1. Chapter 2 Polar Covalent Bonds; Acids and Bases Organic Chemistry

2. Chapter Objectives Take an in-depth look at polarity of molecules Use formal charges to designate the distribution of electrons Represent molecules with resonance structures by ‘pushing’ electrons Examine the acid-base behavior of molecules Predict acid-base reactions from pKa values

3. Electronegativity electronegativity – (EN) a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound The difference in electronegativity values for two atoms will indicate whether the two atoms form an ionic bond or a polar or nonpolar covalent bond.

4. Electronegativity Table ionic bonding - 2.0 <  EN polar covalent bonding -.5 ≤  EN ≤ 2.0 non-polar covalent bonding -  EN <.5

5. Bond Formation Ionic bonding involves the loss of an electron due to a large difference in electronegativity (2.0 <  EN ) Covalent bonding involves the sharing of electrons Equal sharing: non-polar bond (  EN <.5) Unequal sharing: polar bond (.5 ≤  EN ≤ 2.0)

6. Polarity If one side is more electronegative, it tends to have a partial negative charge (δ-) [electron-rich] The other side tends to have a partial positive charge (δ+) [electron-poor] The δ- and δ+ difference along a bond is called a dipole moment δ-δ- δ+δ+

7. Ionic Character

8. Electrostatic Potential Maps Red – electron rich (δ-) Blue – electron poor (δ+)

10. Electrostatic Potential Maps Red – electron rich (δ-) Blue – electron poor (δ+)

11. Electrostatic Potential Maps Red – electron rich (δ-) Blue – electron poor (δ+)

12. Electrostatic Potential Maps Red – electron rich (δ-) Blue – electron poor (δ+) You describe it… What molecule do you think it is? Take a guess…

13. Inductive Effect 2.1 2.5 3.5 2.5 3.5 acetic acid (ethanoic acid) (this is a carboxylic acid due to the –OH off the carbonyl group) inductive effect – the shifting of electrons in a σ (sigma) bond in response to the electronegativity of nearby atoms. Let’s discuss what it means to be an acid.

14. Inductive Effect Why would HCN allow the H + to be released (proton donor – acid), thus categorizing HCN as an acid, when CH 4 is not usually categorized as an acid?

15. Dipole Moment Calculations dipole moment (μ – Greek mu) – the magnitude of the charge (Q) at either end of the molecular dipole times the distance (r) between the charges measured in debyes (D) μ = Q x r Just be familiar with magnitude of values and that the D following the value is the unit. Section 2.2

16. Dipole Moments 2.1 2.5 3.5 2.5 3.5 acetic acid (ethanoic acid) overall dipole moment = 1.70 D

17. Dipole Moments acetic acid (ethanoic acid) overall dipole moment = 1.70 D

18. Water and Ammonia

19. Dipole Moment Values

20. Non-Polar Molecules

21. You Try It. Draw the complete Lewis Structure for the alcohol, methanol (methyl alcohol). Show the general direction of its dipole moment. (μ =1.70)

22. You Try It. Determine if the following molecules are polar or non polar. Show any bond dipoles. (a) (b) (c)

23. You Try It. Draw Lewis Structures for each of the following molecules and predict whether each has a dipole moment. If you expect a dipole moment, draw it in the correct direction. (a) C 2 HF(b) CCl 4 (c) CH 3 CHO

24. Formal Charges formal charges – these charges do not imply the presence of actual ionic charges …instead they give insight into the distribution of electrons calculating the formal charges of each atom in a molecule will help you determine the best, most favorable structure (lowest energy) Section 2.3

25. General Rules of Stability Lewis structures that approximate the actual molecule most closely are those that have: maximum number of covalent bonds minimum separation of unlike charges formal charges of zero are ideal placement of any negative charges on the most electronegative atom (or any positive charge on the most electropositive atom) Ex. Oxygen would rather 1- then 1+

26. Formal Charges formal charge is calculated in the following manner: If it violates HONC 1234, then it will have a formal charge on it.

27. DMSO (dimethyl sulfoxide)

29. Formal Charges Summary

30. Nitromethane Determine any formal charges on nitromethane, CH 3 NO 2

31. Nitromethane Determine any formal charges on nitromethane, CH 3 NO 2

32. Formal Charges Give the formal charges for any atom on each of the following compounds Recall, having an overall + charge means that there is one less electron CH 4 H 3 O + NH 3 BH 3

33. Formal Charges Give the formal charges for any atom on each of the following compounds Recall, having an overall + charge means that there is one less electron H 2 C=N=NO 3 [H 2 CNH 2 ] + (This has resonance structures.) (1 very likely, 1 less likely, 1 very unlikely)

34. Resonance Structures Some molecules cannot be represented by a single structure. In these cases we draw 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 The resonance forms are connected by a double-headed arrow. Section 2.4

35. Resonance Structures When two or more structures are possible, the molecule will show characteristics of each structure. Experiments show that these two structures are equivalent…both C-O bonds are same length and strength. Since both structures are equally likely, the real structures is most likely a perfect blend of each of these. This is not always the case.

36. Resonance Structures Draw resonance structures for NO 3 - The “real” structure is a resonance hybrid Each oxygen has a partial negative charge

37. Resonance Structures The “real” structure is said to have its electrons delocalized and is represented by a dotted bond Remember, different resonance structures are not always equivalent. Why would the one on the left be least influential?

38. Resonance Structures In some cases, one resonance form is more stable than another (one accommodates formal charges better)

39. Rules for Resonance Forms When drawing resonance structures, follow these rules: 1. Individual resonance forms are imaginary, not real 2. Resonance forms differ ONLY in the placement of their pi or non-bonding electrons 3. Different resonance forms of a substance don’t have to be equivalent 4. All resonance forms must be valid Lewis structures and obey normal rules for valency 5. The resonance hybrid is more stable than any individual resonance form Section 2.5

40. When drawing resonance structures, follow these rules: 1. 2. Resonance forms differ ONLY in the placement of their pi or non-bonding electrons Rules for Resonance Forms

42. When drawing resonance structures, follow these rules: 1. 2. 3. Different resonance forms of a substance don’t have to be equivalent Rules for Resonance Forms

44. When two resonance forms are nonequivalent, the actual structure of the resonance hybrid is closer to the more stable form than to the less stable form. Rules for Resonance Forms

45. When drawing resonance structures, follow these rules: 1. 2. 3. 4. All resonance forms must be valid Lewis structures and obey normal rules for valency 5. Rules for Resonance Forms

47. When drawing resonance structures, follow these rules: 1. 2. 3. 4. 5. The resonance hybrid is more stable than any individual resonance form Rules for Resonance Forms

48. In general, the larger the number of resonance forms, the more stable (lower in energy…less reactive) a substance is because electrons are spread out over a larger part of the molecule and are closer to more nuclei. The more the negative charge is spread out, the better! Rules for Resonance Forms

49. General Trends + C - C + N/O - N/O

50. Remember Curved arrows always represent the movement of electrons, not atoms. Electrons always move towards the more electronegative element or positive charges. Electron pairs can only move to adjacent positions. The Lewis structures that result must be valid and have the same net charges.

51. You Try It! Do the Resonance and Formal Charges Worksheet.

52. Radicals radical - (free radical) a neutral substance that contains a single, unpaired electron in one of its orbitals, denoted by a dot ( · ) leaving it with an odd number of electrons. Radicals are highly reactive! (octet rule) Radicals can form from stable molecules and can also react with each other. Section 2.6

53. Radical Resonance Resonance forms for radicals will depend upon three-atom groupings that contain a multiple bond next to a p-orbital.

54. Pentadienyl Radical

57. You try it. Show all the resonance forms for the straight chained C 7 H 9. radical (in line angle)..

58. Chapter 2 Polar Covalent Bonds; Acids and Bases Part II Organic Chemistry

59. Define and describe acids and bases based on the Brønsted-Lowry and Lewis definitions Use the curved-arrow formalism to show movement of electrons between Lewis acids and bases Determine conjugate acid-base pairs Predict strength of acids and bases based on size, electronegativity, resonance stabilization, hybridization, and induction Predict reactions using pKa values Section 2.7-2.11 Objectives

60. Why Study Acids & Bases? At a deeper level, acid-base strength allows us to predict reactivity. Compounds tend to react in such a way that they become more stable (in the long run). Compounds considered “strong” are called that (technically) because they dissociate completely, but (practically) also because they tend to react quickly. (This occurs because of LOW stability.) Compounds considered “weak” tend not to react quickly or completely because they are stable. (“happy” where they are )

61. Everything wants to be at the lowest possible energy. (most stable) Stability

62. Acids and Bases Definitions of acids and bases: Brønsted-Lowry definition Acids donate protons. (H + ) (proton donor) Hint for recognizing the acid – look for Hs! Bases accept protons. (proton acceptor) Lewis definition Acids accept electrons. (electrophile) Bases donate electrons. (nucleophile) Hint for recognizing the base – look for electrons! Either a lone pair or pi bonded electrons seek out electrophiles. Sections 2.7 & 2.11

63. Lewis Acids and Bases Acid-base reactions can take place with or without a proton. Lewis bases are species that are able to donate a pair of electrons - called nucleophiles (lover of nuclei). Lewis acids are species that can accept this same pair of electrons – called electrophiles (lover of electrons). The reactions are drawn using curved arrow formalism (movement of electrons represented with arrows). Section 2.11

64. Lewis Acids and Bases

65. Group 3A elements, such as B (BF 3 ) and Al (AlCl 3 ), are Lewis acids because they have an unfilled valence p-orbital and can accept electron pairs from Lewis bases. The Lewis definition of acidity includes metal cations, such as Mg 2+ They accept a pair of electrons when they form a bond to a base. Transition-metal compounds, such as TiCl 4, FeCl 3, ZnCl 2, and SnCl 4, are Lewis acids

66. Some Lewis Bases

67. Morphine – Acid and Base

68. Acid Reactions So, according to Brønsted-Lowry definition, if something loses a hydrogen, it has acted as an acid. It then has the capability to accept a proton. Therefore, what is it called at this point? A BASE! Acids will donate a proton (react) to become a conjugate base. Bases will accept a proton (react) to become a conjugate acid.

69. General Acid-Base Reaction When writing reactions, we show the movement (called an “attack”) of electrons with an arrow. Full headed arrow – both electrons Half headed arrow – one electron

70. (electrophile) (nucleophile)

71. BE CAREFUL! Only move electrons!!!! Never start your arrow on an atom. Start it on a pair of electrons. …and since electrons don’t move towards other electrons, be careful where you point your arrow…Point your arrow onto an atom (NOT electrons).

72. Acid-Base Reactions

73. Lewis Acids and Bases Strong nucleophiles are usually very strong Brønsted-Lowry bases.

74. Brønsted-Lowry Theory What is the acid, base, conjugate acid, and conjugate base? acidbase conjugate conjugate acid base (electrophile) (nucleophile)

75. Brønsted-Lowry Theory baseacid conjugate conjugate acid base What is the acid, base, conjugate acid, and conjugate base? (nucleophile) (electrophile)

76. Dual Personality amphoteric – a substance, that depending on the circumstances, can act like an acid or a base (like water!) What makes hydrogen sulfate ion amphoteric?

77. Conjugate Acid-Base Pairs acidconjugate acid baseconjugate base acidbase conjugate acid conjugate base (electrophile) (nucleophile) (nucleophile) (electrophile)

78. Conjugate Acid-Base Pairs The stronger the acid, the weaker the conjugate base. The weaker the acid, the stronger the conjugate base. The stronger the base, the weaker the conjugate acid. The weaker the base, the stronger the conjugate acid.

79. You Try It Write the acid-base reaction for each. Label the acid, base, conj. acid, conj. base. Label the electrophile and nucleophile. Show the curved arrow formalism CH 3 CH 2 OH and NaNH 2 CH 3 COOH and NaOCH 3 CH 3 CH 2 OH and HCl

80. You Try It What is the conjugate base of the following acids? 1. CH 3 COOH 2. CH 3 CH 2 NH 3 + 3.

81. Messing With Stability If I take something that is stable and change it by taking something away from it, what happens? It becomes unstable. Is this good or bad? Ex. CH 3 OH  a weak reagent Pretty stable (How do I know this?) If I remove an “H” – CH 3 O - Not stable at all  a strong reagent

82. Acid and Base Strength Up to this point, the terms we have been using to describe acid-base strength have been very relative. Actual numerical values exist. Recall the K a value Section 2.8

83. The Equilibrium Expression (Law of Mass Action) The relationship between the concentration of products and reactants at equilibrium can be expressed by K

84. Acid Dissociation Constant What if it is a reaction for the dissociation of an acid? What does the size of K a mean? High K a = strong acid Low K a = weak acid acid dissociation constant

85. Acid Strength IN CHEMICAL REACTIONS, THE ARROW USUALLY FAVORS THE PRODUCTION OF A WEAKER ACID AND BASE!!! Why? What favors a weak acid over a strong one? Weak acids and bases are more STABLE. If they weren’t stable, they would react to become stable…that’s why they are weak! Strong acid Weak acidStrong base Weak base What kind of values do you expect?

86. K a vs. pK a Acids with a greater K a value are stronger than acids with a smaller K a value Problem with K a  relatively inconvenient because K a values are usually on a negative power of ten Example: 1.0 x 10 -4 To make things easier, the value pK a is used:

87. Calculating pK a Determine the pK a of Hydrofluoric acid: K a = 3.5 x 10 -4 Phosphoric Acid: K a = 7.5 x 10 -3 ~Which of these acids is stronger? H 3 PO 4 pK a of HF: pK a of H 3 PO 4 : What do you notice about pK a value compared to acid strength? 3.5 2.1 The smaller the pKa, the stronger the acid.

88. pKa and Acid Strength The smaller the pKa, the stronger the acid. (Remember this!!!)

89. Table 2.3, page 51

90. Predicting Acid-Base Reactions from pKa Values Do all acids react with all bases? NO!!!!!!!!!!!! How do we know when an acid will react with a particular base? pKa values Weak acids won’t produce strong acids…so, compare the pKa values of the acid and the conjugate acid to reveal which reaction proceeds, the forward or the reverse. Section 2.9

91. Compare the strengths of the acid and the conjugate acid. Predicting Acid-Base Reactions from pKa Values

92. You Try It

93. Will the following reaction occur? pKa = 49pKa = 16

94. You Try It Write the products of the reaction and determine if it will occur.

95. You Try It

96. Predicting Acid-Base Strength without pKa values Use the pK a values if they are handy. Otherwise… 5 major factors exist which affect acid strength 1. Electronegativity 2. Size 3. Resonance stabilization (delocalization) 4. Hybridization 5. Induction

97. Predicting Acid-Base Strength without pKa values

98. 1. Electronegativity

99. Which will give up a hydrogen ion (proton) more readily? CH 4, NH 3, H 2 O, HF HF is most electronegative therefore the HF bond is shared unequally and easier to break THE MORE ELECTRONEGATIVE THE CONJUGATE BASE, THE STRONGER THE ACID

100. 2. Size

101. Which is most reactive? HF, HCl, HBr, HI Recall: If the negative charge is spread out more, it is a more stable conjugate base…therefore…

102. 3. Resonance Stabilization Again…if the charge on the conjugate base is spread out more than it is a more stable conjugate base. Therefore, the original acid is a stronger acid. …so, how does resonance help this? The negative charge of a conjugate base may be delocalized over several atoms thus making it more stable.

103. 3. Resonance Stabilization The negative charge of a conjugate base may be delocalized over several atoms thus making it more stable.

104. You Try It Which is the strongest acid? CH 3 CH 2 OH, CH 3 COOH, CH 3 SO 3 H

105. 4. Hybridization H 3 C-CH 3 < H 2 C=CH 2 < HC≡CH worst acidbest acid sp 3 sp 2 sp 25%-s 33%-s 50%-s The more percent “s” in character, the closer the electrons are to the nucleus, therefore the more polarized the structure…the H becomes more positive due to the pull of e - towards the C – makes a better acid

106. 4. Hybridization

107. 5. Inductive Effects – e - Withdrawing Electronegative elements “take away” electron density from a negative charge: Stability increases

108. 5. Inductive Effects – e - Withdrawing

109. 5. Inductive Effects – e - Donating hyperconjugation - Donation of a pair of bonding electrons into an unfilled or partially filled orbital

110. 5. Inductive Effects – e - Donating Which is the most stable conjugate base? somewhat destabilizing very destabilizing!

113. Molecular Models Section 2.12

114. Noncovalent Interactions Intermolecular Forces of Attraction dipole-dipole interactions (polar molecules) London forces or dispersion forces (All molecules have this but it is the only force present in nonpolar molecules) hydrogen bonding (polar molecules with F, O, or N bonded to a H) hydrophilic – water loving (attracted to water) hydrophobic – water fearing (not attracted to water) Section 2.13

115. Dipole-dipole Forces Dipole-dipole interactions are attractions between the partial negative ends and partial positive ends of two different molecules (polar molecules).

116. Dispersion Forces London forces or dispersion forces are attractions due to temporary dipoles. All molecules have this but it is the only force present in nonpolar molecules.

117. Hydrogen Bonding Hydrogen bonding is the attraction between a low electronegative H on one molecule and a high electronegative F, O, or N bonded to an H on another molecule.

118. Hydrogen Bonding