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In the solid state, the particles are in fixed positions. When those positions are arranged in a repeating pattern, the solid is called crystalline. When the pattern does not repeat, the solid is amorphous.

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The smallest repeating volume of a solid is a unit cell. One type of unit cell is a face-centered cubic. A face-centered cubic (fcc) is the same length on each side and has a particle at each corner and one on each face (Figure 10.4). A body-centered cubic (bcc) is also the same length on each side with a particle at each corner. Instead of particles on the faces, it has a particle in the center of the cube (Figure 10.5A). A simple cubic just has particles on each corner, and the length of every edge is the same (Figure 10.5B).

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Because ionic solids are extended arrays of anions and cations, their formula is always the simplest ratio of those ions. Because a unit cell is representative of the entire structure, the ratio of ions in the unit cell is the same as the ratio in the overall structure. Since unit cells are traditionally drawn with the center of a particle in each corner, only one-eighth of a corner particle is actually in the unit cell. Particles on the edge count one-fourth, particles on the face count one-half and particles completely within the unit cell count as one.

The size (edge length) of a unit cell depends on the size of the ions and their arrangement. Anions are usually larger than cations, so they normally determine the overall structure with the cations fit in between. In a face-centered cubic, anions actually touch in a diagonal along the face. The edge length can be determined from the radii of these anions. Alternately, the edge length can be used to determine the ionic radii. In a body-centered cubic, ions touch in a diagonal from one corner, across the center of the cell to the opposite corner. In a simple cubic, the particles touch along the edge.

Another way of describing the arrangement of particles in a solid is the close packing model. In this description, particles are arranged as close together as possible and to make maximum use of the space. Each layer has the particles arranged hexagonally around each other and the layers offset (Figure 10.6). A cubic close-packed (ccp) structure has three different layers before the pattern repeats. A cubic close-packed structure makes a face-centered cubic unit cell. A hexagonal close-packed (hcp) structure repeats every other layer. Both of these patterns are very efficient and use 74% of the space. In ionic compounds, the larger ion (normally the anion) is close packed and the smaller ion (cation) is in the holes (the other 26%). A simple cubic packing structure repeats every layer. This is less efficient packing, using only 52% of the available space, but useful for larger ions.

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There are two types of holes in a close-packed structure. Octahedral holes are made from six anions arranged in an octahedral pattern. These are larger than the tetrahedral holes made from four atoms arranged in a tetrahedral. In a unit cell of a closed-packed arrangement there are eight tetrahedral holes and four octahedral holes. Cations fill holes based on their size and the stoichiometry of the structure.

Some structures are so common that they are named, usually for the most common example of that structure. The rock salt or halite structure (NaCl) has anions arranged in a face-centered cubic unit cell with the cations in all the octahedral holes (Figure 10.4). The zinc blende or sphalerite (ZnS) structure is a face-centered cubic of anions with cations in half the tetrahedral holes (Figure 10.10). The fluorite structure (CaF₂) has cations in a face-centered cubic and anions in all the tetrahedral holes. (Figure 10.11) An antifluorite structure just reverses the positions of the anions and cations of the flourite structure. The cesium chloride structure is a simple cubic of anions superimposed over a simple cubic of cations (Figure 10.12).

Ionic solids are held together by the electrostatic attraction of anions and cations. Another type of solid, a network solid, is held together by an extended network of covalent bonds. The formula for these solids is also the simplest ratio of atoms. The arrangement of these solids determines the physical properties. For example, the allotropes (Chapter 1) of carbon are graphite, diamond, and fullerene. Each is a different arrangement of carbon atoms with very different physical properties. There are also many varied arrangements of silicon and oxygen. These are polymorphs, having the same chemical formula and different chemical and physical properties. In silica (SiO₂), each silicon is bonded tetrahedrally to four oxygens, and each oxygen is bonded to two silicon atoms. If an electron replaces the silicon bond of an oxygen atom, that oxygen will have a –1 formal charge, and an ion called a silicate is formed. There are many types of silicates, whose properties depend on the number of oxygen atoms that are charged instead of bonded and the cations that pair with the negative charge.

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A third type of solid is a metallic solid. There are two theories to explain bonding in metals. In one theory, the metal atom acts a cation and an electron. The electron attracts the cation part of other metal atoms, and that cation attracts other electrons. The metal atoms are held together by this "sea of electrons" or "electron glue." The other bonding theory is based on molecular orbital theory. However, since metals solids have an enormous number of atoms (rather than two), there are practically an infinite number of molecular orbitals. These orbitals are so close together in energy that electrons move freely within that band of energy. Therefore this theory of bonding is called band theory. Band theory can be used to explain the bonding of any type of solid and is especially useful in explaining conductivity.

The colors of many solids are due to the crystal field splitting of the d orbitals. When a transition metal is surrounded by anions (as in a crystal), the d orbitals are no longer all the same energy. The anions increase the energy of the orbitals closest to it. The other orbitals must then decrease in energy. If the cation is in an octahedral hole, the and orbitals increase in energy and the d_xy, d_xz, and d_yz orbitals decrease in energy. The energy difference is called the crystal field splitting energy (_o). In a tetrahedral hole, the orbitals split in the opposite way with the and orbitals decreasing in energy and the d_xy, d_xz, and d_yz orbitals increasing. (The crystal field splitting is _t in this case.) If the anions were arranged in a square planar orientation, the octahedral splitting is split again with the and the d_xy slightly higher than in an octahedral orientation.

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Since _t is a fairly small energy, it requires less energy for an electron to go to the higher-energy d level than it does for the electron to pair up in the lower level. This is sometimes true for an octahedral field. Those compounds where the electron prefers the higher energy level to pairing in the lower level are called high-spin (or weak field) compounds. If the electron pairs up in lower d orbital, it is called a low-spin (strong field) compound. Compounds in a tetrahedral orientation are always high spin; compounds in a square planar configuration are always low spin.

The energy required for an electron to move from the lower d orbital to the higher one is often in the same range as visible light. If that is the case, the electron will absorb the light that is the same energy as the crystal field splitting. The color seen is a mixture of the colors that are not absorbed.