Liquids and Solids

1
Liquids and Solids

• Chapter 13

2
Solids, Liquids, and Gases

• Solids liquids are condensed states
• atoms, ions, molecules are close to one another
• highly incompressible
• Liquids gases are fluids
• easily flow
• Intermolecular attractions in liquids solids
  are strong
• Look at Table 13-1 for a general description of
  characterisitics

3
Kinetic-Molecular Description of Liquids and
Solids

• When a gaseous sample is cooled or compressed the
  rapid, random motion of the molecules decrease.
  The attraction between molecules becomes
  significant. When the attractive forces overcome
  the reduced kinetic energies, condensation occurs
  (or the gas turns to a liquid).
• In the liquid phase the particles are close.
  Very little space is unoccupied (discuss).
• The particles, however, still have sufficient
  KEs to partially overcome the attractive forces.
  Therefore, liquids are considered fluids and can
  take the shape of a container.

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Kinetic-Molecular Description of Liquids Solids

• Liquids that diffuse into one another are
  miscible (i.e. one liquid is soluble in the
  other)
• Water and methanol/gasoline and motor oil
• Immiscible liquid do not diffuse into one another
  (i.e. one liquid is not soluble in the other)
• Water and hexane/water and gasoline

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Kinetic-Molecular Description of Liquids Solids

• Cooling the liquid lowers the KE even more.
  Shorter-range attractive force become important
  and the liquid solidifies. Particles in a solid
  cannot move freely past one another as they can
  in a liquid. This is why solids have definite
  shapes and volumes. Solids are essentially
  incompressible.
• If diffusion occurs in solids, it is very slow.

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Intermolecular Attractive Forces

• Intermolecular attractive forces are the forces
  between individual particles of a substance.
• Generally, these forces are very weak compared to
  intramolecular forces (e.g. covalent and ionic
  bonds)
• Covalent bonding and attractive forces in H2O(l)
• Important physical properties such as boiling
  points, vapor pressure, heat of vaporation,
  melting points, and heat of fusion depend on the
  strength of these intermolecular attractive
  forces.
• If intermolecular attractive forces did not
  exist, solids and liquids would not exist (only
  gases)

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Ion-Ion Interactions

• The force of attraction between two charged
  particles can be determined from Coulombs Law
• where q and q- are the charges on the particles
  and d is the distance between them.

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Ion-Ion Interactions

• The energy of attraction between charged
  particles is given by
• The energy of attraction is large for ionic
  compounds due to the charged particles that are
  close together when substance is a solid.
• The melting points for ionic compounds is
  relatively high.
• For a given substance, the separation between
  particles in the solid is less than the
  separation in liquids.
• Energy of attraction is greater in the solid
  phase (generally)

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Ion-Ion Interactions

• Ionic compounds that possess multiply-charged
  ions (e.g. Al3 and O2-) usually have higher
  melting points than ionic compounds that posses
  singly-charged particles. Why (two reasons)?
• Arrange the following ionic compounds in the
  expected order of increasing melting and boiling
  points.
• NaF, CaO, CaF2

10
Dipole-Dipole Interactions

• Permanent dipole-dipole interactions occur
  between polar covalent molecules because of the
  attraction between ?- and ? on different
  molecules.
• Generally, these forces are not as strong as
  ion-ion interactions
• Attraction between partial charges
• Dipole-dipole forces vary as 1/d4 instead of 1/d2
• Decrease faster

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Dipole-Dipole Interactions
The partial positive charge on the hydrogen
atoms is attracted to the partial negative charge
on the nitrogen atoms. Note Dipole-dipole
interactions are dependent on temperature. Why?
12
Hydrogen Bonding

• Hydrogen bonding is a special type of strong
  dipole-dipole interaction
• Occurs between covalent molecules containing H
  and of the three small, highly electronegative
  elements-F, O, or N.
• One molecule most possess a H atom attached to
  one of these highly electronegative atoms.
• The other molecule most possess one these highly
  electronegative atoms.

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Hydrogen Bonding
The partial negative charge of one molecule is
attracted to the partial positive charge of
another molecule. The small sizes of F, O, and N
and their high electronegativities concentrate
the electrons around these atoms.
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Hydrogen Bonding

• Typical hydrogen-bond energies are greater than
  dipole-dipole energies
• 15-20 kJ/mol for hydrogen-bond energies
• 4 kJ/mol for dipole-dipole energies
• 400 kJ/mole for ion-ion interaction energies
• Hydrogen bonding is responsible for the high
  boiling points of water and methanol.

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Dispersion Forces

• These are attractive forces that are present in
  all types of molecules.
• Dispersion forces are weak in small molecules.
• They are important at extremely small distances
  which vary as 1/d7.
• They are the only attractive force present in
  symmetrical nonpolar substances such as Cl2 and
  monatomic species.

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Dispersion Forces

• Dispersion forces result from the attraction of a
  positively charged nucleus to the electron cloud
  of another atom in nearby molecules. As a
  result, temporary dipoles are induced in the
  neighboring atoms or molecules.
• The magnitude of the temporary dipole increases
  with increasing size of the electron cloud (or
  size of the molecule). The larger electron cloud
  is more diffuse and easily distorted. Adjacent
  molecules are polarized by the adjacent nuclei.
  Polarizibility increases with increasing sizes of
  molecules.
• Dispersion forces are, therefore, stronger for
  molecules that are larger or have more electrons.

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Dispersion Forces for Argon
Dispersion forces can also exist between
cations/anions and polarizable atom or molecule.
Examine Figure 13-6.
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Dispersion Forces

• Some trends observed in increasing boiling points
  can largely be attributed to dispersion forces
  (Figure 13-5).
• CH4, SiH4, GeH4, and SnH4
• HCl, HBr, and HI
• Understand Table 13-3.
• Heat of vaporization measures the energy required
  to overcome attractive forces in the liquid.

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The Liquid State

• Viscosity - the resistance to flow of a liquid.
• Generally, the higher the attractive forces in a
  liquid, the greater the viscosity.
• Water versus honey or Karo syrup.
• Pentane versus dodecane
• Viscosity decreases with increasing temperature.
  Why?

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The Liquid State

• Surface Tension - measure of the inward forces
  that must be overcome to expand the surface area
  of a liquid.
• Molecules at the surface are attracted unevenly.
• Water bugs and floating razor blades (or needles)
• Demo Razor blade or needle

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Liquid State

• Capillary action - ability of a liquid to rise or
  fall in a tube.
• Cohesive forces forces holding a liquid
  together.
• Adhesive force forces between a liquid and
  another surface
• Capillary action
• Capillary rise implies adhesive forces gt cohesive
  forces
• Capillary fall implies cohesive forces gt adhesive
  forces
• Water is attracted to a glass capillary tube due
  to attractive forces between the partial negative
  oxygens on the surface of the glass and the
  partial positive charges on the hydrogen atoms.

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Liquid State

• Capillary Action
• The meniscus of water has a concave shape due to
  the strong adhesive forces between the water
  molecules and the glass graduated cylinder.
• The meniscus of Hg has a convex shape because the
  cohesive forces are much stronger than the
  adhesive forces (Figure 13-9).

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Liquid State

• Evaporation process by which molecules escape
  from the surface of a liquid.
• Only the molecules at the surface with sufficient
  KE will be able to escape the attractive forces
  present in the liquid. Therefore, the rate of
  evaporation is proportional to ________.

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Liquid State

• Evaporation
• As the faster molecules leave the liquid, the
  flower molecules are left behind. What will this
  do to the temperature?
• This process is termed as the cooling by
  evaporation.
• DEMO Thermometer in acetone

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Liquid State

• Evaporation and condensation
• In an open container/beaker, all the water that
  is present will eventually escape into the
  gaseous phase (Figure 13-10).
• What if the beaker is sealed? What will happen?
  After molecules enter the gas phase they may be
  recaptured by the liquid by collisions. This
  process is called _______. At some point in
  time, the amount of gaseous molecule leaving the
  gaseous phase will equal the amount reentering
  the liquid phase. This is termed as dynamic
  equilibrium.

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Liquid State

• Vapor pressure the partial pressure of vapor
  molecules above the surface of a liquid at
  equilibrium.
• Vapor pressure increases with temperature. Why?
  Look at Figure 13-13.
• Vapor pressure decreases with increasing
  attractive forces. Why? Look at Table 13-4.
• Hydrogen bonding, dipole-dipole, and dispersion
  forces.
• DEMO H2O and ethyl ether (add H2O to ethyl
  ether)

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Liquid State

• As the temperature is increased, the vapor
  pressure increases until the liquid boils.
• Boiling point temperature at which the vapor
  pressure equals the external pressure.
• The boiling point of H2O is less in Rexburg.
  Why?
• Normal boiling point temperature at which the
  vapor pressure of a liquid is equal to 1 atm.
• Boiling point at 1 atm
• What if the external pressure is lower over
  water?
• DEMO Evacuate a flask filled with water

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Liquid State

• Distillation a method used to separate
  components in a solution based on differences in
  boiling point temperatures.
• Separation of liquids in a solution
• Purification of water (distilled water)
• Separate impurities (e.g. ions) from tap water
• Describe the basics of operation (Figure 13-14)
• Of course, this can be related to vapor pressures.

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The Liquid State

• Heat transfer involving liquids
• Specific heat or molar heat capacity of a liquid
  is the amount of heat that must be added to 1
  gram or 1 mole to raise the temperature by 1?C.
• How much heat is released by 200 g of H2O as it
  cools from 85.0oC to 40.0oC? The specific heat
  of water is 4.184 J/goC.
• Review in Chapter 1 if needed.

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The Liquid State

• The specific heat equation involves changing the
  temperature without a change in state. For a
  liquid, the temperature increases until the
  boiling point is acquired. The temperature
  remains at the boiling point temperature until
  all the liquid has been converted into the gas.
• Boiling water is always 100?C at sea level
• If the temperature is not increasing, what is all
  the heat/energy being used for when the liquid is
  boiling?
• Boiling water on a hotplate or a stove

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The Liquid State

• Answer The heat is providing the energy
  necessary to break up the intermolecular
  attractive interactions present in the liquid.
• Molar heat of vaporization amount of heat that
  must be added to one mole of the liquid at its
  boiling point to convert it to vapor
• There is no change in temperature
• How does heat of vaporization relate to
  intermolecular attractive forces and vapor
  pressure (Table 13-5)
• How does this relate to the cooling effect of
  perspiration?

32
The Liquid State

• Condensation is the reverse of vaporization/evapor
  ation.
• Heat of condensation is the amount of heat that
  must be removed from a vapor to condense it.
• Heat is released when the substance condenses.

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Liquid State

• Calculate the amount of heat necessary to convert
  125 grams of water at 25.0?C to steam at 100?C.
• How many joules of energy must be absorbed by 500
  g of H2O at 50.0oC to convert it to steam at
  120oC? The molar heat of vaporization of water
  is 40.7 kJ/mol and the molar heat capacities of
  liquid water and steam are 75.3 J/mol oC and 36.4
  J/mol oC, respectively.
• If 45.0 g of steam at 140oC is slowly bubbled
  into 450 g of water at 50.0oC in an insulated
  container, can all the steam be condensed?

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Liquid State

• Many trends in physical properties can be
  explained form the strength of intermolecular
  attractive forces in the liquid.
• Table 13-6 (understand these trends based on
  intermolecular attractive forces)
• Arrange the following substances in order of
  increasing boiling points.
• C2H6, NH3, Ar, NaCl, AsH3
• Discuss trends on Figure 13-5 (others as
  required)
• Generally, if a group of molecules possess the
  same type of attractive forces, the boiling point
  increases with size/molecular weight of the
  compound.

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The Clausius-Clapyron Equation

• The equation relates the vapor pressure of
  liquids at different temperatures.
• The vapor pressure is equal to 760 at the normal
  boiling point.
• Appendix E and Table 13-5 produce some normal
  boiling point temperatures (measured at 1 atm).
  These values could be used in the equation above.
• Also, the vapor pressure at the boiling point is
  equal to the atmospheric pressure

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The Clausius-Clapyron Equation

• At what temperature is the vapor pressure of
  water equal to the atmospheric pressure at sea
  level? Would the vapor pressure of boiling water
  in Rexburg be higher or lower than boiling water
  at sea level? Why?
• In Rexburg, the normal atmospheric pressure is
  652 torr. At what temperature does water boil in
  Rexburg?
• What will be the vapor pressure of ethyl alcohol
  at room temperature (298 K)?
• Remember to use the appropriate R.

37
The Solid State

• The melting point/freezing point of a substance
  is the temperature at which the solid and liquid
  phases coexist in equilibrium.
• The normal melting point is the melting point at
  one atmosphere
• Changes in pressures have very small effects on
  melting point

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Heat Transfer Involving Solids

• When heat is added to a solid the temperature
  increases until the melting point is acquired.
  Additional heat is required to convert the solid
  to a liquid (melting). Why? Heat that is added
  during the melting process does not increase the
  temperature. Temperature will rise after the
  solid is all converted to the liquid.

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Heat Transfer Involving Solids

• Molar heat of fusion (?Hfus kJ/mol) is the
  amount of heat required to melt one mole of a
  solid at its melting point. This can be
  converted to heat of fusion which is (kJ/g)
• Heat of fusion depends on _________(Table 13-7)
• Molar heat of solidification/crystallization is
  equal in magnitude (but opposite sign) to the
  molar heat of fusion.
• 6.02 kJ is absorbed when 1 mole is melted and
  6.02 kJ is released when 1 mole is
  frozen/solidifies.

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Heat Transfer Involving Solids
41
Heat Transfer Involving Solids

• Calculate the amount of heat required to convert
  150.0 g of ice at -10.0oC to water at 40.0oC.
  The specific heat of ice is 2.09 J/goC
• Calculate the amount of heat required to convert
  75.0 g of solid ethanol at 117.0?C to gaseous
  ethanol at 95.0?C.
• Table 13-5, 13-7 and and Appendix E

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Sublimation and Vapor Pressures of Solids

• Sublimation is the process by which a solid forms
  a gas (vaporizes) without passing through the
  liquid phase.
• Dry ice (CO2)
• Deposition is the reverse process by which a
  vapor forms a solid without passing through the
  liquid phase.
• Discussion with chemical vapor deposition
• Note Solids have vapor pressures, but they are
  generally very small

43
Phase Diagrams (P versus T)

• Phase diagrams illustrate a particular phase or
  state that is present under specific
  temperature-pressure conditions.
• Temperature is on the y-axis and pressure is on
  the x-axis.
• Using a phase diagram, changes in state can be
  determined when changing pressure and/or
  temperature.

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Phase Diagrams (P versus T)

• Phase boundaries
• Line AC liquid and gas phases coexist at
  equilibrium
• Line AB liquid and solid phases coexist at
  equilibrium
• Line AD gas and solid phases coexist at
  equilibrium
• Triple point (A) at this T and P all three
  phases coexist at equilibrium

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Phase Diagrams (P versus T)

• Notice the negative AB line slope. This is
  unique to H2O. What happens upon compressing a
  solid (up from point J)
• Is there a pressure where sublimation could
  occur?
• Travel a few other lines.

46
Phase Diagrams (P versus T)

• When the temperature is increased at normal
  pressures the solid goes directly to the gas
  phase. At what pressure would the solid go
  through the liquid phase.
• Notice the positive slope on the solid-liquid
  phase boundary. What does this mean?

47
Amorphous and Crystalline Solids

• Amorphous solids do not possess well-ordered
  structures.
• Some can also be characterized as glasses since
  they flow slowly.
• Windows
• Melting extend over a large range.
• Rubber, plastics, and amorphous sulfur
• Crystalline solids have well-defined structures
  consisting of repeating units.
• Repeating unit can be observed upon shattering
  the crystal.
• Possess distinct melting points.

48
Structures of Crystals

• Crystals contain regularly repeating structures.
• Unit cell is the smallest repeating unit of a
  crystal. A unit cell is the fundamental box that
  describes the arrangement of particles in a
  crystal. These unit cells are stacked in three
  dimensions to produce a crystal. The arrangement
  of these unit cells fit into one of seven crystal
  systems. (Table 13-9).
• Crystals have the same symmetry as the unit cells
  since the crystals are built from multiple units
  of these cells.

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Structures of Crystals
Look on page 513. Notice that the structures of
the unit cells are very similar to the actual
structure of the crystal. Why?
50
Structures of Crystals

• Creating the crystal from the unit cells
• Consider the corner of a unit cell as a lattice
  point. In three dimensions, this lattice point
  is shared by eight unit cells. If an object were
  present at this lattice point, it would be
  equally shared by all eight unit cells (1/8 in
  each). How many objects would be in each unit
  cell?
• DEMO Use unit cell building blocks to
  illustrate.

51
Structures of Crystals

• In a simple or primitive lattice structure
  (discussed previously), only the corners were
  occupied by objects. In other crystal types,
  objects may also occupy other positions in the
  unit cell.
• Simple cubic (no additional objects)
• Body-centered cubic (bcc) another object
  occupies the center of the unit cell
• Face-centered cubic (fcc) another six object
  occupies the middle of each of the six square
  faces of the cube

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Structures of Crystals
How many objects/atoms per unit cell for each
crystal type? Discuss Look at Figure 13-24.
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Bonding in Solids

• Categories of crystalline solids
• Metallic solids
• Ionic solids
• Molecular solids
• Covalent solids
• Table 13-10 summarizes properties of each
  category type.

54
Metallic Solids

• Positively charged nuclei are surrounded by a sea
  of delocalized valence electrons. This is the
  reason why most metals are good conductors.
• The nuclei occupy lattice sites
• Lattice types for metals
• Body-centered cubic (bcc)
• Face-centered cubic (fcc)
• Hexagonal close-packed (hcp)
• Examples Li, K, Ca, and Au

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Metallic Solids

• Obtaining the close-packed structures (fcc and
  hcp)
• Hexagonal close packed structure has an ABA
  arrangement.
• The different letters correspond to different
  planes (Figure 13-27a).
• Face-centered cubic (or cubic close-packed)
  structure has an ABC arrangement.
• Figure 13-27b
• For the the close-packed structures approximately
  74 of the volume is occupied.
• The body-centered cubic structures has much less
  of its volume occupied by metal spheres.

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Metallic Solids

• Manganese has a simple cubic unit cell. The
  atomic radius is 3.15 Å. What is the shortest
  distance between neighboring Mn atoms? How many
  nearest neighbors does each atoms have?
• Nickel crystals are face-centered cubic. The
  radius of the nickel atom in the metal is 1.24Å.
  What is the distance between centers of the two
  closest Ni atoms. What is the length of the cell
  edge. How many nearest neighbors does each atoms
  have?
• Calculate the density of metallic nickel.
  Determine the percentage of space that is
  occupied by the nickel atoms.

57
Metallic Solids

• A group IVA element with a density of 11.35 g/cm3
  crystallizes in a face-centered cubic lattice
  whose unit cell edge length is 4.95 A. Calculate
  the elements atomic weight. What is the atomic
  radius of this element?

58
Ionic Solids

• Most salts crystallize as ionic solids with ions
  occupying the unit cell. The most common example
  is sodium chloride, which has a face-centered
  cubic arrangement.
• Many other salts that have the same charge on
  both the anion and cation have the same fcc
  arrangement (LiCl and MgO).
• Even though the solid compound possesses charge,
  it does not conduct electricity. Why?

59
Ionic Solids

• How many Cl- and Na ions are in the fcc unit
  cell? Remember, a sodium ion is in the center.
• Illustrate models and look at 13-13 on demo CD.
• Sodium iodide crystallizes in the fcc structure
  (like NaCl). The I- ion radius is 2.20 Å. The
  I- ions at the corners of the unit cell are in
  contact with those at the centers of the faces.
  Determine the length of the unit cell. Calculate
  the radius of the Na ion assuming anion-cation
  contact.

60
Molecular Solids

• Molecules occupy the lattice positions of the
  unit cell.
• There are covalent bond within the molecules but
  only intermolecular attractive forces between
  molecules. What are the types?
• These solids tend to have lower melting points
  since the forces holding the molecules together
  are weaker.
• Look at page 527 in book and 13-14 on demo CD.

61
Covalent Solids (network solids)

• These can be considered giant molecules with the
  atom bond covalently in a crystalline network.
• Structures are usually very hard with high
  melting points. Most are poor thermal and
  electrical conductors.
• Examples Diamond, graphite, and quartz
• What is the bonding in diamond (hydridization)
  and graphite?
• 13-15 on demo. CD

62
Band Theory of Metals

• Recall that the bonding in metals is due to
  delocalized, mobile electrons that belong to the
  solid as a whole.
• Responsible for the conduction (electrical and
  thermal) of most metals.
• Band theory of metals is used to explain
  properties of metals and other materials.
• According to the MO theory, atomic orbitals
  overlap to produce a set of molecular orbitals.
  The number of generated molecular orbitals is
  equal to the number of overlapping atomic
  orbitals. There is a very large number of atomic
  orbitals in a metal!!!

63
Band Theory of Metals

• In a Na metal, the 3s atomic orbitals overlap to
  produce a very large set of molecular orbitals
  that are very closely spaced in energy. These
  closely spaced orbitals are called a band of
  orbitals. Since there is only one electron in
  the 3s atomic orbital, the molecular orbitals are
  only half filled. Only a small amount of energy
  is needed for the highest energy electrons in the
  3s band to jump into a vacant orbital at a
  slightly higher energy.
• As a result, electrons will flow with an applied
  field.

64
Band Theory of Metals
What if the 3s band is filled? In the Group IIA
metals, the 3s band is filled, but the metals
still conduct.
65
Band Theory of Metals

• For Mg, the 3s and 3p band overlap since there is
  a range of energies for the molecular orbitals.
• Without this overlap Mg metal would not conduct.
• The highest energy electrons are able to move
  into the vacant orbitals in the 3p band.

66
Band Theory of Metals

• Why dont some materials conduct electricity?
• The highest energy electrons in insulators occupy
  filled bands of molecular orbitals that are well
  separated from the lowest empty band.
• For semiconductors, the filled bands are only
  slightly below the empty bands.
• Generally, increasing the temperature or doping
  the material will produce conduction (page
  518-519).

67
Band Theory of Metals