Be very, very careful what you put into that head,
because you will never, ever get it out.
Cardinal Wolsey (1475?-1530)

Bad Chemistry

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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@chemvax.princeton.edu.

Brought to you by: Kevin Lehmann and the Princeton Section of the American Chemical Society.

There have been visitors to this page since March 10, 1996

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 Centigrade.

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 the water.

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!

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