Chapter 9

Molecular Geometry and Bonding Theories

Chapter 9 suggested problems
10th Ed. and 11th Ed.: 21, 27, 35, 47, 49, 51, 76, 96

Class Notes

1. Hybrid orbitals
1. Orbital shapes
1. s
2. p
3. d

2. Orbital overlap is essential for chemical bonding to occur; the greater the overlap the stronger the bond
1. Show overlap between hydrogen and oxygen in water
2. Overlap is an important factor in bond strength but certainly not the only factor

3. If carbon has four valence electrons (2s² 2p²), how can it form four bonds?
1. In the 1930s Linus Pauling suggested that one of the 2s electrons is promoted to the empty 2p orbital, resulting in the formation of four hybrid orbitals that are intermediate in energy and shape
2. Diagram of orbitals vs. energy
3. This hybridization theory can be easily and accurately used to describe bonding in many molecules

   atom	molecule	valence configuration	hybridization	remaining orbitals	geometry	bond angles
   Be	BeH2	2s2	sp	2 p orbitals	linear	180
   B	BF3	2s2 2p1	sp2	1 p orbitals	trigonal planar	120
   C	CH4	2s2 2p2	sp3	none	tetrahedral	109.5
   P	PCl5	3s2 3p3	dsp3	4 d orbitals	trigonal bipyramidal	120, 90
   S	SF6	3s2 3p4	d2sp3	3 d orbitals	octahedral	90

4. Multiple bonds
1. Double bonds - consist of one sigma bond and one pi bond
1. Sigma bonds
1. Cylindrical probability distribution around the bond axis
2. Sigma bonds are created by the overlap of hybridized orbitals

2. Pi bonds
1. Probability out of the plane (above and below) the bond axis
2. Pi bonds are created by the overlap of the unhybridized p orbitals

2. Triple bonds consist of one sigma bond and two pi bonds

2. VSPER - Valence Shell Paired-Electron Repulsion
1. There is an excellent correlation between the type of hybridization that occurs in a molecule, the number of bonds formed, the number of nonbonding pairs of electrons on the central atom, and the geometry of the molecule
2. The mutual repulsion of pairs of electrons cause them to arrange themselves as far apart spatially as possible
3. Lone pairs take up slightly more space than bonding pairs and are more repulsive as a consequence
4. Multiple bonds count as one bonding pair - see Tables 9.2, 9.3

   total pairs	bonding pairs	nonbonding pairs	geometry	bond angles	hybridization	example
   2	2	0	linear	180	sp	BeH2
   3	3	0	trigonal planar	120	sp2	BF3
   	2	1	bent		sp2	NO2
   4	4	0	tetrahedral	109.5	sp3	CH4
   	3	1	pyramidal		sp3	NH3
   	2	2	bent		sp3	H2O
   5	5	0	trigonal bipyramidal	120, 90	dsp3	PCl5
   	4	1	seesaw		dsp3	SF4
   	3	2	t-shaped		dsp3	ClF3
   	2	3	linear		dsp3	XeF2
   6	6	0	octahedral	90	d2sp3	SF6
   	5	1	square pyramidal		d2sp3	BrF5
   	4	2	square planar		d2sp3	XeF4

3. Polar bonds and polar molecules
1. Polar covalent bonds - bonding electrons are not shared equally due to differences in electronegativity between the bonding atoms
1. If the bonding atoms are the same (e.g. H₂, N₂, O₂, F₂, etc.) the bonding atoms are shared equally (50%/50%)
2. If the bonding atoms are not the same the bonding electrons spend more time around the more electronegative atom
3. This causes partial charges to arise on the bonding nuclei

2. As a general rule, molecules with one or more polar bonds are themselves polar which affects interactions with other molecules
3. The polarity of molecules is measured in Debye units (D)
4. Polar molecules are said to have a dipole moment, the magnitude of which indicates the magnitude of the partial charges in the molecule
5. The exceptions: those molecules with polar bonds in which symmetry results in negation of molecular polarity
6. Lone pairs ruin this, negating symmetry unless all axial or equatorial positions are filled with lone pairs
7. Examples
1. BF₃ vs. BF₂H
2. CCl₄ vs. CCl₃H
3. NH₃, ClF₃, XeF₂, BrF₅ and XeF₄