Chapter 10 notes part 2 (1)

Ch 4 ch 2 cl 2 ccl 2 f 2 ccl 4 here are some general

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CH 4 CH 2 Cl 2
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CCl 2 F 2 CCl 4 Here are some general guidelines to follow when determing the polarity of a molecule: - In relative, qualitative terms: the more electronegative atom the more the bond will be polarized and therefore a bigger vector should be drawn (C-F vector will be bigger than the C-Cl vector). - You can think of the C-H bond as being relatively non-polar. - If a molecule has no lone pairs (E = 0) around the central atom (A) and all ligands (X) are the same, then the molecule will always be NON-POLAR. In all such cases the vectors will cancel and the net dipole moment in the molecule will be zero. - If the central atom has one or more lone pairs it will always be POLAR. ** Given any molecule from the geometries we’ve discussed, you should be able to determine if it is polar or non-polar. 10.3 Valence Bond Theory - a quantum mechanical explanation of the covalent bond (with out the math). Basics: 1. An orbital of one atom comes to occupy a portino of the same region of space as an orbital of another atom and therefore overlap. Look at the formation of H 2 2. The total # of electrons in both orbital is no greater than 2. Hybrid orbitals : orbitals for electrons in covalent bonds used to describe bonding obtained by combining atomic orbitals (s and p, s and d, s, p and d, etc).
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The electron configuration of C is 1s 2 2s 2 2p 2 and can be drawn on an energy diagram, showing the relative energies of each orbital.
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