ANALYTICAL CHEMISTRY
392 used as routine analysis ( 2 ) . Counting is carried out with 4chloro-2-methylphenoxyacetanilide as described in this note. ..I series of duplicate determinations has been performed, and Table I1 gives the deviations between these. Calculation of these results shows that a single determination is performed n i t h a standard deviation of 0.79%. The small difference from the statistical counting error (0.61%) shows t h a t no other factors have a serious influence.
ACKNOWLEDGhIEYT
The assistance of Jytle J@rn-Jensenand Yi.vian Larsson is gmtefully acknowledged. LITER4TURE CITED
B., Ottesen, J., and Zerahn, Ii.,Acta Physiol. Scand., 10, 195 (1945). (2) Sorensen, Poul, -4b.i~. CHEY.,26, 1581 (1954). (3) Ib& 27, 388 (1955). . ( 1 ) dmbrosen, J., Madsen.
R E C E I V EforD r e v i e r July 19, 1954. Accepted S o v e m b e r 23, 1954
Reduction of Nitroguanidine by Titanium(ll1) Chloride WARREN W. BRANDT’, JOHN E. DEVRIES, and E. ST. CLAIR GANTZ Analytical Chemistry Branch,
U. S. N a v a l O r d n a n c e
Test Station, Inyokern, China Lake, Calif.
The reduction of nitroguanidine by titanium(II1) chloride in 1 to 1 hydrochloric acid w-as studied with the addition of iron(II), in a 20 to 1 ratio to nitroguanidine. The redox reaction is 98% complete for consumption of 8 equivalents of titanium(II1). The route and mechanism of the iron “catalyzed” reaction was also studied. The end products of the reaction are guanidine and ammonia. It is proposed that the function of the iron(I1) is to stabilize, by complexation, hydroxylaminoguanidine and to make this intermediate susceptible to further reaction with titanium(II1). The reduction of nitroguanidine using a 20 to 20 to 1 ratio of titanium(III), iron(II), and nitroguanidine is quantitative and could be used for assay of nitroguanidine.
T
HE reduction of nitro groups with titanium( 111) has be-
come a standard analytical method for their quantitative determination. The usual procedure involves adding a n excess of titaniuni(III), boiling for a short period, and back-titrating with iron(II1) ( 3 ) . Nitro groups in compounds such as nitrobenzene require 6 equivalents of titanium per mole. I n general, the reduction is carried out in strongly acidic solution. The extension of this method to the nitramine group revealed that in these cases the nitro group required only 4 equivalents of titanium(II1). Kouba and coworkers ( 2 ) determined nitroguanidine quantitatively by this method. However, when they attempted to determine RDX (hexahydro-1,3,5-trinitro-s-triazine) by the same procedure the 4-equivalent reduction per nitro group was only 60% complete. They found t h a t by introducing iron(I1) to the reaction mixture the reduction was within lyOof theory for 4 equivalents. Zimmerman and Lieber ( 7 ) extended the investigation of titanium( 111) reductions to include several nitroammonocarbonic acids. When iron(I1) was introduced into thi! reduction of nitroguanidine and nitroaminoguanidine, they found t h a t the reaction then approached a total of 8 equivalents per mole instead of the usual 4. A new path of reduction was proposed to ekplain this phenomenon. Their proposal involved reductive cleavage of the nitroguanidine as the first step, and subsequent reduction of the nitramino group to a hydrazine. NHSO?
/
PiHN02
/
1 Present address, Chemistry Department, Purdue University, Lafayette. Ind.
iVHNO* +
C=SH H‘
6(H) -+ C=SH
+ 2H2O
(2)
H‘
Recently, Sternglantz ( 5 ) has ahon-n t h a t under weakly acidic conditions, using citrate buffer, nitroguanidine consumes 6 equivalents of titanium(II1) per mole. The current investigation was undertaken in order to study in more detail the iron(I1)-catalyzed reduction of nitroguanidine by titanium(II1) and to attempt to demonstrate the mechanism of the function of the iron(I1). APPARA‘TU S
Because of the instability of titanium(II1) in air, all reagent, solutions were kept under an atmosphere of carbon dioxide both in storage and in the burets. The reaction flasks were all connected to reflux condensers by ground-glass joints. T h e former were also equipped with a small inlet side arm for introducing a stream of carbon dioxide over the surface of the solution. This flow of carbon dioxide over t,he surface was started following the initial degassing of the acid to be used and continued through the final t,itration with iron( 111) without interruption. All carbon dioxide passing into the react,ion flasks was passed through bubblers, so that the rate of flow was readily visible at all times. RE4GE\TS ; i0 . 4 s titanium(II1) solution n a s prepared in 1 to 1 hvdrochloric acid. It is convenient to use titanium hydride (SIetal Hydrides, Inc., Beverly, Mass.) as recommended by Wagner et al. ( 6 ) . A 1 . O N iron(I1) solution was prepared in 1 to 1 hydrochloric acid from reagent grade iron(I1) sulfate. Reagent grade iron(II1) alum [Fe(XH4)(S04):.12H*0]x a s used to prepare a 0.3A\r iron(II1) solution in 1 t o 1 hydrochloric acid. All hydrochloric acid was Baker and idamson C.P. reagent, 3T to 38%. Nitroguanidine was prepared by recrystallization of a commercial sample from water. Potassium nitrate was Baker and .idamson reagent grade. Sitrosoguanidine (melting point 165’ C.) was prepared by the procedure described by Davis ( 1). Aminoguanidine hydrochloride was prepared from Eastman’s white label aminoguanidine bicarbonate. The titanium( 111) and iron( 11) solutions were standardized against r a t i o n a l Bureau of Standards potassium dichromate using sodium diphenylbenzidine sulfonate indicator. The iron(111)was standardized against the titanium( 111).
REDUCTIO? PROCEDURE
The weighed sample u a s dissolved in 50 ml. of 1 to 1 hydrochloric acid and the solution degassed with carbon dioxide for 5 to T minutes. The titanium(II1) and iron(I1) were then added and the solution was refluxed. The mixture was then cooled in a water bath and titrated with iron(II1) to a thiocyanate end point. The time of reflux !vas measured from the start of vigorous bubbling. In the reactions of potassium nitrate and nitrosoguanidine the
V O L U M E 27, NO. 3, M A R C H 1 9 5 5 1 to 1 hydrochloric: acid was degassed and the titanium(II1) and iron(I1) were added before addition of the weighed sample. All results reported have been corrected by means of a blank containing all the reagents except the compound being reduced. I n general, the blank encountered n-as equal to approximately 0.1 equivalent of reduction per mole. RESULTS AND DISCUSSIOX
Over-all Reduction of Nitroguanidine. The reduction of nitroguanidine in 1 to 1 hydrochloric acid was found t o requirt, 4 equivalents per mole in the presence of a 20 to 1 mole ratio of titanium(II1) to nitroguanidine when refluxed for 30 minutes as reported ( 7 ) . The addition of a 20 to 1 excess of iron(I1) caused the predicted increase in reduction (7'). Although Zimmerman and Lieber reported a maximum of i . 6 equivalents per mole, the iibove conditions gave consistent results averaging T.8 equivalents of reduction per mole. This represents approximately 08'34 reduction, which \vas deerned sufficiently close to qumtitative for the current investigation. A 25% decrease in the amount of titaniuni(II1) caused a decrease of about 0.2 equivalrnt per mole. Increasing the iron( 11) concentrat,ion or t h e time of i ~ f l u scaused no appreciable change. In ortlw to determine the function of the iron(II), a stepwise investigation of the reduction of nitroguanidine was undertaken. I n these experiments a calculated amount of titaniuni(II1) was added in order to provide a n integral number of equivalents per mole of nitroguanidine. I n one series, titanium(II1) was refluxed alone with t,he nitroguanidine for 5 minutes, followed by the addition of 20 to 1 iron(I1) and titanium(II1) and refluxing for Itn additional 20 minutes. I n the other series, the usual 20 to 1 excess of iron(I1) \vas added with the titanium(", refluxed for 5 minutes, then escess titanium(II1) was added followed by a n additional reflux period of 20 minutes. T h e over-all equivalents of tit.aniuni(II1) consumed per mole were determined. The results, shown in Table I, demonstrate t h a t t'he iron(I1) is not accomplishing any reduction itself a t any step in the process. I n the second series a 20 to 1 excess of iron(I1) is always present, but the total amount of reduction corresponds to the concentration of t,itanium(III) added initially.
Table 1.
Stepwise Reduction of Nitroguanidine
T i I I I I ) Alonr for Initial Reflux. Eq./Mole ______ Ti added Consumed
Ti(II1) a n d Fe(I1) Initial Reflux, Eq./.\Iolr Ti added Consullled-
for -
End Products of the Reduction. The second step in the investigation was the identification of the products of the 8-equivalent, reduction of nit,roguanidine. -1method was developed for the determinat,ion of any ammonia which might be formed during the reartion. First, crystalline potassium iodate was added until the disappearance of the last traces of iodine coloration. Sodium sulfit,e crystals were then added until the solution was free of iodine monochloride. Sitrogen gas was bubbled through the solution for 15 minutes in order to remove the excess sulfur dioxide. The solution was then transferred t o a Kjeldahl flask and ammonia distilled as in the usual Kjeldahl procedure. T h e cletermination was completed by titration with standard alkali. By utilizing the above met,hod it was possible to demonstrate the production of appreciable quantities of ammonia during the reaction. T h e results n-ere too variable to permit a formulation of the eract number of moles of ammonia for each mole of starting material, b u t the value was approximately unity. Ammonia would be expected as one of the products in the mechanism proposed by Zimmerman and Lieber (see Equation 1). It might
393 also be formed by hydrolytic degradation (Equation 3 ) ; however, the reduction is 90% complete or better, and therefore the latter reaction could not explain the quantity of ammonia formed unless the product of reduction were itself cleaved. This possibility was tested by putting aminoguanidine hydrochloride through the above procedure as a blank. Only traces of ammonia were found.
2- +
HSO, C'=SH H,O
+
CO,
\
+ 2SH:3 + K20
(3)
SH,
I n order to isolate other products of the reaction, a typical 8-equivalent-per-mole reduction was carripd out' and titrated. The iron was oxidized with hydrogen peroxide and extracted with isopropyl ether. The solution was then neutralized with sodium hydroxide until the hydrous titanium(1V) oxide precipitated. After filtration, a saturated aqueous solution of pirric acid !vas added to the filtrate. h heavy yellon- precipitate settled out on standing. -4fter being twice recrystallized from ethyl alcohol the solid melted a t 333" to 334" C. X-ray diffraction patterns agreed n-ith the published results for guanidine picrate (4)and with prepared standards. This demonstrates the presence of guanidine as a n end product of the 8-equivalent reduction of nitroguanidine. T h e amount formed was sufficient. t o classify it as a major product and not the result of a side reaction. T h e effect of hydrogen peroxide in 1 to 1 hydrochloric acid and subsequent neutralization n-as checked with aminoguanidine in a simulated reaction mixture, since aminoguanidine was the end product proposed by Zimmerman and Lieber (7'). S o appreciable quantit,? of guanidine \vas formed. T h e presence of guanidine does not fit into the mechanism proposed by Zimmerman and Lieher. A cleavage of the X-K hond at, some point during the reduction of nitroguanidine is required. T h e simple possibilities for reduction of nitroguanidine to guanidine plus ammonia are presented in Figure 1.
Figure 1. Possible reactions i n reduction of nitroguanidine to guanidine and ammonia 1-H sH
x-
SH-C-SHSOZ I
S" I
NH,-C-SH
I\ITH2-(!-SH2
(1)
KH,
H2O (10)
+ HSOI
KH
I'
~TH,-~--NH
+ KH~OH
394 T h e reduction of nitroguanidine to aminoguanidine followed by reductive cleavage to ammonia and guanidine (Reactions 1,2,3, and 11, Figure 1) would seem to be a plausible path. I t has been reported that aminoguanidine is not reduced by titanium(II1) ( 7 ) , but the reaction in the presence of iron(I1) had not been checked. A sample of aminoguanidine was introduced into the usual reduction misture for an 8-equivalent reduction and refluxed 30 minutes. T h e average of several runs gave a 1% reduction. It was therefore assumed that the path of reduction did not proceed through aminoguanidine. This eliminates Reactions 3, 10, and 11, Figure 1, and requires cleavage of a reduction product of niiroguanidine prior to the aminoguanidine stage; presumably a t the 2- or 4electron step. I n addition to generating guanidine, such a cleavage of a nitroguanidine reduction product would furnish a compound which is an intermediate in the reduction path from niti,ate ion to ammonia. Therefore, the reduction of nitrate ion in 1 to 1 hydrochloric acid by titanium( 111) was investigated. Over-all Reduction of Nitrate. I n t.hese reactions, the potassium nitrate used as the source of nitrate ion was added after all reagents were present. The results of closely agreeing determinations with titanium(II1) alone as the reductant gave 3.8 equivalents per mole, which was int,erpreted as a 4-equivalent reduction (Reactions 12, 13, Figure 1). I n the presence of a 20 to 1 ratio of iron(I1) in addition t o the t,itanium(III) the result$ of four determinations gave an average of 7.3 equivalents per mole, which was interpreted as an %equivalent reduction (Reactions 12-15, Figure 1). Considerable effervescence occurs early in the reduction, which might explain the relatively poor percentage of reduction observed. These results indicate that iron(I1) is effective in causing the reduction of nitrate by titanium(II1) to go beyond the 4-electron point in a fashion similar to its influence on the reduction of nitroguanidine. Two-Equivalent Reduction of Nitroguanidine and Nitrate. Thus far it appeared that iron(I1) was producing some species in the course of the reduction of nitroguanidine which titanium(111) was able to reduce. The determination of where this phenomenon takes place-Le., at, what stage in the reduction-was the next point investigated. I n order to enable a decision to be made concerning where the nitroguanidine reduction joins that of nitrate (Reactions 12 t o 15, Figure l), the latter was studied as a comparison. I n these reactions the previously described scheme of adding an integral number of equivalents of titanium(II1) per mole was utilized. I n some cases the iron(I1) was included for the initial reflux; in others it was added afterward with the excess titanium( 111). Upon refluxing a mixture of 2 equivalents of titanium(II1) per mole of nitrat,e ion with excess iron(I1) for 10 minutes, a reproducible 3-equivalent reduction takes place. This demonstrates the ability of iron(I1) to reduce the 2-equivalent reduction product of nitrate ion-i.e., nitrite ion. The solution is an intense brown, indicating the expected nitrosyl iron( 11) complex ion. When these results are compared to those in Table I, it becomes apparent that nitroguanidine and nitrate are completely different after reduction b y 2 equivalents of titanium(II1) per mole. This comparison eliminates Reactions 4, 5, and 6 , Figure 1, as possible steps in the reduction of nitroguanidine. If nitrate ion is refluxed 5 minutes with 2 equivalents of titanium(II1) per mole (Reaction 12, Figure l ) and this is followed with excess titanium(II1) and iron( 11), the reduction proceeds a total of 2.2 moles where 8 (Reactions 12-15, Figure 1) would be expected if the 2-equivalent product (nitrite) were stable. The instability of the nitrite ion is further demonstrated by repeating the experiment and changing the initial reflux time from 5 minutes t o 1 minute, and finally to merely standing 5 minutes without heating. T h e total equivalents of reduction obtained increase with the order of treatment listed, reaching a maximum of 6.4
ANALYTICAL CHEMISTRY equivalents per mole for the initial 5-minute cold standing period followed by reflux with excess titanium(II1) and iron(I1). X similar study with nitroguanidine points out further dissimilarities b e h e e n the 2-equivalent reduction products of nitroguanidine and nitrate. An initial 2 equivalents of titanium(II1) per mole refluxed for 5 minutes and followed by escess titanium(II1) for 20 minutes gave an average of 3.6 equivalents of reduction per mole, indicating escellent stability of the system with 2 equivalents of titanium(II1) per mole. However, when the initial 2 equivalents of titanium(II1) per mole are refluxed 5 minutes and followed with titanium(II1) and iron(II), the total equivalents per mole observed are only 4.9 instead of close to 8.0. This appears to contradict the immediately preceding conclusion concerning good stability. The most reasonable 2-equivalent reduction product of nitroguanidine would appear to be nitrosoguanidine (Reaction 1, Figure 1). I n order to compare the observed phenomena in the reduction of nitroguanidine, a sample of nitrosoguanidine was prepared. The decomposition of nitrosoguanidine in 1 t o 1 hydrochloric acid is so rapid that it was possible to obtain only an average of 1.5 for the expected 2-equivalent reduction with titanium(II1) and 5.1 with titanium(II1) and iron(I1). If nitrosoguanidine is added to iron(I1) alone in 1 to 1 hydrochloric acid, an intense brown color is formed immediately, indicating a very rapid cleavage of nitrosoguanidine to nitrite ion and guanidine (Reaction 6, Figure 1). I n the light of these results, it is apparent that the addition of 2 equivalents per mole of titanium( 111)to nitroguanidine does not produce appreciable quantities of nitrosoguanidine. This requires a postulation that the first reaction of nitroguanidine with titanium(II1) is essentially a &equivalent reduction (Reactions 1, 2, Figure 1). This postulate finds some support in the previous discussion of the stability of the products of the 2-equivalent reduction with titanium(II1). The observed contradictory results would appear reasonable if it is assumed that the first 2 equivalents of titanium(II1) reduce half of the nitroguanidine a total of 4 equivalents and this 4equivalent product is unstable. Further titanium would thus reduce the remaining nitroguanidine, giving a good over-all percentage reduction, whereas the titanium-iron mixture would have only half of the nitroguanidine remaining t o work on, and an over-all reaction of 6 equivalents would be the maximum b possible. h confirmation of the initial 4-equivalent reduction of nitroguanidine was obtained polarographically. It was first determined that in 1 to 1 hydrochloric acid nitroguanidine gave a reduction TTave which could be used. Then a series of three samples was examined. The first contained just nitroguanidine in 1 t o 1 hydrochloric acid and gave a 142 pa. wave a t Eli2 = -0.24 volt us. the mercury pool. T h e second solution contained an identical amount of nitroguanidine which had been refluxed 10 minutes with 2 equivalents of titanium( 111). This solution gave a wave of 71 pa. a t Eli2 = -0.30 volt. The third was identical except that 4 equivalents of titanium were used. It gave = -0.40 volt. T h e titanium(1V) a wave of 6.6 pa. a t present in the last two 'samples gave a wave a t -0.05 volt which did not interfere with the accuracy of the desired measurements. The mercury pool became contaminated by products of the reaction in the last two samples, which might explain the shifting potential. These data demonstrate the presence of one half of the nitroguanidine after the initial 2-equivalent reduction since the concentration is proportional to the wave height in microamperes. Thus an investigation of the 4equivalent reduction product became a necessity. Four-Equivalent Reduction of Nitroguanidine and Nitrate. Nitroguanidine was refluxed initially with 4 equivalents per mole of titanium(II1) and excess iron(I1). I n some cases this reaction was followed by the addition of excess titanium and a second reflux period of 20 minutes. The initial reflux time was
395
V O L U M E 27, N 0 . * 3 , M A R C H 1 9 5 5 Table 11. Four-Equivalent Reduction of Nitroguanidine“ No Subsequent Reduction Average Time of reflux, total min. eq./mole 2 4 10 30
3.6 3.8 3.9 4.0
Excess Ti(II1) Added for Second Reflux Period Time of Averaae initial ,reflux, total min. eq /inole 2.5 10 15 30
7.1 6.7 6.3 6.2
a Four equivalents of titaniuni(II1) and excess iron(I1) initially present in all runs.
Table 111.
Four-Equivalent Reduction of Nitrate Ion“
N o Subsequent Reduction
Time of reflux, min.
Average total eq./mole
2.5 10
3.9 4.0
Excess of Ti(II1) Added for a Second Reflux Period Time of Average initial ,reflux, total min. eq./mole 5.1 2.0 4.4 5 10
4.0
Four equivalents of titanium(II1) and excess iron(I1) initially present a t all runs
Table IV. Volume of Gas Evolved in 4-Equivalent Reduction of Nitroguanidine and Nitrate Ion Xitroguanidine, 1\11 a Fe(II) Fe(1I) absent present 35.4 35.8 a
12.9 13.6
Nitrate Ion. M I a Fe(I1) Fe(II) absent present 40.4 40.6
39,4 41.6
1.92 millimoles of compound used with 10 minutes‘ boiling in all cases.
varied in order to determine the relative speed of the reduction reaction and the decomposition of the product. The results are shown in Table 11. These results show t h a t in the presence of iron( 11) the 4-equivalent reduction product of nitroguanidine is stable in refluxing 1 t o 1 hydrochloric acid. For comparison, the same type of study was conducted with nitrate ion. I n this case the product of a 4-equivalent reduction was shown to be rather unstable. The results are summarized in Table 111. T h e r demonstrate that although the reduction is completed very quickly, the product is unstable and has completely disappeared after 10 minutes’ reflux. This is considerably different from nitroguanidine (Table 11). A further demonstration of the difference beh-een the 4equivalent reduction of nitroguanidine and nitrate ion is given by the following experiments. These also served to confirm the previously discussed difference in behavior of the 4-equivalent reduction product of nitroguanidine in the presence and absence of iron(I1). This series of experiments involved collecting the gas evolved during the course of reduction. The gas was swept through the svstem with carbon dioxide and collected over 30% potassium hydroxide. The results are given in Table IV. The difference in volume of gas produced when iron(I1) is present or absent showed a lack of identity for the 4-equivalent reduction products of nitroguanidine and nitrate. Thus Reactions 4 to 8, Figure 1, are eliminated. CONCLUSIONS
I t would no\v appear that the following mechanism takes place in the reduction of nitroguanidine. A 4-equivalent reduction occurs initially (Reactions 1 and 2, Figure l ) , and this is followed by a reductive cleavage of hydroxylaminoguanidine (Reaction 9, Figure 1). T h e latter reaction is effectively simultaneous with the subsequent reduction of the hydroxylamine (Reaction 15, Figure 1). The reduction of hydroxylamine hydrochloride was investigated separately and found to consume two equivalents of reducing agent. The reaction with titanium(II1) proceeded rapidly. The reaction with iron(I1) proceeded slowly. Although a small amount of reduction of hydroxylamine would be
expected a t the 6 and 7 equivalent steps in Table I, it apparently is not formed in any appreciable quantity. This could mean t h a t the reaction did not go through hydroxylamine, or t h a t it utilizes the last 4 equivalents in a concurrent fashion, producing no appreciable concentration of hydroxylamine. To eliminate hydroxylamine from the reduction path would require that no cleavage of nitroguanidine or its simple reduction products could occur a t an S--PI’ bond, since these would all give precursors of hydroxylamine. This leaves only a complex procedure in which a coupled product is formed which cleaves to give guanidine and ammonia and allows an over-all 8-equivalent reduction to occur. I t seems more reasonable to assume t h a t the reduction goes through hydroxylamine rapidly without producing a final measurable concentration. The 4-equivalent product is the essential point in the function of the iron(I1). Without the iron(II), titanium(II1) is unable to bring about further reduction and the product is rapidly decomposed, probably by hydrolytic cleavage to hyponitrous acid which also decomposes rapidly. This fact was confirmed by the data in Table IV and by mass spectrographic analysis of the gas from a 4-equivalent reduction of nitroguanidine by titanium( 111). With iron(I1) present, the compound is stabilized to hydrolytic cleavage, and altered in such fashion as to permit reductive cleavage by additional titanium(II1). Although it is a hypothetical compound, hydroxylaminoguanidine is the simplest 4equivalent reduction product t o propose. It is expected to be unstable; thus it would be necessary only to have this compound stabilized by the presence of iron( 11)in order to make its postulation logical. I t has been shorm that large excesses of iron(I1) are necessary in order to approach an 8-equivalent reduction. The total equivalents of reduction increases asymptotically toward 8 as the concentration of iron(I1) increases. This would indicate also that the function of the iron(I1) is not catalytic. It appears rat.her t o be the type of function which would be expected if the excess iron(I1) were affecting a displacement of equilibrium. The most reasonable stabilization of the hydrosylaminoguanidine vould be through some interaction compound, possibly of the nature of a metal-organic complex ion. I t would be expected that such a species would he quite unstable in 1 to 1 hydrochloric acid and that large excessep of iron(I1) would be necessary t,o promote its format,ion. I t is not unreasonable to assume that such interaction could modify the electronic distribution in the hydroxylaminoguanidine sufficiently to facilitate a reduct,ive cleavage of t,he S--S bond by titanium(II1). Attempted isolation of the postulated hydroxylaminoguanidine \\-as unsuccessful. Reaction mixtures, upon neutralization, eff erveeced considerably, and guanidine waF: identified in appreciable quantities. This is the expected product of hydrolytic cleavage of hydroxylaminoguanidine. The compound will probably have to be isolated in combination with iron(II), the only form in which it has been found to be stable. ACKNOWLEDGMENT
This paper is published with the permission of W. B. McLean, technical director of the U. S. Naval Ordnance Test Station. LITERATURE C1TE.D (1) Davis, T. L., ”Chemistry of Powder and Explosives,” p. 392, Wiley, New York, 1943. (2) Kouba, D. L.; Kicklighter, R. C., and Becker, \T’ W., A 4 ~ ~ ~ . CHEM.,20, 948 (1948). (3) Siggia, S., “Quantitative Organic Analysis via Functional Groups,” p. 82, Wiley, Kew York, 1949. (4) Soldate, A. M., and Noyes, R. M , , ANAL.CHEM.,19, 442 (1947).
( 5 ) Sternglanta, P. D., Thompson, R. C., and Savell, W. L., Ibid., 25, 1111 (1953). (6) Wagner, C. D., Smith, R. H., and Peters, E. D., I b i d . . 19, 982-4 (1947). (7) Zimmerman, R. C., and Lieber, E.. I b i d . , 22, 1151 (1950). RECEIVED for review September 7 , 1954. Accepted Sovember 29, 1954
