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Chemical Bonding

Chemical Bonding. Chapter 6. Types of Chemical Bonds. Chemical Bond: mutual electrical attraction b/ the nuclei and valence e - of different atoms Atoms make bonds b/c they become more stable. Types of Chemical Bonds. Ionic Bonding:

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Chemical Bonding

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  1. Chemical Bonding Chapter 6

  2. Types of Chemical Bonds • Chemical Bond: mutual electrical attraction b/ the nuclei and valence e- of different atoms • Atoms make bonds b/c they become more stable

  3. Types of Chemical Bonds • Ionic Bonding: • Results from the electrical attraction between cations and anions • Atoms completely give up e-s to other atoms

  4. Types of Chemical Bonds • Covalent Bonding: • Results from sharing of e- pairs between two atoms • Shared e- are “owned” equally by two atoms

  5. Determining the Type of Bond • Bonds fall somewhere b/ purely Ionic and purely Covalent • Depends on electronegativity – how much is the atom pulling the e-s • Calculate the difference in electronegativities

  6. Determining the Type of Bond • Types: • Ionic: electronegativity dif: 1.7 – 3.2 • Polar-Covalent: dif: 0.3 – 1.7 (1.7 – ionic) • Nonpolar-Covalent: dif: 0 – 0.3 (0.3 – NP)

  7. Determining the Type of Bond • Ionic: • One atom is so electronegative it strips the other atom of electrons making cation and anion • NaCl

  8. Determining the Type of Bond • Polar-Covalent: • Bond where atoms have unequal attraction for shared e-s • More electronegative atom has stronger attraction for e-s

  9. Determining the Type of Bond • Nonpolar-Covalent: • Bond where atoms have equal sharing of e-s, balanced distribution of charge • Bonds b/ two atoms of the same element

  10. Determining the Type of Bond • Determine the type of bond b/ these elements and sulfur, which is more electronegative element? • H, Cs, Cl

  11. Covalent Bonding

  12. Covalent Bonding • Molecule: neutral group of atoms that are held together by covalent bonds. • Molecular compound – covalent compound

  13. Covalent Bonding

  14. Covalent Bonding • Chemical formula: NaCl, MgCl2, H2O (Ionic or Covalent) • Molecular formula: only for molecules (covalent)

  15. Covalent Bonding

  16. Covalent Bond • Bond length: average distance b/ two bonded atoms • Forming bonds – atoms release energy • Same amount of energy needed to break bond  bond energy (kJ/mol)

  17. Forming Covalent Bonds • Share e-s to get noble gas configuration

  18. Octet Rule • Def: atoms gain, lose or share e-s to have octet (8) of electrons in outer energy level • H – exception, only needs 2 e-s • Ex: HF

  19. Exceptions to Octet • B – happy with 6 e-s in outer level • Other elements can have more than 8 valence e-s – PF5, SF6 – d orbitals invovled in bonding

  20. Electron-Dot Notation • Use element symbol and dots to indicate valence e-s. • Period 2:

  21. Lewis Structures • Def: Use electron-dot notation for molecules • H2 • F2 – shared pair with dash (-) • lone pair – unshared pair

  22. Lewis Structures • Structural formula: shows “structure” w/o lone pairs • H – F , H – H • Single bond: covalent bond where one pair of e- being shared b/ two atoms

  23. Lewis Structures • Draw Lewis Structure: • NH3 • H2S

  24. Multiple Covalent Bonds • C, N, and O can share more than 1 pair of e-s • Double bond: two pairs, 4 e-s, being shared • C2H4

  25. Multiple Covalent Bonds • Triple Bond: three pairs, 6 e-s, being shared • N2 • Lengths of multiple bonds? • More bonds – shorter – more energy to break • p.187

  26. Resonance Structures • Molecules can’t be shown with one Lewis structure • Ex: O3

  27. Lewis Structure • CH2O • CH3Br • C2HCl • SO3

  28. Recap Ch. 6 • Bonds: Ionic and Covalent • Ionic, Polar-Covalent, and Nonpolar-Covalent • Drawing Lewis Structures • C2H2 • Type of bonds?

  29. Ionic Bonding

  30. Ionic Compounds • Def: (+) and (-) ions that combine so charges balance out • Crystalline solids • Formula Unit: simplest unit of ionic compound where charges are balanced • NaCl: Na+ Cl- • Video (68)

  31. Forming Ionic Compounds • NaCl – use electron dot diagrams • Compound with Ca and F:

  32. Characteristics of Ionic Bonds • Ions in crystal lattice are more stable – lower potential energy • Lattice energy: energy released when 1 mole of gaseous ions form a lattice • More negative energy = more stable = stronger bonds

  33. Ionic vs. Covalent • Ionic • stronger bonds b/ formula units than b/ molecules in covalent compounds • HIGHER melting and boiling points • hard but brittle • conduct electricity in molten or dissolved state

  34. Ionic vs. Covalent • Covalent • Weak bonds b/ molecules • Most compounds are gases at room temp. • LOW boiling and melting points

  35. Polyatomic Ions • Def: charged group of covalently bonded atoms • Result from excess or lack of electrons in bonding

  36. Metallic Bonding • Excellent conductors in solid state – due to highly mobile valence electrons • Filled outermost sublevel is s • p orbitals empty and some d • Vacant orbitals overlap and valence e-s are free to move throughout entire metal

  37. Metallic Bonding • Valence e-s are delocalized, do not belong to any one atom but move freely • Sea of electrons around metal atoms

  38. Metallic Properties • High electrical and thermal conductivity • Malleability – hammered into sheets • Ductility – drawn into wires

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