It is informative, however, to consider the possible mechanisms for dry CO oxidation. Again the approach is to consider the explosion limits of a stoichiometric, dry CO–O₂ mixture. However, neither the explosion limits nor the reproducibility of these limits is well defined, principally because the extent of dryness in the various experiments determining the limits may not be the same. Thus, typical results for explosion limits for dry CO would be as depicted in Figure 3.5.

Figure 3.5 reveals that the low-pressure ignition of CO–O₂ is characterized by an explosion peninsula very much like that in the case of H₂–O₂. Outside this peninsula one often observes a pale-blue glow, whose limits can be determined as well. A third limit has not been defined, and if it exists, it lies well above 1 atm.

As in the case of H₂–O₂ limits, certain general characteristics of the defining curve in Figure 3.5 may be stated. The lower limit meets all the requirements of wall destruction of a chain propagating species. The effects of vessel diameter, surface character, and condition have been well established by experiment [2].

Under dry conditions, the chain initiating step is

This reaction is slow, but could build up in supply. Ozone (O₃) is the metastable species in the process (like HO₂ in H₂–O₂ explosions) and could initiate chain branching, thus explaining the explosion limits. The branching arises from the reaction

Ozone destruction at the wall to form oxygen molecules would explain the lower limit. Lewis and von Elbe explain the upper limit by the third-order reaction

However, O₃ does not appear to react with CO below 523 K. Since CO is apparently oxidized by the oxygen atoms formed by the decomposition of ozone (the reverse of reaction (3.42)), the reaction must have a high activation energy (>120 kJ/mol). This oxidation of CO by O atoms was thought to be rapid in the high-temperature range, but one must recall that it is a three-body recombination reaction.

Analysis of the glow and emission spectra of the CO–O₂ reaction suggests that excited carbon dioxide molecules could be present. If it is argued that O atoms cannot react with oxygen (to form ozone), then they must react with the CO. A suggestion of Semenov was developed further by Gordon and Knipe [11], who gave the following alternative scheme for chain branching:

Because these mechanisms did not explain shock tube rate data, Brokaw [10] proposed that the mechanism consists of reaction (3.41) as the initiation step with subsequent large energy release through the three-body reaction (3.47) and

The rates of reactions (3.41), (3.47) and (3.48) are very small at combustion temperatures so that the oxidation of CO in the absence of any hydrogen-containing material is very slow. Indeed, it is extremely difficult to ignite and have a flame propagate through a bone-dry, impurity-free CO–O₂ mixture.

Very early, from the analysis of ignition, flame speed, and detonation velocity data, investigators realized that small concentrations of hydrogen-containing materials would appreciably catalyze the kinetics of CO–O₂. The H₂O-catalyzed reaction essentially proceeds in the following manner:

If H₂ is the catalyst, the steps

At high pressures or in the initial stages of hydrocarbon oxidation, high concentrations of HO₂ can make reaction (3.50) competitive to reaction (3.49), however reaction (3.50) is rarely as important as reaction (3.49) in most combustion situations [4]. Nevertheless, any complete mechanism for wet CO oxidation must contain all the H₂–O₂ reaction steps. Again, a complete mechanism means both the forward and backward reactions of the appropriate reactions in Appendix C. In developing an understanding of hydrocarbon oxidation, it is important to realize that any high-temperature hydrocarbon mechanism involves H₂ and CO oxidation kinetics, and that most, if not all, of the CO₂ that is formed results from reaction (3.49).

The very important reaction (3.49) actually proceeds through a four-atom activated complex [12,13] and is not a simple reaction step like reaction (3.21). As shown in Figure 3.6, the Arrhenius plot exhibits curvature [12]. And because the reaction proceeds through an activated complex, the reaction rate exhibits some pressure dependence [14].

Just as the fate of H radicals is crucial in determining the rate of the H₂–O₂ reaction sequence in any hydrogen-containing combustion system, the concentration of hydroxyl radicals is also important in the rate of CO oxidation. Again, as in the H₂–O₂ reaction, the rate data reveal that reaction (3.49) is slower than the reaction between hydroxyl radicals and typical hydrocarbon species; thus, one can conclude—correctly—that hydrocarbons inhibit the oxidation of CO (see Table 3.1).

It is apparent that in any hydrocarbon oxidation process CO is the primary product and forms in substantial amounts. However, substantial experimental evidence indicates that the oxidation of CO to CO₂ comes late in the reaction scheme [15]. The conversion to CO₂ is retarded until all the original fuel and intermediate hydrocarbon fragments have been consumed [4,15]. When these species have disappeared, the hydroxyl concentration rises to high levels and converts CO to CO₂. Further examination of Figure 3.6 reveals that the rate of reaction (3.49) does not begin to rise appreciably until the reaction reaches temperatures above 1100 K. Thus, in practical hydrocarbon combustion systems whose temperatures are of the order of 1100 K and below, the complete conversion of CO to CO₂ may not take place.

As an illustration of the kinetics of wet CO oxidation, Figure 3.7 shows the species profiles for a small amount of CO reacting in a bath of O₂ and H₂O at constant temperature and constant pressure. The governing equations were described previously in Chapter 2. The induction period during which the radical pool is formed and reaches superequilibrium concentrations lasts for approximately 5 ms. Shortly after the CO starts to react, the radicals obtain their maximum concentrations and are then consumed with CO until thermodynamic equilibrium is reached approximately 30 s later. However, 90% of the CO is consumed in about 80 ms. As one might expect, for these fuel-lean conditions and a temperature of 1100 K, the OH and O intermediates are the most abundant radicals. Also note that for CO oxidation, as well as H₂ oxidation, the induction times and ignition times are the same. Whereas the induction time describes the early radical pool growth and the beginning of fuel consumption, the ignition time describes the time for the onset of significant heat release. It will be shown later in this chapter that for hydrocarbon oxidation the two times are generally different.

Solution of the associated sensitivity analysis equations (Figure 3.8) gives the normalized linear sensitivity coefficients for the CO mass fraction with respect to various rate constants. A rank ordering of the most important reactions in decreasing order is

The reverse of reaction (3.49) has no effect until the system has equilibrated, at which point the two coefficients ∂ln Y_CO/∂ln k_49f and ∂ln Y_CO/∂ln k_49b are equal in magnitude and opposite in sense. At equilibrium, these reactions are microscopically balanced, and therefore the net effect of perturbing both rate constants simultaneously and equally is zero. However, a perturbation of the ratio (k_49f/k_49b = K₄₉) has the largest effect of any parameter on the CO equilibrium concentration. A similar analysis shows reactions (3.21) and (3.24) to become balanced shortly after the induction period. A reaction flux (rate-of-production) analysis would reveal the same trends.