APPENDIX A: Molecular Structure and Intermolecular Forces
INTRODUCTION:
The electronic structure of a molecule determines the three dimensional shape (molecular structure) of
the molecule, and its polarity. Knowledge of molecular structure enables us to understand both physical and
chemical characteristics of compounds.
The shape of a molecule and placement of it’s various atoms
determines whether it will be polar or non-polar, or capable of hydrogen bonding. Non-polar molecules are
capable of only London dispersion forces with other molecules. Polar molecules are capable of both London
dispersion and dipole-dipole interactions with other molecules. Molecules that have N-H and O-H bonds are
capable of hydrogen bonding, in addition to London dispersion and dipole-dipole attractive forces. Thus, the
structure of a molecule is responsible for its intermolecular forces with other molecules, which in turn controls
the compound’s physical and chemical characteristics. In chromatography the intermolecular forces between an
analyte and the solvent and the analyte and the stationary phase control how the analyte will migrate.
Polarity, Dipole Moment and Dielectric Constant
A bond or molecule is said to be polar if the electron density within the bond molecule is
unevenly distributed. The electron density in a bond or molecule can become unevenly distributed due to the
attraction that highly electronegative atoms have for electrons. Chlorine is much more electronegative than
hydrogen. The uneven distribution of electrons in H-Cl leads to the development of partially positive (δ+) and
partially negative (δ-) ends of the bond. If the electron distribution in the bond or molecule is symmetrical, then
the bond or molecule is non-polar (Cl-Cl).
δ+
δ-
H
Cl
Cl
μ = 1.08 D
Cl
μ = 0.00 D
More quantitative measures of polarity are dipole moment and the dielectric constant. If
charges q are separated by a distance r, the dipole moment (μ) is defined as the product between q and r. The
dipole moment is represent as an arrow going from the positive to the negative. The dipole moment of some
common organic solvents are listed in Table 1.
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Table 1. Solvent properties.
Solvent
MF
MW
Hexane
CH3(CH2)4CH3
Toluene
C6H5CH3
o
Hazards*
Dipole
Dielectric
constant
C6H14
86.17
Bp ( C)
Density
(g/mL)
68.7
0.659
Flammable
Toxic
0.08
1.89
Elution
Stength
(ε)
0.01
C 7H 8
92.13
110.6
0.867
Flammable
Toxic
0.31
2.4
0.22
C4H10O
74.12
34.6
0.713
Flammable
Toxic, CNS
Depressant
1.15
4.34
0.29
CH3CH2OCH2CH3
Dichloromethane
CH2Cl2
CH2Cl2
84.94
39.8
1.326
1.14
8.93
0.32
Ethyl Acetate
CH3CO2CH2CH3
C 4H 8O 2
88.10
77.1
0.901
Toxic, Irritant
Cancer
suspect
Flammable
Irritant
1.88
6.02
0.45
Acetone
CH3COCH3
C 3H 6O
58.08
56.3
0.790
Flammable
Irritant
2.69
20.7
0.43
Butanone
CH3CH2COCH3
C 4H 8O
72.10
80.1
0.805
Flammable
Irritant
2.76
18.5
0.39
1-Butanol
C4H10O
74.12
117.7
0.810
Flammable
Irritant
1.75
15.8
0.47
CH3CH2CH2CH2OH
Propanol
CH3CH2CH2OH
C 3H 8O
60.09
82.3
0.785
Flammable
Irritant
1.66
19.92
0.63
Ethanol
CH3CH2OH
C 2H 6O
46.07
78.5
0.789
Flammable
Irritant
1.70
26.0
0.68
Methanol
CH3OH
CH4O
32.04
64.7
0.791
Flammable
Toxic
1.7
32.7
0.73
Water
HOH
H 2O
18.02
100.0
0.998
1.87
80.1
>1
Diethyl ether
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The polarity of solvents is also often described by the dielectric constant, ε (Table 1). The dielectric
constant is a ratio and therefore does not have a unit. In general, the dielectric constant of a solvent measures its
capacity to cause particles of opposite charge to separate from one another (Figure 1). Very polar molecules,
such as water (ε = 80.1) can orient themselves around the particles and shield the attraction between the
oppositely charged particles. Non-polar solvents, such as hexane (ε = 2.0) have no net dipole and cannot
effectively shield charged particles from each other. Nonpolar solvents are therefore are poor at dissolving salts
or very polar compounds. Dielectric constants for many commonly used solvents for chromatographic
separation of organic compounds are given in Table 1.
δ+
δ−
δ
δ−
+
δ
+
−
δ−
δ−
δ+
δ+
δ+
-
δ−
δ−
δ+
δ
+
δ+
δ−
Figure 1. Ion-dipole forces between ions and a polar solvent.
Hydrogen Bonding
Of the elements most commonly found in organic compounds (C, H, N, O, Cl, Br, S, P), oxygen
is the most electronegative element and hydrogen is the least electronegative.
The large difference in
electronegativities between hydrogen and oxygen makes the O-H bond very polar. The very positive hydrogen
end of the bond is actually able to form another partial bond (a hydrogen bond) with a partial negatively charged
nitrogen or oxygen of another molecule. The molecule that provides the hydrogen is called the H-bond donor
and the one that accepts the hydrogen is called the acceptor. Not all H-bond acceptors need be H-bond donors.
For example, the partially negatively charged oxygen atoms of aldehydes, ketones, ethers, esters and nitrocompounds can accept H-bonds, but not donate them. The ability to H-bond has far reaching consequences in
organic chemistry. Compounds that can donate and accept H-bonds (those with N-H or O-H bonds) have higher
boiling and melting points than compounds with similar molecular mass, but without N-H or O-H bonds. Also,
N-H or O-H containing compounds are generally more soluble in water than compounds without these bonds.
H-bonding is also critical in helping to determine the structure and function of numerous biomolecules, such as
proteins, DNA and RNA, to mention a few. In chromatography, inclusion of a solvent with hydrogen bonding
capabilities can often change elution order of certain analytes.
Methanol is a very polar molecule. It is capable of H-bonding both with itself and with other
oxygen or nitrogen containing molecules, like water. Thus, methanol is completely soluble in water. However,
with a single methyl group, methanol has only weak London dispersion forces with itself and with other
43
molecules. As the size of the alkyl chain of an alcohol increases (as in ethanol, propanol and butanol), the
importance of London dispersion forces between molecules increases and may become as great or greater than
the stronger dipole-dipole and H-bonding intermolecular forces in alcohols. As a result, methanol is insoluble in
alkanes, such as hexane and heptane, whereas ethanol and propanol of completely miscible with these alkanes.
Because they are immiscible, a combination of methanol and hexane is not a good chromatographic solvent
system. However, hexane plus ethanol or propanol are often very good solvent systems.
+
δ− δ
O H
δ−
O
acceptor
donor
+
δ− δ
N H
δ−
O
δ − δ+
O H
donor
acceptor
donor
+
δ− δ
N H
δ−
N
acceptor
donor
δ−
N
acceptor
Figure 2. Hydrogen Bonding
H-Bond donors and acceptors
R ..
O H
water
alcohols
phenols
R N H
H
R N H
R
O ..
R C O
.. H
1o amines
2o amines
carboxylic
acids
O
..
R C N H
R
..
..
..
..
..
OH
..
..
..
..
..
H O
.. H
amides
ethers
aldehydes
ketones
..
..
O
R C R
..
..
R ..
O R
O
R C H
O ..
R C O
.. R
esters
44
..
R N R
R
3o amines
R
.
.O
..
N .. +
O
..
..
..
..
..
H-Bond acceptors
nitro compounds
APPENDIX B: Physical Constants
It is very important to know physical constants of all compounds used in the lab, before starting the lab.
Knowing the density and boiling point of liquids, the melting point of solids and the toxicity and hazards of all
chemicals is crucial to success in the laboratory. This information can be found in a variety of sources, a few of
which are: the CRC Handbook of Chemistry and Physics, the Merck Index and the Aldrich Catalog. There are
now numerous web sites with physical constants, such as http://hazard.com/, www.chemexper.com, as well as
web sites of chemical companies such as www.sigmaaldrich.com and www.fishersci.com.
In most handbooks, chemicals are listed alphabetically, but they may be found under different
names.
For example, aspirin is found under the name Aspirin in the Merck Index, under the name
Acetylsalicylic acid in the Aldrich Catalog and under Benzoic Acid, 2-(acetyloxy) in the CRC Handbook, 76th ed.
If you know the molecular formula (C9H8O4 for aspirin) you may also use the formula index. Often, several
compounds will have the same molecular formula so you must take care to find the correct entrée.
The information provided by these sources generally includes the molecular formula, molecular
weight (MW) or molecular mass (MM), boiling point (bp), melting point (mp), density (d) and hazards, although
other physical constants may also be listed as well, depending on the source.
The molecular weight is given in grams per mole (g/mol).
The boiling point is given in oC. However, the boiling point is dependent on the pressure at
which it is taken. The “normal boiling point” is determined at 1 atmosphere (760 torr or 760 mm Hg). The
pressure at which the boiling point was determined is usually given as a subscript. For example, for methanol
the Merck Index lists: bp76064.7oC and bp20034.8oC.
The melting point is given in oC. However, the melting point is dependent on the purity of a
compound. Most often, melting ranges are given. Very pure compounds have sharp melting ranges (within 1oC).
For example, the Aldrich Chemical Catalog lists the melting range of acetanilide as:
mp 113-115oC.
The density of liquids is reported in grams per millilter (g/mL). However, the density of a
liquid is dependent upon temperature. The Merck Index lists the density of methanol as: d0 0.8100 and d20
0.7915. The superscript indicates that temperature at which the density was measured, and the subscript 4
indicates the the density of water at 4oC (1.000g/mL) was used as the reference.
You should always be aware of the toxicity and hazards or the chemicals before you use them.
Often many chemicals will have very long MSDS listings. You should try to focus on whether the compound is
flammable, toxic or an irritant. Table 1 provides all of the important physical constants for most of the organic
solvents we will use during the semester. It also lists flammability and permissible exposure limits, which
serves as a measure of the toxicity of the solvent.
45