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The Basics
Bonding and Structure I
1 Energies of Electrons in Atoms and Molecules
• Bonding, structure, shape and reactions of organic molecules are determined PRIMARILY by the Energies
of the electrons in ATOMIC and MOLECULAR ORBITALS
• In organic chemistry we use MODELS, generally the most useful model is the simplest one that explains what
we are trying to understand, in this case the relative energies of elecrons
• Here we introduce a simple model that summarizes the basic factors that determine electron energies in atoms.
Later we will develop more sophisticated models for orbitals and bonding to understand electron energies
1.1 Energies of Electrons in Atomic Orbitals
• Simple picture of factors controlling electron energies (a better description is coming later…..)
electron stabilized by
nucleus, held "tightly"
close to nucleus,
relatively low energy
electron
6+
electron in larger 2s A.O.,
further from nucleus and
shielded by two 1s
electrons, not held tightly,
relatively high energy
electron
3+
electron in even larger 2p
A.O., but, higher positive
charge on the nucleus (6+),
thus outer electrons held
reasonably tightly by
nucleus, moderate energy
even higher positive charge
on the nucleus (8+), but now
atomic orbital has another
negatively charged electron,
outer electrons held
reasonably tightly by
nucleus, moderate energy
8+
1.2 Quantitative Energies of Electrons in Atoms
• Quantitative information about the relative energies of electrons is obtained from measurements of Ionization
Energies, or Ionization Potentials (IPs)
• The "first" IP is the energy required to completely remove the highest energy electron from an atom or
molecule, we are interested mainly in the energies of these highest energy electrons since these are the ones that
are involved in chemical reactions, when we talk about I.P. we are always talking about the FIRST I.P.
Example from General Chemistry (energies in eV, i.e. electron Volts) (do NOT memorize these!)
Energy
energy of an electron that is infinitely far from any nucleus
2s
1s
2p
1s
2p
2p
2p
2s
1s
2p
2p
2p
2p
2s
2s
1s
1s
2p
– "valence" electrons, are involved in reactions/bonding
– "core" electrons, not involved in reactions/bonding
• VALENCE ELECTRONS are in the outer shell, they are highest in energy and get involved in bonding
• CORE ELECTRONS are in the inner shells, are not involved in bonding
Bonding and Structure
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• Electrons in atoms are low in energy because they are negatively charged and are stabilized by the positively
charged nucleus
• Electrons that are held "less tightly" by the nucleus are HIGH in energy, and thus require less energy to remove
from an atom, and thus have a low IP
• Many factors influence atomic I.P.s, (orbital size, nuclear charge, orbital occupancy, etc.), which looks
intimidating at first, however, a detailed understanding is not necessary at this point,
• What you SHOULD KNOW at this point is that I.P. increases (electron energy decreases) roughly with
increasing electronegativity (i.e., left to right across the periodic table and from bottom to top, going up the
periodic table)
hydrogen
helium
1
2
H
He
1.0079
4.0026
lithium
beryllium
boron
carbon
nitrogen
oxygen
fluorine
3
4
5
6
7
8
9
6.941
9.012
10.811
sodium
magnesium
aluminium
11
12
Li Be
B
C
13
Na Mg
14.007
silicon
phosphorus
14
Al
N
12.0107
Si
15
P
O
15.999
sulfur
16
S
neon
10
F
Ne
18.998
20.180
chlorine
argon
17
18
Ar
Cl
22.990
24.306
26.912
28.086
30.974
32.067
35.453
39.948
potassium
calcium
scandium
titanium
vanadium
chromium
manganese
iron
cobalt
nickel
copper
zinc
gallium
germanium
arsenic
selenium
bromine
krypton
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
39.098
40.078
44.956
47.867
50.942
51.996
54.938
55.845
39.098
58.693
63.546
65.39
69.723
72.61
74.922
78.96
79.904
rubidium
strontium
yttrium
zirconium
niobium
molybdenum
technetium
ruthenium
rhodium
palladium
silver
cadmium
indium
tin
antimony
tellurium
iodine
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
85.468
87.62
91.224
92.906
95.94
[98.91]
101.07
85.468
106.42
107.87
112.41
114.818
118.71
121.760
127.60
126.904
cesium
barium
lutetium
hafnium
tantalum
tungsten
rhenium
osmium
iridium
platinum
gold
mercury
thallium
lead
bismuth
polonium
astatine
55
56
71
72
73
74
75
76
77
78
79
80
81
82
83
84
85
132.905
137.32
174.97
178.49
186.21
190.23
132.905
195.08
196.97
200.59
204.383
207.2
208.980
[208.98]
[209.99]
K Ca
Sc Ti
Rb Sr
Y
88.906
Cs Ba
V
36
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br
Zr Nb Mo Tc Ru Rh Pd Ag Cd
Lu Hf Ta W Re Os Ir
180.95183.84
Pt Au Hg
In Sn Sb Te
Kr
83.80
xenon
54
I
Tl Pb Bi Po At
Xe
131.29
radon
86
Rn
[222.08]
1.3 Energies of Electrons in Molecules, e.g. Hydrogen (more)
• We can also measure energies of electrons in molecules as Ionization Potentials
Energy
energy of an electron that is infinitely far from any nucleus
IP ~ 13.6
IP ~ 15.4
lower energy
IN A BOND
H atom
1s atomic orbital
H2 molecule
σ molecular orbital
• IP in MOLECULAR hydrogen is LARGER than in atomic hydrogen
• the electrons in MOLECULAR hydrogen are thus LOWER in energy
Question. Why are the electrons lower in energy in molecular hydrogen compared to H atom?
Answer Because they are IN A BOND - this is really important
each electron
stabilized by
two nuclei
electron stabilized
by one nucleus
H atom
1s atomic orbital
H2 molecule, covalent bond
σ molecular orbital
• In the molecule, the nuclei are shielded from each other by the two electrons
• In the molecule there is an electrostatically stable configuration for the two negatively and two positively
charged particles (the electrons and the protons)
• this is the STANDARD explanation for why the energies of electrons are lowered when in a bond, later we will
learn a better explanation
We have already learned the TWO MOST CRITICAL CONCEPTS for UNDERSTANDING organic chemistry:
1. Forming bonds stabilizes (lowers the energies of) electrons
2. Higher energy electrons are MORE CHEMICALLY REACTIVE
Bonding and Structure
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2 Introduction to Bonding in Molecules: The Lewis Structure Model
2.1 Atomic Valence and Lewis Structures
Making a Molecule of H2 as a Simple Example
1 valence electron
H
1s
H electron configuration
H•
shared electrons: covalent bond
H
H
•H
Lewis "dot" structure
H
H
Lewis or Kekule
structure
• each hydrogen atom makes one bond, because in doing so it lowers the energy of its one electron
• hydrogen can only make one bond because it only has one VALENCE electron
Making a Molecule of CH4 as a Simple Example
2 core
electrons
H
4 valence electrons
1s2 2s22p2
C
H
C electron configuration
H
C
H shares 8
C H
shares 2
H
H
Lewis "dot" structure
H
H
H C
H
H
Lewis or Kekule
structure
• carbon has four valence electrons, therefore it lowers the energy of four electrons by making four bonds
• after making four bonds, according to the Lewis dot model its outer shell is now "full" with 8 electrons, it obeys
the "filled shell", sometimes call "octet rule" for second row elements, but really what it did was lower the energies
of as many electrons as it could, the 4 valence electrons in this case
• the hydrogens also have a "full" first shell with 2 electrons
• The filled shell rule and the pictures of the Lewis dot structures are MODELS (remember, we use the simplest
model that is useful)
Making a Molecule of NH3 as a Simple Example
H
X
2 core
electrons
5 valence electrons
1s2 2s22p3
N
N electron configuration
H
H
N
shares 2
H N
H
no more bonds
H
shares 8
Lewis "dot" structure
H
non-bonding
electrons
H N H
H
Lewis or Kekule
structure
• nitrogen has five valence electrons, but can only lower the energy of three electrons to make bonds before
"filling" the outer electron shell, attempting to make another bond "overfills" the shell and violates the "filled shell
rule" (later we will show that attempting to make another bond, violating the filled shell rule, increases instead of
decreasing electron energy)
• Nitrogen in this case has two NON-BONDING electrons in the outer shell
• The octet rule and the pictures of the atoms and molecules as Lewis dot and Lewis structues are MODELS (we
use the simplest useful model)
UNFORTUNATELY, there is MORE THAN ONE USE of the word VALENCE
1) The outer shell is the VALENCE SHELL, the electrons in the outer shell are the VALENCE ELECTRONS
2) the number of bonds that an atom normally makes without violating the "filled shell rule" is the NORMAL
VALENCE
Bonding and Structure
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• "Normal" Bonding Patterns (normal VALENCE) for various atoms obtained from the atomic configuration and by
using the filled shell rule (again, a better model is coming soon!)
• do NOT memorize this table, learn the Normal Valences by working with them by building organic structures
atom
1st shell
2nd shell
3rd shell
electron
configuration
# valence
electrons
maximum possible
number of electrons in
outer shell
normal valence
(normal # of
covalent bonds)
H
1s
1
2
1
B
1s2 2s22p1
3
8
3
C
1s2 2s22p2
4
8
4
N
1s2 2s22p3
5
8
3
O
1s2 2s22p4
6
8
2
F
1s2 2s22p5
7
8
1
P
1s2 2s22p63s23p3
5
8
3 or 5!
S
1s2 2s22p63s23p4
6
8
2,4 or 6!
• # of valence electrons = # of electrons in highest energy shell
• normal valence = # of electrons required to fill shell rule = # of bonds atoms "normally" makes
• Boron does not have enough electrons to fill the shell, even if all are involved in bonding
• Filled shell or Octet rule doesn't really work for 3rd row elements, e.g. phosphorus and sulfur, see why later
• Obeying the filled shell "rule" is the same as obeying the normal RULES OF VALENCE (normal number of
bands), i.e. 4 bonds to each C, 3 bonds to each N, 2 bonds to each O etc.
Example Problem 1 Draw a Lewis "dot" and a Lewis (Kekule) structure for N2H4
The filled shell "rule"/Normal Valence requires….
• 3 bonds to each nitrogen, N is TRIvalent (each nitrogen wants to make THREE bonds)
• 1 bond to each hydrogen, H is MONOvalent (each hydrogen wants to make ONE bond)
non-bonding electrons
(lone pairs)
octet rule satisfied for N
H
H
H
H
N
N
N
N
H
H
H
H
Lewis "dot" structure
Lewis
structure
• Each N has 8 electrons associated with the formal outer (valence) shell
• Each N "shares" 6 electrons in the three bonds, each N is considered to "own" 1 of each shared pair in each
bond, thus each N "owns" 3 of the bonding electrons
• Each N "has" 2 (or one pair of) nonbonding electrons
• Thus each N shares 8 electrons and "owns" 5 (3 bonding + 2 non-bonding = 5)
• note that the NON-BONDING electrons are shown in the Lewis structure
• LEWIS DOT STRUCTURES ARE TEDIOUS TO DRAW, THIS IS THE LAST TIME WE WILL USE THEM!
Example Problem 2
Draw two different Lewis structures for C2H6O
The filled shell "rule" /Normal Valence Requires…..
• 4 bonds to each TETRAvalent carbon (each carbon wants to make FOUR bonds)
• 2 bonds to each DIvalent oxygen (each oxygen wants to make TWO bonds)
• 1 bond to each MONOvalent hydrogen (each hydrogen wants to make ONE bond)
Bonding and Structure
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• The oxygen "shares" 4 electrons in the 2 bonds, "has" 4 non-bonding electrons, thus "owns" 6 electrons
• Lewis/Kekule structures indicate connectivity of atoms, orientation doesn't matter (for now!)
• Different Lewis (Kekule) structures are ISOMERS, more on isomers later…..
Example Problem 3
Draw a Lewis structure for HNO2, satisfying the normal rules of valence
• 3 bonds to each TRIvalent nitrogen
• 2 bonds to each DIvalent oxygen
• 1 bond to each MONOvalent hydrogen
7 electrons
around each O
O
H
6 electrons on N
7 on one O
N
O
H
N
H O
O
O
O
N
8 electrons on
both O and N
• N "shares" 6 electrons and "has" 2 electrons (obeys octet rule)
• H "shares" 2 electrons"
• EACH O "shares" 4 electrons in the 2 bonds, "has" 4 non-bonding electrons (obeys octet rule)
• NOTE, the nitrogen-oxygen DOUBLE bond is required to satisfy normal rules of valence for N and O
2.2 Condensed Structures
• Equivalent to abbreviated Lewis structures, the order of atom connectivity/bonding implied by the "written" order
but the bonds are not explicitely included
Example 1: Convert the provided condensed formula into a Lewis/Kekule structure
• Obeying the normal rules of valence for C (4 bonds), H (1 bond) and O (2 bonds) can generate only one
possible Lewis structure
carbon with 3 H's connected to
carbon with 2 H's connected to
carbon with 2 H's connected to
oxygen with 1 H
CH3CH2CH2OH
condensed
H
H
H
H
C
C
C
H
H
H
O
Lewis
H
• note that the non-bonding electrons ARE shown in Lewis structures
• H3C- is a common structural motif in organic structures (end of a chain)
Example 2: Convert the provided condensed formula into a Lewis/Kekule structure
• Obeying the normal rules of valence for C (4 bonds), H (1 bond) and O (2 bonds) can generate only one
possible Lewis structure
parenthesis means repeated unit
parenthesis means "off" the main chain
CH3C(CH3)2CH(OH)CO2CH3
functional group
H
H C H
H
H C
C
only possible
H H C H
connectivity order
H
O
C
H
H
"double bond"
O
C
2 bonds to C
H
O
C H
H
-CO2CH3
=
C
H
O
O
C H
H
• note carbon to oxygen double bond, required to satisfy the normal rules of valence for all atoms
• note TWO DIFFERENT USES OF PARENTHESES
1) Indicates a part of the structure that comes "off" the main chain
2) Indicates repeating units along the main chain
Bonding and Structure
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Example 3. Convert the provided condensed formula into a Lewis/Kekule structure
• Obeying the normal rules of valence for C (4 bonds), H (1 bond) and O (2 bonds) can generate only one
possible Lewis structure
parenthesis
means 2 -CH2- groups in a row
CH3(CH2)2CCCOCH2CO2H
acid functional group
H H
H
H C
H
C
C
H
H
C
C
O
H
C
C
H
O
C
must be a triple bond here
O
H
• Note the use of parentheses to indicate repeating units (the other use is to represent a group "off" the main
chain). These two uses are easily distinguishable using the normal rules of valence, only one will "work" in a
particular context
2.3 Line-angle Structures
• Sometimes also called SKELETAL structures
• The lines show bonds between carbon atons, no H's shown unless part of a functional group (see later for
definition)
Example 1
H
H
O
C
H
C
(CH3)3CCOCH2OCH3
H C
C
H
C H
H
condensed
O C H
H C H
Lewis structure
H
H
H
H
O
O
Line angle
(skeletal)
• there is an atom (carbon unless otherwise specified) at each "end" of each line, each line is a covalent bond
• Full Lewis structures sometimes just take too long to draw, and we will usually use line-angle (skeletal)
structures instead (especially for rings)
Example 2
Draw TWO DIFFERENT Lewis (Kekule), condensed and line angle structures for C4H6O (i.e. draw two ISOMERS,
defined in detail below)
Lewis structure
H
C
H
O
C
C
H
H
H C
C
H
H C
H
C
H
Condensed structure
Line Angle structure
O
H
H
C O
CH2CHCH2CHO
this H is part of
the aldehyde
functional group
H2C
H2C
CH
CHO
O
H
H
Bonding and Structure
H
line angle good for rings
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• condensed formulas are not very good for drawing rings
• line-angle structures are very good for rings
• where reasonable, draw the angles roughly correct for the molecular shape, see later
Q. Which Kind of Structure to Draw (condensed, Lewis, line-angle etc)?
A. Depends upon the context, we will use mainly line angle or a line angle/Lewis mix
3 Isomers
• Isomers are different compounds with same molecular formula (we have already seen some of these)
• We meet two kinds of isomers in this course, structural isomers and stereoisomers
3.1 Structural (Constitutional) Isomers
• These are isomers that differ in the order in which the atoms are connected (connectivity of the atoms), they
differ in the order in which the atoms are bonded together
• Their physical and chemical properties are DIFFERENT, they are different structures, different molecules,
different chemicals
Example 1: Structural isomers for C4H10
C4H10
TWO isomers are
possible
H3C CH2 CH2 CH3
H3C
CH
CH3
CH3
isobutane b.p. = -11.7°
butane b.p. = -0.5°
Example 2: Structural isomers for C5H12
4
3
CH2 CH3
C5H12
THREE isomers
H3C CH2 CH2 CH2 CH3
(5) 2
H3C CH
4 CH3
3 CH2
CH3
1
2 CH
1 CH3
are possible
(5) CH3
CH3
H3C C CH3
CH3
same structure NOT isomers
• Note that the direction in which the bonds "point" is irrelevant, structural isomers are generated by connecting
atoms together in a different order only
• There are thus THREE different structural isomers that have the molecular formula C5H12
Example 3: Generate ALL Structural isomers for C6H14
• useful strategy is often to start with the longest possible chain, and progressively "branch"
CH3
CH2 CH2 CH2 CH2 CH3
• We find that there are FIVE possible structural isomers
• the direction in which the bonds "point" is irrelevant, structural isomers are generated by connecting atoms
together in a different order only
Bonding and Structure
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Example 4: How many DIFFERENT isomers are there shown below?
CH3
CH3 CH2 CH
1
2
3
1
2
3
CH3 CH2 CH
CH2 CH3
4
5
4
5
CH2 CH3
1 2
H3C CH2
CH3
5 CH3
H3C 1
2 3 4
CH2 CH CH2
1
CH3
5 CH
3
2
3
4
CH2 CH CH2
CH3
3
HC CH3
5
4
H3C CH2
CH3
there are no structural isomers here, this is the SAME structure, the order of connectivity of he
atoms is the same in each, the structure is simply drawn differently on paper
Visualize Isomers, or Lack Of!
3.2 Stereoisomers, Specifically cis/trans or Geometrical Isomers
• These are isomers of structures that contain C=C double bonds, that have the same atom connectivity (the
atoms are connected together in the same order), but that differ in the orientation of at least some of the atoms in
space (the direction in which they point)
• Their physical and chemical properties are DIFFERENT, they are different structures, different molecules,
different chemicals
• There are different NAMES for stereoisomers! This can be confusing but we will try to use the simplest
terminology
• The isomers that we are talking about in this section of the notes are called geometrical isomers, but that is an
older term that is seldom used now, they are also called configurational isomers, this is sometimes still used, but
the term cis/trans isomers is very commonly used and is very explicit (it is easy to understand) this is the term that
we will mainly use
• Cis/trans geometrical isomers that can occur when a structure contains a C=C double bond
• The term diastereomers also applies to this kind of stereoisomers, and we will use it to distinguish this from
another kind of stereoisomer that we will meet later in the course. It is better to introduce the term diastereomer
now (it will be less confusing later), although we will probably mainly use the term cis/trans isomers.
H3C
(opposite sides)
H
C
H
C
stereoisomers
(diastereomers)
b.p. 0.8°
H
C
H
Bonding and Structure
H
C
H3C
CH3
trans-butene
Br
H
b.p. 3.7°
C
(same side)
CH3
cis-butene
Br
C
Br
C
Br
stereoisomers
(diastereomers)
H
8
H
Br
C
C
H
H
structural isomers
C
Br
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BOND ROTATION (OR LACK OF) IS RESPONSIBLE FOR CIS/TRANS ISOMERS, AND IS THE REASON
THAT THERE ARE NO EQUIVALENT STEREOISOMERS FOR STRUCTURES THAT DO NOT HAVE A C=C
DOUBLE BOND
Compare rotation around SINGLE and DOUBLE bonds
Br
CH3
H C
C H
single bond
H
rotation
H3C
same structure
• rotation around a single bond occurs READILY and does NOT generate isomers
• rotation around single bonds generates conformers of the same structure, to be discussed in more detail later
Visualize Isomerization and Single Bonds
Br
CH3
Br
=
C
H3C
C
trans-
H
NO!
double bond
rotation
H
Br
C
H3C
DIFFERENT structures
stereoisomers
=
C
cis-
Br
CH3
• rotation around a double bond is RESTRICTED (essentially we can assume that it does not occur under normal
conditions), STEREOISOMERS CAN thus form in suitable structures that have C=C double bonds (although
not all structures with C=C double bonds will have isomers)
• Here we have one TRANS-stereoisomer where the two -CH3 groups are on OPPOSITE SIDES of the double
bond, and one CIS-stereoisomer where the two -CH3 groups are on the SAME SIDE of the double bond
Visualize Isomerization and Double Bonds
Bonding and Structure
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• Note that here we have not really explained WHY bond rotation occurs (rapidly) for C-C single bonds
and not for C=C double bonds, for now we will just accept it as a fact, and we will revisit the question of exactly
why later in the course
3.3 Superimposibility
• identical structures are SUPERIMPOSABLE
Example: Are the following two structures the same or are they isomers?
30°
rotate around
rotate entire
indicated bond
structure 30°
same shape
superimposable
• remember, rotation around a single bond occurs readily and does not generate a new structure
• bond rotation followed by complete molecule rotation (by 30°) generates a new version of the original that is
SUPERIMPOSABLE on the second structure
• the structures are superimposable, therefore they are the same structure (just drawn two different ways)
Visualize Bond Rotation with a Model
Example: Are the following two structures the same or are they isomers?
C
H
H
Cl
C
C
H
cis-diastereomer
H
C
Cl
trans-diastereomer
Cl
Cl C
HC
H
H
NON-Superimposable!!
X
Cl
X
Cl
NON-Superimposable!!
C Cl
C Cl
H
• trying to superimpose the cis- structure on top of the trans- fails, the structures are non-superimposable, they
are this isomers, in this case they are stereoisomers
• Note: It is NOT NECESSARY to confirm lack of superimposability when identifying isomers, often this is
not necessary (overkill!), but superimposibility is the most critical test for isomers that we will resort to in difficult
cases, and especially when we encounter more subtle forms of stereoisomers later
3.4 Generating Structural Isomers Using Degrees of Unsaturation
• The degrees (ALSO KNOWN AS ELEMENTS) of unsaturation is an important piece of information about a
molecule/structure that is obtained from the molecular formula
Consider structures consisting of linear chains of carbon atoms joined by single bonds.
How many hydrogen atoms can be attached to these chains of carbon atoms?
Bonding and Structure
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3 carbons - 8 hydrogens
H
H
H
H
C
C
C
H
H
H
4 carbons - 10 hydrogens
add 1 carbon to
the chain
H
H
H
H
H
H
C
C
C
C
H
H
H
H
H
2 "end" hydrogens
each added carbon also adds 2
hydrogens
• every time ONE CARBON is added to a chain, TWO HYDROGEN ATOMS are also added
• that maximum # of H's any organic molecule can have is (2 x # of carbons) + 2 (the "end" H atoms)
• a molecule that has the maximum number of hydrogen atoms is referred to as SATURATED
• a molecule that does not have the maximum number of hydrogen atoms is UNSATURATED
• a molecule can be unsaturated by having double or triple bonds and rings, each "cost" hydrogen atoms
H
H
H
H
H H
C
H H H H
H
C
C
C
H C
H
H
H
C
H
H C
C
C
C
C
C
C
C
C
H
H
H
H H H H H H
H H
H
H
H
H
C5H12
(5 x 2) + 2 = 12
0 degrees of unsaturation
C5H10
double bond "costs" 2xH
1 degree of unsaturation
C5H12
ring "costs" 2xH
1 degree of unsaturation
Calculating Degrees (Elements) of Unsaturation
(Max # of H atoms) - (Actual # of H atoms)
degrees (elements) of unsaturation =
2
(Max # of H atoms = (# carbons x 2)
What if the molecule contains other elements?
Halogens are monovalent, they take the place a hydrogen and are counted as a hydrogen
H
H
X
H
C
C
C
X = F, Cl, Br, I
H
H
H H
Oxygen is divalent, when an oxygen atom is added to a chain zero hydrogens are added, oxygen is IGNORED
2 carbons - 6 hydrogens
H
H
add 1 oxygen to
the chain
H
C
C
H
H
2 carbons and 1 oxygen - STILL 6 hydrogens
H
H
H
H
O
C
H
C
H
zero H atoms added to the chain
H
Nitrogen is trivalent, when a nitrogen atom is added to a chain, one hydrogen atom is also added. Thus one
nitrogen atom is equivalent to half a carbon atom as far as adding hydrogen atoms to the chain is concerned, N is
counted as HALF a carbon atom
2 carbons and 1 nitrogen - 7 hydrogens
2 carbons - 6 hydrogens
H
H
H
C
C
H
H
Bonding and Structure
add 1 nitrogen to
the chain
H
H
C
H
H
11
each N adds 1 H to the chain
H
H
N
C
H
H
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Example 1
C5H9OCl
Max # of H atoms = (5 x 2) + 2 = 12
Actual # of H atoms = 9 + 1 = 10 (Cl counted as H)
degrees of unsaturation = (12 - 10) / 2 = 1 degree e.g.
C5H9OCl MUST have
O
Cl
HO
etc
Cl
double bond
one ring or one double bond but NOT both
ring
• Two POSSIBLE structures are shown, other isomers are also possible for this molecular formula
Example 2 C4H7NO
C4H7NO
H
Max # of H atoms = (4 x 2) + 1 + 2 = 11
Actual # of H atoms = 7
degrees of unsaturation = (11 - 7) / 2 = 2 degrees
C4H7NO MUST have
O
e.g.
2 rings or 2 double bonds or 1 ring and 1 double bond
N
Cl
etc
N
H
• Two POSSIBLE structures are shown, other isomers are also possible for this molecular formula
• knowing the degrees (elements) of unsaturation helps in drawing Lewis structures, from the molecular formula
you can determine the number of double bonds, rings etc.
Example 3 generate as many line-angle structures as you can that have the molecular formula C5H10
degrees (elements) of unsaturation =
(5 x 2) + 2) - (10)
= 1 degree
2
• all structures must have ONE double bond or a ring, but not both….
• approach the problem systematically, start with the longest chain and progressively branch
• where rings are possible, start with the largest ring and progressively get smaller
same
=
• A maximum of TEN possible STRUCTURAL ISOMERS can be drawn, AND one PAIR OF STEREOISOMERS
trans-isomer (from above)
cis-isomer
4 Formal Charges
• The Lewis structure model attempts to obey the filled shell/octet rule for as many atoms as possible by getting
as many electrons as possible into bonds (remember, forming bonds lowers electron energy)
• Sometimes this can result in structures that disobey the normal rules of valence, resulting in an atom "owning"
either more or less electrons than its valence electrons
• This results in an atom having a FORMAL CHARGE
• atoms SHARE electrons in bonds, atoms "own" half of the shared electrons and their non-bonding electrons
Bonding and Structure
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Simple Example
H
# of valence electrons
H=1
C=4
H
H
H C H
H
H
C
H "sees" 2 electrons
"owns" 1 electron, is neutral
C "sees" 8 valence electrons
"owns" 4 valence electrons, is neutral
H
• the carbon "sees" the two shared electrons in each of the four bonds
• the carbon "owns" half of each electron pair in each bond
• the carbon owns 4 electrons, neutral carbon has four valence electrons, the carbon has NO FORMAL CHARGE
Example
# of valence electrons
O=6
N=5
H
CH3NO2
H C
H
O
N
O sees 8, owns 6 electrons, neutral
N sees 8 electrons, owns 4, positive charge
O
O sees 8, owns 7 electrons, negative charge
• this is the only reasonable way to draw a Lewis structure for CH3NO2
• the "filled shell/octet rule" is obeyed for all atoms, but the normal rules of valence are not (4 bonds to the
nitrogen instead of 3, and 1 bond to the lower oxygen instead of two)
• the central nitrogen "sees" 8 electrons, i.e. the ones that are shared in the 4 bonds, and "owns" 4 electrons (half
of the shared electrons), but requires 5 valence electrons to be neutral, the charge is +1
• the LOWER oxygen "sees" 8 electrons (6 non-bonding and the shared pair in the bond), "owns" 7 electrons (the
non-bonding ones and half the shared pair), but requires 6 valence electrons to be neutral, the charge is -1
The Calculation of formal charge using a formula:
formal charge = (# valence e's- (# non-bonding e's) - 1/2 (# e's in bonds) (for TWO electrons per bond)
for CH3NO2 (above)
C = (4) - (0) - 1/2 (8) = 0 (zero formal charge)
Each H = (1) - (0) - 1/2 (2) = 0 (zero formal charge)
N = (5) - (0) - 1/2 (8) = +1 (one positive charge)
O(upper) = (6) - (4) - 1/2 (4) = –0 (zero formal charge)
O(lower) = (6) - (6) - 1/2 (2) = –1 (one negative charge)
• Later we will see that it is easier to determine formal charges on organic structures by NOT using this formula,
but it is useful at this point as a way of making progress in understanding structures, so it is OK to use it for now
Example
Na 5 valance electrons : 5 electrons to be NEUTRAL
Na formally "owns" 1 electron for each bond = 4
Na has zero non-bonding electrons
"owns" one eless lectron than needed for neutrality : POSITIVE CHARGE
H
CH2N2
C
H
Na
Nb
Nb = (5) - (4) - (4) = –1 (one negativecharge)
valence
in bonds
non-bonding
• the "filled shell/octet" rule is obeyed in all cases, but the normal rules of valence are not
• formal charges are associated with structures in which the normal rules of valence are not obeyed
Bonding and Structure
13
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But
# non-bonding
zero formal charge
Charge = 4 - 2 - (4/2) = 0
C
H
H
1/2 # bonding
# valence
• this is methylene (you do not have to know this), it is very reactive (its lifetime is less than 0.0000000001 second
in solution), but we can still draw a valid Lewis structure for it
• not ALL atoms that disobey the normal valence rules have a formal charge!
5 Functional Groups
• Organic molecules consist, in general, of a carbon/hydrogen (hydrocarbon) "skeleton", that mainly determines
the size and shape of the molecule, and……
• the FUNCTIONAL GROUPS, generally involving atoms that are more electronegative than carbon (so that polar
bonds result), such as O, N, S etc.
• chemistry takes place at the FUNCTIONAL GROUPS. When we start to discuss reactions, we will divide them
into those characteristic of the various functional groups
Example of an organic molecule
O
HO
functional groups
functional groups
OH
O
cortisone
anti-inflammatory
carbon/hydrogen "skeleton"
O
functional groups
Alkane (hydrocarbon)
Not really a functional group, but the important "backbone" or "skeleton" of many organic molecules
C-C and C-H bonds tend to be strong and relatively unreactive
H
H
H
C
H
means "the rest of the
molecule is attached here"
C
C
H
H
R stands for any alkyl chain, e.g. R–OH could be:
H3C–OH
or
or
an alkyl chain
CH3CH2OH
NH
OH
etc.
N
bupivacaine, an
epidural anesthetic
O
Some Notation
• a group with ONE substituents is referred to as PRIMARY (1°), the substituent is often an alkyl chain (-R) or an
aryl (aromatic) group (-Ar)
• a group with TWO substituents is referred to as SECONDARY (2°), the substituents are often alkyl chains (-R)
or aryl groups (Ar)
• a group with THREE substituents is referred to as TERTIARY (3°), the substituents are often alkyl chains (-R) or
aryl groups (-Ar)
Bonding and Structure
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H
Z
H
H
Primary
or 1°
R
R
Z
R
R
Secondary
or 2°
Z
R
R
Tertiary
or 3°
Alkene Functional Group
carbon-carbon double bond
OH
C
alkene
diene
alkene
C
linalool, used in the perfume industry
NOT aromatic (see next section)
does not have alternating double/single bonds
Aromatic Functional Group
Alternating C-C and C=C bonds in a ring - be careful to distinguish from Alkene
orange B, food coloring
EtO
aromatic
O
N
N
N N
SO3H
O
aromatic
HO3S
Ar
stands for any aromatic ring system, e.g. Ar–OH could be:
OH
OH
Cl
or
N
OH
etc.
or
Alkyne Functional Group
carbon-carbon triple bond
HO
C
mestranol, the estrogen used in many oral
contraceptives
C
O
Bonding and Structure
15
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Amine Functional Group
contain a nitrogen with at least one alkyl or aryl group, here R1, R2 etc stands for any alkyl chain that may or may
not be the same
R2
R1
R1
R1
N
N
N
R3
R2
3°
H
H
2°
secondary
tertiary
HO
R or Ar
H
O
1°
primary
S
O
amine
NH2
taurine, supposedly active ingredient in
energy drinks e.g. Red Bull
Ether Functional Group
oxygen between 2 alkyl or aryl groups
(R, Ar)
O
O
ether
(R, Ar)
tetrahydrofuran, THF,
common organic solvent
Epoxide Functional Group
3-membered ring containing oxygen
4,5-benzo[a]pyrene oxide,
highly carcinognic
O
epoxide
O
Alcohol
oxygen with 1 R or Ar group and 1 hydrogen
O
(R, Ar)
OH
H
bisabolol is used to aid in the
transfer of drugs through the skin
alcohol
Halide Functional Group
aryl or alkyl group with fluoride, chloride, bromide, iodide...
chloride
R or Ar
R
R
F
R
Cl
Cl
Cl
Cl
chloride
Br
R
Cl
I
chloride
a polychlorinated biphenyl (PCB),
many industrial uses but toxic,
bioaccumulates in animals
chloride
Ketone Functional Group
C=O double bond with 2 akyl or aryl groups
(R, Ar)
O
C
ketone
acetone, the simplest
ketone, common organic
solvent, nail-polish remover
O
(R, Ar)
Aldehyde Functional Group
C=O double bond with 1 R or Ar group and 1 H
O
(R, Ar)
O
C
aldehyde
H
H
HO
vanillin, main extract fron
vanilla bean
OMe
Bonding and Structure
16
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Carboxylic Acid Functional Group
C=O with 1 R or Ar group and 1 –OH
O
(R, Ar)
C
carboxylic acid
O
sorbic acid, food
preservative
OH
OH
Ester Functional Group
C=O double bond with –OR or -OAr group
ester
O
(R, Ar)
C
O
methyl paraben, food
preservative, cosmetics
additive, anibacterial/fungal
HO
O (R, Ar)
O
Amide Functional Group
C=O with -NR2 (R or Ar)
amide
O
C
3°
O
N
R
R
C
N
O
N
R
H
2°
C
1°
N
O
OH
H
H
Cl
loperamide, anti-diarrhea drug
Acid Halide Functional Group
C=O double bond with a halide, X means any halide, Cl, Br etc.
acid chloride
O
(R, Ar)
C
O
acetyl chloride, many
useful reactions
Cl
X
Nitrile Functional Group
Aryl or alkyl group with carbon nitrogen triple bond
O
N
(Ar, R) C
N
nitrile
C N
citalopram,
antidepressant drug
(Ar,R) means aryl or alkyl
Here are some functional groups incorporated into line-angle and condensed structures that you will see and
need to understand
Bonding and Structure
17
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aldehyde
alcohol
carboxylic acid
O
O
H
O
O
H
or
or
CHO
N
N
H
OH
or
nitrile
and amine
amide
H
or
CO2H
or
O
OH
N H
C
NC
NH
NH2
Example problems: Indicate all functional groups, do not include alkanes
ester
(not ketone or ether)
O
O
N
O
ether
H
aromatic (not alkene)
ether
N
OH
aromatic (not alkene)
O
amine
O
O
ester
(not ketone or ether)
alkene
ketone
aromatic
fluoride
CF3
Fluoxetine (Prozac)
Heroin
Bonding and Structure
alcohol
amine
18
O
alkene
testosterone
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