Transcript Alkanes – Molecules w/o functional Groups

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
2o ALCOHOLS
R
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
–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 (1.8g/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
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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+