slide1 n.
Download
Skip this Video
Loading SlideShow in 5 Seconds..
Instructor: Prof. Duncan Wardrop Time/Day: T, 4:30-8:15 p.m. September 9, 2010 PowerPoint Presentation
Download Presentation
Instructor: Prof. Duncan Wardrop Time/Day: T, 4:30-8:15 p.m. September 9, 2010

play fullscreen
1 / 73

Instructor: Prof. Duncan Wardrop Time/Day: T, 4:30-8:15 p.m. September 9, 2010

130 Views Download Presentation
Download Presentation

Instructor: Prof. Duncan Wardrop Time/Day: T, 4:30-8:15 p.m. September 9, 2010

- - - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

  1. Welcome to CHEM 494 Instructor: Prof. Duncan Wardrop Time/Day: T, 4:30-8:15 p.m. September 9, 2010

  2. Course Website http://www.chem.uic.edu/chem494 • Syllabus • Course Policies • Other handouts • Announcements (Course News) • Course Calendar

  3. Keys to Success • 1. Attend all lectures and discussion sections. • 2. Don’t fall behind. Organic chemistry is easy, but each topic builds on the previous. • 3. Don’t memorize. Organic chemistry is conceptual. • 4. Always ask yourself “Why” for everything you read or hear. You may not always find the answer, but just asking will help you to find connections and remember more.

  4. Origins and Examples of Organic Chemistry Introduction

  5. Vitalism Living Nonliving. vs. • posses vital force • compounds derived from are “organic” (coined by J. Berzelius, 1807) • could not be synthesized in the laboratory • termed “inorganic” • derived from nonliving matter • can be synthesized in the lab

  6. Wohler Synthesis Debunks Vitalism F. Wohler (Germany, 1828) AgOCN(aq) + NH4Cl(aq) AgCl(s) + NH4OCN(aq) X Expected Product Actual Product ammonium cyanate (inorganic) urea (organic)

  7. Self Test Question What is the minimum requirement, today, for a chemical substance to be classified as organic? A. derived from living matter B. contains carbon C. cannot be synthesized D. bought at Whole Foods E. combustion yields SO2

  8. Vitalism Lives On. . . terpinen-4-ol sold as tea tree oil or Melaleuca oil fructose “fruit sugar” also in HFCS

  9. Organic Chemistry Everywhere

  10. Natural Products Order: coleoptera Family: coccinellidae Order: Tetraodontoformes

  11. Pharmaceuticals Morphine (analgesic) Tamiflu (influenza) Celebrex (arthritis)

  12. Optics - Transition Lenses? does not absorb visible light UV light heat does absorb visible light

  13. Atomic Structure A General Chemistry Review

  14. Atomic Composition - + + + + - - - Bohr Model

  15. Atomic Number & Mass Number X A A = protons + neutrons Z = protons Z Li 6 - + + + 3 - For neutral molecules, the number of electrons equals the number of protons. -

  16. Self Test Question How many neutrons are in the following atom: 14 C 6 A. 14 B. 6 C. 8 D. 20 E. cannot be determined

  17. Tenets of Schrӧdinger Equation • electrons have wave properties • wave equation gives energy of electron at a location • solutions to wave equation are wave functions (Ψ); a.k.a orbitals • probability distribution = Ψ2 (Heisenburg uncertainty principle)

  18. Probability Distribution vs. Boundary Surface 1s 1s Probability distribution (Ψ2) (“electron cloud”) Boundary Surface (1s orbital) (where the probability = 90-95%)

  19. Wave Function Values:Quantum Numbers • Shrodinger equation → wave function (orbital, Ψ) • many solutions for Ψ, energies of electrons in atom

  20. s-Orbitals • spherically symmetric • possible for n≥1 • 1 s-orbitals for each value of n • 1s = no nodes, 2s = 1 nodes, 3s = 2 nodes, etc. • probability of finding s-electron at nodal surface = 0 • s-orbital energy increases with increasing nodes (as n increases)

  21. p-Orbitals • shaped like dumbells • not possible for n = 1 (n≥2) • 3 p-orbitals for each value of n; they are degenerate (equal in energy) • wave function changes sign at the nucleus (node) • probability of finding p-electron at nodal plane = 0 • higher in NRG than s-orbitals of the same shell

  22. Relative Energies of Orbitals 3d 3d 3d 3d 3d 4s n = 3/4 3p 3p 3p 3s n = 2 Energy 2p 2p 2p 2s n = 1 1s

  23. Order of Orbitals from Periodic Table Read left to right starting at the top left; just like a typewriter. 1s 2s 2p 3s 3p 4s 3d 4p 4d 5p 5s 6p 6s 5d 7s 6d 7p 4f 5f

  24. Electron Configuration: Filling Orbitals 3d 3d 3d 3d 3d 4s 3p 3p 3p n = 3/4 3s spin (ms = +1/2) n = 2 Energy 2p 2p 2p spin (ms = –1/2) 2s • Aufbau principle: fill lowest energy orbitals first. • Pauli exclusion principle: two electrons with same quantum numbers cannot occupy a single orbital. Electrons must have opposite spins in the same orbital. • Hund’s rule: degenerate orbitals filled singly first. n = 1 1s

  25. Self Test Question Which of the following orbitals is highest in energy? A. 3f B. 3p C. 2s D. 3s E. 2d

  26. Self Test Question Which is the correct electron configuration for carbon? A. 1s2, 2s8 B. 1s2, 2s4 C. 1s2, 2s2, 2s2 D. 1s2, 2s2, 2p2 E. none of the above

  27. Bonding and Molecular Structure

  28. What is a Covalent Chemical Bond? valent = bonding covalent = electrons shared between two nuclei • Forces involved: • electron-electron repulsion • nucleus-nucleus repulsion • electron-nucleus attraction electron-electron repulsion < electron-neutron attraction

  29. Three Models of Bonding • Lewis Model: • atoms gain, lose or share electrons in order to achieve a closed-shell electron configurations (all orbitals fully occupied) • closed-shell for row 2 = 8 electrons (octet rule) • only valence electrons (in outermost shell) are involved in bonding Molecular Lewis Structures Molecular Lewis Structures Atomic Structures H–H F–F

  30. Three Models of Bonding • Valence Bond Model: • in-phase overlap of two half-filled orbitals • in-phase = constructive interference • increases probability electrons between two nuclei Boundary Surfaces H–H H H Electrostatic Potential maps: red = negative blue = postive

  31. Three Models of Bonding • Molecular Orbital (MO) Model: • combine atomic orbitals (AO) of all atoms, then extract molecular orbitals • number of atomic orbitals in equals the number of molecular orbitals out • combination possibilites: • additive = produces bonding molecular orbitals • subtractive = produces antibonding molecular orbitals

  32. Molecular Orbital Diagram antibonding (σ*) MO nodal plane e- probability = 0 = no bond H H 1s AO 1s AO Rules Still Apply: 1. Aufbau 2. Pauli 3. Hund bonding (σ) MO

  33. Self Test Question Which is the correct Lewis dot structure for CH4O? D. A. B. E. none of the above C.

  34. Structural, Bond-line and Condensed Formulas CH3CH2CH2CHO CH3CH(OH)CH3 (CH3)2CHCH2CH2CO2H

  35. VSEPR: Quick Review

  36. Self Test Question Which bond-line drawing correctly represents the following condensed formula? CH3CH(OCH2CH3)CH2Br D. A. B. E. C.

  37. Functional Groups

  38. Memorize These Functional Groups NOW • functional group: a defined group of atoms with a specific connectivity • responsible for properties • predictable reactivity • well-defined nomenclature Memorizing: “What is the minimum number of atoms needs to define a functional group and in what order are they bonded?” Flashcards! Table 4.1, page 140

  39. Formal Charge - Example 1 • My logic (no memorizing equations): • Recognize that opposing charges cancel each other • Determine difference in protons vs. electrons by asking: • How many valence (outer shell) electrons does the atom have? • How many valence (outer shell) electrons does the atom “want” (groups #)? • If atom hasmore than it “wants,” negatively charged. • If atom hasless than it “wants,” positively charged. O 8 Has? = 16 2– Wants? = 6 8 Difference? = 2 extra

  40. Formal Charge - Example 2 • My logic (no memorizing equations): • Recognize that opposing charges cancel each other • Determine difference in protons vs. electrons by asking: • How many valence (outer shell) electrons does the atom have? • How many valence (outer shell) electrons does the atom “want” (groups #)? • If atom hasmore than it “wants,” negatively charged. • If atom hasless than it “wants,” positively charged. 5 Has? = Wants? = 6 Difference? = 1 less “formal” = count one electron for each bond

  41. Shortcut for Some 1st and 2nd Period Atoms • For each additional bond = +1 • For each “missing” bond = –1

  42. Self Test Question What is the formal charge on the carbon atom in red? A. -2 B. -1 C. 0 D. +1 E. +2

  43. Self Test Question According to VSEPR, what is the bond angle between 2 H-atoms in methane (below)? • 105º • 109.5º • 107º • 180º • 120º

  44. Structure and Bonding: Resonance Section: 1.8, 1.11 You are responsible for Section 1.10.

  45. The Contradictory Case of Ozone Lewis Structure suggests ozone is non-symmetrical electrostatic potential map red = negative blue = positive Microwave spectroscopy shows that ozone is symmetrical (C2v)

  46. Curious Case of Benzene Predicted Actual resonance C–C bond length: 150 pm C=C bond length: 134 pm All bonds = 140 pm

  47. Resonance Provides the Solution • Problems: • Lewis structures fail to describe actual atom electron densities for molecules with more than one possible electron distribution. • Lewis formulas show electrons as localized; they either belong to a single atom (lone pair) or are shared between two atoms (covalent bond). • Electrons are not always localized; delocalization over several nuclei leads to stabilization (lower energy).

  48. Physical Proof of Resonance in Amides C-N (amide) Bond Length = 133 pm C-N (amine) Bond Length = 147 pm C=N (imine) Bond Length = 129 pm

  49. A Question of Arrows Equilibrium between distinct species Reaction from one species to another “Movement” of electrons from donor to acceptor Indicates that 2 species are contributing resonance structures

  50. Resonance • Solutions: • Actual molecule is considered a resonance hybrid (weighted average) of contributing Lewis structures. (Note: double headed arrow does NOT indicate interconversion.) • Dashed-line notation is sometimes used to indicate partial bonds. • Not all structures contribute equally.