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Chapter 2
Representative Carbon Compounds:
Functional Groups, Intermolecular
Forces and Infrared (IR) Spectroscopy
The ability of carbon to form covalent bonds to other carbon atoms, as
well as atoms such as hydrogen, oxygen, nitrogen and others, leads to
an enormous variety of structures. Many of these are found in
biological systems.
Carbon-carbon Covalent Bonds
• Carbon forms strong covalent bonds to other
carbons and to other elements such as
hydrogen, oxygen, nitrogen and sulfur
• This accounts for the vast variety of organic
compounds possible
• Organic compounds are grouped into
functional group families
• A functional group is a specific grouping of atoms
(e.g. carbon- carbon double bonds are in the family
of alkenes)
• An instrumental technique called infrared (IR)
spectroscopy is used to determine the presence of
specific functional groups
2
Hydrocarbons: Representative Alkanes, Alkenes Alkynes, and
Aromatic Compounds
• Hydrocarbons contain only carbon and hydrogen atoms
• Subgroups of Hydrocarbons:
•
•
•
•
Alkanes contain only carbon-carbon single bonds
Alkenes contain one or more carbon-carbon double bonds
Alkynes contain one or more carbon-carbon triple bonds
Aromatic hydrocarbons contain benzene-like stable structures
(discussed later)
• Saturated hydrocarbons: contain only carbon-carbon
single bonds e.g. alkanes
• Unsaturated hydrocarbons: contain double or triple
carbon-carbon bonds e.g. alkene, alkynes, aromatics
• Contain fewer than maximum number of hydrogens per
carbon
• Capable of reacting with H2 to become saturated
3
Methane
This first member of the alkane family is the major component in
natural gas (60-80%) found in geological deposits. A large industry is
based on the recovery and distribution of natural gas that is widely
used as a fuel for homes and industry. It is also used as a primary
energy source in power plants to produce electricity. About 10% of
the electricity produced in the United States comes from natural gasfired power plants. It is the cleanest of the fossil fuels (coal,
petroleum, natural gas).
Methane is produced by bacteria called methanogens under anaerobic
(oxygen-free) conditions. These microorganisms are found in ocean
trenches, mud, sewage and in the stomach of cows. Very large deposits
of methane hydrates (methane trapped in ice crystals) have been
discovered in deep ocean trenches.
H
109.5o
C
H
H
H
The four equivalent
hydrogens in methane are
located in the corners of a
tetrahedron.
Representative Hydrocarbons
• Alkanes
• Principle sources of alkanes are natural gas and
petroleum
– Smaller alkanes (C1 to C4) are gases at room temperature
• Methane is
– A component of the atmosphere of many planets
– Major component of natural gas
– Produced by primitive organisms called methanogens found in
mud, sewage and cows’ stomachs
5
Alkenes
Hydrocarbons that contain a carbon-carbon double bond are called
alkenes. The two simplest alkenes, ethene (C2H4) and propene (C3H6),
are important industrial chemicals used to prepare polymers.
CH2=CH2
polyethylene
ethene
(30 billion pounds per year in the United States)
CH3CH=CH2
propene
polypropene
(15 billion pounds per year in the United States)
• Alkenes
• Ethene (ethylene) is a major industrial feedstock
– Used in the production of ethanol, ethylene oxide and the polymer
polyethylene
• Propene (propylene) is also very important in industry
– Molecular formula C3H6
– Used to make the polymer polypropylene and is the starting
material for acetone
• Many alkenes occur naturally
7
Alkynes
Hydrocarbons that contain the carbon-carbon triple bond are called
alkynes. The simplest alkyne, ethyne (C2H2), is an important industrial
chemical used as a chemical intermediate (to make other important
chemicals) and as a fuel in the oxyacetylene torch for welding.
HC CH
ethyne
other organics
• Alkynes
• Many alkynes are of biological interest
– Capillin is an antifungal agent found naturally
– Dactylyne is a marine natural product
– Ethinyl estradiol is a synthetic estrogen used in oral
contraceptives
9
• Benzene: A Representative Hydrocarbon
• Benzene is the prototypical aromatic compound
– The Kekulé structure (named after August Kekulé who formulated
it) is a six-membered ring with alternating double and single
bonds
10
• Benzene: A Representative Hydrocarbon
• Benzene does not actually have discreet single and double
carbon-carbon bonds
– All carbon-carbon bonds are exactly equal in length (1.38 Å)
– This is between the length of a carbon-carbon single bond and a
carbon-carbon double bond
• Resonance theory explains this by suggesting there are two
resonance hybrids that contribute equally to the real
structure
– The real structure is often depicted as a hexagon with a circle in
the middle
11
• Molecular orbital theory explains the equal bond lengths of
benzene by suggesting there in a continuous overlap of p
orbitals over the entire ring
• All carbons in benzene are sp2 hybridized
– Each carbon also has an empty p orbital
• Each p orbital does not just overlap with one adjacent p but
overlaps with p orbitals on either side to give a continuous bonding
molecular orbital that encompasses all 6 carbons
• All 6 p electrons are therefore delocalized over the entire ring and
this results in the equivalence of all of the carbon-carbon bonds
12
Polar Covalent Bonds
• Polar covalent bonds occur when a covalent
bond is formed between two atoms of differing
electronegativities
• The more electronegative atom draws electron
density closer to itself
• The more electronegative atom develops a partial
negative charge (d-) and the less electronegative
atom develops a partial positive charge (d+)
• A bond which is polarized is a dipole and has a
dipole moment
• The direction of the dipole can be indicated by a
dipole arrow
– The arrow head is the negative end of a dipole, the
crossed end is the positive end
13
Bond Dipoles
The polarization of electrons in a covalent bond leads to a permanent
separation of charge:
d
d
A
B
B is more electronegative
A permanent bond dipole (two centers of charge) is formed where d
(delta) means partial. This electrical imbalance leads to
electropositive and electronegative regions within the molecule.
Examples
d d
H
F
dH
dH
O d
• Example: the molecule HCl
• The more electronegative chlorine draws electron
density away from the hydrogen
– Chlorine develops a partial negative charge
• The dipole moment of a molecule can be
measured experimentally
• It is the product of the magnitude of the charges (in
electrostatic units: esu) and the distance between the
charges (in cm)
• The actual unit of measurement is a Debye (D) which is
equivalent to 1 x 10-18 esu cm
15
• A map of electrostatic potential (MEP) is a way to
visualize distribution of charge in a molecule
• Parts of the molecule which are red have relatively more
electron density or are negative
– These region would tend to attract positively charged species
• Parts of the molecule which are blue have relatively less
electron density or are positive
– These region would tend to attract negatively charged species
• The MEP is plotted at the van Der Waals surface of a molecule
– This is the farthest extent of a molecule’s electron cloud and
therefore indicates the shape of the molecule
• The MEP of hydrogen chlorine clearly indicates that the
negative charge is concentrated near chlorine
– The overall shape of the molecule is also represented
16
Molecular Dipole
• In diatomic molecules a dipole exists if the two
atoms are of different electronegativity
• In more complicated molecules the molecular
dipole is the sum of the bond dipoles
• Some molecules with very polar bonds will have no
net molecular dipole because the bond dipoles
cancel out
– The center of positive charge and negative charge
coincide in these molecules
17
• Examples
• In carbon tetrachloride the bond dipoles cancel and the
overall molecular dipole is 0 Debye
• In chloromethane the C-H bonds have only small dipoles
but the C-Cl bond has a large dipole and the molecule is
quite polar
18
An unshared pair of electrons on atoms such as oxygen and
nitrogen contribute a great deal to a dipole Water and
ammonia have very large net dipoles
19
• Some cis-trans isomers differ markedly in their
dipole moment
• In trans 1,2-dichloroethene the two carbon-chlorine
dipoles cancel out and the molecular dipole is 0
Debye
• In the cis isomer the carbon-chlorine dipoles
reinforce and there is a large molecular dipole
20
Molecular Dipole Moments in Polyatomic Molecules
In polyatomic molecules, the molecular dipole moment is a vector sum
of the bond dipole moments. From measurements of molecular dipole
moments and structural studies of numerous polyatomic molecules,
approximate bond dipole moments have been calculated.
Some Approximate Bond Dipole Moments (D)
C-C
H-C
H-N
H-O
0.0
0.30
1.31
1.53
H-F 1.98
H-Cl 1.03
H-Br 0.78
H-I
0.38
C-F
C-Cl
C-Br
C-I
1.51
1.56
1.48
1.29
C-O
C=O
C-N
C=N
0.86
2.40
0.40
0.90
from L. N. Ferguson, The Modern Structural Theory of Organic Chemistry
The Importance of Molecular Symmetry
Carbon Tetrachloride, CCl4
Each C-Cl bond is polar covalent. But carbon tetrachloride does not
have a molecular dipole moment. Because of the tetrahedral
symmetry, there is no net molecular dipole moment and CCl4 is a
nonpolar molecule .
The molecular dipole moment
is the vector sum of all the bond dipoles.
Cl 1.56 D
1.56 D
1.56 D
Cl C Cl
Cl1.56 D
tetrahedral geometry
overall =1.56 D - 3 x 1.56 D x .3333 = 0 D
+z
-z
Evaluate the bond
dipole vectors in the
-z direction.
+z
1.56 D
C
Cl
109.5
o
C
Cl
o
70.5
-z
The vector contribution of the
bond dipole in the -z direction is
o
1.56 D x COS 70.5 = 1.56 D x .3333
Methyl Chloride, CH3Cl
The molecular dipole moment of methyl chloride is 1.86 D.
The analysis assumes a
tetrahedral geometry.
0.30
Cl 1.56
0.30
H C H
H 0.30
The net dipole moment in
the +Z direction is
1.56 (C-Cl)
0.30
3 x (C-H)
1.86 D
Note: The direction of the C-H bond moment is determined
from this and similar analyses to be H-C.
Nonbonding Electron Pairs
The contribution of nonbonding electron pairs to the molecular
dipole moment is seen from the following comparison.
NF3
 = 0.24 D
NH3
 = 1.46 D
The much larger molecular dipole moment for NH3 is, at first,
surprising, because of the similar electronegativity differences of the
atoms in the covalent bonds, H-F and N-H. Opposite directions
would be expected for the molecular dipole moments.
3.0
3.0
4.0 F
N
F 4.0
2.1 H
N
F
H
4.0
2.1
H 2.1
Contributions from Nonbonding Electrons
When dipole contributions from the nonbonding electron pairs are
included, the different magnitudes for the molecular dipole moments
of NF3 and NH3 are understandable.
1.46 D
0.24 D
F
N
F
F
Partial cancellation of
the bond dipoles leads to
a small molecular dipole
moment either
or .
H
N
H
H
Bond dipoles are reinforcing
leading to a larger molecular
dipole moment.
Functional Groups
• Functional group families are characterized by the
presence of a certain arrangement of atoms called a
functional group
• A functional group is the site of most chemical
reactivity of a molecule
– The functional group is responsible for many of the
physical properties of a molecule
• Alkanes do not have a functional groups
– Carbon-carbon single bonds and carbon-hydrogen bonds
are generally very unreactive
28
• Alkyl Groups and the Symbol R
• Alkyl groups are obtained by removing a hydrogen from an alkane
• Often more than one alkyl group can be obtained from an alkane by removal of different
kinds of hydrogens
• R is the symbol to represent a generic alkyl groups
–
The general formula for an alkane can be abbreviated R-H
29
• A benzene ring with a hydrogen removed is called a
phenyl and can be represented in various ways
• Toluene (methylbenzene) with its methyl hydrogen
removed is called a benzyl group
30
• Alkyl Halides
• In alkyl halides, halogen (F, Cl, Br, I) replaces the
hydrogen of an alkane
• They are classified based on the carbon the halogen
is attached to
– If the carbon is attached to one other carbon that carbon
is primary (1o) and the alkyl halide is also 1o
– If the carbon is attached to two other carbons, that
carbon is secondary (2o) and the alkyl halide is 2o
– If the carbon is attached to three other carbons, the
carbon is tertiary (3o) and the alkyl halide is 3o
31
Primary Alkyl Halide
H
R C
H
primary
(1o)
H
CH3 C
H
ethyl
H
CH3 C Cl
H
ethyl chloride
a primary alkyl chloride
Secondary Alkyl Halide
R
R C
H
secondary
(2o)
CH3
CH3 C
H
isopropyl
CH3
CH3 C Cl
H
isopropyl chloride
a secondary alkyl chloride
Tertiary Alkyl Halide
R
R C
R
t ert iary
(3o )
CH 3
CH 3
C Cl
CH 3
t ert iary but y l
CH 3
CH 3
C Cl
CH 3
t ert iary but y l chloride
a tertiary alkyl chloride
What kind of a carbon is this?
Alcohols
• In alcohols the hydrogen of the alkane is replaced by the
hydroxyl (-OH) group
– An alcohol can be viewed as either a hydroxyl derivative of an
alkane or an alkyl derivative of water
34
Alcohols
• Alcohols are also classified according to the carbon the
hydroxyl is directly attached to
35
Classification of Alcohols
Alcohols are classified as primary, secondary or tertiary
according to the structure around the carbon to which the
hydroxyl group is attached.
ethyl alcohol
isopropyl alcohol
CH3CH2OH
(CH3)2CHOH
H
CH3 C OH
H
a 1o alcohol
CH3
CH3
C OH
H
a 2o alcohol
tertiary-butyl alcohol
(CH3)3COH
CH3
CH3 C OH
CH3
a 3o alcohol
Quiz Chapter 2 Section 7
Name the following organic compounds using
common alkyl group names.
CH3Cl
CH3CHBrCH3
CH3CH2CH2OH
methyl chloride
isopropyl bromide
propyl alcohol
• Ethers
• Ethers have the general formula R-O-R or R-O-R’ where R’
is different from R
– These can be considered organic derivatives of water in which
both hydrogens are replaced by organic groups
– The bond angle at oxygen is close to the tetrahedral angle
38
• Amines
• Amines are organic derivatives of ammonia
– They are classified according to how many alkyl groups replace
the hydrogens of ammonia
– This is a different classification scheme than that used in alcohols
39
Classifications and Names of Amines
Simple amines are named using the alkyl group name followed by
"amine." When more than one R group is present, the prefixes "di"
and "tri" are included. If there are R and R' groups, the alkyl groups
are named in alphabetical order.
CH3CH2NH2
ethylamine
(a 1o amine)
CH3CHCH3
NH2
isopropylamine
(a 1o amine)
(CH3)2CHNHCH(CH3)2
diisopropylamine
(a
2o
amine)
CH2CH3
CH3CH2NCH2CH3
triethylamine
(a 3o amine)
CH3
CH3CCH3
NH2
tertiary-butylamine
(a 1o amine)
H2
C
H2C
CH2
H2C
CH2
N
H
piperidine
(a 2o amine)
(CH3CH2)2NCH3
diethylmethylamine
(a 3o amine)
N
H
The Structure and Base Property of Amines
:
VSEPR theory
predicts a
tetrahedral
geometry.
:
Amines have a trigonal pyramidal structure similar to ammonia.
N
H
H
H
N
H3C
CH3
CH3
107o
ammonia
108.7o
trimethylamine
Amines are bases.
:
(CH3)3N
base
+
HCl
acid
H ClH3C N CH3
CH3
salt
Quiz Chapter 2 Section 9
Identify the following amines as primary, secondary or tertiary types,
and give their alkyl group names.
(CH3)2NH
(CH3)3CN(CH3)2
CH3CH2CH2CH2NH2
(a 1o amine)
butylamine
(a 2o amine)
dimethylamine
(a 3o amine)
tert-butyldimethylamine
• Aldehydes and Ketones
• Both contain the carbonyl group
• Aldehydes have at least one carbon attached to the carbonyl group
• Ketones have two organic groups attached to the carbonyl group
• The carbonyl carbon is sp2 hybridized
–
It is trigonal planar and has bond angle about 120o
43
Aldehydes and Ketones
The Carbonyl Functional Group
C O
The carbon-oxygen double bond can be viewed as the combination
of sp2 hybridized carbon and oxygen atoms.
atomic O
atomic C
2p
2s
1s
valence
level
2p
2s
1s
Examples of Aldehydes and Ketones
Aldehydes
One of the two groups attached to the carbon atom is H.
R
general structure
C O
H
R may be alkyl
or aryl
The simplest aldehyde is formaldehyde with R= H.
Ketones
The two groups attached to the carbon are R and R'. No H.
R
general structure
C O
R'
R and R' are
alkyl or aryl
The simplest ketone is acetone with R, R'= CH3.
Aldehydes
O
H
O
H
formaldehyde
O
H
acetaldehyde
H
benzaldehyde
Ketones
O
acetone
O
O
ethyl methyl ketone
carvone
• Carboxylic Acids, Esters and Amides
• All these groups contain a carbonyl group bonded to
an oxygen or nitrogen
• Carboxylic Acids
– Contain the carboxyl (carbonyl + hydroxyl) group
47
Examples of Carboxylic Acids
formic acid
(R = H)
O
H
propanoic acid
(R = ethyl)
O
HCO2H
HCOOH
O
O
benzoic acid
(R = phenyl)
H
H
CH3CH2CO2H
CH3CH2COOH
O
O
H
C6H5CO2H
C6H5COOH
Derivatives of Carboxylic Acids
There are several families of organic compounds that are considered to
be derivatives of carboxylic acids because their structures can be
related to carboxylic acids through reaction with water (hydrolysis).
O
H
R
O
carboxylic acid
hydrolysis
H2O
O
R
X
a derivative
• Esters
– A carbonyl group is bonded to an alkoxyl (OR’) group
50
• Amide
– A carbonyl group is bonded to a nitrogen derived from
ammonia or an amine
• Nitriles
• An alkyl group is attached to a carbon triply bonded to
a nitrogen
– This functional group is called a cyano group
51
Summary of Important Families of Organic
Compounds
52
Summary (cont.)
53
Quiz Chapter 2 Section 13
Identify the functional group in the following carbonyl compounds.
O
O
O
H
aldehyde
O
ester
O
HCOOH
carboxylic acid
ketone
O
NH2
ketone
amide
Physical Properties and Molecular Structure
Intermolecular forces have a major influence on the physical
properties of compounds. These forces arise from permanent and
transient (temporary) electrical charges within structures that operate
over short distances (several angstroms) and, therefore, are only
important in the condensed phases (solids and liquids).
Ion-Ion Forces
Ionic materials generally exist in well-ordered solid state structures of
alternating positive and negative charges. The strong electrostatic
attractive forces in the solid state are expressed as the Lattice Energy as
defined in the example of sodium chloride.
+
Na(g)
+
Cl(g)
NaCl(s)
Hlattice = -787 kJ/mol
energy
The lattice energy is the energy released when one mole of the
ions condense from the gas phase into a solid structure.
When ionic materials melt, the well-ordered crystalline solid changes
into the more random (less ordered) liquid state with the partial loss of
strong (stabilizing) ion-ion attractive forces. To achieve the less stable
liquid state, considerable thermal energy must be added, resulting in
high melting points for these materials.
Ion-Ion Forces
The introduction of polyatomic ions into the structure tends to lower the
melting point of ionic materials, because these large groups with
covalent bonds increase the distance between the ions, decreasing the
attractive forces.
mp (oC)
ionic materials
ions
sodium chloride
Na+ Cl-
801
sodium acetate
Na+ CH3CO2-
324
ammonium acetate
NH4+ CH3CO2-
114
stronger
ion-ion
forces in
the solid
state
Boiling points of ionic materials are very high and decomposition
often occurs before vaporization.
Dipole-Dipole Forces
Most organic compounds do not have ionic bonds. They possess
covalent bonds. Because of electronegativity differences, many
covalent bonds are polar leading to permanent molecular dipole
moments, when the internal bond dipole moments do not cancel.
A molecule with a permanent dipole moment is polar.
Intermolecular Dipole-Dipole Forces
In the condensed states (liquids and solids), molecules are sufficiently
close together (within a few angstroms) for the dipoles of different
molecules to interact. The negative pole of one dipole attracts the
positive pole of another dipole, while similarly charged poles of
different dipoles repel.
d-
d+ d-
d+
d-
d+
d-
d+
d-
d - d+
d+ d - d+
These dipole-dipole interactions introduce order and stabilize the
solid and liquid states.
Dipole-dipole interactions are stronger in the more ordered solid
state compared with the less ordered liquid state. Polar materials
tend to have higher melting and boiling points than nonpolar
materials of comparable size.
Physical Properties and Molecular Structure
• The strength of intermolecular forces (forces
between molecules) determines the physical
properties (i.e. melting point, boiling point and
solubility) of a compound
• Stronger intermolecular forces result in high
melting points and boiling points
– More energy must be expended to overcome very strong
forces between molecules
• The type of intermolecular forces important for a
molecule are determined by its structure
• The physical properties of some representative
compounds are shown on the next slide
60
Indentify the intermolecular interaction for each molecule
61
• Ion-Ion Forces
• Ion-ion forces are between positively and negatively charged
ions
• These are very strong forces that hold a solid compound
consisting of ions together in a crystalline lattice
– Melting points are high because a great deal of energy is required to
break apart the crystalline lattice
• Boiling points are so high that organic ions often decompose
before they boil
• Example: Sodium acetate
62
• Dipole-Dipole Forces
• Dipole-dipole forces are between molecules with
permanent dipoles
– There is an interaction between d+ and d- areas in each
molecule; these are much weaker than ion-ion forces
– Molecules align to maximize attraction of d+ and dparts of molecules
– Example: acetone
63
• Hydrogen Bonds
• Hydrogen bonds result from very strong dipoledipole forces
• There is an interaction between hydrogens bonded
to strongly electronegative atoms (O, N or F) and
nonbonding electron pairs on other strongly
electronegative atoms (O, N or F)
64
• Example
• Ethanol (CH3CH2OH) has a boiling point of +78.5oC; its
isomer methyl ether (CH3OCH3) has a boiling point of 24.9oC
– Ethanol molecules are held together by hydrogen bonds
whereas methyl ether molecules are held together only by
weaker dipole-dipole interactions
• A factor in melting points is that symmetrical
molecules tend to pack better in the crystalline lattice
and have higher melting points
65
The Influence of Hydrogen Bonds on Physical Properties
Alcohols have higher boiling points than other polar compounds of
comparable size because they form hydrogen bonds in the liquid state
that are lost in the gas state.
=
O
CH3CH2OH
CH3CH
46
44
dipole moment
1.69 D
2.69 D
MP
-115 oC
-121 oC
BP
78.5 oC
20 oC
MW
solubility in water miscible
miscible
Hydrogen bonding raises the boiling point because a network of intermolecular
bonds (1-9 kcal/mol) must be broken in the transition from the liquid to the gas
state. Note this network is not broken in the transition from the solid to the
liquid state, as revealed by the similar melting points of acetone and
acetaldehyde.
Quiz Chapter 2 Section 14
In each pair of structures below, circle the one with the higher boiling
point, and indicate the intermolecular attractive force responsible for the
difference.
O
H
dipole-dipole
O
OH
OH
O
hydrogen bonding
hydrogen bonding
Influence of Molecular Shape
In addition to polar influences, the melting point of organic
compounds is affected by molecular shape. Molecules that are
more symmetrical tend to pack better in the solid state, and have
higher melting points than comparable molecules with other
shapes.
Alcohol Isomers of C4H10O
CH3
CH3-C-OH CH3CH2CH2CH2OH
CH3
tert-butyl
alcohol
MP 25 oC
butyl alcohol
-90 oC
CH3
CH3
CH3CHCH2OH CH3CH2CHOH
isobutyl alcohol
-108 oC
sec.-butyl alcohol
-114 oC
Additional Examples
Alkanes and Cycloalkanes
CH3CH2CH2CH3
CH3CH2CH3
MP
-188 oC
BP
oC
-42
-127 oC
-32 oC
-138 oC
-0.5 oC
-50 oC
12.5 oC
The more compact cycloalkanes have melting points significantly
higher than the comparable alkanes. The boiling points for these
small sized alkanes and cycloalkanes are not very different.
• van der Waals Forces (London or Dispersion
Forces)
• Van der Waals forces result when a temporary dipole
in a molecule caused by a momentary shifting of
electrons induces an opposite and also temporary
dipole in an adjacent molecule
– These temporary opposite dipoles cause a weak attraction
between the two molecules
– Molecules which rely only on van der Waals forces generally
have low melting points and boiling points
71
Polarizability predicts the magnitude of van der Waals Interactions
• Polarizability is the ability of the electrons on an atom to respond to a changing
electric field
• Atoms with very loosely held electrons are more polarizable
• Iodine atoms are more polarizable than fluorine atoms because the outer shell
electrons are more loosely held
• Atoms with unshared electrons are more polarizable (a halogen is more
polarizable than an alkyl of similar size)
All things being equal larger and heavier molecules have higher boiling points
• Larger molecules need more energy to escape the surface of the liquid
• Larger organic molecules tend to have more surface area in contact with each
other and so have stronger van der Waals interactions
• Methane (CH4) has a boiling point of -162oC whereas ethane (C2H6) has a boiling
point of -88.2oC
72
• Solubilities
• Water dissolves ionic solids by forming strong dipoleion interactions
– These dipole-ion interactions are powerful enough to
overcome lattice energy and interionic interactions in the
solid
75
• Generally like dissolves like
• Polar solvents tend to dissolve polar solids or polar liquids
• Methanol (a water-like molecule) dissolves in water in all
proportions and interacts using hydrogen-bonding to the
water
• A large alkyl group can overwhelm the ability of the
polar group to solubilize a molecule in water
– Decyl alcohol is only slightly soluble in water
– The large alkyl portion is hydrophobic (“water hating”) and
overwhelms the capacity of the hydrophilic (“water loving”)
hydroxyl
76
Hydrogen Bonding Between Solute and Solvent
Hydrogen bonding increases the intermolecular attraction between solute and
solvent, promoting solubility. For example, alcohols with small R groups are
miscible (dissolve in all proportions) in water.
d
H
R d d
O H Od
+
H2O
ROH
H d
d
H
intermolecular
intermolecular
etc.
d O
hydrogen bonding
hydrogen bonding
Hd R
hydrogen bonding between
O d
alcohols and water when R is
H d
small
As the size of R increases, the solubility of an alcohol in water decreases. A
long hydrocarbon chain is hydrophobic ("water fearing") while the
hydroxyl group is hydrophilic ("water loving"). The final solubility is a
balance of these two factors.
Guidelines for Water Solubility
The solubility limit of a material is usually given in grams of solute per 100mL of solvent. An organic molecule is considered to be water soluble if at
least 3 g of the solute dissolves in 100-mL of water.
• Organic compounds with one oxygen atom and one to three
carbon atoms are water soluble.
• Compounds with one oxygen and 4 or 5 carbon atoms have
limited solubilities.
• Compounds with one oxygen and 6 carbons or more are usually
insoluble in water.
Spectroscopic Methods for Structure Determination
One of the most powerful tools of science today is spectroscopic methods
of structure analysis. Many of these methods also are used for chemical
analysis to determine the amount and purity of a material.
All of these methods depend on the interaction of a molecule with
electromagnetic radiation. Molecules absorb radiant energy in very well
defined ways that are characteristic of the molecular structure.
Absorption of Electromagnetic Radiation
The absorption of electromagnetic radiation by a molecule may cause
any of several different changes depending on the magnitude of the
energy uptake, E. In all cases, the increase in energy is considered to
be associated with a transition from a lower to higher energy state, so
the uptake in energy is an exact value. Further, only those values of
energy are absorbed that match the energy gaps between states in a
particular molecule.
Molecular Changes
.
energy
higher
state
E = h
lower
state
M+ + eM + hn
ionization
M*
electronic excitation
M*
vibrational/rotational
excitation
M*
nuclear spin excitation
Characterization of Electromagnetic Radiation
Electromagnetic radiation has the properties of a wave, a repeating
mathematical function describing a series of crests and troughs.
The wave describes oscillations in the magnitude of
electric and magnetic fields of electromagnetic
radiation moving in the x direction
moving wave
electric or
magnetic field
x

cycles per
second

wavelength
one complete cycle
one complete cycle
Electromagnetic radiation is characterized either by:
(A) the length of one complete cycle of the wave which is the
wavelength,d
(B) the number of complete cycles passing a point per second
which is the frequency, n (nu), called cycles/second, or hertz or s-1 .
Some Important Relationships
Wavelength and Frequency
The wavelength and frequency of
electromagnetic radiation are related:
 x  = c
where C is the "speed of light"
3.0 x 1010 cm/s (vacuum)
This relationship means that an electromagnetic wave may be
characterized by either the wavelength or frequency of the radiation.
Some fields of spectroscopic study, by tradition, identify the
electromagnetic radiation by wavelength and others use the frequency.
Energy Content of Electromagnetic Radiation
The energy content of electromagnetic
radiation, which defines the increase in
energy when radiation is absorbed by a
molecule, is given by
E = h = hc
where h is Planck's constant.
The Electromagnetic Spectrum
The electromagnetic spectrum is a continuum of radiation energy.
Different sources will emit radiation in specific regions of the spectrum,
and some sources (such as atoms) may emit only very narrow bands or
lines of radiation energy.
The Electromagnetic Spectrum
increasing frequency, 

Hz
1015
1019
cosmic x-rays vacuum
and
uv
-rays

near visible
uv
near
ir
infrared microwave
200 nm 400 nm 800 nm 2 mm
0.1 nm
108
1013
radio
50 mm
increasing wavelength,
energy
equivalent
in kcal/mol
molecular
change
~106
ionization
~102
electronic
excitation
~100
vibrationalrotational
transition
~10-5
nuclear spin
transition
Problem
What is the energy equivalent in kcal/mol of radiation absorbed by a
molecule at 300 nm (nanometers)?
Solution
Wavelength in cm = 300 nm x 10-9 m/nm x 102 cm/m = 3.00 x 10-5 cm
Frequency  = c/ = (3.00 x 1010 cm/s)/(3.00 x 10-5 cm) = 1.00 x 1015 s-1
E = h = (6.62 x 10-27 erg-sec/molecule) x 1.00 x 1015 s-1
= 6.62 x 10-12 ergs/molecule
E = (6.62 x 10-12 ergs/molecule) x 1.44 x 1016 (conversion factor)
= 9.55 x 104 cal/mol
E = 9.55 x 104 x (1 kcal/103 cal) = 95.5 kcal/mol
Absorption of this radiation in the near ultraviolet causes an electronic
transition within the molecule.
Infrared Spectroscopy: An Instrumental Method for Detecting
Functional Groups
Almost all organic molecules are "infrared active," meaning they
absorb radiation in the infrared region. These absorptions are
associated with transitions from lower to higher vibrational levels
within molecules. The energy increase due to absorption of an infrared
photon is typically 1 - 10 kcal/mol.
higher vibrational
state
E = h
lower vibrational
state
The energy gap between the states
depends on the atoms present, the bond
strength, and the details of the molecular
change.
Infrared Spectroscopy: An Instrumental Method for
Detecting Functional Groups
• Electromagnetic radiation in the infrared (IR)
frequency range is absorbed by a molecule at
certain characteristic frequencies
• Energy is absorbed by the bonds in the molecule and
they vibrate faster
• The bonds behave like tiny springs connecting the atoms
– The bonds can absorb energy and vibrate faster only when the
added energy is of a particular resonant frequency
• The frequencies of absorption are very characteristic of
the type of bonds contained in the sample molecule
• The type of bonds present are directly related to the
functional groups present
• A plot of these absorbed frequencies is called an IR
spectrum
88
Spectral Presentation: A Preview
Infrared spectra are recorded by passing a broad spectral range of
the radiation through a sample, and noting the wavelengths that are
absorbed. The pattern of wavelengths absorbed provides structural
information on the organic compound.
5
6
7 10
m (10-4 cm)
15
100
0
4000
.4
bands indicate
absorption of IR
radiation
.8
absorbance
50
.1
8
transmittance (%)
2.5
wavelength
3
4
3000
2000
1000

wavenumber, cm-1
In addition to the spectral wavelength presented in micrometer or
micron units (10-4 cm or 10-6 m), the spectrum is displayed in
wavenumber units (). This unit is linear with frequency with
 = 1/ (with l in cm) = /c.
A Mechanical Model: Hooke's Law
Two atoms connected by a covalent bond behave like two weights on
the ends of a spring. When the weights are displaced from their rest
position and released, they vibrate only with certain frequencies that
depend on the masses of the weights, and a constant related to the
stiffness of the spring..
The mathematical

-1) =  x c =
1

(s
relationship between the
2p M M
MA + MB
A B
allowed frequency (n), the
masses and the constant
where  is the force constant in
(k) is called Hooke's Law..
dynes/cm or g-cm s-2/cm.
In a similar way, the frequency of the radiant energy absorbed during a
vibrational transition depends only on the masses of the atoms and a
force constant related to the strength of the covalent bond .
Vibrational States in Molecules
In molecules, the vibrational states are quantized in a series of allowed
states of increasing energies Evib = (n + 1/2)h where n = 0, 1, 2, 3...etc.
The allowed transitions are between states where n = + 1, which
means only adjacent states. Accordingly, in an allowed transition
between vibrational states, the change in energy is E = h, with the
frequency expressed by Hooke's Law.
+h
lower vibrational
state
(s-1)
= x c =
higher vibrational
state
1
2p

MAMB
MA + MB
where MA and MB are the
masses of the atoms
Some Vibrational Transitions in Diatomics
The infrared spectra of the H-X compounds show a single absorption
band due to a vibrational transition from state n = 0 to state n = 1
with an increase in internal energy of h.
molecule
 (cm-1)
 (s-1)
(measured)

(dynes/cm)
(calculated)
H-F
3958
1.187 x 1014
8.8 x 105
H-Cl
2885
8.655 x 1013
4.8 x 105
H-Br
2559
3.8 x 105
H-I
2230
7.677 x 1013
6.690 x 1013
2.9 x 105
From the experimentally measured values of , and the masses of
the atoms, the force constants may be calculated by Hooke's Law.
The decreasing values of  from H-F through H-I parallel the
decreasing bond strengths of the H-X compounds .
Additional Meaningful Trends
There are several additional meaningful correlations between
observed infrared transitions and covalent bond strengths that follow
from Hooke's Law. In each example below, the atoms (and therefore
the masses) connected by the covalent bond remain the same.
Therefore, the trend in the values of  is due to changing values for
the force constants (bond strengths).
C-H Bond Stretch
characteristic ranges
 (cm-1)
alkanes
alkenes
3000-2850
3100-3000
alkynes
~3300
increasing force constants
(bond strengths)
Carbon-Carbon Bond Stretch
characteristic ranges
 (cm-1)
C
C
800-1200
C
C
1600-1680
increasing force constants
(bond strengths)
C
C
2100-2250
Carbon-Nitrogen Bond Stretch
characteristic ranges
C
N
 (cm-1)
1000-1350
C
C
N
1640-1690
N
2240-2260
increasing force constants
(bond strengths)
Other Normal Modes of Vibration
The normal modes of vibration refers to the possible motions of
atoms with no change in the center of mass of the molecule. The two
fundamental types of vibrations are stretching and bending. As
shown above, stretching is the back and forth movement of atoms
along the bond axis. Bending is the opening and closing of a bond
angle. Various names such as twisting, scissoring and rocking are
used to describe the bending motions of groups of atoms .
Polyatomic Molecules
The only normal mode in a diatomic molecule is the stretch.
Larger Polyatomics
The number of normal modes
in larger molecules depends
on the number of atoms.
polyatomics
3n - 6
linear polyatomics
3n - 5
where n is the number of atoms
• There are 2 basic different types of stretching and
bending vibrations induced by the absorption of infrared
energy
96
Selection Rule for Light Absorption
The interaction between the vibrating molecule and the radiation
field requires a change in dipole moment during the vibrational
transition. This condition must be met for the vibrational mode to
be "infrared active."
Examples
A
B
+h
Cl
H
Cl
H
stretch
B > A
Because the bond length increases in the higher vibrational state, the dipole
moment is larger and this vibrational transition is infrared active .
+h
O
H
H
A
O
symmetrical stretch
H
H
B
B > A
The molecular dipole moment increases in the higher vibrational state because of
longer bonds, so this vibrational transition is infrared active .
N
N
= 0 D
stretch
N
= 0 D
N
Because there is no change
in dipole moment during the
vibrational transition, N2 is
infrared inactive.
Diagram of a Dual Beam Spectrometer
Io
Io
source of
ir radiation
Io
reference
cell
sample
cell
Io'
I
Io'()
wavelength
selector
The intensities of the transmitted
radiation Io' and I are compared
at different l. When Io' = I , there
is no net absorption by the sample.
100% transmission.
I ()
photometer
recorder
or
computer
99
• The actual relative frequency of vibration can be
predicted
– Bonds with lighter atoms vibrate faster than those with heavier
atoms
100
• Triple bonds (which are stiffer and stronger) vibrate at
higher frequencies than double bonds
– Double bonds in turn vibrate at higher frequencies than
single bonds
• The IR spectrum of a molecule usually contains many
peaks
– These peaks are due to the various types of vibrations
available to each of the different bonds
– Additional peaks result from overtone (harmonic) peaks
which are weaker and of lower frequency
– The IR is a “fingerprint” of the molecule because of the
unique and large number of peaks seen for a particular
molecule
101
102
Interpreting IR Spectra
• Generally only certain peaks are interpreted in the IR
– Those peaks that are large and above 1400 cm-1 are most
valuable
• Hydrocarbons
• The C-H stretching regions from 2800-3300 cm-1 is
characteristic of the type of carbon the hydrogen is
attached to
• C-H bonds where the carbon has more s character are
shorter, stronger and stiffer and thus vibrate at higher
frequency
– C-H bonds at sp centers appear at 3000-3100 cm-1
– C-H bonds at sp2 centers appear at about 3080 cm-1
– C-H bonds at sp3 centers appear at about 2800-3000 cm-1
• C-C bond stretching frequencies are only useful for multiple
bonds
– C-C double bonds give peaks at 1620-1680 cm-1
– C-C triple bonds give peaks at 2100-2260 cm-1
– These peaks are absent in symmetrical double and triple bonds
103
• Example: octane
104
• Example: 1- hexyne
105
• Alkenes
• The C-H bending vibration peaks located at 6001000 cm-1 can be used to determine the
substitution pattern of the double bond
106
• Example: 1-hexene
107
• Aromatic Compounds
• The C-C bond stretching gives a set of characteristic
sharp peaks between 1450-1600 cm -1
• Example: Methyl benzene
108
CH3CH2CH2CH2CH2CH2CH2CH3
octane
CH3CH2CH2CH2CH=CH2
1-hexene
CH3CH2CH2CH2C CH
1-hexyne
CH3CH2CH2CH2C CH
1-hexyne
110
Other Functional Groups
• Carbonyl Functional Groups
• Generally the carbonyl group gives a strong peak
which occurs at 1630-1780 cm-1
– The exact location depends on the actual functional
group present
111
Alcohols and Phenols
• The O-H stretching
absorption is very
characteristic
• In very dilute solutions,
hydrogen bonding is absent
and there is a very sharp
peak at 3590-3650 cm-1
• In more concentrated
solutions, the hydroxyl
groups hydrogen bond to
each other and a very
broad and large peak
occurs at 3200-3550 cm-1
• A phenol has a hydroxyl
group directly bonded to an
aromatic ring
112
A Comparison of Aliphatic and Aromatic Alcohols
OH
CH3CHCH2CH3
sec-butyl alcohol
CH2-OH
benzyl alcohol
Some Characteristic Infrared Absorptions
Bond
Compound Type
Frequency Range, cm-1
R C H
alkanes
2850-2960
1350-1470
C C H
alkenes
3020-3080
675-1000
C C H
alkynes
3300
C C H
alkenes
1640-1680
C C H
alkynes
2100-2260
R O H
alcohols
3610-3640 (monomeric)
3200-3600 (H-bonded)
C=O
R N H
carbonyl stretch
aldehydes
ketones
esters
1690-1740
1680-1750
carboxylic acids
1735-1750
amides
1630-1690
amines
3300-3500
1735-1750
Carboxylic Acids
• The carbonyl peak at 1710-1780 cm-1 is very
characteristic
• The presence of both carbonyl and O-H stretching
peaks is a good proof of the presence of a
carboxylic acid
• Example: propanic acid
115
• Amines
• Very dilute solution of 1o and 2o amines give sharp
peaks at 3300-3500 cm-1 for the N-H stretching
– 1o amines give two peaks and 2o amines give one peak
– 3o have no N-H bonds and do not absorb in this region
• More concentrated solutions of amines have broader
peaks
• Amides have amine N-H stretching peaks and a
carbonyl peak
116