Sept., 1960
EFFECTS OF CHLORATE ON THERMAL STABILITY OF AMMONIUM PERCHLORATE1309
can be explained only by assuming translational motion of polymer chains in the crystals. Acknowledgments.-The authors wish to thank Mr. D. R. Carpenter for performing a large part of the n.m.r. measurements. They are also in-
debted to Dr. A. C. Martellock for supplying the silicone sample. The helpful discussions with Drs. Bruno Zimm and F. P. Price are appreciated. Miss Ann Warner was most helpful in some of the numerical calculations.
THE EFFECTS OF A CHLORATE IMPURITY ON THE THERMAL STABILITY OF AM3XONIUM PERCHLORATE' B Y J.
c. PETRICCIANI, s. E. WIBERLEY,w.H. BAUERAND T. w.CLAPPER
Department of Chemistry, Rensselaer Polytechnic Institute, Troy, N e w Y o r k , Research Department, Amerzcan Potash and Chemical Corporation, Henderson, h'evada Recezved Aprzl 9, 1960
Differential thermal analysis was used to determine the effect of chlorates on the thermal decomposition of ammonium perchlorate. When 0.1% potassium or sodium chlorate was added as an impurity, the decomposition temperature of ammonium perchlorate xas lowered approximately 150". The differential thermal analysis technique is sensitive to 0.00175 chlorate impurity in ammonium perchlorate. The amount of chlorate present in the ammonium perchlorate also can be determined colorimetrically with the same degree of accuracy. Chlorate impurities in ammonium perchlorate also can be determined by infrared absorption techniques; but the minimum amount detectable is in the region of 0.87,. A general theory is presented to account for the marked effect that chlorates have upon the thermal decomposition of ammonium perchlorate.
Introduction Ammonium perchlorate is an extremely strong oxidizing agent and is used as the crystalline inorganic oxidizing agent in the composite propellant type of solid rocket fuels. Since the presence of chlorates was thought to cause a premature decomposition of ammonium perchlorate, the main objective of this investigation was to determine the effect of a chlorate impurity on the thermal stability of ammonium perchlorate. Materials and Methods Ammonium perchloratr obtained from the Fisher Scicntific Company was analyzed colorimetrically and found to contain, less than 0.001% chlorate impurity.z The Fisher ammonium prrchlorate %*asused as the starting material aithout further purification since it contained a negligible amount of chlorate as compared with the amount of chlorate d d e d in prepariilg the samplrs. Samples of ammonium prrchlorate containing 0 01, 0.05, 0.10, 0.50, 1 00, 5 00 :tnd 10.OO~ochlorate impurity wrre prepared from the st:trting material. The chlorate was designated as weight p e r cent. potassium chloratr in ammonium prrchlorate. The htrnplrs were prc,pared h v introduving the chlorate, as a potassium chlorate solution, to :I sollition of the starting material. The final solution w:ts slom~lv evaporated to (Ir\.ness. T h r final drv wright was approsimatelv 20 g. It was thought desirable to have a sample of ammonium !wrrhlorate whirh was free from even a trace of a chlorate impurity to be used as the basis for the comparison with ammonium perchlorate containing chlorates as an impurity. .ipproximately 40 g. of the Fisher ammonium perchlorate, rontaining less than 0.001To chlorate, was rerrvstallized thrre times from water. The ammonium perchlorate. thus obtained was analyzed colorimetrically and found t o be free of anv trace of a chlorate impurity.2 Table I lists the percentage of chlorate impurity in each sample as found by colorimetric analysis.2 The apparatus and procedure used were similar to those developed by Gordon and CampbelL3 Finely divided silica (1) Rased on a thesis by John C. Petricciani submitted t o t h e Department of Chemistry, Rensselaer Polytechnic Institute, in partial fulfillment of t h e requirements for the degree of Bachelor of Science, May, 1958. a n d it Research Project a t American Potash and Chemical Corporation, August, 1957. (2) C. Eger, Anal. Chern.. 27, 1199 (14;:). (3) S. Gordon a n d C. Campbell, Anal. Chem., 27, 1102 (1955).
TABLE I JVEIGHT
% KClOa IN
SAMPLES
Sample no.
% KClOi added
% KC101 analyzed
1 2
0.000 .002 ,010 .050 ,100 ,500 I ,000 5,000 10.000
0.000
3 4 5 (i v
8
0
,001
,009 ,040 ,082 . :370
s70
...
...
which previously had bren heated to 600" : ~ n dcoolcd to room temperature was used as the reference substance. A length of thin-walled Pyrex glass tubing drawn out a t one end was used as a thermowell. The copper block was nickelplated in order t o prevent oxidation a t the high temperat u r e ~ . The ~ differential temperature was plotted by :& Sargent recorder as a function of time. The sample, temperature, measured by mrans of a thermocouple and a potentiometer, was recorded a t intervals which corresponded to respective timing marks of the chart on which the differential temperature was recorded. The recorder was set a t a total span of 5 millivolts. A chart speed of 0.2 inch per minute was found to be suitable. A constant heating rate of about 7" per minute was attained with this equipment. The limit of error in reading the potentiometer was &0.1 mv., or & 2 O . The graph on the Sargent recorder could be read to the nearest 0.01 mv., an error of ~ k 0 . 2 " . The infrared spectra of the samples containing chlorates, pure ammonium perchlorate and pure potassium chlorate were determined with a Perkin-Elmer Model 21 Recording Infrared Spectrophotometer using the KBr technique.5 Several low percentage chlorate samples were measured using as much as 7 mg. in 100 mg. of KBr in an attempt t o get the chlorate absorption band t o appear, but the perchlorate absorption band merely spread out and covered any increased absorption due to the chlorate ion. The 0.5 mg. per 100 mg. of KBr concentration was used for the subsequent spectra since the resolution of the perchlorate band was very good a t that concentration. (4) R. A. Powers a n d N. Hackerman, THISJ O U R Y X L66, , 187 (1952). ( 5 ) hl. M. Stimson a n d M. J . O'Donnell, J . Am. Cheni. Soc., 74, 1803 (1952).
J. C. PETRICCIANI, S. E. WIBERLEY, R. H. BAUERAND T. W. CLAPPER
1310 t 20
0.000%
-4
0.001%
-2:-
2
i
V
J
I! c
0.082 9.
5a
0-
W lL
k !
6
D
0
11-11 I
0
I00
200
SAMPLE
300
TEMPERATURE
400
t
D
("C),
Fig. 1 -Differential thermal anal centage of KCIOI and the sample number are shown at the beginning and end of each curve, respectivelv. Tht. limit of the A T scale for each curve is +20 and -20 above and below the baseline, respectively.
Experimental Results The representative D.T.A. curves presented in E'ig. 1 clearly indicate the presence of an iiicreasingly strong exothermal reaction after lattice transition (24.2') as the amount of chlorate impurity increases. Detectable decomposition occurs only after the lattice transition in all the samples tested. In the region of 0.1% potassium chlorate impurity, the heat evolved after lattice transition is great enough to initiate complete thermal decomposition of the sample. This is an effective 150" lowering of the thermal decomposition temperature of pure ammonium perchlorate, whivh decomposes at -1.00". A comparison of the D.T.-I. curve for pure ammonium perchlorate with that shown by Gordon aiid Campbell3 indicates that their saniple mag have contained chlorate as :HI impurity. AIarkowitz6 also has pointed out that the decwmposition of potassium perchlorate proceeds partially through a chlorate iiitcrmediatc. Decomposition is indicated hy the abrupt chaiige iii the differential temperature at the elid of each rurve in Fig. 1. Curves 4 aiid 5 of Fig. 1 are discontinuous because the heat evolution after lattice transition was too rapid to follow accurately. Since the furnace was heating a t a coiihtaiit rate of 7" per minute and did not reach the sample temperature for several niinutes, the sample cooled somewhat after reaching its peak temperature for those cases. Thc curve was resumed a t the point (b)
31. 31. NarbonitL,
1111sJ U L I I \ ~ L ,
61, W i (l!l57).
Vol. 64
where the sample temperature began to increase again. The infrared spectra of ammonium perchlorate, potassium chlorate and a 5% potassium chlorate sample are presented, in Fig. 2. The absorption band due to the chlorate ion (990-960 cm.-l) does iiot interfere with that of the perchlorate ion (11601060 cm.-l). The 3% potassium (ahlorate sample shows the chlorate absorption band along with the ammonium perchlorate spectra. However, the appearance of the chlorate absorption hand in a sample of ammoiiium perchlorate containing chlorates did not occur below 0.8% chlorate. This precludes using the infrared method as a rapid analytical technique for the determinatioii of trace amounts of chlorate in ammonium perchlorate. A comparison of the infrared spectrum obtained for ammonium perchlorate with that presented by Miller and Wilkiiis shows that the absorption bands are more distinct and sharper using a KBr window than with Suj01.7 The most distinct difference is the resolution of the perchlorate ahsorptioii hand of Miller aiid TTilkins into three components
Discussion h discussion of the decomposition of pure aninionium perchlorate is presented by Bircumshaw and Xewman, and an electron transfer mechanism for the decomposition is suggested.8 The region of 0 to 1% potassium chlorate is considered in the following discussion and a solid solution of potassium chlorate in ammonium perchlorate i\ assumed. As in the case of the perchlorate ion, :in electron transfer from a chlorate ion to an interstitial ammonium ion seems probable a h the initial step 111 decomposition
+
C 1 0 ~
YHd+intcrstitizI
SH,
---+ S H , SlI
--+
+ ('10~ +H
Although Bircumshan- and Nc\vniaii iugge.st that the chlorate radiral is stable within the crystal lattice and may function as an elet-tron trap similar to a c~n~espoiiding perrhlorate radical, it seems quite uiilikely that a chlorate radical n o d d irot deconipose in the temperature range of the lattice trail-itioii of ammonium perchlorate in view of the fact that ammoiiium chlorate itself ii iiiistahle in the region of 100". Thc followiiig step'; iwc tentatively propwed as likely in the dwompoiition 01 the chlorate radical, I\ here rcnc.1 io11 I i\ ahoilt t hernially lieu t r d
+ + 1 1 2 0 +H2CIO + ('IO +1 * ( 7 0 >
r10>--+ ClO ('10 H2('1O2
(1)
0 2
+
Ht I
( t$ )
'l'hc l i 2 ( )111 reachtioil 2 appear\ :tb :t d e ( ~ ) m p o ~ l t ~ o ~ l pioducat of perch1or:tte aiid amnioiiiiim 1 0 1 1 ~iil the niechanisni of Birmnishaw aiid Neil nian. Thk, then, is a possible route for the formation of both chloric aiid hydrochloric acids hich are liited among the decompositioii product 5.' Xnother possibility is the reacatioii of two C10 i a d i c d ~to form hIiller and C I1 U'ilhins
I n n / ChPm
24, 1283 i l ' ! i L )
1, L Rirorimshair arid B H. Nenriinn P r o , R o v Soc ( L o r d o r i ) . am, 11s (ICGZJ, A227, 11: ( I O X ) (8)
Sept., 1960
EFFECTS OF
C H L O R A T E ON
THERMAL STABILITY
OF
1311
AMMONIUM PERCHLOR-4TE
chlorine and oxygen, both of which are also listed among the decomposition products.8 ('10
+ c10
+c1,
+ 0,
7.1 hcal
-1lthough complete thermal data for all the reactions are not available, a moderate over-all exoposition scheme for pure ammonium perchlorate as presented by Bircumshaw aiid Semman. and that just described for ammonium perchlorate containing a chlorate impurity results in these conclusions: 1. The production of a chlorate radical requires three steps in pure ammonium perchlorate, while only one step i q needed in the cases where chlorates are preqeiit. 2 . Sublimation competes with decompositioii in both the pure aiid chlorate-contaminated case, hut sublimation is reduced iii the latter because of the poqsihility of the potassium ion replacing the ammonium ion. i n which case proton transfer is impossible. 3. Decomposition occurs mainly on the surface of the crystal in pure ammonium perchlorate with the elimination of a perchlorate radical. X chlorate radical is not required to be on the surface in order to decompose according to the decomposition whenie just described. More pronounced decomposition within the crystal should he evident in the chlorate containing case, and the heat lost to the outside will be less. The above consideration3 suggest that decomposition should be more proiiouiiced in ammonium perchlorate containing a chlorate impurity than in the pure compound. The decomposition reaction appears to take place nnly after lattice transition. The transition is from the orthorhombic to the cubic form and sub5equently the lattice distances are r e d u ~ e d . Dur~ ing the lattice transition the ammonium ions, especially thwe which are interstitial, may have a hetter chance of interacting with both the per(ahlorate and chlorate ions, which initiate. decompniition as previouily indiratrd. Curve 1 of 1;ig. 1 qhox1-s oiily a slight tendency toward an exothermic reaction after lattice transition, as indicated hy the gentle rise of the differential temperature baseline ( A T T = approximately 0.03). Thiq, howver, may he attributed t o ;t. rhanpe i l l the qpecific heat of the arnmoiiiuni perchloratt. after kit tire tr:tnsitioii. -1 movement of the diilferential temperature baseline greater t hail about, 3'' m-cr :t ianiple temperature range of looo n-as thought to he -igiiificant in view of thc :iho ye. Curre 2 of Fig, 1 iiidica t hn t 0.00 1% c lilora t e impurity ran bc- detected by ll.T.AL A c.liaiige in -pecific heat of the sample after lattice transition would account for an increasing baseline, hut the tmseline of curve 2 increases more rapidly than n-ould be expected from the specific heat effect 1vhic.h was evident in curve 1. Noreover, a maxi-
, i
i
NH4CI04 * 5% KCIO3
0
i>
1
2
Fig. ?.--Infrared
I -b L
-Ap-.
-
~
6 8 IO WAVELENGTH I N MICRONS.
~
12
L
-
14
spectra of S€LC104,KCIOI and NH4C10a containing 5(7, KClOn.
mum is reached in curve 2 which can be explained only on the basis of an exothermic reaction in the sample, especially in view of the fact that iii samples containing increasing concentrations of chlorate impurity a more and more pronounced heat evolution appears in the same region as indicated in curve 2 . It is assumed that decomposition in the 200300" region occurs largely through elimination of perchlorate radicals produced on the hurface of the crystals, and only slightly through the mechanism which would produce a chlorate radical in the interior. Decomposition iiivolving the chlorate radical is suggested for curves 2 through 5 of Fig. 1. The exothermic tendency increases with the amount of chlorate impurity as would be expected nccordiiig to the scheme presented. Finally, in the region of 0.17; chlorate impurity, the heat evolved due to the cvmhiiied effects of chlorate and perchlorate decompoqition hecomes sufficient to induce thermal breakdown of the ammonium and perchlorate inns themselvei, and complete decomposition results as is evident in Curve 6 of Fig. 1. S o attempt \\as made to determine any of tlie thermodynamic properties of any of the decomposition intermediates, but such data would he helpful in a critical analysis and evaluation of the decomposition scheme proposed.
