Dissociation constants, acetic acid indicators

There is less abundant information about HBr and HI, but if we apply Robinson s treatment to the vapour pressure data we find piC(HBr)= -8, pK(Hl)= -9, compared with pR (HCl) = -6, suggesting that the strengths of the three acids are approximately in the ratio 1 10 10. Further, investigations in non-aqueous solvents often show that HBr is a considerably stronger acid than HCl. Thus conductivity measurements in anhydrous acetic acid indicate a ratio of about 20 between the dissociation constants of these two acids, while in acetonitrile we have piC(HBr) = 5.5 and pK(HCl) = 8.9. It is also frequently found that HBr is a much more effective acid catalyst than HCl under conditions where both acids are undissociated. It is of interest that studies... 

Frequently the statement is found that acetic acid is a stronger acid in liquid ammonia than it is in water. It is, however, not true that the dissociation constant of acetic acid is higher in liquid ammonia than it is in water. Comparison of dissociation constants, which are indicative of the acidity of a compound in a particular solvent, is only meaningful in the same solvent. Comparison of dissociation constants in different solvents just have no meaning. In water HsO -ions are produced by the addition of a quantity of acetic acid just as NH4+-ions are produced in liquid ammonia in the presence of acetic acid, but the differences in dielectric constants make the actual concentration of NH4+-ions in liquid ammonia lower than that of H30+-ions in water at the same concentration. The H30" -ion shows stronger acidic properties than the NH4+-ion. 

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... 

The first indication that A-acyloxy-A-alkoxyamidcs reacted by an acid-catalysed process came from preliminary H NMR investigations in a homogeneous D20/ CD3CN mixture, which indicated that A-acetoxy-A-butoxybenzamide 25c reacted slowly in aqueous acetonitrile by an autocatalytic process according to Scheme 4 (.k is the unimolecular or pseudo unimolecular rate constant, K the dissociation constant of acetic acid and K the pre-equilibrium constant for protonation of 25c).38... 

The larger the acid dissociation constant, the stronger is the acid. Hydrochloric acid has an acid dissociation constant of 10, whereas acetic acid has an acid dissociation constant of only 1.74 x 10 . For convenience, the strength of an acid is generally indicated hy its pA a value rather than its A a value The of hydrochloric acid, strong acid, is —7, and the pA a of acetic acid, much weaker acid, is 4.76. 

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)). 

In the following sections we consider the equilibria involved in the titration of a base B with perchloric acid, a reaction of practical interest the titration of acids with acetic acid as solvent is unimportant. The acidity of the solutions can be determined by an indicator or by the change in potential of an electrode responsive to free solvated protons. In water these two methods give the same results. In acetic acid, measurements of potential depend on the extent of formation of free solvated protons, whereas indicators respond to the extent of proton formation whether ion-paired or not thus two different measures of acidity are possible. The constant for ion-pair generation sometimes may be large, but the constant for dissociation to free ions never is. We discuss here the change in acidity during a titration as it would be obtained potentio-metrically and consider the behavior of indicators in Section 4-11. 

Acidity scales for highly acidic solutions in media of low dielectric constant have been proposed, such as for ethanol-water mixtures and glacial acetic acid-acetic anhydride. Ion pairing is a complicating factor in these solvents. The extent of formation of ion pairs and of free ions differs for various indicators and salts in particular, dissociation constants are sensitive to the size of the anion. The Hammett postulate therefore cannot easily be extended to include media of low dielectric constant. 

Figure 1.18 Transition states for the acetate-catalysed, acetic acid-catalysed and water reactions in the mutarotation of tetramethyl glucose. The additional waters for the acetate and acetic acid reactions are drawn to indicate solvation, rather than a change in bonding that would alter fractionation factors. Isotope effects are taken from ref 34 and fractionation factors calculated from their data using 1.0, rather than 1.23, for the fractionation factor of the anomeric hydroxyl. The latter was based on an implausible, equilibrium isotope effect of 4.1 on the acid dissociation constant of tetramethyl glucose. ...

Evidence for the structure (CXXIII) of the hemiacetal is based on the extremely hindered nature of the derived aldehyde (CXXV) and carboxylic acid (CXXVII). Thus, the aldehyde exhibited a negative Cotton effect in methanol, which remained unchanged upon the addition of hydrochloric acid, indicating great resistance toward acetal formation. Attempts to prepare carbonyl derivatives of this aldehyde were unsuccessful. The acid CXXVII was prepared by oxidation of alcohol CXXIV with chromium trioxide in acetic acid. Comparison of the apparent dissociation constant of this acid (pX cs 9.45) with that for... 

The zirconium tetrahalides react with esters to form ZrX4 2 ester adducts (302, 303, 330, 407-410, 412) in which, coordination number six is attained. On the basis of dipole moments (Table XIII), it is concluded that the adducts have the cis structure. This has been supported, at least in the case of ZrCl4 2011300 0 02115, by the infrared spectrum (330). Oryoscopic studies in benzene solution of the 2 1 adducts of zirconium tetrachloride and ethyl formate, ethyl acetate, and ethyl butryrate show that these complexes tend to decompose to the 1 1 species, the extent of dissociation increasing with the number of carbon atoms in the acid radical. The estimated dissociation constant is about 5 x 10", whereas for the ethyl acetate adduct of zirconium tetrabromide it is only 2 X 10". The approximate dissociation constant of the complex zirconium tetraiodide 2 ethyl acetate is 3.5 x 10". The 1 1 species were synthesized by direct reaction in benzene with strictly stoichiometric ratios of the reactants. Oryoscopic determination of molecular weights of the 1 1 complexes indicate that these complexes generally... 

The dissociation constant of CMC is 5.0 X 10, indicating that it is a moderately strong acid, comparable to acetic acid in this property. The molecular weight, as determined by osmotic-pressure determinations, is 6,400 1,000. It is insoluble in water in the acid form, which accounts for its use as a salt. Sodiupi CMC is supplied commercially in several grades, depending upon viscosity and purity. 

Useful exploratory studies of acid-base behaviour in solvents of low dielectric constant have been made by conductance " and potentio-metric " titrations. Association constants are usually obtained from spectrophotometric measurements. The strengths of various bases can be compared by means of their association with an indicator acid like 2,4-dinitrophenol. If both acid and base are colourless, a competition for the base can be established between the acid and an indicator acid like bromophthalein magenta In solvents like benzene, other reactions than simple 1 1 association between B and RX may occur. Self-association of the acid or base is one such auxiliary reaction. A classic example is the dimerisation of carboxylic acids in benzene. If allowance is not made for this, constant values of the quotient [BRX]j[B][RX] will not be obtained. (Variations in the quotient cannot be attributed to interionic forces or other nonideal behaviour BRX is scarcely dissociated into ions at all and in spectrophotometric work very low concentrations of B and RX can be used.) Evidence for association ratios other than 1 1 can be obtained from indicator studies. The method developed by Kolthoff and Bruckenstein for studying reactions in anhydrous acetic acid fails for reactions in benzene and similar solvents because more than one acid molecule reacts with the indicator to give complexes of the form/w J r"(HX)yi. In such studies it is generally a good approximation... 

Attribution of chemical reactions to experimentally accessible dissociation constants is indicated by Equations (l)-(7) (shown here for oxal-acetic acid, the same assignments are used for a-ketoglutaric acid) ... 

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