1. PHAR 1123 PHARMACEUTICAL CHEMISTRY II STRUCTURE, BONDING, AND ORGANIC REACTIONS Faculty of Pharmacy Cyberjaya University College of Medical Sciences

2. STRUCTURE AND BONDING

3. Learning Objectives 1.To discuss electronic structure of atoms. 2.To differentiate shells and orbitals. 3.To write the ground state electron configuration of a given element. 4.To name the two theory of chemical bonds. 5.To discuss valence bond theory and hybrid orbitals. 6.To draw skeletal structures. 7.To discuss polar covalent bonds.

4. ORGANIC CHEMISTRY The study of the compounds of CARBON The chemistry of carbon and only a few other elements (H, O, N, S, P and halogen)

5. Video on organic chemistry http://www.youtube.com/watch?v=JgNg5IQnGhM

6. Electronic Structure of Atoms Atoms are composed of 3 principal kinds of subatomic particles: - protons - neutrons - electrons At the center of an atom, is the nucleus, which is a very tiny, extremely dense core. Electrons Nucleus (protons + neutrons)

7. Shells Electrons do not move freely. Confined to regions of space called principle energy levels, or shells. Numbered 1, 2, 3, and so forth from the inside out. Each shell contain 2n 2 electrons, n = number of the shells. Nucleus 1 2 3

8. Orbitals Shells are divided into subshells: s, p, d, and f. Within this subshells, electrons are grouped in orbitals. Orbital: a region of space that can hold 2 electrons.

10. Different p Orbitals The 3 different p orbitals within a given shell are oriented in space along mutually perpendicular directions, (p x, p y, p z ). The two lobes of each p orbital are separated by a region of zero electron density (node). The separated lobes have different algebraic signs, + and –. The different signs of the lobes have important consequences with respect to chemical bonding and chemical reactivity.

12. Electron Configurations Ground-state electron configuration (lowest energy arrangement) can be predicted by following 3 rules. Aufbau principle The lowest-energy orbitals fill up first. Pauli exclusion principle Electrons act as if they are spinning, which have 2 orientations, up and down. The two electrons that occupy an orbital must be of opposite spin. Hund’s rule If two or more empty orbitals of equal energy are available, one electron occupies each with spins parallel until all orbitals are half full.

13. E.g. Ground-state configuration for carbon (6 electrons): 1s 2 2s 2 2p x 1 2p y 1 or 1s 2 2s 2 2p 2

14. Chemical Bonding Theory Atoms bond together because the compound that results is lower in energy, more stable than separate atoms. Energy (heat) always flows out of the chemical system when a chemical bond forms. The bonds are not oriented randomly, they have specific spatial directions. E.g. methane – the 4 hydrogen to which carbon is bonded sit at the corners of a regular tetrahedron, with carbon in the center.

15. Solid lines: bonds in the plane of the page. Heavy wedge line: bond coming out of the page towards the viewer. Dashed line: bond receding back behind the page, away from the viewer. Solid Heavy wedge Dashed

16. Electron Octet Electron octet in the atom’s valence shell impart special stability to the noble gas elements. The chemistry of main-group elements is governed by their tendency to take on the electron configuration of the nearest noble gas. Group 1 and group 17 will each lose and gain 1 electrons, forming ions. The ions are held together in compounds (NaCl) by an electrostatic attraction that is called ionic bond.

17. For elements closer to the middle of the periodic table, e.g. carbon, it would take too much energy to gain or lose 4 electrons to achieve a noble gas configuration. Thus, carbon bonds to other atoms by sharing the electrons. The shared-electron bond is called a covalent bond. The neutral collection of atoms held together by covalent bonds is called a molecule.

18. Lewis and Kekule Structures LewisKekule Electron-dot structuresLine-bond structures Valence electrons are represented as dots Two-electron covalent bond is indicated as a line drawn between atoms

19. The number of covalent bonds an atom forms depends on how many additional valence electrons it needs to reach a noble gas configuration. ElementNoble gasBonds formed Hydrogen (1s)Helium (1s 2 )1 bond Carbon (2s 2 2p 2 )Neon (2s 2 2p 6 )4 bonds Nitrogen (2s 2 2p 3 )Neon3 bonds Oxygen (2s 2 2p 4 )Neon2 bonds Halogens1 bond

20. Valence electrons that are not used for bonding are called lone-pair electrons. E.g. nitrogen atom in ammonia.

21. Chemical Bonds Chemical bonds theory are to explain the forming of bonds between atoms by electron sharing. 2 models: valence bond theory molecular orbital theory Both has their own strengths and weaknesses, thus they are used interchangeably depending on the circumstances. Valence bond theory is the more easily visualized of the two.

22. Valence Bond Theory A covalent bond forms when two atoms approach each other closely and a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom. The electrons are now paired in the overlapping orbitals and are attached to the nuclei of both atoms, thus bonding the atoms together.

23. E.g. Bonding in a hydrogen molecule

24. The overlapping orbitals have the elongated egg shape if two spheres were pressed together. If a plane were to pass through the middle of the bond, the intersection of the plane and the overlapping orbitals would be a circle. The H – H bond is cylindrically symmetrical. Such bonds are called sigma (σ) bonds.

25. During the bond forming reaction, 436 kJ/mol energy is released. The product has less energy than the starting atoms. The product is more stable than the reactant. Bond strength : 436 kJ/mol. There is also an optimum distance between nuclei that leads to maximum stability, called the bond length.

26. Stability of Covalent Bonds An electron pair occupies the region between two nuclei. This arrangement will shield the repulsive forces from one positively charged nucleus to the other nucleus. At the same time, electron pair attracts both nuclei. The internuclear distance is fixed to within very narrow limits. This distance is the bond length and every covalent bond has a definite bond length.

27. sp 3 Hybrid Orbitals How are sp 3 hybrid orbitals formed? An s orbital and 3 p orbitals can combine (hybridized) to form 4 equivalent atomic orbitals with tetrahedral orientation. The tetrahedrally oriented orbitals are called sp 3 hybrids.

29. Why should sp 3 hybrid orbitals formed? The sp 3 hybrid orbitals are unsymmetrical about the nucleus. One of the two lobes is much larger than the other and can therefore overlap more effectively with an orbital from another atom when it forms a bond. Thus sp 3 hybrid orbitals form stronger bonds than do unhybridized s or p orbitals.

30. The asymmetry of sp 3 orbitals arises because of the difference in algebraic signs of each lobes. When a p orbital hybridizes with an s orbital, the positive p lobe adds to the s orbital but the negative p lobe subtracts from the s orbital. E.g. the formation of methane.

31. Each C – H bond has a strength of 436 kJ/mol and a length of 109 pm. The bond angle is 109.5 o.

32. Ethane The same kind of orbital hybridization for methane also accounts for the bonding together of carbon atoms into chains and rings. This makes possible the many millions of organic compounds. Ethane is the simplest molecule containing a carbon-carbon bond.

33. Each C – H bond has a strength of 423 kJ/mol and a length of 109 pm. Each C – C bond has a strength of 376 kJ/mol and a length of 154 pm. The bond angles is ~109.5 o.

34. sp 2 Hybrid Orbitals Another possibility of hybridization. 2s orbital combines with only two of the three available 2p orbitals. Three sp 2 hybrid orbitals result, one 2p orbital remains unchanged. The three sp 2 orbitals lie in a plane at angles of 120 o to one another, the remaining p orbital perpendicular to the sp 2 plane. E.g. ethylene.

36. sp 2 hybrid carbon atom

37. When two sp 2 hybridized carbons approach each other, they form a σ bond by sp 2 -sp 2 head-on overlap. The unhybridized p orbitals approach with the correct geometry for sideways overlap, forming pi (Π) bond. The combination of an sp 2 -sp 2 σ bond and a 2p-2p Π bond results in the sharing of four electrons and the formation of a carbon-carbon double bond.

39. H atoms form s bonds with four sp 2 orbitals. H–C–H and H–C–C bond angles of about 120°. C=C double bond in ethylene is shorter and stronger than the single bond in ethane. Ethylene C=C bond length is 134 pm (C–C 154 pm). The carbon-carbon double bond is less than twice as strong as a single bond. This is due to the overlap in the π part of the double bond is not as effective as the overlap in the σ part.

40. sp Hybrid Orbitals A carbon 2s orbital hybridizes with only a single p orbital. Two sp hybrid orbitals result, two p orbitals remain unchanged. The two sp orbitals are oriented 180 o apart (linear molecule) on the x- axis, while the remaining two p orbitals are perpendicular on the y-axis and the z-axis. E.g. acetylene.

41. Two sp hybrid orbitals from each C form sp–sp s bond. p z orbitals from each C form a p z –p z  bond by sideways overlap and p y orbitals overlap similarly.

43. Hybridization video http://www.youtube.com/watch?v=SJdllffWUqg&feature=related

44. Molecular Orbital Theory Covalent bond formation arises from a combination of atomic orbitals on different atoms to form molecular orbitals. Molecular orbital describes a region of space in a molecule where electrons are most likely to be found.

45. Drawing Chemical Structures 2-Methylbutane Kekule structure: Condensed structure: CH 3 CH 2 CH(CH 3 ) 2 Skeletal structure:

46. Rules for drawing skeletal structures: 1.Carbon atoms are not usually shown. Instead, a carbon atom is assumed to be at each intersection of two lines (bonds) and at the end of each line. Occasionally, a carbon atom might be indicated for emphasis or clarity. 2.Hydrogen atoms bonded to carbon are not shown. Since carbon always has a valence of 4, we mentally supply the correct number of hydrogen atoms for each carbon. 3.Atoms other than carbon and hydrogen are shown.

47. Electronegativity The two known chemical bonds: ionic and covalent. However, most bonds are neither fully ionic nor fully covalent but are somewhere between the two extremes. Called polar covalent bonds, which means the bonding electrons are attracted more strongly by one atom than the other so that the electron distribution between atoms is not symmetrical.

48. The symbol δ means partial charge, either partial positive for the electron-poor atom or partial negative for the electron rich atom.

49. Bond polarity is due to differences in electronegativity (EN). EN: the intrinsic ability of an atom to attract the shared electrons in a covalent bond. Fluorine is the most electronegative element (EN = 4.0) and cesium is the least electronegative element (EN = 0.7). Inductive effect: the shifting of electrons in a σ bond in response to the electronegativity of nearby atoms.

50. Dipole Moments Molecules as a whole are also often polar. The polarity caused by: - the net sum of individual bond polarities - lone-pair The measure of net molecular polarity is called the dipole moment, , expressed in debyes (D). 1 D = 3.336 x 10 -30 coulomb meters (C. m)  = Q x r, where Q = magnitude of the charge at either end of the molecular dipoles r = distance between the charges

51. Formal Charges

52. Resonance The true structure is intermediate between the two, they are called resonance forms. The only difference between resonance forms: - the placement of the Π bond. - nonbonding valence electrons.

53. ORGANIC REACTIONS

54. Learning Objectives 1.To classify organic reactions into different kinds and mechanism. 2.To identify nucleophile and electrophile. 3.To discuss electrophilic addition reaction.

55. Chemical Reactions All reactions (in lab or living organisms), follow the same “rules”. Reactions in living organisms look more complex, with the involvement of enzymes. The principles governing all reactions are the same. E.g. biosynthesis of prostaglandin H 2.

57. Organic Reactions ORGANIC REACTIONS What kind?How? - Addition - Elimination - Substitution - Rearrangements Mechanism: - Radical - Polar

58. Addition Reaction Two reactants add together to form a single product with no atoms left-over.

59. Elimination Reaction A single reactant split into two products, often with formation of a small molecule, e.g. water. The opposite of addition reaction.

60. Substitution Reaction Two reactants exchange parts to give two new products.

61. Rearrangement Reaction A single reactant undergoes a reorganization of bonds and atoms to yield an isomeric (compounds that have the same molecular formula but different structure) product.

62. Mechanisms Reaction mechanism: an overall description of how a reaction occurs. A complete mechanism include: - what takes place at each stage of a chemical transformation. - the rate of each steps. - all reactants used and products formed.

63. Bond Breaking Arrowheads with a “half” head (“fish-hook”) indicate homolytic and homogenic steps (called ‘radical processes’). Arrowheads with a complete head indicate heterolytic and heterogenic steps (called ‘polar processes’). Symmetrical bond-breaking Unsymmetrical bond-breaking

64. Bond Making Symmetrical Unsymmetrical

65. Radical Reactions Not as common as polar reactions. Important in some industrial processes and in numerous biological pathways. A radical is highly reactive – contains an atom with an odd number of electrons (usually seven) in its valence shell. A radical can achieve a valence-shell octet through several ways: - radical substitution reaction - radical addition reaction

67. In industry, radical substitution reaction is used for the chlorination of methane. The substitution reaction is the first step in the preparation of dichloromethane and chloroform. 3 types of steps: initiation, propagation, and termination.

68. Initiation Homolytic formation of a few reactive chlorine radicals by irradiation of a small number of chlorine molecules with ultraviolet light.

69. Propagation Reaction with molecule to generate radical. The overall process is called a chain reaction. A radical will collide with a methane molecule and abstract a hydrogen atom. Cycles back and repeats the first propagation step.

70. Termination Combination of two radicals to form a stable product. Occur infrequently.

71. Polar Reactions Polar reactions occur because of the electrical attraction between positive and negative centers on functional groups in molecules. Bond polarity is the result of an unsymmetrical electron distribution in a bond. Due to the difference in electronegativity of the bonded atoms.

72. * Carbon is always positively polarized except when bonded to a metal

73. Polar bonds can also result from the interaction of functional groups with acids or bases. E.g. methanol. In neutral methanol, the C atom is electron-poor due to the electronegative O that attracts the electrons in the C – O bond. On protonation of the methanol oxygen by an acid, a full positive charge on oxygen attracts the electrons in the C – O bond much more strongly. This makes the C much more electron-poor.

74. The fundamental characteristic of all polar organic reactions is that electron-rich sites react with electron-poor sites. Bonds are made when an electron-rich atom shares a pair of electrons with an electron-poor atom, and bonds are broken when one atom leaves with both electrons from the former bond.

75. Has a negatively polarized, electron-rich atom. Can form a bond by donating a pair of electrons to a positively polarized, electron-poor atom. Neutral or negatively charged. E.g. ammonia, water, hydroxide ion, chloride ion. Has a positively polarized, electron poor atom. Can form a bond by accepting a pair of electrons from a nucleophile. Neutral or positively charged. E.g. acids, alkyl halides, carbonyl compounds. NucleophileElectrophile

76. Some nucleophiles and electrophiles. Electrostatic potential maps identify the nucleophilic (red) and electrophilic (blue) atoms.

77. Polar Reaction: Addition of HBr to Ethylene A typical polar reaction – addition reaction of an alkene with hydrogen bromide.

78. The reaction is an example of electrophilic addition reaction. The reaction begins when the alkene donates a pair of electrons from its C=C bond to HBr to for a new C – H bond plus Br -. One curved arrow begins at the middle of the double bond (source of electron pair) and points to the H atom in HBr (the atom to which the bond will form). A second curved arrow begins in the middle of the H-Br bond and points to the Br, indicating that the H-Br bond breaks. The electrons remain with the Br atom, giving Br -.

79. When one of the alkene carbon atoms bonds to the incoming hydrogen, the other carbon atom (having lost its share of the double- bond electrons) is left with six valence electrons, thus it is positively charged. The positively charged species is called a carbocation, is an electrophile that can accept an electron pair from nucleophilic Br - anion. The curved arrow shows the electron pair movement from Br - to the positively charged carbon. This will form a C-Br bond and yield the addition product.

82. Using Curved Arrows Curved arrows are a way to keep track of changes in bonding in polar reaction. The arrows track electron movement. Electrons always move in pairs. Charges change during the reaction. One curved arrow corresponds to one step in a reaction mechanism.

83. Rules for Using Curved Arrows: 1.The arrow goes from the nucleophilic source (Nu: or Nu: - ) to the electrophilic sink (E or E + ). 2. The nucleophilic site can be neutral or negatively charged. 3. The electrophilic site can be neutral or positively charged. 4. Octet rule must be followed.

84. QUESTIONS?? THANK YOU