George Mason University General Chemistry 212 Chapter 15 Organic Chemistry Acknowledgements

1. George Mason University General Chemistry 212 Chapter 15 Organic Chemistry Acknowledgements Course Text: Chemistry: the Molecular Nature of Matter and Change, 6th ed, 2011, Martin S. Silberberg, McGraw-Hill The Chemistry 211/212 General Chemistry courses taught at George Mason are intended for those students enrolled in a science /engineering oriented curricula, with particular emphasis on chemistry, biochemistry, and biology The material on these slides is taken primarily from the course text but the instructor has modified, condensed, or otherwise reorganized selected material.Additional material from other sources may also be included. Interpretation of course material to clarify concepts and solutions to problems is the sole responsibility of this instructor.

2. Organic Chemistry • Life on earth is based on a vast variety of reactions and compounds based on the chemistry of Carbon – Organic Chemistry • Organic compounds contain Carbon atoms, nearly always bonded to other Carbon atoms, Hydrogen, Nitrogen, Oxygen, Halides and selected others (S, P) • Carbonates, Cyanides, Carbides, and other carbon-containing ionic compounds are NOT organic compounds • Carbon, a group 4A compound, exhibits the unique property of forming bonds with itself (catenation) and selected other elements to form an extremely large number of compounds – about 9 million • Most organic molecules have much more complex structures than most inorganic molecules

3. Organic Chemistry • Bond Properties, Catenation, Molecular Shape • The diversity of organic compounds is based on the ability of Carbon atoms to bond to each other (catenation) to form straight chains, branched chains, and cyclic structures – aliphatic, aromatic • Carbon is in group 4 of the Periodic Chart and has 4 valence electrons – 2s22p2 • This configuration would suggest that compounds of Carbon would have two types of bonding orbitals each with a different energy • If fact, all four Carbon bonds are of equal energy • This equalization of energy arises from the hybridization of the 2s & 2p orbitals resulting in 4 sp3 hybrid orbitals of equal energy

4. Organic Chemistry • Hybrid orbitals are orbitals used to describe bonding that is obtained by taking combinations of atomic orbitals of an isolated atom • In the case of Carbon, one “s” orbital and three “p” orbitals, are combined to form 4 sp3 hybrid orbitals • The Carbon atom in a typical sp3 hybrid structure has 4 bonded pairs and zero unshared electrons, therefore, Tetrahedral structure AXaEb (a + b) 4 + 0 = AX4 • The four sp3 hybrid orbitals take the shape of a Tetrahedron

5. Organic Chemistry 2p sp3 sp3 C-H bonds 2s Energy 1s 1s 1s C atom (ground state) C atom (hybridized state) C atom (in CH4)

6. Organic Chemistry Shape of sp3 hybrid orbital different than either s or p

7. Organic Chemistry • The bonds formed by these 4 sp3 hybridized orbitals are short and strong • The C-C bond is short enough to allow side-to-side overlap of half-filled, unhybridized p orbitals and the formation of “multiple” bonds • Multiple bonds restrict rotation of attached groups • The properties of Organic molecules allow for many possible molecular shapes

8. Organic Chemistry • Electron Configuration, Electronegativity, and Covalent Bonding • Carbon ground-state configuration – [He 2s22p2] • Hybridized configuration – 4 sp3 • Forming a C4+ or C4- ion is energetically very difficult (impossible?): • Required energy • Ionization Energy for C4+ - IE1<IE2<IE3<IE4 • Electron Affinity for C4- - EA1<EA2<EA3<EA4 • Electronegativity is midway between metallic and most nonmetallic elements • Carbon, thus, shares electrons to bond covalently in all its elemental forms

9. Organic Chemistry • Molecular Stability • Silicon and a few other elements also catenate, but the unique properties of Carbon make chains of carbon very stable • Atomic Size and Bond strength • Bond strength decreases as atom size and bond length increase, thus, C-C bond strength is the highest in group 4A • Relative Heats of Reaction • Energy difference between a C-C Bond(346 kJ/mol) vs C-O Bond (358 kJ/mol) is small • Si-Si (226 kJ/mo) vs Si-O (368 kJ/mol) difference represents heat lost in bond formation • Thus, Carbon bonds are more stable than Silicon

10. Organic Chemistry • Orbitals available for Reaction • Unlike Carbon, Silicon has low-energy “d” orbitals that can be attacked by lone pairs of incoming reactants • Thus, Ethane (CH3-CH3) with its sp3 hybridized orbitals is very stable, does not react with air unless considerable energy (a spark) is applied • Whereas, Disilane (SiH3 – SiH3) breaks down in water and ignites spontaneously in air

11. Organic Chemistry • Chemical Diversity of Organic Molecules • Bonding to Heteroatoms (N, O, X, S, P) • Electron Density and Reactivity • Most reactions start (a new bond forms) when a region of high electron density on one molecule meets a region of low electron density of another • C-C bond: “Nonreactive” – The electronegativities of most C-C bonds in a molecule are equal and the bonds are nonpolar • C-H bond: “Nonreactive” – the bond is nonpolar and the electronegativities of both H(2.1) & C(2.5) are close • C-O bond: “Reactive” – polar bond • Bonds to other Heteroatoms: Bonds are long & weak, and thus, reactive

12. Carbon Geometry The combination of single, double, and triple bonds in an organic molecule will determine the molecular geometry sp3 sp2 sp sp Tetrahedral trigonal planar linear linear AX4 AX3 AX2 AX2 Review Chapter 11 – Multiple bonding in carbon compounds

13. Hydrocarbons • Compounds containing only C and H • Saturated Hydrocarbons: Alkanes only single () bonds • Unsaturated Hydrocarbons: • AlkenesAlkynes Double (=) Bonds Triple () bonds • Aromatic Hydrocarbons (Benzene rings) (6-C ring with alternating double and single bonds)

14. Hydrocarbons • A close relationship exists among Bond Order, Bond Length, and Bond Energy • Two nuclei are more strongly attracted to two shared electrons pairs than to one: The atoms are drawn closer together and are more difficult to pull part • For a given pair of atoms, a higher bond order results in a shorter bond length and a higher bond energy, i.e., A shorter bond is a stronger bond

15. Alkanes (Aliphatic Hydrocarbons) Normal-chain: linear series of C atoms C-C-C-C-C-C- Branched-chain: branching nodes for C atoms Cycloalkanes: C atoms arranged in rings Hydrocarbons Methyl Propane Cyclohexane

16. Hydrocarbons • Alkanes: CnH2n+2 • Straight Chained Alkanes H H H H H C H H C C C H H H H H Propane Methane H H H H H H H C C C H C H C C H H H H H H H Ethane Butane

17. Hydrocarbons • Branched Chained Alkanes • Cycloalkanes 3-Ethyl-4-MethylHexane Cyclobutane Methylcyclopropane

18. Hydrocarbons • Molecular Formulas of n-Alkanes • Methane: C-1: CH4 • Ethane: C-2: CH3CH3 • Propane: C-3: CH3CH2CH3 • Butane: C-4: CH3CH2CH2CH3 • Pentane: C-5: CH3CH2CH2CH2CH3 • Hexane:C-6: CH3(CH2)4CH3 • Heptane: C-7: CH3(CH2)5CH3 • Octane: C-8: CH3(CH2)6CH3 • Nonane:C-9: CH3(CH2)7CH3 • Decane: C-10: CH3(CH2)8CH3

19. Hydrocarbons Straight Chain (n) Alkanes

20. Hydrocarbons Petroleum Fractions

21. Hydrocarbons Cycloalkanes: CnH2n H H H H C H C C H H C C H C H H H H H H Cyclohexane H H H C C H C H H H C C H C C H H H H Cyclopropane Cyclobutane

22. Hydrocarbons • Structural Isomers • Structural (or constitutional) isomers are compounds with the same molecular formula, but different structural formulas. Created by branching, etc. C H H H 3 H C C C H H C C C C H 3 3 3 3 H H H Butane Isobutane C4H10 C4H10

23. Hydrocarbons • Structural Isomers of Pentane C5H12

24. Hydrocarbons • Chiral Molecules & Optical Isomerism • Another type of isomerism exhibited by some alkanes and many other organic compounds is called Stereoisomerism • Sterioisomers are molecules with the same arrangement of atoms but different orientations of groups in space • Optical Isomerism is a type of stereoisomerism, where two objects are mirror images of each other and cannot be superimposed (also called enantiomers) • Optical isomers are not superimposable because each is asymmetric: there is no plane of symmetry that divides an object into two identical parts

25. Hydrocarbons • Chiral Molecules & Optical Isomerism • An asymmetric molecule is called “Chiral” • The Carbon atom in an optically active asymmetric (l) organic molecule (the Chiral atom) is bonded to four (4) different groups. • Mirror images • 1C1 & 1C2 of molecule 1 (left) can be moved to the right to sit on top of2C1 & 2C2 of molecule 2, i.e., • 1C & 2C groups can be superimposed • But, the two groups on C3 are opposite •  The two forms are optical isomers and cannot be superimposed, i.e., no plane of symmetry to divide molecule into equal parts • C-3 is the “Chiral” Carbon Optical Isomers of3-methylhexane

26. Hydrocarbons • Optical Isomers • In their physical properties, Optical Isomers differ only in the direction each isomer rotates the plane of polarized light • One of the isomers – dextrorotary isomer - rotates the plane in a clockwise direction (d or +) • The other isomer – levorotary isomer - rotates the plane in a counterclockwise direction (l or -) • An equimolar mixture of the dextrorotary (d or +) and levorotary (l or -) isomers: recemic mixture does not rotate the plane of light because the dextrorotation cancels the levoratation

27. Hydrocarbons • Optical Isomers • In their chemical properties, optical isomers differ only in a chiral (asymmetric) chemical environment • An optically active isomer is distinguished by the chiral atom being attached to 4 distinct groups If the attached groups are not distinct the molecule is NOT optically active • An isomer of an optically active reactant added to a mixture of optically active isomers of an another compound will produce products of different properties – solubility, melting point, etc.

28. Nomenclature of Alkanes • Determine the longest continuous chain of carbon atoms. The base name is that of this straight-chain alkane. • Any chain branching off the longest chain is named as an “alkyl” group, changing the suffix –ane to –yl • For multiple alkyl groups of the same type, indicate the number with the prefix di, tri, … Ex. Dimethyl, Tripropyl, Tertbutyl • The location of the branch is indicated with the number of the carbon to which is attached Note: The numbering of the longest chain begins with the end carbon closest to the carbon with the first substituted chain or functional group

29. Nomenclature Example CH3 HC CH2 CH3 H2C CH CH3 CH2 CH3 HC CH3 (Con’t)

30. Nomenclature Example • Determine the longest chain in the molecule • 7 Carbons CH3 HC CH2 CH3 H2C CH CH3 CH2 CH3 HC CH3 Substituted Heptane (7 C) (Con’t)

31. Nomenclature Example • The base chain is 7 carbons – Heptane • Add the name of each chain substituted on the base chain “methyl” groups at Carbon 3 and Carbon 5 “ethyl” group at Carbon 4 CH3 HC CH2 CH3 H2C CH CH3 3 2 1 CH2 4 CH3 HC 5 6 7 CH3 3,5-dimethyl-4-ethylheptane

32. Nomenclature Example • Guidelines for numbering substituted carbon chains • The numbering scheme used in developing the name of a organic compound begins with the end carbon closest to the carbon with the first substituted group or functional group

33. Hydrocarbons • Reactions of Alkanes • Combustion (reaction with oxygen) – Burning C5H12(g) + 8 O2(g)  5 CO2(g) + 6 H2O(l) • Substitution (for a Hydrogen) C5H12(g) + Cl2(g)  C5H11Cl(g) + HCl(g)

34. Hydrocarbons • Alkenes • When a Carbon atom forms a double bond with another Carbon atom, it is now bonded to 2 other atoms instead of 3 as in an Alkane • The Geometry now changes from 4 sp3 orbitals (Tetrahedral AX4E0) to 3 sp2 hybrid orbitals and 1 unhybridized 2p orbital (AX3E0 Trigonal Planar) lying perpendicular to the plane of the trigonal sp2 hybrid orbitals Review Chapter 10 - Geometry

35. Hydrocarbons • Alkenes • Two sp2 orbitals of each carbon form C – H sigma () bonds by overlapping the 1 s orbitals of the two H atoms • The 3rd sp2 orbital forms a C-C () bond with the other Carbon • A Pi () bond forms when the two unhybridized 2p orbitals (one from each carbon) overlap side-to-side, one above and one below the C-C sigma bond • A double bond always consists of 1  and 1  bond

36. Hydrocarbons • Alkenes: CnH2n • Alkenes substitute the single sigma bond () with a double bond – a combination of a sigma bond and a Pi () bond • The double-bonded (-C=C-) atoms are sp2 hybridized • The carbons in an Alkene structure are bonded to fewer than the maximum 4 atoms • Alkenes are considered: unsaturated hydrocarbons H H H H C C C C H H H CH3 Ethene or Ethylene Propene

37. Hydrocarbons • Molecular Formulas of Alkenes • Ethene: CH2=CH2 • Propene: CH2=CHCH3 • Butene: CH2=CHCH2CH3 • Pentene: CH2=CHCH2CH2CH3 • Decene: CH2=CH(CH2)7CH3 • Conjugated Molecules Alkene (or aromatic) with alternating Sigma bonds and Pi bonds) Ex. 2,5-Dimethyl-2,4-Hexadiene CH3CH3=CH-CH=C(CH3CH3)

38. Hydrocarbons • Reactions of Alkenes • Addition Reactions CH3CH=CH2 + HBr  CH3CHBrCH(H2) • Why does the Bromine (Br) attach to the middle carbon? Markownikov’s Rule: When a double bond is broken, the H atom being added adds to the carbon that already has the most Hydrogens CH2→ CH3

39. Hydrocarbons • An addition reaction occurs when an unsaturated reactant (alkene, alkyne) becomes saturated( bonds are eliminated) • Carbon atoms are bonded to more atoms in the “Product” than in the reactant (Ethene is reduced) • Addition Reaction – Heat of Formation Reaction is Exothermic Formation of two strong  bonds from a single  bond and a relatively weak  bond Reactants (bonds brokenProduct (bonds formed) 1 C = C = 614 kJ 1 C – C = – 347 kJ 4 C – H = 1652 kJ 5 C – H = – 2065 kJ 1 H – C = 427 kJ1 C – Cl = – 339 kJ Total = 2693 kJ Total = – 2751 kJ

40. Hydrocarbons • Elimination Reactions • The reverse of “addition reaction”: A saturated molecule becomes “unsaturated” • Typical groups “Eliminated” include: • Pairs of Halogens – Cl2, Br2, I2 • H atom and Halogen – HCL, HBr • H atom and Hydroxyl (OH) – • Driving force – Formation of a small, stable molecule, such as HCl, H2O, which increases ntropy of the system

41. Hydrocarbons • Substitution Reactions • A substitution reaction occurs when an atom (or group) from an added reagent substitutes for an atom or group already attached to a carbon • Carbon atom is still bonded to the same number of atoms in the product as in the reactant • Carbon atom may be saturated or unsaturated • “X” & “y” may be many different atoms (not C) • Reaction of “Acetyl Chloride” and “isopentylalcohol” to form “banana oil”, an ester

42. Hydrocarbons • Nomenclature of Alkenes • Alkenes(-C=C-) are named just as alkanes, except that the –ane suffix is changed to –ene • Alkynes(-CC-)are named in the same way, except that the suffix –yne is used • In either case, the position of the double bond is indicated by the number of the carbon

43. Hydrocarbons • Nomenclature of Alkenes - Example • First, find the longest carbon chain containing the double bond CH2CH3 CH2CHCH3 H3CHC C CH2CH2CH3 CH2CH3 6 7 CH2CHCH3 H3CHC C 1 2 3 4 5 3-propyl-5-methyl-2-heptene CH2CH2CH3

44. Hydrocarbons • Alkenes – Geometric Isomerism • In Alkanes, the C-C bond allows rotation of bonded groups; the groups continually change relative positions • In Alkenes with the C=C bond, the double bond restricts rotation around the bond • Geometric isomers are compounds joined together in the same way, but have different geometries • The similar groups attached to the two carbon atoms of the C=C bond are on the same side of the double bond in one isomer and on the opposite side for the other isomer H3C CH3 H3C H C C C C CH3 H H H trans-2-butene cis-2-butene

45. Hydrocarbons • Alkynes • General Formula - CnH2n-2 • The Carbon-Carbon (-C-C-) bond is replaced by a triple bond • Each Carbon of an Alkyne structure (-CC-) can only bond to one other Carbon in a linear structure • Each C is sp hybridized (sp – linear geometry) • Alkyne compound names are appended by thesuffix “yne” • The  electrons in both alkenes (-C=C-) and alkynes (-CC-) are “electron rich” and act as functional groups • Alkenes and alkynes are much more “reactive” than alkanes

46. Hydrocarbons • Alkynes Ethyne or Acetylene H C C H Propyne A Terminal Acetylene H C C C H 3 CH2 CH2 H3C C C CH3 3-Hexyne

47. Aromatic Hydrocarbons • Aromatic Hydrocarbons are “Planar” molecules consisting of one or more 6-carbon rings • Although often drawn depicting alternating  and  bonds, the 6 aromatic ring bonds are identical with values of length and strength between those of –C-C– & –C=C – bonds • The actual structure consists of 6  bonds and 3 pairs of  electrons “delocalized” over all 6 carbon atoms • The bond between any two carbons “resonates” between a single bond and a double bond The orbital picture shows the two “lobes” of the delocalized  cloud above and below the hexagonal plane of the -bonded carbon atoms

48. Aromatic Hydrocarbons Molecular Orbitals of Benzene

49. Aromatic Hydrocarbons H H H C H H C H C C C C C C C C H C H H C H H H benzene benzene Condensed Resonance Form of Benzene

50. Aromatic Hydrocarbons Substituted Benzenes CH3 CH3 CH3 C2CH3 Methylbenzene (Toluene) 3,4-dimethyl-ethylbenzene m,p-dimethyl-ethylbenzene