1. Unit 8: Chemical Bonding and Molecular Structure Cartoon courtesy of NearingZero.net

2. IQ 1 1)What are valence electrons? 2)Why are they important in bonding? 3)Explain the difference between a cation and an anion. 4)What is electronegativity and what is the periodic trend for it? 5)How can we determine the # of valence electrons by looking at the periodic table? 6)What is an octet? What group of elements have one?

3. Chemical Bonding Problems and questions — How is a molecule or polyatomic ion held together? Why are atoms distributed at strange angles? Why are molecules not flat? Can we predict the structure? How is structure related to chemical and physical properties?

4. Types of Bonds Forces that hold groups of atoms together and make them function as a unit. o Forces that hold groups of atoms together and make them function as a unit. –bonds form in order to… decrease potential energy (PE) increase stability Ionic bonds –  Ionic bonds – transfer of electrons.  Covalent bonds  Covalent bonds – sharing of electrons. 2 types 2 types 1.Non-Polar Covalent 2.Polar Covalent

5. IONICCOVALENT Bond Formation Type of Structure Solubility in Water Electrical Conductivity Other Properties e - are transferred from metal to nonmetal high yes (solution or liquid) yes e - are shared between two nonmetals low no usually not Melting Point crystal lattice true molecules Types of Bonds Physical State solid liquid or gas odorous

6. The type of bond can usually be calculated by finding the difference in electronegativity of the two atoms that are going together. Types of Bond cont.

7. Bond Polarity Most bonds are a blend of ionic and covalent characteristics. Difference in electronegativity determines bond type.

8. Electronegativity Difference If the difference in electronegativities is between: – 2.0 to 4.0: Ionic – 0.4 to 2.0: Polar Covalent – 0.0 to 0.4: Non-Polar Covalent Example: NaCl Na = 0.8, Cl = 3.0 Difference is 2.2, so this is an ionic bond!

9. Ionic Bonding: The Octet Rule Ionic compounds tend to form so that each atom, by gaining or losing electrons, has an octet of electrons in its highest occupied energy level.

10. Ionic Bonding: Compounds Aluminum metal and the nonmetal bromine react to form an ionic solid, aluminum bromide. The electrostatic forces that hold ions together in ionic compounds are called ionic bonds. 7.2

11. Formation of Ionic Compounds –Animation 8 Take an atomic-level look at the formation of KCl. ChemASAP/dswmedia/rsc/asap1_chem05_cman0708.html

12. Ionic Bonding: The Formation of Sodium Chloride  Sodium has 1 valence electron Cl: 1s 2 2s 2 2p 6 3s 2 3p 5 Na: 1s 2 2s 2 2p 6 3s 1  Chlorine has 7 valence electrons  An electron transferred gives each an octet

13. Ionic Bonding: The Formation of Sodium Chloride Cl - 1s 2 2s 2 2p 6 3s 2 3p 6 Na + 1s 2 2s 2 2p 6 This transfer forms ions, each with an octet:

14. Ionic Bonding: The Formation of Sodium Chloride Cl - Na + The resulting ions come together due to electrostatic attraction (opposites attract): The net charge on the compound must equal zero!!! Cl - Na +

15. Formation of Ionic Compounds NaCl is the chemical formula for sodium chloride. 7.2

16. Examples of Ionic compounds All salts, which are composed of metals bonded to nonmetals, are ionic compounds and form ionic crystals. Examples (balance these charges): Ba S Mg Cl Na O K I Ca O Li F

17. Monatomic Cations Name H+H+ Hydrogen Li + Lithium Na + Sodium K+K+ Potassium Mg 2+ Magnesium Ca 2+ Calcium Ba 2+ Barium Al 3+ Aluminum

18. Monatomic Anions Name F-F- Fluoride Cl - Chloride Br - Bromide I-I- Iodide O 2- Oxide S 2- Sulfide N 3- Nitride P 3- Phosphide

19. Properties of Ionic Compounds Structure: Crystalline solids Melting point: Generally high Boiling Point: Generally high Electrical Conductivity: Excellent conductors, molten and aqueous Solubility in water: Generally soluble

20. The Crystal Lattice Ionic compounds form solids at normal temperatures. Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions.

21. Representation of Components in an Ionic Solid Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance.

22. Properties of Ionic Compounds The orderly arrangement of component ions produces the beauty of crystalline solids. 7.2

23. Electron Dot Notation (Shows the number of valence electrons)

24. Lewis Structures Electron distribution is depicted with Lewis StructuresElectron distribution is depicted with Lewis Structures This is how you decide how many atoms will bond covalently! (In ionic bonds, it was decided with charges)This is how you decide how many atoms will bond covalently! (In ionic bonds, it was decided with charges) G. N. Lewis 1875 - 1946

25. Each dot represents a valence electron Each dash represents one covalent bond Lewis Structures

26. Octet Rule –Most atoms form bonds in order to obtain 8 valence e - –Full energy level stability ~ Noble Gases Ne

27. Lewis Structures Valence electrons are distributed as shared or bond pairs and unshared or lone pairs.Valence electrons are distributed as shared or bond pairs and unshared or lone pairs. HCl lone pair (LP) shared or bond pair This is called a Lewis structure.

28. Covalent Compounds- The Octet Rule Covalent compounds tend to form so that each atom, by sharing electrons, has an octet of electrons in its highest occupied energy level. Diatomic Fluorine

29. Cl Each chlorine atom wants to gain one electron to achieve an octet An example of covalent bonding and electron dot notation

30. Cl Neither atom will give up an electron – chlorine is highly electronegative. What’s the solution – what can they do to achieve an octet?

31. Cl

35. octet

36. Cl octet

37. Cl The octet is achieved by each atom sharing the electron pair in the middle.

38. Cl This is the bonding pair

39. Cl It is a single bonding pair

40. Cl It is called a SINGLE BOND

41. Cl Single bonds are abbreviated with a dash

42. Cl This is the chlorine molecule, Cl 2

43. –Hydrogen  2 valence e - –Groups 1,2,3 get 2,4,6 valence e - –Expanded octet  more than 8 valence e - (e.g. S, P, Xe) –Radicals  odd # of valence e - Exceptions: Octet Rule F B F F H O HN O Very unstable!! F F S F F

44. Drawing Lewis Diagrams Find total # of valence e -. Arrange atoms - singular atom is usually in the middle. Form bonds between atoms (2 e - ). Distribute remaining e - to give each atom an octet (recall exceptions). If there aren’t enough e - to go around, form double or triple bonds.

45. Drawing Lewis Diagrams Ex. #1- CF 4 1 C × 4e - = 4e - 4 F × 7e - = 28e - 32e - F F C F F - 8e - 24e -

46. Drawing Lewis Diagrams BeCl 2 1 Be × 2e - = 2e - 2 Cl × 7e - = 14e - 16e - Cl Be Cl - 4e - 12e -

47. Drawing Lewis Diagrams CO 2 1 C × 4e - = 4e - 2 O × 6e - = 12e - 16e - O C O - 4e - 12e -

48. Multiple Covalent Bonds: Double bonds Two pairs of shared electrons

49. Multiple Covalent Bonds: Triple bonds Three pairs of shared electrons

50. Writing Lewis Structures for Polyatomic Ions To find total # of valence e - : –Add 1e - for each negative charge. –Subtract 1e - for each positive charge. Place brackets around the ion and label the charge.

51. Polyatomic Ions ClO 4 - 1 Cl × 7e - = 7e - 4 O × 6e - = 24e - 31e - O O Cl O O + 1e - 32e - - 8e - 24e -

52. NH 4 + 1 N × 5e - = 5e - 4 H × 1e - = 4e - 9e - H H N H H - 1e - 8e - - 8e - 0e - Polyatomic Ions

53. Resonance Structures Molecules that can’t be correctly represented by a single Lewis diagram. Actual structure is an average of all the possibilities. Show possible structures separated by a double-headed arrow.

54. Resonance Structures O O S O O O S O O O S O o SO 3

55. Bond Length and Bond Energy BondLength (pm)Energy (kJ/mol) C - C154346 C=C134612 CCCC 120835 C - N147305 C=N132615 CNCN 116887 C - O143358 C=O120799 COCO 1131072 N - N145180 N=N125418 NNNN 110942

56. 1) Non-polar Covalent Bond (0 – 0.4) –e - are shared equally –symmetrical e - density –usually identical atoms (diatomic) Types of Covalent Bonds

57. ++ -- 2) Polar Covalent Bond (0.4 - 2.0) –e - are shared unequally –asymmetrical e - density –results in partial charges (dipoles)

58. ++ -- ++ Covalent Bonds and Lewis Structures Non-polar Covalent - no charges Polar Covalent - partial charges

59. Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world. Models can be physical as with this DNA model Models can be mathematical Models can be theoretical or philosophical

60. VSEPR Theory Types of e - Pairs –Bonding pairs - form bonds –Lone pairs - nonbonding e - Lone pairs repel more strongly than bonding pairs!!!

61. VSEPR Theory Lone pairs reduce the bond angle between atoms. Bond Angle

62. Draw the Lewis Diagram. Tally up e - pairs on central atom. –double/triple bonds = ONE pair Shape is determined by the # of bonding pairs and lone pairs. Know the 8 common shapes & their bond angles! Determining Molecular Shape

63. Common Molecular Shapes 2 total 2 bond 0 lone AX 2 LINEAR 180° BeH 2 Central atom Outer (bonded) atoms

64. 3 total 3 bond 0 lone TRIGONAL PLANAR 120° BF 3 Common Molecular Shapes

65. 3 total 2 bond 1 lone AX 2 E BENT <120° SO 2 Lone Pair

66. 4 total 4 bond 0 lone TETRAHEDRAL 109.5° CH 4 Common Molecular Shapes

67. 4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107° NH 3 Common Molecular Shapes

68. 4 total 2 bond 2 lone BENT 104.5° H2OH2O Common Molecular Shapes

69. 5 total 5 bond 0 lone TRIGONAL BIPYRAMIDAL 120°/90° PCl 5 Common Molecular Shapes

70. 6 total 6 bond 0 lone OCTAHEDRAL 90° SF 6 Common Molecular Shapes

71. PF 3 4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107° F P F F VESPR Examples

72. CO 2 O C O 2 total 2 bond 0 lone LINEAR 180° Examples

73. TypeGeometryExample Compound Example Structure Bond Angle AX 2 LinearBeCl 2 BeCl 180 o AX 3 Trigonal planar BF 3 B F FF 120 o AX 4 Tetrahedron CH 4 C H H H H 109.5 o

74. TypeGeometryExample Compound Example Structure Bond Angle AX 5 Trigonal Bipyramid PCl 5 Cl P 90 o ; 120 o AX 6 Octahedron SF 6 S F F F F F F 90 o

75. TypeGeometryExample Compound Example Structure Bond Angle AX 4 Tetrahedron CH 4 C H H H H 109.5 o AX 3 E Trigonal Pyramid NH 3 C H H H 107.3 o

76. TypeGeometryExample Compound Example Structure Bond Angle AX 2 E 2 AXE 3 Bent Linear H2OH2O HCl O HH 104.5 o H Cl