Chapter 10 Outline

10.1 Artificial Sweeteners: Fooled by Molecular Shape

10.2 VSEPR Theory: The Five Basic Shapes

VSEPR (Valence Shell Electron Pair Repulsion) is a simple tool to determine the geometry (shape) of a molecule. Understanding the shape of a molecule is a key part of understanding how it will behave. You need to know the names and angles for the structures discussed in this section.

10.3 VSEPR Theory: The Effect of Lone Pairs

10.4 VSEPR Theory: Predicting Molecular Geometries

10.5 Molecular Shape and Polarity

Bond Polarity combines molecular shape and dipole moments. The idea is that dipoles are vectors that add together. Depending upon the geometry they can cancel (as in CO2) or not cancle (as in H2O). You should be able to look at a simple three dimensional structre and determine if the dipoles will cancel.

10.6 Valence Bond Theory: Orbital Overlap as a Chemical

At this point the models that we have used to describe bonding have been very simple. This chapter introduces two new models to describe bonding. These models are more complex but they are able to explain and predict more of the molecular structure and chemical behavior of a compound. It is important to keep in mind that these are models, not the truth. They are tools that let you predict chemistry. They are powerful tools, but they are based on assumptions. It is important to know what these assumptions are so that you do not apply the tool to a problem that it can not solve. For now, I will be careful to clearly state what model I would like you to apply to a system.

10.7 Valence Bond Theory: Hybridization of Atomic Orbitals

The covalent bond was previously introduced as a sharing of electrons between two atoms. Since electrons are in orbitals, a bond is formed by the overlap of these orbitals. The problem with this simple model is that the atomic orbitals that we have discussed do not have the correct orientation for the molecular geometry. Since it is possible to find this geometry from experimental evidence, the theory does not fit. This section adds to the theory by introducing the concept of hybridization. The basic premise is to rearange the atomic orbitals into new hybrid orbitals that will accomidate the molecular geometry. You will need to know the names and geometries for the sp, sp2, and sp3 hybrid orbitals.

10.8 Molecular Orbital Theory: Electron Delocalization

Eventually valence bond theory also falls short. It is unable to explain a number of experimental observations. The classic example is the paramagnetism of O2. Molecular orbital theory is a more complex theory of bonding that is able to explain many of the experiemnts where valence bond theory fails. The basic premise of molecular orbital theory is that when atomic orbitals from atoms overlap to make a bond, they make a new molecular orbital. Since orbitals must be conserved, the overlap of two atomic orbitals makes two new molecular orbitals. This is the step that was overlooked with valence bond theory. These molecular orbitals have names (sigma and pi; bonding, non-bonding, and antibonding) and energy levels. Electrons fill in these orbitals in order of increasing energy. You need to be able to use an MO energy diagram to determine the electronic structure of a simple homonuclear diatomic molecule. Then be able to use this energy diagram to predict the bond order, geometry, and paramagnetic behavior of the molecule.
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