1. Unit 6
Covalent Bonding

2. Lesson 1:Covalent Bonding
Covalent bonds: atoms held together by sharing electrons.
Molecules: neutral group of atoms joined together by covalent bonds.
Diatomic molecule: molecule consisting of 2 atoms. Remember them: F2, Cl2, I2, Br2, H2, N2, O2
Molecules tend to have lower melting and boiling points than ionic compounds.
YouTube - Making Molecules with Atoms

3. Molecular Formula
Shows how many atoms of each element a molecule contains.
Naming binary molecular compounds
Composed of two nonmetals; often combine in more than one way. Ex. CO and CO2
Prefixes are used to name binary molecular compounds.
Prefix
Mono-
Di-
Tri-
Tetra-
Penta-
Hexa-
Hepta-
Octa-
Nona-
Deca-
Number 1 2 3 4 5 6 7 8 9 10

4. Binary Compounds Containing Two Nonmetals
To name these compounds:
give the name of the less electronegative element first with the Greek prefix indicating the number of atoms of that element present
After give the name of the more electronegative non-metal with the Greek prefix indicating the number of atoms of that element present and with its ending replaced by the suffix –ide.
Do not use the prefix mono- if required for the first element.

5. Binary Molecular Compounds
N2O dinitrogen monoxide
N2O3 dinitrogen trioxide
N2O5 dinitrogen pentoxide
ICl iodine monochloride
ICl3 iodine trichloride
SO2 sulfur dioxide
SO3 sulfur trioxide
YouTube - Naming molecular compounds

6. Binary Molecular Compounds Containing Two Nonmetals
As2S3
________________ diarsenic trisulfide
________________ sulfur dioxide
P2O5 ____________________
________________ carbon dioxide
N2O5 ____________________
H2O ____________________
SO2
diphosphorus pentoxide
CO2
dinitrogen pentoxide
dihydrogen monoxide

7. Naming Binary Compounds
Yes
Metal Present? No
Yes
Molecule
Use Greek
Prefixes
Does the metal form
more than one cation? No
Yes
Ionic compound (cation has more
than one charge) Determine the
Charge of the cation; use a Roman
numeral after the cation name.
Ionic compound (cation has
one charge only)
Use the element
name for the cation.
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 98

8. Classwork #1: Do handout “Naming Molecules”

9. Lesson 2: The Nature of Covalent Bonding
Introduction with balloon activity
octet rule: electron sharing occurs usually so that atoms attain the electron configurations of noble gases.
Single covalent bond: two atoms held together by sharing a pair of electrons. Shown as two dots or as a long dash.
A pair of valence electrons that is not shared between atoms is called an unshared pair.

10. O H H O H O H or H

11. Double bonds: covalent bond formed by sharing two pairs of electrons
Triple bonds: covalent bond formed by sharing three pairs of electrons.

12. A dash may replace a pair of dots.
Covalent bonding with equal sharing of electrons occurs in diatomic molecules formed from one element.
chlorine
iodine
hydrogen
nitrogen
A dash may replace a pair of dots.

13. Classwork: introduction to lewis structures.

14. Lesson 3:Molecular Structure
Structural formula: uses symbols and bonds to show relative position of atoms.
Steps to determine Lewis structures for molecules
Predict the location of certain atoms.
Hydrogen is always an end atom
The least electronegative atom is the central atom (usually the one closer to the left on periodic table)
Find the total number of electrons available for bonding. (# of valence electrons of atoms in molecule)
Determine the number of bonding pairs by dividing the total number of electron by 2

15. Place one bonding pair (single bond) between central atom and terminal atoms.
Subtract pairs used in step 4 from bonding pairs in step 3. Place lone pairs around each terminal atom bonded to the central atom to satisfy the octet rule. Any remaining pairs are assigned to the central atom.
If the central atom does not have an octet, convert one or two of the lone pairs on the terminal atoms to a double or a triple bond between central and terminal atom. Some elements like Be, B, Al do not form a complete octet, S and P can have more than 8 valence electrons.

16. Ex. 1 Draw the lewis structure for ammonia, NH3
Hydrogen is an end atom and nitrogen is the central atom.
Total number of valence electrons:
(1 nitrogen x 5 valence electrons)+ (3 hydrogens x 1 valence electron)= 8 valence electrons.
Total number of bonding pairs= 8/2 = 4
Draw single bond from each H to N N H

17. Ex. 1 Draw the lewis structure for ammonia, NH3
Subtract the number of pairs of electrons used from the total pairs of electrons: =1 pair available
One lone pair remains, hydrogen can have only one bond, assign the lone pair to the central atom, N. N H

18. Ex. 2 Draw the lewis structure for carbon dioxide, CO2
Oxygen atoms are end atoms and carbon is the central atom.
Total number of valence electrons:
(1 carbon x 4 valence electrons)+ (2 oxygen x 6 valence electron)= 16 valence electrons.
Total number of bonding pairs= 16/2 = 8
Draw single bond from each C to O C O

19. Ex. 1 Draw the lewis structure for carbon dioxide, CO2
Subtract the number of pairs of electrons used from the total pairs of electrons: =6 pair available
Add three pairs of electrons to each oxygen. C O

20. Ex. 1 Draw the lewis structure for carbon dioxide, CO2
No lone pairs remain for carbon. Carbon does not have an octet, use a lone pair from each oxygen to form a double bond with the carbon atom. C O C O

21. CW: lewis structures handout part 1

22. Lesson :4 Exception to octet rule
Some molecules have an odd number of valence electrons and cannot form an octet around each atom.
Some molecules form with fewer than eight electrons present around an atom. Ex. Boron
Some compounds have central atoms with more than 8 electrons. This is called an expanded octet. Ex. S, Xe and P

23. Ex. 3 Draw the lewis structure for XeF4 (exception octet rule)
F is an end atom and nitrogen is the central atom.
Total number of valence electrons:
(1 xenon x 8 valence electrons)+ (4 fluorines x 7 valence electron)= 36 valence electrons.
Total number of bonding pairs= 36/2 = 18
Draw single bond from each F to Xe Xe F

24. Ex. 1 Draw the lewis structure for XeF4 (exception octet rule)
Subtract the number of pairs of electrons used from the total pairs of electrons: =14 pairs available
14 lone pairs remain, place them around each fluorine so that each fluorine has 8 valence electrons Xe F

25. Ex. 1 Draw the lewis structure for XeF4 (exception octet rule)
There are 2 pairs of electrons still available, place around Xe which is capable of having more than 8 valence electron. Xe F

26. Molecular Shape VSEPR (Valence shell electron pair repulsion) Model
The repulsion between electron pairs in a molecule result in atoms existing at fixed angles from each other. (Remember balloon activity)
Shared electron pairs repel each other
A greater repulsion occurs between unshared electron pairs and shared electron pairs.

27. Bonding and Shape of Molecules: Count number of bonds and unshared pairs of electrons AROUND CENTRAL ATOM and then use table below to determine shape of molecule.
Number
of Bonds
Number of
Unshared Pairs
Covalent
Structure
Shape
Examples 2 3 4 1 2
-Be-
Linear
Trigonal planar
Tetrahedral
Pyramidal
Bent
BeCl2
BF3
CH4, SiCl4
NH3, PCl3
H2O, H2S, SCl2 B C N : O :

28. Use table on last slide to determine shape of molecule.
SO2 H O
Shape: bent
2 bonds and 2 unshared pairs

29. Carbon tetrachloride Cl Cl C Cl Cl Cl C CCl4 Shape: Tetrahedral
4 bonds and 0 unshared pairs.
Carbon tetrachloride – “carbon tet” had been used as dry cleaning solvent
because of its extreme non-polarity.

30. Classwork: do in your notebook
Determine the shape for the following molecules (first draw the lewis structure for the molecule and then use the table on slide #7 to determine the shape taking in consideration the number of bonds and unshared pairs of electrons around the CENTRAL ATOM.)
BF OCl CF4
4. NH BeI2