Lewis Combustion, Flames and Explosions of Gases

2. Auflage 2013
ISBN: 978-1-4832-5839-3
Verlag: Elsevier Science & Techn.
Format: EPUB
Kopierschutz: 6 - ePub Watermark

E-Book, Englisch, 752 Seiten

Combustion, Flames, and Explosions of Gases, Second Edition focuses on the processes, methodologies, and reactions involved in combustion phenomena. The publication first offers information on theoretical foundations, reaction between hydrogen and oxygen, and reaction between carbon monoxide and oxygen. Discussions focus on the fundamentals of reaction kinetics, elementary and complex reactions in gases, thermal reaction, and combined hydrogen-carbon monoxide-oxygen reaction. The text then elaborates on the reaction between hydrocarbons and oxygen and combustion waves in laminar flow. The manuscript tackles combustion waves in turbulent flow and air entrainment and burning of jets of fuel gases. Topics include effect of turbulence spectrum and turbulent wrinkling on combustion wave propagation; ignition of high-velocity streams by hot solid bodies; burners with primary air entrainment; and description of jet flames. The book then takes a look at detonation waves in gases; emission spectra, ionization, and electric-field effects in flames; and methods of flame photography and pressure recording. The publication is a valuable reference for readers interested in combustion phenomena.

Autoren/Hrsg.

Lewis, Bernard

Leseproben

CHAPTER II The Reaction between Hydrogen and Oxygen
Publisher Summary
This chapter discusses the thermal reaction between hydrogen and oxygen. There are many possible elementary reactions in the hydrogen-oxygen system, but their relative importance changes radically over diverse ranges of experimental variables. As conditions are altered, some reactions become insignificant while others come to the fore. The reaction mechanism is, therefore, circumscribed by the range of experimental conditions to which it applies. The chapter discusses that for the chain-initiation reaction, the dissociation of hydrogen in bimolecular collisions has been considered in the past. On the face of it, this reaction appeared plausible because it gave an explanation for the marked dependence of the rate of water formation on the partial pressure of hydrogen, and it appeared consistent with the observed overall activation energy in excess of 100 kcal. As neither the dissociation of hydrogen or of oxygen nor the gas-phase reaction between them can be admitted as the chain-initiation reaction, attention is focused on the experimentally established presence of H2O2 in the reacting mixture. If H2O2 is both formed and destroyed in a chain reaction, it attains a steady state concentration that is independent of the chain carrier concentration. 1 The Thermal Reaction
As has already been discussed in the opening paragraph of the previous chapter, a complex reaction consists of an interplay of various elementary reactions. It is the task of the investigator to choose the experimental conditions so that the controlling reactions can be identified and studied from the observed response of the over-all reaction to the experimental variables. The reaction of hydrogen and oxygen in heated vessels furnishes an outstanding illustration of the possibilities and pitfalls of this type of research. Numerous investigators have experienced the difficulty of obtaining meaningful data on the hydrogen-oxygen system. Over the years the principal features of the reaction have become rather well established, but questions of detail still persist and suggest further attention from research workers. There are many possible elementary reactions in the hydrogen-oxygen system, but their relative importance changes radically over diverse ranges of the experimental variables. As conditions are altered, some reactions become insignificant while others come to the fore. The reaction mechanism is therefore circumscribed by the range of experimental conditions to which it applies. If some plausible elementary reaction is not incorporated in the mechanism, it is not implied that the reaction does not occur, but only that it is insignificant under the specified conditions. The mechanism that is described here has been developed by the authors1 for a particularly wide range of parameters. In its principal features it is consistent with the excellent work of Russian authors2 and of Hinshelwood and co-workers3 at Oxford, where the first extensive systematic studies of the reaction were made. In regard to details of the reaction scheme and the values of the kinetic data, some uncertainties and discrepancies with other work still persist. However, since these questions cannot as yet be resolved satisfactorily, and since the mechanism provides a consistent description of the hydrogen-oxygen system over a wide range of observations, it is reproduced here substantially without change. A THE REACTION MECHANISM
In a mixture of hydrogen and oxygen it is a priori plausible that the presence of a free valence in the form of a radical OH or atom H should result in the following reaction cycle: +H2=H2O+H (I) (I) +O2=OH+O (II) (II) +H2=OH+H (III) (III) The role of the free valence is analogous to that in the hydrogen-chlorine reaction (page 3) and also in the hydrogen-bromine reaction. In the latter, however, the step Br + H2 = HBr + H is endothermic by about 20 kcal., so that this reaction is possible only in strongly energized collisions which, require high temperatures to become frequent. At low temperatures a Br atom may be expected to survive many collisions and migrate to the vessel wall where it is adsorbed and ultimately removed by recombination. A similar situation exists with respect to reaction (II) above, which is endothermic by 17 kcal. Thus, at room temperature and considerably above, a mixture of hydrogen and oxygen is very stable even if hydrogen atoms are introduced from an outside source. The free valence terminates ultimately at the wall. Above some temperature, which depends on many variables of the system and cannot be even approximately generalized, the chain-branching reaction (II) becomes sufficiently frequent compared to the rate of removal of H atoms, to cause multiplication of free valences and explosion. The explosion occurs above some critical pressure which is generally very low, and while its value cannot be generally stated, in glass vessels it rarely exceeds the order of a millimeter. Evidently a set of conditions of pressure, temperature, mixture composition, and vessel factors defines the condition a = ß which exists at the explosion boundary. If the pressure is the only variable, then below the critical pressure, free valences diffuse to and are destroyed at the wall at a rate exceeding their rate of formation and a steady-state reaction ensues which, however, under these circumstances, is so slow that the rate is practically unmeasurable. Experimentally, it is observed that an abrupt transition occurs from practically no reaction to explosive reaction on increasing the pressure to the critical. In order to describe the phenomenon the reactions I to III must be supplemented by suitable equations describing the process of diffusion, adsorption, and surface recombination of H, O, and OH. Russian work has shown2,4 that in general it is possible to neglect the diffusion of O and OH because of the high rates of their gas phase reactions. These authors have studied the first explosion limit extensively and have been able to obtain satisfactory correlations between experiment and theory. In this text we shall confine ourselves to experimental conditions sufficiently remote from the first explosion limit where the destruction of H, O, and OH is kinetically insignificant. As the pressure is increased above the first explosion limit the reaction remains explosive until a second critical pressure is reached above which a steady-state reaction is observed. Just above this second explosion limit the reaction rate is very small and increases with pressure until a third explosion limit is reached. The three explosion limits are illustrated in the temperature-pressure diagram of Fig. 1 for a stoichiometric mixture in a 7.4-cm. spherical Pyrex vessel coated with KCl. With increasing temperature the pressure along the first and third explosion limits decreases and along the second explosion limit it increases. The region between the first and second limits is usually referred to as the explosion peninsula. The existence of the second explosion limit is readily explained if the three-body reaction
FIG. 1 Explosion limits of a stoichiometric hydrogen-oxygen mixture in a spherical KCl-coated vessel of 7.4 cm. diameter. First and third limits are partly extrapolated. First limit is subject to erratic changes. +O2+M=HO2+M (VI) (VI) is added to the scheme. In this reaction the symbol M denotes any third molecule that stabilizes the combination of H and O2. If the free radical HO2 is thought to be relatively unreactive so that it is able to diffuse to the wall, it becomes a vehicle for the destruction of free valences and reaction (VI) is a chain-breaking reaction. Since with increasing pressure the frequency of ternary collisions H + O2 + M increases relative to the frequency of binary collisions H + O2, there exists a pressure above which the rate of removal of free valences by reaction (VI) exceeds the rate of formation of free valences by the branching reaction (II) and the condition a < ß is established. The limit itself is defined by the condition a = ß which becomes k2=k6[M]e (1) (1) or M]e=2k2/k6 (2) (2) where the subscript e refers to the explosion limit. The concentration [M] being proportional to the total pressure, equation (2) defines the pressure at the second explosion limit. Since reaction (II) has a large activation energy relative to reaction (VI), a decrease of temperature causes the ratio 2k2/k6 and therefore the explosion pressure to decrease. For the first explosion limit the condition a = ß results in an equation of the form k2[O2]=f(DH, DO, DOH, ?H, ?O, ?OH, d) (3) (3) where the function on the right-hand side contains diffusion and chain-breaking efficiency terms and the vessel diameter, d. In equation (3) the chain-branching term on the left-hand side increases with pressure while the chain-breaking term either...