2. Covalent Bonding Molecular Compounds

3.  Describe a covalent bond in terms of the difference in electronegativity of the atoms and the energy changes in bond formation.  Correlate bond energy and bond length for single, double and triple bonds.  State the octet rule and families of elements involved.  Write the Lewis electron dot structures of compounds and ions containing covalent single, double and triple bonds.  Explain resonance and apply it to determine the bond orders of relevant compounds. Objectives

4.  Describe the shapes of molecules using Valence Shell Electron Pair Repulsion (VSEPR) Theory.  Identify appropriate molecular shapes (linear, trigonal planar, tetrahedral, pyramidal and angular/bent) and specify the corresponding bond angles.  Describe the effect of unshared electron pairs on bond angles.  Describe the formation of σ (sigma) and π (pi) bonds and identify them in molecules.  List the types of intermolecular forces and rank them by strength.  Determine whether a bond is polar or non-polar.  Determine whether a molecule is polar or non-polar.  Write the name of a binary covalent compound given its formula, or its formula given its name.

5. LET’S FIRST REVIEW IONIC BONDING

6. In an IONIC bond, electrons are lost (metals) or gained (nonmetals), resulting in the formation of IONS in ionic compounds. FK

7. FK

8. FK

9. FK

10. FK

11. FK

12. FK

13. FK + _

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

15. FK + _ The ionic bond is the attraction between the positive K + ion and the negative F - ion

16. So what are covalent bonds?

17. A type of bonding between non- metals. This involves the sharing of electrons by 2 or more atoms. The electrons are shared, not transferred as with ionic bonds. Covalent Bonding

18. Covalent Compounds Are made up of two NONMETALS. (Both elements are from the right of the step line (green elements shown)).

19. When nonmetals bond together a covalent bond is created and we call them molecules or molecular compounds!

20. Molecules Molecules are neutral atoms that are joined together by covalent bonds Molecular formula - shows you how many atoms of each element is in a substance Example: CO 2, NH 4

21. Formation of a Covalent Bond

22. The electrons move freely about in either orbital.

23. Bond length - the distance between the two atoms nuclei. Bond energy - the amount of energy required to break a chemical bond.

24. Octet Rule and Covalent Bonding An octet is 8 valence electrons that want to achieve a noble gas configuration! Molecules want the same thing, but they share their valence electrons to achieve the octet rule.

25. EXAMPLES INCLUDE THE BONDS BETWEEN: H 2 F 2 Br 2 Cl 2 HCl H 2 O A type of bonding between non- metals. Other than hydrogen these elements are on the right side of the periodic table.

26. 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

27. Lewis Dot Structures Formulas in which atomic symbols represent the element and all inner-shell electrons, dots represent valence electrons and dashes between two atomic symbols represent electron pairs in covalent bonds. Cl

28. Question In the conductivity lab, why did ionic compounds conduct electricity and molecular compounds did not conduct electricity?

29. Question In the conductivity lab, why did ionic compounds conduct electricity and molecular compounds did not conduct electricity? Answer Formation of ions with ionic compounds

30. Single Covalent Bonds When atoms share one pair of electrons they form a single covalent bond

31. Cl 2 Chlorine forms a covalent bond with itself

32. Cl How will two chlorine atoms react?

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

34. Cl what can they do to achieve an octet? Share unpaired electrons!

35. Cl

39. octet

40. Cl octet

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

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

43. Cl This is the bonding pair (shared pair of electrons)

44. Cl It is a single bonding pair

45. Cl It is called a SINGLE BOND

46. Cl Single bonds are abbreviated with a dash Normally in the final structure the valence electrons are not drawn

47. Cl This is the chlorine molecule, Cl 2

48. Try these on your own: HCl 1) Hydrogen and chlorine form a single bond to yield hydrochloric acid 2) Two hydrogen atoms combine with one oxygen atom to yield water H2OH2O

49. Classwork Page 244 Questions 1, 2, 3, 4, 5

50. Double Bonds Sharing of two pairs of electrons between two atoms

51. O2O2 Oxygen is also one of the diatomic molecules

52. How will two oxygen atoms bond? OO

53. OO Each atom has two unpaired electrons

54. OO

55. OO

56. OO

57. OO

58. OO

59. OO

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

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

62. OO

63. OO

64. OO

65. OO

66. O O Both electron pairs are shared.

67. 6 valence electrons plus 2 shared electrons = full octet O O

68. 6 valence electrons plus 2 shared electrons = full octet O O

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

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

71. O O = This is the oxygen molecule, O 2

72. Triple Bonds Atoms that share three pairs of electrons: Example: N 2

74. Bond length - the distance between the two atoms nuclei. Bond energy - the amount of energy required to break a chemical bond.

75. Multiple Bonds Double bonds are stronger and shorter than single bonds. Triple bonds are even stronger and shorter than both double and single bonds. Common between carbon, oxygen and nitrogen atoms.

76. Bond Lengths and Bond Energies

77. Try these on your own: NH 3 CH 4 H 2 O 2 PCl 3

78. Try these on your own: CH 2 O CO 2 HCN

79. Properties of Covalent Compounds Switch PowerPoint Presentation – slide 80

80. Covalent bonds can be polar or non-polar. If a bond is non-polar, that means that there is an equal sharing of electrons between atoms (Cl 2 ). If a bond is polar, that means that the electrons are not shared equally, making one side of the bond more negative (where the electrons end up) and the other side more positive (where the electrons were). Example is water (H 2 O)

81. Electronegativity This chart will help you determine if it is polar or nonpolar 0.0-0.3 Non polar covalent 0.3-1.7 polar covalent >/= 1.8 Ionic

82. Are formed when 2 atoms with different electronegativities form a covalent bond. The atoms DO NOT SHARE electrons equally. The electrons are more strongly attracted to the MOST electronegative atom. This uneven sharing of electrons creates a dipole force (tiny magnet) and is indicated with partial negative and positive charges. atom. electronegativity Polar Compounds

83. Bond polarity Since, atoms are sharing within a covalent bond… If they share unequally they are a polar covalent bond Examples: H 2 O

84. Homework Page 275 Problems 118, 119 and 120

85. Molecular Geometry Switch PowerPoint Presentation – slide 88

86. End of Covalent Bonding

87. Intermolecular Forces

88. What makes water (H 2 O) so ideally suited to sustain life on earth? Why is water unique compared to a similar compound (H 2 S)? Intermolecular forces Intermolecular Forces

89. Boiling Points: water = 100 o C (liquid) H 2 S = -60 o C (gas) Water is one of the few compounds where the density of the solid (ice) is less than the density of the liquid. Ice floats – so why is this important!

90. Intermolecular Forces – the forces of attraction between molecules. Not actually bonds. These forces are weaker than covalent bonds, ionic bonds and metallic bonds.

91. Dipole – is created by opposite charges that are separated by a short distance. The dipole’s direction is from the positive pole to the negative pole. The dipole (magnet) is created due to electronegativity differences between atoms. H Cl (H= 2.1; Cl = 3.0)

92. The negative region in one polar molecule attracts the positive region in adjacent molecules. The forces of attraction between polar molecules are known as dipole-dipole forces.

93. Dipole forces between I-Cl molecules Electronegativity of I = 2.7; Cl = 3.0

94. This intermolecular attraction between molecules causes the molecules to be held together tighter causing increasing boiling points. Example: boiling point of I-Cl is 97 o C; boiling point of F-F is -188 o C

95. Examples: molecules

96. A force keeping molecules attracted to each other which always includes hydrogen. (hence this name). This is a strong type of dipole-dipole force. When hydrogen bonds to certain highly electro- negative atoms such as oxygen, nitrogen or fluorine, the bonds are very polar. This leaves the hydrogen rather positive, wishing for some negativity in the form of electrons. Since it can’t pull its electron back from these atoms, it actually is attracted to the UNSHARED PAIRS of electrons from other nearby molecules.. Hydrogen Bonding

97. The + signs in blue are the hydrogen atoms. See how they can be attracted to the blue UNSHARED PAIRS of electrons of a nearby water molecule? That is hydrogen bonding. the hydrogen bond is only about 5% the strength of a covalent bond, but much stronger than a dipole weak bond

98. Hydrogen Bonding Example: boiling point of H 2 S: -61 o C; boiling point of H 2 O: 100 o C Hydrogen bonds to oxygen, nitrogen or fluorine

99. Induced Dipole Forces A temporary dipole created when a polar molecule interacts with a nonpolar molecule. Very weak forces

100. Van Der Waal’s Forces Also called dispersion force, which is due to the movement of molecules and to the number of electrons around. They create temporary but real attractions between molecules.

101. When you have a lot of electrons and they are all buzzing around quickly from time to time a temporary condition can be created where there are more electrons in one area. This creates a temporary but real electrostatic force of attraction between the (+) and the (-) zones created. The more electrons you have (Iodine has lots more than fluorine for example) the more opportunity for this temporary condition to exist. It’s weak but real and in group 17 at STP the results make for some gases, a liquid and a solid at constant STP.

102. Homework