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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
ALCOHOLS
H
R C OH
H
1o CH3CH2CH2CH2OH Butanol
2o ALCOHOLS R C OH
R
H
CH3CH2CHCH3
OH
2-Butanol
1-
R C OH3o ALCOHOLS
R
R
CH3 C CH3
CH3
OH
tert- Butanolor
2-Methyl 2-Propanol
1o AMINES R NH2 CH3CH2CH2 NH2 Propylamine
2o AMINES R NHR
CH3CH2 NH
CH3
Ethyl,Methylamine
3o AMINES R N RR
CH3 N CH3
CH3
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– A B A + B C + 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 + B C rate: k=[A][B] [mol/Ls]– Dependent on 2 molecules “second order”
• A B 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+
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