1. LECTURE 2

2. Structure and bonding
in organic compounds

3. Atomic Orbitals
Quantum mechanics: describes electron energies and locations by a wave equation
Wave function is solution of wave equation
Each wave function is an orbital
Orbital is a part of space where electron is the most likely to be
Orbital has no specific boundary so we show the most probable area

4. Shapes of Atomic Orbitals for Electrons
s and p orbitals most important in organic and biological chemistry
s orbitals: spherical, nucleus at center
p orbitals: dumbbell-shaped, nucleus at middle
d orbitals: cloverleaf-shaped, nucleus at center

5. 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

6. Orbitals and 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 maximum two electrons

7. Orbitals and Shells

8. Orbitals and Shells
First shell contains one s orbital, denoted 1s, holds only 2 electrons
Second shell contains one s orbital (2s) and three p orbitals (2p), capacity - 8 electrons
Third shell contains an s orbital (3s), three p orbitals (3p), and five d orbitals (3d), capacity - 18 electrons

9. PERIODIC CHART
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10. Ground state electronic configurations of H, C, N, O2p H N 1s 2s 1s 2p 2p C O 2s 2s 1s 1s

11. Atoms form bonds because the compound that results is more stable than the separate atoms
Ionic bonds in salts form as a result of electron transfer
Organic compounds have covalent bonds from sharing electrons (G. N. Lewis, 1916)

12. Lewis structures (electron dot) show valence electrons of an atom as dots
Hydrogen has one dot, representing its 1s electron
Carbon has four dots (2s2 2p2)
Kekule structures (line-bond structures) have a line drawn between two atoms indicating a 2 electron covalent bond.
Stable molecule results at completed shell, octet (eight dots) for main-group atoms (two for hydrogen)

13. Number of covalent bonds formed by H, C, N, O, halides

14. Lewis and Kekule structures
Valence electrons not used in bonding are called
nonbonding electrons, or lone-pair electrons

15. Molecular Orbital Theory – σ bond
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) MO is higher energy

16. Molecular Orbital Theory –  bond
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

17. The Nature of Covalent Bonds Molecular orbital theory
H–H bond results from the overlap of two singly occupied
hydrogen 1s orbitals
H-H bond is cylindrically symmetrical, sigma (σ) bond

18. Bond Energy Reaction 2 H·  H2 releases 436 kJ/mol (104 kcal/mol)
Product has 436 kJ/mol less energy than two hydrogens
H–H has bond strength of 436 kJ/mol. (1 kJ = kcal;
1 kcal = kJ)

19. 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

20. Atomic orbitals of carbon (hybridization)2p C 2s
Atoms surround carbon at corners of a tetrahedron

21. Atomic orbitals of carbon (hybridization) 4 equivalent sp3 orbitals
Ground state carbon One electron excited
hybridization
Carbon sp3
4 equivalent sp3 orbitals

22. Carbon sp3 hybridization
sp3 hybrid orbitals: s orbital and three p orbitals combine to form four equivalent, unsymmetrical, tetrahedral orbitals (sppp = sp3), Pauling (1931)

23. The Structure of Methane
sp3 orbitals on C overlap with 1s orbitals on 4 H atoms to form four identical C-H bonds
Each C–H bond has a strength of 436 kJ/mol (104 kcal/mol) and length of 109 pm (1.09Å)
Bond angle: each H–C–H is 109.5°, the tetrahedral angle.

24. The Structure of Ethane
Two carbons bond to each other by σ overlap of an sp3 orbital from each
Three sp3 orbitals on each C overlap with H 1s orbitals to form six C–H bonds (σ)

25. Carbon sp2 hybridization
sp2 hybrid orbitals: 2s orbital combines with two 2p orbitals, giving
3 orbitals (spp = sp2)
sp2 orbitals are in a plane with120° angles
Remaining p orbital is perpendicular to the plane

26. Orbitals and the Structure of Ethylene

27. The Structure of Ethylene
H atoms form σ 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 134 pm (C–C 154 pm)

28. Carbon sp hybridization
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

29. Orbitals of Acetylene
Two sp hybrid orbitals from each C form sp–sp σ bond
pz orbitals from each C form a pz–pz  bond by sideways overlap and py orbitals overlap similarly

30. The Structure of Acetylene
Sharing of six electrons forms C≡C (1σ and 2π)
Two sp orbitals form σ bonds with hydrogens

31. Comparison of bond energy and bond length in hydrocarbons

32. Molecular Orbitals and Structure of Benzene
All its C-C bonds are the same length: 139 pm — between single (154 pm) and double (134 pm) bonds
Electron density in all six C-C bonds is identical
Structure is planar, hexagonal
C–C–C bond angles 120°
Each C is sp2 and has a p orbital perpendicular to the plane of the six-membered ring

33. Molecular Orbitals of Benzene
Six sp2 hybridized carbons linked by σ bonds form a flat ring
The 6 p-orbitals combine to give
Three bonding orbitals with 6  electrons,
Three antibonding with no electrons
Orbitals with the same energy are degenerate

34. Molecular Orbitals of Benzene

35. Hybridization of Nitrogen
H–N–H bond angle in ammonia (NH3) 107.3°
C-N-H bond angle is °
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 σ bonds to H and CH3

36. Hybridization of Oxygen
H–O–H bond angle in water (H2O) 104.5°
O’s orbitals (sppp) hybridize to form four sp3 orbitals
Two sp3 orbitals are occupied by lone pairs of electrons, and two sp3 orbitals have one electron each, forming bonds to H

37. Polar Covalent Bonds: Electronegativity
Covalent bonds can have ionic character
These are polar covalent bonds
Bonding electrons attracted more strongly by one atom than by the other
Electron distribution between atoms is not symmetrical
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38. The Periodic Table and Electronegativity

39. 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
Electronegativities are based on an arbitrary scale
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 = EN of H = 2.1

40. Bond Polarity and Inductive Effect
Nonpolar Covalent Bonds: atoms with similar EN
C-C C-H
Polar Covalent Bonds: Difference in EN of atoms < 2
Ionic Bonds: Difference in EN > 2
C-O, C-X bonds (more electronegative elements) are polar
Bonding electrons shifted toward electronegative atom
C acquires partial positive charge, +
Electronegative atom acquires partial negative charge, -
Inductive effect: shifting of electrons in a bond in response to EN of nearby atoms

41. Electrostatic Potential Maps
Electrostatic potential maps show calculated charge distributions
Colors indicate electron-rich (red) and electron-poor (blue) regions
Arrows indicate direction of bond polarity

42. 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 are soluble in polar solvents like water; nonpolar substances are insoluble in water

43. Polar Covalent Bonds: Dipole Moments
Dipole moment () - Net molecular polarity results from summation of individual bond polarities, formal charges and lone pair contributions
 - magnitude of charge Q at end of molecular dipole times distance r between charges
 = Q  r, in debyes (D), 1 D =  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.

44. Comparison of Dipole Moments

45. Polar organic molecules Formaldehyde Chloromethaneδ- δ+ δ+ δ-
Formaldehyde Chloromethane
μ = 2.33 D
μ = 1.87 D

46. Absence of Dipole Moments
In symmetrical molecules, the dipole moments of each bond has one in the opposite direction
The effects of the local dipoles cancel each other

47. Noncovalent Interactions (Intermolecular weak forces)
Dipole-dipole forces
Dispersion forces
Hydrogen bonds

48. Dipole-Dipole Interactions
• Occur between polar molecules as a result of electrostatic interactions among dipoles
• Forces can be attractive of repulsive depending on orientation of the molecules

49. Dispersion Forces
• Occur between all neighboring molecules and arise because the electron distribution within molecules are constantly changing

50. Hydrogen Bond Forces
• Most important noncovalent interaction in biological molecules
• Forces are result of attractive interaction between a hydrogen bonded to an electronegative O or N atom and lone electron pair on another O or N atom

51. Hydrogen Bond Forces

52. Acids and Bases: The Brønsted–Lowry Definition
“Brønsted-Lowry” is usually shortened to “Brønsted”
A Brønsted acid is a substance that donates a proton (H+) in the reaction (it is proton donor)
A Brønsted base is a substance that accepts a proton (H+)
in the reaction (it is proton acceptor)
“proton” is a synonym for H+ - loss of an electron from hydrogen atom (H) leaving the bare nucleus - a proton

53. The Reaction of Acid with Base

54. Acid and Base Strength
The equilibrium constant (Keq) for the reaction of an acid (HA) with water is a measure of the strength of the acid
Stronger acids have larger Keq
Brackets [ ] indicate concentration, moles per liter, M.

55. Ka – the Acidity Constant
Ka ranges from 1015 for the strongest acids
to very small values (10-60) for the weakest

56. pKa – the Acid Strength Scale
pKa = -log Ka
The free energy in an equilibrium is related to –log of Keq (ΔG = -RT log Keq)
A smaller value of pKa indicates a stronger acid and is proportional to the energy difference between products and reactants
The pKa of water is 15.74

57. Relative Acidity

58. Organic Acids
characterized by the presence of positively polarized hydrogen atom
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59. Organic Bases
Have an atom with a lone pair of electrons that can bond to H+
Nitrogen-containing compounds derived from ammonia are the most common organic bases
Oxygen-containing compounds can react as bases with a strong acid or as acids with strong bases

60. Acids and Bases: The Lewis Definition
Lewis acids are electron pair acceptors and Lewis bases are electron pair donors
Brønsted acids are not Lewis acids because they cannot accept an electron pair directly (only a proton would be a Lewis acid)
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 pKa

61. Lewis Acids (electrophiles)
The Lewis definition of acidity includes metal cations, such as Mg2+
They accept a pair of electrons when they form a bond to a base
Compounds such as BF3 and AlCl3, are Lewis acids because they have unfilled valence orbitals and can accept electron pairs from Lewis bases
Transition-metal compounds, such as TiCl4, FeCl3, ZnCl2, and SnCl4, are Lewis acids
Organic compounds that undergo addition reactions with Lewis bases are called electrophiles or Lewis Acids
The combination of a Lewis acid and a Lewis base can be shown with a curved arrow from base to acid

62. Example of Lewis Acid-Base Reaction

63. Lewis Bases (nucleophiles)
Lewis bases can accept protons as well as Lewis acids, therefore the definition encompasses that for Brønsted bases
Most oxygen- and nitrogen-containing organic compounds are Lewis bases because they have lone pairs of electrons
Some compounds can act as both acids and bases, depending on the reaction

64. Lewis Bases