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: 0.4 – 1.6 Nonpolar-Covalent: dif: 0 – 0.3

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 HCl

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, F 2 C-H

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.  Molecule: neutral group of atoms that are held together by covalent bonds.  Molecular = “covalent”

13. Covalent Bonding

14.  Chemical formula: Ionic or Covalent NaCl, MgCl 2, H 2 O  Molecular formula: only for molecules (covalent)

15. Covalent Bonding

16. Forming Covalent Bonds  Share e - s to get noble gas configuration

17. 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: F 2

18. Exceptions to Octet  B – happy with 6 e - s in outer level  Other elements can have more than 8 valence e - s – PF 5, SF 6 – d orbitals invovled in bonding

19. Electron-Dot Notation  Use element symbol and dots to indicate valence e - s.  Period 2: 1A2A3A4A5A6A7A8A LiBeB CNOFNe

20. Lewis Structures  Def: Use electron-dot notation for molecules  H 2  F 2 – shared pair with dash (-)  lone pair – unshared pair

21. Lewis Structures  Single bond: covalent bond where one pair of e - being shared b/ two atoms  H-H and HCl :

22. Lewis Structures  Draw Lewis Structure: 1. Start with e-dot diagram of each 2. Atom with most bonding sites in middle 3. Circle unpaired e - s to make bonds 4. Replace circles with dashes NH 3 H 2 S

23. 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  C 2 H 4  Triple Bond: three pairs, 6 e - s, being shared  N 2

24. Bond Lengths  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)  Lengths of multiple bonds? More bonds – shorter – more energy to break p.187

25. Resonance Structures  Molecules can’t be shown with one Lewis structure  Ex: O 3

26. Lewis Structure  CH 2 O  CH 3 Br  C 2 HCl  SO 3

27. Recap Ch. 6  Bonds: Ionic and Covalent  Ionic, Polar-Covalent, and Nonpolar-Covalent  Drawing Lewis Structures C 2 H 2 Type of bonds?

28. Ionic Bonding

29. 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)

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

31. 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 energy released = lower potential energy = more stable = stronger bonds

32. 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

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

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

35. Metallic Bonding  Excellent conductors in solid state – due to highly mobile valence electrons  Filled outermost sublevel is s Vacant p and d orbitals overlap - valence e - s are delocalized, do not belong to any one atom but move freely

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

37. Molecular Geometry  3-D arrangement of molecules

38. VSEPR Theory  Valence-shell, electron-pair repulsion  Def: repulsion b/ valence e - pairs around atom causes them to be as far apart as possible

39. Shapes  NO lone pairs on CENTRAL atom  Symmetrical  Linear  Trigonal-Planar  Tetrahedral  Trigonal- bipyramidal  Octahedral  WITH lone pairs on CENTRAL atom  Non- symmetrical  Trigonal-pyramidal  Bent (angular)

40. Shapes – NO lone pairs on central atom 1. Linear (AB 2 ): A – central atom B-bonded atoms - 3 atom molecules CO 2 - 2 atom molecules, O 2, HCl, etc. - bond angles: 180 o

41. Shapes – NO lone pairs on central atom 2. Trigonal Planar (AB 3 ): - BCl 3 - bond angles: 120 o

42. Shapes – NO lone pairs on central atom 3. Tetrahedral (AB 4 ): - CCl 4 - bond angles: 109.5 o

43. Shapes – NO lone pairs on central atom 4. Trigonal-bipyramidal (AB 5 ): - PCl 5 - bond angles: 120 o and 90 o

44. Shapes – NO lone pairs on central atom 5. Octahedral (AB 6 ): - SF 6 - bond angles: 90 o

45. Shapes – WITH lone pairs on central atom 6. Trigonal-Pyramidal (AB 3 E): A – central atom B – bonded atoms E – lone pair - NH 3 - triangular sides - bond angles: 107 o

46. Shapes – WITH lone pairs on central atom 7. Bent or Angular (AB 2 E 2 ): - H 2 O - bond angles: 105 o

47. Molecular Polarity  Polarity of each bond  Molecular polarity

48. Molecular Polarity 1. Has ALL bonds NONPOLAR  nonpolar molecule 2. Has bonds nonpolar AND polar  polar molecule 3. Has ALL bonds POLAR  depends on shape Symmetrical shape (linear - octahedral)  NONPOLAR Non-symmetrical shape (bent & trigonal pyramidal)  POLAR

49. Molecular Polarity Examples  CCl 4  PH 3  CBr 3 H

50. Intermolecular Forces

51.  “between molecule” forces  Generally weaker than bonds b/ atoms  Boiling point – good to measure intermolecular forces

52. Dipole-Dipole Forces  Dipole- equal but opposite charges separated by a short distance  Video 124

53. Dipole-Dipole Forces  Induced Dipole: polar molecule makes a dipole on a nonpolar molecule  Ex: O 2 dissolved in H 2 O  Weaker than regular dipole forces

54. Hydrogen Bonding  Type of dipole-dipole force  Def: H-atom bonded to highly e - neg atom is attracted to lone pair of the e - neg atom in nearby molecule  Ex: HF, H 2 O, NH 3

55. Hydrogen Bonding

56. London Dispersion Forces  Def: constant motion of e - s and creation instantaneous dipoles  Video 133