STRUCTURES OF SOLID

CLASSIFICATIONS OF SOLIDS Metallic solids are held together by a delocalized “sea” of

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.

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.

Crystalline and Amorphous Solids

Solids in which atoms are arranged in an orderly repeating pattern are 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).

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.

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

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.

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.

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.

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

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.

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).

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.

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.

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.