1. Ionic, Covalent and Metallic structures of solids

2. Introduction
Where does the concept of bonding come from and how do we know what structures will be formed??
The Atom
Valence electrons
Now we know bonding refers to the forces which hold atoms together
There are three types of bonding Ionic, Covalent and Metallic

3. Vocab: Crystalline Solid: Amorphous solid: Alloy: Conductivity:
Lattice:
Unit Cell:
Amorphous solid:
Alloy:
Conductivity:
Boiling Point:

4. Ionic Ionic compounds are Solids at room temp Metal and Nonmetal
Held together by electrostatic attraction between cations and anions
Crystaline solid … crystal lattice structure

5. Ionic 
Ionic lattices are extremely difficult structures to break apart. As a result, all ionic substances are solids with high melting points.
Potassium iodide, magnesium chloride and calcium oxide. All of these ionic compounds have melting points over 500 °C.
There is one method of breaking up the lattice - dissolve the ionic compound in water. Water has the ability to separate the ions from the lattice and allow them to move freely as a solution.

6. Ionic
Ionic Solids
Shape of Unit Cell determines breaking angle of crystal.
3 unit cells: simple cubic, body centered cubic, face centered cubic
High Lattice Energy makes them rigid, high melting and boiling points, etc.
Rigid structure of charges means they do not conduct electricity when solid (but do when dissolved in water or liquid)

8. Metallic
Only Metals
Held together by electrostatic attraction between positive nucleus and a sea of mobile valence electrons
Also called “delocallized covalent”
Covalent = sharing
Delocallized = not staying in one area
Mobile electrons = good
conductors of electricity
and heat

9. Metallic Bonding: Close Packing Strong and nondirectional
Difficult to separate metal atoms
High melting and boiling points
Solid at room temperature (except mercury)
Easy to move their relative position
Malleable and ductile
Can be Amorphous Solid

10. Metallic Alloys: 2 or more metallic elements Substitutional Alloy
All similar size
Structure of Alloy similar to each pure metal
Ex: brass = copper and zinc, sterling = Ag & Cu
Interstitial Alloy
Very different sized elements
Holes between larger atoms filled by smaller atoms
Ex: steel = iron & carbon; harder, stronger, less ductile than pure iron

11. Only nonmetal elements
Covalent molecules
Only nonmetal elements
Covalent compounds have no ions. Their atoms join up by sharing electrons with their neighbours in small groups or clusters of atoms called molecules.
For example:
You will find that covalent compounds have low melting points. They exist either as gases (methane), liquids (water) or as easily melted solids (paraffin wax).

12. Covalent molecules
The covalent bonds inside the molecules are very strong. The molecules don't break apart easily. However the forces attracting neighbouring molecules to each other are very weak. It is therefore very easy to separate molecules from one another: e.g. ammonia
When a covalent compound melts or boils, it is the forces between the molecules which are broken. Very little energy is needed to make this happen, so covalent substances have low melting and boiling points.

13. Covalent Network Structures
Only Nonmetal atoms
Giant network of localized covalent bonds
Rigid like ionic crystal lattices (crystalline structures)
Bonded by electrostatic attraction between shared, localized valence electrons and positive nucleus.
Since e- are localized
Poor conductors of heat
Insulators (do not conduct electricity)
Rigid (will break before it changes shape)

14. Covalent network structures
Diamond, graphite, the element silicon and silicon dioxide are examples
Diamond and Silica (glass)

15. Covalent network structures
These covalent network structures need a large amount of energy to break all of the strong covalent bonds in them, so these substances have very high melting points.
Unlike the ionic lattices, the covalent networks are insoluble in water - they are not broken apart by trying to dissolve them.

16. Covalent network structures
Diamond vs Graphite
Diamond: crystalline
hard, colorless, electrical insulator
Bonded in 6-membered rings
Graphite: amorphous-like
slippery, black, electrical conductor
Bonded in sheets of trigonal planar arrangements
Pi bonds between the layers account for the conductivity

17. Semiconductors Silicon network structures
Metallic: temp & conductivity inversely proportional
Silicon Network Structures: temp & conductivity directly proportional
Doping increases conductivity
Replace a small fraction of silicon with elements containing 1 more valence electron (n-type) or 1 less valence electron (p-type)

18. Compare Properties
Sulphur is easily ground, it is brittle, so it is a nonmetal. In fact it is a covalent molecule
Tin can be flattened but not broken.
Magnesium is not easily broken up

19. Summary Melting/boiling points Strength of bonds Examples Ionic High
The ionic bond is strong
Sodium Chloride
Covalent
Low
The covalent bond is strong intramolecular
Weak intermolecular
Water
Covalent network
Covalent bonds are strong
Diamond
Metallic
Varies
The Strength of the metallic bond varies
Magnesium