1. 1 Chemical Bonds The Formation of Compounds From Atoms Chapter 11 Hein and Arena

2. 2 11.2 Lewis Structures of Atoms

3. 3 Metals form cations and nonmetals form anions to attain a stable valence electron structure.valence

4. 4 This stable structure often consists of two s and six p electrons. These rearrangements occur by losing, gaining, or sharing electrons.

5. 5 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

6. 6 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.

7. 7 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

8. 8 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.

9. 9 11.4 Lewis Structures of the first 20 elements.

10. 10 11.3 The Ionic Bond Transfer of Electrons From One Atom to Another

11. 11 The chemistry of many elements, especially the representative ones, is to attain the same outer electron structure as one of the noble gases.

12. 12 With the exception of helium, this structure consists of eight electrons in the outermost energy level.

13. 13 After sodium loses its 3s electron, it has attained the same electronic structure as neon.

14. 14 After chlorine gains a 3p electron, it has attained the same electronic structure as argon.

15. 15 Formation of NaCl

16. 16 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.

17. 17 Formation of MgCl 2

18. 18 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.

19. 19 NaCl is made up of cubic crystals.In the crystal each sodium ion is surrounded by six chloride ions.

20. 20 In the crystal each chloride ion is surrounded by six sodium ions. 11.5

21. 21 The ratio of Na + to Cl - is 1:1 There is no molecule of NaCl 11.5

22. 22 11.6 Electronegativity

23. 23 electronegativity: The relative attraction that an atom has for a pair of shared electrons in a covalent bond.

24. 24 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.

25. 25 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.

26. 26 : 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.

27. 27 A scale of relative electronegativities was developed by Linus Pauling.

28. 28 Electronegativity decreases down a group for representative elements. Electronegativity generally increases left to right across a period.

29. 29 The electronegativities of the metals are low. The electronegativities of the nonmetals are high. 11.1

30. 30 The polarity of a bond is determined by the difference in electronegativity values of the atoms forming the bond.

31. 31 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.

32. 32 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

33. 33 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

34. 34 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

35. 35 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

36. 36 A dipole is a molecule that is electrically asymmetrical, causing it to be oppositely charged at two points. A dipole can be written as + -

37. 37 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.

38. 38 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.

39. 39 11.11 Relating Bond Type to Electronegativity Difference.

40. 40 11.7 Lewis Structures of Compounds

41. 41 In writing Lewis structures, the most important consideration for forming a stable compound is that the atoms attain a noble gas configuration.

42. 42 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.

43. 43 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.

44. 44 Procedures for Writing Lewis Structures

45. 45 AtomGroupValence Electrons Cl7A7 H1A1 C4A4 N5A5 S6A6 P5A5 I7A7 Valence Electrons of Group A Elements

46. 46 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.

47. 47 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.

48. 48 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).

49. 49 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 : : ::

50. 50 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.

51. 51 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 ::

52. 52 Step 4. Distribute pairs of electrons (pairs of dots) around each atom (except hydrogen) to give each atom a noble gas configuration.

53. 53 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.

54. 54 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.

55. 55 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.

56. 56 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 ::

57. 57 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

58. 58 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

59. 59 11.8 Complex Lewis Structures

60. 60 There are some molecules and polyatomic ions for which no single Lewis structure consistent with all characteristics and bonding information can be written.

61. 61 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.

62. 62 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 :

63. 63 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.

64. 64 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

65. 65 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. -

66. 66 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

67. 67 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 : : : : -

68. 68 11.9 Compounds Containing Polyatomic Ions

69. 69 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.

70. 70 Sodium nitrate, NaNO 3, contains one sodium ion and one nitrate ion. sodium ion Na + nitrate ion N O : : O : : : O : : : : - Na +

71. 71 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.

72. 72 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.

73. 73 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.

74. 74 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.

75. 75 11.10 Molecular Shape

76. 76 The 3-dimensional arrangement of the atoms within a molecule is a significant feature in understanding molecular interactions.

77. 77 11.12

78. 78 11.11 The Valence Shell Electron Pair (VSEPR) Model

79. 79 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.

80. 80 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.

81. 81 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.

82. 82 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.

83. 83 Ball and stick models of methane, CH 4, and carbon tetrachloride, CCl 4. 11.13

84. 84 Ammonia, NH 3, has four electron pairs around nitrogen. The arrangement of the electron pairs is tetrahedral.

85. 85 The shape of the NH 3 molecule is pyramidal. One of its electron pairs is a nonbonded (lone) pair.

86. 86 Water has four electron pairs around oxygen. The arrangement of electron pairs around oxygen is tetrahedral.

87. 87 The H 2 O molecule is bent. Two of its electron pairs are nonbonded (lone) pairs.

88. 88 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.

89. 89