Acid-dissociation constant cation reaction with

From the acid dissociation constants of EDTA (1,2) it can be calculated that between pH 7,4 and 6,6 an EDTA solution is represented to an extent greater than 99,8<7oby the forms HY and HjY . If the EDTA is in sufficient excess with respect to the cation studied, the pH of the solution, on complexing, will remain between the above limits and the experimental reaction will be described unambigously by the equation ... 

We can now make sensible guesses as to the order of rate constant for water replacement from coordination complexes of the metals tabulated. (With the formation of fused rings these relationships may no longer apply. Consider, for example, the slow reactions of metal ions with porphyrine derivatives (20) or with tetrasulfonated phthalocyanine, where the rate determining step in the incorporation of metal ion is the dissociation of the pyrrole N-H bond (164).) The reason for many earlier (mostly qualitative) observations on the behavior of complex ions can now be understood. The relative reaction rates of cations with the anion of thenoyltrifluoroacetone (113) and metal-aqua water exchange data from NMR studies (69) are much as expected. The rapid exchange of CN " with Hg(CN)4 2 or Zn(CN)4-2 or the very slow Hg(CN)+, Hg+2 isotopic exchange can be understood, when the dissociative rate constants are estimated. Reactions of the type M+a + L b = ML+(a "b) can be justifiably assumed rapid in the proposed mechanisms for the redox reactions of iron(III) with iodide (47) or thiosulfate (93) ions or when copper(II) reacts with cyanide ions (9). Finally relations between kinetic and thermodynamic parameters are shown by a variety of complex ions since the dissociation rate constant dominates the thermodynamic stability constant of the complex (127). A recently observed linear relation between the rate constant for dissociation of nickel complexes with a variety of pyridine bases and the acidity constant of the base arises from the constancy of the formation rate constant for these complexes (87). 

According to the Brpnsted definition, the acidity of a molecule is associated with its capacity to give up a proton Ph—NH2 — Ph—NH +H+. The change of standard enthalpy or free energy of this deprotonation reaction is a measure of the intrinsic acidity. As discussed above, in solution, the propensity of an aniline derivative is to accept a proton. The measured dissociation constant (pATa) is related to the basicity of the neutral molecule (or the acidity of the anilinium cations). As a consequence, relatively little is known about their acidity and/or the anilinide anions. However, the NH acidities have been well established in hydroxamic acids even though the latter usually behave as O-acids134. It is therefore of interest to get some insight into the deprotonation of aniline in the gas phase. 

This procedure of expressing the acids as oxo-hydroxo complexes of water is in full agreement with the corresponding reactions occurring with, e.g., the transition metal cations in aqueous solutions. Table 8.5 provides the dissociation constants for acids expressed in this unconventional way. °... 

The new ligand (6) also bears a 3— charge when it is fully ionized, but a recent kinetic study of its formation reactions in fact dealt with 2-h cations, specifically of Ca, Mg, Mn, and Cd. This primarily thermodynamic study reveals some tantalizing hints on kinetic behavior, for example the period of about 5 hours mentioned for the Mn " plus (6) system to come to equilibrium/ Lanthanide(III) cations react with (7) about a hundred times faster, but with (8) about ten times slower, than with the much-studied >NCH2C0 ligand (3)/ Rate constants for dissociation of lanthanide(III) complexes of (7) and (8) are, in the present context, rather rapid—of the order of 10 to 10 s in neutral solution, with the usual acid-catalyzed pathway/ ... 

When this is the case, the heat of reaction must be quite independent of the nature of the anion and of the cation, aa these are not affected by the reaction. This is clearly true for nitric and hydrochloric acids with all the bases given in the table. For sulphuric and carbonic acids, however, the conditions for the validity of the theory are apparently not fulfilled. In the first case, the heat of dilution of sulphuric acid amounts to 2000 cal., and this amount must be subtracted from the figure given in the table, as it is evolved when the alkali and acid are mixed. In the second case, carbonic acid is so weak an acid that it is practically undissociated. The heat necessary for the dissociation into ions therefore uses up part of the heat of neutralisation. From the table it follows that the electrolytic dissociation of J mol. HgCOg requires 13700 — 10200 = 3500 calories. The constant heat of neutrahsation 13700 cal. is the heat of ionisation of water, i.e, the quantity of heat required for the dissociation of water, and liberated on the combination of its ions. 

Angelici and co-workers have ranked the acidity and the bond dissociation energies of about 50 cationic hydrides by protonating neutral metal complexes in CH2CICH2CI (DCE) with triflic acid (HOSO2CF3) and measuring the enthalpy of the reaction [53, 54]. These fall in the range of 10 to 40 kcal/mol. These results will be anion-dependent because of the low dielectric constant of this solvent. 

Introduction. Solutions of f-BuOK in DMSO are highly basic because the solvent strongly complexes with potassium cations, producing activated ligand-separated and dissociated f-butoxide anions in a medium of high dielectric constant. This base/ solvent system is capable of deprotonating weakly acidic carbon and other acids.It is widely used to effect /3-elimination reactions and isomerizations of unsaturated systems. DMSO K+ is present in low concentrations in f-BuOK/DMSO solutions. ... 

Photoprotolytic reactions in micellar solutions were systematically studied by the authors of the present review with coworkers [15, 16, 66, 120-122]. Effective dissociation rate constants ki for a set of hydroxyaromatic compounds in anionic, cationic and uncharged micelles were determined from the dependence of fluorescence quantum yields on the concentration of micelles [Eq. (22)] or on the concentration of acid [Eq. (49)]. The latter method was also used for determining the effective equilibrium constants and effective rate constants of the reverse protonation reaction [Eq. (49)]. The data obtained are presented in Table 3. 

Further kinetic results relevant to the ligand replacement reactions discussed in this section are those for dissociation of lanthanide(m)-cydta complexes. Nearly all the tervalent rare-earth cations were included in this investigation, along with the closely related complexes of scandium(ra) and yttrium(m). Dissociation rates vary with pH, but are unaffected by the addition of copper(n) ions. Thus there is no direct transfer of cydta from Ln + to Cu + in the manner of equation (7) (next section). The rate constants for the acid-dependent dissociation path show a large variation with the nature of the lanthanide cation they vary from 1291 mol s (at 25 °C, fjb = 0.1 mol 1 ) for lanthanum(ra) to 0.0171 mols for lutecium(m) the trend of rate constants is one of regular decrease as the ionic radius decreases. The rate constant of 0.019 1 mol s for dissociation of the scandium(in) complex is, however, much higher than would be expected from the ionic radius of Sc +. Activation parameters for dissociation of these cydta complexes are reported for the La", Gd, Dy Tm ", Lu , and Y compounds. ... 

Monovalent cations such as Na" and form weak complexes with the mono- and difunctional organic acid anions. Stability constants are typically in the range of log K = —1.1 to —1.5 (for the dissociation reaction Martell and Smith 1977). In any aqueous speciation model a small percentage (less than about 10%) of the available organic acid anion will form monovalent cation complexes. These complexes should be included in chemical models for completeness. A small error will result if the complexes are ignored and this may or may not be of concern depending on the application. Kharaka et al. (1986) report that the Na and K salicylate complexes are more stable, however, and inclusion of these species will be necessary in waters having substantial salicylic acid concentrations. 

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