Valence Bond Theory
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This document provides an overview of chemical bonding concepts including the four main types of bonds (ionic, covalent, metallic, and polar covalent) and their properties. It also discusses electronegativity and bond polarity, Lewis structures, the octet rule, valence shell electron pair repulsion theory (VSEPR) and molecular geometry. Hybridization of orbitals is explained using examples of sp, sp2, and sp3 hybridization. The document concludes by noting some exceptions and limitations of the valence bond theory for more complex molecules and transition metals.
Valence Bond Theory
- 1. Chemical Bonds and Properties What we will learn: • 4 types of bonding and their properties • Electronegativity and polar bonds • Valence Bond Theory • VSEPR • Hybridization of Orbitals • Sigma- and pi-bonds • Oxidation Numbers • Formal Charge • Resonance
- 2. How to study effectively ?
- 4. O3 • What different kind of electrons are in this molecule ? • Is this molecule stable ? • Does this molecule have a charge ? • Is this molecule linear or bent ? • Is the bond strength higher, the same or lower than in O2 ?
- 5. H2O • Is this molecule linear or bent ? • How many different kinds of electrons are in this molecule ? • Why is this molecule a liquid at RT, but H2S is a gas ? • Why is this molecule more stable than Hydrogenperoxide ?
- 6. CO • Is this molecule stable ? • Is this molecule more or less reactive than CO2 ?
- 7. • Is this molecule stable ? • Is this molecule planar ? • Why is this molecule a weak acid ?
- 8. Part 1 What is a chemical bond ?
- 9. Concept Map
- 10. After this lesson, we should understand: • Valence electrons • Covalent and ionic bond • σ- and π-bonds, lone pairs • Lewis Structure • Electronegativity • Dipole moment • Oxidation number • Formal Charge • Octet Rule • Hybridization • Basic shapes of simple molecules
- 11. Ionic Bond Normally between a metal and a non-metal: They exchange electrons and become ions (charged atoms) which attract each other by electrostatic force. A pair of ions does not stay alone but form crystals
- 12. Covalent Bond Two non-metals share (valence) electrons: (Remark: Transition metals can form covalent bonds also !)
- 13. Polar Covalent Bond Two non-metals share electrons unevenly because of electronegativity difference. Electrons are closer to one atom than the other. This results on partially negative and positive charges on the atoms
- 14. Metallic Bond Metal atoms share all their valence electrons, which freely move between all atoms which form a network. Therefore all metals can conduct electricity and look shiny
- 15. 4 Types of Bond
- 16. Bond Polarity
- 17. Polar Bonds Uneven sharing of electrons due to differences in Electronegativity The “pull” an atom has for electrons
- 22. Polar Molecules Electrons are not equally shared in a bond, which can lead to a dipolmoment of the whole molecule
- 23. Polar Bonds and Geometry
- 25. Which of these molecules have the greatest dipole moment ?
- 26. Which bond type ? (exception: Transition metals !)
- 30. (1) Formal Charge Split all bonds in the middle => “real” charge on atoms (2) Octet Rule Count all bonding electrons for one atom => 8 is most stable (3) Oxidation Number Give all bonding electrons to the more electronegative atom
- 31. Special Cases “Extended octet” Especially P and S can use d-orbitals to make more than 3 resp. 2 bonds ! 6 VE: Especially common for B and Al !
- 32. Part 2: Valence Bond Theory (VB) “Valence Electrons are located in bonds and lone pairs”
- 35. Sigma bonds
- 36. Pi Bonds
- 37. The VB Theory works well for diatomic molecules. But for more complex molecules we need an extension !
- 38. VSEPR Valence Shell Electron Pair Repulsion Lewis 2D 3D structure
- 39. Most common geometrical structures
- 45. Draw 3D from Lewis Formula • CO2 vs. SO2 • O2 vs. O3 • H2O vs. H2O2 • SOCl2 vs. HCHO
- 46. “Resonance” Write the resonance formula for OZONE ! Does the molecule have a charge ?
- 48. Important exception: Carbon Monoxide !
- 50. Why another theory ?
- 51. Hybrid Orbitals All orbitals in an atom involved in sigma-bonds hybridize (mix) into orbitals of equal energy. Pi-Bonds are still formed by p-orbitals. Lone electron pairs count as “single bonds” and are part of the hybridization !
- 56. sp Hybridization
- 57. Which hybrids form and when ? + lone pairs
- 58. Tasks • Draw the Lewis Structures and the Hybrid Orbitals for Ethane, Ethene and Ethyne (mark the hybrid orbitals) • Which hybridization has the central atom in: H2O, O2, NH3, NH4+, N in pyridine, O in THF, S in SOCl2, C in HCHO
- 59. Inclusion of d-orbitals Elements in row 3 and up, hybridization can include also d-electrons. Typical example: SF6 with 6 sigma bonds
- 60. sp3d
- 61. Special Case: Transition Metal compounds Bonds in transition metal compounds are either ionic like in FeCl3 but can be covalent as well as in Fe(CO)5 => the VB Theory is not very suitable to explain the bonding for transition metals ! This can be done with the “Crystal Field Theory” (soon to come ….. )