Base dissociation constant calculation

The initial goal of the kinetic analysis is to express k as a function of [H ], pH-independent rate constants, and appropriate acid-base dissociation constants. Then numerical estimates of these constants are obtained. The theoretical pH-rate profile can now be calculated and compared with the experimental curve. A quantitative agreement indicates that the proposed rate equation is consistent with experiment. It is advisable to use other information (such as independently measured dissociation constants) to support the kinetic analysis. 

Chemists have calculated the extent to which most acids and bases will dissociate in water. This mathematical value is called the acid dissociation constant (Ka) for acids and the base dissociation constant (Kb) for bases. The higher the value for Ka or Kb, the more the acid or base dissociates in water and the stronger it is. 

The characteristic taste of tonic water is due to the addition of quinine. Quinine is a naturally occurring compound that is also used to treat malaria. The base dissociation constant, Kb, for quinine is 3.3 x 10 . Calculate [OH ] and the pH of a 1.7 x 10 mol/L solution of quinine. 

Pyridine, C5H5N, is used to manufacture medications and vitamins. Calculate the base dissociation constant for pyridine if a 0.125 mol/L aqueous solution has a pH of 9.10. 

The base-dissociation constant expression can be used to calculate the concentrations of ions just like the acid-dissociation expression. The problem-solving strategy is identical. 

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... 

These three values for y permit the calculation of a concentration-based dissociation constant from the thermodynamic constant of 7.1 X 10- (see Appendix 3) ... 

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. 

From what we have said above, it follows that the acid-base equilibrium in the solutions containing metal cations and oxide ions in different sections of the titration curve is described either by the dissociation constant (in unsaturated solutions) or by the values of solubility product (in saturated solutions). In Refs. [175, 330] we proposed a method based on the analysis of the scatter in the calculated equilibrium parameters corresponding to the titration process. Indeed, in the unsaturated solution section there is no oxide precipitation and the calculated value of the solubility product increases monotonously (the directed shift) whereas the calculated value of the dissociation constant fluctuates about a certain value, which is the concentration-based dissociation constant of the studied oxide. 

The extent to which a weak base accepts a proton from water to form OH is expressed by a base-dissociation constant, K. Bronsted-Lowry bases include NH3 and amines and the anions of weak acids. All produce basic solutions by accepting from water, which yields OH and thus makes [HaO" ] < [OH ]. A solution of HA is acidic because [HA] [A ], so [HsO ] > [OH j. A solution of A is basic because [A ] >> [HA], so [OH ] > [H3O ]. By multiplying the expressions for Kg of HA and Kb of A , we obtain K. This relationship allows us to calculate either Kg of BH, the cationic conjugate acid of a molecular weak base B, or Kb of A , the anionic conjugate base of a molecular weak acid HA. 

Although the acid-dissociation constant for phenol (CgHjOH) is listed in Appendix D, the base-dissociation constant for the phenolate ion (CgHsO ) is not. (a) Explain why it is not necessary to list both for phenol and Ky for the phenolate ion. (b) Calculate Ky for the phenolate ion. (c) Is the phenolate ion a weaker or stronger base than ammonia ... 

Alternatively, a classification series resulting from the reaction of bases with their standard partner water following the general scheme Bi + H2O Ai + OH can also be given. The corresponding base dissociation constant /Cb can be calculated using Eq. (2.23) ... 

This expression enables us to calculate the value of the degree of hydrolysis from the dissociation constants of the acid and the base. 

The above examples assume that the strong base KOH is completely dissociated in solution and that the concentration of OH ions was thus equal to that of the KOH. This assumption is valid for dilute solutions of strong bases or acids but not for weak bases or acids. Since weak electrolytes dissociate only slightly in solution, we must use the dissociation constant to calculate the concentration of [H" ] (or [OH ]) produced by a given molarity of a weak acid (or base) before calculating total [H" ] (or total [OH ]) and subsequendy pH. 

If the agent is an acid or a base its degree of ionization will depend on the pH. If its acid dissociation constant,is known, the degree of ionization at any pH may be calculated or determined by reference to published tables. 

Figure 6.16 Energy level diagram for the two-step inhibition of dihydrofolate reductase by methotrexate. The AGbinding were calculated at 30°C base on the dissociation constants reported by Williams et al. (1979).

Dissociation constant of silicic acid calculated according to the a + = [(Kacx Kw)/c]1/2fbrmula for dissociation of salts formed from weak acid and strong base a+ is the activity of protons (from pH), K w is the ionization constant of water, and c is the concentration of silicate solution. 

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-... 

T0 calculate the retardation factors for ionizable compounds such as acids and bases, the fraction of unionized acid (aj or base (at) needs to be determined (see Dissociation constant). According to Guswa et al. (1984), if it is assumed only the un-ionized portion of the acid is adsorbed onto the soil, the retardation factor for the acid becomes ... 

C. Reduction Potentials of Lewis Base Complexes and the Calculation of Dissociation Constants for the Fe(II) Complexes... 

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... 

In gradient elution of weak acids or bases, gradients of organic solvent (acetonitrile, methanol, or tetrahydrofuran) in buffered aqueous-organic mobile phases are most frequently used. The solvent affects the retention in similar way as in RPC of nonionic compounds, except for some influence on the dissociation constants, but Equations 5.8 and 5.9 usually are accurate enough for calculations of gradient retention volumes and bandwidths, respectively. 

There are several ways of identifying whether a compound is an acid or a base, depending on what it does with protons and electrons pH and pOH calculations, along with the values of dissociation constants K and /CJ, can help chemists determine the properties of these acids and bases. 

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