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Thermodynamics of the reduction of iron oxide powders with hydrogen
J. Chem. Thermodynamics 1972, 4, 57-64
Thermodynamics of the reduction of iron oxide powders with hydrogen a H. RAU Philips Forschungslaboratorium Aachen GmbH, Aachen, Germany (Received 5 July 1971) Static type measurements are described of the reduction equilibria of the iron oxides with hydrogen carried out in a sealed apparatus provided with a palladium membrane. By this means the H2 pressure could be measured during the reaction. A quartz spiral gauge enabled the total pressure (Ha + H20) to be measured at the same time. Such an apparatus could also be used to advantage for similar measurements on other oxide systems. The results show no measurable influence of the particle size or of doping with tin on the reduction equilibria. During the heat treatment an appreciable growth of the particles takes place, accompanied by a strong reduction in the reaction velocity. From the results the following thermodynamic quantities were calculated: AH°(Fe304, 298.15 K) ~ - 266.7kcalth tool-1, S°(FeaO4, 298.15 K) -----35.03 calth tool -1, AH~(ct-FeaO3, 298.15 K) ------ 197.1 kcalth mol-1, and the eutectoid temperature of FeO was 570 °C.
1. Introduction Thermodynamic data for the iron oxides have been the subject of numerous investigations by equilibrium or calorimetric measurements, especially from the metallurgical point of view. tl} On the other hand, reduction with hydrogen plays an important role in the preparation of magnetic recording materials (tapes). These reductions are often performed at relatively low temperatures and on very fine powders, so that the question arises whether the published thermodynamic data can be extrapolated to these experimental conditions without loss of accuracy. In this study, static type measurements are described which were performed with a sealed apparatus provided with a palladium membrane, by which the H 2 partial pressure could be measured during the reduction process. A quartz spiral gauge enabled measurements to be made of the total pressure (H 2 + H 2 0 ) at the same time. The results show no measurable influence of the particle size or of doping with tin on the equilibrium constants. During the heat treatment, growth of the particles occurs, regardless of their individual oxygen content, accompanied by a strong decrease in the reaction velocity. So, after treatment at higher temperatures the a Presented at the International Symposium on Metallurgical Chemistry, July 1971, Brunel University, U.K.
58
H. RAU
measurements at lower temperatures (300 to 400 °C) could not .be repeated. The mechanism of this particle growth is still unclear.
2. Experimental PREPARATION OF THE SAMPLES For the preparation of powders of iron oxide particles ofacicular shape (for magnetic recording tapes) ~-FeOOH is a suitable starting material. This material can be converted into other compounds (e.g. v-Fe203) by pseudomorphic reactions. Particles of ~-FeOOH of length 0.1 t o 2.0 t t m a n d diameter o f about 1/10 of the length can be prepared by a crystallization and oxidation process of iron hydroxides in aqueous solution.(2) To a solution of FeSO4 • 7H20 in water, acidified with H2SO4 to a pH of about 2, a solution of NaOH in water (0.25 to 2.5 M) is added.'f The suspension of Fe(OH)2 is then treated with oxygen or air at room temperature until no bivalent iron is present. The suspension is filtered, washed with water and acetone, and dried a t 60 °C. By changing the reaction conditions the shape of the particles can be controlled. The ~-FeOOH was dehydrated at a temperature of about 300 °C in vacuum. Some samples were at the same time reduced to Fe30 4 or Fe with hydrogen at about 310 °C. Since fine particles of iron are pyrophoric, these samples were treated with oxygen under reduced pressure or with alcohol at room temperature. When alcohol had been used, the organic residue had to be removed before the equilibrium measurements by treatment with hydrogen at 350 °C for some days. The ~-FeOOH powders were not very pure. They contained some contaminations of about 0.01 mass per cent, e.g. Mn, 0.05; Mg, 0.03; Ni, 0.03; Na, 0.03; and Si, 0.08 mass per cent. One sample contained no tin, the other was doped with 0.36 mass per cent of Sn. The samples had different oxygen contents corresponding to the two-phase regions between Fe and FeO, FeO and Fe304, and Fe304 and Fe203, respectively. EXPERIMENTAL PROCEDURE The samples (about~ lg) were added to the apparatus shown in figure 3 and then baked in-vacuum at 300 °C for some hours to remove any water. The apparatus consisted of two parts completely separated by a palladium tube (4 mm x 0.2 ram) closed at one end by welding. The inner part, made of quartz glass, was connected to a quartz spiral(differential) gauge. Thus the total pressure (H2 + H20)could be measured by counterbalancing this gauge by. air pressure on the outside of the spiral. From the outer part pure hydrogen could be admitted or drawn away via the palladium membrane, the pressure being measured by~a mercury manometer or a McLeod gauge. The partial pressure of the H 2 0 above the sample was therefore calculated as the difference between the pressure readings in the inner and-the outer part, of the apparatus. Since the partial pressure of hydrogen pressure in the case of the Fe203 + Fe304 equilibrium was in the range of several cTorr, any water desorbing t Throughout this paper M = mol dm-8; Tort = (101.325/760) kPa; atm = 101.325 kPa; caleb = 4.184 J.
F I G U R E 1. Powder of u-FeOOH, the starting material.
F I G U R E 2. The same powder as in figure 1 after the equilibrium measurements. Oxygen mole fi'action approximately 0.30. lfacing p. 58
REDUCTION OF IRON OXIDE WITH HYDROGEN
59
pure hydroge --------~/dt McLe°dI I mercurypressure glass-to-metal ~-~'-~!] gauge ~_._]/gauge seal d ~ ~ ~ hard soldere l~ quartz-to-glass seal -@.~ heatingwire .j ~
~diffusionpump ~ m o l e c u l a r sieves
~ ,,- thermocouple palladiumtube 1(~ ~ ] ~ l - - ~ closedby ~ ~[~ ~[I [I~I~PI welding )1"~_~.1[lt~'-~[,,| I~ \V IH/ reaction
/f~
1~~ temperature
tomercury /~--J~/ [ pressure ~ . Ht IQ~'~_ ~ 4 quartzspiral_
~ ().J~
]
di erent,al
FIGURE 3. The apparatus used for the equilibriummeasurements.
from the glass walls or from the stopcock grease had to be removed. This could be achieved by the use of molecular sieves, pre-baked in vacuum at 300 °C. Temperatures were measured by a Ni-to-NiCr thermocouple previously checked at the freezing temperatures of Ag, Sb, Zn, and Sn. At nearly each temperature the equilibrium was measured starting from both an excess and a deficiency of hydrogen compared with the equilibrium value. The hydrogen was admitted to the outer part and diffused into the inner part via the palladium tube. The total pressures varied between 300 and 800 Torr. During the runs, the apparent equilibrium constant at different times was calculated and plotted against time. The apparent equilibrium constant is defined as the ratio of the apparent partial pressure of H20 (= total pressure minus partial pressure of H2) and the partial pressure of H2 at the same time. This ratio approaches the true equilibrium constant after sufficient time. At the lower temperatures, e.g. below 400 °C for the Fe + Fe304 mixtures, complete equilibrium could not be reached even after some days. There still remained some difference in the equilibrium constants calculated from the measurements starting with hydrogen excess or deficiency. At temperatures above 600 °C the reaction was very fast and reaction times of some minutes were sufficient. The diffusion of hydrogen through the palladium tube seems not to be the limiting factor, the diffusion being sufficiently fast even a t 300 °C.
60
H. RAU
In the case of the Fe20 3 -k FeaO4 mixture, complete equilibrium could not be reached even after 3 d. Because of the low reaction velocity here, only measurements above 500 °C were performed. On the other hand, this reaction is slow only near equilibrium. When the hydrogen content in the reaction gas is high, the reaction is fast even at 300 °C, but it becomes increasingly slower when the hydrogen content is below 0.1 mole per cent. 3.
Results
The results of the measurements on mixtures of Fe and Fe304 are shown in figure 4 and table 1. Open and full symbols mean that the equilibrium was reached starting from lower or higher hydrogen content than corresponds to the final value. The equilibrium studied is given by the equation: ¼Fe304 + H2 = ¼Fe + HzO(g ).
(1)
Figure 5 and table 1 present the results of the measurements on the reaction: 3~-FezOa + H 2 = 2FeaO4+H20(g),
(2)
t/°C
800
700
--I
I
600
500
450
t
l
1
~
400 ......
I
350
300
I
~ - ~ - "
"
\ X
'~FeO " ~ 8
Fe3()4
x~'x= ~v. ~xX~,7°C ..~ %
-o.5
-Fo 0
-1.0
v l
1.0
t
I
................
1.2
I
1.4
__l
I
I.
_
1.6
103 K / T FIGURE 4. Experimental results. Open and full symbols mean that the equilibrium was reached from lower or higher hydrogen contents compared with the equilibrium value. The results of Emmett and Shultz(1°~are included as crosses. O, O, FeO= for 1.2 < x < 1.3, no tin; II, [], FeO~ for 0.4 < y < 0.6, no tin; T, V, FeO~ for 1.2 < z < 1.3, with tin 0.36 mass per cent.
1.8
REDUCTION OF IRON OXIDE WITH HYDROGEN
61
TABLE 1. Experimental results of the equilibrium measurements. The values given represent the last ones measured, i.e. the values found when the partial pressures p(Ha) and p ( H 2 0 ) of H2 and H20 had remained nearly constant over a long period of time (up to 30 h). They are upper or lower bounds for the true equilibrium constants Kp at that temperature T for the measurements started either with water or hydrogen excess, respectively. Keynumbers: I, 2, 3, 4, 5, Sign:
3Fe2Oa + H2 = 2FeaO4 + H20, undoped; ¼FeaO4 + H2 = ~Fe + H~O, undoped; same as 2, doped with 0.36 mass per cent of Sn; ½FeaO4 + H2 = " F e e " + H20, undoped; " F e e " + Ha = Fo + HaO, undoped;
+ , measurements started with hydrogen excess; - - , measurements started with hydrogen deficiency.
T
,o(1-120)
pffI~)
K
Torr
Torr
767 767 810 810 840 840 583 583 583 611 612 641 641 683 684 730.5 731 790.5 790.5 839 839 810 810 573 624 621 653 707 712 750 750 750 795 795 882 882 882 859 859 860
790 787.5 835.5 833.5 855 851 447.3 477.6 622.7 577.6 455.1 597.8 456.3 465.4 404.9 533.8 354.7 584.0 369.8 525.4 339.4 463.5 333.6 441.4 493.3 363.0 480.5 487.2 40L1 404.0 561.4 402.7 543.5 439.1 272.5 175.3 507.7 371.3 310.4 448.9
0.024 0.020 0.023 0.0305 0.022~ 0.0262 24.1 26.5 30.0 36.4 34.9 53.9 45.0 61.4 55.2 99.2 68.1 151.0 98.0 172.2 111.0 127.5 96.2 18.6 37.9 27.4 46.4 75.1 65.1 81.9 112.4 84.6 142.0 118.1 139.2 95.2 249.4 145.2 125.2 152.9
logl0Kp
4.517 4.596 4.560 4.437 4.580 4.591 -- 1.269 -- 1.256 -- 1.317 --1.201 --1.113 --1.045 --1.006 --0.879 --0.865 --0.730 --0.716 --0.588 --0.577 --0.485 --0.486 --0.561 --0.541 --1.375 --1.114 --1.123 --1.015 --0.812 --0.789 --0.694 -- 0.698 --0.678 --0.583 --0.573 --0.291 --0.268 --0.307 --0.407 --0.394 --0.467
Keynumber
1 1 1 1 1 1 2 2 2 2 2 2 2 2 2 2 2 2 2 2 2 2 2 3 3 3 3 3 3 3 3 3 3 3 4 4 4 4 4 5
Sign
+ -q-+ --+ + --+ -----+ -+ --+ + + -+ + -+ + -+ ---+ + -q7
62
H. RAU TABLE 1--continued T
p(H20)
p(H~)
V.
Torr
Torr
860 885 884 920 920 942.5 942 964 966 978
340.9 344.3 290.6 494.9 442.7 445.4 319.4 310.7 274.8 399.5
117.0 126.0 107.4 191.7 171.2 182.5 132.2 133.3 118.1 174.6
Keynumber
Ioglo Kp
5 5 5 5 5 5 5 5 5 5
--0.458 --0.437 --0.432 --0.412 --0.413 --0.387 --0.383 --0.367 --0.365 --0.359
Sign
÷ ÷ ÷ + + q-
t/°C
1000 900 800 t
700
600
500
I
L
T
I
\
-,(N
-5
x\
"\
- Fe203
\
-10
\
-15 I;%O4 -20
I
.1
0.7
0.9
1.1
1.3
103K/T F I G U R E 5. Oxygen dissociation pressures of Fe2Oa. Flengas; m . . . . . , Kodera et al. ~5~
, this work; . . . . .
, Charette and
REDUCTION OF IRON: OXIDE WITH HYDROGEN
63
transformed to partial pressures of 02 using the known equilibrium constants for the formation of water, ca) In the same figure some data from the literature are given. Charette and Flengas ~4) performed e.m.f, measurements with a ZrO2 + CaO electrolyte, Kodera et aL ~5) measured the dissociation pressure directly in a McLeod gauge and analysed the gas evolving by a mass spectrometer. For the calculation of thermodynamic quantities from the results the values of Charette and Flengas were also taken into account. Using the "second-law" method, the equilibrium according to equation (2) gave for the enthalpy of reaction: AH2(298.15 K) = 0.15 kcalth mol-1, and for the entropy of reaction: AS2(298.15 K) = 21.284 calt~ K - 1 mol- 1. From this and heat capacities, enthalpies of transformation, and the standard entropy of 0t-Fe203 given in the literature, ~6' 7) the entropy of Fe304 was calculated to be S°(Fe304, 298.15 K) = 35.03 calth K -1 mo1-1. This value is nearly identical with that given by Westrum and Gronvold ta) calculated from low temperature heat capacities. It was not clear whether a zero temperature entropy of Fe30~ exists. To investigate this, Westrum and Gr~nvold suggested the measurement of reaction entropies such as that of equation (1). On the other hand, the equilibrium of equation (2) is much better for this purpose than equation (1), since the entropy of Fe304 enters the equilibrium constant of reaction (2) with an eightfold weight compared with reaction (1). The entropy of FeaO4 given here is in agreement with the assumption that no zero temperature entropy is present and that the entropy given by Kubaschewski et aLta) is high. Using this entropy of Fe304, the Cp and S(298.15 K) values for the other substances in the equilibrium, and the enthalpy of formation of H20(g), the results of the measurements on equilibrium (1) give the enthalpy of formation of Fe304 as AH~'(Fe304, 298.15 K) = -266.7 kcalth mol-1, which is to be compared with the value: AH~(Fe304, 298.15 K) = -266.9 kcalth mol-1, given by Kubaschewski et aL ~3) From the enthalpy change of reaction (2) the enthalpy of formation of a-Fe203 is found to be AH~(0~-Fe203, 298.15 K) = - 197.1 kcalth mol- 1, while Roth and Wienert <9) found by calorimetric measurements: AH~(~-Fe203, 298.15 K) = -(195.2+0.2) kcalth mol-1. The equilibrium constants measured are about 60 per cent higher than those estimated by Richardson and Jeffes,~x°) who assumed an accuracy of their estimate of about :]:: 5 kcalth mol- 1, corresponding to a factor of about 20 in the equilibrium constants. Since the equilibrium of equation (2) was never measured directly, it seems now to be much better established than before.
64
H. RAU
It is meaningless to evaluate thermodynamic quantities for FeO from the measure. ments above the eutectoid (570 °C), since the composition of FeO in equilibrium with Fe or F%O4 is a function of the temperature itself. The values found agree well with those given by Emmett and Schultz, (~2) although the eutectoid was found somewhat higher (570 °C instead of 559 °C).
4. Conclusions It is shown by the measurements reported that there is no measureable effect of particle size or of doping with tin on the equilibrium constants. Since the particle size changed appreciably during the experiments, the particles being very fine at the beginning and growing gradually much larger, the equilibrium curve measured should show an unusual dependence on the temperature in this case, if particle size had any role to play. In similar studies R. Fricke and G. Weitbrecht (~I) found an influence of the particle size on the equilibrium constants. However, their powders were much finer than those studied here. Thus these different results are not in contradiction. REFERENCES I. Kubaschewski, O.; Catterall, J. A. Thermochemical Data o f Alloys. Pergamon Press: London and New York. 1956. 2. van der Giessen, A. A.; Klomp, C. J. IEEE Trans. Magnet. 1969, 5, 317. 3. Kubaschewski, O.; Evans, E. L.; Alcock, C. B. Metallurgical Thermochemistry, fourth edition. Pergamon Press: London and New York. 1967. 4: Charette, G. G.; Flengas, S. N. J. Electroehem. Soc. 1968, 115, 796. 5. Kodera, K.; Kusanoki, I.; Watanabe, M. Bull. Chem. Soe. Japan 1969, 42, 3036. 6. Coughlin, J. P.; King, E. G.; Bennickson, K. R. J. Amer. Chem. Soc. 1951, 75, 3891. 7. Gr~nvold, F.; Westrum, F. W. J. Amer. Chem. Soc. 1959, 81, 1780. 8. Westrum, E. F.; Gr~nvold, F. J. Chem. Thermodynamics 1969, 1, 543. 9. Roth, W. A.; Wienert, F. Arch. Eisenhilttenwesen 1934, 7, 460. 10. Richardson, F. D.; Jeffes, J. H. E. J. 1ton Steellnst. 1948, 160, 261. 11. Fricke, R.; Weitbrecht, G. Z. Elektrochem. 1941, 47, 487; 1942, 48, 82, 106. 12. Emmett, P. H.; Shultz, J. F. J. Amer. Chem. Soe. 1933, 55, 1376.
Thermodynamics of the reduction of iron oxide powders with hydrogen a H. RAU Philips Forschungslaboratorium Aachen GmbH, Aachen, Germany (Received 5 July 1971) Static type measurements are described of the reduction equilibria of the iron oxides with hydrogen carried out in a sealed apparatus provided with a palladium membrane. By this means the H2 pressure could be measured during the reaction. A quartz spiral gauge enabled the total pressure (Ha + H20) to be measured at the same time. Such an apparatus could also be used to advantage for similar measurements on other oxide systems. The results show no measurable influence of the particle size or of doping with tin on the reduction equilibria. During the heat treatment an appreciable growth of the particles takes place, accompanied by a strong reduction in the reaction velocity. From the results the following thermodynamic quantities were calculated: AH°(Fe304, 298.15 K) ~ - 266.7kcalth tool-1, S°(FeaO4, 298.15 K) -----35.03 calth tool -1, AH~(ct-FeaO3, 298.15 K) ------ 197.1 kcalth mol-1, and the eutectoid temperature of FeO was 570 °C.
1. Introduction Thermodynamic data for the iron oxides have been the subject of numerous investigations by equilibrium or calorimetric measurements, especially from the metallurgical point of view. tl} On the other hand, reduction with hydrogen plays an important role in the preparation of magnetic recording materials (tapes). These reductions are often performed at relatively low temperatures and on very fine powders, so that the question arises whether the published thermodynamic data can be extrapolated to these experimental conditions without loss of accuracy. In this study, static type measurements are described which were performed with a sealed apparatus provided with a palladium membrane, by which the H 2 partial pressure could be measured during the reduction process. A quartz spiral gauge enabled measurements to be made of the total pressure (H 2 + H 2 0 ) at the same time. The results show no measurable influence of the particle size or of doping with tin on the equilibrium constants. During the heat treatment, growth of the particles occurs, regardless of their individual oxygen content, accompanied by a strong decrease in the reaction velocity. So, after treatment at higher temperatures the a Presented at the International Symposium on Metallurgical Chemistry, July 1971, Brunel University, U.K.
58
H. RAU
measurements at lower temperatures (300 to 400 °C) could not .be repeated. The mechanism of this particle growth is still unclear.
2. Experimental PREPARATION OF THE SAMPLES For the preparation of powders of iron oxide particles ofacicular shape (for magnetic recording tapes) ~-FeOOH is a suitable starting material. This material can be converted into other compounds (e.g. v-Fe203) by pseudomorphic reactions. Particles of ~-FeOOH of length 0.1 t o 2.0 t t m a n d diameter o f about 1/10 of the length can be prepared by a crystallization and oxidation process of iron hydroxides in aqueous solution.(2) To a solution of FeSO4 • 7H20 in water, acidified with H2SO4 to a pH of about 2, a solution of NaOH in water (0.25 to 2.5 M) is added.'f The suspension of Fe(OH)2 is then treated with oxygen or air at room temperature until no bivalent iron is present. The suspension is filtered, washed with water and acetone, and dried a t 60 °C. By changing the reaction conditions the shape of the particles can be controlled. The ~-FeOOH was dehydrated at a temperature of about 300 °C in vacuum. Some samples were at the same time reduced to Fe30 4 or Fe with hydrogen at about 310 °C. Since fine particles of iron are pyrophoric, these samples were treated with oxygen under reduced pressure or with alcohol at room temperature. When alcohol had been used, the organic residue had to be removed before the equilibrium measurements by treatment with hydrogen at 350 °C for some days. The ~-FeOOH powders were not very pure. They contained some contaminations of about 0.01 mass per cent, e.g. Mn, 0.05; Mg, 0.03; Ni, 0.03; Na, 0.03; and Si, 0.08 mass per cent. One sample contained no tin, the other was doped with 0.36 mass per cent of Sn. The samples had different oxygen contents corresponding to the two-phase regions between Fe and FeO, FeO and Fe304, and Fe304 and Fe203, respectively. EXPERIMENTAL PROCEDURE The samples (about~ lg) were added to the apparatus shown in figure 3 and then baked in-vacuum at 300 °C for some hours to remove any water. The apparatus consisted of two parts completely separated by a palladium tube (4 mm x 0.2 ram) closed at one end by welding. The inner part, made of quartz glass, was connected to a quartz spiral(differential) gauge. Thus the total pressure (H2 + H20)could be measured by counterbalancing this gauge by. air pressure on the outside of the spiral. From the outer part pure hydrogen could be admitted or drawn away via the palladium membrane, the pressure being measured by~a mercury manometer or a McLeod gauge. The partial pressure of the H 2 0 above the sample was therefore calculated as the difference between the pressure readings in the inner and-the outer part, of the apparatus. Since the partial pressure of hydrogen pressure in the case of the Fe203 + Fe304 equilibrium was in the range of several cTorr, any water desorbing t Throughout this paper M = mol dm-8; Tort = (101.325/760) kPa; atm = 101.325 kPa; caleb = 4.184 J.
F I G U R E 1. Powder of u-FeOOH, the starting material.
F I G U R E 2. The same powder as in figure 1 after the equilibrium measurements. Oxygen mole fi'action approximately 0.30. lfacing p. 58
REDUCTION OF IRON OXIDE WITH HYDROGEN
59
pure hydroge --------~/dt McLe°dI I mercurypressure glass-to-metal ~-~'-~!] gauge ~_._]/gauge seal d ~ ~ ~ hard soldere l~ quartz-to-glass seal -@.~ heatingwire .j ~
~diffusionpump ~ m o l e c u l a r sieves
~ ,,- thermocouple palladiumtube 1(~ ~ ] ~ l - - ~ closedby ~ ~[~ ~[I [I~I~PI welding )1"~_~.1[lt~'-~[,,| I~ \V IH/ reaction
/f~
1~~ temperature
tomercury /~--J~/ [ pressure ~ . Ht IQ~'~_ ~ 4 quartzspiral_
~ ().J~
]
di erent,al
FIGURE 3. The apparatus used for the equilibriummeasurements.
from the glass walls or from the stopcock grease had to be removed. This could be achieved by the use of molecular sieves, pre-baked in vacuum at 300 °C. Temperatures were measured by a Ni-to-NiCr thermocouple previously checked at the freezing temperatures of Ag, Sb, Zn, and Sn. At nearly each temperature the equilibrium was measured starting from both an excess and a deficiency of hydrogen compared with the equilibrium value. The hydrogen was admitted to the outer part and diffused into the inner part via the palladium tube. The total pressures varied between 300 and 800 Torr. During the runs, the apparent equilibrium constant at different times was calculated and plotted against time. The apparent equilibrium constant is defined as the ratio of the apparent partial pressure of H20 (= total pressure minus partial pressure of H2) and the partial pressure of H2 at the same time. This ratio approaches the true equilibrium constant after sufficient time. At the lower temperatures, e.g. below 400 °C for the Fe + Fe304 mixtures, complete equilibrium could not be reached even after some days. There still remained some difference in the equilibrium constants calculated from the measurements starting with hydrogen excess or deficiency. At temperatures above 600 °C the reaction was very fast and reaction times of some minutes were sufficient. The diffusion of hydrogen through the palladium tube seems not to be the limiting factor, the diffusion being sufficiently fast even a t 300 °C.
60
H. RAU
In the case of the Fe20 3 -k FeaO4 mixture, complete equilibrium could not be reached even after 3 d. Because of the low reaction velocity here, only measurements above 500 °C were performed. On the other hand, this reaction is slow only near equilibrium. When the hydrogen content in the reaction gas is high, the reaction is fast even at 300 °C, but it becomes increasingly slower when the hydrogen content is below 0.1 mole per cent. 3.
Results
The results of the measurements on mixtures of Fe and Fe304 are shown in figure 4 and table 1. Open and full symbols mean that the equilibrium was reached starting from lower or higher hydrogen content than corresponds to the final value. The equilibrium studied is given by the equation: ¼Fe304 + H2 = ¼Fe + HzO(g ).
(1)
Figure 5 and table 1 present the results of the measurements on the reaction: 3~-FezOa + H 2 = 2FeaO4+H20(g),
(2)
t/°C
800
700
--I
I
600
500
450
t
l
1
~
400 ......
I
350
300
I
~ - ~ - "
"
\ X
'~FeO " ~ 8
Fe3()4
x~'x= ~v. ~xX~,7°C ..~ %
-o.5
-Fo 0
-1.0
v l
1.0
t
I
................
1.2
I
1.4
__l
I
I.
_
1.6
103 K / T FIGURE 4. Experimental results. Open and full symbols mean that the equilibrium was reached from lower or higher hydrogen contents compared with the equilibrium value. The results of Emmett and Shultz(1°~are included as crosses. O, O, FeO= for 1.2 < x < 1.3, no tin; II, [], FeO~ for 0.4 < y < 0.6, no tin; T, V, FeO~ for 1.2 < z < 1.3, with tin 0.36 mass per cent.
1.8
REDUCTION OF IRON OXIDE WITH HYDROGEN
61
TABLE 1. Experimental results of the equilibrium measurements. The values given represent the last ones measured, i.e. the values found when the partial pressures p(Ha) and p ( H 2 0 ) of H2 and H20 had remained nearly constant over a long period of time (up to 30 h). They are upper or lower bounds for the true equilibrium constants Kp at that temperature T for the measurements started either with water or hydrogen excess, respectively. Keynumbers: I, 2, 3, 4, 5, Sign:
3Fe2Oa + H2 = 2FeaO4 + H20, undoped; ¼FeaO4 + H2 = ~Fe + H~O, undoped; same as 2, doped with 0.36 mass per cent of Sn; ½FeaO4 + H2 = " F e e " + H20, undoped; " F e e " + Ha = Fo + HaO, undoped;
+ , measurements started with hydrogen excess; - - , measurements started with hydrogen deficiency.
T
,o(1-120)
pffI~)
K
Torr
Torr
767 767 810 810 840 840 583 583 583 611 612 641 641 683 684 730.5 731 790.5 790.5 839 839 810 810 573 624 621 653 707 712 750 750 750 795 795 882 882 882 859 859 860
790 787.5 835.5 833.5 855 851 447.3 477.6 622.7 577.6 455.1 597.8 456.3 465.4 404.9 533.8 354.7 584.0 369.8 525.4 339.4 463.5 333.6 441.4 493.3 363.0 480.5 487.2 40L1 404.0 561.4 402.7 543.5 439.1 272.5 175.3 507.7 371.3 310.4 448.9
0.024 0.020 0.023 0.0305 0.022~ 0.0262 24.1 26.5 30.0 36.4 34.9 53.9 45.0 61.4 55.2 99.2 68.1 151.0 98.0 172.2 111.0 127.5 96.2 18.6 37.9 27.4 46.4 75.1 65.1 81.9 112.4 84.6 142.0 118.1 139.2 95.2 249.4 145.2 125.2 152.9
logl0Kp
4.517 4.596 4.560 4.437 4.580 4.591 -- 1.269 -- 1.256 -- 1.317 --1.201 --1.113 --1.045 --1.006 --0.879 --0.865 --0.730 --0.716 --0.588 --0.577 --0.485 --0.486 --0.561 --0.541 --1.375 --1.114 --1.123 --1.015 --0.812 --0.789 --0.694 -- 0.698 --0.678 --0.583 --0.573 --0.291 --0.268 --0.307 --0.407 --0.394 --0.467
Keynumber
1 1 1 1 1 1 2 2 2 2 2 2 2 2 2 2 2 2 2 2 2 2 2 3 3 3 3 3 3 3 3 3 3 3 4 4 4 4 4 5
Sign
+ -q-+ --+ + --+ -----+ -+ --+ + + -+ + -+ + -+ ---+ + -q7
62
H. RAU TABLE 1--continued T
p(H20)
p(H~)
V.
Torr
Torr
860 885 884 920 920 942.5 942 964 966 978
340.9 344.3 290.6 494.9 442.7 445.4 319.4 310.7 274.8 399.5
117.0 126.0 107.4 191.7 171.2 182.5 132.2 133.3 118.1 174.6
Keynumber
Ioglo Kp
5 5 5 5 5 5 5 5 5 5
--0.458 --0.437 --0.432 --0.412 --0.413 --0.387 --0.383 --0.367 --0.365 --0.359
Sign
÷ ÷ ÷ + + q-
t/°C
1000 900 800 t
700
600
500
I
L
T
I
\
-,(N
-5
x\
"\
- Fe203
\
-10
\
-15 I;%O4 -20
I
.1
0.7
0.9
1.1
1.3
103K/T F I G U R E 5. Oxygen dissociation pressures of Fe2Oa. Flengas; m . . . . . , Kodera et al. ~5~
, this work; . . . . .
, Charette and
REDUCTION OF IRON: OXIDE WITH HYDROGEN
63
transformed to partial pressures of 02 using the known equilibrium constants for the formation of water, ca) In the same figure some data from the literature are given. Charette and Flengas ~4) performed e.m.f, measurements with a ZrO2 + CaO electrolyte, Kodera et aL ~5) measured the dissociation pressure directly in a McLeod gauge and analysed the gas evolving by a mass spectrometer. For the calculation of thermodynamic quantities from the results the values of Charette and Flengas were also taken into account. Using the "second-law" method, the equilibrium according to equation (2) gave for the enthalpy of reaction: AH2(298.15 K) = 0.15 kcalth mol-1, and for the entropy of reaction: AS2(298.15 K) = 21.284 calt~ K - 1 mol- 1. From this and heat capacities, enthalpies of transformation, and the standard entropy of 0t-Fe203 given in the literature, ~6' 7) the entropy of Fe304 was calculated to be S°(Fe304, 298.15 K) = 35.03 calth K -1 mo1-1. This value is nearly identical with that given by Westrum and Gronvold ta) calculated from low temperature heat capacities. It was not clear whether a zero temperature entropy of Fe30~ exists. To investigate this, Westrum and Gr~nvold suggested the measurement of reaction entropies such as that of equation (1). On the other hand, the equilibrium of equation (2) is much better for this purpose than equation (1), since the entropy of Fe304 enters the equilibrium constant of reaction (2) with an eightfold weight compared with reaction (1). The entropy of FeaO4 given here is in agreement with the assumption that no zero temperature entropy is present and that the entropy given by Kubaschewski et aLta) is high. Using this entropy of Fe304, the Cp and S(298.15 K) values for the other substances in the equilibrium, and the enthalpy of formation of H20(g), the results of the measurements on equilibrium (1) give the enthalpy of formation of Fe304 as AH~'(Fe304, 298.15 K) = -266.7 kcalth mol-1, which is to be compared with the value: AH~(Fe304, 298.15 K) = -266.9 kcalth mol-1, given by Kubaschewski et aL ~3) From the enthalpy change of reaction (2) the enthalpy of formation of a-Fe203 is found to be AH~(0~-Fe203, 298.15 K) = - 197.1 kcalth mol- 1, while Roth and Wienert <9) found by calorimetric measurements: AH~(~-Fe203, 298.15 K) = -(195.2+0.2) kcalth mol-1. The equilibrium constants measured are about 60 per cent higher than those estimated by Richardson and Jeffes,~x°) who assumed an accuracy of their estimate of about :]:: 5 kcalth mol- 1, corresponding to a factor of about 20 in the equilibrium constants. Since the equilibrium of equation (2) was never measured directly, it seems now to be much better established than before.
64
H. RAU
It is meaningless to evaluate thermodynamic quantities for FeO from the measure. ments above the eutectoid (570 °C), since the composition of FeO in equilibrium with Fe or F%O4 is a function of the temperature itself. The values found agree well with those given by Emmett and Schultz, (~2) although the eutectoid was found somewhat higher (570 °C instead of 559 °C).
4. Conclusions It is shown by the measurements reported that there is no measureable effect of particle size or of doping with tin on the equilibrium constants. Since the particle size changed appreciably during the experiments, the particles being very fine at the beginning and growing gradually much larger, the equilibrium curve measured should show an unusual dependence on the temperature in this case, if particle size had any role to play. In similar studies R. Fricke and G. Weitbrecht (~I) found an influence of the particle size on the equilibrium constants. However, their powders were much finer than those studied here. Thus these different results are not in contradiction. REFERENCES I. Kubaschewski, O.; Catterall, J. A. Thermochemical Data o f Alloys. Pergamon Press: London and New York. 1956. 2. van der Giessen, A. A.; Klomp, C. J. IEEE Trans. Magnet. 1969, 5, 317. 3. Kubaschewski, O.; Evans, E. L.; Alcock, C. B. Metallurgical Thermochemistry, fourth edition. Pergamon Press: London and New York. 1967. 4: Charette, G. G.; Flengas, S. N. J. Electroehem. Soc. 1968, 115, 796. 5. Kodera, K.; Kusanoki, I.; Watanabe, M. Bull. Chem. Soe. Japan 1969, 42, 3036. 6. Coughlin, J. P.; King, E. G.; Bennickson, K. R. J. Amer. Chem. Soc. 1951, 75, 3891. 7. Gr~nvold, F.; Westrum, F. W. J. Amer. Chem. Soc. 1959, 81, 1780. 8. Westrum, E. F.; Gr~nvold, F. J. Chem. Thermodynamics 1969, 1, 543. 9. Roth, W. A.; Wienert, F. Arch. Eisenhilttenwesen 1934, 7, 460. 10. Richardson, F. D.; Jeffes, J. H. E. J. 1ton Steellnst. 1948, 160, 261. 11. Fricke, R.; Weitbrecht, G. Z. Elektrochem. 1941, 47, 487; 1942, 48, 82, 106. 12. Emmett, P. H.; Shultz, J. F. J. Amer. Chem. Soe. 1933, 55, 1376.