Proton solvation, solution acidity

The extension of analytical mass spectrometry from electron ionization (El) to chemical ionization (Cl) and then to the ion desorption (probably more correctly ion desolvation ) techniques terminating with ES, represents not only an increase of analytical capabilities, but also a broadening of the chemical horizon for the analytical mass spectrometrist. While Cl introduced the necessity for understanding ion—molecule reactions, such as proton transfer and acidities and basicities, the desolvation techniques bring the mass spectrometrist in touch with ions in solution, ion-ligand complexes, and intermediate states of ion solvation in the gas phase. Gas-phase ion chemistry can play a key role in this new interdisciplinary integration. 

An acid is defined as a proton donor within the Lowry-Brpnsted theory (see Chapter 6). Molecules of acid ionize in aqueous solution to form an anion and a proton, both of which are solvated. An acid such as ethanoic acid (VI) is said to be weak if the extent to which it dissociates is incomplete we call it strong if ionization is complete (see Section 6.2). 

For the acidic proton transfer of Eqn. 3-44, the proton solvation processes of Eqns. 3-32 and 3-42 are represented by the proton level versus concentration curves of Eqns. 3-39 and 3-43, respectively, as shown in Fig. 3-19. In this proton level diagram, the proton level in an acetic acid solution is given by the intersecting point (mH,o - where cross each other the occupied proton level versus concentration curve of H3O ion and the vacant proton level versus concentration curve of Ac" ion, as expressed in Eqn. 3-46 ... 

Although solvated electrons were used to trigger chain reactions like SrnI reactions by primary scission of bonds [309], most reports in the literature mention the need for a proton donor. As a matter of fact, to render electron exchanges between ens and organic substrates irreversible, it appears necessary to add to the solvent some proton donor known to react very slowly with the solvated electron. Mixtures of HMPA-ethanol and HMPA-acetic acid have been used successfully [310] as proton rich solutions. 

The standard state might be chosen in various ways (e.g., as the state at inhnitely diluted solution). The resulting standard acidity scale is characterized by the activity of the proton solvated by the given solvent, HS, according to... 

If a catalyst has a different physical state from the reaction mixture that it is catalysing, it is said to be a heterogeneous catalyst. An example of this would be a metal catalyst, such as platinum, which catalyses the hydrogenation of gaseous alkenes. If, however, the catalyst has the same physical state, it is said to be a homogeneous catalyst, for example, the solvated proton in an acid catalysed reaction that is occurring in solution. 

From the definition of acidity in solution as presented in Eq. (7.4), it is clear that in order to compute AG° one must know the proton solvation energy. However, from the previous discussion on the structural models for the hydrated proton one may anticipate some difficulties. For example, how many water molecules should be considered in the calculation of the proton solvation energy In other words How large should one take the H (H20) cluster One reasonable approach should be to examine the convergence of... 

Rustad JR, Hay BP (1995) A molecular-dynamics study of solvated orthosilicic acid and orthosilicate anion using parameterized potentials.. Geochim Cosmochim Acta 59 1251-1257 Rustad JR, Hay BP, Halley JW (1995) Molecular dynamics simulation of iron(III) and its hydrolysis products in aqueous solution. J Chem Phys 102 427-431 Rustad JR, Wasserman E, Felmy AR (1999b) Molecular modeling of the surface charging of hematite - II Optimal proton distribution and simulation of surface charge versus pH relationships. Surf Sci 424 28-35... 

In all proton containing solvents acid-base phenomena can be described in terms of the Bronsted-Lowby theory. All of these solvents have the solvated proton in common as the solvent cation, and this determines to a considerable extent the chemistry in their solutions. Bronsted acids are usually characterized by their acidic strength in water, e.g. by the acidity constant in this solvent. Thus acetic acid and hydrofluoric acid both behave as moderately weak acids in water with at room temperature. When acetic acid is dissolved in liquid hydrogen fluoride, the former is successfully competing for the protons, so that acetic acid acts as a base ( acetic-base ) in this medium just as it does in nitric acid ... 

One of the most important insights we can gain about acid-base chemistry is the ability to predict what is the stronger acid or base when confronted with a comparison. Here, this will be a completely thermodynamic analysis, and we leave it until Chapter 9 to discuss the kinetics of proton transfers, In order to be able to make a sound prediction, we wilt cover correlations between gas phase and solution acidities. Then, numerous factors that control acidity will be covered—namely, solvation, resonance, electronegativity, inductive effects, etc. These are all topics that we have covered in Chapters 1-4, and hence this chapter serves as a nice recap. 

In another solvent the activity of the proton will, in general, be greatly different from that in water. Thus the solvated proton in liquid ammonia is NH4 , while the solvated proton in sulphuric acid is the H3S04 cation, and solutions of NH4 in liquid ammonia or of H3804 in sulphuric acid have vastly different acidities from solutions of H3O in water. Comparison of the acidities of such widely different systems is a difficult problem that is discussed later. We simply point out here that an approximate pH scale may be defined... 

H3O" is strictly the oxonium ion actually, in aqueous solutions of acid this and Other solvated-proton structures exist, but they are conveniently represented as... 

Procedures to compute acidities are essentially similar to those for the basicities discussed in the previous section. The acidities in the gas phase and in solution can be calculated as the free energy changes AG and AG" upon proton release of the isolated and solvated molecules, respectively. To discuss the relative strengths of acidity in the gas and aqueous solution phases, we only need the magnitude of —AG and — AG" for haloacetic acids relative to those for acetic acids. Thus the free energy calculations for acetic acid, haloacetic acids, and each conjugate base are carried out in the gas phase and in aqueous solution. 

Many organic reactions involve acid concentrations considerably higher than can be accurately measured on the pH scale, which applies to relatively dilute aqueous solutions. It is not difficult to prepare solutions in which the formal proton concentration is 10 M or more, but these formal concentrations are not a suitable measure of the activity of protons in such solutions. For this reason, it has been necessaiy to develop acidity functions to measure the proton-donating strength of concentrated acidic solutions. The activity of the hydrogen ion (solvated proton) can be related to the extent of protonation of a series of bases by the equilibrium expression for the protonation reaction. 

The mode of extraction in these oxonium systems may be illustrated by considering the ether extraction of iron(III) from strong hydrochloric acid solution. In the aqueous phase chloride ions replace the water molecules coordinated to the Fe3+ ion, yielding the tetrahedral FeCl ion. It is recognised that the hydrated hydronium ion, H30 + (H20)3 or HgO,, normally pairs with the complex halo-anions, but in the presence of the organic solvent, solvent molecules enter the aqueous phase and compete with water for positions in the solvation shell of the proton. On this basis the primary species extracted into the ether (R20) phase is considered to be [H30(R20)3, FeCl ] although aggregation of this species may occur in solvents of low dielectric constant. 

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