1. William H. Brown & Christopher S. Foote
Organic Chemistry
William H. Brown & Christopher S. Foote

2. Covalent Bonding & Shapes of Molecules
Chapter 1

3. Organic Chemistry The study of the compounds of carbon
Over 10 million structures have been identified
about 1000 new ones are identified each day!
C is a small atom
it forms single, double, and triple bonds
it is intermediate in electronegativity (2.5)
it forms strong bonds with C, H, O, N, and some metals

4. Structure of Atoms Electronic
small dense nucleus, diameter m, which contains most of the mass of the atom
extranuclear space, diameter m, which contains positively-charged electrons

5. Structure of Atoms Electronic
Electrons are confined to regions of space called principle energy levels (shells)
each shell can hold 2n2 electrons (n = 1,2,3, )

6. Electronic Structure of Atoms
Shells are divided into subshells called orbitals, which are designated by the letters s, p, d, f,
s (one per shell)
p (set of three per shell 2 and higher)
d (set of five per shell 3 and higher) .....

7. Electron Structure of Atoms
Aufbau Principle: orbitals fill from lowest to highest energy
Pauli Exclusion Principle: only two electrons per orbital, spins must be paired
Hund’s Rule: for a set of degenerate orbitals, add one electron in each before a second is added in any one
Example: Write the ground-state electron configuration for each element
(a) Li (b) O (c) Cl

8. Lewis Structures Gilbert N. Lewis
Valence shell: the outermost electron shell of an atom
Valence electrons: electrons in the valence shell of an atom; these electrons are used to form chemical bonds
Lewis structure:
the symbol of the atom represents the nucleus and all inner shell electrons
dots represent valence electrons

9. Lewis Structures
Table 1.4 Lewis Structures for Elements 1-18

10. Formation of Chemical Bonds
Atoms bond together so that each atom acquires an electron configuration the same as the noble gas nearest it in atomic number
an atom that gains electrons becomes an anion
an atom that loses electrons becomes a cation
Two extremes of bonding
ionic bond: a chemical bond resulting from the electrostatic attraction of an anion and a cation
covalent bond: a chemical bond formed between two atoms by sharing one or more pairs of electrons

11. Electronegativity
Electronegativity: a measure of the force of an atom’s attraction for the electrons it shares with another atom in a chemical bond
Pauling scale
increases left to right in a row
increases bottom to top in a column

12. Electronegativity
Table 1.6 Classification of Chemical Bonds

13. Bond Dipoles
Table 1.7 Average Bond Dipoles of Selected Covalent Bonds

14. Lewis Structures To write a Lewis structure
determine the number of valence electrons
determine the arrangement of atoms
connect the atoms by single bonds
arrange the remaining electrons so that each atom has a complete valence shell
show bonding electrons as a single bond (a single line); show nonbonding electrons as a pair of dots
in a single bond atoms share one pair of electrons, in a double bond they share two pairs of electrons, and in a triple bond they share three pairs of electrons

15. Lewis Structures

16. Lewis Structures In neutral molecules hydrogen has one bond
carbon has 4 bonds and no unshared electrons
nitrogen has 3 bonds and 1 unshared pair of electrons
oxygen has 2 bonds and 2 unshared pairs of electrons
halogens have 1 bond and 3 unshared pairs of electrons

17. Formal Charge
Formal charge: the charge on an atom in a molecule or polyatomic ion
To derive formal charge
1. write a correct Lewis structure for the molecule or ion
2. assign each atom all of its unshared (nonbonding) electrons and one-half its shared (bonding) electrons
3. compare this number with the number of valence electrons in the neutral, unbonded atom

18. Formal Charge
if the number assigned to the bonded atom is less than that assigned to the unbonded atom, the atom has a positive formal charge
if the number is greater, the atom has a negative formal charge

19. Formal Charge
Example: Draw Lewis structures and show all formal charges for these ions

20. Exceptions to the Octet Rule
Molecules containing atoms of Group 3A elements, particularly boron and aluminum

21. Exceptions to the Octet Rule
Atoms of third-period elements have 3d orbitals and may expand their valence shells to contain more than 8 electrons
phosphorus may have up to 10

22. Exceptions to the Octet Rule
and S, which may have up to 12 electrons in its valence shell

23. Functional Groups
Functional group: an atom or group of atoms within a molecule that shows a characteristic set of physical and chemical properties
Functional groups are important for three reason; they are
1. the units by which we divide organic compounds into classes
2. the sites of characteristic chemical reactions
3. the basis for naming organic compounds

24. Alcohol
contains an -OH (hydroxyl) group

25. Amine
contains an amino group; a nitrogen bonded to one, two, or three carbon atoms
may by 1°, 2°, or 3°

26. Aldehyde and Ketone
contains a carbonyl (C=O) group

27. Carboxylic Acid
contains a carboxyl (-COOH) group

28. Carboxylic Ester
a derivative of a carboxylic acid in which the carboxyl hydrogen is replaced by a carbon group

29. VSEPR Model

30. VSEPR Model
Example: predict all bond angles for these molecules and ions

31. Polar and Nonpolar Molecules
To determine if a molecule is polar, we need to determine
if the molecule has polar bonds
the arrangement of these bonds in space
Dipole moment (): the vector sum of its individual bond dipole moments in a molecule
reported in debyes (D)

32. Polar and Nonpolar Molecules
these molecules have polar bonds, but each has a zero dipole moment

33. Polar and Nonpolar Molecules
These molecules have polar bonds and a dipole moment greater than zero

34. Resonance
For many molecules and ions, no single Lewis structure provides a truly accurate representation

35. Resonance Linus Pauling - 1930s
many molecules and ions are best described by writing two or more Lewis structures
individual Lewis structures are called contributing structures
connect individual contributing structures by double-headed (resonance) arrows
the molecule or ion is a hybrid of the various contributing structures

36. Resonance
Examples: equivalent contributing structures

37. Resonance
Curved arrow: a symbol used to show the redistribution of valence electrons
In using curved arrows, there are only two allowed types of electron redistribution:
from a bond to an adjacent atom
from an atom to an adjacent bond
Electron pushing is a survival skill in organic chemistry. Learn it well!

38. Resonance All contributing structures must
1. have the same number of valence electrons
2. obey the rules of covalent bonding
no more than 2 electrons in the valence shell of H
no more than 8 electrons in the valence shell of a 2nd period element
a 3rd period element may have up to 12 electrons in its valence shell
3. differ only in distribution of valence electrons
4. have the same number of paired and unpaired electrons

39. Resonance
Examples of ions and a molecule best represented as resonance hybrids

40. Resonance Preference 1: filled valence shells
structures in which all atoms have filled valence shells contribute more than those with unfilled valence shells

41. Resonance Preference 2: maximum number of covalent bonds
structures with a greater number of covalent bonds contribute more than those with fewer covalent bonds

42. Resonance Preference 3: least separation of unlike charge
structures with separation of unlike charges contribute less than those with no charge separation

43. Resonance
Preference 4: negative charge on the more electronegative atom
structures that carry a negative charge on the more electronegative atom contribute more than those with the negative charge on the less electronegative atom

44. Quantum or Wave Mechanics
Albert Einstein: E=hn (energy is quantized)
light has particle properties
Louis deBroglie: wave/particle duality
Erwin Schrödinger: wave equation
 2 is the probability of finding an electron in a given region of space
the standard representation of an atomic orbital is a boundary surface representing 95% probability of finding an electron in that region of space

45. Shapes of 1s and 2s Orbitals

46. Shapes of a Set of 2p Atomic Orbitals

47. Molecular Orbital Theory
electrons in atoms exist in atomic orbitals
electrons in molecules exist in molecular orbitals (MOs)
using the Schrödinger equation, we can calculate the shapes and energies of MOs

48. Molecular Orbital Theory
Rules:
combination of n atomic orbitals gives n MO
MOs are arranged in order of increasing energy
MOs filling is governed by the same rules as for atomic orbitals:
Aufbau principle: fill beginning with LUMO
Pauli exclusion principle: no more than 2e- in a MO
Hund’s rule: filling of degenerate orbitals

49. Molecular Orbital Theory
Terminology
ground state = lowest energy
excited state = NOT lowest energy
 = sigma bonding MO
* = sigma antibonding MO
 = pi bonding MO
* = pi antibonding MO

50. Molecular Orbital Theory
sigma 1s bonding and antibonding MOs

51. Molecular Orbitals
computed sigma bonding and antibonding MOs for H2

52. Molecular Orbitals
pi bonding and antibonding MOs

53. Molecular Orbitals
computed pi bonding and antibonding MOs for ethylene

54. Molecular Orbitals
computed pi bonding and antibonding orbitals for formaldehyde

55. Hybrid Orbitals The Problem: A Solution
bonding by 2s and 2p atomic orbitals would give bond angles of approximately 90°
instead we observe bond angles of approximately 109.5°, 120°, and 180°
A Solution
hybridization of atomic orbitals
2nd row elements use sp3, sp2, and sp hybrid orbitals for bonding

56. Hybrid Orbitals Hybridization of orbitals (L. Pauling)
the combination of two or more atomic orbitals forms a new set of atomic orbitals, called hybrid orbitals
We deal with three types of hybrid orbitals
sp3 (one s orbital + three p orbitals)
sp2 (one s orbital + two p orbitals)
sp (one s orbital + one p orbital)
Overlap of hybrid orbitals can form two types of bonds depending on the geometry of overlap
 bonds are formed by “direct” overlap
 bonds are formed by “parallel” overlap

57. sp3 Hybrid Orbitals
each sp3 hybrid orbital has two lobes of unequal size
the sign of the wave function is positive in one lobe, negative in the other, and zero at the nucleus
the four sp3 hybrid orbitals are directed toward the corners of a regular tetrahedron at angles of 109.5°

58. Bonding in CH4, NH3, and H2O

59. sp2 Hybrid Orbitals

60. Bonding in CH2=CH2

61. Bonding in CH2=O

62. sp Hybrid Orbitals two lobes of unequal size at an angle of 180°
the two unhybridized 2p orbitals are perpendicular to each other and to the line through the two sp hybrid orbitals

63. Bonding in C2H2

64. Hybrid Orbitals

65. Hybrid Orbitals

66. Prob 1.26
Write Lewis structures for these molecules

67. Prob 1.27
Write Lewis structures for these ions

68. Prob 1.28
Complete the structural formula and write the molecular formula of each compound.

69. Prob 1.29
Which structural formulas are incorrect, and why?

70. Prob 1.30
Add valence electrons to complete the outer shell of each atom and assign any formal charges

71. Prob 1.31
Assign all formal charges

72. Prob 1.37
Use the VSEPR model to predict all bond angles

73. Prob 1.38
Use the VSEPR model to predict all bond angles

74. Prob 1.39
Use the VSEPR model to predict the geometry of each ion.

75. Prob 1.47
Identify the functional groups in each compound.

76. Prob 1.48
Which compounds have dipole moments greater than zero? In what direction does it point?

77. Prob 1.55
State the orbital hybridization of each highlighted atom.

78. Prob 1.56
Describe each highlighted bond in terms of the overlap of atomic orbitals.

79. Prob 1.57
Predict all bond angles, the hybridization of each carbon, and the shape of a benzene molecule.

80. Prob 1.58
Complete the Lewis structure and describe each highlighted bond in terms of the overlap of atomic orbitals.

81. Prob 1.61
Draw a Lewis structure for each compound. Show covalent bonds by dashes, and ionic bonds by the charge on each ion.

82. Prob 1.70
Identify the atoms in the organic starting material that change hybridization upon reaction, and what the change is.

83. Prob 1.70
Identify the atoms in the organic starting material that change hybridization upon reaction, and what the change is.

84. Covalent Bonds & Shapes of Molecules
End Chapter 1