Acid dissociation constant weak bases

This relationship between and Kb simplifies the tabulation of acid and base dissociation constants. Acid dissociation constants for a variety of weak acids are listed in Appendix 3B. The corresponding values of Kb for their conjugate weak bases are determined using equation 6.14. 

The principal limitation to using a titration curve to locate the equivalence point is that an inflection point must be present. Sometimes, however, an inflection point may be missing or difficult to detect, figure 9.9, for example, demonstrates the influence of the acid dissociation constant, iQ, on the titration curve for a weak acid with a strong base titrant. The inflection point is visible, even if barely so, for acid dissociation constants larger than 10 , but is missing when is 10 k... 

The plT at which an acid-base indicator changes color is determined by its acid dissociation constant. For an indicator that is a monoprotic weak acid, ITIn, the following dissociation reaction occurs... 

In the discussion of the relative acidity of carboxylic acids in Chapter 1, the thermodynamic acidity, expressed as the acid dissociation constant, was taken as the measure of acidity. It is straightforward to determine dissociation constants of such adds in aqueous solution by measurement of the titration curve with a pH-sensitive electrode (pH meter). Determination of the acidity of carbon acids is more difficult. Because most are very weak acids, very strong bases are required to cause deprotonation. Water and alcohols are far more acidic than most hydrocarbons and are unsuitable solvents for generation of hydrocarbon anions. Any strong base will deprotonate the solvent rather than the hydrocarbon. For synthetic purposes, aprotic solvents such as ether, tetrahydrofuran (THF), and dimethoxyethane (DME) are used, but for equilibrium measurements solvents that promote dissociation of ion pairs and ion clusters are preferred. Weakly acidic solvents such as DMSO and cyclohexylamine are used in the preparation of strongly basic carbanions. The high polarity and cation-solvating ability of DMSO facilitate dissociation... 

An inflection point in a pH-rate profile suggests a change in the nature of the reaction caused by a change in the pH of the medium. The usual reason for this behavior is an acid-base equilibrium of a reactant. Here we consider the simplest such system, in which the substrate is a monobasic acid (or monoacidic base). It is pertinent to consider the mathematical nature of the acid-base equilibrium. Let HS represent a weak acid. (The charge type is irrelevant.) The acid dissociation constant, = [H ][S ]/[HS], is taken to be appropriate to the conditions (temperature, ionic strength, solvent) of the kinetic experiments. The fractions of solute in the conjugate acid and base forms are given by... 

The Henderson-Hasselbalch equation provides a general solution to the quantitative treatment of acid-base equilibria in biological systems. Table 2.4 gives the acid dissociation constants and values for some weak electrolytes of biochemical interest. 

Formazans behave as weak acids as well as weak bases. Salts of formazans have been isolated.26,334,335 The acid dissociation constants of some substituted formazans have been determined from their solution spectra.336... 

Consider how a weak electrolyte is distributed across the gastric mucosa between plasma (pH 7.4) and gastric fluid (pH 1.0). In each compartment, the Henderson-Hasselbalch equation gives the ratio of acid-base concentrations. The negative logarithm of the acid dissociation constant is designated here by the symbol pAa rather than the more precisely correct pK1. 

It is possible to compare the strengths of weak acids by the values of their acid dissociation constants Ka. Figure 3.1 shows the titration curves for acids (HA or BH+) of different Ka values. The ordinate shows poH, which is defined by paH = -loga(SI I)). paH corresponds to the pH in aqueous solutions (see Section 3.2). The poH of non-aqueous solutions can be measured with a glass pH electrode or some other pH sensors (see Sections 3.2.1 and 6.2). For the mixture of a weak acid A and its conjugate base B, poH can be expressed by the Henderson-Hassel-balch equation ... 

The definition of pH is pH = —log[H+] (which will be modified to include activity later). Ka is the equilibrium constant for the dissociation of an acid HA + H20 H30+ + A-. Kb is the base hydrolysis constant for the reaction B + H20 BH+ + OH. When either Ka or Kb is large, the acid or base is said to be strong otherwise, the acid or base is weak. Common strong acids and bases are listed in Table 6-2, which you should memorize. The most common weak acids are carboxylic acids (RC02H), and the most common weak bases are amines (R3N ). Carboxylate anions (RC02) are weak bases, and ammonium ions (R3NH+) are weak acids. Metal cations also are weak acids. For a conjugate acid-base pair in water, Ka- Kb = Kw. For polyprotic acids, we denote the successive acid dissociation constants as Kal, K, K, , or just Aj, K2, A"3, . For polybasic species, we denote successive hydrolysis constants Kbi, Kb2, A"h3, . For a diprotic system, the relations between successive acid and base equilibrium constants are Afa Kb2 — Kw and K.a Kbl = A w. For a triprotic system the relations are A al KM = ATW, K.d2 Kb2 = ATW, and Ka2 Kb, = Kw. 

Terms to Understand acid dissociation constant, K.d fraction of association, a pK weak base... 

The bridging hydroxo ligand is a weak acid as well as an extremely weak base. The acid properties are by far the best investigated, and acid dissociation constants have been reported for singly and doubly bridged dinuclear complexes containing one or two hydroxo bridges (cf. Tables XV and XVI). 

Figure 2.11. Percent of ionogenic (ionizable) species present for weak acids and bases when solution pH is 2 units above or below the acid dissociation constant.

The acid-dissociation constant, Ka, is the equilibrium constant for the ionization of a weak acid to a hydrogen ion and its conjugate base ... 

So far, we ve been looking at the equilibria of weak acids, from which we developed the acid-dissociation constant, Ka. There is a similar process for weak bases. Let s use the weak base ammonia. In solution, ammonia establishes the equilibrium shown below ... 

Very little accurate e.m.f. work has been done on the dissociation constants of bases, chiefly because moderately weak bases are very volatile, while the non-volatile bases, e.g., anilines, are usually very weak. An exception to this generalization is to be found in the aliphatic amino-acids which will be considered in connection with the subject of ampho-... 

The same general conclusions concerning the effect of the dissociation constant of the weak base and the concentration of the salt on the degree of hydrolysis are applicable as for the salt of a weak acid. The results in Table LXV would hold for the present case provided the column headed ka were replaced by h. Further, since the dissociation constants of bases do not vary greatly with temperature, the influence of increasing temperature on the hydrolysis of the salt of a weak base will be very similar to that on the salt of a weak acid. 

This table can be helpful in estimating the pATs of other weak acids from their structures. In using this table it is important to remember that -log [H ] is used in the expression for the acid dissociation constant in terms of pH. To obtain pKs based on -log y (H ) [H ], add 0,0.08, 0.11,0.12, and 0.14 at ionic strengths of 0, 0.05, 0.10,0.15, and 0.25 M, respectively, at 298.15 K as indicated by Table 1.3. PaddedForm rounds the output to two figures to the right of the decimal point. There is a list of full names of reactants in the Appendix of this book. The reactants bpg, nmn, pep, and prpp are bisphosphoglycerate, nicotinamidemononucleotide, phosphoenolpyruvate, and 5-phosphoribosyl-alpha-pyrophosphate, respectively. 

Where Ka is the acid dissociation constant. (Note that the definition of Ka is based on the Bronsted definition.) Values of Ka can vary tremendously (IQis to 10-60). after all anything with at least one proton can be considered an acid under some circumstances with this definition. The common definition of a strong acid is an acid which dissociates completely in a 1 M solution. The common strong acids in aqueous solution, such as sulphuric, nitric and hydrochloric acids have Ka values (for the first dissociation in the case of sulfuric) of 10 to 10. Thus they all dissociate completely (first dissociation only for sulphuric) in aqueous solution, though they wiU have different strengths in some other solvents. Most common organic acids are weak in aqueous solution, having Ka values of lO-s to Note that... 

At the beginning, the solution contains only a weak acid or a weak base, and the pH is calculated from the concentration of that solute and its dissociation constant. 

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