The acidity scale

Where do hydrocarbons lie on the acidity scale As the data in Table 8.1 show, both methane (pKa 60) and ethylene (plC, = 44) are very weak acids and thus do not react with any of the common bases. Acetylene, however, has piCa = 25 and can be deprotonated by the conjugate base of any acid whose pKa is greater than 25. Amide ion (NH2-), for example, the conjugate base of ammonia (pKa - 35), is often used to aeprotonate terminal aikynes. 

Hardly surprisingly the tendency of alkanes to lose proton and form carbanions is not marked, as they possess no structural features that either promote acidity in their H atoms, or are calculated to stabilise the carbanion with respect to the undissociated alkane (cf. carboxylic acids, p. 55). Thus CH4 has been estimated to have a pKa value of 43, compared with 4.76 for MeC02H. The usual methods for determining pKa do not, of course, work so far down the acidity scale as this, and these estimates have been obtained from measurements on the iodide/organo-metallic equilibria ... 

In addition it is possible to explain the inversion about the acidity scale of aliphatic alcohols observed on passing from the gas phase to aqueous solutions. Jam s formalism (53) for solvation energy was successfully used on this occasion. 

In principle, the A 0(H) function is of limited interest for kinetic applications because the indicators are chemically very different from the organic substrates generally used. On the other hand, as the measurements are based on pH determination, the length of the acidity scale is limited by the pA" value of the solvents. However, very interesting electrochemical acidity studies have been performed in HF by Tremillon and co-workers, such as the acidity measurement in anhydrous HF solvent and the determination of the relative strength of various Lewis acids in the same solvent. By studying the variation of the potential of alkane redox couples as a function of acidity, the authors provide a rational explanation of hydrocarbon behavior in the superacid media.48... 

The acidity scale in anhydrous hydrogen fluoride has been the subject of electrochemical investigations by Tremillon and coworkers48 and is presented in Figure 1.9. The figure also indicates the acidity constants of various Lewis acids allowed to buffer the medium to a pH value as calculated by Eq. (1.33), or in dilute solution by Eq. (1.34). 

The acidity of a solution is more quantitatively defined as a pH value. This is a term that expresses the hydrogen ion concentration on a 1-14 scale. Solutions with pH 1-6 show acidic properties, pH 7 is neutral and solutions with pH 8-14 are basic. The lower the acidic scale is, the more acidic the solution alternatively, above 7, the more basic a solution is, the greater the value. 

Where do hydrocarbons lie on the acidity scale As the data ia TahU 8.1 indicate, both metliane pKj 60) and ethylene are -ery... 

IV) a close similarity between HCr(CO)5 and MCr(CO)4(NO) (which is also reflected in the occupation number of the two natural orbitals involving s (Table V)). The relative high ratio between s and 3d 2 in these two systems as compared to HMn(CO)5 and H2Fe(CO)4 accounts for the tendency of the two Cr systems to behave as hydride donors. On the other hand the smaller value of the s / 3d 2 ratio in the HFe(CO)4, HMn(CO)5 and H2Fe(CO)4 systems is in line with their nown acidity. Interestingly the value of this ratio roughly follows the acidity scale of these three systems (24). 

Despite the favorable properties of acetonitrile as a solvent, its use for equilibrium acidity measurements has its definite limitations. The pK range that is tolerable is limited at the high end by onset of solvent deprotonation, and at the low end by substrate autodissociation, as has been implicated for HCo(CO)4 [14a] and TpCr(CO)3H [22b]. These limitations can be overcome by the choice of a less polar solvent, e.g. 1,2-dichloroethane (DCE), dichloromethane, or THE. To make reliable, quantitative comparisons of thermodynamic data obtained in different solvents, it is necessary to link the acidity scales and electrode potential references in the different solvents. This has all too often proven to be a far from trivial task. Although, in principle, 1 1 relationship between the acidity scales in different solvents never exists, pK differences between closely related compounds are often almost constant when compared in different solvents. This is because their solvation properties are similar, because of similarities in size and charge distribution. In less... 

The work described in the foregoing sections is of a preliminary nature. Nevertheless, it offers hope that experimental scales of free hydrogen ion concentration (pcn or pmn) in seawater may be feasible. One need not know pmn or pan to derive meaningful equilibrium data, such as acid-base ratios and solubilities, provided that suitable apparent equilibrium constants are chosen (7). In these cases, the unit selected for the acidity scale disappears by cancellation. Nevertheless, the acidity of seawater is a parameter of broader impact. It plays a role, for example, in the kinetics of organic oxidation-reduction reactions and in a variety of quasi-equilibrium processes of a biological nature. The concentration of free hydrogen ions is clearly understood, and its role in these complex interactions is more clearly defined than that of a quantity whose unit purports to involve the concept of a single-ion activity. 

The proton activity is defined as Brpnsted acidity and the acidity scale is related to the deprotonation of an acid to an anion. The pH scale is normalized through the convention that the dissociation constant for the hydronium ion ( Thw) is set equal to unity where Ky/= 1.0.10 at 25°C. The pH of a solution is usually measured by determining the potential of a cell in which a... 

In particular, for the conjugate acid-base pair Ax/Bx, which is located on the acidity scale over the acid-base range of the solvent L of the first kind (see Fig. 1.1.1), the complete transformation into conjugate acid with the formation of the equivalent concentration of the base of the solvent will be observed. It should be added that the acid formed would possess no acidic properties in the said solvent. Similarly, the conjugate pair A2/B2 is completely transformed into the conjugate base, which shows no basicity in the solvent. Hence, the acid-base ranges in solvents of the first kind are limited on two sides. 

To elucidate the validity of the acidity scales obtained in molten alkali metal halides which belong to the solvents of the second kind, we shall consider the principal acid-base equilibrium, which is achieved by the... 

Another disadvantage of the experimental acidity scales is the absence of serious quantitative data because, as a rule, a number of equilibria are reciprocally affected in the studied solutions. In this case, the estimation of some equilibrium constants based on one point is impossible. Therefore, more precise data on the acidity scales may be obtained by establishing equilibrium constants of acid-base reactions and the effect exerted by the cation and anion composition of an ionic melt on them. The regularities obtained on the basis of these parameters will help us to treat some aspects of the problem in question more correctly. 

Thus, the method of constructing the acidity scales based on determination of the position of acid-base pairs using one experimental point (weight) cannot be considered correct. Also, similar scales for solvents of the first kind are distorted appreciably because of the levelling of acidic properties of strong acids which, according to the said method, should possess the same acidic properties. 

Another approach to this problem was proposed by Tremillon et al. [169-171], Their idea consists in the construction of the general acidity scale for chloride melts and determination of the positions of standard solutions of strong Lux bases on this scale. As media for acid-base reactions, chloride melts should be classified among solvents of the second kind, and, therefore, the positions of such solvents on the acidity scale should be described by only one point—the acidity function for the standard solution of a strong base. 

It is obvious that the expression enclosed in the brackets by the author of the present book is nothing but the primary medium effect of O2- expressed via the difference in the values of the equilibrium constants of equation (1.3.6) for the media compared the molten equimolar KCl-NaCl mixture, which was chosen as a reference melt, and for which pKHa/H20 was found to be 14 at 700 °C, and the melt studied. As to the physical sense of the common acidity function Cl, this is equal to the pO of the solution in the molten equimolar KCl-NaCl mixture, whose acidic properties (oxide ion activity) are similar to those of the solution studied. Moreover, from equation (1.3.7) it follows that solutions in different melts possess the same acidic properties (f ) if they are in equilibrium with the atmosphere containing HC1 and H20 and Phc/Ph2o — constant. This explanation confirms that the f function is similar to the Hammett function. Therefore, Cl values measured for standard solutions of strong bases in molten salts allow the prediction of the equilibrium constants on the background of other ionic solvents from the known shift of the acidity scales or the f value for the standard solution of a strong Lux base in the solvent in question. According to the assumption made in Refs. [169, 170] this value may be obtained if we know the equilibrium constant of the acid-base reaction (1.3.6) in the solvent studied. 

There is no the upper limit to the basicity in the molten KCl-LiCl (0.41 0.59) eutectic mixture, and the shift of the acidity scale against the equimolar KCl-NaCl mixture is approximately equal to 8 [169]. 

The concentration of H" " or OH in aqueous solution can vary over extremely wide ranges, from 1 M or greater to 10 M or less. To construct a plot of concentration against some variable would be very difficult if the concentration changed from, say, 10 M to 10" M. This range is common in a titration. It is more convenient to compress the acidity scale by placing it on a logarithm basis. The pH of a solution was defined by Sprenson as... 

Recall from Chapter 7 that, because the equilibrium constants of the blood buffer systems change with temperature, the pH of blood at the body temperature of 37°C is different than at room temperature. Hence, to obtain meaningful blood pH measurements that can be related to actual physiological conditions, the measurements should be made at 37°C and the samples should not be exposed to the atmosphere. (Also recall that the pH of a neutral aqueous solution at 3TC is 6.80, and so the acidity scale is changed by 0.20 pH unit.)... 

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