Chapter 2

Chapter 2 - Chapter 2: Alkanes, Thermodynamics, and Kinetics

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Unformatted text preview: Chapter 2: Alkanes, Thermodynamics, and Kinetics 2,2,4-Trimethylpentane: An octane All Reactions Are Equilibria "Barrier" kcal/mol Exothermicity ~20 or CH3Cl + Na+ -OH CH4 + O2 CH3OH + Na+ Cl-23.4 kcal/mol high CO2 + 2H2O -213 kcal/mol Equilibrium lies very much to the right. What governs these equilibria? Chemical Thermodynamics and Kinetics 1. Chemical Thermodynamics: Energy changes during reaction, extent of "completion of equilibration," "to the left/right," "driving force." 2. Chemical Kinetics: How fast is equilibrium established; rates of disappearance of starting materials or appearance of products The two principles may or may not go in tandem Equilibria: Two typical cases 1. A K B K= [B] [A] = [products] [reactants] K = equilibrium constant [ ] = concentration in mol L-1 K 2. A +B C+D K= [C][D] [A][B] If K large: reaction "complete," "to the right," "downhill." How do we quantify? Gibbs free energy, !G Gibbs Free Energy, G !G = -2.3 RT logK T in kelvins, K (zero kelvin = -273 C) R = gas constant ~ 2cal deg-1 mol-1 Large K : Large negative !G : downhill Equilibria and Free Energy At 25C (298K): G = - 1.36 logK Enthalpy H and Entropy S !G = !H - T!S Kcal mol-1 cal-1 deg-1 mol-1 or entropy units, e.u. Enthalpy !H = heat of the reaction; for us, mainly due to changes in bond strengths: !H = (Sum of strength of bonds broken) (sum of strengths of bonds made) Example: _ CH3CH2 H + 101 Cl_Cl 58 CH3CH2_Cl + H_Cl 84 103 !H = 159 187 = -28 kcalmol-1 !H negative: called "exothermic" if positive: called "endothermic" !S = change in the "order" of the system. Nature strives for disorder. More disorder = positive !S (makes a negative contribution to !G ) Chemical example: CH2 CH2 + HCl 2 molecules CH3CH2Cl 1 molecule !H = -15.5 kcal mol-1 !S = -31.3 e.u. If # of molecules unchanged, !S small, !H controls ( we can estimate value from bond strength tables) Rates All processes have "activation barriers". Rate controlled by: 1. Barrier height (structure of transition state TS) 2. Concentration (the number of collisions increase with concentration) 3. T (increased T means faster moving molecules; number of collisions increases) 4. "Probability" factor (how likely is a collision to lead to reaction; depends on sterics, electronics) Boltzmann Distribution What supplies the energy to get over the barrier? The average kinetic energy of molecules at room temperature is ~ 0.6 kcal/mol. Reaction Rate Rate measurements : Give Rate Laws, tell us something about TS structure. Most common: 1. If rate = k [A] A B 1st order rate law Unimolecular reaction (TS involves only A) 2. A+B C 2nd order rate law If rate = k [A][B] Bimolecular reaction (TS involves both A and B). How do we measure barrier ? Energy of activation from Arrhenius equation: -Ea k = Ae RT at high T, k = A, "maximum rate" Potential Energy Diagrams [TS] E kf Reactant Ea !G kr !H (when !S small) Product [A] [B] Reaction coordinate = progress of reaction K= [B] [A] = k forward k reverse Rate Determining Transition State Many reactions have many steps, but there is always a rate determining TS (bottleneck). TS Problem: Which is right: On heating, a. Compound A converts to C directly. b. It goes first to B and then to C. c. It stays where it is. B A C Acid-Base Equilibria HA + H2O Acid H3O + A Conjugate Base + - Brnsted and Lowry: Acid = proton donor Base = proton acceptor Acid-Base: Electron "Pushing" and Electrostatics H O H Charge moves: +1 A B -1 + H H + H Cl + H O H H + Cl + - e-pushing arrows HA + H2O [H3O] [A] K= [HA] [H2O] + H3O + A + - - Solvent 55 mol/L - Acidity constant [H3O][A] Ka = K x 55 = [HA] + pKa = -log Ka Acidity Acidity increases with: 1. Increasing size of A (H A gets weaker; charge is better stabilized in larger orbital; down the PT) 2. Electronegativity (moving to the right in PT) 3. Resonance, e.g., CH3OH 15.5 :O : CH3COH 4.3 : : : : : : : : pKa CH3O: :O : CH3 : : : : :O: H2SO4 - -5.0 C O: : : : : :O S OH :O: : : : : : : : : Relative Acid Strengths Strong Weak Very weak Lewis Acids and Bases Lewis acids: e-deficient Lewis bases: Lone e-pairs F B F F 6e R R O R R O R _ _ R S R R _ - R N R F B F F O BF3 + - e-pushing arrows Lewis Acid-Base Electrostatics F F B + F O CH2CH3 CH2CH3 F B F - F + O CH2CH3 CH2CH3 + Alkanes Hydrocarbons without functional groups Line notation: 1 = 10-8 cm Straight chain: CH3CH2CH2CH3 C4H10 Branched: CH3 CH Butane H CH C CH3 CH3 C4H10 2-Methylpropane and are constitutional isomers. Same molecular formula, different connectivity Cyclic: Cyclohexane C6H12 Bicyclic: Bicyclo[2.2.0]octane C8H14 Polycyclic . . . . . . Homologous series: Insert-CH2- groups into C-C bonds. Straight chain CH3(CH2)xCH3 General molecular formula for acyclic systems. Cyclic alkanes: CnH2n Barry Sharpless (Scripps) NP 2001 Date Mon Sep 12 23:56:24 EDT 2005 26,676,640 organic and inorganic substances 56,744,740 sequences Virtual Exploration of the Small-Molecule Chemical Universe below 160 Daltons Tobias Fink, Heinz Bruggesser, and Jean-Louis Reymond* The development of modern medicine largely depends on the continuous discovery of new drug molecules for treating diseases. One striking feature of these drugs is their relatively small molecular weight (MW), which averages only 340. Recently, drug discovery has focused on even smaller building blocks with MW of 160 or less to be used as lead structures that can be optimized for biological activity by adding substituents. At that size it becomes legitimate to ask how many such very small molecules would be possible in total within the boundaries of synthetic organic chemistry? To address this question we have generated a database containing all possible organic structures with up to 11 main atoms under constraints defining chemical stability and synthetic feasibility. The database contains 13.9 million molecules with an average MW of 153, and opens an unprecedented window on the small-molecule chemical universe. Angew. Chem. Int. Ed. 2005, 44, 1504 1508 (edited) The Names of Alkanes are Based on the IUPAC Rules Naming Alkyl Substituents Change ending ane to yl, as in methane / methyl, hexane / hexyl Short notation: Alkane R-H / alkyl R"Lingo": RCH2 "primary" R R R C "secondary" R C "tertiary" H R IUPAC Rules 1. Find the longest chain and name it (Table 2-5) CH3CHCH2CH3 CH3 A (methyl substituted) butane An octane (substituted by ethyl, two methyls) When there are two equal longest chains, choose the one with more substituents 4 substituents 3 substituents 2. Name substituents (as alkyl or halo) Halo: Bromo, fluoro, chloro, iodo a. For straight chain R: Methyl, ethyl, propyl etc. b. For branched chain R: Find longest chain (starting from point of attachment). Name substituents Example: (Methylpropyl) Branched Alkyl Groups c. Multiple same substituents: For R = straight, use prefix di-, tri-, tetra-, penta-, etc.: Dimethylhexane For R = branched, use: bis-, tris-, tetrakis-, etc., and alkyl name in parentheses: Bis(methylpropyl) d. Common names: we will use colloquially isopropyl, tert-butyl, neopentyl 3. Number stem, starting from the end closest to a substituent: 2 1 2 3 4 1 3 4 5 6 7 8 9 If both ends equidistant to the first substituent, proceed until the first point of difference: 5 7 1 2 3 4 6 8 9 Branched substituents: Number from carbon of attachment (C1) Defined as 1 1 2 3 4. Name the alkane in alphabetical (not numerical) order of substituents, location given by number prefix. 8 7 6 5 4 3 2 1 5-Ethyl-2-methyl2-Methylbutane octane Alphabet: Di-, tri-, etc. not counted for main stem R. But: Counted when in branched R 8 7 6 5 4 3 2 1 7 6 4 2 1 5-Ethyl-2,2-dimethyloctane Not counted Counted 8 5 3 5-(1,1-Dimethylethyl)3-ethyloctane { { Problem: I Cl Br Longest chain? 8 7 6 5 4 3 1 2 I Cl Br Substituents? 8 7 6 5 4 3 1 2 Iodo I 1-Chloroethyl Cl Dimethyl Br Bromo Final name? 8 7 6 5 4 3 1 2 Iodo I 1-Chloroethyl Cl Dimethyl Br Bromo 1-Bromo-5-(1-chloroethyl)-7-iodo-2,2-dimethyloctane Physical Properties of Alkanes: Intermolecular Forces Increase With Size Intermolecular Forces Coulomb forces in salts Dipole-dipole interactions in polar molecules Intermolecular Forces Idealized (pentane) Experimental (heptane) London forces: Electron correlation (Polarizability: Deformability of e-cloud) The Rotamers of Ethane Staggered Eclipsed Staggered Newman Projections Note: Newman projection occurs along only one bond. Everything else is a substituent. Rotation with Newman Projections Rotation Around Bonds is Not "Free": Barriers to Rotation Ethane has barrier to rotation of ~3 kcal mol-1. Barrier due to steric and electronic effects. Most stable rotamer is staggered Transition state is eclipsed e-Repulsion antibonding bonding Orbital stabilization Potential Energy Diagrams (TS = transition state) Propane: Methyl Increases Barrier Butane: Isomeric Staggered and Eclipsed Rotamers Rotamers and Energy Diagram ...
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This note was uploaded on 03/21/2012 for the course CHEM 140A taught by Professor Whiteshell during the Fall '04 term at UCSD.

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