1. Covalent Bonding
Results from the sharing of electron pairs between two atoms
B/w nonmetal and a nonmetal
Nonpolar Covalent: a covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge
Polar Covalent: a covalent bond in which bonded atoms have an unequal attraction for the shared electrons

2. Polarity and Electronegativity
Polar: an “uneven” distribution of charge
Remember Electronegativity
What are the trends on the periodic table?
Which element is the most electronegative?

3. Should be 0.3 for nonpolar
Should be for polar
For ionic,

5. Some Examples
Using electronegativity values, designate whether the bonds would be nonpolar covalent, polar covalent, or ionic.
C and H
C and S
O and H
Na and Cl
Cs and S

7. Some Ionic Characteristics
High melting points
Hard, brittle substances (solids with lattice structures)
Good insulators
Conduct electricity in solution (electrolyte)

8. Metallic Bonding
Metals tend to have high melting points and boiling points suggesting strong bonds b/w atoms.
Example: Sodium
All of the 3s orbitals on all of the atoms overlap to give a vast number of molecular orbitals which extend over the whole piece of metal.

9. The electrons are said to be delocalized.
The electrons can move freely within these molecular orbitals, and so each electron becomes detached from its parent atom.
The electrons are said to be delocalized.
The metal is held together by the strong forces of attraction between the positive nuclei and the delocalised electrons. (In a way, they act like “cement” holding the positive metal ions in relatively fixed positions)
described as "an array of positive ions in a sea of electrons".

10. Method of Drawing Lewis Structures
Sum the valence electrons of all atoms. Subtract an electron for a (+) sign & add an electron for a (-) sign.
Put the atom wanting the most bonds in the middle
An atom will form one bond for each electron it wants.
Put the remaining atoms around the central atom, giving them the number of bonds they want.
Fill in pairs of electrons until every atom has eight electrons
Exceptions: H, B, Be

11. Some Covalent Characteristics
Depends on polarity and whether the molecule is a covalent network (these can drastically influence characteristics)
Nonpolars generally are gases with low boiling points, are not good conductors of electricity (nonelectrolyte), and are volatile
Polars generally are liquids with higher boiling points. Highly polar ones can conduct electricity
Covalent networks can have very high melting points and are usually nonconductors

12. Molecular Geometries
VSEPR Theory: “valence shell electron pair repulsion”
Repulsion between the sets of valence level electrons surrounding an atom causing these to be oriented as far apart as possible.

13. 2 atoms bonded to central atom
No lone pairs
Type of molecule: AB2
3 atoms bonded to central atom
No lone pairs
Type of molecule: AB3

14. Let’s look at CCl4 Lewis Structure
However, the Lewis structure provides no information about the shape of the molecule
Atoms Bonded to Central = 4
Lone pairs = 0
Type of molecule = AB4
Carbon tetrachloride is tetrahedral in structure

15. Bent:
2 atoms bonded to the central atom
2 sets of lone pairs

20. Can nonpolar molecules contain polar bonds?
First assign dipoles and + and - signs
Each C-Cl bond is polar
Overall though, CCl4 is non polar, due to its symmetrical shape.
The more polar a molecule is the greater separation of charge
The more polar a molecule is the more attracted it will be to another polar molecule

21. Intermolecular Forces
Intermolecular forces are the forces of attractions that exist between molecules in a compound.
These cause the compound to exist in a certain state of matter: solid, liquid, or gas; and affect the melting and boiling points of compounds as well as the solubilities of one substance in another.

22. The stronger the attractions between particles (molecules or ions), the more difficult it will be to separate the particles.
In a solid the kinetic energy of the molecules is small compared to the strength of the intermolecular forces, so each molecule can only move short distances around a fixed position.
In a liquid the kinetic energy of the molecules is comparable to the intermolecular forces between them, so the molecules can move around, but usually stay within a molecular diameter of each other.
In a gas the kinetic energy of the molecules is much greater than the intermolecular forces between them and the molecules move freely, colliding far less frequently than in a liquid.

23. Gas, Liquid, vs. Solid
For each phase, a molecule or compound would most likely be polar or nonpolar and why?
Solids:
The stronger the intermolecular forces the closer the molecules pull themselves.
More density: greater likelihood of being a solid at room temperature.
Ionic compounds have a very large separation of charge, therefore they have a strong force of attraction → likely to be solid at room temperature.

24. Liquids:
At room temperature the atoms within a liquid are not held as closely as those within solid.
Mostly polar, especially water due to hydrogen bonding.
Gases:
At room temperature gases are nonpolar
There is no separation of charge between most gas molecules
Ex. O2, CH4

25. What types of intermolecular forces are there?
Interacting molecules or ions
Are ions involved?
Are polar molecules involved?
Are H atoms bonded to N,
O, or F atoms?
Hydrogen Bonding
Dipole-Dipole Forces
London Dispersion Forces
Are polar molecules and ions
both present?
Ionic Bonding
Ion-dipole force
Yes No

26. What causes intermolecular forces?
Molecules are made up of charged particles: nuclei and electrons.
When one molecule approaches another there is a multitude of interactions between the particles in the two molecules.
Each electron in one molecule is subject to forces from all the electrons and the nuclei in the other molecule.

27. Ionic Bonding: Ion-Dipole Forces
When an ionic substance dissolves in a polar solvent (that is, a solvent whose molecules have a permanent dipole moment) the majority of the solvent molecules orient themselves with the oppositely charged end of the solvent molecule near an ion.
This attraction between the ions and the solvent molecules can win out over the attraction of the ions to each other, allowing the substance to stray in solution.
For an insoluble ionic substance this is not the case. Ion - dipole forces are responsible for the dissolution of ionic substances in water.

28. Van der Waals Forces Include: Hydrogen Bonding Dipole-dipole
Dipole-Induced Dipole
London Dispersion Forces

29. Covalent Bonding: Dipole-Dipole Forces
If two neutral molecules, each having a permanent dipole moment, come together such that their oppositely charged ends align, they will be attracted to each other.
In a liquid or solid these alignments are favored over those where like-charged ends of the molecules are close together and hence repel each other.

30. Hydrogen Bonding
Hydrogen bonds form only between a limited number of elements bonded in a specific sequence.
So a necessary condition for hydrogen bonding is that one molecule must contain a H bonded to either N, O, or F; and the other molecule must contain either N, O, or F.
If a hydrogen bond can form between a pair of molecules it will be stronger than other intermolecular forces between the molecules.
Hydrogen bonding is responsible for the unexpectedly high boiling point of water

31. Dipole-Induced Dipole Forces
A polar molecule (lower left) carries with it an electric field and this can induce a dipole moment in a nearby non-polar molecule (lower right).
This will cause an attraction between the molecules.
This type of force is responsible for the solubility of oxygen (a non-polar molecule) in water (polar).

32. London Dispersion Forces
Arise from the temporary variations in electron density around atoms and molecules.
Nonpolar molecules have a certain minimum symmetry to their average shape and electron distribution. Picture 1 at left depicts two nonpolar molecules.
However at any instant the electron distribution around an atom or molecule will likely produce a dipole moment (figure 2 on left) which will average out to zero over a period of time.

33. London Dispersion Continued
But even a temporary dipole moment can induce a (temporary) dipole moment in any nearby molecules (picture 3 on left) causing them to be attracted to the first molecule.
Unlike forces between molecules with permanent dipole moments, dispersion forces always act to attract the molecules to each other regardless of the relative orientation of the molecules
Molecules containing large atoms (e.g. bromine or iodine) have large polarisabilities and so give rise to large dispersion forces. This explains the increasing melting and boiling points of the halogens going down that group of the periodic table.