# How can hydrogen and oxygen coexist in high temperature? Where did the oxygen come from?

A gasification worked with the combustion of bio pallets with production of syn gases. Pure oxygen and $$\ce{H2O}$$ were fed into the combustion reactor to combust the pallets with no opening of any syn gas outlets. Once the pallets got ignited enough, then there was no more oxygen input. A kind of gas reformer (catalyst) was equipped in the reformer. The temperature in the reactor reached more than $$\pu{1000 ^{\circ}C}$$. Then all the syn gases were collected and analyzed. What I found was that the composition of the gases were methane (43%), hydrogen (23%), oxygen (20%), and other gases (carbon dioxide and monoxide). I don’t understand how such a high volume of oxygen coexisted with hydrogen together in that high temperature reactor. Is it possible that those gases stay together without an oxidation reaction occurring? Where did the oxygen come from?

• The fact that there is methane and carbon monoxide suggest that there is incomplete combustion. You do not mention if pure oxygen or air was used and was heat alone used or a flame ? – porphyrin May 25 at 11:15
• Was there any H2O? Did the composition change after being at 1000C for different times? – James Gaidis May 25 at 12:47
• @porphyrin I am sorry I didn't describe it well. It was pure oxygen to burn the pallets. Once the pallets was ignited enough, then there was no input of oxygen anymore. I thought that all oxygen added were used for the oxidation of carbon. The reactor was a completely closed system. I am sorry I can't let you know the whole system of the reactor in detail. Also a gas reformer (catalyst) was equipped. The pallets turned to char coals and the syn gas was collected. – Clayweaver May 25 at 19:40
• @James Gaidis I actually tried to measure them at the second round running. I asked for an institute to analyze the syn gases after it cooled down (~300 degree Celsius). They failed to analyze them. Their GC couldn't read the gases. They said that the GC coould read the standard samples but not my gas samples. They made errors to read. However, I found that the reactor produced alcohols (I guess it's methanol) condensed on the bottom. I think that the methanol might be a mixture with H2O. – Clayweaver May 25 at 20:07

I'll answer the question in general of how Hydrogen and Oxygen can coexist in metastable mixtures at higher temperatures, though without hazarding a guess for your specific case. Like you and others, I'd be very surprised if your bulk gas is actually well mixed at $$\pu{900 ^{\circ}C}$$ without exploding! It's far enough beyond the autoignition temperature, that an explosive reaction seems inevitable. However, metastable combinations of methane, hydrogen and oxygen in the $$\pu{550 ^{\circ}C}$$ range have been reported in the literature.

Hydrogen can be a surprisingly picky burner. I was told by a colleague who specialized in rocket propulsion that in between Hydrogen's upper and lower explosive limits are a host of regions where it strongly resists ignition, even above its reported autoignition temperature of ~$$\pu{570 ^{\circ}C}$$. According to her, the reaction between hydrogen and oxygen is quite complex with several important intermediates. Having concentrations in these non-explosive regions apparently allows the hydrogen-oxygen mix to be metastable for long periods. This would of course not extend to long time periods at $$\pu{900 ^{\circ}C}$$!

More in my territory, there's also many chemicals which inhibit the slow combustion (non-explosive) of hydrogen and oxygen up through the $$\pu{500 ^{\circ}C}$$s. The most common is $$\ce{KCl}$$, which strongly inhibits the reaction of oxygen and hydrogen in non-explosive conditions allowing a metastable mixture to exist. Walls of vessels can also similarly inhibit reaction. The reaction mechanism apparently stems from $$\ce{KCl}$$ and other surfaces' abilities to consume reaction intermediates of the combustion such as monatomic Hydrogen:

$$\ce{KCl + H^. -> K^. + HCl}$$ $$\ce{SiO2 + H^. -> SiO(OH)}$$ or direct adsorption on metal surfaces like nickel, palladium, etc.

After doing a bit of digging, methane is apparently a superb inhibitor of hydrogen and oxygen combustion. The effect is explained in a 1959 paper by R. Baldwin, et al. (see following reference). The paper explains that when methane is above a sharp critical concentration it drastically raises the autoignition temperature of hydrogen's explosive reaction and almost completely eliminates slow combustion in non-explosive conditions. In other words, when enough methane is present, the hydrogen and oxygen mixture either does nothing or explodes completely. From data taken in systems between $$\pu{450\! -\! 550 ^{\circ}C}$$, there was a very clear relationship of methane concentration vs. explosive limit taking the form of (values approximate):

Explosive Limit (in $$\pu{K}$$) = $$\frac{\pu{3350K}}{2.3-\log_{10}(X_\ce{CH4})}$$

where $$X_\ce{CH4}$$ is the mole fraction of methane, studied in the 0.01 - 0.1 range (significantly lower than yours). For this particular heuristic, the mole fraction of hydrogen was 0.28, oxygen 0.14, and the balance was nitrogen. The paper suggests the mechanism is non-equilibrium formaldehyde formation.

As a final note, please be very cautious when handling or recreating these mixtures. Just because the hydrogen isn't burning, doesn't mean it can't explode. A lab down the hall from mine was fortunate to loose only a wall to a hydrogen evolving system which ran fine for months and then, quite suddenly, decided to explode.