1. Structure and Bonding
Why this chapter?
• Review ideas from general chemistry: atoms, bonds,
molecular geometry
What is Organic Chemistry?
Living things are made of organic
chemicals
 Proteins that make up hair
 DNA, controls genetic make-up
 Foods, medicines
 Examine structures below

2
Origins of Organic Chemistry
Foundations of organic chemistry from mid-1700’s.
Compounds obtained from plants, animals hard to isolate,
and purify.
Compounds also decomposed more easily.
Torben Bergman (1770) first to make distinction between
organic and inorganic chemistry.
It was thought that organic compounds must contain
some “vital force” because they were from living
sources.
3
Because of “Vital force”, it was thought that organic
compounds could not be synthesized in laboratory like
inorganic compounds.
1816, Chervrut showed that not to be the case, he could
prepare soap from animal fat and an alkali
1828, Wholer showed that it was possible to convert
inorganic salt ammonium cyanate into organic substance
“urea”
4
• Organic
chemistry is study of carbon
compounds.
• Why is it so special?
- 90% of more than 30 million chemical
compounds contain carbon.
- Examination of carbon in periodic chart
answers some of these questions.
- Carbon is group 4A element, it can share 4
valence electrons and form 4 covalent
bonds.
5
1.1 Atomic Structure

Structure of an atom
◦ Positively charged nucleus (very dense, protons and neutrons) and
small (10-15 m)
◦ Negatively charged electrons are in a cloud (10-10 m) around
nucleus

Diameter is about 2  10-10 m (200 picometers (pm)) [the unit
angstrom (Å) is 10-10 m = 100 pm]
6
The Atomic Symbol
A = Atomic mass
=# p
C = Charge
+#n
= + or - values
A
Z
Z = Atomic #
#p=#e
C
X
#
# = Number of atoms
in a formula.
The Atomic Symbol
A = Atomic mass
6
6
= # protons
+ # neutrons
12
6
Z = Atomic number
X
- +
+
-
= # protons = # electrons
+
-
+
+
+
The Atomic Symbol
A = Atomic mass
= # protons + # neutrons
12
6
C
-
-
+
+
Z = Atomic number
+
-
= # protons = # electrons
+
+
-
+
The Atomic Symbol
A = Atomic mass
C = Charge = +1
= 11
p + 12
n = 23
23
1+
Na
11
Z = Atomic #
= p = 11
Sodium
# = 1 atom in formula.
Why is the atomic weight on the tables not a whole #?
47
Silver
Ag
107.87
Atomic number
Name of the element
Elemental Symbol
Atomic mass (weight)
Atomic weight = The average, relative
mass of an atom in an element.
Isotopes of Hydrogen

Isotopes = Atoms of the same element but
having different masses.
1
2
3
1H
1 H
1 H
+
Protium
99.99%
-
+
Deuterium
0.01%
+
Tritium
Trace %
Isotopes of Hydrogen

Isotopes = Atoms of the same element but
having different masses.
1
2
3
1H
1 H
1 H
+
-
+
+
Average Atomic weight of Hydrogen
= 1.00794 amu
Isotopes of Carbon
12
C
6
-
-
+
+
13
C
6
-
-
-
+
+
+
+ +
-
98.89%
-
+
+
+ +
-
-
+
14
C
6
- -
1.11%
+
-
-
+
-
+
+
+ +
Trace %
Average Atomic weight of C= 12.011 amu
-
Radioactive Isotopes
3
1
H-3
H
C-14
-
-
-
+
+
Nucleus is unstable
So falls apart (decays)
Giving radioactive particles
14
6
+
C
-
+
+
+ +
-
-
Electronic arrangement
A new layer is
added for each
row or period in
the table.
Electron arrangement
24
Mg
12
32
18
2
8
Shells = Energy levels
Electrons
fill layers
around
nucleus
Low  High
IA
IIA
1
H
1
7
Li
3
4
He
2
9
Be
4
2, 1
2, 2
IA
IIA
IIIA
1
H
1
7
Li
3
9
Be
4
2, 1
2, 2
11
5
B
2, 3
IIIA
11
5
IVA
12
6
B
2, 3
VA
13
7
C
2, 4
N
2, 5
IA
IIA
1
H
1
2, 1
23
11
4
He
2
9
Be
4
7
Li
3
Na
2, 8, 1
VIIIA
2, 2
24
12Mg
2, 8, 2
20
10
Ne
2, 8
40
18Ar
2, 8, 8
1
1
H
1
2
9
Be
4
7
Li
3
2, 1
23
11
Valence electrons
Where most chemical
Reactions occur.
Na
2, 8, 1
3
11
B
5
2, 2
24
12Mg
2, 8, 2
2, 3
27
13Al
2, 8, 3
1
8
1
H
1
2
9
Be
4
7
Li
3
2, 1
23
11
Octet Rule
Na
2, 8, 1
2, 2
24
12Mg
2, 8, 2
4
He
2
20
10
Ne
2, 8
40
18Ar
2, 8, 8
The octet rule
Atoms are most stable if they have a filled
or empty outer layer of electrons.
Except for H and He, a filled layer
contains 8 electrons - an octet.
Atoms gain, lose or share electrons to
make a filled or empty outer layer.
Atoms gain, lose or share electrons based
on what is easiest.
1.2 Atomic Structure: Orbitals

Quantum mechanics: describes electron
energies and locations by a wave equation
◦ Wave function solution of wave equation
◦ Each wave function is an orbital,
26
Orbitals

A plot of y 2 describes where electron most
likely to be

Electron cloud has no specific boundary so we
show most probable area
s (1)
p (3)
d (5)
f (7)
An f orbital


4 different kinds of orbitals for e-s
s and p orbitals most important in organic & biochem
Orbitals and Shells



Orbitals are grouped in shells of increasing size and energy
Different shells contain different numbers and kinds of orbitals
Each orbital can be occupied by two electrons
28
Orbitals and Shells



1st shell contains one s orbital, denoted 1s, holds 2 e-s
2nd shell contains one s orbital (2s) and three p orbitals
(2p), 8 electrons
3rd shell contains an s orbital (3s), three p orbitals (3p),
and five d orbitals (3d), 18 electrons
29
p-Orbitals

In each shell there are three perpendicular p orbitals,
px, py, and pz, of equal energy

Lobes of a p orbital are separated by region of zero
electron density, a node
30
1.3 Atomic Structure: Electron Configurations

Ground-state electron configuration (lowest energy
arrangement) of an atom lists orbitals occupied by its electrons.
Rules:

1. Lowest-energy orbitals fill first: 1s  2s  2p  3s  3p  4s
 3d (Aufbau (“build-up”) principle)

2. Electrons act as if they were spinning around an axis. Electron
spin can have only two orientations, up  and down . Only two
electrons can occupy an orbital, and they must be of opposite spin
(Pauli exclusion principle) to have unique wave equations

3. If two or more empty orbitals of equal energy are available,
electrons occupy each with spins parallel until all orbitals have one
electron (Hund's rule).
31
1.3 Atomic Structure: Electron Configurations
The Aufbau (“build-up”) principle
 Electrons fill from the
f
low  high.
fill n = 1 before n = 2 ,
fill s before p ...
d
n=4
p
s
d
Pauli Exclusion principle
 Each subshell contains orbitals
which can hold a max of 2 e-s
 e’s in same orbital must be
of opposite spins: 1  & 1 .
n=3
p
s
p
s
s
n=2
n=1
Hund’s Rule
 Electrons don’t share same orbital unless they need to.
(i.e. no pairing until each orbital of the set has an e-)
5d
Major trends in electron filling

Exceptions:
4f
◦ Fill 4s before 3d
◦ Fill 5s before 4 d
◦ Fill 5p before 4f
5p
4d
5s
4p
This is why
transition metals
are assigned as
B group elements.
3d
4s
3p
3s
2p
2s
1s
Electron Configuration
6s __
5s __
5d __ __ __ __ __
4f __ __ __ __ __ __ __
5p __ __ __
7
4d __ __ __ __ __ Li
3
4p __ __ __
3d __ __ __ __ __
4s __
3px __ y__ z __
1s22s1
3s __
2px __ y__ z __
2s __
1s __
Electron Configuration
6s __
5s __
5d __ __ __ __ __
4f __ __ __ __ __ __ __
5p __ __ __
16
4d __ __ __ __ __8 O
4p __ __ __
3d __ __ __ __ __
4s __
3p __ __ __
1s22s22p4
3s __
2px __ y __z __
2s __
1s __
Electron Configuration
6s __
5s __
5d __ __ __ __ __
4f __ __ __ __ __ __ __
5p __ __ __
4d __ __ __ __ __
30Zn
4p __ __ __
3d __ __ __ __ __
4s __
3px __ y __z __
3s __
1s22s22p63s23p64s23d10
2px __ y__ z __
[Ar] 4s23d10
2s __
1s __
[Ar] 3d104s2
s
Classification by sublevels
p
H
He
d
Li Be
Na Mg
K Ca Sc Ti
Rb Sr
Y
V
B
C
N
O
F
Ne
Al
Si
P
S
Cl
Ar
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
Cs Ba Ls Hf
Ta W Re Os
Ir
I
Xe
Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
f
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
3 - 25
Using the periodic table
to find sublevels
s
1
H
He
2
Li
Be
3
Na Mg
4
p
d
2
B C N O F Ne
3
Al Si P S Cl Ar
Ca
3
Sc
Ti V Cr Mn Fe Co Ni Cu Zn
4
Ga Ge As Se Br Kr
5
Rb Sr
4
Y
Zr Nb Mo Tc Ru Rh Pd Ag Cd
5
In Sn Sb Te I Xe
6
Cs Ba
5
Ls
Hf Ta W Re Os Ir Pt Au Hg
6
Tl Pb Bi Po At Rn
7
Fr
6
Ac
K
Ra
4
f
5
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
Learning Check:
Which of the following represents the ground-state
electronic configuration for carbon (atomic number 6)?
1.
2.
3.
4.
5.
1s22s22px22py02pz0
1s22s22px12py12pz0
1s22s12px12py1 2pz1
1s12s22px12py1 2pz1
1s22s22px02py02pz2
39
Solution:
Which of the following represents the ground-state
electronic configuration for carbon (atomic number 6)?
1.
2.
3.
4.
5.
1s22s22px22py02pz0
1s22s22px12py12pz0
1s22s12px12py1 2pz1
1s12s22px12py1 2pz1
1s22s22px02py02pz2
40
1
H
1
Octet Rule
1s2
1s1
9
Be
4
7
Li
3
1s2,
23
Na
11
4
He
2
20
10Ne
1s2, 2s2
2s1
24
12
1s2, 2s2 2p6
40
18
Mg
Ar
1s2, 2s2 2p6, 3s2 3p6
1s2, 2s2 2p6, 3s1
1s2, 2s2 2p6, 3s2
[Ne] 3s1
1
H
1
7
Li
3
23
Na
11
Lewis Structures
H
Li
Na
K
Show only
Valence
Electrons
Nonmetals Share e-s
1
H
with other nonmetals
2
3
4
5
Be
B
C
N O
Na Mg
Al
Si P
Li
6
S
7
8
He
F Ne
Cl Ar
Metals give e-s to nonmetals
K
Ca Ga Ge As Se Br Kr
Common ions
Representative Elements
1+
H
4+
2+
4-
3+
Transition Elements
Li Be
3-
2-
1-
He
Ne
B
C
N
O
F
Al
Si
P
S
Cl Ar
Variable
Na Mg
K
Ca Sc Ti
Rb Sr
Y
V
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
Cs Ba Ls Hf Ta W Re Os Ir
Fr Ra Ac
I
Xe
Pt Au Hg Tl Pb Bi Po At Rn
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
Formation of NaCl
e- moves from Metal  Nonmetal
Stable octets
Na + Cl
+
Na
Metal
Cation
_
+ Cl
Nonmetal
Anion
+ and - ions attract to form an ionic bond.
Ionic compounds
•
•
•
•
Not individual molecules
Form crystal arrays
Ions touch many others
Formula represents the average ion ratio
Na Cl
Cl Na Cl Na
NaCl
sodium chloride
Nonmetals Share e-s
1
H
with other nonmetals
2
3
4
5
Be
B
C
N O
Na Mg
Al
Si P
Li
6
S
7
8
He
F Ne
Cl Ar
Metals give e-s to nonmetals
K
Ca Ga Ge As Se Br Kr
Covalent Bonds
H
+
H
H H
Cl
+
Cl
Cl Cl
O
+
O
O O
N
+
N
N
N
Covalent Bonds
H H
H-H
H2
Cl Cl
Cl-Cl
Cl2
O O
O=O
O2
N
N
N N N2
Covalent Bonds
Carbon monoxide
C
O
CO
C O
CO2
O=C=O
Carbon dioxide
O C O
May modify rules to improve sound.
ie - monoxide not monooxide.
Review:
Naming covalent compounds
CO
carbon monoxide
CO2
carbon dioxide
N 2O 5
dinitrogen pentoxide
SiO2
silicon dioxide
ICl3
iodine trichloride
P2O5
diphophorous pentoxide
carbon tetrachloride
CCl4
May modify rules to improve the sound.
Example - use monoxide not monooxide.
Shape
CO2
O C
O
O=C=O
Linear Shape (180o)
Shape
BF3
F
F
F B F
B
F
F
(120o)
Trigonal Planar
1.4 Development of Chemical Bonding Theory

Kekulé and Couper independently observed that carbon
always has four bonds

van't Hoff and Le Bel proposed that the four bonds of
carbon have specific spatial directions
◦ Atoms surround carbon as corners of a tetrahedron
60
Shape
e-’s in 2 directions = 180o
O=C=O
Linear
e-’s in 3 directions = 120o
O
H
C
Trigonal planar
H
Shape
e-’s in 4 directions = 109.5o
H N H H N H
H
H
H
H
N
H
Pyramidal
(109.5o)
Tetrahedral Configuration of Electrons
Trigonal Pyramid Configuration of Atoms
Tetrahedral electron-pair
Geometries
Tetrahedral
Pyramidal
Bent
71
Non-bonding electrons

Valence electrons not used in bonding are called
nonbonding electrons, or lone-pair electrons
◦ Nitrogen atom in ammonia (NH3)
 Shares six valence electrons in three covalent bonds
and remaining two valence electrons are nonbonding
lone pair
72
Learning Check

How many hydrogen atoms does phosphorus
bond to in forming phosphine, PH?
Hint: Draw the electron dot structure and then
use Hydrogen to provided the needed bonded
electrons
73
Solution

How many hydrogen atoms does phosphorus
bond to in forming phosphine, PH?
H
H P
H
H
3
H P
H
Hint: Draw the electron dot structure and then
use Hydrogen to provided the needed bonded
electrons
74
Learning Check

How many hydrogen atoms does Carbon need
in forming chloroform, CH?Cl
H
C
Cl
75
Solution

How many hydrogen atoms does Carbon need
in forming chloroform, CH?Cl
H
H
Cl
C
H
H
Cl
C
3
H
H
H C Cl
C
H
HH
Cl
76
1.5 The Nature of Chemical Bonds:
Valence Bond Theory


Covalent bond forms when two atoms
approach each other closely so that a singly
occupied orbital on one atom overlaps a singly
occupied orbital on the other atom
Two models to describe covalent bonding.
Valence bond theory, Molecular orbital theory
Valence Bond Theory:
 Electrons are paired in the overlapping orbitals
and are attracted to nuclei of both atoms
◦ H–H bond results from the overlap of two
singly occupied hydrogen 1s orbitals
◦ H-H bond is cylindrically symmetrical, sigma
(s) bond
77
Bond Energy


Reaction 2 H·  H2 releases 436 kJ/mol
Product has 436 kJ/mol less energy than two atoms:
H–H has bond strength of 436 kJ/mol.
(1 kJ = 0.2390 kcal; 1 kcal = 4.184 kJ)
78
Bond Length



Distance between
nuclei that leads to
maximum stability
If too close, they
repel because both
are positively
charged
If too far apart,
bonding is weak
79
1.6 sp3 Orbitals and the Structure of Methane
Carbon has 4 valence electrons (2s2 2p2)
 In CH4, all C–H bonds are identical (tetrahedral)
 sp3 hybrid orbitals: s orbital and three p orbitals
combine to form four equivalent, unsymmetrical,
tetrahedral orbitals (sppp = sp3), Pauling (1931)

80
The Structure of Methane



sp3 orbitals on C overlap with 1s orbitals on 4 H atoms
to form four identical C-H bonds
Each C–H bond has a strength of 436 (438) kJ/mol and
length of 109 pm
Bond angle: each H–C–H is 109.5°, the tetrahedral angle.
81
1.7 sp3 Orbitals and the Structure of Ethane





Two C’s bond to each other by s overlap of an sp3 orbital
from each
Three sp3 orbitals on each C overlap with H 1s orbitals to
form six C–H bonds
C–H bond strength in ethane 423 kJ/mol
C–C bond is 154 pm long and strength is 376 kJ/mol
All bond angles of ethane are tetrahedral
82
1.8 sp2 Orbitals and the Structure of
Ethylene



sp2 hybrid orbitals: 2s orbital combines with two 2p
orbitals, giving 3 orbitals (spp = sp2). This results in a
double bond.
sp2 orbitals are in a plane with120° angles
Remaining p orbital is perpendicular to the plane
83
Bonds From sp2 Hybrid Orbitals
Two sp2-hybridized orbitals overlap to form a s bond
 p orbitals overlap side-to-side to formation a pi () bond
 sp2–sp2 s bond and 2p–2p  bond result in sharing four
electrons and formation of C-C double bond
 Electrons in the s bond are centered between nuclei
 Electrons in the  bond occupy regions are on either side
of a line between nuclei

84
Structure of Ethylene




H atoms form s bonds with four sp2 orbitals
H–C–H and H–C–C bond angles of about 120°
C–C double bond in ethylene shorter and stronger than single bond
in ethane
Ethylene C=C bond length 134 pm (C–C 154 pm)
85
1.9 sp Orbitals and the Structure of
Acetylene
C-C a triple bond sharing six electrons
 Carbon 2s orbital hybridizes with a single p orbital giving
two sp hybrids
◦ two p orbitals remain unchanged
 sp orbitals are linear, 180° apart on x-axis
 Two p orbitals are perpendicular on the y-axis and the zaxis

86
Orbitals of Acetylene


Two sp hybrid orbitals from each C form sp–sp s bond
pz orbitals from each C form a pz–pz  bond by sideways
overlap and py orbitals overlap similarly
87
Bonding in Acetylene

Sharing of six electrons forms C C

Two sp orbitals form s bonds with hydrogens
88
89
1.10 Hybridization of Nitrogen and
Oxygen
Elements other than C can have hybridized orbitals
 H–N–H bond angle in ammonia (NH3) 107.3°
 C-N-H bond angle is 110.3 °
 N’s orbitals (sppp) hybridize to form four sp3 orbitals
 One sp3 orbital is occupied by two nonbonding electrons,
and three sp3 orbitals have one electron each, forming
bonds to H and CH3.

90
1.11 Molecular Orbital Theory

A molecular orbital (MO): where electrons are most likely to
be found (specific energy and general shape) in a molecule

Additive combination (bonding) MO is lower in energy

Subtractive combination (antibonding) MO is higher energy
91
Molecular Orbitals in Ethylene



The  bonding MO is from combining p orbital lobes with
the same algebraic sign
The  antibonding MO is from combining lobes with
opposite signs
Only bonding MO is occupied
92
1.12 Drawing Structures
Drawing every bond in organic molecule
can become tedious.
 Several shorthand methods have been
developed to write structures.
 Condensed structures don’t have C-H or
C-C single bonds shown. They are
understood.
e.g.

93
3 General Rules:
1) Carbon atoms aren’t usually shown. Instead a carbon
atom is assumed to be at each intersection of two lines
(bonds) and at the end of each line.
2) Hydrogen atoms bonded to carbon aren’t shown.
3) Atoms other than carbon and hydrogen are shown (See
table 1.3).
94
Summary






Organic chemistry – chemistry of carbon compounds
Atom: positively charged nucleus surrounded by negatively
charged electrons
Electronic structure of an atom described by wave equation
◦ Electrons occupy orbitals around the nucleus.
◦ Different orbitals have different energy levels and different
shapes
 s orbitals are spherical, p orbitals are dumbbell-shaped
Covalent bonds - electron pair is shared between atoms
Valence bond theory - electron sharing occurs by overlap of
two atomic orbitals
Molecular orbital (MO) theory, - bonds result from
combination of atomic orbitals to give molecular orbitals, which
belong to the entire molecule
95
Summary (cont’d)




Sigma (s) bonds - Circular cross-section and are formed by
head-on interaction
Pi () bonds – “dumbbell” shape from sideways interaction of p
orbitals
Carbon uses hybrid orbitals to form bonds in organic molecules.
◦ In single bonds with tetrahedral geometry, carbon has four sp3
hybrid orbitals
◦ In double bonds with planar geometry, carbon uses three
equivalent sp2 hybrid orbitals and one unhybridized p orbital
◦ Carbon uses two equivalent sp hybrid orbitals to form a triple
bond with linear geometry, with two unhybridized p orbitals
Atoms such as nitrogen and oxygen hybridize to form strong,
oriented bonds
◦ The nitrogen atom in ammonia and the oxygen atom in water
are sp3-hybridized
96
Learning Check:
Based on the octet rule, which line-bond structures
is/are correct?
H H
H C C C H
H
A
1.
2.
3.
4.
5.
H
H C C C H
H
B
H H
H C C C H
H
C
H H H
H C C C H
H
D
A and B
A and C
B and D
Only B
Only C
97
Solution:
Based on the octet rule, which line-bond structures
is/are correct?
H H
H C C C H
H
A
1.
2.
3.
4.
5.
H
H C C C H
H
B
H H
H C C C H
H
C
H H H
H C C C H
H
D
A and B
A and C
B and D
Only B
Only C
98
Learning Check:
Which of these statements concerning p-orbitals is
false?
1.
2.
3.
4.
5.
They consist of two equivalent lobes.
They are absent from the first shell of atomic
orbitals.
They can form π bonds.
They only participate in bonding on carbon
atoms.
They can hold a maximum of two electrons.
99
Solution:
Which of these statements concerning p-orbitals is
false?
1.
2.
3.
4.
5.
They consist of two equivalent lobes.
They are absent from the first shell of atomic
orbitals.
They can form π bonds.
They only participate in bonding on carbon
atoms.
They can hold a maximum of two electrons.
100
Learning Check:
What is the hybridization of the oxygen-bonded
carbon atom in the following molecule?
H3C
1.
2.
3.
4.
5.
s
p
sp
sp2
sp3
N C O
101
Solution:
What is the hybridization of the oxygen-bonded
carbon atom in the following molecule?
H3C
1.
2.
3.
4.
5.
s
p
sp
sp2
sp3
N C O
102
Learning Check:
What is the hybridization of nitrogen atom in
trimethylamine :N(CH3)3?
1.
2.
3.
4.
5.
sp
sp2
sp3
dsp2
unhybridized
103
Solution:
What is the hybridization of nitrogen atom in
trimethylamine :N(CH3)3?
1.
2.
3.
4.
5.
sp
sp2
sp3
dsp2
unhybridized
104
Learning Check:
What type of orbital is not used in constructing the
molecule shown below?
O
1.
2.
3.
4.
5.
s
p
sp
sp2
sp3
O
105
Solution:
What type of orbital is not used in constructing the
molecule shown below?
O
1.
2.
3.
4.
5.
s
p
sp
sp2
sp3
O
106
Learning Check:
What is the molecular formula of the molecule
shown below?
O
1.
2.
3.
4.
5.
C8H12O
C8H11O
C8H7O
C7H7O
C7H10O
107
Solution:
What is the molecular formula of the molecule
shown below?
O
1.
2.
3.
4.
5.
C8H12O
C8H11O
C8H7O
C7H7O
C7H10O
108
Learning Check:
Electrons cannot occupy an antibonding molecular
orbital.
1.
2.
True
False
109
Solution:
Electrons cannot occupy an antibonding molecular
orbital.
1.
2.
True
False
110
Learning Check:
How many hydrogen atoms are present in the naturallyoccurring terpene α-terpinene shown below?
1.
2.
3.
4.
5.
14
15
16
17
18
111
Solution:
How many hydrogen atoms are present in the naturallyoccurring terpene α-terpinene shown below?
1.
2.
3.
4.
5.
14
15
16
17
18
112
Learning Check:
Select the best condensed structural formula
for the following molecule:
HO OH
H
O
1.
2.
3.
4.
5.
(CH3)2CHCH2COHOHCOH
CH3CH3CHCH2C(OH)2CHO
(CH3)2CHCH2C(OH)2CHO
(CH3)2CHCH2C(OH)2COH
CH3CHCH3CH2C(OH)2CHO
113
Solution:
Select the best condensed structural formula
for the following molecule:
HO OH
H
O
1.
2.
3.
4.
5.
(CH3)2CHCH2COHOHCOH
CH3CH3CHCH2C(OH)2CHO
(CH3)2CHCH2C(OH)2CHO
(CH3)2CHCH2C(OH)2COH
CH3CHCH3CH2C(OH)2CHO
114
Learning Check:
What is the correct order of carbon-carbon bond
lengths in ethane, ethylene and acetylene?
1.
2.
3.
4.
5.
6.
ethane < ethylene < acetylene
ethane < acetylene < ethylene
ethylene < ethane < acetylene
ethylene < acetylene < ethane
acetylene < ethane < ethylene
acetylene < ethylene < ethane
115
Solution:
What is the correct order of carbon-carbon bond
lengths in ethane, ethylene and acetylene?
1.
2.
3.
4.
5.
6.
ethane < ethylene < acetylene
ethane < acetylene < ethylene
ethylene < ethane < acetylene
ethylene < acetylene < ethane
acetylene < ethane < ethylene
acetylene < ethylene < ethane
116
Learning Check:
What kind of orbital and with how many nodal
planes is shown in the picture below?
1.
2.
3.
4.
5.
σ antibonding, two nodal planes
σ antibonding, one nodal plane
π bonding, one nodal plane
π antibonding, two nodal planes
π antibonding, one nodal plane
117
Solution:
What kind of orbital and with how many nodal
planes is shown in the picture below?
1.
2.
3.
4.
5.
σ antibonding, two nodal planes
σ antibonding, one nodal plane
π bonding, one nodal plane
π antibonding, two nodal planes
π antibonding, one nodal plane
118
Learning Check:
What is the best description of bond length?
1.
2.
3.
4.
5.
a distance between nuclei that yields the best orbital
overlap
a distance between nuclei that yields the smallest
nuclear-nuclear repulsion
a distance between nuclei that yields the smallest
electron-electron repulsion
a distance between nuclei that yields the largest
electron-nuclei attraction
a distance between nuclei that is a compromise of all
electrostatic interactions
119
Solution:
What is the best description of bond length?
1.
2.
3.
4.
5.
a distance between nuclei that yields the best orbital
overlap
a distance between nuclei that yields the smallest
nuclear-nuclear repulsion
a distance between nuclei that yields the smallest
electron-electron repulsion
a distance between nuclei that yields the largest
electron-nuclei attraction
a distance between nuclei that is a compromise of all
electrostatic interactions
120
Learning Check:
What is the O-C-O bond angle in potassium
carbonate, K2CO3?
1.
2.
3.
4.
5.
60°
90°
109.5°
120°
180°
121
Solution:
What is the O-C-O bond angle in potassium
carbonate, K2CO3?
1.
2.
3.
4.
5.
60°
90°
109.5°
120°
180°
122
Learning Check:
What is the relative orientation of two π
orbitals in acetylene?
1.
2.
3.
4.
5.
60°
90°
109.5°
120°
180°
123
Solution:
What is the relative orientation of two π
orbitals in acetylene?
1.
2.
3.
4.
5.
60°
90°
109.5°
120°
180°
124
Learning Check:
In relation to σ bonds, which statement about π
bonds is correct?
1.
2.
3.
4.
5.
π bonds are weaker and π electrons have higher energy
π bonds are stronger and π electrons have higher
energy
π bonds are weaker and π electrons have lower energy
π bonds are stronger and π electrons have lower
energy
π bonds and σ bonds have equal strengths, and their
electrons have the same energies
125
Solution:
In relation to σ bonds, which statement about π
bonds is correct?
1.
2.
3.
4.
5.
π bonds are weaker and π electrons have higher energy
π bonds are stronger and π electrons have higher
energy
π bonds are weaker and π electrons have lower energy
π bonds are stronger and π electrons have lower
energy
π bonds and σ bonds have equal strengths, and their
electrons have the same energies
126
Learning Check:
Which of the orbitals shown below can be occupied in a
hydrogen atom?
A
1.
2.
3.
4.
5.
B
C
D
A
B
C
D
Any one of these
127
Solution:
Which of the orbitals shown below can be occupied in a
hydrogen atom?
A
1.
2.
3.
4.
5.
B
C
D
A
B
C
D
Any one of these
128
Learning Check:
Which statement about the antibonding orbital in
H2 is false?
1.
2.
3.
4.
5.
It is higher in energy than the corresponding
bonding orbital
It is a molecular orbital
It has a nodal plane (a node for short)
It can hold up to two electrons
The two lobes have different charges
129
Solution:
Which statement about the antibonding orbital in
H2 is false?
1.
2.
3.
4.
5.
It is higher in energy than the corresponding
bonding orbital
It is a molecular orbital
It has a nodal plane (a node for short)
It can hold up to two electrons
The two lobes have different charges
130
Learning Check:
What is the hybridization of nitrogen in :NCCH3?
1.
2.
3.
4.
5.
sp
sp2
sp3
sp3d
nitrogen is not hybridized
131
Solution:
What is the hybridization of nitrogen in :NCCH3?
1.
2.
3.
4.
5.
sp
sp2
sp3
sp3d
nitrogen is not hybridized
132
Each shell (floor of the Hotel)
Has subshells (s,p,d,f)
Hotel Model
f
n = 4 (4th floor)
d
p
32 es
n = 3 (3rd floor)
d
18 e-
p
s
n = 2 (2nd floor)
p
8 es
n = 1 (1st floor)
s
2 e-