1. University of Isfahan Department of Chemistry FUNDAMENTALS OF ORGANIC CHEMISTRY FOR BIOLOGY STUDENTS Dr. Gholam Ali Koohmareh نيمسال اول 91-90

2. 2 Text Book: Fundamentals of Organic Chemistry By: John. Mc. Murry مباني شيمي آلي نوشته جان مك موري ترجمه دكتر عيسي ياوري

3. نصايحي چند در مورد مطالعه شيمي آلي

4. ماده مفقوده ساختار دانش در شيمي آلي ساختارها شبيه هرم متعادلي است كه روي نوك ايستاده است

5. كل ساختار خراب ميشود !!! اگر چيزي گم شده باشد.....

6. براي درك ساختار مواد بايد مطالعه كنيد. هرمطلبي را كه مسلط نميشويد چندين بار تكرار كنيد. ما از مباحث ساده شروع ميكنيم و وارد بحثهاي پيچيده تر ميشويم.

7. اگر مطلبي را متوجه نميشويد هر چه زودتر و قبل از انباشته شدن مطالب كمك بگيريد.... ميتوانيد بعد از كلاس يا در طول ساعات اداري در اطاق من در گروه شيمي و يا از طريق ايميل مشكل درسي خود را مطرح كنيد Email: Koohmareh@gmail.com هر زماني كه وقت داشتيد تمرين حل كنيد. انجام تمرينات درك شما را بيشتر ميكند و اين مستلزم آن است كه هر چه ياد ميگيريد بكار گيريد. تمرينات آخر فصل بسيار مفيدند. يادگيري فعال بسيار موثرتر ازمطالعه اجباري است !

8. %90 چيزي را كه ميگوييم و بكار ميگيريم %70 چيزي را كه ميگوييم و صحبت ميكنيم %50 چيزي را كه ميبينيم و ميشنويم ما حفظ ميكنيم %30 چيزي را كه ميبينيم %20 چيزي را كه ميشنويم %10 چيزي را كه ميخوانيم ….. يك مطالعه روانشناسي عجيب ولي واقعي فعال اجباري

9. فصل اول. ساختار و پيوند

10. 10 Organic Chemistry “Organic” – until mid 1800’s referred to compounds from living sources (mineral sources were “inorganic”) Wöhler in 1828 showed that urea, an organic compound, could be made from a non-living source:

11. 11 Organic Chemistry Today, organic compounds are those based on carbon structures and organic chemistry studies their structures and reactions Includes biological molecules, drugs, solvents, dyes, food additives, pesticides, and others. Does not include metal carbonate salts and other simple ionic compounds (inorganic)

12. 12 Organic Chemistry

13. 13 1.1 Atomic Structure Structure of an atom Positively charged nucleus (very dense, protons and neutrons) and small (10 -15 m) Negatively charged electrons are in a cloud (10 -10 m) around nucleus Diameter is about 2  10 -10 m (200 picometers (pm)) [the unit angstrom (Å) is 10 -10 m = 100 pm]

14. 14 1.1 Atomic Structure

15. 15 Atomic Number & Mass Number The atomic number (Z) is the number of protons in the atom's nucleus The mass number (A) is the number of protons plus neutrons All the atoms of a given element have the same atomic number

16. 16 Atomic Number and Atomic Mass All the atoms of a given element have the same atomic number Isotopes are atoms of the same element that have different numbers of neutrons and therefore different mass numbers The atomic mass (atomic weight) of an element is the weighted average mass in atomic mass units (amu) of an element’s naturally occurring isotopes

17. 17 1.2 Atomic Structure: Orbitals Quantum mechanics: describes electron energies and locations by a wave equation Wave function is a solution of the wave equation Each Wave Function describes an orbital,  A plot of  2 describes the probabililty of finding the electron The electron cloud has no specific boundary so we show the most probable volume. The boundary is called a “probablility surface”

18. 18 Shapes of Atomic Orbitals for Electrons For the existing elements, there are four different kinds of orbitals for electrons based on those derived for a hydrogen atom. Denoted s, p, d, and f s and p orbitals most important in organic chemistry (carbon atoms have no d or f orbitals)

19. 19 Shapes of Atomic Orbitals for Electrons s orbitals: spherical, nucleus at center p orbitals: dumbbell-shaped, nucleus at middle

20. 20 Orbitals and Shells Orbitals are grouped in shells of increasing size and energy As shells increase in size, they contain larger numbers and kinds of orbitals Each orbital can be occupied by two electrons (Pauli Exclusion Principle)

21. 21 Orbitals and Shells First shell contains one s orbital, denoted 1s, holds only two electrons Second shell contains one s orbital (2s) and three p orbitals (2p): eight electrons Third shell contains an s orbital (3s), three p orbitals (3p), and five d orbitals (3d): 18 electrons

22. 22 Orbitals and Shells

23. 23 p-Orbitals In each shell there are three perpendicular p orbitals, p x, p y, and p z, of equal energy Lobes of a p orbital are separated by region of zero electron density, a node

24. 24 1.3 Atomic Structure: Electron Configurations Ground-state electron configuration of an atom lists orbitals occupied by its electrons. 1. Lowest-energy orbitals fill first: 1s  2s  2p  3s  3p  4s  3d (Aufbau (“build-up”) principle) 2. Electron spin can have only two orientations, up  and down . Only two electrons can occupy an orbital, and they must be of opposite spin (Pauli exclusion principle) to have unique wave equations 3. If two or more empty orbitals of equal energy are available, electrons occupy each with spins parallel until all orbitals have one electron (Hund's rule).

25. 25 Electron Configurations

26. 26 Electron Configurations

27. 27 1.4 Development of Chemical Bonding Theory Kekulé and Couper independently observed that carbon always has four bonds van't Hoff and Le Bel proposed that the four bonds of carbon have specific spatial directions Atoms surround carbon as corners of a tetrahedron Note that a wedge indicates a bond is coming forward Note that a dashed line indicates a bond is behind the page

28. 28 Tetrahedral Carbon:

29. 29 1.5 The Nature of the Chemical Bond 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 transfers Organic compounds have covalent bonds from sharing electrons (G. N. Lewis, 1916)

30. 30 1.5 The Nature of the Chemical Bond Lewis structures shown valence electrons of an atom as dots Hydrogen has one dot, representing its 1s electron Carbon has four dots (2s 2 2p 2 ) Stable molecule results when the outermost (valence) shell is completed: octet (eight dots) for main-group atoms (two for hydrogen)

31. 31 Number of Covalent Bonds to an Atom Atoms with one, two, or three valence electrons form one, two, or three bonds Atoms with four or more valence electrons form as many bonds as they need electrons to fill the s and p levels of their valence shells to reach a stable octet

32. 32 Number of Covalent Bonds to an Atom

33. 33 Valences of Carbon Carbon has four valence electrons (2s 2 2p 2 ), forming four bonds (CH 4 )

34. 34 Valences of Oxygen Oxygen has six valence electrons (2s 2 2p 4 ) but forms two bonds (H 2 O)

35. 35 Valences of Nitrogen Nitrogen has five valence electrons (2s 2 2p 3 ) but forms only three bonds (NH 3 )

36. 36 Non-bonding electrons Valence electrons not used in bonding are called nonbonding electrons, or lone-pair electrons Nitrogen atom in ammonia ( NH 3 ) Shares six valence electrons in three covalent bonds and remaining two valence electrons are nonbonding lone pair

37. 37 Some simple molecules:

38. 38 1.6 Valence Bond Theory Covalent bond forms when two atoms approach each other closely so that a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom

39. 39 1.6 Valence Bond Theory Electrons are paired in the overlapping orbitals and are attracted to nuclei of both atoms H–H bond results from the overlap of two singly occupied hydrogen 1s orbitals H-H bond is cylindrically symmetrical, sigma (  ) bond

40. 40 Bond Energy Reaction 2 H·  H 2 releases 436 kJ/mol Product has 436 kJ/mol less energy than two atoms: H–H has bond strength of 436 kJ/mol. (1 kJ = 0.2390 kcal; 1 kcal = 4.184 kJ)

41. 41 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

42. 42 What is a Bond? An Energy Minimum

43. 43 1.7 Hybridization: sp 3 Orbitals and the Structure of Methane Carbon has 4 valence electrons (2s 2 2p 2 ) In CH 4, all C–H bonds are identical (tetrahedral) sp 3 hybrid orbitals: s orbital and three p orbitals combine to form four equivalent, unsymmetrical, tetrahedral orbitals (sppp = sp 3 ), Pauling (1931)

44. 44 sp 3 Hybridization

45. 45 Tetrahedral Structure of Methane sp 3 orbitals on C overlap with 1s orbitals on 4 H atom to form four identical C-H bonds Each C–H bond has a strength of 438 kJ/mol and length of 110 pm Bond angle: each H–C–H is 109.5°, the tetrahedral angle.

46. 46 Tetrahedral Structure of Methane

47. 47 1.8 Hybridization: sp 3 Orbitals and the Structure of Ethane Two C’s bond to each other by  overlap of an sp 3 orbital from each Three sp 3 orbitals on each C overlap with H 1s orbitals to form six C–H bonds

48. 48 1.8 Hybridization: sp 3 Orbitals and the Structure of EthaneEthane C–H bond strength in ethane 420 kJ/mol C–C bond is 154 pm long and strength is 376 kJ/mol All bond angles of ethane are tetrahedral

49. 49 1.9 Hybridization: sp 2 Orbitals and the Structure of EthyleneEthylene

50. 50 1.9 Hybridization: sp 2 Orbitals and the Structure of Ethylene sp 2 hybrid orbitals: 2s orbital combines with two 2p orbitals, giving 3 orbitals (spp = sp 2 ) sp 2 orbitals are in a plane with120° angles Remaining p orbital is perpendicular to the plane 90  120 

51. 51 Bonds From sp 2 Hybrid Orbitals Two sp 2 -hybridized orbitals overlap to form a  bond p orbitals overlap side-to-side to formation a pi (  ) bond sp 2 –sp 2  bond and 2p–2p  bond result in sharing four electrons and formation of C-C double bond Electrons in the  bond are centered between nuclei Electrons in the  bond occupy regions are on either side of a line between nuclei

52. 52 Bonds From sp 2 Hybrid Orbitals

53. 53 Structure of EthyleneEthylene H atoms form  bonds with four sp 2 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 133 pm (C–C 154 pm)

54. 54 1.10 Hybridization: sp Orbitals and the Structure of Acetylene C-C a triple bond sharing six electrons

55. 55 1.10 Hybridization: sp Orbitals and the Structure of Acetylene 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

56. 56 Orbitals of Acetylene Two sp hybrid orbitals from each C form sp–sp  bond p z orbitals from each C form a p z –p z  bond by sideways overlap and p y orbitals overlap similarly

57. 57 Bonding in AcetyleneAcetylene Sharing of six electrons forms C  C Two sp orbitals form  bonds with hydrogens

58. 58 Comparison of C-C & C-H bonds:

59. 59 Comparison of C-H bonds: MoleculeBondEnergy (kcal)Length (pm) Ethane C sp3 -H100110 Ethylene C sp2 -H106108 Acetylene C sp -H132106

60. 60 Comparison of C-C bonds: MoleculeBondEnergy (kcal)Length (pm) Ethane C sp3 -C sp3 90154 Ethylene C sp2 -C sp2 146133 Acetylene C sp -C sp 200120

61. 61 1.11 Hybridization of Nitrogen and Oxygen Elements other than C can have hybridized orbitals H–N–H bond angle in ammonia (NH 3 ) 107.3° N’s orbitals (sppp) hybridize to form four sp 3 orbitals One sp 3 orbital is occupied by two nonbonding electrons, and three sp 3 orbitals have one electron each, forming bonds to H

62. 62 Hybridization of Oxygen in Water The oxygen atom is sp 3 -hybridized Oxygen has six valence-shell electrons but forms only two covalent bonds, leaving two lone pairs The H–O–H bond angle is 104.5°

63. 63 1.12 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.

64. 64 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

65. 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  The reverse is also a Brønsted acid– base reaction of the conjugate acid and conjugate base

66. 66 The Reaction of HCl with H 2 O

67. 67

68. 68 Quantitative Measures of Acid Strength The equilibrium constant (K e ) 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 e Note that brackets [ ] indicate concentration, moles per liter, M.

69. 69 K a – the Acidity Constant The concentration of water as a solvent does not change significantly when it is protonated in dilute solution. The acidity constant, K a for HA equals K e times 55.6 M (leaving [water] out of the expression) K a ranges from 10 15 for the strongest acids to very small values (10 -60 ) for the weakest

70. 70 K a – the Acidity Constant

71. 71 1.13 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

72. 72 pK a – the Acid Strength Scale pK a = -log K a (in the same way that pH = -log [H+] The free energy in an equilibrium is related to –log of K eq (  G = -RT log K eq ) A larger value of pK a indicates a stronger acid and is proportional to the energy difference between products and reactants

73. 73 pK a – the Acid Strength Scale The pK a of water is 15.74

74. 74 pK a – the Acid Strength Scale

75. 75 1.14 Predicting Acid–Base Reactions from pK a 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, and can be used to calculate the extent of transfer

76. 76 Predicting Acid–Base Reactions from pK a Values

77. 77 Predicting Acid–Base Reactions from pK a Values

78. 78 Prob’s 2.14 & 2.15: Will these reactions take place as written? pK a = 19 pK a = 36

79. 79 1.15 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

80. 80 Some organic acids:

81. 81 Organic Acids Those that lose a proton from O–H, such as methanol and acetic acid Those that lose a proton from C–H, usually from a carbon atom next to a C=O double bond (O=C–C–H)

82. 82 Organic Acids

83. 83 Conjugate bases:

84. 84 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

85. 85 Organic Bases

86. 86 1.16 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 pK a

87. 87 Lewis Acids and the Curved Arrow Formalism 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 Group 3A elements, such as BF 3 and AlCl 3, are Lewis acids because they have unfilled valence orbitals and can accept electron pairs from Lewis bases Transition-metal compounds, such as TiCl 4, FeCl 3, ZnCl 2, and SnCl 4, are Lewis acids

88. 88 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

89. 89 Illustration of Curved Arrows in Following Lewis Acid-Base Reactions

90. 90 BF 3 as a Lewis Acid: BF 3 O(CH 3 ) 2 F 3 B-O(CH 3 ) 2

91. 91 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

92. 92 Lewis Bases

93. 93 Prob. : Identify acids & bases

94. 94 Summary Organic chemistry – chemistry of carbon compounds Atom: positively charged nucleus surrounded by negatively charged electrons Electronic structure of an atom described by wave equation Electrons occupy orbitals around the nucleus. Different orbitals have different energy levels and different shapes s orbitals are spherical, p orbitals are dumbbell-shaped Covalent bonds - electron pair is shared between atoms Valence bond theory - electron sharing occurs by overlap of two atomic orbitals

95. 95 Summary Sigma (  ) bonds - Circular cross-section and are formed by head-on interaction Pi (  ) bonds – “dumbbell” shape from sideways interaction of p orbitals Carbon uses hybrid orbitals to form bonds in organic molecules. In single bonds with tetrahedral geometry, carbon has four sp 3 hybrid orbitals In double bonds with planar geometry, carbon uses three equivalent sp 2 hybrid orbitals and one unhybridized p orbital Carbon uses two equivalent sp hybrid orbitals to form a triple bond with linear geometry, with two unhybridized p orbitals Atoms such as nitrogen and oxygen hybridize to form strong, oriented bonds The nitrogen atom in ammonia and the oxygen atom in water are sp 3 -hybridized

96. 96 Summary A Brønsted(–Lowry) acid donatea 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