1. Iran University of Science & Technology
Organic Chemistry
M. R. Naimi-Jamal
Faculty of Chemistry
Iran University of Science & Technology

2. Polar Covalent Bonds Acids and Bases
Chapter 1. Continue
Polar Covalent Bonds Acids and Bases
Based on: McMurry’s Fundamental of Organic Chemistry, 4th edition, Chapter 1

3. 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 in not symmetrical

4. 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
Arbitrary scale. As shown in next figure, electronegativities are based on an arbitrary scale
F is most electronegative (EN = 4.0), Cs is least (EN = 0.7)

5. The Periodic Table and Electronegativity

6. Bond Polarity and Electronegativity
Metals on left side of periodic table attract electrons weakly: lower electronegativities
Halogens and other reactive nonmetals on right side of periodic table attract electrons strongly: higher electronegativities
Electronegativity of C = 2.5

7. Bond Polarity and Inductive Effect
Nonpolar Covalent Bonds: atoms with similar electronegativities
Polar Covalent Bonds: Difference in EN of atoms < 2
Ionic Bonds: Difference in electronegativities > 2 (approximately).
Other factors (solvation, lattice energy, etc) are important in ionic character.

8. Bond Polarity and Inductive Effect
Bonding electrons are pulled toward the more electronegative atom in the bond
Electropositive atom acquires partial positive charge, +
Electronegative atom acquires partial negative charge, -
Inductive effect: shifting of electrons in a bond in response to the electronegativities of nearby atoms

9. Electrostatic Potential Maps
Electrostatic potential maps show calculated charge distributions
Colors indicate electron-rich (red) and electron-poor (blue) regions

10. Polar Covalent Bonds: Dipole Moments
Molecules as a whole are often polar, from vector summation of individual bond polarities and lone-pair contributions

11. Polar Covalent Bonds: Dipole Moments
Dipole moment - Net molecular polarity, due to difference in summed charges
 - magnitude of charge Q at end of molecular dipole times distance r between charges
 = Q  r, in debyes (D)
1 D = 3.34  1030 coulomb meter

12. Dipole Moments in Water and Ammonia
Large dipole moments
Electronegativities of O and N > H
Both O and N have lone-pair electrons oriented away from all nuclei

14. Practice: Suggest an explanation
Ammonia (NH3) has a dipole moment of 1.46 D, while the dipole moment of NF3 is only 0.24 D. Why?

15. 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

16. Cis- & trans-1,2-dichloroethylenes:

17. Formal Charges
Sometimes it is necessary to have structures with formal charges on individual atoms
We compare the bonding of the atom in the molecule to the valence electron structure
If the atom has one more electron in the molecule, it is shown with a “-” charge
If the atom has one less electron, it is shown with a “+” charge

18. Formal Charges

19. Nitromethane:

21. Formal Charges

22. Resonance
Some molecules have structures that cannot be shown with a single Lewis representation
In these cases we draw Lewis structures that contribute to the final structure but which differ in the position of the  bond(s) or lone pair(s)
Such a structure is delocalized and is represented by resonance forms

23. Resonance
The resonance forms are connected by a double-headed arrow

24. Resonance Hybrids
A structure with resonance forms does not alternate between the forms
Instead, it is a hybrid of the two resonance forms, so the structure is called a resonance hybrid
For example, benzene (C6H6) has two resonance forms with alternating double and single bonds
In the resonance hybrid, the actual structure, all of the C-C bonds are equivalent, midway between double and single bonds

25. Resonance Hybrids

26. Resonance Hybrids

27. Rules for Resonance Forms
Individual resonance forms are imaginary - the real structure is a hybrid (only by knowing the contributors can you visualize the actual structure)
Resonance forms differ only in the placement of their  or nonbonding electrons
Different resonance forms of a substance don’t have to be equivalent
Resonance forms must be valid Lewis structures: the octet rule usually applies
The resonance hybrid is more stable than any individual resonance form would be

28. Curved Arrows and Resonance Forms
We can imagine that electrons move in pairs to convert from one resonance form to another
A curved arrow shows that a pair of electrons moves from the atom or bond at the tail of the arrow to the atom or bond at the head of the arrow

29. Curved Arrows and Resonance Forms

30. Drawing Resonance Forms
Any three-atom grouping with a multiple bond has two resonance forms

31. Different Atoms in Resonance Forms
Sometimes resonance forms involve different atom types as well as locations
The resulting resonance hybrid has properties associated with both types of contributors
The types may contribute unequally

32. Resonance in the acetone enolate
The “enolate” derived from acetone is a good illustration, with delocalization between carbon and oxygen.

33. 2,4-Pentanedione The anion derived from 2,4-pentanedione
Lone pair of electrons and a formal negative charge on the central carbon atom, next to a C=O bond on the left and on the right
Three resonance structures result

34. Practice: Draw three resonance forms:

35. Solution:

36. Practice: Draw three resonance forms:

37. Solution:

38. Solution:

39. Acids and Bases: The Brønsted–Lowry Definition
Brønsted–Lowry theory defines acids and bases by their role in reactions that transfer protons (H+) between donors and acceptors.
“proton” is a synonym for H+ - loss of an electron from H leaving the bare nucleus - a proton. Protons are always covalently bonded to another atom.

40. Brønsted Acids and Bases
“Brønsted-Lowry” is usually shortened to “Brønsted”
A Brønsted acid is a substance that donates a hydrogen ion, or “proton” (H+): a proton donor
A Brønsted base is a substance that accepts the H+: a proton acceptor

41. The Reaction of HCl with H2O
When HCl gas dissolves in water, a Brønsted acid–base reaction occurs
HCl donates a proton to water molecule, yielding hydronium ion (H3O+) and Cl
The reverse is also a Brønsted acid–base reaction of the conjugate acid and conjugate base

42. The Reaction of HCl with H2O

44. Quantitative Measures of Acid Strength
The equilibrium constant (Keq) for the reaction of an acid (HA) with water to form hydronium ion and the conjugate base (A-) is a measure related to the strength of the acid
Stronger acids have larger Keq
Note that brackets [ ] indicate concentration, moles per liter, M.

45. Ka – the Acidity Constant
The concentration of water as a solvent does not change significantly when it is protonated in dilute solution.
The acidity constant, Ka for HA equals Keq times 55.6 M (leaving [water] out of the expression)
Ka ranges from for the strongest acids to very small values (10-60) for the weakest

46. Ka – the Acidity Constant

47. Acid and Base Strength
The ability of a Brønsted acid to donate a proton to is sometimes referred to as the strength of the acid.
The strength of the acid can only be measured with respect to the Brønsted base that receives the proton
Water is used as a common base for the purpose of creating a scale of Brønsted acid strength

48. pKa – the Acid Strength Scale
pKa = -log Ka (in the same way that
pH = -log [H+]
The free energy in an equilibrium is related to –log of Keq (DG = -RT log Keq)
A larger value of pKa indicates a stronger acid and is proportional to the energy difference between products and reactants

49. pKa – the Acid Strength Scale
The pKa of water is 15.74

50. pKa values for some acids:

51. Predicting Acid–Base Reactions from pKa Values
pKa values are related as logarithms to equilibrium constants
The difference in two pKa values is the log of the ratio of equilibrium constants, and can be used to calculate the extent of transfer

52. Predicting Acid–Base Reactions from pKa Values

53. Predicting Acid–Base Reactions from pKa Values

54. Organic Acids and Organic Bases
The reaction patterns of organic compounds often are acid-base combinations
The transfer of a proton from a strong Brønsted acid to a Brønsted base, for example, is a very fast process and will always occur along with other reactions

55. Organic Acids
Those that lose a proton from O–H, such as methanol (CH3OH) and acetic acid (CH3COOH)
Those that lose a proton from C–H, usually from a carbon atom next to a C=O double bond (O=C– C– H)

56. Organic Acids

57. Organic Acids:

58. Carboxylic Acids:

59. Conjugate Bases:

60. 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 when with a strong acid or as acids with strong bases

61. Organic Bases

62. Acids and Bases: The Lewis Definition
Lewis acids are electron pair acceptors;
Lewis bases are electron pair donors
The Lewis definition leads to a general description of many reaction patterns but there is no quantitatve scale of strengths as in the Brønsted definition of pKa

63. Lewis Acids and the Curved Arrow Formalism
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
Group 3A elements, 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

64. Lewis Acids and the Curved Arrow Formalism
Organic compounds that undergo addition reactions with Lewis bases (discussed later) are called electrophiles and therefore Lewis Acids
The combination of a Lewis acid and a Lewis base can shown with a curved arrow from base to acid

65. Illustration of Curved Arrows in Following Lewis Acid-Base Reactions

66. Some Lewis Acids:

67. Lewis Bases
Lewis bases can accept protons as well as other 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 either acids or bases, depending on the reaction

68. Some Lewis Bases

69. Lewis Acid-Base Reactions

70. Imidazole:

71. Drawing Chemical Structures
Condensed structures: C-H and C-C and single bonds aren't shown but understood
If C has 3 H’s bonded to it, write CH3
If C has 2 H’s bonded to it, write CH2; and so on.
Horizontal bonds between carbons aren't shown in condensed structures - the CH3, CH2, and CH units are assumed to be connected horizontally by single bonds, but vertical bonds are added for clarity

72. 2-methylbutane Structures

73. Skeletal Structures
Minimum amount of visual information but unambiguous
C’s not shown, but assumed to be at each intersection of two lines (bonds) and at end of each line
H’s bonded to C’s aren't shown – whatever number is needed will be there to fill out the four bonds to each carbon.
All atoms other than C and H are shown

75. Practice: How many H’s on each carbon?

76. Solution:

77. Problem: How many H’s on each carbon?

78. Molecular Models
We often need to visualize the shape or connections of a molecule in three dimensions
Molecular models are three dimensional objects, on a human scale, that represent the aspects of interest of the molecule’s structure (computer models also are possible)
Drawings on paper and screens are limited in what they can present to you

79. Molecular Models
Framework models (ball-and-stick) are essential for seeing the relationships within and between molecules – you should own a set
Space-filling models are better for examining the crowding within a molecule
Framework
Space-filling

80. Summary
Organic molecules often have polar covalent bonds as a result of unsymmetrical electron sharing caused by differences in the electronegativity of atoms
The polarity of a molecule is measured by its dipole moment, .
(+) and () indicate formal charges on atoms in molecules to keep track of valence electrons around an atom
Some substances must be shown as a resonance hybrid of two or more resonance forms that differ by the location of electrons.

81. Summary A Brønsted(–Lowry) acid donates a proton
A Brønsted(–Lowry) base accepts a proton
The strength Brønsted acid is related to the -1 times the logarithm of the acidity constant, pKa. Weaker acids have higher pKa’s
A Lewis acid has an empty orbital that can accept an electron pair
A Lewis base can donate an unshared electron pair

82. Summary In condensed structures C-C and C-H are implied
Skeletal structures show bonds and not C or H (C is shown as a junction of two lines) – other atoms are shown
Molecular models are useful for representing structures for study

83. Chapter 1-2, Questions
42, 43, 44, 46, 47, 50, 54, 55, 57