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I. Introduction to Bonding (p. 161 – 163)

Ch. 6 & 7 - Chemical Bonding. I. Introduction to Bonding (p. 161 – 163). 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.

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I. Introduction to Bonding (p. 161 – 163)

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  1. Ch. 6 & 7 - Chemical Bonding I. Introduction toBonding(p. 161 – 163)

  2. 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

  3. 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

  4. A. Types of Bonds RETURN

  5. A. Types of Bonds RETURN

  6. A. Types of Bonds Ionic Bonding - Crystal Lattice RETURN

  7. A. Types of Bonds Covalent Bonding - True Molecules Diatomic Molecule RETURN

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

  9. B. Vocabulary • Chemical Bond • electrical attraction between nuclei and valence e- of neighboring atoms that binds the atoms together • bonds form in order to… • decrease PE • increase stability

  10. B. Vocabulary CHEMICAL FORMULA IONIC COVALENT Formula Unit Molecular Formula NaCl CO2

  11. B. Vocabulary COMPOUND more than 2 elements 2 elements Binary Compound Ternary Compound NaCl NaNO3

  12. B. Vocabulary ION 2 or more atoms 1 atom Monatomic Ion Polyatomic Ion Na+ NO3-

  13. C. Bond Polarity • Most bonds are a blend of ionic and covalent characteristics.

  14. C. Bond Polarity • Nonpolar Covalent Bond • e- are shared equally • symmetrical e- density • usually identical atoms

  15. - + C. Bond Polarity • Polar Covalent Bond • e- are shared unequally • asymmetrical e- density • results in partial charges (dipole)

  16. C. Bond Polarity • Nonpolar • Polar • Ionic View Bonding Animations.

  17. C. Bond Polarity • Electronegativity • Attraction an atom has for a shared pair of electrons. • higher e-neg atom  - • lower e-neg atom +

  18. C. Bond Polarity • Electronegativity Trend (p. 151) • Increases up and to the right.

  19. C. Bond Polarity • Difference in the elements’ e-negs determines bond type

  20. Ch. 6 & 7 - Chemical Bonding II. Molecular Compounds(p. 164 – 172, 211 – 213)

  21. A. Energy of Bond Formation • Potential Energy • based on position of an object • low PE = high stability

  22. no interaction increased attraction A. Energy of Bond Formation • Potential Energy Diagram attraction vs. repulsion

  23. increased repulsion balanced attraction & repulsion A. Energy of Bond Formation • Potential Energy Diagram attraction vs. repulsion

  24. Bond Energy Bond Length A. Energy of Bond Formation • Bond Energy • Energy required to break a bond

  25. A. Energy of Bond Formation • Bond Energy • Short bond = high bond energy

  26. X 2s 2p B. 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

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

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

  29. 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.

  30. 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

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

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

  33. C. Molecular Nomenclature • The Seven Diatomic Elements Br2 I2 N2 Cl2 H2 O2 F2 H N O F Cl Br I

  34. Ch. 6 & 7 - Chemical Bonding III. Ionic Compounds(p. 176 – 180, 203 – 211)

  35. A. Energy of Bond Formation • Lattice Energy • Energy released when one mole of an ionic crystalline compound is formed from gaseous ions

  36. B. Lewis Structures • Covalent – show sharing of e- • Ionic – show transfer of e-

  37. B. Lewis Structures • Covalent – show sharing of e- • Ionic – show transfer of e-

  38. C. Ionic Nomenclature Ionic Formulas • Write each ion, cation first. Don’t show charges in the final formula. • Overall charge must equal zero. • If charges cancel, just write symbols. • If not, use subscripts to balance charges. • Use parentheses to show more than one polyatomic ion. • Stock System - Roman numerals indicate the ion’s charge.

  39. C. Ionic Nomenclature Ionic Names • Write the names of both ions, cation first. • Change ending of monatomic ions to -ide. • Polyatomic ions have special names. • Stock System - Use Romannumerals to show the ion’s charge if more than one is possible. Overall charge must equal zero.

  40. C. Ionic Nomenclature • Consider the following: • Does it contain a polyatomic ion? • -ide, 2 elements  no • -ate, -ite, 3+ elements  yes • Does it contain a Roman numeral? • Check the table for metals not in Groups 1 or 2. • No prefixes!

  41. C. Ionic Nomenclature Common Ion Charges 1+ 0 2+ 3+ NA 3- 2- 1-

  42. C. Ionic Nomenclature • potassium chloride • magnesium nitrate • copper(II) chloride  KCl • K+ Cl- • Mg2+ NO3-  Mg(NO3)2  CuCl2 • Cu2+ Cl-

  43. C. Ionic Nomenclature • NaBr • Na2CO3 • FeCl3 • sodium bromide • sodium carbonate • iron(III) chloride

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