1. Chapter 10 Chemical Bonding II

2. Valence Bond Theory
Linus Pauling and others applied the principles of quantum mechanics to molecules
they reasoned that bonds between atoms would arise when the orbitals on those atoms interacted to make a bond
the kind of interaction depends on whether the orbitals align along the axis between the nuclei, or outside the axis
Valence Bond Theory: A quantum mechanical model which shows how electron pairs are shared in a covalent bond.
Bond forms between two atoms when the following conditions are met:
Covalent bonds are formed by overlap of atomic orbitals, each of which contains one electron of opposite spin.
Each of the bonded atoms maintains its own atomic orbitals, but the electron pair in the overlapping orbitals is shared by both atoms.
The greater the amount of overlap, the stronger the bond.

3. Orbital Interaction
In some cases, atoms use “simple” atomic orbital (e.g., 1s, 2s, 2p, etc.) to form bonds.
In other case, they use a “mixture” of simple atomic orbitals known as “hybrid” atomic orbitals.
as two atoms approached, the partially filled or empty valence atomic orbitals on the atoms would interact to form molecular orbitals
the molecular orbitals would be more stable than the separate atomic orbitals because they would contain paired electrons shared by both atoms

4. Valence Bond Theory - Hybridization
one of the issues that arose was that the number of partially filled or empty atomic orbital did not predict the number of bonds or orientation of bonds
C = 2s22px12py12pz0 would predict 2 or 3 bonds that are 90° apart, rather than 4 bonds that are 109.5° apart
to adjust for these inconsistencies, it was postulated that the valence atomic orbitals could hybridize before bonding took place
one hybridization of C is to mix all the 2s and 2p orbitals to get 4 orbitals that point at the corners of a tetrahedron

5. Valence Bond Theory Main Concepts
the valence electrons in an atom reside in the quantum mechanical atomic orbitals or hybrid orbitals
a chemical bond results when these atomic orbitals overlap and there is a total of 2 electrons in the new molecular orbital
the electrons must be spin paired
the shape of the molecule is determined by the geometry of the overlapping orbitals

6. Types of Bonds
a pi (p) bond results when the bonding atomic orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nuclei
between unhybridized parallel p orbitals
the interaction between parallel orbitals is not as strong as between orbitals that point at each other; therefore s bonds are stronger than p bonds
a sigma (s) bond results when the bonding atomic orbitals point along the axis connecting the two bonding nuclei
either standard atomic orbitals or hybrids
s-to-s, p-to-p, hybrid-to-hybrid, s-to-hybrid, etc.

7. Tro, Chemistry: A Molecular Approach

8. Hybridization some atoms hybridize their orbitals to maximize bonding
hybridizing is mixing different types of orbitals to make a new set of degenerate orbitals
sp, sp2, sp3, sp3d, sp3d2
more bonds = more full orbitals = more stability
better explain observed shapes of molecules
same type of atom can have different hybridization depending on the compound
C = sp, sp2, sp3

9. Hybrid Orbitals H cannot hybridize!!
the number of standard atomic orbitals combined = the number of hybrid orbitals formed
the number and type of standard atomic orbitals combined determines the shape of the hybrid orbitals
the particular kind of hybridization that occurs is the one that yields the lowest overall energy for the molecule
in other words, you have to know the structure of the molecule beforehand in order to predict the hybridization

10. sp3 Hybridization of C

11. sp3 Hybridization atom with 4 areas of electrons tetrahedral geometry
109.5° angles between hybrid orbitals
atom uses hybrid orbitals for all bonds and lone pairs

12. Methane Formation with sp3 C
Ammonia Formation with sp3 N

13. sp2 atom with 3 areas of electrons trigonal planar system
C = trigonal planar
N = trigonal bent
O = “linear”
120° bond angles
flat
atom uses hybrid orbitals for s bonds and lone pairs, uses nonhybridized p orbital for p bond

14. 3-D representation of ethane (C2H4)

15. Bond Rotation
because orbitals that form the s bond point along the internuclear axis, rotation around that bond does not require breaking the interaction between the orbitals
but the orbitals that form the p bond interact above and below the internuclear axis, so rotation around the axis requires the breaking of the interaction between the orbitals

16. sp atom with 2 areas of electrons linear shape 180° bond angle
atom uses hybrid orbitals for s bonds or lone pairs, uses nonhybridized p orbitals for p bonds

17. sp3d atom with 5 areas of electrons around it trigonal bipyramid shape
See-Saw, T-Shape, Linear
120° & 90° bond angles
use empty d orbitals from valence shell
d orbitals can be used to make p bonds

18. sp3d2 atom with 6 areas of electrons around it octahedral shape
Square Pyramid, Square Planar
90° bond angles
use empty d orbitals from valence shell
d orbitals can be used to make p bonds

19. Five types of hybrid are shown below
# of e- groups around central atom
Hybrid orbitals used
Orientation of Hybrid Orbitals 2 sp 3
sp2 4
sp3 5
sp3d 6
sp3d2

20. Predicting Hybridization and Bonding Scheme
Start by drawing the Lewis Structure
Use VSEPR Theory to predict the electron group geometry around each central atom
Use Table 10.3 to select the hybridization scheme that matches the electron group geometry
Sketch the atomic and hybrid orbitals on the atoms in the molecule, showing overlap of the appropriate orbitals
Label the bonds as s or p

21. Examples:
Predict the Hybridization and Bonding Scheme of All the Atoms in
Then sketch a σ framework and a π framework
CH3CHO
CH2NH
H3BO3

22. Problems with Valence Bond Theory
VB theory predicts many properties better than Lewis Theory
bonding schemes, bond strengths, bond lengths, bond rigidity
however, there are still many properties of molecules it doesn’t predict perfectly
magnetic behavior of O2