The Reduction of Ferric Oxide
to
Iron
JANE NASH
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Northwestern
University, Evanston, Illinois
OF THE most familiar experiments1 in general ONE chemistry is the determination of the equivalent weight of a metal by reduction of its oxide in hydrogen. At Northwestern the metal oxides are issued to the students in the form of unknowns. Very satisfactory results have been obtained with the oxides of copper and of nickel, but students often make large errors in the reduction of ferric oxide. This has led to the supposition that there must be some ambiguity in the laboratory directions. The purposes of this work were to find directions which would give results with a minimum of error, and at the same time to examine the products formed during the course of the reduction. It is well known that the reduction of ferric oxide passes through several stages, involving, perhaps, solid solutions of the various reduction products. Emphasis was, therefore, placed on the time and temperature of heating of the ferric oxide, C, p. ferric oxide was first weighed, strongly heated, then weighed again to determine its moisture content, which proved to be negligible. A sample of the oxide was then weighed in a porcelain boat which could be strongly heated in an atmosphere of hydrogen. Temperatures were determined by a single junction chromel-alumel thermocouple. Final weighing of the porcelain boat and reduction product was sometimes done in a hydrogen-filled test tube because it was discovered that when the oxide had reached certain stages of reduction it reoxidized immediately on exposure to air. In the first test to be described, the oxide was reduced under conditions paralleling as closely as possible the standard laboratory directions for this experiment. About 0.5 g. of ferric oxide was heated in a stream of hydrogen at about 50U°C. for five minutes. During this time the oxide quickly turned black and then, more slowly, just began to turn gray. It was cooled in hydrogen, then removed for weighing. The moment that the boat was removed from the hydrogen atmosphere the material began to change color to a brownish orange. The boat became very hot. This material weighed almost the same as the original ferric oxide. One important source of error in this experiment now becomes clear. At certain stages of reduction the material suffers spontaneous reoxidation in air. Students Selwood, This Journal, 19, 375 (1942). 1
46
in
hurry to complete this experiment have sometimes been criticized for trying to weigh the porcelain boat and contents while they are still warm. In at least a
cases, it
becomes evident, the boat
cool before being removed from the hydrogen but spontaneous reoxidation took place immediately on exposure to air. In the next experiment the oxide was heated twice as long at the same temperature. During the first few minutes the powder turned a deep black. Then it gradually changed to a fairly light gray. The loss of oxygen was 30.8 per cent, compared with a theoretical oxygen content of 30.00 per cent in Fe203. The reduction product, which must have been practically pure iron, did not change in the least when it was gently heated in air. The main purpose of this work had, therefore, been achieved. The complete reduction of ferric oxide may be obtained by extending the reduction period to about ten minutes at about 500°C. It is essential that the first black reduction product be completely converted to gray. These directions apply to a 0.5 g. sample of oxide. Several tests were now made in an effort to learn more of the composition of the pyrophoric and partially reduced oxide. The products were weighed in hydrogen. The results are summarized in Table 1. some
now
TABLE Test 1
2 3 4 5
% Loss of 7.47 1.02 21.6 8.87 14.U
was
i
Activity toward Keoxidatioti O1
in Air
Color
Slightly
gray
Slightly
gray
Black
Black Black
Spontaneous on gentle heating Spontaneous at 180° C Sj)ontaneous on gentle healing Spontaneous at 100°C. Spontaneous at 80°C.
It is clear that activity toward re oxidation is found over widely differing oxygen content. All the partially reduced oxides were attracted to a magnet, and in some 'cases the reoxidized product was also ferromagnetic. We are unable to state precisely what condition leads to maximum pyrophoricity. Metallic iron appears not to be responsible. Possibly ferrous oxide is the pyrophoric substance and this starts the oxidation of Fe304 which must also be present. The ferromagnetism of the final reoxidation product may be due to y-Fe^Os. This work was done at the suggestion, and under the direction, of Professor P. W. Selwood.
