Physical Chemistry
The application of physical methods as aids to chemical science began in early times, and some of these, such as the determinations of gas and vapor densities, specific heats, and crystalline forms have been mentioned already in this article. Within recent times physical chemistry has greatly developed and a few of its important achievements will now be described.
Molecular Weight Determinations.—Gas and vapor densities in connection with Avogadro’s principle, formed the only basis for molecular weight determinations until comparatively recent times. The early methods of Gay-Lussac and Dumas for vapor density were supplemented in 1868 by the method of Hofmann, whereby vapors were measured under diminished pressure over mercury. In 1878 Victor Meyer introduced a simpler method depending upon the displacement of air or other gas by the vapor in a heated tube. As refractory tubes, such as those of porcelain or even iridium, could be used in this method, molecular weights at extremely high temperatures were determined with interesting results. For instance, it was found that iodine vapor, which shows the molecule I2 at lower temperatures, gradually becomes monatomic with rise in temperature, that sulphur vapor dissociates from S8 to S2 under similar conditions, and that most of the metals, including silver, have monatomic vapors.
In 1883 and later it was pointed out by Raoult that the molecular weights of substances could be found from the freezing points of their solutions, but this method was complicated from the fact that salts, strong acids and strong bases behaved quite differently from other substances in this respect, and allowances had to be made for the types of substances used. The complication was afterwards explained by the ionization theory of Arrhenius. Better apparatus for this method was soon devised by Beckmann, who introduced also a method depending upon the boiling points of solutions, and these two methods are still the standard ones for determining molecular weights in solution. They are very extensively employed by organic chemists.
It has been found that the majority of substances when dissolved have the same molecular weight as in the gaseous condition, provided that they can be volatilized at comparable temperatures. For instance, sulphur in solution has the formula S8, iodine is I2 and the metals are monatomic.
Van’t Hoff’s Law and Arrhenius’s Theory of Ions.—Modern views on solutions date largely from 1886, when van’t Hoff called attention to the relations existing between the osmotic pressure exerted by dissolved substances and gas pressure.
Pfeffer, a botanist, was the first to measure osmotic pressure (1877). Basing his conclusions chiefly upon Pfeffer’s determinations, van’t Hoff formulated a new and highly important law, which may be stated as follows: The osmotic pressure exerted by a substance in solution is equal to the gas pressure that the substance would exert if it were a gas at the same temperature and the same volume. Further investigations have fully established the fact that molecules in dilute solution obey the simple laws of gases.
It was pointed out by van’t Hoff that salts, strong acids and strong bases showed marked exceptions to his law in exerting much greater osmotic pressures than those calculated for them.
The next year in 1887, Arrhenius explained this abnormal behavior of salts, strong acids and strong bases by assuming that they dissociate spontaneously into ions when they dissolve, and that these more numerous particles act like molecules in producing osmotic pressure. He showed that these exceptional substances all conduct electricity in solution, while those conforming with van’t Hoff’s law do not, and according to his theory the ions become positively or negatively charged when they are formed, and these charged ions conduct the current. For example a molecule of sodium chloride was supposed to give the two ions Na+ and Cl-, thus exerting twice as much osmotic pressure as a single molecule.
Determinations of osmotic pressure or related values, such as depression of the freezing point and of electric conductivity, indicated that ionization could not be regarded as complete in any case except in exceedingly dilute solutions, and that the extent of ionization varied with different substances. The fact that osmotic pressures and electric conductivities gave closely agreeing results in regard to the extent of ionization in various cases, is the strongest evidence in support of the theory.
It was difficult at first for many chemists to believe that atoms, such as those of sodium and chlorine, and groups such as NH4 and SO4 could exist independently in solution, even though electrically charged. However, the theory rapidly gained ground and is now accepted by nearly every chemist as a satisfactory explanation of many facts.
During recent years, many investigations relating to osmotic pressure and ionization have been carried out in the United States, but only the work of Morse, A. A. Noyes, and the late H. C. Jones can be merely alluded to here. It should be mentioned that the eminent author of the ionic hypothesis gave the Silliman Memorial course of lectures at Yale in 1911 on Theories of Solution.
Colloidal Solutions.—Graham, an English chemist, in 1861 was the first to make a distinction between substances forming true solutions, which he called crystalloids, and those of a gummy nature resembling glue, which in solution do not diffuse readily through parchment membranes, as crystalloids do, and which he called colloids. The separation of colloids by means of parchment was called dialysis, and this process has come into extensive use in preparing pure colloidal solutions. Slow diffusion is now regarded as characteristic of colloids rather than their gummy condition.
Colloidal solutions occupy an intermediate position between true solutions and suspensions, resembling one or the other according to the kind of colloid and the fineness of division. By preparing filters with pores of varying degrees of fineness, Bechold has been able to separate colloids from each other in accordance with the size of their particles. It has also been possible to prepare different solutions of a colloid varying gradually from one in which the particles were undoubtedly in suspension to one which had many of the properties of a true solution.
Beginning in 1889, Carey Lea described in the Journal (37, 476, 1889 et seq.) a variety of methods for preparing colloidal solutions of the metals, consisting in general of treating solutions of metallic salts with mild reducing agents. His work on colloidal silver was particularly extensive and interesting. Solutions of this kind have recently yielded some extremely interesting results by means of the ultra-microscope, an apparatus devised by Zsigmondy and Siedentopf. A very intense beam of light is passed through the solution and observed at right angles with a powerful microscope. Under these conditions, particles much too small to be seen by other means, reveal their presence by reflected light. It has been possible in a very dilute solution of known strength to count the particles and thus to calculate their size. The smallest colloidal particles measured in this way were of gold and were shown to have approximately ten times the diameter, or 1000 times the volume, attributed to ordinary molecules. It is of interest that the particles appear in rapid motion corresponding to the well-known Brownian movement.
The chemistry of colloids has now assumed such importance that it may be considered as a separate branch of the science. It has its own technical journal and deals largely with the chemistry of organic products. All living matter is built up of colloids, and hæmoglobin, starch, proteins, rubber and milk are examples of colloidal substances or solutions. Among inorganic substances, many sulphides, silicic acid, and the amorphous hydroxides, like ferric hydroxide, frequently act as colloids.
Law of Mass Action.—Berthollet about the beginning of the last century was the first chemist to study the effect of mass, or more correctly, the concentration of substances on chemical action. His views summarized by himself are as follows: “The chemical activity of a substance depends upon the force of its affinity and upon the mass which is present in a given volume.” The development of this idea, which is fundamentally correct, was greatly hindered by the fact that Berthollet drew the incorrect conclusion that the composition of chemical compounds depended upon the masses of the substances combining to produce them, a conclusion in direct contradiction to the law of definite proportions, and since this view was soon disproved by Proust and others, Berthollet’s law in its other applications received no immediate attention. Mitchell, however, pointed out in the Journal (16, 234, 1829) the importance of Berthollet’s work, and Heinrich Rose in 1842 again called attention to the effect of mass, mentioning as one illustration the effect of water and carbonic acid in decomposing the very stable natural silicates. Somewhat later several other chemists made important contributions to the question of the influence of concentration upon chemical action, but it was the Norwegians, Guldberg and Waage, who first formulated the law of mass action in 1867.
This law has been of enormous importance in chemical theory, since it explains a great many facts upon a mathematical basis. It applies particularly to equilibrium in reversible reactions, where it states that the product of the concentrations on the one side of a simple reversible equation bears a constant relation to the products of the concentrations on the other side, provided that the temperature remains constant. In cases of this kind where two gases or vapors react with two solids, the latter if always in excess may be regarded as constant in concentration, and the law takes on a simpler aspect in applying only to the concentrations of the gaseous substances. For example, in the reversible reaction
3Fe + 4H2O ⇄ Fe3O4 + 4H2, which takes place at rather high temperatures, a definite mixture of steam and hydrogen at a definite temperature will cause the reaction to proceed with equal rapidity in both directions, thus maintaining a state of equilibrium, provided that both iron and the oxide are present in excess. If, however, the relative concentrations of the hydrogen and steam are changed, or even if the temperature is changed, the reaction will proceed faster in one direction than in the other until equilibrium is again attained.
The principle of mass action also explains why it is sometimes possible for a reversible reaction to become complete in either direction. For instance, in connection with the reaction that has just been considered, if steam is passed over heated iron and if hydrogen is passed over the heated oxide, the gaseous product in each case is gradually carried away, and the reaction continually proceeds faster in one direction than in the other until it is complete, according to the equations
3Fe + 4H2O → 3Fe3O4 + 4H2, and Fe3O4 + 4H2 → 3Fe + 4H2O.
Many other well-known and important facts, both chemical and physical, depend upon this law. It explains the circumstance that a vapor-pressure is not dependent upon the amount of the liquid that is present; it also explains the constant dissociation pressure of calcium carbonate at a given temperature, irrespective of the amounts of carbonate and oxide present; in connection with the ionic theory, it furnishes the reason for the variable solubility of salts due to the presence of electrolytes containing ions in common; and it elucidates Henry’s law which states that the solubilities of gases are proportional to their pressures.
Ostwald, more than any other chemist, has been instrumental in making general applications of this law, and he made particularly extensive use of it in connection with analytical chemistry in a book upon this subject which he published.
The Phase Rule.—In 1876 Willard Gibbs of Yale published a paper in the Proceedings of the Connecticut Academy of Science on the “Equilibrium of Heterogeneous Substances,” and two years later he published an abstract of the article in the Journal (16, 441, 1878). He had discovered a new law of nature of momentous importance and wide application which is called the “Phase-Rule” and is expressed by a very simple formula.
The application of this great discovery to chemical theory was delayed for ten years, partly, perhaps, because it was not sufficiently brought to the attention of chemists, but largely it appears because it was not at first understood, since its presentation was entirely mathematical.
It was Rooseboom, a Dutch chemist, who first applied the phase-rule. It soon attracted profound attention, and the name of Willard Gibbs attained world-wide fame among chemists. When Nernst, who is perhaps the most eminent physical chemist of the present time, was delivering the Silliman Memorial Lectures at Yale a few years ago, he took occasion to place a wreath on the grave of Willard Gibbs in recognition of his achievements.
To understand the rule, it is necessary to define the three terms, introduced by Gibbs, phase, degrees of freedom and component.
By the first term, is meant the parts of any system of substances which are mechanically separable. For instance, water in contact with its vapor has two phases, while a solution of salt and water is composed of but one. The degrees of freedom are the number of physical conditions, including pressure, temperature and concentration, which can be varied independently in a system without destroying a phase. The exact definition of a component is not so simple, but in general, the components of a system are the integral parts of which it is composed. Any system made up of the compound H2O, for instance, whether as ice, water or vapor, contains but one component, while a solution of salt and water contains two. Letting P, F, and C stand for the three terms, the phase-rule is simply
F = C + 2 − P
that is, the number of degrees of freedom in a system in equilibrium equals the number of components, plus two, minus the number of phases. The rule can be easily understood by means of a simple illustration. In a system composed of ice, water and water vapor, there are three phases and one component and therefore
F = 1 + 2 − 3 = 0
Such a system has no degrees of freedom. This means that no physical condition, pressure or temperature can be varied without destroying a phase, so that such a system can only exist in equilibrium at one fixed temperature, with a fixed value for its vapor-pressure.
For instance, if the system is heated above the fixed temperature, ice disappears and if the pressure is raised, vapor is condensed. If this same system of water alone contains but two phases, for instance, liquid and vapor, F = 1 + 2 − 2 = 1, or there is one degree of freedom. In such a system, one physical condition such as temperature can be varied independently, but only one, without destroying a phase. For instance, the temperature may be raised or lowered, but for every value of temperature there is a corresponding value for the vapor-pressure. One is a function of the other. If both values are varied independently, one phase will disappear, either vapor condensing entirely to water or the reverse. Finally if the system consists of one phase only, as water vapor, F = 2, or the system is divariant, which means that at any given temperature it is possible for vapor to exist at varying pressures.
The illustration which has been given relates to physical equilibrium, but the rule is applicable to cases involving chemical changes as well. In comparing the phase-rule with the law of mass action, it will be noticed that both have to do with equilibrium. The great advantage of the former is that it is entirely independent of the molecular condition of the substances in the different phases. For instance, it makes no difference so far as the application of the rule is concerned, whether a substance in solution is dissociated, undissociated or combined with the solvent. In any case, the solution constitutes one phase. On the other hand, the rule is purely qualitative, giving information only as to whether a given change in conditions is possible. The law of mass action is a quantitative expression so that when the value of the constant is once known, the change can be calculated which takes place in the entire system if the concentration of one substance is varied. The law, however, requires a knowledge of the molecular condition of the reacting substances, which may be uncertain or unknown, and chiefly on this account it has, like the phase-rule, often only a qualitative significance.
The phase rule has served as a most valuable means of classifying systems in equilibrium and as a guide in determining the possible conditions under which such systems can exist. As illustrations of its practical application, van’t Hoff used it as an underlying principle in his investigations on the conditions under which salt deposits have been formed in nature, and Rooseboom was able by its means to explain the very complicated relations existing in the alloys of iron and carbon which form the various grades of wrought iron, steel and cast iron.
Thermochemistry.—This branch of chemistry has to do with heat evolved or absorbed in chemical reactions. It is important chiefly because in many cases it furnishes the only measure we have of the energy changes involved in reactions. To a great extent, it dates from the discovery by Hess in 1840 of a fundamental law which states that the heat evolved in a reaction is the same whether it takes place in one or in several stages. This law has made it possible to calculate the heat values of a large number of reactions which cannot be determined by direct experiment.
Thermochemistry has been developed by a comparatively few men who have contributed a surprisingly large number of results. Favre and Silbermann, beginning shortly after 1850, improved the apparatus for calorimetric determinations, which is called the calorimeter, and published many results. At about the same time Julius Thomsen, and in 1873 Berthelot, began their remarkable series of publications which continued until recently. Thomsen’s investigations were published in 1882 in 4 volumes. It is probably safe to say that the greater part of the data of thermochemistry was obtained by these two investigators. The bomb calorimeter, an apparatus for determining heat values by direct combustion, was developed by Berthelot. The recent work of Mixter at Yale, published in the Journal, and of Richards at Harvard should be mentioned particularly. Mixter’s work in this field began in 1901 (12, 347). Using an improved bomb calorimeter, he has developed a method of determining the heats of formation of oxides by combustion with sodium peroxide. By this same method as well as by direct combustion in oxygen, he has obtained results which appear to equal or excel in accuracy any which have ever been obtained in his field of work. Richards’s work has consisted largely of improvements in apparatus. He developed the so-called adiabatic calorimeter which practically eliminates one of the chief errors in thermal work caused by the heating or cooling effect of the surroundings. This modification is being generally adopted where extremely accurate work is required.