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Structure & Reactivity Alkanes – Molecules w/o functional Groups • Hydrocarbons – Alkanes, Alkenes, Alkynes. • Functional Groups; Aromatics – Polar bonds create chemical reactivity – Haloalkanes, Alcohols,Phenols, Ethers, Carbonyls, Aldehydes, Ketones, Carboxylic Acids, Anhydrides, Esters, Amides, Nitriles, Amines, Thiols • “R” – residue (Alkyl Group) – R-OH – an alcohol – R-NH2 – an amine H 1o ALCOHOLS R C OH CH3CH2CH2CH2OH 1-Butanol OH CH3CH2CHCH3 2-Butanol H R 2o ALCOHOLS R C OH H CH3 R 3o ALCOHOLS R C OH CH3 2o AMINES 3o AMINES R NH 2 R R NH R R N CH 3 OH R 1o AMINES C CH3CH2CH2 NH 2 tert- Butanol or 2-Methyl 2-Propanol Propylamine CH3 CH3CH2 NH Ethyl,Methylamine CH3 R CH3 N CH 3 Tri-methylamine Alkanes – Only single bonds, C, H – Straight chained, branched, cyclic – IUPAC Nomenclature “International Union of Pure & Applied Chemistry” – Homologous series of Alkanes CH3(CH2)nCH3 • -(CH2)- methylene group – Constitutional Isomers (branched alkanes) Types of Carbon in Organic Molecules • Primary C – connected to only one additional C (Methyl group) • Secondary C - connected to two additional C (-CH2-) • Tertiary C - connected to three additional C (Isopropyl group) • Quaternary C - connected to four additional C (tert-Butyl group) Alkanes • Bond angles, Molecular Shapes Alkanes • Physical Properties – Gases – liquids – solids Intermolecular Forces • A: Ionic compounds (salts) – very strong Coulomb attraction • B:Polar compounds (e.g. Haloalkanes) – Dipole-dipole interaction • C:Nonpolar compounds (alkanes) – Very weak London forces Bond Rotation - Conformations • Freedom of rotation about a C-C single bond • Newman Projection Formulas • Potential Energy Diagrams of Bond Rotation Bond Rotation - Conformations • Newman Projection Formulas Bond Rotation - Conformations • Newman Projection Formulas Bond Rotation - Conformations • Potential Energy Diagrams of Bond Rotation in Ethane Bond Rotation - Conformations • Potential Energy Diagrams of Bond Rotation in Propane Bond Rotation - Conformations • Potential Energy Diagrams of Bond Rotation of Butane Kinetics & Thermodynamics • Chemical Thermodynamics – Changes in energy during a reaction, determines the extent to which a reaction goes to completion • Chemical Kinetics – Velocity, rate of a reaction (change in concentration of reactants/product) • Reaction may be under thermodynamic or kinetic control Equilibrium • State of a reaction when there is no more change in reactant and product conc. • Equilibrium constant K –AB A+BC+D – K = [B]/[A] K = [C][D]/[A][B] – Large k value, reaction goes to completion Gibbs Standard Free Energy Change • Go = -RT ln K (in kcal/mol) • Negative Go - release of energy • Free energy change – changes in bond strength (enthalpy H) & degree of order (entropy S) • Go = Ho – T So Enthalpy Change Ho • Sum of strength of bonds broken – sum of strength bonds formed • Negative Ho - heat releasing, exothermic • Positive Ho - heat absorbing, endothermic • CH4 + 2O2 CO2 + 2H2O Ho = -213 kcal/mol – 1 mol methane = 16g – 213 kcal/16g = 13.3 kcal/g – Fats: 9 kcal/g – Alcohol: 7 kcal/g – Sugars: 4 kcal/g Entropy Change S • Value of S increases with increasing disorder • Nitroglycerin • 4 C3H5N3O9 6N2 + 12 CO2 + 10 H2O + O2 + energy (lots of it! as heat!) Activation Energy • Most exothermic reactions do not occur spontaneously • Bond breaking precedes bond formation • Reaching of Transition State requires Activation Energy (input) – E.g. gasoline, wood, H2/O2 Reaction Rates k = rate constant • A+BC rate: k=[A][B] [mol/Ls] – Dependent on 2 molecules “second order” • AB rate: k[A] [mol/Ls] – Dependent on 1 molecule “first order” Temperature Effects on Rx rates • Arrhenius Equation • k = A e-Ea/RT (A = max. rate constant) • More molecules have sufficient energy to overcome Ea • Approx. 10oC increase 2-3x increased rate • At extremely high temperature Ea/RT approaches 0, e-Ea/RT = 1 • A maximum rate of particular reaction Review of Acids & Bases • BrØnsted & Lowry Definition: – Acid = H+ donor – Base = H+ acceptor • Water (can behave as both) pure H2O is “neutral” • H2O + H2O H3O+ + OH• Kw = [H3O+][OH-] = 10-14 mol2/L2 • [H3O+]= 10-7 mol/L (1.8g/l water = 0.00000018%) – 1.8 parts per trillion • pH = -log [H3O+]= 7 Review of Acids & Bases • Acidity of Acids – HA + H2O H3O+ + A– K = [H3O+][A-]/[HA][H20] – In aqueous solution [H2O] constant 55 mol/L – Acidity constant Ka – Ka = K[H20] = [H3O+][A-]/[HA] – pKa = -log Ka ( pKa = pH + pA- -pHA) – pKa = pH where 50% of acid is dissociated [A] = [HA] – “weak acids” pKa > 4 Review of Acids & Bases • • • • • • Basicity of Bases A- + H2O OH- + HA K’ = [OH-][HA]/[A-][H20] Kb = K’[H2O] = [OH-][HA]/[A-] Ka x Kb = Kw = 10-14 NH3: pKb = 4.74 pKa: 9.26 Reasons for Acid/Base Strengths • Increasing size of anion A- allows better distribution of negative charge – HI>HBr>HCl>HF • Electronegativity of the element to which H is attached: – HF>H2O>H3N>H4C • Resonance favors dissociation – Acetic acid, sulfuric acid Review of Acids & Bases • Lewis Acids-Bases • Electron Pair Acceptors – Acids – BH3, Carbocation, AlCl3, MgCl2 • Electron Pair Donators – Bases – OH-, R-OH, RNH2 • Important concept for many organic Rx – Conversion of a Haloalkane in to an Alcohol: – (CH3)3C-Cl (CH3)3C+ (carbocatioin) + Cl – – (CH3)3C+ + H2O (CH3)3C-OH + H+