Q1 Essentials of Organic Chemistry Romy Kiffen 2021
Chapter 1
Particles and Symbols of the Atom
The number of protons in the nucleus is called
the atomic number (Z)
The s Atomic Orbitals
The s atomic orbitals begin at n = 1 and are
spherically shaped. There is a single orbital
within each ns subshell Within a particular shell,
s orbitals are lower in energy than other orbitals
As n increases, additional sign changes (nodes)
appear in the shape of the orbital
The p Atomic Orbitals
The p atomic orbitals begin at n = 2 and are
dumbbell shaped. There are 3 orbitals within each
np subshell, which correspond to the three
Cartesian directions. The np orbitals all have the
same energy; according to Hund’s rule electrons
are left unpaired when filling until they must be
paired. All p orbitals contain a node at the nucleus.
Filling up the orbitals
First 1s is filled up and then the 2s, Next 1 2px, 2py, 2pz, then the
second of those consecutively. After this 3s is filled up ect.
Ionic Bonding and Electron Transfer
One can imagine ionic bonds as arising from the transfer of an
electron from a metal to a nonmetal For example, gaseous
sodium can transfer an electron to gaseous chlorine.
• Gaseous Na and Cl spontaneously form solid NaCl
, Q1 Essentials of Organic Chemistry Romy Kiffen 2021
Covalent Bonding
A neutral hydrogen atom needs one more valence electron to achieve a
full valence shell Two hydrogens sharing their valence electrons achieve
a full n = 1 shell
In H2 , the electron configuration of each hydrogen atom is analogous to that of
helium, the first noble gas
A neutral fluorine atom needs one more valence electron to achieve a full
valence shell Two fluorines sharing their valence electrons achieve a full n
= 2 shell containing 8 electrons (is called an octet)
In F2 , the electron configuration of each hydrogen atom is analogous to
that of neon, the second noble gas
The six electrons on the periphery of each fluorine are nonbonding electron pairs or
lone pairs.
The octet rule: in stable molecules,
Lewis Symbols and Structures second-row atoms share electrons
In Lewis structures, electrons are represented as dots or lines. A line until they achieve an octet
denotes a pair of electrons (2) shared in a covalent bond; atomic
symbols denote atoms + core e− ’s
Not all electrons will be involved in covalent bonding; unshared pairs are drawn as dots on the
edges of their associated atoms
• Key premise: electrons are localized on a single atom or between two atoms
• Double bond: four electrons are shared (2 × 2)
• Triple bond: six electrons are shared (3 × 2)
Polar Covalent Bonds
Two atoms of different electronegativities share electrons
unequally in a covalent bond. The result is a polar
covalent bond The greater the difference in
electronegativity, the more polarized the bond
Polar covalent bonds tend to be sites of reactivity in
organic molecules.
Formal Charge
In this case, the atom has nonzero formal charge.
To calculate formal charge, subtract the valence
electron count of the neutral atom (VEC) from the
formal valence electron count of the covalently
bound atom (FEC)
To calculate the formal valence electron count of an atom, count all unshared electrons and half of
all bonding electrons Refer to the periodic table (group number) for the number of valence
electrons in the neutral atom
FC = FEC - VEC
,Q1 Essentials of Organic Chemistry Romy Kiffen 2021
Levels of Organic Structure
• The composition of a molecule is captured
by its molecular formula
• The constitution (connectivity) of a molecule
is captured by a structural formula showing
how atoms are connected
• The configuration of a molecule is captured
by a structural drawing showing the
positions of atoms in space
Compounds can have identical structures on a
broad level but differ on a more specific level; such
compounds are called isomers.
Resonance
Molecules containing delocalized electrons have multiple Lewis-structural
representations called resonance structures.
The compound does not convert between A and B. it is in reality a weighted sum of structures A
and B.
Rules of Resonance
• Rule 1. The connectivity and positions of the atoms must remain the same in all resonance
structures
• Rule 2. Each contributing structure must have the same number of electrons and the same
net charge
• Rule 3. Each contributing structure must have the same number
of unpaired electrons
• Rule 4. Atoms of second-row elements must not violate the octet
rule
Which resonance form is most important?
• Rule 5. Resonance structures containing more covalent bonds
are typically more important
• Rule 6. Resonance structures with minimal separation of
opposite charges are most important
• Rule 7. Resonance structures with negative charge on the most
electronegative atom and positive charge on the most
electropositive atom are most important
The Importance of Resonance
Recognizing resonance and drawing valid resonance structures are important skills for at least
two reasons:
1. Resonance indicates that molecules are stabilized due to the delocalization of charge.
2. Resonance forms containing formal charges can reveal hidden points of reactivity in
organic
A full-headed arrow implies movement of a pair
of electrons; arrow starts at the source and ends
at the sink.
, Q1 Essentials of Organic Chemistry Romy Kiffen 2021
Brønsted-Lowry Acids and Bases
Acid: a species that donates a proton to a base
Base: a species that accepts a proton from an acid
Species related by the addition or removal of a single proton are called conjugates
• When an acid surrenders a proton in an acid-base reaction, its conjugate base is formed
• When a base gains a proton in an acid-base reaction, its conjugate acid is formed
The Acid Dissociation Constant
Water reacts with an acid HA to form H3O+ and A−. The
equilibrium constant for this process is called the acid
dissociation constant (Ka)
The larger Ka is, the stronger the acid Ka for an acid is related to Kb for
the conjugate base of the acid
𝐾𝑎 𝐾𝑏 = 𝐾𝑤 = 1.0 ∙ 10−14
pKa : A Brain-friendly Acidity Constant
Ka values span a massive range and are usually far less than 1, so we usually work with
𝑝𝐾𝑎 = −log10 𝐾𝑎
Thanks to the negative sign, a smaller pKa is associated with a stronger acid
Structural Factors Influencing Acidity
There are four structural factors that exert a profound influence on the acidity of a proton
1. The strength of the bond to the atom from which the proton is lost
2. The electronegativity of the atom from which the proton is lost
3. Inductive effects of nearby atoms
4. Electron delocalization in the conjugate base
Factor 1: H–A Bond Strength
As the size of the atom to which H is bonded increases, the H–A
bond strength decreases and acidity increases. Greater
polarizability of larger conjugate base anions also contributes to
greater acidity. The hydrohalic acids illustrate this effect nicely.
Factor 2: Electronegativity of A in H–A
As the electronegativity of A increases, the polarization of the H–A
bond increases and acidity increases. The partial positive charge
on H increases, increasing the likelihood that it will be transferred
as H+ (i.e., HA will act as an acid). The covalent hydrides of
second-row nonmetals illustrate this effect nicely. Acidity increases
moving to the right in the periodic table.
Chapter 1
Particles and Symbols of the Atom
The number of protons in the nucleus is called
the atomic number (Z)
The s Atomic Orbitals
The s atomic orbitals begin at n = 1 and are
spherically shaped. There is a single orbital
within each ns subshell Within a particular shell,
s orbitals are lower in energy than other orbitals
As n increases, additional sign changes (nodes)
appear in the shape of the orbital
The p Atomic Orbitals
The p atomic orbitals begin at n = 2 and are
dumbbell shaped. There are 3 orbitals within each
np subshell, which correspond to the three
Cartesian directions. The np orbitals all have the
same energy; according to Hund’s rule electrons
are left unpaired when filling until they must be
paired. All p orbitals contain a node at the nucleus.
Filling up the orbitals
First 1s is filled up and then the 2s, Next 1 2px, 2py, 2pz, then the
second of those consecutively. After this 3s is filled up ect.
Ionic Bonding and Electron Transfer
One can imagine ionic bonds as arising from the transfer of an
electron from a metal to a nonmetal For example, gaseous
sodium can transfer an electron to gaseous chlorine.
• Gaseous Na and Cl spontaneously form solid NaCl
, Q1 Essentials of Organic Chemistry Romy Kiffen 2021
Covalent Bonding
A neutral hydrogen atom needs one more valence electron to achieve a
full valence shell Two hydrogens sharing their valence electrons achieve
a full n = 1 shell
In H2 , the electron configuration of each hydrogen atom is analogous to that of
helium, the first noble gas
A neutral fluorine atom needs one more valence electron to achieve a full
valence shell Two fluorines sharing their valence electrons achieve a full n
= 2 shell containing 8 electrons (is called an octet)
In F2 , the electron configuration of each hydrogen atom is analogous to
that of neon, the second noble gas
The six electrons on the periphery of each fluorine are nonbonding electron pairs or
lone pairs.
The octet rule: in stable molecules,
Lewis Symbols and Structures second-row atoms share electrons
In Lewis structures, electrons are represented as dots or lines. A line until they achieve an octet
denotes a pair of electrons (2) shared in a covalent bond; atomic
symbols denote atoms + core e− ’s
Not all electrons will be involved in covalent bonding; unshared pairs are drawn as dots on the
edges of their associated atoms
• Key premise: electrons are localized on a single atom or between two atoms
• Double bond: four electrons are shared (2 × 2)
• Triple bond: six electrons are shared (3 × 2)
Polar Covalent Bonds
Two atoms of different electronegativities share electrons
unequally in a covalent bond. The result is a polar
covalent bond The greater the difference in
electronegativity, the more polarized the bond
Polar covalent bonds tend to be sites of reactivity in
organic molecules.
Formal Charge
In this case, the atom has nonzero formal charge.
To calculate formal charge, subtract the valence
electron count of the neutral atom (VEC) from the
formal valence electron count of the covalently
bound atom (FEC)
To calculate the formal valence electron count of an atom, count all unshared electrons and half of
all bonding electrons Refer to the periodic table (group number) for the number of valence
electrons in the neutral atom
FC = FEC - VEC
,Q1 Essentials of Organic Chemistry Romy Kiffen 2021
Levels of Organic Structure
• The composition of a molecule is captured
by its molecular formula
• The constitution (connectivity) of a molecule
is captured by a structural formula showing
how atoms are connected
• The configuration of a molecule is captured
by a structural drawing showing the
positions of atoms in space
Compounds can have identical structures on a
broad level but differ on a more specific level; such
compounds are called isomers.
Resonance
Molecules containing delocalized electrons have multiple Lewis-structural
representations called resonance structures.
The compound does not convert between A and B. it is in reality a weighted sum of structures A
and B.
Rules of Resonance
• Rule 1. The connectivity and positions of the atoms must remain the same in all resonance
structures
• Rule 2. Each contributing structure must have the same number of electrons and the same
net charge
• Rule 3. Each contributing structure must have the same number
of unpaired electrons
• Rule 4. Atoms of second-row elements must not violate the octet
rule
Which resonance form is most important?
• Rule 5. Resonance structures containing more covalent bonds
are typically more important
• Rule 6. Resonance structures with minimal separation of
opposite charges are most important
• Rule 7. Resonance structures with negative charge on the most
electronegative atom and positive charge on the most
electropositive atom are most important
The Importance of Resonance
Recognizing resonance and drawing valid resonance structures are important skills for at least
two reasons:
1. Resonance indicates that molecules are stabilized due to the delocalization of charge.
2. Resonance forms containing formal charges can reveal hidden points of reactivity in
organic
A full-headed arrow implies movement of a pair
of electrons; arrow starts at the source and ends
at the sink.
, Q1 Essentials of Organic Chemistry Romy Kiffen 2021
Brønsted-Lowry Acids and Bases
Acid: a species that donates a proton to a base
Base: a species that accepts a proton from an acid
Species related by the addition or removal of a single proton are called conjugates
• When an acid surrenders a proton in an acid-base reaction, its conjugate base is formed
• When a base gains a proton in an acid-base reaction, its conjugate acid is formed
The Acid Dissociation Constant
Water reacts with an acid HA to form H3O+ and A−. The
equilibrium constant for this process is called the acid
dissociation constant (Ka)
The larger Ka is, the stronger the acid Ka for an acid is related to Kb for
the conjugate base of the acid
𝐾𝑎 𝐾𝑏 = 𝐾𝑤 = 1.0 ∙ 10−14
pKa : A Brain-friendly Acidity Constant
Ka values span a massive range and are usually far less than 1, so we usually work with
𝑝𝐾𝑎 = −log10 𝐾𝑎
Thanks to the negative sign, a smaller pKa is associated with a stronger acid
Structural Factors Influencing Acidity
There are four structural factors that exert a profound influence on the acidity of a proton
1. The strength of the bond to the atom from which the proton is lost
2. The electronegativity of the atom from which the proton is lost
3. Inductive effects of nearby atoms
4. Electron delocalization in the conjugate base
Factor 1: H–A Bond Strength
As the size of the atom to which H is bonded increases, the H–A
bond strength decreases and acidity increases. Greater
polarizability of larger conjugate base anions also contributes to
greater acidity. The hydrohalic acids illustrate this effect nicely.
Factor 2: Electronegativity of A in H–A
As the electronegativity of A increases, the polarization of the H–A
bond increases and acidity increases. The partial positive charge
on H increases, increasing the likelihood that it will be transferred
as H+ (i.e., HA will act as an acid). The covalent hydrides of
second-row nonmetals illustrate this effect nicely. Acidity increases
moving to the right in the periodic table.
