Chapter 11

Structure of solids continued

Structure and Bonding in Metals

• Metals have:– High thermal and electrical conductivity– Are malleable– Are Ductile

• The reason for this is they are like small spheres packed together and bonded equally with other metal atoms in all directions.

Body-centered & Face-centered Crystal Lattice

                      

              

Closest Packing

• The structural model has uniform spheres as atoms packed in a manner that most efficiently uses the available space.

• The top layer does not lie directly on the spheres below but in the spaces available.

Hexagonal close packing

• When the atoms in the third layer lay over the atoms in the first layer.

• The unit cell here is body centered.

Other examples in nature of Hexagonal Close Packing

Cubic Close packing

• When the first and the fourth layer line up with one another.

• The unit cell shown is face centered cubic.

Bonding Model for Metals

• Metals qualities are best explained by the electron sea model.

• This envisions a regular array of organized cations surrounded by delocalized sea of electrons.

• This allows the movement of electrical current, and the metal ions can be easily moved around as a metal is hammered into a shape.

Metal Strength

• Sodium, potassium and lithium are soft metals that may be cut with a spoon! They have only one valence electron each.

• Chromium and iron are much harder metals each with 6 and 8 valence electrons respectively.

• What about mercury?

Discussion

• Mercury hangs on to its valence 6s electrons very tightly. Mercury-mercury bonding is very weak because its valence electrons are not shared readily. (In fact mercury is the only metal that doesn't form diatomic molecules in the gas phase).

• Hg 200.59 [Kr] 4d10 4f14 5s2 5p6 5d10 6s2

Other notes:

• Metal alloys are a substance that contains a mixture of elements and has metallic properties.

• There are two types of alloys:– Substitutional alloy– Interstitial alloy

Substitutional Interstitial

Bonding in Molecular Solids• Molecular solids are held together by

intermolecular forces.

• London forces, Dipole-dipole and hydrogen bonding.

• The properties of the molecular solids depends not only on the strength of these forces but also on the ability of the molecules to closely pack.

• Examples: Ar, CO2, and H2O

Network Solids

• Many atomic solids form strong directional covalent bonds. This allows the formation of “giant” molecules.

• Silicon and Carbon form some of the most important network solids.

• Diamond and graphite are both made of carbon. Yet diamond is a poor conductor and graphite can conduct electricity.

Why?

• Diamond is carbon bound in a tetrahedral shape to other carbons (sp3). This localizes the electrons and prevents conduction.

• Graphite is layers of 6 carbon rings with some delocalized electrons between the sheets of rings. Aka. sp2 hybridization with pi-bonds.

This is why!

Silica

• Silica (SiO2) crystal when heated to 1600 °C and cooled rapidly an amorphous solid called glass is formed.

Ionic Solids

• These are stable high melting substances held together by strong static forces between oppositely charged ions.

• Most are binary solids and can be modeled by closest packing spheres.

• The smaller cations fit in the holes created by closely packing the anions.

• The packing is done to maximize the oppositely charged particles and minimize the repulsions by ions with the same charge.

Shapes

• There are three types of holes in closest packed structures.

• Trigonal holes formed by three sphere in the same layer

• Tetrahedral holes formed when a sphere sits in the dimple of three spheres in an adjacent layer.

• Octahedral holes are formed by two sets of three spheres of the closest packed structure.

• The relative size of the wholes is : Trigonal<tetrahedral<octahedral