Vol. 18, KO.3
INDUSTRIAL A N D ENGINEERING CHEMISTRY
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Oxidation of Metallic Iron by a Current of Air in Presence of Iron Salts’ By Peter Fireman MAGNETIC PIGMENT Co., TRENTON, N J.
H E conversion of iron into a hydrated ferric oxide can be directly and expeditiously accomplished by blowing air through a bundle of strips of the metal suspended in a solution of a ferrous salt. A successful procedure for achieving this result was first described by Penniman and Zoph.2 A few observations made by the writer in the study of this process may be of interest. Process The oxidation is conveniently carried out on a large scale. Operations are conducted in tanks of 15,000 to 18,000-gallon capacity. Inside each tank is placed a tub or box of nearly the same height. The two vessels communicate a t the upper ends through perforations in the inner one. The whole is a l e d with the solution of a ferrous salt and one of the vessels is filled with iron scrap. The latter rests on a perforated bottom through which the air is admitted. Through the upper lateral openings and those in the false bottom a good circulation is kept up between the inner and outer vessels by the air coming in from a compressor. Having filled the tank with the solution of a ferrous salt and the receptacle for iron with scrap, the operation is started by introducing a rapid current of air and live steam. A temperature of 55” to 65” C. is conducive to rapid oxidation. The yield, however, is greatly determined by the concentration of the solution of the ferrous salt. A specific gravity of 1.08 to 1.09 is ample in the case of ferrous sulfate and 1.06 to 1.07 for the ferrous chloride. The process is continuous and is kept up without interruption for a number of days. Only from time to time scrap metal must be added to replace the iron used up-the amount so consumed may reach a ton in 24 hours. The iron solution, on the other hand, lasts for a very long time, being used over and over again, in one charge after another. The hydrated ferric oxide produced carries 2 to 3 per cent of sulfur trioxide and to the corresponding extent the iron salt is used up when operating with the sulfate. Like losses take place when the chloride is used. A charge is considered completed when the desired color is attained. Color Changes during Process A remarkable feature of the transformation of iron into hydrated ferric oxide under the conditions described is the continuous change of color which the product undergoes. At first the pigment is yellow of varying shades; after a while it turns reddish brown, deepening more and more in the course of time; a t other times the yellow passes to yellowish brown, then to darker brown, and thereupon warms up to a decided red brown. These changes of color are more characteristic of the hydrated ferric oxides made in an iron sulfate solution. In an iron chloride solution like changes take place, but they are all darker a t the same stages of the process and end with purplish browns. The red-brown pigment forms a very fine powder, but very heavy although its specific gravity is only about 4.
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1 Presented before the Division of Physical and Inorganic Chemistry a t the 70th Meeting of the American Chemical Society, Los Angeles, Calif., August 3 to 8 , 1925. 1 U. S. Patent 1,327,061 (January 6 , 1920).
In oil it grinds out to a brownish yellow paint. On ignition it turns purplish red. The final red-brown product of an 8 days’ run in which ferrous sulfate had been used as the reaction medium was found to have the following composition on drying to a constant weight in the water bath: Per cent 83.79 2.59 0.07 13.63
Fez08 SO8 SiOa Hi0
1oo.08
Some light is thrown on the nature of the reaction taking place in the process by the fact that during the oxidation some of the ferrous salt is converted into ferric, which remains in solution. Usually there is found from 1 to 2 per cent, occasionally up to 3 per cent, of ferric iron in solution. This suggests that the ferrous hydroxide of the partly hydrolyzed ferrous salt is first attacked by the oxygen of the air, causing the separation of the hydrated ferric oxide and releasing a corresponding amount of acid, which goes to form a certain amount of ferric salt. The latter is then reduced by the iron, which in dissolving restores the ferrous salt. Thus, the new incoming air finds the conditions of the solution exactly the same as the air which preceded it and repeats the same oxidizing action. Nature of Reddish Brown Pigment The question now arises as to the nature of the water which the new reddish brown pigment carries. I n this connection Posnjak and Merwin3 recently proved “rather conclusively that no series of hydrates of ferric oxide exists among the natural minerals. The only existing hydrate is ferric oxide monohydrate.” A like conclusion they draw with reference to the synthetic hydrates of ferric oxide. The question then is-is the reddish brown pigment ferric oxide monohydrate or merely ferric oxide holding a quantity of adsorbed water? The latter supposition suggests itself in view of another statement by Posnjak and M e r ~ i n . Speaking ~ of the synthetic preparations, they say: It seems certain t h a t only two distinct types of “amorphous” hydrated ferric oxide e x i s t - o n e yellow and the other reddish brown. The yellow is apparently essentially ferric oxide monohydrate, while the reddish brown substance may hold its water in either a dissolved or a n adsorbed condition (or both).
It appeared of interest to examine whether or not the new reddish brown product is a true hydrated ferric oxide, by studying its dehydration curve. The heating a t each temperature was continued until a constant weight was reached. Starting with material dried in the water bath the following data were obtained: Time of heating Hours 70 58
Temperature O
c.
125 ( * 7 ) 145 ( * 7 )
Loss Per cent 1.53
0.76 2.29 8.27
155 to 185“ 0.37 210 to 220 0.I9 230 to 250 a All except 0.44 per cent of this loss took place at 155O to 175’ C. During the last few days 10’ to 15’ C. of heat were added to expedite the dehydration.
436 (18days) 44 63 ~~
* A m . J . Sci., 47, 311 (1919).
March, 1926
I S D U S T R I A 4 LA S D E N G I N E E R I S G C H E M I S T R Y
The copious loss of water, 8.27 per cent, given off within a range of about 30 degrees of temperature may well be considered as due to the decomposition of a true hydrate. On the assumption of a ratio of one molecule of ferric oxide to one of water the loss of water should have been 9-45 per cent. In view of the likely interference in this case by adsorption phenomena, such a divergence is not to be considered as excessive. Thus it may well be accepted that the reddish brown oxide under examination is a monohydrate of ferric oside. Further, a specific gravity determination gave a t 25" C. a value of 3.835. Since the hydrate contained 4.18 per cent
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of water in excess of that requisite for its being a monohydrate, a corresponding correction may be made, with the result that the specific gravity becomes 4.376. Both the corrected and the uncorrected values of the specific gravity are fairly in accord with those of the various types of natural monohydrates of ferric oxide, crystalline as well as amorphous, examined by Posnjak and Merwin. Consequently, while these authors may be right in their finding that none of the water carrying reddish brown oxides known to them, as described in the chemical literature, is ferric oxide monohydrate, the reddish brown pigment examined by the writer appears to be a true monohydrate of ferric oxide.
Deflagration Products of Smokeless Powder' A n Analytical and Physiological Study By H . C. Knight and D. C. Walton THE RESEARCH LABORATORIES. EDGEWOOD ARSENAL,LID.
HE Research Laboratories of Edgewood Arsenal were requested by the Navy Department to furnish certain data on the effects of the burning of smokeless powder in a semiclosed space simulating a turret-that is, one in which no excessive pressure could develop, but in which there was no free access of air for complete combustion. No quantitative data on this subject were available. Consequently a series of experiments was undertaken in order to determine (1) the composition of' the gases formed, (2) the physiological effects of these gases, and (3) how long these gases would persist in toxic concentrations under varying conditions.
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Apparatus
The deflagration chamber adapted to this work was a horizontal cylindrical tank, as shown in Figure 1, made of 3/8inch steel, 8 feet in diameter and 18 feet long, having a capacity of 910 cubic feet. The chamber was placed facing the north, and in this end a rectangular opening 27 inches wide and 37 inches high was cut. The tank was provided with sampling tubes as shown in Figure 1 and with a track running from the closed end of the tank out through the opening to a distance of 100 feet. In the center of the tank, straddling the track, mas a steel table 26 by 37 inches and 33 inches high. -4flat sheet of steel, grooved to fit the track, served as a carriage for introducing and removing steel cages containing animals. Materials
The powder used was a single-base Navy powder, 6 inch, 53 caliber, which had been inspected on November 30, 1924.
It was placed in muslin bags in 35-pound lots, two bags constituting a charge for firing. The powder was ignited by means of an electric squib placed in a bag containing 5 ounces of black powder, which had previously been removed from the gun charge as received. Methods of Testing and Analysis
Samples of gases were taken through tubes in the side of the chamber into 250-cc. Pyrex sampling tubes evacuated to about 2 mm. of mercury. Samples were analyzed according to the scheme of Clemens Tho Bureau of Mines apparatus described by Burrel13was used. The figures 1 Received August 5, 1925. Published by permission of the Chief, Chemical Warfare Service. * Scott, Standard Methods of Chemical Analysis, 3rd ed., p. 1238. * THISJ O U R N A L , 4, 297 (1912).
given in the analyses are for 760 mm. and 27" C. No account was taken of the water formed during combustion. Temperatures were taken by a Hoskins thermoelectric pyrometer, type P. A. The thermocouple was always put through the side of the chamber a t the same height a t which gas samples were taken. The maximum pressure only was taken by a water manometer, inserted a t point 1. This manometer consisted of a vertical U-tube open a t the ends. Prior to firing it was filled with water until the outer end just overflowed. Any displacement due to pressure within the tank caused an overflow and loss of water, so that the lowering of the water level, multiplied by 2, gave the maximum pressure in inches of water. Some of the animals were examined before and after exposure to determine the percentage of hemoglobin in the blood as an index of the degree of lung edema developed. They were also examined to determine the degree of saturation of the blood with carbon monoxide. For hemoglobin the Sahli method was used, a sample of blood being drawn into a pipet and then diluted with 0.1 N hydrochloric acid and distilled water till it matched the standard. For carbon monoxide the Sayer method was used. Twotenths of a cubic centimeter of the blood was drawn up in a pipet and diluted to 2 cc. with distilled water. The diluted blood was then put in a test tube containing 0.02 gram each of powdered tannic and pyrogallic acids, shaken, allowed to stand 8 minutes, and then compared with t.he standard scale of color tubes. All animals were carefully observed after exposure and symptoms noted. I n some instances animals that appeared normal were sacrified and autopsied to be sure that no injury had been received. The organs of practically all the animals dying or sacrificed were sectioned and examined microscopically. Experimental Details
METHOD-TWO35-pound bags of smokeless powder were placed on the table and the black powder charge pinned to them. Embedded in this black powder charge was a small electric squib, which was connected by wires led through D to a blasting machine placed about 30 feet from the tank. After collecting meteorological data the signal to fire was given. Deflagration of the powder began a t once, causing a
