Advanced Organic Chemistry
HC1
Exam material:
Charge stability, orbitals and molecular shape
- Charge stability → how stable a molecule or intermediate is when it carries a positive or
negative charge, what influences charge stability:
• Electronegativity: More electronegative atoms stabilize negative charges better
• Resonance: If a charge can be delocalized over several atoms, it becomes more stable.
• Inducive effects: Nearby atoms can pull or push electron density through bonds
• Hybridization: Orbitals with more s-character stabilize negative charge better (sp > sp² >
sp³)
- Orbitals → describe where electrons are likely to be found
• The s orbitals is spherical, the p orbitals are dumbbell shaped (have 2 lobes)
• S and P orbitals combine to form hybrid orbitals (sp, sp², sp³)
▪ sp³ → tetrahedral (~109.5°) (sigma bonds)
▪ sp² → trigonal planar (~120°) (pi bonds)
o One C-C bond is the sigma bond, the other C-C bond results from the
overlap of the two unhybridized p orbitals
▪ sp → linear (180°) (pi bonds)
o one C-C bonds is the sigma bond, the other two C-C bonds are the pi
bonds
- How atoms form covalent bonds → As the two orbitals start to overlap to form the covalent
bond, energy is released (and stability increases) because the electron in each atom is
attracted both to its own nucleus and to the positively charged nucleus of the other atom
• At the bond length of the new covalent bond, there is maximum stability and minimum
potential energy
• Bond dissociation energy: the energy required to break the bond
Substitution Mechanisms (SN1 and SN2) → there are two different mechanisms by which a
substitution reaction can take place, in both mechanisms, the nucleophile replaces the leaving group
- SN2 (substitution nucleophilic bimolecular)
• One-step reaction
▪ The nucleophile attacks from the backside, simultaneously the LG leaves
• The rate depends on the concentration of both the alkyl halide and the nucleophile
• As the alkyl group becomes larger, the rate becomes slower (steric hindrance)
• Backside attack: the reaction requires the nucleophile to hit the carbon on the side
opposite the side that is bonded to the leaving group → you only get one product, the
inversion product
• Key features: favored by strong nucleophiles and primary carbons (C bonded to only
one other C)
, - SN1 (substitution nucleophilic unimolecular)
• Two-step reaction:
1) leaving group leaves & carbocation forms
2) nucleophile attacks the carbocation
• The rate depends only on the concentration of the alkyl halide
• Only tertiary alkyl halides undergo SN1 reactions with poor nucleophiles, such as water
and alcohols
• In step 2 the nucleophile can attack from the front or back → you get two products, one
with inversion and one with retention
• Key features: there is a carbocation intermediate, favored by tertiary carbons (C bonded
to three other C’s)
Transition state (and reaction coordinate) → A transition state is the highest-energy point during a
chemical reaction where bonds are partially broken and partially formed
- The activation energy = the energy needed to reach the transition state
- Reaction coordinate → shows the energy vs. reaction progress
Periodic table
,Lecture slides
When trying to understand a new reaction, apply electronegativity to understand electron densities
(delta plus or minus), not formal charge
- Electronegativity = a chemical property measuring the atom’s ability to attract and hold shared
electrons within a chemical bond
- If two atoms are equally electronegative, the bonding pair of electrons will be half way between
the atoms:
- If one atom (B) is more electronegative than atom A, the B end of the bond has more than its
fair share of electron density and so becomes slightly negative
Valence electrons (property of the atom)→ are electrons in the outermost shell of an atom, and that
can participate in the formation of a chemical bond
- You can determine the number of valence electrons from the group (column) in the
periodic table for main-group elements:
• For main group elements, groups 1-2→ valence electrons = group number
• For main group elements, groups 13-18 → valence electrons = group number - 10
The octet rule (electron arrangement in molecules)→ the tendency of atoms to prefer to have eight
electrons in the valence shell
- When atoms have fewer than eight electrons, they tend to react and form more stable
compounds
- The rule works for the main-group elements in the 2th period of the periodic table (e.g. C, N, O,
F)
- It does not work for e.g. Hydrogen and Helium, which only need
2 electrons in their valance shell for stability
Octet rule vs. charge → The octet rule only counts how many electrons are around the atom, not who
they “belong” to.
- Formal charge looks at electron ownership:
• If an atom is assigned fewer electrons than its normal valence electrons → positive
charge
• If it is assigned more electrons → negative charge
- So an atom can have 8 electrons around it (octet rule satisfied), but shill carry a charge
because of how the electrons are distributed in bonds
Reaction examples (when looking at electronegativity/octetrule/charge):
- Example with CH3O and H3O
• Oxygen has 6 valence electrons
• Oxygen 1- → octet rule, one bond too little
• Oxygen 1+ → octet rule, one bond too much
- How do they react
• Oxygen has a higher electronegativity compared to hydrogen
, • Bond between oxygen and hydrogen → electrons are not equally shared: oxygen is
pulling electrons stronger, has a higher electron density. This gives oxygen a δ minus,
hydrogen δ plus (polar bond)
• When oxygen has a positive charge → higher pulling of electrons
• The electrons of CH3O go to the δ plus on the hydrogen, the positive oxygen pulls the
electrons of the carbon
- Example with CH3OH and CH₃–C⁺(OH)–CH₃
• Oxygen again has a higher electronegativity compared to hydrogen, it pulls the
electrons from the double bond to itself
More examples about where the formal charge “lies” about electron density
- NH4 → nitrogen has a higher electronegativity compared to hydrogen, nitrogen is δ minus and
hydrogen δ plus
• Left: bases react with NH4 at H, not N
• Right: nucleophiles react with the ion at the carbon (δ plus), not N
Charged atoms can be stabilized/destabilized by neighboring atoms
- Positive charges are stabilized through donation of electron density by neighboring atoms
• Neighboring atoms or groups can donate electron density through bonds (inductive
effect or resonance), This reduces the electron deficiency of the positively charged
atom
- High charge density is unstable
• If the charge can spread out (delocalize) over multiple atoms or bonds, the charge
density decreases, making the molecule more stable
- Electron-withdrawing groups destabilize positive charge
• If a neighboring group removes electron density from a positively charged atom → this
destabilizes it (higher charge density)
Polarizability = how easily an atom’s electron cloud can distort or spread out.
- Smaller ions with tightly held electrons (like F⁻) → less polarizable
- Larger ions with more diffuse electrons (like I⁻) → more polarizable
Charge
1. The less charge the better → a molecule is more stable if it is unreactive, the more charged the
molecules are the less stable it is
HC1
Exam material:
Charge stability, orbitals and molecular shape
- Charge stability → how stable a molecule or intermediate is when it carries a positive or
negative charge, what influences charge stability:
• Electronegativity: More electronegative atoms stabilize negative charges better
• Resonance: If a charge can be delocalized over several atoms, it becomes more stable.
• Inducive effects: Nearby atoms can pull or push electron density through bonds
• Hybridization: Orbitals with more s-character stabilize negative charge better (sp > sp² >
sp³)
- Orbitals → describe where electrons are likely to be found
• The s orbitals is spherical, the p orbitals are dumbbell shaped (have 2 lobes)
• S and P orbitals combine to form hybrid orbitals (sp, sp², sp³)
▪ sp³ → tetrahedral (~109.5°) (sigma bonds)
▪ sp² → trigonal planar (~120°) (pi bonds)
o One C-C bond is the sigma bond, the other C-C bond results from the
overlap of the two unhybridized p orbitals
▪ sp → linear (180°) (pi bonds)
o one C-C bonds is the sigma bond, the other two C-C bonds are the pi
bonds
- How atoms form covalent bonds → As the two orbitals start to overlap to form the covalent
bond, energy is released (and stability increases) because the electron in each atom is
attracted both to its own nucleus and to the positively charged nucleus of the other atom
• At the bond length of the new covalent bond, there is maximum stability and minimum
potential energy
• Bond dissociation energy: the energy required to break the bond
Substitution Mechanisms (SN1 and SN2) → there are two different mechanisms by which a
substitution reaction can take place, in both mechanisms, the nucleophile replaces the leaving group
- SN2 (substitution nucleophilic bimolecular)
• One-step reaction
▪ The nucleophile attacks from the backside, simultaneously the LG leaves
• The rate depends on the concentration of both the alkyl halide and the nucleophile
• As the alkyl group becomes larger, the rate becomes slower (steric hindrance)
• Backside attack: the reaction requires the nucleophile to hit the carbon on the side
opposite the side that is bonded to the leaving group → you only get one product, the
inversion product
• Key features: favored by strong nucleophiles and primary carbons (C bonded to only
one other C)
, - SN1 (substitution nucleophilic unimolecular)
• Two-step reaction:
1) leaving group leaves & carbocation forms
2) nucleophile attacks the carbocation
• The rate depends only on the concentration of the alkyl halide
• Only tertiary alkyl halides undergo SN1 reactions with poor nucleophiles, such as water
and alcohols
• In step 2 the nucleophile can attack from the front or back → you get two products, one
with inversion and one with retention
• Key features: there is a carbocation intermediate, favored by tertiary carbons (C bonded
to three other C’s)
Transition state (and reaction coordinate) → A transition state is the highest-energy point during a
chemical reaction where bonds are partially broken and partially formed
- The activation energy = the energy needed to reach the transition state
- Reaction coordinate → shows the energy vs. reaction progress
Periodic table
,Lecture slides
When trying to understand a new reaction, apply electronegativity to understand electron densities
(delta plus or minus), not formal charge
- Electronegativity = a chemical property measuring the atom’s ability to attract and hold shared
electrons within a chemical bond
- If two atoms are equally electronegative, the bonding pair of electrons will be half way between
the atoms:
- If one atom (B) is more electronegative than atom A, the B end of the bond has more than its
fair share of electron density and so becomes slightly negative
Valence electrons (property of the atom)→ are electrons in the outermost shell of an atom, and that
can participate in the formation of a chemical bond
- You can determine the number of valence electrons from the group (column) in the
periodic table for main-group elements:
• For main group elements, groups 1-2→ valence electrons = group number
• For main group elements, groups 13-18 → valence electrons = group number - 10
The octet rule (electron arrangement in molecules)→ the tendency of atoms to prefer to have eight
electrons in the valence shell
- When atoms have fewer than eight electrons, they tend to react and form more stable
compounds
- The rule works for the main-group elements in the 2th period of the periodic table (e.g. C, N, O,
F)
- It does not work for e.g. Hydrogen and Helium, which only need
2 electrons in their valance shell for stability
Octet rule vs. charge → The octet rule only counts how many electrons are around the atom, not who
they “belong” to.
- Formal charge looks at electron ownership:
• If an atom is assigned fewer electrons than its normal valence electrons → positive
charge
• If it is assigned more electrons → negative charge
- So an atom can have 8 electrons around it (octet rule satisfied), but shill carry a charge
because of how the electrons are distributed in bonds
Reaction examples (when looking at electronegativity/octetrule/charge):
- Example with CH3O and H3O
• Oxygen has 6 valence electrons
• Oxygen 1- → octet rule, one bond too little
• Oxygen 1+ → octet rule, one bond too much
- How do they react
• Oxygen has a higher electronegativity compared to hydrogen
, • Bond between oxygen and hydrogen → electrons are not equally shared: oxygen is
pulling electrons stronger, has a higher electron density. This gives oxygen a δ minus,
hydrogen δ plus (polar bond)
• When oxygen has a positive charge → higher pulling of electrons
• The electrons of CH3O go to the δ plus on the hydrogen, the positive oxygen pulls the
electrons of the carbon
- Example with CH3OH and CH₃–C⁺(OH)–CH₃
• Oxygen again has a higher electronegativity compared to hydrogen, it pulls the
electrons from the double bond to itself
More examples about where the formal charge “lies” about electron density
- NH4 → nitrogen has a higher electronegativity compared to hydrogen, nitrogen is δ minus and
hydrogen δ plus
• Left: bases react with NH4 at H, not N
• Right: nucleophiles react with the ion at the carbon (δ plus), not N
Charged atoms can be stabilized/destabilized by neighboring atoms
- Positive charges are stabilized through donation of electron density by neighboring atoms
• Neighboring atoms or groups can donate electron density through bonds (inductive
effect or resonance), This reduces the electron deficiency of the positively charged
atom
- High charge density is unstable
• If the charge can spread out (delocalize) over multiple atoms or bonds, the charge
density decreases, making the molecule more stable
- Electron-withdrawing groups destabilize positive charge
• If a neighboring group removes electron density from a positively charged atom → this
destabilizes it (higher charge density)
Polarizability = how easily an atom’s electron cloud can distort or spread out.
- Smaller ions with tightly held electrons (like F⁻) → less polarizable
- Larger ions with more diffuse electrons (like I⁻) → more polarizable
Charge
1. The less charge the better → a molecule is more stable if it is unreactive, the more charged the
molecules are the less stable it is
