NEW HIGH-TEMPERATURE FIXATION REACTIONS O F NITROGEN BY NORMAN
w.
KRASE AND BILL MACKEY*
The remarkable growth of industries based on the fixation of atmospheric nitrogen during the past fifteen years has few parallels in history. In 1910a total of 9000 tons of atmospheric nitrogen was fixed; the production rising to 600,000 tons in 1925. Of the 1925 production, 6.7 percent were fixed by the arc process, 30.3 percent by the cyanamide process and 63 percent by the direct synthetic ammonia process. Present developments clearly show that the last named will further increase in popularity. A careful consideration of the synthetic ammonia process will incite in one a feeling of doubt that the last word has been said or that future developments will be only improvements in the details of existing processes. Research and development have carried the operations into the field of high temperatures and extremely high pressures. Enormous technical difficulties have been overcome and continued success seems certain. When it is realized that the micro-organisms in the soil fix atmospheric nitrogen at normal temperature and pressure on a scale far beyond what we may hope to attain commercially in the near future, it seems well to give thought to other methods of attacking the general problem. I n particular is it desirable to exhaust the obvious possibilities of existing data as a guide to future research. Practically the only avenue by which we can approach the general problem of nitrogen fkation is that of thermodynamics. The principles and methods of this subject have received wide exploitation in several branches of chemistry and its potential use in enabling predictions to be made with regard to chemical reactions has been ably demonstrated.’ The well known procedure of deriving new equilibrium data by the combination of several known equilibria is the method employed. An attempt has been made to indicate to what extent a given conclusion is unreliable by pointing out, for example, the range of extrapolation of data and the assumptions involved in a given calculation. A consideration of nitrogen-fixing reactions may well begin with those in which oxygen is concerned. The many important compounds such as nitric acid and its salts are of major importance in modern industry. A glance a t the data showing the free energy of formation of the nitrogen oxides reveals the fact that all are unstable and tend to decompose and liberate nitrogen. The data a t ordinary temperature and pressure are as follows:(I) I/Z (2) I/Z
(3) Nz
Nz Nz
+
+ I/.
+
202
0 2
O2 = NO(g) = NOz(g) = Nz04(g)
AFo298 =
20850
= 11920 AFOm= 22640
AF02g8
* Department of Chemistry, University of Illinois. 1
Lewis and Randall: “Thermodynamics and the Free Energy of Chemical Compounds”.
SEW HIGH-TEMPERATURE FIXATION REACTIONS OF NITROGEN
I 489
Investigation of these reactions over a wide range of temperature showed that N O becomes more stable with respect to its elements a t very high temperatures. The arc process of nitrogen fixation demonstrates the feasibility of establishing equilibrium a t a high temperature and by sudden cooling produces a gas a t ordinary temperature containing vastly more than its equilibrium concentration of NO. The slow rate of decomposition of N O at room temperature makes it for all ordinary purposes a stable compound. The arc process using air as a source of nitrogen and oxygen yields a gas containing from two to three percent NO by volume. This dilute gas is scrubbed with water or alkali and the process yields dilute nitric acid or a nitrate. Much of the cost of producing fixed nitrogen by the arc process results from the necessary use of dilute nitrose gas for absorption in water or alkali. Proposals to increase the concentration of NO in arc gas by the use of enriched air or by the use of Nz and O2in the stoichiometric proportions have been made. Calculations on reaction ( I ) can best be made by use of the data obtained by Nernstl using air. Table I shows the concentration of nitric oxide in percent by volume at equilibrium as a function of temperature. The results from 1811' to 2675' absolute are experimental values of Nernst using air. The calculations a t higher temperatures and those for stoichiometric proportions of nitrogen and oxygen were made using the relation2: (4) AFO = and the relation : (5)
21600
-
2.50T
Fa = -RTlnK
Equation (4)is developed3from the data of Nernst, Thomsen and Berthelot's value of 2 1 6 0 0 calories as the heat of reaction, and the assumption that the heat capacities of 02, Ns and NO are the same. The extrapolation is over a range of about 800'. TABLE I Temp. I811
Percent NO by volume Air I /2 N t I /2 O2
+
.37
2033
.64
2195
.95
2580
2.05
Temp.
-
2675
.80
2.62
Percent NO by volume Air I /2 Nt 112 0%
+
2.23
2.86
3000
3.47
4.52
3200
4.27
5.51
3500
5.52
7.23
Before considering these results any further we shall discuss some results obtained with another reaction for the formation of NO. 1Z. anorg. Chem., 49,213 (1906). 2"Our attention haa been called by one of the authors of recent data on the nitric oxide equilibrium, (Briner, Boner and Rothen: J. Chim. phys., 23, 788 (1926)) to the fact that
a new free energy equation has been developed by them on the basis of the new data. This calculation, it appears, involves the determination of the constant in the Nernst equation. In our recalculation, we have followed the more recent terminology of thermodynamics aa illustrated in the publications of G. N. Lewis, and have pointed out the effect of the new data on several important related equilibria. This latter was the primary purpose in publishing the recalculated values." 3 Lewis and Randall: loc. cit., p. 560.
NORMAN W. KRASE AND BILL MACKEY
I490
This calculation is given in more detail than subsequent ones. If we combine reaction (I)with the reaction expressing the dissociation of COZwe get the reaction desired :
+ +
1/2 Nz 1 / 2 0 2 = NO = CO (6) COz ( 7 ) I/Z NZ CO2 = NO (I)
ACp(1) = o ACp (CO ACp (Cod ACp ( 7 ) =
+
AHz%= 21600; AF'zQ~ =
+ 1/2 Oz AH2q8 = 68100;AF'zs
+ CO
1/2 0 2 )
2.75
-
20850
= 61750
A H z Q=~ 89700; AFOZ~E = 82600
= 9.75 = 7.00
,0056T
+
T
.OOI~
+ .ooiI T - .ooooo185 T2
+ .00000186T2
From which AH, = 89137 AF = AH, - 2.75 TlnT .0028 T? - . O O O O O O ~ ITS IT When T = 298, F = 82600 and I = - 7.09 - .00000031 T3 - 7.09 T Then A F " = 89137 - 6.33 T log T +
+
+
Substituting in values of T which are of interest from the standpoint of the arc process we find the results shown in Table 11.
TABLEI1 T 2000
3000 3500
AF"
41920 18580 6602
K 2.64 X IO-^ ,0444 .3872
70 NO .32
10.88 23.75
The data in Tables I and I1 and subsequent ones are graphically shown in Fig. I where the equilibrium percentage of NO is plotted against the absolute temperature. It is seen that the use of Nz and 0 2 in stoichiometric proportions results in comparatively little improvement over the use of air in the arc process. The use of C 0 2 and N z in the arc, however, should yield a gas containing a t least twice as great a concentration of NO as that obtainable from Nzand 02. This is contrary to the experimental results of Muthmann and Schaidhauf' who found that mixtures of N2 and COz gave the same yield of NO as mixtures of NPand 02. The conditions of their experiments, however, were not exactly those obtaining in an arc and further tests seem desirable. It should be pointed out that the conclusions of thermodynamic calculations do not indicate that a given result will be obtained-they merely state what result is the true one based on establishment of true equilibrium. LZ.Electrochemie, 17,497 (1911).
NEW HIGH-TEMPERATURE FIXATION REACTIONS OF NITROGEX
1491
Nitrogen, Hydrogen and Carbon The reactions of nitrogen and carbon with and without the presence of hydrogen are of considerable interest. The formation of cyanogen and hydrocyanic acid are represented by the following equations : (8) K2
+
2
(9) I / Z K2
C(gr.1 = (22x2 (g) I/Z HP C (gr.) = HCN(g)
+
+
Petcent
by
AF02ps = ~ Z O O O AF"zs8 = 28910
Volume
FIG.I
The data are taken from Lewis and Randall'. The data for reaction (8) involve some uncertainties for the value of the heat of reaction which is finally estimated as 32,000. The data for reaction (9) are obtained by combination of about twelve equilibrium measurements and compare well with the direct experimental values of von Wartenburg*. The use of reaction (8) seems t o be without promise a t any temperature and accurate calculations of the change of free energy with the temperature cannot be made. Attempts by von Wartenburg to form cyanogen in an arc operating in nitrogen have failed entirely. Apparently cyanogen is completely unstable a t all temperatures. 1
Loc. cit., pp. 590, 592. Z. anorg. Chem., 52, 299 (1907).
NORMAN W. RRASE AND BILL MACHEY
I492
The measurements of von Wartenburg on nitrogen, hydrogen and carbon permit the formulation of the free energy-temperature equation on the assumption that ACp is zero. This equation is':
- 10:3 T
AF" = 32000
(IO)
The experimental data of von Wartenburg are given in Table 111.
TABLE I11 T 1908
Percentage HCN
1.95 3.1 4.7
2025
2148
Using equation ( I O ) we have calculated the equilibrium concentrations a t temperatures up to 3500°K, an extrapolation of almost 1500'. The calculated results are given in Table IV.
TABLEIV T 298
28910
8.1
I000
2 I 700
I.g
2000
11400
3000 3500
%HCN
K
AF"
x x
,8321
I IO0
,00093
10-6
,0572
- 4050
-
10-22
1 ' 785
2.78 29.4 47 ' 2
Apparently the use of Nf and H 2in an arc struck between carbon electrodes or the use of mixed Nz and H2 passed through graphite used as the resistor of an electric furnace should result in the production of gas containing significant quantities of HCN. Nitrogen and Methane
The use of nitrogen and hydrocarbons at elevated temperatures such as are obtainable in the electric arc furnishes further interesting speculation. I n the case of methane reliable equilibrium data exist for its formation from the elements. Pring and Fairlie' have measured the equilibrium partial pressures for the reaction: (11)
C (graphite)
+ 2H2 = CHI (9)
These results have been considered by Lewis and Randall2 and the' free energy-temperature relation developed. If we proceed as follows:
+
2H2 (12) CH, = C (graphite) C (graphite) = HCK (g) (9) I / Z Nz 1/2 H z CHI = HCN 3/2 H2 (13) 1 / 2 N2
+ +
1 2
J. Chem. SOC., 101, 91 (1912) loc. cit., pp. 571-572.
+
+
NEW HIGH-TEMPERATURE FIXATION REACTIONS OF S I T R O G E S
I
493
The addition of the free energy equations for reactions (12) and (9) gives the free energy-temperature equation for reaction (13) as follows: (14)
AF" = 48,300 - 15.18 T log T .0000002 T3 15.7 T
+
+
.0008
T2
Calculation at a number of temperatures yields the data shown in Table V.
TABLE V T 298
K 1.8 X IO-~O
AF" 41720 17800 -22118 - 64748 -85268
1000 2000
3000 3500
%HCN
-
,000I 3 2
1.0
269.2 54,830 223J900
39.4 40 40
It is seen that, if equilibrium is established a t 2 0 0 0 " or above, a gas containing 40 percent of HCN and 60 percent of H2will result. This temperature permits the reaction to be carried on in other apparatus than the arc furnace and makes possible the employment of catalysts for the establishment of equilibrium. Much work has been reported on different methods of causing nitrogen and hydrocarbons to react and on suitable apparatus for this purpose. There is no question but that the reaction offers extremely interesting possibilities. Nitrogen and Acetylene Sufficient data for estimating the results of the high temperature reaction of nitrogen and acetylene' exist. Combination of the following relations permits calculation of the yield of HCN: (IS) (9)
1/2
(16)
1/2
1/2
+
AF" = - 2 7 2 0 0 5.4 T C2Hz = c'+ 1/2 Hz C (gr) = HCN (gas) AF" = 32,000 - 10.3 T Hz 1/2 AF" = 4800 - 4.9 T C2H2 1 / 2 N B = HCN (gas)
+
+
+
Substitution of appropriate temperature values yields the data in Table VI. TABLE VI Temp. 298
KP ,06363
1000
I .Oj
2000
3.52
%HCN ,018 34.4 63.8
Temp. 3000 3500
KP
%HCS
5.25
72.3 74.7
5.88
This reaction is of obvious interest and deserves attention from workers in the field of nitrogen fixation. Francis: Ind. Eng. Chem., 20, 27y (1928)
I494
NORMAN W. KRASE AND BILL MACKEY
Summalzl Several new high-temperature fixation reactions of nitrogen have been investigated with the aid of thermodynamics. I n the case of nitrogen and oxygen only a comparatively small increase in the concentration of nitric oxide results when stoichiometric proportions are used as compared with the result using air. The reaction of nitrogen and carbon dioxide in the arc should produce a concentration of nitric oxide more than double that when air is used. Three reactions of nitrogen yielding hydrocyanic acid appear to be promising. Nitrogen, hydrogen, and carbon should yield concentrations of HCN exceeding fifteen percent at the usual arc furnace temperature. The interaction of nitrogen and methane to form HCW and hydrogen is complete at zoooo absolute. The combination of nitrogen and acetylene to form HCK is quite marked a t relatively low temperatures and should yield a gas containing almost 7 0 percent HCN by volume a t the arc furnace temperature.
