Chapter 1
Carbon Compounds and Chemical
Bonds
Organic Chemistry
Organic chemistry is the chemistry of compounds containing carbon.
The name mistakenly implies that all organic compounds are
associated with living systems. While many carbon- containing
compounds are the basis of life, there are many organic compounds
that have nothing to do with living systems.
The practice of organic chemistry goes back thousands of years. A
discovery of a wine residue in a Sumerian jug shows that the technology
of winemaking is almost 7500 years old.
sugar (C12H22O11)
sweet principle of
sugar cane
fermentation
ethyl alcohol (C2H5OH)
Whence Wine?
A recent article in the Chronicle of
Higher Eduction (August 15, 2003)
details the research of Patrick E.
McGovern, a chemistry professor at
the University of Pennsylvania.
Blending chemistry and archeology,
Prof. McGovern is an expert on the
origins of winemaking. His research
suggests that the art of fermenting
grapes began some 9,000 years ago
when human society was
transforming from hunter-gatherer
to communities based on agriculture
and animal husbandry.
The research of Prof. McGovern
and others points to the widespread
production of wine throughout the
middle east in the region of the
countries of Iran, Iraq and Turkey
during the Sumerian and later
civilizations. It is believed that the
domestication of grapes for
winemaking began somewhere in
this region, and that all modern
grapes for winemaking are
descended from that original
domesticated wild subspecies.
Demise of the Life Force Idea
Early in the 19th century, numerous "organic compounds" were
prepared from nonliving sources. The first such example was the
synthesis of urea (a constituent of urine) by evaporation of a solution of
ammonium cyanate, an "inorganic compound."
1828: Friedrich Wohler
evaporation
NH4OCN
ammonium cyanate
H2NCONH2
urea
Structural Isomerism
In addition to disproving vitalism, the Wohler experiment shows the
properties of molecules depend on the way atoms are linked. The
structure is important. This idea was further developed in the late 19th
century.
Empirical and Molecular Formulas
During the 19th century, quantitative methods for determining chemical
composition were developed by Justis Liebig, J.J. Berzelius, and J.B.
Dumas. The chemical analyses could only provide empirical formulas, the
simplest possible ratio of the elements.
In 1860, at the famous Karlsruhe Conference called by August Kekulé,
the Italian chemist Stanislao Cannizzaro used the earlier arguments of
Amedeo Avogadro (1811) to assign proper molecular formulas to many
simple organic compounds.
Cannizzaro showed that while the same empirical formula of CH2 is
found in ethylene, cyclopentane, and cyclohexane, their molecular
formulas are C2H4, C5H10, and C6H12, respectively.
• Structural Theory
• Central Premises
• Valency: atoms in organic compounds form a
fixed number of bonds
• Carbon can form one or more bonds to
other carbons
In all cases, the tetravalency of carbon is maintained.
7
• Isomers
• Isomers are different molecules with the same
molecular formula
• Many types of isomers exist
• Example
• Consider two compounds with molecular formula
C2H6O
• These compounds cannot be distinguished based on
molecular formula; however they have different
structures
• The two compounds differ in the connectivity of their
atoms
8
• Constitutional Isomers
• Constitutional isomers are one type of isomer
• They are different compounds that have the
same molecular formula but different
connectivity of atoms
• They often differ in physical properties (e.g.
boiling point, melting point, density) and
chemical properties
9
QUIZ Chapter 1 Sections 2 and 3
Assemble the follow ing atoms into plausible structures consistent
w ith the valencies of the atoms.
(A) C3H 8O
(B) C3H 6
H H
H H H
H C C C O H
H
H H H
H H H
H C C C H
H O H
H
H H
H
H
H
C
H
H
C
C
H
H
H
H C C O C H
H H
H C C C H
t w o const it ut ional isomers
H
t hree const it ut ional isomers
Comment on their physical properties
• Three Dimensional Shape of Molecules
• Virtually all molecules possess a 3-dimensional shape
which is often not accurately represented by drawings
• It was proposed in 1874 by van’t Hoff and le Bel that the
four bonds around carbon where not all in a plane but
rather in a tetrahedral arrangement i.e. the four C-H
bonds point towards the corners of a regular
tetrahedron
11
Chemical Bonds
The first descriptions of the nature of chemical bonds were put
forward in 1916 by Gilbert Newton Lewis (1875-1946) and Walther
Kossel (1888-1956). (Remember, the structure of the atom had been
deduced only a few years earlier by Ernest Rutherford and Niels
Bohr. )
Two types of chemical bonds were recognized:
(A) Ionic bonds are formed by the transfer of one or more
valence-level electrons from one atom to another.
-e +e
-
Li+
:F :
: :
: :
.Li
. F:
Li+Flithium fluoride
(B) Covalent bonds are formed when atoms share
valence-level electrons.
: :
: :
: :
: :
: F. + . F :
:F : F :
fluorine
Both types of bonds lead to electronic configurations around atoms
that mimic a noble gas--filled outermost electronic levels.
• Chemical Bonds: The Octet Rule
• Octet Rule
• Atoms form bonds to produce the electron
configuration of a noble gas (because the
electronic configuration of noble gases is
particularly stable)
• For most atoms of interest this means
achieving a valence shell configuration of 8
electrons corresponding to that of the
nearest noble gas
• Atoms close to helium achieve a valence
shell configuration of 2 electrons
• Atoms can form either ionic or covalent
bonds to satisfy the octet rule
13
Ionic Bonds
When atoms lose or gain electrons, they develop charges and
are called ions. Positively charged ions are cations, and
negatively charged ions are called anions.
Li +
.
lithium cation
Ffluoride anion
Electronegativity
Electronegativity measures the tendency of an atom to attract
electrons. Electronegativity increases in going left to right across a row
in the Periodic Table. In some groups, it decreases in going down a
column.
Row 2
Li Be B C N O F
increasing electronegativity
Ionic bonds are formed when atoms of
very different electronegativities interact.
Group 7
F
decreasing
Cl
electronegativity
Br
I
most electronegative
least elect ronegat iv e
Linus Pauling (1901-1994)
Nobel Prize in Chemistry 1954
Nobel Peace Prize 1962
Electronegativity is the ability of an atom to attract electrons
• Ionic Bonds
• When ionic bonds are formed atoms gain or lose
electrons to achieve the electronic configuration of the
nearest noble gas
• In the process the atoms become ionic
• The resulting oppositely charged ions attract and form
ionic bonds
• This generally happens between atoms of widely
different electronegativities
• Example
• Lithium loses an electron (to have the configuration
of helium) and becomes positively charged
• Fluoride gains an electron (to have the configuration
of neon) and becomes negatively charged
• The positively charged lithium and the negatively
charged fluoride form a strong ionic bond (actually
in a crystalline lattice)
16
• Covalent Bonds
• Covalent bonds occur between atoms of similar
electronegativity (close to each other in the periodic table)
• Atoms achieve octets by sharing of valence electrons
• Molecules result from this covalent bonding
• Valence electrons can be indicated by dots (electron-dot
formula or Lewis structures) but this is time-consuming
• The usual way to indicate the two electrons in a bond is to
use a line (one line = two electrons)
17
Multiple Covalent Bonds and the Octet Rule
Covalent bonds between atoms may involve 4 or 6 electrons.
These are called multiple covalent bonds.
Example N2
Each nitrogen atom has 5 outermost or valence-level electrons.
.
:N. .
The covalent bonding in N2 involves 6 electrons:
N
N
or
:N
N:
The Octet Rule
By sharing 6 electrons, each nitrogen achieves the electron
configuration of neon with an octet of electrons in the valence level.
This tendency to reach 8 electrons in the valence level is called the
octet rule.
Examples of the Octet Rule
N
N
nitrogen
: :
: :
H
H:C :H
H
methane
H:O:H
H:F:
water
hydrogen
fluoride
The above structures are called Lewis structures in honor of G.N.
Lewis. In a Lewis structure, all the valence-level (outermost)
electrons are shown as dots. The Lewis structure of an atom is the
chemical symbol with the valence-level electrons shown as dots.
Gilbert Newton Lewis (1875-1946)
A leading physical chemist at UC Berkeley
during the early decades of the 20th century.
Several simple rules allow us to draw proper Lewis Structures:
1. Lewis structures show the connections between atoms in a
molecule or ion using only the valence electrons or the atoms
involved. Valence electrons are those of an atom’s outermost shell.
2. For main group elements, the number of valence electrons a
neutral atom brings to a Lewis structure is the same as its group
number in the periodic table. Carbon, for example, is in Group
IVA and it has four valence electrons; the halogens (e.g., fluorine)
are in Group VIIA and each has seven valence electrons;
hydrogen is in Group IA and has one valence electron.
3. If the structure we are drawing is a negative ion (an anion), we
add one electron for each negative charge to the original count of
valence electrons. If the structure is a positive ion (a cation), we
subtract one electron for each positive charge.
Several simple rules allow us to draw proper Lewis Structures:
4. In drawing Lewis structures we try to give each atom the electron
configuration of a noble gas. To do so, we draw structures where
atoms share electrons to form covalent bonds or transfer electrons
to form ions.
a. Hydrogen forms one covalent bond by sharing its electron
with an electron of another atom so that it can have two
valence electrons, the same number as in the noble gas helium.
b. Carbon forms four covalent bonds by sharing its four
valence electrons with four valence electrons from other
atoms, so that it can have eight electrons (the same as the
electron configuration of neon, satisfying the octet rule).
c. To achieve an octet of valence electrons, elements such as
nitrogen, oxygen, and the halogens typically share only
some of their valence electrons through covalent bonding,
leaving others as unshared electron pairs.
• Writing Lewis Structures (short form)
• Atoms bond by using their valence electrons
• The number of valence electrons is equal to the group
number of the atom
• Carbon is in group 4A and has 4 valence electrons
• Hydrogen is in group 1A and has 1 valence electron
• Oxygen is in group 6A and has 6 valence electrons
• Nitrogen is in group 5A and has 5 valence electrons
• To construct molecules the atoms are assembled with
the correct number of valence electrons
• If the molecule is an ion, electrons are added or
subtracted to give it the proper charge
• The structure is written to satisfy the octet rule for
each atom and to give the correct charge
• If necessary, multiple bonds can be used to satisfy the
octet rule for each atom
23
• Example
• Write the Lewis structure for the chlorate ion (ClO3-)
• The total number of valence electrons including
the electron for the negative charge is calculated
• Three pairs of electrons are used to bond the
chlorine to the oxygens
• The remaining 20 electrons are added to give
each atom an octet
24
• The carbonate ion with 24 valence electrons
and two negative charges must incorporate a
double bond to satisfy the octet rule for
every atom
• The organic molecules ethene (C2H4) and
ethyne (C2H2) must also use multiple bonds
to satisfy the octet rule for each atom
25
• Exceptions to the Octet Rule
• The octet rule applies only to atoms in the
second row of the periodic table (C, O, N, F)
which are limited to valence electrons in the 2s
and 2p orbitals
• In higher rows other orbitals are accessible and
more than 8 electrons around an atom are
possible
• Example: PCl5 and SF6
26
Electron-Deficient Molecules
There are many exceptions to the octet rule among
compounds derived from early main group elements in
Period 2 (Be and B). The chemical reactivity of these
compounds reflects their electron-deficiency.
Examples
: :
: :
: :
: :
+ 2 :Cl .
:Cl :Be:Cl: or :Cl -Be-Cl:
only 4 valence electrons
around Be
: :
:F :B:F:
:F:
: :
+ 3 :F.
or :F -B-F:
:F:
:
: :
: :
: :
.
.B.
: :
BF3
boron
trifluoride
.Be.
: :
BeCl 2
berylium
chloride
only 6 valence electrons
around B
Electron-Deficient Molecules
Berylium chloride exists as a monomer only in the gas phase.
In the solid state, a polymeric structure forms:
Cl
Be
Cl
Cl
Be
Cl
Be
Cl
Cl
Be
there is an octet of electrons
in the valence-level around Be
Boron trifluoride reacts rapidly with chemical species
that donate an electron pair:
fast
: :
: : : :
: :
: :
-
BF3 + :F:
there is an octet
F
F:
:
of electrons in the
:F :B:F: or F B F
valence-level
F
:F:
around B
QUIZ Chapter 1 Sections 4, 5 and 6
Which electronic structure below is the correct Lewis structure of
H2CO, where the hydrogen atoms are bonded to the carbon atom?
solution: There are 12 valence electrons.
H.. .. .
.C
. ....O .
H
B
H.. .. ..
..C..O. .
H
C
Only C is consistent with the Octet Rule.
H..
.C
. :O:
H
D
: :
choices
H.. ....
.C
. ....O
H
A
• Formal charge
• A formal charge is a positive or negative charge
on an individual atom
• The sum of formal charges on individual atoms
is the total charge of the molecule or ion
• The formal charge is calculated by subtracting
the assigned electrons on the atom in the
molecule from the electrons in the neutral
atom
• Electrons in bonds are evenly split between
the two atoms; one to each atom
• Lone pair electrons belong to the atom itself
30
• Examples
• Ammonium ion (NH4)+
• Nitrate ion (NO3)-
31
Examples of Formal Charges
: :
: :
formal
charge
H
1
-
1
0
C
4
-
4
0
atom
: :
NH3
atom
: :
CH4
H
H C H
H
valence electrons in:
free atom bonded
state
HN H
H
valence electrons in:
free atom bonded
state
formal
charge
H
1
-
1
0
N
5
-
(3 + 2)
0
Examples of Formal Charges
atom
: :
: :
NH4
H
HN H
H
: :
::
: :
CO 32-
formal
charge
H
1
-
1
0
N
5
-
4
+1
atom
:O:
:O:C :O:
valence electrons in:
free atom bonded
state
valence electrons in:
free atom bonded
state
formal
charge
C
4
-
4
0
O
6
-
6
0
O
6
- (6 + 1)
-1
• A Summary of Formal Charges
34
• Resonance
• Often a single Lewis structure does not accurately
represent the true structure of a molecule
• The real carbonate ion is not represented by any
of the structures 1,2 or 3
• Experimentally carbonate is known not to have
two carbon-oxygen single bonds and one double
bond; all bonds are equal in length and the
charge is spread equally over all three oxygens
35
• The real carbonate ion can be represented by a
drawing in which partial double bonds to the
oxygens are shown and partial negative charge
exists on each oxygen
• The real structure is a resonance hybrid or mixture
of all three Lewis structures
• Double headed arrows are used to show that the
three Lewis structures are resonance contributors
to the true structure
• The use of equilibrium arrows is incorrect since
the three structures do not equilibrate; the true
structure is a hybrid (average) of all three Lewis
structures
36
• One resonance contributor is converted to another by the use
of curved arrows which show the movement of electrons
• The use of these arrows serves as a bookkeeping device
to assure all structures differ only in position of electrons
• A calculated electrostatic potential map of carbonate clearly
shows the electron density is spread equally among the three
oxygens
• Areas which are red are more negatively charged; areas
of blue have relatively less electron density
37
• Rules for Resonance:
• Individual resonance structures exist only on paper
• The real molecule is a hybrid (average) of all contributing
forms
• Resonance forms are indicated by the use of doubleheaded arrows
• Only electrons are allowed to move between resonance
structures
• The position of nuclei must remain the same
• Only electrons in multiple bonds and nonbonding
electrons can be moved
• Example: 3 is not a resonance form because a hydrogen atom
has moved
38
• Rules for Resonance:
• All structures must be proper Lewis
structures
39
• The energy of the actual molecule is lower than the
energy of any single contributing form
• The lowering of energy is called resonance stabilization
• Equivalent resonance forms make equal contributions to
the structure of the real molecule
• Structures with equivalent resonance forms tend to be
greatly stabilized
• Example: The two resonance forms of benzene
contribute equally and greatly stabilize it
• Unequal resonance structures contribute based on their
relative stabilities
• More stable resonance forms contribute more to the
structure of the real molecule
40
• Rules to Assign Relative Importance of
Resonance Forms
• A resonance form with more covalent bonds is
more important than one with less
• Example: 6 is more stable and more
important because it has more total
covalent bonds
41
• Resonance forms in which all atoms have a
complete valence shell of electrons are more
important
• Example: 10 is more important because all
atoms (except hydrogen) have complete
octets
42
• Resonance forms with separation of charge are less
important
• Separation of charge cost energy and results in a
less stable resonance contributor
• Example: 12 is less important because it has
charge separation
• Forms with negative charge on highly
electronegative atoms are more important
• Those with positive charge on less
electronegative atoms are also more important
43
• Example
• The nitrate ion is known to have all three
nitrogen-oxygen bond lengths the same and the
negative charge spread over all three atoms
equally
44
• Example
• Resonance theory can be used to produce three
equivalent resonance forms
• Curved arrows show the movement of
electrons between forms
• When these forms are hybridized (averaged)
the true structure of the nitrate ion is obtained
45
Quantum Mechanics
The development of quantum mechanics in the late 1920s
led to a better understanding of the relationship between
electronic and molecular structures.
Louis de Broglie, as a graduate student working on his Ph D in physics at
the Sorbonne in 1924, reasoned: If radiant energy under some conditions
behaved as a stream of particles, could not particles under some conditions
show the properties of radiant energy...a wave?
He proposed that an electron in an orbit around a nucleus (a Bohr orbit)
would have wave properties. There would be a characteristic wavelength
(l) associated with an electron in a particular Bohr orbit (an electron of
particular energy):
l = h/mv
where h is Planck's constant,
m is the mass of the electron,
and v is its velocity.
Wave Mechanics
A wave may be a traveling wave moving in a direction (x), or
a standing or stationary wave between two points.
A wave is described mathematically as a repeating function (, psi)
with positive and negative numerical values that repeat as a function
of distance (X) and time.

(+)
(+)
wave
function
(-)
(-)
x
A wave is a series of crests and troughs that repeat.
Standing Waves
A standing wave is a wave oscillating perpendicular to an axis between
two fixed points. An example is the string of a musical instrument such
as a violin. When plucked or bowed, the string vibrates in standing or
stationary waves. The waves do not travel along the string. Only an
integral number of half waves may exist between the fixed points.
Standing Waves
2 nodes
The points where the wave
function Y remains zero is a node.
In addition to each end where the
string is attached, there may be 0,
1,2, etc. nodes along the string.

1 node
0 nodes
• Quantum Mechanics
• A mathematical description of bonding that
takes into account the wave nature of
electrons
• A wave equation is solved to yield a series of
wave functions for the atom
• The wave functions psi () describe a series
of states with different energies for each
electron
• Wave Equations are used to calculate:
• The energy associated with the state of
the electron
• The probability of finding the electron in
a particular state
49
• Phase sign: Wave equations, when solved, may be
positive, negative or zero
• In analogy to a wave in a lake, when the wave is
above the average lake level, the sign is positive
( = +) ; when it is below the lake level it is
negative ( = - )
• When the wave is exactly at average lake level it
has a sign of 0 and this is called a node ( = 0)
• Wave equations can reinforce each other if they have
the same sign or interfere with each other if they
have different signs
50
:
Erwin Schrodinger (1887-1961)
:
Schrodinger obtained his Ph D in Physics at the
University of Zurich in 1910 and then served in the
Austrian Army during WWI. After the war he
resumed his career in science. In 1925, while at the
University of Zurich, he published his famous wave
equation that helped launch the field of quantum
mechanics. In later years, he was forced to flee
Austria with the rise of Nazism, and lived and
worked in England and Ireland..
Atomic Orbitals
:
Erwin Schrodinger (University of Zurich) in 1926 worked out the
mathematical equations describing the motions of an electron as a
standing wave around the nucleus in terms of its energy. Only certain
circular orbits around the nucleus have a correct circumference to
accommodate an integral number of waves, a complete standing wave.
Other circular orbits would not provide a good fit for an electron
wave and are not allowed.
a complete standing
electron wave
a mismatch
allowed
not allowed
This required "fit" for the electron wave seemed consistent with the
already accepted quantization of electron energies in the Bohr model.
Atomic Orbitals
:
The Schrodinger Equation
:
For the system of one electron and one proton (the hydrogen atom), Schrodinger
worked out the complicated mathematical expression for the wave function (Y)
that describes the position of the electron in three-dimensional space (x,y,z
coordinates) around the nucleus.
When the wave function is evaluated for different values of the total
energy of the electron, a series of solutions result. Each solution
consists of a specific wave function (Y) associated with a different value
of energy (E) for the electron.
The wave functions give numerical values (positive and negative phasing) for
regions in space around the nucleus for electrons of specified energies. These
mathematical solutions by themselves do not have physical meaning. However,
Y 2 for a particular location (x, y, z) expresses the PROBABILITY of finding an
electron for a particular location in space.
Plots of Y 2 in three dimensions generate the shapes of the familiar
s, p, and d atomic orbitals (AOs), which we use as our models for
atomic structures.
• Atomic Orbitals (AOs)
• The physical reality of  is that when
squared ( 2) it gives the probability of
finding an electron in a particular location in
space
• Plots of  2 in three dimensions generate
the shape of s, p, d and f orbitals
• Only s and p orbitals are very important in
organic chemistry
• Orbital: a region in space where the
probability of finding an electron is large
• The typical representation of orbitals are
those volumes which contain the electron
90-95% of the time
54
• 1s and 2s orbitals are spheres centered around
the nucleus
• Each orbital can accommodate 2 electrons
• The 2s orbital is higher in energy and
contains a nodal surface ( = 0) in its center
55
• Each 2p orbital has two nearly touching spheres (or
lobes)
• One sphere has a positive phase sign and the other
a negative phase sign; a nodal plane separates the
spheres
• There are three 2p orbitals which are perpendicular
(orthoganol) to each other
• Each p orbital can accommodate 2 electrons for a
total of 6 electrons
• All three p orbitals are degenerate (equal in energy)
• The 2p orbitals are higher in energy than the 1s or 2s
56
• The sign of the wave function does not
indicate a greater or lesser probability of
finding an electron in that location
• The greater the number of nodes in an
orbital the higher its energy
• 2s and 2p orbitals each have one node
and are higher in energy than the 1s
orbital which has no nodes
57
• Atoms can be assigned electronic
configuration using the following rules:
• Aufbau Principle: The lowest energy orbitals
are filled first
• Pauli Exclusion Principle: A maximum of two
spin paired electrons may be placed in each
orbital
• Hund’s Rule: One electron is added to each
degenerate (equal energy orbital) before a
second electron is added
58
Examples: Some Second Row Elements
Hund's rule
2p
2s
1s
B
C
N
O
F
Pauli exclusion
principle
What is an example of the Aufbau Principle?
Ne
Quiz Chapter 1 Section 10
Provide the chemical symbol (including charge) of the chemical species
with the following electron configurations:
1s2 2s2 2p6 3s2 (no charge)
Mg
1s2 2s2 2p6 (charge of +2)
Mg 2+
1s2 2s2 2p6 (charge of -2)
O 2-
electronic energy
Add the missing electrons in the diagram below as arrows (up or down
to indicate spin) for the ground electronic state of atomic carbon.
2p
2s
1s
Molecular Orbitals
When atoms combine to form molecules, the electrons reside in
molecular orbitals. Only two electrons may be placed in each
molecular orbital with paired spins.
Two important assumptions are:
(1) Each pair of electrons resides in a localized region near the
nuclei of the atoms that are forming the molecular orbital.
(2) The shapes of the localized molecular orbital are related in
a simple way to the shapes of the combining atomic orbitals.
• Molecular Orbitals (MOs)
• A simple model of bonding is illustrated by forming
molecular H2 from H atoms and varying distance:
• Region I: The total energy of two isolated atoms
• Region II: The nucleus of one atom starts attracting the
electrons of the other; the energy of the system is
lowered
• Region III: at 0.74 Å the attraction of electrons and
nuclei exactly balances repulsion of the two nuclei; this
is the bond length of H2
• Region IV: energy of system rises as the repulsion of the
two nuclei predominates
62
• This simple model of bonding does not take into
account the fact that electrons are not stationary
but constantly moving around
• Heisenberg uncertainty principle: the position
and momentum of an electron cannot
simultaneously be known
• Quantum mechanics solves this problem by talking
about the probability ( 2) of finding an electron at
a certain location in space
• As two atoms approach each other their atomic
orbitals (AOs) overlap to become molecular
orbitals (MOs)
• The wave functions of the AOs are combined to
yield the new wave functions of the MOs
• The number of MOs that result must always equal
the number of AOs used
63
• Example: H2 molecule
• As the hydrogen atoms approach each other their 1s
orbitals (1s) begin to overlap
• The MOs that form encompass both nuclei
• The electrons are not restricted to the vicinity of one
nucleus or another
• Each MO has a maximum of 2 spin-paired electrons
• Addition of wave functions of the two atoms leads to
a bonding molecular orbital
• Subtraction of wave functions of the two atoms leads
to an anti-bonding molecular orbital
• The mathematic operation by which wave functions
are added or subtracted is called the linear
combination of atomic orbitals (LCAO)
64
• Bonding Molecular Orbitals (molec)
• AOs combine by addition (the AOs of the same
phase sign overlap)
• The wave functions reinforce
• The value of  increases between the two nuclei (b)
• The value of 2 (electron probability density) in the
region between the two nuclei increases (a)
• The two electrons between the nuclei serve to
attract the nuclei towards each other
• This is the ground state (lowest energy state) of the
MO
65
• Antibonding molecular orbital ( *molec)
• Formed by interaction of AOs with opposite phase signs
• Wave functions interfere and a node is produced ( = 0)
• In the region between the two nuclei
• A node is produced
• On either side of the node  is small
•  2 (electron probability density) is small
66
• Electrons in the antibonding orbital avoid the
region between the two nuclei
• Repulsive forces between the nuclei predominate
and electrons in antibonding orbitals make nuclei
fly apart
67
• The energy of electrons in the bonding orbitals is
substantially less than the energy of electrons in
the individual atoms
• The energy of electrons in the antibonding
orbitals is substantially more
• In the ground state of the hydrogen molecule
electrons occupy the lower energy bonding orbital
only
68
Examples of Covalent Bonds
2 H.
H
H
1s
1s
2 F.
F
F
2p
2p
H2
H
H
 bond
bond length 0.74 Å
BDE 435 kJ/mol
F2
F
F
2p bond
bond length 1.42 Å
BDE 159 kJ/mol
Bonding and Antibonding MOs
combining atomic orbitals
molecular orbitals
s *
antibonding
+
s
s
s
bonding
p *
antibonding
+
p
p *
antibonding
+
p
p
p
bonding
p
p
bonding
• The Structure of Methane and Ethane: sp3
Hybridization
• The structure of methane with its four identical
tetrahedral bonds cannot be adequately
explained using the electronic configuration of
carbon
• Hybridization of the valence orbitals (2s and 2p)
provides four new identical orbitals which can
be used for the bonding in methane
• Orbital hybridization is a mathematical
combination of the 2s and 2p wave functions to
obtain wave functions for the new orbitals
71
Hybrid Orbitals and Molecular Shape
sp3 hybridization
CH4
The ground state electron configuration of C is
1s2 2s2 2p2
or
1s
2s
2px
valence level
This electron configuration explains
neither the bonding capacity of C
nor the shape of CH4 .
2py
2pz
Hybrid Orbitals
Step One
Electron Promotion 2s
2p
+energy
1s
2s
2px
2py
2pz
1s
2s
ground state of C
Step Two
1s
2s 2px 2py 2pz
This set of valence
orbitals does not
explain the molecular
shape of methane.
2px 2py 2pz
excited state
Orbital Hybridization
1s
sp3 sp3 sp3 sp3
hybrid orbitals
Overview
orbital hybridization
electron promotion
+ energy
1s
2s
2px 2py 2pz
ground state for C
1s
2s 2px 2py 2pz
excited state
2px 2py 2pz
2s
2s
1s
1s
sp3 sp3 sp3 sp3
hybrid orbitals
sp3 sp3 sp3 sp3
energy
2px 2py 2pz
1s
1s
The Three-Dimensional Symmetry of sp3 Hybrid Orbitals
2px orbital
z
y
x +
2s orbital
2py orbital
2pz orbital
yields
mixing of the
2s, 2px,2py and 2pz orbitals
four sp3 hybrid
orbitals directed
towards the corners
of a tetrahedron
• When one 2s orbital and three 2p orbitals are
hybridized four new and identical sp3 orbitals are
obtained
• When four orbitals are hybridized, four orbitals
must result
• Each new orbital has one part s character and 3
parts p character
• The four identical orbitals are oriented in a
tetrahedral arrangements
• The antibonding orbitals are not derived in the
following diagram
• The four sp3 orbitals are then combined with the 1s
orbitals of four hydrogens to give the molecular
orbitals of methane
• Each new molecular orbital can accommodate 2
electrons
76
77
• An sp3 orbital looks like a p orbital with one lobe
greatly extended
• Often the small lobe is not drawn
• The extended sp3 lobe can then overlap well with the
hydrogen 1s to form a strong bond
• The bond formed is called a sigma () bond because
it is circularly symmetrical in cross section when view
along the bond axis
78
• A variety of representations of methane show its
tetrahedral nature and electron distribution
• a. calculated electron density surface b. ball-andstick model c. a typical 3-dimensional drawing
79
• Ethane (C2H6)
• The carbon-carbon bond is made from overlap
of two sp3 orbitals to form a  bond
• The molecule is approximately tetrahedral
around each carbon
80
• The representations of ethane show the
tetrahedral arrangement around each carbon
• a. calculated electron density surface b. balland-stick model c. typical 3-dimensional
drawing
• Generally there is relatively free rotation about 
bonds
• Very little energy (13-26 kcal/mol) is required
to rotate around the carbon-carbon bond of
ethane
81
• The Structure of Ethene (Ethylene) : sp2 Hybridization
• Ethene (C2H2) contains a carbon-carbon double bond
and is in the class of organic compounds called
alkenes
• Another example of the alkenes is propene
• The geometry around each carbon is called trigonal
planar
• All atoms directly connected to each carbon are in
a plane
• The bonds point towards the corners of a regular
triangle
• The bond angle are approximately 120o
82
• There are three  bonds around each carbon of ethene
and these are formed by using sp2 hybridized orbitals
• The three sp2 hybridized orbitals come from mixing one s
and two p orbitals
• One p orbital is left unhybridized
• The sp2 orbitals are arranged in a trigonal planar
arrangement
• The p orbital is perpendicular (orthoganol) to the
plane
83
The Hybrid Orbital Model
The structural and electronic properties of ethene are consistent with the
hybrid orbital model using carbons that are sp2 hybridized.
A set of valence level hybrid orbitals is obtained by mathematically
combining the 2s and two 2p atomic orbitals of carbon. This process can
be broken down into two steps, electron promotion and hybridization or
mixing.
+ energy
2px 2py 2pz electron 2px 2py 2pz hybridization
promotion
2s
energy
2s
2pz
sp2 sp2
sp2
valence-level orbitals
1s
1s
ground state
of C
excited state
of C
1s
This process is overall energetically favorable
(despite electron promotion) because it allows
more complete use of the valence electrons in
forming bonds.
Spatial Arrangement of the Hybrid Orbitals
pz
sp2 hybridized carbon
valence
level
2pz
sp2 sp2
sp2
sp2
C
sp2
sp2
bonding orbitals
1s
Three equivalent sp2 orbitals
are in the X-Y plane. The idealized
inter-orbital angle is 120o.
The Carbon-Carbon Double Bond
The carbon-carbon double bond results from valence level orbital
interaction between two sp2 hybridized carbons as shown below. In
this picture, the two atomic centers are brought together along an
axis that allows overlap of sp2 hybrid orbitals.
pz
pz
sp2
C
sp2
C


C
C

The double bond is made
up of a  bond from
overlapping sp2 orbitals,
and a  bond from
overlapping pz orbitals.
• Overlap of sp2 orbitals in ethylene results in formation of
a  framework
• One sp2 orbital on each carbon overlaps to form a
carbon-carbon  bond; the remaining sp2 orbitals form
bonds to hydrogen
• The leftover p orbitals on each carbon overlap to form a
bonding  bond between the two carbons
• A  bond results from overlap of p orbitals above and
below the plane of the  bond
87
• The bonding  orbital is lower in energy than the
antibonding orbital
• In the ground state two spin paired electrons are in
the bonding orbital
• The antibonding *orbital can be occupied if an
electron becomes promoted from a lower level ( e.g.
by absorption of light)
88
•
•
The bonding  orbital results from overlap of p orbital lobes of the same sign
The antibonding * orbital results from overlap of p orbital lobes of opposite sign
• The antibonding orbital has one node connecting the two nuclei and another
node between the two carbons
89
• The  orbital is lower in energy than the 
orbital
• The ground state electronic configuration of
ethene is shown
90
Complete Set of MO’s for Ethene
Molecular Orbitals of Ethene
*
atomic orbitals
atomic orbitals
pz
pz
sp2
1s
H H
pz
H

sp2
1s
C
C
sp2
H H
pz
H
sp2
C
C
H
H
pz atomic orbitals
combine to form
bonding and
antibonding  MOs
Molecular Orbitals of Ethene
*c-c
*
atomic orbitals
atomic orbitals
pz
pz
sp2
1s
H H
pz
H

1s
C
C
sp2
c-c
C
H
sp2
H H
pz
H
sp2
C
sp2 atomic orbitals
combine to form
bonding and antibonding
C-C MOs
H
Molecular Orbitals of Ethene
*C-H
combine to form
 (C-H) bonding
and antibonding MOs
combine to form
 (C-H) bonding
and antibonding MOs
*c-c
*
atomic orbitals
atomic orbitals
pz
pz
sp2
1s
H H
pz
H
sp2

1s
C
C
c-c
sp2
H H
pz
H
sp2
C
C
H
H
C-H
(4 filled MOs)
Molecular Orbitals of Ethene
Approximate MO
Scheme for Ethene
atomic orbitals
atomic orbitals
*C-H
*c-c
pz
pz
*
sp2
sp2
1s
H H
pz
H
1s

C
C
c-c
sp2
C
(4 filled MOs)


H
C
H
C
H
H

pz
H
sp2
C-H
H
H H
C
H
• Restricted Rotation and the Double Bond
• There is a large energy barrier to rotation (about
264 kJ/mol) around the double bond
• This corresponds to the strength of a  bond
• The rotational barrier of a carbon-carbon single
bond is 13-26 kJ/mol
• This rotational barrier results because the p orbitals
must be well aligned for maximum overlap and
formation of the  bond
• Rotation of the p orbitals 90o totally breaks the 
bond
95
Restricted Rotation Around the Double Bond
The picture of the carbon-carbon double bond as separate  and 
bonds is consistent with the observed large barrier (approx. 251
kJ/mol) to rotation around the C-C axis. While -bonds are
cylindrically symmetrical around the axis of the bond (meaning that
rotation around the bond axis does not cause loss of orbital overlap),
-bonds are not.
A rotation of one carbon relative to the other uncouples (loss of orbital
overlap) the p-orbitals. When the p-orbitals are 90o apart, there is no
overlap and the electronic energy is the same as for an electron
residing in an isolated p-orbital.
C C
or
rotation
C C
o
90
C
C
localized
p-orbitals
o
90
view along the
C-C axis
C C
loss of orbital overlap
-bond
The Energy Barrier to Rotation
C
C C
C C
o
o
localized
p-orbitals
90
-bond
250
An energy barrier of close
to250 kJ/mol is estimated
during rotation around the
carbon-carbon double bond.
kJ/mol
-electron potential energy
-bond
90
C
0
45
90
degree of rotation
135
180
Cis-Trans Isomerism in Alkenes
The barrier of close to 251 kJ/mol means that rotation around the
carbon-carbon double bond does not occur under ordinary conditions.
This restricted rotation leads to a type of isomerism: the existence of
different compounds with the same formula.
The two compounds below with the same formula of C2H2Cl2
illustrate this isomerism.
Cl
Cl
H
H
cis-1,2-dichloroethene
"cis" (Latin: "on this side")
Cl
H
H
Cl
trans-1,2-dichloroethene
"trans" (Latin: "across")
-80o C
60o C
-50o C
48o C
dipole 1.90 D
moment
0D
MP
BP
These isomers are sometimes called geometric isomers.
• Cis-trans isomers
• Cis-trans isomers are the result of restricted
rotation about double bonds
• These isomers have the same connectivity of
atoms and differ only in the arrangement of atoms
in space
• This puts them in the broader class of
stereoisomers
• The molecules below do not superpose on each
other
• One molecule is designated cis (groups on same
side) and the other is trans (groups on opposite
side)
99
• Cis-trans isomerism is not possible if one
carbon of the double bond has two identical
groups
100
Quiz Chapter 1 Section 13
Identify the covalent bonds in ethene by type ( or
) and indicate the atomic orbitals from the atoms
that form the covalent bonds.
bond type

C p
C p
H
H
C
C
H
H
ethene
bond type

combining
atomic orbitals
combining
atomic orbitals
C sp 2
2
C sp
bond type

combining
atomic orbitals
C sp 2
H s
• The Structure of Ethyne (Acetylene): sp
Hybridization
• Ethyne (acetylene) is a member of a group of
compounds called alkynes which all have carboncarbon triple bonds
• Propyne is another typical alkyne
• The arrangement of atoms around each
carbon is linear with bond angles 180o
102
• The carbon in ethyne is sp hybridized
• One s and one p orbital are mixed to form
two sp orbitals
• Two p orbitals are left unhybridized
103
• The two sp orbitals are oriented 180o
relative to each other around the carbon
nucleus
• The two p orbitals are perpendicular to
the axis that passes through the center of
the sp orbitals
104
• In ethyne the sp orbitals on the two carbons
overlap to form a  bond
• The remaining sp orbitals overlap with hydrogen
1s orbitals
• The p orbitals on each carbon overlap to form two
 bonds
• The triple bond consists of one  and two  bonds
105
• Depictions of ethyne show that the electron
density around the carbon-carbon bond has
circular symmetry
• Even if rotation around the carbon-carbon
bond occurred, a different compound
would not result
106
• Bond Lengths of Ethyne, Ethene and Ethane
• The carbon-carbon bond length is shorter as
more bonds hold the carbons together
• With more electron density between the
carbons, there is more “glue” to hold the
nuclei of the carbons together
107
• The carbon-hydrogen bond lengths also get shorter with more
s character of the bond
• 2s orbitals are held more closely to the nucleus than 2p
orbitals
• A hybridized orbital with more percent s character is held
more closely to the nucleus than an orbital with less s
character
• The sp orbital of ethyne has 50% s character and its C-H
bond is shorter
• The sp3 orbital of ethane has only 25% s character and its
C-H bond is longer
108
Orbital Hybridization, Bond Lengths
and Bond Strengths
The greater the degree of s-character in a hybrid orbital that
overlaps with another atomic orbital to form a covalent bond, the
shorter the covalent bond and the stronger the bond.
hydrocarbon
H
H
H
H
H
H
hybridization
sp3
bond lengths
C-C
C-H
Ho (C-H)
1.54 Å 1.10 Å
410
(kJ/mol)
(sp3 - 1s)
H
H
H
H
sp2
H
H
sp
1.34 Å 1.09 Å
452
(sp2- 1s)
1.20 Å 1.06 Å
(sp- 1s)
523
Orbital Hybridization Influences on
C-C Bond Lengths and Bond Strengths
C
CH3
orbitals
C-C Bond Length
Ho (C-C)
(kJ/mol)
H3C CH3
sp3
sp3
1.54 Å
368
CH3
sp2
sp3
1.50 Å
385
CH3
sp
sp3
1.46 Å
490
CH2=CH
HC C
Shorter bonds are generally stronger bonds.
Quiz Chapter 1 Section 14
The longest C-H bond among those
labeled in the structure above is
The strongest C-H bond among
those labeled in the structure
below is
B
A
H C
C
H
C
H
C
H
C
H
C
D
H
Section 15--A Summary of Important Concepts
That Come from Quantum Mechanics
1.
An atomic orbital (AO) corresponds to a region of space about the nucleus of a single
atom where there is a high probability of finding an electron. Atomic orbitals called s
orbitals are spherical; those called p orbitals are like two almost-tangent spheres.
Orbitals can hold a maximum of two electrons when their spins are paired. Orbitals
are described by a wave function, , and each orbital has a characteristic energy. The
phase signs associated with an orbital may be (+) or (-).
2.
When atomic orbitals overlap, they combine to form molecular orbitals (MOs). Molecular
orbitals correspond to regions of space encompassing two (or more) nuclei where electrons
are to be found. Like atomic orbitals, molecular orbitals can hold up to two electrons if
their spins are paired.
3.
When atomic orbitals with the same phase interact, they combine to form a bonding
molecular orbital:
The electron probability density of a bonding molecular orbital is large in the region of
space between the two nuclei where the negative electrons hold the positive nuclei together
4.
An antibonding molecular orbital forms when orbitals of opposite phase sign overlap:
An antibonding orbital has higher energy than a bonding orbital. The electron probability
density of the region between the nuclei is small and it contains a node--a region where
 = O. Thus, having electrons in an antibonding orbital does not help hold the nuclei
together. The internuclear repulsions tend to make them fly apart.
5.
The energy of electrons in a bonding molecular orbital is less than the energy of the electrons
in their separate atomic orbitals. The energy of electrons in an antibonding orbital is greater
than that of electrons in their separate atomic orbitals.
6.
The number of molecular orbitals always equals the number of atomic orbitals from which
they are formed. Combining two atomic orbitals will always yield two molecular orbitals-one bonding and one antibonding.
7.
Hybrid atomic orbitals are obtained by mixing (hybridizing) the wave functions for orbitals
of different types (i.e., s and p orbitals) but from the same atom.
8.
Hybridizing three p orbitals with one s orbital yields four sp3 orbitals. Atoms that are sp3
hybridized direct the axes of their four sp3 orbitals toward the corners of a tetrahedron.
The carbon of methane is sp3 hybridized and tetrahedral.
9.
Hybridizing two p orbitals with one s orbital yields three sp2 orbitals. Atoms that are sp2
hybridized point the axes of their three sp2 orbitals toward the corners of an equilateral
triangle. The carbon atoms of ethane are sp2 hybridized and trigonal planar.
10. Hybridizing one p orbital with one s orbital yields two sp orbitals. Atoms that are sp
hybridized orient the axes of their two sp orbitals in opposite directions (at an angle of
180o). The carbon atoms of ethyne are sp hybridized and ethyne is a linear molecule.
11. A sigma () bond (a type of single bond) is one in which the electron density has circular
symmetry when viewed along the bond axis. In general, the skeletons of organic molecules
are constructed of atoms linked by sigma bonds.
12. A pi () bond, part of a double and triple carbon-carbon bonds, is one in which the electron
densities of two adjacent parallel p orbitals overlap sideways to form a bonding pi molecular
orbital
Valence Shell Electron Pair Repulsion Theory
VSEPR theory provides a simple way to predict molecular
geometries around a central atom. These predictions are consistent
with hybrid molecular orbital theory..
Key Features of VSEPR Theory
(1) Molecules or ions may be analyzed where a central atom is
covalently bonded to two or more atoms or groups.
(2) All pairs of valence electrons around the central atom are counted:
bonding and nonbonding.
(3) Because of electron-electron repulsion, pairs of electrons tend to stay
as far apart as possible. Repulsion due to nonbonding pairs of
electrons is greater than repulsion due to bonding pairs.
(4) The preferred geometry has all pairs of valence electrons as far
apart as possible to minimize repulsion among electron pairs.
Examples
H
H C H
H
Methane
:
:
:
:
There are four pairs of bonding electrons in
the valence level around the central carbon.
Maximum separation of the four pairs is
achieved with a tetrahedral geometry where
each electron pair points to the corner of a
tetrahedron.
H
H C 109.5o
H
H
an idealized
tetrahedral
geometry
• The number of molecular orbitals formed equals the number of the
atomic orbitals used
• Hybridized orbitals are obtained by mixing the wave functions of
different types of orbitals
• Four sp3 orbitals are obtained from mixing one s and three p
orbitals
• The geometry of the four orbitals is tetrahedral
• This is the hybridization used in the carbon of methane
• Three sp2 orbitals are obtained from mixing one s and two p
orbitals
• The geometry of the three orbitals is trigonal planar
• The left over p orbital is used to make a  bond
• This is the hybridization used in the carbons of ethene
• Two sp orbitals are obtained from mixing one s and one p orbital
• The geometry of the two orbitals is linear
• The two leftover p orbitals are used to make two  bonds
• This is the hybridization used in the carbons of ethyne
• Sigma () bonds have circular symmetry when viewed along the
bond axis
• Pi () bonds result from sideways overlap of two p orbitals
• Molecular Geometry: The Valence Shell Electron Pair
Repulsion (VSEPR) Model
• This is a simple theory to predict the geometry of
molecules
• All sets of valence electrons are considered including:
• Bonding pairs involved in single or multiple bonds
• Non-bonding pairs which are unshared
• Electron pairs repel each other and tend to be as far
apart as possible from each other
• Non-bonding electron pairs tend to repel other
electrons more than bonding pairs do (i.e. they are
“larger”)
• The geometry of the molecule is determined by the
number of sets of electrons by using geometrical
principles
119
• Methane
• The valence shell of methane contains four
pairs or sets of electrons
• To be as far apart from each other as
possible they adopt a tetrahedral
arrangement (bond angle 109.5o)
• The molecule methane is therefore
tetrahedral
120
• Ammonia
• When the bonding and nonbonding electrons are
considered there are 4 sets of electrons
• The molecule is essentially tetrahedral but the actual
shape of the bonded atoms is considered to be
trigonal planar
• The bond angles are about 107o and not 109.5o
because the non-bonding electrons in effect are
larger and compress the nitrogen-hydrogen bond
121
• Water
• There are four sets of electrons including 2 bonding pairs
and 2 non-bonding pairs
• Again the geometry is essentially tetrahedral but the
actual shape of the atoms is considered to be an angular
arrangement
• The bond angle is about 105o because the two “larger”
nonbonding pairs compress the electrons in the oxygenhydrogen bonds
122
• Boron Trifluoride
• Three sets of bonding electrons are farthest
apart in a trigonal planar arrangement (bond
angle 120o)
• The three fluorides lie at the corners of an
equilateral triangle
• Beryllium Hydride
• Two sets of bonding electrons are farthest
apart in a linear arrangement (bond angles
180o)
123
• Carbon Dioxide
• There are only two sets of electrons around the
central carbon and so the molecule is linear (bond
angle 180o)
• Electrons in multiple bonds are considered as one
set of electrons in total
• A summary of the results also includes the geometry
of other simple molecules
124
Quiz Chapter 1 Section 16
Use VSEPR theory to predict the geometry around the carbon atom in
each of the following chemical species.
: :
H2C=O
formaldehyde
H3C+
methyl cation
-
H3C:
methyl anion
H
H
C O
H
1 set (4) of
bonding electrons
and two pairs
of bonding electrons
trigonal planar
C H
H
three pairs
of bonding electrons
trigonal planar
H
C
H
H
three pairs
of bonding electrons
and one pair
of nonbonding electrons
tetrahedral
(trigonal pyramidal)
• Representations of Structural Formulas
• Dot formulas are more cumbersome to
draw than dash formulas and condensed
formulas
• Lone-pair electrons are often (but not
always) drawn in. They are drawn when
they are crucial to the chemistry being
discussed
126
• Dash formulas
• Each dash represents a pair of electrons
• This type of representation is meant to emphasize
connectivity and does not represent the 3-dimensional
nature of the molecule
• The dash formulas of propyl alcohol appear to have
90o angles for carbons which actually have
tetrahedral bond angles (109.5o)
• There is relatively free rotation around single bonds so
the dash structures below are all equivalent
127
• Constitutional isomers
• Constitutional isomers have the same
molecular formula but different connectivity
• Propyl alcohol (above) is a constitutional
isomer of isopropyl alcohol (below)
128
•
Condensed Structural Formulas
• In these representations, some or all of the dash lines are omitted
• In partially condensed structures all hydrogens attached to an
atom are simply written after it but some or all of the other bonds
are explicitly shown
• In fully condensed structure all bonds are omitted and atoms
attached to carbon are written immediately after it
• For emphasis, branching groups are often written using vertical
lines to connect them to the main chain
129
• Bond-Line Formulas
• A further simplification of drawing organic molecules is to
completely omit all carbons and hydrogens and only
show heteroatoms (e.g. O, Cl, N) explicitly
• Each intersection or end of line in a zig-zag represents a
carbon with the appropriate amount of hydrogens
• Heteroatoms with attached hydrogens must be drawn
in explicitly
130
• Cyclic compounds are condensed using a drawing of
the corresponding polygon
• Multiple bonds are indicated by using the
appropriate number of lines connecting the atoms
131
• Three-Dimensional Formulas
• Since virtually all organic molecules have a 3dimensional shape it is often important to be able to
convey their shape
• The conventions for this are:
• Bonds that lie in the plane of the paper are
indicated by a simple line
• Bonds that come forward out of the plane of the
paper are indicated by a solid wedge
• Bonds that go back out of the plane of the paper
are indicated by a dashed wedge
• Generally to represent a tetrahedral atom:
• Two of the bonds are drawn in the plane of the
paper about 109o apart
• The other two bonds are drawn in the opposite
direction to the in- plane bonds but right next to
each other
132
Three-Dimensional Representations
Three-dimensional representations of structures are often important.
By convention, a solid wedge is a bond projected out of the plane of
paper or screen towards the viewer. A dashed wedge is a bond
projected out of the plane away from the viewer. A solid line is a
bond in the plane.
towards viewer
in plane
away from viewer
Examples
H
Br
H
C
H
CH3Br
H
Cl
C H
H
OH
CH3OH
H3 C
H
C
H
CH3CH2Cl
134
• Trigonal planar arrangements of atoms can
be drawn in 3-dimensions in the plane of
the paper
• Bond angles should be approximately
120o
• These can also be drawn side-on with the
central bond in the plane of the paper,
one bond forward and one bond back
• Linear arrangements of atoms are always
best drawn in the plane of the paper
135
Quiz Chapter 1 Section 17
Are the following pairs of structural formulas the same or different
compounds?
(CH3)2CHCH2CH3
CH3CH2CHCH3
CH3
CH3CH2CHClCHCH3
CH3CH2CHClCHCH2CH3
CH2CH3
CH3
CH3CHCH2CHCH2CH3
CH3
Cl
(CH3)2CHCHClCH2CH2CH3
same
same
different
Chapter 1
137