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Chemical Bonding

Chemical Bonding. Chemical Bonds, Lewis Symbols, and the Octet Rule. Chemical bond : attractive force holding two or more atoms together. Covalent bond results from sharing electrons between the atoms. Usually found between nonmetals.

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Chemical Bonding

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  1. Chemical Bonding

  2. Chemical Bonds, Lewis Symbols, and the Octet Rule • Chemical bond: attractive force holding two or more atoms together. • Covalent bond results from sharing electrons between the atoms. Usually found between nonmetals. • Ionic bond results from the transfer of electrons from a metal to a nonmetal. • Metallic bond: attractive force holding pure metals together.

  3. Figure 8.3: Ionic Bonding

  4. Figure 8.5: Covalent Bonding

  5. Chemical Bonds

  6. Strengths of Covalent Bonds

  7. Chemical Bonds, Lewis Symbols, and the Octet Rule Lewis Symbols

  8. Chemical Bonds, Lewis Symbols, and the Octet Rule • The Octet Rule • All noble gases except He has an s2p6 configuration. • Octet rule: atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs). • Caution: there are many exceptions to the octet rule.

  9. Bond Polarity and Electronegativity • Electronegativity • Electronegativity: The ability of one atoms in a molecule to attract electrons to itself. • Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F). • Electronegativity increases • across a period and • down a group.

  10. Figure 8.6: Electronegativities of Elements Electronegativity

  11. Bond Polarity and Electronegativity • Figure 8.7: Electronegativity and Bond Polarity • There is no sharp distinction between bonding types. • The positive end (or pole) in a polar bond is represented + and the negative pole -. HyperChem

  12. Drawing Lewis Structures • Follow Step by Step Method (See Ng Web-site) • Total all valence electrons. [Consider Charge] • Write symbols for the atoms and guess skeleton structure [ define a central atom ]. • Place a pair of electrons in each bond. • Complete octets of surrounding atoms. [ H = 2 only ] • Place leftover electrons in pairs on the central atom. • If there are not enough electrons to give the central atom an octet, look for multiple bonds by transferring electrons until each atom has eight electrons around it. HyperChem CyberChem video

  13. Lewis Structures - Examples

  14. Exceptions to the Octet Rule • Less than an Octet • Relatively rare. • Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A. • Most typical example is BF3. • Formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds.

  15. Exceptions – Central Atoms - Less than an Octet

  16. Exceptions to the Octet Rule • More than an Octet • This is the largest class of exceptions. • Atoms from the 3rd period onwards can accommodate more than an octet. • Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density. HyperChem

  17. Exceptions – Central Atoms - Greater than an Octet

  18. Molecular Shapes • There are five fundamental geometries for molecular shape:

  19. Figure 9.3 HyperChem

  20. Molecular Shapes – 3D Notations

  21. Summary of VSEPR Molecular Shapes HyperChem CyberChem video See Ng Web-site

  22. Examples: VSEPR Molecular Shapes

  23. VSEPR Model • The Effect of Nonbonding Electrons • By experiment, the H-X-H bond angle decreases on moving from C to N to O: • Since electrons in a bond are attracted by two nuclei, they do not repel as much as lone pairs. • Therefore, the bond angle decreases as the number of lone pairs increase HyperChem

  24. VSEPR Model HyperChem • Figure 9.10: Shapes of Larger Molecules • In acetic acid, CH3COOH, there are three central atoms.

  25. Figure 8.10: Drawing Lewis Structures Resonance Structures

  26. HyperChem Figure 9.12

  27. Figure 9.11: Molecular Shape and Molecular Polarity HyperChem

  28. Figure 9.13: Molecular Shape and Molecular Polarity HyperChem

  29. Covalent Bonding and Orbital Overlap • Lewis structures and VSEPR do not explain why a bond forms. • How do we account for shape in terms of quantum mechanics? • What are the orbitals that are involved in bonding? • We use Valence Bond Theory: • Bonds form when orbitals on atoms overlap. • There are two electrons of opposite spin in the orbital overlap.

  30. Figure 9.14: Covalent Bonding and Orbital Overlap

  31. VSEPR Model (Figure 9.6) • To determine the electron pair geometry: • draw the Lewis structure, • count the total number of electron pairs around the central atom, • arrange the electron pairs in one of the above geometries to minimize e--e- repulsion, and count multiple bonds as one bonding pair.

  32. VSEPR Model

  33. Drawing Lewis Structures • Formal Charge • Consider: • For C: • There are 4 valence electrons (from periodic table). • In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. • Formal charge: 4 - 5 = -1.

  34. Drawing Lewis Structures • Formal Charge • Consider: • For N: • There are 5 valence electrons. • In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. • Formal charge = 5 - 5 = 0. • We write: CyberChem video

  35. Chemical Bonding Lewis VSEPR shapes AXE notation Polarity

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