1 Chapter 13 The Chemistry of Solids Jeffrey Mack California State University, Sacramento
2 Metallic & Ionic Solids
3 Crystal Lattices Regular 3-D arrangements of equivalent LATTICE POINTS in space. Lattice points define UNIT CELLS Unit cells are the smallest repeating internal unit that has the symmetry characteristic of the solid.
4 Properties of Solids 1. Molecules, atoms or ions locked into a CRYSTAL LATTICE. 2. Particles are CLOSE together. 3. These exhibit strong intermolecular forces 4. Highly ordered, rigid, incompressible ZnS, zinc sulfide
5 Types of Solids Type: Examples: Forces: Ionic Compounds NaCl, BaCl 2, ZnS Ion-Ion (ionic bonding) Metals Fr, Al Metallic Molecular Ice, I 2, C 12 H 22 O 11 Dipole-Dipole ot Induced Dipoles Network Diamond, Graphite Extended Covalent bonds Amorphous Glass, Coal Covalent; directional electron-pair bonds
6 Network Solids Diamond Graphite
7 Cubic Unit Cells There are 7 basic crystal systems, but we will only be concerned with CUBIC form here. All sides equal length All angles are 90 degrees 1/8 of each atom on a corner is within the cube 1/2 of each atom on a face is within the cube 1/4 of each atom on a side is within the cube
11 Simple Cubic Unit Cell Each atom is at a corner of a unit cell and is shared among 8 unit cells. Each edge is shared with 4 cells Each face is part of two cells.
12 Atom Packing in Unit Cells Assumes atoms are hard spheres and that crystals are built by PACKING these spheres as efficiently as possible.
13 Atom Packing in Unit Cells
14 Crystal Lattices Packing of Atoms or Ions FCC is more efficient than either BC or PC. Leads to layers of atoms.
15 Crystal Lattices Packing of Atoms or Ions Packing of C 60 molecules. They are arranged at the lattice points of a FCC lattice.
16 Classifications of Solids Solids can be classified on the basis of the bonds that hold the atoms or molecules together. This approach categorizes solids as either: molecular Network (covalent) ionic metallic
17 Molecular Solids Molecular solids are characterized by relatively strong intramolecular bonds between the atoms that form the molecules The intermolecular forces between these molecules are much weaker than the bonds. Because the intermolecular forces are relatively weak, molecular solids are often soft substances with low melting points. Examples: I 2 (s), sugar (C 12 H 22 O 11 ) and Dry Ice, CO 2 (s)
18 Network (Covalent) Solids In Network solids, conventional chemical bonds hold the chemical subunits together. The bonding between chemical subunits is identical to that within the subunits resulting in a continuous network of chemical bonds. Two common examples of network solids are diamond (a form of pure carbon) and quartz (silicon dioxide). In quartz one cannot detect discrete SiO 2 molecules. Instead the solid is an extended threedimensional network of...-si-o-si-o-... bonding.
19 Ionic Solids Ionic solids are salts, such as NaCl, that are held together by the strong force of attraction between ions of opposite charge. q( + ) q( -) F» 2 r Because this force of attraction depends on the square of the distance between the positive and negative charges, the strength of an ionic bond depends on the radii of the ions that form the solid. As these ions become larger, the bond becomes weaker.
20 Metallic Solids In Molecular, ionic, and covalent solids the electrons in these are localized within the bonding atoms. Metal atoms however don't have enough electrons to fill their valence shells by sharing electrons with their immediate neighbors. Electrons in the valence shell are therefore shared by many atoms, instead of just two. In effect, the valence electrons are delocalized over many metal atoms. Because these electrons aren't tightly bound to individual atoms, they are free to migrate through the metal. As a result, metals are good conductors of electricity.
21 Bonding in Ionic Compounds: Lattice Energy The energy of an ion pair (cation/anion) is described by Coulombs law: U ion pair ( n + e - )( n - e + ) = C d n + = cation charge, n = anion charge d = distance between ion centers lattice U is the energy of formation of one mole of the solid crystaline compound from its ions in the gas phase. + - M ( g) + X ( g) MX( s)
22 Lattice Energy The Lattice Energy of a salt is dependant upon the charge and size of the ions. U ion pair = C ( n + e - )( n - e + ) d
23 Lattice Energy Calculation of lattice energy via the Born Haber cycle, an application of Hess s law.
24 Problem: Calculate the molar enthalpy of formation, f H, of solid lithium fluoride from the lattice energy and following thermochemical data.
25 Problem: Calculate the molar enthalpy of formation, f H, of solid lithium fluoride from the lattice energy and following thermochemical data. Solution: Approach this problem using Hess s Law. You need to find the enthalpy for the reaction: Li(s) + ½ F 2 (g) LiF(s)
26 Problem: Calculate the molar enthalpy of formation, f H, of solid lithium fluoride from the lattice energy and following thermochemical data. Start by drawing the Born-Haber cycle for the reaction:
27 Problem: Calculate the molar enthalpy of formation, f H, of solid lithium fluoride from the lattice energy and following thermochemical data. Start by drawing the Born-Haber cycle for the reaction: Li + (g) + F (g) IE EA Li(g) F(g) sub H D o Li(s) + ½ F 2 (g) LiF(s)
28 Problem: Calculate the molar enthalpy of formation, f H, of solid lithium fluoride from the lattice energy and following thermochemical data. Using Hess s Law, the enthalpy of formation is found by: Li + (g) + F (g) IE EA Li(g) F(g) sub H D o Li(s) + ½ F 2 (g) LiF(s) f H o = sub H + I 1 + D o + EA + lattice U
29 Problem: Calculate the molar enthalpy of formation, f H, of solid lithium fluoride from the lattice energy and following thermochemical data. f H o = sub H + IE + D o + EA + lattice U Li(s) Li(g) sub H = kj/mol Li(g) Li + (g) + e IE = kj/mol ½ F 2 (g) F(g) D o = kj/mol F(g) + e F (g) EA = kj/mol Li + (g) + F (g) LiF(s) lattice U = 1037 kj/mol f H = = 607 kj/mol
30 Phase Changes Involving Solids Melting: Conversion of Solid into Liquid The melting point of a solid is the temperature at which the lattice collapses into a liquid. Like any phase change, melting requires energy, called the enthalpy of fusion. Energy absorbed as heat on melting = enthalpy of fusion fusion H (kj/mol) Energy evolved as heat on freezing = enthalpy of crystallization fusion H (kj/mol) Enthalpies of fusion can range from just a few thousand joules per mole to many thousands of joules per mole.
31 Enthalpies of Fusion Are a Function of Intermolecular Forces
32 Phase Changes Involving Solids Sublimation: Conversion of Solid into Vapor Molecules can escape directly from the solid to the gas phase by sublimation Solid Gas Energy required as heat = sublimation H Sublimation, like fusion and evaporation, is an endothermic process. The energy required as heat is called the enthalpy of sublimation.
33 Sublimation Sublimation entails the conversion of a solid directly to its vapor. Here, iodine (I 2 ) sublimes when warmed.
34 Transitions Between Phases: Phase Diagrams Phase diagrams are used to illustrate the relationship between phases of matter and the pressure and temperature.
35 Phase Diagram for Water Liquid phase Solid phase Gas phase
36 Phase Equilibria Water Solid-liquid Gas-Liquid Gas-Solid
37 Triple Point Water At the TRIPLE POINT all three phases are in equilibrium.
38 Phases Diagrams: Water T( C) P(mmHg) Normal boil point (at 1atm): Normal freeze point (at 1atm): Triple point:
39 Phases Diagrams: Water Water has its maximum density at 4 C, in the liquid phase. Most substances have a maximum density in the solid phase. Hydrogen bonding accounts for water s deviation from normal behavior.
40 Phases Diagram: Water At constant temp, an increase in pressure can bring about a phase change from solid to liquid!
41 Phases Diagrams: Water At constant temp, an increase in pressure can bring about a phase change from solid to liquid! This occurs when the blade of an ice skate runs on the ice. Ice skaters actually ride on a film of water, not the ice!
42 CO 2 Phase Diagram Notice the CO 2 has a forward slope of the solid/liquid boundary. This is seen because CO 2 does not exhibit hydrogen bonding.
43 CO 2 Phases Separate phases Increasing pressure More pressure Supercritical CO 2
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