1. Chemical Bonding Chapter 8 & 9

2. I. Overview/Types of Compounds

3. Overview of Bonds Chemical Bonds are attractions between atoms or ions in a compound A bond forms when valence electrons are either shared or transferred from one atom to another Most elements are more stable when bonded to other atoms, which is why bonds are formed.

4. Compounds Before we start learning details about bonding, let’s review compounds!

5. A. Definition of a Compound Remember, a compound consists of 2 or more elements chemically bonded together. It is always neutral. EX: H 2 0, NaCl, Sb 3 (PO 4 ) 5

6. B. Types of Compounds 2 types of chemical compounds

7. 1. Molecular or Covalent Compound: consists of only non-metal atoms that share electrons which form a covalent bond. –Examples: H 2 0, C 8 H 18 –smallest unit is called a molecule –molecules also include diatomics (ex. H 2 )

8. 2. Ionic Compound: a compound that is formed from the attraction between ions of opposite charges. It is held together by an ionic bond and usually contains a metal and nonmetal.

9. Ionic Compounds Examples: NaCl, KBr smallest unit is called a formula unit Cations are + ions. This happens when an atom loses electrons. (ca+ion) Anions are – ions. This happens when an atom gains electrons. (a negative ion)

10. Practice 1. Determine whether the following are ionic or molecular covalent compounds a. N 2 O 5 b. PbNO 5 c. KFd. AgCl e. PCl 3 ionic molecular ionic

11. always 1+ 2+3+3-2-1-

12. Practice 2. Write the ion formed by these elements a.iodine b.potassium c.sulfur d.gallium e.strontium S -2 I-I- Sr 2+ Ga +3 K+K+

13. C. Properties of Compounds ionic and molecular covalent compounds typically have very different properties this stems from the fact that the ionic attraction between formula units is far stronger than the attraction between covalently- bonded molecules

14. Ionic Bond (strong bond) = Mr. & Mrs. Masculino Covalent Bond (weaker bond) = Mrs. Masculino & Justin Bieber

15. Comparison of Properties CharacteristicMolecular Compounds Ionic Compounds representative unit MoleculeFormula Unit type of elementsNonmetals onlyMetals and Nonmetals type of bondingCovalent (sharing) Ionic (transfer) physical state at RT S, L, or GSolid only melting point<300 °C **> 300 °C hardnesssoft  hardhard conduct electricity? noyes (when dissolved in water) ExamplesCandle waxRock salt (NaCl) **would melt in a bunsen burner

16. Ionic Compounds = transfer of electrons Na Cl + - NaCl Is formed

17. Molecular Compounds = share electrons H O H2OH2O H

18. II. Types of Bonds We will learn about 4 different types of bonds: ionic, non-polar covalent, polar covalent, and metallic Bonds between two elements can be categorized based on the difference between the electronegativities of those two elements.

19. A. Ionic Bonds

20. Ionic Bonds Ionic bonds are the result of the transfer of electron. It holds an ionic compound together. Remember, an ionic compound is a compound formed from the attraction between a cation and anion, which usually includes a metal and non-metal.

21. Ionic Bonds After a cation and anion are bonded to form an ionic compound, the net charge is zero The Electronegativity difference between the two atoms bonded is >1.7

22. [ ] - Ionic Bonds Na Na + Cl

23. Ionic Compounds = transfer of electrons Na Cl + - NaCl Is formed

24. Drawing ionic compounds [ ] Cl + - Na Remember when drawing ionic compounds, only the anion is in brackets & has its valence electrons showing

25. Another Example 2+ [ ] O 2- Ca

26. Ionic Bonds  Strongest type of bond between atoms. This results in high melting points and high boiling points.  Formula units form a crystalline structure

27. B. Covalent Bonds

28. A covalent bond is formed by the sharing of electron-pairs within molecules Covalent bonds are formed between non-metals Ex. H O H

29. Covalent Bonds Two atoms can share more than one covalent bond between them –Single bonds form between two atoms that SHARE 1 PAIR of electrons. –Double bonds form between two atoms that SHARE 2 PAIRS of electrons. –Triple bonds form between two atoms that SHARE 3 PAIRS of electrons.

30. Triple Bonds have the shortest bond length Bond LengthBond Energy H LongestWeakest (163 kJ/mol) O Medium(418 kJ/mol) N ≡ NShortestStrongest (945 kJ/mol)

31. There are two types of covalent bonds: non-polar covalent and polar covalent

32. Non-Polar Covalent Bonds electronegativity difference between atoms: 0-0.3 even distribution of electrons symmetrical

33. observed in all diatomic atoms (BrINClHOF elements) BrIClNHO F

34. Use multiple resources to answer the following questions 1.Define polar and non-polar covalent bonds 2.Which has a dipole? What is a dipole? 3.Provide at least 2 examples of each & draw pictures too 4.How can you identify polar from non-polar molecules? 5.What are the ranges for differences in electronegativities?

35. Polar vs Non-Polar Covalent Bonds

36. Polar Covalent Bonds electronegativity difference between atoms: 0.3-1.7 polar means “having opposite ends” unequal sharing of charges

37. Remember Electronegativity? © 2003 Prentice Hall

38. Polar Covalent Bonds A dipole is formed the partial charges can be shown with the following symbols: δ - and δ+

39. Unequal sharing of electrons © 2003 Prentice Hall

40. Polar Covalent Bonds electrons are more attracted to the more electronegative element Opposites attract concept is observed

41. Polar or Non-Polar? Has an even distribution of charge Non-Polar!!

42. Polar or Non-Polar? Has a Δ EN = 0 Non-Polar!!

43. Polar or Non-Polar? Has a δ + and δ - end Polar!!

44. Polar or Non-Polar? I 2 Non-Polar!!

45. Polar Means: “Having Opposite Ends” (like N and S Poles)

46. C. Metallic Bonds

47. A metallic bond is the bond that holds two metals together. The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons.

48. Metallic Bonds The bonding electrons are shared over the whole metal sample. This is known as electron delocalization.

49. Because of the nature of metallic bonding, metals are malleable and ductile. -malleable = ability of a substance to be hammered or beaten into thin sheets -ductile = ability of a substance to be drawn or pulled through a small opening to produce a wire

50. Summary Chart Don’t forget your general rules of Ionic compounds = Metal + Non-Metal Type of Bond (between 2 atoms) ΔEN (Paulings) Metallic bond not quantified using EN Ionic (transfer of electrons) >1.7 Polar Covalent (unequal sharing of electrons) 0.4 – 1.7 Non-Polar Covalent (equal sharing of electrons) 0 – 0.4 diatomics Use a table/chart with electronegativity values to determine the difference

51. Electronegativity Values

52. 0.31.7 © 2003 Prentice Hall

53. H – F ΔEN 1.9 H – Cl ΔEN 0.9 H – Br ΔEN 0.7 H – I ΔEN 0.4 Most PolarLeast Polar

54. Although a compound is an ionic compound, it can still possess covalent bond characteristics and vice versa Ex. H – F Ionic Character Increases Covalent Character Increases 2.1 4.0 ΔEN = 1.9 While this is a molecular compound, it possess ionic bond characteristics

55. III. Lewis Structures

56. Only for molecular compounds! Recall drawing electron dot diagrams using the number of valence electrons (Roman Numerals on PT)

57. Lewis Structures are drawings!! Symbol = Nuclei Dots = Valence Electrons Dashes = Shared Electron Pair Electron Pair = unshared (aka “lone pair”)

58. = electron pair 1 bond = 2 shared electrons

59. Electron pairs that are shared can be replaced with a line to represent a covalent bond. Up to three bonds can be shared between two elements

60. Single bonds are formed between two elements that have 2 shared electrons (Ex. ) Double bonds are formed between two elements that share 4 valance electrons (Ex. ) Triple bonds are formed between two elements that share 6 valence electrons (Ex: )

61. Triple Bonds have the shortest bond length Bond LengthBond Energy N LongestWeakest (163 kJ/mol) N Medium(418 kJ/mol) N ≡ NShortestStrongest (945 kJ/mol)

62. Steps to Drawing Lewis Structures Choose the center atom (Tips: choose the one you have less of, and carbon loves to be the center of attention) Draw an electron dot structure for the central atom Draw the remaining electron dot structures near the central atom’s unpaired electrons Create bonds between the unpaired electrons Rearrange your structure so it looks “neat” (spread the atoms out) Add lone pairs so elements 1-5 have “duets” and all other elements have “octets” Disclaimer: These rules actually only work for the examples we’re doing in our class. There are far more extensive rules that we will ignore for the sake of simplification.

63. Octet Rule Octet Rule: chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level -Elements 1-5 need to be surrounded by 2 electrons -All other elements need to be surrounded by 8 electrons

64. Same as

65. © 2003 Prentice Hall

66. Examples CCl 4 C Cl

67. Examples C2H2C2H2 C H H C C H H C Are the duets & octets complete?

68. Examples CO 2 O O C O O C Are the duets & octets complete?

69. C O O Examples CO 2 O O C Are the duets & octets complete?

70. Exceptions to the octet rule (incomplete octets) Be H H B F F F BF 3 BeH 2

71. Exceptions to the Octet Rule BF 3 3 valence & 3 bonded atoms BeH 2 2 valence & 2 bonded atoms

72. What does polar and non-polar mean again?

73. Review on how to draw structures (Ex. NH 3 ) N H H H

74. Polar N H H H S S Non-Polar EN = 2.1 EN = 3.0 EN = 2.1

75. Polar & Non-Polar Molecules Like bonds, molecules can be polar (uneven distribution of electrons) or non-polar (even distribution) Non-polar molecules are symmetrical and have no lone pairs around the center atom Polar molecules have lone pairs around the central atom or an uneven distribution of electrons

76. NP O O C Practice Identify each molecule as polar or non-polar NP

77. B F F F B F F F

78. P Practice Identify each molecule as polar or non-polar NP

79. C 2 HCl Practice Identify each molecule as polar or non-polar P C Cl H C

80. Polar & Non-Polar Molecules © 2003 Prentice Hall

81. Dipoles Some molecules have dipoles that cancel out and have a zero net dipole The poles can be combined to form a net dipole H H O Positive end Negative end

82. Polar or Non-polar Molecules? CCl 4 C Cl EN = 2.5 EN = 3.0

83. C2H2C2H2 C H H C Polar or Non-polar Molecules? CO 2 O O C Polar or Non-polar Molecules?

84. N H H H A better way to draw it N H H H Br

85. B F F F Draw BF 3 B F F F

86. IV. VSEPR Theory

87. VSEPR Theory VSEPR stands for Valence Shell Electron Pair Repulsion VSEPR is used to describe the 3D orientation of the electron regions. Shared electrons that make up covalent bonds and lone pairs will repel each other as much as possible

88. B F F F N H H H BF 3 vs NH 3 Shape: Trigonal Planar Shape: Trigonal pyramidal

89. Lone pairs repel more than bonds do BF 3 vs NH 3 Shape: Trigonal Planar Shape: trigonal pyramidal Look at the chart on page 7 of your notes

90. Ahhh! Memorization!? You will be responsible for predicting shapes with a * next to the name of shape without any notes You should be able to determine the shape of non - * compounds given this chart

92. 1 lone pair Trigonal planar * /bent* Linear *

93. 1 lone pair 2 lone pairs tetrahedral * /bent *