1. Chemical Bonding b Chapter 6

2. Chemical bond b The force (electrical attraction) that binds two atoms together

3. Types of bonds: b Ionic bonds b Covalent bonds b Metallic bonds

4. lonic Bond b A chemical bond between cations (+) and anions (-).

5. Covalent Bond b one or more pairs of electrons are shared by two atoms.

6. Two types of covalent bonds: b polar covalent bond b nonpolar covalent bond

7. Polar covalent bonds: b unequal sharing of e-’s b results in partial charges b occur between atoms of differing electronegativities

9. Nonpolar covalent bonds: b equal sharing of e-’s b no partial charges b electronegativity differences are ≤ 0.4

10. Electronegativity differences: b Ionic bonds: ≥1.7 b Polar covalent bonds: 0.4 to1.7 b Nonpolar covalent bonds: ≤ 0.4

11. Electronegativity :

12. Electronegativity:

13. Types of bonds: Each O-H bond: 3.44 3.44-2.20 1.24 1.24 two polar covalent bonds

14. Determine Bond Types for: b H-O b K-F b Na-F b C-H b N-H b O-Cl

15. Octet Rule Exceptions b Period 1, Li & Be: only 2 valence electrons is a full shell b B: often has only 6 valence electrons b P & S: may have 10 or 12 valence electrons

16. Molecule b atoms held together by covalent bonds.

17. Diatomic Molecules b 2 atoms of one element bonded together b H 2 N 2 O 2 F 2 Cl 2 Br 2 I 2

18. Lewis Electron Diagrams b Dots represent valence electrons. b written on 4 sides around symbol b 2 electrons on same side are paired

19. Forming a bond:

20. b Single Bond: one pair of shared electrons (2 e-) b Double Bond: two pairs of shared electrons (4 e-) b Triple Bond: three pairs of shared electrons (6 e-)

21. Lewis Structures b a pair of electrons is represented with a dash instead of 2 dots is single bond is single bond = is a double bond ≡ is a triple bond

22. Draw Lewis Structures of: b PH 3 b H 2 S b HCl b CCl 4 b SiH 4

23. Sigma Bond, σ b a single covalent bond b electron pair is shared by the overlap of bonding orbitals

24. Pi Bond, π b a bond formed when parallel orbitals overlap to share electrons

25. Resonance

26. Resonance b A condition that occurs with multiple valid Lewis structures b A combination of all possible structures

27. Bond (dissociation) Energy b The energy it takes to break a bond.

28. Bond lengths: b Distance between nuclei at lowest potential energy.

29. Bond lengths: b Smaller atoms have shorter bond lengths. b Multiple bonds are shorter than single bonds. b Shorter bonds are stronger than longer bonds

31. lonic Bond b A type of chemical bond between cations (+ charge) and anions (- charge).

32. For every compound, the total positive charge MUST equal the total negative charge

33. X 2 Y: has 2 X’s and 1 Y b a subscript following a symbol shows relative numbers of that ion or atom in a compound

34. Combine these ions in all possible ways to form a binary ionic compound: Au 3+ Pb 2+ O 2– Cl – Au2O3PbO AuCl3PbCl2

35. Formula Unit (Empirical Formula) b the smallest whole- number ratio of ions that make up a compound

37. Coordination number: number of ions of opposite charge surrounding the ion in a crystal

39. Ionic Compounds b high melting points which show strong bonds b brittle b most dissolve in water and dissociate into ions, which conduct electricity

40. Electrolyte b An ionic compound that conducts an electric current when dissolved in water.

41. Lattice Energy b The energy needed to separate ions in an ionic compound

42. b Monoatomic Ions: ions from only one atom b Polyatomic Ions: ions from more than ions from more than one atom one atom

43. Draw Lewis Structures for: b NH 3 b NH 4 + b SO 4 2 ‒ b PO 4 3 ‒

44. Metallic Bond b Attraction of metal cations and delocalized electrons

45. Metallic Bond Strength increases with: b more mobile valence electrons b smaller atom sizes

46. VSEPR (Valence Shell Electron Pair Repulsion) b Pairs of molecular valence electrons stay as far away from other pairs as possible

47. Linear Linear (AX 2 ) b Bonded atoms form a straight line b Bond angle = 180°

48. Trigonal Planar Trigonal Planar (AX 3 ) b the bonded atoms form a triangle. b Bond angles = 120°

49. Tetrahedral Tetrahedral (AX 4 ) b shape around central atom (A) that has 4 single bonds(X 4 ) Bond angles =109.5° Bond angles =109.5°

50. (Trigonal) Pyramidal (Trigonal) Pyramidal (AX 3 E) b Central atom has 3 single bonds (X 3 ) and one lone pair of electrons (E). Bond angles = 107.3° olar, Bond angles = 107.3° olar,

51. Bent Bent (AX 2 E 2 ) b Central atom has two single bonds (X 2 ) and two unshared pairs of electrons (E 2 ). Bond angles = 104.5° Bond angles = 104.5°

52. Shapes & bond angles? b HCN b AsI 3 b H 2 Te b PF 3 b CBr 4

53. Hybrid Orbitals b Orbitals of bonded valence electrons containing a combination of properties of the original orbitals.

54. b sp hybrid orbital – from an s and a p orbital (linear) b sp 2 hybrid orbital – from 1 s and 2 p orbitals (trigonal planar) b sp 3 hybrid orbital – from 1 s and 3 p orbitals (bent, tri.pyramidal, tetrahedral)

55. Dipole: b An asymmetrical distribution of positive and negative charge within a molecule. b The central position of + charge is not the same as the central position of ‒ charge.

56. Polar bonds, polar molecule

57. Polar bonds, nonpolar molecule

58. b For a molecule to be polar, at least one bond must be polar.

59. b If there is no dipole (symmetrical), the molecule is nonpolar b If there is a dipole (asymmetrical), the molecule is polar

60. H...... │.. :Cl−C −Cl: :Cl−C −Cl: ˙ ˙ ˙ ˙ ˙ ˙ │ ˙ ˙ H H Polar or nonpolar?

61. b Bond polarity: electronegativity difference b Molecule polarity: 1) must have a polar bond 2) must have a dipole

62.  Intramolecular force:  Intramolecular force: A force that exists within a molecule  Intermolecular force: A force that exists between molecules

63. b Intramolecular & intermolecular forces:

64. Intramolecular & intermolecular forces:

65. b b (London) Dispersion forces b b Dipole-dipole forces b b Hydrogen bonds Intermolecular Forces:

66. Induced Dipole: b A dipole created by the presence of a neighboring dipole.

67. b intermolecular force of attraction between induced dipoles. These increase with molecular size and mass. (London) Dispersion Force:

68. b b intermolecular force of attraction between neighboring permanent dipoles. Dipole-Dipole Force:

69. b b intermolecular forces between the hydrogen atom of one molecule and a highly electronegative atom of another. Hydrogen Bonds:

70. Water (s) Water (s)

71. Review #1: What determines if a molecule is polar? MUST have a dipole

72. Review #2: What is an intramolecular force? A force that exists within a molecule

73. Review #3: What are intermolecular forces? The forces that exist between molecules: dispersion, dipole- dipole, and hydrogen bonds

74. If a compound starts with a metal or “NH 4 ”, what kind of compound is it? If a compound starts with a metal or “NH 4 ”, what kind of compound is it? Ionic Review #4:

75. If a compound starts with any other nonmetal, what kind of compound is it? Molecule: covalent bonds Review #5: