Structures of Actual Ionic Solids

Closest packed structures contain twice as many tetrahedral holes as packed spheres. Closest packed structures contain the same number of octahedral holes as packed spheres. 

In this section we will consider some specific binary ionic solids to show how these solids illnstrate the ideas of ion packing. Becanse an ionic solid must be neutral overall, the stoichiometry of the compound (the ratio of the numbers of anions to cations) is determined by the ion charges. On the other hand, the structure of the compound (the placement of the ions in the solid) is determined, at least to a first approximation, by the relative sizes of the ions. 

Before we consider specific compounds, we need to consider the locations and relative numbers of tetrahedral and octahedral holes in the closest packed structures. The location of the tetrahedral holes in the face-centered cubic unit cell of the ccp structure is shown in Fig. 16.41(a). Note from this figure that there are eight tetrahedral holes in the unit cell. Recall from the discussion in Section 16.4 that there are four net spheres in the face-centered cubic unit cell. Thus there are twice as many tetrahedral holes as there are packed spheres in the closest packed structure. 

The location of the octahedral holes in the face-centered cubic unit cell is shown in Fig. 16.42(a). The easiest octahedral hole to find in this structure is the one at the center of the cube. Note that this hole is surrounded by six spheres, as is required to form an octahedron. Since the remaining octahedral holes are shared with other unit cells, they are more difficult to visualize. However, it can be shown that the number of octahedral holes in the ccp structure is the same as the number of packed spheres. 

Using these ideas, we will now consider the structures of some specific ionic solids. 

This means that for the solid MX, the M+ ions occupy the octahedral holes in the range where 

TABLE 16.6 Guidelines for Filling Various Types of Holes for the Ionic Solid MX 

For large M+ ions r+ 0.732R —), the solid switches to the simple cubic arrangement just described. 

The guidelines for filling the tetrahedral, octahedral, and cubic holes are summarized in Table 16.6. 

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In the first chapter, we defined the nature of a solid in terms of its building blocks plus its structure and symmetry. In the second chapter, we defined how structures of solids are determined. In this chapter, we will examine how the solid actually occurs in Nature. Consider that a solid is made up of atoms or ions that are held together by covalent/ionic forces. It is axiomatic that atoms cannot be piled together and forced to form a periodic structure without mistakes being made. The 2nd Law of Thermodynamics demands this. Such mistakes seriously affect the overall properties of the solid. Thus, defeets in the lattice are probably the most important aspect of the solid state since it is impossible to avoid defects at the atomistic level. Two factors are involved ... 

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. 

The MCAT does not directly test your knowledge of the structure of solids beyond Ionic add molecular solids however, it is good to at least be aware dial atoms can form substances in many ways- Molecular solids are actually less common than other types of solids. There has been an MCAT passage on this topic. 

The halogen hydracids are all weaker than perchloric acid. Actually they are not ideal strong electrolytes, although they a pproach this behavior when water is used as the solvent. Certainly, these compounds differ distinctly from typical strong electrolytes such as potassium chloride and other neutral salts. The difference probably originates in the structure of the solid form. Neutral salts in the solid crystalline state possess a coordination lattice. Simple molecules do not exist in this type of lattice since the constituents of the salt are present. solely in the ionic form. Each ion is surrounded in a uniform manner by a definite number of other ions of opposite charge. Indeed it is no longer correct to speak of undissociated molecules in the solid state. 

The effective charge on a vesicle is only partly determined by the fixed charges, if any, at its membrane surface. Since a vesicle is not a solid colloidal particle we must consider also the effective charge due to the potential that arises as a consequence of the semipermeability of the boundary membrane and the difference in ionic activities between the interior of the vesicle and the external cytoplasmic medium. The charge can then be estimated from the actual structure of the electric field set up across the membrane and extending into the diffuse boundary layer outside the membrane wall. 

Structures of Ionic Compounds We write the formula of an ionic compound such as lithium fluoride simply as LiF, but this is really the empirical, or simplest, formula. The actual solid contains huge and equal numbers of... 

The experimental approach discussed in this article is, in contrast, particularly amenable to investigating solvent contributions to the interfacial properties 131. Species, which electrolyte solutions are composed of, are dosed in controlled amounts from the gas phase, in ultrahigh vacuum, onto clean metal substrates. Sticking is ensured, where necessary, by cooling the sample to sufficiently low temperature. Again surface-sensitive techniques can be used, to characterize microscopically the interaction of solvent molecules and ionic species with the solid surface. Even without further consideration such information is certainly most valuable. The ultimate goal in these studies, however, is to actually mimic structural elements of the interfacial region and to be able to assess the extent to which this may be achieved. 

Translation of ions within crystals is less frequently observed than is rotation. Perhaps one of the most interesting cases is that of silver iodide which may actually be said to melt in halves. When this solid is heated to 145.8° C, the crystal structure then changes and the ionic conductivity increases tremendously the iodide ions are hexagonally closest-packed below the transition temperature but at this temperature they rearrange to form a more open structure, and the silver atoms are allowed to move within the lattice. At 555° C, the network of iodide ions collapses, and the compound becomes a liquid. The solids Cul and Ag2Se show similar behavior. 

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