Lattice energy table

Some properties of selected vanadium compounds are Hsted in Table 1. Detailed solubiUty data are available (3), as are physical constants of other vanadium compounds (4). Included are the lattice energy of several metavanadates and the magnetic susceptibiUty of vanadium bromides, chlorides, fluorides, oxides, and sulfides (5). 

These trends are apparent In the values of lattice energy that appear in Table Notice, for example, that the lattice energies of the alkali metal chlorides decrease as the size of the cation increases, and the lattice energies of the sodium halides decrease as the size of the anion increases. Notice also that the lattice energy of MgO is almost four times the lattice energy of LiF. Finally, notice that the lattice energy of Fc2 O3, which contains five ions in its chemical formula, is four times as large as that of FeO, which contains only two ions in its chemical formula. 

C08-0081. Refer to the lattice energy values in Table Predict the lattice energy of MgS, and explain your prediction. 

The table shows the lattice energy for some ionic compounds. Based on these data, which of these compounds would require the most energy to separate the bonded ions ... 

The ligand field stabilization is expressed in the lattice energies of the halides MX2. The values obtained by the Born-Haber cycle from experimental data are plotted v.v. the d electron configuration in Fig. 9.5. The ligand field stabilization energy contribution is no more than 200 kJ mol-1, which is less than 8% of the total lattice energy. The ionic radii also show a similar dependence (Fig. 9.6 Table 6.4, p. 50). 

There is another use of the Kapustinskii equation that is perhaps even more important. For many crystals, it is possible to determine a value for the lattice energy from other thermodynamic data or the Bom-Lande equation. When that is done, it is possible to solve the Kapustinskii equation for the sum of the ionic radii, ra + rc. When the radius of one ion is known, carrying out the calculations for a series of compounds that contain that ion enables the radii of the counterions to be determined. In other words, if we know the radius of Na+ from other measurements or calculations, it is possible to determine the radii of F, Cl, and Br if the lattice energies of NaF, NaCl, and NaBr are known. In fact, a radius could be determined for the N( )3 ion if the lattice energy of NaNOa were known. Using this approach, which is based on thermochemical data, to determine ionic radii yields values that are known as thermochemical radii. For a planar ion such as N03 or C032, it is a sort of average or effective radius, but it is still a very useful quantity. For many of the ions shown in Table 7.4, the radii were obtained by precisely this approach. 

As defined earlier, the lattice energy is positive while the solvation of ions is strongly negative. Therefore, the overall heat of solution may be either positive or negative depending on whether it requires more energy to separate the lattice into the gaseous ions than is released when the ions are solvated. Table 7.7 shows the heats of hydration, AH °hy(, for several ions. 

The lattice energies of these compounds are easily calculated or found in tables. When the values are substituted in Eq. (9.103), it is found that... 

It is also interesting to note that metal ions having low polarizability (Al3+ Be2+ etc.) are those that are acidic (as shown in Eq. (9.17)). Also, in Chapter 7 we discussed how the polarization of ions leads to a lattice energy that is higher than that predicted on the basis of electrostatic interactions alone. The polarizability data shown in the table make it easy to see that certain ions are much more polarizable than others. Although we will not visit again all of the ramifications of electronic polarizability, it is a very useful and important property of molecules and ions that relates to both chemical and physical behavior. 

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... 

The early agreement between calculated and experimental heats for fluorides was fortuitous because the high value given to D(F2) was compensated by the large electron affinity value [A/f F ) = lD(Ft) -.EA(F)]. The drop in value of AH°(F >) over the years (see Table I) vitiates some of the more elaborate lattice-energy calculations and Ka-pustinskii s semiempirical method seems adequate (138), but see reference (126). 

Table 1.12 Lattice energies of olivine end-members values in kJ/mole (Ottonello, 1987).

The CFSE contribution to lattice energy is almost insignificant for meta- and orthosilicates in which normal and distorted octahedral coordinations are present (Ml and Ml sites, respectively). As shown in table 1.20, the CFSE gap between normal and distorted octahedral fields is in fact only a few kJ/mole. 

Table 1.19 Lattice energy and CFSE for selected spinel compounds. Values of U, E, and from Ottonello (1986). CFSE values calculated with values proposed by McClure (1957). Data in kJ/mole x = degree of inversion.

The MEG model has been extensively used to determine lattice energies and interionic equilibrium distances in ionic solids (oxides, hydroxides, and fluorides Mackrodt and Stewart, 1979 Tossell, 1981) and defect formation energies (Mack-rodt and Stewart, 1979). Table 1.21 compares the lattice energies and cell edges of various oxides obtained by MEG treatment with experimental values. 

Table 1.21 MEG values of lattice energy and lattice parameter for varions oxides, compared with experimental values. Source of data Mackrod and Stewart (1979). C/ is expressed in kJ/mole (in A) corresponds to the cell edge for cnbic snbstances, whereas it is the lattice parameter in the a plane for AI2O3, Fe203, and Ga203 and it is the lattice parameter parallel to the sixfold axis of the hexagonal unit cell in mtiles CaTi03 and BaTi03.

Table 4.2 Defect energies in forsterite and fayalite based on lattice energy calcnlations. 7 = ionization potential E = electron affinity = dissociation energy for O2 A77 = enthalpy of defect process (adapted from Ottonello et af, 1990).

Table 5.12 reports a compilation of thermochemical data for the various olivine components (compound Zn2Si04 is fictitious, because it is never observed in nature in the condition of pure component in the olivine form). Besides standard state enthalpy of formation from the elements (2) = 298.15 K = 1 bar pure component), the table also lists the values of bulk lattice energy and its constituents (coulombic, repulsive, dispersive). Note that enthalpy of formation from elements at standard state may be derived directly from bulk lattice energy, through the Bom-Haber-Fayans thermochemical cycle (see section 1.13). 

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