1. 4.4 Metallic bonding
4.4.1 Describe metallic bond as the electrostatic attraction between a lattice of positive ions surrounded by delocalized valence electrons.
4.4.2 Explain the electrical conductivity and malleability of metals
Students should appreciate the economic importance of these properties and the impact that the large-scale production of iron and other metals has made on the world.

2. Metallic bond Occurs between atoms with low electronegativities
Metal atoms pack close together in 3-D, like oranges in a box.
Close-packed lattice formation

3. Many metals have an unfilled outer orbital
In an effort to be energy stable, their outer electrons become delocalised amongst all atoms
No electron belongs to one atom
They move around throughout the piece of metal.
Metallic bonds are not ions, but nuclei with moving electrons

4. Physical Properties Conductivity
Delocalised electrons are free to move so when a potential difference is applied they can carry the current along
Mobile electrons also mean they can transfer heat well
Their interaction with light makes them shiny (lustre)

5. Malleability
The electrons are attracted the nuclei and are moving around constantly.
The layers of the metal atoms can easily slide past each other without the need to break the bonds in the metal
Gold is extremely malleable that 1 gram can be hammered into a sheet that is only 230 atoms thick (70 nm)

6. Melting points
Related to the energy required to deform (MP) or break (BP) the metallic bond
BP requires the cations and its electrons to break away from the others so BP are very high.
The greater the amount of valence electrons, the stronger the metallic bond.
Gallium can melt in your hand at 29.8 oC, but it boils at 2400 oC!

7. Alloys
Alloying one metal with other metal(s) or non metal(s) often enhances its properties
Steel is stronger than pure iron because the carbon prevents the delocalised electrons to move so readily.
If too much carbon is added then the metal is brittle.
They are generally less malleable and ductile
Some alloys are made by melting and mixing two or more metals
Bronze = copper and zinc
Steel = iron and carbon (usually)

8. Economic importance
Iron is found by certain percentages in minerals, such as iron oxides like of magnetite (Fe3O4), hematite (Fe2O3), and many others.
Hematite- up to 66% pure could be put in a blast furnace directly for the production of iron metal
98% of iron production is destined for making steel

9. Who needs it? Where does it come from?
China, then Japan, then Korea are the world’s largest consumer's of iron
Where does it come from?
Iron rich minerals are commonly found everywhere in the world, however China, Brazil and Australia are the highest producers of iron ore mining
The main constraint is the position of the iron ore relative to market, the cost of rail infrastructure to get it to market and the energy cost required to do so.

10. Exercise:
Use the commonly accepted model of metal bonding to explain why:
The boiling points of metals in the 3rd period increase from sodium to magnesium to aluminum.
Most metals are malleable
All metals conduct electricity conduct electricity in the solid state.

11. Reading on pages
Page 375 # 9.70, 9.74, 9.72

12. 4.5 Physical Properties
4.5.1 Compare and explain the following properties of substances resulting from different types of bonding: melting and boiling points, volatility, conductivity and solubility.
Look at how impurities affect these properties
Solubilities of compounds in polar and non-polar solvents
Solubilities of alcohols in water being related to chain length

13. General physical properties
Depend on the forces between the particles
The stronger the bonding between the particles, the higher the M.P and BP
MP tends to depend on the existence of a regular lattice structure

14. Impurities and Melting points
An impurity disrupts the regular lattice that its particle adopts in the solid state, so it weakens the bonding.
They always LOWER melting points
Its often used to check purity of a known molecular covalent compound because its MP will be off, proving its contamination

15. How would this ideal heat curve look different if the substance was contaminated?

16. Volatility
A qualitative measure of how readily a liquid or solid is vaporised upon heating or evaporation
It is a measure of the tendency of molecules and atoms to escape from a liquid or a solid.
Relationship between vapour pressure and temperature (B.P)
Mostly dealing with liquids to gas, however can occur from solid directly to gas (dry ice).
The weaker the intermolecular bonds, the more volatile

17. Conductivity C12H22O11(s)  C12H22O11(aq)
Generally molecules have poor solubility in polar solvents like water, but if they do dissolve they do not for ions
There are no charged particles to carry the electrical charge across the solution.
Example: sugar dissolves in water
C12H22O11(s)  C12H22O11(aq)

18. Dissolving sugar (covalent compound)
It takes energy to break the bonds between the C12H22O11 molecules in sucrose crystal structure.
It also takes energy to break the hydrogen bonds in water so that one of these sucrose molecules can fit into solution.
In order for sugar to dissolve, there must be a greater release of energy when the dissolution occurs than when the breaking of bonds occur.

19. Ionic compounds
The energy needed to break the ionic bond must be less than the energy that is released when ions interact with water.
The intermolecular ion-dipole force is stronger than the electrostatic ionic bond
Breaks up the compound into its ions in solution.

20. Soluble salt in water breaks up as
NaCl (s)  Na+ (aq) + Cl- (aq)

23. Ionic compounds Held together by strong 3-d electrostatic forces.
They are solid at room temperature and pressure
If one layer moves a fraction, the ions charges are off and now repulsion occurs. This is the reason they are strong, yet brittle.

24. Molten or dissolved ionic compounds conduct electricity
Insoluble in most solvents, yet H2O is polar and attracts both the + and – ions from salts

25. Covalent bonding properties
Giant covalent
Ex: diamond, silicon dioxide
Very hard
Very high MP (>1000oC)
Does not conduct
Insoluble in all solvents
Molecular covalent
Ex: CO2, alcohols, I2
Usually soft, malleable
Low MP (<200oC)
Does not conduct
More soluble in non-aqueous solvents, unless they can h-bond

26. Solubility of methanol in water
Alcohols generally become less soluble, the longer the carbon chain due to the decreasing tendency for hydrogen bonding to occur intermolecularly.

27. States of matter Physical state depends on intermolecular forces
The weaker the attraction, the more likely it’s a gas, while stronger attractions indicate solid.

28.
Metallic bonding review
Solubility review
History involved with dissolving ionic compounds