ELECTROLYTIC FORMATION OF PERCHLORATE BY
E. I,. MACK
The usual conception of the mechanism involved in the anodic formation of perchlorates from chlorates is largely due to Oechslil and his theory has gradually found its way into the text-books of electrochemistry with the result that it is a t the present time commonly accepted as representing the actual mechanism of the process. According to this theory the reaction at the anode during chlorate electrolysis is not to be considered a direct addition of oxygen to C103’ with the formation of Clod’ but rather as a more complicated process, during which we must suppose the liberation of the chlorate ion. Briefly, the various steps involved are as follows : Chlorate ion is liberated by the current and, assuming that it acts in a manner analogous to other strongly acid anions, reacts with water a t the anode with the formation of free chloric acid and liberation of oxygen,
+ H 2 0 = 2HC103 + 0.
&lo3
(1)
The free chloric acid generated in this way is probably at a very high concentration in the film of electrolyte which is in immediate contact with the anode. Since in this condition it is known to be very instable, spontaneous decomposition takes place, perchloric and chlorous acids being formed, 2HC103 = HClOi
+ HClOz.
(2)
The free chlorous acid thus formed very evidently cannot remain in the solution as such, since in contact with the chloric acid present it would immediately evolve chlorine dioxide : HC103
+ HClOz = HzO + ClOz
(3)
This evolution of chlorine dioxide has never been observed to take place during this process, and the theory, therefore, assumes that the chlorous acid formed as shown in ( 2 ) is immediately oxidized to chloric acid by the oxygen from the Zeit. Elektrochemie, 9, 807 (1903).
Electrolytic Formation o j Perchlorate
239
decomposition of water by the liberated C103’ ions, as shown in ( I ) . The sum total of the reactions which, according to this theory, takes place at the anode may then be represented by: 2C103’ HzO 2F = HC104 HClOa Oechsli, in order to satisfactorily explain the formation of perchlorate, finds it necessary to assume the liberation of the chlorate ion mainly for the following reasons: Oxygen when discharged under conditions such that the electrode is reversible probably requires a lower voltage than that needed for the liberation of oxy-acid anions,l but if we deal with an electrode which is not reversible toward oxygen, that is, where the discharge of oxygen requires a considerable overvoltage, a point may be reached a t which the discharge of the oxy-acid anion will become easier, or require a lower voltage, than that for oxygen and will consequently take place. From this standpoint, assuming that C103’ is discharged, the formation of perchlorate could be expected to take place only at an anode where the overvoltage for oxygen liberation is sufficient to allow the preferential discharge Qf the chlorate ion. It is held that this hypothesis is supported by the following facts: (I) Perchlorate formation goes on with a markedly higher efficiency at a smooth platinum anode, where oxygen overvoltage is known to be high, than at an anode of platinized platinum, where oxygen discharge takes place much more easily. ( 2 ) If the solution is made alkaline a t the anode the efficiency of perchlorate production falls off rapidly; a fact which is explained by the increased ease of oxygen discharge in a solution containing a high concentration of hydroxyl ions. The result is a discharge of oxygen along with the chlorate ions and a consequent drop in the anode efficiency. (3) This hypothesis best explains the rise from a comparatively low efficiency a t the beginning of the electrolysis to a higher value as the electrolysis proceeds. When the oxidation is taking place a t a high efficiency and the current is interrupted for a
+
+
Bose: Zeit. Elektrochemie, 5, 169 (1898).
+
E . L.Mack
240
short time, it is found, upon closing the circuit, that the efficiency is low just as it was during the first few minutes of the experiment. The idea is that during the break in the current flow the concentration of free chloric acid a t the anode falls, largely through diffusion, and that a certain length of time is necessary for the chloric acid to accumulate to that‘ concentration in which it is unstable. During this time oxygen is evolved and the current efficiency necessarily falls. (4) Changing from a smooth platinum anode to one of platinized platinum brings about a marked decrease in efficiency and this is best explained by considering it equivalent to a decrease in current density because of the larger surface presented by the platinized electrode. At this lower current density the efficiency must be lower because we would have more oxygen liberated in proportion to the number of chlorate ions than at the higher current density. Furthermore, many other electro-chemical reactions which are commonly considered as being due to anion liberation show this same characteristic of decreased efficiency with platinized anodes. The formation of persulphuric acid and Kolbe’sl synthesis of ethane are cited as being processes of this type. ( 5 ) An increase of temperature, a t a fixed current density, brings about a decrease in perchlorate formation and a corresponding increase in oxygen discharge. This is considered to be the result of increase in hydroxyl ion concentration caused by the increased dissociation of water a t the higher temperature. (6) Finally, perchlorate formation must be conditioned by chlorate ion discharge since the reaction has not been duplicated chemically; that is, chlorate has not been oxidized t o perchlorate by chemical oxidizing agents. This theory is open to the objection that it rests upon several assumptions which may or may not be true. We have no direct evidence that chloric acid, when sufficiently concentrated, decomposes with the formation of perchloric and chlorous acids. The visible products of the decomposition are known to be perchloric acid, chlorine dioxide, chlorine 1
Liebig’s Ann., 69,2 5 7 (1849).
Electrolytic Forvtzatiovt of Perchlorate
241
and oxygen. While the intermediate formation of chlorous acid. is possible it has not, up to this time, been demonstrated satisfactorily. If we are to reason from analogy to the reaction which takes place when alkali chlorates are heatedl it would appear much more probable that the primary decomposition products would be perchloric and hydrochloric acids. Granting, however, that the liberated chlorate ion and free chloric acids would act in the manner assumed, the theory is still unsatisfactory in that it does not explain the decrease in yield of perchlorate on substitution of platinized for smooth platinum anodes or the marked fall in efficiency with rising temperature. In the former case, using a platinum anode a current efficiency of 97.0 percent was noted,2 while with a platinized anode, the other conditions remaining unchanged, the efficiency dropped t o 2 . 7 percent. It is very evident that it would be necessary to assume a tremendous decrease in current density to account for the marked efficiency drop a t the platinized anode since a succeeding experiment shows that, under like conditions, using a smooth platinum anode and a current density approximately one-fourth that in the first experiment cited, the efficiency fell only to 92.0 percent. Regarding the temperature effect it may be stated with certainty that the increase in the dissociation of water with a temperature rise of 73O is insufficient to account for a decrease in efficiency of 97 percent. The efficiency a t 7' is 9 7 percent while no perchlorate is formed a t 80" C, other conditions remaining unchanged. It has seemed advisable, therefore, to postulate that the electrolytic formation of perchlorate does not depend upon the liberation of the chlorate ion but rather is to be considered as a direct addition of oxygen to the chlorate ion. The object of this paper is to show that the phenomena observed during perchlorate formation are most easily and satisfactorily accounted for from this viewpoint. It is further Cf. Scobai: Zeit. phys. Chem., 44, 3 1 9 (1903).
E. L. M a c k
2 42
proposed to furnish independent proof that chlorate may be oxidized to perchlorate by strictly chemical means; that is, under conditions where the liberation of chlorate ion cannot be assumed. Let us consider for a moment the conditions existing a t an insoluble anode during electrolysis of solutions containing oxygen or hydroxyl ions. Recent investigations1 have shown that t h e action of the electrolyzing current is to liberate oxygen in the active or atomic form. This oxygen unless depolarized by some substance present a t the anode will be converted into molecular or gaseous oxygen, the reaction being represented by : 201 = 0 2
Further, it has been pointed out that the rate of this reaction, and consequently the concentration of active oxygen or 01, which is present a t the anode a t any given time, is dependent upon, (I) the nature of the anode material, (2) the current density, (3) temperature, (4) nature of the solution, and ( 5 ) the elapsed time of electrolysis. Schoch2 has pointed out that only an exceedingly small voltage is required to discharge oxygen or any other ion into an electrode which is absolutely free from electromotively active material, and it follows from this and the above statement that the potential of the anode, a t any given time, in a solution capable of discharging oxygen, depends primarily upon the concentration of active oxygen then existent a t the electrode. Thus we see that by properly varying the conditions we may have oxygen discharged a t any potential above a certain minimum value which is determined by the equilibrium value of the equation 201 = 0 2
a t the particular anode and under the particular conditions in question. It may be well to point out here that in order to oxidize a given substance present in the electrolyte, chlorate for 1 2
Bennett and Thompson. Trans. Am. Electrochem. SOC., 29, 15 (1916). Jour Phys. Chem., 14, 665 (1910).
Electrolytic Forvnation o j Perchlorate
243
instance, it is necessary that the potential of the anode should rise not to that point necessary to bring about the liberation of gaseous oxygen, as is commonly believed, but only to that point necessary to produce a concentration o j active oxygen suficient to oxidize the chlorate ion. That this is the case and that the potential necessary for the latter process is very much less than that required for the former is conclusively shown by the work of Sch0ch.l Using a N / 3 potassium chlorate solution and an iron anode i t was found that the formation of perchlorate began when the anode potential reached + 0 . 0 2 3 volt,* while a succeeding experiment showed that at an iron 1.5 anode oxygen is not evolved until an anode potential of volts is reached. This experiment, even without the support of further evidence, shows that the formation of perchlorate is not determined by the liberation of chlorate ion since no one would claim for this ion a discharge potential as low as $0.023 volt. Most of the attempts a t measurement have obtained values around I .37 volts.3 Having thus shown that perchlorate can be produced at an anode potential far below that required for the liberation of oxygen as well as much below that commonly assigned for C103’ discharge, let us consider the phenomena actually observed during the electrolysis of a chlorate solution, assuming the formation of perchlorate to be the result of a direct addition of oxygen to the chlorate ion. Using a smooth platinum anode in a strong neutral chlorate solution with a current density of, say, 4 to 6 amperes per square decimeter, we find the efficiency of perchlorate formation to be high, possibly as high as 95 percent. The anode potential is also found to be high and from the standpoint taken we consider that we now have a t the anode active oxygen
+
+
Jour. Phys. Chem., 14, 735 (1910). All potentials given in this paper are the observed readings against the normal calomel electrode. The signs are those observed in the experimental arrangement. This is in accord with the suggestion of Luther. See Le Blanc: Textbook of Electrochemistry, 4th Ed., 24j. Le Blanc: Zeit. phys. Chem., 8, 299 (1891).
E . L. ,!$Tack
244
in such concentration that it is able to oxidize the chlorate ion rapidly and with good efficiency. We know that we are here dealing with an anode which exhibits a high overvoltage. This means that the concentration of active oxygen is above the equilibrium value. If we decrease ,the current density below the value first taken, we find a decrease in current efficiency with respect to perchlorate formation and a corresponding decrease in anode potential. This becomes intelligible when we consider that a t moderate current densities oxygen overvoltage, and consequently concentration of active oxygen, decreases1 with decrease of current density. Again, if the temperature be raised we find the current efficiency decreases markedly. This, again, is in accord with the view taken since we have seen that the concentration of active oxygen present a t the anode a t any given time is dependent upon the state of the reaction 201 = 0 2
Increased temperature increases the velocity of this reaction and tends to keep the concentration of active oxygen down to the equilibrium value. We would expect, therefore, a lower active oxygen concentration and consequently a lower oxidizing power. A sufficiently high current density a t the higher temperature would, however, tend to increase the rate of production of active oxygen and increase its concentration in the electrode. As the current density is increased we should, therefore, expect to obtain a consequent rise in efficiency of chlorate oxidation tending to overcome the inhibitory effect of the high temperature. This is actually realized experimentally.2 The current density used was 16 amperes per square decimeter and with the solution maintained at 80' gave a current efficiency of 13.5-14.8 percent. This was raised to 40.0-42.5 percent when the current density was increased to 20.8 amperes per square decimeter, the temperature remaining constant a t 80'. 1 2
Bennett and Thompson: Trans. Am. Electrochem. SOC.,29, 15 (1916). Oechsli: Zeit. Electrochemie, 9, 817 (1903).
Electrolytic Formatiow OJ Perchlorate
245
If the solution is made alkaline the current efficiency falls off rapidly, a fact which is explained without difficulty when we recall that the discharge potential for oxygen a t platinum is lower1 in alkaline than in acid solutions. In other words, the active oxygen concentration is lower. Furthermore, oxidizing power in general is lower in alkaline than in neutral or acid solutions. This is to be concluded from the fact that many oxidations which take place readily in acid solution are either entirely prevented or proceed a t a very much decreased rate when the solution is made alkaline. It is noted that on starting the electrolysis both the potential of the anode and the efficiency are somewhat below their normal values. Both, however, steadily increase to a maximum with increasing time. This is entirely to be expected from the fact that reliable measurements of anodic overvoltage2 always indicate a low concentration of active oxygen during the first few moments of oxygen discharge. Finally, on substituting an anode of platinized platinum for one of polished platinum, other conditions remaining constant, we find a notable decrease in anodic efficiency. This, once more, is satisfactorily explained from our viewpoint, since the work of Coehn and Osaka3 interpreted from the viewpoint which we have taken shows that a t a given current density we must have a lower concentration of active oxygen at a platinized anode than at one of smooth platinum. Their results4 gave for the potential of oxygen at smooth platinum f 1 . 3 9 volts, and a t platinized platinum f 1 . 1 9 volts. From what has been said above, it becomes evident that we can satisfactorily account for the anodic formation of perchlorate without assuming anything a t all about the liberation of the chlorate ion, by considering that the oxidation is Smale: Zeit. phys. Chem., 14,j 7 7 (1894). Foerster and Piguet: Zeit. Elektrochemie, IO, 714 (1904). Zeit. anorg. Chem., 34, 86 (1903). The measurements given by them were made against a normal hydrogen electrode but for the sake of consistency in this paper these have been converted to the normal calomel electrode values.
246
E . L. Mack
dependent only upon the presence at the anode of a sufficient concentration of active oxygen. It has been satisfactorily shown that this concentration corresponds to an anodic potential far below that required for liberation of gaseous oxygen and nearly as far below that required for the liberation of the chlorate ion, if we are to accept for this the values given by the best measurements available. Although it may be readily admitted that the potential of the anode during the electrolysis of a chlorate solution is a t all times probably above that necessary for chlorate discharge, it must be kept in mind that these reactions will not take place as long as we have chlorate present in sufficient concentration. This follows since it has been shown that the direct oxidation of the C103’ is the more easily accomplished process, that is, the one requiring the lowest potential and as such will take place first. This does not mean that the oxidation will take place under all conditions with high efficiency since we have seen that those factors which unfavorably affect this reaction are exactly those which tend to lower the concentration of active oxygen at the electrode. The direct result of this decrease is a decrease in its ability to oxidize chlorate and a corresponding increase in the ease of oxygen evolution and the consequent drop in current efficiency with respect to perchlorate formation. Experimental Although the theory outlined above is well supported by known facts, it has seemed desirable to have independent proof of the fact that chlorates may be directly oxidized to perchlorates by chemical means and under conditions such that the liberation of the C103’ ion cannot be assumed. We have evidence that this is true in the fact that chlorates when heated to moderate temperatures produce perchlorate, one portion of the chlorate furnishing oxygen for the oxidation of another portion; a reaction which may be represented by: 4NaC103 = NaCl
+ 3NaC104
The possibility of a direct oxidation was even more
Electrolytic Forvnation o j Perchlorate
247
strikingly shown by Fowler and Grant.l They found on heating chlorate with silver oxide that the chlorate was completely converted to perchlorate without the loss of oxygen, metallic silver being the other product. It was desired, however, to show that the oxidation could be brought about in an aqueous solutibn and for this purpose several oxidizing agents were selected as the most useful and their,action on sodium chlorate investigated as described below. Method of Analysis The method used for determining the amount of the various chlorine acids in the presence of each other was essentially that described by Treadwell-Hall2 and is briefly as follows : In a solution containing chlorides, chlorates and perchlorates, or the corresponding free acids, the amount of chlorine as chloride was determined directly by titration of one sample ( I ) . A separate portion ( 2 ) titrated, after reduction with ferrous sulphate with an excess of dilute sulphuric acid, gave the amount of chlorine as chlorate plus that as chloride. A third portion ( 3 ) of the solution after evaporation to dryness with sodium carbonate was fused in a tall platinum crucible, the mouth being closed by a plug of loosely packed asbestos fiber. This sample showed the total amount of chlorine present in all forms. The amount of chloride present in the original solution is, then, shown by ( I ) , that as chlorate by the difference between that found in ( 2 ) and ( I ) , and that present as perchlorate by the difference between the amount shown by ( 3 ) and ( 2 ) . The amount of chlorine in these samples was in all cases determined3 by addition of an excess of tenth normal silver nitrate solution, filtering off the precipitated silver chloride and titrating the excess silver nitrate with a tenth normal solution of ammonium sulphocyanate, using ammonium ferric Jour. Chem. SOC.,57, 2 7 2 (1890). Analytical Chemistry, p. 463 (1911). Sutton: Volumetric Analysis, 9th Ed., p.
172.
2 48
E . L.i i a c k
alum as indicator. This method was found to be rapid and gave consistent and very accurate results. Duplicate samples containing 0.3 gram chlorine frequently gave results differing by only 0.3 milligram. In addition to the quantitative determination all solutions were examined for the presence of perchlorates by the microchemical test described by van Breukeleveen. Experiments with Persulphate The first oxidizing agent studied was sodium persulphate. This is commonly considered as a very powerful oxidizing agent and it seemed probable that it might convert chlorate to perchlorate even in aqueous solution. This was established as true by the following experiments. Preliminary Experimewt.-A solution containing about 5 percent of sodium chlorate was boiled for I O minutes with an excess of sodium persulphate. After evaporation to dryness concentrated hydrochloric acid was added and the solution again evaporated, on a water bath, to convert all remaining chlorate to chloride. After dissolving the residue in water, silver nitrate was added until all chloride present was removed. After filtering the solution was evaporated, the residue fused with sodium carbonate, dissolved in dilute nitric acid and silver nitrate added. A heavy precipitate of silver chloride indicated that a portion of the chlorate originally present had been converted to perchlorate. Experimenls 8-12 .-One gram portions of sodium chlorate, representing 0,3324 gram chlorine as chlorate, were boiled with solutions of sodium persulphate, the conditions being as shown below. In all cases a considerable oxidation t o perchlorate was found. Tests of the several solutions for chlorides, after boiling with persulphate, showed the absence of chlorine in this form. We have to deal with chlorine as chlorate and perchlorate only. The latter was determined by difference in the first experiments. Rec. Trav. chim. Pay-Bas., 17, 94 (1898).
Electrolytic Formation o j Perchlorate I1
Chlorine (as chlorate) taken 0.3324 Water, cc. 1 IO0 Sodium persulphate (grams) IO Length of time heated (minutes) 30 Chlorine (as chlorate) found ' 0.2710 Percentage oxidized to I perchlorate 1847
IO0 IO
0.2705 18.32
i
IO0
25
25
IO
IO
10
30
30
0.2736
0.2730
0.275 I
17.69
17.87
17.24
Chlorine (as chlorate) taken Chlorine (as chlorate) found after treatment Oxidized to perchlorate
1
I2
0.3324 0.3324
~
1
249
Price: Per-A'cids and Their Salts, p, 3 j (1912).
0.3324 gram 0,2565 gram 0.0759 gram
250
E. L.M a c k
Chlorine (as chlorate) taken Chlorine (as chlorate) found after treatment Oxidized to perchlorate Oxidized to perchlorate
0.3324 gram gram gram 33.75 percent 0.2202 0.1122
Electrolytic Forniatiox oJ Perchlorate
Chlorine (as chlorate) taken Chlorine (as chlorate) found after treatment Total chlorine found after treatment
2.51
3324gram o 2886 gram 0.3312 gram
0
25 2
E . L. AWack
Chlorine (as chlorate) found after treatment Oxidized to perchlorate
0 . 3 2 6 2 gram
o
0062
gram
Electrolytic Formation o j Perchlorate
253
It seemed probable that a gas with higher ozone concentration might be more effective in obtaining the desired oxidation. An apparatus was, therefore, constructed which was capable of producing oxygen containing a very high percentage of ozone. The ozone was produced by electrolysis of sulphuric acid solutions at low temperatures. The apparatus employed was essentially a modification of that described by Fischer and Massenay.l KO attempt was made to control the concentration of ozone produced by this apparatus but it continuously furnished a gas of very much higher ozone content than that from the Siemens tube. One analysis showed a concentration of 21 percent ozone by volume. Exflerimelzts 38 and 39.-Solutions containing one gram sodium chlorate and 2 cc sulphuric acid (sp. gr. 1.82) in 25 cc water were used and were maintained at 100' by means of a water bath. Ozonized oxygen produced by the apparatus described was allowed to bubble thl;ough these solutions for two hours. As in the previous cases a slight oxidation of the chlorate was obtained. ANALYSIS 39
Chlorine (as chlorate) found after treatment 0.3300 gm Oxidized to perchlorate 0.0024 gm Oxidized t o perchlorate 0.7170
0.3324 gm 0.3296 gm o.0028 gm 0 .s 2 7 G
The experiments with ozone show, thereiore, that while ozone is capable of oxidizing acid solutions of chlorates, the oxidation is not at all efficient. In all cases described the amount of ozone used was in excess of the theoretical amount necessary to oxidize all chlorate present. It may be pointed out that while the amount oxidized is small it is nevertheless certain that perchlorate was produced since the Zeit. anorg. Chem., 52,
202,
z z g (1907).
254
E . L.Mack
microchemical test showed, unquestionably, the presence of perchlorate in all solutions after treatment with ozone. Since acid solutions of permanganates are known to be strong oxidizing agents, this was the next material studied. Experiment 45.-A solution containing one gram sodium chlorate, I cc sulphuric acid (sp. gr. 1 . 8 2 ) and one gram potassium permanganate in IOO cc water was boiled for 30 minutes. The solution showed no perchlorate present. ANALYSIS
Chlorine (as chlorate) taken Chlorine (as chlorate) found after treatment
0.3324 gram 0,3326 gram
This experiment shows that acid solutions of potassium permanganate will not oxidize chlorate to perchlorate. The next oxidizing agent studied was aqueous sodium peroxide. Experiment 46 .-A solution containing one gram sodium chlorate and I O grams sodium peroxide in 2 5 cc was boiled for 15 minutes. After cooling and diluting to IOO cc the solution was acidified with dilute nitric acid. ANALYSIS
Chlorine (as chlorate) taken Chlorine (as chlorate) found after treatment
0,3324 gram 0.3317 gram
The slight discrepancy noted here between the amount of chlorate taken and that found after treatment is probably to be regarded as experimental error since the microchemical test used failed to show the presence of perchlorate. The amount represented by the difference in the two chlorine determinations is probably within the experimental error in this case since the presence of large amounts of sodium salts in the solution made the analysis somewhat less accurate than in previous cases. Experiments with Hydrogen Peroxide Hydrogen peroxide is probably to be considered as one of the strongest of oxidizing agents since its decomposition
Electrolytic Formation of Perchlorate
Chlorine (as chlorate) taken Chlorine (as chlorate) found after treatment
255
0.3324 gram ‘0.3327 gram
E . L.Mack
256
It is undoubtedly true that oxidizing agents, in general, are more powerful in the presence of acid than in neutral or alkaline solutions. Therefore, it was decided to investigate the action of hydrogen peroxide on acid solutions of chlorates. Some work has been done along this line by Tanatarl who investigated the action of hydrogen peroxide on halogen oxyacids. He found that hydrogen peroxide had no effect on either neutral or acid solutions of chlorates, while bromates were quickly reduced to bromide with evolution of some bromine. These conclusions are certainly in error in so far as they state that acid chlorate solutions are unaffected by hydrogen peroxide. This will be seen from the following experiments. Experiment @.-One gram sodium chlorate was dissolved in 25 cc of 30 percent hydrogen peroxide which had previously been acidified by I cc sulphuric acid (sp. gr. 1 . 8 2 ) . The ' solution was boiled for one hour. Soon after the solution had reached the boiling point a yellow gas was evolved. This was a t first thought to be chlorine but more careful examination showed it to be a mixture of chlorine dioxide and chlorine. The microchemical test showed the presence of perchlorate in the solution and analysis gave the following results: ANALYSIS
Chlorine (as chlorate) taken Chlorine (as choride) found after treatment Chlorine (as chlorate) found after treatment Total chlorine found after treatment
0.3324 gram 0.2916 gram 0.0030 gram 0.2927 gram
This experiment shows that chlorate, through the action of acid solutions of hydrogen peroxide, is largely converted to chloride. A considerable amount of chlorine and chlorine dioxide is evolved at the same time. Experiment 49.-The conditions in this experiment were the same as in. the previous case with the exception that 3 percent hydrogen peroxide was used. The solution was boiled for 1 2 hours. 1
Ber. deutsch. chem. Ges., 32, I013 (1899).
Electrolytic Formation of Perchlorate
Chlorine (as chlorate) taken Chlorine (as chloride) found after treatment Chlorine as chlorate, found after treatment
1
This will be discussed more fully in a later paper.
257
0.3324 gram 0 . 0 1 2 2 gram 0.3162 gram
258
E . L. Mack
which was unstable under ordinary conditions. Since the acid next higher than chloric acid is perchloric acid and this was shown to be perfectly stable under the conditions existing in the experiments it was necessary to assume the formation of an acid still higher than perchloric. The further assumption necessary was that this hypothetical higher acid would decompose spontaneously with the liberation of oxygen and formation of hydrochloric acid. Assuming for the moment the existence of this compound it seemed highly probable that during its formation from chloric acid we would pass through the perchlorate stage or, in other words, that it should be formed equally well from perchloric acid. A preliminary experiment showed, however, that this theory was untenable. A 5 percent solution of potassium perchlorate acidified with sulphuric acid was added to an equal volume of 30 percent hydrogen peroxide and the whole boiled for I O minutes. No chlorides were present in the solution after this treatment. This experiment is supported by the work of Tanatarl who found that perchloric acid is not affected by hydrogen peroxide. It became necessary, therefore, to abandon the position first taken. The possibility next considered was the existence of a molecular compound between chloric acid and hydrogen peroxide. No assumptions were made as to the nature of this hypothetical compound but for the time being it was considered analogous to the possible compounds between ferrous sulphate and hydrogen peroxide described by Mummery. The hypothesis which we are now considering postulated, of course, that this compound HC103.xH202, if existent, would decompose with formation of hydrochloric acid. If this compound were actually formed on mixing solutions of chloric acid and hydrogen peroxide it seemed probable that its presence would be revealed by a measurable thermal change. Accordingly the heat of dilution of solutions of Ber. deutsch. chem. Ges., 32, 1013 (1899). *Jour. SOC.Chem. Ind., 32, 889 (1913).
Electrolytic Formatiom of Perchlorate
259
chloric acid with hydrogen peroxide was determined as follows : Experiment jo.-The apparatus used consisted of a silvered Dewar flask of about zoo cc capacity, fitted with a stopper carrying a small funnel tube and a 100"thermometer, graduated in 0.1'. The stopper was also cut away at one side to permit the introduction of a glass stirring rod. Twenty cc of a 5 percent chloric acid solution was introduced into the flask and the temperature allowed to become constant. 5 cc 30 percent hydrogen peroxide solution were then brought to the same temperature and introduced into the flask. The temperature of the mixture rose rapidly and reached a value where it remained constant for several minutes. The thermal effect on diluting 20 cc of 5 percent chloric acid with 5 cc of water was next determined, the procedure being exactly similar to that used in the previous case. The data obtained were as follows: Initial temperature of chloric acid 24.70; Initial temperature of hydrogen peroxide 24. 70' Maximum temperature of mixture 24.95 Rise in temperature of mixture 0.25' Initial temperature of chloric acid 24.55: Initial temperature of water 24.55 Maximum temperature of mixture 24.80' Rise in temperature of mixture 0.25' These experiments are sufficient to indicate that there is little possibility of the formation, at room temperature, of a compound between chloric acid and hydrogen peroxide. One possibility remained to be considered, however, before abandoning this hypothesis. It has been mentioned that the apparent reducing action of hydrogen peroxide on chloric acid does not become evident until a temperature of about 80" is reached. The possibility existed, therefore, that the intermediate compound just considered might be formed at a point somewhat above room temperature. If this were true the formation of the compound would be indicated by a break in the heating curve of a mixture of chloric acid and hydrogen peroxide. An investigation of the heating curve of a 5 percent chloric acid solution mixed with an equal volume
E . L.Mack
26 0
of 30 percent hydrogen peroxide failed to show any break which would indicate the formation of a compound. The conclusion was, therefore, reached that the apparent reduction of chlorate to chloride by hydrogen peroxide could not be assigned to the formation of an instable compound between the two. The most satisfactory explanation of this phenomenon, however, was found in the fact that in a dilute solution of chloric acid we have a small amount of chlorine liberated through spontaneous decomposition of the acid. This chlorine, by its action on the hydrogen peroxide, immediately forms hydrochloric acid. The hydrochloric acid thus produced would, of course, immediately attack the remaining chlorate, the products of this reaction being hydrochloric acid, chlorine, chlorine dioxide and oxygen. The reaction is, therefore, auto-catalytic, and a small amount of hydrochloric acid is sufficient to start the decomposition. It has been mentioned that the formation of chlorides in a solution of chloric acid and hydrogen peroxide, does not begin to be appreciable until the temperature is raised to about 80'. This is in accord with the work of Sand2 who investigated the reaction between chloric acid and hydrochloric acid. He found that the velocity of the reaction' became appreciable a t temperatures above 70'. This explanation of the apparent reducing action of hydrogen peroxide on chloric acid has the advantage that it makes no assumptions whatever but uses only well-known facts to account for the phenomenon. There is scarcely room for doubt that the explanation put forward represents the true mechanism of the reaction. It also shows that hydrogen peroxide does not here act as a reducing agent, in the true sense of the term. Experiments with Activated Oxygen The fact that oxygen is activated or ionized by ultra1 2
Fairley. Jour. Chem. Soc., 31, 2 2 (1877). Zeit. phys. Chem., 50, 465 (1904).
Electrolytic Formation of Perchlorate
261.
violet light rays is well estab1ished.l In this condition it is a very powerful oxidizing agent, probably stronger than either ozone or hydrogen peroxide. The action of activated oxygen on chloric acid was, therefore, investigated. It is frequently stated that solutions of chloric acid are slowly oxidized on standing in sunlight. While this is probably true, the necessity for the presence of oxygen has not been recognized. A solution was made up containing 0.4784 gram sodium chlorate and 0.25 cc sulphuric acid (sp. gr. 1.820) in 45 cc. This represented 0.1582 gram chlorine as chlorate. The solution was placed in a transparent quartz flask and a neutral atmosphere maintained by continuously passing in nitrogen which had been freed from traces of oxygen by washing with alkaline pyrogallol solution. Ultraviolet light was obtained by means of a Cooper Hewitt quartz tube mercury arc lamp, the flask containing the solution being suspended about 5 cm from the lamp. The experiment continued for 28 hours. No oxidation of the chlorate took place, as is shown by the analysis. ANALYSIS Chlorine (as chlorate) taken Chlorine (as chlorate) found after experiment
0.1582 gram 0.1580 gram
This experiment shows conclusively that, in the absence of oxygen, dilute chloric acid solutions are not photochemically oxidized to any appreciable extent. The next experiment was similar to the preceding one, except that oxygen was continuously bubbled through the solution. The flask held 50 cc of solution containing 1.000 gram of sodium chlorate and 0.50 cc sulphuric acid (sp. gr. I ,820). This represents 0.3305 gram chlorine as chlorate. The solution was subjected to the action of the ultraviolet rays for 8 hours. After the treatment, the microchemical test showed perchlorate to be present. This was confirmed by analysis. Sheppard: Photo-chemistry, p. 253 (1914).
E . L.1Vack
262
ANALYSIS
Chlorine (as chlorate) taken Chlorine (as chlorate) found after experiment Oxidized to perchlorate Oxidized to perchlorate ' .
0,3305 gram 0.3269 gram 0.0036 gram I .og percent
Time elapsed (hours)
Sodium chlorate found per I O cc solution
0
0.0072
24
0.1158 0.1160 0.1157
72
96
It will be noted that during the first 24 hours a small amount of the chlorate was oxidized. This represents about 1.1 percent of the total amount present. During the remaining time no further oxidation took place. It will also be noted that these results are practically identical with those obtained when ozonized oxygen was used as the oxidizing agent, as well as with those obtained by Oechslil with ozone. It has proved impossible to account satisfactorily for the fact that the oxidation of the chlorate stops when approximately one percent has been converted to perchlorate. The possibility of perchloric acid being instable under ultraviolet rays was considered and investigated. If this were true, after exposing Zeit. Elektrochemie, 9 , 807 (1903).
Electrolytic FormatFon of Perchlorate
2 63
the acid to ultraviolet light we should expect to find some of the chlorine present as one of the lower oxy-acids, presumably as chlorate. A solution containing about 2 percent perchloric acid was placed in a quartz test-tube and oxygen slowly bubbled through the liquid. The tube was exposed to the rays from the mercury arc lamp for several hours. Examination of the solution showed that all of the chlorine remained in the form of perchlorate. The hypothesis mentioned is, therefore, untenable. The possibility remains that active oxygen is less stable in the presence of perchloric acid than under ordinary conditions. This is equivalent to saying that perchloric acid catalyzes the reaction 2 0 1 = O2 and hence tends to keep the concentration of active oxygen down. We know that under the conditions maintained -the concentration of this substance is very low, at best, and the presence or formation in the solution of any substance which tended to convert it to the molecular form might force its concentration below the point necessary for chlorate oxidation. There seems to be no satisfactory way of testing this hypothesis, however, and it must stand simply as a suggestion which, if true, might account satisfactorily for the facts observed. This phase of the problem should be further investigated with a view to determining what stops the formation of perchlorate when a certain amount has been produced and whether the position of the equilibrium can be shifted by varying the concentration of active oxygen used. It seems probable that if the amount of active oxygen present or rather the rate of its production could be increased a further oxidation of chlorate would result. In the experiments described in this paper, with the exception of those in which persulphate was used, the concentration of active oxygen attainable was necessarily low and not comparable, in degree, to that present at the anode during electrolysis. However, it is held that the conditions maintained duplicate qualitatively those a t the anode during chlorate electrolysis and show that perchlorate can be formed by the direct addition of oxygen to the chlorate ion.
E . L.Mack
264
From what has been said above i t may be concluded that: I . Chlorate may be oxidized to perchlorate by persulphuric acid, ozone, and hydrogen peroxide in acid solutions. 2 . In the case of hydrogen peroxide the hydrochloric acid formed interferes with the reaction. 3. This oxidation may also be carried on with oxygen activated by ultraviolet light. 4. The reaction is considered to take place according to: HC108
+
0 1=
HC104,
the quantity of oxidation depending on the concentration of active oxygen present. 5 . Since active oxygen chemically produced oxidizes chloric acid to perchloric acid and since active oxygen is formed a t the anode it follows that the electrochemical formation of perchlorate is the result of direct oxidation. 6. Perchlorate is formed at the anode at a potential far below that necessary for the continuous discharge of any ion present in the solution. 7. The conditions realized in the experiments described above duplicate, qualitatively, those existing a t the anode during electrolysis. Cornell University
