STABILITY OF FERROUS HYDROXIDE PRECIPITATES
Nov., 1956
1519
STABILITY ‘OF FERROUS HYDROXIDE PRECIPITATES BYF. J. SHIPICO AND DAVIDL. DOUGLAS’ Knolls Atomic Power Laboratory,2 Schenectady, New York Received A p r i l 86,1966
The stability of ferrous hydroxide precipitates a t temperatures up to 316” has been investigated. Pure ferrous hydroxide in contact with solutions of potassium chloride containing either excess ferrous ion or excess hydroxyl ion and with rigorous exclusion of oxygen were found to be completely stable for periods up to six months a t temperatures below 100”. When precipitated with nickel hydroxide the spontaneous reaction 3Fe(OH)2 FeaOl Ht 2H20 takes lace a t ordinary temperatures with an activation energy of 10.2 f 1 kcal./mole. Excess hydroxyl ion corn letely inhiiits the reaction. At temperatures in the range 150-210” pure ferrous hydroxide decomposes according to the agove reaction with a conveniently measurable rate. The activation energy was measured and found t o be 30 i3 kcal./mole. I n this temperature range silica and hydroxyl ion in the supernate inhibit the decomposition. By X-ray structure analysis and microscopic examination a-iron was identified in the magnetite crystals formed in the thermal decomposition of the ferrous hydroxide at 178’ and above.
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I. Introduction Many references to the importance ,of the chemistry of ferrous hydroxide to iron corrosion can be found in the literature.a In addition this compound is thought to play a part in the chemistry of the Edison nickel-iron a c c ~ m u l a t o r . While ~ the room temperature stability of precipitates of pure ferrous hydroxide is fairly well established (see below), the literature contains virtually no data as to its stability above 100”. In 1933 Schikorr5 studied the decomposition of ferrous hydroxide, which he assumed to take place according to the reaction 3Fe(OH)z e Fe304
+ Hz + 2H20
(1)
by measuring the amount of hydrogen evolved from the precipitated solid over long periods of time. I n his experiments air was carefully excluded and the precipitate was not separated from the supernatant solution. Schikorr observed a slow evolution of hydrogen when the supernate contained excess ferrous ion, but in the presence of excess base no reaction occurred. Although some of his experiments ran for as long as 75 days, the maximum yield was about 16% of the theoretical amount of hydrogen calculated from reaction (1). As a result of these studies the reaction (1) is referred to in the literature as “Schikorr’s reaction.’’ Fricke and Rihl,6 in a review of the properties of ferrous hydroxide, reported that the pure white precipitate is oxidized by water a t room temperature. Evans and Wanklyn7 investigated the reaction further in 1948. Using a somewhat more elaborate method of preparing the precipitates of ferrous hydroxide, these investigators found them to be stable a t all temperatures up t o looo, even in the presence of excess ferrous ion in the supernate. However, a reaction resulting in the evolution of hydrogen and the formation of magnetite was found to take place when certain salts were present. (1) Now a t the Chemistry Research Department, Research Laboratory, General Electric Co., Schenectady, New York. (2) Operated b y the General Electric Company for the U. S. Atomio Energy Commission. (3) (a) U. R. Evans, ”Metallic Corrosion, Passivity and Protection,” Edward Arnold and Co., London, 194G, p. 209; (b) U. R. Evans, Engzneering, ( M a y 8, 1953). (4) (a) J. T. Crennell and F. M. Lea, “Alkaline Accumulators,” Longmans, Green and Company, London, 1928, p. 83; (b) 0. Glesmer and J. Einerhand, 2. Elektrochem., 64, 302 (1950). (5) G. Schikorr, Z. anorg. allgen. Chem., 212, 33 (1933). (6) R. Fricke and S. Rihl, ibid., 261, 414 (1943). (7) U. R. Evans and J. N. Wanklyn, Nature, 162,27 (1948).
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NiS04, PtClz and NazS “catalyzed” the reaction when present in the solution. The finely divided metals platinum, copper and nickel were also effective “catalysts.” Leussing and Kolthoff ,* in a recent study of the solubility of ferrous hydroxide, report that the precipitate remains pure white at room temperature for periods of time up to several weeks. The decomposition of the dried precipitate a t 200” was observed by Goldbergg in 1914. Huttig and Moldner’O obtained similar results in a more thorough study in 1931. I n the present paper we report the results of a study of the thermal decomposition of slurried ferrous hydroxide a t temperatures up to 316’. This covers the temperature range of interest in considerations of the mechanism of corrosion of iron in high-temperature water. I n addition, we describe a brief study of the very interesting “catalysis” of this reaction by coprecipitated nickelous hydroxide. 11. Experimental Pure white ferrous hydroxide was precipitated by mixing ferrous chloride and potassium hydroxide solutions in the absence of oxygen. These solutions were prepared by the method of Evans and Wanklyn,’ Le., by the electrolysis of a one molar potassium chloride solution using an iron anode (99.99’% Fe) and a platinum cathode. The water used to prepare the potassium chloride was purified by passing it through a bed of deoxygenating resin and then a bed of deionizing resin.Il The oxygen content of the water was O% NiClz-Ni(OH)r precioitated first: exot. 80. 10% NiClp-Fe(OH)~Drecioitated . before addition of ‘NiClzi expi. 89, 50% ’NiCIi..
TABLE I
EVOLUTION OF HYDROGEN FROM FERROUS HYDROXIDE COPRECIPITATED WITH 10% NICKEL RATECONSTANTS FOR Temp., OC.
THE
k, CC. H P AP/AL.O
mm./min.
V , cc. ; : 8 ‘
ZidO*
log k
-1.932 55.6 1.17 0.1 0.16 0.95 -2.024 55.3 0.1 0.13 5.43 -1.265 24.95 1.05 39.3 47.0 6.93 -1.159 24.9 1 12 -1.300 64.5 5.01 24.85 0.59 0.50 56.0 3.68 -1.434 24.8 -0.006 54.8 12.4 40.4 1.72 9.34 -1.025 2.75 25.8 40.4 21.7 -0,664 54.7 3.74 44.1 -0.562 52.2 27.4 54.7 3.99 a Pressures have been corrected to 273°K. and corrected for vapor pressure of water.
IO
40
50
80
70
60
TIME -MINUTES.
90
Fig. 3.-Effect of temperature on the rate of hydrogen evolution from ferrous hydroxide coprecipitated with 10% nickel hydroxide (0.0003 mole NiClz 0.003 mole FeCln precipitated with 0.006 mole KOH).
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The reaction was inhibited, but not completely so, when the catalyst was added after precipitation of the ferrous hydroxide. Addition of NiClz immediately after the ferrous hydroxide had been precipitated yielded 0.5 cc. of hydrogen in one hour (Fig. 2), while a precipitate that had aged for two hours before the addition of NiC12produced only 0.4 cc. of hydrogen in 20 hours. When ferrous chloride was added after nickel hydroxide had been precipitated, the reaction was retarded also as is shown in Fig. 2. The addition of hydrogen to the I I 32 3~I.+,Xr9. 34 33 36 cell immediately after coprecipitation of nickel and iron hydroxides had no affect on the yield. Fig.4.-Arrhenius plot of nickel “catalyzed” decomposiAn attempt was made to increase the yield of tion of ferrous hydroxide to magnetite and hydrogen; 0.0003 mole NiCls + 0.003 mole FeCI?; coprecipitated with the reaction by homogeneous pre~ipitation‘~ of 0.006 mole KOH; units of IC are cc. (i3TP) of Hz per min. the ferrous hydroxide and nickel hydroxide. I n the first experiment the anode compartment solu- which the precipitate became greenish to white. tion was added to a cell containing 0.5 g. of urea Precipitation continued for some hours with no and the usual 10% (0.0003 mole) of NiC12. After further magnetite or hydrogen being formed. It is sealing, the cell was heated in a water-bath a t estimated that the magnetite yield was about 5% 90-100”. Precipitation accompanied by hydrogen of the total iron. Increasing the proportion of evolution began after about one hour, the first nickel to 50% in a second experiment produced precipitate being black and magnetic. This be- substantially the same result, the final gas being havior continued for approximately an hour after composed of 0.6 cc. of Hz and 28 cc. of Con. The ferrous hydroxide Drecipitate was quite dense and (14) L. Gordon, Anal. Cham., 24, 459 (1952). I
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F. J. SHIPKOAND DAVIDL. DOUGLAB
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Vol. GO
settled quickly-suggesting homogeneous precipitation as a good method of preparing this solid. Thermal Decomposition.-The results obtained in the study of the decomposition of ferrous hydroxide a t elevated temperatures are shown in Figs. 5 and 6. Below 100" the evolution of
energy of 30 kcal./mole. The lack of precision in the results makes the uncertainty in this energy equal to about 3 kcal./mole. Figure 6 summarizes some experiments carried out t o study the possible effect of incomplete precipitation on the rate of decomposition of ferrous hydroxide slurries. I n addition, results showing the marked inhibiting effect of silica are given. At the lower temperatures an excess of either 20 hydroxyl ion or ferrous ion remaining in solution 18 inhibits the decomposition markedly. Very likely 16 this accounts for some of the spread in the results 14 of the thermal decomposition experiments, since \a.1 2 loss of a few per cent. of either the FeClz or KOH ry' 10solutions was possible by holdup in the glass system. At the higher temperature, 316", excess ferrous ion 8does not appreciably affect the decomposition. 6I n another experiment the presence of an atmos4phere of hydrogen rather than helium in the auto2clave did not inhibit the reaction. / I l l 24 26 28 x7 32 34 36 38 40 42 4 4 Solid reaction product from the high temperature Fig. 5.-Thermal decomposition of ferrous hydroxide at experiments was collected by filtering and washing elevated temperatures. The curves show total hydrogen and an X-ray powder diagram (Debye-Schemer) (cc. STP) evolved as a function of time and temperature. made. A full interpretation of the patterns could be given in terms of magnetite alone for runs a t temperatures of 168" and below and in terms of magnetite plus small amounts of or-iron for reactions carried out a t 178" and above. I n order to I confirm the presence of iron, photomicrographs were made of polished sections of the well formed magnetite crystals mounted in Lucite. A few ferrite grains were found among the magnetite crystals. I n all cases the ferrite was in contact with magnetite. This interesting phenomenon will be discussed in the next section. The thermal decomposition of dried Fe(OH)2 2 6 8 10 12 14 16 18 20 22 24 26 28 30 32 34 36 YJ 10 42 44 was studied briefly. A ferrous hydroxide precipiHWRS. Fig. 6.-Thermal decomposition of ferrous hydroxide at tate was collected on a glass frit in an evacuated high temperatures; effect of silica, ferrous ion and hydroxyl cell, washed with water and vacuum dried. The ion. precipitate did not decompose a t a measurable hydrogen was undetectable, while at temperatures rate below loo", but a t temperatures of 150-200" above about 220" the reaction went to completion hydrogen and water were readily produced. The too rapidly to be followed conveniently. As can solid reaction product was magnetite. X-Ray be seen in Fig. 5, the reaction rate at 178 and 208" analysis failed to show the presence of iron in the decreased with time. The same may be true a t magnet,ite. the lower temperatures but no long term experiIV. Discussion ments were carried out to investigate this point. From the results obtained in this study as well I n this plot the data points represent two or more as those of previous investigators it is clear that experiments a t each temperature except the lowest, the decomposition of ferrous hydroxide results in 105" and the highest, 316". The apparent dis- the formation of magnetite and hydrogen. This placement of the start of the reaction from zero is to be expected from the thermodynamics of the time is due to the time required to heat the auto- system. Standard free energy changes correspondclave to temperature in the oven. This time ing to various possible decomposition reactions of amounted to from 1 to 1.5 hours depending on the ferrous hydroxide are given in Table 11. These temperature. It is likely that some of the spread were calculated from available standard free enerin the data is due to this limitation in the method. gies of formation of the compounds.1ss16 Lack of The low yield a t 316" probably was caused by a reliable heat capacity data prevented calculation loss in preparation of the ferrous hydroxide slurry. of the free energy changes a t higher temperatures, Making use of the data given in Fig. 5 a rough although rough estimates indicated that at 500°K. determination of the activation energy of the they are still negative by a few kilocalories. reaction was made. Rates of hydrogen evolution It appears certain that ferrous hydroxide prewere estimated graphically. I n the case of the cipitates can be kept indefinitely at temperatures two highest temperatures, 178 and 208", the rate (15) K. H. Kelley, "Selected Values of Chemical Thermodynamic in the initial stage of the reaction was used. An Properties," U. S. Bureau of Standards Circular No. 500 (1949). Arrhenius plot made using these slopes and the (16) C. J. Smithells, "Metals Reference Book," Interscience Pubcorresponding temperatures yields an activation lishers Inc., London, 1955, p. 676. I
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STABILITY OF FERROUS HYDROXIDE PRECIPITATES
Nov., 1956 TABLEI1
STANDARD FREEENERGY CHANGES OF REACTIONS OF FERROUS HYDROXIDE AGO, kcal.
Reaction
+ +
+ + + +
(1) BFe(0H)dc) tFeaOdc) Hz(d 2HzO(1) (2) 2Fe(OH)dc) FezOa(c) Hdg) HzO(1) (3) Fe(0H)dc) FezOdc) tFeaO4c) HzO(1) (4) 4Fe(OH)z(c) FeaOdc) Fe(c) 4H20( 1)
+
+
298’K. (25’)
-9.9
-3.0 -6.9
-7.7
below 100” if oxygen is rigorously excluded, and if the iron is sufficiently pure. One is inclined to ascribe Schikorr’s original results, in which ferrous hydroxide very slowly decomposed partially to magnetite and hydrogen, to an impure iron. At this time it is possible to say very little about the actual mechanism of the reaction which takes place in the presence of nickel (copper and cobalt). I n view of the results of the homogeneous precipitation experiments, in which the formation of magnetite ceases even though ferrous hydroxide is continuously precipitated in the presence of nickel ion, it seems unlikely that the effect of nickel is one of a catalyst in the usual sense of the word. For this reason quotation marks have been used in this paper. A possible explanation of the peculiar low yields in this reaction is that it is very sensitive to p H . The marked effect of excess base in completely inhibiting the reaction supports this hypothesis. This would require that the reaction be heterogeneous, i.e., the reaction take place on the surface of the precipitate with some constituent of the solution participating. The evidence suggests that coprecipitation of the iron and nickel as hydroxides is necessary for the reaction. Adding the nickel chloride after precipitation of the ferrous hydroxide results in some reaction as nickel ion is adsorbed on the surface yielding a coprecipitated iron-nickel hydroxide. It is possible that the peculiar aging effects noted in other reactions of ferrous hydroxide” may be involved in the low yields of the “catalyzed” reaction. The results of the homogeneous precipitation experiments seem to contradict this but may prove not to do so when more is known about the exact changes taking place in the solution during precipitation. The thermal decomposition of the pure ferrous hydroxide slurries at high temperatures shows some of the same features as the room temperaturenickel “catalyzed” reaction, e.g., excess hydroxyl (17) 0. Bauditsoh and L. A. Welo, Naturwissenschajlen, 21, 659
(1933).
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ion inhibits both. The inhibiting effect of silica may be due to a formation of ferrous silicate on the surface of the hydroxide or to an increase in p H resulting from the dissolution of the glass. In any event the thermal reaction appears to be heterogeneous also. It is instructive to consider these results in the light of previous studies of the corrosion of pure iron in high temperature water.13 I n these it was found that a t 240” the corrosion of iron is described by a linear law, indicating that an interface reaction is rate controlling. This reaction may well be the decomposition of ferrous hydroxide to magnetite and hydrogen. One method of establishing whether this is the case would be to measure the activation energy of the corrosion reaction and compare it with that of the thermal decomposition of ferrous hydroxide. This would not prove that Schikorr’s reaction is involved in iron corrosion, but would make it very likely. The iron found in the magnetite produced by the thermal decomposition of ferrous hydroxide indicates that the reaction 4Fe(OH)a @ Fe
+ FeaOc + 4Hz0
is proceeding simultaneously. That this reaction may take place at appropriate temperatures was suggested by Pourbaixll* his speculation being based on work by Thiesse.lg The decomposition of ferrous oxide to magnetite and iron, viz.
+ Fe
4Fe0 tFea04
has been studied by Chaudron and co-workers.Z0 They report the transformation to take place in the temperature range 250400”. In the course of a study of the chemistry of ferrous hydroxide a t Battelle Memorial Institute the formation of ferrite in the course of the thermal decomposition has been observed This reaction merits further study t o elucidate the effects of temperature, time of exposure, method of preparation and purity of the ferrous hydroxide and of dissolved salts, The assistance of the KAPL Metallurgical and Chemical Analysis Activity in performing various analyses is gratefully acknowledged, and particularly the assistance of C. F. Pachucki and L. R. Yetter for mass spectrometer analyses. We are also indebted to Mr. L. M. Osika of Physical Met,allurgy, Technical .Department, for the X-ray structure analyses. (18) M . J. N. Pourbaix, “Thermodynamics of Dilute Aqueous Solutions,” Trans. J. N. Agar, Edward Arnold and Co., London, 1949,p. 89. (19) Thiesse, Thesis, Nancy, 1937. (20) R. Collongues and G. Chaudron, Compt. rend., 234, 728 (1952); R. Collongues, R. Sifferlen and G. Chaudron, Reu. Met., SO, 727 (1953). ( 2 1 ) Private communication from Dr. P. D. Miller, Physical Chemistry and Corrosion Technology Division, Battelle Memorial Institute, Columbus, Ohio.
