1. Types of chemical bonds
Bond: Force that holds groups of two or more atoms together and makes the atoms function as a unit.
Example: H-O-H
Bond Energy: Energy required to break a bond.
Ionic Bond: Attractions between oppositely charged ions.
Example: Na+ Cl-

2. Types of chemical bonds
Ionic Compound: A compound resulting from a positive ion (usually a metal) combining with a negative ion (usually a non-metal).
Example: M X-  MX
Covalent Bond: Electrons are shared by nuclei.
Example: H-H
Polar Covalent Bond: Unequal sharing of electrons by nuclei.
Example: H-F
Hydrogen fluoride is an example of a molecule that has bond polarity.

3. Lewis structures
Lewis Structure: Representation of a molecule that shows how the valence electrons are arranged among the atoms in the molecule.
Bonding involves the valence electrons of atoms.
Example: Na● H-H

4. Lewis structures of elements
Dots around elemental symbol
Symbolize valence electrons
Thus, one must know valence electron configuration

6. Lewis Structures of molecules
Single Bond: Two atoms sharing one electron pair.
Example: H2
Double Bond: Two atoms sharing two pairs of electrons.
Example: O2
Triple Bond: Two atoms sharing three pairs of electrons.
Example: N2
Resonance Structures: More than one Lewis Structure can be drawn for a molecule.
Example: O3

7. Rules for Lewis structures of molecules
Write out valence electrons for each atom
Connect lone electrons because lone electrons are destabilizing
Become two shared electrons
Called a “bond”
Check to see if octet rule is satisfied
Recall electron configuration resembling noble gas
In other words, there must be 8 electrons (bonded or non-bonded) around atom
Non-bonded electron-pair
Called “lone pair”

8. Let’s do some examples on the board
Duet rule F2
Octet rule O2 N2

9. Lewis structures Example
Write the Lewis Structure for the following molecules:
H2O
CCl4
Where does the carbon go & why?
PH3
H2Se
C2H6

10. Lewis structures continued
C2H4
C2H2
SiO2

11. Polyatomic ions If positive charge on ion If negative charge on ion
Take away electron from central species
If negative charge on ion
Add electron to central species
Example:
H3O+

12. Your turn
NH4+
ClO-
OH-

13. Resonance structures
When structures can be written in more than one way O3
Actual molecule is “in-between”
Resonance hybrid
Another example
HCO3-
What would its resonance hybrid look like?

14. Practice
NO2-
NO3-

15. Practice
H2O2
H3O+

16. Aberrant compounds
Odd-electron species NO
NO2

17. Aberrant compounds
Incomplete octet
BH3

18. Aberrant compounds Expanded octet
Some central atoms can exceed an octet
Third period and higher elements can do this
E.g., Al, Si, P, S, Cl, As, Br, Xe, etc.
d-orbitals can accommodate extra electrons

19. Examples
AsI5
XeF2

20. Practice
SCl6
XeF4

21. Practice
SO32-
PO33-
SO2
SO3
H2SeO4

22. Electronegativity
Electronegativity: The relative ability of an atom in a molecule to attract shared electrons to itself.
Example: Fluorine has the highest electronegativity.
Similar electronegativities between elements give non-polar covalent bonds ( )
Different electronegativities between elements give polar covalent bonds ( )
If the difference between the electronegativities of two elements is about 2.0 or greater, the bond is ionic

23. Electronegativity Example
For each of the following pairs of bonds, choose the bond that will be more polar.
Al-P vs. Al-N
C-O vs. C-S

24. Dipole moment Dipole Moment
A molecule that has a center of positive charge and a center of negative charge
Will line up on electric field
In Debye units
1 D = 3.34 x C  m

25. Examples F2
CO2
H2O
NH3
BF3
CCl4

26. Molecular polarity Net-dipole moment leads to molecular polarity
Thus the following two that have net-dipole moments are polar:
H2O
NH3

27. Molecular structure
Molecular Structure: or geometric structure refers to the three-dimensional arrangement of the atoms in a molecule.
Bond Angle: The angle formed between two bonds in a molecule.

28. Molecular structure: VSEPR
The VSEPR Model: The valence shell electron pair repulsion model is useful for predicting the molecular structures of molecules formed from nonmetals.
The structure around a given atom is determined by minimizing repulsions between electron pairs.
The bonding and nonbonding electron pairs (lone pairs) around a given atom are positioned as far apart as possible.

29. Molecular Structure: VSEPR
Steps for Predicting Molecular Structure Using the VSEPR Model
1. Draw the Lewis structure for the molecule.
2. Count the electron pairs and arrange them in the way that minimizes repulsion (that is, put the lone pairs as far apart as possible).
3. Determine the positions of the atoms from the way the electron pairs are shared.
4. Determine the name of the molecular structure from the positions of the atoms.

31. Example
Br2
CO2
CF4
PF3

32. Your turn
NH4+
XeF4
AsI5
SF3 +
I3 -