1. About this Lecture Introduction to Organic Chemistry

2. About this lecture: Professor: Ping Lu （吕萍） Chemistry Department, Zhejiang Univ., Room 726, Building 8, Yuquan campus 0571-87952543(O), 13777847196 pinglu@zju.edu.cn Organic Chemistry Web address: http://jpkc.zju.edu.cn/k/146/ Platform course: 基础有机化学教学大纲 Textbook: McMurry, Organic Chemistry, 6th edition, Thomson press 王彦广等，有机化学，第二版，化工出版社

3. Tentative syllabus for 2008 fall (2 credits): WeekDateConcepts Teaching Hours 19.2-9.5Introduction Aliphatic Hydrocarbon 4 29.9-9.12Aliphatic Hydrpcarbon Stereochemistry 4 39.16-9.19Stereochemistry Unstaturated Hydrocarbon 4 49.23-9.28Unsaturated Hydrocarbon Radical Reaction 6 510.7-10.10Aromatic Hydrocarbon4 610.14-10.14Aromatic Hydrocarbon2 710.21-10.24Spectroscopy Alkyl Halide 4 810.28-10.31Alkyl Halide Alcohol, Phenol and Ether 4

4. Tentative syllabus for 2008 winter (2 credits): WeekDateConcepts Teaching Hours 111.11-11.14Aldehyde and Ketone4 211.18-11.21Aldehyde and Ketone Carboxylic Acid and its derivative 4 311.25-11.28Carboxylic Acid and its derivative4 412.2-12.5Carboxylic Acid and its derivative Nitrogen Containing Compounds 4 512.9-12.12Nitrogen Containing Compounds4 612.16-12.19Heterocyclic Compounds4 712.23-12.26Carbohydrates4 812.30-1.2Amino Acids Introduction to Organic Synthesis 4

5. Grading: Homework: 10 % Quiz: 30 % Final: 60 % Difficulties for students to learn: numerous compounds; numerous functional groups; numerous reactions. Suggestions for students to pass: preview the lecture; solve the problem in each section; write when you study; learn by teaching and studying; use the molecular model when you study. Monitor???

6. Introduction to Organic Chemistry

7. Requirements: 了解有机化学的发展史以及有机化学与生命科学的关系； 有机分子的结构：共价键、碳原子的特性； 有机化合物分子的表示法：结构式、投影式； 有机化合物中共价键：碳原子的杂化轨道、 σ 键和 π 键； 共价键的属性；键长、键角、键能、极性和极化度； 有机化合物结构和物理性质的关系，分子间作用力对溶解度、沸 点、熔点、比重的影响。 Reading Chapters: McMurry: Chapter 1 and 2 王彦广： Chapter 1 Homework: McMurry: 1.46, 1.47 王彦广：问题 1-2

8. 1.Organic compound and organic chemistry 1.1 Organic compound 1.2 Representation of organic structure 1.3 Fields in organic chemsitry 2.Structural Theory 2.1 History of structural theory 2.2 Atomic orbital theory 2.3 Hybrid orbital theory 2.4 Molecular orbital theory 3. Polar Covalent Bonds, Acid and Base 3.1 Polar Covalent Bonds: Electronegativity 3.2 Intermolecular attractive forces 3.3 Brønsted acid and base 3.4 Lewis acid and base 4. Classification of Organic Compounds 4.1 Carbon skeleton 4.2 Functional Groups

9. 1. Organic Compound and Organic Chemistry 1.1 Organic Compound Elements in organic compounds: C, H, O, X, N, P, S, Si… CaC 2, CO 2, CO… inorganic compound Definition of related terms: carbon compounds, hydrocarbons, isomerism Characters of organic compounds: combustion, solubility, volatile with relative lower b.p., relative lower m.p., covalent bond vs ionic bond, van der waals interaction vs electrostatic force, reaction speed and by-products F, Cl, Br, I

10. Circle of carbon compounds in nature

11. 1.2 Representation of organic structure Lewis formula: Kekule structure: Condensed formula ： Octet Rule Bond structure:

12. Space-fillingFramework Molecular Models

13. 1.3 Fields in organic chemsitry Basic organic chemistry —— source, structure, property, preparation and application of organic compounds; Organic synthesis —— propose a synthetic way and prepare it; Physical organic chemistry —— theoretically explain why and how the oragnic reaction goes; Organic structure analysis —— determine organic structure Natural products —— structural analysis and total synthesis Organometallic chemistry —— dealing with the compounds containing C-M bond Solid support for materials, life science, environmental…..

14. In 1828, Friedrich Wöhler

16. 2. Structural Theory structure —— property; composition —— connection 2.1 History of structural theory In 1852, Frankland: Element valence In 1858, Kekule and Couper: C is tetravalent, and could be connected in linear, branch and cyclic. single bond, double bond and triple bond In 1861, Structure and Property In 1874, van't Hoff and Le Bel: tetrahedron structure of C Note that a wedge indicates a bond is coming forward Note that a dashed line indicates a bond is behind the page

17. 2.2 Atomic orbital theory Quantum mechanics: describes electron energies and locations by a wave equation – Wave function solution of wave equation – Each Wave function is an orbital,  A plot of  2 describes where electron most likely to be Electron cloud has no specific boundary so we show most probable area

18. Four different kinds of orbitals for electrons based on those derived for a hydrogen atom Denoted s, p, d, and f s and p orbitals most important in organic chemistry s orbitals: spherical, nucleus at center p orbitals: dumbbell-shaped, nucleus at middle

19. Orbitals are grouped in shells of increasing size and energy Different shells contain different numbers and kinds of orbitals Each orbital can be occupied by two electrons

20. In each shell there are three perpendicular p orbitals, px, py, and pz, of equal energy Lobes of a p orbital are separated by region of zero electron density, a node

21. C: 1S 2 2S 2 2P 2 Atomic Structure: Electron Configurations Ground-state electron configuration: 1. Lowest-energy orbitals fill first: 1s  2s  2p  3s  3p  4s  3d (Aufbau (“build-up”) principle) 2. Electron spin can have only two orientations, up  and down . Only two electrons can occupy an orbital, and they must be of opposite spin (Pauli exclusion principle) to have unique wave equations 3. If two or more empty orbitals of equal energy are available, electrons occupy each with spins parallel until all orbitals have one electron (Hund's rule).

22. Valence Bond Theory Covalent bond forms when two atoms approach each other closely so that a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom Electrons are paired in the overlapping orbitals and are attracted to nuclei of both atoms – H–H bond results from the overlap of two singly occupied hydrogen 1s orbitals – H-H bond is cylindrically symmetrical, sigma (  ) bond

23. Bond Energy Reaction 2 H·  H 2 releases 436 kJ/mol bond strength

24. Bond Length Distance between nuclei that leads to maximum stability If too close, they repel because both are positively charged If too far apart, bonding is weak

25. 2.3 Hybrid orbital theory Carbon has 4 valence electrons (2s 2 2p 2 ) In CH 4, all C–H bonds are identical (tetrahedral) sp 3 hybrid orbitals: s orbital and three p orbitals combine to form four equivalent, unsymmetrical, tetrahedral orbitals (sppp = sp 3 ), Pauling (1931)

26. sp 3 orbitals on C overlap with 1s orbitals on 4 H atom to form four identical C-H bonds Each C–H bond has a strength of 438 kJ/mol and length of 110 pm Bond angle: each H–C–H is 109.5°, the tetrahedral angle.

27. sp 3 Orbitals and the Structure of Ethane

28. sp 2 Orbitals and the Structure of Ethylene 90  120 

29. sp Orbitals and the Structure of Acetylene

30. Summary of sp3, sp2 and sp: Shape Bond length Bond angle Bond strength

31. Hybridization of Nitrogen and Oxygen Elements other than C can have hybridized orbitals H–N–H bond angle in ammonia (NH 3 ) 107.3° N’s orbitals (sppp) hybridize to form four sp 3 orbitals One sp 3 orbital is occupied by two nonbonding electrons, and three sp 3 orbitals have one electron each, forming bonds to H

32. A molecular orbital (MO): where electrons are most likely to be found (specific energy and general shape) in a molecule Additive combination (bonding) MO is lower in energy Subtractive combination (antibonding) forms MO is higher 2.4 Molecular orbital theory

33. Molecular Orbitals in Ethylene The  bonding MO is from combining p orbital lobes with the same algebraic sign The  antibonding MO is from combining lobes with opposite signs Only bonding MO is occupied

34. 3.1 Polar Covalent Bonds: Electronegativity Covalent bonds can have ionic character These are polar covalent bonds 3. Polar Covalent Bonds, Acid and Base

35. Bond Polarity and Electronegativity Electronegativity (EN): intrinsic ability of an atom to attract the shared electrons in a covalent bond Differences in EN produce bond polarity F is most electronegative (EN = 4.0), Cs is least (EN = 0.7) Metals on left side of periodic table attract electrons weakly, lower EN Halogens and other reactive nonmetals on right side of periodic table attract electrons strongly, higher electronegativities EN of C = 2.5

36. Bond Polarity and Inductive Effect Nonpolar Covalent Bonds: similar EN Polar Covalent Bonds: Difference in EN of atoms < 2 Ionic Bonds: Difference in EN > 2 » C–H bonds, relatively nonpolar C-O, C-X bonds (more electronegative elements) are polar Bonding electrons toward electronegative atom  + and  - Inductive effect: shifting of electrons in a bond in response to EN of nearby atoms C——Cl -- ++

37. Polar Covalent Bonds: Dipole Moments Molecules as a whole are often polar: from vector summation of individual bond polarities and lone-pair contributions Strongly polar substances soluble in polar solvents like water; nonpolar substances are insoluble in water. Dipole moment - Net molecular polarity, due to difference in summed charges –  - magnitude of charge Q at end of molecular dipole times distance r between charges  = Q  r, in debyes (D), 1 D = 3.336  10  30 coulomb meter length of an average covalent bond), the dipole moment would be 1.60  10  29 C  m, or 4.80 D.

38. 3.2 Intermolecular attractive forces 3.2.1 Ion-Ion forces NaCl – electrostatic forces

39. Dynamics 70K 3.2.2 Hydrogen Bonds and dipole dipole forces

40. 3.2.3 van der Waal forces (London forces, dispersion forces)

41. Summary of Intermolecular Attractive Forces

42. Attraction Strength Ion-ion ~300 kJ/mol Induced dipole-dipole 50~200 kJ/mol Dipole-dipole 5~50 kJ/mol Hydrogen bonds 4~120 kJ/mol van der Waals < 5 kJ/mol, variable Summary of Intermolecular Attractive Forces

43. 3.3 Brønsted Acids and Bases “Brønsted-Lowry” is usually shortened to “Brønsted” A Brønsted acid is a substance that donates a hydrogen ion (H + ) A Brønsted base is a substance that accepts the H +

44. Quantitative Measures of Acid Strength The equilibrium constant (K e ) for the reaction of an acid (HA) with water to form hydronium ion and the conjugate base (A - ) is a measure related to the strength of the acid Stronger acids have larger K e Note that brackets [ ] indicate concentration, moles per liter, M. 55.6 M Molecular weight of water: 18 1 L water or 1000g water at 25 o C acidity constant, Ka

45. pK a – the Acid Strength Scale pK a = -log K a  G = -RT log K eq Smaller pKa, stronger acidity The pK a of water is 15.74

46. Organic Acids and Organic Bases The reaction patterns of organic compounds often are acid-base combinations The transfer of a proton from a strong Brønsted acid to a Brønsted base, for example, is a very fast process and will always occur along with other reactions Organic Acids: Organic Bases:

47. 3.4 Lewis acid and base _ Definition Lewis acids are electron pair acceptors and Lewis bases are electron pair donors The Lewis definition leads to a general description of many reaction patterns but there is no scale of strengths as in the Brønsted definition of pK a

48. 4. Classification of Organic Compounds 4.1 Carbon skeleton a)Acyclic or open chain compounds: These contain alkanes, alkenes, alkynes and their derivatives. These are also called aliphatic compounds. b) Cyclic or closed chain compounds: Cyclic compounds whose rings are made up of only one kind of atoms, i.e. carbon atoms are called homocyclic or cabocyclic compounds. Aliphatic cyclic compounds are called alicyclic compounds eg cyclopropane, cyclobutane etc

49. Organic compounds containing one or more fused or isolated benzene rings and their functinalized derivatives are called benzenoids or aromatic compounds, eg benzene, toluene, naphthalene, anthracene etc. Cyclic compounds containing one or more heteroatoms (usually O,N, S etc) are called hetrocyclic compounds eg ethylene oxide, tetrahydrofuran (THF), furan, pyrole etc

50. Functional group - collection of atoms at a site within a molecule with a common bonding pattern The group reacts in a typical way, generally independent of the rest of the molecule For example, the double bonds in simple and complex alkenes react with bromine in the same way 4.2 Functional Group

51. Alkenes have a C-C double bond Alkynes have a C-C triple bond Arenes have special bonds that are represented as alternating single and double C-C bonds in a six-membered ring Alkyl halide: C bonded to halogen (C-X) Alcohol: C bonded O of a hydroxyl group (C-OH) Ether: Two C’s bonded to the same O (C-O-C) Amine: C bonded to N (C-N) Thiol: C bonded to SH group (C-SH) Sulfide: Two C’s bonded to same S (C-S-C) Aldehyde: one hydrogen bonded to C=O Ketone: two C’s bonded to the C=O Carboxylic acid:  OH bonded to the C=O Ester: C-O bonded to the C=O Amide: C-N bonded to the C=O Acid chloride: Cl bonded to the C=O

52. Functional group Secondary suffix Functional group Secondary suffix Alcohols (-OH)-olAldehydes (-CHO)-al Ketones (>C=O)-one Carboxylic acids (- COOH) -oic acid Amines (-NH 2 )-amine Acid amides (- CONH 2 ) -amide Acid chlorides (- COCl) -oyl chlorideEsters (-COOR)-oate Nitrites (-C = N)-nitriteThioalcohols (-SH)-thiol *Delete e from alkane, adding with the following suffix

53. Functional group Prefix Functional group Prefix - COOHCarboxy-CHOFormyl -COOR Alkoxy cabonyl or Carbalkoxy >CO Oxo or Kelo -COCLChloroformyl-OHHydroxy -CONH 2 Carbamoyl-SHMecaplo -CNCyano-NH 2 Amino -ORAkoxy=NHImino -XHalo-NO 2 Nitro When the functional groups act as substituents, they are named as:

54. Organic chemistry – chemistry of carbon compounds Atom: positively charged nucleus surrounded by negatively charged electrons Electronic structure of an atom described by wave equation Electrons occupy orbitals around the nucleus s orbitals are spherical, p orbitals are dumbbell-shaped Covalent bonds - electron pair is shared between atoms Molecular orbital (MO) theory - bonds result from combination of atomic orbitals to give molecular orbitals, which belong to the entire molecule  bonds - Circular cross-section and are formed by head-on interaction  bonds – “dumbbell” shape from sideways interaction of p orbitals Alkane - sp3 hybrid orbitals Alkene - sp2 hybrid orbitals and one unhybridized p orbital Alkyne - sp hybrid orbitals with two unhybridized p orbitals Summary

55. Organic molecules often have polar covalent bonds as a result of unsymmetrical electron sharing caused by differences in the electronegativity of atoms The polarity of a molecule is measured by its dipole moment, . Intermolecular attractive forces: ion-ion electrostatic forces, hydrogen bonding, dipole dipole interaction, van der Waals forces A Brønsted(–Lowry) acid donatea a proton A Brønsted(–Lowry) base accepts a proton The strength Brønsted acid is related to the -1 times the logarithm of the acidity constant, pKa. Weaker acids have higher pKa’s A Lewis acid has an empty orbital that can accept an electron pair A Lewis base can donate an unshared electron pair In condensed structures C-C and C-H are implied