2. Chapter 6 Bonding

3.  Write the electron configuration of ions of representative elements.  Define ionic bonds in terms of the difference in electronegativity of the atoms.  Write the Lewis electron dot structures of atoms, ions and ionic compounds.  Write the name of an ionic compound (binary, polyatomic and transition metal) given its formula, or its formula given its name.  Describe the properties of ionic, metallic and covalent molecules and identify the forces holding them together. Objectives

4. Valence Electrons Mendeleev ordered his periodic table according to chemical behavior. A pattern he saw was that elements in the same group behaved the same chemically. This is due to its valence electrons. Valence electrons are the outermost electrons and are involved in bonding.

5. Calculating valence electrons Groups on the periodic table can help! What’s left out on the periodic table?

6. Transition Metals Transition metals can have different amounts of valence electrons depending on how they bond with nonmetals

7. Decide how many valence electrons each one has? 1.Li 2.F 3.Kr 4.Mg 5.Group 13 6.Carbon 7.Phosphorus 1. 1 valence electron 2. 7 valence electron 3. 8 valence electron 4. 2 valence electron 5. 3 valence electron 6. 4 valence electron 7. 5 valence electron

8. Gilbert Lewis Lewis developed the idea of the octet and coined the term Lewis dot structures He was nominated 35 times for the Nobel Prize in chemistry, but never won.

9. Electron Dot structures A diagram that shows only the valence electrons around the atom as dots Some rules: Each side can only have 2 dots, for a maximum amount of 8 dots An Exception to group 18 is Helium, it will only have 2

10. Electron Dot structures Draw the correct Lewis dot structure for each given: 1 V.E2 V.E3 V.E 4 V.E5 V.E 6 V.E7 V.E8 V.E

11. Try these on your own:

12. Electron dot structures

13. Think back to Nobel Gases Nobel gases are unreactive… What do you notice about their electron dot structures??

14. Octet Rule Atoms want to achieve stability or noble gas configuration by attaining 8 valence electrons. Maximum number in the s and p orbitals. Atoms do not want to gain or lose more than 3 electrons when bonding!

15. Octet Rule for F 2

16. Octet Rule for HCl

17. Which is the only group that has full octets? NOBLE GASES!

18. Ions

19. Ions are formed when an atom loses or gains electrons Ions can gain or lose just one or multiple electrons to bond and achieve stability Ions do not want to gain or lose more than 3 electrons though when bonding

20. Cations Cations are formed when an atom loses electrons and becomes positive Metals form cations Examples: magnesium, potassium, aluminum

21. Anions Anions are formed when an atom gains electrons and becomes negative Nonmetals form anions Examples: fluorine, oxygen, phosphorus

22. How do we know if atoms lose or gain electrons? Let’s look at their Valence electrons Every atom wants to be at 8 valence electrons Sodium has 1 VE, so its easier to lose 1 and have a +1 Chlorine has 7 VE, so its easier to gain 1 and would have a – 1

23. Charge or Oxidation state The charge corresponds with the amount of valence electrons:

24. Did it lose or gain? Oxide ion Cesium ion Aluminum ion Bromide ion Phosphide ion 1. O -2 : gained 2 electrons 2. Cs +1 : lost 1 3. Al +3 lost 3 electrons 4. Br -1 gained 1 electrons 5. P -3 gained 3 electrons

25. Valence electrons: Gain or lose, and how many? 1.Strontium 2.Hydrogen 3.Sulfur 4.Xenon 5.Iodine 1. Lose 2 2. Lose 1 3. Gain 2 4. Neither, neutral 5. Gain 1

26. Polyatomic Ions

28. Ions Review

29. Predicting Ionic Charges Group 1: Lose 1 electron to form 1+ ions H+H+H+H+ Li + Na + K+K+K+K+

30. Predicting Ionic Charges Group 2: Loses 2 electrons to form 2+ ions Be 2+ Mg 2+ Ca 2+ Sr 2+ Ba 2+

31. Predicting Ionic Charges Group 13: Loses 3 electrons to form 3+ ions B 3+ Al 3+ Ga 3+

32. Predicting Ionic Charges Group 15: Gains 3 electrons to form 3- ions N 3- P 3- As 3- Nitride Phosphide Arsenide

33. Predicting Ionic Charges Group 16: Gains 2 electrons to form 2- ions O 2- S 2- Se 2- Oxide Sulfide Selenide

34. Predicting Ionic Charges Group 17: Gains 1 electron to form 1- ions F 1- Cl 1- Br 1- Fluoride Chloride Bromide I 1- Iodide

35. Predicting Ionic Charges Group 18: Stable Noble gases do not form ions!

36. Predicting Ionic Charges Groups 3 - 12: Many transition elements have more than one possible oxidation state. Iron(II) = Fe 2+ Iron(III) = Fe 3+

37. Predicting Ionic Charges Groups 3 - 12: Some transition elements have only one possible oxidation state. Zinc = Zn 2+ Silver = Ag +

38. Chapter 6 Types of Bonding

39. Atoms seldom exist in nature as independent particles. Nearly all substances are made up of a combination of atoms that are held together by chemical bonds. Chemical Bond – a mutual electrical attraction between the nuclei and valence electrons of different atoms that bind the atoms together.

40. When atoms bond, their valence electrons are redistributed in ways that make the atoms more stable. The ways in which the electrons are redistributed determines the type of bonding. Metals tend to lose electrons to form positive ions, or cations, and nonmetals tend to gain electrons to form negative ions, or anions.

41. Chemical bonding that results from the electrical attraction between cations and anions is called ionic bonding. Chemical bonding that results from the sharing of electrons between atoms is covalent bonding. In covalent bonds the shared electrons are owned equally by the two bonded atoms.

43. You can estimate the type of bonding (ionic or covalent) between elements by the difference in electronegativities (Chapter 5).

44. 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

45. Polar Covalent vs. Non-polar Covalent Blue shading represents electron density

46. Use electronegativity differences and Figure 6-2 to classify bonding between sulfur, S, and the following elements: hydrogen, H; cesium, Cs; and chlorine, Cl as either ionic, polar covalent or non-polar covalent. Sample Problem Page 163

47. electronegativity difference bond type S and H 2.5 - 2.1 = 0.4 polar S and Cs 2.5 - 0.7 = 1.8 ionic S and Cl 3.0 - 2.5 = 0.5 polar

48. Ionic Bonds

49. Opposites attract! An ionic compound is composed of positive and negative ions that are combined so that the charges are equal. Cations will combine with anions to form ionic compounds or salts. (metal and nonmetal) This electrostatic force that holds them together is called an ionic bond!

50. Na + + Cl - NaCl Example: sodium chloride consists of a positive ion (Na) with a +1 charge and negative ion (Cl) with a -1 charge. Draw electron configurations and Lewis Dot Structures.

51. Ionic Bonds Cation + Anion  Ionic bond (Ionic compound) Salt (NaCl)  Ionic bond between a metal and non-metal Look at the electron dot structures. Note how the sodium donates its electron to chlorine, now chlorine has an octet and a negative charge

52. Structure of Sodium Chloride Crystal

54. Ionic bonds Show the ionic bonds of the following using the electron dot structure method. 1.Magnesium and oxygen (MgO) 2. Potassium and Sulfur (K 2 S) 3. Calcium and chloride (CaCl 2 )

55. Properties of ionic compounds

56. Electrically neutral Hard, but brittle Most form crystal lattices at room temperature Generally, have high melting points and Boiling Points Fluorite (CaF 2 ) Cinnabar (HgS)

57. Properties of ionic compounds Conduct electricity when dissolved in water or as a liquid. Solids do not conduct electricity.

58. Ionic Compounds in Water When ionic compounds are placed in water, they will dissociate. This means the anions and cations will split apart and form weak bonds with the water molecules.

59. Sodium chloride vs sugar NaCl Sodium is a metal Chloride is a nonmetal Melting point ~800 o C C 12 H 22 O 11 C, H, and O are nonmetals Melting point ~185 o C Which one (or both) is/are an ionic compound(s)?

60. Crystal Lattice Ionic compounds form in repeating patterns of anions and cations. Their crystal structure will be the same for a particular compound.

61. Ionic crystals Ionic compounds that are crystals are made out of small pieces called unit cells A unit cell is the simplest repeating unit in a crystal NaCl’s unit cell:

62. Lesson check: True or False: 1.Ionic compounds have low melting points 2.Nitrogen and oxygen form an ionic compound 3.Ionic Compounds are normally liquids 4.Ionic Compounds conduct electricity when dissolved 1. False 2. False 3. False 4. True

63. Metallic Bonding

64. What are some properties of metals? Think back to Chapter 1

65. Properties of metals Good conductors of electricity Ability to be drawn into a wire (Ductile) Ability to be hammered, without breaking (Malleable) Not brittle

66. Metallic bonds Metals are composed of closely packed cations held together by their outer electrons. These outer electrons can be referred to as a sea of electrons. The electrons move freely between metal atoms. This “sea of electrons” is what allows metals to be molded, hammered, or stretched.

67. Metallic bonding Metallic Bonds are the forces or attraction between those free floating outside electrons and the positively charged metal ions

68. Metal alloys An alloy is a mixture of 2 or more elements (one must be a metal) These are uniform throughout, so a homogeneous mixture Examples: Brass (copper and zinc); Sterling silver (silver and copper); Bronze (copper and tin)

69. Why have alloys? Alloys are important because they are combining properties and are often superior compared to the pure elements Typically, more inexpensive than the pure element: Sterling silver vs pure silver $0.95 vs $1.68

70. Think about this: A bronze statue is beginning to turn green; bronze is an alloy made of copper and tin Which element is causing it to become green? Hint: Think statue of liberty

71. Covalent Molecules

72. Molecules and Molecular Compounds Ionic bonds are a + and -, but what about CO 2 ? What is type of elements are C and O? When nonmetals bond together a covalent bond is created and we call them molecules or molecular compounds!

73. Molecules Molecules are neutral atoms that are joined together by covalent bonds Molecular Compound another way a saying molecule Molecular formula shows you how many atoms of each element is in a substance Example: CO 2, NH 4

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

75. Sharing electrons Recall that ionic bonds give and take electrons… Molecules share their electrons between the 2 atoms. When they share their valence electrons, a covalent bond is made

76. Single covalent bonds When atoms share one pair of electrons they form a single covalent bond Example: H 2 Let’s draw it:

77. Diatomics are molecules Do we remember the catch phrase for diatomics? BrINClHOF

78. Show these diatomics: Cl 2 Br 2 I 2 F 2 What about H 2 O?

79. Structural Formula Electron dot structure represents bonds as 2 dots coming together: A structural formula represents covalent bonds as dashes

80. What did we call those 2 dots next to one another? Lone pairs or unshared pairs They do NOT participate in bonding, but you must show them!

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

82. Double Bonds Atoms that share two pairs of electrons Example: CO 2

83. Double Bonds Draw O 2 :

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

85. Properties of Covalent Molecules

86. Properties of covalent molecules Made out of nonmetals Can be a solid, liquid, or gas at room temperature Low melting point and boiling points Poor to non- conductors of heat and electricity

87. H 2 O vs NaCl Liquid water Solid water

88. Strengths of covalent bonds vs. ionic bonds

90. Bonding Theories

91. How do we decide where to find electrons? The modern atomic theory tells us that they are most probable at certain locations

92. Molecular Orbitals When two atoms combine, their atomic orbitals combine and overlap to produce molecular orbitals A molecular orbital belongs to the whole of the molecule Each orbital can only contain 2 electrons When an orbital overlaps and participates in a covalent bond, it can be classified as a bonding orbital

93. Sigma bonding (σ) The first bond between sharing atoms is classified as sigma bonds

94. Pi bonding (π) The second type of bonding can be a pi Remember the first is a sigma and the rest can be pi bonds

95. Molecular Geometry VSEPR Theory

96. Valence-Shell-Electron-Pair-Repulsion theory This theory helps us understand the 3D structure of molecules and their properties. Bonding and unshared pairs of valence electrons become very important to us within VSEPR theory! The shapes of molecules are determined because electron pairs want to be far apart from each other (repulsion).

97. AXE A represents the central atom X represents the bonding atoms E represents the lone pairs the central atom has – Method to represent compounds

98. Compounds with no lone pairs

99. Linear Draw or build CO 2 (AX 2 ) AX 2 Meaning 1 central atom, 2 bonded atoms It has a linear shape No lone pairs A bond angle of 180 o Bonding pairs are far apart from each other

100. Trigonal Planar Draw or build BF 3 This has a trigonal planar AX 3 Meaning 1 central, 3 bonded Bond angle: 120 o Bonds pointing to the corners of a triangle

101. Tetrahedral Draw or build CH 4 This has a tetrahedral AX 4 Meaning 1 central, 4 bonded Bond angle: 109.5 o

102. Use VSEPR theory to predict the shape of: aluminum trichloride, AlCl 3 Problem hydrogen iodide, HI carbon tetrabromide, CBr 4 dichloromethane, CH 2 Cl 2

103. Compounds with lone pairs Lone pairs occupy space, but only bonded atoms determine the name

104. Bent Draw or build H 2 O This has a bent shape AX 2 E 2 Meaning 1 central, 2 bonded, 2 lone pairs Bond angle: 105 o Similar angles to the tetrahedral bond angles

105. Trigonal pyramidal Draw or build NH 3 This has a trigonal pyramidal AX 3 E Meaning 1 central, 3 bonded, 1 lone pair Bond angle: 107 o Similar angles to the tetrahedral bond angles

106. What shape is each one? BeCl 2 OF 2 AlCl 3 PCl 3 CF 4 Linear Bent Trigonal Planar Trigonal pyramidal Tetrahedral

107. Bond Polarity Switch presentations – slide 80

108. Chemcatalyst __.__: What does it mean to share equally? What does it mean to share unequally?

109. Sharing electrons Think of tug of war… In covalent bonding, sometimes molecules share their electrons unevenly

110. Bond polarity Since, atoms are sharing within a covalent bond… If they share equally they are a nonpolar covalent bond Examples: Diatomic atoms are nonpolar because they pull on each other evenly

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

112. Bond polarity Since, polar bonds are unequally sharing we will have dipoles. But how will we decide polarity?? Electronegativity!!

113. Dipoles The more electronegative will have the arrow point towards it and have a slightly negative charge The less electronegative will have a slightly positive charge

114. Electronegativity This chart will help you determine if it is polar or nonpolar 0.0-0.4 Non polar covalent 0.4-1.0 Slightly polar covalent 1.0-2.0 Very polar covalent >/= 2.0 Ionic

115. Decide the polarity of the following: N-H F-F Ca- Cl Al- Cl 0.9 slightly polar 0 Nonpolar 2.0 ionic 1.5 very polar