On the explosion limits of various combustible gases using nitrous oxide as oxidizer

On the Explosion Limits of Various Combustible Gases us&g Nitrous Oxide as Oxidizer M. DESTRIAU L~Lboratoire de Chimie Gdndrule, Facult.d des Sciences de Paris, Sorbonne (Received

October 1962)

The explosion limits o/ methane, ethane, propane and butane have been studied by a static method using nitrous oxide as oxidizer. A comparison is drawn with similar results obtained with olefins (ethylene) and other /uels such as hydrogen and ammonia. A n interpretation o / t h e data obtained has been attempted in terms o / t h e thermal theory.

ALTHOUGH a certain number of investigations have been carried out in the past on the explosion limits of various mixtures of hydrocarbons and nitrous oxide 1, very little precise information on these systems is available if one seeks a systematic and comparable set of data with. particular reference to identical experimental conditions. This work deals with the comparison of the spontaneous ignition of a certain number of fuels, namely the paraffinic hydrocarbons from methane to n-butane, ethylene as an example of olefins, and hydrogen and ammonia as examples of carbon-free combustibles, all of them in the presence of nitrous oxide as oxidizer. A comparison of these three cases in terms of the thermal theory has been successfully attempted.

Nitrous oxide (N20) was of commercial anaesthesia grade with a few per cent nitrogen. The various fuels used were: Methane Propane n-Butane

t

Ethane Ethylene

"/~Phillipsresearch grade, 99'9 per cent J

Phillips pure grade, 99"0 per cent

Hydrogen:

commercial electrolytic compressed gas; treated by potassium pyrogallate for removal of oxygen traces

Ammonia:

commercial pure 99"9 per cent.

compressed

gas,

Spontaneous Ignition Limits of Methane, Propane and Butane In order to obtain a useful comparison we have taken the stoichiometric mixtures of each hydrocarbon, namely 4N20+CH4; 10N20+C~H~, 13N20+C4H~0 and a series of nitrous oxide rich mixtures containing respectively 1, 2, 3, etc., times the proportion of nitrous oxide for the stoichiometric mixture 2. F i g u r e I is a pressure/temperature curve for propane-nitrous oxide. In each case the limit point is always taken as the last pressure, operating by pressure reduction, for which a flame occurs. F i g u r e 2 is derived from F i g u r e I and from similar curves with methane and butane, and shows the pressure limit variation as a function of nitrous oxide concentration for a given temperature (900°C). The limiting pressures are slightly higher for methane than for propane and for

Experimental The spontaneous ignition limits have been determined as a function of pressure for a given set of mixture concentrations and temperatures, b y visual observation of the flame phenomenon, and occasionally using a manographic recording. The reaction vessel of fused silica is 80 m m long and of 1 9 m m inside diameter. It lies inside a classical well-lagged, electrically heated furnace. The temperature, according to the mixtures studied, has been made to vary from 750 ° to 1 000°C. Conventional pumping devices and control apparatus have been used through~ut. 347

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Vol. 6

M. Destriau

propane higher than for butane. They tend to give practically identical values for over-rich nitrous oxide mixtures. t~

t,.

2

\\

\\

nitrous oxide mixtures* while F i g u r e 6, for comparison, has been drawn for ammoniaoxygen mixtures. I

E E

600

~~.--~

,oo

800

850 900 Temperature

400

9 0 °C

Figure I. Limiting pressure variations as functions of temperature [or a series of mixtures of nitrous oxide and propane. Curves I, 2 3, 4, 5 and 6 relate to mixtures containing respectively 2, 5, 7, 13, 35 and 50 times the proportion of nitrous oxide in the stoichiometric mixture

200

• :__

..........

900oC Spontaneous Ignition Limits of Ethane and Ethylene F i g u r e 3 relates to ethane and ethylene mixtures with nitrous oxide considered as above (6N20 + Coil, = stoichiometric mixture for ethylene and 7N20+ C2H ~= stoichiometric mixture for ethane) '~. The limiting pressures appear slightly higher for ethane than for ethylene. Here again they seem to merge for the higher contents of nitrous oxide. I

O~

"7"

E250 200

0

f I r 2 3 4 Concentration

1

i 5 x N20 rich

Figure 3. Limiting pressure variations as [unctions of concentration /or mixtures: X nitrous oxideethane, and Q) nitrous oxide-ethylene, at 850 ° and 900°C

At this stage of this investigation and considering the importance of nitric oxide (NO) in combustion mechanisms either as an inhibitor of radical supported reactions or as a promoting agent of sensitized reactions we found it appropriate to extend part of this work to a study of

150

~200 I

100 50 I

o

i

;

2

I

I

/

£

Concentration x N20 rich Figure 2. Limiting pressure variations as [unctions of concentration [or mixtures: X nitrous oxidemethane, ~ nitrous oxide-propane, and + nitrous oxide-butane, at 900°C

Spontaneous Ignition Limits of Hydrogen and Ammonia F i g u r e 4 concerns various mixtures of hydrogen and nitrous oxide ~. F i g u r e 5 relates to a m m o n i a -

O/

~

750

;

I

800

i

1

i

850

Temperature

°C

Figure 4. Limiting pressure variations as functions of temperature for a series of mixtures of nitrous oxide and hydrogen. Curve I - - N 2 0 + H2, 2--

2 N.~O+I-t2, and 3--3 N20+H2

the influence of nitric oxide on some of the mixtures considered 5. F i g u r e 7 relates to the addition of definite proportions of nitric oxide

December 1 9 6 2

Explosion limits of combustible gases using nitrous oxide as oxidizer

to one definite mixture (C2H6+ 14N20 ) at a given temperature (850°C). Bearing in mind that for each point the partial pressure of nitric oxide has been subtracted, the influence of the addition of nitric oxide appears obvious.

349

O~ :~ 200 E

ck 100 I

I

I

I

i

I

I

1

2

3

4

5

6

7

NO

Figure 7--Limiting pressure variations [or mzxture 14 N._,O-C.2H~; at 850°C as a function of NO added to the mixture

00 100~

moles

]

I

I

850

~ ~ . . . ~

900 Temperature

]

F i g u r e 9 relates to mixtures of ammonia and I

nitric oxide without nitrous oxide.

l

950

°C

Figure 5. Limiting pressure variations as ]unctions of temperature for a series of mixtures of nitrous oxide and ammonia. Curve 1--NH3+ 1- N._O (×),

2--NH3+½ N20 (~), 3--NH~+5/12 N20 (X), 4-NHa+½ N20 (@), 5--NH3+I N~.O (X), 6--NH2+3 N_~O (X), 7--NH3+9/2 N20 (---O--), 8--NH~+6 N20 (A), 9--NH.~+ 18 N.,O ('.~) F i g u r e 8 pertains to one ammonia-nitrous oxide mixture (NH~ + 6N.~O) with various amounts of nitric oxide. However, in this case the pressure shown is the total pressure without subtraction of the nitric oxide partial pressure. The promoting effect of nitric oxide is clearly seen.

,~=Q E .......

r" 'u

2~ R T -

in which Q is the heat evolved by a given mixture at temperature T under pressure p, v is the reaction rate, E is the activation energy, r is the reaction vessel diameter, /t is the thermal conductivity of the gaseous mixture, and ,~ is a

5o0 E E

Theoretical Considerations We have attempted to interpret our results in terms of the thermal theory. In this theory when the reaction heat cannot be totally evacuated through the vessel walls, a critical condition is established which leads to the spontaneous ignition of the mixture. This critical condition can be written as follows 6

2 400

~z, oo I E 300

300

2°°l 100 L

,, \

\ ,\'~ .\ \

.2oo t

~

L

I

950 1000 Temperature

- -

2

--

,oo

i

oC

Figure 6. Limiting pressure variations as ]unctions of temperature for a series of mixtures o] ammonia and oxygen. Curve I--NH3+0"5 O2 (×), 2--NH3+ 1"5 02 (+), 3--NH3+2'4 O2 (C)), 4--NH3+S'0 02 ( X )

0

I 800

]

!

L 900

L

Temperature

L

°C

FLgure 8. Limiting pressure variations as functions o] temperature. Curve I--NH3+6 N20, 2--NH3+

6 N20 + 4 NO, 3--NH3 + 6 N.~O+ 2 NO

350

M. D e s t r i a u

physically dimensionless quantity equal to two for a cylindrical vessel at the limiting pressure p. Having measured pressures and not rates, the latter have been deduced from the molecular collision theory for C oH4-NoO and C,oH,-N20 40O

-

E 300 -

~% \

200 10

I

I

r

1100 Temperature

I

]

1150 °C

Figure 9. Limiting pressure variations as ~unctions o[ temperature. Curve I--NH3+NO, 2--NH3+3"3 NO,

3--NH3 + 6 1NO

mixtures. In the assumption where the initiating mechanism lies in the dissociation of the oxidizer N20, the rate is fundamentally dependent on this dissociation process N~O + N,_O > N20 + NoO activated and therefore on the number of molecular collisions between N 2 0 molecules per cm a per sec.

The activation energy for the dissociation of N20 is that given by W. M. GRAVEN7, 52 kcal. The thermal conductivities of C2H 6 and N20 at high temperatures are not known. However, the conductivity A is related to the viscosity '7 (which is better known) by h:k, l

this viscosity depending on the temperature through the following relation ,7= k'T"; l o g ~ = l o g k ' + n l o g T This gives ~ by linear extrapolation. At 750°C, we have found for ~ values ranging from 0"8 (for the mixture C2H,+35N20 ) to 3.6 (for the mixture C2H ~+ N~O) compared with the theoretical value 2. At 950°C, these values range from 0.4 to 1.9

Vol. 6

for the mixtures C~H4-N20 (according to the composition). For NH3-N20 mixtures we have deduced rates from the experiments of A. VOLDERS and A. VAN TIGGELENa without making any assumptions about the mechanism of reaction. This has led us to values ranging from 1 to 7 for the critical parameter ~ for various concentrations and temperatures ranging from 900 ° to 950°C. Details of these calculations can be found under refs. 3 and 4. Conclusions In conclusion one may say that with hydrocarbons, the initiating mechanism may well originate in the dissociation reaction of nitrous oxide, as this is shown not only by the theoretical considerations above, but also by the experimental evidence given in Figure 2 in which the limiting pressures for all the paraffinic hydrocarbons are shown and in which it can be seen that methane lies close to the other two hydrocarbons, This would not have been so if the thermal ignition of methane had had to be initiated by the cracking of the methane molecule. This, however, does not preclude any other initiation mechanism which may lead to similar experimental evidence. References I SMITH,E. J. J. chem. Soc. 1953, 1271 KENWRIGHT, 1:{., ROBINSON, P. L. a n d TRENWITH, A. B. J. chem. Soc. 19S$, 660 FENIMOR~, C. P. a n d KELSO, J. R. J. Amer. chem. Soc. 1949, 71, 3706 ~OZLOVSKII, A. I. J. chem. Phys. Moscow, 1956, 30, 912 a n d 1444 2 DESTRIAU, M. J. Chim. phys. 1960, 57, 69 3 DESTRIAU, M. a n d DELPECH, F, J. Chim. phys. 1960, 57, 1006 •t DESTRIAU, M. a n d CLADE, D. J. Chim. phys. 1959, ~, 936 5 DESTRIAU, M. a n d LAFFITTE, P. C. R. Acad. Sci., Paris, 1960, 250, 3022 DESTRIAU, M. a n d LAFFITTE, P. C. R. Acad. Sci., Paris, 1961, 252, 4003 6 FRANK-KAMENETSKY, D. A. Diffusion and Heat Exchange in Chemical Kinetics, p 205. P r i n c e t o n U n i v e r s i t y P r e s s : N e w Jersey, 1955 7 GRAVEN, W . M. J. Amer. chem. Soc. 1959, 817 6190 s VOLDERS, A. a n d VAN TIGGELEN, A. Bull. Soc. chim. Belg. 1955, 64, 736