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William H. Brown
Christopher S. Foote
Brent L. Iverson
Eric Anslyn
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Chapter 1
Covalent Bonding and Shapes of Molecules
William H. Brown • Beloit College
1-1
Organic Chemistry
 The
study of the compounds of carbon.
 Over 10 million compounds have been identified.
• About 1000 new ones are identified each day!
C
is a small atom.
• It forms single, double and triple bonds.
• It is intermediate in electronegativity (2.5).
• It forms strong bonds with C, H, O, N, and some
metals.
1-2
Schematic View of an Atom
• A small dense nucleus,
diameter 10-14 - 10-15 m,
which contains positively
charged protons and
most of the mass of the
atom.
• An extranuclear space,
diameter 10-10 m, which
contains negatively
charged electrons.
1-3
The Concept of Energy
 In
the ground state of carbon, electrons are
placed in accordance with the quantum
chemistry principles (aufbau, Hund’s rule, Pauli
exclusion principle, etc.) that dictate the lowest
energy form of carbon.
 If we place the electrons in a different manner (as
for example with one electron in the 2s and three
electrons in the 2p) we would have a higher
energy level referred to as an excited state. When
the electrons are rearranged back to the ground
state, energy is released.
1-4
The Concept of Energy
 Electrons
in the lowest energy orbital, 1s, are
held tightest to the nucleus and are the hardest
to remove from the atom.
 First ionization energy: The energy needed to
remove the most loosely held electron from an
atom or molecule.
1-5
Shapes of Atomic s and p Orbitals
• All s orbitals have the
shape of a sphere with
the center of the sphere
at the nucleus.
• Figure 1.8 (a)
Calculated and (b)
cartoon representations
showing an arbitrary
boundary surface
containing about 95%
of the electron density.
1-6
Shapes of Atomic s and p Orbitals
• Figure 1.9 (a) Three-dimensional representations of the
2px, 2py, and 2pz atomic orbitals computed using the
Schrödinger equation. Nodal planes are shaded.
1-7
Shapes of Atomic s and p Orbitals

Figure 1.9(b) Cartoon representations of the 2px, 2py, and
2pz atomic orbitals.
1-8
Molecular Orbital Theory
 MO
theory begins with the hypothesis that
electrons in atoms exist in atomic orbitals and
electrons in molecules exist in molecular
orbitals.
1-9
Molecular Orbital Theory
 Figure
1.10 MOs derived from combination by (a)
addition and (b) subtraction of two 1s atomic
orbitals.
1-10
VB: Hybridization of Atomic Orbitals
Energy
• The number of hybrid orbitals formed is equal to the
number of atomic orbitals combined.
• Elements of the 2nd period form three types of hybrid
orbitals, designated sp3, sp2, and sp.
• The mathematical combination of one 2s atomic orbital
and three 2p atomic orbitals forms four equivalent sp3
hybrid orbitals.
2p
sp3
2s
sp 3 Hybridization, with electron population for
carbon to form four single bonds
1-11
VB: Hybridization of Atomic Orbitals
1.12 sp3 Hybrid orbitals. (a) Computed and
(b) cartoon three-dimensional representations.
(c) Four balloons of similar size and shape tied
together, will assume a tetrahedral geometry.
 Figure
1-12
VB: Hybridization of Atomic Orbitals
 Figure
1.13 Orbital overlap pictures of methane,
ammonia, and water.
1-13
VB: Hybridization of Atomic Orbitals
1.14 sp2 Hybrid orbitals and a single 2p
orbital on an sp2 hybridized atom.
 Figure
1-14
VB: Hybridization of Atomic Orbitals
mathematical combination of one 2s atomic
orbital and one 2p atomic orbital gives two
equivalent sp hybrid orbitals.
Energy
 The
2p
2p
sp
2s
sp Hybridization, with electron
population for carbon to form
triple bonds
1-15
VB: Hybridization of Atomic Orbitals
 Figure
1.16 sp Hybrid orbitals and two 2p orbitals
on an sp hybridized atom.
1-16
Combining VB & MO Theories
double bond uses sp2 hybridization.
 Consider ethylene, C2H4. Carbon and other
second-period elements use a combination of
sp2 hybrid orbitals and the unhybridized 2p
orbital to form double bonds.
A
1-17
Combining VB & MO Theories
A
carbon-carbon triple
bond consists of one s
bond formed by
overlap of sp hybrid
orbitals and two p
bonds formed by the
overlap of parallel 2p
atomic orbitals.
1-18
Covalent Bonding of Carbon
Groups
Bonded
to Carbon
Orbital
Hybridization
4
sp 3
3
sp
2
2
sp
Predicted
Bond
Angles
Types of
Bonds
to Each Carbon
Example
Name
H H
H-C- C- H
H H
Ethane
109.5°
four s bonds
120°
three sbonds
and one p bond
H
H
C C
H
H
Ethene
180°
two s bonds
and two p bonds
H- C C- H
Ethyne
1-19
Polar and Nonpolar Molecules
 To
determine if a molecule is polar, we need to
determine
• if the molecule has polar bonds and
• the arrangement of these bonds in space.
 Molecular
dipole moment (): The vector sum of
the individual bond dipole moments in a
molecule.
• reported in Debyes (D)
1-20
Electrostatic Potential (elpot) Maps
 Relative
electron density distribution in
molecules is important because it allows us to
identify sites of chemical reactivity.
• Many reactions involve an area of relatively high
electron density on one molecule reacting with an area
of relatively low electron density on another molecule.
• It is convenient to keep track of overall molecular
electron density distributions using computer
graphics.
1-21
Electrostatic Potential (elpot) Maps
 In
electrostatic potential maps (elpots)
• Areas of relatively high calculated electron density are
shown in red.
• Areas of relatively low calculated electron density are
shown in blue.
• Intermediate electron densities are represented by
intermediate colors.
1-22
Polar and Nonpolar Molecules
 These
molecules have polar bonds, but each
molecule has a zero dipole moment.
Cl
F
O
C
B
O
F
F
Carbon dioxide
=0D
Boron trifluoride
=0D
C
Cl
Cl
Cl
Carbon tetrachloride
=0D
1-23
Polar and Nonpolar Molecules
 These
molecules have polar bonds and are polar
molecules.
direction
of dip ole
moment
N
O
H
H
Water
 = 1.85D
H
H
Ammonia
 = 1.47D
H
direction
of dip ole
moment
1-24
Polar and Nonpolar Molecules
• Formaldehyde has polar bonds and is a polar
molecule.
direction
of dip ole
moment
O
H
C
H
Formaldehyde
 = 2.33 D
1-25
Resonance
many molecules and ions, no single Lewis
structure provides a truly accurate
representation.
:
:
 For
O:
:O :
H3 C C
and
O:
-
:
:
:O:
H3 C C
Ethanoate ion
(acetate ion)
1-26
Resonance

Examples: equivalent contributing structures.
CH3
O:
CH3
O:
C
: O :-
:
N itrite ion
(equivalent con trib uting
s tru ctures)
C
:
:O :-
:
:
O:
: O :-
O:
:N
:
:N
:
:
:
:O: -
A cetate ion
(equ ivalen t contributin g
s tru ctures)
1-27
Resonance
 Curved
arrow: A symbol used to show the
redistribution of valence electrons.
 In using curved arrows, there are only two
allowed types of electron redistribution:
• from a bond to an adjacent atom.
• from a lone pair on an atom to an adjacent bond.
 Electron
pushing is a survival skill in organic
chemistry.
• learn it well!
1-28
Resonance
 All
contributing structures must
1. have the same number of valence electrons.
2. obey the rules of covalent bonding:
• no more than 2 electrons in the valence shell of H.
• no more than 8 electrons in the valence shell of a 2nd
period element.
3. differ only in distribution of valence electrons; the
position of all nuclei must be the same.
4. have the same number of paired and unpaired electrons.
1-29
Resonance
 The
carbonate ion
• Is a hybrid of three equivalent contributing structures.
• The negative charge is distributed equally among the
three oxygens as shown in the elpot.
1-30
Resonance
 Preference
4: negative charge on the more
electronegative atom.
• Structures that carry a negative charge on the more
electronegative atom contribute more than those with
the negative charge on a less electronegative atom.
O
(1)
C
H3 C
O
O
CH3
(a)
Less er
con trib ution
(2)
C
H3 C
CH3
(b)
Greater
contribu tion
C
H3 C
CH3
(c)
S hould n ot
be d raw n
1-31
Lewis Dot Structures
 Gilbert
N. Lewis
 Valence shell:
• The outermost occupied electron shell of an atom.
 Valence
electrons:
• Electrons in the valence shell of an atom; these
electrons are used to form chemical bonds and in
chemical reactions.
 Lewis
dot structure:
• The symbol of an element represents the nucleus and
all inner shell electrons.
• Dots represent electrons in the valence shell of the
atom.
1-32
Lewis Dot Structures
Table 1.4 Lewis Dot Structures for Elements 1-18
1A
Na
4A
5A
6A
7A
.
8A
He
.
.
Be
Mg
:
:
.
B
.
Al
:
.
C:
.
.
. N. :
.
:.:
.
:..F :
:
.
Si :
.
.
. .P :
.
: S. :
.
O
:Cl :
:
: :
Li
3A
:N e :
: :
H
2A
:

:A r :
1-33
Lewis Model of Bonding
 Atoms
interact in such a way that each
participating atom acquires an electron
configuration that is the same as that of the
noble gas nearest it in atomic number.
• An atom that gains electrons becomes an anion.
• An atom that loses electrons becomes a cation.
• The attraction of anions and cations leads to the
formation of ionic solids. This ionic interaction is often
referred to as an ionic bond.
• An atom may share electrons with one or more atoms
to complete its valence shell; a chemical bond formed
by sharing electrons is called a covalent bond. Bonds
may be partially ionic or partially covalent; these
1-34
bonds are called polar covalent bonds
Electronegativity
 Electronegativity:
• A measure of an atom’s attraction for the electrons it
shares with another atom in a chemical bond.
 Pauling
scale
• Generally increases left to right in a row.
• Generally increases bottom to top in a column.
1-35
Covalent Bonds
 The
simplest covalent bond is that in H2
• The single electrons from each atom combine to form
an electron pair.
H•
+
•H
H-H
H0 = -435 kJ (-104 kcal)/mol
• The shared pair functions in two ways simultaneously;
it is shared by the two atoms and fills the valence shell
of each atom.
 The
number of shared pairs
• One shared pair forms a single bond
• Two shared pairs form a double bond
• Three shared pairs form a triple bond
1-36
Polar and Nonpolar Covalent Bonds
 Although
all covalent bonds involve sharing of
electrons, they differ widely in the degree of
sharing.
 We divide covalent bonds into
• nonpolar covalent bonds and
• polar covalent bonds.
D i fference in
El ectron eg ati vity
Betw een Bo nded Ato ms
Less than 0.5
0.5 to 1.9
Greater than 1.9
Typ e of Bond
N on pol ar cov alent
Pol ar co valent
Io ns f orm
1-37
Polar and Nonpolar Covalent Bonds
• An example of a polar covalent bond is that of H-Cl.
• The difference in electronegativity between Cl and H is
3.0 - 2.1 = 0.9.
• We show polarity by using the symbols d+ and d-, or by
using an arrow with the arrowhead pointing toward the
negative end and a plus sign on the tail of the arrow at
the positive end.
d+
H
dCl
H
Cl
1-38
Polar Covalent Bonds
 Bond
dipole moment ():
• A measure of the polarity of a covalent bond.
• The product of the charge on either atom of a polar
bond times the distance between the two nuclei.
• Table 1.7 shows average bond dipole moments of
selected covalent bonds.
Bond
Dipole
Bond (D )
Bond
Dipole
Bond (D )
Bond
D ipole
Bond (D)
H-C
H-N
H-O
H-S
C-F
C-Cl
C-Br
C-I
C-O
C=O
C-N
-C=N
0.3
1.3
1.5
0.7
1.4
1.5
1.4
1.2
0.7
2.3
0.2
3.5
1-39
Lewis Structures
 To
write a Lewis structure
•
•
•
•
Determine the number of valence electrons.
Determine the arrangement of atoms.
Connect the atoms by single bonds.
Arrange the remaining electrons so that each atom has
a complete valence shell.
• Show a bonding pair of electrons as a single line.
• Show a nonbonding pair of electrons (a lone pair) as a
pair of dots.
• In a single bond atoms share one pair of electrons, in a
double bond they share two pairs of electrons and in a
triple bond they share three pairs of electrons.
1-40
Lewis Structures - Table 1.8
H-O-H
H 2 O (8)
Water
H
H
H
C2 H 4 (12)
Ethylene
•
•
•
•
•
H-N-H
H
N H 3 (8)
Ammonia
H
C C
 In
H
H-C-H
H
CH 4 (8)
Meth ane
H-Cl
HCl (8)
Hyd rogen ch loride
O
H
H-C C-H
C O
C2H 2 (10)
Acetylen e
H
CH 2O (12)
Formald ehyde
H
O
C
O
H
H 2CO 3 (24)
Carbonic acid
neutral molecules
hydrogen has one bond.
carbon has 4 bonds and no lone pairs.
nitrogen has 3 bonds and 1 lone pair.
oxygen has 2 bonds and 2 lone pairs.
halogens have 1 bond and 3 lone pairs.
1-41
Formal Charge
 Formal
charge: The charge on an atom in a
molecule or a polyatomic ion.
 To derive formal charge
1. Write a correct Lewis structure for the molecule or ion.
2. Assign each atom all its unshared (nonbonding)
electrons and one-half its shared (bonding) electrons.
3. Compare this number with the number of valence
electrons in the neutral, unbonded atom.
Formal
charge
N umber of
= valence electrons
in th e neutral,
un bonded atom
All
One h alf of
un shared + all sh ared
electrons
electrons
4. The sum of all formal charges is equal to the total
charge on the molecule or ion.
1-42
Apparent Exceptions to the Octet Rule
 Molecules
that contain atoms of Group 3A
elements, particularly boron and aluminum.
B
Boron trifluoride
: Cl
Al
: Cl :
:
:
:F:
: Cl :
: :
: :
:F
:
:
: F:
6 electrons in the
valence shells of boron
and aluminum
Aluminum chloride
1-43
Apparent Exceptions to the Octet Rule
 Atoms
of third-period elements, such as S and P,
are often drawn with more bonds than allowed by
the octet rule.
:
• The P in trimethylphosphine obeys the octet rule by
having three bonds and one unshared pair.
• A common depiction of phosphoric acid, however, has
five bonds to P, which is explained by invoking the use
of 3d orbitals to accommodate the additional bonds.
: Cl :
:O:
: Cl
Cl :
P
H- O-P- O-H
CH3 -P- CH3
Cl :
: Cl
CH3
O-H
:
:
: :
: : : :
: : : :
:
:
Trimethylphosphine
Phosphorus
pentachloride
Phosphoric
acid
1-44
Apparent Exceptions to the Octet Rule
 However,
the use of 3d orbitals for bonding is in
debate.
 An alternative representation that gives P in
phosphoric acid an octet has four bonds and a
positive formal charge on P. The oxygen involved
in the double bond of the alternative depiction
has one bond and a negative formal charge.
O
HO
P
OH
formal
charges
O
OH
HO
P
OH
OH
1-45
Apparent Exceptions to the Octet Rule
H 3C
S
CH3
H 3C
S
:
H
CH3
Dimethyl sulfoxide
:
Hydrogen
sulfide
OH
:O :
HO
:O :
:
HO
: :
: :
S
S +2
Sulfuric acid
::
:O:
:O :
: :
S
formal
charges
:O :
:O :
:
H
: :
:
• Sulfur, another third-period element, is commonly
depicted with varying numbers of bonds. In each of the
alternative structures sulfur obeys the octet rule, and
has one or more positive formal charges.
OH
formal
charges
1-46
Functional Groups
 Functional
group: An atom or group of atoms
within a molecule that shows a characteristic set
of physical and chemical properties.
 Functional groups are important for three
reason; they are:
1. the units by which we divide organic compounds into
classes.
2. the sites of characteristic chemical reactions.
3. the basis for naming organic compounds.
1-47
Alcohols
 Contain
an -OH (hydroxyl) group bonded to a
tetrahedral carbon atom.
H H
:
-C-O-H
:
Fu nctional
group
H-C-C-O-H
H H
Ethan ol
(an alcohol)
 Ethanol
may also be written as a condensed
structural formula.
CH3 -CH2 -OH
or
CH3 CH2 OH
1-48
Alcohols
• Alcohols are classified as primary (1°), secondary (2°),
or tertiary (3°) depending on the number of carbon
atoms bonded to the carbon bearing the -OH group.
H
CH3 -C-OH
H
A 1° alcohol
H
CH3 -C-OH
CH3
A 2° alcohol
CH3
CH3 -C-OH
CH3
A 3° alcohol
1-49
Alcohols
• There are two alcohols with molecular formula C3H8O.
HHH
H-C-C-C-O-H
or CH3 CH2 CH2 OH
H HH
a 1° alcohol
H
HOH
H C-C-C-H
HH H
or
OH
CH3 CHCH3
a 2° alcohol
1-50
Amines
an amino group; an sp3-hybridized
nitrogen bonded to one, two, or three carbon
atoms.
 Contain
• An amine may by 1°, 2°, or 3°.
Methylamine
(a 1° amine)
CH3 N H
CH3
Dimethylamine
(a 2° amine)
:
H
:
:
CH3 N H
CH3 N CH3
CH3
Trimethylamine
(a 3° amine)
1-51
Aldehydes and Ketones
 Contain
a carbonyl (C=O) group.
: O:
C
O
H
CH3 -C- H
Functional Acetaldehyde
(an aldehyde)
group
:O:
O
C
CH3 -C- CH3
Functional Acetone
(a ketone)
group
1-52
Carboxylic Acids
 Contain
a carboxyl (-COOH) group.
:
Fu nctional
group
:O:
CH3 -C-O-H
:
O
C O H
or CH3 COOH or CH3 CO2 H
Acetic aci d
(a carboxy li c acid )
1-53
Carboxylic Esters
 Ester:
A derivative of a carboxylic acid in which
the carboxyl hydrogen is replaced by a carbon
group.
O
C O
Functional
group
O
CH3 - C-O- CH 2 -CH3
Ethyl acetate
(an ester)
1-54
Carboxylic Amide
 Carboxylic
amide, commonly referred to as an
amide: A derivative of a carboxylic acid in which
the -OH of the -COOH group is replaced by an
amine.
O
C N
Fu nctional
group
O
CH3 -C-N-H
H
Acetamid e
(a 1° amid e)
• The six atoms of the amide functional group lie in a
plane with bond angles of approximately 120°.
1-55
Covalent
Bonds &
Shapes of
Molecules
End Chapter 1
1-56