Organic Chemistry

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

2. ORAGANIC CHEMISTRY PART I 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. 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. Atomic Structure

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

7. Atomic Orbitals Electrons surrounding atoms are concentrated into regions of space called atomic orbitals. • 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. 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. 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. 1S2 2S2 3S2 4S2 5S2 6S2 7S2 Electronic Configurations of Atoms 2p6 3d10 3p6 4p6 4f14 4d10 5d10 5p6 5f14 6d10 6p6 6f14 7d10 7p6

11. Write electron configurations of Carbon atom 126C………………………………………………....... 2p 2s 1s

12. Molecular Orbitals • Covalent bond Electrostatic Interactions

13. Valences of Carbon • Carbon has four valence electrons (2s2 2p2), forming four bonds (CH4)

14. Valences of Nitrogen • Nitrogen has five valence electrons (2s2 2p3) but forms only three bonds (NH3)

15. Non-bonding electrons • Valence electrons not used in bonding are called nonbonding electrons, orlone-pair electrons • Nitrogen atom in ammonia (NH3) • Shares six valence electrons in three covalent bonds and remaining two valence electrons are nonbonding lone pair

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 (s) bond

17. Bond Energy • Reaction 2 H· H2 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. Bond energy • พลังงานพันธะ คือ พลังงานที่ใช้ในการสลายพันธะระหว่างอะตอมของธาตุภายในโมเลกุลที่อยู่ในสถานะก๊าซออกเป็นอะตอมเดี่ยว ๆ • เช่น • พลังงานพันธะเฉลี่ย คือ พลังงานเฉลี่ยที่ใช้สลายพันธะแต่ละพันธะในคู่อะตอมเดียวกัน D(C-H) = (1660/4) kJ/mol = 415 kJ/mol

19. Bond energy

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. Bond Lengths

22. Bond Lengths

23. Bond length and Bond strength

24. electron configurations of Carbon atom 126C Not CH2 CH4 2p 2s 1s Why? Hybridization

25. sp3 Hybridizationof Carbon Ground state Excited state sp3-hybridization state Hybridization Promotion of electron

26. Hybridization: sp3 Orbitals • sp3 hybrid orbitals:s orbital and three p orbitals combine to form four equivalent, unsymmetrical, tetrahedral orbitals (sppp = sp3), Pauling (1931)

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. The Structure of Ethane • Two C’s bond to each other by s overlap of an sp3 orbital from each • Three sp3 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. Hybridization of Nitrogen • Elements other than C can have hybridized orbitals • H–N–H bond angle in ammonia (NH3) 107.3° • N’s orbitals (sppp) hybridize to form four sp3 orbitals • One sp3 orbital is occupied by two nonbonding electrons, and three sp3 orbitals have one electron each, forming bonds to H

30. Hybridization of Oxygen • The oxygen atom is sp3-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. sp2 Hybridizationof Carbon Ground state Excited state sp2-hybridization state Hybridization Promotion of electron

32. 120 90 Hybridization: sp2 Orbitals • sp2 hybrid orbitals: 2s orbital combines with two 2p orbitals, giving 3 orbitals (spp = sp2) • sp2 orbitals are in a plane with120° angles; trigonal planar • Remaining p orbital is perpendicular to the plane

33. Bonds From sp2 Hybrid Orbitals • Two sp2-hybridized orbitals overlap to form a s bond • p orbitals overlap side-to-side to formation a pi () bond • sp2–sp2s bond and 2p–2p bond result in sharing four electrons and formation of C-C double bond • Electrons in the s bond are centered between nuclei • Electrons in the  bond occupy regions are on either side of a line between nuclei

34. The Orbital of Ethene

35. Bonding in Ethylene • H atoms form s bonds with four sp2 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. Hybridization: spOrbitals • 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. Orbitals of Acetylene • Two sp hybrid orbitals from each C form sp–sps bond • pz orbitals from each C form a pz–pz bond by sideways overlap and py orbitals overlap similarly

38. Orbitals of Acetylene

39. Bonding in Acetylene • Sharing of six electrons forms C ºC • Two sp orbitals form s bonds with hydrogens

40. Bond Polarity

41. H2, Cl2: HCl: Polarity • Polarity refers to a separation of positive and negative charge. • In a nonpolar bond, the bonding electrons are shared equally. • In a polar bond, electrons are shared unequally.

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

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

44. Bond Polarity

45. Molecule Polarity

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. Intramolecular forces

48. Stearic replusion

49. Intramolecular Hydrogen Bond

50. Intermolecular forces