1. Chapter 1

2. ORGANIC CHEMISTRY STUDYOFCARBON CONTAINING COMPOUNDS Compounds from Nature Synthetic compounds: invented by organic chemists and prepared in their laboratories Friedrich Woehler’s urea synthesis Ammonium isocyanate + heat ------> urea NH 4 CNONH 2 CONH 2 “I have been able to make urea without aid of kidney of man or dog”. 1828

3. Some organic chemicals DNA Essential oils Medicines Active Pharmaceutical Ingredients Excipients Materials Fuels Pigments

4. WITH ITSELF No limit AND

5. Electronic Structure of Atoms Structure of atoms – a small dense nucleus, diameter 10 -14 - 10 -15 m, which contains positively charged protons, neutrons and most of the mass of the atom – extranuclear space, diameter 10 -10 m, which contains negatively charged electrons

7. Notice: one s orbital in each principal shell three p orbitals in the second shell (and in higher ones) five d orbitals in the third shell (and in higher ones)

8. Rules for Electron Configurations Capacities of shells (n) and subshells (l)

9. Electronic Structure of Atoms Electrons are confined to regions of space called principle energy levels (shells) – each shell can hold 2n 2 electrons (n = 1, 2, 3, 4......)

10. Electronic Structure of Atoms Shells are divided into subshells called orbitals, which are designated by the letters s, p, d,........ – s (one per shell) – p (set of three per shell 2 and higher) – d (set of five per shell 3 and higher).....

11. Electronic Structure of Atoms Rule 1: Rule 1: orbitals fill from lowest energy to highest energy Rule 2: Rule 2: only two electrons per orbital, spins must be paired Rule 3: Rule 3: for a set of orbitals with the same energy, add one electron in each before a second is added in any one

12. “Periodic” Behavior of Elements Flame tests: elements with low first ionization energies are excited in a flame, and often emit in the visible region of the spectrum Atoms emit energy when electrons fall from higher to lower energy states BaSrCa KNaLi

13. Atomic Spectrum of Hydrogen

14. Electronic Structure of Atoms The pairing of electron spins

16. Lewis Structures Gilbert N. Lewis Valence shell Valence shell: the outermost electron shell of an atom Valence electrons Valence electrons: electrons in the valence shell of an atom; these electrons are used in forming chemical bonds Lewis structure Lewis structure – the symbol of the atom represents the nucleus and all inner shell electrons – dots represent valence electrons For Nitrogen atom: Valence shell of Nitrogen= 3 Number of valence electrons of Nitrogen = 5

17. Lewis Structures Lewis structures for elements 1-18 of the Periodic Table For Nitrogen atom: Valence shell of Nitrogen= 3 Number of valence electrons= 5

18. Lewis Model of Bonding Atoms bond together so that each atom in the bond acquires the electron configuration of the noble gas nearest it in atomic number anion – an atom that gains electrons becomes an anion cation – an atom that loses electrons becomes a cation – Ionic bond – Ionic bond: a chemical bond resulting from the electrostatic attraction of an anion and a cation – Covalent bond – Covalent bond: a chemical bond resulting from two atoms sharing one or more pairs of electrons We classify chemical bonds as ionic, polar covalent, and nonpolar covalent based on the difference in electronegativity between the atoms

19. Electronegativity Electronegativity Electronegativity: a measure of the force of an atom’s attraction for the electrons it shares in a chemical bond with another atom Pauling scale – increases from left to right within a period – increases from bottom to top in a group

20. Electronegativity Electronegativity of atoms (Pauling scale)

22. Electronegativity Electronegativity and chemical bonding Example: NaCl Na = 0.8, Cl = 3.0 Difference is 2.2, so this is an ionic bond!

23. Coulomb’s Law “The energy of interaction between a pair of ions is proportional to the product of their charges, divided by the distance between their centers”

24. What forces that hold atom together within molecules?

25. Covalent Bonding Forces  Electron – electron repulsive forces repulsive forces  Proton – proton repulsive forces repulsive forces  Electron – proton attractive forces attractive forces

26. Bond Length Diagram Net repulsionNet attraction Scientists can determine the internuclear distances that correspond to the lowest energy states of molecules http://ch301.cm.utexas.edu/simulations/bond-strength/BondStrength.swf

27. Bond Length and Energy Bonds between elements become shorter and stronger as multiplicity increases

28. Covalent Bonds A covalent bond forms when electron pairs are shared between two atoms whose difference in electronegativity is 1.9 or less – an example is the formation of a covalent bond between two hydrogen atoms – the shared pair of electrons completes the valence shell of each hydrogen.

29. Polar Covalent Bonds In a polar covalent bond  - – the more electronegative atom has a partial negative charge, indicated by the symbol  -  + – the less electronegative atom has a partial positive charge, indicated by the symbol  + in an electron density model – red indicates a region of high electron density – blue indicates a region of low electron density

30. Polar and Nonpolar Molecules – ammonia and formaldehyde are polar molecules – acetylene is a nonpolar molecule

31. Carbon – Intro and Review Atomic Structure – Atoms – made up of protons, neutrons, electrons – Isotopes – same # protons; different # neutrons Electronic Structure – Electrons determine structure give rise to bonding behave like waves orbitals (s, p)

32. Orbital overlap to form σ bonds.

33. Orbital overlap to form  bonds.

34. Electron Probabilities and the 1s Orbital The 1s orbital looks very much like a fuzzy ball, that is, the orbital has spherical symmetry The electrons are more concentrated near the center Spherical symmetry; probability of finding the electron is the same in each direction. The electron cloud doesn’t “end” here … … the electron just spends very little time farther out.

35. Electron Probabilities and the 2s Orbital The region near the nucleus is separated from the outer region by a spherical node - a spherical shell in which the electron probability is zero The 2s orbital has two regions of high electron probability, both being spherical

36. The Three p Orbitals 2p

37. The Five d Orbitals 3d

38. Rules for Electron Configurations Subshell filling order... Each subshell must be filled before moving to the next level 1s 2 2s 2 2p 6 3s 2 3p 6... 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s

39. The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins (Hund’s rule). C 1s 2 2s 2 2p 2 N 1s 2 2s 2 2p 3 O 1s 2 2s 2 2p 4 F 1s 2 2s 2 2p 5 Ne 1s 2 2s 2 2p 6

40. Periodic Relationships The valence shell is the outermost occupied shell The period number = principal quantum number, n, of the electrons in the valence shell

41. Atomic Orbitals 1s – 1 st orbital – s type (spherical) – 1s, 2s, 3s

42. Atomic Orbitals 2s orbital (spherical)

43. Atomic Orbitals p (2p, 3p…) – 3 orbitals oriented perpendicular to each other – have node (region of 0 e - density) nodal plane 2p orbital

44. Atomic Orbitals p (2p, 3p…) – 3 orbitals oriented perpendicular to each other – have node (region of 0 e - density) nodal plane – shape dumbbell

45. Chapter 1 Electronic Configuration of Atoms Aufbau – Fill lowest energy orbital 1 st Hund’s Rule – 1 e - into each orbital of = energy Pauli Exclusion Principle Pauli Exclusion Principle Electrons in the same orbital are spin paired Electrons in the same orbital are spin paired  

46. Electronic Configurations

47. WHY DO HYBRIDS ?? 1. Electron pair repulsions are minimized (= lower energy) 2. Stronger bonds (= lower energy) are formed 3. Hybrids have better directionality for forming bonds

48. Shapes of Atomic Orbitals All s orbitals have the shape of a sphere, with its center at the nucleus – of the s orbitals, a 1s orbital is the smallest, a 2s orbital is larger, and a 3s orbital is larger still

49. Shapes of Atomic Orbitals – A p orbital consists of two lobes arranged in a straight line with the center at the nucleus

50. Orbital Overlap Model A covalent bond forms when a portion of an atomic orbital of one atom overlaps a portion of an atomic orbital of another atom – in forming the covalent bond in H-H, for example, there is overlap of the 1s orbitals of each hydrogen

51. Hybrid Orbitals We will study three types of hybrid atomic orbitals sp 3 sp 3 (one s orbital + three p orbitals give four sp 3 orbitals) sp 2 sp 2 (one s orbital + two p orbitals give three sp 2 orbitals) sp sp (one s orbital + one p orbital give two sp orbitals) Overlap of hybrid orbitals can form two types of bonds, depending on the geometry of the overlap  bonds  bonds are formed by “direct” overlap  bonds  bonds are formed by “parallel” overlap

52. sp 3 Hybrid Orbitals – Each sp 3 hybrid orbital has two lobes of unequal size – The four sp 3 hybrid orbitals are directed toward the corners of a regular tetrahedron at angles of 109.5°

55. sp 3 Hybrid Orbitals – orbital overlap bonding in water, ammonia, and methane

56. sp 2 Hybrid Orbitals An sp 2 hybrid orbital has two lobes of unequal size – the three sp 2 hybrid orbitals are directed toward the corners of an equilateral triangle at angles of 120° – the unhybridized 2p orbital is perpendicular to the plane of the three sp 2 hybrid orbitals

58. – a carbon-carbon double bond consists of one sigma (  ) bond and one pi (  ) bond sp 2 Hybrid Orbitals

59. – a carbon-oxygen double bond also consists of one sigma (  ) bond and one pi (  ) bond

60. sp Hybrid Orbitals Each sp hybrid orbital has two lobes of unequal size – the two sp hybrid orbitals lie in a line at an angle of 180° – the two unhybridized 2p orbitals are perpendicular to each other and to the line through the two sp hybrid orbitals

62. sp Hybrid Orbitals – a carbon-carbon triple bond consists of one sigma (  ) bond and two pi (  ) bonds

63. Hybrid Orbitals Summary of orbitals and bond types

64. Examples of sigma σ bonds formed from sp 3 hybrid orbitals

65. Orbital overlap to form σ bonds.

66. Orbital overlap to form  bonds.

67. ..

69. Examples of natural acyclic compounds, their sources (in parentheses), and selected characteristics

70. Examples of natural heterocyclic compounds having a variety of heteroatoms and ring sizes.

71. Examples of natural carbocyclic compounds with rings of various sizes and shapes.

72. Isomerism The Molecular Formula of a substance gives the number of different atoms present. The Structural Formula indicates how those atoms are arranged. Isomers are molecules with the same number and kinds of atoms but different arrangements of the atoms. Structural (or Constitutional) isomers have the same molecular formula but different structural formulas.

73. Constitutional Isomerism – the potential for constitutional isomerism is enormous World population is about 6,000,000,000

74. 74 Condensed Structural Formulas

75. 75 Cyclic Molecules

76. 76 Bond-line Formulas

77. 77

78. In this representation, bonds that project upward out of the plane of the paper are indicated by a wedge, those that lie behind the plane are indicated with a dashed wedge, and those bonds that lie in the plane of the page are indicated by a line. 78 Three-Dimensional Formulas

79. writing structural Formulas In a continuous chain, atoms are bonded one after another. In a branched chain, some atoms form branches from the longest continuous chain.

80. Abbreviated Structural Formulas

81. Formal Charge Here, some molecules one or more atoms maybe charged +ve or –ve which comes from the chemical reactions. Its important to know how to tell where the charge is located. H3O+H3O+H3O+H3O+

82. Formal Charge minus The formal charge on an atom in a covalently bonded molecule or ion is the number of valence electrons in the neutral atom minus the number of covalent bonds to the atom and the number of unshared electrons on the atom.

86. Resonance Resonance structures Resonance structures of a molecule or ion are two or more structures with identical arrangements of the atoms but different arrangements of the electrons. resonance hybrid If resonance structures can be written, the true structure of the molecule or ion is a resonance hybrid of the contributing resonance structures.

87. Resonance

88. Physical measurements tell us that none of the foregoing structures accurately describes the real carbonate ion. 1.31 Å 1.20 Å 1.41 Å Experimentally, It was found that all three carbon–oxygen bond lengths are identical: 1.31 Å. This distance is intermediate between the normal C=O (1.20 Å) and C-O (1.41 Å) resonance hybrid The real carbonate ion has a structure that is a resonance hybrid of the three contributing resonance structures