2. CLASSIFICATIONS OF
SOLIDS held together by a delocalized ―sea‖ of
 Metallic solids are
collectively shared valence electrons.
 This form of bonding allows metals to conduct electricity.
 It is also responsible for the fact that most metals are relatively strong
without being brittle.
 Ionic solids are held together by the mutual attraction
between cations and anions.
 Differences between ionic and metallic bonding make the electrical
and mechanical properties of ionic solids very different from those of
metals.
 Covalent-network solids are held together by an extended
network of covalent bonds.
 This type of bonding can result in materials that are extremely
hard, like diamond, and it is also responsible for the unique properties
of semiconductors.

3.  Polymers contain long chains of atoms, where the atoms
within a given chain are connected by covalent bonds and
adjacent chains held to one another largely by weaker
intermolecular forces.

 Polymers are normally stronger and have higher melting
points than molecular solids, and they are more flexible than
metallic, ionic, or covalent-network solids.
 Nanomaterials are solids in which the dimensions of
individual crystals have been reduced to the order of 1–100
nm.
 As we will see, the properties of conventional materials
change when their crystals become this small.

5. Crystalline and Amorphous

Solidsatoms are arranged in an orderly repeating pattern are
Solids in which
called crystalline solids.
 These solids usually have flat surfaces, or faces, that make definite
angles with one another.
 The orderly arrangements of atoms that produce these faces also cause
the solids to have highly regular shapes
 Examples of crystalline solids include sodium chloride, quartz, and
diamond.
 Amorphous solids (from the Greek words for ―without form‖) lack the
order found in crystalline solids. At the atomic level the structures of
amorphous solids are similar to the structures of liquids, but the
molecules, atoms, and/or ions lack the freedom of motion they have in
liquids. Amorphous solids do not have the well-defined faces and shapes
of a crystal.
 Familiar amorphous solids are rubber, glass, and obsidian (volcanic
glass).

6. Unit Cells and Crystal
Lattices
 In a crystalline solid there is a relatively small repeating
unit, called a unit cell, that is made up of a unique
arrangement of atoms and embodies the structure of the
solid.
 The structure of the crystal can be built by stacking this unit
over and over in all three dimensions.
 Thus, the structure of a crystalline solid is defined by (a) the
size and shape of the unit cell and (b) the locations of atoms
within the unit cell.
 The geometrical pattern of points on which the unit cells are
arranged is called a crystal lattice.
 The crystal lattice is, in effect, an abstract (that is, not real)
scaffolding for the crystal structure.

8.  To understand real crystals, we must move
from two dimensions to three.
 In three dimensions, a lattice is defined by
three lattice vectors a, b, and c
 These lattice vectors define a unit cell that is
a parallelepiped (a six-sided figure whose
faces are all parallelograms) and is described
by the lengths a, b, c of the cell edges and
the angles α β γ between these edges.
 There are seven possible shapes for a three
dimensional unit cell are as shown

9.  If we place a lattice point at each corner of a unit cell, we get
a primitive lattice.
 All seven lattices that are primitive lattices.
 It is also possible to generate centered lattices by placing
additional lattice points in specific locations in the unit cell.
 This is illustrated for a cubic lattice for body-centered cubic
lattice has one lattice point at the center of the unit cell in
addition to the lattice points at the eight corners.
 A face-centered cubic lattice has one lattice point at the
center of
 each of the six faces of the unit cell in addition to the lattice
points at the eight corners.
 Centered lattices exist for other types of unit cells as well.
Examples include bodycentered tetragonal and face-centered
orthorhombic.
 Counting all seven primitive lattices as well as the various
types of centered lattices, there are a total of 14 three
dimensional lattices.

13. METALLIC SOLIDS
 Metallic solids, also simply called metals, consist entirely
of metal atoms.
 The bonding in metals is too strong to be due to dispersion
forces, and yet there are not enough valence electrons to
form covalent bonds between atoms.
 The bonding, called metallic bonding, results from the fact
that the valence electrons are delocalized throughout the
entire solid.
 That is, the valence electrons are not associated with
specific atoms or bonds but are spread throughout the
solid.
 We can visualize a metal as an array of positive ions
immersed in a ―sea‖ of delocalized valence electrons.

14. Electron-Sea Model
 A simple model for characteristics of metals is the electron-sea
model,which pictures the metal as an array of metal cations in a ―sea‖ of
valence electrons
 The electrons are confined to the metal by electrostatic attractions to the
cations, and they are uniformly distributed throughout the structure.
 The electrons are mobile, however, and no individual electron is confined
to any particular metal ion.
 When a voltage is applied to a metal wire, the electrons, being negatively
charged, flow through the metal toward the positively charged end of the
wire.
 The high thermal conductivity of metals is also accounted for by the
presence of mobile electrons.
 The movement of electrons in response to temperature gradients permits
ready transfer of kinetic energy throughout the solid.
 The ability of metals to deform (their malleability and ductility) can be
explained by the fact that metal atoms form bonds to many neighbors.
 Changes in the positions of the atoms brought about in reshaping the
metal are partly accommodated by a redistribution of electrons.

17. IONIC SOLIDS
 Ionic solids are held together by the electrostatic
attraction between cations and anions—ionic bonds.
 The high melting and boiling points of ionic compounds
are a testament to the strength of the ionic bonds.
 The strength of an ionic bond depends on the charges
and sizes of the ions. the attractions between cations
and anions increase as the charges of the ions go up.
 Thus NaCl, where the ions have charges of and , melts
at 801 °C, whereas MgO, where the ions have charges
of and , melts at 2852 °C.
 The interactions between cations and anions also
increase as the ions get smaller

18.  Although ionic and metallic solids both have
high melting and boiling points, the differences
between ionic and metallic bonding are
responsible for important differences in their
properties.
 Because the valence electrons in ionic
compounds are confined to the anions, rather
than being delocalized, ionic compounds are
typically electrical insulators.
 They tend to be brittle, a property explained by
repulsive interactions between ions of like
charge.

20. COVALENT-NETWORK
SOLIDS
 Covalent-network solids consist of atoms held
together in large networks by covalent bonds.
 Because covalent bonds are much stronger than
intermolecular forces, these solids are much
harder and have higher melting points than
molecular solids.
 Diamond and graphite, two allotropes of
carbon, are two of the most familiar covalent-
network solids.
 Other examples are silicon, germanium, quartz
(SiO2), silicon carbide (SiC), and boron nitride
(BN).

21.  In diamond, each carbon atom is bonded
tetrahedrally to four other carbon atoms
 The structure of diamond can be derived from
the zinc blende structure if carbon atoms
replace both the zinc and sulfide ions.
 The carbon atoms are sp3 -hybridized and
held together by strong carbon–carbon single
covalent bonds.
 The strength and directionality of these bonds
make diamond the hardest known material.
 The stiff, interconnected bond network is also
responsible for the fact that diamond is one of
the best-known thermal conductors.
 Not surprisingly, diamond has a high melting
point, 3550 °C.

22.  In graphite, the carbon atoms form covalently bonded layers that
are held together by intermolecular forces.
 The layers in graphite are the same as the graphene sheet
 Graphite has a hexagonal unit cell containing two layers offset so
that the carbon atoms in a given layer sit over the middle of the
hexagons of the layer below.
 Each carbon is covalently bonded to three other carbons in the
same layer to form interconnected hexagonal rings.
 Electrons move freely through the delocalized π orbitals, making
graphite a good electrical conductor along the layers
 conducting electrode in batteries.
 These sp2-hybridized sheets of carbon atoms are separated by
3.35 A from one another, and the sheets are held together only by
dispersion forces.
 Thus, the layers readily slide past one another when rubbed, giving
graphite a greasy feel.

23.  This tendency is enhanced when impurity
atoms are trapped between the layers, as
is typically the case in commercial forms of
the material.
 Graphite is used as a lubricant and as the
―lead‖ in pencils. The enormous differences
in physical properties of graphite and
diamond—both of which are pure carbon—
arise from differences in their three-
dimensional structure and bonding.