1 / 38

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:

thea
Télécharger la présentation

Chemical Bonding

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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

More Related