1. Covalent Bonding

2. The Covalent Bond Octet Rule Covalent Bonds molecule
The chemical bond that results from the sharing of valence electrons.
Shared electrons count for both atoms
Mostly in nonmetals
molecule

3. Naturally occur as diatomic molecules
Hydrogen - H2
Nitrogen – N2
Oxygen – O2
Fluorine – F2
Chlorine – Cl2
Bromine – Br2
Iodine – I2

4. Nature of Bonding Lewis Structures Single covalent bonds
Use electron-dot diagrams to show how electrons are arranged in molecules
Single covalent bonds
1 pair of shared electrons
Double covalent bonds
2 pair of shared electrons
Triple covalent bonds
3 pair of shared electrons

5. Sigma Bonds (s) Pi bonds (p) Electrons shared in between atoms
Single covalent bond
One in every molecule
Pi bonds (p)
Parallel orbitals overlap to share electrons
In multiple bonds all but one

6. Structural formulas (Lewis Dot Structures)
Predict the location of certain atoms.
Hydrogen is always terminal.
Least electronegative is the central atom
Find the total # of valence electrons
Determine the # of bonding pairs
Place one bonding pair between the central atom and all of the terminal atoms
Determine the # of remaining electron pairs. Place lone pairs around terminal atoms to satisfy the octet rule. Remaining pairs go to the central atom.
Convert to multiple bonds in order to fulfill the octet rule if necessary.

7. Naming Binary Molecular Compounds
The first name in the formula is always named first, using the element name from the periodic table.
The second element in the formula is named using the root of the element and adding the suffix –ide.
Prefixes are used to indicate the number of atoms of each type present in the compound.
Mono- is not used on the first element.

8. Naming Binary Acids Use the prefix hydro- for the hydrogen part
Use the root of the second part of the acid and add the ending –ic.
End by adding the word acid to the end.

9. Naming Oxyacids Do not write anything for the hydrogen.
Use the root of the oxyanion
If the oxyanion ends in –ate replace it with –ic.
If it ends in –ite replace it with –ous.

10. Laws Relating to Naming
Law of definite proportions
For any chemical compound, the masses of the elements are always in the same proportions.
Always simple whole number ratios
Law of multiple proportions
Different masses of one element combine with the same masses of another element in a ratio of small whole numbers
Form different compounds

11. Bond theories Molecular orbitals VSEPR theory
Overlapped atomic orbitals in two atoms that are combined
Orbitals apply to the molecule as a whole
VSEPR theory
Repulsion between pairs of electrons adjust to provide the greatest distance between electron pairs

12. Hybridization
Atomic orbitals mix
the result is hybrid orbitals

13. Polarity Sharing is not necessarily equal Table 8.3 page 238
Equal is nonpolar
Unequal is polar
Slight charge variations
Table 8.3 page 238
Polar Molecules
Two poles - dipole

14. Intermolecular Forces
Weaker than Ionic or Covalent
Determine State of molecular compound
Van der Waals forces
Two weakest intermolecular
Dispersion Forces
Weakest of all
Moving electrons
Greater in with more electrons

15. Hydrogen Bonds Dipole Interactions Strongest intermolecular
Polar molecules
Charges in dipoles
Hydrogen Bonds
Strongest intermolecular
5% strength of a covalent bond
Polar bonds with hydrogen