1. June 9, 2009 – Class 35 and 36 Overview
11.3 Hybridization of Atomic Orbitals
Hybrid orbitals properties including energy and shape, types (sp, sp2, sp3, sp3d, sp3d2) found in common molecules, hybrid orbitals and VSEPR theory.
11.4 Multiple Covalent Bonds
Bonding using hybridization, sigma and pi bonds. 1

2. What a Bonding Theory Should Do
Lewis theory is simple and structures can be determined rapidly.
It does not account for odd-electron species, resonance structures or the magnetic and spectral properties of molecules.
Example: Why is O2 paramagnetic?
VSEPR theory allows shape predictions
Neither yield quantitative information about bond lengths or energies
What a Bonding Theory Should Do 2

3. Introduction to Valence Bond Theory
Valence-bond method: treats a covalent bond in terms of the overlap of pure or hybridized orbitals. Electron probability (or electron charge density) is concentrated in the area of overlap.
This theory tells us what a covalent bond is and correlates molecular shapes to the interactions of atomic orbitals.
The basic principle of valence bond theory is that a covalent bond forms when half filled orbitals on two different atoms (atomic orbitals) overlap. Example: H2
Introduction to Valence Bond Theory 3

4. Introduction to Valence Bond Theory
Localized electron model: according to valence bond theory, core electrons and lone-pair electrons retain the same orbital locations as in the separated atoms.
Charge density of the bonding electrons is concentrated in regions of orbital overlap. Example: Bonding in H2S.
Note: (+) and (-) signs denote phase signs, not charges!
Introduction to Valence Bond Theory 4

5. Hybridization of Atomic Orbitals
In order to account for many molecular shapes, atomic orbitals must be hybridized.
Hybridization: refers to combining pure atomic orbitals to generate hybrid orbitals in the valence bond approach to covalent bonding.
Hybrid orbital: is one of a set of identical orbitals reformulated from pure atomic orbitals and used to describe certain covalent bonds.
Example: Tetrahedral carbon
Hybridization of Atomic Orbitals 5

6. Hybridization of Atomic Orbitals
Hybrid orbital properties
The number of hybrid orbitals equals the total number of atomic orbitals that are combined.
Hybridization rationalizes experimentally determined shape, it is not an actual physical phenomenon.
Atomic orbital energy is conserved upon hybridization.
Example: For tetrahedral C, the p orbitals each move down ¼ of the energy difference between the s and p orbitals, while the s orbitals move up by ¾.
Hybridization of Atomic Orbitals 6

7. Hybridization of Atomic Orbitals
sp3 hybrid orbital: one of the four orbitals formed by the hybridization of one s and three p orbitals. The angle between any two of the orbitals is the tetrahedral angle, 109.5o.
Hybridization of Atomic Orbitals 7

8. Hybridization of Atomic Orbitals8

9. Hybridization of Atomic Orbitals
Bonding and structure of tetrahedral methane (CH4) – an sp3 hybridized molecule. 9

10. Hybridization of Atomic Orbitals
Bonding and structure of trigonal pyramidal methane (NH3) – an sp3 hybridized molecule.
Note that hybrid orbitals can accommodate lone pair electrons as well as bonding pairs. 10

11. Hybridization of Atomic Orbitals
sp2 hybrid orbital: one of the three orbitals formed by the hybridization of one s and two p orbitals. The angle between any two of the orbitals is 120o.
This hybridization scheme is common to most boron containing compounds.
Hybridization of Atomic Orbitals 11

12. Hybridization of Atomic Orbitals12

13. Hybridization of Atomic Orbitals
Bonding and structure of trigonal planar BF3 – an sp2 hybridized molecule. 13

14. Hybridization of Atomic Orbitals
sp hybrid orbital: one of the pair of orbitals formed by the hybridization of one s and one p orbital. The angle between the two orbitals is 180o.
This hybridization scheme is common to most beryllium containing compounds.
Hybridization of Atomic Orbitals 14

15. Hybridization of Atomic Orbitals15

16. Hybridization of Atomic Orbitals
Bonding and structure of linear BeCl2 – an sp hybridized molecule. 16

17. Hybridization of Atomic Orbitals17

18. Hybridization of Atomic Orbitals
sp3d hybrid orbital: one of the five orbitals formed by the hybridization of one s, three p, and one d orbital. The five orbitals are directed to the corners of a trigonal bipyramid.
Example: Hybridization of phosphorus, P.
Hybridization of Atomic Orbitals 18

19. Hybridization of Atomic Orbitals
Bonding and structure of trigonal bipyramidal PCl5 – an sp3d hybridized molecule.
This hybridization scheme also accounts for the shapes of seesaw, t-shaped and some linear molecules. 19

20. Hybridization of Atomic Orbitals
sp3d2 hybrid orbital: one of the six orbitals formed by the hybridization of one s, three p and two d orbitals. The six orbitals are directed to the corners of a regular octahedron.
Hybridization of Atomic Orbitals 20

21. Hybridization of Atomic Orbitals
Bonding and structure of octahdral SF6 – an sp3d2 hybridized molecule.
This hybridization scheme also accounts for the shapes of square pyramidal and square planar molecules. 21

22. Hybridization of Atomic Orbitals22

23. Hybridization of Atomic Orbitals23

24. Hybrid orbitals and VSEPR theory
VSEPR method uses empirical data to give an approximate molecular geometry, whereas the valence bond method relates to the orbitals used in bonding based on a given geometry.
We can choose the likely hybridization scheme for a central atom in a structure in the valence-bond method by:
writing a plausible Lewis structure for the species of interest
using VSEPR theory to predict the probable electron-group geometry of the central atom.
selecting the hybridization scheme corresponding to the electron-group geometry.
Hybrid orbitals and VSEPR theory 24

25. Hybrid orbitals and VSEPR theory
Problem: Predict the shape of the following molecules and a hybridization scheme consistent with this prediction
SiF4
XeF4
Problem: Describe the molecular geometry and propose a plausible hybridization scheme for the central atom in the ion:
Cl2F+
BrF4+
Hybrid orbitals and VSEPR theory 25

26. Multiple Covalent Bonds
sigma (s) bonds: results from the end-to-end overlap of simple or hybridized atomic orbitals along the straight line joining the nuclei of the bonded atoms.
pi (p) bonds: results from the side-to-side overlap of p orbitals, producing a high electron charge density above and below the line joining the bonded atoms.
Multiple Covalent Bonds 26

27. Multiple Covalent Bonds
C2H6 - Molecule should be tetrahedral (sp3 hybridized) about each C atom.
σ bonding; end-to-end overlap of an sp3
hybridized orbital from each carbon 27

28. Multiple Covalent Bonds
C2H6 - Molecule should be tetrahedral (sp3 hybridized) about each C atom.
C2H4 - Molecule should be trigonal planar (sp2 hybridized) about each C atom. 28

29. Multiple Covalent Bonds – C2H4 (ethylene)
Molecular shape is determined by the orbitals forming the s-bonds (s-framework).
Rotation about the double bond is severely restricted, and the double bond is rigid.
Twisting one of the –CH2 groups out of plane would reduce the amount of p-orbital overlap and weaken the p bond. 29

30. Multiple Covalent Bonds – C2H4 (ethylene)30

31. Multiple Covalent Bonds
C2H6 - Molecule should be tetrahedral (sp3 hybridized) about each C atom.
C2H4 - Molecule should be trigonal planar (sp2 hybridized) about each C atom.
C2H2 - Molecule should be linear (sp hybridized) about each C atom. 31

32. Multiple Covalent Bonds – C2H2 (acetylene)32

33. Multiple Covalent Bonds – C2H2 (acetylene)33

34. Multiple Covalent Bonds34

35. Multiple Covalent Bonds
Problem: Describe the types of bonds and orbitals present for:
(a) acetone (CH3)2CO
(b) HCN
(c) CO2
Problem: Hydrazine, N2H4, and carbon disulfide, CS2, form a cyclic molecule with the following Lewis structure. How do shape and hybridization about C and N change when hydrazine and CS2 form this product?
Multiple Covalent Bonds 35