T H E POTENTIAL OF T H E IRON ELECTRODE* BY W. H. HAMPTOA-
The standard electrode potentials have been accurately determined for most of the common metals from electrodes possessing a high degree of constancy and reproducibility. However, one of the greatest difficulties met in potential measurements has been to obtain such an electrode from a form of iron that will represent the most stable state of this metal and which can be used in a reversible cell. There is described in thie paper a series of electrodes and equilibrium measurements from which a value has been determined for the standard electrode potential of iron. This value can hardly be in error by more than a few millivolts and is much more accurate than anything published heretofore. There has been only a small amount of work published dealing with a direct study of the electrode potential of iron, although much discussior, occurs in the literature regarding the phenomena that are observed when this metal is involved in electrolytic processes. The most important work connected with the potential measurements is that of Richards and Behr,l mho studied the e. m. f . of the cell, Fe 1 FeS04(o,glLI),KCl(o.IM), HgClI Hg. They used various kinds of iron and their results varied several hundredths of a volt, depending both upon the kind of iron and its previous treatment. By the use of finely divided reduced iron, constancy was obtained within about one centivolt, but because of the large liquid potential involved and because of the unknown activity of the Fe++ ion in such a concentrated solution, it is impossible to calculate from these measurements a valae for the standard electrode potential. Experimental Part The cells used in the following experiments consisted of the elements Fe I Fe Clz (o.IM), HgCl I Hg and contained no liquid junction. Various concentrations of ferrous chloride were employed, but, for the most part, the measurements were made in one-tenth molal solutions and, iinless otherwise stated, this mill be the concentration referred to. To prevent oxidation, an H-cell was used, which could be filled and kept free from air as much as possible. The two half-cells were provided with a glass stopcock between them. The procedure in filling was to add about an inch of mercury to the calomel side and then evacuate this portion. Ferrous chloride solution saturated with calomel together with calomel paste was added by simply letting the vacuum draw in these materials. The other half of the cell was filled iinder an atmosphere of hydrogen and then stoppered to exclude the air. Complete connection was now made by opening the stopcock between the two cells. * Contribution from the Chemical Laboratory of the University of California. Carnegie Inst. Pub., No. 61.
POTER'TIAL O F THE IRON ELECTRODE
981
Standardized ferrous chloride solutions were prepared by dissolving powdered iron in hydrochloric acid following a similar procedure to that used by Richard and Behr in preparing their ferrous sulfate solutions. For the metal electrodes, three different varieties of iron were tried, an iron amalgam, finely divided iron and ordinary annealed iron.
Iron Anialgrlrn: Measurements were made on amalgams which were prepared by electrolyzing a solution of ferrous sulfate, a method which was first discovered by Jou1e.l As iron is but very slightly soluble in mercury, it deposits as a suspended solid phase completely taken up and wetted by the mercury. Amalgam mixtures can thus be obtained varying in structure from liquid to almost completely solid in form, depending upon the amount of iron present. The iron in the solid phase is very finely divided as it gives no granular feeling whatever when rubbed between the fingers. The metal in these amalgams shows its ordinary reactivity, for on exposure to air a black film of oxide will immediately form on the surface of the mixture and, when water is present, rust will appear. The sample for cell KO. I was taken from an amalgam containing only a little solid phase in the liquid mercury. Cells Xo. 2 and No. 3 had for their electrodes amalgams of about the same consistency as soft putty, while the iron-mercury mixture for cell No. 4 was somewhat crumbly in structure.
TABLE I Electrodes prepared from Iron Amalgams CELLKO. I CELLS o . 2 CELLNo. 3 CELLNo. 4 Amalgam Maximurn Amalgam Maximum Amalgam Maximum Amalgam untouched. after untouched. after untouched. after untouched. stirring. stirring stirring.
Time 0
min. 2 hrs. I day 2 days 3 days 4 days j days 6 days 13 days 20 days 40 days
.7970
.8198
.go63
,818j .8160 ,8137
.go74
.8100 .8147 .81o1 .go63 .So28
20
,7880 .7682 ,7933 .go01
.794z
8012 8100 8106 810j
*Sew solution added. J. Chem. Soo., 16, 378 (1863)
,79:3 .7927 .7907
,8101
.801j
.go84
,7890
.go62
.7983
,8115
.7970 *.7981 7994 '
982
W. H. HAMPTON
SAMPLENo. 5-cell
on which a series of measurements were made at different concentrations
Concentration ,009 j Molal .os0
),
.0865 .IO
50 2.245 5.03
Potential
. 8 3 0 Max. Value .8278 =i.-o o o j for period 1 5 days. ,8136 & ,0005 ” ,, 3 0 ,, ,8087 i .ooo5 ” 5 ,, ,)
1,
‘
,, ”
.7478 i .ooo; .6668 i ,0005 ,5760 i .ooo5
11
”
’I
”
” 11
I3 6
8
11
” 17
It can be seen that the potential fell continually in all the cells on standing. Such an action results probably from the iron being used up on the surface and on bringing fresh material to the surface layer by stirring a higher potential is again obtained. As would be expected, the e.m.f. does not fall off so rapidly with the more solid amalgams. These cells as a whole, however, show the typical variations of potential measurements of iron which are hard to fully explain. Cell Xo. 2 and Cell No. 3, for example, were made up to be as nearly alike as possible. The best explanation of differences encountered above seems to be that based on solubility phenomena. As the mercury completely covers the solid phase, the activity shown must be that of the metal in solution. The surface layer must always be a saturated solution to give the true activity of pure iron. We are confronted both with the possibility of a very slow rate of solution and with the exceedingly small solubility of iron in mercury. There is no means of telling directly whether or not any particular sample has for its surface film a saturated, unsaturated] or, perhaps, a super-saturated solution of iron. However, if all of the possibilities exist, some rough equilibrium should appear from which the potentials of an unsaturated solution would gradually decrease, and those of a super-saturated solution approach that equilibrium arising from the reaction of iron with water. The experiments support this idea somewhat. From these considerations, the best sort of amalgam would be one which had as large a percentage of iron as possible with just sufficient mercury to cover the surface. To obtain this condition, a portion of the material used in cell Xo. 4 was heated under a vacuum, and the mercury distilled off until the residue began to assume a dark gray color differing markedly from the bright lustre of mercury. This material (sample No. 5 ) gave one of the best cells for reaching a constant value, and a series of measurements were made in various concentrations of ferrous chloride. The value of .8087 volts for 0.1molal solution checks that of .8136 volts for .086 j molal solution when the potentials are calculated for the same concentration. The results for the amalgams in general indicate that the solid phase is not any compound formation of iron with mercury but is a finely divided suspension of the former metal, since the value .8087 volts checks very closely that of iron produced by reducing Fez03 with hydrogen.
POTENTISL O F THE IRON ELECTRODE
*
983
Finely Divided Iron: Table I1 contains a series of measurements made on electrodes prepared from finely divided iron. The samples for cells numbers 6, 7 and 8 were prepared by reducing F e z 0 3with hydrogen which had been freed of all traces of oxygen and water vapor. For the material used in cell No. 6 the reduction was carried out a t about 800°C for 24 hours. The reduced iron was withdrawn from the electric furnace, crumbled up, and hurridly packed around the platinum head at the bottom of the cell. There was considerable evidence of oxidation as the e.m.f. was quite low a t first and then gradually increased. This was apparently due t o a thin, oxidized film formed on the metal while being removed from the furnace and placed in the cell which was slowly dissolved, though not entirely, as the cell did not reach the accepted equilibrium value of the amalgam variety. The sample used in cell No. 7 was made byreducing F e z 0 3under the same conditions as used for cell No. 6 except that an iron wire was placed in the Fe203and on reduction this held the spongy mass of iron formed. Instead of powdering this up and placing it in the bottom of the cell, it was hung from the top immediately after withdrawing from the furnace. Cell No. 7 had the exceptionally constant value of .8088 volts which did not vary more than a tenth of a millivolt for thirty dzys. During this time, the solution was removed over the iron three times and each solution of ferrous chloride had been separately prepared. The electrode of cell No. 8 was a sample reduced a t 850°C which gave a high initial value. Richard and Behr attribute such high initial values t o an active state of hydrogen occluded by the metal. However, the potentials TABLE I1 Electrodes prepared from finely divided iron Time
Cell K c . 6 Cell F o . 7 Cell S o . 8 Cell X0.9 Cell KO.I O Cell No. I I Cell No. 12 .So17 ,7884 .7962 .8390 .8108 ,8316 .8462 max. ,8202 max.
. 8082 ,8084
,8216
,8048
,8018
, 8 15 2
.80j8
,808j
,8115 ‘ 8099
,8061 .So65
.So24 .So31
.8085 ,8086
.So70
,8048
,8077
,8048
,8087 ,8087 .8088 .8088 .So88 .So88
.SO35
.7997
.8102
.?$IO8
.80j2
,7906
.So32
.so52
.So90
.8088 .So90 .So62 .So76 .SI01
984
717. H. HAMPTON
given by iron produced by reducing Fe?03which carbon monoxide (cells No. I I and No. 12) had similar high values at the beginning. This iron contained considerable carbon after reduction and it is not known what effect this might have had. It seems improbable that carbon thus contained would increase the potential in any way. The writer is inclined to believe that these high initial potentials are due to exceedingly small particles of iron which are used up until the variation of activity with the size of particles has become negligible, and an equilibrium value around .8090 volts is established. The electrodes of cells KO. 9 and No. I O were a variety of Kahlbaum’s C. P. iron powder which had been reduced a second time before placing in the cell. These agreed fairly well with the constant cell, Both the electrodes mere of the same material and given as nearly as possible the same treatment. This again illustrates the type of differences in measurements which are constantly met and hard to account for. Annealed Iron: Turning now to the series of measurements made with annealed iron as electrodes, it will be unnecessary to explain any of the experiments in detail as all the plain wire electrodes behaved similarly, giving considerably lower values than the finely divided form. No equilibrium value can be assigned to this form of iron, but, in general, the potential was about .05 volts lower than the other varieties. At times, samples approach the value of the finely divided form wikhin .03 volts. Making the iron the cathode by simply leaving the working battery of the potentiometer slightly unbalanced against it, the e. m. f . would rise at a much faster rate than could be accounted for by polarization resulting from concentration changes, for, by placing the same electromotive force (one or two tenths of a milivolt) on several types of cells not in connection with this work, there was no effect whatever. In like manner, by making the electrode serve as an anode, the potential would decrease. This seemed to confirm the idea that iron in the annealed form has always a more or less passive state when immersed in water solutions and that the true activity may be very near that of the reduced iron. Furthermore, the same sample gave values more nearly that of finely divided iron in solutions of higher concentrations, differing .03 volts and, at times, approaching the higher value within about 01.volts. For a final test of this idea, which seems quite conclusive, it was proposed to use some ionizing agent other than water with which the iron would not react. The only available medium was an eutectic mixture of lithium and potaseium chloride. Ferrous chloride previously dehydrated and dried under an atmosphere of hydrogen chloride was added to the eutectic mixture. The resulting mixture of fused salts was further dried by first passing dry hydrogen chloride through the mixture and then hydrogen. Annealed iron wire served as one of the electrodes and finely divided reduced iron as the other. TWO cells were thus made up using separately prepared electrodes. One of these showed finely divided iron to be more active by 2.3 f .3 millivolts. The temperature of this cell was maintained at 38ooC by a fused nitrate bath. The other showed a slightly greater difference, remaining constant at 3.3 f .3 millivolts and the temperature was varied from 35oOC to 45oOC. We would
POTENTIAL O F THE IROR ELECTRODE
985
naturally expect to get small differences of this order of magnitude, which are attributed to the relative state of subdivision. The true activity of annealed iron is therefore very near that of finely divided iron, and the large difference shown in aqueous solutions must be due to passivity of the annealed form incurred by water. Table I11 contains a series of measurements using iron wires for the electrodes. The electrode of cell Xo. 13 was a plain iron mire that had been heated for 36 hours at 85ooC in the reduction furnace and allowed to cool slowly. The wire in cell No. 14 was superficially oxidized by heating and cooling in air for several times and then reduced in the furnace. The wire in cell No. 15 had been set in water until a large amount of rusting had taken place. It was then reduced by hydrogen. The electrode of cell No. 16 mas dipped in dilute hydrochloric acid for several minutes, It was then washed in 0.1 M ferrous chloride and placed in the cell. The wire in ccll No. 17 was simply sand-papered before placing in the cell.
TABLE I11 Electrodes prepared from iron wire Time o days 2
4 6
8 IO
,) ,, ”
Cell S o . 13 .7190 ’ 7409 ‘7484 . 7 j06
”
,,
Cell KO. 14
Cell No. 1 5 8052 7797 7773 7773 7774
’7703 ‘7798 .7834 ,7835 ’ 7830
Cell Yo. 16 ’ 7486 ,7800 ‘7793 ‘
.7532
’)
,,
2
4 6
8
,7575 . 7 580 ‘7592 .7696
,, ,) ,) ”
.7140 .7095 .7106 .7153
”
. 7 2 70
IO
))
I2
’,
16
”
,7803 .7719
,7515 .7578 Cell No. 18
0
.7490 ,7912
.7800
I2
16 24
7807
Cell S o . 17
7330 .7302 .7350 ~
Cell S o . 19 .6864 .7189 ,7160
Cell No.
20
.7475 ,7592 ‘ 7492 .7484 ’ 7484
.7377 . -/ -9 2 9 ,7186
-
The electrodes of cells Xo. 18 and S o . 19 received the same previous treatment as those of No. 16 and No. 17 respectivelj-, but the measurements were made in ,500M ferrous chloride solucion. Cell No. 2 0 is shown for comparison having for its electrode reduced iron oxide in .so0 M ferrous chloride solution.
986
W. H. HAMPTON
Measurements at various concentrations: Only a few cells were measured in concentrations other than one tcnth molal. In all cases, the nature of the variations was the same. In the more concentrated solutions the variations as a whole were not quite so great. However, for Folutions as dilute as .OI M no samples would give a constant value within any reasonable time. The tendency in dilute solutions was to decrease continually for a long period of time. Calculation of Results
Choice of results f o r calculation: The measurements given comprise only a small amount of the total number made, but they are selected to show the typical variations as well as to present all the peculiarities involved in this study of irop. In many ways, they duplicate the observation of others, but additional experiments were carried out t o test new ideas which made it possible to arrive at a more truly representative value of the activity of iron. The most constant value given by any sample was that of cell S o . 7 , namely .go88 volts in . I M ferrous chloride solution, Several cells remained constant €or long periods of time within a few tenths of a millivolt of this potential. Furthermore it was approached from both directions. For constant values at other concentrations, we have the series of measurements with sample No. j. There is no check on the results in concentration above .I M for there are no data at hand to make the necessary calculations. However, the values for .086 j M, .o j iLI and ,009 j 1 4 can be used and are tabulated in Table IV. Calculation of Results: When two faradays of electricity pass through the cell Fe 1 FeClz(aq),HgCl 1 Hg, the following reaction takes place Fe+z HgCl=Fe+++zCl-+z Hg. (1) The electromotive force is given by the equation E = E o - RT/zF In 4yam3 =E" - 0.08873 log + ' r y m (2) in which E" is the potential when the activity of the ferrous chloride is unity, m is the molal concentration and y the activity c0efficient.l
As there are no data at hand from which the activity coefficients of ferrous chloride can be calculated, the values for barium chloride mere taken at corresponding concentrations. The justification for doing this arises from the fact that the chlorides of the bivalent salts as a class with the exception of HgClz and CdClz behave very similarly. For example, the solubilities of thallous chloride in . I to .OI molal concentrations of ferrous chloride mere found to be the same as that given by Bray and Winninghof2 for barium chloride. In column z of Table IV are given the activity coefficients obtained by plotting the values which have been accurately calculated by Lewis and Randall for barium chloride from freezing point data against concentration and interpolating for various concentrations. 2
Lewis and Randall : "Thermodynamics," Chapter XXIX (1923). J. Am. Chem. SOC., 33, 1663 (1911).
POTEXTIAL O F THE IRON ELECTRODE
987
TABLE IV m
Y
E
E"
.IO0 ,086;
I501 . j16 .j63 .719
,8088
'7113 '7II7 ' 7095 ' 7090
.os0 '
009 5
,8136 .82
78
,8830
From a review of the discussion of the different measurements, the most probable value for E" is that calculated for . I M and we may then write the equation : Fe I FeClz (as.), HgCl 1 Hg : E" = .7 I 13 (3 1
The Sormal Electrode Poteiatial: potential, we make use of the cell1
To calculate the normal electrode
Hz/HC1 (o.oIM), HgClIHg: E = . ~ I o ~ (4) and using for the activity coefficient*of 0.01 M HCl, y = ,924 we obtain: Hz HCl, HgCl I Hg : E" = . 2 7 0 0 (5) Since all the factors of equations (3) and ( 5 ) are of activity one we may write them in the form Fe I Fe++ j / C1-, HgClj Hg E " = . 7113 H, I H+ C1-, HgCl I Hg E"= . 2 7 0 0 Combining these two equations gives ~
Fe Fe++, H+ I Hz E O = 4413 (6) Since (HZI H+) is taken as zero, we arrive a t the value Eo= .4413 for the normal electrode potential of iron. The reliability of this value will be discussed further after considering the results obtained by equilibrium measurements which are to follow. ~
Equilibrium Measurements Since the potential of the thallium electrode is near that of iron, it seemed possible to find some equilibrium that involved mutual oxidation and reduction of these metals. \Vith data already at hand the equilibrium Fe+2 TlCl=Fe+++z Cl-+z T1 would permit an easy calculation of the iron electrode potential provided the concentration and activity of the ferrous chloride could be determined. Preliminary experiments showed that the reaction would proceed in either direction although the rate was quite slow. From these observations, a one-liter flask was fitted with four IOO cc containers connected to the neck so that samples could be withdrawn for analysis. T o the flask were added 40 grams of thallous chloride, 8 grams of powdered reduced iron, and 700 cc of water. The filling was carried out under an atmosphere of hydrogen to prevent any oxidation. The vessel was then evacuated and sealed from the air. This mixture was placed in a thermostat a t 2 j"C and continually agitated by a mechanical shaker. Lewis, Brighton and Sebastian: J. Am. Chem. SOC.,39, 2245 (1917). Chapter XXVI (1923).
* L r w s and Randall: "Thermodgnzmics,"
988
W. H. HAMPTON
The sample tubes were removed from time to time and portions of the solution were rapidly weighed out nnd added to dilute hydrochloric acid, A titration was then made with .05 N potassium permanganate, a method used by A. A. Xoyesl for the quantitative determination of thallium. The end point was somewhat masked by the yellow color of ferric iron but by having a comparison solution present it could be judged within two drops of .05 N potassium permanganate. Check experiments showed this method to be accurate within 0.5 per cent. This gave the total number of equivalents to oxidize both the iron and the thallium. The total chloride ion content was determined in a second portion of the solution acidified with nitric in-
stead of hydrochloric acid. From a result of these two determinations we can calculate the separate concentrations of ferrou? chloride and dissolved thallous chloride from the relations: Fe+-++Tl+= total concentration of C1Fe+++ zT1+= total no, of equi. K M n 0 4
z
I t was found that the solubility of thallous chloride in solutions of ferrous chloride is practically the same as in solutions of barium chloride of like concentrations. The curve in Fig. I represents the solubility of thallous chloride in varying concentrations of barium chloride2 and the crosses that in ferrous chloride. B y the use of this curve and the titration with potassium permanganate the composition of sample portions could be determined with sufficient accuracy for calculations which are to follow later. Table V gives a series of analyses made upon the mixture described above, and it shows how slowly the reaction proceeds toward equilibrium. Most of the reaction took place within the first nine days, at the end of which time 2. physik. Chem., 9, 6c3 (1892). Brajr and Winninghof: J. Am. Chem. Soc., 33, 1663 (1911).
POTENTIAL O F THE IRON ELECTRODE
989
the concentration of fercous chloride became .0306 M and only increased to .0453 M after a period of three months. When all the sample flasks had been removed from the larger vessel, the remaining contents were removed with as little exposure to air as possible and placed in a smaller container. The last two analyses were made on portions taken from this vessel when the amount of solution became exhausted. TThile these analyses indicate equilibrium, it is doubtful whether this point was yet reached for the shaking had been intermittent during the last period. The rate of reaction of this type depends very largely upon the agitation. There was no exception to the behavior of iron in these experiments of showing apparently different activities as in the electrode measurements. Several mixtures in smaller amounts were made up and these behaved differently. After equal intervals of time, in some mixtures there appeared to be no reaction a t all, while varying amounts of thallium mere formed in others. None of these however appeared to be more reactive than the one for which Ihe analysis is given in Table V. There was even more difficulty in approaching the equilibrium by the reduction of ferrous chloride with metalic thallium, and, while the reaction is known to proceed, the rate is very slow. The speed of such a reaction necessarily depends to n large extent upon the amount of metalic surface exposed
TABLE V Mixture made up with 40 grams TlC1, 8 grams iron and Time of analyses o days 9 ” 21
31 60 83 90 I02
,, ,’
Equi. of KMn04 per 1000gm. sol.
-
” ” ”
,,
.0524 .053I ‘0533
,0308 ,0333 ,0388 ,0416 043 5 ’
’0057
I’
.07IO
-
-
,0042
,0453
Equilibrium Measurements, FeC12
22
’ 0053 ’0047 ,0046 ’ 0043
.
TABLE VI Sample Number I . Equi. of KRlnOl Time in I O C O grams sol. o days ,0950 8 ,’ .0865 I 5 71 .0768
cc water
Molal cone. TI C1 ,01607
Molal cone, Fe C12
.0423 .0437 ,0478
700
+ T1 2
Sample Number 2. Equi. of KMn04 ‘rime in 1000 grams sol. o days .IO10 6 ” .0999 14 ” .0968 2I ” .0942 28 ” ,0920
990
U'. H. HAMPTON
and, for a soft metal like thallium, it is hard to produce a finely divided condition. Experiments were tried with small shavings of metallic thallium and a solid material from a saturated thallium amalgam. These gave only more or less qualitative results and nothing definite enough to be used in calculations. Later a method of approaching equilibrium from this direction was tried but time did not permit carrying this work to completion. Mixtures were prepared by reducing thallous chloride with excess iron powder and undissolved salt. The thallium separates out as small particles and remains for the most part in this condition. Table VI gives the results of two experiments on such mixtures beginning with one tenth molal ferrous chloride. Sample No. I was the residue from the run given In Table V and No. 2 was that of a separately prepared mixture. The latter was in a flask and did not permit mechanical agitation and was only shaken from time to time by hand.
It is evident that neither sample had reached equilibrium when the last analyses were made. At this time the solution over the mixture in Xo. I became exhausted and the work was discontinued. Calculation of Results Equation (7) can be expressed in the following cell reaction: Fe FeC12 (aq), TlCl I T1 ~
the e. m. f. of which is given by equation
(2)
E = E " - .08873 log. 1.j88ym Assuming equilibrium when the last analysis was made of the mixture given in Table V, we then have E"=.08873 log. 1.588y .0453 (8) To determine y we take the activity coefficient of barium chloride of the same ionic strength as the solution of ferrous chloride and thallous chloride which corresponds to a molality of ,0467. The activity coefficient for this concentration is ,j74. Sctbstituting this value for equation (8) and solving for E" we obtain FeIFeClz (aq), TlCllTl : E " = - , 1 2 2 8 (9) Combining this with the equation' T1 I TIC1, HgCl I Hg : E = A246
we obtain FeIFeCL (as) HgClIHg : Eo=.70zz Considering the equilibrium now as approached from right to left and making similar calculations for the last analysis of samle Xo. I given in Table VI, we have FeIFeClz(aq), HgCljHg : E O = . 7 1 3 5 Gerke: J. Am. Chem. SOC., 44, 1684 (1922).
POTEKTIAL O F T H E IRON ELECTRODE
99 1
Discussion of Results
It has been shown that the reaction expressed by equation KO. 7 has been made to approach equilibrium in both directions. Assuming equilibrium a t the times of the last analyses, the E" values for equation (3) differ by .0113 volts, but it is certain that in neither case had the reaction ceased. Furthermore, all the observed values of Table IT7lie between these two. The following values have been given for comparison: (a) From the equilibrium measurement: zTIC1+Fe:Eo= . 7 0 2 2 (b) From electromotive force: E"= . ? I 1 3 (c) From the equilibrium measurement: FeC12+ zT1: E" = . 7 135 By examination of Table V and VI we would more likely expect the value of (c) to be too high and that of (a) to be too low. Although predictions of this kind may seem a little questionable when so many factors have to be considered as those involved in this equilibrium, nevertheless electromotive force measurements on iron in a finely divided condition, whether the amalgam variety or the reduced oxide, show quite conclusively that the value of E" can be but very little if any lower than that given for (b). I n round numbers, then, the normal electrode potential is taken as .++IO volts for all ordinary forms of iron, a value which is in all probability correct within three millivolts. Summary
'
The potential of several varieties of iron was measured in the cell Fe FeClz (aq), HgCl I Hg in which there was no liquid junction. An amalgam variety prepared by electrolyzing a solution of ferrous sulfate, using a mercury cathode, gave .go90 volts as the best equilibrium value in 0. I M ferrous chloride. Finely divided iron prepared by reducing Fez03gave ,8088 volts as the best equilibrium value in 0.1M ferrous chloride. Potentials of annealed iron in water solutions were much lower than those obtained with the other varieties. By measuring the potential between finely divided iron and the annealed form in a solution free from water and consisting of ferrous chloride in a fused eutectic mixture of lithium and potasPium chloride it was found that there was very little difference between the activity of each. The difference in water solutions was attributed to passivity of compact iron in contact with water, The value .4413 volts was calculated from the above measurements as the most probable value of the normal electrode potential, and it is not considered in error by more than three millivolts. Measurements made on the equilibrium Fe+ a TlCl= Fe+++z C1-+ a T1 substantiated this conclusion.
