Chapter 15 Organic Chemistry

1. Life is chemistry. Source: Science, Jan 9, 2009, “On the origins of life on earth” Chapter 15Organic Chemistry

2. Organic chemistry is enough to drive one mad. - Friedrich Wöhler What is an organic compound? Organic compound – contains carbon, nearly always bonded to other C and H, and often other elements

3. Vitalism A major misconception that stifled organic chemistry research in early 19th Century. Resulted in the basic distinction between organic and inorganic substances An unobservable spiritual energy existed within organic compounds of living things, making them impossible to synthesize and fundamentally different from inorganic compounds (compounds of the “mineral world” – mostly, what we have studied so far) Wöhler’s experiment changed that Vitalism: Organic molecules were yhought to arise spontaneously (Spontaneous Generation) and could not be synthesized from inorganics Urea synthesized from ammonium cyanate. (2 compounds – same molecular formula) Reading for today: Did life originally arise from inorganic chemicals?

4. Classes of organic molecules I. Hydrocarbons – simplest type of organic compound -functional groups & reactivity -polymers

5. Classes of organic molecules II. Biomolecules – natural polymers -polysaccharides, proteins, nucleic acids

6. Section 15.1: What’s so special about Carbon? Atomic properties of carbon (and bonding behavior) make it special. Structural complexity of organic compounds C’s location in the periodic table tells you a lot Always bonds covalently – moderate EN makes formation of C ions energetically impossible under ordinary conditions. Why? – REVIEW Ionization energy (IE) – Chap 8 Energy required for the complete removal of 1 mole of e- from 1 mole of gaseous atoms or ions (E to overcome attraction between protons & e-)

7. Structural complexity of organic compounds C is small and forms 4 covalent bonds Ionization energy (IE) – Chap 8 As size decreases, more E to remove an e-

8. Ionization energy (IE) – Energy required to get C4+ ion = IE1 + IE2 + IE3 + IE4 A lot of energy to remove an e- ……..and to add e-’s  Electron affinity (EA): The energy change accompanying the addition of 1 mole of e-’s to 1 mole of gaseous atoms or ions. Energy is required to get C4- ion = EA1 + EA2 + EA3 + EA4 EA1 is negative (exothermic): Energy released EA2 – EA4 are positive (endothermic): Energy required

9. Structural complexity of organic compounds C forms short, strong bonds: Small size allows close approach of another atom Carbon has the ability to catenate – form chains of atoms (= large, complex molecules) Due to the sp3 hybridization: C forms 4 bonds in nearly all of its compounds

10. Structural complexity of organic compounds Carbon easily forms double and triple bonds: C – C bond is short enough to allow side-to-side overlap Double bond Triple bond

11. Structural complexity of organic compounds Double and triple bonds: Restricts rotation = MORE variety

12. Structural complexity of organic compounds C’s location in the periodic table tells you a lot – Periodic Trends So why don’t Si, Ge and Sn also form organic compounds? In same Group 4A as C. (1) Atomic size and bond strength i.e. C – C bonds = 347 kJ/mol Si – Si bonds = 226 kJ/mol (2) ∆Hreaction i.e. C – C (347), C – O (358) Si – Si (226), Si – O (368) (3) Orbitals available for reaction i.e. C has s and p orbitals Si has s, p, and d orbitals d orbitals can be attacked by lone e- pairs of incoming reactants Ethane (CH3-CH3): Stable in water and air Disilane (SiH3-SiH3): Breaks down in water, spontaneous ignition in air

13. Chemical diversity of organic compounds CRC Handbook of Physics and Chemistry - # of C-based compounds dwarfs the # of compounds formed from all of the other elements combined Chemical diversity also a result of atomic and bonding behavior of carbon. Bonding to heteroatoms: Organic compounds contain atoms other than C and H (also N, O, S, P and halogens) Example: 23 organic molecules 4 singley bonded C 1 O Filled in with H

14. Chemical diversity of organic compounds Electron density and reactivity Most chemical reactions start (and new bonds form) when a region of high e- density on one molecule meets a region of low e- density on another Regions of high e- density can be due to: (1) Multiple bonds (2) Partial charges (3) Lone pairs 4 bonds commonly found in organic molecules: C – C: Generally, unreactive – EN values equal and bond is nonpolar C – H: Largely unreactive – EN values close (C = 2.5, H = 2.1) – bond is short (strong)  C and H are both small atoms C – O: Reactive – Highly polar (∆ EN = 1.0) O end of bond is e- rich Bonds to other heteroatoms (S, P, Br): Reactive – bonds longer (S, P, Br large relative to H)

15. Chemical diversity of organic compounds Functional Group – a specific combination of bonded atoms that reacts in a characteristic way, no matter what organic molecule it occurs in In fact, reactions in organic molecule nearly always take place at functional groups. Example: Structure of amino acids

16. 20 amino acids differ only by functional group

17. Section 15.2: Hydrocarbons Organic Molecule-Animal Analogy for Hydrocarbons: • C – C bonds form the skeleton • H atoms are the skin covering the skeleton • Functional groups are limbs protruding from body ready to “grab” (react with) reactants Hydrocarbons – a large group of organic compounds containing only H and C atoms Example: Natural gas and gasoline are hydrocarbon mixtures

18. Section 15.2: Hydrocarbons As the number of carbon atoms increases, the number of arrangements increases. Carbon skeletons – What different possible arrangements exist for C atoms? For example: If you have two carbon atoms, there is one possible arrangement C – C

19. Section 15.2: Hydrocarbons Practice Drawing Hydrocarbons Purpose: Get a sense of the number of possibilities for a given formula (i.e. C6H14) Steps: #1: Are there single, double, or triple bonds? How many of each? #2: Figure out the arrangement of C atoms #3: Add the H skeleton • Six C atoms, no multiple bonds, no rings • (2) Four C atoms, one double bond, no rings • (3) Four C atoms, no multiple bonds, one ring

20. Section 15.2: Hydrocarbons • Hydrocarbon classification – 4 main groups: • Alkanes – single bonds • Alkenes – double bonds • Alkynes – triple bonds • Aromatic Hydrocarbons - rings Alkanes – CnH2n+2 Each carbon is sp3hybridized Each C is bonded to the maximum number of other atoms – saturated hydrocarbons Naming: Each chain, branch or ring has a name based on the number of carbons Prefix + root + suffix Root: # of carbon atoms in the longest continuous chain in the molecule (Table 15.1) Suffix: type of organic compound (identifies key functional group)  -ane for alkanes Prefix: groups attached to the main chain Example: Table 15.2

21. Section 15.2: Hydrocarbons Different ways to depict molecules

22. Section 15.2: Hydrocarbons Cyclic Hydrocarbons – Rings Cycloalkanes – 2 H’s are lost when ring forms from straight chain – CnH2n

23. Section 15.2: Hydrocarbons Isomers – two or more compounds with the same molecular formula but with different properties Constitutional Isomers – different arrangements of bonded atoms

24. Section 15.2: Hydrocarbons Physical Properties of Alkanes Why do we see this trend in boiling point?

25. Section 15.2: Hydrocarbons Often indicated with L and D: L-alanine D-alanine Most naturally proteins are composed of L-amino acids: L-leucine, L-glutamine. Opposite for naturally occuring carbohydrates: D-glucose metabolized, L-glucose excluded Chiral Molecules and Optical Isomerism Optical isomers – molecules are mirror images of each other

26. Section 15.2: Hydrocarbons Alkenes – CnH2n Each carbon is sp2hybridized Each C is bonded to fewer than max # of other atoms – unsaturated hydrocarbons Naming: Each chain, branch or ring has a name based on the number of carbons Prefix + root + suffix Root: # of carbon atoms in the chain that contains the double bonds (even if not longest) Suffix: type of organic compound (identifies key functional group)  -ene for alkenes Prefix: groups attached to the main chain Name these alkenes:

27. Section 15.2: Hydrocarbons Geometric Isomers:Cis-Trans Isomerism – because π bonds restrict rotation

28. Section 15.2: Hydrocarbons Alkynes – CnH2n-2 Each carbon is sp hybridized Alkanes 1 σ bond Alkenes 1 σ bond 1 π bond Alkynes 1 σ bond 2 π bonds

29. Section 15.2: Hydrocarbons Benzene is simplest example Aromatic hydrocarbons – one or more rings of 6 carbons atoms Naming = attached groups + -benzene suffix

30. Section 15.3: Organic Reaction Types Functional Group – a specific combination of bonded atoms that reacts in a characteristic way, no matter what organic molecule it occurs in Notation: R – CH2 – Br where R is an alkyl group (a saturated hydrocarbon chain) Three main reaction types: 1) Addition reactions: unsaturated reactant  saturated product Generic reaction Example: Ethylene Characteristics: • common for double and triple bonded C’s, and C = O bonds • π bonds break, σ bonds remain • reaction occurs b/c it is energetically favorable

31. Show why is this reaction energetically favorable

32. Section 15.3: Organic Reaction Types Reactants Bond Energy Products Bond Energy Addition Reaction example 2693 kJ 3098 kJ Elimination Reaction example 4410 kJ 4373 kJ What is wrong with this picture? 2) Elimination reactions: opposite of addition reactions saturated reactant  saturated product Example Generic reaction Characteristics: • Typically eliminates: 2 halogens (i.e. Cl2), H and halogen (i.e. HBr), or H and –OH group (i.e. H2O) • Driving force of this reaction is formation of small, stable molecules

33. Thermodynamics in a Nutshell Reactants Bond Energy Products Bond Energy Addition Reaction example 2693 kJ 3098 kJ Elimination Reaction example 4410 kJ 4373 kJ G – Gibbs free energy – in chemistry, the “force” that causes chemical reactions – can tell us whether or not a reaction will occur H – enthalpy – keeps track of the quantity of energy – in chemical reactions, it is the energy change during a reaction (∆Hreaction, ∆Hlattice) You can ask: Will the reaction occur spontaneously? ∆H is negative  exothermic (energy lost) = more stable = YES ∆H is positive  endothermic (energy required) = less stable = NO

34. Thermodynamics in a Nutshell Heat flow Dissolution (Chap12) Diffusion Hot Cold Proton Pump (non-spontaneous) G – Gibbs free energy – in chemistry, the “force” that causes chemical reactions – can tell us whether or not a reaction will occur S – entropy – keeps track of the distribution of energy in a system Rule: Energy becomes distributed more uniformly (more disordered) with time

35. Summary: Thermodynamics in a Nutshell G – Gibbs free energy – in chemistry, the “force” that causes chemical reactions – can tell us whether or not a reaction will occur H – enthalpy – keeps track of the quantity of energy – in chemical reactions, it is the energy change during a reaction (∆Hreaction, ∆Hlattice) You can ask: Will the reaction occur spontaneously? ∆H is negative  exothermic (energy loss as heat) = more stable = YES ∆H is positive  endothermic (energy needs to be added) = less stable = NO S – entropy – keeps track of the distribution of energy in a system – energy becomes distributed more uniformly (more disordered) with time You can ask: Will the reaction occur spontaneously? uniformity/disorder increases  YES uniformity/disorder decreases  NO

36. Reactants Bond Energy Products Bond Energy Addition Reaction example 2693 kJ 3098 kJ Elimination Reaction example 4410 kJ 4373 kJ

37. Section 15.3: Organic Reaction Types 3) Substitution reactions: Generic reaction Example Characteristics: • C involved in bonding can be saturated or unsaturated (involved in double, triple bonds)

38. Section 15.3: Redox Process in Organic Reactions More EN atom takes e- density from C (oxidation) Example: C – C bonds replaced with C – O bonds 2 CH3-CH3 + 7 O2 4 CO2 + 6 H2O Less EN atom gives e- density to C (reduction) Example: C – H bonds replaces a C – O bond CH3O  CH4 In O Chem: Focus is usually on the organic reactant only. Oxidation: C forms more bonds to O, Br, F, etc or fewer to H Reduction: C forms fewer bonds to O, Br, F, etc or more bonds to H Oxidation-reduction reactions in O-chem: Do NOT monitor change in O.N. of various C atoms in a compound. Rather, note movement of e- density around C based on # of more/less EN atoms Nature’s Redox: Respiration (Oxidation) Photosynthesis (Reduction)

39. Section 15.4: Properties & Reactivities of Functional Groups The distribution of e- density in the functional group affects the reactivity • Functional groups with single bonds only • alcohols, haloalkanes, amines (2) Functional groups with double bonds alkenes, carbonyl group (aldehydes & ketones) (3) Functional groups with both single and double bonds carboxylic acid, ester, amide (4) Functional groups with triple bonds nitrile, alkynes

40. Section 15.5: Monomers & Polymers – Synthetic Macromolecules Polymers – many monomer units bonded together

41. Section 15.5: Monomers & Polymers – Synthetic Macromolecules Petroleum-based products – there will be a shortage of raw materials soon bisphenol A (BPA) - used in synthesizing DGEBA, a building block for an epoxy resin

42. Section 15.5: Monomers & Polymers – Synthetic Macromolecules Addition polymers – as each monomer adds to the chain, it forms a new reactive site.

43. Section 15.6: Monomers & Polymers – Biological Macromolecules

44. Section 15.6: Polysaccharides Glucose is a monosaccharide – alcohol and aldehyde groups react to make cyclic forms Polysaccharide chains formed from cyclic forms that undergo dehydration reactions. Different disaccharides formed from different monosaccharides: sucrose (table sugar): glucose (C-1) + fructose (C-2) lactose (milk sugar): glucose (C-1) + galactose (C-4) maltose (beer): glucose (C-1) + glucose (C-4)

45. Section 15.6: Polysaccharides 3 main groups of polysaccharides: Cellulose – most abundant organic chemical on earth, structural function (plant cell walls), long chains of glucose, humans cannot digest this (cows, sheep, termites) Starch – energy storage in plants (amylose and amylopectin) Glycogen – energy storage in animals (C6H10O5)n

46. Section 15.6: Amino Acids and Proteins

47. Section 15.6: Amino Acids and Proteins tyrosine valine leucine tyrosine Ionic bonds Hydrogen bonds Disulfide bonds (covalent) + hydrophobic interactions between –CH3

48. Section 15.6: Nucleic Acids, DNA, and RNA Pyrimidines: Thymine (T) [Uracil (U)], cytosine (C) Purines: Guanine (G), Adenine (A) A – T(U) , G – C