ELECTRODE POTENTIALS IN FUSED SYSTEMS
Nov., 1954
fusion of the number of moles of water per mole of hydrate is obtained. Division of this quantity by the heat of fusion of a mole of ice (1430 cal./g. mole) yields the number of moles of water per mole of hydrate or 16.93 0.53. This result compares favorably with the formula, CHClzF.17Hz0, predicted by Claussen.636
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(5) W. F. Claussen, J . Chem. Phys., 19, 259 (1951). (6) W. F. Claussen, ibid., 19, 662 (1951).
965
Acknowledgment.-This research was partially supported by the Office of Naval Research under contract N8-onr-77100. Appreciation is expressed to the Office of Naval Research and to the University of Oklahoma Research Institute for administration of the contract as Research Project Number 46. The authors wish to thank Dr. V. G. Meadors for his helpful suggestions in the preparation of this paper.
ELECTRODE POTENTIALS IN FUSED SYSTEMS. I. SODIUM HYDROXIDE 'i2
BY KURTH. STERNAND JACKK. CARLTON The Institute of Science and Technology, University of Arkansas, Fayettewille, Arkansas Received January 16, 1964
Potentials of silver, copper, nickel, cobalt and tungsten were measured in fused sodium hydroxide over the temperature range 340-600°,using gold as reference electrode. Potential-time curves were obtained. Potential changes are related t o reactions occurring a t the electrodes.
Introduction This study was undertaken to determine if e.m.f. measurements offered a convenient method of detecting reactions between metals and molten compounds. Sodium hydroxide was selected because its reactions with various metals had been investigated previously2 and because some work is currently being carried on in this Laboratory on the reaction between nickel and sodium hydroxide a t high temperatures. I n addition, we hoped to gain information of more general interest on the potentials of metals in foreign ion systems. Virtually no work of this kind has been reported in the literature although several workers have studied foreign ion cells in aqueous solution^.^-* The work of MacGillavry and co-workers on the potential of nickel in foreign ion solutions, including aqueous sodium hydroxide, is of interest and will be referred to later. Lux has studied reactions in various melts, particularly on the addition of oxides to a KzS04-LizS04 eutectic a t 950°.9 He found that gold electrodes could be used to follow the progress of a reaction by measuring the change in potential between a reference and an indicator electrode. Recently Hill and Porteriohave studied the electrode potentials of various metals in a KzSO4-Li2SO*melt in the presence of oxygen, using an oxygen electrode as reference.'O I n that case, however, the system (1) Presented in part at the Southeast-Southwest Regional Conclave of the American Chemical Society, New Orleans, December 1953. (2) The authors wish to acknowledge the financial support of the U. 9. Air Force under sub-contract AF 18(600)-960. (3) M. Le Blanc and L. Bergman. Ber., 42, 4728 (1909). (4) A. R. Tourcky and S. E. 8. El Wakkad, J . Cham. Boc., 740,749 (1948). (5) L. Colombier, Compl. rend., 199, 273, 408 (1934). (6) D. MacGillavry, J. H. Rosenbaum and R. W. Swenson, J . EEectrochem. SOC.,99, 22 (1951). (7) D. MacGillavry, J. Chem. Phys., 19, 1499 (1951). ( 8 ) D. MacGillavry, J. J. Singer and J. H. Rosenbaurn, ibid., 19, 1195 (1951). (9) H. Lux, Nalurwiss., 26, 92 (1940): 2. Eleklrochem., 62, 220 (1948); ibid., 63, 43 (1949). (10) D. Cr. Hill and B. Porter, Tars JOVRNAL, in press.
is reversible with respect to oxide ion and no longe.. constitutes a true foreign ion melt. I n the present work, oxygen was excluded from the system to preclude the possibility of reaction between the metal electrodes and atmospheric oxygen. Experimental Part Materials.-Metal electrodes were obtained from various sources. All were spectroscopically pure. Copper, tun sten and. silver electrodes were made from No. 10 B. and gage wire. Each nickel electrode consipted of a solid cylinder l/1 inch in diameter and 1 inch long, welded t o a rigid nickel wire. Cobalt electrodes were prepared from cobalt rondelles into which nickel wires were hammered to provide connection to the external leads. The sodium hydroxide used in most of the measurements was Mallinckrodt Analytical Reagent Grade containing about 2% sodium carbonate. I n a few measurements a special grade of sodium hydroxide containing less than 0.1% sodium carbonate was used. This was kindly given to us by Dr. Overh o l m of the Oak Ridge National Laboratory. Experimental Methods.-The construction of the furnace is shown in Fig. 1. All openings in the glass cover were made as small as possible to prevent contamination of the system by air. The only openings in the furnace cover were those for the two electrodes, thermocouple wires and heater wires. These openings also served as exit for the gas assed through the system. A i measurements were made in a gold crucible since gold is the only metal known not to react with sodium hydroxide a t high temperatures. A heavy (No. 10) gold wire was fused to the crucible for use as electrode. The procedure in carrying out most of the measurements was as follows. The gold crucible was placed in the furnace and filled with 40-50 g. of sodium hydroxide. The furnace was then closed and the heating current turned on, while argon which had been bubbled through a potassium yrogallate solution, concentrated sulfuric acid, and passed &rough a drying tube filled with calcium rhloride, was passed through the cell. After the desired temperature had been reached, the metal electrode under study was lowered into the melt and the potential between it and the gold electrode measured by means of a Leeds and Northrup K-2 otentiometer. The temperature of the melt was measyedP by a chromel-alumel thermocouple whose hot junction was kept in the top of the furnace and could be lowered into the melt. Since the measured electrode potentials were relatively insensitive to temperature changes no attempt was made to control the temperature more accurately than f3'.
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KURTH. STERNAND JACK K. CARLTON
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GOLD ELECTRODE
Vol. 58
ELECTRODE
PYREX COVER .HEATER WIRE
.GOLD CRUCIELE , FURNACE CORE
100 2oo
I
01
0
BRICK PYREX CONTAINER 'METAL CAN
I Fig. 1.-Apparatus
50 100 150 200 250 300 350 400 450 Time, min. Fig. 2.-Electrode potentials of cobalt in molten sodium hydroxide: 1, "C.; 0, 350; A, 430; V,535.
900 800
I for the measurement of electrode potentials.
When a gold wire was inserted in the melt small potentials (15-20 mv.) were measured, indicating a thermal gradient between the center and wall of the crucible. No corrections were made for this thermogalvanic potential or for any of the junction potentials in the system. Several runs were made to establish the reproducibility of the results. Using both copper and tungsten electrodes, the steady-state potentials of each of these metals could be duplicated to within f 0 . 5 mv. Potential values were not affected by the pretreatment of the electrodes. Polishing with steel wool or emery paper, followed by washing in distilled water and drying with tissue paper, was sufficient. At the end of each run the melt was dissolved in water and the crucible cleaned by soaking in concentrated nitric acid for several hours, washing with distilled water and drying at 110'. Analyses for metal content of the dissolved melts were carried out in many of the runs. Figures 2 and 5 show the time-potential curves for cobalt, copper, nickel and silver. All the potentials shown are negative relative to gold. Our observations on the individual metals are given below. Nickel.-Up to 500" nickel does not react with molten sodium hydroxide if oxygen is excluded from the system. In the presence of even small concentrations of oxygen a black oxide coating forms on the electrode. During the formation of this oxide, probably NiO, which takes place very rapidly, the electrode potential drops from its initial value to zero. Analysis of the melt by the usual analytical methods shows less than 0.5 mg. of nickel ion, indicating that the potential drop is due only to reaction(s) occurring on the electrode surface. It is interesting to note that the shape of the time-potential curves is quite similar to that of similar curves reported by MacGillavry and co-workers8 for nickel in aqueous hydroxide solution. However, in aqueous solution the curve rises over a longer time interval. Cobalt.-The behavior of cobalt is essentially like that of nickel. However, the metal reacts with sodium hydroxide at somewhat lower temperatures than does nickel. The slow decrease of potential at 535' is accompanied by a gradual accumulation of sludge in the melt. Cobalt is more resistant to air oxidation than nickel as shown by the relatively slight change in potential when oxygen is allowed to leak into the system. Silver.-The potential of silver in sodium hydroxide is more sensitive to very low concentration of impurities in the melt than that of any other metals investi ated. The heavy curves of Fig. 5 show the behavior of %e metal in Mallinckrodt Analytical Reagent Grade sodium hydroxide. The potential decreases so rapidly from its initial value that no accurate measurements were possible. Thus the otential a t zero time is quite uncertain. It was noted, Rowever, that a t the end of each run the electrode was coated with a dull white substance. Since the Malliuckrodt reagent con-
r
1
700 600
.--TRACE
OF Op
500 I
"1 100 0
I
0
50
100 150 200 250 300 Time, min. Fig. 3.-Electrode potentials of nickel in molten sodium hydroxide: t, 'C.; 0, 420; V, 450; A, 470; 0, A, 580.
tains 2% sodium carbonate it was considered possible that the coating was silver carbonate. Measurements in a special grade of sodium hydroxide containing 0.01% ' sodium carbonate, kindly given to us by the Oak Ridge National Laboratory, gave the time potential-curve shown as the light curve in Fig. 5. Fisher Certified Reagent Grade, containing 0.1% sodium carbonate gave similar results. However, the addition of sodium carbonate to a melt of this reagent did not change the shape of the curve although it raised the potential. It was then noted that the Mallinckrodt Reagent contained 0.01% chloride whereas the Fisher reagent contained 0.000%. When a small amount of chloride is added to the Fisher reagent the results are the same as for Mallinckrodt sodium hydroxide. Hence, we conclude that the coating on the electrode responsible for the rapid decrease of potential is silver chloride. Tungsten.-Tungsten reaches its final potential very quickly. Its value is 1.05 volts, almost independent of temperature. The metal behaves the same in both the presence and absence of oxygen. The addition of sodium tungstate to the melt produces no change in potential. Up to 500' tungsten dissolves slowly or not at all in sodium hydroxide since less than 0.5 mg. of tungsten is found in the melt. An ap reciable amount of current can be drawn from the cell WP/NaOH/Au without changing its potential. In one experiment six milliamperes were drawn from the cell for an hour. The cell regained its initial potential a few minutes after the circuit was opened.
ELECTRODE POTENTIALB IN FUSED SYSTEMS
Nov., 1954
Copper.-This metal does not react with sodium hydroxide below 400" as shown by the steady potentials measured. At 535" the metal reacts slowly. Potentials are observed to fall slowly and a blue color appears in the melt. These observations are in agreement with those of Bergman and LeBlanc' who re ort no reaction st 435" but find the melt becoming dark b&e a t 568". In the presence of trace amounts of oxygen (such as are present in unpurified tank argon) copper shows rapid blackening due to oxide formation. This is shown by the light curve in Fig. 4. Several measurements of this particular phenomenon show that the shape of the potential curve:accompanying oxidation is quite characteristic of copper. In addition to the initial relatively flat part of the curve which may last up to one hour and probably corresponds to the buildup of a critical thickness of oxide, the decreasing part of the curve shows an inflection a t approximately 0.3 volt. So far we have no evidence as to the meaning of thiszinflection.
700
r
600
\ h \
BECOMES CONSTANT AT 500 M V . 7
967
700 600 500
400 300 I
2
200
II
2
100 0
- 100 -200 100 150 200 250 Time. min. Fig. 5.-Electrode potentials of silver in sodium hydroxide: t, "c.; 0, 350; A, 400; V, 500. 50
3
100 150 200 150 300 Time, min. Fig. 4.-Electrode potentials of copper in molten sodium hydroxide: t, "C.; 0, 340; A, 400; 0 , 535; -, in the presence of traces of 0 2 at 500".
50
It is very difficult to determine the electrode reaction(s) responsible for the measured potentials, primarily because analysis of the melts, at the temperature under investigation, cannot be carried out. Hence we do not know what species are present in the system when the potential is being measured. It is only possible to dissolve the melt in water and analyze the solut.ion. For example, one attempt was made to study the potential dependence of the copper electrode on the copper ion concentration in the melt. Various concentrations of anhydrous CuBrz in NaOH in the range 0.005 to 0.05 molal were prepared and the otential of the copper electrode measured as above. Quayitatively, the potential decreased with increasing copper ion concentration. However, CuBrz reacts with molten NaOH according to the equation CuBr -+ 2NaBr HgO CuO. Hence, the 2NaOH copper ion concentration decreases slowly from its initial value and the potential finally reaches the value it has in the absence of added CuBrz. Analysis of solutions initially free from copper showed that they contained copper concentrations from 0.0001 to 0.003 molal with no change in measured potential. When copper is made the anode of the cell Cu/NaOH/Au by imposing an external potential it is found that only about 10% of the total current passed results in the solution of copper from the electrode. For example, in one run 3 X 10-4 equivalent of copper was found in the melt or lated faraday orelecout on the gold cathode when 4 X tricity had been passed. However, some gas evolution was observed at the anode. Similarly, when gold was made the anode, the resulting melt was less than 10-6 molal in gold.
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Discussion From the above observations the following generalizations may be made :
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1. Reproducible electrode potentials of copper, nickel, cobalt and tungsten can be measured in sodium hydroxide, using gold as reference electrode. No constant potential was obtained for silver but the highest value on a time-potential curve is reproducible. 2. Any change in potential may be due either to film formation on the electrode or to the solution of the metal in the melt. The potential change indicative of film formation is much more rapid than for solution. 3. Up to 400" none of the metals dissolves in sodium hydroxide to an appreciable extent in the absence of oxygen. At 500" none of the metals except copper does, and the latter only slowly. These results are in general agreement with the data of Bergman and LeBlanca and a concurrent study of the high temperature reaction between nickel and sodium hydroxide carried on in our laboratory. l1 4. The metals do not constitute polarizable electrodes in the sense that mercury in contact with an aqueous solution is polarizable. For example, when a steady potential has been reached the cell will recover from brief charges of several tenths of a volt in a few minutes. Hence, a mechanism for charge transfer across the metal-solution interface must exist. 5 . Metals which are capable of formi.ng films in sodium hydroxide, ie., whose films are insoluble in the medium, also show an "induction period," a time interval between the initial immersion of the metal in the melt and the time at which a steady potential is reached. This time interval is longer a t lower temperatures and may extend to several hours near the melting point of sodium hydroxide. The phenomenon seems to be connected with the establishment of a reproducible surface on the metal. When a piece of metal has been immersed (11) Unpublished results.
HELGA BOEDTKER AND PAUL DOTY
Vol. 58
No theory has yet been formulated t o account for the potentials observed in fused sodium hydroxide. One difficulty preventing such a formulation is our lack of knowledge as to what species exist in molten sodium hydroxide. For example, evidence exists" that above 600" Na2Ni02,H20 and H2 are formed in the reaction between nickel and sodium hydroxide. Below 600" the reaction is so slight that products cannot be isolated. It is entirely possible, however, that very low concentrations of these compounds, while not detectable by ordinary chemical methods may be potential determining. Some evidence for this view rests on the fact that only a small part of the current passed through the system, as described above, results in solution of the metal. It is also possible that the potential is determined by more than one electrode reaction. For example, the observation of the plating out of copper at the TABLEI cathode and simultaneous gas evolution a t the ELECTRODE POTENTIALS OF METALS IN SODIUM HYDROXIDEanode indicate that the cells are not reversible with respect t o the metal ions involved, or a t least that AT 400' M++ 2e- are reother reactions besides M Potential (v.) Meta relative t o gold sponsible for the observed potentials. It is possible, Copper 0.61 for example, that a t the anode 0- --t O2 2e-, 2e--+ Hz 20-. while a t the cathode 20H0.67 Nickel It is likely that our understanding of potentials Silver 0.68 in the sodium hydroxide system must await the Cobalt 0.70 study of potentials in simpler electrolytes. Tungsten 1.05
in the melt for some time the surface appears to have been "polished" more brightly and uniformly than would be possible by purely mechanical means. The more the metal has been polished before immersion the shorter is the induction period. Thus it appears likely that the induction period represents the removal of surface impurities. The final electrode potential is not affected by the previous treatment of the metal. On the other hand the potential of nickel in aqueous sodium hydroxide depends on the careful reduction and outgassing of the metal before immersion in the solution. Probably the solution of surface impurities occurs rapidly only at temperatures higher than the melting point of sodium hydroxide. 6. In Table I are listed the steady-state electrode potentials of the metals studied.
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A STUDY OF GELATIN MOLECULES, AGGREGATES AND GELS' BY HELGABOEDTKER~ AND PAUL DOTY Contribution from Gibbs Laboratory, Chemistry Department, Harvard University, Cambridge, Mass. Receiued February 4* 196.4
Samples of moderately high molecular weight gelatin were found to be molecular dispersed in 0.15 M NaCl at 40' and in 2 M KCNS. Light-scattering measurements of the molecular size showed the gelatin molecules to be random coils with mean configurations comparable to those of typical synthetic polymers. Gelatin solutions form gels when the temperature is lowered, provided that the gelatin concentration is sufficiently high and particular substances such as KCNS and uare are absent. Light scattering and viscometric studies a t 18' show that aggregate formation occurs instead of gelation when the gelatin concentration is below a critical value. The size and weight of the aggregates increase with further lowering 3f the temperature and with the gelatin concentration. The distribution of mass within the aggregates resembles a Gaussian distribution. Despite this resemblance to randomly kinked, chain molecules, the aggregates are cross-linked. The nature of these cross-links is shown by the fact that these aggregates dissociate slowly and only to a limited extent upon dilution although, like gelatin gels and crystalline polymers in general, they dissociate completely in a narrow temperature range. This and other information indicates that the cross-links are very small crystallites. I n gelatin gels a t 18" thr! intensity of scattered light was found to be substantially greater than in the corres onding solutions at 40"; however, the scattering in the gels decreased with concentration in contrast to the behavior o r the solutions. Evidence is presented that the scattering centers in the gels are density fluctuations similar in weight and volume to the aggregates studied in solutions and falling off in scattering does not necessarily indicate smaller scattering centers. The gelation of gelatin is attributed to the simultaneous growth and interlocking of aggregates arising from crystallite formation.
An examination of the extensive literature on gelatin shows that little is known concerning the relation of gelatin molecules, which can exist independently in solution, to the gel, which they can form a t sufficiently high concentrations and low temperatures. Since gels fail to form upon cooling when the concentration is sufficiently low, we thought that in this situation gelatin molecules would, nevertheless, come together to form aggregates. Such aggregates were indeed readily detected and their study has served to bridge the gap between the independent molecule and the gel (1) Presented in part at the 123rd Meeting of the American Chemical Society, Lo8 Angeles, March, 1953. (2) Public Health Service Predoctoral Fellow, 1852-1953.
because they display the same forces that hold the gelatin molecules together in the gel state on a scale accessible to study in solution. Our work had two other motives as well. From the point of view of light scattering, it was difficult to believe the wide-spread assertion that gelatin ge!s do not scatter light except near the isoelectric point in the absence or near absence of added electrolyte. Moreover it was of interest to see how the light scattering of the gel, provided it scattered, differed from the corresponding solution, to see if an adequate explanation could be offered and to see what conclusions could be drawn about the structure of the gel. The second motive arose from the renewed interest in the molecular struc-
