1. 1 Chemical Bonds The Formation of Compounds From Atoms Chapter 11 Hein and Arena Eugene Passer Chemistry Department Bronx Community College © John Wiley and Sons, Inc Version 2.0 12 th Edition

2. 2 Chapter Outline 11.1 Periodic Trends in Atomic PropertiesPeriodic Trends inAtomic Properties 11.2 Lewis Structures of AtomsLewis Structures of Atoms 11.3 The Ionic Bond: Transfer of Electrons From One Atom to AnotherThe Ionic Bond: Transfer ofElectrons From One Atom to Another 11.7 Lewis Structures of CompoundsLewis Structures ofCompounds 11.8 Complex Lewis StructuresComplex Lewis Structures 11.4 Predicting Formulas of Ionic CompoundsPredicting Formulas ofIonic Compounds 11.5 The Covalent Bond: Sharing ElectronsThe Covalent Bond:Sharing Electrons 11.6 ElectronegativityElectronegativity 11.11 The Valence Shell Electron Pair (VSEPR) ModelThe Valence ShellElectron Pair (VSEPR)Model 11.9 Compounds Containing Polyatomic IonsCompounds ContainingPolyatomic Ions 11.10 Molecular ShapeMolecular Shape

3. 3 11.1 Periodic Trends in Atomic Properties

4. 4 Characteristic properties and trends of the elements are the basis of the periodic table’s design.

5. 5 These trends allow us to use the periodic table to accurately predict properties and reactions of a wide variety of substances.

6. 6 Metals and Nonmetals

7. 7 Chemical Properties of Metals metals tend to lose electrons and form positive ions. nonmetals tend to gain electrons and form negative ions. Chemical Properties of Nonmetals When metals react with nonmetals, electrons are usually transferred from the metal to the nonmetal.

8. 8 Physical Properties of Metals lustrous malleable good conductors of heat good conductors of electricity nonlustrous brittle poor conductors of heat poor conductors of electricity Physical Properties of Nonmetals

9. 9 Metalloids have properties that are intermediate between metals and nonmetals.

10. 10 The Metalloids 1.boron 2.silicon 3.germanium 4.arsenic 5.antimony 6.tellurium 7.polonium

11. 11 Metals are found to the left of the metalloids Nonmetals are found to the right of the metalloids. 11.1

12. 12 Atomic Radius

13. 13 Atomic radii increase down a group. For each step down a group, electrons enter the next higher energy level. 11.2

14. 14 Radii of atoms tend to decrease from left to right across a period. For representative elements within the same period, the energy level remains constant as electrons are added. This increase in positive nuclear charge pulls all electrons closer to the nucleus. 11.2 Each time an electron is added, a proton is also added to the nucleus.

15. 15 Ionization Energy

16. 16 The ionization energy of an atom is the energy required to remove an electron from an atom. Na + ionization energy → Na + + e -

17. 17 The first ionization energy is the amount of energy required to remove the first electron from an atom. He + first ionization energy → He + + e - He + 2,372 kJ/mol → He + + e - The second ionization energy is the amount of energy required to remove the second electron from an atom. He + + 5,247 kJ/mol → He 2+ + e - He + + second ionization energy → He 2+ + e -

18. 18 The first ionization energy is the amount of energy required to remove the first electron from an atom. He + first ionization energy → He + + e - He + 2,372 kJ/mol → He + + e - The second ionization energy is the amount of energy required to remove the second electron from an atom. He+ + 5,247 kJ/mol → He ++ + e - He + + second ionization energy → He ++ + e -

19. 19 As each succeeding electron is removed from an atom, ever higher energies are required.

20. 20 Periodic relationship of the first ionization energy for representative elements in the first four periods. 11.3 Ionization energies gradually increase from left to right across a period. IA IIA IIIA IVA VA VIA VIIA Noble Gases 1 2 3 4

21. 21 Periodic relationship of the first ionization energy for representative elements in the first four periods. 11.3 Ionization energies of Group A elements decrease from top to bottom in a group. IA IIA IIIA IVA VA VIA VIIA Noble Gases Distance of Outer Shell Electrons From Nucleus nonmetalsmetals nonmetals have higher ionization potentials than metals

22. 22 11.2 Lewis Structures of Atoms

23. 23 Metals form cations and nonmetals form anions to attain a stable valence electron structure.valence

24. 24 This stable structure often consists of two s and six p electrons. These rearrangements occur by losing, gaining, or sharing electrons.

25. 25 Na with the electron structure 1s 2 2s 2 2p 6 3s 1 has 1 valence electron. The Lewis structure of an atom is a representation that shows the valence electrons for that atom.valence Fluorine with the electron structure 1s 2 2s 2 2p 5 has 7 valence electrons

26. 26 valence electrons: the electrons that occupy the outermost energy level of an atom. valence electrons are responsible for the electron activity that occurs to form chemical bonds.

27. 27 The Lewis structure of an atom uses dots to show the valence electrons of atoms. The number of dots equals the number of s and p electrons in the atom’s outermost shell. B Paired electrons Unpaired electron Symbol of the element 2s22p12s22p1

28. 28 The number of dots equals the number of s and p electrons in the atom’s outermost shell. S 2s22p42s22p4 The Lewis structure of an atom uses dots to show the valence electrons of atoms.

29. 29 11.4 Lewis Structures of the first 20 elements.

30. 30 11.3 The Ionic Bond Transfer of Electrons From One Atom to Another

31. 31 The chemistry of many elements, especially the representative ones, is to attain the same outer electron structure as one of the noble gases.

32. 32 With the exception of helium, this structure consists of eight electrons in the outermost energy level.

33. 33 After sodium loses its 3s electron, it has attained the same electronic structure as neon.

34. 34 After chlorine gains a 3p electron, it has attained the same electronic structure as argon.

35. 35 Formation of NaCl

36. 36 The 3s electron of sodium transfers to the 3p orbital of chlorine. Lewis representation of sodium chloride formation. A sodium ion (Na+) and a chloride ion (Cl - ) are formed. The force holding Na + and Cl - together is an ionic bond.

37. 37 Formation of MgCl 2

38. 38 Two 3s electrons of magnesium transfer to the half-filled 3p orbitals of two chlorine atoms. A magnesium ion (Mg 2+ ) and two chloride ions (Cl - ) are formed. The forces holding Mg 2+ and two Cl - together are ionic bonds.

39. 39 NaCl is made up of cubic crystals.In the crystal each sodium ion is surrounded by six chloride ions.

40. 40 In the crystal each chloride ion is surrounded by six sodium ions. 11.5

41. 41 The ratio of Na + to Cl - is 1:1 There is no molecule of NaCl 11.5

42. 42 Relative Size of Sodium Ion to Chloride Ion

43. 43 A sodium ion is smaller than a sodium atom because: (1) the sodium atom has lost its outermost electron. (2) The 10 remaining electrons are now attracted by 11 protons and are drawn closer to the nucleus. 11.6

44. 44 A chloride ion is larger than a chlorine atom because: (1) the chlorine atom has gained an electron and now has 18 electrons and 17 protons. (2) The nuclear attraction on each electron has decreased, allowing the chlorine to expand. 11.6

45. 45 Metals usually have one, two, or three electrons in their outer shells. When a metal reacts it: –usually loses one, two, or three electrons –attains the electron structure of a noble gas –becomes a positive ion. The positive ion formed by the loss of electrons is much smaller than the metal atom.

46. 46 Nonmetals usually have one, two or three electrons in their outer shells. When a nonmetal reacts it: –usually gains one, two, or three electrons –attains the electron structure of a noble gas –becomes a negative ion. The negative ion formed by the gain of electrons is much larger than the nonmetal atom.

47. 47

48. 48 11.4 Predicting Formulas of Ionic Compounds

49. 49 In almost all stable chemical compounds of representative elements, each atom attains a noble gas electron configuration.

50. 50 Metals will lose electrons to attain a noble gas configuration. Nonmetals will gain electrons to attain a noble gas configuration. Barium and Sulfur Combine. –sulfur gains two electrons from barium and attains an argon configuration. –barium loses two electrons to sulfur and attains a xenon configuration. S [Ne]3s 2 3p 4 Ba [Xe]6s 2 Ba → [Xe] + 2e - S + 2e - → [Ar] Ba + S → BaS

51. 51 Because of similar electron structures, the elements of a family generally form compounds with the same atomic ratios.

52. 52

53. 53 10.17 The group numbers for the representative elements are equal to the total number of outermost electrons in the atoms of the group. The elements of a family have the same outermost electron configuration except that the electrons are in different energy levels.

54. 54 The atomic ratio of the alkali metal sodium to chlorine is 1:1 in NaCl. The atomic ratios of the other alkali metal chlorides can be predicted to also be 1:1. LiCl, KCl, CsCl, FrCl

55. 55 The atomic ratio of hydrogen to nitrogen is 3:1 in ammonia (NH 3 ). Nitrogen is the first member of group 5A. The atomic ratio of hydrogen when combined with other group 5A elements can be predicted to also be 3:1. PH 3, AsH 3, SbH 3, BiH 3

56. 56 11.5 The Covalent Bond: Sharing Electrons

57. 57 A covalent bond consists of a pair of electrons shared between two atoms. In the millions of chemical compounds that exist, the covalent bond is the predominant chemical bond.

58. 58 Substances which covalently bond exist as molecules. Carbon dioxide bonds covalently. It exists as individually bonded covalent molecules containing one carbon and two oxygen atoms. 11.7

59. 59 The term molecule is not used when referring to ionic substances. Sodium chloride bonds ionically. It consists of a large aggregate of positive and negative ions. No molecules of NaCl exist. 11.7

60. 60 Covalent bonding in the hydrogen molecule Two 1s orbitals from each of two hydrogen atoms overlap. Each 1s orbital contains 1 electron. The orbital of the electrons includes both hydrogen nuclei. The most likely region to find the two electrons is between the two nuclei. The two nuclei are shielded from each other by the electron pair. This allows the two nuclei to draw close together. Two 1s orbitals from each of two hydrogen atoms overlap. 11.8

61. 61 11.9 Covalent bonding in the chlorine molecule Each unpaired 3p orbital on each chlorine atom contains 1 electron. Two 3p orbitals from each of two chlorine atoms overlap. The orbital of the electrons includes both chlorine nuclei. The most likely region to find the two electrons is between the two nuclei. The two nuclei are shielded from each other by the electron pair. This allows the two nuclei to draw close together. Two 3p orbitals from each of two chlorine atoms overlap. Each chlorine now has 8 electrons in its outermost energy level.

62. 62 hydrogen chlorine iodine nitrogen Covalent bonding with equal sharing of electrons occurs in diatomic molecules formed from one element. A dash may replace a pair of dots.

63. 63 11.6 Electronegativity

64. 64 electronegativity: The relative attraction that an atom has for a pair of shared electrons in a covalent bond.

65. 65 If the two atoms that constitute a covalent bond are identical, then there is equal sharing of electrons. This is called nonpolar covalent bonding. Ionic bonding and nonpolar covalent bonding represent two extremes.

66. 66 If the two atoms that constitute a covalent bond are not identical, then there is unequal sharing of electrons. This is called polar covalent bonding. One atom assumes a partial positive charge and the other atom assumes a partial negative charge. –This charge difference is a result of the unequal attractions the atoms have for their shared electron pair.

67. 67 : HCl ++ -- Shared electron pair. : The shared electron pair is closer to chlorine than to hydrogen. Partial positive charge on hydrogen. Partial negative charge on chlorine. Chlorine has a greater attraction for the shared electron pair than hydrogen. Polar Covalent Bonding in HCl The attractive force that an atom of an element has for shared electrons in a molecule or a polyatomic ion is known as its electronegativity.

68. 68 A scale of relative electronegativities was developed by Linus Pauling.

69. 69 Electronegativity decreases down a group for representative elements. Electronegativity generally increases left to right across a period.

70. 70 The electronegativities of the metals are low. The electronegativities of the nonmetals are high. 11.1

71. 71 The polarity of a bond is determined by the difference in electronegativity values of the atoms forming the bond.

72. 72 If the electronegativity difference between two bonded atoms is greater than 1.7-1.9, the bond will be more ionic than covalent. If the electronegativity difference is greater than 2, the bond is strongly ionic. If the electronegativity difference is less than 1.5, the bond is strongly covalent.

73. 73 HH Hydrogen Molecule If the electronegativities are the same, the bond is nonpolar covalent and the electrons are shared equally. The molecule is nonpolar covalent. Electronegativity 2.1 Electronegativity 2.1 11.10 Electronegativity Difference = 0.0

74. 74 If the electronegativities are the same, the bond is nonpolar covalent and the electrons are shared equally. Cl Chlorine Molecule Electronegativity 3.0 Electronegativity 3.0 The molecule is nonpolar covalent. Electronegativity Difference = 0.0 11.10

75. 75 If the electronegativities are not the same, the bond is polar covalent and the electrons are shared unequally. HCl Hydrogen Chloride Molecule Electronegativity 2.1 Electronegativity 3.0 The molecule is polar covalent. ++ -- Electronegativity Difference = 0.9 11.10

76. 76 Sodium Chloride Na + Cl - If the electronegativities are very different, the bond is ionic and the electrons are transferred to the more electronegative atom. Electronegativity 0.9 Electronegativity 3.0 The bond is ionic.No molecule exists. Electronegativity Difference = 2.1 11.10

77. 77 A dipole is a molecule that is electrically asymmetrical, causing it to be oppositely charged at two points. A dipole can be written as + -

78. 78 An arrow can be used to indicate a dipole. The arrow points to the negative end of the dipole. HClHBrH O H Molecules of HCl, HBr and H 2 O are polar.

79. 79 A molecule containing different kinds of atoms may or may not be polar depending on its shape. The carbon dioxide molecule is nonpolar because its carbon-oxygen dipoles cancel each other by acting in opposite directions.

80. 80 11.11 Relating Bond Type to Electronegativity Difference.

81. 81 11.7 Lewis Structures of Compounds

82. 82 In writing Lewis structures, the most important consideration for forming a stable compound is that the atoms attain a noble gas configuration.

83. 83 The most difficult part of writing Lewis structures is determining the arrangement of the atoms in a molecule or an ion. In simple molecules with more than two atoms, one atom will be the central atom surrounded by the other atoms.

84. 84 Cl 2 O has two possible arrangements. Cl-Cl-O The two chlorines can be bonded to each other. Cl-O-Cl The two chlorines can be bonded to oxygen. Usually the single atom will be the central atom.

85. 85 Procedures for Writing Lewis Structures

86. 86 AtomGroupValence Electrons Cl7A7 H1A1 C4A4 N5A5 S6A6 P5A5 I7A7 Valence Electrons of Group A Elements

87. 87 Step 1. Obtain the total number of valence electrons to be used in the structure by adding the number of valence electrons in all the atoms in the molecule or ion. –If you are writing the structure of an ion, add one electron for each negative charge or subtract one electron for each positive charge on the ion.

88. 88 Step 1. The total number of valence electrons is eight, two from the two hydrogen atoms and six from the oxygen atom. Write the Lewis structure for H 2 O.

89. 89 Step 2. Write the skeletal arrangement of the atoms and connect them with a single covalent bond (two dots or one dash). –Hydrogen, which contains only one bonding electron, can form only one covalent bond. –Oxygen atoms normally have a maximum of two covalent bonds (two single bonds, or one double bond).

90. 90 Step 2. The two hydrogen atoms are connected to the oxygen atom. Write the skeletal structure: Write the Lewis structure for H 2 O. Place two dots between the hydrogen and oxygen atoms to form the covalent bonds. H O H or H O H : : ::

91. 91 Step 3. Subtract two electrons for each single bond you used in Step 2 from the total number of electrons calculated in Step 1. –This gives you the net number of electrons available for completing the structure.

92. 92 Step 3. Subtract the four electrons used in Step 2 from eight to obtain four electrons yet to be used. Write the Lewis structure for H 2 O. H O H ::

93. 93 Step 4. Distribute pairs of electrons (pairs of dots) around each atom (except hydrogen) to give each atom a noble gas configuration.

94. 94 Step 4. Distribute the four remaining electrons in pairs around the oxygen atom. Hydrogen atoms cannot accommodate any more electrons. Write the Lewis structure for H 2 O. These arrangements are Lewis structures because each atom has a noble gas electron structure. H O H or H O H : : :: : : : : The shape of the molecule is not shown by the Lewis structure.

95. 95 Step 1. The total number of valence electrons is 16, four from the C atom and six from each O atom. Write a Lewis structure for CO 2.

96. 96 Step 2. The two O atoms are bonded to a central C atom. Write the skeletal structure and place two electrons between the C and each oxygen. O C O :: Write a Lewis structure for CO 2.

97. 97 Write a Lewis structure for CO 2. Step 3. Subtract the four electrons used in Step 2 from 16 (the total number of valence electrons) to obtain 12 electrons yet to be used. O C O ::

98. 98 O C O :: Step 4. Distribute the 12 electrons (6 pairs) around the carbon and oxygen atoms. Three possibilities exist. Many of the atoms in these structures do not have eight electrons around them. Write a Lewis structure for CO 2. O C O :: :: :: : : : : : : : : : : 4 electrons 6 electrons 6 electrons 6 electrons : : : : : : 6 electrons IIIIII

99. 99 Write a Lewis structure for CO 2. O C O :::: : : : : Step 5. Remove one pair of unbonded electrons from each O atom in structure I and place one pair between each O and the C atom forming two double bonds. O C O : : :: : : : : : : :: : : : : Each atom now has 8 electrons around it. Carbon is sharing 4 electron pairs. double bond

100. 100 11.8 Complex Lewis Structures

101. 101 There are some molecules and polyatomic ions for which no single Lewis structure consistent with all characteristics and bonding information can be written.

102. 102 Step 1. The total number of valence electrons is 24, 5 from the nitrogen atom and 6 from each O atom, and 1 from the –1 charge. Write a Lewis structure for NO 2.

103. 103 Step 2. The three O atoms are bonded to a central N atom. Write the skeletal structure and place two electrons between each pair of atoms. Write a Lewis structure for NO 2. O N O :: O :

104. 104 Step 3. Subtract the 6 electrons used in Step 2 from 24, the total number of valence electrons, to obtain 18 electrons yet to be placed. O N O :: O : Write a Lewis structure for NO 2.

105. 105 O N O O Step 4. Distribute the 18 electrons around the N and O atoms. Write a Lewis structure for NO 2. : ::: :: : : : :: : electron deficient

106. 106 O : ::: :: : : : :: : O N O Step 4. Since the extra electron present results in nitrate having a –1 charge, the ion is enclosed in brackets with a – charge. Write a Lewis structure for NO 2. -

107. 107 Write a Lewis structure for NO 2. Step 5. One of the oxygen atoms has only 6 electrons. It does not have a noble gas structure. Move the unbonded pair of electrons from the N atom and place it between the N and the electron-deficient O atom, making a double bond. N O : : O : : : O : : : : - : : electron deficient

108. 108 A molecule or ion that shows multiple correct Lewis structures exhibits resonance. Write a Lewis structure for NO 2. Step 5. There are three possible Lewis structures. N O : : O : : : O : : : : - : N O : : O : : O : : : : - Each Lewis structure is called a resonance structure. N O : : O : : : O : : : : -

109. 109 11.9 Compounds Containing Polyatomic Ions

110. 110 A polyatomic ion is a stable group of atoms that has either a positive or negative charge and behaves as a single unit in many chemical reactions.

111. 111 Sodium nitrate, NaNO 3, contains one sodium ion and one nitrate ion. sodium ion Na + nitrate ion N O : : O : : : O : : : : - Na +

112. 112 The nitrate ion is a polyatomic ion composed of one nitrogen atom and three oxygen atoms. N O : : O : : : O : : : : - Na + It has a charge of –1 One nitrogen and three oxygen atoms have a total of 23 valence electrons.

113. 113 The –1 charge on nitrate adds an additional valence electron for a total of 24. N O : : O : : : O : : : : - Na + The additional valence electron comes from a sodium atom which becomes a sodium ion.

114. 114 Sodium nitrate has both ionic and covalent bonds. N O : : O : : : O : : : : - Na + Ionic bonds exist between the sodium ions and the carbonate ions. covalent bond ionic bond Covalent bonds are present between the carbon and oxygen atoms within the carbonate ion.

115. 115 When sodium nitrate is dissolved in water the ionic bond breaks. N O : : O : : : O : : : : - Na + The sodium ions and nitrate ions separate from each other forming separate sodium and nitrate ions. N O : : O : : : O : : : : - Na + The nitrate ion, which is held together by covalent bonds, remains as a unit.

116. 116 11.10 Molecular Shape

117. 117 The 3-dimensional arrangement of the atoms within a molecule is a significant feature in understanding molecular interactions.

118. 118 11.12

119. 119 11.11 The Valence Shell Electron Pair (VSEPR) Model

120. 120 The VSEPR model is based on the idea that electron pairs will repel each other electrically and will seek to minimize this repulsion. To accomplish this minimization, the electron pairs will be arranged as far apart as possible around a central atom.

121. 121 BeCl 2 is a molecule with only two pairs of electrons around beryllium, its central atom. Its electrons are arranged 180 o apart for maximum separation.

122. 122 BF 3 is a molecule with three pairs of electrons around boron, its central atom. Its electrons are arranged 120 o apart for maximum separation. This arrangement of atoms is called trigonal planar.

123. 123 CH 4 is a molecule with four pairs of electrons around carbon, its central atom. An obvious choice for its atomic arrangement is a 90 o angle between its atoms with all of its atoms in a single plane. However, since the molecule is 3-dimensional, the molecular structure is tetrahedral with a bond angle of 109.5 o.

124. 124 Ball and stick models of methane, CH 4, and carbon tetrachloride, CCl 4. 11.13

125. 125 Ammonia, NH 3, has four electron pairs around nitrogen. The arrangement of the electron pairs is tetrahedral.

126. 126 The shape of the NH 3 molecule is pyramidal. One of its electron pairs is a nonbonded (lone) pair.

127. 127 Water has four electron pairs around oxygen. The arrangement of electron pairs around oxygen is tetrahedral.

128. 128 The H 2 O molecule is bent. Two of its electron pairs are nonbonded (lone) pairs.

129. 129 Structure Determination Using VSEPR 1.Draw the Lewis structure for the molecule. 2.Count the electron pairs and arrange them to minimize repulsions. 3.Determine the positions of the atoms. 4.Name the molecular structure from the position of the atoms.

130. 130