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Honors Chemistry

Honors Chemistry. Chapter 9: Chemical Bonding I. 9.1 Lewis Dot Diagrams. Symbol surrounded by dots for valence e- Separate dots as much as possible. 9.2 Ionic Bonding. Electrostatic force holding ions together Ions formed by electron transfer Low E ion loses e - , high e.a. gains e -

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Honors Chemistry

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  1. Honors Chemistry Chapter 9: Chemical Bonding I

  2. 9.1 Lewis Dot Diagrams • Symbol surrounded by dots for valence e- • Separate dots as much as possible

  3. 9.2 Ionic Bonding • Electrostatic force holding ions together • Ions formed by electron transfer • Low Eion loses e-, high e.a. gains e- • Na + Cl  Na+ Cl- • Mg + S  Mg2+ S2- • Li + S  2 Li+ S2-

  4. 9.3 Lattice Energy • Energy that holds ionic compounds together in a crystal lattice • Transfer of e- requires energy (Eion) and releases energy (e.a.) • In general, the cation requires more energy than the anion releases, which makes bond formation unstable • Lattice energy releases additional energy, making bond formation stable

  5. 9.4 The Covalent Bond • Covalent bond = shared pair of electrons • F + F  F – F • Shared pair – shared electrons, bond • Lone pair – electrons not involved in bond

  6. 9.4 Lewis Structures • Representation of covalent compounds using dots for e- and lines for bonds • Octet rule • atoms bond in such a way as to gain 8 e- in valence shell • Exceptions – H and He • Multiple Bonds • Double bond – share 2 pairs; eg, O2 • Triple bond – share 3 pairs; eg, N2

  7. 9.4 Bonding Summary

  8. 9.4 Ionic / Covalent Properties • Intermolecular attractive forces • Ionic – strong, covalent – weak • Consider phase, density, solubility, conductivity • Ionic Covalent • Solid Liquid or gas • High density Low density • Usually soluble Often insoluble • Good conductor Poor conductor

  9. 9.5 Electronegativity • Element’s relative attraction for shared e-

  10. 9.5 Pauling Electronegativities

  11. 9.5 Electronegativity and Atomic Radius

  12. 9.5 Bond Character • Degree of sharing of the bonded e- • Depends on the difference in electronegativity • Small electronegativity difference • Equal sharing of bonded e- • True covalent bond (nonpolar covalent) • Moderate electronegativity difference • Unequal sharing of bonded e- • Polar covalent bond • Large electronegativity difference • Ionic bond

  13. 3.0 DEN 2.0 0.0 9.5 Bond Character

  14. 9.6 Writing Lewis Structures • Draw a reasonable skeletal structure for the compound • Count the total valence electrons available • Draw single bonds between all atoms and use remaining electrons to fulfill the octet rule • If there are not enough electrons to fulfill the octet rule, form double or triple bonds • Draw Lewis structures for NH3, O3, CO32-

  15. 9.7 Formal Charges • Difference between the number of e- an atom has in a Lewis structure and the number of e- it has as a free element • Assigning formal charges • Atom counts all its nonbonding e-’s • Atom counts 1 e- from each bond • Total formal charge must add up to the total charge of the molecule / ion

  16. 9.7 Formal Charges • The most plausible structures have: • The fewest formal charges • Formal charges of smallest magnitude • Negative formal charges on the most electronegative elements • No adjacent charges of the same type

  17. 9.8 Resonance • Some molecules have more than one plausible Lewis structure • Resonance –use of both structures to represent a molecule • Reality is that bonds are delocalized

  18. 9.8 Resonance • Draw resonance structures for • N2O • HSCN • NO3- • CO32- • Extreme resonance – C6H6

  19. 9.9 Exceptions to the Octet Rule • Incomplete Octets • Not enough electrons to make an octet • Usually occurs with Groups IIA and IIIA • Examples: BeH2, BF3 • Consider resonance in BF3 • BF3 + NH3 F3BNH3 • Coordinate covalent bond • One atom donates both shared electrons

  20. 9.9 Exceptions to the Octet Rule • Odd Electron Molecules • Often called radicals • Examples: NO, NO2 • Single e- goes on element with lower EN • Very reactive • Tend to form dimers

  21. 9.9 Exceptions to the Octet Rule • Expanded Octets • More than 8 e- on central atom • Requires the atom to have a d orbital • Can’t happen with periods 1 and 2 • Examples: SF6, XeF4, ClO4-

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