Undissociated acid

Addition of excess H ions to this solution will cause the equilibrium to move towards undissociated acid thereby decreasing the concentration of Ac . This effect is known as the common-ion effect and is of considerable practical importance. Thus, e.g. in the precipitation of metal ions as insoluble sulphides. 

Results can sometimes be unexpected. The first study of this type made use of labeled Aerosol OTN [111], an anionic surfactant, also known as di-n-octylsodium sulfosuccinate. The measured F was twice that in Eq. III-93 and it was realized that hydrolysis had occurred, that is, X + H2O = HX + OH , and that it was the undissociated acid HX that was surface-active. Since pH was essentially constant, the activity of HX was just proportional to C. A similar behavior was found for aqueous sodium stearate [112]. 

The inhibitory activity of sorbates is attributed to the undissociated acid molecule. The activity, therefore, depends on the pH of the substrate. The upper limit for activity is approximately pH 6.5 in moist appHcations the degree of activity increases as the pH decreases. The upper pH limit can be increased in low water activity systems. The following indicates the effect of pH on the dissociation of sorbic acid, ie, percentage of undissociated sorbic acid at various pH levels (76,77). 

It will be noted that pyrrole-3-carboxylic acid (154) is an appreciably weaker acid than benzoic acid and this is attributed to the stabilization of the undissociated acid by electron release from nitrogen. The 2-carboxylic acids of furan, thiophene, selenophene and tel-lurophene are all stronger acids than benzoic acid, tellurophene-2-carboxylic acid (pisTa 4.0) being the weakest acid in this series (77AHC(21)119). 

Organic acids convert the blue mesomerically stabilized phenolate anion to the red undissociated acid. Reductones (e.g. ascorbic acid) reduce the reagent to a colorless salt. 

Evidence for the undissociated acid comes from spectroscopic data but the solutions cannot be concentrated without decomposition. HCIO2 is a moderately strong acid Ka(25°C) 1.1 x 10 (cf H2Se04 Ka. 2x 10 2, H4P2O7 Ka 2.6 x 10 2). 

This equation says that the logarithm of the concentration of dissociated acid [A-] divided by the concentration of undissociated acid [HA] is equal to the pH of the solution minus the pKa of the acid. Thus, if we know both the pH of the solution and the pKa of the acid, we can calculate the ratio of [A-] to [HA]. Furthermore, when pH = pKai the two forms HA and A- are present in equal amounts because log 1 = 0. 

If the dissociation constant of the acid HA is very small, the anion A- will be removed from the solution to form the undissociated acid HA. Consequently more of the salt will pass into solution to replace the anions removed in this way, and this process will continue until equilibrium is established (i.e. until [M + ] x [A-] has become equal to the solubility product of MA) or, if sufficient hydrochloric acid is present, until the sparingly soluble salt has dissolved completely. Similar reasoning may be applied to salts of acids, such as phosphoric(V) acid (K1 = 7.5 x 10-3 mol L-1 K2 = 6.2 x 10-8 mol L-1 K3 = 5 x 10 13 mol L-1), oxalic acid (Kx = 5.9 x 10-2 mol L-K2 = 6.4 x 10-5molL-1), and arsenic)V) acid. Thus the solubility of, say, silver phosphate)V) in dilute nitric acid is due to the removal of the PO ion as... 

The activity coefficient ya of the undissociated acid is approximately unity in dilute aqueous solution. Expression (24) thus becomes ... 

Kresge and Chiang480 measured the rate coefficients for detritiation of [1-3H]-2,4,6-trimethoxybenzene in acetate buffers and found the first-order rate coefficient (lO7 ) to increase from 2.5 at 0.01 M acetic acid to 8.3 at 0.1 M acetic acid, whereas if the reaction was specific acid-catalysed no change in rate should have been observed. A similar technique to that described above for separation of the rate coefficients due to hydronium ions and other acids was used, the values for the former being obtained using dilute hydrochloric acid at which acidities no undissociated acid was present (Table 131). Rate coefficients were then measured... 

The kinetics of desulphonation of sulphonic acid derivatives of m-cresol, mesitylene, phenol, p-cresol, and p-nitrodiphenylamine by hydrochloric or sulphuric acids in 90 % acetic acid were investigated by Baddeley et a/.701, who reported (without giving any details) that rates were independent of the concentration of sulphuric acid and nature of the catalysing anion, and only proportional to the hydrogen ion concentration. The former observation can only be accounted for if the increased concentration of sulphonic acid anion is compensated by removal of protons from the medium to form the undissociated acid this result implies, therefore, that reaction takes place on the anion and the mechanism was envisaged as rapid protonation of the anion (at ring carbon) followed by a rate-determining reaction with a base. 

From the above equations that relate to [H ] and to the concentrations of undissociated acid and its conjugate base, when... 

The bulk of root products are C compounds derived from products of photosynthesis. The root products that are not C compounds are few (H", inorganic ions, water, electrons) but nevertheless are deemed to be highly significant. Both H and electrons may be secreted as C compounds in the form of undissociated acids and reducing agents, respectively. 

For ionisation of m- and p-substituted benzoic acids (44), the hydroxylic solvent is capable of solvating both the undissociated acid (44) and the carboxylate anion (45) obtained from its ionisation. 

By plotting the measured rate constant versus the undissociated acid concentration, one obtains for this type of catalysis a straight line with intercept kx and slope ax = (/cHA + kA-/q). If the procedure is repeated for other ratios, enough information is obtained to permit evaluation of /cHA and kA-. The hydrogen and hydroxide ion concentrations corresponding to a given ratio q may be determined from equation 7.3.12 and the dissociation constant for water. 

A reaction catalyzed by undissociated acid will have the dependence of log k on pH shown in Figure 3g. Specific acid and specific base catalysis are presumed to be absent. If specific and general acid and base catalysis are both operative, one is able to obtain a variety of interesting log k versus pH curves, depending on the relative contributions of the different terms in various pH ranges. Curves i and j of Figure 7.3 are simple examples of these types. 

Since the spontaneous reaction term dominates in the pH range 4 to 6, studies of the reaction in this range are particularly suitable for measuring the small catalytic effects of weak acids and weak bases. If one employs the more general expression for k involving contributions from the undissociated acid and the anion resulting from dissociation, determine the coefficients of these terms from the data below. 

It is of interest to consider the form of the Bronsted plot or Eigen plot to be expected for reaction of a series of related intramolecularly hydrogen-bonded acids with hydroxide ion by the mechanisms in Schemes 5 and 6. The effect of a substituent on the value of the dissociation constant of an intramolecularly hydrogen-bonded acid (8 log K) will be two-fold. The stability of the undissociated acid will be modified because of a substituent effect on the... 

We first determine the concentration of undissociated acid. We then use this value in a set up that is based on the balanced chemical equation. 

But we need to be careful when talking about the magnitudes of Consider the case of sodium ethanoate dissolved in dilute mineral acid the reaction occurring is, in fact, the reverse of that in Equation (4.45), with a proton and carboxylate anion associating to form undissociated acid. In this case, = 1 mol before the reaction occurs, and its value decreases as the reaction proceeds. In other words, we need to define our reaction before we can speak knowledgeably about it. We can now rewrite our question, asking Why is < 1 for a weak acid ... 

Theoretical considerations show that the free energy of dissociation of an acid in water, and hence the dissociation constant, is governed by the algebraic sum of the free energies for the solution of the undissociated acid in water, for vaporisation of the acid, for the formation of a free proton and an anion from the molecule of acid in the gas phase, and for hydration of the proton and anion. Thus the true acidity, given by the third of these... 

There are certain reactions, e.g. inversions of sucrose and methane etc. in which the rate of reactions were found to be proportional to the concentration of H+ ions. Similarly, there are reactions which are catalyzed by OH ions, e.g. conversion of acetone into diacetone alcohol or decomposition of nitroso-triacetoneamine. These are known as specific hydrogen ion catalyzed or specific hydroxyl-ion catalyzed reactions. Also there are some reactions in which both H+ and OFF ions act as catalysts probably along with water. The undissociated acid or base have negligible effect on the rate of reaction. The hydrolysis of ester is an example in which both H+ and OH ions act as catalyst... 

Another way to obtain HENRY S constant H of undissociated acid is from high concentration vapor-liquid equilibria where dissociation is negligible. Using NRTL equation for the representation of the data of BROWN and EWALD (6) at high concentration in acetic acid (10 2 < x < 1), he finds the limiting activity coefficient of undissociated acid at 100°C... 

H Henry s constant of undissociated acid or base (atm. kg/mole )... 

The indicator is in equilibrium between the undissociated acid, which is one colour, and its conjugate base, which is a different colour. 

The dissociation constant of acetic acid is 1.754 x 10. Calculate the degree of dissociation of 0.01-molar acid in the presence of 0.01-molar NaCl. Use the Debye-Hiickel limiting law to calculate the activity coefficients of the ions. Take the activity coefficient of the undissociated acid as unity. In your calculation neglect the concentration of in comparison with the concentration of Na+. 

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