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December 1, 2009 •T H E “ M A K E U P ” L E C T U R E •T O P I C S Molecular Polarity (8.7) Introduction to Bonding Theories (9.1) Valence Bond Theory (9.2) Molecular Polarity Polarity = uneven distribution of charge Bond polarity = Electrons drawn closer to the more electronegative atom Molecular polarity = Molecule as a whole has a net separation of charge A polar molecule must have polar bonds A molecule is polar if the directions of the polar bonds don’t cancel/offset eachother Nonpolar Examples: CH4, CO2 Polar Examples: H2O, NH3 Nonpolar Examples: CH4, CO2 Polar Examples: H2O, NH3 Why is Polarity Important? Polarity dictates many molecular properties Physical state (solid, liquid, gas) CO2 (44 g/mol) vs. H2O (18 g/mol) Solubility Chapter 9- Chemical Bonding Theories Valence Bond Theory: Uses Lewis Structures Bonds form using shared electrons between overlapping orbitals on adjacent atoms. Orbitals arrange around central atom to avoid each other. Two types of bonds: sigma () and pi (). Qualitative, visual- good for many atom systems in ground state Molecular Orbital Theory: Uses MO Diagrams Orbitals on atoms “mix” to make molecular orbitals, which go over 2 or more atoms. Two electrons can be in an orbital. Quantitative- needed to describe excited states Sigma () Bonding Orbitals on bonding atoms overlap directly between bonding atoms Sigma () Bonding Consider VSEPR Shapes and bonding: What’s wrong with this picture? Atoms bond by having their valence orbitals overlap Bonding orbitals are not the same shape as atomic orbitals 2pz 2px 2py 2s Orbitals in CH4 Electron configurations: H = 1s1 C = 1s22s22p2 Atomic orbitals change shape when they make molecules Hybrid Orbitals