1. Chapter 6.2-6.3 – Fundamentals of Chemical Bonding
CHM1111 Section 04
Instructor: Dr. Jules Carlson
Class Time: M/W/F 1:30-2:20
Wednesday, October 19th

2. How to Draw Electron Density and Overlap
Determine the valence electrons for each atom.
Look at the electronegativity difference and assess if bond is covalent, polar-covalent, or ionic.
Draw the outermost electrons for each atom, remember s orbitals are spherical and p orbitals are lobe-shaped.
Note: Position the orbitals in such a way to maximize the amount of overlap.
4. Show charge separation if any, δ means partial charge. Show partial charges for Δχ > 0.5 between atoms.

3. Orbital Shapes – s orbitals
s orbitals are spherical, and have one node fewer than n.
Nodes occur where the wave has a minimal amplitude (see circles).
Waves can also have different phases (see with p orbitals)

4. Orbital shapes – p orbitals
p orbitals are lobed-shaped
The three p orbitals are oriented across the three axes
There is a node which separates each p-orbital into 2 phases, the electron occupies both phases
Oriented along z-axis
Oriented along y-axis
Oriented along x-axis

5. Lewis Structures
Lewis structure: Drawing of a molecule showing how atoms are bonded together and reveals the distribution of bonding and non-bonding valence electrons.

6. Lewis Structure Conventions
Each atom is represented by its elemental symbol.
Only valence electrons are shown.
A line joining two elemental symbols represents one pair of electrons shared in a bond between two atoms (can have double, triple bonds too).
Dots placed next to an elemental symbol represent non-bonding electrons.

7. Bonding
Outer atoms only bond to one atom, and inner atoms bond to more than one atom.
Hydrogen atoms are always outer atoms
More electronegative atoms are outer atoms
Order of atom appearance in formula is often the bonding pattern (e.g HCN is H-C-N) – not true for oxoacids.

8. Rules for Building Lewis Structures
Count the valence electrons.
Assemble the bonding framework, placing 2 electrons per bond.
Place three non-bonding pairs of electrons on each outer atom except H.
Assign the remaining electrons to inner atoms to complete their octets.
Optimize electron configurations of the inner atoms.
Identify equivalent or near-equivalent Lewis structures.

9. First Example Dichloromethane CH2Cl2
Count the valence electrons. If anion, add one electron for each negative charge. If cation, remove one electron for each positive charge.
C – group 14: s2p2, H- group 1: s1
Cl – group 17: s2p5
So 7 for Cl x 2 Cl + 4 for C + 1 for H x 2 H = 20 e-
Assemble the bonding framework:
Remember H’s on outside and electronegative atoms.

10. First Example Continued
Place 3 non-bonding pairs on each outer atom except H.
Non-bonding pairs are also called lone pairs.
Assign the remaining valence electrons to inner atoms

11. Example Where Optimizing is Needed
Work through as we did with dichloromethane to step 4 for Formaldehyde H2CO.
Look at number of bonds to Carbon, how can we adjust this?

12. I Clicker Question Draw the Lewis Structure for PBr3
Which of the following statements are true:
Each bromine has 3 lone pairs.
Phosphorus has no lone pairs.
Phosphorus has one lone pair.
Both (a) and (b)
Both (a) and (c)

13. Example with multi-inner atoms
Draw the Lewis Structure for (CH3)2CCH2

14. Where the Octet Rule Does Not Apply
Draw Lewis Structures for the following compounds:
PBr5
O2SCl2
KrF4
Also, go through example 6.10 for Lewis Structure of an Oxoanion (p. 318)

15. Resonance Structures Draw the Lewis Structure for the nitrate ion.
Draw the Lewis Structure for ozone, O3.

16. Resonance Structures

17. Hints for Lewis Structures
Octet rule always applies for elements with n < 3, may apply elsewhere too.
Carbon forms 4 bonds.
Hydrogen typically forms one bond to other atoms.
When multiple bonds are forming, they are usually between C, N, O or S.
Nonmetals can form single, double, and triple bonds, but not quadruple bonds.
Always account for single bonds and lone pairs before forming multiple bonds.
Try to get all formal charges = 0, if impossible, place negative charges on electronegative atoms.
Look for resonance structures.