Philip Dutton University of Windsor, Canada N9B 3P4 Prentice-Hall © 2002 General Chemistry Principles and Modern Applications Petrucci Harwood Herring 8 th Edition Chapter 11: Chemical Bonding II: Additional Aspects

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 2 of 47 Contents 11-1What a Bonding Theory Should Do 11-2Introduction to the Valence-Bond Method 11-3Hybridization of Atomic Orbitals 11-4Multiple Covalent Bonds 11-5Molecular Orbital Theory 11-6Delocalized Electrons: Bonding in the Benzene Molecule 11-7Bonding in Metals Focus on Photoelectron Spectroscopy

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 3 of What a Bonding Theory Should Do Bring atoms together from a distance. –e - are attracted to both nuclei. –e - are repelled by each other. –Nuclei are repelled by each other. Plot the total potential energy versus distance. –-ve energies correspond to net attractive forces. –+ve energies correspond to net repulsive forces.

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 4 of 47 Potential Energy Diagram

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 5 of Introduction to the Valence-Bond Method Atomic orbital overlap describes covalent bonding. Area of overlap of orbitals is in phase. A localized model of bonding.

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 6 of 47 Bonding in H 2 S

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 7 of 47 Example 11-1 Using the Valence-Bond Method to Describe a Molecular Structure. Describe the phosphine molecule, PH 3, by the valence-bond method.. Identify valence electrons:

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 8 of 47 Example 11-1 Sketch the orbitals: Overlap the orbitals: Describe the shape: Trigonal pyramidal

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 9 of Hybridization of Atomic Orbitals

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 10 of 47 sp 3 Hybridization

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 11 of 47 sp 3 Hybridization

Methane sp 3 hybrid orbitals General ChemistrySlide 12 /69 Tetrahedron in a cube Coordinates: sxyz Gábor I. Csonka © 2013

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 13 of 47 Bonding in Methane

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 14 of 47 sp 3 Hybridization in Nitrogen

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 15 of 47 Bonding in Nitrogen

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 16 of 47 sp 2 Hybridization

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 17 of 47 Orbitals in Boron

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 18 of 47 sp Hybridization

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 19 of 47 Orbitals in Beryllium

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 20 of 47 sp 3 d and sp 3 d 2 Hybridization

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 21 of 47 Hybrid Orbitals and VSEPR Write a plausible Lewis structure. Use VSEPR to predict electron geometry. Select the appropriate hybridization.

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 22 of Multiple Covalent Bonds Ethylene has a double bond in its Lewis structure. VSEPR says trigonal planar at carbon.

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 23 of 47 Ethylene

Double bond = 1 C-C  + 1 C-C  bonds. + - Notice that the C-C  MO is antisymmetric to the molecular plane, it is positive above and negative below the plane.

General Chemistry: Chapter 12Slide 25 of 47 Acetylene Acetylene, C 2 H 2, has a triple bond. VSEPR gives linear at carbon. C: sp hybrid There are two  bonds in acetylene (ethyne)

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 26 of Molecular Orbital Theory Atomic orbitals are isolated on atoms. Molecular orbitals span two or more atoms. LCAO –Linear combination of atomic orbitals. Ψ 1 = φ 1 + φ 2 Ψ 2 = φ 1 - φ 2

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 27 of 47 Combining Atomic Orbitals

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 28 of 47 Molecular Orbitals of Hydrogen

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 29 of 47 Basic Ideas Concerning MOs Number of MOs = Number of AOs. Bonding and antibonding MOs formed from AOs. e - fill the lowest energy MO first. Pauli exclusion principle is followed. Hund’s rule is followed

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 30 of 47 Bond Order Stable species have more electrons in bonding orbitals than antibonding. Bond Order = No. e - in bonding MOs - No. e- in antibonding MOs 2

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 31 of 47 Diatomic Molecules of the First-Period BO = (1-0)/2 = ½ H2+H2+ BO = (2-0)/2 = 1 H2H2 BO = (2-1)/2 = ½ He 2 + BO = (2-2)/2 = 0 He 2 BO = (e - bond - e - antibond )/2

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 32 of 47 Molecular Orbitals of the Second Period First period use only 1s orbitals. Second period have 2s and 2p orbitals available. p orbital overlap: –End-on overlap is best – sigma bond (σ). –Side-on overlap is good – pi bond (π).

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 33 of 47 Molecular Orbitals of the Second Period

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 34 of 47 Combining p orbitals

The energy level diagram for MOs General Chemistry: Chapter 12Slide 35 of 47 The AOs are on the left and right (yellow and blue), the MOs are in the center (orage and green). MO labels are  and  for bonding  * and  * for antbonding.

Inversion of the MO energies between N 2 and O 2 General Chemistry: Chapter 12Slide 36 of 47 For diatomic molecules with 7  Z the  2p orbital energy is above  2p molecular orbital energy. The latter is filled first.

General Chemistry: Chapter 12Slide 37 of 47 MO Diagrams of 2 nd Period Diatomics C 2 electron configuration:

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 38 of 47 MO Diagrams of Heteronuclear Diatomics

General Chemistry: Chapter 12Slide 39 of Delocalized Electrons

General Chemistry: Chapter 12Slide 40 of 47 Benzene: delocalized bonding Delocalized bonding is not localized between two atoms: instead, each link has a 'fractional double bond order’. There is a corresponding 'delocalization energy', identifiable with the stabilization of the system compared with a hypothetical alternative in which formal (localized) single and double bonds are present. Bond order = 1.5 Delocalization energy = –263.7 kJ mol -1

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 41 of 47 Benzene, 6 MO energies

Benzene 6 MO shapes Slide 42 /69 Bonding  orbitals 6 e - Non bonding  orbitals

The charge distribution in benzene Általános Kémia, Kötés szerkezetSlide 43 /69 The electrostatic potential surface. Red: negative charge Blue: positive charge

The charge distribution in borazine Általános Kémia, Kötés szerkezetSlide 44 /69 Uneven charge distribution Less stable 6 electrons delocalized  system

Pyridine Slide 45 /69 _ _ _ 6 delocalised electrons

Butadiene Slide 46 /69 4 delocalised electrons

CO 3 2- Slide 47 /69 6 delocalised electrons 6 e 2 e from the 2 negative charge + 4 e from the 2p atomic orbitals of the 3 O and 1 C atoms..

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 48 of 47 Ozone

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 49 of Bonding in Metals Electron sea model –Nuclei in a sea of e -. –Metallic lustre. –Malleability. Force applied

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 50 of 47 Bonding in Metals Band theory. Extension of MO theory. N atoms give N orbitals that are closely spaced in energy. N/2 are filled. The valence band. N/2 are empty. The conduction band.

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 51 of 47 Band Theory

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 52 of 47 Semiconductors

General Chemistry: Chapter 12Slide 53 of 47 Photovoltaic Cells

Solar cell details General Chemistry: Chapter 12Slide 54 of 47

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 55 of 47 Focus on Photoelectron Spectroscopy

Prentice-Hall © 2002General Chemistry: Chapter 12Slide 56 of 47 Chapter 12 Questions 1, 3, 8, 10, 16, 29, 33, 39, 45, 59, 68, 72, 76