1. Chapter 2

2. The Lewis Model of Bonding:The Lewis Model of Bonding: Atoms bond together in such a way that each atom participating in a chemical bond acquires a completed valence-shell electron configuration resembling that of the noble gas nearest it in the Periodic Table. Introduction

3. Ionic Bond –Ionic Bond – is a chemical bond resulting from the electrostatic attraction between an anion and a cation. Na + F Na + + F -....

4. Covalent Bond –Covalent Bond – is a chemical bond formed by the sharing of electron pairs between atoms. 2 H· H 2 Based on the the degree of electron sharing, covalent bonds can be divided into: Nonpolar covalent bondNonpolar covalent bond Polar covalent bondPolar covalent bond

5. I. Polar Covalent Bonds I. Polar Covalent Bonds A.Electronegativity B.Dipole Moments A.Electronegativity B.Dipole Moments

6. Polar Covalent Bond –Polar Covalent Bond – is a covalent bond with ionic character.  Bonding electrons are attracted more strongly by one atom than by the other  Electron distribution between atoms is not symmetrical

7. A.Polar Covalent Bonds: Electronegativity Electronegativity (EN):Electronegativity (EN):  is a measure of the force of an atom’s attraction for electrons it shares in a covalent bond with another atom.  is an intrinsic ability of an atom to attract the shared electrons in a covalent bond

8. Bond Polarity and Electronegativity EN of an atom is related to its ionization energy (IE) and electron affinity (EA).  E = IEAtom in ground state (g) + Energy  Atom + (g) + e -  E = IE  E = EAAtom (g) + e -  Atom -  E = EA bond polarityDifferences in EN produce bond polarity. Bond Polarity ~  EN ~  EN  EN

9. Bond Polarity and Electronegativity EN is based on an arbitrary scale. Linus PaulingThe most widely used scale of EN was devised by Linus Pauling in the 1930’s and is based on bond energies. Important EN values: F is the most electronegative (EN = 4.0) Cs is the least electronegative (EN = 0.7) C has an EN = 2.5

10. The Periodic Table and Electronegativity

11. Metals on left side of periodic table attract electrons weakly, lower EN

12. The Periodic Table and Electronegativity Halogens and other reactive nonmetals on right side of periodic table attract electrons strongly, higher EN

13. Classification of Covalent Bonds Difference in EN between bonded Atoms Type of Bond Similar EN Nonpolar Covalent Less than 2 Polar Covalent Greater than 2 Ionic

14. Classification of Covalent Bonds Difference in EN between bonded Atoms Type of Bond Less than 0.5 Nonpolar Covalent 0.5 to 1.9 Polar Covalent Greater than 1.9 Ionic

15. Classification of Covalent Bonds: Examples C–H bonds are relatively nonpolar  EN = EN C – EN H = 2.5 - 2.1 = 0.4 C-O and C-X bonds (more electronegative elements) are polar  EN = EN O – EN C = 3.5 - 2.5 = 1

16. Bonding electrons are drawn toward electronegative atom  C acquires partial positive charge,  +  Electronegative atom acquires partial negative charge,  - Bond Polarity and Inductive Effect The crossed arrow indicates the direction of the electron displacement

17. Inductive EffectInductive Effect – is the shifting of electrons in a bond in response to EN of nearby atoms  Metals (Li and Mg) inductively donate e-  Reactive nonmetals (O and Cl) inductively withdraw e- Bond Polarity and Inductive Effect

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

19. Practice Problem: Which element in each of the following pairs is more electronegative? a.Li or H b.B or Br c.Cl or I d.C or H

20. Practice Problem: Use the  +/  - convention to indicate the direction of expected polarity for each of the bonds indicated a.H 3 C — Br b. H 3 C — NH 2 c.H 3 C — Li d.H 2 N — H e.H 3 C — OH f.H 3 C — MgBr g. H 3 C — F

21. Practice Problem: Use the electronegativity values to rank the following bonds from least polar to most polar: H 3 C — Li H 3 C — K H 3 C — F H 3 C — MgBr H 3 C — OH

22. Practice Problem: Look at the following electrostatic potential map of methyl alcohol, and tell the direction of polarization of the C-O bond:

23. B.Polar Covalent Bonds: Dipole Moments Molecular PolarityMolecular Polarity  is the tendency of molecules as a whole to be polar  results from vector summation of individual bond polarities and lone-pair contributions “Like dissolves like”“Like dissolves like” Strongly polar substances are soluble in polar solvents like water; nonpolar substances are insoluble in water.

24. To predict whether a molecule is polar, determine: 1. if the molecule has polar bonds, and 2. the arrangements of these bonds in space

25. Dipole moment (  )Dipole moment (  ) – is a measure of net molecular polarity, due to difference in summed charges  = Q  r magnitude of charge Q at either end of molecular dipole distance r between charges It is expressed in Debyes (D) 1 D = 3.336 x 10 -30 Coulomb meter

26. r average covalent bond = 100pm Q electron = 1.60 x 10 -19 C The dipole moment (  ) of an average covalent bond is 4.80 D Calculating the dipole moment of an average bond:

27. r Given that r C-Cl = 178 pm and assuming  C-H is negligible, then   calculated CH 3 Cl = 178 pm x 4.80 D = 8.5 D   measured CH 3 Cl = 1.87D  % ionic character =  measured /  calculated x 100 = 22% Calculating Ionic Character: Chloromethane

28. H 2 O and NH 3 have large dipole moments:  EN O and EN N > EN H  Both O and N have lone-pair electrons oriented away from all nuclei Dipole Moments in Water and Ammonia

29. H 2 O and NH 3 have large dipole moments:  EN O and EN N > EN H  Both O and N have lone-pair electrons oriented away from all nuclei

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

31. Practice Problem: Carbon dioxide, CO 2, has zero dipole moment even though carbon-oxygen bonds are strongly polarized. Explain.

32. Practice Problem: Make three-dimensional drawings of the following molecules, and predict whether each has a dipole moment. If you expect a dipole moment, show its direction. a.H 2 C = CH 2 b.CHCl 3 c.CH 2 Cl 2 d.H 2 C = CCl 2

33. II. Formal Charges and Resonance II. Formal Charges and Resonance A.Formal Charges B.Resonance A.Formal Charges B.Resonance

34. A.Formal Charges Formal Charge -Formal Charge - is the charge on an atom in a molecule or polyatomic ion

35. No formal charge Comparing the bonding of the atom in the molecule to the valence electron structure

36. Calculating the formal charge

37. 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 dipolarNeutral molecules with both a “+” and a “-” are dipolar

39. Practice Problem: Dimethyl sulfoxide, a common solvent, has the structure indicated. Show why dimethyl sulfoxide must have formal charges on S and O.

40. Practice Problem: Calculate formal charges for the nonhydrogen atoms in the following molecules: a.Diazomethane, b.Acetonitrile oxide, c.Methyl isocyanide

41. Practice Problem: Organic phosphates occur commonly among biological molecules. Calculate formal charges on the four O atoms in the methyl phosphate ion.

42. B.Resonance Some molecules have Lewis structures that cannot be shown with a single representation 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)

43. Resonance Forms Resonance formsResonance forms – are Lewis structures of the same molecule whose only difference is the placement of  and nonbonding valence electrons (= delocalized). The atoms occupy the same place in the different forms The connections between atoms are the same. double-headed arrowThe resonance forms are connected by a double-headed arrow

44. 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

45. Resonance Hybrids: Benzene For example, benzene (C 6 H 6 ) has two resonance forms with alternating double and single bonds –In the resonance hybrid, the actual structure, all its C-C bonds are equivalent, midway between double and single

46. Rules for Resonance Forms 1 1.Individual resonance forms are imaginary 1.Individual resonance forms are imaginary - the real structure is a hybrid (only by knowing the contributors can you visualize the actual structure)

47. Rules for Resonance Forms 2 2.Resonance forms differ only in the placement of their  or nonbonding electrons

48. Rules for Resonance Forms 3 3.Different resonance forms of a substance don’t have to be equivalent

49. Rules for Resonance Forms 4 4.Resonance forms must be valid Lewis structures: the octet rule applies

50. Rules for Resonance Forms 5 5.The resonance hybrid is more stable than any individual resonance form Resonance leads to stability. The larger the # of the resonance forms, the more stable the substance

51. Rules for Resonance Forms 1.Individual resonance forms are imaginary - the real structure is a hybrid (only by knowing the contributors can you visualize the actual structure) 2.Resonance forms differ only in the placement of their  or nonbonding electrons 3.Different resonance forms of a substance don’t have to be equivalent 4.Resonance forms must be valid Lewis structures: the octet rule applies 5.The resonance hybrid is more stable than any individual resonance form

52. 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 arrowA 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

53. 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 The “enolate” derived from acetone is a good illustration, with delocalization between carbon and oxygen

54. Drawing Resonance Forms Any three-atom grouping with a p orbital on each atom has two resonance forms X, Y, Z can be C, N, O, P or S

55. Resonance Structures: 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

56. Relative Importance of contributing Structures The most important contributing structures have: 1.filled valence shells 2.a maximum number of covalent bonds 3.the least separation of unlike charge, and 4.any negative charge on a more electronegative atom and/or any positive charge on a less electronegative atom

57. Practice Problem: Draw the indicated number of resonance structures for each of the following species: a.The nitrate ion, NO 3 - (3) b.The allyl cation, H 2 C=CH-CH 2 + (2) c.Hydrazoic acid, (2) d.(2)

58. III. Acids and Bases 1 III. Acids and Bases 1 A.The Brønsted-Lowry Definition B.Acid and Base Strength C.Predicting Acid-Base Reactions from pKa Values A.The Brønsted-Lowry Definition B.Acid and Base Strength C.Predicting Acid-Base Reactions from pKa Values

59. III. Acids and Bases 2 III. Acids and Bases 2 D.Organic Acids and Organic Bases E.The Lewis Definition D.Organic Acids and Organic Bases E.The Lewis Definition

60. “acid”“base”  The terms “acid” and “base” can have different meanings in different contexts. For that reason, we specify the usage with more complete terminology There are three definitions of acid-base: Arrhenius, Bronsted-Lowry, and Lewis Introduction

61. ArrheniusAccording to the Arrhenius definition: An Arrhenius acidH +An Arrhenius acid - is a substance that donates H + in aqueous solutions An Arrhenius baseOH -An Arrhenius base - is a substance that donates OH - in aqueous solutions ArrheniusThe Arrhenius definition is not useful in organic chemistry.

62. A. The Brønsted-Lowry Definition Brønsted–Lowry theoryBrønsted–Lowry theory defines acids and bases by their role in reactions that transfer protons (H + ) between donors and acceptors “Brønsted-Lowry” “Brønsted”“Brønsted-Lowry” is usually shortened to “Brønsted”

63. Brønsted Acids and Bases A Brønsted acid is a substance that donates a hydrogen ion (H + ) A Brønsted base is a substance that accepts the H + “proton” is a synonym for H + (i.e. Loss of the valence electron from a neutral H leaves only the hydrogen nucleus - a proton)

64. Brønsted acid–base Reactions

65. The Reaction of HCl with H 2 O When HCl gas dissolves in water, a Brønsted acid–base reaction occurs HCl donates a proton to water molecule, yielding hydronium ion (H 3 O + ) and Cl -

66. The reverse is also a Brønsted acid–base reaction of the conjugate acid and conjugate base Other Brønsted acid–base Reactions

67. Practice Problem: Nitric acid (HNO 3 ) reacts with ammonia (NH 3 ) to yield ammonium nitrate. Write the reaction, and identify the acid, the base, the conjugate acid product, and the conjugate base product.

68. B. Acid and Base Strength The equilibrium constant (K eq ) 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 K eq

69. K a – the acidity constant The concentration of water as a solvent does not change significantly when it is protonated  Molecular weight (H 2 O) = 18 g/mol.  Density (H 2 O) = 1000 g/l  [H 2 O] ~ 55.6 M at 25°C

70. K a – the acidity constant The acidity constant, K a for HA is K a ranges from 10 15 for the strongest acids to very small values (10 -60 ) for the weakest

71. Acid-Base Strength The “ability” of a Brønsted acid to donate a proton is sometimes referred to as the strength of the acid. The strength of the acid is 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

72. pK a – The acid strength scale Acid strength is expressed using pKa values: pKa = - log Ka The free energy in an equilibrium is related to - log of K eq :  G = -RT ln K eq A larger value of K a indicates a stronger acid and is proportional to the energy difference between products and reactants

73. pK a – The acid strength scale The pK a of water is 15.74

74. strongerweake weakerstrongerThe stronger the acid, the weaker its conjugate base. The weaker the acid, the stronger the conjugate base.

75. Practice Problem: Formic acid, HCO 2 H, has pKa = 3.75, and picric acid, C 6 H 3 N 3 O 7, has pKa = 0.38. Which is the stronger acid?

76. Practice Problem: Amide ion, H 2 N -, is a much stronger base than hydroxide ion, HO -. Which would you expect to be a stronger acid, NH 3 or H 2 O? Explain.

77. C. Predicting Acid-Base Reactions from pKa Values pK a values are related as logarithms to equilibrium constants The difference in two pK a values is the log of the ratio of equilibrium constants:  pKa = log K eq2 /K eq1   pKa can be used to calculate the extent of H + transfer

78. H + will always go from the stronger acid to the stronger base

79. The product acid must be weaker and less reactive than the starting acid The product base must be weaker and less reactive than the starting base

80. The stronger base holds the proton more tightly

81. Practice Problem: Will either of the following reactions take place as written, according to the pKa data? a.HCN + CH 3 CO 2 - Na +  Na + - CN + CH 3 CO 2 H b.CH 3 CH 2 OH + Na +- CN  CH 3 CH 2 O - Na + + HCN ? ?

82. Practice Problem: Ammonia, NH 3, has pKa = 36 and acetone has pKa = 19. Will the following reaction take place?

83. Practice Problem: What is the Ka of HCN if its pKa = 9.31?

84. D. Organic Acids and Organic Bases The reaction patterns of organic compounds often are acid-base combinations The transfer of a H + from a strong Brønsted acid to a Brønsted base –is a very fast process –will always occur along with other reactions

85. Organic Acids Organic acids lose a proton either from: O–H, such as methanol and acetic acid C–H, usually from a carbon atom next to a C=O double bond (O=C–C–H)

86. Organic Acids The conjugate base resulting from loss of H + is stabilized by having its negative charge on a highly electronegative oxygen atom  Acidity

87. Organic Acids

88. Organic Bases an atom with a lone pair of electronsOrganic 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

89. Organic Bases

90. E. The Lewis Definition A Lewis acidA Lewis acid is an electron pair acceptor; it has a low-energy empty orbital A Lewis baseA Lewis base is an electron pair donor

91. Brønsted acids are not Lewis acids because they cannot accept an electron pair directly protonLewis acid –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 pK a

92. Lewis Acids Lewis acids include: H +H + Metal cations They accept a pair of electrons when they form a bond to a baseMetal cations, such as Mg 2+ : They accept a pair of electrons when they form a bond to a base Group 3A elements They have unfilled valence orbitals and can accept electron pairs from Lewis basesGroup 3A elements, such as BF 3 and AlCl 3 : They have unfilled valence orbitals and can accept electron pairs from Lewis bases Transition-metal They have unfilled valence orbitals and can accept electron pairs from Lewis basesTransition-metal compounds, such as TiCl 4, FeCl 3, ZnCl 2, and SnCl 4 : They have unfilled valence orbitals and can accept electron pairs from Lewis bases

93. Lewis Acids

94. electrophilesLewis AcidsOrganic compounds that undergo addition reactions with Lewis bases are called electrophiles and therefore Lewis Acids The combination of a Lewis acid and a Lewis base can be shown with a curved arrow from base to acid

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

96. Lewis Bases 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

97. Lewis Bases

98. Practice Problem: Using curved arrows, show how the species in part (a) can act as Lewis bases in their reactions with HCl, and show how the species in part (b) can act as Lewis acids in their reaction with OH - a.CH 3 CH 2 OH, HN(CH 3 ) 2, P(CH 3 ) 3 b.H 3 C +, B(CH 3 ) 3, MgBr 2

99. Practice Problem: Explain by calculating formal charges why the following acid-base reaction products have the charges indicated: a.F 3 B — O — CH 3 | CH 3 - + b.Cl 3 Al — N — CH 3 | CH 3 | - +

100. Practice Problem: The organic compound imidazole can act as both an acid and a base. Look at the following electrostatic potential map, and identify the most acidic hydrogen atom and the most basic nitrogen in imidazole

101. IV. Chemical Structures IV. Chemical Structures A.Drawing Chemical Structures B.Molecular Models A.Drawing Chemical Structures B.Molecular Models

102. A. Drawing Chemical Structure Chemists use shorthand ways for writing structures. Organic molecules are usually drawn as: Condensed structures Skeletal structures

103. C-H and C-C single bonds aren't shown but understood - If C has 3 H’s bonded to it, write CH 3 - If C has 2 H’s bonded to it, write CH 2 ; and so on. Horizontal bonds between carbons aren't shown in condensed structures—the CH 3, CH 2, and CH units are simply placed next each other but vertical bonds are added for clarity Condensed Structures

104. Skeletal Structures Minimum amount of information but unambiguous Rules for Skeletal Structures:Rules for Skeletal Structures: 1.C’s are not shown 1.C’s are not shown, assumed to be at each intersection of two lines (bonds) and at end of each line 2.H’s bonded to C’s aren't shown 2.H’s bonded to C’s aren't shown – whatever number is needed will be there 3.All atoms other than C and H are shown

106. Practice Problem: Tell how many hydrogens are bonded to each carbon in the following compounds, and give the molecular formula of each substance:

107. Practice Problem: Propose skeletal structures for compounds that satisfy the following molecular formulas (there is more than one possibility in each case): a.C 5 H 12 b.C 2 H 7 N c.C 3 H 6 O d.C 4 H 9 Cl

108. Practice Problem: The following molecular model is a representation of para-aminobenzoic acid (PABA), the active ingredient in many sunscreens. Indicate the positions of the multiple bonds, and draw a skeletal structure (gray = C, red = O, blue = N, ivory = H)

109. B. Molecular Model We often need to visualize the shape or connections of a molecule in three dimensions Molecular modelsMolecular 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) Space-filling Framework

110. Space-filling Framework Drawings on paper and screens are limited in what they can present to you 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

111. Chapter 2