Unit 8: Valence Bond Theory

Introduction:

Unit 8,"Valence Bond Theory" is the last unit in the Chem 211 series. Valence Bond Theory was developed by Linus Pauling and others in the late 1930's to explain bonding and molecular geometry with one underlining theory. The electron pair bonding theory of G.N, Lewis had provided a frame work for explaining molecular bonding but it could not discuss the "shapes" of molecules and had difficulty with certain "resonance" structures which seem to have "fractional" bond orders.The everyday discussion of bonding in organic chemistry is a blend of Lewis structures and the ideas of the valence bond method.

It would be a good to review Unit 7 before beginning this unit. VSEPR was the system used to predict molecular geometry. Using the valence bond approach, the geometry of the molecule is determined by which "hybrid" atomic orbitals form the hybrid orbital set used by any particular atom for bonding. The most important atoms to consider are C, N and O, so most of the introduction to the "VB" Theory will focus on these atoms. As promised, Unit 8 is much shorter then the previous ones. Unit 8 begins by recalling some of the concepts of VSEPR and than introduces "hybrid" orbitals, sp^3, sp^2 and sp as well as dsp^3 and d^2sp^3 and there respective bonding geometries. Just like Unit 7, Unit 8 is highly pictorial and some what difficult in this medium.

Unit Objectives: After completing Unit 8, you should;

1. to describe molecular bonding using valence bond concepts of hybridized atomic orbitals and bonds arising from the mutual"overlap" of atomic and/or hybrid orbitals.

2. to define a "sigma" bond and a "pi" bond using the valence bond method

3. to discuss the predicted geometry around certain atoms in molecules based on their orbital hybridizations. (See the table on page 489).

4. to make sketches of the hybrid orbitals, sp, sp^2 and sp^3 and show how a sigma or pi bond can be formed from the appropriate orbital overlap.

Reading and Study Assignment:

Text: Chapter 10, "Bonding and Molecular Structure: Orbital Hybridization etc." pages 457 to 475. Much of the discussion in the text and ROM is summarized in Fig.10.10 on page 468. Keep this figure handy as you work on the problems and quiz.

CD-ROM Review Screen 9.16 "Determining Molecular Shape" and the "models" file. Chapter 10 , Screens 10.3, 10.4 10.5 10.6 and 10.7 are assigned for this unit.

Study Hints:

1. A good idea to determine the possible hybridization of some atomic center in a given structure is to consider the VSEPR model for the geometry about that atomic center and then pick the appropriate hybrid orbital consistent with the pre-determined geometry,

2. If you have difficulty assigning hybrid orbitals to atoms in organic structures recall that the "first" bonding pair form a"sigma" bond and then if other bonding pairs are involved, they a "pi" bonds. The "sp^x" hybridization is "sp" with "p " raised to the number of sigma bonds and/ or lone pairs minus 1. For example in ethylene, H2C=CH2, each "C" atom is connected to three sigma bonds, one to each hydrogen and one toe the neighboring C atom, and one "pi" bonding pair between the carbons. Therefore around each C is 3 sigma bonds so the hybridization is sp^2 because the hybrid orbital is given by sp^3-1 = sp^2.

Start with Screen 10.3 "Valence Bond Theory" Do question # 2, only.

Screen 10.4 "Hybrid Orbitals" Remember just as many atomic orbitals are used to form the same number of "hybrid orbitals" For example, an "s" an a "p" atomic orbital hybridize to form two "sp" hybrid orbitals and an "s" and three "p" orbitals hybridize to form an "sp^3" hybrid and so on. Omit the question for Screen 10.4.

Screen 10.5 "Sigma Bonding" This represent one of the principle concepts of valence bond thinking. Do the questions associated with this screen and these will be answered in the Answer Book.

Screen 10.6 "Determining Hybrid Orbitals" Answer all the questions for this screen and do Exercise 10.1 on page 10-3 of the workbook.

Last screen, Screen 10.7 "Multiple Bonding" Here you need to understand the difference in the "overlap" for a pi bond vs a sigma bond. Do question # 2 and the Exercise 10.2 on page 10-3 of the workbook.

No e-mail assignment instead this unit quiz is 50 points an will be a brief review for the final.

Workbook problems: Chapter 10, 1 (15), 2 (37), 3 (39), 4 (41) and 5 (43 ) plus these questions,

a) What is the minimum number and maximum number of hybrid orbitals that can be formed by a carbon atom. Briefly explain.

b) What are the approximate angles between the "lobes" of electron density in "sp" , "sp^2" and "sp^3" hybrid orbital sets ?

c) If an atom is "sp" hybridized, how many unhybridized p atomic orbitals remain ? How many "pi" bonds could be formed by sp hybridized atoms ?

d) What is the hybridization of the methane molecules surrounding Joe Chemist ?

Orbital Discussion - Professor Jay Siegel at UCSD

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