Chemistry 331 - Winter 1996
Elements of Organic Chemistry I

Professor Carl C. Wamser

Chapter 1 - Structure and Bonding

Tues, Jan 9

Why organic chemistry?

• carbon has unique chemistry
• bonds to every other element
• bonds to itself in long chains
• organic chemistry involves enormous variety
• many possible structures
• many possible reactions

Why learn organic chem?

• what do organic chemists do?
• understanding structures, reactions
• correlation of structures with properties
• synthesis of compounds with specific properties
• who else needs organic?
• basis of all life processes
  - the great variety of structures and reactions make life possible

How to handle variety

• nomenclature
• clear methods for naming structures and reactions
• structures
• organized by functional groups
• reactions
• organized by reaction types (what happens?)
• organized by reaction mechanisms (how does it happen?)

What should you get out of organic?

• from complex names, be able to derive a structure
• from complex structures, be able to identify functional groups, predict characteristic properties
• work through reaction mechanisms - what is a molecule likely to do under certain conditions
• "think like a molecule"

A reaction example

• CH3OH + HCl --> CH3Cl + H2O
• Reaction type ?
• acid/base, oxidation/reduction, addition/elimination, substitution, rearrangement
• substitution - of Cl for OH (identify bonds broken and made)
• Reaction mechanism ?
• how does the reaction occur (step-by-step)

A reaction mechanism

• Break the molecules into atoms
• Reassemble into products
• What's wrong with this mechanism?
• It would take too much energy (unnecessarily)

A better mechanism

• Balance bond breaking with bond making
• First an acid-base reaction:
  CH3OH + HCl --> CH3OH2+ + Cl-
• Then a substitution:
  CH3OH2+ + Cl- --> CH3Cl + H2O

The Periodic Table

• atomic number (defines element)
• atomic weight (isotopes)
• electron shells (rows)
• groups (similar properties)
• filled shells (the noble gases)
• valence electrons (for bonding)

Lewis Structures

• the octet rule
  - atoms strive to get 8 electrons (full shell)
• ionic bonding
  - gaining or losing electrons
  - stabilized by Coulombic attractions
• covalent bonding
  - sharing electrons
  - most common bonding in organic compounds

Atomic Orbitals

• wavefunctions
  - describe location of electrons
• s orbital (spherical)
• p orbitals (three: x,y,z)
  - (dumbbell shape - 2 lobes)
• d orbitals (4 lobes)
  - not usually needed for organic chemistry
• hybrid orbitals
  - combination orbitals

Hybrid Orbitals

• sp hybrids (one s plus one p)
  - makes two identical orbitals (linear)
• sp2 hybrids (one s plus two p)
  - makes three identical orbitals (trigonal)
• sp3 hybrids (one s plus three p)
  - makes four identical orbitals (tetrahedral)

Bonding

• attraction between negative electrons and positive nuclei
  - repulsions between electrons
  - repulsions between nuclei
• bonding is a balance between the attractions and repulsions
• characteristic bond lengths and strengths

Molecular Orbitals

• overlap of atomic orbitals
• electrons are close to two nuclei
• bonding and antibonding combinations

Sigma Bonds

• cylindrical symmetry
• formed by end-on overlap

Pi Bonds

• nodal plane through bond axis
• formed by side-on overlap

Writing Organic Structures

• Lewis structures
  - all electrons shown
• Kekule structures
  - show bonds as lines
  - lone pairs sometimes omitted
• line structures
  - omit lone pairs
  - omit hydrogens on carbons
  - omit carbons
  (assumed to be at the end of every bond)

3-Dimensional Structures

• dotted-line / wedge
• ball-and-stick
• space-filling

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Visualizing chemical structures

• Name (common or systematic)
• Condensed Formula (as usually typed out)
• Lewis Structure (all atoms and bonds shown)
• Line Structure (omit hydrogens, assume carbons at vertices)
• 3-D Structure (show bond orientations)
• Ball-and-Stick Structure (like a molecular model you could make)
• Space-Filling Model (approximates full size of electron distribution)

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methane: CH4

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benzene: C6H6

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penicillin:

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Try a 3-D image:

1) Click at the bottom after you read these instructions
2) Cross your eyes to see two double images
3) Make the middle images merge

Try it!

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Thurs, Jan 11

Typical Valence

• H 1 valence electron 1 bond
• C 4 valence electrons 4 bonds
• N 5 valence electrons 3 bonds + 1 lone pair
• O 6 valence electrons 2 bonds + 2 lone pairs
• F 7 valence electrons 1 bond + 3 lone pairs

Molecular Orbitals

• H + H ---> H-H
• s + s ---> sigma bond
• H-H bond length 0.74 Å
• H-H bond strength 104 kcal/mole

Methane

• C + 4 H ---> CH4
• sp3 + s ---> four sigma bonds
• C-H bond length 1.10 Å
• C-H bond strength 104 kcal/mole
• orientation - tetrahedral
• bond angles 109.5°

Why Hybrid Orbitals?

• good shape (directional)
• allows high overlap in bonding
• maximizes electron density between atoms
• good orientation
• minimizes repulsions between orbitals

Ethane

• CH3CH3
• includes a C-C sigma bond
• overlap of sp3 - sp3 orbitals
• C-C bond length 1.54 Å
• C-C bond strength 88 kcal/mole
• C-H bond length 1.10 Å
• C-H bond strength 98 kcal/mole
• orientation - two tetrahedra

Ethylene

• CH2=CH2
• includes a C=C double bond (alkene family)
• overlap of sp2 - sp2 orbitals (sigma bond)
• overlap of p - p orbitals (pi bond)
• C-C bond length 1.33 Å
• C-C bond strength 152 kcal/mole
• C-H bond length 1.08 Å
• C-H bond strength 103 kcal/mole
• orientation - planar
• bond angles 120°

Acetylene

• HC=CH
• includes a C=C triple bond (alkyne family)
• overlap of sp - sp orbitals (sigma bond)
• overlap of p - p orbitals (pi bonds - two)
• C-C bond length 1.20 Å
• C-C bond strength 200 kcal/mole
• C-H bond length 1.06 Å
• C-H bond strength 125 kcal/mole
• orientation - linear

Bonding Trends

Bond Lengths in Angstroms (Bond Strengths in kcal/mole)
Compound Hybrid C-C bond C-H bond
Ethane sp3 1.54 (88) 1.10 (98)
Ethylene sp2 1.33 (152) 1.08 (103)
Acetylene sp 1.20 (200) 1.06 (125)

• multiple bonds are stronger
• shorter bonds are stronger

Electronegativity

• tendency of an atom to attract electrons in a covalent bond
• in the Periodic Table, electronegativity increases to the right and up
• F > O > Cl ~ N > Br > C > H > metals

Polar Covalent Bonds

• electrons in a covalent bond may not be equally shared
  H-Cl is actually d+ H-Cl d-
• where d+ or d- represents a partial charge

Polar Bonds to Carbon

• C-C bonds are nonpolar
• C-H bonds are generally considered nonpolar
• C-X bonds are polarized with carbon d+
  for X = F, Cl, Br, I, O, S, N
• C-M bonds are polarized with carbon d-
  for M = metals

Bronsted-Lowry Acid/Base

• acid - donates H+
• base - accepts H+
  NH3 + H2O <==> NH4+ + OH-
  base + acid <==> acid + base
• note conjugate acid-base pairs
  (differ by H+)

Acidity Constant (Ka)

• HA + H2O <==> H3O+ + A-

usually simplified to
HA <==> H+ + A-

• Ka = [H+] [A-] / [HA]

Acid Strength (pKa)

• stronger acids have higher Ka
  for HCl, Ka = 107
  for CH3COOH, Ka = 10-5
  (acetic acid, found in vinegar)
  
  
• pKa = - log Ka
• stronger acids have a lower pKa
  for HCl, pKa = -7
  for CH3COOH, pKa = 5

pH and pKa

• Ka = [H+][A-]/[HA]
• pKa = pH - log([A-]/[HA])
• for pH = pKa, [A- ] = [HA]
• for pH < pKa, HA predominates
• for pH > pKa, A- predominates
  e.g., for acetic acid at pH = 7
  [CH3COO-] > [CH3COOH]

Acid-Base Reactions

CH3COOH + OH- <==> CH3COO- + H2O
acid base base acid
pKa = 5 pKa = 15.7
stronger acid
reaction favored to the right

Acid-Base Reactions

CH3COOH + H2O <==> CH3COO- + H3O+
acid base base acid
pKa = 5 pKa = 1.7
reaction favored to the left

Acid-Base Reactions

HC=CH + Na+NH2- <==> HC=C-Na+ + NH3
acid base base acid
pKa = 25 pKa = 35
reaction favored to the right

Acid-Base Reactions

CH3OH + CN- <==> CH3O- + HCN
acid base base acid
pKa = 16 pKa = 9
reaction favored to the left

Lewis Acids

• acid - accepts an electron pair
• base - donates an electron pair
• (making a new covalent bond)

H+ + H2O <==> H3O+
L acid L base new O-H bond
BF3 + NH3 <==> F3B-NH3
L acid L base new B-N bond

Tues, Jan. 16

Formal Charge

• #valence e- - (# bonds + # lone e-)
• (How many e- is normal for this atom?)
  - (How many e- in this compound?)

Resonance

• more than one possible Lewis structure for a compound
• What's the best Lewis structure?
• follow the octet rule
• electronegativity determines the best place to locate charges
  carbon monoxide
  nitromethane