2. Chapter-2-1 Chemistry 281, Winter 2015, LA Tech CTH 277 10:00-11:15 am Instructor: Dr. Upali Siriwardane E-mailE-mail: upali@latech.edu Office: 311 Carson Taylor Hall ; Phone: 318-257- 4941; Office Hours: MTW 8:00 am - 10:00 am; Th,F 8:30 - 9:30 am & 1:00-2:00 pm. January 13, 2015 Test 1 (Chapters 1&,2), February 3, 2015 Test 2 (Chapters 2 & 3) February 26, 2015, Test 3 (Chapters 4 & 5), Comprehensive Final Make Up Exam: March 3 Chemistry 281(01) Winter 2015

3. Chapter-2-2 Chemistry 281, Winter 2015, LA Tech Molecular structure and bonding Lewis structures 2.1 The octet rule 2.2 Resonance 2.3 The VSEPR model Valence-bond theory 2.4 The hydrogen molecule 2.5 Homonuclear diatomic molecules 2.6 Polyatomic molecules Molecular orbital theory 2.7 An introduction to the theory 2.8 Homonuclear diatomic molecules 2.9 Heteronuclear diatomic 2.10 Bond properties

4. Chapter-2-3 Chemistry 281, Winter 2015, LA Tech Lewis Theory of Bonding Octet Rule All elements except hydrogen ( hydrogen have a duet of electrons) have octet of electrons once they from ions and covalent compounds.

5. Chapter-2-4 Chemistry 281, Winter 2015, LA Tech Noble gas configuration The noble gases are noted for their chemical stability and existence as monatomic molecules. Except for helium, They share a common electron configuration that is very stable. This configuration has 8 valence-shell electrons. All other elements reacts to achieve Noble Gas Electron Configurations. valence e- He 2 Ne 8 Ar 8 Kr 8 Xe 8 Rn 8 valence e- He 2 Ne 8 Ar 8 Kr 8 Xe 8 Rn 8

6. Chapter-2-5 Chemistry 281, Winter 2015, LA Tech The octet rule Atoms are most stable if they have a filled or empty outer layer of electrons. Atoms are most stable if they have a filled or empty outer layer of electrons. Except for H and He, a filled layer contains 8 electrons - an octet. Except for H and He, a filled layer contains 8 electrons - an octet. Two atoms will Two atoms will gain or lose (ionic compounds) share (covalent compounds) Many atoms with fewer electrons will share (metallic compounds) share (metallic compounds)

7. Chapter-2-6 Chemistry 281, Winter 2015, LA Tech What changes take place during this process of achieving closed shells? a) sharing leads to covalent bonds and molecules b) gain/loss of electrons lead to ionic bond c) Sharing with many atoms lead to metallic bonds

8. Chapter-2-7 Chemistry 281, Winter 2015, LA Tech Lewis Electron Dot symbols X Basic rules Draw the atomic symbol. Treat each side as a box that can hold up to two electrons. Count the electrons in the valence shell. Start filling box - don’t make pairs unless you need to.

9. Chapter-2-8 Chemistry 281, Winter 2015, LA Tech Lewis symbols LiBeBC NOFNe Lewis symbols of second period elements

10. Chapter-2-9 Chemistry 281, Winter 2015, LA Tech What is a Lewis Structure (electron-dot formula) of a Molecule? A molecular formulas with dots around atomic symbols representing the valence electrons A molecular formulas with dots around atomic symbols representing the valence electrons All atoms will have eight (octet) of electrons All atoms will have eight (octet) of electrons (duet for H) if the molecule is to be stable. (duet for H) if the molecule is to be stable.

11. Chapter-2-10 Chemistry 281, Winter 2015, LA Tech Single covalent bonds HH H C HH H Do atoms (except H) have octets? FF

12. Chapter-2-11 Chemistry 281, Winter 2015, LA Tech Lewis structures This is a simple system to help keep track of electrons around atoms, ions and molecules - invented by G.N. Lewis. This is a simple system to help keep track of electrons around atoms, ions and molecules - invented by G.N. Lewis. If you know the number of electrons in the valence-shell of an atom, writing Lewis structures is easy. If you know the number of electrons in the valence-shell of an atom, writing Lewis structures is easy. Lewis structures are used primarily for s- and p-block elements. Lewis structures are used primarily for s- and p-block elements.

13. Chapter-2-12 Chemistry 281, Winter 2015, LA Tech Add all valence electrons and get valence electron pairs Pick the central atom: Largest atom normally or atom forming most bonds Connect central atom to terminal atoms Fill octet to all atoms (duet to hydrogen) How do you get the Lewis Structure from Molecular formula?

14. Chapter-2-13 Chemistry 281, Winter 2015, LA Tech Lewis Structure of H 2 O

15. Chapter-2-14 Chemistry 281, Winter 2015, LA Tech Types of electrons Bonding pairs Two electrons that are shared between two atoms. A covalent bond. Unshared (nonbonding ) pairs A pair of electrons that are not shared between two atoms. Lone pairs or nonbonding electrons. H Cl oo Bonding pair Unshared pair

16. Chapter-2-15 Chemistry 281, Winter 2015, LA Tech 2 bond pairs= 2 x 2 = 4 2 lone pairs = 2 x 2 =4 Total 8 = 4 pairs Bond pairs: an electron pair shared by two atom in a bond. E.g. two pairs between O--H in water. Lone pair : an electron pair found solely on a single atom. E.g. two pairs found on the O atom at the top and the bottom. Lewis Structure of H 2 O

17. Chapter-2-16 Chemistry 281, Winter 2015, LA Tech Lewis Structure of H 2 S

18. Chapter-2-17 Chemistry 281, Winter 2015, LA Tech Lewis Structure of CCl 4

19. Chapter-2-18 Chemistry 281, Winter 2015, LA Tech CO 2 CO 2 NH 3 (PH 3 ) NH 3 (PH 3 ) PCl 3 (PF 3, NCl 3 ) PCl 3 (PF 3, NCl 3 ) What is the Lewis Structure?

20. Chapter-2-19 Chemistry 281, Winter 2015, LA Tech Lewis structure and multiple bonds O=C=O This arrangement needs too many electrons. How about making some double bonds? That works! O C O = is a double bond, the same as 4 electrons

21. Chapter-2-20 Chemistry 281, Winter 2015, LA Tech Multiple bonds So how do we know that multiple bonds really exist? The bond energies and lengths differ! BondBond LengthBond energy typeorder pm kJ/mol C C 1 154 347 C C 2 134 615 C C 2 134 615 C C 3 120 812 C C 3 120 812

22. Chapter-2-21 Chemistry 281, Winter 2015, LA Tech Formal Charges Formal charge = valence electrons - assigned electrons There are two possible Lewis structures for a molecule. Each has the same number of bonds. We can determine which is better by determining which has the least formal charge. It takes energy to get a separation of charge in the molecule There are two possible Lewis structures for a molecule. Each has the same number of bonds. We can determine which is better by determining which has the least formal charge. It takes energy to get a separation of charge in the molecule (as indicated by the formal charge) so the structure with the least formal charge should be lower in energy and thereby be the better Lewis structure (as indicated by the formal charge) so the structure with the least formal charge should be lower in energy and thereby be the better Lewis structure

23. Chapter-2-22 Chemistry 281, Winter 2015, LA Tech Formal Charge Calculation Formal charge = group number in periodic table number of bonds number of unshared electrons – – An arithmetic formula for calculating formal charge.

24. Chapter-2-23 Chemistry 281, Winter 2015, LA Tech Electron counts" and formal charges in NH 4 + and BF 4 - "

25. Chapter-2-24 Chemistry 281, Winter 2015, LA Tech What is Resonance Structures? Several Lewis structures that need to be drawn for molecules with double bonds Several Lewis structures that need to be drawn for molecules with double bonds One Lewis structure alone would not describe the bond lengths of the real molecule. One Lewis structure alone would not describe the bond lengths of the real molecule. E.g. CO 3 2-, NO 3 -, NO 2 -, SO 3 E.g. CO 3 2-, NO 3 -, NO 2 -, SO 3

26. Chapter-2-25 Chemistry 281, Winter 2015, LA Tech Sometimes we can have two or more equivalent Lewis structures for a molecule. O - S = O O = S - O O - S = O O = S - O They both- satisfy the octet rule - have the same number of bonds - have the same number of bonds - have the same types of bonds - have the same types of bonds Which is right? Resonance structures

27. Chapter-2-26 Chemistry 281, Winter 2015, LA Tech They both are! O - S = O O = S - O O - S = O O = S - O O S O O S O This results in an average of 1.5 bonds between each S and O. Resonance structures of SO 2

28. Chapter-2-27 Chemistry 281, Winter 2015, LA Tech Resonance structures of CO 3 2- ion

29. Chapter-2-28 Chemistry 281, Winter 2015, LA Tech Resonance structures of NO 3 - ion

30. Chapter-2-29 Chemistry 281, Winter 2015, LA Tech Resonance structures of SO 3

31. Chapter-2-30 Chemistry 281, Winter 2015, LA Tech Resonance structures of NO 2 - ion

32. Chapter-2-31 Chemistry 281, Winter 2015, LA Tech Resonance structures of C 6 H 6 Benzene, C 6 H 6, is another example of a compound for which resonance structure must be written. Benzene, C 6 H 6, is another example of a compound for which resonance structure must be written. All of the bonds are the same length. All of the bonds are the same length. or

33. Chapter-2-32 Chemistry 281, Winter 2015, LA Tech Exceptions to the octet rule Not all compounds obey the octet rule. Three types of exceptions Three types of exceptions Species with more than eight electrons around an atom. Species with more than eight electrons around an atom. Species with fewer than eight electrons around an atom. Species with fewer than eight electrons around an atom. Species with an odd total number of electrons. Species with an odd total number of electrons.

34. Chapter-2-33 Chemistry 281, Winter 2015, LA Tech Atoms with more than eight electrons Except for species that contain hydrogen, this is the most common type of exception. Except for species that contain hydrogen, this is the most common type of exception. For elements in the third period and beyond, the d orbitals can become involved in bonding. For elements in the third period and beyond, the d orbitals can become involved in bonding.Examples 5 electron pairs around P in PF 5 5 electron pairs around P in PF 5 5 electron pairs around S in SF 4 5 electron pairs around S in SF 4 6 electron pairs around S in SF 6 6 electron pairs around S in SF 6

35. Chapter-2-34 Chemistry 281, Winter 2015, LA Tech An example: SO 4 2- 1. Write a possible arrangement. 2. Total the electrons. 6 from S, 4 x 6 from O add 2 for charge total = 32 3. Spread the electrons around. SO O O O - - || SO O O O

36. Chapter-2-35 Chemistry 281, Winter 2015, LA Tech Atoms with fewer than eight electrons Beryllium and boron will both form compounds where they have less than 8 electrons around them.

37. Chapter-2-36 Chemistry 281, Winter 2015, LA Tech Atoms with fewer than eight electrons Electron deficient. Species other than hydrogen and helium that have fewer than 8 valence electrons. They are typically very reactive species. F|B|FF|B|F F - + H | :N - H | H F H | F - B - N – H | F H

38. Chapter-2-37 Chemistry 281, Winter 2015, LA Tech What is VSEPR Theory Valence Shell Electron Pair Repulsion This theory assumes that the molecular structure is determined by the lone pair and bond pair electron repulsion around the central atom

39. Chapter-2-38 Chemistry 281, Winter 2015, LA Tech What Geometry is Possible around Central Atom? What is Electronic or Basic Structure? What is Electronic or Basic Structure? Arrangement of electron pairs around the central atom is called the electronic or basic structure Arrangement of electron pairs around the central atom is called the electronic or basic structure What is Molecular Structure? What is Molecular Structure? Arrangement of atoms around the central atom is called the molecular structure Arrangement of atoms around the central atom is called the molecular structure

40. Chapter-2-39 Chemistry 281, Winter 2015, LA Tech Possible Molecular Geometry 1. Linear (180) 2. Trigonal Planar (120) 3. T-shape (90, 180) 4. Tetrahedral (109) 5. Square palnar ( 90, 180) 6. Sea-saw (90, 120, 180) 7. Trigonal bipyramid (90, 120, 180) 8. Octahedral (90, 180)

41. Chapter-2-40 Chemistry 281, Winter 2015, LA Tech Molecular Structure from VSEPR Theory H 2 O H 2 O Bent or angular Bent or angular NH 3 NH 3 Pyramidal Pyramidal CO 2 CO 2 Linear Linear

42. Chapter-2-41 Chemistry 281, Winter 2015, LA Tech Molecular Structure from VSEPR Theory SF 6 SF 6 Octahedral Octahedral PCl 5 PCl 5 Trigonal bipyramidal Trigonal bipyramidal XeF 4 XeF 4 Square planar Square planar

43. Chapter-2-42 Chemistry 281, Winter 2015, LA Tech What is a Polar Molecule? Molecules with unbalanced electrical charges Molecules with unbalanced electrical charges Molecules with a dipole moment Molecules with a dipole moment Molecules without a dipole moment are called non-polar molecules Molecules without a dipole moment are called non-polar molecules

44. Chapter-2-43 Chemistry 281, Winter 2015, LA Tech How do you a Pick Polar Molecule? Get the molecular structure from VSEPR theory Get the molecular structure from VSEPR theory From  (electronegativity) difference of bonds see whether they are polar-covalent. From  (electronegativity) difference of bonds see whether they are polar-covalent. If the molecule have polar-covalent bond, check whether they cancel from a symmetric arrangement. If the molecule have polar-covalent bond, check whether they cancel from a symmetric arrangement. If not molecule is polar If not molecule is polar

45. Chapter-2-44 Chemistry 281, Winter 2015, LA Tech H 2 O Bent or angular, polar-covalent bonds, asymmetric molecule-polar NH 3 Pyramidal, polar-covalent bonds, asymmetric molecule-polar CO 2 Linear, polar-covalent bonds, symmetric molecule-polar Which Molecules are Polar

46. Chapter-2-45 Chemistry 281, Winter 2015, LA Tech What is hybridization? Mixing of atomic orbitals on the central atoms valence shell (highest n orbitals) Bonding: s p d sp, sp 2, sp 3, sp 3 d, sp 3 d 2 PxPx PyPy PzPz dz2dz2 d x 2 - y 2

47. Chapter-2-46 Chemistry 281, Winter 2015, LA Tech What is hybridization? Mixing of atomic orbitals on the central atom Bonding a hybrid orbital could over lap with another (  ) atomic orbital or (  ) hybrid orbital of another atom to make a covalent bond. possible hybridizations: sp, sp 2, sp 3, sp 3 d, sp 3 d 2

48. Chapter-2-47 Chemistry 281, Winter 2015, LA Tech What is Valence Bond Theory Describes bonding in molecule using atomic orbital Describes bonding in molecule using atomic orbital orbital of one atom occupy the same region with a orbital from another atom orbital of one atom occupy the same region with a orbital from another atom total number of electrons in both orbital is equal to two total number of electrons in both orbital is equal to two Be Cl 2

49. Chapter-2-48 Chemistry 281, Winter 2015, LA Tech sp 2 and sp 3 Hybridization BF 3

50. Chapter-2-49 Chemistry 281, Winter 2015, LA Tech What are  and  bonds  bonds single bond resulting from head to head overlap of atomic orbital  bond double and triple bond resulting from lateral or side way overlap of atomic orbitals

51. Chapter-2-50 Chemistry 281, Winter 2015, LA Tech How do you tell the hybridization of a central atom? Get the Lewis structure of the molecule Get the Lewis structure of the molecule Look at the number of electron pairs on the central atom. Note: double, triple bonds are counted as single electron pairs. Look at the number of electron pairs on the central atom. Note: double, triple bonds are counted as single electron pairs. Follow the following chart Follow the following chart

52. Chapter-2-51 Chemistry 281, Winter 2015, LA Tech Kinds of hybrid orbitals Hybridgeometry# of orbital sp linear2 sp linear2 sp 2 trigonal planar3 sp 3 tetrahedral4 sp 3 d trigonal bipyramid5 sp 3 d 2 octahedral6

53. Chapter-2-52 Chemistry 281, Winter 2015, LA Tech Hybridization involving d orbitals Co(NH 3 ) 6 3+ ion Co 3+ : [Ar] 3d 6 Co(NH 3 ) 6 3+ ion Co 3+ : [Ar] 3d 6 Co 3+ : [Ar] 3d 6 4s 0 4p 0 Co 3+ : [Ar] 3d 6 4s 0 4p 0 Concentrating the 3d electrons in the d xy, d xz, and d yz orbitals in this subshell gives the following electron configuration hybridization is sp 3 d 2 Concentrating the 3d electrons in the d xy, d xz, and d yz orbitals in this subshell gives the following electron configuration hybridization is sp 3 d 2

54. Chapter-2-53 Chemistry 281, Winter 2015, LA Tech Molecular Orbital Theory Molecular orbitals are obtained by combining the atomic orbitals on the atoms in the molecule. Molecular orbitals are obtained by combining the atomic orbitals on the atoms in the molecule.

55. Chapter-2-54 Chemistry 281, Winter 2015, LA Tech Bonding and Anti-bobding Molecular Orbital

56. Chapter-2-55 Chemistry 281, Winter 2015, LA Tech Basic Rules of Molecular Orbital Theory The MO Theory has five basic rules: The number of molecular orbitals = the number of atomic orbitals combined The number of molecular orbitals = the number of atomic orbitals combined Of the two MO's, one is a bonding orbital (lower energy) and one is an anti-bonding orbital (higher energy) Of the two MO's, one is a bonding orbital (lower energy) and one is an anti-bonding orbital (higher energy) Electrons enter the lowest orbital available Electrons enter the lowest orbital available The maximum # of electrons in an orbital is 2 (Pauli Exclusion Principle) The maximum # of electrons in an orbital is 2 (Pauli Exclusion Principle) Electrons spread out before pairing up (Hund's Rule) Electrons spread out before pairing up (Hund's Rule)

57. Chapter-2-56 Chemistry 281, Winter 2015, LA Tech Bond Order Calculating Bond Order Calculating Bond Order

58. Chapter-2-57 Chemistry 281, Winter 2015, LA Tech Homo Nuclear Diatomic Molecules Period 1 Diatomic Molecules: H 2 and He 2

59. Chapter-2-58 Chemistry 281, Winter 2015, LA Tech Homo Nuclear Diatomic Molecules Period 2 Diatomic Molecules and Li 2 and Be 2

60. Chapter-2-59 Chemistry 281, Winter 2015, LA Tech Homo Nuclear Diatomic Molecules

61. Chapter-2-60 Chemistry 281, Winter 2015, LA Tech Molecualr Orbital diagram for O 2, F 2 and Ne 2

62. Chapter-2-61 Chemistry 281, Winter 2015, LA Tech Molecualr Orbital diagram for B 2, C 2 and N 2

63. Chapter-2-62 Chemistry 281, Winter 2015, LA Tech Homonuclear Diatomic Molecules 2 nd Period

64. Chapter-2-63 Chemistry 281, Winter 2015, LA Tech Electronic Configuration of molecules When writing the electron configuration of an atom, we usually list the orbitals in the order in which they fill. Pb: [Xe] 6s 2 4f 14 5d 10 6p 2 We can write the electron configuration of a molecule by doing the same thing. Concentrating only on the valence orbitals, we write the electron configuration of O 2 as follows. O 2 : ( 2  2s  2 (2  * 2s ) 2 (2  2p  2 (2  2p  4 (2  * 2p  2 ( 2  * 2p )

65. Chapter-2-64 Chemistry 281, Winter 2015, LA Tech Electronic Configuration and bond order

66. Chapter-2-65 Chemistry 281, Winter 2015, LA Tech Hetero Nuclear Diatomic Molecules HF molecule

67. Chapter-2-66 Chemistry 281, Winter 2015, LA Tech Hetero Nuclear Diatomic Molecules Carbon monoxide CO

68. Chapter-2-67 Chemistry 281, Winter 2015, LA Tech Metallic Bonding Metals are held together by delocalized bonds formed from the atomic orbitals of all the atoms in the lattice. Metals are held together by delocalized bonds formed from the atomic orbitals of all the atoms in the lattice. The idea that the molecular orbitals of the band of energy levels are spread or delocalized over the atoms of the piece of metal accounts for bonding in metallic solids. The idea that the molecular orbitals of the band of energy levels are spread or delocalized over the atoms of the piece of metal accounts for bonding in metallic solids.

69. Chapter-2-68 Chemistry 281, Winter 2015, LA Tech Bonding Models for Metals Band Theory of Bonding in Solids Band Theory of Bonding in Solids Bonding in solids such as metals, insulators and semiconductors may be understood most effectively by an expansion of simple MO theory to assemblages of scores of atoms Bonding in solids such as metals, insulators and semiconductors may be understood most effectively by an expansion of simple MO theory to assemblages of scores of atoms

70. Chapter-2-69 Chemistry 281, Winter 2015, LA Tech Linear Combination of Atomic Orbitals

71. Chapter-2-70 Chemistry 281, Winter 2015, LA Tech Linear Combination of Atomic Orbitals

72. Chapter-2-71 Chemistry 281, Winter 2015, LA Tech

73. Chapter-2-72 Chemistry 281, Winter 2015, LA Tech Types of Materials A conductor (which is usually a metal) is a solid with a partially full band A conductor (which is usually a metal) is a solid with a partially full band An insulator is a solid with a full band and a large band gap An insulator is a solid with a full band and a large band gap A semiconductor is a solid with a full band and a small band gap A semiconductor is a solid with a full band and a small band gap Element Band Gap C 5.47 eV Si 1.12 eV Ge 0.66 eV Sn 0 eV Element Band Gap C 5.47 eV Si 1.12 eV Ge 0.66 eV Sn 0 eV

74. Chapter-2-73 Chemistry 281, Winter 2015, LA Tech

75. Chapter-2-74 Chemistry 281, Winter 2015, LA Tech Superconductors When Onnes cooled mercury to 4.15K, the resistivity suddenly dropped to zero When Onnes cooled mercury to 4.15K, the resistivity suddenly dropped to zero

76. Chapter-2-75 Chemistry 281, Winter 2015, LA Tech The Meissner Effect Superconductors show perfect diamagnetism. Superconductors show perfect diamagnetism. Meissner and Oschenfeld discovered that a superconducting material cooled below its critical temperature in a magnetic field excluded the magnetic flux.Results in levitation of the magnet in a magnetic field. Meissner and Oschenfeld discovered that a superconducting material cooled below its critical temperature in a magnetic field excluded the magnetic flux.Results in levitation of the magnet in a magnetic field.

77. Chapter-2-76 Chemistry 281, Winter 2015, LA Tech Theory of Superconduction BCS theory was proposed by J. Bardeen, L. Cooper and J. R. Schrieffer. BCS suggests the formation of so-called 'Cooper pairs' BCS theory was proposed by J. Bardeen, L. Cooper and J. R. Schrieffer. BCS suggests the formation of so-called 'Cooper pairs' Cooper pair formation - electron- phonon interaction: the electron is attracted to the positive charge density (red glow) created by the first electron distorting the lattice around itself.

78. Chapter-2-77 Chemistry 281, Winter 2015, LA Tech High Temperature Superconduction BCS theory predicted a theoretical maximum to Tc of around 30-40K. Above this, thermal energy would cause electron- phonon interactions of an energy too high to allow formation of or sustain Cooper pairs. BCS theory predicted a theoretical maximum to Tc of around 30-40K. Above this, thermal energy would cause electron- phonon interactions of an energy too high to allow formation of or sustain Cooper pairs. 1986 saw the discovery of high temperature superconductors which broke this limit (the highest known today is in excess of 150K) - it is in debate as to what mechanism prevails at higher temperatures, as BCS cannot account for this. 1986 saw the discovery of high temperature superconductors which broke this limit (the highest known today is in excess of 150K) - it is in debate as to what mechanism prevails at higher temperatures, as BCS cannot account for this.