Quantitative Evidence.

—Some of the most exact quantitative evidence bearing on these relations, such as the results of investigations, by Griffith and by Taylor, on the freezing-point depressions of solutions of electrolytes, may be briefly considered. The depression of the freezing-point of a given solvent by a solute is proportional to the concentration of the solute or proportional to its osmotic pressure. Further, according to the Van 't Hoff Hypothesis (p. 15), the osmotic pressure at a constant temperature is dependent only on the number of molecules present in unit volume, and not on the nature or composition of the molecules: the freezing-point of the solvent is depressed, likewise, proportionally to the total concentration of the solute, irrespective of the fact whether the solution contains only one, or more than one molecular species. The ratio, observed depression / concentration,116 or Δ / C, should be constant,117 therefore, in a given solvent, for dilute solutions of all kinds of solutes, simple or mixed. Griffith118 found, for a solution of cane sugar, a non-electrolyte, in water, the ratio of the freezing-point depression to the concentration to be 1.858°. For instance, the freezing-point of a 0.01 molar solution of cane sugar (3.42 grams of cane sugar per liter; C12H22O11 = 342) is found to be −0.01858°, and 0.01858 / 0.01 = 1.858. This ratio should be the same, as stated above, according to van 't Hoff's theory of solutions, for dilute aqueous solutions of all solutes. But the ratio Δ / C for an aqueous solution of potassium chloride, an electrolyte, was found to increase slowly and continuously until in 0.0003 molar solution the ratio 3.72 was found, which is exactly twice the value obtained with cane sugar. The result indicates, therefore, a gradual dissociation of the potassium chloride with increasing dilution, and a dissociation, ultimately, of each molecule of the salt into two new molecules, in all respects exactly as demanded by the theory of Arrhenius.

Loomis119 found in a similar way a ratio of 3.61 for HCl, 3.71 [p069] for KOH, 3.60 for KCl, 3.67 for NaCl, 3.73 for HNO3, etc., when 0.01 molar aqueous solutions were used. For similar solutions of calcium chloride CaCl2, magnesium chloride MgCl2, and sodium sulphate Na2SO4, the value 5.07 was found as the ratio between the depressions of the freezing-point and the concentration of the salts in extremely dilute solutions—a result showing, plainly, a dissociation of each salt into three smaller molecules. The limit 5.67 for such a dissociation is not quite reached in these cases, because salts of the types Me″X2 and Me2′Y″ ionize less readily than do the electrolytes Me′X′, a fact also shown by their conductivities.

We thus find that the most exact work on molecular weight determinations in dilute aqueous solutions agrees excellently, as does the conductivity of such solutions, with the demands of the theory of ionization, a fact which is particularly impressive because osmotic pressure and electrical conductivity are in no wise fundamentally related phenomena, and yet each, as a measure of ionization or electrolytic dissociation, leads independently to the same conclusion.120

The Chemical Composition of the Ions of Electrolytes.

—Accepting the theory of Arrhenius, we may now inquire more closely than heretofore, first, what compounds are subject to electrolytic dissociation, and then, what the chemical composition of their ions is and how it is determined.

The compounds which are dissociated into ions, by solvents which cause ionization (p. 62), comprise the salts, the acids and the bases; chemists are, in fact, more inclined now to invert the statement and say that those substances which have long been known as salts, acids and bases, owe the essential characteristics, which led to their classification, to the fact that they are ionizable (see pp. 72–82). The composition of the ions, formed from the simpler of these compounds,121 may be expressed by saying that the metal component or metal-like component (hydrogen, ammonium) forms the positive ion (cation, metal ion) and all the rest of the [p070] molecule forms the negative ion (anion, acid ion122). Thus, sodium chloride, nitrate, sulphate, phosphate yield the sodium-ion, Na+, and the chloride (Cl−), nitrate (NO3−), sulphate (SO42−), or phosphate (PO43−) ions; cupric nitrate Cu(NO3)2 dissociates into the cupric ion (Cu2+) and the nitrate ion, calcium sulphate CaSO4 into the calcium-ion (Ca2+) and the sulphate-ion, aluminium sulphate Al2(SO4)3 into the aluminium-ion (Al3+) and the sulphate-ion.

But the question arises, as to how we know that the salts mentioned produce ions of the given composition; why, for instance, should sodium nitrate be considered to dissociate into sodium, Na+, and nitrate ions, NO3−, the nitrogen atom carrying all of the oxygen atoms with it in the negative ion?

Fig. 13.

The composition of the ions of a salt can be determined experimentally123 by devices of which the U-tube experiment (p. 45) may be considered to be a simple type. For instance, if we wish to determine the composition of the ions of sodium nitrate, we could cover a solution of sodium nitrate with a solution, say, of hydrochloric acid, pass a current through the liquids, and determine the composition of the components that have moved to the negative and positive poles, respectively. In practice, the device could be elaborated for the sake of convenience. Stopcocks, for instance, might be placed in the U-tube, at the points of separation of the nitrate solution and the hydrochloric acid (see Fig. 13), the stopcocks being opened only during the passage of the current. Or porous plates or cells might be used, in place of stopcocks, at these [p071] points. Now, if we assume sodium nitrate to be dissociated, not into Na+ and NO3−, but let us say into positive ions NaO+ and negative ions NO2−, the changes which would result from the passage of a current would be as follows: starting with the action at the positive pole, we should find chloride ions discharged and chlorine evolved at the pole (the evolution of chlorine could be avoided, if considered desirable, by the use of a silver anode, which would absorb the liberated chloride-ion to form insoluble silver chloride on the electrode). At the same time, hydrogen ions would move out of the space P, being repelled by the positive pole, and attracted by the negative. At the boundary between the sodium nitrate solution and the hydrochloric acid, the negative ions of sodium nitrate, which we are supposing to have the composition NO2−, would move up from B toward the positive pole (cf. exp., p. 45), being attracted by its charge; at any moment we should have in any part of P as many negative ions (Cl− and NO2−), as there are hydrogen ions, the solution showing no excess of free electricity at any point. Now, if NO2− were the ion that moved up into the space P, then we should presently find nitrous acid around the positive pole in space P, H+ and NO2− combining to form nitrous acid, HNO2. But, as a matter of experiment, although the tests for nitrous acid belong to the most sensitive ones in chemistry, no trace of this acid is found there; what we do find is nitric acid, HNO3, resulting obviously from the presence in space P of both hydrogen ions and nitrate ions, NO3−, which have moved up from space B. It is clear, that the presence of nitric acid in the region around the positive pole means that the nitrogen atoms must have carried with them all three of the oxygen atoms of the nitrate—in a word, that the composition of the negative ion of sodium nitrate is NO3− and not, say, NO2−. Similarly, considering what happens in space N, round the negative pole, we have here an evolution of hydrogen, a migration of some chloride ions out of N into the space C, and, at the same time, a migration of the positive ions of space C into the division N. On examining the solution in N, we now find sodium chloride, with unchanged hydrochloric acid, exactly what we should expect from the migration of the ion Na+ toward the negative electrode.124 If the positive ion were, say, [p072] NaO+, we should expect to obtain either some of the hypochlorite NaOCl (NaO+ + Cl−), or, at least, an evolution of oxygen in this place, since sodium chloride is formed. As a matter of experiment, no oxygen is evolved here, and no trace of hypochlorite is found in N round the negative pole, although the tests for hypochlorites are extremely sensitive.

Comparatively simple methods, in principle of the nature outlined, enable us, then, to determine experimentally the composition of the ions into which ionizable compounds, salts, acids and bases, dissociate. Whenever any doubt may exist about the composition of the ions of a given electrolyte, this device may be employed to settle the matter, and there will presently be occasion to employ the U-tube for such a purpose.

Ionization and Chemical Activity.

—The fact that the theory of ionization gives us adequate explanations of the conductivity shown by dissolved electrolytes and of their abnormally high osmotic pressures, would have been in itself of interest to chemists; but, if its applications were limited to these phenomena, we should not be considering it in connection with qualitative chemical analysis, nor would the theory, presumably, have greatly affected the development of chemistry, as it has done. It is the fact that the electrolytic dissociation of an electrolyte into its ions involves chemical changes of the most profound nature, and most intimately affects chemical reactivity, that has made it play, in the last two decades, such a leading rôle in the development of chemistry, and that makes it necessary to include its consequences in the consideration of analytical problems, if one would understand, as far as present knowledge permits, the reactions involved in chemical analysis.

Hydrogen chloride, as a perfectly dry gas, is a non-conductor of electricity and, at the same time, it is found to be chemically inactive—it does not combine, for instance, with dry ammonia125 or act upon dry calcium carbonate126 or on dry litmus. Hydrogen chloride, subjected to great pressure at a low temperature, is liquefied. The liquid is also a very poor conductor of [p073] electricity127 and does not show the chemical activity of ordinary, aqueous hydrochloric acid; it does not combine with calcium oxide or attack marble, zinc, iron or even magnesium.127

A solution of hydrogen chloride in a poorly ionizing medium, like benzene or toluene, is an extremely poor conductor. There is an extremely small conductivity indicating only a trace of ionization.128

Exp. A solution of hydrogen chloride, prepared by passing the dried gas into benzene or toluene (thiophene-free benzene will not become discolored), and kept anhydrous by means of fused calcium chloride, is tested for its conductivity, by dipping into it electrodes connected with a lighting circuit and a galvanometer.

Such a solution behaves chemically, also, quite differently from the aqueous solutions of hydrogen chloride with which we are familiar: dry steel nails,129 dropped into it, will remain almost unchanged—there is no marked evolution of hydrogen (exp.). Perfectly dried marble, added to it, will not give rise to the evolution of carbon dioxide129 (exp.). We find thus, in all the cases discussed—the nonconducting dry gas, the anhydrous liquefied hydrogen chloride and the anhydrous benzene solution—an absence of ionization,130 as indicated by the lack of conductivity, and, along with this, a lack of the familiar action of hydrochloric acid as an acid. If we dissolve the gas in water, we obtain a well-conducting solution (exp.), in which, according to molecular weight determinations, the [p074] hydrogen chloride is more or less largely ionized, and this same solution has all the well-known chemical properties of hydrochloric acid—it evolves hydrogen liberally when given an opportunity to act upon zinc or iron (exp.), it evolves carbon dioxide copiously when marble is brought into contact with it (exp.). In such an aqueous solution we have both the ions of the acid and the more or less non-ionized hydrogen chloride, the action (HCl ⇄ H+ + Cl−) being reversible.

Now, since in those cases in which we have admittedly only non-ionized hydrogen chloride, there is no vigorous chemical action, we are bound to conclude that, in the aqueous solution where we have both the non-ionized and the ionized substance, it must be the new components, the ions of the acid, which give this solution its new qualities, the well-known properties of a pronounced acid.131 This conclusion, that the acid properties of hydrogen chloride in aqueous solution are due to the ionized hydrogen chloride, rather than to the hydrogen chloride itself, is one of fundamental importance.

Dry Salts and their Aqueous Solutions.

—If the study of the relation of ionization to chemical activity be extended, it is found that a dry salt, such as, for example, silver nitrate, in crystals or finely pulverized, is not perceptibly ionized, for it is a non-conductor (exp.). The same result is obtained with potassium chromate. If the dry powders are intimately mixed, there is no chemical action between them, no perceptible change occurs. The aqueous solutions of the salts are excellent conductors (exp.), as are the aqueous solutions of almost all salts; the dissolved salts are therefore largely ionized. As soon as the ionizing medium, water, is added to the dry, yellow mixture of silver nitrate and potassium chromate, instantly a chemical change results—red silver chromate, Ag2CrO4, is precipitated (exp.). Now, in the aqueous solution of these salts we have both non-ionized molecules and their ions:
AgNO3 ⇄ Ag+ + NO3−,
K2CrO4 ⇄ 2 K+ + CrO42−.

Since there is no interaction when the dry salts, containing only the non-ionized substances, are mixed, and since there is interaction [p075] when the solutions are mixed, in which both the non-ionized and the ionized salts are present, one must conclude again that the formation of silver chromate is the result of the action of the silver ions on the chromate ions in the solution. In point of fact, there could hardly fail to be an action, since the positive silver ions and the negative chromate ions, moving in all directions through the solution, must collide and be discharged, or combine, to form molecular silver chromate. This salt happens to be very difficultly soluble, and to be colored red, as well, so that silver chromate is precipitated and is immediately recognizable.

The two dry powders in the experiment were allowed to be in contact for only a few moments. It is important to note, therefore, that dry sodium acid carbonate and dry potassium acid tartrate are also nonconductors (exp.), and that the intimate mixture of these two powders is kept for years in the well-known form of baking powders without appreciable decomposition—yet best in tin vessels, to exclude moisture. The aqueous solutions, however, are good conductors (exp.), and, when dissolved, these salts are more or less ionized. The addition of water to the mixed salts (exp.) leads at once to the well-known action, carbon dioxide being liberated and sodium-potassium tartrate or Rochelle salt being formed.

Behavior of Fused Salts.

—It may be objected that there are common cases, where dry salts are known to act upon each other; barium sulphate is fused with sodium carbonate to convert the former into the carbonate, BaSO4 + Na2CO3 ⇄ Na2SO4 + BaCO3. Before one decides that this must be an instance of the action of non-ionized salts on each other, the conductivity of dry salts under the conditions of the experiment, namely at an elevated temperature, must be examined. There is no difficulty in recognizing that while dry sodium carbonate or potassium nitrate at ordinary temperatures does not conduct a current, and is not perceptibly ionized, each salt, when fused, becomes an excellent conductor (exp.,132 with potassium nitrate). It is, in fact, well known that, in many electrolytic operations, fused salts133 are used in place [p076] of solutions. It must be added that the heat, not the change of state, causes the ionization, careful work having shown that conductivity begins to be appreciable below the point of fusion.

Exp. Two platinum wires, fused, one inch apart, into a glass rod, are connected with a sensitive galvanometer and the lighting circuit. When the glass is warmed, a current is found to pass.

The action between barium sulphate and sodium carbonate at a high temperature does not mean, then, that the non-ionized salts interact; on the contrary, we find that, under such conditions, coincident with the evidence of reactivity, we have also decided conductivity—again indicating decided ionization.

Dry Salts at Ordinary Temperatures.

—Inasmuch as ordinary temperatures are still far removed from the absolute zero, one must suppose that dry salts must be ionized, minimally at least, even at room temperature, and should therefore react with each other. Presumably they do, only so slowly, as a result of the minimal degrees of ionization, and of the few chances of collision between ions of opposite charges, owing to the restricted range of the molecular motions, that the total change is imperceptible. Critical work on the question is most desirable. In this connection it maybe said that Spring134 found that dry salts do interact at ordinary temperature, when subjected to great pressures, provided the volume of the products is smaller than the volume of the initial substances. Whether this action is due to the ionization of the salts, minimal as it is, or whether we have here a case of interaction of non-ionized molecules, has not, it seems, been determined; it would require difficult quantitative work to settle the question.

Influence of Light and Heat.

—It is apparent that heat, a form of energy, contributes to the dissociation of ionogens, and it is natural that we should consider other forms of energy, e.g. light, to have the same power. One may speculate about the possibility of light inducing chemical action (e.g. in starting the combination of hydrogen and chlorine, or in photography) by its ionizing power, and about the possibility that rapid combination of oxygen and hydrogen follows the application of a flame to the mixture, as the result of increased ionization of the components at the elevated temperature of the flame and of the burning gases. The experimental evidence shows that, in some actions of this nature, ionization is an important factor, while in other instances it appears to be negligible.135 [p077]

Conclusions.

—It appears that we must accept the conclusion, that the reactions of salts, acids and bases (ionogens) in aqueous solution (the so-called "salt reactions") are the reactions of the ions and not of the non-ionized molecules. This conclusion is of the greatest and most practical importance in our science. It is the natural inference from the results of the qualitative experiments described in the three preceding sections. Its final adoption, however, is based on the existence of a great mass of quantitative evidence, and to the consideration of some of this we now turn.

Quantitative Relations.

—If the active components in aqueous solutions of acids, bases and salts are the ions, rather than the undissociated compounds, then quantitative data supporting such a conclusion should be found. Such quantitative confirmation, from the point of view of chemical activity, is not lacking. Only a small part of the data can be considered here.

Conductivity measurements, and the lowering of freezing-points and the elevation of boiling-points, show that there are very decided differences in the degrees of ionization of different acids and bases in solutions of equivalent concentration. For instance, potassium hydroxide is somewhat more ionized than is barium hydroxide, and decidedly more so than ammonium hydroxide, a fact that can readily be demonstrated by the conductivities of the solutions136:

Exp. Equivalent solutions (1 / 10 normal) of the three bases are introduced into three vertical tubes, containing electrodes connected, in parallel, with a lighting circuit and with small electric lamps. When the two electrodes in each of the three tubes are at equal distances from each other, the lamp connected with the potassium hydroxide solution glows most brightly, that connected with the barium hydroxide solution a little less brightly, and the lamp connected with the ammonium hydroxide solution does not glow at all—not enough current is carried through the ammonium hydroxide solution to heat the filament in the corresponding lamp sufficiently to make it red. Now, the current, for a given fall of potential, is proportional to the conductivity of a solution (p. 48) and, in the equivalent solutions137 the conductivity depends on the proportion of charged particles (the degree of ionization) of the base. It is clear then, that ammonium hydroxide is very much less ionized than are the two other bases. [p078]

The resistance in a tube may be reduced, and the conductivity increased, by reducing the distance through which the current must be carried, i.e. by bringing the electrodes closer together. In the solution of barium hydroxide, we find that we must reduce the distance between the electrodes to about five-sixths the corresponding distance in the potassium hydroxide solution before we obtain, approximately, as bright a lamp from the current passing through it, and in the case of ammonium hydroxide, we must bring the electrodes so close together that they almost touch, the distance being only one or two hundredths of the distance between the electrodes in the potassium hydroxide solution.

For the degrees of ionization of the three bases we have, approximately, the relation αK : αBa : αNH4 :: dK : dBa : dNH4, if we indicate by dK, dBa, dNH4 the distances between the electrodes in the three solutions when the lamps are of uniform brightness, i.e. when the same quantity of current passes through each solution. In this deduction, the conductivities of the bases at infinite dilution (Λ∞) are taken to be the same, which is roughly true.

The experiment gives us a rough measure of the relative conductivities and the relative degrees of ionization of the three bases. It shows that potassium hydroxide is somewhat more ionized than is barium hydroxide, in equivalent solution, and decidedly more than is ammonium hydroxide.

Limiting the further discussion, at this moment, to potassium hydroxide and ammonium hydroxide, we should find that, since in equimolar solutions, a larger portion of the former is ionized than of the latter, the potassium hydroxide solution must contain the larger proportion or concentration of hydroxide-ion, HO−, which is the characteristic ion of bases. It should, therefore, show the chemical characteristics of a base much more decidedly than the ammonium hydroxide solution. That such is the case can be very simply shown by adding equal quantities (0.1 c.c.) of the 0.1 molar solutions to equal volumes (50 c.c.) of water138 containing some phenolphthaleïn. This is an indicator for bases and acids, like litmus, but it is less sensitive to hydroxide-ion than is litmus. We find that the potassium hydroxide causes a very decided change, producing a deep red color with the phenolphthaleïn, whereas the ammonium hydroxide only produces a pink hue.139 [p079]

In all the chemical changes produced by these alkalies, the same difference in intensity of action is shown, that is here exhibited towards indicators. If, for example, we measure the rate of change in an action, which is slow enough to be measured [p080] and which proceeds quantitatively in proportion to the concentration of hydroxide-ion, we find that the measured rates of change indicate the same ratio in the concentrations of hydroxide-ion in potassium and ammonium hydroxide solutions, as is indicated by quantitative conductivity measurements. An action suitable for the purpose is the saponification of an ester, such as ethyl acetate. Under the influence of an alkali, like potassium hydroxide, ethyl acetate is decomposed, more or less rapidly, into an acetate and alcohol: we have, for instance,

CH3CO2C2H5 + KOH → CH3CO2K + C2H5OH.

The rate of saponification is found to be proportional to the concentration of hydroxide-ion, and not to the total concentration of the base, and the action may be formulated more accurately as follows:

CH3CO2C2H5 + K+ + HO− → CH3CO2− + K+ + C2H5OH
or CH3CO2C2H5 + HO− → CH3CO2− + C2H5OH.

For ammonium hydroxide we have similarly,

CH3CO2C2H5 + NH4+ + HO− → CH3CO2− + NH4+ + C2H5OH.

Now, Arrhenius140 proved that the rate of saponification of ethyl acetate by ammonium hydroxide, which is very much slower than the rate of saponification by potassium hydroxide of equivalent concentration, does agree quantitatively, indeed, with the rate demanded by the theory of ionization, when the hydroxide-ion is considered the active component of the bases, to which the saponification is due.

Exp. A rough idea of the difference in the chemical actions of the two bases may be obtained by observing their effects on the ester, methyl acetate, which is decomposed into an acetate and methyl alcohol rather rapidly. To 50 c.c. of (CO2 free) water containing some phenolphthaleïn, 10 c.c. of 0.1 molar potassium hydroxide is added; a similar mixture with 10 c.c. of 0.1 molar ammonium hydroxide solution is prepared. To each of the solutions, 2 c.c. (an excess) of methyl acetate is added (to the ammonium hydroxide solution first), and the mixtures are shaken for a moment. At room temperature, the mixture containing potassium hydroxide will become pale pink in a few minutes, and colorless soon thereafter, while the mixture [p081] containing ammonium hydroxide will still be deep red at the end of 45 minutes.141

In the following tables are summarized some of the results which have been obtained in comparing the activity of bases, in saponifying methyl acetate, and the concentrations of the hydroxide-ion, in the solutions of the bases, as determined by conductivity measurements. The comparisons are made by representing the activity of the hydroxide-ion in a solution of lithium hydroxide by 100 and by expressing the ratio of the activity of a given base to that of the lithium hydroxide in percentages of the activity of the latter. All the bases were used in 0.025 molar concentration, and their degrees of ionization are given in the last column of the table.

Chemical Activity of Bases and Their IonizationA

Base. Activity. Relative
Concentration
of HO−.
Lithium hydroxide 100 97  
Potassium hydroxide 98 97  
Sodium hydroxide 98 97  
Ammonium hydroxide 2 2.5
Ethyl ammonium hydroxide 12 16. 

[A] Whetham, Theory of Solutions, p. 338 (1902). (Cf. Walker, Introduction to Physical Chemistry, p. 277 (1899).)

An ester is decomposed also under the influence of acids, in aqueous solution, into an organic acid and an alcohol, and cane sugar is similarly decomposed into glucose and fructose (grape sugar and fruit sugar): C12H22O11 + H2O → C6H12O6 + C6H12O6. Both actions are found to be caused by the influence of the hydrogen ions of the acids used and to proceed, at a given temperature, with a velocity proportional to the concentration of the hydrogen ions. Now, in 0.1 molar solution, acetic acid is very little ionized (1.3%), as compared with hydrochloric acid (91%), the degrees of ionization being determined by conductivity measurements (p. 50); the relation may easily be demonstrated with the aid of the conductivity apparatus used to show the difference in ionization between potassium hydroxide and ammonium hydroxide. In the presence of 0.1 molar hydrochloric acid, the decomposition of cane sugar actually proceeds at 79 times the rate that it does in the presence of 0.1 molar acetic acid. The ratio of the concentrations of the hydrogen-ion in the two [p082] solutions is, in fact, 70 : 1. There is, therefore, close agreement142 between the relative chemical activity of the two acids and the relation demanded, if we assume, on the basis of the theory of ionization (p. 77), that the chemically active components of the acids are their ions and particularly their hydrogen ions.

In the following table, the relative activities of acids, in accelerating the decomposition of methyl acetate by water, are contrasted, in a similar fashion, with their relative conductivities. The conductivities of acids depend to such an extent on the concentration of hydrogen-ion, which moves five times, or more, faster than the anions and carries therefore the greater part of the current, that the conductivities of acids, in equivalent concentrations, may be considered an approximate measure of their relative degrees of ionization and of the concentrations of hydrogen-ion. For the purpose of comparison the activity and the conductivity of molar hydrochloric acid are both represented by 100. All the acids were used in normal solutions.

Chemical Activity of Acids and Their Ionization.A

Acid. Activity. Conductivity.
Hydrochloric acid 100   100  
Nitric acid 92   100  
Sulphuric acid 74   65  
Acetic acid 0.3 0.4
Formic acid 1.3 1.7
Chloracetic acid 4.3 4.9
Tartaric acid 2.3 2.3

[A] Whetham, Theory of Solutions, p. 338.

In the following chapters and, indeed, throughout our further work, we shall continually meet additional instances of the quantitative relation between chemical activity and ionization. In fact, the results obtained in the field of quantitative measurement of chemical action, of which the above are single instances, have demonstrated, more than anything else, the value of the theory of ionization to chemistry and the necessity of taking it into account in expressing the results of chemical action in mathematical terms. [p083]

Summary.

—Measurements made, then, in three great and independent fields of investigation, electrical conductivity, osmotic pressure and the allied relations, and chemical activity, bring independent testimony to the correctness of the fundamental assumptions of Arrhenius's theory of ionization. In the quantitative study of solutions of electrolytes, wherever secondary disturbing influences are eliminated or are taken into account as far as possible, the three lines of investigation give results which agree satisfactorily143 as to the degree of ionization of the electrolytes under examination. With the aid of this theory, predictions of the course of chemical action may now be made more definitely and with more assurance than ever before in the history of chemistry.

Chemical Activity of Non-ionized Molecules.

—The conclusion reached, on the basis both of qualitative and, particularly, of quantitative evidence, that the reactions of salts, acids and bases (ionogens) in aqueous solutions (the so-called salt reactions) are the actions of the ions and not of the non-ionized molecules, does not necessarily mean that the non-ionized molecules are altogether inactive chemically. Some chemists believe that ions have only the advantage of an enormously greater degree of reactivity.144 If the atoms in a molecule carry electric charges even prior to their separation (ionization), as expressed for instance by the formula H+Cl− for hydrogen chloride (p. 43), it can be readily seen that an action145 resulting from the collision of a molecule of hydrogen chloride with a molecule of potassium hydroxide H+Cl− + HO−K+ ⇄ K+Cl− + HO−H+ would, in some respects, resemble the ionic action H+ + Cl− + K+ + HO− ⇄ K+ + Cl− + HOH. The latter would probably have the advantage of a considerably smaller resistance to the action, and consequently of a far greater speed. But the question of supreme importance and interest to chemistry is the question as to which actions are found experimentally to be of moment in any given case. Now, a great mass of corroborative evidence shows that for the interactions of ionogens in aqueous solutions the ionic actions, probably on account of their enormous speeds, are the important ones. [p084]

On the other hand, there are large numbers of compounds, especially among organic substances, which do not appear to ionize to a measurable extent and whose actions, in large part at least, appear to be the actions of non-ionized molecules. It is characteristic that most of these actions take place at slow, very frequently easily measurable, rates of speed. Critical study shows that even for such actions ionization often plays a very important rôle, at least in some of their stages, and throughout the field of organic chemistry the intimate relations between electrical phenomena and chemical activity can be readily recognized.146 But these relations are not obvious ones and do not, as yet, play a dominant rôle. It is because analytical chemistry deals predominantly with the reactions of ionogens, that the study of the reactions of ions will demand our extended attention.

Reactions in Non-aqueous Solutions.

—Kahlenberg147 has made some interesting and important, although not conclusive, contributions to the problem of chemical activity of ionogens in non-aqueous solutions. He has found, for instance, that zinc, left in contact with a benzene solution of hydrogen chloride carefully freed from moisture, displaces hydrogen.148 Cadmium, aluminium and magnesium, on the other hand, do not evolve hydrogen in such a solution.149 The solution shows an enormous resistance to the passage of an electric current, and the conclusion is drawn by Kahlenberg that the liberation of hydrogen is due to the action of non-ionized hydrogen chloride on the zinc.

It is evident,150 from Walden's equation showing the relation between the ionizing power and the dielectric constant of a solvent (p. 63), that the presumption is that hydrogen chloride in benzene solution is not absolutely non-ionized, but rather that it is ionized in traces.151 No exact measurements of the [p085] degrees of ionization of hydrogen chloride in benzene solution have been made; that the solution shows an enormous resistance to the passage of the electric current and can be, at best, very little ionized, is all that has been established. In default of exact data, the semiquantitative determination by Kablukoff, showing that a 0.25 molar solution of hydrogen chloride in benzene has a resistance of 120 × 106 ohms, is of interest. From the meager data concerning the dimensions of the electrodes used, one may calculate (with the aid of a not unreasonable assumption as to the limiting value of the conductivity, at infinite solution) that the degree of ionization of the acid in the solution is perhaps of the order 5E−9, and the concentration of hydrogen-ion,152 consequently, roughly 10−9. Now, the evolution of hydrogen by means of zinc, in aqueous solutions, takes place according to the equation Zn ↓, + 2 H+ ⇄ Zn2+ + H2 ↑, and depends on a ratio of the concentrations of zinc-ion and hydrogen-ion153 (Chapters XIV and XV, q.v.). Even if the concentration of hydrogen-ion is very small, zinc will liberate hydrogen, provided the conditions are such that the concentration of zinc-ion cannot reach a large enough value to satisfy the equilibrium ratio, and stop the action. Now, in an alkaline solution, zinc-ion is converted into zincate-ion (Zn2+ + 4 HO− ⇄ ZnO22− + 2 HOH) and a large concentration of zinc-ion cannot accumulate. The consequence is that zinc liberates hydrogen freely even from alkaline solutions, for instance from molar solutions of potassium hydroxide, in which the concentration of hydrogen-ion, roughly 10−14, is very much smaller than that calculated for the benzene solution of hydrogen chloride (namely, 10−9). Now, although the values of the solution-tension constants of elements change most decidedly with a change of solvent, it seems likely154 that their ratios, on which their mutual displacement depends, will not be found materially altered. Zinc chloride being insoluble in benzene, the ratio for equilibrium may not be fulfilled for zinc in contact with a benzene solution of hydrogen chloride. Hence, with that solvent, the evolution of hydrogen may, so far as it goes, very well be due to precisely the same machinery as that operating in aqueous solution. The liberation goes on until the metal is protected against any further action by a film of the solid chloride. It seems, therefore, at least possible, that the evolution of hydrogen observed by Kahlenberg and his collaborators155 is a purely ionic action, the same as the similar [p086] action in aqueous solution has been proved to be by quantitative measurements.156 Only exact measurements, comparable with those made in aqueous solutions, can settle the question at issue, and until such quantitative evidence is forthcoming, a definite conclusion that the action of hydrogen chloride in benzene solution is or is not an ionic action is not warranted by the facts.

Equally interesting are Kahlenberg's observations of interaction between hydrogen chloride in benzene solution and a similar solution of copper oleate. Each solution shows absence of appreciable conductivity, yet, when the solutions are mixed, precipitation of copper chloride occurs instantly. Whether we have here an instantaneous action between non-ionized molecules, as claimed by the observer, or whether the minimal ionization157 of the hydrogen chloride and copper oleate, the existence of which we have a right to assume, is sufficient to account for this rapid action, both components being in solution and intimately mixed, is a question of the greatest interest. But until quantitative measurements of all the factors involved in chemical actions in benzene solution are obtained, a very difficult, but necessary task, which the discoverer of the action omitted to perform, no definite conclusion whatever can be based on such results, interesting as they are. Water, although it is only minimally ionized (tables, Chapter VI), hydrolyzes salts like potassium cyanide and aluminium chloride almost instantly, and it has been rigorously proved that the resulting condition of equilibrium involves the ions of water158 (Chapter X). With the possibility that the well-known enormous speeds of action of ions may completely offset the tremendous reduction in concentration of the hydrogen-ion, in a benzene solution of hydrogen chloride as compared with an aqueous solution, further analysis of the relations is imperative.159 Until such investigations have been carried out, we must consider [p087] it possible that the reactions of hydrogen chloride in benzene solution may be reactions of its non-ionized molecules or reactions of its ions. In view of the undoubted minute concentrations of the latter, as compared with aqueous solutions, and in view of the inertness of non-ionized hydrogen chloride in aqueous solutions as compared with the activity of its ions, a benzene solution of hydrogen chloride should show far less ionic activity than an aqueous solution, and that such is the case is brought out clearly by Kahlenberg's interesting experiments.