1. Chemical Bonding

2. Chemical Bonds Compound are formed from chemically bound atoms or ions Bonding only involves the valence electrons

3. Chemical Bonds Defn – force holding two atoms together How are they formed? Atoms gain, lose, or share valence electrons Why does bonding occur? Stability – achieve octet rule

4. Lewis Structures Electron Dot Diagrams –show valence e - as dots –distribute dots like arrows in an orbital diagram –4 sides = 1 s-orbital, 3 p-orbitals –EX: oxygen 2s2p O X

5. Lewis Structures Octet Rule –Most atoms form bonds in order to obtain 8 valence e - –Full energy level stability ~ Noble Gases Ne

6. Electron Dot Structure Shows valence electrons around atomic symbol hydrogen nitrogen chlorine (group 5) (group 7) (group 1) H N Cl

7. Types of Chemical Bonds 3 Types – ionic bond –covalent bond – metallic bond

8. Ionic Bonding

9. Ionic Bonding Vocabulary Ionic compounds are referred to as Formula Units. Compounds are composed to two or more elements. Binary Compound – 2 elements - NaCl Ternary Compound – 3 or more elements – NaHCO 3 Ion – A charged atom Monatomic Ion – 1 atom Na 1+ Polyatomic Ion - 2 or more atoms NO 3 1-

10. Ionic Bond Defn – force holding cations and anions together ABA+A+ B-B- Ionic bond Cations – positively charged ions Anions – negatively charged ions

11. Ionic Bond Where are these bonds found? In Ionic Compounds

12. Ionic Bonding Properties Bond FormationElectrons are transferred from metal to nonmetal Type of StructureCrystal Lattice Physical StateSolid Melting PointHigh Solubility in WaterYes Electrical ConductivityYes – in solutions or liquid

13. Ionic Bonding What’s going on? If I gave you a compound, how can you tell if it is ionic or not? combo of metal + nonmetal giving/taking of valence electrons

15. Formation of Ionic Bonds NaCl NaCl + Na 1+ + Cl 1- 2s 2 2p 6 3s 1 3s 2 3p 5 2s 2 2p 6 3s 2 3p 6 8 v.e.

17. Formation of Ionic Bonds CaBr 2 CaBr + Br Ca 2+ + Br 1- Br 1-

18. Ionic Formulas 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.

19. Ionic Formula Names Write the names of both ions, cation first. Change ending of monatomic anions 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.

20. 1+ 2+3+ 4+/4- 3-2-1- 0 Common Ion Charges (Oxidation States)

21. zpotassium chloride zmagnesium nitrate zcopper(II) chloride K + Cl  Mg 2+ NO 3  Cu 2+ Cl   KCl  Mg(NO 3 ) 2  CuCl 2 Ionic Formulas

22. zNaBr zNa 2 CO 3 zFeCl 3 sodium bromide sodium carbonate iron(III) chloride Ionic Formula Names

23. Metallic Bonding

24. Metallic Bonding Defn – attraction of metallic cations Occurs only in metals

25. Metallic Bonding Properties Bond FormationDelocalized among metal atoms Type of StructureElectron Sea Physical StateSolid Melting PointVery High Solubility in WaterNo Electrical ConductivityYes

26. Metallic Bonding Defn – bond formed from attraction between positive nuclei and delocalized electrons –holds metals together Delocalized Electrons – electrons detached from parent atom –“lost electron away from home”

27. Electron Sea Model Defn – electrons move freely within other molecular orbitals

28. Properties of Metals Electron sea model gives metals certain physical properties 1)Shiny – due to photoelectric effect 2)Conduct electricity and heat – electrons move easily from one place to another 3)Malleable (pound into sheets) 4)Ductile (put into wires)

29. Why malleable and ductile? atoms can also move from one place to another and still remain in contact with and bonded to the other atoms and electrons around them shifted atoms Shape #1 Shape #2

30. Covalent Bonding

31. Covalent Bond Defn – two atoms share one pair of electrons ABA B Electrons shared

32. Covalent Bonds Where are these bonds found? - molecules (molecular compounds) - polyatomic ions

33. Covalent Bonding What’s going on? Molecule – formed when 2 or more atoms bond covalently Sharing of electrons

34. Covalent Bonding Properties Bond FormationElectrons are shared between two nonmetal atoms Type of StructureTrue molecules Physical StateLiquid or gas, brittle solids Melting PointLow Solubility in WaterUsually not Electrical ConductivityNo

35. Two Types of Covalent Bonds i) nonpolar covalent – equal sharing of e - ii) polar covalent – UNequal sharing of e -

36. Nonpolar vs. Polar NONPOLARPOLAR

37. Nonpolar vs. Polar

39. Single Bond Defn – one pair (2) of e - shared Lewis Structures – represents how atoms in molecules are arranged –atoms MUST obey octet rule (except hydrogen)

40. Lewis Structures bonded electrons – occur between bonded atoms A B AB single bond or

41. Lewis Structures Unshared or Lone Pairs – electron pairs NOT involved in bonding AB A B lone pairs

42. Lewis Structures Examples H 2 O H H (8 valence e - or 4 pairs) O O H H O H H

43. Lewis Structures Examples NHF 2 (20 v.e. or 10 pairs) N F F H F N F H N F F H

44. Multiple Covalent Bonds Double Bond – two pairs (4) e - shared A B AB O O2O2 O (12 v.e. = 6 pairs) O O O O O O

45. Multiple Covalent Bond Triple Bond – three pairs (6) e - shared A B AB N2N2 (10 v.e. = 5 pairs) N N N N N N N N

46. Comparing single, double, and triple bonds Bond Strength Bond Length Triple > Double > Single Single > Double > Triple The shorter the bond, the stronger it is

47. Polyatomic Ions Defn – CHARGED group of atoms covalently bonded - ex: SO 4 2-, NH 4 1+, NO 3 1-

48. Polyatomic Ions SO 4 2- (32 v.e. = 16 pairs) O O O O S 2- O O O O S 2-

49. Polyatomic Ions NH 4 1+ (8 v.e. = 4 pairs) H H HHN 1+ H H HH N 1+

50. Using electronegativity to determine bond type Recall electronegativity: how much an atom wants electrons Each atom is assigned a number between 0-4.0 to determine electronegativity strength

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

52. ++ -- ++ Lewis Structures Nonpolar Covalent - no charges Polar Covalent - partial charges

53. Bond Polarity Electronegativity –Attraction an atom has for a shared pair of electrons. –higher e - neg atom   - –lower e - neg atom   +

54. Bond Polarity Electronegativity Trend Increases up and to the right.

56. We know 3 types of bonds: - nonpolar covalent - polar covalent - ionic To determine bond type, subtract electronegativity values and see scale Using electronegativity to determine bond type

57. Scale 04.0 1.70.3 polar covalent nonpolar covalent ionic Using electronegativity to determine bond type

58. H and Cl3.0 – 2.1= 0.9 polar covalent C and S2.5 – 2.5= 0 nonpolar covalent Na and F4.0 – 0.9= 3.1 ionic Using electronegativity to determine bond type

59. Dipole Moment defn – imbalance of electron density in a covalent bond –Due to electronegativity of atoms  - ( partial negative) = signifies more EN atom  + (partial positive) = signifies less EN atom = shows direction of dipole moment

60. Examples H = 2.2 C = 2.6 N = 3.0 Cl = 3.2 O = 3.4 F = 4.0 HO ++ -- ClC -- ++ NH -- ++ CF ++ --

61. Naming Covalent Compounds 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.

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

63. zCCl 4 zN 2 O zSF 6 carbon tetrachloride dinitrogen monoxide sulfur hexafluoride Covalent Compound Names

64. zarsenic trichloride zdinitrogen pentoxide ztetraphosphorus decoxide AsCl 3 N2O5N2O5 P 4 O 10 Covalent Formulas

65. NOF Cl Br I H C. Molecular Nomenclature The Seven Diatomic Elements Br 2 I 2 N 2 Cl 2 H 2 O 2 F 2

66. Intermolecular Forces

67. Defn – attractive forces between 2 molecules

68. Intermolecular Forces Dipole-Dipole – attraction between oppositely charged polar molecules ++ -- ++ -- ++ --

69. Intermolecular Forces London Dispersion Forces – very weak, very brief dipole moment created in nonpolar molecules

70. Electrons evenly distributed Temporary dipole London force

71. Intermolecular Forces Hydrogen Bonding – strong bond between H and N,O, or F of another molecule - Water is prime example O H H O ++ ++ --

72. O H H O O H H O O H H O O H H O O H H O hydrogen bond ++ ++ -- ++ ++ --

73. Strength Ranking Hydrogen > dipole-dipole > London

74. VSEPR Valence Shell Electron Pair Repulsion Defn – determines the shape of molecule Electron pairs try to stay far away as possible

75. # atoms bonded to central atom # lone pairs shape 40tetrahedral

76. Tetrahedral

77. # atoms bonded to central atom # lone pairs shape 4 3 1 0tetrahedral trigonal pyramidal

78. Trigonal Pyramidal

79. # atoms bonded to central atom # lone pairs shape 4 3 1 2 0 2 tetrahedral trigonal pyramidal bent

80. Bent

81. # atoms bonded to central atom # lone pairs shape 4 3 1 3 2 0 2 0 tetrahedral trigonal pyramidal bent trigonal planar

82. Trigonal Planar

83. # atoms bonded to central atom # lone pairs shape 4 3 1 3 2 2 0 2 0 0 tetrahedral trigonal pyramidal bent trigonal planar linear

84. Linear