1. Chemical Bonding Chapter 6

2. Chemical bond A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds them together

3. Types of bonds Ionic electrons are transferred from one atom to another, forming ions Electrical attraction of large numbers of cations and anions holds the compound together Covalent – electron pairs are shared between the atoms

4. Ionic Bonds Ions get very close together and are attracted to each other. Example: NaCl Na + Cl -

5. Covalent Bonds The atoms share electrons, and their electron clouds overlap. Example: carbon monoxide C O

6. Electronegativity difference Can be used to predict how two atoms will react with each other. Has no units, because it is a comparison. When it is high, the bond is ionic. When it is low, the bond is covalent.

7. Ionic – covalent continuum Real bonds are usually sort of ionic and sort of covalent. See figure 6-2 on page 162. Difference of 1.7 or less is considered covalent 50% ionic character or less Difference of more than 1.7 is considered ionic More than 50% ionic character

8. Nonpolar-covalent bond Bonding electrons are shared equally Balanced distribution of electrical charge 0 to 5% ionic character 0 to 0.3 electronegativity difference

9. Polar Having uneven charge distribution

10. Polar-covalent bond The bonded atoms have unequal attraction for the shared electrons 5 to 50% ionic character 0.3 to 1.7 electronegativity difference

11. Polar and nonpolar bonds Electron density is greater around Cl than H because it has greater electronegativity  + and  - indicate partial charge Electron cloud drawings nonpolar polar  

12. Comparing bond types Comparison between ionic and covalent bondingComparison between ionic and covalent bonding

13. Discuss Page 163 Sample problem Practice problem Section review 1 and 2

14. Molecule A neutral group of atoms that are held together by covalent bonds A single molecule is capable of existing on its own May consist of atoms of the same element or different elements

15. Molecular compound Compound whose simplest units are molecules

16. Chemical formula Uses atomic symbols and subscripts to indicate the relative numbers of atoms of each kind in a chemical compound.

17. Molecular formula Shows the types and numbers of atoms in a single molecule of a molecular compound H 2 O O 2 C 12 H 22 O 11

18. Diatomic molecule Contains only two atoms

19. Bonding Bonding makes atoms be at lower potential energy levels, which is favored by nature. Attractive forces (nucleus-electron) are balanced by repulsive forces (nucleus-nucleus and electron-electron).

20. Bond length Average distance between bonded nuclei at their lowest potential energy

21. Bond energy The amount of energy that must be added to break a chemical bond and form neutral isolated atoms. The same amount of energy that the atoms had to release when then bonded.

22. 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 The outermost s and p sublevels want to be full

23. Exceptions to octet rule Hydrogen only needs two electrons, not eight Boron tends to only be surrounded by six electrons. Some atoms can have more than eight when combining with fluorine, oxygen, and chlorine. This expanded valence also includes the d orbitals.

24. Electron dot diagrams Usually only the valence electrons are involved in chemical reactions. Electron dot diagrams allow us to draw these electrons around the symbol for an element.

25. Drawing electron dot diagrams 1.The element’s symbol represents the nucleus and all the electrons in the inner energy levels. 2.Write the noble gas configuration of the element. Add the superscripts of the s and p orbitals to get the number of valence electrons.

26. 3.Draw dots on the sides to represent electrons.

27. Examples

28. Molecules Can be represented by dot diagrams Example: HCl Shared pairs can be replaced by long dashes

29. Lone pair Unshared pair of electrons Not involved in bonding

30. Discuss Read sample problem 6-2 on page 170

31. Lewis diagram Formulas in which atomic symbols represent the nuclei and inner-shell electrons Dot diagrams with dots or dashes for shared electrons

32. Structural diagrams Leaves out the unshared pairs Indicates the kind, number, arrangement, and bonds of atoms

33. Single bond Covalent bond produced when one pair of electrons is shared

34. Discuss Sample problem 6-3 on pages 171-172

35. Double bond Covalent bond produced when two pairs of electrons are shared Example: CO 2

36. Triple bond Covalent bond formed when three pairs of electrons are shared Example: N 2

37. Multiple bonds Double or triple bonds Have higher bond energies Have shorter bond lengths

38. Resonance structures Bonding in molecules or ions that cannot be correctly represented by a Lewis structure A double-headed arrow is placed between the two possibilities Example: ozone

39. Discuss Covalent Bonds More Covalent Bonds Sample problem 6-4 on page 174

40. Examples Draw a Lewis Structure and the structural formula for each of the following: H 2 O CH 4 CH 4 O O 2 C 2 H 2 C 2 H 4

41. Discuss Section Review on page 175

42. Ionic compound Composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal Usually crystalline solids Not composed of independent units like molecules Ionic bonding More ionic bonding

43. Chemical formula For ionic compounds Shows the simplest ratio of the ions that gives no net charge

44. Formula unit The simplest collection of atoms from which an ionic compound’s formula can be established. For sodium chloride, NaCl The simplest way to show a one to one ratio

45. Ratio of ions Depends on the charges of the ions combined. Na + and Cl - form NaCl Mg 2+ and F - form MgF 2

46. Forming ionic compounds Examples: NaCl MgF 2

47. Crystal lattice Ions minimize their potential energy by forming an orderly arrangement

48. Lattice energy The energy released when one mole of an ionic crystalline compound is formed from gaseous ions. Used to compare bond strengths

49. Molecular vs. ionic compounds Forces between molecules are weaker than ionic bonding forces Molecular compounds have lower melting points Ionic compounds are hard but brittle. Ionic compounds conduct electricity when melted or dissolved in solution

50. Polyatomic ions A charged group of covalently bonded atoms. Form ionic bonds just like other ions. Examples Ammonium ion NH 4 + Sulfate ion SO 4 2-

51. Discuss Use electron-dot notation to demonstrate the formation of ionic compounds involving Li and Cl Ca and I What basic unit are molecular compounds composed of? Ionic compounds?

52. Metals High electrical conductivity. Metals form crystals and their orbitals overlap. All the atoms have the same attraction for their electrons, so the electrons can move easily from one atom to another.

53. Valence electrons Metals have very few Their p sublevels are empty Transition metals also have many vacant d sublevels

54. Metallic bonding Delocalized electrons – don’t belong to one specific atom Sea of electrons formed around the metal atoms Metallic bonding results from the attraction between metal atoms and the surrounding sea of electrons

55. Metallic properties High electrical and thermal conductivity Ability to absorb a wide range of light frequencies Electrons enter excited states and then return to ground states This makes metals look shiny

56. Metallic properties Metals bond the same way in all directions Their layers can slide past each other without breaking Malleabiltiy - the ability to be hammered or beaten into thin sheets Ductility – the ability to be drawn, pulled or extruded into a thin wire

57. Metallic bond strength When there is a bigger nucleus and more electrons, the bond is stronger. Stronger bonds have higher heats of vaporization

58. Discuss Describe the electron-sea model of metallic bonding. What is the relationship between metallic bond strength and heat of vaporization? Explain why most metals are malleable and ductile but ionic crystals are not.

59. Molecular polarity Uneven distribution of molecular charge Determined by the polarity of each bond and the geometry of the molecule

60. Molecular geometry There are two equally successful theories One accounts for molecular bond angles The other describes the orbitals that contain the valence electrons

61. VSEPR theory Valence-shell, electron-pair repulsion The repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible To predict geometry, one must consider the location of all electron pairs around the bonded atoms.

62. Linear molecules A is the central atom B is an atom bonded to it Formulas like AB 2 form linear molecules, with the B atoms on opposite sides of the A atom

63. Trigonal Planar AB 3 A is in the center. The 3 Bs make a triangle around it. The molecule is flat (planar)

64. Tetrahedral AB 4 A pyramid with A at the center and a B at each corner

65. When there are lone pairs The repel the other electrons just like shared pairs do Use an E to represent the unshared pairs

66. Bent or angular AB 2 E Like trigonal planar, but without the top of the triangle

67. Trigonal pyramidal AB 3 E Like tetrahedral, but without the top of the pyramid A is at the center, the Bs are the base, and E is at the top

68. Bent or angular AB 2 E 2 Like a tetrahedral, but only the atoms (not the unshared pairs) determine the shape

69. Trigonal bipyramidal AB 5 Like two pyramids on top of each other

70. Octahedral AB 6 8 sided shape

71. VSEPR Double and triple bonds are treated the same way as single bonds. Use Lewis structures and table 6-5 on page 186 to predict molecule shapes

72. Hybridization Explains how orbitals get rearranged when atoms form covalent bonds Describes how atoms are bonded

73. Hybrid orbitals Orbitals of equal energy produced by the combining of two or more orbitals on the same atom The number of hybrid orbitals equals the number of orbitals they were made from

74. Example: Methane CH 4 To get 4 equal orbitals, the s and p orbitals combine. They form 4 sp 3 orbitals

75. Other hybrid orbitals

76. Intermolecular forces The forces of attraction between molecules Usually weaker than covalent, ionic, and metallic bonds Can be measured by boiling point Higher boiling point means stronger intermolecular forces

77. Dipole Equal but opposite charges separated by a short distance Direction is from the positive pole to the negative pole

78. Dipole-dipole forces Forces of attraction between polar molecules Strongest intermolecular forces Only act over short distances

79. More than two atoms Polarity of the molecule depends on the polarity and orientation of each bond

81. Induced dipoles A polar molecule can induce a dipole in a nonpolar molecule by attracting its electrons. It is temporary

82. Hydrogen bonding A hydrogen atom in a compound with a very electronegative atom is attracted to an unshared pair of electrons in a nearby molecule. Examples: Water (H 2 O) Hydrogen fluoride (HF) Ammonia (NH 3 )

83. Hydrogen bonding Causes higher boiling points

84. Polarity and Hydrogen Bonding

85. London dispersion forces The intermolecular attractions caused by the constant motion of electrons and the creation of instantaneous dipoles Act between all atoms and molecules The only intermolecular forces among noble gas atoms and nonpolar molecules Low boiling points

86. discuss What is the difference between intramolecular forces and intermolecular forces? What is a dipole? In what kind of molecule are dipoles common? What is hydrogen bonding? What are London dispersion forces? Molecular Geometry review