1. Organic Chemistry MS.SUPAWADEE SRITHAHAN DEPARTMENT OF CHEMISTRY MAHIDOL WITTAYANUSORN SCHOOL

2. 2 CONTENTS  INTRODUCTION  CLASSIFICATION NAMING AND PROPERTIES OF ORGANIC COMPOUND  BONDING OF ORGANIC COMPOUND  ALKANE & CYCLOALKANE  ALKENE & CYCLOALKENE  ALKYNE & CYCLOALKYNE

3. INTRODUCTION Structure and Bonding

4. 4 Organic Chemistry “Organic” – until mid 1800’s referred to compounds from living sources (mineral sources were “inorganic”) Wöhler in 1828 showed that urea, an organic compound, could be made from a minerals Today, organic compounds are those based on carbon structures and organic chemistry studies their structures and reactions Includes biological molecules, drugs, solvents, dyes Does not include metal salts and materials (inorganic) Does not include materials of large repeating molecules without sequences (polymers)

5. 5 Atomic Structure

6. 6 Shells 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

7. 7 Atomic Orbitals Electrons surrounding atoms are concentrated into atomic orbitals. regions of space called atomic orbitals. s, p, d, and f Four different kinds of orbitals ; 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

8. 8 p-Orbitals 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

9. 9 Electron Configurations Ground-state electron configuration of an atom lists orbitals occupied by its electrons. Rules: 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).

10. 10 Electronic Configurations of Atoms 1S 2 2S 2 3S 2 4S 2 5S 2 6S 2 7S 2 2p 6 3p 6 4p 6 5p 6 6p 6 7p 6 3d 10 5d 10 6d 10 7d 10 4d 10 4f 14 6f 14 5f 14

11. 11 Write electron configurations of Carbon atom 12 6 C………………………………………………....... 1s1s 2s2s 2p2p

12. 12 Molecular Orbitals Covalent bond Electrostatic Interactions

13. 13 Valences of Carbon Carbon has four valence electrons (2s 2 2p 2 ), forming four bonds (CH 4 )

14. 14 Valences of Nitrogen Nitrogen has five valence electrons (2s 2 2p 3 ) but forms only three bonds (NH 3 )

15. 15 Non-bonding electrons Valence electrons not used in bonding are called nonbonding electrons, or lone-pair electrons Nitrogen atom in ammonia (NH 3 ) Shares six valence electrons in three covalent bonds and remaining two valence electrons are nonbonding lone pair

16. 16 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

17. 17 Bond Energy Reaction 2 H·  H 2 releases 436 kJ/mol Product has 436 kJ/mol less energy than two atoms: H–H has bond strength of 436 kJ/mol. (1 kJ = 0.2390 kcal; 1 kcal = 4.184 kJ)

18. 18 Bond energy พลังงานพันธะ คือ พลังงานที่ใช้ในการสลาย พันธะระหว่างอะตอมของธาตุภายในโมเลกุลที่อยู่ ในสถานะก๊าซออกเป็นอะตอมเดี่ยว ๆ เช่น พลังงานพันธะเฉลี่ย คือ พลังงานเฉลี่ยที่ใช้สลาย พันธะแต่ละพันธะในคู่อะตอมเดียวกัน D(C-H) = (1660/4) kJ/mol = 415 kJ/mol

19. 19 Bond energy

20. 20 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

21. 21 Bond Lengths

22. 22 Bond Lengths

23. 23 Bond length and Bond strength

24. 24 electron configurations of Carbon atom 12 6 C Not CH 2 CH 4 1s1s 2s2s 2p2p Why? Hybridization

25. 25 sp 3 Hybridization of Carbon Ground stateExcited statesp 3 -hybridization state Hybridization Promotion of electron

26. 26 Hybridization: sp 3 Orbitals sp 3 hybrid orbitals: s orbital and three p orbitals combine to form four equivalent, unsymmetrical, tetrahedral orbitals (sppp = sp 3 ), Pauling (1931)

27. 27 Tetrahedral Structure of Methane sp3 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.

28. 28 The Structure of Ethane Two C’s bond to each other by  overlap of an sp 3 orbital from each Three sp 3 orbitals on each C overlap with H 1s orbitals to form six C–H bonds C–H bond strength in ethane 420 kJ/mol C–C bond is 154 pm long and strength is 376 kJ/mol All bond angles of ethane are tetrahedral

29. 29 Hybridization of Nitrogen 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

30. 30 Hybridization of Oxygen The oxygen atom is sp 3 -hybridized Oxygen has six valence-shell electrons but forms only two covalent bonds, leaving two lone pairs The H–O–H bond angle is 104.5°

31. 31 sp 2 Hybridization of Carbon Ground stateExcited statesp 2 -hybridization state Hybridization Promotion of electron

32. 32 Hybridization: sp 2 Orbitals sp 2 hybrid orbitals: 2s orbital combines with two 2p orbitals, giving 3 orbitals (spp = sp 2 ) trigonal planar sp 2 orbitals are in a plane with120° angles; trigonal planar Remaining p orbital is perpendicular to the plane 90  120 

33. 33 Bonds From sp 2 Hybrid Orbitals Two sp 2 -hybridized orbitals overlap to form a  bond p orbitals overlap side-to-side to formation a pi (  ) bond sp 2 –sp 2  bond and 2p–2p  bond result in sharing four electrons and formation of C-C double bond Electrons in the  bond are centered between nuclei Electrons in the  bond occupy regions are on either side of a line between nuclei

34. 34 The Orbital of Ethene

35. 35 Bonding in Ethylene H atoms form  bonds with four sp 2 orbitals H–C–H and H–C–C bond angles of about 120° C–C double bond in ethylene shorter and stronger than single bond in ethane Ethylene C=C bond length 133 pm (C–C 154 pm)

36. 36 Hybridization: sp Orbitals C-C a triple bond sharing six electrons Carbon 2s orbital hybridizes with a single p orbital giving two sp hybrids two p orbitals remain unchanged sp orbitals are linear, 180° apart on x-axis Two p orbitals are perpendicular on the y-axis and the z-axis

37. 37 Orbitals of Acetylene Two sp hybrid orbitals from each C form sp–sp  bond pz orbitals from each C form a p z –p z  bond by sideways overlap and p y orbitals overlap similarly

38. 38 Orbitals of Acetylene

39. 39 Bonding in Acetylene Sharing of six electrons forms C  C Two sp orbitals form  bonds with hydrogens

40. 40 Bond Polarity

41. Polarity Polarity refers to a separation of positive and negative charge. In a nonpolar bond, the bonding electrons are shared equally. HCl:  In a polar bond, electrons are shared unequally. H 2, Cl 2 :

42. Electronegativity Electronegativity refers to the ability of an atom in a molecule to attract shared electrons. The Pauling scale of electronegativity:

43. Bond Polarity A polar bond can be pictured using partial charges:  = 0.9 Electronegativity Difference Bond Type 0 - 0.5 Nonpolar 0.5 - 2.0Polar 2.0  Ionic 2.1 3.0 ++  H Cl

44. 44 Bond Polarity

45. 45 Molecule Polarity

46. 46 Types of Interactions 1. Intramolecular force Covalent bond Ionic bond Metallic bond Stearic replusion Intramolecular Hydrogen Bond 2. Intermolecular force Van de Waals force Hydrogen bond

47. 47 Intramolecular forces

48. 48 Stearic replusion

49. 49 Intramolecular Hydrogen Bond

50. 50 Intermolecular forces

51. 51 Ion – Dipole Forces * Between a charged ion and polar molecule

52. 52 Dipole-Dipole forces * Between neutral polar molecule

53. 53 London dispersion forces * Between non polar/non polar molecules

54. 54 Hydrogen bonding * Between hydrogen and an electronegative atom such as F, O or N

55. 55 Structural of organic compounds 1. Dot structure 2. Dash Formula 3. Condensed formula 4. Partial Condensed Formula CH 3 CH 2 CH 2 COOH

56. 56 Structural of organic compounds 5. Line-angle formula or bond line formula CH 3 CH 2 CH 2 COOH

57. 57 6. Three-dimensional formulas Bonds that project upward out of the plane of the paper Bonds that lie behind the plane Bonds that lie in the plane of the page

58. 58 Sample Problem Rewrite each of the following condensed structural formulas, as dash formulas as :

59. 59 Write dash formulas for each of the following bond-line formulas: