Stability of ion pairs

The basic effects responsible for the properties of electrolyte solutions are ion solvation, ion 2issociation to ion pairs and higher ion aggregates with and without inclusion of solvent molecules. FTIR (Fourier transform infrared) and MW (microwave) spectra are a valuable source of information on ion-solvent and ion-ion interactions and yield factual knowledge on the structure and dynamics of electrolyte solutions. The efficiency of these methods is exemplified for solvation in aptotic and protic solvents, hydrophobic solvation, association to charged and neutral ion aggregates, and stability of ion pairs. 

After reading the pertinent discussion in Langmuir (1979), contrast the applicability of the electrostatic, Bjerrum, and Fuoss thermodynamic models for predicting the stabilities of ion pairs. 

For assay of active substances both as raw material and from pharmaceutical dosage formulations, the best reliability is assured by using ISME. ISME can be successfully used for uniformity content tests as well as for in vitro and in vivo dissolution tests of tablets.152 160 The ISME proposed for assay of pharmaceutical products is based on the ion pair complexes, but the responses of ISME are highly affected by the stability of ion pair complexes. The ISME run is explained by the multilayer configuration of the membrane.161... 

It is obvious from this theoretical approach that the energetic stabilization of ion-pairs induced by interaction vith the electric field E becomes increasingly important as the size of the ions and their dipole moments (jj) increase. The more polar ion-pairs are more stabilized by E, clearly increasing from tight to loose ion-pairs, i.e. vith their dissociation and polarity. 

The relative stabilities of ion pairs and charged entities in solution also depend on the nature of both salt and solvent. Ion pairs tend to form according to the Coulombic attractive force, especially when the charge/radius ratio is high... 

The extent of the ionization produced by a Lewis acid is dependent on the nature of the more inert solvent component as well as on the Lewis acid. A trityl bromide-stannic bromide complex of one to one stoichiometry exists in the form of orange-red crystals, obviously ionic. But as is. always the case with crystalline substances, lattice energy is a very important factor in determining the stability and no quantitative predictions can be made about the behaviour of the same substance in solution. Thus the trityl bromide-stannic bromide system dilute in benzene solution seems to consist largely of free trityl bromide, free stannic bromide, and only a small amount of ion pairs.187 There is not even any very considerable fraction of covalent tfityl bromide-stannic bromide complex in solution. The extent of ion pair and ion formation roughly parallels the dielectric constant of the solvents used (Table V). The more polar solvent either provides a... 

On comparing results obtained in various solvents, it is practically only thermodynamic effects which can be explained and taken into account, whereas specific solvent effects which depend on the structure of the solvent molecule can only be dealt with empirically. Furthermore the dielectric constant changes with the solvent, and this affects the existence of ion pairs. This effect will always be troublesome for complexes in which a positively charged donor molecule is stabilized by a counter-ion. 

Section 3.2 includes an extensive discussion on the formation of odd-electron bonds, ion pairing, and the distonic stabilization of ion-radicals at the expense of separation between their spins and charges. Section 3.3 deals with ion-radicals from the class of even spin-charge distribution. This class occnrred more frequently in scientific works of past decades. However, the reader will find newly developed manifestations of the principle of the released electron, concerning spread conjugation and the fates of ion-radical precursors with increased dimensionality. 

When the ion-pair partitioning is indicated in the quadrant diagram (below) it becomes obvious that a circle of equilibria is present. Knowing the octanol pKa, the log P and the aqueous pKa should allow one to calculate the partition coefficient of the ion pair. From these equilibria one can write that the difference in log P between the acid and its salt is the same as the difference between the pKa s (Equation 9). The closer the pKa s, the more lipid soluble the ion pair will be, relative to the acid. Internal hydrogen bonding or chelation that stabilizes an ion pair will affect the octanol stability more than the aqueous stability, where it is less needed, and so will decrease the delta pKa. Chelation should therefore favor biolipid solubility of ion pairs. Ultimate examples are available in some ionophores. This is one of the properties of some of the herbicides I pointed out earlier. 

The first reaction is the slow, or rate-determining, ionization of the substrate to form a carbocation intermediate. The products of this first step will tend to be stabilized best in polar solvents. The SnI type of reaction can also lead to the formation of ion pair intermediates, as shown in the following reaction scheme ... 

Figure 18 (a) Mechanism of ion pair transport mediated by a cation carrier (b) plot of initial transport rates (V) of cation picrates as a function of the logarithm of the stability constants... 

This calculation is still hypothetical, in that the actual substance formed when sodium metal reacts with difluorine is solid sodium fluoride, and the standard enthalpy of its formation is -569 kJ mol-1. The actual substance is 311 kJ mol-1 more stable than the hypothetical substance consisting of ion pairs, Na+F (g), described above. The added stability of the observed solid compound arises from the long-range interactions of all the positive Na+ ions and negative F ions in the solid lattice which forms the structure of crystalline sodium fluoride. The ionic arrangement is shown in Figure 7.5. Each Na+ ion is octahedrally surrounded (i.e. coordinated) by six fluoride ions, and the fluoride ions are similarly coordinated by six sodium ions. The coordination numbers of both kinds of ion are six. 

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