Separating Cobalt From Nickel (Cobalt From Li-ion Batteries - Part 4) 

These annoying advertisings of the same shitty internet sites (brilliant, squarespace, nordvpn, audible, jlcpcb or yfood) really suck. Big time. I'll start giving negative votings for this crap. I miss the times when RU-vid videos weren't a business for everyone. It's just annoying.

That's awesome. Glad you enjoyed it!

I’ve always loved organic chemistry without even stopping to wonder what inorganic chem really was

Hmmm... interesting. While that would probably work for this particular instance, I'm thinking there are a couple of issues with that kind of separation that I don't really like: 1) When atomically mixed, the cobalt in the cobalt/nickel salts might have a hard time dissolving properly, as the nickel can kind of 'protect' it from touching the solvent (similar to the reasons for requiring inquartation for dissolving gold alloys). With such a low fraction of nickel, this is probably not much of an issue here though. 2) I'm also worried that even small amounts of water (even leftover hydration of the chlorides) could impact the process and bring a bit of nickel into solution. Honestly, I'm not very sure about either of these points though. Do you have any further insight to an acetone separation like this?

I'm going to try that myself. Thanks for the tip!

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I love how your comments just get more ridiculous and nonsensical every time. lol

This is an open source and access to all platform, not like your industrial robber barons.

The cathode of a lithium ion battery, in this day and age, does not contain cobalt alone. Generally, it contains a mixed oxide of cobalt, nickel, and magnanese, with a certain amount of lithium ions intercalated depending on the state of charge. Normally, this is all deposited on an aluminium substrate. As such, dissolving up cathodes will give you a mixture of all these metals (Co, Ni, Mn, Al, and Li) in solution. This is exactly what we saw in this series of videos. If using hot acetic acid is all you've done, I can guarantee that you don't have a pure solution of cobalt. What I show in these videos is the amount of work that would be required to actually separate the components of your product.

The only trouble there (aside from the difficulty of that distillation) is that the lithium and aluminium would have to be in the form of metals to do that, and turning them into their metallic forms is much harder than separating them chemically, haha. (Also note that in the batteries themselves, lithium isn't actually in its metallic form either). When in solution, it's actually pretty straightforward to precipitate the aluminum as the hydroxide, leaving the lithium behind (though I suppose I actually tried that in the second episode and it didn't work, hmmm...)

@@ScrapScience Just thought that the H2 in excess would do the job of converting lithium to metal for you.

Hey at least with that much cobalt you get a life time supply of B12 vitamin

As interesting as it would be, lead dioxide and lead-based electrodes are something I'd much prefer to steer well clear of these days. The toxicity is simply far too much for me to work with for my current setup.

I'll have to look around and try to find someone who will run a sample for me or something, haha.

@@ScrapScience You can send in to me, I'm in Ukraine, lmao

Haha. I'm afraid the process is very far away from being economically viable at this stage. For starters, the cost of electricity for electrowinning the transition metals is already more expensive than the metals themselves, and that doesn't even take into account the costs of the reagents I've used so far, including sulfuric acid, hydrochloric acid, calcium hydroxide, butanone, ammonium persulfate, sodium thiocyanate, and sodium/potassium carbonate. If you can somehow make the electrowinning step more efficient, find some way of effectively recovering the butanone solvent, and get your batteries for free, it could conceivably be useful, but you'd be putting a lot of work into it for a very small profit if so.

@@ScrapScience I wouldn't even attempt such an undertaking - way over my pay grade, as they say. No, I was asking if this could be close to commercially viable at scale. I don't think the cost of electricity should come into it - just the maintenance and outgoings of a suitable PV array. As for the reagents - can these possibly be recovered using cheap PV electricity?

Yeah I can definitely see that if you can get free electricity from photovoltaics it might have a better shot. With some development you might be able to recover the butanone from the solvent extraction pretty well, but recovering any of the other reagents (acid, hydroxides, carbonates, thiocyanates) is going to be near impossible since they'll all be used up in the reactions they're involved in. If a business could get free battery waste and free electricity, I could see the possibility of this being viable on a large scale. Still, a lot of research these days is now going towards new efficient processes for separating cobalt and nickel, because this solvent extraction process really doesn't seem to play nicely on large scales (I think).

@@MiniLuv-1984 solar is orders of magnitude too diffuse for such a thing. There is, however, a way to use the light of previous suns (well, colliding neutron stars, at any rate) in the form of uranium and similar chonky bois. With that, of which Australia has plenty under the ground and vastly more in the ocean that surrounds it, you do not have to worry about destroying a bunch of delicate ecosystem by shading away the light it depends on with square kilometers of PV panels.

@@davidfetter Well that may be a way forward, but solar panels are proven and here now. A fission plants is at least 10 years or more before we see an energetic electron from it and fusion is always going to be 20 years away. I hope my negative portrayal of fusion is misplaced however. Even if they cracked it tomorrow, it will still be decades before it delivers any power. As for fission, dirty, limited resource fossil fuels got us into this shit to begin with, so I really don't want to see any fusion power plants doing the same thing.

When I was making this video, it wasn't an issue at all - there were no glassware laws in my location. However, I recently moved to a different state and the laws are a bit different. I'm still looking into the specifics, but I basically need to make my channel a registered business and sign some forms to say I won't be making drugs if I ever want to own any glassware with ground glass joints over here.

@@ScrapScience tyvm. Basically the US DEA compiled a list of laboratory equipment that can conceivably be used for manufacturing "illegal" drugs. So just about anything ends up on a list. About ten years ago I had to buy some very basic equipment in order to prepare metal etchants for determining the microstructure. Basically, we had to register the company with the DEA, list the equipment that we would be purchasing, describe what it would be used for and where it would be located and utilized.

One day, yes! I don’t have current plans for it, and it might not be for a very long while, but I want to give this a go one day.

That might be an idea to visit at some point. What kind of dessicant do you think would work? I'm under the impression that chlorides are a bad idea because I wouldn't really want to be making chlorates or chlorine in the electrolysis, and hydroxides will be a little wasteful since they'll start absorbing CO2 and turning into carbonates.

@@ScrapScience what about pure sulpheric acid? It's a powerful hygroscopic electrolyte. Maybe you can easilly make a Sodium silicate bonded fiberglass mesh to provide a high surface area substrate over which you may blow air. Apparently melamine sponges worked for about a weak for researchers, who settled on sintered glass foam and platinum electrodes.

Hmm, sulfuric acid is definitely a good plan. Do you know what concentration the sulfuric acid will dilute to due to its hygroscopic nature? I'm under the impression that high concentrations of sulfuric acid aren't very conductive and ideally we need it to dilute to a concentration of less than 70% (I'm guessing) to make electrolysis proceed nicely.

@@ScrapScience a few things that may become apparent on reflection. The acid concentration will depend on an equilibrium that depends on surface area, mass of air flow, humidity, temperature and the rate water is electrolized depending on many of the same factors as well as the voltage applied and the distance current must flow from one plate to the other. As you pointed out, sulpheric acid is not a great conductor at high concentrations, but as the humidity increases, more water dilutes the sulpheric acid which in turn conducts more current. And The applied voltage effects the current flow. The researches point out this works down to around 4% relative humidity. And a mention that over the concentrations of their work the electrolyte had a -30 something Celsius freezing point sugesting it may work in very cold conditions. There is an article over on nature that graphes the researchers results testing a glass sintered foam with a 15 miliamps cm² electric current with a near linear relationship between concentration and relative humidity at 25c, that ranges from 62% concentration at 20% relative humidity to what appears to be 26.8% concentration at 80% humidity.

@@ScrapScience I'll send the link to you in an email. 😁

The channel has some videos on Sodium (Na) and Calcium (Ca). Here are they: Sodium: ru-vid.com/video/%D0%B2%D0%B8%D0%B4%D0%B5%D0%BE-wb65RZh1ptU.html Calcium: ru-vid.com/video/%D0%B2%D0%B8%D0%B4%D0%B5%D0%BE-L7X_v4dZEQw.html

I'll be making magnesium metal in an upcoming video. It's all filmed and everything, but there'll be a wait since there's quite a backlog of videos to get through first. A seawater extraction is definitely not a bad idea, and it's actually something I was considering doing for the magnesium, haha. We'll get there one day hopefully.

Glad you like the separation! I've always been somewhat keen to try making hydrochloric acid with the chlorine and hydrogen from a chloralkali cell, but it's a dangerous enough experiment that I don't have any current plans on doing it in the near future. It's definitely something that will always be in the back of my mind though, so I'll likely get to it one day. Using the chlorine for chlorinations is a little beyond my expertise, since I'm not particularly comfortable with organic reactions, and there are much better ways to make chlorine gas for something like that. Either way, I'm hoping to try hydrochloric acid production at some point, but I can't make any promises at this stage.

I don't think these particular tetrahedral complexes can be crystallised, at least as far as I'm aware. Maybe with the right counter-ion, it might be possible.

Yeah, it's definitely a bit weird. Pretty much everyone else I've seen do this kind of extraction has had nickel be the most abundant species in solution after dissolving up the materials. It seems like I've either got batteries which contain an unusually low fraction of cobalt, or maybe my heat treatment and dissolution process is somehow selective for dissolving cobalt over nickel. The latter is unlikely though, since I found that the acid seemed to dissolve almost all of the electrode mass I started with.

@@ScrapScience I guess you meant to say unusually high fraction of Cobalt, but that would also be highly unlikely unless it was a custom made ultra high end Li-ion battery with a price tag I even don't want to speculate. I'm working with Li-ion battery manufacturing and have been in the business for 15 years and I never heard of ratios like that =)

Oh yeah whoops - unusually high is what I meant there. These are mostly just batteries out of old iPhones so I doubt they're anything special. Seems like a bit of a mystery. If I ever figure out what's going on, I'll let you know lol.

brilliant

lithium is hiding from you lulz

electrolytic Silicon metal from SiCl4 route (SiO2 + C -> Si +CO2, then purify the silicon through molten SiCl4 electrorefining route)

Also, where did you find those solubility statistics of the coordination complexes in organic solvents? I have personally never encountered them.

I'd imagine that it's likely for other organic solvents to work, but it's definitely something that would require some testing. I only used butanone because the document I'm following said that it works. I'm afraid I don't have much in terms of solubility tables for inorganic salts/complexes in organic solvents either (there's too much information there to compile neatly, I'd imagine). While limited, old solubility dictionaries like this one can be useful (though they are limited): archive.org/details/dictionaryofchem00comeuoft/

@@ScrapScience Thanks a lot for your reply! I wonder if it is possible to separate ferric ions with cobalt ions by this way (since they both form the thiocyanate complexes). Will the ferric complex dissolve into the organic layer?

I'm fairly certain it predates the discovery of fission. They were using similar methods to separate the lanthanides in the 1800's.

Nuclear nerd fren

Did you watch the video?

yes

@@NitinKumar-ez3cc It's butanone for this particular separation. I talk about it at 7:40 and 11:09, among other places in the video.