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Chapter 9: Structure and Energetics of Solids

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    440862
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    In the chemistry of molecular compounds, we are accustomed to the idea that properties depend strongly on structure. For example we can rationalize the polarity of the water molecule based on its shape. We also know that two molecules with the same composition (e.g., ethanol and dimethyl ether) have very different properties based on the bonding arrangements of atoms. It should come as no surprise that the properties of extended solids are also connected to their structures, and so to understand what they do we should begin with their crystal structures. Most of the metals in the periodic table have relatively simple structures and so this is a good place to begin.

    clipboard_e32e3b74ea300fc2ec237429e9f87a79b.png
    Figure \(\PageIndex{1}\): Over \(\frac{2}{3}\) of the elements in the periodic table are metals. All elemental metals (except the three - Cs, Ga, Hg - that are liquid at or near room temperature) are crystalline solids at room temperature, and most have one of three simple crystal structures.

    With ionic solids we "graduate" from simple metal structures based on sphere packings to more complex structures where there are both anions and cations. In this chapter we will try to systematize the structures of inorganic solids - metal oxides, halides, sulfides, and related compounds - and develop some rules for which structures to expect based on electronegativity differences, hard-soft acid-base rules, and other periodic trends. We will see that many of these structures are based on the simple packing of spheres.

    clipboard_eb85164193d4eb90f6592c32ba78b455c.png
    Figure \(\PageIndex{2}\): The morphology of twinned crystals of iron pyrite (FeS2) is related to the underlying cubic symmetry of the unit cell. Like NaCl, the pyrite crystal structure can be thought of as a face-centered cubic array of anions (S22-) with cations (Fe2+) occupying all the octahedral holes.

    Inorganic solids often have simple crystal structures, and some of these structures are adopted by large families of ionic or covalent compounds. Examples of the most common structures include NaCl, CsCl, NiAs, zincblende, wurtzite, fluorite, perovskite, rutile, and spinel. We will develop these structures systematically from the close packed and non-close packed lattices.

    Learning Objectives
    • Determine the formula of an ionic solid based on the unit cell
    • Predict relative lattice energies based on ion sizes and charges
    • Understand how lattice energy plays a role in solubility
    • Explain how particle size affects the material properties in nanomaterials

    Thumbnail image shows the unit cell of the zincblende crystal structure

    • Section 9.1: Crystal Structures and Unit Cells
      Solid state structures can be visualized as spheres packed into a box. There are two closest packed arrangements which fill the largest possible space in the box, as well as arrangements where the spheres are packed less densely. The structure of a larger crystal can be represented by a unit cell, the smallest repeating unit.
    • Section 9.2: Energetics of Ionic Solids- Lattice Energy
      The energetics of an ionic solid can be approximated reasonably well by considering the attraction and repulsion between ions in the crystal
    • Section 9.3: Lattice Enthalpies and Born Haber Cycles
      Lattice enthalpy is a measure of the strength of the forces between the ions in an ionic solid. The greater the lattice enthalpy, the stronger the forces. Those forces are only completely broken when the ions are present as gaseous ions, scattered so far apart that there is negligible attraction between them. You can show this on a simple enthalpy diagram.
    • Section 9.4: Lattice Energy and Solubility
      Lattice energies can also help predict compound solubilities. For a solid to be soluble, the energy required to break the lattice must be offset by the solvation of the ions or molecules in solvent
    • Section 9.5: Bonding in Metals and Semicondoctors
      Band theory is a theory to describe bonding in metals and other extended solids. It is an extension of molecular orbital theory, where large numbers of atomic orbitals mix together to form very closely spaced groups of molecular orbitals called bands.
    • Section 9.6: Nanomaterials
      Ordinary elements and inorganic compounds acquire different properties when their dimensions are less than ~100 nm. The properties of these nanomaterials, such as color and reactivity are strongly affected by the size of the particle.
    • Section 9.1b: Ionic Radii and Radius Ratios


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