Kinetic Molecular Theory • Particles are always in motion. • Temperature is a measure of average kinetic energy of particles. • Intermolecular forces hold particles together. Stronger forces require more energy (higher temp.) to overcome.
Intermolecular Forces • Always weaker than chemical bond • Affect structure and state of matter
Dipole-Dipole Forces • Positive and negative ends of polar molecules attract each other. • About 1% as strong as covalent or ionic bonds • Weaken as distance between molecules increases
Hydrogen Bonding • Especially strong dipole-dipole force • Occurs when H bonds to a strongly electronegative atom—O, N, or F • Very strong because 1) molecule is very polar & 2) small size of H
H Bonding • Example—water • More pronounced in molecules formed from small atoms (dipoles can come closer) • High boiling point
London Dispersion Forces • Forces that exist in all atoms and molecules but that are significant only among Noble gases and nonpolar molecules • Result from temporary dipoles formed when electrons distribute themselves unevenly—can induce a dipole in a neighboring atom • VERY WEAK
London Dispersion Forces • Stronger in larger atoms or molecules due to the greater chance of the formation of instantaneous dipoles.
Physical Properties • Melting and boiling points are higher when IM attractions are stronger • More energy required to separate molecules
Which would have the higher boiling point & Why? • Cl2 or F2? • H2O or H2S? • SiBr2 or SBr2? • CH4 or C10H22? • O2 or NO?
States of Matter • Gases—weak IM forces (like London Dispersion Forces) • Liquids—intermediate IM forces • Solids—strong IM forces
Liquids & IM Forces • Surface tension—result of IM forces that resist an increase in surface area • Capillary action—result of cohesive forces within liquid and adhesive forces between liquid and tube
Liquids (cont’d) • Viscosity—the ability of a liquid to resist flow (resist change in shape) • All effects are higher with more polar molecules.
Solids • Amorphous—without definite structure • Crystalline—definite structures
Solids (Crystals) • Ionic solids—made of charged particles; ions at lattice points • Molecular solids—made of neutral particles; molecules at lattice points • Atomic solids—made of neutral particles; atoms at lattice points; 3 types
Ionic Solids • Ions at lattice points • Closest packed spheres • Arranged to minimize repulsions and maximize attractions • Conducts only when melted
Molecular Solids • Lattice positions occupied by molecules • Internal covalent bonds are strong, but intermolecular forces are weak • IM force: dipole/dipole if polar covalent bond; London dispersion forces (larger in larger molecules)
Atomic Solids • Network—directional covalent bonds; forms giant molecules (diamond, graphite,and silicon); highest melting points • Metallic—delocalized covalent bonds; atoms have closest packing structure; high melting points
Atomic Solids • Group 8A—Noble gases—London dispersion forces only; low melting points.
Network Atomic Solids • Strong, directional bonds • Form giant “molecules” • Typically brittle & poor conductors • Examples—carbon and silicon
Carbon Network • Follows a molecular orbital (not atomic orbital) model
Diamond • Tetrahedral--sp3 hybridized bonds stabilize structure • Large gaps exist between filled and unfilled molecular orbitals—hard for electrons to move—no conductivity
Graphite • Fused carbon rings form sheets • Trigonal planar—sp2 hybridized (1 p orbital remains unhybridized) • Delocalized electrons in orbital causes graphite to be conductive
Figure 10.22: The structures of diamond and graphite. In each case only a small part of the entire structure is shown.
Closest Packed Solids • aba pattern—alternating layers—atoms in 3rd layer lie directly above atoms in 1st layer—hexagonal unit cell—body centered • abca pattern—atoms in 1st and 4th layers are in line; 2nd & 5th layer; 3rd & 6th layer—face-centered cubic cell
Density of Closest Packed Solids • To calculate density, you need to know: • MASS • VOLUME
Mass • Figure out how many atoms in one unit cell • Multiply by molar mass • Divide by Avogadro’s number • You now know a mass in grams
Face-Centered Cubic Unit Cell If these are atoms of calcium, what is the mass of the cell?
Volume • Determine the length of one side of the cube by using the atomic radius (varies depending on type of unit cell) • Cube the side length.
Simple Cubic--aaa If the atomic radius of this atom is 122 pm, what is the volume?
Body-Centered Cubic--aba If the atomic radius of the atom is 246 pm, what is the volume of the cell?
Face-Centered Cubic--abca If the atomic radius is 291 pm, what is the volume of the cell?
Sample Problem • Silver crystallizes in a face-centered cubic closest packed structure. The radius of the silver atom is 144 pm. Calculate the density of silver.
Bonding in Metals • Strong, non-directional bonds • Atoms are hard to separate but easy to move. • “Electron sea” model • Mobile electrons carry heat or electricity easily
Band Model or Molecular Orbital Model • Electrons travel around metal crystal in a molecular (instead of atomic) orbitals • Result is a continuum of levels that eventually merge to form a band.
Figure 10.19: The molecular orbital energy levels produced when various numbers of atomic orbitals interact.
Band or MO Model • Empty orbitals close in energy exist. • Electrons are very mobile into and out of these similar-energy orbitals--CONDUCTIVITY.
Semiconductors • Some electrons can cross the “energy gap” between molecular orbitals—somewhat conductive • Higher temperatures result in more electrons’ being able to reach conductive bands
Doping • Adding other elements with one more or one less electron than a semiconductor can increase conductivity
n-type semiconductor • An element with one more valence electron is added • More valence electrons are available to move into conduction bands • What could be used to dope Si?
p-type semiconductor • An element with one less valence electron is added • The absence of a valence electron creates a hole through which electrons can travel