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Bonding and Molecular Structure

Bonding and Molecular Structure

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Bonding and Molecular Structure

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  1. Bonding and Molecular Structure Chapter 10 Chapter 10

  2. Covalent Bonding and Orbital Overlap • Lewis structures and VSEPR do not explain how a bond forms. • VSEPR predicts the shape of a molecule, but it does not explain how the molecule is put together. • One method to explain bonding would be Valence Bond Theory: • Bonds form when atomic orbitals on atoms overlap. • Two electrons are shared by the orbital overlap. Chapter 10

  3. Covalent Bonding and Orbital Overlap Chapter 10

  4. Hybrid Orbitals • sp Hybrid Orbitals • Consider BeF2 • - Be has a 1s22s2 electron configuration. • There is no unpaired electron available for bonding. • We could promote and electron from the 2s orbital on Be to the 2p orbital to get two unpaired electrons for bonding. • The F-Be-F bond angle is 180 (VSEPR theory). • BUT the geometry is still not explained. Chapter 10

  5. Hybrid Orbitals • sp Hybrid Orbitals • We can solve the problem by allowing the 2s and one 2p orbital on Be to mix or form two hybrid orbitals (process called hybridization). • - The hybrid orbitals come from an s and a p orbital and is called sp hybrid orbital. Chapter 10

  6. Hybrid Orbitals • sp Hybrid Orbitals • The two sp hybrid orbitals are 180 apart. • Since only one of the Be 2p orbitals has been used in hybridization, there are two unhybridized p orbitals remaining on Be. Chapter 10

  7. Hybrid Orbitals Hybrid Orbitals - Other hybrids can be formed by mixing additional p and/or d orbitals Chapter 10

  8. Hybrid Orbitals Chapter 10

  9. Hybrid Orbitals Chapter 10

  10. Hybrid Orbitals • Summary • To assign hybridization: • draw a Lewis structure • assign the electron pair geometry using VSEPR theory • from the electron pair geometry, determine the hybridization Chapter 10

  11. Multiple Bonds • -Bonds - electron density lies on the axis between the nuclei. • -Bonds - electron density lies above and below the plane of the nuclei. • A double bond consists of one -bond and one -bond • A triple bond has one -bond and two -bonds • -bonds come from unhybridized p orbitals. Chapter 10

  12. Multiple Bonds Chapter 10

  13. Multiple Bonds Ethylene, C2H4 Chapter 10

  14. Multiple Bonds Acetylene, C2H2 Chapter 10

  15. Molecular Orbitals • Some aspects of bonding are not explained by Lewis structures, VSEPR theory and hybridization. • For these molecules, we use Molecular OrbitalTheory (MO theory). • Just as electrons in atoms are found in atomic orbitals, electrons in molecules are found in molecular orbitals. Chapter 10

  16. Molecular Orbitals • The Hydrogen Molecule • When two AOs overlap two MO’s form. • One molecular orbital has electron density between nuclei (, bonding MO); • One molecular orbital has little electron density between nuclei (*, antibonding MO). Chapter 10

  17. Molecular Orbitals The Hydrogen Molecule Chapter 10

  18. Molecular Orbitals The Hydrogen Molecule Chapter 10

  19. Molecular Orbitals The Hydrogen Molecule Chapter 10

  20. Molecular Orbitals • The Hydrogen Molecule • MO diagram shows the energies and orbitals in an MO diagram. • The total number of electrons in all atoms are placed in the MO’s starting from lowest energy (1s) and ending when you run out of electrons. • Note that electrons in MO’s have opposite spins. • H2 has two bonding electrons. • He2 has two bonding electrons and two antibonding electrons. Chapter 10

  21. Molecular Orbitals • Bond Order Define Bond Order = ½(bonding electrons - antibonding electrons). Bond order = 1 for single bond. Bond order = 2 for double bond. Bond order = 3 for triple bond. Bond order for H2 = ½(bonding electrons - antibonding electrons) = ½(2 - 0) = 1. Therefore, H2 has a single bond. • Bond order for He2 = ½(bonding electrons - antibonding electrons) = ½(2 - 2) = 0. • Therefore He2 is not a stable molecule. Chapter 10

  22. Second-Row Diatomic Molecules • AO’s combine according to the following rules • The number of MO’s equals the number of AO’s • AO’s of similar energy combine • As overlap increases, the energy of the MO decreases • Pauli: each MO has at most two electrons • Hund: for degenerate orbitals, each MO is first occupied singly. Chapter 10

  23. Second-Row Diatomic Molecules • Molecular Orbitals from 2p Atomic Orbitals • There are two ways in which two p orbitals overlap • end on so that the resulting MO has electron density on the axis between nuclei ( orbital) • sideways so that the resulting MO has electron density above and below the axis between nuclei ( orbital). • The p-orbitals must give rise to 6 MO’s: one  and two ’sone * and two *’s • The relative energies of these  and -orbitals can change. Chapter 10

  24. Second-Row Diatomic Molecules Molecular Orbitals from 2p Atomic Orbitals Chapter 10

  25. Second-Row Diatomic Molecules Electron Configurations for B2 through Ne2 Chapter 10

  26. Second-Row Diatomic Molecules • Electron Configurations for B2 through Ne2 • For B2, C2 and N2 the 2p orbital is higher in energy than the 2p. • For O2, F2 and Ne2 the 2p orbital is lower in energy than the 2p. Chapter 10

  27. Second-Row Diatomic Molecules Electron Configurations for B2 through Ne2 Chapter 10

  28. Second-Row Diatomic Molecules Electron Configurations for B2 through Ne2 Chapter 10

  29. Second-Row Diatomic Molecules • Electron Configurations and Molecular Properties Two types of magnetic behavior: • Paramagnetism, unpaired electrons in molecule • Diamagnetism, no unpaired electrons in molecule Chapter 10

  30. Second-Row Diatomic Molecules Electron Configurations and Molecular Properties Chapter 10

  31. Homework 10, 12, 24, 47 Chapter 10