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Chapter 9 Molecular Geometries and Bonding Theories

Chapter 9 Molecular Geometries and Bonding Theories. Multiple Bonds. Formation of two π bonds in acetylene. Fig 9.26. Describing σ and π bonds in a molecule. σ. σ. formaldehyde. σ. π. Fig 9.27 Formation of σ and π bonds in formaldehyde, CH 2 O.

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Chapter 9 Molecular Geometries and Bonding Theories

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  1. Chapter 9Molecular Geometriesand Bonding Theories

  2. Multiple Bonds

  3. Formation of two πbonds in acetylene Fig 9.26

  4. Describing σ and π bonds in a molecule σ σ formaldehyde σ π Fig 9.27 Formation of σ and π bonds in formaldehyde, CH2O

  5. How many s and p bonds are in the acetic acid (vinegar) molecule CH3COOH? H H C H C O O H Sigma (s) and Pi Bonds (p) 1 sigma bond Single bond 1 sigma bond and 1 pi bond Double bond Triple bond 1 sigma bond and 2 pi bonds s bonds = 6 + 1 = 7 p bonds = 1

  6. Molecular Orbital (MO) Theory In MO theory, we invoke the wave nature of electrons • If waves interact constructively, the resulting orbital is lower in energy: a bonding molecular orbital. • If waves interact destructively, the resulting orbital is higher in energy: an antibonding molecular orbital.

  7. MO Theory • In H2 the two electrons go into the bonding molecular orbital. • The bond order is one half the difference between the number of bonding and antibondingelectrons: Bond order = ½ (no. of bonding e− – no. of antibonding e−) Here: ½ (2-0) = 1

  8. MO Theory Fig 9.35 • In the case of He2, the bond order would be: Here: ½ (2-2) = 0 • Therefore, MO theory predicts that He2 does not exist, which we know to be true.

  9. MO Theory He2+ • In the case of He2+, the bond order would be: ½ (2-1) = 1/2 • Therefore, MO theory predicts that He2+ does exist and it will be relatively stable

  10. MO Theory – Second-Row Diatomics • Consider only homonuclear diatomic molecules • Number of MOs = number of AOs combined • AOs combine most effectively with other AOs of similar energy • The greater the overlap of AOs, the lower the energy of MO • Each MO can hold max of 2 electrons (Pauli exclusion) • Hund’s rule applies (same spin in degenerate orbitals)

  11. MOs for Li2 and Be2 Fig 9.37 Energy-level diagram for the Li2 molecule

  12. MOs from 2p Atomic Orbitals Fig 9.38 • For atoms with both s and porbitals, there are two types of interactions: • The porbitals that are head to head overlap in  fashion. • The other two sets of porbitals overlap in  fashion.

  13. MO Theory – Second-Row Diatomics Fig 9.43 • There are both s and p bonding molecular orbitals and s* and * antibonding molecular orbitals • Diagram fits only O2 and F2

  14. MO Theory Fig 9.45 Fig 9.44 • The smaller p-block elements in the second period have a sizeable interaction between the s and p orbitals: • This flips the order of the  and  molecular orbitals in O2 and F2

  15. Second-Row MO Diagrams

  16. Fig 9.48 Paramagentism of O2

  17. Fig 9.48 Paramagentism of O2 Figure 09.48

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