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Published byCharles Reed Modified over 9 years ago
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Molecular Orbital Theory (What is it??) Better bonding model than valence bond theory Electrons are arranged in “molecular orbitals” Dealing with valence shell electrons
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Molecular OrbitalsMolecular Orbitals Combination of atomic orbitals 2 atomic orbitals------2 molecular orbitals (one bonding, one antibonding) Molecular region where electrons are likely to be found within a chemical compound Mathematically based Behave like molecules Deals with electrons arranged in molecules, NOT atoms
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Bonding Molecular Orbital ( σ ) Contain electrons involved in chemical bonding Contribute to bond strength Increase stability Lower energy level for electrons More electrons present, more stability within molecule Increased electron density between atoms
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Anti-bonding Molecular Orbitals ( σ *) Contain electrons NOT involved in bonding Electrons hang out away from bond Decreased stability and bond strength Higher energy level for electrons Instability within molecule Decreased electron density between atoms
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**Electrons want to be at a LOW energy level SO---generally pair up and reside in bonding molecular orbitals.
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How are electrons placed in molecular orbitals? 1)Electrons want to be in the lowest-energy molecular orbitals as possible. 2)Only 2 electrons found in each molecular orbital. 3)Electrons are placed in molecular orbitals by themselves (parallel spins) unless they have to be paired up (opposite spins).
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Bond OrderBond Order Determined by molecular orbitals = (# electrons in bonding MO) – (# electrons in anti-bonding MO) 2
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Ex. 1 Molecular orbital energy-level diagram H2H2
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Example 2: Molecular orbital energy-level diagram H2+H2+
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What happens when “p” atomic orbitals combine? Each “p” orbital combines with another “p” orbital—2 molecular orbitals produced Of the 2 molecular “p” orbitals— 1 lower energy bonding orbital 1 higher energy anti-bonding orbital One p orbital produces orbital overlap σ p, σ * p Other 2 p orbitals overlap in parallel π p, π* p
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Example 3: O 2
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What is polarity?What is polarity? Focus on covalent bonds Contributes to the properties of chemical compounds Based on electronegativity difference between atoms within chemical bonds
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Electronegativity and Bond Polarity Covalent Bond Polar Covalent Ionic Bond (0—0.5)(0.5—1.9) (1.9—3.5)
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Chose your polarityChose your polarity 1)Nonpolar Covalent small to same electronegativity difference Occurs among same or similar atoms bonded together EQUAL sharing of electrons, equal distribution of electron density 2)Polar Covalent difference in electronegativity Most electronegative atom pulls electrons towards it UNEQUAL sharing of electrons, electron density displaced toward electronegative atom
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Partial ChargesPartial Charges Exist in polar covalent molecules Demonstrates tendency of one end to be “slightly” negative or positive NO charge on the whole molecule
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Polar vs. Nonpolar Molecules Polar Exhibit difference in electronegativity—partial charges Based on polar covalent bonds Nonpolar No electronegativity difference No partial charges
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Dipoles Bond Dipoles Charge separation within the chemical bond Only applies to a specific chemical bond Molecular Dipole Charge separation among WHOLE molecule Includes ALL chemical bonds Equal to bond dipole in a diatomic molecule (ex. HCl, HI)
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Example 1: HClExample 1: HCl
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What about polyatomic molecules ?
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Distinguishing between Polar/Nonpolar Molecules 1)Draw Lewis Structure 2)Identify polar bonds 3)Determine molecular shape 4)Draw dipoles
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Example 2: CO 2 Polar bonds? Dipoles cancel?
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Example 3: H 2 OExample 3: H 2 O Polar bonds? Dipoles cancel?
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Example 4: BF 3 Polar bonds? Dipoles cancel?
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Example 5: CCl 4 Polar bonds? Dipoles cancel?
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Example 6: CH 3 ClExample 6: CH 3 Cl Polar bonds? Dipoles cancel?
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Example 7: NH 3 Polar bonds? Dipoles cancel?
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Homework pp. 427-429 #25, 29, 37-38, 45-46, 55, 63
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