Physical Chemistry


Heat

In the seventeenth and eighteenth centuries the worlds of chemistry and physics seemed well separated. Chemistry was the study of those changes that involved alterations in molecular structure. Physics was the study of those changes that did not involve such alterations.

In the early nineteenth century, while Humphry Davy was altering the molecular arrangement of inorganic compounds and Pierre Eugene Marcelin Berthelot was altering the molecular arrangement of organic compounds, physicists were studying the flow of heat. This study of the flow of heat is called thermodynamics (from Greek words for "heat movement").

Prominent in this field were the English physicist James Prescott Joule (1818-1889) and the German physicist Julius Robert von Mayer (1814-1878) and Hermann Ludwig Ferdinand von Helmholtz (1821-1894). By the 1840's their work made it clear that those changes or variations undergone by heat and other forms of energy, no energy was either created nor destroyed. This principle is called the law of conservation of energy or the first law of thermodynamics.

The work of the French physicist Nicolas Leonard Sadi Carnot (1797-1832), the English physicist William Thomson, later Lord Kelvin (1824-1907), and the German physicist Rudolf Julius Emanuel Clausius (1822-1888) went further. It was shown that heat flowed spontaneously from a point at higher temperature to one of lower temperature, and that work could be obtained from heat only when such a heat flow across a temperature-difference existed. This inference could be generalized to apply to any form of energy flowing from a point of higher intensity to one of lower.

Clausius, in 1850, invented the term entropy for the ration of the heat content of an isolated system to its absolute temperature. He showed that in any spontaneous energy change the entropy of the system would increase. This principle is called the second law of thermodynamics.

But this sort of advance in physics could not be isolated from chemistry. After all, apart from the sun, the major sources of heat in the nineteenth century world lay in chemical reactions such as the burning of wood, coal, or oil.

Other chemical reactions also evolved heat, the neutralization of acids by bases, for instance. In fact, all chemical reactions involved some sort of heat transfer, either the emission of heat (sometimes light) to the outside world, or the absorption of heat (sometimes light) from the outside world.

It was in 1840 that the worlds of physics and chemistry met and began to fuse in the work of a Swiss-Russian chemist, Germain Henri Hess (1802-1850). He announced the results of careful measurements he had made of the actual quantity of heat evolved in the chemical reactions, of fixed quantities of some substances. He was able to demonstrate that the quantity of heat produced (or absorbed) in going from one substance to another was the same no matter by what chemical route the change occurred, or in how many stages. Because of this generalization (Hess's law), Hess is sometimes considered the founder of thermochemistry (heat-chemistry).

Hess's law made it seem highly likely that the law of conservation of energy applied to chemical changes as well as to physical changes. Indeed, to generalize further, the laws of thermodynamics very likely held in chemistry as in physics.

This line of experiment and reasoning made it seem that chemical reactions, like physical processes, had an inherent spontaneous direction in which entropy was increased. Entropy is a difficult quantity to measure directly, however, and chemists sought other and simpler criteria that would serve as the measure of the "driving force".

In the 1860's, Berthelot, who had done such important work in organic synthesis turned to thermochemistry. He devised methods for conducting chemical reactions within a closed chamber surrounded by water at known temperature. From the rise in the temperature of the surrounding water at the conclusion of the reaction, the quantity of heat evolved by the reaction could be measured.

Using such a calorimeter (from the Latin for heat-measure), Berthelot ran careful determinations of the quantity of heat evolved by hundreds of different chemical reactions. Independently, the Danish chemist Hans Peter Jorgen Julius Thomsen (1826-1909) did similar experiments.

Berthelot felt that reactions that give off heat were spontaneous, while those that absorbed heat were not. Since every reaction that gave off heat had to absorb heat when forced into reverse (Lavoisier and Laplace were the first to hold such views), this meant that any chemical reaction would move spontaneously in only one direction and would give off heat while doing so.

As an example, when hydrogen and oxygen combine to form water, the reaction gives off a great deal of heat. This reaction is spontaneously, and, once started, goes rapidly to completion - sometimes with explosive violence.

The reverse reaction of water breaking down into hydrogen and oxygen requires an input of energy. The energy may be in the form of heat or electricity. Such a breakdown of the water molecule is not spontaneous. It does not seem to occur at all until energy is supplied and, even then, the reaction ceases the moment the energy input is interrupted.

But Berthelot's generalization, however plausible it seems on the surface, is flawed. In the first place, not all spontaneous reactions give of heat. Some absorb heat so that, as they proceed, the temperature of the surroundings actually drops.

Secondly, there are reversible reactions. In these, substances A and B can react spontaneously and be converted to substances C and D, while C and D can, just as spontaneously, react back to A and B. And all this happens despite the fact that if heat is given off in the reaction occurring in one direction it must be absorbed in the reverse reaction. A simple example is that of hydrogen iodide, which breaks down to a mixture of hydrogen and iodine. The mixture is capable of recombining to hydrogen iodine. This can be written in equation form:

2HI <----> H2 + I2

The double arrow indicates a reversible reaction.

Reversible reactions were already known in Berthelot's time. They were first carefully studied, in 1850, by Williamson in the course of the work which led to his conclusions concerning ethers. He found situations in which, beginning with a mixture of A and B, the substances C and D were formed. If he began instead with a mixture of C and D, substances A and B were formed. In either case, there would be a mixture of A, B, C, and D in the end, with the proportions apparently fixed. The mixture would be at an equilibrium.

Williamson, however, did not believe that because the composition of the mixture was apparently fixed, nothing was happening. He felt that A and B were reacting to C and D, while C and D were reacting to A and B. Both reactions were in constant progress but neutralized each other's effects, giving the illusion of rest. This condition was dynamic equilibrium.

Williamson's work marked the beginning of the study of chemical kinetics - the study of the rates of chemical reactions. It was quite evident from Williamson's work that something more than the mere evolution of heat dictated the spontaneity of the chemical reaction. This "something more" was already being worked out while Berthelot and Thomsen were making their numerous calorimetric measurements. But, unfortunately, the matter remained buried in a little-known language.

Chemical Thermodynamics

In 1863, the Norwegians chemists Cato Maximillian Guldberg (1836-1902) and Peter Waage (1833-1900) published a pamphlet dealing with the direction of spontaneous reactions. They returned to a suggestion made half a century before by Berthollet, that the direction taken by a reaction depended upon the mass of the individual substances taking part in the reaction.

To Guldberg and Waage, it seemed that mass alone was not the entire answer. Rather it was a question of the amount of mass of a particular substance crowded into a given volume of the reacting mixture, or the concentration of the substance, in other words.

Suppose A and B can react to form C and D, while C and D can react to form A and B. This double reaction can be represented as follows:

A + B <----> C + D

The situation symbolized is an example of one of Williamson's reversible reactions, and it reaches an equilibrium under conditions in which A, B, C, and D all exist in the system. The point of equilibrium depends on the rate at which A and B react (rate 1) as compared with that at which C and D react (rate 2).

Suppose rate 1 is much higher than rate 2. In that case, A and B are reacting quickly, producing a considerable quantity of C and D; While C and D react slowly and produce a small amount of A and B. Before long, most of the A and B has changed over to C and D and little has changed back. When the reaction reaches equilibrium, then, C and D dominate the mixture. Looking at the equation just above, we would say that the equilibrium point is "far to the right".

The reverse is true when rate 2 is much higher than rate 1. In that case, C and D would react to produce A and B much more quickly than A and B would react to produce C and D. At equilibrium, A and B would dominate the mixture. The equilibrium point is then "far to the left".

But rate 1 depends on how frequently a molecule of A happens to collide with a molecule of B, for only upon such a collision can reactions occur, and not always even then. Again, rate 2 depends on how frequently a molecule of C collides with a molecule of D.

Suppose, then, that additional A or B (or both) is added to the system without changing its volume. The concentration of A or B (or both) is increased and there is now a greater likelihood of collisions among them (just as there is a greater likelihood of automobile collisions when a highway is crowed at rush hour than when it is relatively empty at midmorning.)

Increasing the concentration of A or B or both thus increases rate 1; decreasing the concentration will decrease the rate. Similarly, an increase in the concentration of C or D or both, will increase the concentration of rate 2. By altering rate 1 or rate 2, one can alter the composition of the equilibrium mixture. If the concentration of any of the participating substances is altered, therefore, the position of the equilibrium point is changed.

Though the concentrations of A, B, C, and D at equilibrium would shift as one or more of these components were added to or taken from the mixture, Guldberg and Waage found they could cling to one unchanging factor. The ratio of the product of the concentrations of substances on one side of the double arrow to the product of the concentrations on the other side of the double arrow, at equilibrium, remains constant. (The law of chemical equilibrium)

Suppose we represent the concentration of a given substance by placing brackets about its symbol. We can say, then, in connection with the reaction we have been discussing, that, at equilibrium:

[C][D]/[A][B] = K

The symbol K, represents the equilibrium constant, which is characteristic for any given reversible reaction being run at a fixed temperature.

Guldberg and Waage's law of mass action was an adequate guide to the understanding of reversible reactions, much more so then Berthelot's suggestion. Unfortunately, Guldberg and Waage published in Norwegian, and their work went unnoticed until 1879, when it was translated into German.

In the meantime, an American physicist, Josiah Willard Gibbs (1839-1903), was systematically applying the laws of thermodynamics to chemical reactions. He published a number of long papers on the subject between 1876 and 1878.

Gibbs evolved the notion of free energy, a quantity which incorporated within itself both heat content and entropy. When a chemical reaction occurred, the free energy of the system changed. When the free energy decreased, the entropy always increased, and the reaction was spontaneous. (The value of the free energy lay in the fact that its change was easier to measure than the change in entropy.) The change in heat content depended on the exact amount by which free energy decreased and entropy increased. Usually, the heat content also decreased in a spontaneous reaction so that heat was given off. Occasionally, though, the change in free energy and entropy was such that the heat content increased and then a reaction, though spontaneous, absorbed energy.

Gibbs also showed that the free energy of a system changes somewhat with changes in the concentration of the chemicals making up that system. Suppose that the free energy of A + B is not much different from that of C + D. Then, the small changes introduced by changes in concentration might be enough to make the free energy of A + B more than that of C + D at some concentrations and less at others. The reaction could move spontaneously in one direction at one set of concentrations and in the opposite direction (but just as spontaneously) at another set.

The rate at which free energy changes as the concentration of a particular substance changes is the chemical potential of that substance, and Gibbs could show that it was the chemical potential that acted as the "driving force" behind chemical reactions. A chemical reaction moved spontaneously from a point of high chemical potential to one of low, as heat flowed spontaneously from a point of high temperature to one of low.

In this way, Gibbs gave meaning to the law of mass action for he showed that at equilibrium the sum of the chemical potentials of all the substances involved was at a minimum. If one began with A + B, it moved down the chemical potential "hill" and C + D was formed. If one began with C + D, it moved downward as A + B was formed. At equilibrium, the bottom of the "energy valley" between the two "hills" had been reached.

Gibbs went on to apply thermodynamic principles to equilibria between different phases (liquid, solid, and gas) included within a particular chemical system. For instance, liquid water and water vapor (one component, two phases) could exist together at different temperatures and pressures, but if the temperature was fixed, the pressure was fixed also. Liquid water, water vapor, and ice (one component, three phases) could exist all together at only one particular temperature and pressure.

Gibbs worked out a simple equation, the phase rule, which enabled one to predict the manner in which temperature, pressure, and the concentrations of various components could be varied under all combinations of components and phases.

Thus was founded chemical thermodynamics in such detail and with such thoroughness that little was left to be done by those who came after Gibbs. (An example of one important addition was that introduced by the American chemist Gilbert Newton Lewis (1875-1946). In 1923, in a classic book on thermodynamics, he introduced the concept of activity. The activity of a substance is not identical with its concentration but is related to it. The equations of chemical thermodynamics can be made more accurate over a wider range, if activity is substituted for concentration.) Nevertheless, despite the fundamental importance and remarkable elegance of Gibb's work, it did not at once receive recognition in Europe, since it was published in an American journal that was ignored by the European leaders in the field.

Catalysis

In the final quarter of the nineteenth century, Germany was leading the world in the study of physical changes associated with chemical reactions. The outstanding worker in this field of physical chemistry was the Russian-German chemist Friedrich Wilhelm Ostwald (1853-1932). It was thanks to him, more than to any other individual, that physical chemistry came to be recognized as a discipline in its own right. By 1887, he had written the first textbook on the subject and founded the first journal to be devoted exclusively to the field.

Ostwald was among the first Europeans to discover and appreciate Gibbs's work. He translated Gibbs's papers on chemical thermodynamics into German in 1892. Ostwald proceeded to put Gibbs's theories to use almost at once in connection with the phenomenon of catalysis.

Catalysis (a word suggested by Berzelius in 1835) is a process whereby the rate of a particular chemical reaction is hastened, sometimes enormously so, by the presence of small quantities of a substance which does not itself seem to take part in the reaction. Thus, powdered platinum will catalyze the addition of hydrogen to oxygen and to a variety of organic compounds, as Davy (the isolator of sodium and potassium) discovered in 1816. Again, acid will catalyze the breakdown to simpler units of a number of organic compounds, as G.S.Kirchhoff first showed in 1812. At the conclusion of the reaction, the platinum or the acid is still present in its original quantity.

Ostwald prepared, in 1894, a summary of someone else's paper on the heat of combustion of foods, this summary to appear in his own journal. He disagreed strongly with the conclusions of the writer, and to buttress his disagreement discussed catalysis.

He pointed out that the theories of Gibbs made it necessary to assume that catalysts hastened reactions without altering the energy relationships of the substances involved. The catalyst, he maintained, must combine with the reacting substance to form an intermediate that breaks up to give the final products. The breakup of the intermediate released the catalyst, which thus resumed its original form.

Without the presence of this catalyst-combined intermediate, the reaction would proceed much more slowly, sometimes so slowly as to be imperceptible. Hence, the effect of the catalyst was to hasten the reaction without itself being consumed. Furthermore, since a molecule of catalyst was used over and over, a small quantity of catalyst was sufficient to hasten a great deal of reaction.

This view of catalyses is still held today. It has helped to explain the activity of the protein catalysts (or enzymes) which control the chemical reactions in living tissue.

Ostwald was a follower of the principles of the Austrian physicist and philosopher Ernst Mach (1838-1916), who believed that scientists should deal only with matters that could be directly measured, and should not create "models" based only on indirect evidence. For this reason, Ostwald refused to accept the reality of atoms, since there was no direct evidence for their existence. He was the last important scientist to resist the atomic theory (though he did not deny its usefulness).

Here the matter of Brownian motion came up. This phenomenon, involving the rapid, irregular motion of small particles suspended in water, was first observed in 1827 by a Scottish botanist, Robert Brown (1773-1858).

The German-Swiss physicist Albert Einstein (1879-1955) showed, in 1905, that this motion could be attributed to the bombardment of the particle by molecules of water. Since, at any given moment, more molecules might be striking from one direction than from the other, the particles would be pushed now here, now there. Einstein worked out an equation which could be used to calculate the actual size of the water molecules once certain properties of the moving particles were measured.

A French physicist, Jean Baptiste Perrin (1870-1942), made the necessary measurements in 1908 and produced the first hard and fast estimate of the diameter of molecules and, therefore, of atoms. Since the Brownian motion was a reasonably direct observation of the effects of individual molecules, even Ostwald had to abandon his opposition to the atomic theory.

Nor was Ostwald the only one in the 1890's to recognize the worth of Gibbs. The Dutch physical chemist, Hendrik Willem Bakhuis Roozeboom (1854-1907), publicized Gibbs's phase rule throughout Europe and did so most effectively.

Then, too, Gibbs's work was translated into French in 1899 by Henri Louis Le Chatelier (1850-1936). Le Chatelier, a physical chemist, is best known today for his enunciation of a rule, in 1888, that is still called Le Chatelier's principle. This rule may be stated: Every change of one of the factors of an equilibrium brings about a rearrangement of the system in such a direction as to minimize the original change.

In other words, if a system in equilibrium is placed under increased pressure, it rearranges itself so as to take up as little room as possible and thus decrease the pressure. If the temperature is raised, it undergoes a change that absorbs heat and lowers the temperature and so on. As it turned out, Gibbs's chemical thermodynamics explained Le Chatelier's principle neatly.

The late discovery of Gibbs by Europeans did not delay the development of physical chemistry as much as it might have, for many of Gibbs's findings were worked out independently, during the 1880's, by Van't Hoff, who had previously presented the world of chemistry with the tetrahedral carbon atom.

Van't Hoff was second only to Ostwald as an important worker in the field of physical chemistry. He worked on the problems of solutions in particular. By 1886 he was able to show that the molecules of dissolved substances, moving randomly through the body of the liquid in which they were dissolved, behaved, in some ways, according to rules analogous to those which described the behavior of gases.

Nor did the new study of physical chemistry connect chemical reactions with heat alone; it was rather with energy generally. Electricity could be produced by chemical reactions and could in turn bring about chemical reactions.

Walther Hermann Nernst, a German (1864-1941), applied the principles of thermodynamics to the chemical reactions proceeding in a battery. In 1889, he showed how the characteristics of the current produced could be used to calculate the free energy change in the chemical reaction producing the current.

Light was still another form of energy that could be produced in a chemical reaction and, as was discovered even before the nineteenth century, it could in turn induce chemical reactions. In particular, light could break down certain silver compounds, liberating black grains of metallic silver.

The study of such light-induced reactions is termed photochemistry ("light-chemistry").

In the 1830's, the action of light on silver had been developed into a technique for allowing sunlight to paint a picture. A layer of silver compound upon a glass plate (later, upon a flexible film) is briefly exposed, by way of a focusing lens, to a sunlit scene. Different areas of the silver compound are exposed to different amounts of light, according to how much was reflected from this point or that point in the scene. The brief exposure to the light increases the tendency of the silver compound to break down to metallic silver; the brighter the light, the more sharply increased the tendency.

The silver compound is then treated with chemicals that bring about such a breakdown to metallic silver. The region exposed to bright light completes the breakdown much more rapidly. If the "development" is stopped at the right point, the glass plate is covered by a pattern of dark (silver grains) and light (unchanged silver compound) that complements the pattern of the original scene.

Through further optical and chemical processes that need not be described here, a realistic portrayal of the scene is eventually obtained. The process is termed photography ("light-writing"). A number of men contributed to the new technique, including the French physicist Joseph Nicephore Niepce (1765-1833), the French artist Louis Jacques Mandes Daguerre (1789-1851), and the English inventor William Henry Fox Talbot (1800-1877).

Particularly interesting was the manner in which light behaved almost as a catalyst. A small quantity of light could induce a mixture of hydrogen and chlorine to react with explosive violence where, in the dark, no reaction at all would occur.

The explanation for this drastic difference in behavior was finally advanced by Nernst in 1918. A small quantity of light suffices to break the chlorine molecule apart into two single chlorine atoms. One chlorine atom (much more active in itself than as part of a molecule) snatches a hydrogen atom from the hydrogen molecule, to form a hydrogen chloride molecule. The remaining hydrogen atom, isolated, snatches a chlorine from a chlorine molecule; the remaining chlorine atoms snatches a hydrogen from a hydrogen molecule, and so on.

The original bit of light is thus responsible for a photochemical chain reaction, which leads to the formation of a great many hydrogen chloride molecules at an explosive rate.

Ionic Dissociation

Added to Ostwald and Van't Hoff was another master of early physical chemistry, the Swedish chemist Svante August Arrhenius (1859-1927). As a student, Arrhenius turned his attention to electrolytes; that is, to those solutions capable of carrying an electric current.

Faraday had worked out the laws of electrolysis, and from those laws it had seemed that electricity, like matter, might well exist in the form of tiny particles. Faraday had spoken of ions, which might be considered as particles carrying electricity through a solution. For the next half century neither he nor anyone else had ventured to work seriously on what the nature of those ions might be. This did not mean to no valuable work was done. In 1835, the German physicist Johann Wilhelm Hittorf (1824-1914) pointed out that some ions traveled more rapidly than others. This observation led to the concept of transport number, the rate at which particular ions carried the electric current. But even calculation of this rate still left the nature of ions an open question.

Arrhenius found his entry into the field through the work of the French chemist Francois Marie Raoult (1830-1901). Like Van't Hoff, Raoult studied solutions. His studies were climaxed in 1887 with his establishment of what is now call Raoult's law: The partial pressure of solvent vapor in equilibrium with a solution is directly proportional to the mole fraction of the solvent.

Without going into the definition of mole fraction, it is sufficient to say that this rule made it possible to estimate the relative number of particles (whether of atoms, molecules or the mysterious ions) of the substance which is dissolved (the solute) and of the liquid in which it is dissolved (the solvent).

In the course of this research, Raoult had measured the freezing points of solutions. Such freezing points were always "depressed"; that is, were lower than the freezing point of the pure solvent. Raoult was able to show that the freezing point was depressed in proportion to the number of particles of solute present in the solution.

But here a problem was created. It was reasonable to suppose that when a substance dissolved in, say, water it broke up into separate molecules. Sure enough, in the case of non-electrolytes such as sugar, the depression of the freezing point fit that assumption. However, when an electrolyte like common salt (NaCl) was dissolved, the depression of the freezing point was twice as great as it should have been. The number of particles present was twice the number of salt molecules. If barium chloride (BaCl2) was dissolved, the number of particles present was three times as great as the number of molecules.

A molecule of sodium chloride is made up of two atoms, and a molecule of barium chloride is made up of three atoms. It seemed to Arrhenius that when certain molecules were dissolved in a solvent such as water, those molecules broke down into the individual atoms. Furthermore, since those molecules, once broken down, conducted an electric current (whereas molecules such as sugar, which did not break apart, did not carry an electric current), Arrhenius further suggested that the molecules did not break down (or "dissociate") into ordinary atoms, but into atoms carrying an electric charge.

Faraday's ions, Arrhenius proposed, were simply atoms (or groups of atoms) carrying either a positive or a negative electric charge. The ions were either the "atoms of electricity" or they carried those "atoms of electricity". (The latter alternative eventually proved correct.) Arrhenius used his theory of ionic dissociation to account for many facts of electrochemistry.

Arrhenius's ideas, advanced as his Ph.D. thesis in 1884, met with considerable resistance; his thesis was almost rejected. However, Ostwald, impressed, offered Arrhenius a position and encouraged him to continue work in physical chemistry.

In 1889, Arrhenius made another fruitful suggestion. He pointed out that molecules, on colliding, need not react unless they collide with a certain minimum energy, and energy of activation. When this energy of activation is low, reactions proceed quickly and smoothly. A high energy of activation, however, might keep a reaction from proceeding at more than an infinitesimal rate.

If, in the latter case, the temperature were raised so that a number of molecules received the necessary energy of activation, the reaction would then proceed suddenly and quickly, some times with explosive violence. The explosion of a hydrogen-oxygen mixture when the ignition temperature is reached is an example.

Ostwald used this concept profitably in working out his theory of catalysis. He pointed out that the formation of a catalyst-combined intermediate required a smaller energy of activation than the direct formation of the final products required.

More on Gases

The properties of gases came under new and refined scrutiny during the burgeoning of physical chemistry in the late nineteenth century. Three centuries earlier, Boyle had advanced Boyle's law, stating that the pressure and volume of a given quantity of gas varied inversely (provided, as was later shown, that temperature is held constant.)

This law turned out to be not exactly true. The German-French chemist Henri Victor Regnault (1810-1878) made many careful measurements of gas volumes and pressures in the mid-nineteenth century and showed that, especially as pressure was raised or temperature was lowered, gases did not quite follow Boyle's law.

At about the same time, the Scottish physicist James Clerk Maxwell (1831-1879) and the Austrian physicist Ludwig Boltzmann (1844-1906) had analyzed the behavior of gases on the assumption that they were an assemblage of a vast number of randomly moving particles (the kinetic theory of gases). They were able to derive Boyle's law on this basis, provided they made two further assumptions;

  1. that there were no attractive forces between gas molecules.
  2. that the gas molecules were of zero size.

Gases that fulfill these assumptions are perfect gases.

Neither assumption is quite correct. There are small attractions between gas molecules, and though these molecules are exceedingly small, their size is not zero. No actual gas is quite "perfect", therefore, although hydrogen and the later-discovered helium come close.

Taking these facts into account, the Dutch physicist Johannes Diderik Van der Waals (1837-1923), in 1873, worked out an equation that related pressure, volume, and temperature of gases. This equation included two constants, a and b (different for each gas), the existence of which allowed for the size of the molecules and the attractions among them.

The better understanding of the properties of gases helped to solve the problem of liquefying them.

As early as 1799, the gas ammonia was liquefied by being cooled while it was under pressure. (Raising the pressure raises the temperature at which a gas will liquefy and makes the liquefaction process that much easier.) Faraday was particularly active in this field of investigation, and by 1845 had been able to liquefy a number of gases, chlorine and sulfur dioxide among them. Once a liquefied gas is released from pressure, it begins to evaporate rapidly. The process of evaporation absorbs heat and the temperature of the remaining liquid drops drastically. Liquid carbon dioxide will under such conditions freeze to solid carbon dioxide (dry ice). By mixing solid carbon dioxide with ether, Faraday could obtain temperatures of -78oC.

Foiling his best efforts, however, were such gases as oxygen, nitrogen, hydrogen, carbon monoxide, and methane. No matter how high the pressures he worked at, Faraday could not liquefy them. These substances came to be termed "permanent gases".

In the 1860's, however, an Irish chemist, Thomas Andrews (1813-1885), was working with carbon dioxide which he had liquefied by pressure along. Slowly raising the temperature, he noted the manner in which the pressure had to be increased to keep the carbon dioxide in the liquid state. He found that at a temperature of 31oC., no amount of pressure sufficed. At that temperature, in fact, the gas and liquid phase seemed to melt together, so to speak, and become indistinguishable. Therefore, Andrews suggested (in 1869) that for each gas there was a critical temperature, above which no amount of pressure along could liquefy it. It followed that permanent gases were simply those with critical temperatures well below those reached in the laboratory.

Meanwhile, Joule and Thomson in their studies on heat had discovered that gases could be cooled by allowing them to expand. If gases were allowed to expand, then compressed under conditions which did not allow them to regain the lost heat, and expanded once more, and so on over and over, then very low temperatures could be achieved. Once a temperature below the critical temperature for that gas was reached, application of pressure would liquefy it.

Using this technique, the French physicist Louis Paul Cailletet (1832-1913) and the Swiss chemist Raoul Pictet (1846-1929) were able to liquefy such gases as oxygen, nitrogen, and carbon monoxide by 1877. Hydrogen, however, still balked their efforts.

As a result of Van der Waals's work, it became clear that in the case of hydrogen the Joule-Thomson effect would work only below a certain temperature. Its temperature had to be lowered before the cycle of expansion and contraction could be started.

In the 1890's, the Scottish chemist James Dewar (1842-1923) began work on the problem. He prepared liquid oxygen in quantity and stored it in a Dewar flask. This is a double-walled flask with a vacuum between the walls. The vacuum will not transmit heat by conduction or convection, since both phenomena require the presence of matter. Heat is transmitted across a vacuum only by the comparatively slow process of radiation. By silvering the walls so that radiated heat would be reflected rather than absorbed, Dewar slowed down that process even further. (Our "thermos" of today is simply a Dewar flask with a stopper.)

Hydrogen liquefied at 20oK., a temperature but twenty degrees above absolute zero. (The concept of absolute zero, the lowest temperature possible, was first advanced by Thomson (Lord Kelvin) in 1848. In recognition of this proposal, the absolute temperature scale (based on Kelvin's concept) is symbolized as oK. In 1905, Nernt showed that entropy was zero at absolute zero (the third law of thermodynamics). From this, one could deduce that though a temperature of absolute zero could be approached as closely as desired, it could never actually be reached.) This is not a record low liquefaction point, by any means. In the 1890's, the inert gases had abeen discovered, and one of these, helium, liquified at the still lower temperature.

The Ductch physicist Heike Kamerlingh Onnes (1853-1926) overcame the last obstacle when, in 1908, he cooled helium first in a bath of liquid hydrogen, then applied the Joule-Thomson effect. He produced liquid helium at a temperature of 4oK.