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Chapter 9.4-9.8 Bonding Theory

Chapter 9.4-9.8 Bonding Theory

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Chapter 9.4-9.8 Bonding Theory

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  1. Chapter 9.4-9.8Bonding Theory Complete Ch 9 problems # 29, 36, 38, 39, 42, 45, 49, 52, 55, 63, 75, 77, 85, 87 Valence Bond Theory Molecular Orbital Theory

  2. Let’s Focus on Valence Bond Theory As we saw earlier, a single bond formed between two atoms is formed with orbital overlap.

  3. Let’s Focus on Valence Bond Theory Polyatomic molecules form via hybrid orbitals. Carbon, for example, yields four sp3 hybrid orbitals, allowing a molecule like CH4 to have a tetrahedral shape.

  4. Hybrid Orbitals As we have discussed before, once you know the electron-domain geometry, you know the hybridization state of the atom.

  5. Valence Bond Theory designates two bond types arising from orbital overlap and hybrid orbitals. Sigma Bonds (σ) And Pi Bonds (π)

  6. Sigma () Bonds • Sigma bonds are characterized by • Head-to-head overlap. • Cylindrical symmetry of electron density about the internuclear axis. In other words, electron density concentrated between the two bonded nuclei.

  7. Pi () Bonds • Pi bonds are characterized by • Side-to-side overlap. • Electron density above and below the internuclear axis.

  8. Single Bonds Single bonds are always  bonds, because  overlap is greater, resulting in a stronger bond and more energy lowering.

  9. Multiple Bonds In a multiple bond one of the bonds is a  bond and the rest are  bonds.

  10. Multiple Bonds • In a molecule like formaldehyde (shown at left) an sp2 orbital on carbon overlaps in  fashion with the corresponding orbital on the oxygen. • The unhybridized p orbitals overlap in  fashion.

  11. Multiple Bonds In triple bonds, as in acetylene, two sp orbitals form a  bond between the carbons, and two pairs of p orbitals overlap in  fashion to form the two  bonds.

  12. Delocalized Electrons: Resonance When writing Lewis structures for species like the nitrate ion, we draw resonance structures to more accurately reflect the structure of the molecule or ion.

  13. Delocalized Electrons: Resonance • In reality, each of the four atoms in the nitrate ion has a p orbital. • The p orbitals on all three oxygens overlap with the p orbital on the central nitrogen.

  14. Delocalized Electrons: Resonance This means the  electrons are not localized between the nitrogen and one of the oxygens, but rather are delocalized throughout the ion and thus accounting for the equal bond lengths observed throughout the structure.

  15. Resonance The organic molecule benzene has six  bonds and a p orbital on each carbon atom.

  16. Resonance • In reality the  electrons in benzene are not localized, but delocalized. • The even distribution of the  electrons in benzene makes the molecule unusually stable.

  17. Example • Consider acetonitrile: • Predict the bond angles around each carbon atom • Give the hybridization at each of the carbon atoms • Determine the number of σ and π bonds in the molecule

  18. Example Which of the following molecules or ions will exhibit delocalized bonding? SO3, SO32-, H2CO, O3, NH4+

  19. Valence Bond Theory in a Nut Shell… • All covalent bonds are formed by the sharing of two or more electrons. All bonds will have at least one of these bond pair concentrated in the space between the two interacting nuclei forming a sigma bond. • The appropriate set of hybrid orbitals used to form these sigma bonds is determined by the geometry of the molecule. • Atoms that share more than one pair of electrons will do so in pi bonds, which lie above and below the nuclei. • Molecules with two or more resonance structures can have pi bonds that extend beyond the two bonded atoms, allowing electrons to become delocalized.

  20. Molecular Orbital (MO) Theory Though valence bond theory effectively conveys most observed properties of ions and molecules, there are some concepts better represented by molecular orbitals.

  21. 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.

  22. Molecular Orbital (MO) Theory If waves interact destructively, the resulting orbital is higher in energy: an antibonding molecular orbital.

  23. MO Theory • MO Theory begins with the construction of molecular orbitals. • Whenever two atomic orbitals overlap, two molecular orbitals form. • Let’s consider H2. Notice, two orbitals will overlap, and thus, two molecular orbitals formed.

  24. Let’s turn this into a diagram…

  25. 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 antibonding electrons. • Recall, bond order indicates the stability of the bond formed.

  26. 1 2 (2 - 0) = 1 MO Theory For hydrogen, with two electrons in the bonding MO and none in the antibonding MO, the bond order is

  27. 1 2 (2 - 2) = 0 MO Theory • In the case of He2, the bond order would be • Therefore, He2 does not exist.

  28. Example Write the electron configuration in terms of MOs and determine the bond order of H2- and H2+.

  29. Example According to molecular orbital theory, would either Be2 or Be2+ be expected to exist? Explain.

  30. MO Theory • For atoms with both s and p orbitals, there are two types of interactions: • The s and the p orbitals that face each other overlap in  fashion. • The other two sets of p orbitals overlap in  fashion.

  31. MO Theory • The resulting MO diagram looks like this. • There are both s and p bonding molecular orbitals and s* and * antibonding molecular orbitals.

  32. MO Theory • 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 s and p molecular orbitals in these elements.

  33. Second-Row MO Diagrams

  34. Paramagnetism vs. Diamagnetism Molecular orbital theory allows one to predict the behavior of a substance in the presence of a magnetic field. Molecules with one or more unpaired electrons are attracted into a magnetic field. This behavior is called paramagnetism. Substances with no unpaired electrons are weakly repelled from a magnetic field…they exhibit diamagnetism.

  35. Example Predict the magnetic properties and bond order of the peroxide ion, O22- and the acetylide ion, C22-.

  36. Molecular Orbital Theory • This model represents electrons in allowed energy states called molecular orbitals which can spread out across all the atoms of a molecule. • MO diagrams represent the combination and relative energies of molecular orbitals. • The diagrams allow for easy calculation of bond order, indicating the stability of the bond formed. • Bonding and antibonding MOs formed by the combination of s orbitals are called sigma molecular orbitals while the combination of p orbitals are called pi molecular orbitals. • The MO model correctly predicts properties like paramagnetism and diamagnetism.

  37. SAMPLE INTEGRATIVE EXERCISE Putting Concepts Together Elemental sulfur is a yellow solid that consists of S8 molecules. The structure of the S8 molecule is a puckered eight-membered ring (Figure 7.30). Heating elemental sulfur to high temperatures produces gaseous S2 molecules: (a) With respect to electronic structure, which element in the second row of the periodic table is most similar to sulfur? (b) Use the VSEPR model to predict the S—S—S bond angles in S8 and the hybridization at S in S8. (c) Use MO theory to predict the sulfur–sulfur bond order in S2. Is the molecule expected to be diamagnetic or paramagnetic? (d) Use average bond enthalpies (Table 8.4) to estimate the enthalpy change for the reaction just described. Is the reaction exothermic or endothermic?