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Solids. Structures of Solids. Crystalline vs. Amorphous Crystalline solid : well-ordered, definite arrangements of molecules, atoms or ions. Most solids are crystalline: iron, sand, salt, sugar Amorphous solid : random arrangements of molecules, atoms, or ions.
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Structures of Solids Crystalline vs. Amorphous • Crystalline solid: well-ordered, definite arrangements of molecules, atoms or ions. • Most solids are crystalline: iron, sand, salt, sugar • Amorphous solid: random arrangements of molecules, atoms, or ions. • A few solids are amorphous: glass, rubber, cotton candy, butter • Amorphous solids usually solidify too fast for an orderly arrangement to form
Structures of Crystalline Solids Unit Cells • Crystalline solid: well-ordered, definite arrangements of molecules, atoms or ions. • Crystals have an ordered, repeated structure. • The smallest repeating unit in a crystal is a unit cell. • Unit cell is the smallest unit with all the symmetry of the entire crystal. • Three-dimensional stacking of unit cells is the crystal lattice.
Structures of Solids Unit Cells
Structures of Solids Three Types of Unit Cells • Primitive cubic • atoms at the corners of a simple cube • each atom shared by 8 unit cells; so 1 atom per unit cell • Body-centered cubic (bcc), • atoms at the corners of a cube plus one in the center of the body of the cube • corner atoms shared by 8 unit cells, center atom completely enclosed in one unit cell; so 2 atoms per unit cell
Structures of Solids Three Types of Unit Cells • Face-centered cubic (fcc), • atoms at the corners of a cube plus one atom in the center of each face of the cube • corner atoms shared by 8 unit cells, face atoms shared by 2 unit cells; so 4 atoms per unit cell
Structures of Solids Three Types of Unit Cells
Structures of Solids Crystal Structure of Sodium Chloride • Face-centered cubic lattice. • Two equivalent ways of defining unit cell: • Cl- (larger) ions at the corners of the cell, or • Na+ (smaller) ions at the corners of the cell. • The cation to anion ratio in a unit cell is the same for the crystal. In NaCl each unit cell contains same number of Na+ and Cl- ions. • Note the unit cell for CaCl2 needs twice as many Cl- ions as Ca2+ ions.
Structures of Solids Crystal Structure of Sodium Chloride
Structures of Solids X-Ray Diffraction • When waves are passed through a narrow slit they are diffracted. • When waves are passed through a diffraction grating (many narrow slits in parallel) they interact to form a diffraction pattern (areas of light and dark bands). • Efficient diffraction occurs when the wavelength of light is close to the size of the slits. • The spacing between layers in a crystal is 2 - 20 Å, which is the wavelength range for X-rays.
Structures of Solids X-Ray Diffraction
Structures of Solids X-Ray Diffraction • X-ray diffraction (X-ray crystallography): • X-rays are passed through the crystal and are detected on a photographic plate. • The photographic plate has one bright spot at the center (incident beam) as well as a diffraction pattern. • Each close packing arrangement produces a different diffraction pattern. • Knowing the diffraction pattern, we can calculate the positions of the atoms required to produce that pattern. • We calculate the crystal structure based on a knowledge of the diffraction pattern.
Bonding in Solids • There are four types of solid: • Molecular (formed from molecules): usually soft with low melting points and poor conductivity. • Covalent network (formed from atoms): very hard with very high melting points and poor conductivity. • Ions (formed form ions): hard, brittle, high melting points and poor conductivity. • Metallic (formed from metal atoms): soft or hard, high melting points, good conductivity, malleable and ductile.
Bonding in Molecular Solids • Intermolecular forces: dipole-dipole, London dispersion and H-bonds. • Weak intermolecular forces give rise to low melting points. • Room temperature gases and liquids usually form molecular solids at low temperature. • Efficient packing of molecules is important (since they are not regular spheres).
Bonding in Covalent Network Solids • Covalent Bonding. • Atoms held together in large networks. • Examples: diamond, graphite, quartz (SiO2), silicon carbide (SiC), and boron nitride (BN). • In diamond: • each C atom has a coordination number of 4; • each C atom is tetrahedral; • there is a three-dimensional array of atoms. • Diamond is hard, and has a high melting point (3550 C).
Bonding in Covalent Network Solids • In graphite • each C atom is arranged in a planar hexagonal ring; • layers of interconnected rings are placed on top of each other; • the distance between C atoms is close to benzene (1.42 Å vs. 1.395 Å in benzene); • the distance between layers is large (3.41 Å); • electrons move in delocalized orbitals (good conductor).
Bonding in Ionic Solids • Ions (spherical) held together by electrostatic forces of attraction: • The higher the charge (Q) and smaller the distance (d) between ions, the stronger the ionic bond. • There are some simple classifications for ionic lattice types:
Bonding in Ionic Solids • NaCl Structure • Each ion has a coordination number of 6. • Face-centered cubic lattice. • Cation to anion ratio is 1:1. • Examples: LiF, KCl, AgCl and CaO. • CsCl Structure • Cs+ has a coordination number of 8. • Different from the NaCl structure (Cs+ is larger than Na+). • Cation to anion ratio is 1:1.
Bonding in Ionic Solids • Zinc Blende Structure • Typical example ZnS. • S2- ions adopt a fcc arrangement. • Zn2+ ions have a coordination number of 4. • The S2- ions are placed in a tetrahedron around the Zn2+ ions. • Example: CuCl. • Fluorite Structure • Typical example CaF2. • Ca2+ ions in a fcc arrangement. • There are twice as many F- per Ca2+ ions in each unit cell. • Examples: BaCl2, PbF2.
Bonding in Metallic Solids • Metallic solids have metal atoms in hcp, fcc or bcc arrangements (only Po is simple cubic packing). • Coordination number for each atom is either 8 or 12. • Problem: the bonding is too strong for London dispersion and there are not enough electrons for covalent bonds. • Resolution: the metal nuclei float in a sea of electrons. • Metals conduct because the electrons are delocalized and are mobile.
