1. Properties of bonding
Mrs. Kay

2. Properties of Ionic bonding
Determined by their crystalline structures (how the crystals form)
Solid at room temperature (no movement)
High melting points = strong bonds
Very hard and brittle

3. Molten compounds conduct electricity
Solid structure does not conduct electricity because of rigidity and no ionic movement for electricity to pass through.

4. Strength Depends on the radii and charges on the ions
Increasing metallic charge= stronger bonds = highest melting points
Which would produce stronger bonds Na+, Ca+2, or Al+3?
Na+ < Ca+2 < Al+3

5. Solubility
Most are soluble (can dissolve into ions) in polar solvents (ex: water, ammonia)
These solutions do conduct electricity because of mobile ions (electrolytes)

6. Hydration
The process which polar solvent molecules interact with ions in the crystal lattice and cause the ionic crystal to dissolve, releasing ions into solution
Water surrounds the ion (ion-dipole interactions)

7. Properties of simple covalent molecules
Covalent molecules exist as s, l, or g
Usually soft
Evaporate easier than ionic
Low melting and boiling points

8. Do not conduct electricity in liquid or solid state
Not soluble in polar solvents, but may be soluble in nonpolar solvents (CCl4 or gasoline)
Napthalene (smells like moth balls)

9. Note:
Molecules: any electrically neutral group of atoms that are bonded tightly together to be considered one particle.
Ex: Cl2, NH3, H2O
Ionic compounds are not molecules!!!
NaCl is not one molecule but a crystal lattice structure with attractive forces holding them together.

10. Metallic Bonding Name 4 Characteristics of a Metallic Bond.
What is a Metallic Bond?
- A metallic bond occurs in metals. A metal consists of positive ions surrounded by a “sea” of mobile electrons.
Good conductors of heat and electricity
Great strength
Malleable and Ductile
Luster
This shows what a metallic bond might look like.

11. 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

12. 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

13. 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)

14. 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)

15. 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!

16. 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)

17. Network Covalent Molecules: Allotropes of carbon
elements can exist in two or more different forms because the element's atoms are bonded together in a different manner
Carbon has 3 allotrophes
Diamond
Graphite
Fullerenes (C60)
Nanotubes
Buckminster Fullerene

18. Diamonds
carbon atoms are bonded together in a tetrahedral lattice arrangement (3D framework)
Giant covalent structure
Very strong, so they require a lot of energy to break them
M.P is 3820 K
Does NOT conduct electricity
4x harder than any other natural mineral

19. Graphite
has a sheet like structure where the atoms all lie in a plane and are only weakly bonded to the sheets above and below. (2D framework)
Much softer, conducts electricity. (delocalised electron)
The C-C bonds are still quite strong.
Each carbon bonded to 3 other carbon.

20. Fullerene C60
consists of 60 carbon atoms bonded in the nearly spherical configuration
C60 is highly electronegative, meaning that it readily forms compounds
Low solubility, low conductivity (greater than diamond, but much lower than graphite)
Buckminster Fullerene (photo) made up of hexagon and pentagon carbon formations
Also includes nanotubes (cylindrical) Made up of hexagons of carbon