Where to put the carbon?

You may have seen a few articles lately about a nuisance of a chemical called carbon dioxide, namely, that it's the waste product of a number of very common chemical reactions and doesn't itself react with much (other than plants, but that reaction isn't fast or extensive enough to keep up with our current production rate) meaning it accumulates in the environment.

So, we're trying to make sure less of it gets into the environment. One class of methods which you may have heard of is carbon capture and sequestration, where after production it's captured, compressed, and often pumped deep underground—sometimes into retired oil wells, sometimes into the deep ocean, or many other places.

Before it can be stored, however, it has to be captured. Scrubbing can be highly effective at removing CO₂ from smokestacks and other concentrated sources, traditionally with amine solutions. Then there was a new discovery about the CO₂ absorption of polyethylenimine, which was what caused me to start researching this post.

One of the issues with a reaction that is very effective at grabbing a chemical out of the air is making it let go again. Most of the chemicals that are good at grabbing CO₂ are too expensive to use only once. Polyethylenimine is of great interest because it releases the CO₂ easily by heating it up, which can let the CO₂ be collected in concentrated form for use elsewhere.

Both causing CO₂ release (to a container, of course, not to the atmosphere) and converting it to something useful does take energy, which usually results in more CO₂being released, so while it can help as long as they remove more than they produce, it isn't quite there yet.

More recently yet, a new chemical reaction was discovered that reacts favourably with CO₂ - another potential sink for captured CO₂, and one that produces energy instead of consuming it as well as producing useful products:

\[3\mathrm{CO}_2 (g) + 4\mathrm{Li}_3\mathrm{N} (s) \rightarrow \mathrm{C}_3\mathrm{N}_4 (s) + 6\mathrm{Li_2O} (s)\] \[\mathrm{CO}_2 (g) + 2\mathrm{Li_3N} (s) \rightarrow \mathrm{Li_2CN_2} (s) + 2\mathrm{Li_2O} (s) \]

The tests described ranged from 250C to 400C starting temperature; it doesn't look like any room-temperature tests were reported. But this isn't another case of burning energy to dispose of CO₂—the reaction is strongly exothermic; once started a properly managed continuous reactor could maintain its own ideal temperature, and would only need an external heat source to jump-start the reaction. Assuming, of course, that the 250C temperature is in fact the lowest temperature at which these reactions take place. If it produces excess energy, it might even be able to heat the CO₂ capture solution to release the CO₂, which would be nice.

The reagent, Li₃N is consumed, however, which means it has to be continuously supplied. From the description in the link, it doesn't sound like a fun chemical to work with. ("May ignite on contact with water or moist air"? While it's fun to make stuff burn or blow up, trying to work with something that can spontaneously ignite is a bit nerve-wracking.)

Li₃N itself is easily made from metallic Li and N₂ gas. However, metallic lithium has to be purified, which takes energy. It'll be an interesting exercise to see if the cost of creating lithium, then Li₃N, is low enough to make this commercially feasible.