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 geometry by VSEPR
- electron pairs repel one another, maximize their separation
- BeH2 is linear (2 pairs of electrons, both in covalent bonds)
- BH3 is trigonal planar (3 pairs of electrons, all in covalent bonds)
- CH4 is tetrahedral (4 pairs of electrons, all in covalent bonds)
- atoms that satisfy the octet rule have 4 pairs of electrons - overall tetrahedral
- H2O molecular structure is bent (2 covalent bonds, 2 lone pairs)
- NH3 molecular structure is pyramidal (3 covalent bonds, one lone pair)
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-F is polarized with excess electron density closer to F
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)