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Chemical Bonding

Chemical Bonding. Introduction to Bonding. Vocabulary. Chemical Bond attractive force between atoms or ions that binds them together as a unit bonds form in order to… decrease potential energy (PE) increase stability. Vocabulary. CHEMICAL FORMULA. IONIC. COVALENT. Formula Unit.

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Chemical Bonding

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  1. Chemical Bonding Introduction toBonding

  2. Vocabulary • Chemical Bond • attractive force between atoms or ions that binds them together as a unit • bonds form in order to… • decrease potential energy (PE) • increase stability

  3. Vocabulary CHEMICAL FORMULA IONIC COVALENT Formula Unit Molecular Formula NaCl CO2

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

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

  6. 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 odorous

  7. 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

  8. Types of Bonds Ionic Bonding

  9. Types of Bonds Covalent Bonding

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

  11. Types of Bonds Ionic Bonding - Crystal Lattice RETURN

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

  13. Bond Polarity • Most bonds are a blend of ionic and covalent characteristics. • Difference in electronegativity determines bond type.

  14. Bond Polarity • Electronegativity • Attraction an atom has for electrons. • higher e-neg atom  - • lower e-neg atom +

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

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

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

  18. Bond Polarity Because chlorine is more electronegative than hydrogen, the electron density is greater around the chlorine atom.

  19. Bond Polarity • Nonpolar • Polar • Ionic

  20. Bond Polarity • Examples: • Cl2 • HCl • NaCl 3.0-3.0=0.0 Nonpolar 3.0-2.1=0.9 Polar 3.0-0.9=2.1 Ionic

  21. Bond Polarity In a polar covalent bond, the bonding electrons will spend a greater amount of time around the atom that has the stronger affinity for electrons. A good example of a polar covalent bond is the hydrogen-oxygen bond in the water molecule. Close

  22. Chemical Bonding II. Molecular Compounds

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

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

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

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

  27. 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.

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

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

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

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

  32. Chemical Bonding III. Ionic Compounds

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

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

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

  36. 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.

  37. 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 Roman numerals to show the ion’s charge if more than one is possible. Overall charge must equal zero.

  38. 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!

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

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

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

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