T H E OXIDATION POTEIC’TIALS O F SOME HYPOCHLORITE SOLUTIOi$S* BY V. H. RESIINGTOS AND H. M. TRIYBLE
Berthollet,‘ as early as 1 ~ 8 jshowed , that solutions of chlorine in caustic potash do not weaken linen fabrics as do its aqueous solutions. Crossley? found that the hydrogen ion concentration of a solution has an important effect upon the oxidizing properties of hypochlorites toward cellulose, noncellulose impurities in fibers of vegetable origin, and certain dyes, Weissenbach and R;lestrezat3found that the effectiveness of Dakin-Daufresne hypochlorite solution in killing certain bacteria was much increased by rendering it slightly acid. Dunstan and his coworkers4 have found that the secret of success in the use of hypochlorites in refining certain petroleum products lies in controlling the alkalinity of the hypochlorite solutions within rather narrow limits. Schwalbe and Wenzlj found that, in the course of bleaching with alkaline hypochlorites, the bicarbonates which form retard the reaction, and may even bring it to a standstill. It seems clear from these few examples that hypochlorites possess different oxidizing powers in different media; but our knowledge of their relative activities under various conditions is distinctly limited. An oxidizing materid is, in general, characterized by its power to yield oxygen, either directly or indirectly. This power may not extend t o visible evolution of gas but, in any case, the oxidizing power of an oxidizing agent is greater the greater its tendency to yield oxygen. According to a theory which was first proposed by Nernst,6 an electrode which is not attacked, when immersed in a solution of an oxidizing agent, adsorbs oxygen from it and acquires a charge. The potential so set up, then, is directly proportional to the osmotic pressure of oxygen in the solution, and so also t o the oxidizing power of the solution. This theory accounts satisfactorily for the setting up of a potential under these circumstances. It seemed t o us that it should be possible to get a measure of the oxidizing powers of solutions of hypochlorites by determining their oxidation potentials. A survey of the literature shows that but little work has been done upon the oxidation potentials of hypochlorite solutions, and that this work has given only disconnected and fragmentary data. * Contribution from the Chemistry Department of the Oklahoma Agricultural and Mechanical College. 1 hlellor: “Treatise on Inorganic and Theoretical Chemistry,” 2, 243 (1922). * A . IT.Crossley: Pharm. J. 115, 693-5 (192j). 3 R. J. Weissenbach and 3Iestrezat: Compt. rend. SOC. Biol.. 81, 93-6 (1918). A . E. Dunstan: J. Inst. Petroleum Tech., 10, 201-1 j (1924). 5 C. G. Schwalbe and H. M-enzl: Z. angew. Chem., 36, 302-4 (1923). 6 W ,Sernst: “Theoretical Chemistry,” Fourth Ed., p, 731 (1903).
OXIDATION POTENTIALS OF HYPOCHLORITE SOLUTIONS
42 5
Apparatus and Reagents The apparatus used in measuring the potentials consisted of an ordinary student potentiometer with the usual accessories. h tenth normal calomel electrode prepared according to directions given by Clark’ was used as reference electrode. Electrodes of smooth platinum were used in this work. Measurements were standardized by means of a certified standard Weston cell. All pipettes, volumetric flasks and burettes used in the research were calibrated and the necessary corrections were applied. h special type of electrode vessel, shown in Fig. I , was developed for this work. Zonite, a commercial preparation of sodium hypochlorite, was used in some of the earlier experiments. A11 other chemicals were of the best C. P. grade. The conductivity water used in making up all solutions was prepared by redistilling a good quality of distilled water from alkaline permanganate. The hypochlorous acid which was used in most of the work was prepared in the following manner. Chlorine from a cylinder was passed into a suspension of mercuric oxide, yielding a mixture of salts of mercury and hypochlorous acid. This solution was diluted so as to contain about one percent of hypochlorous acid, and then distilled under the reduced pressure produced by a filter pump. A slow current of air was drawn through the solution during the distillation. To secure a product yet more nearly free from hydrochloric acid the solation was distilled a second time in the same manner. The final product was then analyzed by adding a known quantity of the solution to an acidified potassium iodide solution, and titrating the iodine which was set free with standardized sodium thiosulphate solution. The solutions were then diluted to approximately the concentration desired for our experiments with conductivity water and kept in brown glass bottles in a cool place until needed. Hypochlorous acid thus prepared contained only a very small trace of chloride. No solution more than two days old was ever used. The hypochlorous acid solubions were always analyzed just before they were used in experiments, and adjusted to the exact concentration which was wanted. The Experiments .Ilarge part of this investigation was of necessity given over to a study of the use of platinum electrodes in determining the oxidation potentials of hypochlorites in solution since it seems that this problem has never been satisfactorily solved. We have been unable to get any consistent results with platinized electrodes. We believe that they take up, either in preparation or while being used, impurities which act as “poisons,” and which cannot be removed by any method of treatment which wc have tried. It was found that when a n untreated electrode of smooth platinum, is put in contact with a solution of hypochlorite the potential usually changes regularly but it shows no indication of ever reaching any equilibrium value. Results bvith a given electrode are rarely or never twice the same. The potentials tend to fall lower and lower in successive experiments. Different electrodes seem t o I%-.11.Clark: “The Determination of Hydrogen Ions,” 2nd E d , Chap S V I I (1925’
426
V. H. REI\IIXGTOS A S D H. M , TRIMBLE
behave quite differently. The electrodes were treated a t different time> with chromic acid, concentrated and fuming nitric acid, solut,ions of caustic alkalies and alkaline solutions of potassium permanganate, followed by washing with conductivity water, without improving their performance. Treatments with alkaline solutions follo,wed by concentrated nitric acid gave no better results. Evidently these treatments failed to clean the electrodes. If we accept Sernst’s theory as to the setting up of an oxidation potential in t,hese experiment,s, it follows that any foreign substance Tyhich may be present upon the electrode a t the start may hinder the adsorption of oxygen
A - calomel half c e l l . 8 -opening forburet2.e tip. C -electrode. FIG.I Ele2trode Yeseel
and so retard the establishing of a final equilibrium and cause the results to be irregular. Acting upon this suggestion, we cleaned an electrode by immersing it for some time in concentrated nitric acid, washed it thoroughly with conductivity water and heated it to a cherry red in the flame of a bunsen burner. It was allowed to cool for an instant and then immersed in the solution, and readings were taken a t once. The results were much becter than previously. An approximate equilibrium was soon reached and results in successive experiments were fairly consistent. The beneficial effects of degassing the electrodes each time before using them were obvious. TThen an electrode is heated in the flame of a bunsen burner it may become contaminated by the gascs of the flame. Then, too, the glass inseal which holds it’ is often cracked in the process. T o overcome these difficulties we constructed an eiectrode which could he heated by passing a n electric ciu’rent thrwgh it,. This consisted simply of a loop of platinum wire about 0.4 niillim:,ter ;i: diameter the cJncis nE which mere 5ealed through the ends ~ in tiicim&r hown in C c;f Fig, I . \Tire iarger than o . . mm.
OXIDATIOS POTESTIALS O F HYPOCHLORITE SOLCTIOSS
42 7
proved unsuitable, as it cracked the glass seal which held it when it was heabed. At first copper wires were fused to the ends of the platinum wire and extended up the tubes, serving as leads. We found, that a pot'ential was set up at the copper-platinum contact surface, and that, this contact potential varied by some millivolts from time to time. This difficulty was removed by using platinum wire connections xelded to the wire of the loop. We found, however, that making contact with the loop by means of a column of mercury Jvhich filled the tube gave the same potential as did connections which mere all of platinum. provided that the mercury was clean. This method of making the contact was used in all our careful measurements. A detailed account of all our efforts to evolve a standard method to use in treating our electrodes cannot be given at this point. As the result of much study, however, we found the following procedure to give sat'isfactory results. Place the electrode in hot, but not boiling, concentrated nitric acid I. for ten to fifteen minutes. Kash thoroughly with conductivity water after cooling and then dry 2. the glass shank of the electrode with a clean lintless towel. 3 . Heat the platinum wire to dull redness for one or two minutes by pasing a suitable electric current through it. 1. After cooling, place it at once in the solution. The necessity for careful preparation of the electrode before each experiment, holding rigidly to t'he same routine method each time, cannot be too forcibly emphasized. Even with such treatment the useful life of an electrode is limited. It was found in many cases t,hat when electrodes began to give erratic results the tips of the lead glass inseals had cracked. h very tiny crack, which apparently did not reach through to the mercury \vas sufficient to cause the trouble. Khenever a crack appeared the electrode was useless for further work. K e believe that these cracks took up nitric acid during the treatment of the electrode and later yielded it to the solution at the electrode surface, thus causing the potential as measured to be abnormal and inconstant. I t was found that electrodes made from new platinum wire were usually much better than those made from wire which had been used previously. We at first believed that this might be due to pitting of the wire by the hypochlorite solution, but careful examination under the microscope failed to reveal any apparent difference between new and used wire. The alteration in the surface which occurred in service, if any, was certainly very slight. Kew wire only was used in making up electrodes for our careful experiments. Results were accepted as trustworthy in this work only when they had been secured iyith electrodes which (a) were free from mechanical defects, (b) which came to equilibrium normally and regularly as shown by measured potentials and (c) which had never given results which were erratic or abnormal. K e believe that t,hese are safe criteria to employ.
V. H. REYIK'GTON A S D H. M. TRIMBLE
428
Our first experiments with eonite solutions exposed to the air, using our best electrodes, never gave an equilibrium potential even in six to ten hours, though the pot'ential rose only slowly after about thirty minutes. This constant rise in potential was not due to a conversion of .hypochlorite to chlorate, for some experiments showed that solutions of sodium chlorate a t concentions comparable to those which might result from our solutions showed much lower potentials than those given by zonite. It was found that solutions fresh from the original container always had a lower potential than those which had been in contact with the atmosphere. It seemed probable that this effect might be due to reaction of carbon dioxide of the air with
1.10
1.00
$ .go
3
k
t
I ; 0 /o PO n n m minutes.
30
YO
50
60
FIG.2 Effect of eliminating CO, from above Alkaline Hypochlorite Solutions
the sodium hypochlorite t o liberate hypochlorous acid. To test this matter an apparatus such as is shown in Fig. I was used. h stream of oxygen was first passed through a soda-lime tube to free it from carbon dioxide and then it was led into the cell over the solution. A fresh solution of zonite showed a slow rise in potential a t first, and the potential became constant after about 14 minutes. At the end of 2 4 minutes, a current' of COZ was substituted for the oxygen, and the potential rose 0.219 volts in the next 8 . j minutes. Carbon dioxide-free oxygen was then again passed, and the potenti 1 came t o a halt almost at once. The results of this experiment are shown graphically in Fig. 2 . Merely interrupting the stream of oxygen after equilibrium had been reached, and breathing into the vessel caused the potential to rise abruptly by &s much as o.oj to 0.1 volt in several experiments. The necessity for excluding carbon dioxide from our alkaline solutions was apparent.
OXIDATION POTENTIALS O F HYPOCHLORITE SOLUTLOSS
429
Effect of adding a Base or an Acid to Hypochlorite Solutions With the technique in hand, we first studied in a qualitative way the effect of progressively adding a base to a solution of hypochlorous acid, and then, reversing the process, the effect of progressively adding an acid to a solution of sodium hypochlorite. Twnety-five cubic centimeters of hypochlorous acid containing only a trace of hydrochloric acid were placed in the flask. h burette, filled with 0.50 S . NaOH was placed so that its tip entered the cell through a hole in the stopper as shown in Fig. I , B. Carbon dioxide-free oxygen was passed during these experiments. Potential readings of the cell were taken every
FIG.3 Relation of Oxidation Potential to Acid or Base added
few minutes and after the lapse of ten minutes when approximahe constancy had been attained a final reading was taken and recorded. Without removing either electrode from the cell, a measured quantity of base was admitted and the cell was swirled as much as connections would permit in order ho facilitate mixing. Ten minut,es later !Then approximate equilibrium had been reached another reading was taken and more alkali added as before. These operations were repeated until several cubic centimeters of base had been added in excess of the amount calculated as necessary to just neutralize the hypochlorous acid present. At this point a burette containing 0 .j o S . HC1 was substituted for that contsining alkali. The cell was not otherwise disturbed.
430
V. H. REMIXGTOS AXD H.M. TRIMBLE
Hydrochloric acid was then added step by step to acidify the solution, employing the same routine method of addition as was used with the alkali, and the potentials corresponding to various quantities of acid added were determined. The results of the above experiments are shown graphically in Fig. 3. Since the curves have been carefully drawn to scale, tables giving the data will not be included. The graph is representative of several experiments which checked fairly well. The lower curve indicates the changes during the addition of alkali; the upper, those found during the addition of acid. The two “breaks” in these curves, we find, occur respectively a t the points where free hydrochloric acid and hypochlorous acid are just neutralized when adding the base; or first introduced when adding the acid. We have carried out similar titrations using tenth normal reagents and find that electrometric titrations of hypochlorite solutions to determine their hypochlorite content and the amount of added acid or base can be carried out using our technique. The process is tedious, because it is necessary to wait for ten minutes before taking the final readings, and somewhat inconvenient because of the necessity for excluding carbon dioxide; but it is fairly accurate. We hope to study it in greater detail in the near future.
Experiments in Buffer Solutions This set of experiments was, of course, qualitative rather than quantitative in nature, so far as a coordination between oxidation potential and hydrogen ion concentration is concerned. F e next measured the oxidation potentials in buffer solutions, keeping the concentration of hypochlorite constant. A few experiments were first tried using a buffer mixture made up with KC1 and HC1 solutions. The maximum oxidat,ion potential of HOCl in such mixtures was reached within a very few minutes, five or six a t the most, and remained constant over a period of thirty minutes or more. Although these experiments were very satisfactory from that standpoint, ClzO and C1, were evolved quite vigorously. The concentration of HOCl necessarily diminishes rather rapidly under these condit;ons and the gases evolved have a serious destructive action upon the apparatus used. Solutions so buffered were, therefore, not investigated further. It seemd best to work with solutions which are more stable and, which correspond more closely to those which are commonly used in a commercial way. Phosphate buffer solutions were prepared after the manner outlined by q .. ,orensen.l He made up his buffer mixtures using fifteenth molal phosphate solutions. I n our case, however, it was necessary to prepare more concentrated solutions so that after mixing with the HOCl solution thc phosphates present xould have the concentration specified by him. Therefore, we used fifth molal solutions of Y a 2 H P 0 4 and KH,POI which were standardized by precipitating and weighing the phosphate as ammonium phosphomolybdate. Table I presents information concerning the composition of the 1
W. hl. Clark: “The Determination of Hydrogen Ions,” p. 114.
OXIDATIOS POTENTIALS O F HYPOCHLORITE BOLUTIOSS
431
mixtures used. Column 2 gives the pH to be established; columns 3 and 1 give the amounts of .2 31. KH3POI and T\Ta2HP04 respectively, to be taken; and column 6 gives the amount of water necessary t o bring the volume of the solution to 50 c.c., the constant volume to which all solutions were diluted. For example, in making up buffer mixture No. I , we mixed 5 cc. of HOCl solution, 0.41 cc. of K a 2 H P 0 4 ,16.25 cc. of KH2PO4 and 28.34 cc. of water. This final solution is of the same concentration in buffer agents as the first given by Sorensen and has a pH of j.29. Walbuml states that the alteration of pH with temperature is for the most part negligible for phosphate buffer mixtures. Our experiments were carried out a t room temperatures, which averaged about 2 4 T . TABLE
1
Composition of Mixtures following Sorensen’s Table Solutions mixed. HOCl Cont. , 0 1 0 2 gm. per cc. Sa2HP04 . zoo molal , 2 0 0 molal KHnPOn C.C.
C.C.
SO.
PH
NalHPOl
KHxPOI
HOCl
H20
0.41 0.83 3.33 11.68
16.25 15.83 13.33
5.00
5.00
28.31 28.34 28.34
5.00
5.00
28.34
15.00
1.66 0.83
5.00
28.34
5.00
28.34
C.C.
I .
5.29
2 .
559
3.
6.24 7.16 7.73 8.04
4. j.
6.
C.C.
15.83 TABLE
j.00
11
Composition of Mixtures following Auerbach and Pick’s Table Solutions mixed. HOCl cont. . O I O Z gm. per cc. SarC03 ,400 molal ,400molal XallCO, C.C.
C.C.
PH
T\’a2COJ
SaHC08
HOCl
H20
8.35 8.90
0.00
25.00
5.00
20.00
8.
I . 2 j
23.75
5.00
20,oo
9.
9.15
2.50
22.50
5,oo
20.00
IO.
1 0 . IO
12.50
5.00
20.00
11.
10.65 11.59
20.00
12.50 5,oo
5.00
20.00
25.00
0.00
j . 00
20.00
C.C.
so. I ‘
12,
C.C.
To cover a more alkaline range of pH, Auerbach and Pick2 used buffer solutions consisting of mixtures of sodium carbonate and sodium bicarbonate, K e prepared .4molal solutions of Na?C03 and SaHCO3 and from them made
’ W. M. Clark: “The Determination of Hydrogen Ions,” p. 116. :IF’.
31. Clark: “Determination of Hydrogen Ions,” p. 323.
V. H. REMINGTON AND H. M. TRIMBLE
432
up sohitions of the concentrations given by Auerbach and Pick. Table 11, similar to that on the phosphates, furnishes information concerning the composition of each of our solutions for the pH values given. The method which we used in making up thsee solutions has been given above. I n these buffered solutions the normality of the hypochlorous acid was about 0.02, while the combined molal concentration of the salts of the buffer mixtures was 0 . 2 M in the carbonate, and 0.067 M. in the phosphate buffer mixtures. 1.36
1.20
FIG.4 Oxidation Potentials of Hypochlorite a t Various Hydrogen Ion Concentrations
Kow hypochlorous acid is a very weak acid. Noyes and Wilson' find that its ionization constant in ,001 normal solution is 6.79 X IO-IO. Sandz found the value 3.7 X IO-*. Results by other workers are in substantial agreement with these. Since the free hydrochloric acid present in our stock solutions was never more than sufficient in amount to give a very faint opalescence when silver nitrate solution was mixed with the solution of hypochlorous acid, it seems safe to assume that the concentration of hydrogen ion in them was very small. We believe, therefore, that the hydrogen ion concentrations in the solutions whose oxidation potentials were determined were not sensibly different from the hydrogen ion concentrations characteristic for the buffer mixtures themselves. W.A. Xoyes and Thomas A. Wilson: J. .Zm. Chem. SOC.,44, 1634 ( 1 ~ 2 2 ) . J. Sand: 2. physik. Chem., 48, 610 (1904).
OXIDATIOS POTEKTIALS O F HYPOCHLORITE SOLrTIOXS
433
The usual course of bhe building up of a potential with hypochlorite in a buffer solution, using an electrode prepared as previously described, may be set forth a t this point. h rapid rise of 0.01volt, or a little more, took place in the first few minutes, and the potential approached constancy after 8 or 9 minutes. The next five minutes showed an increase of only I or z millivolts. At the end of this time there was no further change in potential in most cases. In some instances a further rise occurred, but this rise was never more than one millivolt in the ensuing 15 minutes. We therefore, took the potential at the end of fifteen minutes as the equilibrium potential for the solution which was being examined. In this work we did not depend upon results found with any one electrode or set of electrodes, but we used several in rotation, cleaning them carefully between determinations, and we always repeated the measurements two or three times with different sets of newly constructed electrodes. The results of our work with buffer solutions are given in the Tablei I11 and IV and shown graphically in Fig. 4. Each value is the mean of those found in many determinations. The potentials as measured differed from these values in every case by not more than i.004volts. The concentration of HOC1 in these buffered solutions was in every case .OOIOZ grams per cubic centimeter.
TABLE I11 Variation of Potential with pH with Sorensen’s Phosphate Mixtures PH
Cell Potential
Oxidation Potential
5.29
0.934
I . 272
5.59
0.912
6.24
0.873
1.250 1.211
PH
7.16 7.73 8.04 TABLE
Cell Potential
Oxidation Potential
0.820
I . 158
0.798 0.790
1,135 I . 127
IT
Variation of Pot’entialwith p H with Auerbach and Pick’s Carbonate Mixtures P €1
8.35 8.90 9.15
Cell Potential
Oxidation Potential
PH
Cell Potential
Oxidation Potential
0.777
I.IIj
I O . IO
0.623
0.961
0.723
1.061
10.6j
0.584
0.922
0.695
1,033
1 1 . j9
0 .j46
0.884
We are a t a loss to explain adequately the discontinuity of the curves as found with phosphate and with carbonate buffer mixtures. Certainly the small quantity of hypochlorous acid added could not shift the hydrogen ion concentrations of the buffer mixtures in such manner as to give these discrepancies. We believe that the reason for this failure to give a continuous curve lies in some intrinsic difference in the actions of phosphate and carbonate buffer mixtures;-that is, in a specific “salt effect.’’ Other buffer solutions cover the range which we have investigated, but those whose pH values are well established contain organic components and so could not be used in conjunction with hypochlorites.
434
v. n. REMINGTON
AND H. M. TRIYBLE
It was planned to repeat these experiments with buffered solutions, using sodium hypochlorite, but. lack of t,ime prevented. X few preliminary experiments, however, showed that solutions of sodium hypochlorite and hypochlorous acid a t the same concentration, in terms of available chlorine, gave nearly the same potent,ials in a solution of given pH, those with sodiuni hypochlorite being, in general, slightly lower. Our results, though they only partially corer the field, show, we believe, that the oxidizing power of hypochlorite solutions increases as one passes from alkaline to acid solutions. This is in agreement with the information upon this subject which is now available in the literature. This work seems t o indicate, in a rough way, that the oxidizing power of a hypochlorite solution is directly proportional to its hydrogen ion concentration. S-arY The oxidation potentials of hypochlorous acid and sodium hypochlorite in various solutions have been determined. h great part of the work was given over to a study of the method of measuring these potentials at platinum electrodes, and an apparatus for their measurement has been evolved. A method of treating the electrodes which gives reproducible potentials has been worked out. The oxidizing power of hypochlorites in solution as measured indirectly by potentiometric methods increases with acidity, that is, with increase in hydrogen ion concentration. I t has been found possible to determine roughly the hypochlorite content and the amount of added acid or base in hypochlorite solutions by electrometric titration. Buffered solutions of the hypochlorites are more stable and give much more readily reproducible potentials than do unbuffered solutions. Slightly alkaline solutions of the hypochlorites are very readily decomposed by action of COa of the air, with a marked increase in oxidizing power. Stillwater, Oklahoma. September 4, 1928.
