Ammonia planet- atmosphere

Question

On my previously asked question about the biochemistry of life on my ammonia. it leads me to my next question to ask.

First things first, the life on this planet is boron based. They breathe methane and drink ammonia. Due to their biochemistry being boron based. I realized that the gases such as bromine, chlorine, oxygen and fluorine can't be in large quantities as they form strong bonds with boron making life difficult to form. That however, leads to a big problem with complex intelligent boron life and their ability to make fire. In order to form fire you need three things. Fuel, an ignition source, and an oxidizing agent. Without those 4 main gases being in high quantity in this atmosphere, I lack a common oxidizing agent for life to use.

My question is, what gas in this alien atmosphere could be used as an alternative to those 4 oxidizers that wouldn't pose a threat to the formation of boron based life.

let me know if I got any information wrong in my scenario. I hope this is clear enough to be answerable.

Atmospheric composition if needed- nitrogen- 93% methane- 6% hydrogen- 0.2% other trace gases- ethane, diacetylene, methylacetylene, acetylene, propane, cyanoacetylene, hydrogen cyanide, carbon dioxide, carbon monoxide, cyanogen, argon and helium 0.8%

Answer 1

There really isn't one. What you have there is a heavily reducing atmosphere; any strongly oxidizing gasses will have very short lifetimes, existing only in small quantities and needing to be constantly replenished, either by geochemical or biological activity. No oxidizing gas will support fire in an environment that is otherwise as you have described it.

But, that doesn't mean you can't have fire at all. You just need to flip the roles of oxidizer and fuel. Organisms on this world may produce oxidizers as part of their own metabolism, such that particular tissues are more or less chemically oxidizing, and can react with the reducing gasses in the air; i.e., atmospheric methane and hydrogen may react with biologically-produced hydrazine, nitrogen oxides, peroxides, etc., present in biological tissues.

Alternately, or additionally, both oxidizing and reducing agents may be available in liquid and/or solid form, such that no atmospheric gasses are required for the reaction at all. In other words, the inhabitants of this world may produce materials analogous to guncotton, gunpowder, or various solid rocket fuels which most obviously could be contained in bombs (quite useful for some purposes, not so much for things like cooking or metallurgy), but which could also be used to construct torches or smolder-furnaces for intense directed heat or long-term lower-power heat production, respectively.

Answer 2

Your oxidizer is acetylene. Your fuel is hydrogen.

It would be different in that the roles are be reversed, but it is a reducing planet where it is all reversed. Gathering fuel would be gathering acetylene, here the oxidizer. You would "burn" it with the atmospheric hydrogen, reducing it to methane. Hydrogenation of acetylene is an energetically favorable process (see below) but only a quarter as energetic as oxidizing methane to CO2. But you would get heat and hot CH4 (flame!) and so you could call it fire.

Reducing fires might require strong air currents to provide adequate H2 - methods like a bellows or chimney could work.

https://www.chemteam.info/Thermochem/HessLawIntro3.html

Example #1: Hydrogenation of double and triple bonds is an important industrial process. Calculate (in kJ) the standard enthalpy change ΔH for the hydrogenation of ethyne (acetylene) to ethane:

H−C≡C−H(g) + 2H2(g) ---> H3C−CH3(g)

Bond enthalpies (in kJ/mol): C−C (347); C≡C (839); C−H (413); H−H (432)

Solution:

1) You have to put energy into a bond (any bond) to break it. Bond breaking is endothermic. Let's break all the bonds of the reactants:

one C≡C ⇒ +839 kJ two C−H ⇒ 413 x 2 = +826 kJ two H−H ⇒ 432 x 2 = +864 kJ The sum is +2529 kJ

Note there are two C−H bonds in one molecule of C2H2 and there is one H−H bond in each of two H2 molecules. Two different types of reasons for multiplying by two.

2) You get energy out when a bond (any bond) forms. Bond making is exothermic. Let's make all the bonds of the one product:

one C−C ⇒ −347 kJ six C−H ⇒ −413 x 6 = −2478 The sum is −2826 kJ

3) ΔH = the energies required to break bonds (positive sign) plus the energies required to make bonds (negative sign):

+2529 + (−2825) = −296 kJ/mol