Oxidation and Reduction in Organic Chem- istry from the Standpoint of Potential Differences The System Hydroquinone and Quinone By FREDERIC STEARNS GRANGER, Ph.D. COLUMBIA UNIVERSITY PRESS 1920 All rights reserved Oxidation and Reduction in Organic Chem- istry from the Standpoint of Potential Differences i The System Hydroquinone and Quinone By FREDERIC STEARNS GRANGER, PhJX fork COLUMBIA UNIVERSITY PRESS 1920 All rights reserved Copyright, 1920 By COLUMBIA UNIVERSITY PRESS Printed from type, March 1920 TO MY MOTHER 435324 ACKNOWLEDGMENT The author wishes to express to Prof. John M. Nelson his apprecia- tion of his kindly encouragement and valued advice and aid in the carry- ing out of this work. ABSTRACT OF DISSERTATION. (1) What was attempted? (2) In how far were the attempts successful? (3) What contribution actually new to the science of Chemistry has been made? (1) It has been undertaken (a) to determine whether the tendency of a typical organic oxidation from one non-electrolyte to another, or its reversal, to take place gives rise to a constant, measurable and reproduci- ble potential difference, when the concentrations of all of the substances entering into the electrochemical equation, by which the process may be theoretically represented, are fixed; and (b) to determine whether this potential varies with the concentrations of these substances in the same way that inorganic potentials have been found to vary, viz., approxi- mately in accordance with the law of Mass Action as expressed in the van't Hoff isotherm. (2) A typical case was chosen which was practically adapted to such an investigation, namely, the oxidation of hydroquinone to quinone, or its reverse. It was found that whenever the concentrations could be fixed or determined to a reasonable degree of certainty, the deviation from the van't Hoff equation was less than the average of the deviations obtained in parallel work on inorganic electrolytes. (3) It has been shown experimentally in the specific instance investi- gated that the electrochemical theory of oxidation and reduction, in its full quantitative significance, is applicable to organic non-electrolytes in the same way and to the same extent as to inorganic electrolytes; that there is no fundamental difference between the two classes of substances, in this respect; that nothing but practical difficulties, peculiar to a par- ticular case in our opinion stand in the way of extending these relation- ships to other organic systems. Incidentally, it has been shown that the solubility product law, which has been developed in connection with the ionic dissociation of difficultly soluble electrolytes and has already been applied to quinhydrone by Luther and Leubner, although apparently applying with a fair degree of accuracy to the dissociation of quinhydrone in the presence of small excesses of hydroquinone, breaks down, evidently, near the point at which the solution is saturated with respect to hydroquinone, and is actually reversed at the saturation point; and that this reversal is prob- ably due, not to a change in the true solubility of the quinhydrone, but to a rapid increase in the dissociation constant in the neighborhood of the saturation point for hydroquinone. PRELIMINARY NOTE. A certain amount of freedom has been exercised in this work, in the use of cer- tain terms. An electrolyte is generally defined as a solution which conducts electricity. For lack of a better word, however, the author has used this term to refer to the dissolved substance to which the solution owes its conductivity. The sense in which the term is used will be apparent from the context. No terms could be found which fitted the case exactly. Again, the author has occasionally used the terms potential, oxidizing potential and reducing potential instead of the usual expression potential difference. The latter term, in this connection, has a purely electrical significance, whereas, when the author omits the word "difference," he is referring to a different thing, although it may have the same numerical value. He is using the term "potential" in a chemical, rather than in a purely electrical sense. In either case, however, it refers to the so-called "inten- sity" or "head" factor of the form of potential energy in question. In this work the directly observed electromotive force measurements are, for the sake of brevity, omitted. Only the so-called single or pole potential differences are given. Neglecting the unknown potential difference at the boundary between the two solutions making up the cell, the electromotive force of the cell, in the direction from organic electrode through external circuit to calomel electrode, may be regarded as the difference between the organic pole potential difference and the calomel pole poten- tial difference, both taken in the direction from pole to solution. The organic pole potential difference is therefore derived by algebraically adding the calomel pole poten- tial difference value, taken as standard, to the observed electromotive force all values being taken in the directions above mentioned. The pole po l ential difference, for a given system measures the so-called oxidizing "power" of that system since it measures its tendency to give up positive electricity. These oxidizing "powers," expressed in this electrical way, are sometimes called oxi- dizing potentials or simply potentials, because of their analogy to electrical and mechanical po entials, and because they can be balanced by electrical potential d fferences, by which means they are measured, and therefore expressed in he s:me units. Oxi- dizing potential differs from e!ectrical potential difference, however, in that it is an inde- pendent specific property of the system in question for the given concentrations and temperature, involving nothing external to the system, instead of a relation between it and its environment, imparted to it by temporary external conditions. It is there- fore spoken of as a potential rather than as a potential difference. The relati nship between oxidizing "power" or potential and the electrical potential difference by which it is measured is analogous to the relationship between the pressure of a steam boiler, which is an independent property of the system, wet steam, for the conditions existing within the boiler, and the mercury level difference of a manometer by which it might be measured. In the present work we are concerned with the chemical significance of the values in question, rather than with their electrical significance. When the term potential is used alone, in this work, instead of potential difference, it is used in the chemical sense in the sense of oxidizing potential. The reducing potential of a system, i. e., its ten- dency to give up negative electricity, is simply its oxidizing potential with the sign changed. INTRODUCTION. It is customary at the present time, as is evident from the contents of our text-books on organic chemistry, to treat the subject of oxidation- reduction of organic compounds in a manner similar to that followed in the field of inorganic chemistry previous to the electro-chemical method. In other words, more attention is directed towards the selection of the most suitable oxidizing or reducing agent for the particular reaction at hand, than to a consideration of the dynamics or energy relationships involved. The oxidizing agent and the temperature, acidity, alkalinity and concentrations of the solution, etc., which furnish the maximum yield of the oxidation product, have been selected from trial experiments rather than from reasoning based upon recognized general principles underlying reactions of this kind. Furthermore, some claim 1 that this electrochemical method of studying the energy changes in reactions like this is only applicable .to electrolytes, and that it is entirely too hypo- thetical to view oxidation and reduction, in the case of organic com- pounds, in this way. It is quite generally accepted that the electro- chemical way for accounting for oxidation and reduction, in the case of electrolytes, is far superior to the older way, based upon such ideas as sub- stitution of oxygen for hydrogen, removal of hydrogen, addition of oxy- gen, etc. This is very often emphasized in the lecture room by the familiar experiment known as "oxidation at a distance." 2 It therefore follows that if it is not permissible, until more experimental data are available, to use the electrochemical way for explaining oxidation and reduction in the case of non-electrolytes, the older way must be employed. Thus we will have two theories for oxidation-reduction, one for inorganic and one for organic chemistry. PREVIOUS WORK IN ORGANIC CHEMISTRY. In the field of organic chemistry, the very little work that has been done at all, in the investigation of Oxidizing Potentials, has been only of the most empirical character. Bancroft, in i892, 3 included, among the large number of more or less common oxidizing and reducing mixtures with which he worked, two organic reducing agents, namely, alkaline solu- tions of hydroquinone and of pyrogallol, of indefinite composition. The proportions used in making up the organic solutions were not given, but had they been, they would not have thrown much light upon the composi- tion of the resulting mixture since hydroquinone and pyrogallol not only 1 Bray and Branch, /. Am. Chem. Soc., 35, 1440 (1913); Lewis, Ibid., 35, 1448 (1913) . 2 See "Qualitative Chemical Analysis," Stieglitz, p. 253. 3 Z. physik. Ch., 10, 387. 10 an electronic transfer which can be used to develop an e. m. f . ; as the pre- liminary study is certainly in harmony with this idea we may interpret the reaction as follows for the present: CH 2 = CH 2 + C1 2 ^ +CH 2 +CH 2 + 2C1- -> CH 2 C1 CH 2 C1." In the second, a similar experiment was carried out with the cell HCHO | AT/5 NaOH | O 2 . Before adding the formaldehyde, oxygen was bubbled against one of the platinized electrodes, in the U-tube of sodium hydroxide solution, and the electromotive force measured. Every time a drop of formalin solution was added to the other arm, the electromotive force would show an abrupt increase to a comparatively constant value. Similar experiments were carried out with phenolphthalin (reduction product of phenolphthalein) and oxygen and with phenolphthalein and hydrogen. In the first case no perceptible pink color appeared on connecting the poles until a current was forced through the cell. In the second case, however, the solution became lighter without the use of an external current. This concludes the list of all the work that could be found on organic potentials up to the present. If we compare this with what has been done in inorganic chemistry, the contrast is apparent. In inorganic chem- istry, investigation has invariably shown that potentials varied with the concentrations of the substances involved in the electrochemical reaction, to which the oxidation or reduction was theoretically attributable. And when the absence of interfering factors seemed to permit, which was true in nearly every case investigated, this variation was found to be in as good accordance as could be expected with the van't Hoff Isotherm, which is the thermodynamical verification of the Law of Mass Action. 1 Now, in every one of these cases, either the oxidizing or reducing agent, as the case may be, or its product, was an electrolyte. The van't Hoff equation has never been applied to a case where both were non-electro- lytes. In none of the organic cases just reviewed, has any attempt been made to apply this or any other generalization, or to find any new rela- tionship; to establish any connection between the potentials measured and any particular reaction or principle. In all of the organic work, no account has been taken of the products of the reaction. There has been no attempt at coordination of any kind. They are simply isolated and empirical measurements. 1 Nernst, "Theoretical Chemistry" (English Ed.), p. 690; Ibid., p. 785; Peters, Z. physik. Ch., 26, 193 (1898); Friedenhagen, Z. anorg. Ch., 29, 396 (1902); Tower, Z. physik. Ch., 18, 17 (1.895); Luther and Michie, Z. Electroch., 14, 126 (1908); Luther and Sammet, Ibid., n, 293 (1905). Since this equation is based on the perfect gas laws, or, in the case of solutions, on the corresponding relationships for osmotic pressure, it is only applicable to the degree in which these laws hold, in the case in question. T I According to the van't Hoff equation or the principles of chemical equilibrium, a reducing or oxidizing agent, in solution absolutely free from any of its product, should have a potential of infinity (minus or plus). Therefore, since so-called "pure" solutions of oxidizing or reducing agents do not show infinite potentials, the presumption is that they contain quantities of products sufficient to account for the potentials observed, though too small to be detected by analytical means. Actual calcula- tions from existing experimental data will show that the concentrations of these substances can be that small without the theoretical potentials being any larger than those actually found for pure solutions, which are usually of the order of about i volt, and never more than two. This is reasonable enough, for, if the solution were so pure as to give a potential greater than this, its oxidizing or reducing power would be so great that, we would expect it to react with the water, or anything else available producing product until the excessive tension was relieved. Measure- ments on so-called pure agents are therefore to be taken, as quantitative data, with more or less reserve, as we are not justified in assuming that these indeterminately small quantities of products present may not vary greatly, relatively speaking, in different specimens, even though a par- ticular series of measurements has given fairly reproducible results. One more point may be mentioned. Mere potential measurements, without a correlation of some sort, prove nothing even qualitatively. Any two solutions containing electrolytes 'will give a potential difference, due to the relative discharging tendencies of their ions, even when no oxidizing or reducing agent is added, unless, by a rare coincidence they happen to have practically the same oxidizing potential. PURPOSE OF PRESENT WORK. From what has just been said it 'is evident that we cannot consider a potential measurement made upon a "pure" solution of a reducing agent as a characteristic fixed property of that substance (like, for in- stance, a melting point). Such a value, on the contrary, would be an accidental one. If it is only desired to measure the reducing power of that particular solution, for empirical purposes, then such a measurement gives all the information that is required. But it is useless as a basis for general calculations. For it has been found, whenever investigated, that the potential is as much a property of the oxidation product of the reducing agent as of the reducing agent itself, and is a function of the con- centrations of both. In other words, it is a property of the system rather than of the substance. There is, of course, such a thing as the intrinsic tendency of the oxidizing or reducing agent to pass to another stage of oxidation, by virtue of which it owes its oxidizing or reducing properties. This is a characteristic property of the substance, being independent of 12 concentrations and represented by the term RT/wF log e K in the van't Hoff equation, in the form: RT C, W 'CV".. RT * = ~^T log, n , ni , n ,n*r ~^T iogeKo or \^i \^z Mr RT Ci m C 2 W2 7T = log e c m/c /B j ' - + 7T . (l) It determines the equilibrium constant. But we cannot measure this directly, by any means known to us at the present time. We can only determine the potential of the system, under known concentration con- ditions, and for this calculate by means of the above equation the intrinsic oxidizing of reducing tendency, which is sometimes written TT O and some- times A. 1 We must first, therefore, know all the concentrations involved and test our relationship. These facts have been recognized in inorganic chemistry for years (though confused by the tendency to attribute all oxidation to oxygen and all reduction to hydrogen). In organic chemistry, however, as has been shown, no attempt has ever been made, so far as we can discover, to put organic oxidation and reduction on a similar basis. The idea seems to have prevailed that the electrochemical theory of oxidation and reduction was not applicable to non-electrolytic substances, such substances being oxidized or reduced through the medium of "nascent" oxygen or hydrogen liberated by the oxidizing or reducing agent, or by the removal of hydrogen or oxygen, from the non-electrolytic substance by the agent. The purpose of the present investigation is to apply the electrochemical theory to a typical organic oxidation-reduction in a man- ner parallel to that in which it has been applied in inorganic chemistry, and to see whether or not the theory fits the observed facts as satisfac- torily as it has been found to in inorganic chemistry. The conception that an electric charge can be given up by an elec- trically neutral body and still leave an electrically neutral body is, of course easily explained by the giving up of an oppositely charged ion at the same time, as illustrated by the following two equations: CH 3 CH 3 OH ^ CH 3 CHO + 20 + 2H+ (2) C 6 H 4 (OH) 2 - C 6 H 4 2 + 20 + 2 H+ (3) SELECTION OF A SYSTEM. The first problem confronted was the selection of a suitable system, one that would adapt itself to such an investigation. The required quali- fications were the following: (i) It must be a typical reversible organic oxidation. 1 This should not be confused with the use of the symbol A to represent free energy. 13 (2) The reaction must be well defined, that is, the equation must be definitely known, and it must be as free as possible from intermediate stages and tendency to go to further stages of oxidation. (3) It must be possible to determine the concentrations of the sub- stances involved, with sufficient accuracy, by analytical means, or else to fix or control them in some way. (4) The substances involved must all be sufficiently stable so that the system will not become lost in a maze of unknown side reactions before a definite potential can be obtained. And there must be no reactions entering in to throw the potential off and prevent the true reading from being obtained. (5) The substances must all be sufficiently non-volatile so as not to be removed appreciably by the bubbling of an inert gas through the solution to remove the air. (6) They must be appreciably soluble in water. (7) They should not be acids or bases or react in any way with the acid or base used for fixing the hydrogen ion concentration, unless the concen- trations can still be controlled, in spite of this fact. It is evident that a system fulfilling these specifications is not easy to find in organic chemistry. The system hydroquinone/quinone was selected as more nearly meet- ing these requirements than any other system that was thought of. Its advantages lay in the following facts: The solubilities were such that the concentrations could be fixed by using saturated solutions, which would still be fairly dilute. It is a well defined, clean cut, perfectly reversible reaction. Neither substance is very volatile. Quinone is about as vola- tile as water, but the working conditions were not such that this gave any trouble. Neither substance is a pronounced acid or base. The reac- tion is a rapid one. The substances are not very unstable, except in alka- line solution. Quinone can be determined accurately by analytical means, even in the presence of hydroquinone. The system possesses, however, certain disadvantages. Hydroquinone, when added to an alka- line solution, is a strong enough acid to nearly neutralize the alkali, so that a special investigation would have to be undertaken to find a means of accurately determining the hydrogen ion and other concentrations. Furthermore, hydroquinone and quinone are both very unstable in alka- line solution, somewhat so even in pure water and more so in neutral potassium chloride solution, the solution becoming dark brown owing to the formation of a tarry material. Finally, the system is complicated by the fact that hydroquinone and quinone combine directly, in solution and elsewhere, to form the beautiful bronze-green crystalline addition product, quinhydrone. QUINHYDRONE. Quinhydrone is formed immediately whenever quinone is added to a solution of hydroquinone, and vice versa. In solution it exists in mobile equilibrium with its components. In aqueous solution, it is highly dis- sociated (over 90%), but its solubility is so slight that it is impossible for a considerable concentration of one of its constituent substances to exist in the presence of a considerable concentration of the other. There- fore, a solution cannot be simultaneously saturated with hydroquinone and quinone, the molar solubilities of these substances being, respec- tively, about 640 and 125 times that of quinhy drone. In a nearly satura- ted solution of either, the addition of only a small quantity of the other is required to produce a precipitate of quinhydrone. INFLUENCE OF HYDROGEN ION CONCENTRATION. In Equations (2) and (3) given above, hydrogen ion appears as a prod- uct of the oxidation of alcohol and of hydroquinone, and would appear, similarly, in the cases of most organic oxidations. Theoretically, then, acidity should favor the reaction in the oxidizing direction (backwards, as written) and oppose it in the reducing direction. The reducing poten- tial of the system should therefore decrease, that is, the oxidizing poten- tial should increase, with increasing hydrogen-ion concentration. Acidity and alkalinity have long been recognized as important fac- tors in determining the course of oxidation and reduction, especially in organic chemistry. It would therefore seem of importance to test the above theoretical influence experimentally. This never seems to have been done, even qualitatively, in organic chemistry, though the quantita- tive theory has been checked in a number of cases in inorganic chemis- try. But because it is found that the oxidizing potential of the system is favored by acidity, and the reducing potential by alkalinity, it does not follow that alkalinity should always be found to favor the action of hydroquinone as a reducing agent or acidity to favor in every case the reduction of quinone. The influence of hydrogen ion upon the other system, entering into the reaction, must be taken into account. For instance, the oxidation of hydroquinone by a ferric salt should be hindered and the reduction of quinone by an iodide, favored by acidity, for we have: C 6 H 4 (OH) 2 ^ C 6 H 4 2 + 20 + 2H + . 20 + 2Fe + + + ^ 2Fe^ C 6 H 4 (OH) 2 + 2 Fe +++ ^ C 6 H 4 2 + 2 Fe + + and C 6 H 4 2 + 2 H + ^ C 6 H 4 (OH) 2 + 20 20 + 2l~ I 2 C G H 4 2 + 2H+ + 2l~ C 6 H 4 (OH) (4) (5) 15 But the oxidation of hydroquinone to quinone by permanganate should be favored by acidity, because 5 (C 6 H 4 (OH) 2 ^ C 6 H 4 2 + 2H+ + 20) __ 2 (MnQ 4 - + 8H + Mn ++ + 50 + 4H 2 O) 5C 6 H 4 (OH) 2 -f 2 Mn0 4 - + 6H+ ^ 5C 6 H 4 O 2 + 2 Mn ++ + 8H 2 O In the hydroquinone reaction, it will be seen, we have involved only one hydro- gen ion per charge transferred, whereas in the permanganate reaction the ratio is 8 to 5, so that the influence in the latter system outweighs that in the former. It makes a difference also what concentrations are being controlled. For instance the oxidation of chromium from the trivalent to the hexavalent condition by sodium peroxide may be written so as to appear to be either hindered or favo ed by alkalinity. 2 (Cr(OH) 3 + 50H- ^ CrO 4 = + 4H 2 O +30) 3 (2H 2 + O," ^ 4 QH- + 2) 2Cr(OH) 3 + 3O 2 = ^ 2CrO 4 = + 2H 2 O + 2OH~ or 2(Cr + + + + 80H- ^ Cr0 4 = + 4H 2 O + 3) 3 ( 2 H 2 + 2 = ^ 4 OH- + 20) - 4OH~ + 3O 2 = ^ 2CrO 4 ~ + 2H 2 O This simply means that, if we are starting with chromic hydroxide, excess alkali would tend to oppose the oxidation, alkali being one of the products ; whereas, if we are starting with a chromic salt, additional alkali would be used up in the formation of the hy- droxide. (This may not be true, experimentally. Other factors, which we have not considered, which are unknown perhaps, may enter in. For instance, acidity or de- creasing alkalinity would hinder the reaction by removing peroxide ion. But the above would be the interpretation of the equations as they stand ) So we must consider the actual conditions in writing our oxidation equations. In order to determine the influence of hydrogen-ion concentration on the potential, it was hoped that a system could be found in which the concentrations of the other factors could be fixed by the use of saturated solutions of the substances involved, while the acidity varied. In the one selected, however, the formation of quinhydrone made this impossible, as already pointed out. But, fortunately, the concentrations could be fixed, as well, though indirectly, by saturating the solution with quin- hydrone and one of the components. PLAN OF THE WORK. In view of the considerations outlined above, the investigation was carried out according to the following plan: Potential measurements were made on a series of solutions saturated simultaneously with hydro- quinone and quinhydrone, in which the hydrogen-ion concentration was varied and the conductivity furnished by the use of hydrochloric acid and sodium hydroxide, in various concentrations; for the neutral solu- tion, potassium chloride was added as electrolyte. A parallel series was started, with quinone and quinhydrone, but because of the instability of the former, it was not carried any further than normal and tenth- i6 normal acid. To determine the effect of varying other concentrations, a series was run in tenth-normal acid saturated with quinhydrone with varying concentrations of hydroquinone added. In order to know the concentrations it was necessary to determine the solubilities of the sub- stances and the dissociation constant of quinhydrone. The results were compared with the van't Hoff equation, as a convenient standard of reference, in the form RT , (quinone) (H+) 2 RT t = 2-F 10g < (hydroquinone) ' ~ 2 T ^ K ' (9) In neutral and alkaline solutions, because of the instability of hydro- quinone and quinone and other difficulties, data of definite quantitative significance were not obtained. The results were, however, of qualitative interest. In acid solutions, the results paralleled those in inorganic chem- istry. PART I I. POTENTIAL MEASUREMENTS IN ACID SOLUTIONS. Mixtures were made up as described below, and their potentials mea- sured in a half -element vessel, against a saturated potassium chloride calomel electrode, by means of an e. m. f. combination of the type Hg | HgCl sat. KC1 | sat. KC1 | solution A | Pt, employing a sensitive potentiometer and galvanometer. Saturated potassium chloride solution was used as the connecting medium. The cells were kept immersed in a constant temperature bath at 25 C., the temperature remaining constant to a hundredth of a degree. Nitrogen was bubbled through the mixture in the cell, for the first few hours, to insure complete removal of the air and to provide agitation at the start. The gas inlet and outlet tubes were then closed to prevent access of air. The potential of each cell was measured at least once a day over a period of from one to three weeks. The following mixtures were investigated: A. Normal hydrochloric acid, saturated with hydroquinone and quin- hydrone. B. Tenth-normal hydrochloric acid, saturated with hydroquinone and quinhy drone. C. Hundredth-normal hydrochloric acid, saturated with hydroquinone and quinhydrone. D. Normal hydrochloric acid, saturated with quinone and quinhydrone. B. Tenth-normal hydrochloric acid, saturated with quinone and quin- hydrone. F. A/"/ 10 HC1, o.i molar hydroquinone, saturated quinhydrone. G. N/io HC1, 0.05 molar hydroquinone, saturated quinhydrone. H. A/YIO HC1, 0.02 molar hydroquinone, saturated quinhydrone. I. N/io HC1, o.oi molar hydroquinone, saturated quinhydrone. J. N/io HC1, saturated with quinhydrone alone. Two or more cells of each mixture were made up and examined, usually at different times, and sometimes with different lots of materials, in order to determine the reproducibility of the potentials found. The calomel electrode was assigned the value 0.5265 volts, which is the value provisionally adopted by Dr. H. A. Fales 1 for 25 C. Adding this to the electromotive force of the cell, when the mix- ture is the positive electrode, gives the single potential difference 2 or the oxidizing potential of the mixture in question, plus or minus any contact potential differences at the boundaries of the solutions. The 1 Private communication. 2 See Le Blanc, "Text-Book of Electrochemistry," Eng. Trans, of 4th Ger. Ed., p. 234, et seq.; and the articles already cited. i8 usual custom of writing potentials with a minus sign before them when the tendency is for the solution to give up positive electricity to a zero electrode, i. e., to oxidize it, has been departed from because it seems more logical and less confusing to regard such potentials as positive. That is, such a mixture would be said to have a positive oxidizing potential or a negative reducing potential, and we are expressing the values as oxidizing potentials. The question is, of course, an arbitrary one. The results are given in Table I. TABLE I. MIXTURE A. Cell No. 1. Cell No. 2. Cell No. 3. Hours from start. Volts, i Hours. Volts. Hours. Volts. o 0.8884 iVi 0.8889 I2 3 / 4 0.8834 24 l / z 0.8833 23 ! /4 0.8866 29 3 / 4 0.8832 49 J /4 0.8810 5 iVi o . 8847 3I 3 A 0.8830 78 0.8812 77 0.8871 58 3 A 0.8830 97 */4 0.8807 95 3 A 0.8867 7574 0.8832 I2lV2 0.8827 ii9 3 A 0.8851 9472 0.8831 143 V4 . 0.8805 147 0.8841 II 9 74 0.8828 16272 0.8801 16774 0.8817 H472 0.8823 Cell No. 3 (Continued). I 73 1 /2 o . 8804 Hours. Volts. I92 3 /4 0.8797 369 3 / 4 0. 8776 21574 0.8819 384 o. 8769 2 4 l74 0.88II 40774 o. 8767 26 5 3 /4 o . 8805 43274 0. 8765 28972 0.8786 547 o. 8753 31274 0.8790 479V2 o. 8746 33972 0.8782 MIXTURE B. Cell No. 1. Cell No. 2. Cell No. 3. Cell No. 4. Hours. Volts. Hours. Volts. Hours. Volts. Hours. Volts. l / 4 0.8213 I8V4 o . 8309 2'/2 o . 8303 o o. 8091 29 x /2 0.8234 24 J /4 o . 8303 21 0.8306 87 2 0.8293 5lV2 0.82892 42 J /4 0.8319 548/4 0.8315 22V2 0.8299 74 J /2 0.82992 72 0.82762 69 3 A 0.8309 4772 0.8296 97 o. 83282 89 J /2 0.8282 9i 3 A 0.8317 70 0.8295 I2I J /4 0.83I5 2 i M'A 0.8311 93 0.8289 143 O.83II I2l72 0.8285 16774 0.8303 I42 3 /4 0.8290 Cell No. 4 (Continued). i8 7 3 /4 O.83OO 16674 0.8298 Hours. Volts. 2II l /2 0.8294 1 9 0.8298 338 0.8297 235 3 A 0. 289 21474 0.8294 362 1 /4 O.8295 26972 0.8288 24572 o . 8303 386V2 0.8304 297 3 A O.8266 264 0.8300 4I3V4 0.8292 31372 O.8274 29274 0.8306 432 V2 0.8278 331 0.8250 3097'2 o . 8303 1 All potential values given in this table are the single potential differences de- rived from the observed electromotive force of the combination by adding the value (0.5265) of the calomel electrode. 2 Average of several readings taken that day. TA: BUS I (Continued}. MIXTURE C. Cell No. 1: Cell No. 2. Hours. Volts. Hours. Volts. l/ 4 0-7321 Va 0.7753 2O 0.7626 20V2 0.7756 43 3 A 0.7736 50 0.7734 73'A 0.7718 66/4 0.7709 90 0.7714 9I*/2 0.7673 if 4 Va 0.7691 I40V2 0.7697 i6 3 3 A 0.7718 i63 3 A 0.7700 187 0.7725 i87V2 o . 7696 2IO l /2 0.7725 2i7Va o . 7687 240 3 A 0.7708 234V4 0.7685 257V2 0.7712 262 3 / 4 o . 7686 286 0.7738 aftvVi o . 7670 305 V4 0.7741 307 o . 7642 33i 3 A o . 7662 361 o . 7646 380 0.7634 402 3 A 0.7638 MIXTURE D. Cell No. 4. Cell No. 5. Hours. Volts. Hours. Volts. O I 0353 I .OIO2 3 I .0190 l /2 I. 0137 22 l /a 0.9928 2O o . 9944 49 0.9845 23 J /2 0.9924 76 3 A 0.9850 4 2 3 A 0.9829 94V4 0.9847 71 0.9836 I20V2 o . 9840 93'A 0.9841 146 0.9838 116 o . 9868 1 68 0.9826 I4lV4 o . 9860 195 Va 0.9835 I62V4 o . 9842 2I4 l / 0.9831 I86V2 o . 9884 2 3 8 3 /4 0.9835 217 o . 9869 26l 0.9786 239 l /4 o . 9849 28 9 3 /4 0-9791 259 3 A 0.9835 3I2 X /4 0.9781 275 0.9834 334*/a 0.9826 359 3 A 0.9825 38o 3 / 4 0.9853 405V4 0.9862 435V2 0.9869 457 3 A 0.9762 478V2 o . 9808 20 TABLE I (Continued). MIXTURE E. Cell No. Hours. Volts. 0.9517 I5 3 /4 o . 9502 40/t 0.9469 65 3 A 0-9434 94 3 A 0.9390 114 0.9379 I36V4 0-9371 l62 3 /4 0.9348 187 0.9330 233V2 0.9299 26o l / 2 0.9285 291 0.9284 306V4 0.9268 329V2 0.9280 353 3 A 0.9279 378*A 0.9275 406 'A 0.9266 448 l A 0.9307 473*/4 0.9315 492 o . 9307 Cell No. i. Hours. Volts. 4 3 A 0.8695 iSVz o . 8705 38 0.8701 6S'/4 0.8695 85 0.8689 H3 3 /4 0.8685 I42V4 0.8676 167 0.8666 191 0.8663 238V2 0.8625 Cell No. 1. Hours. Volts. ' I8V2 0.8860 42 1 A 0.8850 . 75 J A 0.8837 90 3 /4 0.8828 I23 3 A 0.8807 I47 1 /* 0.8790 173 0.8785 243 0.8754 3I5V4 0.8715 MIXTURE F. MIXTURE G. Hours. O 9 22 1/4 47 3 A 76'A 96 H8'/2 I44 3 A 169 215'A 273 335 3 A 360 1/' 388V, 43Q 3 A 455 V4 Cell. No. 2. Hours. 2 3 /4 15 34 3 A 64 3 A Si 1 /* H0 3 /4 I36 3 A .Cell No. 2. 235 Volts. 0.9524 0.9505 0.9496 o . 9456 0.9412 0.9395 0.9385 o . 9360 0.9340 0.9307 0.9290 0.9293 0.9287 0.9301 0.9306 0.9297 0.9285 0.9318 0.9327 0.9319 o 9330 Volts. o . 8702 0.8689 0.8687 0.8692 0.8681 0.8686 0.8658 0.8671 0.8663 0.8635 Hours. I5 3 A 39 3 /4 72'A 88 I20 3 /4 W/! 309 Cell No. 2. Volts. 0.8860 0.8854 0.8850 O . 8849 0.8839 0.8829 0.8818 0.8790 0.8760 21 TABLE I (Concluded). Hours. Cell No. 1. 72 I0 4 3 /4 154 293 Hours. 33'A 49 8 1 1 /* 105 X A I 4 I 3 / 4 2OI 26 9 3 /4 Cell No. 1. Hours. 4 3 A I8V4 89 l /s 109 "8V* Cell No. 1 . 174'A 202 /2 226V4 251 298V2 Volts. o . 9030 0.9000 0.8992 0.8982 0.8966 0.8951 0.8933 0.8893 0.8839 Volts. 0.9049 o . 9068 0.9062 0.9051 0.9033 0.9022 0.8998 0.8975 Volts. 0.9155 0.9142 0.9132 0.9116 0.9100 o 9099 0.9079 o . 9069 0.9047 0.9024 0.9012 0.8995 0.8952 MIXTURE H. MIXTURE I. MIXTURE J. Hours. v< 21 54 69 3 A 102 Vi 126 i5iVi I73 3 A i94Vi Cell No. 2. Volts. 0.9018 0.9014 0.9007 0.8998 0.8987 0.8978 0.8969 0.8938 0.8921 Hours. Vi 30 3 /4 46'A Cell No. 2. 102 3 / 4 "8V4 i98Vi 267 Hours. I3V4 37 3 A 69 8?Vt ii6 3 / 4 I34V4 i64Vi i88 3 / 4 211 Cell No. 2. Volts. 0.9071 0.9062 o . 9062 0.9055 0.9044 0.9027 o . 8994 0.8971 Volts. 0.9136 0.9130 0.9119 0.9108 0.9094 o . 9082 0.9067 0.9054 o . 9042 0.9007 From the above data, it will be seen that each of the various mixtures has its characteristic potential, which is fairly reproducible, the potentials being obviously, therefore, a function of the relative concentrations of the substances involved, and that the potentials given by the solutions saturated with hydroquinone become "constant" almost from the start, and remain so for a long time, after which they begin to fall off gradually. The hydroquinone solutions which were not saturated with it, however, showed potentials which fell off almost steadily throughout the period during which the cell was kept, at the rate, usually, of about a milli- 22 volt per day, showing the necessity of fixing the concentrations by satura- tion, in order to obtain constant potentials. The potentials given by the solutions saturated with quinone showed the peculiar behavior of falling several hundredths of a volt, and then remaining stationary for a long time. In the tenth-normal acid, this constant value was reached only after a much longer time than in the case of the normal acid, but the fall in the case of the latter was greater. Moreover, it was only when a very large excess of solid quinone was present, that these stationary periods were obtained with D and B at all. The above points and the exact degree of constancy and reproducibility are brought out more clearly by graphical representation in the charts, in which the above potentials are plotted against time. 1 It is convenient to consider the above data in groups or series. 'Series I, consisting of Mixtures A, B, C and Series II, consisting of Mixtures D and K, show the effect of varying hydrogen-ion concentration or acidity on the potentials, the concentrations of the other substances involved being kept approximately constant, in each series. Series III, consisting of mixtures B, F, G, H, I, J and K, shows the effect of varying hydro- quinone and quinone concentrations with approximately constant acidity. II. SOLUBILITIES AND THE DISSOCIATION OF QUINHYDRONE. The purpose in using saturated solutions was to fix the concentrations of the substances in question, throughout the life of the cell, regardless of side reactions taking place. If we wish to know what these concentra- tions are, however, we must know the solubilities of the substances. Solubility of Hydroquinone. No data on the solubility of hydroquinone at 25 C. could be found in the literature, so it was determined by evap- orating 5 cc. of the saturated solution to constant weight, in vacuum, at room temperature, in a small weighed flask. The results were as follows : II. Grams per 100 cc. solution. Moles -- 2 -- . ----- per Solvent. No. 1. No. 2. No. 3. No. 4. No. 5. Average. liter. Water ............. 7-094 7 .091 7.112 7.086 7.10 0.645 o.oi N HC1 ........ 7.060 7.128 7.136 7.028 7.146 7.10 0.645 o.iNHCl .......... 6.978 6.944 6 -96 0.633 NHC1 ............. 5.436 5.442 5-44 0.494 It will be seen that the solubility of hydroquinone in water is decidedly decreased by the presence of hydrochloric acid. Solubility of Quinone. The solubility of quinone in water at 25 C. was determined by Robt. Luther and A. Leubner, 2 using the analytical method of Amand Valeur, 3 which consists in titrating the iodine liberated by 1 See pp. 29 and 32. 2 /. /. prakt. Chem., n. f., 85, 314-321 (1912). 3 Compt. rend., 129, 552 (1899). 23 quinone from a hydrochloric acid solution of potassium iodide, with thio- sulphate. They give the value 1.265 moles per liter in which there is evidently an error in the placing of the decimal point, since this figure is just about ten times too large. Using the same method, we obtained the following: TABLE III. . Moles per liter. > Solvent. No. 1. No. 2. Grams per 100 cc. Water 0.1266 0.1266 1.37 JV/ioHCl 0.1275 0.1275 1-38 NHC1 ' 0.1332 0.1332 1-44 It will be observed that the solubility of quinone in water is decidedly increased by the presence of hydrochloric acid. Solubility and Dissociation of Quinhydrone. The determination of the solubility of quinhy drone is complicated by the fact that it is highly dissociated in aqueous solution into its two components, hydroquinone and quinone. That it is made up of equimolecular proportions of these has been shown by Liebermann, 1 by obtaining the maximum yield from equi- molar proportions upon mixing aqueous solutions of the two; by Hesse 2 by acetylation, which attacked only the hydroquinone, and removing the quinone by evaporation; and by Nietzki, 3 who reduced the quinone with SO 2 and titrated back the excess with iodine. The same fact was also confirmed in the analytical work of Valeur 4 and of Luther and Leub- ner, 4 and similarly, in the present work, in which it was further verified by mixing together equivalent quantities of hydroquinone and quinone, in solution, and determining the percentage of solids in the filtrate which checked that in a solution saturated with quinhydrone directly. In Series III 5 the hydrogen ion and quinhydrone concentrations were kept approximately constant and the hydroquinone concentration .was varied. But since a mobile equilibrium exists in solution between the quinhy- drone and its dissociation products, we have also a varying concentra- tion of quinone to consider. In fact, it is the concentration of quinone rather than that of the quinhydrone in which we are primarily interested. The latter does not enter directly into the reaction whose potential we are measuring, but its unavoidable presence is utilized as a means of regulating the concentrations of the quinone and hydroquinone. In order to know what these are, however, we must know not only the true solubility of the quinhydrone but its equilibrium constant as well. By "true solubility" we mean the concentration of dissolved undissociated 1 Ber., 10, 1615 (1877). 2 A., 200, 248 (1877). 3 Ber., 10, 2000 (1877) and A., 215, 130 (1882). 4 Loc. cit. 6 See page 22. 24 quinhy drone in equilibrium with the solid. The quantity of quinhy - drone which dissolves in saturating the solution we will designate as the ' ' apparent solubility . ' ' In Series I and II it is also necessary to know the solubility, etc., of the quinhy drone, since the solubilities do not remain constant, owing to the effect of the hydrochloric acid upon them. Luther and Leubner 1 undertook to determine the true solubility and dissociation constant of quinhydrone in water by the method used by von Behrend 2 in the case of the phenanthrene picrates, namely, the de- pression of the apparent solubility by an excess of one of the dissociation products. They saturated water and hydroquinone solutions of known concentrations with quinhydrone at 25 C. and determined the total quinone (combined and uncombined) in the filtered solution by Valeur's method, which is perfectly applicable, owing to the complete dissocia- tion of the quinhydrone as the quinone is removed by the iodide. This total quinone represents (formula-weight for formula -weight) the total quinhydrone which has dissolved in saturating the solution, or the ap- parent solubility. They made a number of determinations for each con- centration of hydroquinone taken, the averages of which, as given in Table IV, they used in their calculations as follows: Let s = true solubility of quinhydrone , in formula-weights per liter a = apparent solubility of quinhydrone ' b = excess of hydroquinone added h = actual concentration of hydroquinone f in moles per liter q = actual concentration of quinone Then total hydroquinone (combined and uncombined) = a -}- b = s-f-h and h = a -f- b s total quinone (as determined by titration) = a = s + q, and q = a s Let K = *L*-B f and P = Ks = h X q = (a + b s) (a s) (10) s If the "Law of Mass Action" held perfectly, and s remained constant throughout the range of experiments, K would be a constant, namely, the dissociation constant. If the experimental precision were fine enough, s and K could be calculated from any pair of determinations by means of simultaneous equations of the form of equation (10) below. But comparatively slight deviations from these ideal conditions render this method of calculation inapplicable, so that recourse must be had to a method of trial and approximation which Luther and Leubner car- ried out in the following way: By trying different values for s in the equation (a + b s)(a s) = P (10) 1 Loc. cit. 2 Zeit. f. physik. Chem., 15, 183 (1894). a. P. Deviation. _ P (Average.) a -f b. (s 0.0013.) from mean. = 0.0013 0.01827 0.01827 0.000288 0.000009 0.221 0.01421 O.O242I 0.000296 0.000001 0.227 O.OII50 0.03150 o . 000308 O.OOOOII 0.236 o . 00664 0.05664 0.000296 0.000001 0.227 they found the value for s for which P showed the nearest approach to constancy, that is, for which the sum of the deviations of P from its mean value, were at a minimum. This value for s they found to be 0.0013, for which the mean value of P was 0.000297, and the sum of the devia- tions O. 00002 2. TABLE IV. a. b. O 0.01 0.02 0.05 We repeated the work of Luther and Leubner, extending the range up to the saturation point for hydroquinone, in water, tenth-normal and normal hydrochloric acid. The solutions, with excess of quinhy- drone, were placed in a large test tube fitted with a spiral mechanical stirrer and immersed in the thermostat, and stirred vigorously and con- tinuously. Samples were taken about every fifteen to thirty minutes, by means of a pipette fitted with a filter, until two successive titrations gave the s0me result, which was usually the case with the first two sam- ples. In a number of cases fresh mixtures were made up and tested as a check, and the results in every case were almost identical with the original, so it was not deemed necessary to verify all the solutions in this way, the regularity of the results and the parallelism between the aqueous and acid solutions also serving as a check. TABLE V. Part i. Water. No hydrochloric acid. Moles per liter. quinone d. j3 O +!<*. V C p X 10 for values of s at heads of columns. II III < S H 10 tN CN o 8 00 | | & < X! ' C i 8 i 8 | b. a. b + a. d d d d d o" o d 0.0178 o .0178 272 274 275 282 282 283 284 284 O.OI 0.0135 0.0235 271 273 275 28l 282 282 283 283 0.02 0.0106 .0306 273 274 2 7 6 284 284 285 286 286 0.05 0.00625 o -05625 273 275 _ 2 7 8 290 290 291 292 293 O. I 0.00374 10374 250 256 263 28l 282 284 286 287 O.2 0.00244 .20244 229 240 250 288 290 294 298 300 0-3 0.00189 o .30189 177 193 207 265 268 274 280 283 0.4 O.OOI79 0.40179 197 216 237 312 316 325 333 337 0-5 0.00172 .50172 210, 236 260 356 361 371 38i 385 Sat'd 0.00l8l5 0.645+5 332 364 395 519 526 539 55i 558 Sum of deviations of first four values of P X I0 6 from their mean 2 4" J J J j I I T T Sum of deviations of first J. 1 seven from their mean... 1 86 154 127 35 32 32 32 34 Mean of first seven . . 240 2 ; * 261 282 282 28* 287 288 K for s = 0.00098. 0.289 0.288 0.291 0.297 0.290 0.300 0.280 0.232 0.378 0.550 26 TABLE V (Continued). Part 2. Tenth-Normal Hydrochloric Acid. P X 10 6 for values of s at heads of columns. b. a. b +a. 0.0013. 0.0011. 0.00103. 0.00102. 0.00101. is. lor s = 0.00102. 0.0173 0.0173 256 262 264 265 266 0.260 O.OI 0.0131 0.0231 * 2 5 8 264 267 267 267 0.262 0.02 O.OIO2 o . 0302 257 264 267 268 268 0.263 0.05 o . 00593 0-05593 253 265 269 269 270 0.264 O. I 0.00363 0.10363 238 260 267 268 268 0.263 O.2 0.00237 0.20237 215 255 270 272 274 0.267 0-3 0.00190 0.30190 1 80 241 262 265 268 0.260 0.4 0.00172 0.40172 1 68 240 2 7 6 280 284 0.271 0.5 0.00170 0.50170 200 300 336 341 346 0.329 Sat'd 0.00181 0-633+5 323 450 494 500 506 0.490 Mean of first seven . . . . 237 2SQ 266 268 260 Sum of deviations of first seven from their mean 156 16 Part 3. Normal Hydrochloric Acid. P X 10 6 for values of s at heads of columns. b. a. b + a. 0.0012. 0.0010. 0.00088. < 0.00087. 0.00086. iv. tor s = 0.00087. O O.OI62 O.OI62 225 231 235 235 236 0.270 O.OI O.OIlS 0.0218 219 225 229 229 230 0.263 0.02 0.0091 0.0291 220 228 232 232 232 0.267 0.05 0.0052 0.0552 216 228 235 235 236 0.271 O.I 0.0031 0.1031 194 215 227 228 229 O.262 0.2 0.00202 O . 2O2O2 I6 5 206 230 232 234 0.267 0-3 0.00164 0.30164 134 193 229 232 235 0.267 0-4 O.OOI52 0.40152 128 208 256 260 264 0.299 Sat'd Mean of 0.00159 first seven.. 0-494+s 193 196 292 218 351 231 356 232 361 233 0.409 Sum of deviations of first seven from their mean 191 80 18 13 17 Attention is called to the peculiar fact that in ail three cases at the saturation point of hydroquinone, the apparent solubility of quinhydrone actually increases in- stead of decreasing with increasing excess of hydroquinone. The results are given in Table V. On comparing our results for water solutions with those of Luther and Leubner, we note the following: In the first place, our results are lower than theirs by about 5% of the ap- parent solubilities themselves. We do not know the reason for this. In the second place, the value of s, which gives the minimum sum of devia- tions of P from its mean, for the range comprising our first four solutions, which was as far as their investigation was carried, is 0.00125, which checks their value 0.0013 by less than 5%. But the minimum sum of our deviations for this range is only o . 000002 whereas theirs is o . 000022, eleven times as much. From this it appears that we have been able to obtain a much greater degree of precision. 27 If, however, we consider the entire range up to the saturation point of hydroquinone, we see in the column of s = 0.00125 a decided but con- tinuous decrease in P reaching a minimum when b = 0.3, followed by a continuous increase becoming suddenly abrupt at the saturation point. It was to verify these seemingly abnormal results that determinations were repeated which satisfied us that they were not due to experimental error. If, now, we try smaller values of s, we find that our first seven values of P become more and more uniform, givdng a minimum devia- tion sum of 0.00032 for values of s between o.ooioo and 0.00096. We therefore select the middle value 0.00098 as the most probable value of s in water, according to our experiments. This involves the assump- tion as a working hypothesis, in preference to the only other and much less probable alternative, viz., that they vary inversely to each other, thus giving a constant product, that both s and K remain approximately constant for values of b up to 0.3. Above this value, one or both must increase. A clue as to which one seems to be given to us by our potential measurements, as will be shown in the next section. Perfectly parallel results, it will be noticed, were obtained with tenth- normal and normal hydrochloric acid solutions. The solubilities and dissociation constants of quinhy drone are recapitulated in Table VI. The basis of the last column is to be found in the next section. TABLE VI. Moles per liter. Dissociation constant. Solvent. Water Apparent solubility. . . o 0178 True solubility. 0.00098 O.OOIO2 0.00087 Normal . (Av. 1st 7.) 0.289 0.263 o. 267 Saturated hydroquinone. 0.550 0.490 O. AOQ N/io HC1 O OI7"? #HC1.. O.OI62 III. THEORETICAL SIGNIFICANCE OF THE POTENTIALS. If the potentials observed measure mainly the tendency to take place of a reaction represented by the following equation: C 6 H 4 2 + 2H+ + 2 e ^ C 6 H 4 (OH) 2 (n) then, theoretically, we should find the following relation to hold ap- proximately RT r q X [H+] 2 " 0.0298 log q X h [H+ l 2 - _ log Kol volts for 25 C. (12) or, as this form of equation is usually written, 7T = 0.02 9 8 log q X [ H + ]2 + TTo = 0.0596 log [H + ] + 0.0298 log ; ( I3 ) n 28 in which TT is the observed potential for the concentrations q, [H+] and q X [H + l 2 h, K is the value of ~ ^ when TT = o, i. e., at "equilibrium," and q x [H + l 2 7T is the value of IT when ^ = i, for then the first term 0.0298 log 3 __^ - = o, and TT O = 0.0298 log K . But so far, we have data for h and q at our disposal only in the cases of mixtures F, G, H, I and J. In the last form of the above equation (Equation (13)), the first term TT O is, by definition, a constant, and, in the tenth-normal hydrochloric acid solutions, the second term 0.0596 log [H+] may be assumed to be nearly a constant. Therefore, the differences (next to the last column, TABUJ VII. Difference. b. q=a s. h = a+b s. q/h. 0.0298 log q/h. Calculated. Found. J o 0.01628 0.01628 i. ooo o j I o.oi 0.01208 0.02208 0.547 0.00781! TT O.OO74I O.OO65 H 0.02 0.00918 0.02918 0.315 0.01495} 0.0162 0.0155 G 0.05 0.00491 0.05491 0.0895 0.0312 j v* F o.i 0.00261 0.10261 0.0254 0.0475 ' The figures in this table are derived, as indicated, from those of Table V, Part 2, which see. The column "b" is common to both tables. Table VII) between the values of 0.0298 log q/h may be taken as ap- proximating the theoretical potential differences themselves. That is, for this series, theoretically, TTi 7T 2 = 0.0298 log (q/h)i 0.0298 log (q/h) 2 . (14) The observed potential differences are represented by the distances between the lines in Chart I. Each pair, consisting of a light and a heavy line, represents the duplicate cells of the mixture indicated by the letter at the right-hand end of the pair of lines. The heavy lines connect the points representing the readings of cell No. i and the light lines those of cell No. 2 of the mixtures in question. 1 The distances between the various pairs of lines, it will be seen, correspond quite distinctly to the theoretical differences 2 which are indicated by the distances between the horizontal dotted lines in the chart. The manner in which the absolute values repre- sented by these dotted lines were obtained, will appear a little later in this discussion. Before we can arrive at a definite single figure to be taken from our experimental data as the true representative experimental potential value 1 Except in the case of B in which they are No. 3 and No. 4, respectively. 2 Table VII, "Differences calculated," TTI TTZ, Eq. (14). CHART I -SERIES 3 .94 .92 20 40 60 80 100 120 160 ISO (for purposes of numerical comparison and calculation) for each mixture, several things must be taken into account. In the first place, we notice a well-defined and quite regular sloping off of the potentials, which be- comes more marked as the proportion of quinone increases, indicating one or more side reactions involving quinone, and associated very likely with the increasing brown color which solutions of quinone acquire on standing, and consistent with the general instability or reactivity of qui- none. But what we want is the potential of the mixture as we have made it up, before its original composition has been altered by any side reactions. The natural solution of the difficulty would be to take the average initial or first day's readings, but it will be noticed that the read- ings taken during the first twenty-four hours are usually quite erratic and non-reproducible, indicating the influence of accidental external factors before they have come to equilibrium with the principal system. To obtain the true initial potentials, therefore, we are compelled to re- sort to a sort of graphical extrapolation, by inspection instead of calcula- tion. Without laying down any hard and fast rule, we have selected as the true initial potential for each of the mixtures in question that value which, upon inspection, ignoring the irregularities of the first day or' two, seems to be most consistent with the contours of the curves for that mixture and with those of the neighboring curves. The values so selected are marked on the chart by the short lines extending to the left from the left 30 border of the diagram. The fairness of this method of approximation and its precision of half a millivolt, which is sufficient for the purpose, may be seen on inspection of the chart. The method is not as crude as it might seem, at first glance, and is the only one suitable for the purpose. The values, "TT," in Table IX 1 are these values, and the "Differences found," in the last column of Table VII, are the differences between them. It will be noticed that the observed differences ("found") do not differ greatly from the theoretical differences ("calculated"). Thus our experimental results are found to be in rough agreement with the van't Hoff equation, in the form TT = 7T + 0.0298 log (q/h) + 0.0596 log [H+] (13) in the case of the potentials of those mixtures for which we have the neces- sary concentration data. But in the case of the solutions saturated with hydroquinone (mixtures A, B and C), this concentration data is lacking, owing to the fact already pointed out that the constancy of both s and K, upon which the method of calculation depends, does not continue up to the saturation point of the solution with respect to hydroquinone, so that the method is inapplicable. But if we make the tentative assumption that the above expression (Eq. (13)) holds also in this case, we have a means of calculating the concentration ratio (q/h), and from it the con- centration of quinone (q), the solubility of quinhy drone (s), and the value of the dissociation constant (K). We are then in a position to apply the above equation to the study of the effect of varying acidity in the case of the stable potentials of the saturated hydroquinone solutions (Series I). Inspection of the chart brings out the fact that mixture F gives the most stable potential of the group already considered. So we select it as the basis for our calculations, taking for TTJ?, the initial reading (o . 8695) obtained as explained above, which is also the average of the readings up to the point at which the falling off commences. For TT B we have the average value 0.8300. We have, then, TT F TT B = 0.0298 log (q/h) F 0.0298 log (q/h) B (14) or, 0.8695 0.8300 = 0.0475* 0.0298 log (q/h) B . Solving (q/h) B = -log-{ - 039 o 5 + 9 8 0475 ) = o.oo I202 ; h B = o.633, 2 soq B = 0.001202 X 0.633 = 0.00076; 1 Page 33. * See Table VII. 2 See Table II. S B = a q = o.ooiSi 1 0.00076 = 0.00105; p B = q x h = 0.00076 X 0.633 = 0.000481; P 0.000481 K = - = - - = 0.457. s 0.00105 Now the difference between the value (0.00105) obtained for s, in saturated hydroquinone solution by this method, and the value (0.00102) obtained in the last section 2 for that range over which s and K appear to be approximately constant, is well within the limits of certainty of the method by which the value (0.00105) was obtained. This becomes e vide at when we repeat the above calculation, using each of the other mixtures of the series as the basis, in place of mixture F. Bans. SB. F 0.00105 G 0.00105 H . O.OOIOI I o . 00094 J o . 00092 The average of these values is only 0.00099, but the higher values are based on the mixtures giving the more stable and definite potentials and therefore carry more weight. All that we are really justified in say- ing, then, is that apparently the true solubility (s) of quinhydrone, re- mains about constant and it is only its dissociation "constant" which in- creases suddenly as the saturation point for hydroquinone is approached. For our purpose, therefore, we consider excess of hydroquinone to be without influence on the true solubility of quinhydrone. We have re- calculated our data for Mixture B on this basis (s equals 0.00102 instead of 0.00105). The results are recorded in Table VIII. We did not run a special series of solubility experiments for N/ioo HC1 because there is so little difference between A 7 /io HC1 and water that interpolation would show a negligible difference between water and N/ioo HC1, and we thought that data for water as solvent would be of more general value than data for N/ioo HC1. We also wanted to compare our results with those of Luther and Leubner, who worked only with solutions in pure water. Table VIII is based on the equation TA TB = (0.0298 log (q/h) A .+ 0.0596 log [H+] A ) - (0.0298 log (q/h) B + 0.0596 log [H+] B ). (15) TO cancels out. The values under "IT (observed)" are the observed aver- age values (taken from Chart II, in which all of the cells of the mixtures indicated are plotted from data of Table I, as in Chart I) of the poten- tials, during their period of stability. iSee Table V, Part 2. 2 See Table V. TABLE VIII. Mix- q = ture. a s. Calcu- lated. h. [H+]. (0.0298 log q/h) + (0.0596 log [H+]. A 0.00072 0.4940.790 (0.0845) + (-0.00609) =0.0906 B 0.00079 0.633 0.0921 ( 0.0867) + (-0.0618 )= 0.1485 C 0.000835 0.645 0.00971 ( o.o86i) + (-o.i20o )= 0.2061 J-576 0.0579 Ob- Ob- served, served. - 88 3} 0.053 0.830 0.770! 0-060 In like manner, TT A TT C = 0.1155 (calculated) and o. 113 (observed). Thus the observed potentials are found roughly to be in agreement with the theory in the case of varying hydrogen-ion concentration also CHART E- SERIES 1 8.2 50 300 350 400 450 The values under [H + ] were obtained as follows: Bray and Hunt, 1 by the conductivity method, found for the degree of ionization (a) of hydro- chloric acid at 25, 92.1% in tenth-normal solution, and 97.1% in hun- dredth-normal solution. But no direct data could be found in the litera- ture for normal hydrochloric acid at 25. Kohlrausch, however, 2 gives the following for the equivalent conductivity (A) at 18: Normality. A. 0.01 370 O.I 351 i.o 301 From these figures we find A.i/A.oi = 0.948 at 18 and Hunt's data we find that a.\/a.n = 0.948 at 25. 1 /. A. C. S., 33, 781 (1911)- 2 Landolt-Bornstein Tab., 2410., p. 1104. while from Bray Moreover, accord- 33 ing to Kohlrauscb, 1 the temperature coefficient of conductivity of HC1 solutions varies less and less with increasing concentration, approaching constancy as normal concentration is approached. Thus: Concentration o.ooi o.oi o. i o.s Coefficient 0.0163 0.0158 0.0153 0.0152 (The coefficient for normal hydrochloric acid is not given.) In the absence of definite data, therefore, we assume that (*1\ orf ) = ( } = 0.858 (from Kohlrausch's figures) \A.i/25 \.l/25 \A.i/i8 ai = a.i( ) =0.921X0.858 = 0.790. \A.i/i8 3o we take, as an approximation, the value 0.790 for the hydrogen ion [Concentration in our normal hydrochloric acid solution (A). The method used by the leading workers in this field in inorganic chem- istry 2 for comparing observed potentials with the theory, was to calcu- late, for each observed potential, by means of an equation similar in form to the one we have been using (Equation (12)), the value of the expression RT/nF log e K (designated by Peters and Friedenhagen as A and by Luther and in the present work as TT O ), the degree of constancy of these values measuring the agreement with the theory. Applying this method to our own results, we find the following : 7T = TT 0.0596 log [H+] 0.0298 log q/h (16) TABLE IX. F. G. H. I. J. B. C. A. -0.0596 log [H + ] r, . 0. . 0. 0618 8695 0. 0. 0618 8860 0.0618 0.9015 0.0618 0.9075 0, ,0618 9150 0.0618 0.8300 0.1200 0.7700 0.0061 0.8830 -0.0298 log q/h.. . 0. 0475 0. 0312 0.0149 0.0078 0, 0000 0.0867 0.0861 0.0845 ro . . . 0. 9788 ,9790 0.9782 0.9771 0, ,9768 0.9785 0.9761 0.9736 The deviations (Table IX) are of the respective values of ir from their mean, which is 0.9773. The average deviation, or numerical mean of these deviations is 0.0014 volt, or i .4 millivolts. This compares favor- ably with the closest agreement that has been obtained in parallel work in inorganic chemistry, as can be seen from the following summary: 1 Loc. cit., 244 b. p. 1115. 2 See articles by Peters, Friedenhagen and Luther, already cited, p. 10. 34 System. Ferrous /ferric Ferrous /ferric Ferrocyanide/ferricyanide Manganate/permanganate Uranous/uranyl Iodine /iodate B romine /br ornate Bromide/bromine Iodide/iodine Hy droquinone /quinone nvestigators. Peters (Tab. i) Peters (Tab. 2) Friedenhagen Friedenhagen Luther & Michie Luther & Sammet Luther & Sammet Luther & Sammet Luther & Sammet This work Average deviation. Millivolts. 39 2-7 i.o 4-o I .2 0.8 I.O 0.5 I .O 14 From 0.0298 log K c ,-33 9773> we can calculate the value of K c which is i .6 X 10 We may also use the mean value TT O = o. 9773 as a basis for calculating theoretical potentials of the mixtures under consideration as follows: TABLE X. Mixture. 7TO + .0.0298 log l(q/h) [n+] 2 1 = TT (calculated). TT (observed). A 0.9773 o . 0906 0.8867 0.8830 B 0-9773 0.1485 0.8288 o . 8300 c 0-9773 0.2061 0.7712 0.7700 F 0-9773 0.1093 0.8680 0.8695 G 0.9773 0.0930 0.8843 . 8860 H 0-9773 0.0767 0.9006 0.9015 I 0-9773 o . 0696 0.9077 0.9075 'J 0-9773 0.0618 0.9155 0.9150 The relation of the theoretically calculated potentials to the observed can best be seen by graphical means. They are represented on the charts by the straight horizontal dotted lines. In concluding this section, attention is called to the following points: It was not found possible to determine the concentrations directly or to compute them with mathematical certainty from the experimental data. Instead, the values used have been derived from the experimental data by a method based upon probability. Therefore, it must be admitted that the method by which they were obtained, alone, does not justify us in presenting them as anything more than probable values. But this method has been clearly and definitely recorded, so that the doubt exists only as to their significance. The fact has been established that the values so obtained, when taken as the actual concentrations, fit the elec- trochemical theory to the degree of precision shown. This coincidence is so remarkable that it greatly strengthens the probability that these values approximate the actual concentrations, or at least the effective concentrations. PART H. I. POTENTIALS OF SOLUTIONS SATURATED WITH QUINONE. Only two mixtures saturated with quinone were tried, viz., D and E, already described. In Table I the measurements on two cells of each of these mixtures are given. The two E cells are the only ones that were tried of this mixture. But, in the case of D, the two given are the best specimens out of five. The first three are given in Table XL TABLE XL Days. Cell No. 1. I . 0064 I 0.9903 2 o . 9860 3 0.9855 4 0.9835 5 0.9835 6 o . 9805 7 0.9780 8 o . 9740 9 0.9705 10 o . 9648 ii 0.9565 12 0.9487 13 o . 9390 14 0.9345 15 0.9285 16 0.9235 17 0.9200 18 0.9180 19 0.9156 20 0.9150 21 0.9138 22 0.9125 23 0.9130 24 0.9130 25 0.9125 26 0.9115 27 0.9110 28 0.9110 29 0.9104 30 0.9103 31 0.9100 32 o . 9094 33 o 9092 Cell No. 2. Cell No. 3. Days. Cell No. 1. (cont'd) 0.9901 o . 9842 34 0.9077 o - 9749 0.9761 35 0.9I50 2 o . 9668 o . 962 i 36 0.9597 0.9505 37 38 0.9389 0.9277 39 0.9230 o . 9076 40 0.9120 0.9147 0.8937 41 0.9107 0.8976 0.8661 42 0.8888 0.8310 43 0.8723 0.7885 44 o . 9082 45 0-8573 0-7559 46 o.9io6 3 0.8474 0.7423 47 o 9090 0.8322 0.7238 48 o . 9070 0.8316 0.7175 49 o 9049 0.8275 o . 662 i 50 0.8295 o . 6248 51 o . 9088 52 0.9081 0.8180 0.5903 53 0.9083 54 o . 9068 0-7935 0.5747 55 0.9063 56 o . 9062 0.7849 0.5719 57 0.7715 0.5724 58 59 0.9041 o 7605 0.5764 60 o . 9048 61 0.9041 0-7499 0.5736 0.7489 o.5728 1 -o.955 8 o - 7493 0-9375 0.7452 0.9227 0.7410 0.9108 o . 7409 1 At this point it was found that practically all of the undissolved solids had dis- appeared. So the cell was opened and more quinone and quinhydrone added and nitrogen passed for a few hours. The two readings on that day were taken before and after adding the solids. 2 On this morning it was found that the vibrator on the relay of the temperature 36 The first cell, it will be seen, was measured over a period of two months, during which time it fell steadily, except for occasional periods of apparent equilibrium lasting a few days. In all our cells the electrode was im- bedded in an excess of the solids with which the solution was saturated. In the case of this cell, it was noticed that the yellow quinone particles soon became dark green, and this change was naturally associated with the falling potential. It was thought that the dark product might be, partly at least, quinhy drone. To test this idea the second cell was made up without the addition of quinhydrone, the idea being that if the quinone produced its own quinhydrone the solution would eventually become saturated with this substance, in which event the potential, though originally higher, would, other conditions being equal, become about the same as that of cell No. i. Cell No. 3 was made up similarly to cell No. i, but with less excess solids, and started simultaneously with Cell No. 2. It will be noticed that the potential of No. 2 is from the start intermediate between that of No. i and No. 3. The dark solid was observed in Cell No. 2 on the second day, so its formation begins early. No definite con- clusions on this point could be drawn, owing to unknown conditions, however, and the matter was not considered of sufficient importance to the main investigation to be pursued further at this time. On comparing Cells No. i and No. 3 we see that evidently the excess of solids has a great deal to do with the stability of the potentials. This is strikingly borne out by the great increase in potential to nearly the original value, on adding fresh solids to No. 3 on the 2 9th day. The sudden increase was not due to exposure to the air on opening the cells, as it did not begin until after the solids were added. The fourth and fifth cells were there- fore made up with a much greater excess of solids, as previously men- tioned, expecially quinone, as were also the tenth-normal hydrochloric acid cells (Mixture E), giving the decidedly more stable potentials already recorded in Table I. By means of the equation q[H+ 2 l 7T = 0.0298 log ^-L- -" -f TTo n and the data of the preceding sections, we can calculate a theoretical potential for the mixture B. We have (assuming K and s to be unaffected by an excess of quinone to saturation (0.1275 molar) regulator had stuck during the night permitting the temperature of the thermostat to rise to 39 C. The bath was immediately cooled to 25 again, but this temporary heating seems to have had a permanent elevating effect on the potential. 3 Nitrogen was passed again for a few hours to see what effect the stirring and re- moval of gases might have on the potential. It seems to have had a slight elevating effect. 37 h X q = 0.000268 (Table V, Part 2); h = - - = 0.002102; [H+] = 0.0921 (Table VIII); TT O = 0.9773 (Tables IX and X); o. 1275 X (o.092i) 2 * = 0.9773 + 0.0298 log - VV = 0-968 7 . This value, it will be seen, is higher than even the initial values (0.9517 and 0.9524) observed for this mixture, and considerably higher than the apparent equilibrium (horizontal part of the potential-time curves) values. We have already seen, in Series III as a whole, that there is a distinct falling short of the theory, which becomes more and more marked as the proportion of quinone is increased. Furthermore, the increasing instability of the potentials parallels this deviation from the theory. These facts point to the conclusion that the falling off from the theory is connected in some way with the instability of quinone rather than to the failure of the theory. Considering the length of time required for the ap- parent equilibrium to be reached, in mixtures D and E, in connection with the instability of quinone, it seems most probable that the initial poten- tials are closer to the true potential of the system as it was made up origin- ally than the constant potentials reached later are. Yellow quinone particles could still be distinguished in the solid mass around the elec- trode when cells No. 4 and No. 5, of Mixture D, and the cells of Mixture E, were finally removed, and the odor of quinone was marked. But even though the solution remained saturated, which is not certain, products may have been formed which affected either the potential directly or the hydrogen-ion concentration. It is evident that something affected the potential. Hydrochloric acid solutions of quinone and also of quinhydrone soon acquire a deep rose color on standing, which is not extracted by ether. This color is noticeable even in Vioo normal hydrochloric acid. Accord- ing to Beilstein, dilute hydrochloric acid is without action on quinone. When solid quinone is treated with concentrated hydrochloric acid, it first becomes dark and then dissolves to a colorless solution. The final product is monochlorhydroquinone. 1 That this reaction is not confined to the concentrated conditions described in the literature was easily shown by a simple rough qualitative experiment. A series of portions of quinone were treated, respectively, with hydrochloric acid in concentrations varying from concentrated (12-13 normal) to normal. With the con- centrated acid the reaction took place immediately. The others followed in succession in the order of their concentrations. After twenty-four hours, the first four (including 7 . 5 normal) had become colorless ; the next one (6 normal) was well on the road to colorless; the others were successively 'Wohler, A., 51, 155. 38 darker, the last three (3.5, 2 and i normal) being just at the darkest intermediate stage. After two days more no further changes had taken place. The complete reaction was also obtained by adding concentrated hydrochloric acid to an aqueous solution of quinone. Now this reaction involves two changes: (i) reduction to hydroquinone, and (2) chlorina- tion of the hydroquinone, as represented in the two following equations : (1) C 6 H 4 2 + 2HC1 > C 6 H 4 (OH) 2 + C1 2 (2) C 6 H 4 (OH) 2 + C1 2 > C 6 H 3 (OH) 2 C1 + HC1 That the reaction does probably go through these two steps, one at a time, and that the dark intermediate product is probably mainly quin- hydrone, was demonstrated in the following way: Concentrated hydro- chloric acid was added to a saturated aqueous solution of quinone, drop by drop, until darkening just began. The darkening continued without further addition of acid, and the characteristic bronze-green product soon separated out. It was recrystallized from glacial acetic acid. The product had the appearance characteristic of quinhy drone. The air dried crystals gave a titration, by the Valeur method, corresponding to 92 . 8% quinhy drone. No further confirmation was made at this time. Here we have, evidently, a simple case of reduction by hydrochloric acid. The rose color, mentioned above, was also noticed as an intermediate stage in the reaction. The whole reaction should form an interesting sub- ject for a special investigation, expecially because of the fact that passing chlorine gas into a solution of hydroquinone oxidizes it through quin- hy drone completely to quinone, instead of chlorinating it; that is, it re- verses the first of the above reactions instead of causing the second to take place. From the above it is evident that quinone does react with hydrochloric acid, even in dilute solution, which means a lowering of hydrogen-ion concentration and corresponding lowering of potential. This may ac- count in part for the falling quinone potentials and perhaps also for their low initial values as compared to the theory. The falling off of the potential of Mixture A (normal hydrochloric acid saturated with hydroquinone and quinhydrone) from the theoretical may also, perhaps, be attributable to reaction between the hydrochloric acid and the quinone (from the quinhydrone) lowering the hydrogen-ion concentration. II. POTENTIALS IN NEUTRAL AND ALKALINE SOLUTIONS. As already mentioned, no definite results were obtained in neutral or alkaline solutions. The various cells that were tried will be taken up individually for record and to illustrate the general qualitative behavior. We will call this series, Series IV. All of these solutions were saturated with hydroquinone and quinhydrone. 39 No. i. (Tenth-normal Potassium Chloride.) Days. Volts. 1 Days. Volts. Days. Volts. Days. Volts. O 0-55II IO 0.7136 2O 0.7326 30 0.7538 I 0.5610 II 0.7154 21 0.7342 31 0.7570 2 o . 6640 12 0.7170 22 0.7369 32 3 0.6615 13 0.7183 23 0-7379 33 0-7547 4 0.6727 14 0.7214 24 0-7399 34 5 0.6828 15 0.7240 25 0-7435 35 0-7547 6 0.6932 16 0.7244 26 0.7438 36 0-7545 7 0.7070 17 0.7257 27 0.7461 37 0.7541 8 0.7088 18 0.7282 28 0.7501 38 0-7543 9 0.7114 19 0.7304 29 0.7523 On the 38th day, an apparent equilibrium having been reached, the cell was opened, while a strong stream of nitrogen was bubbled through it to keep out the air, and a few drops of a saturated solution of sodium hydroxide were added. The cell was closed again with a moderate stream of nitrogen left running. Eight minutes later the potential had fallen to 0.4194, and after i hour it was 0.4040. The nitrogen was allowed to run for several hours, and then stopped. The table is continued below. The first column gives days after start, and the second column, days after adding the alkali. In the third column are the potentials. i. n. in. 59 21 0.2829 60 22 61 23 0.2729 62 24 63 25 0.2682 64 26 0.2712. 65 27 0.2635 66 28 0.2607 67 29 0.2783 68 30 0.2667 69 31 0.2659 Here again a constant potential was finally reached. The contents of the cell were, at this point, titrated for alkali, and showed approxi- mately 0.034 normal. Owing to the very dark color of these solutions, it is impossible to obtain an accurate titration of them. It is necessary to use litmus paper, as indicator, and the value obtained is only a rough 1 All potential values given in this section are single potential differences for the solutions in question. They are derived from the observed electromotive forces of the combinations by adding the value (0.5265) of the calomel electrode when the solution forms the positive pole, or by subtracting the observed e. m. f. from 0.5625 when the solution forms the negative pole; for in the former case the solution has a higher and in the latter a lower oxidizing potential than that of the calomel electrode. In which of these two ways a given value was derived can be determined by noting whether it is larger or smaller than 0.5265. I. II. III. I. II. ill. 39 i 0.3623 49 ii 0.3243 40 2 0.3661 50 12 0.3174 4i 3 0.3624 5i 13 0.3106 42 4 0.3630 52 14 0.3079 43 5 0.3640 53 15 44 6 0-3545 54 16 0.2977 45 7 0.3446 55 17 46 8 56 18 0.2944 47 9 0-3345 57 19 48 10 0.3299 58 20 0.2807 40 approximation. Of course, even an accurate titration would give in this case no indication of the hydrogen-ion concentration, as the un- limited supply of hydroquinone present neutralizes most of the hydroxide ion. Relatively very large quantities of hydroquinone are required to saturate these alkaline solutions, where much alkali is present, owing to salt formation (acid ionization) and resinification. A considerable ex- cess of solid was always added. CELL No. 2. (Tenth-normal Potassium Chloride.) Days. Volts. Days. Volts. Days. Volts. Days. Volts. Days. Volts. o . 5409 II 0.6993 22 0.7117 33 0.7420 44 I 0.5571 12 0.7016 23 34 0-7439 45 0-7431 2 0.6287 13 0.7027 24 0.7199 35 o - 7449 46 0-7435 3 0.6576 14 o . 7024 25 36 0-7474 47 0.7444 4 o . 6707 15 0.7049 26 37 5 0.6786 16 0.7071 27 38 0-7450 6 o . 6830 17 28 39 0.7441 7 o . 6869 18 o . 7084 29 0.7305 40 0-7445 8 0.6915 19 0.7097 30 0-7335 4i 0.7425 9 0.6954 20 0.7102 31 0.7387 42 0.7424 10 0.6972 21 0.7104 32 43 0-7433 Here again a constant potential was reached which was only about ten millivolts lower than that reached by Cell No. i. On the forty- seventh day a few drops of saturated sodium hydroxide solution we^e added, as before. Ten minutes later the potential was 0.7424, after l /% hour 0.7362, after 3 hours 0.4409, and after 6 hours 0.4141. The table is continued below i. ii. in. 62 15 0.4731 63 16 0.4763 64 17 o . 4808 65 18 0.4841 66 19 0.4832 67 20 0.4864 68 21 0.4768 The contents of the cell titrated 0.03 normal alkali. In composi- tion, therefore, this cell is almost a duplicate of No. i. Their potentials do not differ greatly during the neutral stage, but ifter the addition of alkali they do not check at all. The erratic nature of the alkaline potentials is further evidenced by the fact that, while those of Cell No. i decrease steadily to the constant potential, those of Cell No. 2 first de- crease and then increase. To see whether the potassium chloride exerted any influence upon the potential, saturated potassium chloride solution was tried instead of the tenth -normal. I. II. ill. I. II. III. 48 i o . 4040 55 8 49 2 0.3898 56 9 o . 4405 50 3 0.3892 57 10 0.4465 5i 4 58 ii 52 5 0.4117 59 12 0-4594 53 6 60 13 54 7 0.4273 61 H 0.4692 Ceu, No. 3. (Saturated Potassium Chloride.) Days. Volts. 0.4884 1 0.4880 2 0.4890 3 0.4900 This potential being almost constant from the start, though very different from the tenth-normal KC1 potentials, three drops of saturated sodium hydroxide solution were added on the third day. Twenty-five minutes later the potential was 0.3623 volt. The table is continued below. CEU< No. 3 (Continued). I. II. III. I. II. III. I. II. III. 4 i 0.3346 ii 8 0.2380 18 15 0.2257 5 2 0.2691 12 9 0.2359 19 16 O.2244 6 3 0.2559 13 10 0-2354 20 17 0.2240 7 4 0.2415 14 ii 0.2327 21 18 0.2209 8 5 0.2486 15 12 O.2294 22 19 0.2209 9 6 0.2417 16 13 0.2285 23 20 0.2212 10 7 0.2396 17 14 O.226O 24 21 0.2213 The. con tents titrated 0.018 normal alkali. This potential, it will be noticed, fell steadily to an apparent equilibrium, which was, however, lower than either No. i or No. 2, though the cell contained less alkali. The neutral potential was also lower. CELL No. 4. (Saturated Potassium Chloride.) Days. Volts. Days. Volts. Days. Volts. Days. Volts. Days. Volts. O 0-4375 H o . 5058 28 42 0.5634 56 o 5857 I o . 4466 15 0.5098 29 43 0.5661 57 o 5790 2 o . 4626 16 0.5157 30 5265 44 0.5654 58 3 0.4738 17 31 o 5533 45 o . 5690 59 o 5913 4 0.4776 18 0.5237 32 46 60 5 0.4874 19 0.5261 33 5525 47 0.5717 61 o 5921 6 0.4911 20 0.5292 34 48 0-5734 62 o 5919 7 o . 4934 21 35 o 5587 49 0-5733 63 5926 8 22 0.5321 36 5601 50 0.5757 64 o 5899 9 23 0-5347 37 o 5601 5i 0.5765 65 5827 10 0.4965 24 0-5349 38 o. 5601 52 66 0. 5897 ii 0.5005 25 0.5522 39 53 0.5801 12 0.5024 26 40 o. 5623 54 0.5802 13 0.5043 27 4i 0. 5626 55 No alkali was added to this cell. It is evident that neutral potentials for this system cannot be determined in potassium chloride solutions. The following alkaline cells were made up without potassium chloride, by adding hydroquinone and quinhy drone to sodium hydroxide solutions 42 of various concentrations, in the electrode vessel, while a rapid stream of nitrogen was bubbled through the solution. The cells *were then closed and the stream of nitrogen continued through the night to facilitate the dissolving of the solids. The inlet and outlet tubes were then closed, as with the other cells, shutting the cell off from the air, in an atmosphere of nitrogen. Nevertheless the solutions became dark immediately upon adding the hydroquinone. CELL No. 5. Days. Volts. Days . Volts. Days . Volts Days. Volts. Days. Volts. O 0.2292 4 O. 1495 8 O. 1420 12 0. 1375 16 0.1335 I 0.2315 5 0. 1475 9 O. 1406 13 0. 1342 17 0.1324 2 0.1607 6 0. 1446 10 0. 1405 14 0. 1297 3 0.1536 7 0. 1441 ii O . 1444 15 O. 1329 This cell titrated 0.4 normal when removed, on the iyth day. CELL No 6. (Titrated i .6 normal, when removed.) Days . Volts. Days. Volts. Days. Volts. Days. Volts. Days. Volts. O o. 1098 6 0.0711 12 O . 0498 i8 2 0.0806 23 O.IO2I I o. 1464 7 0.0673 13 0.0489 19 0.0740 24 O.IOI7 2 0-I449 1 8 o . 062 i 14 0.0475 - 3 O.I5I7 25 0-.0996 3 o . 0980 9 0.0612 15 0.0472 2O O.IoSl 26 0.0971 4 0.0850 10 0.0576 16 0.0466 21 O.IO39 5 0.0782 II 0.0564 17 0.0463 22 O.I04I The following might be suggested as an explanation of the changes in potential following the stopping and resumption of the flow of nitrogen in the above cell. Un- der the catalytic effect of the platinum, hydrogen is produced at the electrode, as fol- lows: C 6 H 4 (OH) 2 C 6 H 4 2 + 2 H+ + 20 2H+ H 2 + 20 C 6 H 4 (OH) 2 ^ C 6 H 4 2 + H 2 This reaction reaches equilibrium when the concentration of hydrogen produced is such that the two systems represented by the two partial or electrochemical reactions have the same potential. While the nitrogen is bubbling through the solution, at the electrode, the hydrogen cannot accumulate. We have then a resultant potential, in- termediate between the respective potentials of the two systems. When the stream of nitrogen is stopped, however, the hydrogen accumulates, and the resultant potential falls towards that of the hydroquinone system. On resuming the flow of nitrogen, the hydrogen is again driven out and the resultant potential rises again. 1 The passage of nitrogen was stopped after taking this reading. 2 The stop-cock in the bridge tube, connecting the vessel with the intermediate KC1 solution, slipped, admitting air to the bridge tube. This was forced out again by means of nitrogen, without air having been admitted to the cell, but the cell contents were disturbed thereby. 3 Passage of nitrogen for three hours was followed by a further rise in potential. 43 CELL No. 7. (Hydroquinone and Quinone Added to o . i N NaOH to Saturation.) Hours. Volts. Hours. Volts. Hours. Volts. Hours. Volts. 19 0.3250 117 0.2736 212 0.2513 307 0.2473 44 0.3201 141 0.2624 236 0.2556 336 0.2576 72 0.3I2I 170 0.2601 26l o . 2460 365 0.2544 93 0.2914 190 0.2581 284 0.2493 38l 0.2498 CELL No. 8. (Duplicate of Cell No. 7-) o 0.3280 6 7 0.2236 162 0.2533 260 0.2504 8 0.2935 98 0.2338 187 0.2544 283 0.2535 20 0.2420 127 0.2363 212 0.2523 44 0.2428 141 0.2396 241 0.2526 It will be noticed that the initial potentials of the above duplicates check quite well, but that after the first day there is no reproducibility. In all probability, therefore, only the initial potential can be taken as representing the simple original known system as it was made up, and this, probably only approximately; and no definite significance, relative to the original system is to be attached to any apparent equilibrium reached later. The later potentials are evidently more or less accidental. The following three cells were made up similarly to the above, with o.oi N sodium hydroxide solution: Cell No. 9. Cell No. 10. Cell No. 11. Cell No. 11 (Cont'd). Days. Volts. Days. Volts. Days. Volts. Days. Volts. 0.4027 O o . 4064 O 0.4065 14 0.5346 I 0.4417 I 0.4396 I 0-4454 15 0.5410 2 0.5464 2 o . 4806 2 0-4939 16 0.5452 3 0.4498 3 0.4726 3 0.4693 17 4 0.4583 4 0.4724 4 0.4883 18 0.5559 5 0.4811 5 0.5H3 19 0-5595 6 0.5199 20 0.5621 7 0.5223 21 0-5579 8 0.5143 22 0.5675 9 0.5164 23 0.5732 10 0.5166 24 0.5606 ii 0.5176 25 0.5598 12 0.5265 26 o . 5504 -' 13 0.5265 27 0.58II It would be interesting to compare these initial potentials wich the van't Hoff equation. But to do this the concentrations must be known, and the determination of these in alkaline solutions is a very different proposition from what it is in acid solutions, owing to the acid properties of hydfoquinone, and perhaps, to some extent, of quinone, and their in- stability toward alkalies. To determine this data would require a special and difficult investigation, beyond the scope of this undertaking. 44 Qualitatively it is shown that the oxidizing potentials decrease, that is, the reducing potentials increase with alkalinity, as would be expected from the theory. The results in neutral and alkaline solutions are given largely as a matter of record. Hardly more has been done than to give an idea of some of the difficulties confronting an attempt to study the system quan- titatively under these conditions. The circumstances here are proba- bly similar to those encountered in the case of the solutions containing quinone in excess (namely, disturbing reactions), but to a more pro- nounced degree; and the main disturbing reactions are probably different. One of the difficulties is the determination of hydrogen ion concentra- tion. In the case of the acid solutions conductivity data were used for this purpose. This involved the tentative assumption that the condition of the acid, or the hydrogen ion concentration, was the same as in a pure solution of hydrochloric acid. We cannot make a similar assumption regarding the sodium hydroxide because of the action of the hydroquinone as an acid, toward alkali. We therefore cannot use the conductivity method in this case. The question naturally arises, why not determine the hydrogen ion concentration by the e. m. f. method with a hydroge electrode. This method was not used at all in this work (although it was realized that the assumption upon which the use of the conductivity values was based might be the source of considerable error), because it was believed that, even if a constant potential could be obtained with a hydrogen electrode, in the presence of another active electrochemical system, this potential still might be very different from the true hydrogen- hydrogen-ion potential because of the influence of the other system. It was felt that this was too big a question to be taken up merely as an inci- dental to this investigation. PART III. DETAILS OF EXPERIMENTAL METHODS. All volumetric apparatus used was Eimer and Amend normal apparatus. Each piece was calibrated and found to be accurate to within the limits of precision possible in its use. Standard hydrochloric acid solutions were made up from Baker's Analyzed hydrochloric acid and standardized at 25 against Baker's Analyzed sodium carbonate, dried to constant weight. The manner in which the solutions were made up for potential measurement is illus- trated by the case of Mixture A. Enough powdered hydroquinone and quinhydrone to just saturate 100 cc. of normal hydrochloric acid was weighed out and placed in a 100 cc. volumetric flask. 50 cc. of 2 normal acid were then added, from a pipette, at 25 C., and the flask gradually diluted to the mark, keeping the contents at 25, with shaking, and later whirling, in the constant temperature bath, so that when the mark was reached practically all of the solids had dissolved, and the solution was just normal to hydrochloric acid and saturated with the solids at 25. The hydroquinone used was partly Merck's and partly Eimer and Amend's. Only one grade was sold and it was claimed to be very pure because of its method of manufacture. No accurate method of deter- mining its purity could be found. It all melted sharply at 169 (uncor- rected). This is the melting point given in Beilstein. 1 The corrected melting point was 173.0. The quinone was Kahlbaum's. Some of it was recrystallized from gasoline (which was found to be an excellent solvent for separating qui- none from its impurities), and some was sublimed, both giving clean, bright yellow products, titrating the same (99.4% of the theory) and melting sharply at 115.7 (corrected) (the melting point given by Beil- stein) . The quinhydrone was made in two ways; some by treating hydro- quinone with ferric chloride, acidified with hydrochloric acid, in aqueous solution, and some by mixing together equivalent quantities of hydro- quinone and quinone, in aqueous solution. The product, in each case, was filtered off and washed with water. The product of the first method titrated 99.2% and that of the second 99.4% of the theoretical, the latter checking the quinone. Some was recrystallized from alcohol v followed by ether, and some from glacial acetic acid. Glacial acetic acid is an excellent solvent for the purpose, being by far the best of the three. All of the recrystallized products titrated 99 . 4% of the theoretical. Only the products giving this titration were used in the experiments. Quin- hydrone decomposes on heating, before reaching its melting point. 1 Hlasiwetz, A., 177, 336. 4 6 Nothing new or unusual, in the way of apparatus or method, was in- volved in the measurement of the potentials. The mixtures whose poten- tials were to be measured were placed in Kales calomel cells. 1 The cell and bridge tube were filled with the solution, leaving about an inch and a half of free space from the top of the cell. Sufficient mixed excess solids were added to loosely fill the two bulbs at the base of the cell, so that the electrode would be immersed in a bed of solids, in order to keep the solu- tion saturated in the region of the electrode. A two-hole rubber stopper, containing the electrode and a drawn out tube for the introduction of nitrogen, both reaching nearly to the bottom of the cell, was inserted. A tapering gas outlet tube, fitted with a pinch-cock, was inserted in the side tube of the cell. Bight cells or less, at a time, one of which was the stand- ard cell, were grouped in a circle around a beaker containing saturated potassium chloride solution, their bridge tubes dipping into the solution. To minimize flow and diffusion the cock in the bridge tube was kept closed with a band of vaseline around the upper and lower thirds and the middle third free for the solution to pass around. The ends of the bridge tubes, dipping into the salt solution, were plugged with wooden plugs wrapped with filter paper. The cluster of cells was immersed in the constant temperature bath. The nitrogen passed from the tank through an alkaline pyrogallol solution, then successively through water, cotton, a large bottle of water, immersed in the constant temperature bath, acting as a humidifier (to prevent evaporation in the cell) and as a manifold, from which the gas was distributed to the various cells. The closing of the bridge tubes, mentioned above, did not increase the resis- tance enough to interfere materially with the sensitiveness of the measure- ment. The standard cell used was the saturated potassium chloride calomel cell, which was found to be the most satisfactory, having by far the great- est constancy and reproducibility. It was compared with a similar cell belonging to Dr. Fales, which had been checked against a number of other standard cells. The comparison extended over a period of one week. The potential differences between the two cells (our cell minus Dr. Fales') are given below : Hours. Volts. Hours. Volts. Hours. Volts. Hours. Volts. o 0.00042 27 0.00033 72 O.OOO22 I2 4 0.00034 I 0.00042 31 0.00030 74 O.OOO25 125 0.00028 2 0.00042 35 0.00030 75 0.00023 127 O.OOO3O 6 0.00042 49 O.OOO3O 78 0.00023 131 O.OOOOS 24 0.00034 54 O.OOO25 103 0.00029 144 0.00034 25 0.00034 57 O.OOO29 I2O 0.00037 147 0.00031 26 0.00033 59 0.00025 122 O.OOO3O 1 Fales and Vosburgh, J. A. C. S., 40, 1305 (1918). 47 Two months later a new standard cell was made up and compared with the old one, which had been in constant use in the meanwhile. The re- sults were as follows (new cell minus old cell) : Days. Volts. Days. Volts. Days. Volts. Days. Volts. 0.00040 2 0.00031 5 o.ooooo 8 0.00016 '/ O.OOOI6 3 O.OOO2O 6 0.00006 9 O.OOOI6 I 0.00017 4 0.00016 7 0.00018 10 O.OOOII The measurements were made with a Leeds and Northrup potentiom- eter, a sensitive galvanometer, and a certified Weston Standard Cell. A normal and a Beckman thermometer were used in the constant temperature bath. They were standardized against a Baudin Normal thermometer (Bur. Int. des Poids et Mesures No. 18,537). The bath was electrically regulated, and, as previously stated, the temperature variation was less than a hundredth of a degree. Details of Sohibility Determinations. The methods have already been outlined. Details of manipulation, etc., are given below. Hydroquinone. In the case of water, as solvent, an excess of hydro- quinone was shaken repeatedly with water, several minutes at a time, during a period of several hours, in a bottle in the constant temperature bath, and allowed to stand in the bath over night. The acid solutions were made up as follows, so as to be at the desired normality after satura- tion with hydroquinone : About 5 grams of hydroquinone were placed in a 100 cc. volumetric flask, 50 cc. of acid of twice the desired normality were added and thoroughly shaken, and then enough water was added with shaking to dissolve all of the hydroquinone. Small portions of powdered hydroquinone were then added, alternately with the water, all at 25, so that the solution was kept saturated with hydroquinone as it was diluted to the mark. It was then transferred to a bottle, with more hydroquinone, shaken intermittently, as in the previous case, and allowed to stand over night, or long enough to leave a clear supernatant solution. Five cc. of the solution were then pipetted off into a weighed 10 cc. distilling flask, fitted with a small rubber stopper and fine drawn out tube, extending to the bottom of the flask, as is customary in vacuum evaporations, to aid ebullition. It was then connected with a vacuum pump and evaporated to constant weight, at room temperature or with gentle warming (not over 35), at a pressure of about 8 mm. Heating much over 40 at this pressure may cause appreciable evaporation of the hydroquinone. The fact that concordant results were obtained from different solutions, shaken different lengths of time, was taken to indi- cate that the solutions were saturated. 4 8 No suitable analytical method for the determination of hydroquinone could be found, but it is possible that the reduction of silver ion in neu- tral nitrate solution might be developed into such a method when the proper conditions are worked out. In ammoniacal solution, quinone will also reduce silver. Quinone and quinhydrone. The way in which the solutions were saturated and sampled has already been completely described under the section on Solubilities, etc. Only details of the analytical method remain to be given. Solutions. Tenth-normal potassium dichromate made from pure, dried potassium dichromate, to be used as the standard. 10% colorless potas- sium iodide solution. Tenth, hundredth and five -hundredth normal sodium thio sulphate solutions standardized against the dichromate solution. The end point is very delicate. Method of Standardization or Analysis. To the sample of solution to be titrated, a number of cc. of 10% KI solution, approximately equal to the number of cc. of Af/io thiosulphate that would be required, were added. Half this number of cc. of concentrated hydrochloric acid were then added and the solution titrated immediately, using starch as indicator, if de- sired. Usually the end point is sufficiently delicate without it. In some cases the exact starch end point is rendered uncertain by the pinkish color probably caused by the action of the hydrochloric acid on the quinone. Except for the last mentioned difficulty, which only causes trouble in the case of the very dilute solutions (quinhydrone in HC1 solutions with considerable excess hydroquinone) and is not very serious then, the method seems to be a very accurate one. Repeated determinations on purified quinone and quinhydrone, varying the concentrations somewhat, with- out going to extremes, always gave the same result, viz., 99.3-99.5% of the theoretical. The remaining 0.6% may be moisture or a constant error in the method. No statement or data as to whether the method is supposed to give theoretical results or not could be found in the litera- ture on the subject. CONCLUSION. It has been found that whenever the concentrations could be con- trolled and side reactions minimized (namely, in the acid solutions with small proportions of quinone) fairly constant and reproducible potentials were obtained with solutions containing hydroquinone and quinone, in equilibrium with their addition compound, quinhydrone. Very likely, with elaborate and specially designed apparatus and carefully studied precautions, still more definite results could have been obtained. The present work being more or less of a pioneer nature, great precision, in this respect, was not attempted, the necessary conditions not having yet been worked out. 49 These potentials were found to vary with the probable approximate concentrations of the substances involved in the electrochemical equa- tion, in fair accordance with the van't Hoff constant temperature equa- tion. In alkaline solutions no quantitative study was possible, owing to in- terfering reactions, but, qualitatively, the oxidizing potential was found to be much lower than in acid solutions, and to fall with increasing alka- linity in accord with the theory. Potentials in neutral solutions were in- termediate between those in acid and in alkaline solutions, but were in- definite and greatly affected by the electrolyte. Although no case could be found better adapted to such an investiga- tion the conditions are nevertheless such that the concentrations of the substances involved, including hydrogen ion, could only be approximated. To this end, the solubilities of hydroquinone, quinone and quinhydrone, and the probable approximate dissociation constants of quinhydrone, in water and in dilute hydrochloric acid solutions, have been determined. Part of this has merely been a more thorough repetition and extension of work already done by others. The new fact, however, was brought out, that the solubility product law, when applied to saturated quinhydrone solutions, in the presence of an excess of hydroquinone, fails to hold for concentrations of hydroquinone approaching saturation, owing evidently to an increase in the dissociation constant, which becomes so abrupt at the saturation point that the law is actually reversed, that is, that addi- tion of hydroquinone causes an increase, instead of a decrease, in the quantity of quinhydrone which will go into solution. This cannot be attributed to error in the determination of the solubility of hydroquinone, since the change, though sudden, is a smoothly accelerated change with a gradual beginning. The abruptness of the change, however, suggests a peculiar condition of what Washburn calls "the thermodynamic environ- ment," 1 connected with the fact of saturation. It was also found that hydrochloric acid decreased the solubility of hydroquinone but increased that of quinone. The confirmation of the van't Hoff equation in a case of this kind per- forms two distinct functions: (i) It verifies the Law of Mass Action for the electrochemical equation, thus furnishing evidence supporting the use of the electrochemical equation. (2) It affords a means of evaluating the intrinsic oxidizing tendency (represented by the constant term of the van't Hoff equation), which may be called, for convenience, "the normal potential," (i. e., the potential for unit concentrations or concentration ratio of 1) and which forms a basis for comparing the system with other systems, independently of concentrations, and for calculating the potential for any given concentrations. 1 Principles of Physical Chemistry. 50 Since the concentrations were only more or less approximated, and since the holding of the ideal gas laws in the case of osmotic pressure, upon which the van't Hoff equation is based, is also only an approximation, very close agreement with the theory could not be expected. As it was, the agree- ment was of the same order as that obtained for parallel work in inorganic chemistry. Thus the quantitative relationships, upon which the electrochemical theory of oxidation is based, have been extended experimentally into the field of organic chemistry, and in this way the chemistry of organic non-electrolytes, has, in the only instance so far investigated, been corre- lated with that of inorganic electrolytes. RETURN TO the circulation desk of any University of California Library or to the NORTHERN REGIONAL LIBRARY FACILITY Bldg. 400, Richmond Field Station University of California Richmond, CA 94804-4698 ALL BOOKS MAY BE RECALLED AFTER 7 DAYS 2-month loans may be renewed by calling (510)642-6753 1-year loans may be recharged by bringing books to NRLF Renewals and recharges may be made 4 days prior to due date. DUE AS STAMPED BELOW SEMtONILL - JUL 6 U. C. BCRKCLEY 12,000(11/95) 35324 fi-7 UNIVERSITY OF CALIFORNIA LIBRARY