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CHEM 1A: VSEPR Theory Now that we have an understanding of covalent bonding and how atoms share electrons to form molecules and polyatomic ions, we will use Lewis dot structures to predict electronic and molecular geometries. In turn, the geometries of the molecules will influence many different physical and chemical properties like melting point, boiling point, surface tension, viscosity, elasticity, and flammability. Background information on VSEPR theory can be found on the course website and at this link: http://c4.cabrillo.edu/chem1a/ex1/ For each molecule below, fill in the table to determine the electron geometry & molecular geometry about the CENTRAL ATOM. Then sketch a perspective drawing of the substance and LABEL ALL BOND ANGLES. 1. HCN 6. NO3 - 11. NH4 + 16. IF6 + 2. H2S 7. NF3 12. SbCl5 17. IF5 3. CO2 8. CO3 2- 13. SbCl4 - 18. IF4 - 4. SO2 9. SO3 2- 14. ClF3 19. IOF5 5. NO2 - 10. H2CO 15. BrCl2 - 20. SOF4 Practice Visualizing Shapes: Some of the shapes we study are less intuitive than others. Our goal is to be able to visualize these shapes in our mind. For practice with the structures above, use the molecular modeling simulation found as a link on the course website for today’s lab or visit this URL to run a helpful simulation: http://phet.colorado.edu/en/simulation/molecule-shapes. For each structure that you draw, confirm the geometry using the simulation. Just don’t let the computer do the thinking for you—the simulation should be a tool for visualizing shapes, not a crutch for completing the table. As you work through these practice problems, keep in mind the following points: The molecular modeling simulation will not give exact bond angles. You must consider when deviations from ideal geometries are expected. These deviations should be reflected in the approximate bond angles you include in the perspective drawings. When predicting bond angles, try to visualize electron density around the central atom. The three factors below will subtly affect the bond angles about an atom: o Lone pairs: Since these nonbonding electrons are not confined between nuclei, lone pairs tend to occupy more space than bonding electrons. This effect produces greater electron repulsions for lone pairs and influences bond angles. o Multiple Bonds: Double or triple bonds have more electrons, which translate to greater electron- electron repulsions, and larger bond angles where these bonds are involved. o Electronegative Elements: Atoms with a strong pull on electrons change the distribution of negative charge in a bond, which affects the amount of electron-electron repulsions and influences bond angles. See the background link above for more information.
