Proton-transfer reaction strengths

A proton transfer reaction involves breaking a covalent bond. For an acid, an H — X bond breaks as the acid transfers a proton to the base, and the bonding electrons are converted to a lone pair on X. Breaking the H — X bond becomes easier to accomplish as the bond energy becomes weaker and as the bonding electrons become more polarized toward X. Bond strengths and bond polarities help explain trends in acidity among neutral molecules. 

Figure 5.35 displays a schematic energy-level diagram for each stationary species along the proton-transfer reaction coordinate. Although Pauling bond strengths... 

Variations in dielectric constant should alter the relative strength of acids of different charge types, since the amount of electrical work involved in a proton-transfer reaction must vary with the dielectric constant of the medium. Since a lowering in dielectric constant increases the work required to separate the ions (for example, to ionize an uncharged acid one must create an anion and a cation), any addition of organic solvent should lead to an increase of pXa values. 

When the reference acid, HA, is water, we can set up a scale of base strengths from the equilibrium constant, Kb, measured for the proton-transfer reaction shown in Equation 23-3 ... 

Finally, let s look at a salt such as (NH4)2C03 in which both the cation and the anion can undergo proton-transfer reactions. Because NH4+ is a weak acid and CO32- is a weak base, the pH of an (NH4)2C03 solution depends on the relative acid strength of the cation and base strength of the anion ... 

A quantitative determination of the strength of Lewis acids to establish similar scales (Ho) as discussed in the case of protic (Br0nsted-type) superacids would be most useful. However, to establish such a scale is extremely difficult. Whereas the Brpnsted acid-base interaction invariably involves a proton transfer reaction that allows meaningful comparison, in the Lewis acid-base interaction, involving for example Lewis acids with widely different electronic and steric donating substituents, there is no such common denominator.25,26 Hence despite various attempts, the term strength of Lewis acid has no well-defined meaning. 

These acid-base pairs are called conjugate pairs. The extent of the reaction depends on the relative strength of the acid and the base. The model developed by Bron-sted has the big advantage that -> acid-base equilibria can be described mathematically in a simple way, and it is used to estimate the pH of solutions. (According to the IUPAC recommendations the term protolysis for proton transfer reactions should be discouraged). 

According to the Arrhenius theory, the strength of an acid depends on the fraction of the acid that ionizes to produce H30+ ions. When defined in terms of the Brpnsted approach, the acid strength is reflected by the magnitude of the equilibrium constant for the proton transfer reaction... 

A proton transfer reaction represents an equilibrium. Because an acid donates a proton to a base, thus forming a conjugate acid and conjugate base, there are always two acids and two bases in the reaction mixture. Which pair of acids and bases is favored at equilibrium The position of the equilibrium depends on the relative strengths of the acids and bases. 

Usually these reactions have been studied in water or in other protic solvents such as the alcohols. Thus, the acid-base properties of the solvent are important in determining the relative strength of acids and bases which are solutes in water. This leads to the definition of two other types of proton transfer reaction, namely, the protolysis reaction,... 

A very convincing support for the existence of solvent controlled proton dissociation reactions in aqueous solutions has risen from the theoretical studies of Ando and Hynes [105-108] who have studied the proton dissociation of simple mineral acids HCl and HF in aqueous solutions. The two acids seem to follow a solvent-controlled proton transfer mechanism with a Marcus-like dependence of the activation energy on the acid strength. Recently, a free energy relationship for proton transfer reactions in a polar environment in which the proton is treated quantum mechanically was found by Kiefer and Hynes [109, 110]. Despite the quite different conceptual basis of the treatment the findings bear similarity to those resulting from the Marcus equation Eq. (12.19) which has been used to correlate the proton transfer rates of photoacids with their piG [ 101,102 ]... 

To summarize, the proton transfer reaction can be broken into three distinct parts Diffusion of the reactants to within the radius of the ionic atmosphere accelerated diffusion to within the encounter distance and subsequent conversion of the encoimter complex to products. For reactions in which the equilibrium is rapidly established within the encounter complex, the rate equations are dominated by the diffusion process. This results in the loss of information about the dynamics of the encounter complex. For such a reaction some information can be obtained about the ionic radius by varying the ionic strength and using an electrostatic theory (such as is done for Deby-Hiickel activity coefficients) to calculate the effect of shielding by the ions. ... 

It is possible to observe intermediates in the proton transfer reactions by using low-polar or non-polar media and by varying the concentration and strength of proton donors with the combination of variable temperature IR and NMR spectroscopies. Under suitable conditions, this approach further enables the determination of the thermodynamic parameters associated to each step from the temperature dependence of the corresponding formation constants. It is now experimentally confirmed that the first step of the interaction indeed consists of the dihydrogen bond formation. 

Understand how the equilibrium position of a proton-transfer reaction relates the strengths of the acids and bases involved. (Section 16.3)... 

TABLE 2.3 1 1 Some Organic Lewis Bases and Their Relative Strengths in Proton-Transfer Reactions ... 

---