Chapter 13
The Chemistry of
Solids
Jeffrey Mack
California State University,
Sacramento
Metallic & Ionic Solids
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.
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
Types of Solids
Type:
Ionic Compounds
Metals
Molecular
Network
Amorphous
Examples:
Forces:
NaCl, BaCl2, ZnS
Ion-Ion (ionic bonding)
Fr, Al
Metallic
Ice, I2, C12H22O11
Dipole-Dipole ot
Induced Dipoles
Diamond, Graphite
Extended Covalent
bonds
Glass, Coal
Covalent; directional
electron-pair bonds
Network Solids
Diamond
Graphite
Cubic Unit Cells
There are 7 basic crystal systems, but we will
only be concerned with CUBIC form here.
All sides
equal length
• 1/8 of each atom on a
corner is within the
cube
• 1/2 of each atom on a
All angles
face is within the cube
are 90 degrees
• 1/4 of each atom on a
side is within the cube
Cubic Unit Cells
Primitive
cubic (PC)
Bodycentered
cubic (BCC)
Facecentered
cubic (FCC)
Cubic Unit Cells
Unit Cells for Metals
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.
Atom Packing in Unit Cells
Assumes atoms are hard spheres and that crystals are built by
PACKING these spheres as efficiently as possible.
Atom Packing in Unit Cells
Crystal Lattices—Packing of Atoms
or Ions
• FCC is more
efficient than either
BC or PC.
• Leads to layers of
atoms.
Crystal Lattices—Packing of Atoms
or Ions
Packing of C60
molecules. They are
arranged at the lattice
points of a FCC
lattice.
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
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:
I2(s), sugar (C12H22O11) and “Dry Ice”, CO2(s)
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 SiO2
molecules. Instead the solid is an extended threedimensional network of ...-Si-O-Si-O-... bonding.
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.
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.
Bonding in Ionic Compounds:
Lattice Energy
The energy of an ion pair (cation/anion) is described
by Coulombs law:
Uion pair
(n +e - )(n -e + )
=C´
d
n+ = cation charge, n = anion charge
d = distance between ion centers
latticeU 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 )
Lattice Energy
The Lattice Energy of a
salt is dependant upon
the charge and size of
the ions.
Uion pair
(n +e - )(n -e + )
=C´
d
Lattice Energy
Calculation of lattice energy via the Born–Haber cycle, an
application of Hess’s law.
Problem:
Calculate the molar enthalpy of formation,
fH°, of solid lithium fluoride from the lattice
energy and following thermochemical data.
Problem:
Calculate the molar enthalpy of formation,
fH°, 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 F2(g)  LiF(s)
Problem:
Calculate the molar enthalpy of formation,
fH°, of solid lithium fluoride from the lattice
energy and following thermochemical data.
Start by drawing the Born-Haber cycle for the reaction:
Problem:
Calculate the molar enthalpy of formation,
fH°, of solid lithium fluoride from the lattice
energy and following thermochemical data.
Start by drawing the Born-Haber cycle for the reaction:
F(g)
+
Li+(g)
EA
IE
F(g)
Li(g)
subH
Li(s)
Do
+
½ F2(g)

LiF(s)
Problem:
Calculate the molar enthalpy of formation,
fH°, of solid lithium fluoride from the lattice
energy and following thermochemical data.
Using Hess’s Law, the enthalpy of formation is found
by:
+
F(g)
Li+(g)
EA
IE
F(g)
Li(g)
subH
Li(s)
Do
+
½ F2(g)

fHo = subH + I1 + Do + EA + latticeU
LiF(s)
Problem:
Calculate the molar enthalpy of formation,
fH°, of solid lithium fluoride from the lattice
energy and following thermochemical data.
fHo = subH + IE + Do + EA + latticeU
Li(s)  Li(g)
Li(g)  Li+(g) + e–
½ F2(g)  F(g)
F(g) + e–  F–(g)
Li+(g) + F–(g)  LiF(s)
∆subH°
IE
Do
EA
∆latticeU
= +159.37 kJ/mol
= +520. kJ/mol
= +78.99 kJ/mol
= –328.0 kJ/mol
= –1037 kJ/mol
fH° =
= –607 kJ/mol
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
fusionH (kJ/mol)
Energy evolved as heat on freezing = enthalpy of
crystallization
fusionH (kJ/mol)
Enthalpies of fusion can range from just a few thousand
joules per mole to many thousands of joules per mole.
Enthalpies of Fusion Are a Function
of Intermolecular Forces
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 =
sublimationH
• Sublimation, like fusion and evaporation, is
an endothermic process.
• The energy required as heat is called the
enthalpy of sublimation.
Sublimation
Sublimation entails the conversion of a solid directly to its
vapor. Here, iodine (I2) sublimes when warmed.
Transitions Between Phases: Phase
Diagrams
Phase diagrams are used to illustrate the relationship between
phases of matter and the pressure and temperature.
Phase Diagram for Water
Liquid phase
Solid phase
Gas phase
Phase Equilibria—Water
Solid-liquid
Gas-Liquid
Gas-Solid
Triple Point—Water
At the TRIPLE POINT all three
phases are in equilibrium.
Phases Diagrams: Water
T(˚C) P(mmHg)
Normal boil point (at 1atm):
100
760
Normal freeze point (at 1atm):
0
760
Triple point:
0.0098
4.58
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.
Phases Diagram: Water
• At constant temp,
an increase in
pressure can
bring about a
phase change
from solid to
liquid!
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!
CO2 Phase Diagram
• Notice the CO2 has a forward slope of the solid/liquid
boundary.
• This is seen because CO2 does not exhibit hydrogen
bonding.
CO2 Phases
Separate
phases
Increasing
pressure
More
pressure
Supercritical
CO2