1. science
2. chemistry

Organic
Chemistry
William H. Brown
Christopher S. Foote
Brent L. Iverson
1-1
Covalent
Bonding &
Shapes of
Molecules
Chapter 1
1-2
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
1-3
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
1-4
Electron Configuration of Atoms
 Electrons
are confined to regions of space called
principle energy levels (shells)
• each shell can hold 2n2 electrons (n = 1,2,3,4......)
Number of
Relative Energies
Electrons Shell
of Electrons
Shell
Can Hold
in These Shells
higher
4
32
3
18
2
8
1
2
lower
1-5
Electron Configuration of Atoms
 Shells
are divided into subshells called orbitals,
which are designated by the letters s, p, d, f,........
• s (one per shell)
• p (set of three per shell 2 and higher)
• d (set of five per shell 3 and higher) .....
Shell Orbitals Contained in That Shell
3
3s, 3px , 3py, 3pz, plus five 3d orbitals
2
2s, 2px , 2py, 2pz
1
1s
1-6
Electron Configuration of Atoms
 Aufbau
Principle:
• orbitals fill in order of increasing energy from lowest
energy to highest energy
 Pauli
Exclusion Principle:
• only two electrons can occupy an orbital and their
spins must be paired
 Hund’s
Rule:
• when orbitals of equal energy are available but there
are not enough electrons to fill all of them, one
electron is added to each orbital before a second
electron is added to any one of them
1-7
Electron Configuration of Atoms
 The
pairing of electron spins
1-8
Electron Configuration of Atoms
 Table
1.3 The Ground-State Electron
Configuration of Elements 1-18
1-9
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 valence electrons
1-10
Lewis Dot Structures

Table 1.4 Lewis Dot Structures for Elements 1-18
1A
2A
3A
H.
.
Si : . P :
: O:
.
.
:S :
.
.
:F :
.
He :
:N e :
: :
Al:
.
.
8A
:
.
.
.
C: N :
: Cl : : A r :
:
N a. M g :
B:
.
7A
:
Be :
5A 6A
:
Li .
.
4A
1-11
Lewis Model of Bonding
 Atoms
bond together so that each atom acquires
an electron configuration 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
• 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 bonds are called polar covalent bonds
1-12
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
1-13
Formation of Ions
A
rough guideline:
• ions will form if the difference in electronegativity
between interacting atoms is 1.9 or greater
• example: sodium (EN 0.9) and fluorine (EN 4.0)
• we use a single-headed (barbed) curved arrow to show
the transfer of one electron from Na to F
F
••
+
Na
••
-
•
•
•
•
+
•
•
Na
••
F
••
• in forming Na+F-, the single 3s electron from Na is
transferred to the partially filled valence shell of F
Na(1s22s 22p63s1 ) + F(1s 22s2 2p5 )
Na+(1s2 2s22p6) + F-(1s2 2s2 2p6 )
1-14
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
1-15
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
• polar covalent bonds
Difference in
Electronegativity
Between Bonded Atoms
Less than 0.5
0.5 to 1.9
Greater than 1.9
Type of Bond
Nonpolar covalent
Polar covalent
Ions form
1-16
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+ dH Cl
H
Cl
1-17
Polar Covalent Bonds
 Bond
dipole moment (m):
• 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 nuclei
• Table 1.7 shows average bond dipole moments of
selected covalent bonds
Bond
Dipole
Bond (D)
H-C
H-N
H-O
H-S
0.3
1.3
1.5
0.7
Bond
Dipole
Bond (D)
C-F
C-Cl
C-Br
C-I
1.4
1.5
1.4
1.2
Bond
Dipole
Bond (D)
C-O
C=O
C-N
-C=N
0.7
2.3
0.2
3.5
1-18
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 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
1-19
Lewis Structures - Table 1.3
H-O-H
H2 O (8)
Water
H
H
H
C2 H4 (12)
Ethylene
•
•
•
•
•
H-N-H
H
NH3 (8)
Ammonia
H
C C
 In
H
H-C-H
H
CH4 (8)
Methane
H-Cl
HCl (8)
Hydrogen chloride
O
H
H-C C-H
C O
C2H2 (10)
Acetylene
H
CH2O (12)
Formaldehyde
H
O
C
O
H
H2CO3 (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
1-20
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
Number of
Formal = valence electrons
charge
in the neutral,
unbonded atom
All
One half of
unshared + all shared
electrons
electrons
1-21
Formal Charge

Example: Draw Lewis structures, and show which atom in
each bears the formal charge
(a) NH2
(b) HCO3
(c) CO3
2-
+
(d) CH3 NH3
(e) HCOO
(f) CH3 COO
1-22
Exceptions to the Octet Rule
 Molecules
containing atoms of Group 3A
elements, particularly boron and aluminum
B
Boron trifluoride
: Cl
Al
: Cl :
:
:
:F:
: Cl :
: :
: :
:F
6 electrons in the
valence shells of boron
and aluminum
:
:
: F:
Aluminum chloride
1-23
Exceptions to the Octet Rule
 Atoms
of third-period elements have 3d orbitals
and may expand their valence shells to contain
more than 8 electrons
• phosphorus may have up to 10
:
H- O-P- O-H
:
Cl :
:
Cl :
:O:
: :
: Cl
P
: : : :
O-H
Phosphoric
acid
:
CH3
Trimethylphosphine
: : : :
:
CH3 -P- CH3
: Cl
: Cl :
Phosphorus
pentachloride
1-24
Exceptions to the Octet Rule
• sulfur, another third-period element, forms compounds
in which its valence shell contains 8, 10, or 12
electrons
H-O- S-O-H
:
CH 3 -S-CH 3
:
:
:
:
H-S- H
:
: O:
:
: O:
:O :
Hydrogen
sulfide
Dimethyl
sulfoxide
Sulfuric
acid
1-25
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
1-26
Alcohols
 contain
an -OH (hydroxyl) group
H H
:
-C-O-H
:
Functional
group
H-C-C-O-H
H H
Ethanol
(an alcohol)
 Ethanol
may also be written as a condensed
structural formula
CH3 -CH2 -OH
or
CH3 CH2 OH
1-27
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
CH3 -C-OH
H
CH3 -C-OH
CH3
CH3 -C-OH
CH3
A 1° alcohol
CH3
A 2° alcohol
CH3
A 3° alcohol
CH3
1-28
Alcohols
• there are two alcohols with molecular formula C3H8O
HHH
H-C-C-C-O-H
H HH
H
HOH
H C-C-C-H
HH H
or CH3 CH2 CH2 OH
a 1° alcohol
or
OH
CH 3 CHCH3
a 2° alcohol
1-29
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)
CH 3 N H
CH 3
Dimethylamine
(a 2° amine)
:
H
:
:
CH 3 N H
CH 3 N CH 3
CH 3
Trimethylamine
(a 3° amine)
1-30
Aldehydes and Ketones
 contain
O
a carbonyl (C=O) group
O
C H
CH3 -C-H
Functional Acetaldehyde
group
(an aldehyde)
O
O
C
CH3 -C-CH3
Functional Acetone
group
(a ketone)
1-31
Carboxylic Acids
 contain
a carboxyl (-COOH) group
:
Functional
group
:O:
CH3 -C-O-H
:
O
C O H
or CH3 COOH or CH3 CO2 H
Acetic acid
(a carboxylic acid)
1-32
Carboxylic Esters
 Ester:
a derivative of a carboxylic acid in which
the carboxyl hydrogen is replaced by a carbon
group
O
: O:
:
Functional
group
CH3 -C-O-CH2 -CH3
:
C O
Ethyl acetate
(an ester)
1-33
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
Functional
group
O
CH3 -C-N-H
H
Acetamide
(a 1° amide)
• the six atoms of the amide functional group lie in a
plane with bond angles of approximately 120°
1-34
VSEPR
 Based
on the twin concepts that
• atoms are surrounded by regions of electron density
• regions of electron density repel each other
H
N
H H
H
H
H
H
H
O
H
H
H
H
H
C
H
H
:
N
H
O H C C H H C N
:
O C
H
C
O
:
C
:
C
: :
2 regions of e- density
(linear, 180°)
C
:
3 regions of e- density
(trigonal planar, 120°)
:
4 regions of e- density
(tetrahedral, 109.5°)
H
1-35
VSEPR Model

Example: predict all bond angles for these molecules and
ions
( a ) N H4 +
( b ) CH3 NH 2
( d ) CH3 OH
( e ) CH 3 CH = CH 2
( f ) H 2 CO 3
( g ) HCO 3 -
( h) CH3 CH O
( i) CH 3 COOH
( j ) BF4 -
1-36
Polar and Nonpolar Molecules
 To
determine if a molecule is polar, we need to
determine
• if the molecule has polar bonds
• the arrangement of these bonds in space
 Molecular
dipole moment (m): the vector sum of
the individual bond dipole moments in a
molecule
• reported in debyes (D)
1-37
Polar and Nonpolar Molecules
 these
molecules have polar bonds, but each has
a zero dipole moment
Cl
F
O
C
B
O
F
F
Carbon dioxide
m=0D
Boron trifluoride
m=0D
C
Cl
Cl
Cl
Carbon tetrachloride
m=0D
1-38
Polar and Nonpolar Molecules
 these
molecules have polar bonds and are polar
molecules
direction
of dipole
moment
N
O
H
H
Water
m = 1.85D
H
H
Ammonia
m = 1.47D
H
direction
of dipole
moment
1-39
Polar and Nonpolar Molecules
• formaldehyde has polar bonds and is a polar molecule
direction
of dipole
moment
H
O
C
H
Formaldehyde
m = 2.33 D
1-40
Resonance
 For
many molecules and ions, no single Lewis
structure provides a truly accurate
representation
O
O
and
H3 C C
O
-
H3 C C
O
Ethanoate ion
(acetate ion)
1-41
Resonance
 Linus
Pauling - 1930s
• many molecules and ions are best described by
writing two or more Lewis structures
• individual Lewis structures are called contributing
structures
• connect individual contributing structures by doubleheaded (resonance) arrows
• the molecule or ion is a hybrid of the various
contributing structures
1-42
Resonance

Examples: equivalent contributing structures
CH3
O:
CH3
O:
C
: O :-
:
Nitrite ion
(equivalent contributing
structures)
C
:
:O :-
:
:
O:
: O :-
O:
:N
:
:N
:
:
:
:O: -
Acetate ion
(equivalent contributing
structures)
1-43
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 an atom to an adjacent bond
 Electron
pushing is a survival skill in organic
chemistry
• learn it well!
1-44
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
• 3rd period elements, such as P and S, may have up to
12 electrons in their valence shells
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
1-45
Resonance
 The
carbonate ion, for example
• a hybrid of three equivalent contributing structures
• the negative charge is distributed equally among the
three oxygens
1-46
Resonance
 Preference
1: filled valence shells
• structures in which all atoms have filled valence shells
contribute more than those with one or more unfilled
valence shells
+
CH 3 O
••
C
H
H
Greater contribution;
both carbon and oxygen have
complete valence shells
••
CH 3 O
••
+
C
H
H
Lesser contribution;
carbon has only 6 electrons
in its valence shell
1-47
Resonance
 Preference
2: maximum number of covalent
bonds
• structures with a greater number of covalent bonds
contribute more than those with fewer covalent bonds
CH3
+
O
••
••
C
H
H
Greater contribution
(8 covalent bonds)
CH3 •O
•
+
C
H
H
Lesser contribution
(7 covalent bonds)
1-48
Resonance
 Preference
3: least separation of unlike charge
• structures with separation of unlike charges contribute
less than those with no charge separation
Greater contribution
(no separation of
unlike charges)
:
:
O:
CH3 -C- CH3
:O: -
CH3 -C- CH3
Lesser contribution
(separation of unlike
charges)
1-49
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 the less electronegative atom
O
(1)
C
H3 C
O
O
CH3
(a)
Lesser
contribution
(2)
C
H3 C
CH3
(b)
Greater
contribution
C
H3 C
CH3
(c)
Should not
be drawn
1-50
Quantum or Wave Mechanics


Albert Einstein: E = hn (energy is quantized)
• light has particle properties
Louis deBroglie: wave/particle duality
h
 mn

Erwin Schrödinger: wave equation
• wave function, : a solution to a set of equations that
depicts the energy of an electron in an atom
• each wave function is associated with a unique set of
quantum numbers
• each wave function occupies three-dimensional space
and is called an orbital
•  2 is the probability of finding an electron at a given
point in space
1-51
Shapes of 1s and 2s Orbitals
distribution (2) for 1s and 2s orbitals
showing an arbitrary boundary surface
containing about 95% of the electron density
 Probability
1-52
Shapes of a Set of 2p Atomic Orbitals
 Three-dimensional
shapes of 2p atomic orbitals
1-53
Molecular Orbital Theory
 Electrons
in atoms exist in atomic orbitals
 Electrons in molecules exist in molecular orbitals
(MOs)
 Using the Schrödinger equation, we can
calculate the shapes and energies of MOs
1-54
Molecular Orbital Theory
 Rules:
• combination of n atomic orbitals (mathematically
adding and subtracting wave functions) gives n MOs
(new wave functions)
• MOs are arranged in order of increasing energy
• MO filling is governed by the same rules as for atomic
orbitals:
• Aufbau principle: fill beginning with LUMO
• Pauli exclusion principle: no more than 2e- in a MO
• Hund’s rule: when two or more MOs of equivalent
energy are available, add 1e- to each before filling
any one of them with 2e1-55
Molecular Orbital Theory
 Terminology
•
•
•
•
•
•
•
•
ground state = lowest energy state
excited state = NOT lowest energy state
 = sigma bonding MO
* = sigma antibonding MO
 = pi bonding MO
* = pi antibonding MO
HOMO = highest occupied MO
LUMO = lowest unoccupied MO
1-56
Molecular Orbital Theory
 Sigma
1s bonding and antibonding MOs
1-57
Molecular Orbital Theory
• MO energy diagram for H2: (a) ground state and (b)
lowest excited state
1-58
Molecular Orbitals
• computed sigma bonding and antibonding MOs for H2
1-59
Molecular Orbitals
 pi
bonding and antibonding MOs
1-60
Molecular Orbitals
• computed pi bonding and antibonding MOs for
ethylene
1-61
Molecular Orbitals
• computed pi bonding and antibonding orbitals for
formaldehyde
1-62
Hybrid Orbitals
 The
Problem:
• bonding by 2s and 2p atomic orbitals would give bond
angles of approximately 90°
• instead we observe bond angles of approximately
109.5°, 120°, and 180°
A
Solution
• hybridization of atomic orbitals
• 2nd row elements use sp3, sp2, and sp hybrid orbitals
for bonding
1-63
Hybrid Orbitals
 Hybridization
of orbitals (L. Pauling)
• the combination of two or more atomic orbitals forms a
new set of atomic orbitals, called hybrid orbitals
 We
deal with three types of hybrid orbitals
sp3 (one s orbital + three p orbitals)
sp2 (one s orbital + two p orbitals)
sp (one s orbital + one p orbital)
 Overlap
of hybrid orbitals can form two types of
bonds depending on the geometry of overlap
 bonds are formed by “direct” overlap
 bonds are formed by “parallel” overlap
1-64
sp3 Hybrid Orbitals
• each sp3 hybrid orbital
has two lobes of
unequal size
• the sign of the wave
function is positive in
one lobe, negative in the
other, and zero at the
nucleus
• the four sp3 hybrid
orbitals are directed
toward the corners of a
regular tetrahedron at
angles of 109.5°
1-65
sp3 Hybrid Orbitals
• orbital overlap pictures of methane, ammonia, and
water
1-66
sp2 Hybrid Orbitals
• the axes of the three sp2 hybrid orbitals lie in a plane
and are directed toward the corners of an equilateral
triangle
• the unhybridized 2p orbital lies perpendicular to the
plane of the three hybrid orbitals
1-67
Bonding in Ethylene
1-68
Bonding in Formaldehyde
1-69
sp Hybrid Orbitals
• two lobes of unequal size at an angle of 180°
• the unhybridized 2p orbitals are perpendicular to each
other and to the line created by the axes of the two sp
hybrid orbitals
1-70
Bonding in Acetylene, C2H2
1-71
Hybrid Orbitals
Groups Orbital Predicted
Bonded HybridBond
to Carbon ization
Angles
4
sp 3
109.5°
Types of
Bonds
to Carbon
Example
Name
4 sigma bonds
HH
H-C-C-H
Ethane
HH
2
2
sp
sp
2
120°
180°
3 sigma bonds
and 1 pi bond
2 sigma bonds
and 2 pi bonds
H
H
C
H
C
Ethylene
H
H-C C-H
Acetylene
1-72
Bond Lengths and Bond Strengths
Orbital
Overlap
Bond Length
(pm)
Bond Strength
[kJ (kcal)/mol]
Name
Formula
Bond
Ethane
HH
H-C-C-H
C-C
sp3 -sp 3
153.2
376 (90)
C-H
sp3 -1s
111.4
422 (101)
H
C-C
sp2-sp2, 2p-2p
133.9
727 (174)
H
C-H
sp 2-1s
110.0
464(111)
C-C
sp-sp, two 2p-2p
C-H
sp-1s
121.2
109.0
966 (231)
556 (133)
HH
H
Ethylene
C C
H
Acetylene H-C C-H
1-73
Covalent
Bonds &
Shapes of
Molecules
End Chapter 1
1-74