drawing organic structures functional groups … of amines and amides •classify these functional...
TRANSCRIPT
STRUCTURE
Dr. Sheppard
CHEM 2411
Spring 2015
Klein (2nd ed.) sections 1.8-1.10, 1.12-1.13, 2.7-2.12, 3.2, 3.4-3.5, 3.8-3.9, 4.6-4.13, 4.14, 8.5, 15.16, 21.3
Topics
• Structure
• Physical Properties
• Hybridization
• Resonance
• Acids and Bases
• Conformations of Alkanes and Cycloalkanes
• Unsaturation
• Alkene Stability
Structure
• Drawing organic structures
• Sigma (s) and pi (p) bonds
• Single bonds = 2e- = one sigma bond
• Double bonds = 4e- = one sigma bond and one pi bond
• Triple bonds = 6e- = one sigma bond and two pi bonds
• Which bond is shortest? Longest? Weakest? Strongest?
• Remember formal charges
Ionic Structures
• Be on the lookout for metals (cations) and ions
• Example: NaOCH3
• This is a Na+ cation and a CH3O- anion
• Example: NH4Cl
• This is a NH4+ cation and a Cl- anion
Classification of Atoms
• C atoms can be classified as:
• Primary (1º) = C bonded to 1 other C
• Secondary (2º) = C bonded to 2 other C
• Tertiary (3º) = C bonded to 3 other C
• Quaternary (4º) = C bonded to 4 other C
Classification of Atoms
• H atoms are classified based on the type of carbon to
which they are attached
Classification of Alcohols and Alkyl Halides
• Alcohols and alkyl halides are classified based on the
type of carbon to which the -OH or –X is bonded
• Classify these alcohols/alkyl halides as 1º, 2º or 3º:
Alkyl Halide Structure
• In addition to primary, secondary, and tertiary, alkyl
halides can be classified as:
• Geminal
• Vicinal
• Vinyl
• Aryl
Classification of Amines and Amides
• Amines and amides are classified based on the number of
C atoms bonded to the N
Classification of Amines and Amides
• Classify these functional groups:
Electronegativity and Bond Polarity
• Electronegativity
• Ability of atom to attract shared electrons (in a covalent bond)
• Most electronegative atom = F
• Differences in electronegativity determine bond polarity
• Bond polarity
• How electrons are shared between nuclei
• Equal sharing of electrons = nonpolar; unequal = polar
Bond Polarity
• Example: C─O
• What atom is more electronegative (C or O)?
• More EN atom has partial negative charge (d-)
• Less EN atom has partial positive charge (d+)
• Arrow shows direction of polarity
• Nonpolar bonds
• Any atom with itself
• C─H
Molecular Dipole Moment
• Overall electron distribution within a molecule
• Depends on bond polarity and bond angles
• Vector sum of the bond dipole moments (consider both magnitude and direction of individual bond dipole moments)
• Lone pairs of electrons contribute to the dipole moment
• Symmetrical molecules with polar bonds = nonpolar
Intermolecular Forces
• Strength of attractions between molecules
• Based on molecular polarity
• Influence physical properties (boiling point, solubility)
1. Dipole-dipole interactions
2. Hydrogen bonding
3. London dispersions (van der Waals)
1. Dipole-Dipole Interactions
• Between polar molecules
• Positive end of one molecule aligns with negative end of
another molecule
• Lower energy than repulsions
• Larger dipoles cause higher boiling points
2. Hydrogen Bonding
• Strongest dipole-dipole attraction
• H-bonded molecules have higher boiling points
• Organic molecule must have N-H or O-H
• The hydrogen from one molecule is strongly attracted to a
lone pair of electrons on the other molecule
3. London Dispersion Forces
• van der Waals forces
• Exist in all molecules
• Important with nonpolar compounds
• Temporary dipole-dipole interactions
• Molecules with more surface area have stronger
dispersion forces and higher boiling points
• Larger molecules
• Unbranched molecules
CH3 CH2 CH2 CH2 CH3
n-pentane, b.p. = 36°C
CH3 CH
CH3
CH2 CH3
isopentane, b.p. = 28°C
C
CH3
CH3
CH3
H3C
neopentane, b.p. = 10°C
Boiling Points and Intermolecular Forces
CH3 O CH3CH3 CH2 OH
dimethyl ether, b.p. = -25°C ethanol, b.p. = 78°C
N CH3H3C
CH3
CH3CH2CH2 N
H
HN CH3CH3CH2
H
ethanol, b.p. = 78°C ethyl amine, b.p. 17°C
trimethylamine, b.p. 3.5°C propylamine, b.p. 49°C ethylmethylamine, b.p. 37°C
CH3 CH2 OH CH3 CH2 NH2
Solubility and Intermolecular Forces
• Like dissolves like
• Polar solutes dissolve in polar solvents
• Nonpolar solutes dissolve in nonpolar solvents
• Molecules with similar intermolecular forces will mix freely
Example
• Which of the following from each pair will have the higher
boiling point?
(a) CH3CH2CH2CH3 CH3CH2CH2OH
(b) CH3CH2NHCH3 CH3CH2CH2NH2
(c) CH3CH2CH2CH2CH3 CH3CH2CH(CH3)2
Example
• Will each of the following molecules be soluble in water?
(a) CH3CO2H
(b) CH3CH2CH3
(c) CH3C(O)CH3
(d) CH2=CHCH3
Example
• Draw the structure of the alkane with the molecular
formula C5H12 that has the lowest boiling point.
Structure of Organic Molecules
• Previously:• Atomic/electronic structure
• Lewis structures
• Bonding
• Now:• How do atoms form covalent bonds?
• Which orbitals are involved?
• What are the shapes of organic molecules?
• How do bonding and shape affect properties?
Linear Combination of Atomic Orbitals
• Bonds are formed by the combination of atomic orbitals
containing valence electrons (bonding electrons)
• Two theories:
• Molecular Orbital Theory
• Atomic orbitals of two atoms interact
• Bonding and antibonding MO’s formed
• Skip this stuff
• Valence Bond Theory (Hybridization)
• Atomic orbitals of the same atom interact
• Hybrid orbitals formed
• Bonds formed between hybrid orbitals
• How many valence electrons? In which orbitals?
• So, both the 2s and 2p orbitals are used to form bonds
• How many bonds does carbon form?
• All four C-H bonds are the same
• i.e. there are not two types of bonds from the two different orbitals
• How do we explain this?
Hybridization
Let’s consider carbon…
Hybridization
• The s and p orbitals of the C atom combine with each
other to form hybrid orbitals before they combine with
orbitals of another atom to form a covalent bond
• Three types we will consider:
• sp3
• sp2
• sp
sp3 hybridization
• 4 atomic orbitals → 4 equivalent hybrid orbitals
• s + px + py + pz → 4 sppp → 4 sp3
• Orbitals have two lobes (unsymmetrical)
• Orbitals arrange in space with larger lobes away from one
another (tetrahedral shape)
• Each hybrid orbital holds 2e-
• The sp3 hybrid orbitals
on C overlap with 1s
orbitals on 4 H atoms to
form four identical C-H
bonds
• Each C–H bond
strength = 439 kJ/mol;
length = 109 pm
• Each H–C–H bond
angle is 109.5°, the
tetrahedral angle
Formation of methane
Motivation for hybridization?
• Better orbital overlap with larger lobe of sp3 hybrid orbital
then with unhybridized p orbital
• Stronger bond
• Electron pairs farther apart in hybrid orbitals
• Lower energy
Another example: ethane
• C atoms bond by overlap of an sp3 orbital from each C
• Three sp3 orbitals on each C overlap with H 1s orbitals
• Form six C–H bonds
• All bond angles of ethane are tetrahedral
• Both methane and ethane have only single bonds
• Sigma (s) bonds
• Electron density centered between nuclei
• Most common type of bond
• Pi (p) bonds
• Electron density above and below nuclei
• Associated with multiple bonds
• Overlap between two p orbitals
• C atoms are sp2 or sp hybridized
Bond rotation
• Single (s) bonds freely rotate
• Multiple (p) bonds are rigid
sp2 hybridization
• 4 atomic orbitals → 3 equivalent hybrid orbitals
+ 1 unhybridized p orbital
• s + px + py + pz → 3 spp + 1 p = 3 sp2 + 1 p
• Shape = trigonal planar (bond angle = 120º)
• Remaining p orbital is perpendicular to hybrid orbitals
Formation of ethylene (C2H4)
• Two sp2-hybridized orbitals overlap to form a C─C s bond
• Two sp2 orbitals on each C overlap with H 1s orbitals (4 C ─ H)
• p orbitals overlap side-to-side to form a p bond
• s bond and p bond result in sharing four electrons (C=C)
• Shorter and stronger than single bond in ethane
• 4 atomic orbitals → 2 equivalent hybrid orbitals
+ 2 unhybridized p orbitals
• s + px + py + pz → 2 sp + 2 p
• Shape = linear (bond angle = 180º)
• Remaining p orbitals are perpendicular on y-axis and z-axis
sp hybridization
Formation of acetylene (C2H2)
• Two sp-hybridized orbitals overlap to form a s bond
• One sp orbital on each C overlap with H 1s orbitals (2 C─H)
• p orbitals overlap side-to-side to form two p bonds
• s bond and two p bonds result in sharing six electrons (C≡C)
• Shorter and stronger than double bond in ethylene
Summary of Hybridization
Hybridization of C sp3 sp2 sp
Example Methane, ethane Ethylene Acetylene
# Groups bonded to C 4 3 2
Arrangement of groups Tetrahedral Trigonal planar Linear
Bond angles ~109.5 ~120 ~180
Types of bonds to C 4s 3s, 1p 2s, 2p
C-C bond length (pm) 154 134 120
C-C bond strength (kcal/mol) 90 174 231
Hybridization of Heteroatoms
• Look at number of e- groups to determine hybridization
• Each lone pair will occupy a hybrid orbital
• Ammonia:
• N’s orbitals (sppp) hybridize to form four sp3 orbitals
• One sp3 orbital is occupied by the lone pair
• Three sp3 orbitals form bonds to H
• H–N–H bond angle is 107.3°
• Water
• The oxygen atom is sp3-hybridized
• The H–O–H bond angle is 104.5°
Example
• Consider the structure of
thalidomide and answer
the following questions:
a) What is the hybridization of each oxygen atom?
b) What is the hybridization of each nitrogen atom?
c) How many sp-hybridized carbons are in the molecule?
d) How many sp2-hybridized carbons are in the molecule?
e) How many sp3-hybridized carbons are in the molecule?
f) How many p bonds are in the molecule?
Example
• Consider the structure of 1-butene:
a) Predict each C─C─C bond angle in 1-butene.
b) Which carbon-carbon bond is shortest?
c) Draw an alkene that is a constitutional isomer of 1-butene.
Resonance
• Multiple Lewis structures for one molecule
• Differ only in arrangement of electrons
• Example: CH2NH2+ ion
• These are resonance structures/forms
• Valid Lewis structures (obey Octet Rule, etc.)
• Same number of electrons in each structure
• Atoms do not move
• Differ only in arrangement of electrons (lone pair and p electrons)
Resonance
Hybrid
• These structures imply that the C─N bond length and
formal charges are different
• Actually not true; these structures are imaginary
• Molecule is actually one single structure that combines all
resonance forms
• Resonance hybrid
• Contains characteristics of each resonance form
• More accurate and more stable than any single resonance form
• Lower energy (more stable) because of charge delocalization
Electron Movement
• Electrons move as pairs
• Can move from an atom to an adjacent bond, or from
bonds to adjacent atoms or bonds
• Use curved arrows to show e- motion (electron pushing)
• Start where electrons are, end where electrons are going
• Connect resonance forms with resonance arrow
• This is not an equilibrium arrow
Contribution to Hybrid Structure
• Resonance forms do not necessarily contribute equally to
the resonance hybrid
• They are not necessarily energetically equivalent
• More stable structures contribute more
1. Filled valence shells
2. More covalent bonds (minimizes charges)
3. Least separation of unlike charges (if applicable)
4. Negative charge on more EN atom (if applicable)
• Which of these is the major contributor to the resonance
hybrid?
Benzene
• Resonance structures:
• Curved arrows?
• Is one structure more stable (contribute more)?
• Resonance hybrid:
• All carbon-carbon bonds are the same length
• Somewhere between C─C and C=C
Acetone
• Resonance structures:
• Curved arrows?
• Which structure is the major contributor?
• Which is(are) the minor contributor(s)?
• Are any structures not likely to form?
• Resonance hybrid:
Drawing Resonance Structures
• Rules:
• Never break a single bond
• Only lone pair or p-electrons can move
• Never exceed an octet for C, O, N, X (or 2e- for H)
• Patterns:
1. p bonds
Drawing Resonance Structures
• Patterns:
2. Allylic charges or lone pairs
Electrons move towards positive charge!
Drawing Resonance Structures
Electrons move away from negative charge!
• Patterns:
2. Allylic charges or lone pairs
Drawing Resonance Structures
• Patterns:
3. Lone pair next to positive charge
Summary of Resonance Structures
Examples
O
CH2 CH CH CH3
CH2 CH CH CH3
Examples
O
H
NH2
H Br
Acids and Bases
• Two types in organic chemistry
1. Brønsted-Lowry
• Acid = proton (H+) donor; base = proton acceptor
• Some molecules can be both (e.g. water) = amphoteric
• Reaction will proceed from stronger acid/base to weaker acid/base
• Acid strength measured by pKa
• Stronger acid = lower pKa
Brønsted-Lowry Acids and Bases
• Electron flow in acid-base (proton-transfer) reactions:
• The reaction “mechanism”
• Example:
Brønsted-Lowry Acids and Bases
• You can predict acid strength without a pKa value
• Strong acids have weak conjugate bases
• Weak conjugate bases are stable structures
Brønsted-Lowry Acids and Bases
• Weak conjugate bases are stable structures
• Have negative charge on EN atom (within a period)
• Have negative charge on a larger atom (within a group)
• Negative charge delocalized by resonance
• Negative charge stabilized by induction
Brønsted-Lowry Acids and Bases
• Weak conjugate bases are stable structures
• Negative charge on sp > sp2 > sp3
Example
• Which is the stronger acid in each pair?
a) H2O or NH3?
b) HBr or HCl?
c) CH3OH or CH3CO2H?
d) CH3CO2H or Cl3CCO2H?
• Circle the most acidic H atom in this molecule:
Acids and Bases
2. Lewis
• Acid = electron pair acceptor, “electrophile”
• Base = electron pair donor, “nucleophile”
• Lewis acid react with Lewis base form a new covalent bond
Lewis Acids
• Incomplete octet (e.g. CR3+, BX3), or
• Polar bond to H (e.g. HCl), or
• Carbon with d+ due to polar bond (e.g. CH3Cl)
Lewis Bases
• Nonbonded electron pair (anything with O, N, anions)
Lewis Bases
• If there is more than one possible reaction site (more than
one atom with a lone pair), reaction occurs so that the
more stable product is formed.
• Example: Which oxygen is protonated when acetic acid
reacts with sulfuric acid?
Molecular Model Kits
• How to use
• Make a model for ethane
• Make a model for butane
• Make a model for cyclohexane
• Use 6 white hydrogens and 6 green hydrogens
• Put 1 green and 1 white hydrogen on each carbon atom
• The green and white hydrogen atoms should alternate (so as you
look at the molecule from the top the H’s should alternate green-
white-green-white-green-white around the ring)
Alkane Three-dimensional Structure
• Methane:
• With 2 or more carbons, 3D arrangement can change due
to C─C bond rotation
• Conformations
• Same molecular formula
• Same atom connectivity
• Different 3D arrangement due to rotation around single bond
• Ethane:
Newman Projections
• Used to better visualize conformations
• View the C─C from the end (look down the C─C bond)
• Represent the C atoms as a dot (front carbon) and circle
(back carbon)
• Show bonds coming out of the circle and dot
• Example:
Ethane Conformations
• Staggered vs. eclipsed
• Staggered is more stable (lower E) due to maximum
separation of electron pairs in covalent bonds
• Eclipsed is less stable (higher E) due to electron repulsions
Dihedral Angle• The degree of rotation between C-H bonds on the front
and back carbons
• Torsional strain• Accounts for energy difference between eclipsed and staggered
• Barrier to rotation
• Caused by electron repulsion
• Overcome by collisions of molecules
Butane Conformations
• Look down C2─C3 bond to draw Newman projections
• Each C has 2 H atoms and 1 CH3 group
• Dihedral angle is angle between CH3 groups
• There are six conformations of butane:
• How many staggered conformations? How many eclipsed?
Strain in Butane Conformations
• Torsional strain
• Barrier to rotation
• Example: eclipsed vs. staggered conformations
• Steric strain
• Repulsive interaction when atoms are forced close together (occupy
the same space)
• Example: CH3-H eclipsed vs. CH3-CH3 eclipsed conformations
• Example: Anti vs. gauche conformations
• So, which conformation is lowest in E? Highest in E?
• What would the plot of energy vs. dihedral angle look like?
Butane Conformations
Interpreting Newman Projections
• Which of the following Newman projections does NOT
represent 2-methylhexane?
Cycloalkane Three-dimensional Structure
• C atoms in cycloalkanes are sp3
• Bond angles are not always 109.5º
• Bond angles are dictated by the number of atoms in the ring
• Angle strain = Forcing angles smaller or larger than 109.5º
• Cycloalkanes can also have torsional strain (eclipsed H’s)
Strain in Cycloalkanes
Cycloalkane Conformations
• Cycloalkanes adopt more stable conformations to relieve
strain
• Cyclopropane
• “Bent” bonds
Cycloalkane Conformations
• Cyclobutane
• Puckered conformation
• Cyclopentane
• Envelope conformation
Cyclohexane
• Most stable cycloalkane
• Most abundant in nature
• No angle strain (109.5º)
• No torsional strain (all H’s staggered)
• Conformation = chair
• Drawing chairs: also see Klein p. 171
Cyclohexane
• Axial and equatorial hydrogens
• Axial = parallel to axis through ring
• Equatorial = perpendicular to axis
• Each C has one axial H and one equatorial H
• Look at molecular model
Cyclohexane
Ring Flip
• Interconversion of two chair conformations
• Try this with your molecular model
• If no substituents, these conformations are equal in energy
Monosubstituted Cyclohexanes
• Two conformations
1. Substituent in axial position
2. Substituent in equatorial position
• These conformations are not equal in energy
• Example: methylcyclohexane
Steric strain =
1,3-diaxial interactions
(“1,3” refers to distance
between groups)
Larger groups have
more steric strain
Disubstituted Cyclohexanes
• The most stable conformation has the most substituents
in the equatorial position
• Conformational analysis
• Look at all chair conformations (cis and trans) and analyze stability
• Example: 1,4-dimethylcyclohexane
Additional Cyclohexane Conformations
• Boat
• No angle strain
• High torsional strain
• High steric strain
• Very unstable
• Twist-boat
• Relieves some torsional and steric strain
• No angle strain
• Lower E than boat
• Higher E than chair
Energy Diagram for Cyclohexane
Conformations
Conformations of Polycyclic Molecules
• Fused rings
• Typically adopt chair conformations
• Norbornane and derivatives locked in boat conformation
Example
• Draw the most stable chair conformation for the following
molecules:
• trans-1,2-dimethylcyclohexane
• trans-1-isopropyl-3-methylcyclohexane
Degree of Unsaturation
• Unsaturated compounds
• Have less than (2n+2) H atoms for (n) C atoms
• Contain elements of unsaturation
• p bonds
• Rings
• Calculating degree of unsaturation
• Index of Hydrogen Deficiency (IHD)
• IHD = C - ½ (H + X) + ½ (N) + 1
• Ex: C6H14 IHD = 6 - ½(14) + 1 = 0 Alkane
• Ex: C6H12 IHD = 6 - ½(12) + 1 = 1 1 p bond or 1 ring
• Ex: C6H10 IHD = 6 - ½(10) + 1 = 2 2 p bonds, 2 rings, or 1 of each
C6H14 C6H12 C6H10
Alkene Stability
• Which alkene is more stable, cis or trans?
• Cis has steric strain between R groups
Alkene Stability
• Stability determined by heats of hydrogenation or
combustion
• Heat of hydrogenation = heat of reaction for addition of H2 (with
metal catalyst) to alkene
• Heat of reaction is proportional to energy of alkene
• Smaller magnitude DH = more stable alkene
Alkene Stability
• Heat of combustion= heat of reaction for combustion of alkene to
CO2 and H2O
• Heat of reaction is proportional to energy of alkene
• Smaller magnitude DH = more stable alkene
Alkene Stability
• Trends in alkene stability
• Trans is more stable than cis
• More substituted C=C is more stable
• Why?
• Hyperconjugation
• Stabilizing effect of adjacent orbital overlap
• Bond strengths
• sp2-sp3 bond more stable than sp3-sp3
Example
• There are four stereoisomers for 2,5-octadiene. Draw all
of the stereoisomers and circle the structure that is the
most stable.