1. Chapter 8: Covalent Bonding

2. The goal of all Chemical Bonds
Elements bond to have a full outer shell like noble gases.
Ionic bonds: lose and gain electrons
Covalent bonds: share electrons

3. Covalent Bonds Occur between two nonmetals
Molecular compounds have covalent bonds.
A bond in which two atoms share ONE pair of electrons between them is a
SINGLE COVALENT BOND

4. Hydrogen gas (H2) Hydrogen gas wants to look like Helium.
To do this it needs to get one more electron.

5. Diatomic Molecules Atoms that exist as a pair in nature
There are 7 diatomic elements
MEMORIZE!

6. Hydrogen = H2 Nitrogen = N2 Oxygen = O2 Fluorine = F2 Chlorine = Cl2
Bromine = Br2
Iodine = I2
Hydrogen gas
Nitrogen gas
Oxygen gas
Fluorine gas
Chlorine gas
Bromine
Iodine

7. Naming Covalent
Use prefixes to distinguish the number of each atom present in the compound
Example:
C2H4 = dicarbon tetra hydride
You still add the ending –ide to the last element

8. Prefixes 1 = mono 2 = di 3 = tri 4 = tetra 5 = penta 6 = hexa
7 = hepta
8 = octa
9 = nona
10 = deca

9. Examples

10. Common Names for some… Water = H2O Ammonia = NH3 Methane = CH4
Ozone = O3

11. Goal: to understand the shape of compounds
Lesson Objective:
To be able to draw Lewis Dot Structures
Goal: to understand the shape of compounds

12. How to draw Lewis Dot Structures
Determine the number of valence electrons for each atom by using the periodic table.
Calculate the total number of electrons
Be sure to multiply the number of valence electrons if there is more than one atom of the element in a compound
Write the symbols for each atom and the correct amount for each element

13. What order do I write the symbols in for a molecule?
Generally the first element in the formula goes in the center.
Good rule of thumb – the least electronegative element goes in the center.
Some elements NEVER go in the center (like: H, halogens)

14. Remember…
Each element only needs 8 electrons
1 dot = 1 electron
Only two dots per side
* Remember, H and He are exceptions, they can have just 2 dots

15. Once you have the symbols written, draw a solid line between each element. This represents the SINGLE COVALENT BOND.
For each line, subtract 2 from your valence electron total.
Each line = 2 electrons

16. Distribute remaining dots remembering to satisfy the octet rule (8 electrons per symbol)
If you run out of dots, you might need to make a double or triple bond.
Double bonds: atoms share TWO PAIRS of electrons. (4 electrons total)
Triple Bonds: atoms share THREE pairs of electrons. (6 electrons total)

17. Double and Triple Covalent Bonds
Double bonds: atoms share TWO PAIRS of electrons. (4 electrons total)
Triple Bonds: atoms share THREE pairs of electrons. (6 electrons total)

18. Example: HBr
# Valence electrons =
Write the symbols:

19. Examples:
H2O
NF3

20. Examples O2
CO2 CO N2

21. Resonance
Sometimes there is more than one correct way to write a Lewis dot structure.
Example: O3

22. Exceptions to the octet rule…
P and S can hold more than 8 electrons in some molecules
Ex) PBr5
Boron only requires 6 electrons to be stable
Ex) BH3

23. Lewis Dot Structures for Ions
Works the same way, BUT
For each + charge subtract an electron.
For each – charge add an electron
Ex) Cl-  8 valence electrons

24. 16.3 – Bond Polarity
Atoms participating in covalent bonds share electrons, but they do not always share equally.
Recall: Electronegativity: the tendency for an atom to attract electrons to itself when chemically combined with another element.

25. Nonpolar covalent bonds
Electrons are shared equally 
Examples: H2, O2, N2
How do we know?
Think about tug of war…

26. Tug of War
Both Hs have an electronegativity of 2.1 (table 14.2, p 405).
Equal electronegativy = share electrons equally = nonpolar bond

27. Polar Bonds Don’t share electrons equally.
Have a difference in electronegativity of .4 – 2
Ex) HCl
H: 2.1
Cl: 3.0
Difference of .9 so polar bond.

28. Polar bonds have partial charges on atoms
Polar Bonds - Notation
More electronegative element
d means partial
O -- H
H -- O
Polar bonds have partial charges on atoms

29. Chemical Bonding Ionic Bonds Covalent Bonds >2 Non polar Polar
0-0.4
0.4-2

30. Non Polar Molecules Symmetric shapes with similar atoms.
Ex) tetrahedral, linear, trigonal planar…
To get this, most times you must draw lewis dot structures.

31. Polar Molecules
Molecules that have partially positive end and one partially negative end.
Dipole: molecule with two charged regions (or “poles”)

32. Remember…Just because it has polar bonds doesn’t make it polar…
Polar bonds can “cancel” each other out and result in a non polar molecule.

33. Intermolecular Forces (IMF)
Attractions BETWEEN MOLECULES
van der Waals Forces
Dispersion
Forces
Dipole
Interactions
Weakest type of intermolecular interactions

34. Dispersion Forces Weakest of all IMFs Caused by motion of electrons
Strength of Dispersion forces increases with increasing size (molecular weight).

35. Dipole Interactions
Occurs when polar molecules are attracted to each other

36. A hydrogen bond is 5% the strength of the average covalent bond.
Hydrogen Bonding
Very strong dipole interaction between H and F, N, or O.
A hydrogen bond is 5% the strength of the average covalent bond.

37. Ionic vs. Covalent ** See table 16.5 ** Melting Point:
ionic > covalent
Electrical Conductivity
Ionic > covalent
Makeup of elements:
Solubility:
Covalent depends on make up of molecule