Cation acidity constant table

Table 1.5 represents acidity constants of organic compounds (AH) and their cation-radicals (AH+ ) calculated for their solutions in DMSO (a very polar solvent) at 25°C. 

Acidity constants are given for a range of compounds in Table 8.1. When more than one ionizable proton is present, pK, pK2, etc. values are given. Cations formed from the indicated compound by protonation are indicated by ( + 1) or (+2) for a dication. For example, the dissociation of 3-acetamidopyridine is reported in Table 8.1 as 4.37(+l). This means dissociation of the compound that is protonated (at the pyridine nitrogen atom). 

In a polar solvent, heterolytic cleavage leading to proton abstraction is usually facilitated because of the favorable solvation energy of the proton, and cation radicals are ordinarily much more acidic than the corresponding neutral compounds. Table 1-5 combines acidity constants of organic compounds (AH) and their cation radicals (AH+ ) calculated for their solutions in dimethylsulfoxide (DMSO, a very polar solvent) at 25°C. 

Charged metals (cations) in water behave as Lewis acids (willing to accept electrons). Water on the other hand, because it is willing to share its two unshared oxygen-associated pair of electrons, behaves as a Lewis base. Strong H2Q-metal (Lewis base-Lewis acid) interactions allow H+ on the water molecule to dissociate, hence, low pH water is produced. The degree of dissociation of water interacting with a cation (Mn+) is described by the metal hydrolysis constant (Table 2A)... 

Table 1 Ionization Constants of Neutral and Cation Acids of type HA and HB respectively, in Water, Amphiprotic, and Dipolar Aprotic Solvents...

In the case of pentavalent cations, the conditional interaction constant (Table I) obtained for Np02" (log P = 4.6) shows a relatively low affinity of the neptunyl cation with the humic acids as it could be predicted from the charge of the ion. The interaction constant obtained for the hexavalent cation (U) with humic acid (Table I) is independent of pH (4-5) in the non-hydrolysis pH-range but some variation with uranium concentration is observed as for trivalent cations. Moreover, the complexation of uranium to humic substances is of the same order of magnitude than the complexation of trivalent actinides which corroborates chemical analogy between both cations. 

It is plain from this table that the ratio of the dissociation constants in water and alcohol is not constant. Uncharged acids possess dissociation constants which in alcohol are 10 to 10 times smaller than in water the difference is much less for monovalent cation acids. 

As mentioned earlier, a method for classifying Lewis acids is based on relative equilibrium constants in water for the formation of complexes. Hydration energies must play an important role. As Table 1.4 shows, F forms a stronger bond to all cationic acids in the gas phase than 1 does. But because the hydration energy of F is 50 kcal greater than for I , inversions in A/Z°, and in formation constants, can occur for acids like Ag+, but not for CH3CO+. 

The most water-like of this class of solvents is methanol, for it maintains much the same nice balance of basic and acidic properties found in water. Its autoprotolysis constant is smaller than that of water (Table 3.3.4) because of its lower dielectric constant. Medium effects for transfer of ionisation equilibria from water to methanol are approximately constant for closely related acids. For six cation acids, the pyri-dinium ion and five methyl derivatives, the average medium ejffect is 0.06 0.02, small because the ionisation of these cations creates no new charge field. For phenol and thirteen of its derivatives the medium effect is 4.32 0.09 smaller values are obtained for nitrophenols, possibly because the anions are stabilised by dispersion interactions with methanol. For 23 carboxylic acids, aliphatic and aromatic, the average medium effect is 4.87 0.15. Values of the medium effect for individual acids are collected in Appendix 3.5.5. 

Table 1 Ionization constants of neutral and cation acids of type HA and HB+, respectively, in water, amphiprotic, and dipolar aprotic solvents.

Now we shall analyze the regularities in rate constants of protolytic photodissociation of hydroxyaromatic compounds in the micellar phase and in the membranes of vesicles ( Tables 1 and 5 ). For one and the same compounds in cationic micelles the acidity constants are usually greater than in the aqueous solution ( pA decreases ), and dissociation rate... 

Recently kinetic data have become available for the nitration in sulphuric acid of some of these hydroxy compounds (table 10.3). For 4-hydroxyquinoline and 4-methoxyquinoline the results verify the early conclusions regarding the nature of the substrate being nitrated in sulphuric acid. Plots of log Q against — (Lf + logioflHao) fo " these compounds and for i-methyl-4-quinolone have slopes of i-o, i-o and 0-97 at 25 C respectively, in accord with nitration via the majority species ( 8.2) which is in each case the corresponding cation of the type (iv). At a given acidity the similarity of the observed second-order rate constants for the nitrations of the quinolones and 4-methoxy-quinoline at 25 °C supports the view that similarly constructed cations are involved. Application of the encounter criterion eliminates the possibilities of a... 

Bases of low polarizabiUty such as fluoride and the oxygen donors are termed hard bases. The corresponding class a cations are called hard acids the class b acids and the polarizable bases are termed soft acids and soft bases, respectively. The general rule that hard prefers hard and soft prefers soft prevails. A classification is given in Table 3. Whereas the divisions are arbitrary, the trends are important. Attempts to provide quantitative gradations of "hardness and softness" have appeared (14). Another generaUty is the usual increase in stabiUty constants for divalent 3t5 ions that occurs across the row of the Periodic Table through copper and then decreases for zinc (15). 

This is a reasonable inference, because site binding is significant only with multivalent cations and strong electrostatic interactions. Under these conditions ion polarization occurs and bonds have some covalent character (Cotton Wilkinson, 1966). This is illustrated by the data of Gregor, Luttinger Loebl (1955a,b). They measured the complexation constants of poly(acrylic acid), 0 06 n in aqueous solution, with various divalent metals, which, as it so happens, are of interest to AB cements (Table 4.1). The order of stability was found to be... 

The bimolecular rate constants were determined (Burke 2001) for the repair of carotenoid radical cations by trolox, ascorbic, ferrulic, and uric acids from the pulse radiolysis studies of carotenoids in aqueous micellar solutions (see Table 14.10). 

Table XIX contains stability constants for complexes of Ca2+ and of several other M2+ ions with a selection of phosphonate and nucleotide ligands (681,687-695). There is considerably more published information, especially on ATP (and, to a lesser extent, ADP and AMP) complexes at various pHs, ionic strengths, and temperatures (229,696,697), and on phosphonates (688) and bisphosphonates (688,698). The metal-ion binding properties of cytidine have been considered in detail in relation to stability constant determinations for its Ca2+ complex and complexes of seven other M2+ cations (232), and for ternary M21 -cytidine-amino acid and -oxalate complexes (699). Stability constant data for Ca2+ complexes of the nucleosides cytidine and uridine, the nucleoside bases adenine, cytosine, uracil, and thymine, and the 5 -monophosphates of adenosine, cytidine, thymidine, and uridine, have been listed along with values for analogous complexes of a wide range of other metal ions (700). Unfortunately comparisons are sometimes precluded by significant differences in experimental conditions.

The situation just discussed probably applies also to the attack of water on other kinds of stabilized carbocations. For example, some of the many transient carbocations studied in recent years by McClelland and Steenken, and their coworkers (e.g. Steenken et al., 1986 McClelland and Steenken, 1988), have relatively constant kOH/kw ratios. For alkyldialkoxy cations [41], oh w = 103 to 104 m-1 and so pKt = 10 to 11 for the trialkoxy analogues [42], kon/kw == 104 to 106m i and p/Cf = 8 to 10 (Table A6.3), suggesting a more acidic transition state for [42], due to the extra oxygen atom. Within each series there is a systematic variation of p/C, since logfc0H correlates with logfcw, with a slope of approximately 0.6 (McClelland and Steenken, 1988), rather than 1. Estimates of the Leffler indices for the two series of... 

For nitration of aromatic hydrocarbons with acetylnitrate, there is a clear linear correlation between the IPs of these hydrocarbons and rate constants relative to benzene (Pedersen et al. 1973). Table 4.4 jnxtaposes spin densities of cation-radicals and partial rate factors of ring attacks in the case of nitration of isomeric xylenes with nitric acid in acetic anhydride. 

Considering that heavy and transition metals may reach subsurface water as hydrated cations at neutral pH, they may behave as acids, due to formation of a hydration shell surrounding the cation. The acidity of hydrated cations depends on the acid dissociation constant (pK ) values. The lower the pK value of the metal, the lower the pH at which precipitates are formed. Values of pK for major heavy metals are presented in Table 5.5. 

It is obvious that the acid-base property of a dissolved oxide may markedly affect the structure of a silicate melt. A dissolved acidic oxide associates the free oxygen, thus displacing reaction 6.4 toward the right, resulting in a marked correlation between the field strength of the dissolved cation and the polymerization constant of the melt. This correlation is shown in the values listed in table 6.2. 

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