Molecular structure and bonding

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1 Chemistry 481(01) Spring 2017 Instructor: Dr. Upali Siriwardane Office: CTH 311 Phone Office Hours: M,W 8:00-9:00 & 11:00-12:00 am; Tu,Th, F 9:30-11:30 a.m. April 4, 2017: Test 1 (Chapters 1, 2, 3, 4) April 27, 2017: Test 2 (Chapters (6 & 7) May 16, 2016: Test 3 (Chapters. 19 & 20) May 17, Make Up: Comprehensive covering all Chapters Molecular structure and bonding Lewis structures 2.1 The octet rule 2.2 Structure and bond properties 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 Chapter-2-1 Chapter-2-2 What changes take place during this process of achieving closed shells? a) sharing leads to covalent bonds and molecules Covalent Bond: each atom gives one electron Coordinative bond: two electron comes from one atom b) gain/loss of electrons lead to ionic bond Cations and anions: Electrostatic attractions c) Sharing with many atoms lead to metallic bonds: delocalization of electrons How do you get the Lewis Structure from Molecular formula? 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) Chapter-2-3 Chapter Draw Lewis structure for SbF 5, ClF 3, and IF 6+ : 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 Chapter-2-5 Chapter-2-6 1

2 What Geometry is Possible around Central Atom? What is Electronic or Basic Structure? Arrangement of electron pairs around the central atom is called the electronic or basic structure What is Molecular Structure? Arrangement of atoms around the central atom is called the molecular structure Possible Molecular Geometry Linear (180) Trigonal Planar (120) T-shape (90, 180) Tetrahedral (109) Square palnar ( 90, 180) Sea-saw (90, 120, 180) Trigonal bipyramid (90, 120, 180) Octahedral (90, 180) Chapter-2-7 Chapter Predict geometry of central atom using VSEPR and the hybridization in problem 1. SbF 5, ClF 3, and IF 6+ : Formal Charges Formal charge = valence electrons - assigned electrons If 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 Chapter-2-9 Chapter-2-10 Formal Charge Calculation Electron counts" and formal " charges in NH 4+ and BF - 4 An arithmetic formula for calculating formal charge. Formal charge = group number in periodic table number of bonds number of unshared electrons Chapter-2-11 Chapter

3 Resonance structures of SO 2 Resonance structures of CO 3 2- ion They both are! O - S = O O = S - O O S O This results in an average of 1.5 bonds between each S and O. Ave. Bond order= total pairs shared/ # bonds= 3/2=1.5 Chapter-2-13 Chapter-2-14 Resonance structures of C 6 H 6 Exceptions to the octet rule 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. or Not all compounds obey the octet rule. Three types of exceptions Species with more than eight electrons around an atom. Species with fewer than eight electrons around an atom. Species with an odd total number of electrons. Chapter-2-15 Chapter-2-16 Valence-bond (VB) theory VB theory combines the concepts of atomic orbitals, hybrid orbitals, VSEPR, resonance structures, Lewis structures and octet rule to describe the shapes and structures of some common molecules. It uses the overlap of atomic orbitals or hybrid orbitals of the to from sigma (s), pi (p) bonds and (d) bonds Linear Combination of Atomic Orbitals Symmetry Adapted Linear Combination of Atomic Orbitals LCAO Atomic orbitals on single atom: Hybridization Atomic orbitals in a molecule with more than one atom: Molecular Orbital (MO) formation General rule Number of Hybrid Orbital produced = # hybridized Number of MO produced = # orbitals combined Chapter-2-17 Chapter

4 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. How do you tell the hybridization of a central atom? 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. Follow the following chart possible hybridizations: sp, sp 2, sp 3, sp 3 d, sp 3 d 2 Chapter-2-19 Chapter-2-20 Kinds of hybrid orbitals Hybrid geometry # of orbital sp linear 2 sp 2 trigonal planar 3 sp 3 tetrahedral 4 sp 3 d trigonal bipyramid 5 sp 3 d 2 octahedral 6 What is hybridization? Mixing of atomic orbitals on the central atoms valence shell (highest n orbitals) Bonding: s p d sp, sp 2, P x P y P z d z 2 d x2 - y 2 sp 3, sp 3 d, Chapter-2-21 sp 3 d 2 Chapter-2-22 Possible hybridizations of s and p sp-hybridization: 1 = 1/ 2 s - 1/ 2 p 2 = 1/ 2 s + 1/ 2 p sp 2 -hybridization: 1 = 1/ 3 s + 1/ 6 px + 1/ 2 py 2 = 1/ 3 s + 1/ 6 px - 1/ 2 py 3 = 1/ 3 s - 2/ 6 px sp 3 -hybridization: 1 = 1/ 4 s + 1/ 4 px + 1/ 4 py + 1/ 4 pz 2 = 1/ 4 s - 1/ 4 px - 1/ 4 py + 1/ 4 pz 3 = 1/ 4 s + 1/ 4 px - 1/ 4 py - 1/ 4 pz 4 = 1/ 4 s - 1/ 4 px + 1/ 4 py -1/ 4 pz Possible hybridizations of s and p sp-hybridization: Chapter-2-23 Chapter

5 What are p and s bonds s bonds single bond resulting from head to head overlap of atomic orbital p bond double and triple bond resulting from lateral or side way overlap of p atomic orbitals d bond double and triple bond resulting from lateral or side way overlap of d atomic orbitals Atoms with more than eight electrons 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. Examples 5 electron pairs around P in PF 5 5 electron pairs around S in SF 4 6 electron pairs around S in SF 6 Chapter-2-25 Chapter-2-26 An example: SO Why hypervalent compounds are formed by elements such as Si, P and S, but not by C,N and O? 1. Write a possible arrangement. O O S O 2. Total the electrons. 6 from S, 4 x 6 from O add 2 for charge total = Spread the electrons around. O O O - S - O O Chapter-2-27 Chapter-2-28 Atoms with fewer than eight electrons Atoms with fewer than eight electrons Beryllium and boron will both form compounds where they have less than 8 electrons around them. Electron deficient. Species other than hydrogen and helium that have fewer than 8 valence electrons. : : : : : Cl:Be:Cl: :F:B:F: :F: : : : : : : They are typically very reactive species. F - F B F + H :N H H F H F - B <- N - H F H Chapter-2-29 Chapter

6 What is a Polar Molecule? Molecules with unbalanced electrical charges Molecules with a dipole moment Molecules without a dipole moment are called non-polar molecules How do you a Pick Polar Molecule? a) Get the molecular structure from VSEPR theory b) From c (electronegativity) difference of bonds see whether they are polar-covalent. c) If the molecule have polar-covalent bond, check whether they cancel from a symmetric arrangement. d) If not molecule is polar Predicting symmetry of molecule and the polarity will be discussed in detail in Chapter 7. Chapter-2-31 Chapter-2-32 Linear Combination of Atomic Orbitals Symmetry Adapted Linear Combination of Atomic Orbitals LCAO Atomic orbitals on single atom: Hybridization Atomic orbitals in a molecule with more than one atom: Molecular Orbital (MO) formation General rule Number of Hybrid Orbital produced = # hybridized Number of MO produced = # orbitals combined 6. Draw a diagram to illustrate each described overlap: a) s bonding overlap of two p orbitals b) d bonding overlap of two d orbitals c) p bonding overlap of a p orbital and a d orbital d) s antibonding overlap of a p and a d orbital e) d antibonding overlap of two d orbitals. Chapter-2-33 Chapter-2-34 What are p and s bonds s bonds p bond What are d bonds d bond double and triple bond resulting from lateral or side way overlap of d atomic orbitals Chapter-2-35 Chapter

7 Kinds of hybrid orbitals Hybrid geometry # of orbital sp linear 2 sp 2 trigonal planar 3 sp 3 tetrahedral 4 sp 3 d trigonal bipyramid 5 sp 3 d 2 octahedral 6 5. Using valence-bond (VB) theory to explain the bonding in the coordination complex ion, Co(NH 3 ) Chapter-2-37 Chapter-2-38 Hybridization involving d orbitals Co(NH 3 ) 3+ 6 ion Co 3+ : [Ar] 3d 6 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 5. What is the oxidation state of metal in (a) Co(NH 3 ) 6 3+ ion (b) PtCl 4 2- ion. a) [Co(NH 3 ) 6 ] 3+ Co 3+ and NH 3 is neutral Oxidation Sate of Co 3+ is +3 and NH 3 is 0 Therefore sum of the oxidation should be equal to +3 +3= Co(NH 3 ) 6 = (Co)3+6((NH 3 )0)= +3 Co is +3 in [Co(NH 3 ) 6 ] 3+ b) Pt is +2 in [PtCl 4 ] 2- because Cl - is -1 Chapter-2-39 Chapter-2-40 Linear Combination of Atomic Orbitals Symmetry Adapted Linear Combination of Atomic Orbitals LCAO Atomic orbitals on single atom: Hybridization Atomic orbitals in a molecule with more than one atom: Molecular Orbital (MO) formation General rule Number of Hybrid Orbital produced = # hybridized Number of MO produced = # orbitals combined Basic Rules of Molecular Orbital Theory The MO Theory has five basic rules: 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) Electrons enter the lowest orbital available The maximum # of electrons in an orbital is 2 (Pauli Exclusion Principle) Electrons spread out before pairing up (Hund's Rule) Chapter-2-41 Chapter

8 Molecular Orbital Theory Bonding and Anti-bobding Molecular Orbital Molecular orbitals are obtained by combining the atomic orbitals on the atoms in the molecule. Chapter-2-43 Chapter-2-44 Bond Order Calculating Bond Order Homo Nuclear Diatomic Molecules Period 1 Diatomic Molecules: H 2 and He 2 Chapter-2-45 Chapter-2-46 Homo Nuclear Diatomic Molecules Homo Nuclear Diatomic Molecules Period 2 Diatomic Molecules and Li 2 and Be 2 Chapter-2-47 Chapter

9 Molecualr Orbital diagram for Molecualr Orbital diagram for B 2, C 2 and N 2 O 2, F 2 and Ne 2 Chapter-2-49 Chapter Using molecular orbital theory and diagrams, explain why, O 2 is a paramagnetic whereas N 2 is diamagnetic. 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 : (2s) 2 (2s*) 2 (2p) 4 (2p*) 2 Chapter-2-51 Chapter-2-52 Electronic Configuration and bond order Electronic Configuration and bond order Chapter-2-53 Chapter

10 Hetero Nuclear Diatomic Molecules Carbon monoxide CO 8. Draw molecular orbital diagrams for HF, CO, NO, NO +. Calculate their bond order and predict magnetic properties. Chapter-2-55 Chapter-2-56 MO Correlation Diagrams ( Walsh Diagrams) The correlation diagram clearly indicates that the molecular orbital energy levels changes as the H 3 changes from linear to cyclic (equilateral triangle) structure. In the case of linear H 3 the overlap between two terminal H is minimal, where as in the case of cyclic H 3 the overlap is substantial. This will bring the lowest MO (bonding) and the highest MO (antibonding) down in energy. At the same time, the non-bonding MO (middle one) will go up in energy, leading to a degenerate set of levels. Thus H 3+ (two electrons) will be triangular. Walsh Diagram for H 3 : Chapter-2-57 Chapter Draw a molecular orbital diagram for triangular H 3+ and describe the bonding. 10. Draw a Walsh diagram (orbital correlation diagram) and show that triangular H 3+ is more stable than linear H 3+. Chapter-2-59 Chapter

11 Conjugated and aromatic molecules trans-1,3-butadiene Allyl radical Cyclopropenium ion: C 3 H 3 + Cyclobutadiene Cyclopentadiene Benzene C 7 H 7+ (tropyllium) and C 8 H 8 2+ trans-1,3-butadiene Chapter-2-61 Chapter-2-62 Allyl radical Cyclopropenium ion: C 3 H 3 + Chapter-2-63 Chapter-2-64 Cyclopentadiene Benzene Chapter-2-65 Chapter

12 Aromatic Rings 11. Using molecular orbital diagrams for pi (p) orbitals explain the relative stabilities of the following: (a) C 3 H 3 and C 3 H 3 + (b) C 4 H 4 and C 4 H 4 + (c) C 5 H 5 and C 5 H 5 - (d) C 6 H 6 and C 6 H 6 + (e) C 7 H 7 and C 7 H 7 + Chapter-2-67 Chapter-2-68 The Isolobal Analogy Different groups of atoms can give rise to similar shaped fragments. Chapter-2-69 Chapter Pick the isolobal fragments among the following: a) Co 3 (CO) 9 Co(CO) 3, Co 3 (CO) 9 PR, Co 3 (CO) 9 CH b) H 3 CCl, Mn(CO) 5 H, Re(CO) 5 Cl c) R 2 SiH 2, Fe(CO) 4 H 2, H 2 CH 2 Metallic Bonding 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. Chapter-2-71 Chapter

13 Linear Combination of Atomic Orbitals Bonding Models for Metals 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 Chapter-2-73 Chapter-2-74 Linear Combination of Atomic Orbitals Chapter-2-75 Chapter-2-76 Band Theory of Metals 13. Describe metallic bonding and properties in terms of: a) Electron-sea model of bonding: b) Band Theory: Chapter-2-77 Chapter

14 14. Draw the s band (molecular orbitals) for ten Na on a line (one dimensional) and show bonding and anti-bonding molecular orbitals and fill electrons. 15. Describe the metallic properties of sodium in terms of band theory. Chapter-2-79 Chapter Using a band diagram, explain how magnesium can exhibit metallic behavior even though its 3s band is completely full. Types of Materials 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 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 Chapter-2-81 Chapter Draw a Band diagram for carbon/silicon/germanium/tin, and label valence band, conduction band and band gap? Chapter-2-83 Chapter

15 18. Draw a band diagrams to show the difference between(band gaps: C = 5.47, Si = 1.12, Ge = 0.66, Sn = 0) Conductor (Sn): 19. Draw a band diagram for thermal/photo (Intrinsic) and doped (Extrinsic) semiconductors and explain the origin of semicondictivity? Thermal/photo (Intrinsic) (Ge): Insulator (C): Doped (Extrinsic) (Si/As): Semiconductor (Ge): Chapter-2-85 Chapter Draw a band diagram for a p-type (Si/Ga) and n- type (Si/As) semiconductors and show holes and electrons that is responsible for semiconductivity. p-type(si/ga): 22. What the difference between a transistor (semiconductor device) and vacuum tube? n-type(si/as): Chapter-2-87 Chapter-2-88 What is a transistor? 21. What is a transistor with emitter (E), collector(c) and base (B), and how it works? Chapter-2-89 Chapter

16 23. Using the diagram explain how a diode work. Superconductors When Onnes cooled mercury to 4.15K, the resistivity suddenly dropped to zero Chapter-2-91 Chapter-2-92 The Meissner Effect Theory of Superconduction 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. 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 - electronphonon interaction: the electron is attracted to the positive charge density (red glow) created by the first electron distorting the lattice around itself. Chapter-2-93 Chapter-2-94 High Temperature Superconduction BCS theory predicted a theoretical maximum to Tc of around 30-40K. Above this, thermal energy would cause electronphonon interactions of an energy too high to allow formation of or sustain Cooper pairs 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. Chapter

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