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Introduction to Bonding and Lewis Dot Diagrams

Get your missing or late work in! This lesson covers the basics of chemical bonding, including covalent, ionic, and metallic bonds, as well as Lewis dot diagrams. Don't miss out on improving your understanding of this important topic. All work is due by 12/19/18.

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Introduction to Bonding and Lewis Dot Diagrams

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  1. OK, so you’re not Einstein, but even he figured it out! Get your missing or late work in! Some points are better than none!All Work due by 12/19/18

  2. Bellringer • The elements in group 2A are called? • The elements in group 7a are called?

  3. I. Introduction toBonding and Lewis Dot Diagrams Ch. 5 & 6 - Chemical Bonding

  4. A. Chemical Bonds • Definition: • An electrical attraction between nuclei and valence e- of neighboring atoms that binds the atoms together • bonds form in order to… • decrease potential energy • increase stability

  5. A. Types of Bonds COVALENT IONIC e- are transferred from metal to nonmetal e- are shared between two nonmetals Bond Formation Type of Structure true molecules crystal lattice Physical State liquid or gas solid Melting Point low high Solubility in Water yes usually not yes (solution or liquid) Electrical Conductivity no Other Properties Weak bonds Strong bonds

  6. A. Types of Bonds METALLIC e- are delocalized among metal atoms Bond Formation Type of Structure “electron sea” Physical State solid Melting Point very high Solubility in Water no yes (any form) Electrical Conductivity malleable, ductile, lustrous Other Properties

  7. A. Types of Bonds Na + Cl Na+ + Cl-

  8. A. Types of Bonds Ionic Bonding - Crystal Lattice

  9. A. Types of Bonds • Seven elements form Diatomic Molecules: Br2 I2 N2 Cl2 H2 O2 F2 H N O F Cl Br I

  10. A. Types of Bonds Metallic Bonding - “Electron Sea”

  11. B. 2 types of covalent bonds: 1. Nonpolar Covalent Bond • e- are shared equally • symmetrical e- density • usually identical atoms (O2, N2 etc.)

  12. - + B. 2 types of covalent bonds: 2. Polar Covalent Bond • e- are shared unequally (For example: H2O ) • asymmetrical e- density • results in partial charges (dipole)

  13. B. Bond Polarity • Nonpolar • Polar • Ionic

  14. B. Bond Polarity • Electronegativity • Attraction an atom has for a shared pair of electrons. • higher e-neg atom  - • lower e-neg atom + • “FONCl BrICS”

  15. B. Bond Polarity • Difference in the elements’ e-negs determines bond type: > 2.0 Ionic 0.4 – 2.0 Polar C < 0.4 Nonpolar C

  16. B. Bond Polarity • Example: • F2 :4.0 – 4.0 = 0 Nonpolar Covalent • H20: 3.44 – 2.10 = 1.34 Polar Cov. • NaCl: 3.16 – 0.93 = 2.23 Ionic

  17. X C. Lewis Dot Diagrams • Electron Dot Diagrams – model for showing ionic and covalent bonds • show valence e- as dots • distribute dots like arrows in an orbital diagram

  18. Gilbert Newton Lewis Invented “Electron-dot” formulas or “Lewis Structures” I’m so tired of writing all those useless inner electrons, in the Bohring models!

  19. Ne C. Lewis Structures • Octet Rule • Most atoms form bonds in order to obtain 8 valence e- • Full energy level stability ~ Noble Gases

  20. C. Lewis Structures • Ionic – Show the transfer of e-

  21. - + + C. Covalent – Electrons shared • Nonpolar Covalent - no charges • Polar Covalent - partial charges

  22. F F F S F F F F B F F H O H N O Very unstable!! C. Octet Rule • Exceptions: • Hydrogen  2 valence e- • Groups 1,2,3 get 2,4,6 valence e- • Expanded octet  more than 8 valence e- (e.g. S, P, Xe) • Radicals  odd # of valence e-

  23. D. Drawing Lewis Diagrams 1. Find total # of valence e-. 2. Arrange atoms - singular atom is usually in the middle. 3. Form bonds between atoms (2 e- shared is a single bond – one pair) 4. Distribute remaining e- to give each atom an octet (recall exceptions). 5. If there aren’t enough e- to go around or an element is not stable, form double(2 pairs of shared e-) or triple (3 pairs of shared e-) bonds.

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