1. Objectives: Explain covalent bonding using correct vocabulary.
Draw Lewis structures to show covalent compounds.
Bellwork:
Get out a sheet of paper to take some notes on.
What kind of compound is CO2? How do you know?

2. Bonds are…
Forces that hold groups of atoms together and make them function as a unit. Two types:
Ionic bonds – transfer of electrons (gained or lost; makes formula unit)
Covalent bonds – sharing of electrons. The resulting particle is called a “molecule”

3. Nonmetals hold on to their valence electrons. Covalent bonds
They can’t give away electrons to bond.
But still want noble gas configuration.
Get it by sharing valence electrons with each other = covalent bonding
By sharing, both atoms get to count the electrons toward a noble gas configuration.
Covalent bonds

4. Fluorine has seven valence electrons (but would like to have 8)
Covalent bonding F

5. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
Covalent bonding F F

6. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons… F F

7. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons… F F

8. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons… F F

9. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons… F F

10. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons… F F

11. F F Covalent bonding …both end with full orbitals
Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons…
…both end with full orbitals F F

12. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons…
…both end with full orbitals F F
8 Valence electrons

13. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons…
…both end with full orbitals F F
8 Valence electrons

14. Molecular Compounds
Compounds that are bonded covalently (like in water, or carbon dioxide) are called molecular compounds
Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds – this is not as strong a bond as ionic

15. Molecular Compounds
Thus, molecular compounds tend to be gases or liquids at room temperature
Ionic compounds were solids

16. - Page 215
These are some of the different ways to represent ammonia:
3. The ball and stick model is the BEST, because it shows a 3-dimensional arrangement.
1. The molecular formula shows how many atoms of each element are present
2. The structural formula ALSO shows the arrangement of these atoms!

17. H O Water Each hydrogen has 1 valence electron
- Each hydrogen wants 1 more
The oxygen has 6 valence electrons
- The oxygen wants 2 more
They share to make each other complete
Water H O

18. Water H O

19. H O H Water So, a second hydrogen attaches
Every atom has full energy levels
Water
Note the two “unshared” pairs of electrons H O H

20. Sometimes atoms share more than one pair of valence electrons.
A double bond is when atoms share two pairs of electrons (4 total)
A triple bond is when atoms share three pairs of electrons (6 total)
Know these 7 elements as diatomic:
Br2 I2 N2 Cl2 H2 O2 F2
Multiple Bonds

21. PH3 Steps to drawing Lewis structures for covalent bonds:
Sum all valence electrons for all atoms in the compound.
Place the least electronegative atom (except H) in the middle unless it’s indicated that another atom is in the center (usually by writing it first in the formula)
Connect all atoms to the center with a single bond.
Complete the octet of all atoms attached to the central atom
Place leftover electrons if any on central atom.
If the central atom doesn’t have at least 8 electrons, try multiple bonds.
PH3

22. The 3 Exceptions to Octet rule
For some molecules, it is impossible to satisfy the octet rule
#1. usually when there is an odd number of valence electrons
NO2 has 17 valence electrons, because the N has 5, and each O contributes 6.
It is impossible to satisfy octet rule, yet the stable molecule does exist
The 3 Exceptions to Octet rule

23. Exceptions to Octet rule
#2 Molecules that have fewer than the octet rule (H wants 2, Be wants 4, B wants 6)
Example: Boron Trifluoride
#3 – Molecule that have more than the octet rule usually having 10 or 12; examples exist because they are in period 3 or beyond; PCl5 or SF6
Exceptions to Octet rule

24. Steps to drawing Lewis structures for covalent bonds:
Sum all valence electrons for all atoms in the compound.
Place the least electronegative atom (except H) in the middle unless it’s indicated that another atom is in the center (usually by writing it first in the formula)
Connect all atoms to the center with a single bond.
Complete the octet of all atoms attached to the central atom
Place leftover electrons if any on central atom.
If the central atom doesn’t have at least 8 electrons, try multiple bonds.

25. Draw the following Lewis structures:
CO2 CO
SCl6
BBr3

26. NONPOLAR GEOMETRY

27. INTERMOLECULAR FORCES Sigma bonds are stronger than Pi bonds
INTERMOLECULAR FORCES- are forces involving molecules and other molecules
HYDROGEN (dipole-dipole)> Dipole-Dipole > Van Der Waals> London Dispersion
INTRAMOLECULAR FORCES- are forces involving the bonds within the molecules such as sigma and pi bonds and polarity
Sigma bonds are stronger than Pi bonds
Polarity influences the shape/geometry of the molecule and the intermolecular forces available.