SOME EXPERIMENTS ON IRON BY LAWRENCE GANE KNOWLTON
Introduction As the title indicates, this investigation has to do with various experiments involving iron. The first three parts of the work deal with new methods for the determination of carbon in iron. The study of the tin-iron alloys was taken up with the object of determining the form of carbon when dissolved in iron. The next division is concerned with the effect of aniline on the rate of solution of iron in acid with the result that it can be classed with a group of compounds which act similarly. Next, some work on the reduction of nitrobenzene which was undertaken originally, t o determine the difference in amount of corrosion of iron in sodium sulfate and sodium chloride solutions, but which eventually led to other results. Lastly, some experiments on passivity which were carried out with tin, but could be duplicated with iron in appropriate concentration of nitric acid.
An Electrolytic Method for Total Carbon When iron dissolves in hydrochloric or sulfuric acids, much of the combined carbon, in the form of iron carbide, reacts with the acid forming hydrocarbons and leaves a small amount of carbon behind. This is due to the presence of nascent hydrogen a t the point where the iron carbide is reacting. Under these conditions the carbon readily combines with the hydrogen. A possible reaction would be as follows: 2Fe3C 12HCl = 6FeC12 C2Hz 5H2 A small amount of the carbon of the iron carbide does not unite with the hydrogen of the acid which accounts for the carbon remaining. That the combined carbon is not reduced to hydrocarbons by nascent hydrogen, was proved by making a piece of iron, high in carbide, cathode in a sodium sulfate solution. KO trace of hydrocarbons was found in the hydrogen evolved. If iron is dissolved in nitric acid, sp. gr.1.2, the concentration used in the determination of graphitic carbon, most of the combined carbon is oxidized to soluble compounds which color the solution brown; a little carbon dioxide and a trace of hydrocarbons are formed. Only graphitic carbon remains behind after the iron dissolves. If it is desired to dissolve the iron sample and have all the carbon remain, as in the determination of total carbon, the potassium copper chloride method may be used. The sample is dissolved in a solution containing the potassium copper chloride and some hydrochloric acid. The concentration of the acid is not high enough t o react with the iron carbide. The reactions involved are as follows: Fe zCuCl? = FeC12 zCuC1 Fe3C 6CuC12 = 6CuCl 3FeC12 f C
+
+
+
+
+
+
+
SOME EXPERIMENTS O N I R O S
I573
Iron made anode in a solution of certain electrolytes should dissolve and leave all its carbon behind as the hydrogen is given off at the cathode and thus has no opportunity to react with the carbon of the iron carbide, as it is broken down. This method is similar to the potassium copper chloride method except that a different method of solution of the iron is used. The electrochemical equations involved are as follows: Fe - zF = Fe++ Fe3C - 6F = 3Fe++ C
+
Justurn* employed this method for determining total carbon. He made a steel rod anode in a dilute hydrochloric acid solution, sp. gr. 1.1, using a platinum cathode. The current was allowed to run over night, I O or 15 grams of steel being dissolved. The carbon could then be filtered off and determined by combustion in the usual manner. No figures are given as to the accuracy of his results. I t would seem, however, that the iron would not dissolve uniformly thus giving cause for error. The method used in these experiments was to place the iron turnings in a platinum dish and make this anode. A strip of copper served as cathode. The current from two storage batteries was used which gives 3.6 volts and .65 amperes a t the start. The amperage changes considerably during the run. As electrolyte, a 10% solution of sodium chloride was first employed. It was found that chlorine was evolved which of course would not be expected as the potential required to make iron dissolve, is less than that to discharge chlorine. To find out the cause of this, some current-voltage readings were taken. It was found that after the current starts, the amount flowing through the circuit gradually decreases. This is explained as due to the increased resistance caused by the iron hydroxide formed. This would form around the iron and increase the resistance between the platinum dish and the iron. If acid is added, or the solution is stirred, the current goes up again. When the iron hydroxide forms around the iron it becomes easier to discharge the chlorine than to force the iron to dissolve. A 10% solution of sodium sulfate was substituted for the sodium chloride and while some oxygen was evolved it does not oxidize the carbon as can be seen from the results. About 1.5 grams of iron turnings were weighed out and placed in the platinum dish. The sodium sulfate solution was then poured in, and the current started. As mentioned before, as the electrolysis proceeds, iron hydroxide appears and settles down on the turnings tending to prevent the dissolving of the iron. Acid could be added but there is danger of adding enough to attack the iron. Instead, the solution containing the iron hydroxide is poured off leaving the undissolved iron behind in the dish. More sodium sulfate solution was added and the electrolysis continued. This process has to be repeated several times. A half day or more is required for all the iron to dissolve. The contents of the platinum dish is then added Chem. News, 41, 17 (1880).
1574
LAWREXCE GANE KNOWLTOP;
to that already poured off and the mixture warmed with 15 or 20 cc. of concentrated hydrochloric acid to dissolve the iron hydroxide. The carbon is then filtered off on a Gooch crucible and determined by combustion in the usual way. The results check very closely with those found by the ordinary method of combustion of the iron sample. h white iron sample was first tried, which contains all the carbon in the combined form, so that if any of the carbon were removed in the form of a gas, it would be noticeable in the results. The other two are gray iron, number three being one analyzed by the U.S. Bureau of Standards. I
Ordinary Combustion Electrolytic Method
2.137/c 2.187,
3
2
3 00% 3.00%
2
19cc
2'27%
Different determinations on the same sample checked closely. I n the case of No. I , 2 . 1 5 , 2.18 and 2 . 2 2 were the results obtained: This method is accurate and would be satisfactory where rapid deterniinations are not required.
An Electrolytic Method for Graphitic Carbon in Iron Haber' states that an iron cathode corrodes in an ammoniacal solution of ammonium nitrate. Here, as the iron carbide dissolves, it would come in contact with nascent hydrogen from the electrolysis and form hydrocarbons. Thus there would be only graphitic carbon left after the solution of the iron. A piece of sheet iron was made cathode in a strong solution of ammoniacal ammonium nitrate and the current density kept low as recommended in the above article. The iron did corrode very slowly but it was found that a piece of iron in a similar solution corroded about as much without the use of current. Heating the solution or increasing the current, does not increase the rate of corrosion appreciably. No determinations of the amount of carbon remaining were made as i t would take an extremely long time for all the iron to dissolve. The current has nothing to do with the corrosion process. A spot of iron or iron carbide dissolves as anode in a local cell and at another point, the cathode, hydrogen is set free. Under these conditions the iron carbide does not come in contact with nascent hydrogen as it is broken down, so no hydrocarbons will be formed. Then, if the iron dissolved completely, most or perhaps all of the carbon should be left behind. When the iron dissolves in either hydrochloric or sulfuric acids, as mentioned before, the carbon does come in contact with the nascent hydrogen and forms the hydrocarbons which are evolved. If the iron should dissolve as cathode, the hydrogen given off a t the same time should react with the combined carbon and only graphitic carbon should remain behind. There is, however, the possibility of the hydrogen 1Z. Elektrochemie, 7, 733 (1901).
I575
SOME EXPERIMESTS ON IRON
reducing the nitrate radical rather than reacting with the combined carbon, and if this were the case, more than the graphitic carbon should remain behind.
Effect of Depolarizing Agents on the Amount of Carbon remaining when Iron dissolves in Acid As stated before, when iron dissolves in hydrochloric or sulfuric acids, the iron carbide reacts with the acid forming hydrocarbons which pass off as a gas. If a depolarizer were present the hydrogen from the above reaction might be oxidized and thus not have an opportunity to react with the carbon of the iron carbide. This would leave more than the usual amount of carbon behind. There is also a possibility that the depolarizer may be able to oxidize the carbon formed, either graphitic or carbide. The effect of the depolarizer may be tested by the amount of carbon remaining after the iron has dissolved. The idea for this work came from the fact that the iron-tin alloy mentioned later, left more carbon behind when dissolved in aqua regia, than in nitric acid of the concentration used for determining graphitic carbon. This effect will be taken up later. Potassium permanganate was chosen as the depolarizer for these experiments. A standard solution, rather concentrated, from .os to .06 grams per cc., was used. This was boiled to remove organic matter and filtered through an asbestos filter which removes manganese dioxide. Samples of the iron turnings, always within .OI grams of 1 . 5 5 grams, were placed in a beaker and a definite volume of the standard permanganate solution added from a burette. The total volume was made up to 50 cc. by adding water. Four cc. of concentrated sulfuric acid were then added and the action usually allowed to proceed without heating. After the first vigorous reaction, the beaker is heated and more sulfuric acid added to be sure that all the iron carbide has dissolved. Iron carbide dissolves less readily than iron. Mellor states that N / I O hydrochloric acts on it at goo, and the N acid gradually dissolves i t at ordinary temperatures. A series of runs was made using two different iron samples. No. I is white iron with a total carbon content of 2 . 1 3 % and no graphitic carbon. No. I1 is a gray iron, total carbon 3.31%~and graphitic carbon 2.54%. The results obtained are as follows: I I1 Gr. KMnOl
I-none 2-none 3-1.5 4-1.8 5-2.
9-2.3
Gr. KMnO.
% Carbon
1.3
3.01
.55
1.5
4
1.9 2.12 2.14
1.7
3.08 3.16
2.
3.11
2.2
3.13
2.13
2.4
3.02
1.64
2.5
2.98
4 4
(in 7 5 cc.)
% Carbon .75
8 4
I
6-2.3 7-2.7 8-2.95 10-2.3
cc. acid 4
4 4 5
2.04
7
1.9
1.25
I576
LAWRENCE GANE KNOWLTON
The first runs were made with white iron and, as can be noted, with the correct concentration of permanganate the total carbon can be obtained, which would suggest using this as a method for total carbon. The runs were then made with the gray iron wit'h less satisfactory results. The maximum percentage of carbon is about .IS%, less than the total carbon. This was found to be due to the oxidation of the graphitic carbon by the permanganate, as will be explained later. I t will be seen from the results, that there is a maximum in the amount of carbon remaining which falls off on either side. This is explained as follows. With a smaller amount of permanganate present, it does not oxidize the hydrogen rapidly enough to prevent some of the hydrogen reacting with the carbon of the iron carbide. This carbon then escapes in the form of hydrocarbons. With too much permanganate present, it oxidizes the carbon. Increasing the amount of acid (9) would be expected to decrease the amount of carbon as the iron would dissolve more rapidly, thus giving the permanganate less opportunity for its action on the hydrogen. Making the solution more dilute (8) would leave less permanganate in contact with the iron to act. This could be overcome by stirring. As mentioned, the maximum amount of carbon that was left, was the same as the total carbon in the case of the white iron but somewhat less with gray iron. Another gray iron gave similar results whose total carbon was 2.19% and graphitic carbon 1.82%. I11 Gr. K M n 0 4
R Carbon
1.8
1.72
2.1
2.00
One would expect the white iron to give the lower results because of the larger amount of combined carbon which has an opportunity to go off as hydrocarbons. It was thought that by changing the conditions during the dissolving of the gray iron, the carbon remaining could be brought up to the total carbon content. The acid content was then reduced to z cc. of concentrated sulfuric acid. Under these conditions some iron salt is precipitated. This does not dissolve readily even when more sulfuric acid is added. Hydrochloric acid, however, will dissolve it. The amount of carbon remaining was not raised by this change. Other runs were made, stirring until the permanganate was used up. This did not change the results appreciably. Better results were obtained if the permanganate was added gradually and stirred during the addition. Gradual addition would tend to prevent the permanganate becoming concentrated enough to oxidize the carbon while stirring would keep the permanganate in contact with the iron, as it dissolved. The explanation of the low results in the case of the gray iron was finally found to be due to the oxidation of the graphitic carbon. To determine whether the permanganate oxidizes the graphitic carbon, the following experiment was carried out. Gray iron, 11, was dissolved in nitric acid, sp. gr. 1.2, which leaves only graphitic carbon undissolved. This carbon was filtered
I577
SOME EXPERIMENTS ON IRON
off on a Gooch crucible, then placed in a beaker with 50 cc. of a solution containing 4 cc. of concentrated sulfuric acid and . 2 5 grams of permanganate. The solution was heated for 15 minutes a t a temperature of 60' - 65OC. The graphitic carbon content had then dropped from 2.54YGto 2 . ~ 8 7 ~ . As mentioned before when iron dissolves, iron carbide goes into solution after the iron. I n the case of the gray iron, with little combined carbon, the graphite will come out a t the first part of the action and thus is exposed to the oxidizing action of the permanganate for a considerable length of time, while the concentration of the permanganate is still high. With white iron the carbon does not come out until the last of the action and is in contact with a more dilute permanganate for a shorter length of time. Thus there is little opportunity for its oxidation. It would then be logical to conclude that the permanganate could be added some time after the acid had begun to act on the iron and be more effective than if added before the action starts. This conclusion was borne out by experiments. The acid was allowed to act for a certain length of time, five to seven minutes, and then the permanganate was dropped in rapidly, about as fast as it was used up. The addition of the permanganate was continued until the action of the acid on the iron was practically complete. With iron 11, this method gave a value of 3.z40j0, . o 7 y 0 less than the total carbon value. With iron II1,it gave 2.17%, .ozOjOlessthan the total carbon and almost within the limit of experimental error. Here are two causes of the lower results in the case of iron 11. First; it has a higher total carbon content thus there is more likelihood of the carbon being oxidized. Second; the turnings are not uniform as is the case with iron 111. Some of the smaller particles of iron would dissolve completely before the permanganate is added, thus losing the combined carbon as hydrocarbons. From the above, i t is evident that the method given would be practical for the determination of total carbon in white iron, as it is accurate and as rapid as any method involving the solution of iron previous to combustion. It would not be reliable in the case of the gray iron, due to the difficulties mentioned. As mentioned before, if iron is dissolved in aqua regia it leaves more than the graphitic carbon. Runs were made using the same iron samples as before. The results follow: KO.of cc.HNO,
No. of cc.HC1
8
21
1.5
0
25
25
8
24
0
8
0
24
No. of CC.H*O
Iron sample used
White I White I Gray I1 Gray I1
% Carbon .a9
.32
2.87 2.56
These results show that more carbon remains when aqua regia is used to dissolve the iron than i: either acid were used alone. This would be explained as follows. When there is a mixture of the two acids present, the nitric acid or the chlorine produced would oxidize the hydrogen preventing
I578
LAWRENCE GANE KNOWLTON
the formation of hydrocarbons. Of course nitric acid alone would be capable of oxidizing this hydrogen. There is an added effect, however; some of the nitric acid is used up in oxidizing the hydrochloric acid thus preventing it from acting on the carbon. This would cause more carbon to be left behind. Some runs were also made to determine the oxidizing effect of hot aqua regia and chlorine on the carbon from the white iron and the graphitic carbon from the gray iron. The carbon from the white iron was produced by treatment with permanganate and acid. The graphitic carbon was produced in the usual manner. The results follow: cc. of HCl
cc. of HS03
cc. of water
Source of Carbon
24 24
8 8 8
15
I5
White iron I Gray iron I11 Gray iron I11
24
0
Before
2.13%
After
1.8
%
36% 1.8 70
1.8
70
1.8yo
1
When chlorine was bubbled into heated water containing the carbon from the white iron, the carbon content went down to 1.28%. Using the graphitic carbon from iron 111, after similar treatment with chlorine for an hour and a half, there was no loss in carbon. Aqua regia and chlorine are similar in that they oxidze the carbide carbon but not the graphitic carbon. It would follow, then, that if chlorine were bubbled into an acid solution in which iron was dissolving, it should raise the amount of carbon as was the case with permanganate. Of course the experiment should be so regulated as to prevent the oxidation of the carbide carbon as much as possible. Runs were made with both white and gray iron. The iron turnings were placed in the acid and heated to 60’ - 80’. Chlorine was bubbled in from a cylinder as long as there was any action. % Carbon
cc. of HCl
cc. of water
Iron sample
IO
40
I1
2.94
30
20
I
1.00
The results do not reach the value for total carbon in either case due either to the escape of hydrocarbons or to the oxidation of some of the carbide carbon. It would probably be possible, by regulating conditions, to bring these values up to some extent. However, it would be more difficult to regulate the concentration of the chlorine than that of the permanganate. The rate of solution of the iron could be readily controlled by the concentration of the acid and the temperature.
Effect of Tin on Carbon in Iron I n Percy’s Metallurgy,* Eyferth has experimented on the action of tin on cast iron. He thought that if 25yGtin was added to and stirred with molten gray cast iron, the whole of the graphitic carbon would separate out and this would be left behind in the crucible when the alloy is poured out. If this were true it ought to have an interesting bearing upon the form of carbon “Metallurgy: Iron and Steel”, 163 (1864).
I579
SOME EXPERIMENTS ON IRON
in the melt. If there is a reversible equilibrium between combined and uncombined carbon in the melt, then tin should displace all the carbon from the melt or if we assume with Jeffries and ilrcher that there is 110 combined carbon in the melt, it should also displace all the carbon. The iron-tin system was first worked out by Isaac and Tammann.' Iron and tin are not miscible in all proportions in the melt: between 50% and 8gyO, two layers are formed in equilibrium a t 1140OC. Below this temperature mixed crystals of 1 9 7 ~tin content deposit from the iron rich layer leaving the melt containing 89Yc tin. At 893'C a reaction occurs between the iron rich mixed crystals and the tin rich melt with the formation of a compound whose formula is not given. The system was further investigated by Wever and Reinecken.2 The diagram of the system was changed somewhat and the composition of the iron-tin compounds separating out was determined. An article on the iron-tin-carbon system by Goerens and EllingenJ3has been published. The diagram is not worked out but the constituents of the system are determined by cooling curves and microscopic examination. The miscibility gap between iron and tin in the melt is increased by the presence of carbon. The solid phases separating out from the iron layer are said to be first, austenite containing tin, then cementite, the iron-tin compound and lastly a ternary eutectic. Analysis of the iron layer is made for tin, total and graphitic carbon. No mention is made of whether carbon is thrown out or not. To verify Eyferth's experiments, various alloys of iron and tin were made up by melting gray iron in the induction furnace and adding tin. The resulting melt was poured into sand. Graphitic carbon was thrown out but by no means the whole amount, as shown by the following table: Tin
Iron for Alloy I Alloy I Iron for Alloys 2-6 Alloy 2 Alloy 2 annealed Alloy 3 Alloy 4 Alloy 5 Alloy 6
I2
.a
2.1
.06
2.1
15
2.3 3.00 2.6
.29 .22 1.75 .34 .26 Loss of
7.5 I3 16 of iron
I
2 2
annealed
3 4 5
6
2.52 2.48 2.48 2.7
3.2 3.00 2.64
Z. anorg. Chem., 53, 281 (1901). * Z. anorg. Chem., 151, 349 (1926). 3
hletallurgie, 7, 72 (1910).
G. C. 2.95 .3 2 ' 54
I5 I5
T. C. on basis
Alloy Blloy Aolly Alloy Alloy Alloy Alloy
T. C. 2.99 2.19 3.31
2.2
G. C. on basis of iron
.35 .07 ,34 .26 1.9 .39 .31
carbon
'5 .8 .8 .6 .I .3
'7
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LAWRENCE GANE KNOWLTON
From the above results it will be seen that .8y0is the most carbon thrown out which amounts to only about a third of the graphitic carbon contained in the original iron. The amount of carbon thrown out of the alloy as graphite, increases as the tin content of the alloy is increased; alloy 4 with 7 . 5 5 tin lost very little while alloys around I jyc tin lost .6-.8$,. The graphitic carbon content of the alloy also decreases with increase in amount of tin as a comparison of alloy 4 with the others shows. The tin content of the iron can only be increased to a certain point, 16%, after which two layers appear, the tinrich layer containing about 9 0 7 , tin. This tin-rich layer contains no carbon. The amount of carbon in the iron layer can not be lowered beyond a certain amount when the iron layer is saturated with tin. The amount of carbon thrown out, however, increases with increasing amounts of tin as more iron would enter the tin layer and this iron would lose all its carbon. This would also effect the accuracy of the results on the basis of the iron present (columns 3 and 4). It would make the amount of carbon thrown out as graphite greater than it apparently is whenever a tin rich layer is formed. As an analogy to this effect of tin in displacing carbon from iron, can be taken the effect of alcohol on a solution of sodium chloride. Sodium chloride is only slightly soluble in alcohol and if alcohol is added to a salt solution, the salt precipitates out, if enough alcohol is added all the salt should crystallize out. In like manner carbon is not soluble in tin and when tin is added to a solution of carbon in iron carbon comes out. If enough tin were present to dissolve all the iron, all the carbon should be thrown out. It will be well to consider here what conclusions can be drawn, from this work and that of others, as to the form of carbon in the melt. There are three possibilities, two of which have already been mentioned. They are I. It may be present as dissolved carbide. 2. It may be present as dissolved carbon. 3. There may be a reversible equilibrium between the two represented as follows Fe3C s 3Fe C. At first thought it would seem possible to determine this from the form in which carbon separates from solution. This line of attack is not conclusive, due t o the fact that the form in which a substance separates from solution is not necessarily the one in which it exists in solution. Any electrolyte dissolved in water gives numerous indications that it is present in the form of ions. However, when the solution is cooled, the substance itself, not its ions, comes out. Thus whether graphite alone, cementite alone, or sometimes one or the other separates, gives no evidence of the form of carbon in the melt. Some authors base their opinion as to this point on such evidence. Desch,’ does not commit himself but indicates that he considers the reversible equilibrium in solution possible to explain the separation of graphite and cementite. Hatfield* has carried out a number of experiments with different iron samples including those containing silicon, in which graphite is generally
+
“Metallogra hy,” 368 (1913). 2 ~ r o c ROY. .
L.,
8 5 ~ 1, (1911).
SOME EXPERIMENTS ON IRON
1581
considered to be the form that separates from solution. Microscopic evidence shows that, in the cases studied, carbide must separate out from solution before carbon can appear and only this structurally free carbide can dissociate. Annealing carbor, is formed by the decomposition of the free carbide from which it is produced. That is, there must be an intermediate formation of cementite before free carbon can be produced. It would be interesting to see if the same results could be obtained with a higher silicon iron than the ones used or also with a high carbon content alone. Here it is more likely that carbon separates from solution. Hatfield does not mention in this article whether he believes his experiments prove the presence of iron carbide in solid solution but this is indicated in his discussion of an article by E. D. Campbell.’ This, as has been shown, would not be necessary as the iron carbide could form as it separated out. Sauveur* states that many evidences point to carbon being dissolved in molten iron as carbide. At this point no statement is made as to what these evidences are. Howeverla a diagram is given for the solubility of carbon in molten iron which has some bearing on the subject. At 1823OC and 6.67y0 carbon there is a slight break in the curve and the melt is said to be liquid Fe3C. Again a t 2 2 2 0 O C and 9.6% carbon liquid Fe2C is said to be present. At this temperature there is another break and above this the Fe2C decomposes into iron and graphite as the solubility becomes less. The latter then separates from solution. There is no evidence to show that these compounds are not decomposed in the melt, even if they should separate as the iron cooled. Also, there is considerable disagreement as to this part of the ironcarbon diagram. Tammann4 refers to another diagram which has no breaks a t the points mentioned and shows no decrease in solubility of carbon above 220ooc.
SauveuIJ also states that if carbon alone is dissolved in solid iron i t should collect in particles large enough to be visible under the microscope. This is not necessary, solutes never segregate in a definite portion of a liquid solution. The separation of “kish”, graphite on the surface of molten iron, also gives no proof of the form of carbon in solution. It might either form by separation from the melt directly, or by rapid decomposition of the iron carbide after this had come Hoyt* gives a summary of some work by E. D. Campbell. This has to do with the form of carbon in solid steel. A series of iron carbides of the general formula C,Fe3, is thought to exist in steel. Their presence is indicated by the different hydrocarbons evolved when the steel is treated with hydroJ. Iron and Steel Inst., 2, 1 2 (1914). “The Metallography and Heat Treatment of Iron and Steel”, 4 2 j (1926). Sauveur: “Xletallography”, 360, Tammann: “Metallography”, 235 (192j). Sauveur: “Metallography”, 394. Hoyt: “Metallography”, Part 11, 143 (1921). Sauveur: “Metallography”, 430. Hoyt: “Metallography”, 186.
1582
LAWREXCE GANE KKOWLTON
chloric acid and also from the color of the nitro derivatives formed when steel is dissolved in nitric acid. A s the temperature increases the carbides in solid solution dissociate into others of lower molecular weight. The carbides also undergo what is termed ionoid dissociation into carbon and iron, with rising temperature. The hardness and other properties of steels are explained as due to the amount and condition of the different carbides present. The suggestion i s made that austenite is a solid solution in which the carbides have undergone almost complete dissociation. This is based on the softness and electrical resistance of the austenitic iron. Campbell’s work is thorough and he meets possible objections to his theory in a satisfactory way. He is dealing with carbon in solid solution, but if the carbides are completely ‘dissociated in austenite they probably would be in the melt. Jeffries and Archer’ believe that only dissolved carbon is present in solution because the iron carbide molecules would be too large to diffuse. “When an intermetallic compound forms and dissolves in one of the component metals it is commonly considered to go into solution as such. For example, carbon is generally held to be in solution in gamma iron as cementite rather than as elementary carbon.” “From evidence available and particularly from a consideration of the phenomena of diffusion, the authors have reached the conclusion that this is not the case, but that the carbon in austenite is present as individual atoms of carbon. These atoms are undoubtedly held strongly to the neighboring iron atoms, but the union is not permanent. Diffusion must consist in a migration of carbon atoms, and not of groups or molecules containing several iron atoms. Such groups could not, on account of their size diffuse through the solid iron.” “According to this view, cementite has no existence except as a crystalline substance which not merely precipitates but is formed on the decomposition of austenite.” Going back to non-metallic solutions there are numerous examples of compounds having large molecules which are soluble and diffuse readily. Sugar and many other compounds can be brought forward to disprove the view of Jeffries and Archer. The effect of tin in throwing out carbon from molten gray cast iron remains as evidence as to the form of carbon in solution. It is evident that, a t the temperature at which the alloys are formed, only a certain amount of the tin will dissolve in the iron, after which two layers are formed. This is shown by the fact that increasing the tin content beyond a certain amount does not increase the amount of carbon thrown out of the iron layer. The carbon coming out of the molten iron indicates the presence of some free carbon. I t might be assumed that the tin could react with the iron carbide forming an iron-tin compound and setting free carbon. This is not possible because the tin would react more readily with the free iron and even if the reaction did occur the iron would react with the carbon again. 1
“The Science of Metals”, 407 (1924).
SOME EXPERIXENTS ON IRON
I583
The action of the tin, then, eliminates the possibility of the carbon being present as iron carbide alone. However, the results can be equally well explained from either of the other points of view. If there is a reversible equilibrium, FerC*3Fe C the tin can displace any amount of carbon as long as the concentration of the tin in the iron can be increased, because the equilibrium shifts to the right as the carbon is thrown out of solution. The fact that all the carbon is displaced from the iron in the tin layer, can also be explained by the shifting of the equilibrium to the right. Thus there is no definite evidence to prove whether the carbon is present in the melt as such, or whether there is a reversible equilibrium between it and iron carbide. To return to some properties of the alloy. From the carbon analyses given, it is evident that the alloys produced have a much higher combined carbon content than the iron which is used to make the alloy. This is not due to rate of cooling as an iron sample melted and cooled under the same conditions as the alloy undergoes no change in the amount of combined carbon. This combined carbon of the alloys is quite stable, as alloy 2 when heated over night at about 900°C showed an increase of only about . 2 5 % in graphitic carbon. The alloys made by Goerens and Ellingen have a high combined carbon content as the following table, taken from the article, shows:
+
Sn added
7
Sn found
T. C.
G. C.
4.89 6.64
3.32
.24
3.11
1.11
I2
9.09
.3.32
‘52
15
8. 8.92
3.2
‘72
6.78 9.24
3.5
9
15
16 19
3.17 2.71
,48 .49
22
I O .5
2.8
.08
30
11.1
2.73
,71
The tin then has the ability to raise the combined carbon content. It would seem logical to believe that the tin has some effect on the silicon in the iron. A white iron can readily be cast if the silicon content is low, but when the silicon content rises to I%-z%, a gray iron will generally be produced. The gray iron used had a silicon content of Z . Z I ~ , , enough to make it evident that the tin had some effect on the silicon in the iron of the alloy. This action of silicon is believed to be due to its effect on the iron carbide, promoting its dissociation into iron and carbon. According to Hoyt’, “The mode of occurrence of silicon in the silicon cast irons is still a matter of conjecture but it seems reasonable to assume that the silicon distributes itself between austenite and cementite, and that the silicon in the cementite makes 1
“Metaliography”, 290.
1584
LAWREXCE GANE KNOWLTON
that constituent break down more easily a t high temperatures.” Hoyt refers to Stead and Hatfield for this statement. Stead’ states that a carbosilicide crystallizes with iron carbide. What the composition of this carbosilicide is is not mentioned but it probably consists of mixed crystals of iron silicide and iron carbide, as mentioned by Gontermann in a reference given later. The diffusion of the silicide leads to the decomposition of the iron carbide. Silicide of iron when heated with pure silicon free white iron, decomposes the carbide in the white iron. Hatfield2 has analyzed and found silicon to be present with iron carbide. When manganese and sulfur, elements which promote the stability of cementite, are present less silcion is found with the carbide. Gontermann3 has investigated the iron-carbon-silicon system. From an iron whose carbon and silicon content corresponds to the iron used for the above experiments, saturated silicon martensite separates out as primary crystals and solidification is concluded by a secondary crystallization of a mixture of saturated martensite and silicon cementite. He is evidently using the term martensite for austenite. The silicon cementite consists of mixed crystals of FeSi and Fe3C. It is evident, then, that silicon is present with cementite in the iron. None of the above authors make an attempt to explain how the presence of ilicon makes the cementite break down into iron and graphite. I n the case h a t Stead mentioned of the decomposition of silicon free white iron by heatng with iron silicide, the silicon would have to diffuse into the cementite in order to make it decompose. That is, silicon is soluble in cementite. If the presence of silicon makes carbon less soluble in iron, the action of the silicon could be explained by saying that it dissolves in the iron, throwing out carbon. Then when iron containing silicon is cooled rapidly, the silicon simply crystallizes with the cementite, but if allowed to cool slowly or reheated to around 1000°C the silicon diffuses into the cementite causing graphite to separate out and giving a gray color to the iron. Silicon does decrease ferro-silicon is dissolved in the solubility of carbon in iron, for when 507~ molten iron a considerable quantity of graphite is thrown out.* I t is also possible that the silicon that is in solid solution, austenite, would lower the solubility of the carbon, and graphite would separate directly from this without undergoing the intermediate formation of cementite. When the action of silicon is explained as above, it would seem that tin should act the same way, as it makes carbon less soluble in the liquid state. However, tin is not soluble in ~ernentite,~ and cementite in itself is quite stable.6 “Searly pure iron-carbon alloys do not graphitize readily.’’ Engineering, 90, j08 (1910). Roy. SOC.,85.4, I (1911). Z. anorg. Chern., 59, 373 (1908). Hague and Turner: J. Iron and Steel Inst., 2, IOO (1910). Desch: “Metallography”, 382, Jeffries and Archer: “The Science of Metals”, 314.
* Proc.
SOME EXPERIMENTS ON IRON
1585
Tamaru’ has investigated the tin-silicon system. The two metals are miscible in all proportions in the melt and silicon separates when the melt is cooled. No compounds or solid solutions are formed. With this in mind it was thought that the tin of the alloy was dissolving the silicon, thus preventing its action on the cementite. To prove this, an alloy was made up to which enough tin was added to insure the formation of a tin layer. This tin layer contained no silicon but the iron layer showed a loss of about .3% silicon. The silicon had probably been thrown out as the alloy cooled. This was verified when silicon was found to be present mixed with the graphite thrown out. Eyferth also mentions that silicon was thrown out from his alloys. One would expect the silicon, as it was soluble in molten tin, to crystallize in the tin layer as it cooled. What has happened is that the silicon separates from the liquid solution because of the difference in density and the probable low viscosity of the molten tin. This is similar to the separationof “kishJJ,graphite onthesurfaceof themoltenmetal,whena high carbon iron is cooled. Also as no eutectic mixture of tin and silicon is formed, the tin layer remains liquid until its solidification temperature is reached, thus making it easier for the silicon to separate. I n the iron layer the iron will solidify first so that the silicon when it crystallizes cannot separate out of the body of the alloy. The tin of the iron layer, as it solidifies last, would hold the silicon in solution until after the cementite had crystallized. The pure cementite would then decompose less readily. This gives rise to a high combined carbon content. The alloy produced is hard and brittle. The brittleness was evident when it was found that the alloy could be readily pounded up to a powder. The alloy could not be turned in a lathe even using stellite as a tool. However, it is not so hard as this would indicate. It will not scratch glass unless heated and quenched. Its effect in the lathe is probably due to a hard constitutent in the alloy, possibly the iron-tin compound, which ruins the cutting edge of the tool. Similar cases are known in which it is impossible to turn soft metals which contain some hard constituent. I n the spheroidizing process, steel which has been cooled rapidly, so that the cementite comes out in the form of spines or network, is annealed. During the process the cementite changes to globules or spheroids which makes the steel much easier to machine. Effect of Aniline on the Rate of Solution of Iron in Acid
It has been stated that in the pickling of iron and steel, acid containing aniline removes the mill scale and does not attack the iron to such a large extent as acid alone. Other substances have been found which act similarly. A number of these substances are mentioned and their effect studied by A. Sieverts and P. Lueg.Z 2. anorg. Chem., 61, 40 (1909).
* Z. anorg. Chem., 126, 193 (1923).
I 586
LAWRENCE GANE KNOWLTON
To test the action of aniline a strip of sheet iron, covered with a coating solution of sulfuric acid containing aniline. of mill scale, was placed in a 107~ The acid was practically saturated with aniline and was kept a t a temperature of 5o0-6oo. The mill scale is removed and much less hydrogen is evolved than if acid alone is used. Another experiment of a more exact nature was the following: A piece of iron was placed in a beaker of 10% sulfuric acid heated to 5 0 O - 6 0 ~ . The hydrogen evolved was collected which amounted to 40 cc. Another piece of the same iron of the same dimensions was placed in 150 cc. of the 10% acid containing I O C C . - I ~ cc. of aniline. This was heated a t the same temperature and for the same length of time as the other. Only about 5 cc. were evolved in this case. The action of aniline and other substances in cutting down the hydrogen 0 evolution may be due to: I. The formation of a film on the metal which would increase the resistance in the local cells on the iron. This would slow up the solution of the iron, as it is an electrolytic process. 2. Through the presence of this film or by some other phenomena, the overvoltage of the hydrogen becomes higher than in acid alone. This makes it more difficult for the iron to go into solution. Some work has been done on the determination of film resistance by means of an oscillograph.' A rise in overvoltage could be determined by means of the usual currentvoltage readings. The decomposition voltage would have a higher value if the overvoltage of hydrogen on iron was raised. A simple potentiometer outfit was set up using a voltmeter to register voltage and a milliammeter for current. When electrolyzed, a sulfuric acid solution of aniline gives aniline black a t the anode so that the curreht begins to flow almost immediately. To prevent this, the anode was placed inside a porous cup. If acid is used around the anode, which is platinum, it increases the tendency for the iron to dissolve so a solution of sodium hydroxide was substituted, zN sulfuric acid is used around the cathode. The iron used was not attacked appreciably at this concentration. The current-voltage run was first made with acid alone around the cathode, then with acid containing aniline, practically saturated. Acid and Aniline
Acid Alone Volts
Milliamps.
Volts
RIilliamps. . I
0
.I
0
1.25
.I
1.25
.I
1.3 I . 3 5 (Decomp. voltage) 1.4 1.5
.I
1.3
. I
.j
1.35
'9 1.8
Bur. Standards Sci. Paper S o . 504 (1925).
I.4
(Decomp. voltage)
. I
.4
1.45
'7
1.55
1.00
1587
SOME EXPERIMENTS ON IROX
It will be seen that the decomposition voltage is about .os of a volt higher when the aniline is present, due to the rise in the hydrogen overvoltage a t the cathode. Current-voltage runs were next made a t the cathode using a calomel electrode. Acid and Aniline
Acid alone Volts
Milliamps.
‘7
‘5
.j 8
,615 ,625 ,645 .66 ,675
Milliamps.
‘5
,48 .525
Volts
I
.j8
3 4
,625 .66 .68
5
’
2
6 7
70
‘72 ‘725
When aniline is present, the voltage, for the same amount of current has risen considerably. This amounts to about .os volt in each case, the same as the rise in the decomposition voltage, which gives conclusive evidence that the rise in overvoltage has occurred at the cathode.
Reduction of Nitrobenzene When iron is placed in sodium chloride solution which would ordinarily contain dissolved oxygen from the air, it corrodes more readily than it would in water containing the same amount of dissolved oxygen. This is due to the type of film formed on the iron. I n water alone, iron reacts producing hydrated iron oxide a t the spot where it dissolves, which tends to cut down the rate of corrosion. I n sodium chloride solution iron reacts producing first ferrous chloride which then diffuses and forms hydrated iron oxide a t the point where it meets the sodium hydroxide coming from the cathodic area. This leaves no protective film over the anodic portion as is the case in water. Metals are known to corrode more readily in sodium chloride solution than in that of other salts. Two papers’ may be found on the corrosion of metal anodes in different solutions. I n the first is a list of salts arranged in “The decreasing order of their metal dissolution.” The halides are placed first then sulfates, etc. Iron should follow this tendency and corrode less in sodium sulfate solution than in sodium chloride. When iron dissolves in water or salt solution, monatomic hydrogen is produced at the cathodic area. If this monatomic hydrogen accumulates, it sets up a back E.M.F. which tends to prevent the solution of the iron. Ordinarily the dissolved oxygen serves as a depolarizer for this hydrogen; other substances might be used instead. I n these experiments nitrobenzene E. P. Schoch and C. P. Randolph: J. Phys. Chem., 14, ;19 (1910);G. R. White: 15, 766 (1911).
1588
LAWRENCE GANE KNOWLTON
serves as a depolarizer. If iron corrodes more readily in a sodium chloride than in sodium sulfate solution, more nitrobenzene should be reduced in the former than in the latter. A mixture of 35 grams of powdered gray iron and 175 cc. of a solution containing P I grams of sodium chloride was placed in a round bottom flask and boiled under a reflux for 9.5 hours. The same procedure was followed for the sodium sulfate using the same number of equivalents of this salt. A run with no dissolved salt gave only traces of aniline. This is due to the nature of the film on the iron. To determine the amount of reduction products the method devised by Allen' was used with some modifications. The reduction products and unchanged nitrobenzene were extracted with benzene and filtered off from the iron and iron oxide on a Buchner funnel, and the extract made up to a definite volume. To determine aniline, I O cc. of the benzene extract were shaken with 150 cc. of dilute sulfuric acid, j to I jo, in three portions. The acid solution was made up to a definite volume and ten cc. portions titrated using the following standard method. A standard solution of potassium bromate was added to the solution of aniline containing potassium bromide. I n the presence of acid the bromate solution liberates bromine and tribromaniline is formed. After the aniline is used up, the bromine color appears. Potassium iodide is added and the iodine liberated by the excess bromine is titrated witha standard solution of sodium thiosulfate. To determine the total amount of unchanged nitrobenzene and reduction products, z cc. portions of the benzene extract were allowed to evaporate a t room temperature. The evaporating dish was weighed a t regular time intervals. At the point where the benzene disappeared, there was a considerable change in evaporation rate. Allen plotted these results in the form of a curve and from the break in the curve the amount of residue was determined. However, there was a sharp enough change in the rate of evaporation so that a definite weight could be determined without plotting the curve which was sufficiently accurate. The unchanged nitrobenzene was determined by evaporating 2 cc. portions, as above. The residue was shaken with a large excess, five or six times as much as required, of alkaline ferrous hydroxide. The shaking was continued for half an hour. This process changes the nitrobenzene to aniline but does not change the other reduction products. The mixture was shaken with benzene and the ferrous hydroxide filtered off and washed with benzene on a Buchner funnel. The benzene extract is shaken with dilute sulfuric acid, as mentioned above and the aniline titrated. Knowing the amount of aniline in the residue the amount coming from the reduction of nitrobenzene can be calculated. Subtracting the sum of the weights of the aniline and the nitrobenzene from the weight of the residue, gives the weight of the other reduction products. Using a known sample of nitrobenzene, it was not possible to recover it as aniline without losing 5 to 107~. This would make J. Phys. Chem., 16, 131 (1912).
I589
SOME EXPERIMENTS ON IRON
the values for nitrobenzene lower than they should be and the values for the other reduction products larger than they actually are. The results for the above mentioned runs follow: Salt
Rt. of aniline Wt. of nitrobenzene Wt. of other reduction products Total n%.recovered
SazSOl
Sac1
12.3
15
5 7 5 5
12.5
5.4 30
28
The sodium chloride is more effective in producing reduction. However the difference between the two is not large. Iron, then corrodes to only a slightly less extent in sodium sulfate solution than in sodium chloride. Snowdon' has done considerable work on reduction of nitrobenzene in a manner similar to the one used in the above experiments. He reduced by stirring a mixture of nitrobenzene and different salt solutions in contact with sheet iron. The mixture was kept at a temperature of about 100' by a bath of boiling water. For most of his work he used ferrous chloride solutions but some runs were made with ferrous sulfate and sodium chloride which are of interest here. Ferrous chloride was found to be considerably more effective than an equivalent amount of ferrous sulfate and in sodium chloride solution there was scarcely any reduction. This does not seem to agree with the results obtained above, so it was thought advisable to try reduction in the presence of ferrous chloride and ferrous sulfate in a manner similar to the one used with the sodium chloride and sodium sulfate solutions. Snowdon used 30 grams of nitrobenzene instead of 35 grams, otherwise the amounts of material are the same. The results are tabulated together with those of Snowdon, in the table below: TABLE I Snowdon
Iron in IO Solution Salt FeClz Tot. wt. recovered 27 . 5 Aniline 11.1 h'itrobenzene 5.7 Other reduction products IO. 7 Time-hours 9.5
Snowdon
10.6
IO
10.6
FeClz
FeSOa
FeS04
I
11.5 13.2
4.7 15.6
9.5
15.5
7
.9 27.4
5.5
6.5
6.5
C1 as in FeC12 NaCl
28
31
18.5
Snowdon
C1 as in FeCL KaC1
7
9.5
7.4
It can be seen that the results do not agree with those of Snowdon, as in ferrous chloride and ferrous sulfate solution practically the same amount of reduction was produced and the sodium chloride is more effective than either of these. J. Phys. Chem., 15, 797 (1911).
'590
LAWREKCE GANE KXOWLTON
Snowdon advanced the theory that the effect of the ferrous chloride was to keep the iron active as the amount of this salt in solution remains practically unchanged. The ferrous chloride is a reducing agent and as reducing agents are known to keep metals active, its effect here would be due to that reason. However, from the results above it seems as if all that is necessary is an electrolyte which serves, as mentioned before, to change the nature of the film on the iron. After running the above experiments it occurred to the author that the 30 grams of iron was not sufficient for the reduction of the 35 grams of nitrobenzene. Very little iron remains after reduction in ferrous chloride and ferrous sulfate solution. A larger amount is left when sodium chloride and sodium sulfate is used. This would seem logical if, in the presence of either of the ferrous solutions, the iron corrodes only to the ferrous condition. Both of these solutions being reducing agents they should tend to keep the iron in the ferrous condition. This would require more iron for a given amount of reduction. I n the case of the sodium chloride and sodium sulfate the iron would corrode to the ferric condition thus requiring less iron for reduction. The oxide produced has a different appearance in the two cases. From the ferrous salt solutions a black oxide is produced while from the sodium salt solutions it is brown, more the color of ordinary ferric hydroxide. Three runs mere then made using 50 grams of iron powder. Otherwise conditions were the same as in the other runs. Appreciable amounts of iron remained at the end of the run even when ferrous chloride and ferrous sulfate were used. A much larger amount was left when sodium chloride was used.
TABLE I1 Salt
Tot. wt. recovered Ani1ine Nitrobenzene Other products Time-hours
FeCL 27
24.8
9
FeSOa
sac1
26 22.6
29
.4
3
I9
3
5
9
9.5
The order of the amount of reduction with the different salts agrees with that of Snowdon. Using ferrous chloride and ferrous sulfate there was practically complete reduction as 24.8 grams of aniline is equivalent to 32.8 grams of nitrobenzene. The amount of reduction using sodium chloride, has also increased but not in a corresponding amount. Here it is not a question of a sufficient amount of iron but the reduction process is speeded up, due to the increased surface of the iron. The results still differ widely from those obtained by Snowdon. It was thought that the kind of iron used might have some effect on the amount of reduction produced. The gray powdered iron used in the preceding experiments would contain most of its carbon in the graphitic form and, as mentioned before, this iron dissolves more readily than iron which contains
1591
SOME EXPERIMENTS ON IRON
iron carbide. Snowdon used sheet iron which would probably contain considerable iron carbide. Strips of sheet iron were used in the following runs which contains all its carbon in the combined form.
TABLE I11 I
Amount of iron Salt Tot. wt. recovered Aniline Nitrobenzene Other products Time-hours
11
40
42
FeCll
FeSOl
20
30
16.9 none
10.8
3 9
I11
30 XaCl
2.2
31 2.9 19.5
10.7 8.5 8
8.6 9.5
VI
V
IV
38 NaCl
9.5
36
45
KaC1
NazSOl
34
33
2.6
I
22.4
.6
26.2
8.9
5.2
I5
9.5
Here the results agree fairly well with those of Snowdon. The most reduction occurs with ferrous chloride, somewhat less with ferrous sulfate, while sodium chloride and sodium sulfate bring very little reduction. The iron strips from the ferrous solutions come out clean while from the other two they come out covered with a deposit of adherent rust. This is an indication of the tendency of the former solutions to keep the iron active while the rust on the iron from the latter would tend to cut down the action. The iron loss in the reduction using ferrous chloride solution was 29 grams which proves the point mentioned above, that 30 grams of iron is not sufficient for reducing the 35 grams of nitrobenzene. When iron strips are used the total surface exposed would be considerably less than if the same amount were present as iron powder. To determine the effect of surface exposed, some of the iron strips were made into turnings and run Is’ made with these. This does not increase the amount of reduction to a very great extent. The greater reduction with the gray iron powder, then, ia not due to its greater surface but to the nature of the iron. This is borne out in actual practice,’ where gray iron filings are used. It is stated that particles of steel which have become mixed with the iron are found unchanged a t the end of the action.
TABLE Is’ Iron Tot. wt. recovered Aniline Nitrobenzene Other products Time-hours 1
I
I1
30
30 28
33
15.5 7
20..
5.5
6
9.5
9
29
17.7 3.8 7.5
16
P. R. Groggins: “Aniline and its Derivatives”.
I11
6.3 8
I592
LAWRENCE GANE KNOWLTON
Some runs were made to see if the iron did eventually go passive in sodium chloride solution. If it goes passive, increasing the time of heating should not increase the amount of reduction; also the iron left after one reduction should not be capable of further action if used again. Gray iron powder was used for these experiments. It will be seen from the results that there is some increase in the amount of aniline produced by boiling for about seven hours longer. The action is slowed up because the iron becomes coated with oxide. The free oxide that accumulates would also tend to settle down on the iron thus retarding its action. I n run I11 the iron left after a run with 50 grams of iron was used again. There was some reduction which indicates that the iron has not gone passive but is not as effective as it was when first used. Two runs were made with strips of the sheet iron and sodium chloride which also have a bearing on the above. In one case glass beads were placed in the flask with the iron strips. These rubbing against the iron should tend to keep it active. In the other run the iron strips were taken out three bimes during the course of the action, cleaned and rubbed with sandpaper. In neither case is there any appreciable increase in the amount of aniline produced. This indicates that the iron remains active, for this treatment should activate it if it had become passive thus increasing the aniline yield. Snowdon also performed another experiment which is not in accordance with those of the author. He states that if strips of iron are placed in a flask and a mixture of boiling ferrous chloride solution, hydrochloric acid and nitrobenzene is kept in contact with it for some time that no appreciable reduction of the nitrobenzene occurs. He explains this as due to the insufficient contact of the nitrobenzene with iron. Better contact is obtained in his other runs by stirring. In the runs with iron strips made by the author, the conditions were exactly the same except for the presence of hydrochloric acid. Here reduction takes place as can be seen from the results in Table 111. The explanation of Snowdon’s results may be found to be due to the position of the iron strips in the flask. If they were not small enough to fit in the bottom of the flask where the nitrobenzene layer is, there would be little contact between the two, therefore little reduction. Passivity of Tin in Nitric Acid Considerable work has been done on the passivity of metals. A good summary of this work is given by Dunstan and Hill.1 The different theories of passivity are mentioned with the conclusion that the oxide theory is the most generally accepted. Their own work favors the oxide theory. The passivity of other metals was also studied. Bancroft favors the oxide theory but believes it to be a higher oxide, not stable in itself, which is adsorbed by the metal. In an article,e Schon states that tin becomes passive in nitric acid, sp. gr. 1 . 4 2 , when in contact with platinum, but without platinum energetic decom2
J. Chem. SOC., 99, 1853 (1911). Z. anal. Chem., 10, 291 (1871).
SOME EXPERIMENTS ON IRON
1593
position takes place. Cadmium behaves similarly, only a more concentrated acid, sp. gr. 1.47, must be used before it will become passive even with platinum. The function of the platinum would be to act as cathode thus liberating hydrogen a t its surface rather than at the surface of the tin where it would tend to reduce the oxide film to which the passivity is due. Hydrogen has a lower overvoltage on platinum which would make it a less powerful reducing agent even if i t did come in contact with the oxide of tin. Faraday‘ has done similar work with iron and finds that it will go passive in nitric acid when in contact with platinum and also if the platinum is simply connected with the iron. It was thought that the passive fdm on the tin might be stannic nitrate as this is produced as an insoluble substance when tin is acted on by concentrated nitric acid.* Two experiments show that this is not the case. Stannic nitrate is soluble in water so if this is the passive film, it should be removed when the tin is placed in water. A rod of tin was made passive by action of nitric acid and then shaken in warm water for a few minutes, dried and put back in nitric acid of such strength that it would be attacked if not passive. I t still remains passive, showing that the film is not soluble stannic nitrate. There is, however, a possibility of the stannic nitrate being adsorbed by the tin so that it would not dissolve in the water. To disprove this, tin was made passive by making i t anode in sodium hydroxide solution. It was washed and dried, then placed in a concentration of nitric acid in which it would ordinarily be attacked. The tin remains passive which indicates the passive film is not stannic nitrate. The method used in the following experiments on passive tin is as follows: Nitric acid, of the desired specific.gravity, is made up by diluting the fuming acid. The strength of the diluted samples was determined by titrating a definite weight of the acid with standard sodium hydroxide solution. T o insure that the tin was active each time it was dipped in the nitric acid, it was allowed to stand in hot hydrochloric acid and then washed and dried before being placed in the nitric acid. Tin, a t ordinary temperatures, will go passive even when not in contact with platinum, in nitric acid, sp. gr. 1.46. At a concentration of sp. gr. 1.456 it is slowly attacked. When a spiral of platinum wire is wound around the tin rod it will go passive in an acid as dilute as sp. gr. 1.426 but begins to be attacked slowly a t sp. gr. 1.423. Platinum should have the same effect in making tin passive if it were connected with a platinum electrode. The platinum electrode is connected to the tin rod and both are immersed in the nitric acid about 2 or 3 cm. apart. Some experiments were tried using nitric acid of the same strength as used with the platinum wire but passivity is not produced. A higher concentration, sp. gr. I .442, will produce passivity. The increased concentration necessary to produce passivity is due to the resistance of the solution. It follows “Experimental Researches”, 2, 240. Mellor: “A Comprehensive Treatise on Inorganic and Theoretical Chemistry”, 7, 330 (1927).
1594
LAWRENCE GAVE XNOWLTON
that if the distance between the tin and the platinum is increased, the tin would have more of a tendency to remain active. The same effect could be produced by putting a resistance in the external circuit. A slide wire resistance was then placed in the circuit. Using about 3 ohms resistance the tin still becomes passive at sp. gr. 1.443. If the resistance is increased to 6 ohms the tin remains active. Tin may also be made passive by making it anode in a nitric acid solution. A cathode of platinum was used and kept z or 3 cm. from the anode. Storage batteries served as the source of current, 9 . 6 volts being theE. M.F.employed. When the current was first applied a large ammeter reading was registered for an instant. This drops quickly back to a very small value, around .oo5 ampere, due to the formation of the passive film. Using this source of current tin went passive in an acid as dilute as sp. gr. 1.344. At sp. gr. 1.34 it remained active. Using 5 . 5 . volts it will not go passive in ordinary concentrated nitric acid, sp. gr. 1.407. The same is true using 3.6 volts. It is evident then, that a rather high E.M.F. must be applied to make a tin anode passive. The function of the current would be to overcome the solution resistance making it easier to discharge hydrogen at the cathode. Also, as the E.M.F. applied is raised it becomes easier to discharge oxygen a t the anode which would aid in the formation of the passive oxide film. This would explain why the smaller E. M. F. does not produce passivity. Krassal deals with a somewhat similar case, the formation of passive iron by making it anode in alkali. In boiling alkalies the passive state is quickly attained with strong currents and no visible alteration of the surface is produced, With weak currents, however, a visible film of oxide is first produced which attains considerable thickness before passivity is arrived at. He explains this as due to the formation of a complete thin film of oxide by the strong current. With weak currents the film is more irregular, so greater thickness is required before it becomes complete. The same idea can be applied to the tin anode in nitric acid. A thin complete film forms with a strong current while with weak currents the film never becomes complete enough to produce passivity. Of course the better film with the stronger current is undoubtedly due to the action of the increased amount of nascent oxygen produced. S-arY
The results obtained in this investigation may be summarized as follows : ( I ) Total carbon may be determined by an electrolytic method, making the iron anode in sodium sulfate solution. (2) Graphitic carbon cannot be determined by making iron cathode in an ammoniacal ammonium nitrate solution. The current does not cause the corrosion and i t is therefore slow and leaves total carbon (not graphitic carbon). 12. Elektrochemie, 15, 490 (1909).
SOME EXPERIMENTS O S IROX
I595
(3) Depolarizers added to iron dissolving in acid raise the amount of carbon remaining. Using potassium permanganate with white iron this can be brought up to the total carbon content, whileit is somewhat less with gray iron due to the oxidation of the graphitic carbon by the permanganate. Chlorine and aqua regia also raise the amount of carbon remaining. (4) z5y0 tin added to molten gray cast iron throws carbon out of the iron but not the total graphitic carbon content as claimed by Eyferth. The formation of two layers with increasing amounts of tin makes the addition of more tin, beyond a certain amount, have little effect. Carbon is present in molten iron either as dissolved carbon alone or there is an equilibrium existing between it and iron carbide. The high combined carbon of the alloy is explained as due to the dissolving of the silicon by the tin. ( 5 ) hniline added to sulfuric acid cuts down the hydrogen evolved by iron to a large extent. This does not prevent the acid from removing mill scale. The presence of the aniline raises the hydrogen overvoltage on the iron. ( 6 ) Gray iron will reduce nitrobenzene to about the same extent when boiled with either sodium chloride or sodium sulfate solutions. This indicates that both are effective in keeping iron active. The results found were not in accordance with those obtained by Snowdon. This was found to be due to the kind of iron used. Iron does not go passive in sodium chloride solution during the course of the action. ( 7 ) Tin goes passive in nitric acid when wrapped with a platinum wire. This is caused by the liberation of the hydrogen from the platinum. Connected to a platinum electrode it goes passive, but in a more concentrated acid, because of the increased solution resistance. It also goes passive when made anode in an acid more dilute than that used with the platinum wire. Passivity in this case is brought about by the discharge of hydrogen at the cathode and the effect of the nascent oxygen a t the anode.
Acknowledgment This investigation was carried out under the supervision of Professor IFr. D. Bdncroft. The writer wishes to take this occasion to express his appreciation of the suggestions and constructive criticisms offered. He feels that it is a rare privilege to have had the opportunity of associating with Professor Bancroft during this work. Cornell Uniuniversity.
