Dissociation constants, acetic acid bases

Ionisation and Dissociation Constants for Acid-Base Equilibria in Anhydrous Acetic Acid at... 

Table I. Over-all Dissociation Constants of Acids, Bases, and Salts in Glacial Acetic Acid (-log autoprotolysis constant of glacial acetic acid, pKg = 14.45). ...

In this way Kolthoff and Bruckenstein59 determined spectrophotometrically at 25° C for acid-base equilibria in glacial acetic acid the following ionization and dissociation constants of the bases ... 

The concept of using the base 10 logarithm to express the magnitude is a widespread practice today. Equilibrium constants of chemical reactions are often noted or compared as pK values where pK = — log 10 (magnitude of equilibrium constant). For example, the extent of dissociation of acetic acid, the acid in vinegar, is quantified by an equilibrium constant of 1.8 x 10-5. Here, then, pK = — log,o (1.8 x 1(T5) = 4.74. 

The equilibrium constant Kn can be obtained by adding equations for (1) the acid dissociation of acetic acid, (2) the base protonation of ammonia, and (3) the reverse of the dissociation of water ... 

Although the first of these reactions has the larger equilibrium constant, we can t consider it to be the principal reaction because the reactants and products are identical. Proton transfer from acetic acid to its conjugate base is constantly occurring, but that reaction doesn t change any concentrations and therefore can t be used to calculate equilibrium concentrations. Consequently, the principal reaction is dissociation of acetic acid. 

When a strong acid or base undergoes a complete ionization in solution, the concentrations of the newly formed ions can be understood using basic stoichiometry principles. This is because essentially all of the acid is converted to ions. With weaker acids and bases, equilibrium is established between the ions, much like the equilibria studied in the last chapter. The concentrations of the ions must be determined by using an equilibrium constant, K. The equilibrium constants used to describes acid-base equilibria are in the same form as Kc from the last chapter. Well use the dissociation of acetic acid to begin our description of the new equilibrium constant. 

An important conclusion has been arrived at by McBam and Kam (loc at) in a paper with the following self-explanatory title The effect of salts on the vapour pressure and degree of dissociation of acetic acid in solution An expenmental refutation of the hypothesis that neutral salts increase the dissociation constants of weak acids and bases ... 

Hall NF, The strength of organic bases in glacial acetic acid solution, JACS, 52,5115-5128 (1930). Cited in Perrin DD, Dissociation Constants of Organic Bases in Aqueous Solution, 1965, Butterwordis, Lond (1965) No. 389. Ref. Hll. Other values, also uncertain (U), are reported in Perrin. 

The ionization eonstant should be a function of the intrinsic heterolytic ability (e.g., intrinsic acidity if the solute is an acid HX) and the ionizing power of the solvents, whereas the dissoeiation constant should be primarily determined by the dissociating power of the solvent. Therefore, Ad is expeeted to be under the eontrol of e, the dieleetrie eonstant. As a consequenee, ion pairs are not deteetable in high-e solvents like water, which is why the terms ionization constant and dissociation constant are often used interchangeably. In low-e solvents, however, dissociation constants are very small and ion pairs (and higher aggregates) become important species. For example, in ethylene chloride (e = 10.23), the dissociation constants of substituted phenyltrimethylammonium perchlorate salts are of the order 10 . Overall dissociation constants, expressed as pArx = — log Arx, for some substanees in aeetie acid (e = 6.19) are perchloric acid, 4.87 sulfuric acid, 7.24 sodium acetate, 6.68 sodium perchlorate, 5.48. Aeid-base equilibria in aeetie acid have been earefully studied beeause of the analytical importance of this solvent in titrimetry. 

The shapes of the titration curves of weak electrolytes are identical, as Figure 2.13 reveals. Note, however, that the midpoints of the different curves vary in a way that characterizes the particular electrolytes. The pV, for acetic acid is 4.76, the pV, for imidazole is 6.99, and that for ammonium is 9.25. These pV, values are directly related to the dissociation constants of these substances, or, viewed the other way, to the relative affinities of the conjugate bases for protons. NH3 has a high affinity for protons compared to Ac NH4 is a poor acid compared to HAc. 

Weak acid with a strong base. In the titration of a weak acid with a strong base, the shape of the curve will depend upon the concentration and the dissociation constant Ka of the acid. Thus in the neutralisation of acetic acid (Ka— 1.8 x 10-5) with sodium hydroxide solution, the salt (sodium acetate) which is formed during the first part of the titration tends to repress the ionisation of the acetic acid still present so that its conductance decreases. The rising salt concentration will, however, tend to produce an increase in conductance. In consequence of these opposing influences the titration curves may have minima, the position of which will depend upon the concentration and upon the strength of the weak acid. As the titration proceeds, a somewhat indefinite break will occur at the end point, and the graph will become linear after all the acid has been neutralised. Some curves for acetic acid-sodium hydroxide titrations are shown in Fig. 13.2(h) clearly it is not possible to fix an accurate end point. 

The authors studied, as they call it, "acid-base equilibria in glacial acetic acid however, as they worked at various ratios of indicator-base concentration to HX or B concentration, we are in fact concerned with titration data. In this connection one should realize also that in solvents with low e the apparent strength of a Bronsted acid varies with the reference base used, and vice versa. Nevertheless, in HOAc the ionization constant predominates to such an extent that overall the picture of ionization vs. dissociation remains similar irrespective of the choice of reference see the data for I and B (Py) already given, and also those for HX, which the authors obtained at 25° C with I = p-naphthol-benzein (PNB) and /f B < 0.0042, i.e., for hydrochloric acid K C1 = 1.3 102, jjrfflci 3 9. IQ-6 an jjHC1 2.8 10 9 and for p-toluenesulphonic acid Kfm° = 3 7.102( K ms 4 0.10-6) Kmt = 7 3.10-9... 

One could go on with examples such as the use of a shirt rather than sand reduce the silt content of drinking water or the use of a net to separate fish from their native waters. Rather than that perhaps we should rely on the definition of a chemical equilibrium and its presence or absence. Chemical equilibria are dynamic with only the illusion of static state. Acetic acid dissociates in water to acetate-ion and hydrated hydrogen ion. At any instant, however, there is an acid molecule formed by recombination of acid anion and a proton cation while another acid molecule dissociates. The equilibrium constant is based on a dynamic process. Ordinary filtration is not an equilibrium process nor is it the case of crystals plucked from under a microscope into a waiting vial. 

There is an important relationship between the dissociation constant for an acid, Ka, and the dissociation constant for its conjugate base, K. Consider acetic acid and its dissociation in water. 

Errors may occur in the Gran titration procedure if weakly acidic species with dissociation constants (expressed as pKd) in the range of the extract pH are present. In particular, curvature or reduction (or both) of the slope of the Gran exponential plot results (24), because weak acid dissociation and titration of released free acidity take place during the portion of the titration used for end-point determination. Fortuitously, some of the common, weak carboxylic acids (e.g., formic and acetic) are not stable toward microbial decomposition when collected in aerosol samples from the atmosphere, so much of the historical data base on strong acid content of aerosols does not suffer from this positive error source, unless of course the microbial processes produce additional strong acids. 

PK. A measurement of the complete ness of an incomplete chemical reaction. It is defined as the negative logarithm ito the base 101 of the equilibrium constant K for the reaction in question. The pA is most frequently used to express the extent of dissociation or the strength of weak acids, particularly fatty adds, amino adds, and also complex ions, or similar substances. The weaker an electrolyte, the larger its pA. Thus, at 25°C for sulfuric add (strong acid), pK is about -3,0 acetic acid (weak acid), pK = 4.76 bone acid (very weak acid), pA = 9.24. In a solution of a weak acid, if the concentration of undissociated acid is equal to the concentration of the anion of the acid, the pAr will be equal to the pH. 

The strength of an acid is determined by its ability to give up protons while the strength of a base is determined by its ability to take up protons. This strength is indicated by the dissociation or equilibrium constant, (pKfl), for the acid or base strong acids have a low affinity for protons, while weak acids have a higher affinity and only partially dissociate (e.g. HC1 (strong) and acetic acid (weak)). 

The acidity constant is a measure of the strength of an acid. If the acidity constant for a particular acid is near 1, about equal amounts of the acid and its conjugate base are present at equilibrium. A strong acid, which dissociates nearly completely in water, has an acidity constant significantly greater than 1. A weak acid, which is only slightly dissociated in water, has an equilibrium constant significantly less than 1. The acidity constant for acetic acid is 1.8 X 10-5—only a small amount of acetic acid actually ionizes in water. It is a weak acid. 

The important difference between the dissociation mechanism of LnMEDTA complexes and the acid catalyzed dissociation of LnEDTA is the formation of mixed LnMEDTA-Ac" complex. The stability constants of such mixed complexes are small (1 to 10) and this may explain why such mixed complexes do not show up in the kinetics of dissociation or exchange of LnEDTA complexes. Further, in the present case [LnMEDTA]0 association with an acetate anion should be more favourable than [LnEDTA]- with an acetate anion based on electrostatic theory. Reference to Table 7.14 shows that Km values increase with increasing atomic number (decreasing radii) for heavy lanthanides but are independent for La-Nd series. 

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