoic acid calculated from conductivity measurements by J. Kendall. The first column contains the acid concentration, the second contains the classical degree of dissociation a multiplied by 100, in the third is found the dissociation constant calculated by the classical method, the fourth gives ac (classical ion concentration), the fifth column shows the true ion concentration ac//x, the sixth contains the concentration of undissociated acid [cHA] = c — (aclf ), and in the last column is found Ka calculated by equation (15). 

In the context of drug-like substances, hydrophobicity is related to absorption, bioavailability, hydrophobic drug-receptor interactions, metabolism and toxicity. Closely related to log P is the octanol-water distribution coefficient (logDpn), accounting for partition of pH-dependent mixture of ionizable species. Ionization of any compound makes it more water soluble and then less lipophilic. The log D can be calculated from log Pand acid dissociation constant pJC, by the following expression [Cronin, Aptula et al, 2002b Livingstone, 2003] ... 

Figure 1.18 Transition states for the acetate-catalysed, acetic acid-catalysed and water reactions in the mutarotation of tetramethyl glucose. The additional waters for the acetate and acetic acid reactions are drawn to indicate solvation, rather than a change in bonding that would alter fractionation factors. Isotope effects are taken from ref 34 and fractionation factors calculated from their data using 1.0, rather than 1.23, for the fractionation factor of the anomeric hydroxyl. The latter was based on an implausible, equilibrium isotope effect of 4.1 on the acid dissociation constant of tetramethyl glucose. ...

The data in this table are presented as values of piC, defined as can be calculated from the equation the negative logarithm of the acid dissociation constant iC for the... 

---