Be very, very careful what you put into that head,
because you will never, ever get it out. Cardinal Wolsey (1475?-1530)
Click on the symbol for its explanation.
This page is intended to be part of the Bad Science pages established
by Alistair B. Fraser. The purpose of this page is to bring to light
commonly mistaught concepts in the field of Chemistry. The intended
audience is secondary school students and their teachers. The page
is at present just beginning, and I would welcome additions which I
would include with due acknowledgement. If you wish to contact me,
please send E-mail to Lehmann@virginia.edu
Brought to you by: Kevin Lehmann.
The Hydrophobic Effect
Theory of Ice skating is all wet!
Ionic solutions don't look like that!
The Hydrophobic Effect does not mean that nonpolar molecules are not
attracted to water!
We have all seen what happens after vinegar and oil salad dressing are
vigorously shook; one does get a mixture of sorts, but after a little
time the ingredients separate with the lighter oil on top and a denser
vinegar/water solution on bottom. This is an illustration of an important
chemistry principle expressed by the rule that 'like dissolves like.'
This refers to the phenomena that when two liquids made of molecules
of similar size and polarities are mixed, they will usually form a single
phase solution, no matter what the relative number of moles of each species.
This is expressed by the jargon that the two substances are miscible
in all proportions. In contrast, when a highly polar substance, such
as water, is mixed with a nonpolar or weakly polar substance, such as
most oils, the substances will separate into two phases. This phenomenon
is usually rationalized in introductory chemistry text books by saying
that oil is hydrophobic, and thus does not make solutions with water,
while polar small organic acids (such as acetic acid from which house
vinegar is made) are hydrophilic, and thus are miscible with water.
This explanation almost universally leads students (and even some professional
chemists) to believe that individual water and oil molecules repel each
other, or at least attract each other very weakly. Nothing can be further
from the case! An individual oil molecule is attracted to a water molecule
by a force that is much greater than the attraction of two oil molecules
to each other.
We can observe the consequence of this greater attraction when we put
a drop of oil on a clean surface of water. Before hitting the surface,
the oil will be in the shape of a spherical droplet. This is because
the oil molecules are attracted to one another and a spherical shape
minimizes the number of oil molecules that are not surrounded by other
molecules. When the oil hits the surface of the water, it spreads out
to form a thin layer. This happens because the attractions between the
oil and water molecules gained by spreading over the surface is larger
than the oil-oil attraction lost in making a large oil surface on top
of the water. If a sufficiently small drop of oil is put on the surface,
it will spread to form a single molecular layer of oil. By measuring
the area produced, one can get a simple estimate for the size of each
oil molecule and thus Avogadro's number.
Given these strong interactions, why does not each oil molecule dive
into the water solution and surround itself with the favorable water
attractions? The reason is that to do so, it must come between water
molecules that are already attracting each other! The strength of water-water
attraction is much higher than water-oil interactions, and thus there
is a net cost of energy in putting the oil molecules into a water solution.
Thus the vast majority of oil molecules stay out of the water, though
as many as will fit will hang on to the surface water molecules that
do not have a full complement of partners.
One can make an analogy to a high school dance, where generally one
will find that the most popular boys and girls will dance with each other
almost exclusively. It is not that the less popular members of the class
do not want to dance with the popular students, it is just that the popular
students have their choice of who to dance with and, since they also
want to dance with popular students, they pair off. The less popular
students, rather than have no partners, will naturally pair off as well.
Common Theory of Ice Skating is all Wet!
Water has many unusual properties. One is that the low pressure solid
form (what we know as ice) has a volume per mole that is ~10% larger
than that of liquid water into which it melts. An everyday consequence
of this fact is that ice cubes float near the surface of water, with
about 10% of their volume above the water-air surface and 90% below.
Another consequence of this decrease in volume upon melting is that the
melting temperature of ice decreases when one increases the pressure
on the ice. This can be rationalized by Le Châtelier's Principle.
An increase in pressure on a sample of ice can be partially removed by
melting the ice and thus lowering the sample volume. However, the effect
is a small one in that it takes a pressure of ~121 atmospheres (1.22
MPa in SI units) to reduce the melting temperature by a mere 1 degree
It is often claimed that one can skate on ice because the pressure of
the skate causes the ice to melt, thus dramatically reducing the friction
between skate and ice. While this makes a good story, it is not quite
correct. If one takes the skater to have a mass of 75 kg (weight of 165
lbs), and the skate to be 3 mm wide and 20 cm long, one can calculate
that entire gravitational force exerted on the area of one skate is only
a pressure of about 12 atmospheres. While one can imagine that the force
is concentrated in a somewhat smaller area, the effect of pressure alone
is clearly enough to shift the melting temperature of the ice by at most
a few tenths of a degree. Since common experience is that ice skating
is possible even when the ambient temperature is well below the normal
freezing point, the pressure induced lowering of the melting point clearly
does not explain this every day observations.
What is responsible then? Scientists have far from a complete understanding
of this everyday phenomenon. It is likely partially related to an effect
known as surface melting. The stability of solids is due to the regular
structure that allows for each molecule to have multiple attractive interactions.
At the surface of a solid, this is not the case, since there are no molecules
'above' the surface to bind to. As a result, the surface molecules will
often distort to make the best of a bad situation by trying to increase
their bonding to each other and those below. This is known as surface
reconstruction. It is also known that the molecules on the surface can
become disordered and liquid like at a temperature below the normal melting
point of a solid, this is the phenomenon known as surface melting. Bringing
up another surface (such as the metal of a skate) will influence this
surface melting, since now the water molecules on the surface can bind
to the metal surface atoms as well. Another important effect is friction,
which can generate enough heat to melt a thin layer of ice in contact
with the skate.
Ionic Solutions Don't Look Like That!
It is widely known that adding salt to water will slightly raise the
boiling temperature. This effect can be easily demonstrated by placing
a small glass of water inside a pan partially filled with highly salted
water and then heating on the stove. If the pan is slowly heated, one
will observe that the water inside the glass will begin to boil before
the water-salt solution. Most introductory chemistry text books explain
that the magnitude of the boiling point increase is proportional to
the fraction of solute particles inside the solution. By boiling point,
I mean the temperature at which the equilibrium vapor pressure of the
solution is equal to the external pressure, commonly due to the atmosphere,
upon the surface of the liquid. It is important to remember that boiling,
the rising of rapidly growing bubbles that rise through the liquid,
is a nonequilibrium phenomenon which begins at an even higher temperature.
Many textbook rationalize this effect by reference to a figure such as
that shown below:
The figure explains the lowering of the equilibrium solvent vapor pressure
of a solution at fixed temperature relative to that of the pure solvent.
The nonvolatile solute molecules are viewed as partially blocking the
escape of solvent molecules from the surface of the solution. In order
to boil the solution, the temperature must be raised to once again
make the solution vapor pressure equal to the external pressure exerted
by the atmosphere upon the surface of the solution.
While this explanation is appealing, it gives a very misleading picture
for what real ionic solutions are like. Ionic salts, such as NaCl, are
only soluble in water because the energy of solvation of the ions is
almost as large as the enormous energy associated with the Coulombic
attraction of the positive and negative ions in the crystal. Because
of this strong solvation, one will essentially never find a Na(+) or
Cl(-) sitting on the surface of the solution since there it would not
have its full solvation shell of water molecules around it. The surface
of even a highly concentrated salt solution is essentially pure water!
Naively following the picture given above, one would be lead to predict
that there is no decrease in vapor pressure when salt is dissolved in
Another failing of the above figure is that it appears to imply that
the extent of the vapor pressure lowering will depend upon the size of
the solute molecules relative to the solvent. In fact, this is not correct.
Using methods of Chemical Thermodynamics such as are typically found
in a Physical Chemistry text book, one can show that for sufficiently
dilute solutions, the magnitude of the vapor pressure decrease depends
only upon the ratio of solute to solvent molecules, and not at all on
the properties of the solute molecules!
Another common chemistry misconception that is related to boiling point
elevation is the reason often given for adding salt to water when boiling
foods. It is often stated that this is to increase the temperature of
the boiling water and thus speed the rate of cooking. It is certainly
true that a small increase in cooking temperature can significantly increase
the rate of cooking; cooking times will typically be only half as long
if the water temperature is raised by 10 - 20 C. However, even if we
make the cooking water as salty as sea water (which requires adding twelve
tablespoons of table salt per gallon of water!), the boiling point will
only increase by 0.6 C which will only decrease the cooking time by a
few percent. If you are such a Type A personality that you feel compelled
to save even this small cooking time, then the last thing you need is
to risk increasing your blood pressure further by consuming so much sodium!