1 Organic Chemistry, Third Edition Janice Gorzynski Smith University of Hawai’i Chapter 1 Lecture Outline Copyright © 2011 The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Prepared by Layne A. Morsch The University of Illinois - Springfield

What is Organic Chemistry? The chemistry of compounds containing ‘Carbon’ element The scope of organic chemistry –Molecular structure: How e - forms bond –Reaction: transformation from one structure to another Reaction mechanism: description of reaction pathway at an electronic level “Relatiobship between structure and reactivity” 2

Chemical Reaction Covalent bond cleavage - homolysis: A:B  A. + B. - heterolysis: A:B  A + + B: - A:B  A + + B: - Covalent bond formation (e - sharing) - A. + B.  A:B - A + + B: -  A:B 3

4 Atomic Structure

5 Atomic number: # of protons in the nucleus Atomic mass: # of protons + # of neutrons 12 6 C : 6 protons & 6 neutrons atomic # is 6, atomic mass is 12. In a neutral atom, # of protons = # of electrons. For ions - cations: positively charged, fewer electrons than its neutral form. - anions: negatively charged, more electrons than the neutral form. Atomic Structure

6 Isotopes Carbon Element (C) Isotopes: the same element, having a different number of neutrons.

Atomic Orbitals 1. A region of space where the probability of finding e - with a particular energy is large 2. Relative energy level 7 E p - 3s p - 2s - 1s Organic Chemistry: the elements of 1 st and 2 nd row (period)

8 s orbital: a sphere of electron density. Atomic Shape p orbital: a dumbbell shape and contains a node (no e - density) at the nucleus.

Elements in the same row: similar in size. 9 The Periodic Table Figure 1.2 Elements in the same column: similar electronic and chemical properties.

10 only one orbital in the first shell. hold a maximum of two e -. 2 elements: H and He Periodic Table The First Row

11 The Second Row Elements 8 elements, obtained by adding electrons to the 2s and three 2p orbitals. Periodic Table Four orbitals (2 e - each) holds max 8 e -. a maximum 8 valence e -.

12 4 orbitals: one 2s orbital, and three 2p orbitals. Periodic Table The Second Row

13 Joining of two atoms in a stable arrangement. Ionic or covalent bonds to attain a complete outer shell of valence e - (octet rule for second row elements). Bonding Ionic bonds: transfer of e - from one element to another. Covalent bonds: sharing of e - between two nuclei.

14 Ionic Bonding Sodium chloride, Na + Cl -

15 Ex. Li + F - - Li loses one e - to make Li + - F gains one e - to make F - Ionic Bonding

16 Covalent Bonding Covalent bond: e- sharing, between elements having similar electronegativity (e.g., CH 4, H 2, Cl 2 ). A compound with covalent bonds is called a molecule. ex) Bonding in molecular hydrogen (H 2 )

17 Second-row elements: max 8 valence e - around them  Octet rule Valence Electrons 1.The outmost electrons 2.Loosely held 3.Participate in chemical reactions 4.The group # of 2 nd row elements=#of valence electrons

18 Elements in Groups 2A and 3A Elements in the Third Row Exceptions to the Octet Rule (2 vacant orbitals) (one vacant orbital)

19 Shared and Unshared Electrons - Shared e - : bonding e - - Unshared e - : ‘nonbonding (nonbonded) e - ’ or ‘lone pairs’

20 Lewis structures: e - dot representations for molecules. ( A solid line = covalent bond (2 e - ). 1.Draw only the valence e -. 2.Second-row elements: no more than 8 e - (octet rule). 3.Hydrogen: 2 e -. Lewis Structures

21 Drawing a Lewis Structure 1. Step 1. Arrange atoms next to each other Step 2. Count the valence e -. Step 3. Arrange the e - around the atoms. Step 4. Assign formal charges to all atoms. Step [1]

22 Step [2] Count the electrons. 1) Count the number of valence electrons from all atoms. Step [3] Arrange the electrons around the atoms. 1) Place a bond between every two atoms (2 electrons). 2) Remaining electrons  lone pairs. Step [4] Assign formal charges to all atoms. Draw a Lewis Structure 2.

23 Multiple Bonds: double & triple bonds If all valence e - are used & an atom does not have an octet, form multiple bonds.

24 Formal Charge Formal charge: the charge assigned to individual atoms in a Lewis structure. # of electrons assigned (owns) = unshared electrons + 1/2 shared (bonding) electrons

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26 Isomers: different molecules having the same molecular formula (different arrangement of atoms) Isomers

27 Resonance Structure 1. 2 Lewis structures having the same placement of atoms but a different arrangement of e -. (Isomers: differ in the arrangement of both atoms and e - ) 2. Major & minor contributor: A major contributor is a more stable form (more bonds and fewer charges).

28 Resonance Hybrids 1.Resonance structures are not real: The true structure is a resonance hybrid 2. Electron pairs are delocalized, & e - delocalization adds stability  resonance stabilized.

29 Drawing Resonance Structures Rule 1. 2 resonance structures differ in the position of multiple bonds and nonbonded e -. The placement of atoms and single bonds always stays the same. Rule 2. 2 resonance structures must have the same number of unpaired electrons. Rule 3. Valid Lewis structure.

30 Curved Arrow Notation Curved arrow shows the movement of an e - pair The head points to where the e - pair “moves.” Example 1: Example 2:

31 Molecular Shape: bond length and bond angle. 1) Decreases across a row of the periodic table. 2) Increases down a column of the periodic table. 1. Bond length

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33 Molecular Geometry Valence Shell Electron Pair Repulsion (VSEPR) theory: Atoms or lone pair e - are as far away from each other as possible.

34 Varying Bond Angles Methane (CH 4 ) Ammonia (NH 3 ) Water (H 2 O) Repulsion of the lone pairs of e -.

35 2 Groups Around an Atom

36 3 Groups Around an Atom

37 4 Groups Around an Atom

38 Drawing Three-Dimensional Structures Solid line, wedge and dashed line (perspective formula)

39 Ammonia (NH 3 ) & Water (H 2 O), - Geometry is a tetrahedron (Nonbonded 2e - is counted as a “Group”)

40 Summary: Predicting Geometry Based on Number of Groups

41 Drawing Organic Molecules 1. Dashed (lone pairs can be presented/omitted) & Condensed Structures (lone pairs are omitted).

42 2. Skeletal Structures A carbon atom at the junction of any two lines or at the end of any line. Draw all heteroatoms and the hydrogens directly bonded to them.

43 Orbitals and Bonding 1. Hydrogen (H 2 ) Sigma bond: cylindrically symmetrical.

44 Orbitals and Bonding: Carbon 1. Electronic configuration of ‘Carbon’ in its ground state, Atomic orbital (1s 2, 2s 2, 2p 2 )

45 1) An electronically excited state (higher energy). 2. Tetravalent Carbon (sp 3 Hybrid Orbital) 2) sp 3 Hybrid orbitals by orbital hybridization (1s 2, 4 x 2sp 3 )

46 Shape and Orientation Carbon Hybrid Orbitals 1.sp 3 Hybrid Orbitals 1) Shape of sp 3 hybrid orbitals. 2) 4 sp 3 hybrid orbitals: tetrahedron

47 CH 4 : sp 3 Hybrid Orbital Overlaping of an sp 3 hybrid orbital of C with a 1s orbital of H.

48 Hybridization and Bonding: CH 3 -CH 3

49 Hybrid Orbitals of NH 3 and H 2 O

50  bond : free rotation around the bond.

51 2. sp 2 & sp Hybrid Orbitals 1) sp hybrid orbital 2) sp 2 hybrid orbital

52 Each carbon is trigonal planar. Each carbon is sp 2 hybridized. sp 2 Hybrid Orbitals

53 Ethylene: sp 2 Hybrid Orbital

54 1)unstable; has a nod 2)Higher reactivity: exposed e - 3)Rigid shape: rotation is restricted.  bond : Rigid, but unstable bond

55 sp Hybrid Orbitals

56 Acetylene (Ethyne): sp Hybrid Orbital

57 All triple bonds are composed of one sigma and two  bonds. Triple Bonds

58 Summary of Bonding in Acetylene

59 Summary of Covalent Bonding

60 Hybridization of Be & B 2 vacant orbitals One vacant orbital

61 Bond Length and Bond Strength 1) As the number of electrons between two nuclei increases, bonds become shorter and stronger. 2) More s character, shorter and stronger bond

62 Hybrid Orbitals of NH 3 and H 2 O

63 Percent s-Character

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65 Electronegativity A measure of an atom’s e - attracting ability in a bond.

66 Electronegativity values: determine e - sharing in a bond equally or unequally between two atoms. Nonpolar bond: e - are equally shared. Polar bond: e - are unequally shared. Bond Polarity

atoms having same or similar electronegativities 2. C-C bond is nonpolar. 3. C–H bonds are considered to be nonpolar: the (electronegativity difference between C and H is small). Nonpolar Bonds

68 1. Bonding between atoms of different electronegativity values. Example: C – O bond Electronegativity of C = 2.5, and of O = 3.4. The bond has a dipole (charge separation). Polar Bonds

69  + : electron deficient.  - : electron rich.  : direction of polarity in a bond. Depicting Polarity

70 Polarity of Molecules Determination of a net dipole of a molecule - Vector of the dipole (direction): decide if individual dipoles cancel or reinforce each other in space.

71 A nonpolar molecule has either no polar bonds, or two or more bond dipoles that cancel. Nonpolar Molecules

72 Electrostatic potential plot of CH 3 Cl