Bonding and Structure
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
- Reaction mechanism
- how does the reaction occur ? (step-by-step)
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 =
ethyl or Et)
- 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
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)
methane: CH
4 
benzene:
C
6H
6
penicillin:
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
Review of Acid-Base Reactions
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)
usually simplified to
HA <==> H+ + A-
Acid Strength (pKa)
- stronger acids have higher Ka
for HCl, Ka = 10E7
for CH3COOH, Ka = 10E-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]
Structural Effects on Acid Strength
- electronegativity
HF > H2O > NH3 > CH4
- weaker bond to H
HI > HBr > HCl > HF
- inductive effects - electron withdrawal
H2SO4 > H2SO3
Cl-CH2-COOH > CH3-COOH
- hybridization with greater s-character
sp C-H > sp2 C-H > sp3 C-H
- delocalization
RCOOH > ROH
Acid-Base Reactions
CH3COOH + OH- <==> CH3COO-
+ H2O
acid (pKa = 5) + base <==> base + acid (pKa = 15.7)
reaction favored for stronger acid
reaction favored to the right
Lewis Acids
- acid - accepts an electron pair
- base - donates an electron pair
(making a new covalent bond)
