What we have learned so far toward molecular structure and properties .

1. What we have learned so far toward molecular structure and properties. Model of the atom The nuclear model of atom Thompson’s model Rutherford’s model Atomic spcetra -- Bohr’s model Quantum theory Periodicity of atomic properties Wave-particle duality Uncertainty principle Wave function – particle in a box Schroedinger equation Atomic radius Ionic radius Ionization energy Electron affinity Periodic table Hydrogen atom (one electron atoms) Principle quantum number Atomic orbitals: radial wave function angular wave function orbital angular momentum magnetic quantum numbers radial distribution function shape of atomic orbitals electron spin Many electron atoms orbital energy split shieldingeffect effectivenuclear charge Pauli exclusion principle Hund’s rule valence shell (electrons)

2. What we have learned so far toward molecular structure and properties. Chemical bond Interaction between two electrons Ionic bond Lewis structure Octet rule Exceptions to octet rule Resonance Formal charge Oxidation number Covalent bond electronegativity Ionic v.s. covalent Dipole moment Polar bond Nonpolar bond Bond strength Bond length IR(infrared) spectroscopy MOLECULAR SHAPE AND STRUCTURE Stability, reactivity, color, size, polarity, solubility, function etc…

3. 3D structure of a molecule is crucial for its property. Sophisticated quantum mechanical calculations are needed to predict the structure. Box 3.1 ⇒Drugsby Design and Discovery 1) Identification of key enzymes 2) Molecular structure determination 3) Hints from Nature --- Natural Products 4) Computer-aided design of molecules with structures fitting into the active site

4. Chapter 3. MOLECULAR SHAPE AND STRUCTURE THE VSEPR MODEL (전자쌍 반발 모델) 3.1 The Basic VSEPR Model 3.2 Molecules with Lone Pairs on the Central Atom 3.3 Polar Molecules VALENCE-BOND THEORY (원자가 결합 이론) 3.4 Sigma and Pi Bonds 3.5 Electron Promotion and the Hybridization of Orbitals (혼성궤도 함수) 3.6 Other Common Types of Hybridization 3.7 Characteristics of Multiple Bonds 2012 General Chemistry I

5. 95 THE VSEPR MODEL (Sections 3.1-3.3) Lewis structure: showing the linkages between atoms and the presence of lone pairs, but not the 3D arrangement of atoms ClF3 CH4 NH3 H2O BF3 BeCl2 SF4 XeF4 IF5 PCl5 SF6

6. Estimating the 3D structure: THE VSEPR MODEL 3.1 The Basic VSEPR Model Valence Shell Electron-Pair Repulsion theory Electron pairs (lone pairs & bonding pairs) repel each other. Proposed by R. J. Gillespie in 1959. Rule 1: Electron pairs move as far apart as possible. VSEPR structures for AXn with no lone pair

7. BF3 CH4 BeCl2

8. SF6 PCl5

9. Rule 2: (Almost) No distinction between single and multiple bonds. BeCl2 CO2 CO32- BF3

10. 98 3.2 Molecules with Lone Pairs on the Central Atom • Rule 3 All regions of high electron density, lone pairs and bonds, are included in a description of the electronic arrangement, • But only the positions of atoms are considered when identifying the shape of a molecule. NH3 CH4

11. 99 • Rule 4 The strength of repulsions are in the order • lone pair-lone pair > lone pair-atom > atom-atom H2O NH3

12. 99 ClF3 SF4 PCl5 axial equatorial axial equatorial T-shaped seesaw shaped more stable

13. 101 • Predicting a molecular shape of XeF4 Step 1 Draw the Lewis structure. Step 2 Assign the electron arrangement around the central atom. Step 3 Identify the molecular shape. AX4E. Step 4 Allow for distortions. Square planar

14. 95

15. AXE method A; central atom X; outside atom E; lone pair

16. 102 3.3 Polar Molecules • Polar molecule: a molecule with a nonzero dipole moment i.e. HCl with a dipole moment of 1.1 D HCl, H2O, CHCl3, cis-dichloroethane, ··· - A polyatomic molecule is polar if it has polar bonds arranged in space in such a way that the dipole moments associated with the bonds do not cancel. polar polar

17. 102 3.3 Polar Molecules • Nonpolar molecule: a molecule with a net zero dipole moment Homonuclear diatomic molecules Polyatomic molecules with symmetry; CO2, BF3, CH4, CCl4, trans-dichloroethane, ··· nonpolar nonpolar

18. 103

19. 105 VALENCE-BOND THEORY (Sections 3.4-3.7) Lewis model of the chemical bond; localized electron model Valence-bond theory; Walter Heitler, Fritz London (1927) Linus Pauling (1931) Quantum mechanical description of the distribution of electrons in bonds Valence electrons are localized either between pairs of atoms or on atoms as lone pairs. Hybridization of atomic valence orbitals with proper symmetry that are localized between pairs of atoms. 2) Placing valence electrons in the hybridized orbitalsas pairs (↑↓) or leaving them localized in lone-pair orbitals on individual atoms in the molecule. VSEPR theory is a simplified one : powerful way of predicting the shape of simple molecules. --- does not explain many things including multiple bond, bond angles ……..

20. 105 3.4 s(Sigma) and p(Pi) Bonds: description of covalent bond Walter Heitler, Fritz London (1927) H2 1) Two hydrogen 1s-orbitals merge (overlap) to form a s-orbitalbetween the two hydrogen atoms. 2) A s-bond is formed as two electrons (↑↓) fill the s-orbital. s-bond ; cylindrically symmetrical with no nodal planes containing the intermolecular axis. i.e. H2, 1s-1s HF, 1s-2pz N2, 2pz-2pz

21. 106 overlap with 1:1 match of orbitals

22. 106 overlap with 1:1 match of orbitals

24. 106 • p-bond • nodal plane containing the interatomic (bond) axis - two cylindrical shapes (lobes), one above and the other below the nodal plane N2 one s-bond with two perpendicular p-bonds - multiple bonds: single bond (one s-bond), double bond (one s- and one p-bond) triple bond (one s- and two p-bonds)

25. 107 Polyatomicmolecules Linus Pauling (1931) BeH2 3.5 Electron Promotion and the Hybridization of Orbitals Why do we have to make hybrid orbitals?

26. not real. localization problem Linear combination of orbitals sp hybrid

27. Linear molecule

28. 107 BH3

29. sp2 hybrid orbitals

30. 107 CH4 promotion hybridization

31. 107 sp3 hybrids h1 = s + px + py + pz h3= s - px + py- pz h2 = s - px - py + pz h4 = s + px - py - pz similar ideas to VSEPR

32. 107 NH3 hybridorbitals can be determined by the steric number based on the VSEPR model. steric number = # of atoms bonded to the central atom + # of lone pairs H2O

33. 109 - sp3d hybrid orbitals in PCl5 - sp3d2 hybrid orbitals in SF6 and XeF4

34. C2H6

35. 111 3.7 Characteristics of Multiple Bonds CO2 Carbon Steric # = 2 Oxygen Steric # = 3

36. 111 3.7 Characteristics of Multiple Bonds CO2

37. 111 3.7 Characteristics of Multiple Bonds - ethene, CH2=CH2 C-C s bond, s(C2sp2, C2sp2) C-C p bond, p(C2p, C2p) each C-H bond formed as s(C2sp2, H1s) restricted rotation

38. 112 - ethyne (acetylene), C2H2 free rotation

39. 112 - benzene, C6H6 Now, delocalization has the meaning ! Still, there are many properties that can not be explained by the current model.

40. 111 CO32- resonance hybrid

41. 3.8 The Limitations of Lewis’s Theory & Valence Bond Theory 113 Paramagnetic O2; unpaired electron(s) Lewis's theory; Valence-bond theory; bond and bond 2 lone pairs on each O occupying the sp2 hybrid orbitals Paramagnetic: tendency to move into the magnetic field. When there are unpaired electrons in the molecule. Diamagnetic: tendency to move out of the magnetic field. When all the electrons in the molecules are paired.

42. 113s 3.8 The Limitations of Lewis’s Theory & Valence Bond Theory Shortcomings of the Valence Bond Model • Inadequate treatment of odd-electron moleculesand resonances O2 N2 2) Magnetism of molecules • Paramagnetic: molecules with unpaired electrons • Diamagnetic: weakly repelled by a magnetic field both are expected to be diamagnetic!!

43. 113 3.8 The Limitations of Lewis’s Theory & Valence Bond Theory Electron deficient diborane, B2H6 does not have enough electrons ! First published by H. C. Lunguet-Higgins, a 2nd year undergraduate student ! At least seven bonds (= 14 electrons) are required, but only 12 valence electrons. - No simple explanation for spectroscopic properties of compounds

44. 113 MOLECULAR ORBITAL THEORY(Sections 3.8-3.12) 3.9 Molecular Orbitals 3.10 Electron Configurations of Diatomic Molecules 3.11 Bonding in Heteronuclear Diatomic Molecules 3.12 Orbitals in Polyatomic Molecules • Molecular Orbital (MO) theory advantages - Addresses all of the above shortcomings of VB theory - Provides a deeper understanding of electron-pair bonds - Accounts for the structure and properties of metals and semiconductors - Universally used in calculations of molecular structures

45. 115s H2+: Prototype Molecular Orbital System • Atomic orbital(AO) theory→ successful for orbital structures of all atoms with both even and odd numbers of electrons • Assume that molecule H2+ ~ as an united atomwith a fragmented nucleus if the nuclei in molecule were fused together •construct the one-electron orbital corresponding to the arrangement of nuclear charges presented by the molecule Coulomb interactions in H2+

46. 115 3.9 Molecular Orbitals Quantum mechanics : the ideal solution to the problem, but……. Even for the smallest molecule, H2 Schroedinger equation will look like….. and way too complex and complicated… So, need simplification ! Simplification 1 (Born - Oppenheimer Approximation)Simplification 2 (Orbital Approximation)Simplification 3 (LCAO Approximation)

47. 115 3.9 Molecular Orbitals The valence-bond (VB) and molecular orbital (MO) theories are both procedures for constructing approximate wavefunctionsof electrons. - In VB theory, bonding electrons are localized on atoms or between pairs of atoms. • Molecular orbitals (MOs) The MO theory can account for electron-deficient compounds, paramagnetic O2, and many other properties by focusing on electrons delocalized over the whole molecule. • MOs formed by linear combination of atomic orbitals (LCAO-MO) Approximate molecular wavefunctions by superimposing (mixing) ofN atomic orbitals cijand Ej are determined by solving the Schrödinger equation

48. Trial wavefunctions for H2using two 1s atomic orbitals of H Increased amplitude in the internuclear region bonding Larger volume for electrons lower kinetic energy (particle-in-a-box) Decreased amplitude in the internuclear region & nodal planeantibonding

49. 115 • Molecular orbital energy-level diagram - relative energies of original AOs and resulting MOs - arrows to show electron spin and location of the electrons - In H2, two 1s-orbitals merge to form the bonding orbital s1s and the antibonding orbital s1s*

50. 116 3.10 Electron Configurations of Diatomic Molecules Building-up principle for MO • Valence electrons in molecular orbitals • 1. Electrons are accommodated in the lowest-energy MO, then in • orbitals of increasingly higher energy. • 2. Pauli exclusion principle: • each MO can accommodate up to two electrons. • If two electrons are present in one orbital, they must be paired. • 3. Hund’s rule: • If more than one MO of the same energy is available, • the electrons enter them singly and adopt parallel spins. H2: The energy of H2 is lower than that of the separate H atoms. Even the energy of H2+ is lower than that of the separate H atoms.