Chem 334 - Fall 1999
Organic Chemistry I
Dr. Carl C. Wamser

Chapter 1 Notes: Overview / Bonding and Structure

Overview of Organic Chemistry

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

Organic chemists

• 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?)
• reactions - organized by reaction mechanisms (how does it happen?)

What should you get out of organic chemistry?

• 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
  • what happens ?
• Reaction mechanism
  • how does the reaction occur ? (step-by-step)

Bonding and Structure

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)

Bonding - Lewis structures

• the octet rule
• ionic bonding
• covalent bonding
  - most common bonding in organic compounds

Typical valence - neutral atoms in normal bonding patterns

• H has 1 valence electron - makes 1 bond
• C has 4 valence electrons - makes 4 bonds
• N has 5 valence electrons - makes 3 bonds + 1 lone pair
• O has 6 valence electrons - makes 2 bonds + 2 lone pairs
• F has 7 valence electrons - makes 1 bond + 3 lone pairs
• Recognize the appearance of these common atoms in correct structures
• Also recognize the common charged forms:

Atom	Bonds	Lone Pairs	Charge
H	1	0	0
H+	0	0	+1
H-	0	1	-1
C	4	0	0
C+	3	0	+1
C.	3	1 e-	0
C-	3	1	-1
N	3	1	0
N+	4	0	+1
N-	2	2	-1
O	2	2	0
O-	1	3	-1
O+	3	1	+1
F	1	3	0
F-	0	4	-1

Writing organic structures

• Lewis structures
  - all electrons shown
• Kekule structures
  - show bonds as lines
  - lone pairs sometimes omitted
• condensed structures
  - common groups are abbreviated (e.g., CH3CH2 )
• 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

• 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)

Your textbook includes a stereoviewer to make this process easier.

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

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

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)

Why hybrid orbitals?

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

Identifying the hybridization of carbon

• identify sigma and pi bonds around the carbon atom
  (you need one sigma bond for each neighboring atom
  and you need a total of four bonds for carbon)

Neighboring Atoms	Sigma Bonds	Pi Bonds	Hybrid	Structure
4	4	0	sp3	tetrahedral
3	3	1	sp2	trigonal planar
2	2	2	sp	linear

Molecular geometry by VSEPR

• electron pairs repel one another, maximize their separation

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 polarized with excess electron density closer to Cl

Polar bonds to carbon

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

Functional Groups

• a specific arrangement of atoms
• define chemical families
• determine chemical properties
• basis of nomenclature
• organization of textbooks
• examples - the simple oxygen families:
  • alcohols, ethers, aldehydes, ketones, carboxylic acids

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 (CO)
    nitromethane (CH3NO2)

Electron-pushing

• keeping track of electrons is crucial in organic chemistry
• curved electron arrows indicate electron pair movement
  - in converting resonance forms, or later, in reactions