1. Solids Chapter
Metals, well you know what they are, but let us describe some properties
On the macroscopic level metals have
High electrical conductivity
High thermal conductivity
Can be bent (ductile) and beaten into shape (malleable)
Shiny
Most are silvery
On the microscopic level metals are composed of
One kind of ion and a number of weakly bound electrons
Classically one can view the electrons as moving (more or less) freely in the metallic lattice
Quantum mechanically one views the electrons as being delocalized in the matrix
This is somewhat similar to the two views of resonance bonding.

2. Melting Points
Metal
Melting Point Cs
28.4 C Ga
29.8 Hg
-38.8 Ni
2651 W
6098
The delocalized electrons hold the metallic lattice together
The strength of the bonds varies dramatically as can be seen from the melting points of different metals
Metals with few (Grp 1) or many (Grp 12) valence electrons have low melting points.
In Grp 1 there is only one electron per metal atom to participate in bonding.
(Figure from Wikipedia)

3. Drude Model
The simplest model of a metal was introduced by Paul Drude in 1900
It is often called the electron sea model, with the metal lattice sitting in a sea of valence electrons that are free to move. It explains the electrical and thermal conductivity of metals as well as their shiny appearance

4. Drude Model
In the Drude model some of the metal atom’s valence electrons move freely in the metal lattice colliding with metal ions. Resistance to flow because of the number of collisions is a measure of the conductivity of the metal
The model assumes that the electrons behave as an ideal gas. This is wrong, even conceptually, the electrons interact, but in the end it provides the right answer.
In an oversimplified way, the interaction of an electron with all the others can be thought of as a drag. Effectively this can be represented as an increase in the mass of the electron slowing it down. The effective mass is written m*
From the Wikipedia
Since with increasing temperature the kinetic energy of the electrons increases the number of collisions will also increase. This means that the conductivity of a metal decreases with increasing temperature.

5. Band Model of Metals
Heitler and London were the first to apply the Schrodinger model of atoms to metals.
This extends the molecular orbital picture we built in Gchem 003 for diatomic molecules with bonding and anti bonding orbitals
If we bring two metal atoms with a single s electron, say Cu [Ar]3d104s1 together we will form bonding and anti bonding orbitals
Now if we bring a third Cu atom together with the two atoms
We can keep on adding atoms A2 A1 A2 A3

6. Building Bands A2 A3 A4 A5 A6 A7 A8
Although it is hard to see (and I am not an artist) as more atoms are added to the one dimensional lattice we are building the states, both bonding and antibonding, get closer together

7. Building Bands
Conduction Band
Valence Band A2 A3 A4 A5 A6 A7 A8
Until at last they form a continuous band. The part of the band filled with the bonding electrons is called the valence band. The upper part which has no valence electrons is called the conduction band. Electrons in the valence band are localized around metal ions.

8. Conductors
In metals there is no gap between the conduction and valence bands
At absolute zero all of the electrons are in the valence band and localized. That means that at 0 K even metals do not conduct electricity.
However, as the metal warms up some of the electrons will have enough energy to be promoted (moved) into the delocalized conduction band where they can conduct electricity.
Conduction Band
Valence Band

9. Insulators A2 A3 A4 A5 A6 A7 A8
Conduction Band
Band Gap
Valence Band
The details are important. Forming the bands can leave an area between the valence and conduction bands where there are no orbitals. That is called a band gap. For electrons to move from the valence to the conduction band they have to jump the band gap. For an insulator the band gap is large compared to thermal energy.

10. Semiconductors
If the bandgap exists, but is small enough that some electrons in the valence band are promoted into the conduction band the material is called a semiconductor. Germanium and silicon are semiconductors, diamond is an insulator.

11. Semiconductors
If a semiconductor is heated it will conduct more easily. The alteration of the number of conduction electrons is one of the reasons that solid state electronics fail at high temperatures.
At 298 K, the thermal energy, RT is about 2.5 kJ/mol so the conduction of pure Si or Ge semiconductors is low because not many electrons will be promoted into the conduction band.

12. Semiconductors
If the temperature increases more electrons will move from the valence to the conduction band in a semiconductor such as silicon, whereas the increased number of collisions of an electron with the metal ions in a metallic conductor will decrease the conductivity.

13. Doping Semiconductors
The conductivity of a semiconductor can be manipulated by introducing impurities that sit interstitually in the metal lattice. These can ether be donors (n-type) that can donate an electron TO the conduction band
Or acceptors (p-type) that can accept an electron FROM the valence band. This leaves what is called a hole in the valence band, and conduction occurs as the hole moves from one metal ion to the next

14. Doping Crystals
The color of gemstones is almost always the result of doping a colorless or clear ionic or covalent crystal with an ionic metal species. For example beryl.
beryllium aluminum silicate,Be3Al2Si6O18, comes in many forms, the one on the left being aquamarine. Looking closely at the circled pillar, one can clearly see the hexagonal nature of the unit cell
VARIETY
COLOR
CHROMOPHORES
Emerald
Chrome-green
Cr, V, Fe2+, Fe3+
Aquamarine
Light blue to sea green
Fe2+, Fe3+
Maxixe beryl`
Blue (fades in sunlight)
Chrysolite
Yellow-green
Golden beryl
Golden-yellow
Heliodor
Greenish-yellow
Morganite
Pink-orange, pale pink
Mn2+, Mn3+
Bixbite (Red Emerald)
Dark-red
Goshenite (Rosterite)
Colorless
None
From the Wikipedia

15. Substitutional Impurities
As the last slide discussed for beryl, substitutional impurities occur when a different atom or ion is substituted into the crystal lattice. In most cases the substituted ion will be of approximately the same size as that of the ion it replaces. Accommodation of ions of different charges requires that in addition to the substituted ion a hole be introduced into the lattice to compensate.

16. Interstitial Impurities
Interstitial impurities occur when a smaller atom or ion becomes trapped into empty spaces in the crystal lattice.
Steel is an example of this with smaller carbon atoms trapped in the iron matrix.
Stainless steel has both interstitial carbon impurities and substitutional chromium and vanadium ions replacing some of the iron ions.

17. Metal Alloys Brass Cu Zn Bronze Sn Alnico Al Ni Co Nichrome Cr SteelFe C
Various
Stainless Steel
C + Various
Solder Pb

18. Polymers (aka Plastics)
Long chain, yuuuuuge organic molecules assembled from smaller molecules called monomers.
Polymers consist of many repeating monomer units
A polymer is analogous to a necklace made from many small beads
A chemical reaction forming polymers from monomers is called polymerization

19. Molecular Weight of Polymers
The length of the polymer chain can vary
To an extent it can be controlled
Many properties of the polymer depend on the chain length
These include such things as density, hardness, malleability, etc.

20. Biopolymers
There are many biologically important polymers, for example cellulose and starch. Cellulose forms the cell walls of plants and is built from beta glucose monomers.
The monomer unit in starch is alpha glucose

21. Nylon

22. Nylon

23. Polyethylene/Polyethene

24. Assignment Working in groups of three
Name 1 other commercial polymer. (Not PET, polystyrene, or nylon)
What is its use?
What monomers is it built out of?
Name 1 other biopolymer. (Not DNA, starch or cellulose)
What is its biological function?