1. PART II

2. LET’S FIRST REVIEW IONIC BONDING

3. In an IONIC bond, electrons are lost or gained, resulting in the formation of IONS in ionic compounds. FK

4. FK

5. FK

6. FK

7. FK

8. FK

9. FK

10. FK + _

11. FK + _ The compound potassium fluoride consists of potassium (K + ) ions and fluoride (F - ) ions

12. FK + _ The ionic bond is the attraction between the positive K + ion and the negative F - ion An ionic bond is a metal and nonmetal

13. So what are covalent bonds?

14. In covalent bonding, atoms still want to achieve a noble gas configuration (the octet rule).

15. In covalent bonding, atoms still want to achieve a noble gas configuration (the octet rule). But rather than losing or gaining electrons, atoms now share an electron pair.

16. In covalent bonding, atoms still want to achieve a noble gas configuration (the octet rule). But rather than losing or gaining electrons, atoms now share an electron pair. The shared electron pair is called a bonding pair

17. In covalent bonding, atoms still want to achieve a noble gas configuration (the octet rule). But rather than losing or gaining electrons, atoms now share an electron pair. The shared electron pair is called a bonding pair A covalent bond is between a nonmetal and a nonmetal

18. Cl 2 Chlorine forms a covalent bond with itself

19. Cl How will two chlorine atoms react?

20. Cl Each chlorine atom wants to gain one electron to achieve an octet

21. Cl Neither atom will give up an electron – chlorine is highly electronegative. What’s the solution – what can they do to achieve an octet?

22. Cl

26. octet

27. Cl octet

28. Cl The octet is achieved by each atom sharing the electron pair in the middle

29. Cl The octet is achieved by each atom sharing the electron pair in the middle

30. Cl This is the bonding pair

31. Cl It is a single bonding pair

32. Cl It is called a SINGLE BOND

33. Cl Single bonds are abbreviated with a dash

34. Cl This is the chlorine molecule, Cl 2

35.  These are atoms that are never, never, never found alone in nature.  H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2

36.  There are 7 diatomic molecules

37.  Start with 7 on the periodic table  That would be Nitrogen (this is one of the diatomic molecules)

38.  Then make a number 7 on the periodic table

39.  Do not forget the 7 th diatomic molecule

41. Rules: 1. The formula is written with the more electropositive element (the one further to the left on the periodic table) placed first, then the more electronegative element (the one further to the right on the periodic table). Example : CO 2 [Important exception: when the compound contains oxygen and a halogen, the halogen is placed first. If both elements are in the same group, the one with the higher period number is named first.] Example: Cl 2 O.

42. 2. The second element’s name ends in –ide 3. Never use the prefix mono with the 1 st word. 4. Do not put two O’s together

43. Naming covalent bonds Prefixes are used for the number of atoms 1 – mono 2 – di 3 – tri 4 – tetra 5 – penta 6 – hexa 7- hepta 8- octa 9- nona 10 - deca

44. Examples: CO 2 – CO – P 4 Cl 3 - Carbon Dioxide Carbon monoxide Tetraphosphorous trichloride

45. How will two oxygen atoms bond? OO

46. O2O2 Oxygen is one of the diatomic molecules

47. OO Each atom has two unpaired electrons

48. OO

49. OO

50. OO

51. OO

52. OO

53. OO

54. Oxygen atoms are highly electronegative. So both atoms want to gain two electrons. OO

55. Oxygen atoms are highly electronegative. So both atoms want to gain two electrons. OO

56. OO

57. OO

58. OO

59. OO

60. O O Both electron pairs are shared.

61. 6 valence electrons plus 2 shared electrons = full octet O O

62. 6 valence electrons plus 2 shared electrons = full octet O O

63. two bonding pairs, O O making a double bond

64. O O = For convenience, the double bond can be shown as two dashes. O O

65. O O = This is the oxygen molecule, O 2 this is so cool! !

66. Triple bonds are when three sets of electrons are shared.

67. Some covalent bonds can have more than one bond type.

68. OF COVALENT COMPOUNDS

70. Since electrons do not like each other, because of their negative charges, they orient themselves as far apart as possible, from each other. This leads to molecules having specific shapes.

71. Atoms bond to form an Octet (8 outer electrons/full outer energy level) Bonded electrons take up less space then un-bonded/unshared pairs of electrons.

73. Number of Bonds = 2 Number of Shared Pairs of Electrons = 2 Bond Angle = 180° EXAMPLE: BeF 2

74. Number of Bonds = 2 Number of Shared Pairs of Electrons = 2 Number of Unshared Pairs of Electrons = 2 Bond Angle = < 120° EXAMPLE: H 2 O

75. Number of Bonds = 3 Number of Shared Pairs of Electrons = 3 Number of Unshared Pairs of Electrons = 0 Bond Angle = 120° EXAMPLE: GaF 3

76. Number of Bonds = 3 Number of Shared Pairs of Electrons = 4 Number of Unshared Pairs of Electrons = 1 Bond Angle = <109.5° EXAMPLE: NH 3

77. Number of Bonds = 4 Number of Shared Pairs of Electrons = 4 Number of Unshared Pairs of Electrons = 0 Bond Angle = 109.5° EXAMPLE: CH 4

78. Number of Bonds = 5 Number of Shared Pairs of Electrons = 5 Number of Unshared Pairs of Electrons = 0 Bond Angle = <120° EXAMPLE: NbF 5

79. Number of Bonds = 6 Number of Shared Pairs of Electrons = 6 Number of Unshared Pairs of Electrons = 1 Bond Angle = 90° EXAMPLE: SF 6

80. Draw the following and state the shape 1.CCl 4 2.O 2 3.SBr 2 4.NI 3