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An abandoned outpost from a long extinct race is discovered. It can easily be repaired and would make an ideal location for a colony, however it has been breached and any usable air has been long lost to space. The outpost has its own power systems that we are capable of turning back on and using, however there is no method on the station to create a breathable environment for humans. It's quite the journey and a colony ship would need to be in transit for more than a decade to reach its destination.

If a colony ship was dispatched to this location, what would it need to bring to generate enough Earth-like air for the population to breathe and use?

My first idea was solid ammonia and liquid oxygen… Burn the ammonia and let it release its nitrogen along with water. Use electrolysis to recover some of the oxygen back as required (or leave it as usable water). The issues here:

  • Way too high of a oxygen to nitrogen ratio… as nice as ammonia is for this, the nitrogen content is relatively low and is created at a 3:1 ratio with oxygen.
  • Need to carry too much liquid oxygen. It would be nice if one chemical compound can be released instead of two separate ones… also not a fan of storing high pressure gas on a long trek colony ship
  • Too much hydrogen… there are a few uses for it (water/fuel), however the ratio of nitrogen to hydrogen is too low and an excess of hydrogen is created in trying to get Earth-like conditions.

Assume any chemical compound can be reasonably manufactured. Liquid gasses require high pressure and have some degree of danger in space transport… solid compounds preferred, but not an end all.

What would you take on your colony ship to create an Earth-like atmosphere in as little amount as storage space possible and as safely as possible?

It is preferable to bring the material on the ship as opposed to trying to harvest it. The tech level on the propulsion systems is limited and the speed to get there depends on some gravitational slingshotting to get it up to speed already, so there isn't the opportunity to stop to harvest anything and then re accelerate.

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  • 1
    $\begingroup$ @Nick2253 See this question $\endgroup$ – DoubleDouble Feb 18 '15 at 22:14
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    $\begingroup$ I figured CO2 could be generated by burning a small amount of the plant matter thats part of the farming operation on the colony ship. Or bring a small dead tree to burn upon arrival. $\endgroup$ – Twelfth Feb 18 '15 at 22:15
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    $\begingroup$ @DoubleDouble A pure oxygen environment is regularly used on Earth for medical purposes. And without something to oxidize, a pure oxygen atmosphere won't be flammable. $\endgroup$ – Nick2253 Feb 18 '15 at 22:17
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    $\begingroup$ What's wrong with frozen air? $\endgroup$ – Oldcat Feb 18 '15 at 22:33
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    $\begingroup$ No one did put Spaceballs referrence here? Am I really the first one to do it? Ok, here you go: youtube.com/watch?v=VptOUWC-Itc $\endgroup$ – Pavel Janicek Feb 19 '15 at 10:02
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Per your question, it sounds like you preferred something that was not a pure oxygen atmosphere. In that case, nitrous oxide ($N_2O$) might be an a good match. At 1 atmos. pressure it becomes liquid at -88.5 °C solid at -90.9 °C easily maintained at a 10 year trip in space.

On arrival, $N_2O$ could be easily decomposed into $O_2$ and $N_2$.

I am not aware of any stable compound that has a ratio closer to the N/O ratio of earth's air. The atmos. pressure in your station would only be 60% of "normal" but that might be acceptable.

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Ammonium nitrate melts at 169 C, and leaves no residue when heated:

$$ \text{NH}_4\text{NO}_3 → \text{N}_2\text{O} + 2\text{H}_2\text{O} $$

Like Gary says $\text{N}_2\text{O}$ can be decomposed into nitrogen and oxygen.

But for the reality check, hydrogen, nitrogen and oxygen are very common elements. The colony ship is bound to find some compounds they could extract it from, although it takes longer.

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    $\begingroup$ Good pt. re: availability. You would also be very likely to want water and food. CHON -- it's what we are made of. I have to assume this question is intended as a startup supply to tide you over before you start mining comets and asteroids. $\endgroup$ – Gary Walker Feb 19 '15 at 19:17
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    $\begingroup$ +1, I gave the check to Gary's answer, but your expnansion here is very useful. Thankyou. $\endgroup$ – Twelfth Feb 20 '15 at 17:57
  • $\begingroup$ I definitely like the advantages that this has over N2O, not only that it's a solid at room temperature (much easier to load into your space ship) but that it also decomposes into N2O and water. $\endgroup$ – Draco18s Jun 21 '18 at 1:29
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People in an enclosed environment don't need anything but oxygen. While it's toxic at high partial pressures, you can account for this by having a lower ambient air pressure than what exists on Earth. For example, an atmosphere with a pressure of .2 bar will have enough oxygen for people to breathe, but won't be toxic. Issues with flammability are also tackled by reducing partial pressure, since it is what determines reaction rates in a gas. Reducing partial pressure will also reduce the structural loads on your habitat, making it easier to patch holes and reducing the rate of air loss if there are any small breaches.

If you don't like liquid oxygen, there's another convenient way to transport it: liquid ozone. Ozone has a higher boiling point than oxygen by about 20 degrees, and is also 50% denser than oxygen. It also self-oxidizes, so you can burn it to produce energy, with the byproduct being oxygen.

The downside is that anything that's a strong enough oxidizer that you can burn it to produce another strong oxidizer is a really good oxidizer which means that a tank of liquid ozone laying around needs to be treated really carefully, or it will explode, even without anything else to burn, because ozone will burn itself. But hey, lower volume for storage and free energy if you treat it right.

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  • $\begingroup$ Assuming there is some co2 added to the mix, would plant life be able to handle this? And it still seems like a relatively high risk of combustibility, even at the lower pressures...am I wrong there? $\endgroup$ – Twelfth Feb 18 '15 at 23:14
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    $\begingroup$ Combustibility is mostly a matter of partial pressure rather than fractional concentration. $\endgroup$ – Mark Feb 19 '15 at 0:21
  • $\begingroup$ I like the answer, but I'm going with the check mark beside the answer that came close to an earth environment. $\endgroup$ – Twelfth Feb 20 '15 at 17:57
  • $\begingroup$ WRT flammability in a pure oxygen atmosphere, remember Apollo 1. $\endgroup$ – jamesqf Nov 9 '16 at 17:02
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    $\begingroup$ Ozone is unstable. It will try to revert back to dioxygen, and will produce lots of energy doing so. Hot ozone is more unstable, but even liquid or solid ozone won't last 10 years. Also if it starts to heat up, it will explode. And the very hot, high-pressure newly created oxygen will be happy violently oxidizing everything it touches, be it by corroding it or setting it on fire. $\endgroup$ – Eth Jun 18 '18 at 15:01
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There are several sustances that can use to generate your atmosphere.
Sice they are quite large, right now I'll post only one of them and when I have more time I'll post the other ones.

You will note that this answer is quite large. That is because I also explain how to produce the sustance, and how to produce the sustances which produce it (maybe that information could be useful for someone who is willing to manufacturate the compound instead of bring it from Earth). If you aren't looking for that, only read up to production title (without include it), In order words, only read the first quarter/fifth of the post.

Also note, that the most important part of my answer is under the tile usage.

Finally, if someone have any question, ask in comment. I'll do my best to address the problem.
Good luck reading!

Nitrous Oxide $\text{N}_2\text{O}$ (Or ammonium nitrate $\text{NH}_4\text{NO}_3$)

$\begin{array}{|c|c} \hline \text{Molar Mass} & 44.013 \text{g/mol} & \\ \text{Density} & 1.977 \text{g/L (gas)} \\ \text{Melting Point} & −90.86 \text{°C} \\ \text{Boiling point} & −88.48 \text{°C} \\ \text{Price} & \text{¿?} \\ \hline \end{array}$

Or ammonium nitrate, see on industrial production.

Introduction: Commonly know as laughing gas or nitrous. At room temperature, it is a colorless non-flammable gas, with a slight metallic scent and taste, and faint, sweet odour. At elevated temperatures, nitrous oxide is a powerful oxidizer similar to molecular oxygen.

Uses: Futhermore build an atmosphere, you can use nitrous oxide in several ways.

  • Recreational (due to euphoric effects upon inhaling).
  • Anaesthetic and pain reducing on surgery and dentistry. (At 50% concentration can be administred by non-professional people)
  • Rocket propellants (used as an oxidizer). Can be use also as a monopropellant rocket with a catalyst.
  • Motor racing to increase the power output of engines.

Advantages:

  • It's non-toxic. (Long-exposure and abuse can produce vitamin B12 defficiency...)
  • It's stable at room temperature -making easy to store and carry on flight-.
  • It's possible to decomposed readily to from breathing air.
  • It has high density and low storage pressure (when maintained at low tempertarue). That enable it to be highly competitive with stored high-pressure gas systems.

Storage:

This can be storaged in 20-25 bars tanks at -20 °C or in high pressure tanks at 45-60 bars.

Usage:

In the presence of a heated catalyst, $\text{N}_2\text{O}$ will decompose exothermically into $\text{N}_2$ and $\text{O}_2$, at a temperature of approximately 577 °C (this even produce energy!).

$$2\text{N}_2\text{O} \Rightarrow 2\text{N}_2 + \text{O}_2 + 82 \text{ kJ/mol}$$ $\begin{array}{|c|c|c|c|c} \hline \text{Examples} & \text{N}_2\text{O} & \text{N}_2 & \text{O}_2 & \text{Energy}\\ \hline \text{by Mass} & 1,000\text{g} & 636.48\text{g} \text{ }(63.64\text{%}) & 363.51\text{g} \text{ }(36.35\text{%}) & 1,863.08\text{ kJ}\\ \text{by Volume} & 1,000\text{cm}^3 & 1,006.33\text{cm} \text{ }(100.63\text{%}) & 502.92\text{cm}^3 \text{ }(50.29\text{%}) & 3,683.32\text{ kJ}\\ \hline \end{array}$

With nitrous oxide you can build and atmosphere of:

$\begin{array}{|cc|r|cc|r|} \hline \text{Chemical} & \text{gr/mol} & \text{Percentage} & \text{Mol Fractal} & \text{Mol} & \text{Partial Pressure} \\ \hline \text{N}_2 & 28.0134 & 63.65\text{%} & 2.272 & 0.666 & 67.550 \text{ kPa} \\ \text{O}_2 & 31.9988 & 36.35\text{%} & 1.135 & 0.333 & 33.775 \text{ kPa} \\ \hline \text{Total} & 60.0122 & 100.00\text{%} & 3.408 & 1.000 & 101.325* \text{ kPa} \\ \hline \end{array}$
101.325 kPa = 1 atm

Okey... that is... fine.

  • A bit high oxygen value: Humans need around 21 kPa of oxygen to "work" properly, you will have 33.775 kPa, your people wouldn't suffer hyperoxia. Hyperoxia is produced when oxygen is above 50 kPa it becomes toxic. You will survive, but I think (complete and subjetive personal opinion) your crew will need a few days to adapt their lungs after stop coughing. Also, (again, my opinion) they will suffer some slightly diseases or deficiencies after some years.
    • You can improve that reducing the overall pressure from $101.325\text{ kPa} \rightarrow 63\text{ kPa}$. So you will have $42\text{ kPa N}_2$ and $21\text{ kPa O}_2$. Good!
      Also, reduce the overall pressure has an advantage: if there is a damage in the ship/station's hull, air will scape slowler from it, giving more time from crew to fix it.

Also, because of the large heat release, the catalytic action rapidly becomes secondary, as thermal autodecomposition becomes dominant (that means you don't need to store much catalyst).
Its catalyst can be any of this (I think): (Note that I don't know about catalysts and their usages)

  • Rhodium: ($ 72.66/g). Density 12.41 g/cm3.

  • Cobalt: ($ 81.50/kg). Density 8.90 g/cm3.

  • Platinum: ($ 28.56/g). Density 21.45 g/cm3. Strong.

  • Palladium: ($ 31.85/g). Density 12.023 g/cm3

  • Cobalt: ($ 81.50/kg). Density 8.90 g/cm3.

  • Copper: ($ 7.20/kg). Density 8.96 g/cm3. Weak.

  • Cerium: ($ 5.51/kg). Density 6.770 g/cm3. Weak.

  • Iron: ($ 71/t for ore). Density 7.874 g/cm3. Weak.

  • Nickel: ($ 15.24/t). Density 8.908 g/cm3. Weak.

Prices from http://www.infomine.com (except cerium). I add density in the list because maybe it could be important for an space ship. Catalyst taken from physicsforums.com.

Production:

Nitrous oxide can be produced by two ways, the industrial way and the laboratoy way. I have already asked on Chemistry SE which is the difference between an "industrial" method and a "laboratory" method (Saldy, the used removes its answer, I'm not sure why).

Industrial production:

Nitrous oxide is prepared on an industrial scale by careful heating of ammonium nitrate at about 250 ºC, which decomposes into nitrous oxide and water vapour. Amonium nitrate density 1.725 g/cm3. The addition of various phosphate salts favours formation of a purer gas at slightly lower temperatures. This reaction may be difficult to control, resulting in detonation. enter image description here

$$\text{NH}_4\text{NO}_3\text{ (solid)} + 36\text{ kJ/mol} \Rightarrow 2\text{H}_2\text{O} \text{ (gas)} + \text{N}_2\text{O} \text{ (gas)}$$ $\begin{array}{|c|c|c|c|} \hline \text{Examples} & \text{NH}_4\text{NO}_3 & \text{H}_2\text{O} & \text{N}_2\text{O} & \text{Energy}\\ \hline \text{Mass} & 1,000\text{g} & 549.86\text{g} \text{ } (54.98\text{%}) & 450.14\text{g} \text{ } (45.01\text{%}) & -449.75\text{ kJ}\\ \text{Volume} & 1,000\text{cm}^3 & 636.57\text{cm}^3 \text{ } (63.65\text{%)} & 260.95\text{cm}^3 \text{ } (26.09\text{%}) & -775.83\text{ kJ}\\ \hline \end{array}$

That means that produce nitrous oxide also produce water to drink, yeah! But, what if we use water to produce more oxygen with electrolisys?:

$$\text{H}_2\text{O} + 241.8\text{ kJ/mol} \Rightarrow \text{H}_2 + \frac{1}{2}\text{O}_2$$

$\begin{array}{|c|c|c|c|} \hline \text{Examples} & \text{NH}_4\text{NO}_3 & \text{O}_2 & \text{N}_2 & \text{H}_2 & \text{Energy}\\ \hline \text{Mass} & 1,000\text{g} & 599.65\text{g} \text{ } (59.96\text{%)} & 349.97\text{g} \text{ } (34.99\text{%)} & 50.36\text{g} \text{ } (5.03\text{%)} & 5.46\text{ MJ}\\ \text{Volume} & 1,000\text{cm}^3 & 723.86\text{cm}^3 \text{ } (72.38\text{%)} & 482.73\text{cm}^3 \text{ } (48.27\text{%)} & 966.49\text{cm}^3 \text{ } (96.64\text{%)} & 9.43\text{ MJ}\\ \hline \end{array}$

$$\text{NH}_4\text{NO}_3\text{ (solid)} + 437.6\text{ kJ/mol} \Rightarrow \text{N}_2 \text{ (gas)} + 2\text{H}_2 \text{ (gas)} + 1\frac{1}{2}\text{O}_2 \text{ (gas)}$$

What do you want to do with all the hydrogen is your problem, not mine (may I suggest nuclear fussion).

With that you can make this atmosphere:

$\begin{array}{|cc|r|cc|r|} \hline \text{Chemical} & \text{gr/mol} & \text{Percentage} & \text{Mol Fractal} & \text{Mol} & \text{Partial Pressure} \\ \hline \text{N}_2 & 28.0134 & 36.85\text{%} & 1.315 & 0.4 & 40.530 \text{ kPa} \\ \text{O}_2 & 31.9988 & 63.15\text{%} & 1.973 & 0.6 & 60.795 \text{ kPa} \\ \hline \text{Total} & 60.0122 & 100.00\text{%} & 3.288 & 1.000 & 101.325* \text{ kPa} \\ \hline \end{array}$
101.325 kPa = 1 atm

Okey... that is... lethal.

  • High oxygen value: Humans need around 21 kPa of oxygen to "work" properly, you will have 60.795 kPa, your people would suffer hyperoxia. Hyperoxia is produced when oxygen is above 50 kPa it becomes toxic. You won't survive mora than 1 or 2 days (maybe less). In this answer I talk of hyperoxia symptoms (also, there is a cute image).
    • You can fix that reducing the overall pressure from $101.325\text{ kPa} \rightarrow 35\text{ kPa}$. So you will have $14\text{ kPa of N}_2$ and $21\text{ kPa of O}_2$. Good!

But now the question is: Where can I find ammonium nitrate?

Ammonium nitrate is a white crystal solid and is highly soluble in water. It is predominantly used in agriculture as a high-nitrogen fertilizer. Its other major use is as a component of explosive mixtures used in mining, quarrying, and civil construction.

Harverst of ammonium nitrate: $\text{NH}_4\text{NO}3$

$\begin{array}{|c|c} \hline \text{Molar Mass} & 80.043 \text{g/mol} & \\ \text{Density} & 1.725 \text{g/cm}^3\text{ (solid)} \\ \text{Melting Point} & 169.6 \text{°C} \\ \text{Boiling point} & 210 \text{°C} \\ \text{Price} & \text{USD } 180-400/\text{ton} \\ \hline \end{array}$

Ammonium nitrate is found as a natural mineral (gwihabaite) [...] often as a crust on the ground and/or in conjunction with other nitrate, iodate, and halide minerals.

Production of ammonium: You can also produce it in several ways, but I will talk about a way because the anothers needs limestones, which currently is not avariable on space, or some components that end falling on a circular reference so quckly. The produciton entails the acid-base reaction of ammonia with nitric acid

$$\text{HNO}_3+\text{NH}_3 \Rightarrow \text{NH}_4\text{NO}_3$$

$\begin{array}{|c|c|c|c|} \hline \text{Examples} & \text{HNO}_3 & \text{NH}_3 & \text{NH}_4\text{NO}_3\\ \hline \text{by Mass} & 787.20\text{g} & 212.77\text{g} & 1,000\text{g} \\ \text{by Volume} & 897.56\text{cm}_3 & 502.78\text{cm}_3\text{ (gas)} & 1,000\text{cm}_3 \\ \text{"} & \text{"} & 0.53\text{cm}_3\text{ (liquid)} & \text{"}\\ \text{"} & \text{"} & 0.44\text{cm}_3\text{ (solid)} & \text{"}\\ \hline \end{array}$

Ammonia is used in its anhydrous form (i.e., gas form) and the nitric acid is concentrated. This reaction is violent owing to its highly exothermic nature.

Ammonia: $\text{NH}_3$

$\begin{array}{|c|c} \hline \text{Molar Mass} & 17.031 \text{g/mol} & \\ \text{Density} & 0.73 \text{g/L (gas)} \\ \text{"} & 681.9 \text{g/L (liquid)} \\ \text{"} & 817 \text{g/cm}^3\text{ (solid)} \\ \text{Melting Point} & −77.73 \text{°C} \\ \text{Boiling point} & −33.34 \text{°C} \\ \text{Price} & \text{USD } 400/\text{ton?} \\ \hline \end{array}$

Sadly, I couldn't find a way to produce ammonia that:

  • Doesn't fall into a circular reference.
  • It isn't made on volcanic region (because there aren't volcanoes on space).
  • Doesn't use coal (there isn't on space).
  • Doesn't use quicklime (I don't think there are quicklime asteroids).
  • Doesn't consume hydrogen (because that need water) nor nitrogen (because that is only harvested from the atmosphere...).

Maybe this could be useful:

[...] ammonia was obtained by the dry distillation of nitrogenous vegetable and animal waste products, including camel dung, where it was distilled by the reduction of nitrous acid and nitrites with hydrogen;

And nitrous acid can be made with nitrine and water, and nitrine... not sure (some ways are circular references, other one are too strange for me).

Nitric Acid: $\text{HNO}_3$

$\begin{array}{|c|c} \hline \text{Molar Mass} & 63.01 \text{g/mol} & \\ \text{Density} & 1.5129 \text{g/L (liquid)} \\ \text{Melting Point} & −42 \text{°C} \\ \text{Boiling point} & 83 \text{°C} \\ \text{Price} & \text{USD ~} 300 \text{-} 375/\text{ton (68% pure)} \\ \hline \end{array}$

Nitric acid is made by reaction of nitrogen dioxide (NO2) with water:

$$3\text{NO}_2 + \text{H}_2\text{O} \Rightarrow 2\text{HNO}_3 + \text{NO}$$

Normally, the nitric oxide (NO) produced by the reaction is reoxidized by the oxygen in air to produce additional nitrogen dioxide $2\text{NO} + \text{O}_2 → 2\text{NO}_2$... maybe a way to dispatch the oxygen spares?

Sadly, the nitrogen dioxide is made with oxygen and nitrogen (which is found on Earth atmosphere).

The other way to make nitric acid is with Hydrogen peroxide and nitrogen dioxide, or with ammonia and oxygen, but produce hydrogen peroxide it too complex and both ways fall on a circular reference. Another way could be with Dinitrogen pentoxide and water, but as you guessed, produce dinitrogen pentoxide make a circular reference.

Laboratory Production

There are several ways to archive this on a lab:

  • I won't talk about the production with urea, nitric acid and sulfuric acid because artifial urea is pointless (use ammonia) and sulfuric acid use oxygen and sulfur, so is also pointless.
  • Ostwal process (Direct oxidation of ammonia with a manganese dioxide-bismuth oxide catalyst) is also pointless because an oxidation obiously need oxygen.
  • Hydroxylammonium chloride reaction with sodium nitrite can't be used because the first compound is very difficult to get (organic).
  • Treating $\text{HNO}_3$ with $\text{SnCl}_2$ and $\text{HCl}$ can't be used because we can't harvest $\text{SnCl}_2$.
  • Hyponitrous acid descomposition can't be used because it's difficult to produce (we need silver hyponitrite).

So our last alternative is the heating of a mixture of sodium nitrate and ammonium sulfate:

$$2\text{NaNO}_3 + (\text{NH}_4)_2\text{SO}_4 \Rightarrow \text{Na}_2\text{SO}_4 + 2\text{N}_2\text{O}+4\text{H}_2\text{O}$$

Sodium nitrate:$\text{NaNO}_3$

$\begin{array}{|c|c} \hline \text{Molar Mass} & 84.9947 \text{g/mol} & \\ \text{Density} & 2.257 \text{g/L (solid)} \\ \text{Melting Point} & 308 \text{°C} \\ \text{Boiling point} & 380 \text{°C} \\ \text{Price} & \text{USD ~} 100 \text{-} 300/\text{ton} \\ \hline \end{array}$

Sodium nitrate can be harvest from Caliche rocks (I don't think it's avariable on asteroids.
Additionally it can be syntethised by a lot of ways:

Sodium nitrate is synthesized industrially by neutralizing nitric acid with sodium carbonate:

$$2\text{HNO}_3 + \text{Na}_2\text{CO}_3 \Rightarrow 2 \text{NaNO}_3 + \text{H}_2\text{O} + \text{CO}_2$$

Or with sodium bicarbonate:

$$\text{HNO}_3 + \text{NaHCO}_3 \Rightarrow \text{NaNO}_3 + \text{H}_2\text{O} + \text{CO}_2$$

Or even with sodium hydroxide (very exothermic) (Also note that sodium hydroixe is made with Chloralkali process which need sea water... or with sodium carbonate and calcium hydroxide making it pointless):

$$2\text{HNO}_3 + \text{NaOH} \Rightarrow \text{NaNO}_3 + \text{H}_2\text{O}$$

Or just sodium reacting with nitric acid (violent reaction):

$$2\text{HNO}_3 + 2 \text{Na} \Rightarrow 2\text{NaNO}_3 + \text{H}_2$$

Sodium carbonate: $\text{Na}_2\text{CO}_3$

$\begin{array}{|c|c} \hline \text{Molar Mass} & 105.9888 \text{g/mol} & \\ \text{Density} & 2.54 \text{g/L (solid)} \\ \text{Melting Point} & 851 \text{°C} \\ \text{Price} & \text{USD ~} 100 \text{-} 280/\text{ton} \\ \hline \end{array}$

This can be hasvert from Natron, again a rock.

Sodium bicarbonate: $\text{NaHCO}_3$

$\begin{array}{|c|c} \hline \text{Molar Mass} & 84.0066 \text{g/mol} & \\ \text{Density} & 2.20 \text{g/L (solid)} \\ \text{Melting Point} & 50 \text{°C} \\ \text{Price} & \text{USD ~} 200 \text{-} 300/\text{ton} \\ \hline \end{array}$

It can be harvest from deposits of nahcolite carbonate mineral.

Sadly it's produced:

  • Using the Solvay process which use sodium chloride, ammonia and carbon dioxide making a circular reference.
  • Also it can me made with carbon dioxide and sodium hydroxide, again making a circular reference.
  • But, it can also be made with the ore trona dissolved in water and treated with carbon dioxide.

    $$\text{Na}_2\text{CO}_3 + \text{CO}_2 + \text{H}_2\text{O} \Rightarrow 2 \text{NaHCO}_3$$

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  • $\begingroup$ Basically the same as HKOB's answer...only a bajillion times more detail. $\endgroup$ – Draco18s Jun 21 '18 at 1:33
  • $\begingroup$ @Draco18s, mmm yes. I like detailled stuff :). When I have more time I'll post another ideas with also detailled description! It was very entertaining do this post! $\endgroup$ – Ender Look Jun 21 '18 at 1:43
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This will be my second post to this quesiton. My first one was this one. Here I'll put more nitrogen and oxygen derivates.

Dinitrogen Pentoxide $\text{N}_2\text{O}_5$ (AKA: Nitrogen Pentoxide)

$\begin{array}{|c|c} \hline \text{Molar Mass} & 108.01 \text{g/mol} & \\ \text{Density} & 1.642 \text{g/L (gas)} \\ \text{Melting Point} & 41 \text{°C} \\ \text{Boiling point} & 47 \text{°C} \\ \text{Price} & \text{¿?} \\ \hline \end{array}$

Introduction: Also known as nitrogen pentoxide, N2O5 is one of the binary nitrogen oxides, a family of compounds that only contain nitrogen and oxygen. It is an unstable and potentially dangerous oxidizer.

Storage: Despite its unstability it can be safestly stored a 8 °C and hide from sunlight. Also, at room temperature it's solid, which means it doesn't need pressurised tanks.

Usage: Dinitrogen Petoxide exists as colourless crystals that sublime slightly above room temperature. The salt eventually decomposes at room temperature into $\text{NO}_2$ and $\text{O}_2$.

$$2\text{N}_2\text{O}_5 \Rightarrow 4\text{NO}_2 + \text{O}_2$$ $\begin{array}{|c|c|c|c|} \hline \text{Examples} & \text{N}_2\text{O}_5 & \text{NO}_2 & \text{O}_2 \\ \hline \text{by Mass} & 1,000\text{g} & 851.87\text{g} \text{ }(85.18\text{%}) & 148.12\text{g} \text{ }(14.81\text{%})\\ \text{by Volume} & 1,000\text{cm}^3 & 744.03\text{cm} \text{ }(74.40\text{%}) & 170.20\text{cm}^3 \text{ }(17.02\text{%})\\ \hline \end{array}$

And nitrogen dioxide can be also broken for more oxygen. (Note that if you didn't want the O2 from the first reaction you could use instead dinitrogen tetroxide). At 150 °C, $\text{NO}_2$ decomposes with release of oxygen via an endothermic process:

$$2\text{NO}_2 + 14\text{ kJ/mol} \Rightarrow 2\text{NO} + \text{O}_2$$

And nitrogen oxide can be also broken for even more oxygen. Since the heat of formation of ·NO is endothermic, NO can be decomposed to the elements exothermically. Catalytic converters in cars exploit this reaction. See my first post for a list of materials which can be used as catalysers.

$$2\text{NO} \Rightarrow \text{O}_2 + \text{N}_2 + 91.29\text{ kJ/mol}$$

So the overal reaction would be:

$$2\text{N}_2\text{O}_5 + 14\text{ kJ/mol} \Rightarrow 5\text{O}_2 + 2\text{N}_2 + 91.29 \text{ kJ/mol} \text{ (Net: }77.29\text{ kJ/mol)}$$ $\begin{array}{|c|c|c|c|} \hline \text{Examples} & \text{N}_2\text{O}_5 & \text{O}_2 & \text{N}_2 & \text{Energy}\\ \hline \text{by Mass} & 1,000\text{g} & 740.64\text{g} \text{ }(74.06\text{%}) & 259.35\text{g} \text{ }(25.93\text{%}) & 715.58\text{ kJ}\\ \text{by Volume} & 1,000\text{cm}^3 & 851.04\text{cm} \text{ }(85.10\text{%}) & 340.53\text{cm}^3 \text{ }(34.05\text{%}) & 1,174.98\text{ kJ}\\ \hline \end{array}$

Which all this stuff we can build an atmosphere of:

$\begin{array}{|cc|r|cc|r|} \hline \text{Chemical} & \text{gr/mol} & \text{Percentage} & \text{Mol Fractal} & \text{Mol} & \text{Partial Pressure} \\ \hline \text{N}_2 & 28.0134 & 25.94\text{%} & 0.925 & 0.285 & 28.950 \text{ kPa} \\ \text{O}_2 & 31.9988 & 74.06\text{%} & 2.314 & 0.714 & 72.325 \text{ kPa} \\ \hline \text{Total} & 60.0122 & 100.00\text{%} & 3.239 & 1.000 & 101.325* \text{ kPa} \\ \hline \end{array}$
101.325 kPa = 1 atm

  • Lethal oxygen value: Humans need around 21 kPa of oxygen to "work" properly, you will have 72.325 kPa, your people will suffer hyperoxia. Hyperoxia is produced when oxygen is above 50 kPa it becomes toxic. You won't survive mora than 1 or 2 days (maybe less). In this answer I talk of hyperoxia symptoms (also, there is a cute image).
    • You can fix that reducing the overall pressure from $101.325\text{ kPa} \rightarrow 29.4\text{ kPa}$. So you will have $8\text{ kPa N}_2$ and $21\text{ kPa O}_2$. Good!
      Also, reduce the overall pressure has an advantage: if there is a damage in the ship/station's hull, air will scape slowler from it, giving more time from crew to fix it.

Another path to proxide oxygen is mixing dinitrogen pentoxide with water to produce nitric acid. But there is no point doing this as explained bellow. $$\text{N}_2\text{O}_5 + \text{H}_2\text{O} \Rightarrow 2\text{HNO}_3$$ With sunlight and heat you could produce nitrogen dioxide and water from nitric acid, the same sustance used in the first path to get oxygen. $$4\text{HNO}_3 \Rightarrow 2\text{H}_2\text{O} + 4\text{NO}_2 + \text{O}_2$$ The water works a catalyst because it's recovered: $$2\text{N}_2\text{O}_5 + 2\text{H}_2\text{O} \Rightarrow 2\text{H}_2\text{O} + 4\text{NO}_2 + \text{O}_2$$ This path is worthless because it produce the same as the path before, but I had to show you.

Dinitrogen Tetraoxide $\text{N}_2\text{O}_4$

$\begin{array}{|c|c} \hline \text{Molar Mass} & 92.011 \text{g/mol} & \\ \text{Density} & 1.44246 \text{g/L (gas)} \\ \text{Melting Point} & -11.2 \text{°C} \\ \text{Boiling point} & 21.69 \text{°C} \\ \text{Price} & \text{¿?} \\ \hline \end{array}$

$$\text{N}_2\text{O}_4 + 57.23 \text{ kJ/mol} \rightleftharpoons 2\text{NO}_2$$ Overall: $$\text{N}_2\text{O}_4 + 71,23\text{ kJ/mol} \Rightarrow 2\text{O}_2 + \text{N}_2 + 91.26\text{ kJ/mol} \text{ (Net: }20.03\text{ kJ/mol)}$$

$\begin{array}{|c|c|c|c|} \hline \text{Examples} & \text{N}_2\text{O}_4 & \text{O}_2 & \text{N}_2 & \text{Energy}\\ \hline \text{by Mass} & 1,000\text{g} & 695.54\text{g} \text{ }(69.55\text{%}) & 304.45\text{g} \text{ }(30.44\text{%}) & 217.69\text{ kJ}\\ \text{by Volume} & 1,000\text{cm}^3 & 702.09\text{cm} \text{ }(70.20\text{%}) & 351.16\text{cm}^3 \text{ }(35.11\text{%}) & 314.01\text{ kJ}\\ \hline \end{array}$

Which produce an atmosphere of:

$\begin{array}{|cc|r|cc|r|} \hline \text{Chemical} & \text{gr/mol} & \text{Percentage} & \text{Mol Fractal} & \text{Mol} & \text{Partial Pressure} \\ \hline \text{N}_2 & 28.0134 & 30.45\text{%} & 1.086 & 0.666 & 33.775 \text{ kPa} \\ \text{O}_2 & 31.9988 & 69.55\text{%} & 2.173 & 0.333 & 67.550 \text{ kPa} \\ \hline \text{Total} & 60.0122 & 100.00\text{%} & 3.259 & 1.000 & 101.325 \text{ kPa} \\ \hline \end{array}$

  • Keeps being lethal due high oxygen value:
    • You can fix that reducing the overall pressure from $101.325\text{ kPa} \rightarrow 31\text{ kPa}$. So you will have $10.5\text{ kPa N}_2$ and $21\text{ kPa O}_2$. Good!

Also, dinitrogen tetraoxide has serveral unuseful paths where water is catalizer (in one of them even oxygen is a catalizer): $$\text{N}_2\text{O}_4 + \text{H}_2\text{O} \Rightarrow \text{HNO}_2 + \text{HNO}_3$$ Subpaths:
• $2\text{HNO}_2 \Rightarrow \text{NO}_2 + \text{NO} + \text{H}_2\text{O}$
• $3\text{HNO}_2 \Rightarrow \text{HNO}_3 + 2\text{NO} + \text{H}_2\text{O}$
• $2\text{HNO}_2 + \text{O}_2 \Rightarrow 2\text{HNO}_3$

The last two path also have: $4\text{HNO}_3 \Rightarrow 2\text{H}_2\text{O} + 4\text{NO}_2 + \text{O}_2$.
Then, $\text{NO}_2$ and $\text{NO}$ can be broken as seen above on the post.

Nitrogen dioxide $\text{NO}_2$ And Nitrogen oxide $\text{NO}$

You've already learn how to break them above!

Also the $\text{NO}_2$ oxygen and nitrogen production, and even the atmospheric composition are exactly the same as with dinitrogen tetraoxide. The only difference is that it doesn't need energy to break down.

$\begin{array}{|c|cc|} \hline \text{} & \text{NO}_2 & \text{NO}\\ \hline \text{Molar Mass} & 46.0055 \text{g/mol} & 30.01 \text{g/mol}\\ \text{Density} & 1.88 \text{g/L (gas)} & 1.3402 \text{g/L (gas)}\\ \text{Melting Point} & -11.2 \text{°C} & -164 \text{°C}\\ \text{Boiling point} & 21.69 \text{°C} & -152 \text{°C}\\ \text{Price} & \text{USD}7-8\text{/kg} & \text{USD}1.48-2.95\text{/kg}\\ \hline \end{array}$

But with nitrogen oxide:

$\begin{array}{|c|c|c|c|} \hline \text{Examples} & \text{NO}_2 & \text{O}_2 & \text{N}_2 \\ \hline \text{by Mass} & 1,000\text{g} & 533.13\text{g} \text{ }(53.31\text{%}) & 466.73\text{g} \text{ }(46.67\text{%})\\ \text{by Volume} & 1,000\text{cm}^3 & 500.00\text{cm}^3 \text{ }(50\text{%}) & 500.17\text{cm}^3 \text{ }(50.01\text{%})\\ \hline \end{array}$

$\begin{array}{|cc|r|cc|r|} \hline \text{Chemical} & \text{gr/mol} & \text{Percentage} & \text{Mol Fractal} & \text{Mol} & \text{Partial Pressure} \\ \hline \text{N}_2 & 28.0134 & 46.68\text{%} & 1.666 & 0.5 & 50.663 \text{ kPa} \\ \text{O}_2 & 31.9988 & 53.32\text{%} & 1.666 & 0.5 & 50.663 \text{ kPa} \\ \hline \text{Total} & 60.0122 & 100.00\text{%} & 3.32 & 1.0 & 101.325 \text{ kPa} \\ \hline \end{array}$

  • Keeps being lethal due high oxygen value:
    • You can fix that reducing the overall pressure from $101.325\text{ kPa} \rightarrow 31\text{ kPa}$. So you will have $21\text{ kPa N}_2$ and $21\text{ kPa O}_2$. Good!

Dinitrogen Trioxide $\text{N}_2\text{O}_3$

$\begin{array}{|c|c} \hline \text{Molar Mass} & 76.01 \text{g/mol} & \\ \text{Density} & 1.447 \text{g/L (liquid)} \\ \text{"} & 1.783 \text{g/L (gas)} \\ \text{Melting Point} & -107 \text{°C} \\ \text{Boiling point} & 3.5 \text{°C} \\ \text{Price} & \text{¿?} \\ \hline \end{array}$

It forms upon mixing equal parts of nitric oxide and nitrogen dioxide and cooling the mixture below −21 °C. This process is reversible so it produce that:

$$\text{N}_2\text{O}_3 \rightleftharpoons \text{NO} + \text{NO}_2$$

Using the process learned above:

$$2\text{N}_2\text{O}_3 + 14\text{ kJ/mol} \Rightarrow 3\text{O}_2 + 2\text{N}_2 + 91.29\text{ kJ/mol} \text{ (Net: }77.29\text{ kJ/mol)}$$ $\begin{array}{|c|c|c|c|} \hline \text{Examples} & \text{N}_2\text{O}_5 & \text{NO}_2 & \text{O}_2 \\ \hline \text{by Mass} & 1,000\text{g} & 631.47\text{g} \text{ }(63.14\text{%}) & 368.54\text{g} \text{ }(36.85\text{%})\\ \text{by Volume} & 1,000\text{cm}^3\text{gas} & 639.42\text{cm} \text{ }(63.94\text{%}) & 429.42\text{cm}^3 \text{ }(42.94\text{%})\\ \text{by Volume} & 1,000\text{cm}^3\text{liquid} & 787.90\text{cm} \text{ }(78.79\text{%}) & 525.44\text{cm}^3 \text{ }(52.54\text{%})\\ \hline \end{array}$

Which produces this atmosphere:

$\begin{array}{|cc|r|cc|r|} \hline \text{Chemical} & \text{gr/mol} & \text{Percentage} & \text{Mol Fractal} & \text{Mol} & \text{Partial Pressure} \\ \hline \text{N}_2 & 28.0134 & 36.85\text{%} & 1.315 & 0.4 & 40.530 \text{ kPa} \\ \text{O}_2 & 31.9988 & 63.15\text{%} & 1.973 & 0.6 & 60.795 \text{ kPa} \\ \hline \text{Total} & 60.0122 & 100.00\text{%} & 3.288 & 1.0 & 101.325 \text{ kPa} \\ \hline \end{array}$

  • Keeps being lethal due high oxygen value:
    • You can fix that reducing the overall pressure from $101.325\text{ kPa} \rightarrow 35\text{ kPa}$. So you will have $14\text{ kPa N}_2$ and $21\text{ kPa O}_2$. Good!

Production:

Sadly, the only way to produce all this usances use the same compounds that they produce or with other very difficult to get, that means a circular reference. So you must bring it from Earth.

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I am kind a fan of a low presure pure oxygen environment which already mentioned.

So you may use solidified oxygen. In the space there is no need to active cooling. You may keep it with simple insulation and a reflective wrapping from radiant heat sources (like sun) and from the rest of your spaceship.

In fact in the failed Apollo 13 misson there is a malfunctioning "heater" that keeps oxygen from solidification.

Or

You may use hydrogen peroxide H2O2. It's a liquid under the room temprature.

You may even transport it in the solid form (it's melting point is very close to water-ice).

For safety reasons it is usually mixed with water.

It's also used as a propellant in the rockets and the jetpacks.

With the silver catalyst, H2O2 very rapidly breaks up (bursts) to water vapour and a oxygen.

Personal Jetpacks used this reaction for a propulsion.

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