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I. The Covalent Bond (p. 240 – 247)

Ch. 8 - Covalent Bonding. I. The Covalent Bond (p. 240 – 247). increased repulsion. balanced attraction & repulsion. A. Why Do Atoms Bond?. Gain stability Lower potential energy Atoms that fulfill the octet rule are more stable. attraction vs. repulsion. H 2 O.

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I. The Covalent Bond (p. 240 – 247)

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  1. Ch. 8 - Covalent Bonding I. The Covalent Bond(p. 240 – 247)

  2. increased repulsion balanced attraction & repulsion A. Why Do Atoms Bond? • Gain stability • Lower potential energy • Atoms that fulfill the octet rule are more stable attraction vs. repulsion

  3. H2O B. What is a covalent bond? • A chemical bond that results from the sharing of electrons • Molecule = two or more atoms that are held together by covalent bonds • Majority of covalent bonds form between nonmetals (CLOSE together on periodic table)

  4. Examples: • Which of the following are covalent compounds? • NaBr • SiO2 • CO2 • AlCl3 • CH4

  5. Properties Table 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 Form electrolytes in solution odorous

  6. Cl2 C. Covalent Bonding Formation • Diatomic molecule = molecule containing only two atoms • Some atoms exist this way because they are more stable than the individual atoms

  7. D. Diatomic Elements • The Seven Diatomic Elements Br2 I2 N2 Cl2 H2 O2 F2 H N O F Cl Br I

  8. X 2s 2p E. Lewis Structures • Electron Dot Diagrams • show valence e- as dots • distribute dots like arrows in an orbital diagram • 4 sides = 1 s-orbital, 3 p-orbitals • EX: oxygen O

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

  10. E. Lewis Structures • Example Electron Dot Notations: • Ca • P • H • C Ca P

  11. E. Lewis Structures • Single Bonds • When atoms share one pair of electrons • The two shared electrons belong to both atoms simultaneously • Lewis Structure dots or a line symbolize a single covalent bond (1 pair of shared e-) H – H

  12. E. Lewis Structures • How would you draw the Lewis structure for fluorine? • Diatomic! • F – F

  13. E. Lewis Structures • Multiple Covalent Bonds • When atoms share more than one pair of electrons • Have higher bond energies and are shorter than single bonds

  14. E. Lewis Structures • Double Bonds • Sharing two pairs of electrons between two atoms • Draw Lewis Structures for 2 oxygen atoms

  15. E. Lewis Structures • Triple Bonds • Sharing three pairs of electrons between two atoms • Draw Lewis Structures for 2 nitrogen atoms

  16. Section 8.2 Naming Molecules

  17. C. Molecular Nomenclature • Prefix System (binary compounds) 1. Less e-neg atom comes first. 2. Add prefixes to indicate # of atoms. Omit mono- prefix on first element. 3. Change the ending of the second element to -ide.

  18. PREFIX mono- di- tri- tetra- penta- hexa- hepta- octa- nona- deca- NUMBER 1 2 3 4 5 6 7 8 9 10 C. Molecular Nomenclature

  19. C. Molecular Nomenclature • CCl4 • N2O • SF6 • carbon tetrachloride • dinitrogen monoxide • sulfur hexafluoride

  20. C. Molecular Nomenclature • arsenic trichloride • dinitrogen pentoxide • tetraphosphorus decoxide • AsCl3 • N2O5 • P4O10

  21. - + + B. Lewis Structures • Nonpolar Covalent - no charges • Polar Covalent - partial charges

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