Organic Chemistry
Vollhardt/Schore 6th Ed.
Chapter 1
Warning: this is not
a “powerpoint course” !!
Version 1.1:
Marc Anderson
Dept of Chemistry and Biochemistry
San Francisco State University
[marc@sfsu.edu]
Chapter 1: Structure and Bonding
in Organic Molecules
[1] introduction
[2] coulomb forces
[3] ionic and covalent bonds
[4] electron dot model of bonding
[5] resonance forms (!!!)
[6] atomic orbitals
[7] molecular orbitals (~) / covalent bonding
[8] hybrid orbitals / hybridization (!!!)
[9] structures, formulas, drawing molecules (!)
[1] Introduction
What do I enjoy about organic chemistry?
It provides a core understanding for a number of
fascinating and highly relevant areas of study:
• medicinal chemistry
• materials: dyes, paints, fibers, polymers
• biochemistry / pharmacology / neuroscience
• fuel chemistry
[2] Coulomb Forces
A simplified model of chemical bonding…
Bonds are made by simultaneous Coulombic attraction
and electron exchange.
When two atoms approach, the electrons of one (-) are
attracted by the protons of the other (+) and vice-versa.
This attraction is defined by Coulomb’s law:
(+) charge  (  ) charge
Attracting Force = constant 
distance
2
The attractive force releases a certain amount of energy
called the bond strength
The optimal distance between the two nuclei is called the
bond length.
There is an optimal balance between these two parameters
ionic and covalent bonds
Covalent Bonds are based on the sharing of electrons. If the
electrons are not shared equally, a polar covalent (partially
ionic) bond is formed, otherwise a pure covalent bond is
formed.
Ionic Bonds are based on the transfer of one or more electrons
from one atom to another. The resulting cation and anion are
electrostatically attracted to each other.
[3] Ionic and Covalent Bonds: The Octet Rule
The periodic table underlies the octet rule.
Electrons in atoms occupy levels or shells of fixed capacity.
The first shell has room for 2, the second 8, and the third 18.
Noble gases have 8 valence electrons (helium has 2) and are
particularly stable.
Other elements lack octets in their outer electron shells and
are driven to form molecules in such a way as to create
a stable octet (“nobel gas”) arrangement
The periodic table underlies the octet rule.
In pure ionic bonds, electron octets can be formed by transfer
of electrons.
Alkali metals react with halogens: one electron is transferred
and then both atoms achieve noble gas configuration
Formation of Ionic Bonds by Electron Transfer
Na 2,8,1 + Cl2,8,7  [Na 2,8 ]+ [Cl2,8,8 ] , or NaCl
Ionization Potential: Na = +119 kcal/mol (endothermic to lose e-)
Electron Affinity: Cl = -83 kcal/mol G (exothermic for Cl
to receive electron)
Electrostatic Attaction = -120 kcal/mol (exothermic attraction)
Eoverall = -84 kcal/mol (very exothermic!)
We can indicated the valence electron shell as “electron
dots”
This is convenient when illustrating electron-transfer:
Note: hydrogen can gain or lose an electron to generate
either “proton” or “hydride” (fundamentally different
species!)
covalency
In covalent bonds, electrons are shared to achieve octet config.
Ionic bonds between identical atoms of the same element do not form.
Instead, the electrons are shared and the bond is covalent!
Does Cl-Cl bond covalently or ionically?
Carbon does NOT form ionic bonds.
Sharing requires loss or gain of 4 e- to achieve noble gas config
Carbon typically achieves neon configuration and bonds covalently
Ionic and
Covalent Bonds:
Examples
In certain cases, one atoms supplies both of the electrons in the
bond:
4 electron (double) and 6 electron (triple) bonds are also formed:
polar covalent bonds
In most organic bonds, the electrons are not shared equally,
thus forming polar covalent bonds.
Pure covalent bonds (perfect sharing of e-) and ionic bonds
(complete transfer of e-) are two extremes.
Many bonds exist between extremes: they are polar covalent
Each element has an electronegativity value which represents
its e- accepting ability. (How strongly it pulls electrons
toward it)
Large difference in electronegativity between atoms will result
in a highly polar covalent bond.
polar covalent bonds
The separation of charges in polar covalent molecules
results in the formation of dipoles:
Symbolize this by an arrow pointing toward the more
electronegative atom (with a “cross” at the other atom):
Note: in symmetrical molecules such as CO2, the
individual dipoles cancel:
Btw: this picture is a computer-generated “electrostatic map”
Red indicates relative presence of electrons;
blue indicates relative absence of electrons.
geometry of atoms: VSEPR theory
Electron repulsion controls the shape (“structure”) of
molecules:
The shapes of molecules can be predicted using the VSEPR
method.
Electron pairs arrange themselves to be as far apart as possible
In the case of:
2 electron pairs: linear geometry
3 electrons pairs: trigonal geometry
4 electron pairs: tetrahedral geometry
… and for organic molecules:
Agenda
[1] introduction
[2] coulomb forces
[3] ionic and covalent bonds
[4] electron dot model of bonding
[5] resonance forms (!!!)
[6] atomic orbitals
[7] molecular orbitals (~) / covalent bonding
[8] hybrid orbitals / hybridization (!!!)
[9] structures, formulas, drawing molecules (!)
[4] Electron-dot model of bonding: Lewis Structures
Simple rules for Lewis Structures:
[1] Draw the molecular skeleton (connect the atoms)
[2] Depict all covalent bonds by a pair of shared electrons
[3] Give each atom a nobel gas configuration (usually a full octet)
[4] Electronegative elements may contain lone pairs of electrons
(N, O, Br, Cl).
More examples:
formal charges
Formal charges need to be determined for all atoms in a
molecule:
FC = (# valence electrons in free, neutral atom) –
- (# lone pair electrons on the atom)
– ½(# covalent bonding electrons surrounding atom)
let’s try these on the board …
We’ll see a simpler formula in a second…
It is useful to memorize the valence electron counts of
B, C, N, and O. (and remember that S is under O).
three main exceptions to the octet rule
[1] Atom has an odd # of electrons:
[2] Central atom has a free p orbital:
+CH ,
3
BeCl2, BH3
[3] Past row 2 of the periodic table, the central atom may be
surrounded by more than 8 electrons (expanded octet).
a simpler method of drawing
organic molecules: Kekulé structures
Covalent bonds can be depicted by straight lines.
Bonding pairs of electrons: represented as straight lines
Lone pairs of electrons: indicated with dots (sometimes omitted)
Structures of this type are called Kekulé structures.
formal charges with Kekulé structures is simplified
“FC = [# valence electrons] – dots – dashes”
(easier!)
Examples: you fill in the charge
DMSO
Nitromethane
(CH3NO2)
Some general rules: memorizing this will save you very
much time, especially on the exams!!!!
“carbocation”
“carbanion”
what is this molecule ???
[5] Resonance Hybrids (***)
The carbonate ion has several correct Lewis structures.
Three equivalent structures must be drawn to accurately represent
the carbonate ion.
The only difference: placement of electrons
Always draw arrows
from a negative
atom!
This will be
emphasized many
many more times…
But what is the true structure of carbonate?
…it can be thought of as of as the average of all three structures
which are each called resonance hybrids.
Each of the resonance structures contributes equally to the overall
structure of this molecule.
The 2 negative charges are
delocalized over all three oxygen
atoms
The resonance structures are
all present simultaneously;
i.e. resonance is not
“interconversion”
resonance hybrids: more examples
Note the direction of the arrows!
resonance hybrids: rules
Not all resonance structures are equal in occurrence.
There are rules that identify prevalence of specific forms:
Rule 1: Structures with a maximum of octets are most important.
Rule 2: Charges should be located on atoms with compatible
electronegativity.
Rule 3.
Separation of charge should be minimized
Example problem
resonance?
Example problem
For each of the following, are they correct or incorrect? why ?
Agenda
[1] introduction
[2] coulomb forces
[3] ionic and covalent bonds
[4] electron dot model of bonding
[5] resonance forms (!!!)
[6] atomic orbitals
[7] molecular orbitals (~) / covalent bonding
[8] hybrid orbitals / hybridization (!!!)
[9] structures, formulas, drawing molecules (!)
[6] Atomic orbital theory
… is a quantum mechanical description of electrons around the
nucleus …
The behavior of electron is described by wave equations.
An electron within an atom can have only certain definite energies
called energy states.
Moving particles such as electrons exhibit a wavelength
determined by the de Broglie relation:
λ=
h
mv
[ h is Plank’s constant, m is the mass of the
electron in kg, and v is the velocity of the electron
in m/s.]
This defines the wavelength of the electron.
(which is strange: we’re not used to a particle having a wavelength)
interference in electron waves.
The electron waves contain nodes, where the amplitude of the wave
changes sign.
The waves can interact with each other, producing either constructive or
destructive interference:
quantum mechanics overview (~)
The wave theory of electron motion is called quantum
mechanics.
The quantum mechanical equations describing the motion of
the electrons are called wave equations. The solutions of
these equations are called wave functions and are
represented by the Greek letter, . (psi)
The square of the wave function, evaluated at a point in
space (x,y,z) represents the probability of finding the
electron at that point at any given time.
Each wave function corresponds to a specific discrete
energy and the system is said to be quantized.
atomic orbitals have characteristic shapes !!
Knowing where electrons are tells us the geometry of the atom!!
(and where bonds form).
This then tells us the possible shape of molecules!
Most importantly: quantum mechanics provides a surface, which
indicates the probability of finding the electron at a given time.
The simplest example is an “s” orbital: (“spherical”)
The next highest energy orbital is the 2s orbital, which also
has a spherical surface:
Of slightly higher energy than 2s are three degenerate 2p
orbitals.
These are shaped like a“dumbell” and point along the three
Cartesian axes.
Organic chemistry deals mostly with the lower s and p orbitals.
Principles describing the asignment electrons to orbitals
[1] Lower energy orbitals are filled before those with higher energy.
(Aufbau principle).
[2] No orbital may be occupied by more than two electrons. (Pauli
Exclusion Principle).
[3] Degenerate orbitals must each receive a single electron of the
same spin before pairing of electrons occurs. (Hund’s rule)
Example: carbon
Examples: N, O, and F
(in their neutral forms: with 5, 6, or 7 valence electrons)
[7] Molecular Orbitals and Covalent Bonding
Atomic orbitals on two different atoms may overlap.
Overlap of two atomic orbitals leads to two molecular orbitals.
(one bonding and one antibonding).
Example: H2 (the simplest possible molecule):
An energy level diagram can now be made of the two
overlapping orbitals
Now we can use the Aufbau process to determine the
electronic configurations of: H2 and He2:
Agenda
[1] introduction
[2] coulomb forces
[3] ionic and covalent bonds
[4] electron dot model of bonding
[5] resonance forms (!!!)
[6] atomic orbitals
[7] molecular orbitals (~) / covalent bonding
[8] hybrid orbitals / hybridization (!!!)
[9] structures, formulas, drawing molecules (!)
[8] Hybridized orbitals: sp3
Orbital hybridization allows us to explain more complex
bonding arrangements that are typically observed
Example: sp3 hybridization explains the shape of tetrahedral carbon
compounds.
One 2s and three 2p orbitals combine to make four sp3 orbitals:
hybridized orbitals: sp2
One 2s and two 2p orbitals combine to make three sp2
orbitals:
one p orbital is left unoccupied! (not clearly shown above)
hybridized orbitals: sp
One 2s and one 2p orbital combine to make two sp orbitals:
(two p orbitals are left unoccupied)
water and ammonia are both sp3
Not all hybrid orbitals participate in bonding. Some contain lone pairs
of electrons.
The bond angles in ammonia are 107.3o and that in water is 104.5o, both close
to 109.5o
multiple bonds: ethene and ethyne
Double and triple bonded molecules utilize two types of
orbitals for bonding:
Single bond:
sp2 or sp hybridized orbital.
Double/triple bond: “extra” p orbital(s) (un-hybridized!)
[9] Structures and formulas in org chem
To establish the identity of a molecule, we determine its
structure.
Empirical formula = what elements present, in what ratios.
(determined via elemental analysis or mass spectrometry).
Substances having the same empirical formula but different
connectivity of atoms are called constitutional or structural
isomers.
[9] Structure identification: spectroscopy
To fully characterize the chemical structure of an organic
molecule we use techniques of spectroscopy (next semester).
1H
NMR:
There are two ways of representing
3D-structures: ball and stick
and space-filling models:
Ethanol (bp= 78.5o C)
Dimethyl Ether (bp= -23o C)
CONSTITUTIONAL ISOMERS: both are C2H6O
condensed and skeletal models
Kekulé
Condensed Skeletal
additional practice problems …
Hint: check your answers using:
Some practice problems.
1. Convert each skeletal structure to Kekule/line‐bond representation
2. Determine the molecular formula Hint: check your answers using:
Challenging practice problems.
Identify examples of sp3, sp2, and sp hybridized atoms, of these complicated
natural products. OH
OH
O
H
N
N
O
O
OH
H
HN
S
O
O
O
O
O
O
N
H
O
Latrunculol A
CH3
Yanucamide A
O