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Gas phase thermal decomposition of ammonium perchlorate
GAS PHASE THERMAL DECOMPOSITION PERCHLORATE
OF AMMONIUM
J. R. MAJER AND M. SMITH Chemistry Department, University of Birmingham, Edgbaston, Birmingham 15 The decomposition of ammonium perchlorate in the gas phase at elevated temperatures has been studied by allowing the vapour from a sample of subliming ammonium perchlorate to pass through a secondary furnace maintained at a higher temperature than that of the subliming ammonium perchlorate. The reaction occurring in the secondary furnace was monitored using a pin-hole leak connected to a mass spectrometer. The effects of temperature and foreign gases have been studied and the major products identified. A reaction scheme is suggested to account for the variation in product yield with variation in experimental parameters.
Introduction BmCUMSHAW and Newman 1'2 made an intensive investigation of the thermal decomposition of ammonium perchlorate (AP) and showed that it decomposes in the solid phase (the low temperature decomposition)and sublimes at temperatures below 300°C. They also found (hat above 350°C a different decomposition reaction takes place. Galwey and Jacobs 3 suggested that this high temperature reaction consisted of dissociation of AP into ammonia and perchloric acid, followed by their evaporation into the gas phase and subsequent decomposition of perchloric acid and oxidation of ammonia. The possibility of making a detailed study of the low temperature thermal decomposition of AP using a mass spectrometer was ~;rst demonstrated by Heath and Majer 4. A solid sample of AP was heated in the ion source of an A.E.I.M.S.2 mass spectrometer. The disadvantage of this method is that the rate of decomposition and sublimation cannot be kept constant and the small size of the furnace and sample did not permit prolonged experiments. This investigation has been extended by other workers s-7; however, their work has been directed towards elucidating the low temperature decomposition. It should be possible to investigate the high temperature reaction by subliming a ~mple of AP and examining the ensuing gas phase reactions by means of a mass spectrometer
coupled to the reaction system. The present paper describes such a system and the results obtained. Experimental
Apparatus Any investigation of the gas phase thermal reactions of AP must involve a procedure which permits differentiation between gas phase and solid phase reactions. The most convenient way of doing this is to pump gaseous species away from the decomposing AP and examine their interactions in a region isolated from tbe decomposing solid. The situation is complicated by the fact that AP will give rise to a solid sublimate at a cool surface in a reaction vessel. Consequently the entire reaction system must be kept at the temperature at which AP is d~omposing to prevent any sublimate being formed. A flow system was used in which gaseous products from the decomposing AP were continuously pumped through a reaction chamber of variable temperature, situated some distance from the solid AP. The gas phase reaction was continuously monitored by a mass spectrometer. A diagram of the system is shown in Figure 1. A pin-hole leak 00045 in. in diameter at the end of 0.1 in. i.d. Pyrex tube led to the ;,on source of an A.E.I.M.S.2 mass spectrometer through 635
ol. 13
J. R. MAJER AIqD M. SMITH
636
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a ~ in. hole drilled in the ion repeller plate of the ion source. The section of the apparatus, in which AP was housed, was heated to a maximum temperature of 320°C by a furnace, mounted on telescopic runners, which enabled it to be slid over the sample. Besides heating the AP, the furnace heated the re~t of the system up to the gas phase reaction chamber and prevented any sublimate being deposited. The gas phase reaction chamber was heated to a maximum temperature of 520°C by a 32 ohm, 28 s.w.g. Nichrome resistance wire wound directly on to the chamber and extending for 3 in. at each end of it. ,The rest of the reaction system between the reaction chamber and the cold trap was heated by Electrothermal heating tapes HT 301. The temperature in this region was measured at a point about two inches above the Housekeeper seal by an iron--constantan thermocouple.
Method About 10 g of AP, in a small Pyrex boat, was first dried and initially decomposed by heating at 300 ° to 320°C for one hour under vacuum in a separate pre-heat apparatus. The sample was then transferred to the reaction system.
The reaction system and mass spectrometer were pumped out and the reaction chamber furnace was set to give a temperature of about 520°C, the region between this and the cold trap baked at about 320°.C and the entire apparatus left pumping in this condition overnight. After setting the long tubular furnace at a temperature of up to 310°C, it was slid over the sample of AP. The temperature of the reaction chamber was next reduced until it was slightly higher than the AP decomposition temperature to ensure that no sublimate of AP was deposited. Analysis of the reactants and products of the gas phase reaction was performed in two ways. First the total mass spectrum of all products and reactants was recorded by scanning the m/e range~ 12-75 and 75-140 separately at a resolution of about 200 with an ionizing electron voltage of 40 eV for the role range 12-18 and 70 eV for the role range 18-140. Alternatively, the ion current at a specific role value could be recorded continuously. When using the total spectrum method of analysis, it was necessary to record spectra at the normal temperature of the reaction chamber before and after obtaining a spectrum at a higher reaction chamber temperature. In the case when the reaction was monitored
December 1969
637
GAS PHASE THERMALDECOMPOSITION OF AMMONIUMPERCHLORATE
by following the change in an individual peak, the peak corresponding to the compound being examined was continuously recorded with the reaction chamber at its normal temperature. The temperature was then raised and the peak height still continuously recorded. After the maximum reaction chamber temperature had been reached, it was returned to its normal value still with the peak height being recorded. The mass spectrometer was calibrated for N20, 02 and NI-Ia. The calibrations were used to calculate the calibration equations for N 2, HCI and NO 2 using thei.~ known a mass spectrometric relative sensitivities. The calibration equations calculated for NH2CI, HCIO, CIO 2 and Ci 2 were only very approximate as their relative sensitivies were estimated.
A peak at m/e 51 was shown to be due to the ion NI-I2CI+ and not CIO + from an examination of the spectrum produced using deuterated AP. Ammonia was first detected at 270°C but was only present in measurable quantities at 300°C. No perchloric acid was detected at any temperature. It was most certainly present because a sublimate of AP was found in the region between the sampling leak and the cold trap of the reaction system. The products of the low temperature decomposition at this optimum temperature of 300°C are given in Table 1 together with their corresponding partial pressures in the region just before the sampling leak. Monochloramine and hypochlorous acid have not previously been reported as products of the decomposition of AP.
Materials AP was supplied by B.D.H. Deuterated AP was prepared by recrystallizing AP twice, from deuterium oxide (KochLight Laboratories, England). 0~ Nitrogen, white spot grade, nitrous oxide, anaesthetic grade, oxygen and hydrogen were obtained from the British Oxygen Co. Ltd, England. Nitric oxide and carbon tetrafluoride were supplied by the Matheson Co. Inc. U.S.A. Ammonia and carbon dioxide were obtained from I.C.I. Ltd, England. Results
The products of the low temperature decomposition and sublimation A sample of AP was heated to temperatures in the range 240° to 300°C with the rest of the reaction system maintained at the same temperature to avoid deposition of a sublimate and a mass spectrum of all the gaseous compounds taken at each temperature. Most of the compolmds were unambiguously identified by their parent ions or isotopic parent ions in the mass spectrum. Ammonia could not be identified by its parent ion NI-I at m/e 17 as the ion OH + from water completely obscured it. Fortunately its fragment ion NH + at m/e 15 was free from major interference by any other ion and was used to show the presence of ammonia 8.
TABLE 1. Products of the low temperature decomposition and sublimation of AP
Product
NH3 H20 Oe HCi N2 N20
Partial pressure, x 1O- 3 t o r t
17"2 83"1 16"1 3"3 0"6 9'7
Product
Partial pressure, × 1O- a torr
NOe NHeC1 CIO2 HCIO CI 2
3'5 1'9 0'1 01 2'4
The products of the gas phase reaction The nature of the products of the gas phase reaction of the compounds from decomposing AP was determined by comparing in another run the mass spectrum of the AP heated at 300° with that obtained when the reaction chamber was raised to 468°C. The difference was a measure of the extent and nature of the gas phase reactions. The partial pressures "of the products of the gas phase reaction as they emerged from the reaction chamber are given in Table 2. Ammonia, which decreased in amount when the reaction chamber temperature was increased, was thus involved in the gas phase reaction. It can be seen from the product analysis that it was converted to N20, NO2 and a small amount of Ne. Nitric oxide was not detected.
J. g. MAJi~
638
TABLE 2. Products of the gas phase decomposition of A P
Product
Partial pressure x 10- a tcr~
H20 NO 2 O2
18"1 8-7
N20
1"4 06
Product
Partial pressure x 10- 3 tort
N2 HCIO C102 0 2
0-3 0-8 0-7 10
1.6
HCI
Initial NH3 pressure 29-3 x 10-3 torr. NH~ reacted 9.6 x 10- 3 tort.
Reaction chamber temperature 468°C.
Effect of temperature The peak height could either be continuously recorded, whilst the rc=~=tion chamber temperature was continuously changing, or be recorded whilst the temperature was increased in steps. Both methods were used for CI[, and the two curves relating changes in peak height with temperature were virtually inL
o10
distinguishable. The more rapid method in which the temperature was continuously changing was therefore used. A number of runs were carried out with a sample of AP heated at 300 ° to 305°C, by following each significant peak as the temperature of the reaction chamber was increased from 330 ° to 530°C. The changes in partial pressure with temperature are given in Figures 2(a) to (d). The change in partial pressure represents the amount of material produced in the gas phase reaction in all cases except ammonia. As separate runs were necessary to follow the change in each product the yields are not inter-related. Figure 2(a) shows that the amount of ammonia reacting increased with temperature. The extent of this can be seen in that approximately 50 per cent of the ammonia entering the reaction chamber had reacted at 500°C. The yields of nearly all the products of the gas phase reaction increased with temperature within the range studied. The exceptions
Initial NH 3 pressure = 14-7x 10-3 torr
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(d) 500
December 1969
GAS FELASE~HERMAL DECOMPOSITION OF AMMONIUM PERCHI~RATE
to this were chlorine dioxide and hypochlorous acid which initially increased but began to decrease at 475 ° and 510°C respectively.
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The effect of an inert gas on the yield of CIOe was examined by observing the change in yield with temperature in the manner described previously, but in the presence of nitrogen. The results are illustrated in Figure 3 in which partial pressure of CIOz is plotted against temperature with and without nitrogen present. It can be seen that with an inert gas present the maximum yield of CIO2 was increased and occurred at a higher temperature.
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Vol. 13 Oxygen was added to the reaction mixture in the manner previously described. The general effect was to decrease the yield of all products in a manner similar to that of an inert gas. In particular it is to be noted that the yields of oxides of nitrogen decreased.
J. R. MAJER AND M. SMITH
640
The effect of increasing amounts of nitric oxide on the reactants and products of the gas phase reaction is shown in Figures 4(a) to (d). The amount of ammonia which had reacted reached a maximum at roughly the same pressure of nitric Oxide as the maximum in the yields of H20, NO2, N 2 0 and HCL At the same pressure of nitric oxide virtually no CIO2 and HCIO were formed. The yield of chlorine reached a maximum at a pressure of nitric oxide slightly lower than that at which maxima for other products were obtained. It can be seen that the yield of oxygen decreased but then increased at a higher pressure of nitric oxide. The increase could be accounted for by the contribution of NtSO + to the peak at m/e 32. TEe effect of nitric oxide on the yield of nitrogen could not be determined, as an impurity of four per cent nitrogen obscured any small change in the yield. The yields of H20, C12, HCI, N 2 0 and NO2 were found to decrease when excess nitric oxide was added. ,_ 10
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F~GUgE 5. Variation in molecular chlorine yields with added ammonia pressure
When ammonia was added to the reaction mixture, it was found that nearly all products decreased, probably due to suppression of the dissociation of AP as shown by Inami et al.9. The exception to this was NH2CI at role 51 in the mass spectrum of the products which showed a rapid increase. At the same time a rapid decrease in the yield of chlortr)e was found as can be seen from Figure 5. The reaction NH3 + Cl2 --, NH2CI + HCI is known to occur at room temperature t°.
Discussion The gaseous products of the low temperature decomposition of AP under the conditions used in this investigation were NH3, HCIO4, H20, N2, 02, HCI, N20, NC~, NH2Cl, Cl 2 and small quantities of HCIO and CIO 2. The product analysis is not entirely in agreement with that of other workers t. Chlorine dioxide has been found which probably arises as a result of thermal decomposition of perchloric acid produced in the dissociation of AP as suggested by Dod~: 1.12. When the temperature of the gases from decomposing AP was increased, the concentration of NI-I3 decreased and that of the other products of the low temperature reaction, except perchloric acid which was not observed, increased. As can be seen from Table 2 the products of this complex gas phase reaction are similar to those of the low temperature decomposition (Table 1). However, NO 2 is the most abundant oxide of nitrogen amongst the products of the gas phase reaction rather than N20 which is the most abundant oxide of nitrogen in the low temperature decomposition. The relative abundance of chlorine oxides is also greater in the products of the gas phase reaction than in the low temperature decomposition products. We will now discuss the significance of these observations and the effect that various parameters have on the yields of both these groups of oxides in the gas phase reaction.
The reactions of oxygen-chlorine compounds Chlorine dioxide has been reported as a product of the thermal decomposition of perchlor/c acid 13. As has been found in the present investigation it appears that perchloric acid from dissociating AP has undergone decomposition probably by the reactions: HO--CIO 3 + M ~ HO + CIO3 + M
[1]
December 1969
GAS PHASE THERMAL DECOMPOSITION OF AMMONIUM PERCHLORATE
CIO3 -~ C102 + O
[2]
It has been suggested that CIO3 decomposes to give CIO CIO3 - , CIO + O2
[3]
Chlorine dioxide may also be formed from CIO4 radicals produced by attack of OH radicals on perchloric acid ~4. HO + HOClO 3 ----,HzO + C104 CI04 ~ CIO~ + 02
[4]
[5]
The decrease in the yield of CIO2 at high temperatures in the present investigation could be due to two ~ypes of reaction removing CIO2 from the system. The first is its unimolecular decompos~.tion which has been shown 7 to give rise to CIO radicals and Cl atoms C l O 2 + M - , ClO + O + M
[6]
or
CIO2 + M --, CI + 02 + M
[7]
The alternative reaction for the removal of C I O 2 is a bimolecular exchange reaction with a
reactive species such as oxyger, atoms ClO,~ + O -~ CIO + 02
[8]
which is known to be a fast reaction is The effect of an inert gas, which increased the yield of CIO2 at high temperatures, throws light on the reaction responsible for the removal of CIO2. As the pressure in the reaction chamber was so low it can be assumed that the rate for the unimolecular decomposition of perchloric acid was below its high pressure limit. Any increase in the total pressure in the system caused by addition of an inert gas would therefore increase the rate of decomposition and consequently give rise to a greatez yield of CIO2. The only reaction which could account for the increase in yield of CIO2 in the presence of an inert gas would be the bimolecular reaction 6. As a reaction of this type would produce CIO radicals and as these arise in the decomposition of perchloric acid (reaction 3) CIO radicals are involved in the gas phase decomposition of AP. Nitric oxgde as a reaction intermediate
Analysis of the products of the gas phase re-
641
action showed that most of the ammonia was converted to nitrogen peroxide. It seems unlikely that the ammonia is oxidized to nitrogen peroxide without nitric oxide first being formed. If this is so then nitric oxide must have reacted with compounds containing oxygen to form nitrogen peroxide. To test this hypothesis nitric oxide was added to the reaction mixture. The yield of CIO2 decreased to zero when sufficient nitric oxide was present, possibly because it reacted directly lsa with nitric oxide to give CIO radicals and NO2. C I O 2 + NO--, C10 + N O 2
[9]
A sin-ilar reaction has been observed between CIO2 and oxygen atoms 15b. Reaction 9 is exothermic (16 _+ 4 kcal/mole) and is energetically favoured. The reaction also accounts for the increase in NO2 when nitric oxide was added. As a consequence of reaction 9 the number of CIO radicals in the system would increa3e. However, these are l~own to undergo a fast reaction with nitric oxide 15 CIO + NO--, C1 + N O 2
[10]
and therefore with nitric oxide present would not be expected to survive long. This reaction would lead to an increase in the number of chlorine atoms and again an increase in the yield o f N O 2. The increase in the yield of Cl 2 when nitric oxide was added follows naturally from the increase in chlorine atom concentration by the recombination reaction C1 + CI + M
--,
CI 2 + M
[11]
Alternatively chlorine atoms may react with nitric oxide to give nitrosyl c h l o r i d e 16 CI + NO + M ~ N O C I + M
[12]
which then reacts with other chlorine atoms to give molecular chlorine and nitric oxide NOCI + CI ---,NO + C12
[13]
The increase in chlorine atom concentration also accounts for the increase in yield of HC1 when nitric oxide was added NH 3 + CI --, NH2 + HCI
[ 14]
This also explains the increase in the amount of ammonia which reacted when nitric oxide was
Vol. 13 L R. MAJ~ AND M. SMITH 642 took place, then their yield would be expected added and gives rise to an increase in the conto drop when excess nitric oxide was added. centration of NI-I2 radicals. The effect of added nitric oxide on the It has been shown that the yield of N 2 0 yields of the products of the gas phase reactions increased when nitric oxide was added to the of AP can be explained in terms of the reactions reaction mixture. suggested above. From these it appears that A possible reason for this is the reaction NO2 can be ~ormed by reaction of the added ] ~ ' I 2 "1" N O 2 --* N 2 0 4- H 2 0 [15] nitric oxide with oxygen-containing species produced in the decomposition of perchloric It ha~ Seen mentioned that there would be an acid. It is therefore suggested that such reactions increase in the concentration of NH,` radicals do take place in the gas phase decomposition of and an increase in the yield of N O 2 with the AP and that nitric oxide, produced by oxidation addition of nitric oxide. An increase in the yield of ammonia, is an active intermediate in the of N20 would therefore follow if the reaction 15 reactions. were t~dng place. It would also explain the concurrent increase in the yield of H20. The formation of the products Further evidence suggesting that 15 is responsible There are two possible modes of formation of for the formation of N20 is that it would reduce HCIO. First the CIO radical could abstract a the concentration of NO 2 produced by reactions hydrogen atom from a hydrogen-containing 9 and 10 and would give rise to a decrease in the compound, in peLrticular NH3, to give rise to yields of N O 2 at some state in the addition of HCIO. nitric oxide as was observed. Finally, the decrease in the yield of oxygen NH 3 + CIO ~ NH 2 + HCIO [18] in the presence of nitric oxide can be expected, Alternatively, reaction between hydroxyl radias nitric oxide reacts with oxygen1 ~. cals and chlorine atoms could account for the 2NO + 02 "-* 2NO2 [16] formation of HCIO. The fall in the yields of H,`O, HCI, Cl 2 and N20 and the decrease in the amount ofammonia reacting when excess nitric oxide was added can be explained by considering the reaction between chlorine atoms and nitric oxide. This reaction and the decomposition of its product, nitrosyl chloride, have been discussed in detail by Ashmore and Spencer Is. They showed that at 300°C the equilibrium NO + CI + M ~- NOCI + M
[17]
resulting from these reactions occurred. If this equilibrium occurred in the present system then it is possible that when excess nitric oxide was added dissociation of nitrosyl chloride could be depressed and the chlorine atom concentration decreased. The yields of C12 and HCI would therefore have decreased with excess nitric oxide as is observed. The amount of ammonia reacted would also ha,Je decreased and so would the concentration of NH2 radicals. If the above process responsible for the formation of N20 and H,`O from N I l , radicals
CI + OH + M --, HCIO + M
[19]
If HCIO were formed by reaction 19 then, as the concentration of chlorine atoms is increased by addition of nitric oxide, an increase in the yield of HCIO would also be expected. However, if reaction 18 were responsible for the formation of HCIO, then as the concentration of CIO radicals is decreased by addition of nitric Oxide, the yield of HCIO would be expected to decrease. The yield is in fact found to decrease under such conditions, so that HCIO is formed by a hydrogen abstraction reaction. There is some evidence that hydrogen abstraction from hydrocarbons by CIO occurs 19'2°. Hydrogen chloride has not been reported as a product of the homogeneous thermal decomposition of perchloric acid but Heath and Majer observed its formation by pyrolysis on a hot wire. In the present system, chlorine atoms may abstract 14 a hydrogen atom from NH3 to give NH 2 radicals and HCI. A fast reaction similar to this occurs 21 at 300°C between methane and chlorine atoms to give CH3 radicals and HC1.
December 1969
GAS PHASE THERMALDECOMPOSITION OF AMMONIUMPERCHLORATE
It has been suggested 14 that OH radicals are formed in the thermal decomposition of perchloric acid and that they react with perchloric acid to give water~. This reaction could take place in the gas phase decomposition of AP. However, OH radicals could also react with other compounds containing ~yd~ogen, ~ particular with ammonia whi:h is present in comparatively large quantities, to give water and NH2 radicals. A similar reaction between CH4 and OH radicals is known "-~-. Oxygen atoms could also abstract a hydrogen atom from ammonia and give rise to more OH radicals "3 NH 3 + O ~ NH 2 + OH
[20]
It appears that NI-I2 radicals are formed from ammonia by hydrogen abstraction with chlorine atoms, oxygen atoms, CIO and OH radicals. It has been shown by Serewicz and Noyes 24 that NH 2 radicals react with nitric oxide to give nitrogen and water. N H 2 -k
NO--,
N 2 -I- H 2 0
[21]
As nitric oxide is probably present as an intermediate, it is possible that it reacts with NH 2 radicals as in reaction 21 ~o produce nitrogen. Addition of nitric oxide to the reaction mixture caused the yield of N20 to increase. Reaction 20 between NO2 and NH2 radicals was suggested as being responsible for the formation of the extra N20. It is possible that this reaction is also responsible for the production of N20 here. Oxidation of ammonia The one remaining problem to be explained is the production of nitric oxide, which has been postulated as an intermediate. It has already been shown by a number of workers 23'25-27 that nitric oxide is produced in the oxidation of ammonia by oxygen atoms and molecular oxygen, but the mechanism is not known. It is probable that in this system ammonia first loses a hydrogen atom by abstraction with chlorine atoms, oxygen atoms, OH radicals and CIO radicals to give NH2 radicals. Further oxidation may take place involving some or all of the radicals, O, OH, CIO and CI to give NO. Molecular oxygen does not seem to be involved as no increase in the yields of oxides of nitrogen
643
was observed when oxygen was added to the system. Conclusions The combLnation of sublimation apparatus, secondary furnace, pin-hole leak and mass spectrometer has been shown to be a satisfactory analytical system for the study of the gas phase decomposition of AP at temperatures between 300° and 520°C. The products of this decomposition have been shown to include the hitherto undetected compound hypochlorous acid (HOCI). The variation of the yields of all products with variation in temperature, pressure and the presence of foreign gases has been studied and the results interpreted in terms of a series of consecutive reactions. An overall reaction scheme which accounts for the presence and variation in yield of all detectable products has been proposed. It involves the sublimation of AP by dissociation into ammonia and perchloric acid followed by dissociation of perchloric acid into hydroxyl radicals and chlorine oxides. The ammonia is attacked by hydroxyl radicals, chlorine oxides, chlorine and oxygen atoms to yield amino radicals. A further oxidation to nitric oxide is followed by secondary reactions to nitrogen and oxides of nitrogen. The end products of the decomposition and subsequent reactions of the chlorine oxides are hydrochloric acid, water and oxygen. These conclusions have largely been predicted by Jacobs 2s.
The authors are indebted to Dr G. A. Heath for helpful advice, to Dr A. R. Hall and Dr G. S. Pearson for discussions of the kinetic data and to the Ministry of Aviation for a maintenance grant for M. Smith. (Received January 1969; revised May 1969)
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j. R. MAJER AND M. SMITH
4 HEATH, G. A. and MAJOr, J. R. Trans. Faraday Soc. 60, 1783 (1965) s PELLETT,G. L. and SAUNDERS,A. R. NASA Third Interagency. Chemical Rocket Propulsion Group Combustion Conference, John F. Kennedy Space Centre, Florida ¢October 1966) o MA¥COCK, J. N., PAt V~XNmmX, V. R. and JACOnS, P. W. M. J. chem. Phys. 46, 2857 (1967) 7 pal Vmth'i~mx, V. R. and MAYCOClL,J. N. J. chem. Phys. 46, 3618 (1967) s American Petroleum Institute, Research Project 44, 'Catalogue of mass spectral data' 9 INAm, S. H., Rossini, W. A. and WLSF.,H. J. phys. Chem. 67, 1077 (1963) to SmLeg, H. H., NL~t, F. T., Daxc, o, R. S. and YANEY,D. J. Amer. chem. Soc. 76, 3909 (1954) it DOl>~, M. C.R. Acad. $ci.. Paris, 200, 63 (1935) 12 DOD]~, M. Bull. Soc. chim. Ft. 5, (5), 170 (1938) 13 F I s m ~ I. P. Trans. Faraday $oc. 63, 684 (1967) 14 LEVY, J. B. J. phys. Chem. 66, 1092 (1962) ts o COXON, J. A. Trans. Faraday $oc. 64, 2118 (1968) is b CLYNE,M. A. A. and CoxoN, J~ A. Trans. Faraday Soc. 62, 1175 (1966)
Vol. 13
to CL~atlL, T. C., CLYNF., M. A. A. and STEDMAN, D. H. Trans. Faraday Soc. 62, 3354 (l~,C~) t7 B O D E N ~ , M. Heir. chim..~cu 18, 743 (1935) t8 ASHMORE,P. G. and SPLICER, M. S. Trans. Faraday Soc. 55, 1868 (1959) 19 PmLLn,S, L. and SHAW,H. Proc. chem. $oc. 294 (1962) 2o HEATH, G. A. and F~.ARSON, G. S. F,lecenth Symposium (International) on Combustion, p 967. The Combustion Institute: Pittsburgh (1967) 21 PRrrCHARD,H. O., PoKe, J. B. and TROTMAN-DICKENSON, A. F. ,7. Amer. chem. $oc. 77, 2629 (1955) 22 GRelNi~, N. R. J. chem. Phys. 46, 295 (1967) 23 WONG, E. L. and Porrea, A. E. .I. chem. Phys. 43, 3371 (1965) z4 S ~ c z , A. and NoY~s JR, W. A. J. phys. Chem. 63, 843 (1959) zs Dau3~oND, L. J. and HlSO~K, S. W. Austral. J. Chem. 20, 825 (1967) ~6 T A L h ' ¢ ~ , T. and Mn~ANA, H. J. chem. Phys. 42, 3737 (1965) 27 Hus~N, D. and No~JsH, R. G. W. Proc. Roy. Soc. A, 273, 145 (1963) za JACOaS, P. W. M. and RUSSELL-JONI~, A. A.LA.A. Jnl, 5, 829 (1967)
OF AMMONIUM
J. R. MAJER AND M. SMITH Chemistry Department, University of Birmingham, Edgbaston, Birmingham 15 The decomposition of ammonium perchlorate in the gas phase at elevated temperatures has been studied by allowing the vapour from a sample of subliming ammonium perchlorate to pass through a secondary furnace maintained at a higher temperature than that of the subliming ammonium perchlorate. The reaction occurring in the secondary furnace was monitored using a pin-hole leak connected to a mass spectrometer. The effects of temperature and foreign gases have been studied and the major products identified. A reaction scheme is suggested to account for the variation in product yield with variation in experimental parameters.
Introduction BmCUMSHAW and Newman 1'2 made an intensive investigation of the thermal decomposition of ammonium perchlorate (AP) and showed that it decomposes in the solid phase (the low temperature decomposition)and sublimes at temperatures below 300°C. They also found (hat above 350°C a different decomposition reaction takes place. Galwey and Jacobs 3 suggested that this high temperature reaction consisted of dissociation of AP into ammonia and perchloric acid, followed by their evaporation into the gas phase and subsequent decomposition of perchloric acid and oxidation of ammonia. The possibility of making a detailed study of the low temperature thermal decomposition of AP using a mass spectrometer was ~;rst demonstrated by Heath and Majer 4. A solid sample of AP was heated in the ion source of an A.E.I.M.S.2 mass spectrometer. The disadvantage of this method is that the rate of decomposition and sublimation cannot be kept constant and the small size of the furnace and sample did not permit prolonged experiments. This investigation has been extended by other workers s-7; however, their work has been directed towards elucidating the low temperature decomposition. It should be possible to investigate the high temperature reaction by subliming a ~mple of AP and examining the ensuing gas phase reactions by means of a mass spectrometer
coupled to the reaction system. The present paper describes such a system and the results obtained. Experimental
Apparatus Any investigation of the gas phase thermal reactions of AP must involve a procedure which permits differentiation between gas phase and solid phase reactions. The most convenient way of doing this is to pump gaseous species away from the decomposing AP and examine their interactions in a region isolated from tbe decomposing solid. The situation is complicated by the fact that AP will give rise to a solid sublimate at a cool surface in a reaction vessel. Consequently the entire reaction system must be kept at the temperature at which AP is d~omposing to prevent any sublimate being formed. A flow system was used in which gaseous products from the decomposing AP were continuously pumped through a reaction chamber of variable temperature, situated some distance from the solid AP. The gas phase reaction was continuously monitored by a mass spectrometer. A diagram of the system is shown in Figure 1. A pin-hole leak 00045 in. in diameter at the end of 0.1 in. i.d. Pyrex tube led to the ;,on source of an A.E.I.M.S.2 mass spectrometer through 635
ol. 13
J. R. MAJER AIqD M. SMITH
636
Thermocouple weft eta
~i ,//(~
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a ~ in. hole drilled in the ion repeller plate of the ion source. The section of the apparatus, in which AP was housed, was heated to a maximum temperature of 320°C by a furnace, mounted on telescopic runners, which enabled it to be slid over the sample. Besides heating the AP, the furnace heated the re~t of the system up to the gas phase reaction chamber and prevented any sublimate being deposited. The gas phase reaction chamber was heated to a maximum temperature of 520°C by a 32 ohm, 28 s.w.g. Nichrome resistance wire wound directly on to the chamber and extending for 3 in. at each end of it. ,The rest of the reaction system between the reaction chamber and the cold trap was heated by Electrothermal heating tapes HT 301. The temperature in this region was measured at a point about two inches above the Housekeeper seal by an iron--constantan thermocouple.
Method About 10 g of AP, in a small Pyrex boat, was first dried and initially decomposed by heating at 300 ° to 320°C for one hour under vacuum in a separate pre-heat apparatus. The sample was then transferred to the reaction system.
The reaction system and mass spectrometer were pumped out and the reaction chamber furnace was set to give a temperature of about 520°C, the region between this and the cold trap baked at about 320°.C and the entire apparatus left pumping in this condition overnight. After setting the long tubular furnace at a temperature of up to 310°C, it was slid over the sample of AP. The temperature of the reaction chamber was next reduced until it was slightly higher than the AP decomposition temperature to ensure that no sublimate of AP was deposited. Analysis of the reactants and products of the gas phase reaction was performed in two ways. First the total mass spectrum of all products and reactants was recorded by scanning the m/e range~ 12-75 and 75-140 separately at a resolution of about 200 with an ionizing electron voltage of 40 eV for the role range 12-18 and 70 eV for the role range 18-140. Alternatively, the ion current at a specific role value could be recorded continuously. When using the total spectrum method of analysis, it was necessary to record spectra at the normal temperature of the reaction chamber before and after obtaining a spectrum at a higher reaction chamber temperature. In the case when the reaction was monitored
December 1969
637
GAS PHASE THERMALDECOMPOSITION OF AMMONIUMPERCHLORATE
by following the change in an individual peak, the peak corresponding to the compound being examined was continuously recorded with the reaction chamber at its normal temperature. The temperature was then raised and the peak height still continuously recorded. After the maximum reaction chamber temperature had been reached, it was returned to its normal value still with the peak height being recorded. The mass spectrometer was calibrated for N20, 02 and NI-Ia. The calibrations were used to calculate the calibration equations for N 2, HCI and NO 2 using thei.~ known a mass spectrometric relative sensitivities. The calibration equations calculated for NH2CI, HCIO, CIO 2 and Ci 2 were only very approximate as their relative sensitivies were estimated.
A peak at m/e 51 was shown to be due to the ion NI-I2CI+ and not CIO + from an examination of the spectrum produced using deuterated AP. Ammonia was first detected at 270°C but was only present in measurable quantities at 300°C. No perchloric acid was detected at any temperature. It was most certainly present because a sublimate of AP was found in the region between the sampling leak and the cold trap of the reaction system. The products of the low temperature decomposition at this optimum temperature of 300°C are given in Table 1 together with their corresponding partial pressures in the region just before the sampling leak. Monochloramine and hypochlorous acid have not previously been reported as products of the decomposition of AP.
Materials AP was supplied by B.D.H. Deuterated AP was prepared by recrystallizing AP twice, from deuterium oxide (KochLight Laboratories, England). 0~ Nitrogen, white spot grade, nitrous oxide, anaesthetic grade, oxygen and hydrogen were obtained from the British Oxygen Co. Ltd, England. Nitric oxide and carbon tetrafluoride were supplied by the Matheson Co. Inc. U.S.A. Ammonia and carbon dioxide were obtained from I.C.I. Ltd, England. Results
The products of the low temperature decomposition and sublimation A sample of AP was heated to temperatures in the range 240° to 300°C with the rest of the reaction system maintained at the same temperature to avoid deposition of a sublimate and a mass spectrum of all the gaseous compounds taken at each temperature. Most of the compolmds were unambiguously identified by their parent ions or isotopic parent ions in the mass spectrum. Ammonia could not be identified by its parent ion NI-I at m/e 17 as the ion OH + from water completely obscured it. Fortunately its fragment ion NH + at m/e 15 was free from major interference by any other ion and was used to show the presence of ammonia 8.
TABLE 1. Products of the low temperature decomposition and sublimation of AP
Product
NH3 H20 Oe HCi N2 N20
Partial pressure, x 1O- 3 t o r t
17"2 83"1 16"1 3"3 0"6 9'7
Product
Partial pressure, × 1O- a torr
NOe NHeC1 CIO2 HCIO CI 2
3'5 1'9 0'1 01 2'4
The products of the gas phase reaction The nature of the products of the gas phase reaction of the compounds from decomposing AP was determined by comparing in another run the mass spectrum of the AP heated at 300° with that obtained when the reaction chamber was raised to 468°C. The difference was a measure of the extent and nature of the gas phase reactions. The partial pressures "of the products of the gas phase reaction as they emerged from the reaction chamber are given in Table 2. Ammonia, which decreased in amount when the reaction chamber temperature was increased, was thus involved in the gas phase reaction. It can be seen from the product analysis that it was converted to N20, NO2 and a small amount of Ne. Nitric oxide was not detected.
J. g. MAJi~
638
TABLE 2. Products of the gas phase decomposition of A P
Product
Partial pressure x 10- a tcr~
H20 NO 2 O2
18"1 8-7
N20
1"4 06
Product
Partial pressure x 10- 3 tort
N2 HCIO C102 0 2
0-3 0-8 0-7 10
1.6
HCI
Initial NH3 pressure 29-3 x 10-3 torr. NH~ reacted 9.6 x 10- 3 tort.
Reaction chamber temperature 468°C.
Effect of temperature The peak height could either be continuously recorded, whilst the rc=~=tion chamber temperature was continuously changing, or be recorded whilst the temperature was increased in steps. Both methods were used for CI[, and the two curves relating changes in peak height with temperature were virtually inL
o10
distinguishable. The more rapid method in which the temperature was continuously changing was therefore used. A number of runs were carried out with a sample of AP heated at 300 ° to 305°C, by following each significant peak as the temperature of the reaction chamber was increased from 330 ° to 530°C. The changes in partial pressure with temperature are given in Figures 2(a) to (d). The change in partial pressure represents the amount of material produced in the gas phase reaction in all cases except ammonia. As separate runs were necessary to follow the change in each product the yields are not inter-related. Figure 2(a) shows that the amount of ammonia reacting increased with temperature. The extent of this can be seen in that approximately 50 per cent of the ammonia entering the reaction chamber had reacted at 500°C. The yields of nearly all the products of the gas phase reaction increased with temperature within the range studied. The exceptions
Initial NH 3 pressure = 14-7x 10-3 torr
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(d) 500
December 1969
GAS FELASE~HERMAL DECOMPOSITION OF AMMONIUM PERCHI~RATE
to this were chlorine dioxide and hypochlorous acid which initially increased but began to decrease at 475 ° and 510°C respectively.
4
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Z
The effect of th~ added gas on the gas phase reaction was L'r~estigated by continuously monitoting invididual peak~ corresponding to reactants and products of the reaction. As the maximum pressure of added gas in the reaction chamber in all experiments was estimated to be less than iO torr and as the decomposing AP was farther from the pumping system than the point of injection of the gases, it was unlikely that the solid phase decomposition would be greatly affected by the added gas. Any change in the partial pressure of the compounds leaving the reaction chamber was therefore considered to be due to the added gas influencing gas phase reactions. LtO
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The effect of an inert gas on the yield of CIOe was examined by observing the change in yield with temperature in the manner described previously, but in the presence of nitrogen. The results are illustrated in Figure 3 in which partial pressure of CIOz is plotted against temperature with and without nitrogen present. It can be seen that with an inert gas present the maximum yield of CIO2 was increased and occurred at a higher temperature.
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Vol. 13 Oxygen was added to the reaction mixture in the manner previously described. The general effect was to decrease the yield of all products in a manner similar to that of an inert gas. In particular it is to be noted that the yields of oxides of nitrogen decreased.
J. R. MAJER AND M. SMITH
640
The effect of increasing amounts of nitric oxide on the reactants and products of the gas phase reaction is shown in Figures 4(a) to (d). The amount of ammonia which had reacted reached a maximum at roughly the same pressure of nitric Oxide as the maximum in the yields of H20, NO2, N 2 0 and HCL At the same pressure of nitric oxide virtually no CIO2 and HCIO were formed. The yield of chlorine reached a maximum at a pressure of nitric oxide slightly lower than that at which maxima for other products were obtained. It can be seen that the yield of oxygen decreased but then increased at a higher pressure of nitric oxide. The increase could be accounted for by the contribution of NtSO + to the peak at m/e 32. TEe effect of nitric oxide on the yield of nitrogen could not be determined, as an impurity of four per cent nitrogen obscured any small change in the yield. The yields of H20, C12, HCI, N 2 0 and NO2 were found to decrease when excess nitric oxide was added. ,_ 10
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F~GUgE 5. Variation in molecular chlorine yields with added ammonia pressure
When ammonia was added to the reaction mixture, it was found that nearly all products decreased, probably due to suppression of the dissociation of AP as shown by Inami et al.9. The exception to this was NH2CI at role 51 in the mass spectrum of the products which showed a rapid increase. At the same time a rapid decrease in the yield of chlortr)e was found as can be seen from Figure 5. The reaction NH3 + Cl2 --, NH2CI + HCI is known to occur at room temperature t°.
Discussion The gaseous products of the low temperature decomposition of AP under the conditions used in this investigation were NH3, HCIO4, H20, N2, 02, HCI, N20, NC~, NH2Cl, Cl 2 and small quantities of HCIO and CIO 2. The product analysis is not entirely in agreement with that of other workers t. Chlorine dioxide has been found which probably arises as a result of thermal decomposition of perchloric acid produced in the dissociation of AP as suggested by Dod~: 1.12. When the temperature of the gases from decomposing AP was increased, the concentration of NI-I3 decreased and that of the other products of the low temperature reaction, except perchloric acid which was not observed, increased. As can be seen from Table 2 the products of this complex gas phase reaction are similar to those of the low temperature decomposition (Table 1). However, NO 2 is the most abundant oxide of nitrogen amongst the products of the gas phase reaction rather than N20 which is the most abundant oxide of nitrogen in the low temperature decomposition. The relative abundance of chlorine oxides is also greater in the products of the gas phase reaction than in the low temperature decomposition products. We will now discuss the significance of these observations and the effect that various parameters have on the yields of both these groups of oxides in the gas phase reaction.
The reactions of oxygen-chlorine compounds Chlorine dioxide has been reported as a product of the thermal decomposition of perchlor/c acid 13. As has been found in the present investigation it appears that perchloric acid from dissociating AP has undergone decomposition probably by the reactions: HO--CIO 3 + M ~ HO + CIO3 + M
[1]
December 1969
GAS PHASE THERMAL DECOMPOSITION OF AMMONIUM PERCHLORATE
CIO3 -~ C102 + O
[2]
It has been suggested that CIO3 decomposes to give CIO CIO3 - , CIO + O2
[3]
Chlorine dioxide may also be formed from CIO4 radicals produced by attack of OH radicals on perchloric acid ~4. HO + HOClO 3 ----,HzO + C104 CI04 ~ CIO~ + 02
[4]
[5]
The decrease in the yield of CIO2 at high temperatures in the present investigation could be due to two ~ypes of reaction removing CIO2 from the system. The first is its unimolecular decompos~.tion which has been shown 7 to give rise to CIO radicals and Cl atoms C l O 2 + M - , ClO + O + M
[6]
or
CIO2 + M --, CI + 02 + M
[7]
The alternative reaction for the removal of C I O 2 is a bimolecular exchange reaction with a
reactive species such as oxyger, atoms ClO,~ + O -~ CIO + 02
[8]
which is known to be a fast reaction is The effect of an inert gas, which increased the yield of CIO2 at high temperatures, throws light on the reaction responsible for the removal of CIO2. As the pressure in the reaction chamber was so low it can be assumed that the rate for the unimolecular decomposition of perchloric acid was below its high pressure limit. Any increase in the total pressure in the system caused by addition of an inert gas would therefore increase the rate of decomposition and consequently give rise to a greatez yield of CIO2. The only reaction which could account for the increase in yield of CIO2 in the presence of an inert gas would be the bimolecular reaction 6. As a reaction of this type would produce CIO radicals and as these arise in the decomposition of perchloric acid (reaction 3) CIO radicals are involved in the gas phase decomposition of AP. Nitric oxgde as a reaction intermediate
Analysis of the products of the gas phase re-
641
action showed that most of the ammonia was converted to nitrogen peroxide. It seems unlikely that the ammonia is oxidized to nitrogen peroxide without nitric oxide first being formed. If this is so then nitric oxide must have reacted with compounds containing oxygen to form nitrogen peroxide. To test this hypothesis nitric oxide was added to the reaction mixture. The yield of CIO2 decreased to zero when sufficient nitric oxide was present, possibly because it reacted directly lsa with nitric oxide to give CIO radicals and NO2. C I O 2 + NO--, C10 + N O 2
[9]
A sin-ilar reaction has been observed between CIO2 and oxygen atoms 15b. Reaction 9 is exothermic (16 _+ 4 kcal/mole) and is energetically favoured. The reaction also accounts for the increase in NO2 when nitric oxide was added. As a consequence of reaction 9 the number of CIO radicals in the system would increa3e. However, these are l~own to undergo a fast reaction with nitric oxide 15 CIO + NO--, C1 + N O 2
[10]
and therefore with nitric oxide present would not be expected to survive long. This reaction would lead to an increase in the number of chlorine atoms and again an increase in the yield o f N O 2. The increase in the yield of Cl 2 when nitric oxide was added follows naturally from the increase in chlorine atom concentration by the recombination reaction C1 + CI + M
--,
CI 2 + M
[11]
Alternatively chlorine atoms may react with nitric oxide to give nitrosyl c h l o r i d e 16 CI + NO + M ~ N O C I + M
[12]
which then reacts with other chlorine atoms to give molecular chlorine and nitric oxide NOCI + CI ---,NO + C12
[13]
The increase in chlorine atom concentration also accounts for the increase in yield of HC1 when nitric oxide was added NH 3 + CI --, NH2 + HCI
[ 14]
This also explains the increase in the amount of ammonia which reacted when nitric oxide was
Vol. 13 L R. MAJ~ AND M. SMITH 642 took place, then their yield would be expected added and gives rise to an increase in the conto drop when excess nitric oxide was added. centration of NI-I2 radicals. The effect of added nitric oxide on the It has been shown that the yield of N 2 0 yields of the products of the gas phase reactions increased when nitric oxide was added to the of AP can be explained in terms of the reactions reaction mixture. suggested above. From these it appears that A possible reason for this is the reaction NO2 can be ~ormed by reaction of the added ] ~ ' I 2 "1" N O 2 --* N 2 0 4- H 2 0 [15] nitric oxide with oxygen-containing species produced in the decomposition of perchloric It ha~ Seen mentioned that there would be an acid. It is therefore suggested that such reactions increase in the concentration of NH,` radicals do take place in the gas phase decomposition of and an increase in the yield of N O 2 with the AP and that nitric oxide, produced by oxidation addition of nitric oxide. An increase in the yield of ammonia, is an active intermediate in the of N20 would therefore follow if the reaction 15 reactions. were t~dng place. It would also explain the concurrent increase in the yield of H20. The formation of the products Further evidence suggesting that 15 is responsible There are two possible modes of formation of for the formation of N20 is that it would reduce HCIO. First the CIO radical could abstract a the concentration of NO 2 produced by reactions hydrogen atom from a hydrogen-containing 9 and 10 and would give rise to a decrease in the compound, in peLrticular NH3, to give rise to yields of N O 2 at some state in the addition of HCIO. nitric oxide as was observed. Finally, the decrease in the yield of oxygen NH 3 + CIO ~ NH 2 + HCIO [18] in the presence of nitric oxide can be expected, Alternatively, reaction between hydroxyl radias nitric oxide reacts with oxygen1 ~. cals and chlorine atoms could account for the 2NO + 02 "-* 2NO2 [16] formation of HCIO. The fall in the yields of H,`O, HCI, Cl 2 and N20 and the decrease in the amount ofammonia reacting when excess nitric oxide was added can be explained by considering the reaction between chlorine atoms and nitric oxide. This reaction and the decomposition of its product, nitrosyl chloride, have been discussed in detail by Ashmore and Spencer Is. They showed that at 300°C the equilibrium NO + CI + M ~- NOCI + M
[17]
resulting from these reactions occurred. If this equilibrium occurred in the present system then it is possible that when excess nitric oxide was added dissociation of nitrosyl chloride could be depressed and the chlorine atom concentration decreased. The yields of C12 and HCI would therefore have decreased with excess nitric oxide as is observed. The amount of ammonia reacted would also ha,Je decreased and so would the concentration of NH2 radicals. If the above process responsible for the formation of N20 and H,`O from N I l , radicals
CI + OH + M --, HCIO + M
[19]
If HCIO were formed by reaction 19 then, as the concentration of chlorine atoms is increased by addition of nitric oxide, an increase in the yield of HCIO would also be expected. However, if reaction 18 were responsible for the formation of HCIO, then as the concentration of CIO radicals is decreased by addition of nitric Oxide, the yield of HCIO would be expected to decrease. The yield is in fact found to decrease under such conditions, so that HCIO is formed by a hydrogen abstraction reaction. There is some evidence that hydrogen abstraction from hydrocarbons by CIO occurs 19'2°. Hydrogen chloride has not been reported as a product of the homogeneous thermal decomposition of perchloric acid but Heath and Majer observed its formation by pyrolysis on a hot wire. In the present system, chlorine atoms may abstract 14 a hydrogen atom from NH3 to give NH 2 radicals and HCI. A fast reaction similar to this occurs 21 at 300°C between methane and chlorine atoms to give CH3 radicals and HC1.
December 1969
GAS PHASE THERMALDECOMPOSITION OF AMMONIUMPERCHLORATE
It has been suggested 14 that OH radicals are formed in the thermal decomposition of perchloric acid and that they react with perchloric acid to give water~. This reaction could take place in the gas phase decomposition of AP. However, OH radicals could also react with other compounds containing ~yd~ogen, ~ particular with ammonia whi:h is present in comparatively large quantities, to give water and NH2 radicals. A similar reaction between CH4 and OH radicals is known "-~-. Oxygen atoms could also abstract a hydrogen atom from ammonia and give rise to more OH radicals "3 NH 3 + O ~ NH 2 + OH
[20]
It appears that NI-I2 radicals are formed from ammonia by hydrogen abstraction with chlorine atoms, oxygen atoms, CIO and OH radicals. It has been shown by Serewicz and Noyes 24 that NH 2 radicals react with nitric oxide to give nitrogen and water. N H 2 -k
NO--,
N 2 -I- H 2 0
[21]
As nitric oxide is probably present as an intermediate, it is possible that it reacts with NH 2 radicals as in reaction 21 ~o produce nitrogen. Addition of nitric oxide to the reaction mixture caused the yield of N20 to increase. Reaction 20 between NO2 and NH2 radicals was suggested as being responsible for the formation of the extra N20. It is possible that this reaction is also responsible for the production of N20 here. Oxidation of ammonia The one remaining problem to be explained is the production of nitric oxide, which has been postulated as an intermediate. It has already been shown by a number of workers 23'25-27 that nitric oxide is produced in the oxidation of ammonia by oxygen atoms and molecular oxygen, but the mechanism is not known. It is probable that in this system ammonia first loses a hydrogen atom by abstraction with chlorine atoms, oxygen atoms, OH radicals and CIO radicals to give NH2 radicals. Further oxidation may take place involving some or all of the radicals, O, OH, CIO and CI to give NO. Molecular oxygen does not seem to be involved as no increase in the yields of oxides of nitrogen
643
was observed when oxygen was added to the system. Conclusions The combLnation of sublimation apparatus, secondary furnace, pin-hole leak and mass spectrometer has been shown to be a satisfactory analytical system for the study of the gas phase decomposition of AP at temperatures between 300° and 520°C. The products of this decomposition have been shown to include the hitherto undetected compound hypochlorous acid (HOCI). The variation of the yields of all products with variation in temperature, pressure and the presence of foreign gases has been studied and the results interpreted in terms of a series of consecutive reactions. An overall reaction scheme which accounts for the presence and variation in yield of all detectable products has been proposed. It involves the sublimation of AP by dissociation into ammonia and perchloric acid followed by dissociation of perchloric acid into hydroxyl radicals and chlorine oxides. The ammonia is attacked by hydroxyl radicals, chlorine oxides, chlorine and oxygen atoms to yield amino radicals. A further oxidation to nitric oxide is followed by secondary reactions to nitrogen and oxides of nitrogen. The end products of the decomposition and subsequent reactions of the chlorine oxides are hydrochloric acid, water and oxygen. These conclusions have largely been predicted by Jacobs 2s.
The authors are indebted to Dr G. A. Heath for helpful advice, to Dr A. R. Hall and Dr G. S. Pearson for discussions of the kinetic data and to the Ministry of Aviation for a maintenance grant for M. Smith. (Received January 1969; revised May 1969)
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