2. 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-NH 2 – an amine

7. Alkanes –Only single bonds, C, H –Straight chained, branched, cyclic –IUPAC Nomenclature “International Union of Pure & Applied Chemistry” – Homologous series of Alkanes CH 3 (CH 2 ) n CH 3 -(CH 2 )- methylene group –Constitutional Isomers (branched alkanes)

11. Types of Carbon in Organic Molecules Primary C – connected to only one additional C (Methyl group) Secondary C - connected to two additional C (-CH 2 -) Tertiary C - connected to three additional C (Isopropyl group) Quaternary C - connected to four additional C (tert-Butyl group)

12. Alkanes Bond angles, Molecular Shapes

13. Alkanes Physical Properties –Gases – liquids – solids

14. 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

15. Bond Rotation - Conformations Freedom of rotation about a C-C single bond Newman Projection Formulas Potential Energy Diagrams of Bond Rotation

16. Bond Rotation - Conformations Newman Projection Formulas

17. Bond Rotation - Conformations Newman Projection Formulas

18. Bond Rotation - Conformations Potential Energy Diagrams of Bond Rotation in Ethane

19. Bond Rotation - Conformations Potential Energy Diagrams of Bond Rotation in Propane

20. Bond Rotation - Conformations Potential Energy Diagrams of Bond Rotation of Butane

21. 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

22. Equilibrium State of a reaction when there is no more change in reactant and product conc. Equilibrium constant K –A  BA + B  C + D –K = [B]/[A]K = [C][D]/[A][B] –Large k value, reaction goes to completion

24. Gibbs Standard Free Energy Change  G o = -RT ln K (in kcal/mol) Negative  G o - release of energy Free energy change – changes in bond strength (enthalpy H) & degree of order (entropy S)  G o =  H o – T  S o

25. Enthalpy Change  H o Sum of strength of bonds broken – sum of strength bonds formed Negative  H o - heat releasing, exothermic Positive  H o - heat absorbing, endothermic CH 4 + 2O 2  CO 2 + 2H 2 O  H o = -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

26. Entropy Change  S Value of S increases with increasing disorder Nitroglycerin  4 C 3 H 5 N 3 O 9  6N 2 + 12 CO 2 + 10 H 2 O + O 2 + energy (lots of it! as heat!)

27. 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, H 2 /O 2

28. Reaction Rates k = rate constant A + B  Crate: k=[A][B] [mol/Ls] –Dependent on 2 molecules “second order” A  B rate: k[A] [mol/Ls] –Dependent on 1 molecule “first order”

29. Temperature Effects on Rx rates Arrhenius Equation k = A e -Ea/RT (A = max. rate constant) More molecules have sufficient energy to overcome E a Approx. 10 o C increase  2-3x increased rate At extremely high temperature Ea/RT approaches 0, e -Ea/RT = 1 A maximum rate of particular reaction

30. Review of Acids & Bases Br Ø nsted & Lowry Definition: –Acid = H + donor –Base = H + acceptor Water (can behave as both) pure H 2 O is “neutral” H 2 O + H 2 O  H 3 O + + OH - K w = [H 3 O + ][OH - ] = 10 -14 mol 2 /L 2 [H 3 O + ]= 10 -7 mol/L (1.8  g/l water = 0.00000018%) –1.8 parts per trillion pH = -log [H 3 O + ]= 7

31. Review of Acids & Bases Acidity of Acids –HA + H 2 O  H 3 O + + A - –K = [H 3 O + ][A - ]/[HA][H 2 0] –In aqueous solution [H 2 O]  constant 55 mol/L –Acidity constant K a –K a = K[H 2 0] = [H 3 O + ][A - ]/[HA] –pK a = -log K a (  pKa = pH + pA - -pHA) –pK a = pH where 50% of acid is dissociated [A - ] = [HA] –“weak acids” pK a > 4

33. Review of Acids & Bases Basicity of Bases A - + H 2 O  OH - + HA K’ = [OH - ][HA]/[A - ][H 2 0] K b = K’[H 2 O] = [OH - ][HA]/[A - ] K a x K b = K w = 10 -14 NH 3 : pK b = 4.74 pK a : 9.26

34. 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>H 2 O>H 3 N>H 4 C Resonance favors dissociation –Acetic acid, sulfuric acid

35. Review of Acids & Bases Lewis Acids-Bases Electron Pair Acceptors – Acids –BH 3, Carbocation, AlCl 3, MgCl 2 Electron Pair Donators – Bases –OH -, R-OH, RNH 2 Important concept for many organic Rx –Conversion of a Haloalkane in to an Alcohol: –(CH 3 ) 3 C-Cl  (CH 3 ) 3 C + (carbocatioin) + Cl – –(CH 3 ) 3 C + + H 2 O  (CH 3 ) 3 C-OH + H +