Electrostatic ion-dipole

Ions not solvated are unstable in solutions between them and the polar solvent molecules, electrostatic ion-dipole forces, sometimes chemical forces of interaction also arise which produce solvation. That it occurs can be felt from a number of effects the evolution of heat upon dilution of concentrated solutions of certain electrolytes (e.g., sulfuric acid), the precipitation of crystal hydrates upon evaporation of solutions of many salts, the transfer of water during the electrolysis of aqueous solutions), and others. Solvation gives rise to larger effective radii of the ions and thus influences their mobilities. 

This indicates a lack of dynamic cohesion within the adducts i.e. the substrate has considerable freedom for reorientation within the receptor. The apparent reason for an absence of mechanical coupling is the nearly cylindrical symmetry of cucurbituril, which allows the guest an axis of rotational freedom when held within the cavity. Hence, the bound substrates show only a moderate increase in tc relative to that exhibited in solution. No relationship exists between values and the thermodynamic stability of the complexes as gauged by K (or K, cf. Tables 1 and 2). It must be concluded that the interior of cucurbituril is notably nonsticky . This reinforces previous conclusions that the thermodynamic affinity within adducts is chiefly governed by hydrophobic interactions affecting the solvated hydrocarbon components, plus electrostatic ion-dipole attractions between the carbonyls of the receptor and the ammonium cation of the ligands. 

Ion-dipole forces are important in solulions of ionic compounds in polar solvents where solvated species such as NatOH,) and F(H 20) (for solutions of NaF in H.O) exist. In the case of some metal ions these solvated species can be sufficiently stable to be considered as discrete species, such as [Co(NHj)6]j+. Complex ions such as the latter may thus be considered as electrostatic ion—dipole interactions, but this oversimplification (Crystal Field Theory sec Chapter 11) is less accurate than are alternative viewpoints. 

In 1979 the results of ab initio calculations at the 4-3IG and 6-3IG level on the same complex as well as a formaldehyde/H+ complex were reported. Structural and energetic ctxnparisons of the two complexes showed that while Li+ prefers a linear geometry for electrostatic ion-dipole bonding, the proton coordinates to the carbonyl through a largely covalent bond, resulting in a bent structure (Cs symmetry)... 

Chemisorption on an MgO surface will be primarily an acid/base interaction. Cation sites are Lewis acids and may interact with donor molecules such as H2O through a combination of electrostatics (ion-dipole attraction) and orbital overlap. Oxide ions also act as basic sites and can interact with acceptors such as H+. In fact one of the most common dissociative reactions is the deprotonation of an adsorbate to produce surface hydroxyl groups. 

The extensive use by Dugan and Magee of trajectory calculations to compute close-collision cross sections for the collision of ions with polar molecules has been reviewed in Section 4.2.2d. For such calculations, a form for only the attractive part of the potential need be assumed and, in this case, a particular value of the ion-molecule separation was used to define a close collision. The form chosen for the potential was the simple, anisotropic, electrostatic ion-dipole potential plus the ion-induced-dipole potential and, for reasons discussed in that section, such a model may only be applied to ion-molecule collisions at thermal energies. 

It is obvious that the major part of the solvation energy is provided by electrostatic ion-dipole interactions only in those systems where the role of the coordinate bonds between the solvent molecules and the dissolved ion is a subordinate one. It follows that it is primarily in the case of anions where agreement can be expected between the experimental solvation energies and those calculated by taking into account exclusively the ion-dipole forces. However, hydrogen bonding may occur here also, giving rise to an increased ion-solvent interaction. 

When classifying these anions, it is important to remember that water, unlike other liquids, exhibits good solvent properties for salts because of its specific structure and the special interaction mechanism between the ion and the water molecule. When an ion is solvated by water, hydrogen bonds are broken (cavity effect) and the water structure is destroyed. The larger the ion, the higher the energy required for the formation of a cavity with molecular dimension. On the other hand, electrostatic ion-dipole interactions occur that lead to the formation of a new structure. Thus, the smaller the ionic radius and the higher the... 

The general trend of a solubility curve can be derived from the Le Chatelier s principle considering the dissolution process as a reaction of breaking and forming new bonds, such as breaking ionic bonds in an ionic crystal (of a salt) and formation of (weak) electrostatic ion-dipole interactions due to hydration of the ions, for example, according to... 

As a more representative example of weaker neutral H-bonded species, let us consider the (HF)2 dimer, which offers a particularly clear contrast to the dipole-dipole expectations of classical electrostatics. The (HF)2 species is bound by about 5 kcal/mol (in the same range as water dimer and many other common H-bonded species) and exhibits a curiously bent equilibrium geometry, as shown in 1/0-9.1. Although HF has a robust dipole moment (calculated as p = 1.92 Debye) and F HF has the linear geometry expected for an electrostatic ion-dipole complex, the nonlinear geometry of (HF)2 clearly differs from the expected linear geometry of a dipole-dipole model. What s going on here ... 

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