1. Chapter 6.5, 7.1, 7.2 – Chemical Bonding – Fundamentals and Theories
CHM1111 Section 04
Instructor: Dr. Jules Carlson
Class Time: M/W/F 1:30-2:20
Friday, October 28th

2. Mid-term Mark Adjustment
Due to the issue with not providing the Ionization Energy for H, Marks for the long answer #2 will be adjusted as follows:
If you got 0/5 your mark is raised to 2/5
If you got 1/5 or 2/5 your mark is raised to 3/5
If you got 3/5 your mark is raised to 4/5
This brings the class average up from 55 % to 61 %.

3. Dipole Moment Example
After determining the Lewis structure and geometry, provide drawings with the dipole moments for the following compounds, and indicate if the dipole moment will be:
large (μ > 1 D), medium (1 D > μ > 0.5 D), or small (μ < 0.5 D)
NO2, PCl5, CH2Cl2, HONH2, SF4

4. SF4 Dipole Moment
Net Dipole moment is only zero if all atoms around the central atom have the same electronegativity (generally only if all the same atom)
AND the bonding arrangement has no lone pairs.
Dipole moment for SF4 is D
Tolles and Gwinn, 1961

5. Pentagonal Bipyrimidal
For Steric number of 7
Has 5 equatorial and 2 axial
bonding locations
Bond angles are 90⁰, 72⁰.
XeF6 would have Pentagonal bipyrimidal geometry with one lone pair, except experimentally shown to be octahedral.
VSEPR does not always apply.

6. Bond Lengths Bond length trends:
The bond length of a covalent bond is the nuclear separation distance where the molecule is most stable.
Bond length trends:
The larger the radii of the atoms, the longer the bond length.
The larger the nuclear charge, the smaller the atomic size, the shorter the bond length
The larger the bond polarity, the larger the separation of charge, the shorter the bond length
Multiple bonds have shorter bond length

7. Average Bond Lengths
Look down H-F, H-Cl, H-Br, increasing due to larger atom size.
Look across C-C, C-N, C-O, C-F, increasing bond polarity, larger nuclear charge reducing atom size, reduced bond length.
C=C (135 pm) shorter than C-C (154 pm).
N≡N (110 pm) shorter than N-N (145 pm).

8. Example 6-18
Which factor accounts for the following differences in bond length?
I2 has a longer bond than Br2.
C-N bonds are shorter than C-C bonds.
H-C bonds are shorter than the C≡O bond.
The carbon-oxygen bond in formaldehyde, H2C=O is longer than that in carbon monoxide C≡O.

9. Bond Energy
Remember that bond energy is the difference in energy in the potential well that a molecule must have to break that bond.
Potential Well

10. Trends in Bond Energy
Bond strength increases as more electrons are shared between the atoms.
Bond strength increases as the electronegativity difference (Δχ) between bonded atoms increases.
Bond strength decreases as bonds become longer.
O—O
Δχ = 0.0
BE = 145 kJ/mol
O — N
Δχ = 0.5
BE = 200 kJ/mol
O —C
Δχ = 1.0
BE = 360 kJ/mol
H—F
92 pm
BE = 565 kJ/mol
H—Cl
127 pm
BE = 430 kJ/mol
H — Br
141 pm
BE = 360 kJ/mol
H — I
161 pm
BE = 295 kJ/mol

11. Delocalization
A framework that can be used to describe bonding is to describe bonds as localized or delocalized.
Localized bonds are described in Lewis structures, where pairs of electrons are shared between two atoms in a bond, or localized on a single atom.
Describing bonds as delocalized indicates that they are not always associated with one atom, and this description is needed to understand some chemical properties (i.e. bonding in NO2)

12. I Clicker Question Which of the following statements are true:
SF6 exhibits octahedral geometry.
N=N bond length is shorter than N-N bond length
Bond angles in tetrahedral compounds always show bond angles of 109.5⁰ or less
Both (a) and (b) are true
All statements are true

13. Orbital Overlap
Since electrons have wave-like properties, we can add or subtract the wavefunctions used to describe their orbitals.
If you sum orbitals, their amplitudes may be added or subtracted depending upon the sign of the waves.
Orbital Overlap: As two hydrogen atoms approach each other, the overlap of their 1s atomic orbitals increases. The wave amplitudes add, generating a new orbital with high electron density between the nuclei.

14. Orbital Overlap Model Conventions
Each electron in a molecule is assigned to a specific orbital.
No two electrons in a molecule have identical descriptions, because Pauli exclusion principle applies to electrons in molecules as well as in atoms.
The electrons in molecules obey the aufbau principle, meaning that they occupy the most stable orbitals available to them.
Even though every atom has an unlimited number of atomic orbitals, the valence orbitals are all that are needed to describe bonding.
Chemistry, Canadian Edition ©2010 John Wiley & Sons Canada, Ltd.
Chapter 9

15. Orbital Overlap for F2 and HF

16. Bonding in Hydrogen Sulphide, Phosphine
Sometimes we need to use different bonding models for different applications, VSEPR does not work for H2S or PH3.
Bond angles in H2S are 92.1⁰ and in PH3 are 93.6⁰, far from tetrahedral geometry.
Bond angles in H2O are 107.5⁰ and in PCl3 are 106.2⁰ (follows VSEPR).
H is small compared to S, or P.

17. Hybridization
An sp3 orbital
So if p bonds are at 90⁰ angles to each other, how do we get tetrahedral geometry?
In methane, valence electrons are 2s, 2p electrons.
The orbitals for these electrons can be combined to from a special set of directional orbitals.
This is referred to as hybridization.
4 sp3 orbitals around C with 1s orbitals from H

18. Hybridization for Tetrahedral Geometry
Look at Methane
Forms sp3 hybrid
orbitals.
If Steric number # = 4, hybridizes sp3.
So how does NH3 bond?
How does H2O bond?
Textbook example 7.3 Methanol:

19. Principles for Hybridization
The number of valence orbitals generated by the hybridization process equals the number of valence atomic orbitals participating in hybridization.
The steric number of an inner atom uniquely determines the number and type of hybrid orbitals.
Hybrid orbitals form localized bonds by overlap with atomic orbitals or with other hybrid orbitals.
There is no need to hybridize orbitals on outer atoms, because atoms do not have limiting geometries. Hydrogen always forms localized bonds with its 1s orbital. The bonds formed by all other outer atoms can be described using valence p orbitals.

20. Types of Hybridized Orbitals
Steric Number
Geometry
Hybridization 2
Linear sp 3
Trigonal Planar
sp2 4
Tetrahedral
sp3 5
Trigonal Bipyrimidal
sp3d 6
Octahedral
sp3d2 7
Pentagonal Bipyrimidal
sp3d3
We will look at the hybridization of each of these in turn with molecule examples.