Physical Organic Chemistry

1. Islamic University in Madinah Department of Chemistry Physical Organic Chemistry Prepared By Dr. Khalid Ahmad Shadid

2. Topics: • What is Physical Organic Chemistry? • What is Mechanism? • Structure and Bonding • Quantum Mechanics & Atomic Structure • Molecular dimensions • Resonance Theory • Electronegativity & Bond Polarity • Intermolecular Forces • Lewis Structures of Common Functional Groups • Carbocations • The Carbanion Chemistry

3. Introduction • What is Physical Organic Chemistry? - Physical organic chemistry is the “ Study of the interrelationships between structure and reactivity of the organic molecules”. - Understanding mechanisms, explaining observed phenomenon (or predicting it) • What is Mechanism? - Lowry and Richardson: “ Theory deduced from available experimental data. The experimental results are the facts; the mechanism is conjecture based upon these facts”. - A mechanism the simplest rationalization that is consistent with all of the available data.

4. Scientific method distinguishes a good scientific • The hypothesis (assumption)• Doing experiments: – confirmation – refutation •“Proof” of a mechanism- a mechanism can never really be ‘proven’- one can propose a hypothetical mechanism- one can conduct experiments designed to refute certain hypotheses- one can retain mechanisms that are not refuted- a mechanism can thus become “generally accepted” • Mechanistic studies• allow a better comprehension of a reaction, its scope and its utility• require kinetic experiments– rate of disappearance of reactants– rate of appearance of products• kinetic studies are therefore one of the most important disciplinesin experimental chemistry• allow a better comprehension of a reaction, its scope and its utility

5. Why Chemical Reactions Happen? • Understanding Chemical reactions lies in: • The structure and bonding of stable compounds. • Reactivity of compounds. • Reactive intermediates and transition states. • Mechanisms through which compounds transform into other compounds. • Kinetics, thermodynamics and the relationship between the two

6. Review of Concepts in Organic Chemistry Quantum Mechanics & Atomic Structure • Wave mechanics & quantum mechanics • Each wave function (y) corresponds to a different energy state for an electron • Each energy state is a sublevel where one or two electrons can reside • Wave functions are tools for calculating two important properties • The energy associated with the state of the electron can be calculated • The relative probability of an electron residing at particular places in the sublevel can be determined • The phase sign y of a wave equation indicates whether the solution is positive or negative when calculated for a given point in space relative to the nucleus • Wave functions y, whether they are for sound waves, lake waves, or the energy of an electron, have the possibility of constructive interference and destructive interference • Constructive interference occurs when wave functions with the same phase sign interact. There is a reinforcing effect and the amplitude of the wave function increases • Destructive interference occurs when wave functions with opposite phase signs interact. There is a subtractive effect and the amplitude of the wave function goes to zero or changes sign

8. Review of Concepts in Organic Chemistry Electrons in simple molecules • The Energy Carve of Diatomic • In an atom, we know the negatively charged electrons are attracted to positively charged nuclei. As two hydrogen atoms approach each other from an infinite distance, the electrons on one atom become attracted to the nucleus of the other atom. This mutual attraction increases as the distance between atoms decreases. As the atoms get closer and closer, the positively charged nuclei and negatively charged clouds began to repel each other. • http://ch301.cm.utexas.edu/simulations/bond-strength/BondStrength.swf

9. Review of Concepts in Organic Chemistry Electrons in simple molecules • There is an optimum distance between two bonded atoms at which the energy is at a minimum. This distance is called the bond length = 74 pm (picometers (pm) which is equal to 10-12 meters ) for the hydrogen molecule H2 or (H-H). • The energy required to break the bond is called the bond strength or (bond dissociation energy). This energy value is 435 kJ/mol (104 kcal/mol) for a hydrogen-hydrogen bond.

10. Review of Concepts in Organic Chemistry Electrons in simple molecules • The two separate hydrogen atoms (H) are higher in energy (less stable) than is the hydrogen molecule (H2). • Atoms and molecules like to be in the lowest energy (most stable) state. Thus when a bond is formed, energy is given off and the resulting molecule has a lower potential energy (is more stable) than the starting individual species has. So a hydrogen molecule H2 is more stable than two separate hydrogen atoms. Which mean that: energy must be added to the hydrogen molecule to break the bond, giving the original separate hydrogen atoms”.

11. Molecular Orbitals • We cannot simultaneously know the position and momentum of an electron • An atomic orbital represents the region of space where one or two electrons of an isolated atom are likely to be found • A molecular orbital (MO) represents the region of space where one or two electrons of a molecule are likely to be found • An orbital (atomic or molecular) can contain a maximum of two spin-paired electrons (Pauli exclusion principle) • When atomic orbitals combine to form molecular orbitals, the number of molecular orbitals that result always equals the number of atomic orbitals that combine An antibonding molecular orbital (y*molec) results when two orbitals of opposite phase overlap A bonding molecular orbital (ymolec) results when two orbitals of the same phase overlap

13. Review of Concepts in Organic Chemistry Electrons in simple molecules • Molecular Orbital for H2 The LCAO-MO Model simply states That: The shape of a molecular orbital is derived from the shape of the atomic orbitals that overlap to form that molecular orbital. when two waves are traveling in opposite directions, and they meet, as in the wakes of two boats, their amplitudes cancel each other. During bonding, atoms do the atomic equivalent—wave functions with the same sign overlap in an in-phase overlap, and wave functions of opposite signs overlap in an out-of-phase overlap”. In-phase overlap is a constructive, or bonding, overlap of atomic orbitals. Out-of-phase is a destructive, or antibonding, overlap of atomic orbitals.

14. To differentiate the antibonding from the bonding orbital, add an asterisk to the σ, giving σ* (sigma star). Bonding molecular orbitals: lower in energy than the atomic orbitals of which it is composed. Electrons in this type of orbital increase the stability of molecule and favor the molecule bonding.

15. Prediction from the MO diagram What if two atoms of Helium combine? Does He2 form? • Antibonding molecular orbitals: higher in energy than the atomic orbitals of which it is composed. Electrons in this type of orbital decrease the stability and will favor the separated atoms. (Unstable but can exist!)

17. Heteronuclear Diatomic Molecules In the Molecular Orbital diagram for NO: The 2s and 2p atomic orbitals of oxygen are slightly lower than those of nitrogen because oxygen is more electronegative than nitrogen.

18. A sigma (s) bond (a type of single bond) is one in which the electron density has circular symmetry when viewed along the bond axis … an antibonding molecular orbital, higher in energy than the AOs. Electron density between the nuclei is decreased. … a bonding molecular orbital, lower in energy than the AOs, and … Two AOs in hydrogen atoms combine to form … Electron density between the nuclei is increased.

19. A pi (p) bond, part of double and triple carbon–carbon bonds, is one in which the electron densities of two adjacent parallel p orbitals overlap sideways to form a bonding pi molecular orbital The two pz orbitals and also the two py orbitals give pi bonding and antibonding MOs. The two px orbitals combine to form sigma bonding and antibonding MOs. What about combination of more than Two p orbital…? conjugated system

20. Electrons in 2s orbitals are next lowest in energy • Electrons of the three 2p orbitals have equal but higher energy than the 2s orbital • Electrons in 1s orbitals have the lowest energy because they are closest to the positive nucleus • Orbitals of equal energy (such as the three 2p orbitals) are called degenerate orbitals

21. Allyl three π molecular Orbital HOMO Allyl Cataion 2e + Allyl Radical 3e Allyl Anion 4e - Allyl systems result from the combination of 3 conjugated p-orbitals.Therefore, this will result in 3 π molecular orbitals

22. π – Molecular orbitals for 1,3-butadiene • π – molecular orbitals for 1,3-butadiene?CH2=CH—CH=CH2How many AO’s in the π system? fourHow many MO’s result? fourHow many electrons in the π system? four

24. The Structure of Methane and Ethane: sp3Hybridization sp3 hybridized carbon covalent bond

25. The Structure of Methane • sp3 orbital • 25% s character, • 75% p character

26. The Structure of Ethane

27. The Structure of Ethene (Ethylene): sp2 Hybridization

28. An sp2-hybridized carbon atom • sp2 orbital • 33% s character, 66% p character A model for the bonding molecular orbitals of ethene formed from two sp2-hybridized carbon atoms and four hydrogen atoms

29. Restricted Rotation and the Double Bond • There is a large energy barrier to rotation associated with groups joined by a double bond • ~264 kJ/mol(strength of the p bond) • To compare: rotation of groups joined by C-C single bonds ~13-26 kJ/mol (cis) (trans) • Restricted rotation of C=C

30. The Structure of Ethyne (Acetylene): sp Hybridization • sp orbital • 50% s character • 50% p character

31. Bond Lengths of Ethyne, Ethene, & Ethane • The carbon–carbon triple bond of ethyne is shorter than the carbon–carbon double bond of ethene, which in turn is shorter than the carbon–carbon single bond of ethane • Reasons: • The greater the s orbital character in one or both atoms, the shorter is the bond. This is because s orbitals are spherical and have more electron density closer to the nucleus than do p orbitals • The greater the p orbital character in one or both atoms, the longer is the bond. This is because p orbitals are lobe-shaped with electron density extending away from the nucleus

32. How to Predict Molecular Geometry: The Valence Shell Electron Pair Repulsion Model • Valence shell electron pair repulsion (VSEPR) model: • We consider molecules (or ions) in which the central atom is covalently bonded to two or more atoms or groups • We consider all of the valence electron pairs of the central atom—both those that are shared in covalent bonds, called bonding pairs, and those that are unshared, called nonbonding pairs or unshared pairs or lone pairs • Because electron pairs repel each other, the electron pairs of the valence shell tend to stay as far apart as possible. The repulsion between nonbonding pairs is generally greater than that between bonding pairs • We arrive at the geometry of the molecule by considering all of the electron pairs, bonding and nonbonding, but we describe the shape of the molecule or ion by referring to the positions of the nuclei (or atoms) and not by the positions of the electron pairs

33. Ammonia Water Angular Trigonal pyramidal A tetrahedral arrangement of the electron pairs explains the trigonal pyramidal arrangement of the four atoms. The bond angles are 107° (not 109.5°) because the nonbonding pair occupies more space than the bonding pairs Methane Tetrahedral

34. Boron Trifluoride Beryllium Hydride Linear Trigonal planar Carbon Dioxide

35. Review of Concepts in Organic Chemistry Molecular dimensions • Atomic radiusionic radius, ri： size of electron cloud around an ion.covalent radius, rc： half of the distance between two atomsof same element bond to each other.van der Waal radius, rvdw： the effective size of atomic cloud around a covalently bonded atoms.

36. Review of Concepts in Organic Chemistry • Bond length measures the distance between nucleus (or the local centers of electron density). • Bond angle measures the angle between lines connecting different nucleus.

37. Resonance Theory： • An extension of valence bond theory for molecules that more than one Lewis structure can be written. Useful in describing electron delocalization, in conjugate system andreactive intermediates.(a)If more than one Lewis structure can be written, which has nuclear positions constant, but differ in assignment of electrons, the molecule is described by a combination ofthese structure (a hybrid of all).(b)The most favorable (lowest energy) resonance structure makes the greatest contribution to the true structure.determining energy： maximum number of covalent bond, minimum separation of unlike charge, placement of negative charge on most electronegative atom (vice versa).(c)Those with delocalized electrons are usually more stable than single localized structure.

38. Resonance Theory： • two equivalent structure charge located equally on two C’s the allyl cation is planar for maximum p interaction restricted rotation around single bond

39. Electronegativity & Bond Polarity • Electronegativity： The power of an atom in a molecule to attract electrons to itself. Pauline 1932 • In complex molecules with many polar bonds involved, electrostatic potential surfaces (from quantum mechanic calculation) are used to view the charge distribution in the whole molecule.

40. red ----negative potentialblue --- positive potentialgreen---neutral

41. Electronegativity & Bond Polarity • when there is a difference in the electronegativities of the bonded atoms the electron pairs will spend more time with one atom than the other, they are not equally shared. This also called the Inductive Effect. The "arrow" indicates the direction of the electron shift. This is a polar bond. • As the more electronegative atom exerts a greater pulling power on the shared electrons and so gains more ‘possession’ of the electron pair • The bond is now unsymmetrical with respect to electron distribution and is said to be polar.

42. The term dipole is often used to indicate presence of two opposite electric charges on bond separated by a distance. • The more electronegative atom, with the greater share of the electrons, has become partially negative or δ– • and the less electronegative atom has become partially positive or δ+

43. Intermolecular Forces • Covalent bonds - Typically the strongest type of bonds - 150-1100 kJ/mol • Intermolecular forces are forces that hold covalent molecules together • electrostatic interactions-dispersion forces-dipole-dipole interactions-hydrogen bonds (H-bonds) • Strength of IM forces depends on:Q - charge on ion µ - dipole momentα - polarizability

45. polarizability ( α): ease of electronic distortion; α increases as number of e– ↑; α increases as size (MW) ↑; highly polarizable molecules: more subject to dispersion forces • London dispersion forces: dipole moment forms when there are more electrons on one side of the molecule; when this happens, a dipole in a neighboring molecule is induced.

48. pi-stacking ‣ Benzene dimer is bound by ~2-3 kcal/mol Edge to Face Face to Face Parallel Displace The interactions can be exploited when designing drugs, which are generally inhibitors of proteins.This is called structure-based drug design. Here is an example of arene/perfluoroarene stacking in an inhibitor of carbonic anhydrase II (PDB code: 1G54):

49. Lewis Structures of Common Functional Groups

50. Carbocations (Carbonium ion) • The carbonium ions are positively charged highly reactive intermediates and have six electrons in the outer shell of central carbon. They may be represented as: • they are very reactive toward any nucleophilic reagent