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TfiE LOW PRESSURE EXPLOSIVE LIMITS IN THE HYDROGEN-OXYGEN COMBINATION FRANK BRESCIA City College of New York, New York City
AN
ARTICLE on the subject of the volume synthesis of water1 has prompted this brief r6sum6 of the combination reaction between hydrogen and oxygen in the gas phase which may be of interest t o the readers of the
JOURNAL. The combination of hydrogen and oxygen between 500" and 600°C. is a chain reaction, the speed of which increases rapidly with pressure. At total pressures (pH, p0J above about 600 mm. the reaction rate is so fast that isothermal conditions cannot be maintained and explosions occur. The velocity of this reaction is catalyzed by the surface of the reaction vessel and by foreign gases, excess hydrogen and oxygen as well as gaseous water. I n general, the velocity is decreased by increasing the surface area available, and the presence of foreign gases produces a positive catalytic action.= This is illustrated in Table 1.
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TABLE 1 Efsct of Surface Area andNitrogen on the Rate of Reaction Between Hydrogen and Oxygen a t Constant Temperature, 555°C. Mixture composition . Time fw 60 per cent reaction in Unpacked Packed H2 0% Ns, silica bulb, silica bulb, mm. mm. mm. Area = 266 cm.' Area = 1246 400 200 ... 3 hrs. 14 min. 40 hrs. 46 min. 200 100 . . . 14 hrs. 0 min. 250 hrs. ... 200 100 379 3 hrs. 53 min. 180 hrs. . ..
termined by the following procedure. Hydrogen is admitted into the reaction vessel contained in a furnace until its pressure is high enough to permit the introduction of a h o w n amount of oxygen without the occurrence of an explosion. The mixture is then slowly withdrawn, the temperatureof the furnacebeing constant and known. When the total pressure falls to a quite reproducible value, a bright flash and audible explosion occur. The amount of hydrogen and oxygen which react during the withdrawal period is negligible. Using this technique the following facts were obtained, with respect to the upper critical pressure, by Thompson and Hin~helwood.~ (a) At constant temperature, it is lowered by an increase in the ratio of hydrogen to oxygen (see Table 2). Effect of Hx/02 Patio on Uppe. Critical Pressure at Constant Temperature, 550°C.
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Total pressure (mm.) Parlial pressures (mm.) at which explosion occurs 8 / O 2 (upper e~itiealp~essure) OS Hs
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A
(b) At constant temperature, it is practically independent of the nature and extent of surface area avail-
The unusual feature of the hy*ogen-oxygen reaction, although exhibited by several other gas-phase reactions involving oxygen, is that a finite reaction rate a t high pressures may become explosive as the pressure is reduced either by the reaction itself or by withdrawing a portion of the mixture. This phenomena of explosive limits was first discovered in 192'8 by Thompson and Hiushelwood2and independently confirmed by Semenoff and his coworker^.^ The general form of the dependence of the reaction rate upon the pressure of hydrogen and oxygen is given in Figure 1. The figure clearly shows the existence o f a lower and an upper critical pressure between which the mixtures are explosive; below and above these explosive limits, the reaction rates'become finite. The pressure value of the upper critical limit was deSCHENBERG, S., THISJOURNAL, 22, 537 (1945). THOMPSON, H. W., AND C. N. HINSHELWOOD, Pmc. Roy. SOC. Total Pmsswa (pH* + p03. London. A122. 610 (19281. Figure 1. General form of dopsnd~ncaof reaction rate upon pres, , sure. The values m w l y indicate the orden of magnitude for a 2 : 1 KOPP,D., A. KOWALSKY, A. SAGULIN, AND N. SEMENOPF, hydrogen-oxygen mirturs by volume at a temperatu.. of 550'C. PluZ.physik. Chem., 6B,307 (1929). tial nrossures of either glu may also be used for the figure.
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able. Of course, if the is ahove the limit, the tures of hydrogen and oGYgen, for which the reaction reaction rate is decreased by increasing the surface rate a t a given temperatu're is negligibly small, become area of the reaction vessel. explosive in the presence of quite small concentrations (c) The upper critical pressure decreases rapidly of NOz, but beyond a definite amount of NOz the reacwith decreasing temperature (see Table 3): tion again becomes very slow. At 410°C. the lower and upper pressure limits are 0.053 mm. and about 10 mm. NO2 for a 2:l hydrogen-oxygen mixture, the total TARTS! 2 Depndence of the Upper Critical Pressure Upon Tempera-. pressure being 150 mm. ture; Ratio of HJOe Constant and Equal to 2.0 Although the reaction between hydrogen and oxygen, as usually expressed, appears to be one of the simplest Total pressure (mm) Temperatum, 'C. (upper critical pressum) reactions, its kinetics are extremely complex and a completely satisfactory mechanism to account for the presence of explosion limits has not yet been proposed. Qualitative explanations are based largely upon the 440 No explosion theoretical work of Christiansen and Kramers6who have shown that reaction chains may he of two types. These (d) Neither the rate of.withdrawa1 nor the length are the so-called stationary and nonstatiouary (or of time the mixture remains in the reaction flask has any branching) chains. A stationary chain or condition effect on the pressure a t which the explosion occurs. results when each chemical change produces only one Hence, a "time-lag" process or the diffusion of the ini- excited molecule capable of continuing the chain. For tial reactants is not involved in the mechanism of the such reactions the rate is finite (but it is possible in exoreaction. thermic reactions for such reactions to go so fast that Semenbff and his coworkers3have studied mainly the isothermal conditions cannot he maintained, leading t o lower explosion limit. The .lower limit was also subse- a thermal explosion). For example the'reactions quently investigated by Hinshelwood and MoelwynH,O* 0, = H,O + O , * hug he^.^ The reaction rate just below the lower limit 2H3 + O%*= H,O* + H,O is practically zero. Unlike the upper limit; the lower pressure limit is practically independent of the relative would lead to a stationary chain. The reactions pressures of the two gases and it increases with deH* Cl = HCI H creasing size of the reaction vessel (see Tahle 4). Ftom H + CL = HCl + Cl A " "
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T A R T S ! d.
Effect of Reaction-Vessel Size on the Lower Pressure Limit at Constant Temperature, 550% Bulb diamete~(cm.)
Total pressure (mm.) (lower pressuve limit)
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typify a case in which a stationary chain results in a thermal explosion (hydrogen-chlorine mixture exposed . .. to light). On the other hand. if the chemical reaction assumed ahove were 2H3
+ OS* = HaO* + HsO*
yielding thereby two excited (or activated) molecules capable of continuing the chain for one excited molecule 550" to 650°C. the influence of tempergture upon the undergoing chemical change, then branching (nonlower limit is negligible but below 550%. the lower limit stationary) chains result. For such reactions the rate increases (see Tahle 5); reaches explosive velocities. The nuclear chemical reaction between UZ3$ or p~~~~ and neutrons is the latest TABLE 5 . example of a nonstationary chain reaction. These reEffect of Tem~eratureU ~ o n the Lower h s s u r b Limit actions become explasive when a t least two neutrons Critical presswe (mm.) . are emitted for every one absorbed (captured). Temperature, T . (total) Hence, a plausible explanation of the low pressure reaction (the presence of explosion pressure limits) is t o ,550 ? . assume that a branching chain mechanism is involved aud that a t pressures below the lower limit the excited molecules lose their energy mainly to the walls of the It was previously mentioned that foreign gases vessel by diffusion. For example, catalyze the hydrogen-oxygen reaction. The effect of 2H1 02' = 2HaO on wall NOa is, however, worthy of further notice since its H+O" = K O- on --. -- wall .. ~ ~ - presence induces explosion limits in nonreactive mixtures of hydrogen and oxygen. It was found5that mix- As the total pressure increases, the rate of diffusion of activated molecules to the walls decreases and the reac' HINSHELWOOD, C . N., AND E. A. MOELWYN-HUGHES, Proc. tion becomes explosive. . At pressures ahove the upper
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Roy. Soc. London, A138, 311 (1932). ibid., A124,219 TKOMPSON, H. W . ,A N D C . N . HINSKELWOOD, (1929).
WHRISTIANSEN, J . A,, AND H. A. KRAMERS, 2.phpik. Chem., 104, 451 (1923).
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limit, it is necessary to assume that the excited molecules lose their energy by some gas-deactivation process. On the basis of this mechanism, the effect of inert gases, which is in general to decrease the pressure value of both explosive limits, can also be explained qualitatively. The presence of foreign molecules may decrease a t low pressures the rate of diiusion of activated molecules to the walls and thereby cause a positive catalytic action. At the higher pressure limit, where the inert gases produce a negative catalytic action, the foreign molecules may accelerate the deactivation of excited species in the gas phase. The mechanism for the low-vressure (ex~losivelimits) reaction proposed by semen& may be k e d to summarize this discussion: 1. Some thermal or catalytic wall reaction yielding 0 as the initial activated centers. 2. Development of branching chains:
+ H, = H,O* H1O* + On = HIO + 0 + 0 0
(1) (2)
3. At pressures below the lower limit diffusion to the wall balances the branching process; destruction or deactivation of chains a t walls: 0
+ wall = ' 1 ~ 0 ~
(3)
4. Deactivation in gas phase a t the upper limit sufficient to balance the branching is brought about by a reaction involving a triple collision O+H,+M=H,O+M
(4)
where M is any inert molecule. Norrish and Griffiths' used NO2 as a photosensitizer for the hydrogen-oxygen reaction and not only confirmed the existence of explosive limits induced in nonreacting mixtures by the presence of NO2 but also found that the rate of reaction for a given concentration of NORRISH, R. G. W., AND J. G. A. GRIPFITHS, Proc. Roy. Sac. London, A139, 147 (1933).
NO2 and reactants is increased by exposing the mixture to the light of a mercury vapor lamp (see Table 6). TARLE fi -. Effect of the Light of a Mercury Vapor Lamp on the N o r Catalyzed Hydrogen-Oxygen Combination; Temperature, 2S1°C~
Pressure of NOz (mm.)
0.04 0.266 0.308
Ratio qf mte of fomtion of water in li htlzn da~k, mm.Yw. 0.002/0.GillZ 0.017/0.0024 Exolosion
I t is known' that under the action ,of 'the light, the following chemical reaction occurs: NO* = P.10 + 0 The atomic oxygen so produced may then serve to develop branching chains (step 2 in the mechanism given above). To explain the explosion limits induced by NOz, the equilibrium NO
+ 0%= NO* + 0
is assumed to exist in the mixture. It then follows that a t low NO. Dressures the ceneration of atomic oxeven is -" favored, Ggile a t the higYher NO; pressures the reverse reaction is favored. This latter action removes atomic oxygen and accounts for the upper explosive limit produced by the higher NOa concentrations. Readers who are interested further in the kinetics and mechanism of the hydrogen-oxygen reaction are referred to:! ., . i ,I J .z, . @ , i ;&j ,L;?, , 1flj l ~
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HINSHELWOOD, C. N., "The Kinetics of Chemical Change," Oxford University Press, New York, 1940. SEMENOFF, N., "Chemical Kinetics and Chain Reeotions." Oxford University Press, New York, 1935. "The Reaction HINSAELWOOD, C. N., AND A. T. WILLIAMSON, Between Hydrogen snd Oxygen," Oxford University Press, New York, 1934. .. ..
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