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Organic Chemistry
PRINCIPLES AND MECHANISMS
Chapter 1: Atomic and Molecular Structure
1.1 What is Organic Chemistry?
• Organic chemistry is the branch of chemistry involving
organic compounds
• In the late 1700s organic compounds were defined as
those compounds obtained from a living organism.
• Inorganic compounds were defined as those that come
from everything else.
• The definition for organic chemistry was redefined in
1828 when Friedrich Wöhler synthesized urea (an organic
compound from inorganic compounds).
Organic Compound Defined
• An organic compound is composed primarily of carbon
and hydrogen.
• This definition is not perfect since some compounds, like
CO2, are also considered inorganic.
Organic Compound Defined
continued…
1.2 Why Carbon?
• The carbon atom is capable of having four covalent
bonds to other atoms
• Carbon atoms can link together in chains of almost any
length (branched chains as well as straight chain)
• To date, there are tens of millions of organic compounds
known an infinite number are possible
1.2 Why Carbon?
continued…
Silicon Won’t Work
• Silicon is found just under carbon on the Periodic Table.
• Silicon can also form four covalent bonds.
• C–C bond (339 kJ/mol) is stronger than Si-Si bond (188
kJ/mol); less energy is needed to break the Si-Si bond
• The strength of typical bonds involving carbon atoms
helps biomolecules remain intact—this is essential for
molecules to store information or provide cellular
structure.
• Electronegativity is a better measure of chemistry.
Silicon is more metallic (electropositive) than carbon
Phosphorus is Better
• Phosphorus is actually a closer
analog given its electronegativity.
• It is a larger and more polarizable
element, and therefore oxidizes
more readily limiting its ability to act
like carbon.
• It is an important element in the
burgeoning field of organic light
emitting diodes (OLEDs)
1.3 Atomic Structure and Ground State
Electron Configurations
• At the center of an atom there is a positively charged
nucleus, composed of protons and neutrons.
1.3 Atomic Structure and Ground State
Electron Configurations
continued…
• Surrounding the nucleus is a cloud of negatively charged
electrons, attracted to the nucleus by simple
electrostatic forces (the forces by which opposite
charges attract one another and like charges repel one
another).
• The space the electron cloud occupies is far larger than
the volume occupied by the nucleus.
The Structure of an Atom
• An atom has no net charge.
• The number of electrons equals the number of protons.
• The number of protons in the nucleus is known as the
atomic number (Z).
• Protons have a +1 charge, electrons have a -1 charge, and
neutrons have no charge.
• If the number of protons and electrons are unequal, then
the entire species bears a charge.
– Cation: More protons than electrons.
– Anion: More electrons than protons.
Atomic Orbitals and Shells
• Electrons in an isolated atom reside in atomic orbitals.
• The exact location of an electron can never be
pinpointed.
• The orbital specifies the region of space where the
probability of finding a given electron is high.
Atomic Orbitals and Shells
continued…
1. Atomic orbitals have different shapes.
2. Atomic orbitals are organized in
shells which are defined by the
principal quantum number, n.
3. Up to two electrons are allowed
in any orbital.
4. With increasing shell number, the size
and energy of the atomic orbital increases.
5. Within a given shell, an atomic orbital’s energy increases in
the following order: s , p , d, etc. In the second shell, for
example, the 2s orbital is lower in energy than the 2p.
Ground State Electron Configurations:
Valence Electrons and Core Electrons
• The way in which electrons are arranged in atomic
orbitals is called the atom’s electron configuration.
• The most stable (i.e., the lowest energy) electron
configuration is called the ground state configuration.
• Electrons occupying the outermost shell (i.e., highest
energy) are valence electrons. The rest are core
electrons.
The Three Rules to Obtaining an Atom’s
Ground State Electron Configuration
1. Pauli’s exclusion principle: No more than two electrons
(i.e., zero, one, or two electrons) can occupy a single
orbital; two electrons in the same orbital must have
opposite spins.
2. Aufbau principle: Each successive electron must fill the
lowest-energy orbital available.
3. Hund’s rule: All orbitals at the same energy must
contain a single electron before a second electron can
be paired in the same orbital.
The Ground State Electron Configuration
For The First 28 Electrons
1.4 The Covalent Bond: Bond Energy
and Bond Length
• A covalent bond is characterized by the sharing of
valence electrons between two or more atoms, as shown
for two H atoms
1.4 The Covalent Bond: Bond Energy
and Bond Length
continued…
• When two H atoms separated by a large distance are
brought together, their total energy begins to
decrease.
• At one particular internuclear distance, the energy of
the molecule is at a minimum.
1.4 The Covalent Bond: Bond Energy and
Bond Length
• Single bonds are the most common type of bond found
in organic molecules.
• The main difference among single, double, and triple
bonds is the number of electrons involved.
• In a single bond, two electrons are shared between
different nuclei; in a double bond, four electrons are
shared; and in a triple bond, six electrons are shared.
• As the number of bonds
increases between a pair
of atoms, bond strength
increases and bond length
decreases.
Average Bond Energies of
Common Single Bonds
Average Bond Energies and Bond Lengths
of Single and Multiple Bonds
• Multiple bonds are shorter and stronger than single
bonds.
Steps to Drawing a Lewis Structure
1. Count the total number of valence electrons in the molecule.
a.
b.
The number of electrons contributed by each atom is the
same as its group number (H = 1, C = 4, N = 5, O = 6, F = 7).
Each negative charge increases the number of electrons by
one; each positive charge decreases the number of electrons
by one.
2. Write the skeleton of the molecule, showing only the atoms
and the single bonds required to hold them together.
a.
If molecular connectivity is not given to you, the central atom
(the one with the greatest number of bonds) is usually the
one with the smallest electronegativity.
Steps to Drawing a Lewis Structure
continued…
3. Subtract two electrons from the total in Step 1 for each
single covalent bond drawn in Step 2.
4. Distribute the remaining electrons as lone pairs.
a.
b.
Start with the outer atoms and work inward.
Try to achieve a filled valence shell on each atom—namely, an
octet on each atom other than hydrogen, and a duet on each
hydrogen atom.
5. If there is an atom with less than a filled valence shell,
convert lone pairs from neighboring atoms into bonding
pairs of electrons, thereby creating double and/or triple
bonds.
Covalent Bonding: Sharing Electrons to
Produce Full Valence Shells
Incomplete Valence Shells and the
Expanded Octet
• Not all atoms have a
complete valence shell.
• In borane (BH3), the B
atom has a share of only
6 valence electrons
(lacks an octet).
Incomplete Valence Shells and
the Expanded Octet
continued…
• There are not enough
valence electrons available
to achieve a complete
valence shell for all atoms
in the molecule.
• Thionyl chloride (SOCl2),
the S atom has a share of
10 electrons.
1.6 Strategies for Success:
Drawing Lewis Dot Structures Quickly
• Drawing Lewis structures can become easier and faster
when you observe patterns in the number of bonds for
certain atoms.
• Each type of atom tends to form a specific number of
bonds and to have a specific number of lone pairs of
electrons.
1.6 Strategies for Success:
Drawing Lewis Dot Structures Quickly
continued…
1.7 Electronegativity, Polar Covalent
Bonds, and Bond Dipoles
• Covalent bonds are characterized by the sharing of
electrons between two atomic nuclei.
• If the atoms are identical, the electrons are shared
equally.
• If the atoms are different, than one nucleus will attract
electrons more strongly than the other.
• That ability to attract electrons in a covalent bond is
governed by the element’s electronegativity (EN).
• The most well-known EN scale was developed by Linus
Pauling, where atoms range from 0 to 4 on
electronegativity.
Electronegativity Trends
• Within the same row, electronegativity values tend to
increase from left to right across the periodic table.
• Within the same column, they tend to increase from
bottom to top.
• Elements with the largest electronegativities tend to be
in the upper right corner of the periodic table (e.g., N, O,
F, Cl, and Br).
– These do not include noble gases
• Elements with the smallest electronegativities tend to be
in the lower left corner (e.g., Rb, Cs, Fr, Ba, and Ra).
Electronegativity Trends
continued…
The Bond Dipole
• In a covalent bond, electrons are likely to be found near
the nucleus of the more electronegative atom and less
likely near the nucleus of the less electronegative atom.
• This creates a separation of partial positive and negative
charges along the bond, called a bond dipole.
The Bond Dipole
continued…
Electrostatic Potential Maps
• Another useful way to illustrate the distribution of charge
along a covalent bond is with an electrostatic potential
map.
• An electrostatic potential map shows a molecule’s
electron cloud in colors that indicate its relative charge.
Electrostatic Potential Maps
continued…
• Red corresponds to a buildup of negative charge,
whereas blue represents a buildup of positive charge.
1.8 Ionic Bonds
• When elements in a compound have large enough differences
in electronegativity, ionic bonding can occur.
• Rather than sharing electrons, the more electronegative atom
acquires electrons given up by the less electronegative atom,
forming oppositely charged ions.
• The electrostatic attraction between the
positively charged cations and the
negatively charged anions constitutes
the ionic bond.
• Sodium chloride (NaCl), for example,
consists of sodium cations (Na+) and
chloride anions (Cl-).
Polyatomic Ions
• Polyatomics can include the hydroxide (OH⁻), methoxide
(CH3O⁻), and methylammonium (CH3NH3+) ion.
• Usually consist only of nonmetals, and their atoms are
held together by covalent bonds.
1.9 Assigning Electrons to Atoms in
Molecules: Formal Charge and Oxidation State
• In an isolated atom or atomic ion, charge is determined by the
difference between the atom’s group number and the actual
number of valence electrons it possesses.
• To get charge, we assign electrons to atoms involved in
covalent bonds by two methods: formal charge or oxidation
state.
• Formal charge: In a given covalent bond, half the electrons
are assigned to each atom involved in the bond.
• Oxidation state: In a given covalent bond, all electrons are
assigned to the more electronegative atom. If the two atoms
are identical, the electrons are split evenly.
Obtaining Formal Charge and
Oxidation State
• Formal charges are used much more often than oxidation
states.
• Unless otherwise stated, you may assume that any
charges that appear in a Lewis structure are formal
charges.
Obtaining Formal Charge and Oxidation State
continued…
1.10 Resonance Theory
• Some species are not
described well by Lewis
structures.
• In HCO2⁻, both of the
C-O bonds are identical
experimentally; (i.e.,
same bond length and
bond strength),
intermediate between
those of a single and a
double bond.
Resonance Theory Rules
• Rule 1 Resonance occurs in species for which there are
two or more valid Lewis structures.
• Rule 2 Resonance
structures are
imaginary; the one,
true species is
represented by the
resonance hybrid.
Resonance Theory Rules
continued…
• Rule 3 The resonance hybrid looks most like the lowest
energy (most stable) resonance structure.
• The two resonance contributors of HCO2⁻ are equivalent.
• As a result, each structure contributes equally to the
resonance hybrid.
• In cases in which resonance structures are inequivalent,
their relative energies must be determined in order to
obtain their relative contributions to the hybrid.
Resonance Theory Rules
continued…
• A resonance structure is lower in energy (i.e., more stable)
with:
• a greater number of atoms with filled valence shells
• more covalent bonds
• fewer atoms with formal charges other than zero
• Rule 4
• Resonance provides stabilization.
• Resonance structures are stabilized as a result of the
delocalization of electrons (electrons have lower energy
when they are less confined).
• Delocalization is lower in energy than localization of
electrons.
Delocalization of Electrons
Resonance Theory Rules
continued…
• Rule 5 Resonance stabilization is usually large where
resonance structures are equivalent.
• If resonance structures are equivalent, then they will
contribute equally to the hybrid, allowing electrons the
greatest possible delocalization!
• Rule 6 All else being equal, the greater the number of
resonance structures, the greater the resonance
stabilization.
Resonance Theory Rules
continued…
1.11 Strategies for Success:
Drawing All Resonance Structures
• It is important to be able to draw all of a species’
resonance structures for two reasons:
 All resonance structures contribute to the features of the
resonance hybrid, and
 The total number of resonance structures is related to
the species’ stability.
1.11 Strategies for Success:
Drawing All Resonance Structures
continued…
• Different resonance structures can be obtained by
movement of electrons.
• We represent movement of electrons by use of curved
arrows.
Lone Pair Adjacent to a Multiple Bond
• The two curved arrows used to interconvert the
resonance structures of HCO2⁻ can be used whenever a
Lewis structure exhibits a lone pair of electrons on an
atom connected to a double bond or triple bond.
• The first curved arrow is used
to convert the lone pair into a
covalent bond.
• The second curved arrow is
used to convert a pair of
electrons from the original
multiple bond into a lone pair.
When Formal Charges
on an Atom Changes
• When a lone pair of electrons on an atom is converted
into a bonding pair, the formal charge of that atom
becomes more positive by 1.
• When a bonding pair involving an atom becomes a lone
pair on that atom, the formal charge of that atom
becomes more negative by 1.
An Incomplete Octet Adjacent to
a Multiple Bond
• Resonance does not always involve lone pairs.
• If an atom with an unfilled valence shell is adjacent to a
double bond or triple bond, then we can convert from
one resonance structure to another using a single
curved arrow.
An Incomplete Octet Adjacent to
a Multiple Bond
• When an atom lacking an octet gains a bond to achieve
an octet, the atom’s formal charge becomes more
negative by 1.
• When a bond is removed from an atom that initially has
an octet, the formal charge of the atom becomes more
positive by 1.
An Incomplete Octet Adjacent to
a Multiple Bond
continued…
A Lone Pair Adjacent to an Atom
with an Incomplete Octet
• An atom with a lone pair attached to an atom lacking an
octet can convert from one resonance structure to
another by involving just two atoms.
A Lone Pair Adjacent to an Atom
with an Incomplete Octet
continued…
A Ring of Alternating Single
and Multiple Bonds
• With a ring of alternating single and multiple bonds, a
pair of electrons from each multiple bond can shift
around the ring—either clockwise or counterclockwise—
to arrive at a new resonance structure
A Ring of Alternating Single and Multiple Bonds
continued…
• We must be careful. If only two or four electrons are
shifted, we end up with a Lewis structure that is not
valid.
1.12 Shorthand Notations
• Lone pairs of electrons are frequently omitted. Be
careful!
• You must have the knowledge of how formal charges
relates to the number of bonds and lone pairs!
1.12 Shorthand Notations
continued…
Shorthand Notations
continued…
• Condensed formulas allow us to include molecules and
molecular ions as part of regular text.
• Each nonhydrogen atom is written explicitly, followed
immediately by the number of hydrogen atoms that are
bonded to it.
• Adjacent nonhydrogen atoms in the condensed formula
are interpreted as being covalently bonded to each other.
Shorthand Notations
continued…
Shorthand Notations
continued…
• Line structures, like condensed formulas, are compact
and can be drawn quickly and easily.
Line Structure Example
• The line structure of CH3CH2CH2CH2CH2CH2CH2NH2 is
drawn below.
1.13 An Overview of Organic Compounds:
Functional Groups
• Specific arrangements of atoms connected by specific types of
bonds tend to react in characteristic ways.
• These structural components are called functional groups.
• A molecule’s reactivity is dictated by the functional groups it
possesses.
• Compounds consisting of only C-C and
C-H single bonds are called alkanes.
• There are several functional groups,
each having their own functional
group name.
Common Functional Groups
Common Functional Groups
Biomolecules, Fundamental Building
Blocks, and Functional Groups
• Organic molecules found exclusively in living organisms
are called biomolecules.
• There are four major classes of biomolecules: proteins,
carbohydrates, nucleic acids, and lipids.
Proteins and Amino Acids
• Proteins are highly versatile biomolecules.
• Actin and myosin are responsible for the mechanical
processes involving muscle tissue.
• Cortactin is responsible for regulating cell shape.
• There are proteins that are classified as enzymes, which
act as catalysts for biological reactions.
• Proteins are constructed from relatively few types of
small organic molecules called a-amino acids.
Proteins and Amino Acids
continued…
Proteins and Amino Acids
continued…
The 20 Naturally Occurring Amino Acids
The 20 Naturally Occurring Amino Acids
continued…
The 20 Naturally Occurring Amino Acids
continued…
The 20 Naturally Occurring Amino Acids
continued…
Carbohydrates and Monosaccharides
• Carbohydrates, also called saccharides, serve a variety of
biological functions:




fuels for metabolic pathways (ex. glycogen)
structure components of cell walls (ex. cellulose)
involved in blood clotting
and can help regulate the immune system
• Carbohydrates are characterized by their chemical
composition:
Carbohydrates and Monosaccharides
continued…
• Carbohydrates are frequently large molecules, called
polysaccharides (which are composed of smaller
molecules known as monosaccharides or simple sugars).
Carbohydrates and Monosaccharides
continued…
Nucleic Acids and Nucleotides
• A nucleic acid is a large molecular chain that is primarily
associated with the storage and transfer of genetic
information.
• A pair of intertwined nucleic acids form the doublehelical deoxyribonucleic acid (DNA), which stores genetic
information, and ribonucleic acid (RNA), which
participates in protein synthesis.
• Nucleic acids are constructed from relatively small
molecular units called nucleotides.
Nucleic Acids and Nucleotides
continued…
Nucleic Acids and Nucleotides
continued…
All nucleotides have three distinct components:
1. an inorganic phosphate (PO4) group,
2. a cyclic monosaccharide (or sugar), and
3. a nitrogenous base.
Nucleic Acids and Nucleotides
continued…
Nitrogenous Bases in RNA and DNA
RNA
• The 4 nitrogenous bases of RNA are uracil, guanine,
adenine, and cytosine (abbreviated U, G, A, and C,
respectively).
DNA
• The 4 nitrogenous bases of DNA are G, A, C, and thymine
(T).
• Three of the bases in DNA are the same as the bases in
RNA.
• Only T (in DNA) is different from U (in RNA).
Nitrogenous Bases in RNA and DNA
continued…
Summary and Conclusions
• Organic chemistry has been redefined to focus solely on
compounds containing carbon atoms and the reactions these
molecules undergo.
• Lewis structures, electronegativity, formal charges, and
resonance structures help us understand the structure and
reactivity of many organic compounds.
• Polar covalent bonds arise when atoms with moderate
differences in electronegativity are bonded together. Ionic
bonds arise when there are large differences in
electronegativity.
• Shorthand notation allows us to conveniently, accurately, and
efficiently draw the structures of simple and complex organic
compounds.
Summary and Conclusions
continued…
• Formal charges and oxidation states reflect different
ways in which valence electrons are assigned to
individual atoms within a molecule.
• Resonance exists when two or more valid Lewis
structures can be drawn for a given molecular species.
• Functional groups are the common bonding arrangement
of relatively few atoms, and these groups dictate the
behavior of the entire molecule.