Physical Properties and Molecular Structure of Organic compound
** So far, we
have said little about one of the most obvious characteristics of organic compounds—
that is, their physical state or phase. Whether a particular substance is a
solid, or a liquid, or a gas would certainly be one of the first observations
that we would note in any experimental work. The temperatures at which
transitions occur between phases— that is, melting points (mp) and boiling
points (bp)—are also among the more easily measured physical properties.
Melting points and boiling points are also useful in identifying and isolating
organic compounds.
** Suppose, for
example, we have just carried out the synthesis of an organic compound that is known
to be a liquid at room temperature and 1 atm pressure. If we know the boiling point
of our desired product and the boiling points of by-products and solvents that
may be present in the reaction mixture, we can decide whether or not simple
distillation will be a feasible method for isolating our product. In another
instance our product might be a solid. In this case, in order to isolate the substance
by crystallization, we need to know its melting point and its solubility in
different solvents.
** The physical
constants of known organic substances are easily found in handbooks and other
reference books. Table (1) lists the melting and boiling points of some of the compounds.
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| Table (1) |
** Often in the
course of research, however, the product of a synthesis is a new compound— one that
has never been described before. In these instances, success in isolating the
new compound depends on making reasonably accurate estimates of its melting
point, boiling point, and solubilities. Estimations of these macroscopic
physical properties are based on the most likely structure of the substance and
on the forces that act between molecules and ions. The temperatures at which
phase changes occur are an indication of the strength of these intermolecular
forces.
(1) Ionic Compounds: Ion–Ion Forces
** The melting
point of a substance: is the temperature at which an equilibrium exists between
the well-ordered crystalline state and the more random liquid state.
** If the
substance is an ionic compound, such as sodium acetate (Table 1), the ion–ion forces
that hold the ions together in the crystalline state are the strong
electrostatic lattice forces that act between the positive and negative ions in
the orderly crystalline structure.
** In Fig (1) each
sodium ion is surrounded by negatively charged acetate ions, and each acetate
ion is surrounded by positive sodium ions.
** A large
amount of thermal energy is required to break up the orderly structure of the
crystal into the disorderly open structure of a liquid. As a result, the temperature
at which sodium acetate melts is quite high, 324 oC.
** The boiling
points of ionic compounds are higher still, so high that most ionic organic
compounds decompose (are changed by undesirable chemical reactions) before they
boil. Sodium acetate shows this behavior.
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| Fig (1): The melting of Sodium acetate |
(2) Intermolecular Forces (van der Waals Forces)
** The forces
that act between molecules are not as strong as those between ions, but they account
for the fact that even completely nonpolar molecules can exist in liquid and
solid states.These intermolecular forces, collectively called van der Waals
forces, are all electrical in nature.
** We will
focus our attention on three types:
(1) Dipole–dipole
forces
(2) Hydrogen
bonds
(3) Dispersion
forces
(A) Dipole–Dipole Forces
** Most organic
molecules are not fully ionic but have instead a permanent dipole moment resulting
from a nonuniform distribution of the bonding electrons.
** Acetone and
acetaldehyde are examples of molecules with permanent dipoles because the
carbonyl group that they contain is highly polarized. In these compounds, the
attractive forces between molecules are much easier to visualize.
** In the
liquid or solid state, dipole–dipole attractions cause the molecules to orient themselves
so that the positive end of one molecule is directed toward the negative end of
another (Fig. 2).
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| Figure (2) |
(B) Hydrogen
Bonds
** Very strong
dipole–dipole attractions occur between hydrogen atoms bonded to small,
strongly electronegative atoms (O, N, or F) and nonbonding electron pairs on
other such electronegative atoms. This type of intermolecular force is called a
hydrogen bond.
** Hydrogen
bonds (bond dissociation energies of about 4 – 38 kJ mol-1) are weaker than ordinary
covalent bonds but much stronger than the dipole–dipole interactions that occur
above, for example, in acetone.
** Hydrogen
bonding explains why water, ammonia, and hydrogen fluoride all have far higher
boiling points than methane (bp -161.6 oC), even though all four
compounds have similar molecular weights.
** One of the
most important consequences of hydrogen bonding is that it causes water to be a
liquid rather than a gas at 25 oC. Calculations indicate that in the
absence of hydrogen bonding, water would have a boiling point near -80 oC
and would not exist as a liquid unless the temperature were lower than that
temperature. Had this been the case, it is highly unlikely that life, as we
know it, could have developed on the planet Earth.
** Hydrogen
bonds hold the base pairs of double-stranded DNA together. Thymine hydrogen
bonds with adenine. Cytosine hydrogen bonds with guanine.
** Hydrogen
bonding accounts for the fact that ethyl alcohol has a much higher boiling point
(78.5 oC) than dimethyl ether (24.9 oC) even though the
two compounds have the same molecular weight. Molecules of ethyl alcohol,
because they have a hydrogen atom covalently bonded to an oxygen atom, can form
strong hydrogen bonds to each other.
** Molecules of
dimethyl ether, because they lack a hydrogen atom attached to a strongly electronegative
atom, cannot form strong hydrogen bonds to each other. In dimethyl ether the
intermolecular forces are weaker dipole–dipole interactions
** A factor (in
addition to polarity and hydrogen bonding) that affects the melting point of
many organic compounds is the compactness and rigidity of their individual
molecules.
** Molecules
that are symmetrical generally have abnormally high melting points. tert-Butyl
alcohol, for example, has a much higher melting point than the other isomeric
alcohols shown here:
(C) Dispersion Forces
** If we
consider a substance like methane where the particles are nonpolar molecules,
we find that the melting point and boiling point are very low: -182.6 oC
and -162 oC, respectively.
** Instead of
asking, “Why does methane melt and boil at low temperatures?” a more appropriate
question might be “Why does methane, a nonionic, nonpolar substance, become a
liquid or a solid at all?” The answer to this question can be given in terms of
attractive intermolecular forces called dispersion forces or London forces.
** An accurate
account of the nature of dispersion forces requires the use of quantum mechanics.
** We can,
however, visualize the origin of these forces in the following way:
(1) The average
distribution of charge in a nonpolar molecule (such as methane) over a period
of time is uniform.
(2) At any
given instant, however, because electrons move, the electrons and therefore the
charge may not be uniformly distributed.
(3) Electrons
may, in one instant, be slightly accumulated on one part of the molecule, and,
as a consequence, a small temporary dipole will occur (Fig. 3).
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| Figure (3) |
(4) This
temporary dipole in one molecule can induce opposite (attractive) dipoles in surrounding
molecules. It does this because the negative (or positive) charge in a portion
of one molecule will distort the electron cloud of an adjacent portion of
another molecule, causing an opposite charge to develop there.
(5) These temporary dipoles change constantly, but the net result of
their existence is to produce attractive forces between nonpolar molecules and
thus make possible the existence of their liquid and solid states.
** Two
important factors determine the magnitude of dispersion forces.
(1) The
relative polarizability of electrons of the atoms involved.
** By polarizability
we mean how easily the electrons respond to a changing electric field. The electrons
of large atoms such as iodine are loosely held and are easily polarized, while
the electrons of small atoms such as fluorine are more tightly held and are
much less polarizable.
** CF4
and CI4 are both nonpolar molecules. But if we were to consider the
intermolecular forces between two CI4 molecules, which contain
polarizable iodine atoms, we would find that the dispersion forces are much
larger than between two CF4 molecules, which contains fluorine atoms
that are not very polarizable.
(2) The
relative surface area of the molecules involved.
** The larger
the surface area, the larger is the overall attraction between molecules caused
by dispersion forces. Molecules that are generally longer, flatter, or
cylindrical have a greater surface area available for intermolecular
interactions than more spherical molecules, and consequently have greater
attractive forces between them than the tangential interactions between
branched molecules.
** This is
evident when comparing pentane, the unbranched C5H12
hydrocarbon, with neopentane, the most highly branched C5H12
isomer (in which one carbon bears four methyl groups). Pentane has a boiling
point of 36.1 oC. Neopentane has a boiling point of 9.5 oC.
The difference in their boiling points indicates that the attractive forces between
pentane molecules are stronger than between neopentane molecules.
** For large
molecules, the cumulative effect of these small and rapidly changing dispersion
forces can lead
to a large net attraction.
(3) Boiling Points
** The boiling
point of a liquid is the temperature at which the vapor pressure of the liquid equals
the pressure of the atmosphere above it.
** The boiling
points of liquids are pressure dependent, and boiling points are always
reported as occurring at a particular pressure, at 1 atm (or at 760 torr), for
example. A substance that boils at 150 oC at 1 atm pressure will
boil at a substantially lower temperature if the pressure is reduced to, for
example, 0.01 torr (a pressure easily obtained with a vacuum pump). The normal
boiling point given for a liquid is its boiling point at 1 atm.
** In passing
from a liquid to a gaseous state, the individual molecules (or ions) of the substance
must separate. Because of this, we can understand why ionic organic compounds often
decompose before they boil. The thermal energy required to completely separate (volatilize)
the ions is so great that chemical reactions (decompositions) occur first.
** Nonpolar
compounds, where the intermolecular forces are very weak, usually boil at low temperatures
even at 1 atm pressure. This is not always true, however, because of other
factors that we have not yet mentioned the effects of:
(1) molecular weight
(2) molecular
shape and
(3) surface
area.
** Heavier
molecules require greater thermal energy in order to acquire velocities
sufficiently great to escape the liquid phase, and because the surface areas of
larger molecules can be much greater, intermolecular dispersion attractions can
also be much larger.
** These
factors explain why nonpolar ethane (bp -88.2 oC) boils higher than
methane (bp -162 oC) at a pressure of 1 atm. It also explains why,
at 1 atm, the even heavier and larger nonpolar molecule decane (C10H22)
boils at 174 oC.
** The relationship between dispersion forces and
surface area helps us understand why neopentane (2,2-dimethylpropane) has a
lower boiling point (9.5 oC) than pentane (36.1 oC), even
though they have the same molecular weight. The branched structure of
neopentane allows less surface interaction between neopentane molecules, hence
lower dispersion forces, than does the linear structure of pentane.
Solved Problem:
Arrange the following compounds according to their expected boiling points,
with the lowest boiling point first, and explain your answer. Notice that the compounds
have similar molecular weights.
Strategy and
Answer:
** Pentane has
no polar groups and has only dispersion forces holding its molecules together.
It would have the lowest boiling point.
** Diethyl
ether has the polar ether group that provides dipole–dipole forces which are
greater than dispersion forces, meaning it would have a higher boiling point
than pentane.
** sec-Butyl
alcohol has an -OH group that can form strong hydrogen bonds; therefore, it would
have the highest boiling point.
(4) Solubility
** Intermolecular
forces are of primary importance in explaining the solubilities of substances.
** Dissolution
of a solid in a liquid is, in many respects, like the melting of a solid. The
orderly crystal structure of the solid is destroyed, and the result is the
formation of the more disorderly arrangement of the molecules (or ions) in
solution.
** In the
process of dissolving, too, the molecules or ions must be separated from each
other, and energy must be supplied for both changes. The energy required to
overcome lattice energies and intermolecular or interionic attractions comes
from the formation of new attractive forces between solute and solvent.
** Consider the
dissolution of an ionic substance as an example:
(1) Here both the lattice energy and
interionic attractions are large. We find that water and only a few other very polar
solvents are capable of dissolving ionic compounds. These solvents dissolve ionic
compounds by hydrating or solvating the ions (Fig. 4).
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| Figure (4) |
(2) Water
molecules, by virtue of their great polarity as well as their very small,
compact shape, can very effectively surround the individual ions as they are
freed from the crystal surface.
(3) Positive
ions are surrounded by water molecules with the negative end of the water
dipole pointed toward the positive ion; negative ions are solvated in exactly
the opposite way.
(4) Because water
is highly polar, and because water is capable of forming strong hydrogen bonds,
the ion dipole forces of attraction are also large.
(5) The energy supplied
by the formation of these forces is great enough to overcome both the lattice energy
and interionic attractions of the crystal.
General rule for solubility
** A general
rule for solubility is that “like dissolves like” in terms of comparable
polarities:
(1) Polar and
ionic solids are usually soluble in polar solvents.
(2) Polar
liquids are usually miscible.
(3) Nonpolar
solids are usually soluble in nonpolar solvents.
(4) Nonpolar
liquids are usually miscible.
(5) Polar and
nonpolar liquids, like oil and water, are usually not soluble to large extents
** Methanol and
water are miscible in all proportions; so too are mixtures of ethanol and water
and mixtures of both propyl alcohols and water. In these cases the alkyl groups
of the alcohols are relatively small, and the molecules therefore resemble
water more than they do an alkane.
** Another
factor in understanding their solubility is that the molecules are capable of
forming strong hydrogen bonds to each other:
Hydrophobic and Hydrophilic groups
** We often
describe molecules or parts of molecules as being hydrophilic or hydrophobic. The
alkyl groups of methanol, ethanol, and propanol are hydrophobic. Their hydroxyl
groups are hydrophilic.
(a) Hydrophobic:
means incompatible with water (hydro, water; phobic, fearing or avoiding).
(b) Hydrophilic:
means compatible with water (philic, loving or seeking).
** Decyl
alcohol, with a chain of 10 carbon atoms, is a compound whose hydrophobic alkyl
group
overshadows its hydrophilic hydroxyl group in terms of water solubility
** An
explanation for why nonpolar groups such as long alkane chains avoid an aqueous
environment—that
is, for the so-called hydrophobic effect—is complex.
** The most
important factor seems to involve an unfavorable entropy change in the water.
Entropy changes have to do with changes from a relatively ordered state to a
more disordered one or the reverse.
** Changes from
order to disorder are favorable, whereas changes from disorder to order are
unfavorable. For a nonpolar hydrocarbon chain to be accommodated by water, the
water molecules have to form a more ordered structure around the chain, and for
this, the entropy change is unfavorable.
** The presence
of a hydrophobic group and a hydrophilic group are essential components of
soaps and detergents.
** The
hydrophobic long carbon chains of a soap or detergent embed themselves in the
oily layer that typically surrounds the thing we want to wash away.
** The
hydrophilic ionic groups at the ends of the chains are then left exposed on the
surface and make the surface one that water molecules find attractive.
** Oil and
water don’t mix, but now the oily layer looks like something ionic and the
water can take it “right down the drain.”
(5) Guidelines for Water Solubility
** Organic
chemists usually define a compound as water soluble if at least 3 g of the
organic compound dissolves in 100 mL of water.
** We find that
for compounds containing one hydrophilic group—and thus capable of forming strong
hydrogen bonds—the following approximate guidelines hold:
(1) compounds
with one to three carbon atoms are water soluble,
(2) compounds
with four or five carbon atoms are borderline
(3) compounds
with six carbon atoms or more are insoluble.
Note
** When a
compound contains more than one hydrophilic group, these guidelines do not
apply.
** Polysaccharides,
proteins and nucleic acids all contain thousands of carbon atoms and many are
water soluble. They dissolve in water because they also contain thousands of
hydrophilic groups.
(6) Intermolecular Forces in Biochemistry
** Later, after
we have had a chance to examine in detail the properties of the molecules that
make up living organisms, we shall see how intermolecular forces are extremely important
in the functioning of cells.
** Hydrogen
bond formation, the hydration of polar groups, and the tendency of nonpolar
groups to avoid a polar environment all cause complex protein molecules to fold
in precise ways—ways that allow them to function as biological catalysts of
incredible efficiency.
** The same
factors (above) allow molecules of hemoglobin to assume the shape needed to
transport oxygen. They allow proteins and molecules called lipids to function
as cell membranes.
** Hydrogen
bonding gives certain carbohydrates a globular shape that makes them highly
efficient food reserves in animals. It gives molecules of other carbohydrates a
rigid linear shape that makes them perfectly suited to be structural components
in plants.
(7) Summary of Attractive Electric Forces
The attractive
forces occurring between molecules and ions that we have studied so far
are summarized
in Table 2.
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| Table (2) |
Reference: Organic chemistry / T.W. Graham Solomons , Craig B.Fryhle , Scott A.snyder , / ( eleventh edition) / 2014.
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