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