hodges
National Hazard
Posts: 525
Registered: 17-12-2003
Location: Midwest
Member Is Offline
|
|
Make Cobalt Chloride from Cobalt Metal?
Was going to get some cobalt (II) chloride, but I see I either have to buy 4 ounces of it (far more than I need) for $23 on eBay, or else order a
smaller amount from a hobby or home schooling supply and pay for UPS order shipping. I see 99.98% cobalt metal is available at $7 for 10 grams on
eBay, shipped from US. Even assuming a poor yield, that would be enough to make more than an ounce of cobalt chloride hexahydrate, which is all I
need.
I figure I should be able to react cobalt metal with HCl. Electromotive series shows cobalt at -0.28V, vs. -0.44 for iron. So I'm guessing the
reaction will be pretty slow, but should be doable (perhaps with some gentle heating).
In order to get pure cobalt chloride (uncontaminated with acid), I'm guessing I will need to neutralize the acid solution with sodium hydroxide to get
cobalt hydroxide, then wash this, and carefully add HCl again until the hydroxide just dissolves.
I could not find any videos that show cobalt reacting with acids online. Supposedly cobalt also forms a +3 ion, similar to iron. But I could not
find any information on it. Not sure if I would have a problem with Co (III) forming due to atmospheric oxygen, as happens when dissolving iron in
HCl.
Does my procedure sound reasonable? Thoughts?
|
|
|
DraconicAcid
International Hazard
Posts: 4334
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
I'd just buy the chloride, personally, and keep the rest of it for another experiment. But you don't really have to worry about getting cobalt(III)-
it's not stable without the right ligands. Without them, aqueous cobalt(III) ions will oxidize water to give the more stable cobalt(II) ion.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
woelen
Super Administrator
Posts: 8014
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
The path through the hydroxide is not pleasant at all. At high pH, cobalt is easily oxidized to its +3 oxidation state, even by oxygen from air. Don't
go that route. I would try dissolving the metal in acid and driving off the HCl by heating. You'll get a very dark blue solution/syrup on heating. On
cooling down you'll get a red solution if the concentration of the acid is not too high. I would perform the crystallization of the solid from a
somewhat acidic solution in order to prevent hydrolysis and formation of basic chlorides.
|
|
|
rockyit98
Hazard to Others
Posts: 283
Registered: 12-4-2019
Location: The Known Universe
Member Is Offline
Mood: no mood is a good mood
|
|
one time i made CoCl2 by opening a Li-Ion rechargeable battery and adding HCl. i had to remove Al and Cu ribbon beforehand.i used it at is ,but maybe
you can filter and let CoCl2 crystallize out from LiCl. KI will reduce Co(iii) back to Co(ii).
be careful Co is toxic. follow proper precautions and cleanup .make sure battery is discharge to prevent fire or explosion.
watch this Extractions&Ire video https://www.youtube.com/watch?v=7nwBwp-TwS8
"A mind is a terrible thing to lose"-Meisner
|
|
|
unionised
International Hazard
Posts: 5126
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
The excess HCl is volatile.
Just leaving the solution somewhere warm would drive it off, however the fumes aren't nice.
Leaving it a larger airtight container (like a sandwich box) with an open shallow container of NaOH will strip out water and HCl.
|
|
|
hodges
National Hazard
Posts: 525
Registered: 17-12-2003
Location: Midwest
Member Is Offline
|
|
Thanks for the suggestions. I ordered the cobalt metal.
Per woelen's advice, will not try going through the hydroxide state. Will plan on drying it as unionised suggested by placing in an airtight
container that has some NaOH in it to absorb the water and remaining HCl. Then, once it is completely dry (and free from HCl), will add water and
evaporate in open air to recrystalize it so I get the standard .6H2O. I can stop evaporation at the point right before it starts turning from pink to
blue if the humidity is low.
|
|
|
UC235
National Hazard
Posts: 565
Registered: 28-12-2014
Member Is Offline
Mood: No Mood
|
|
I just left some Co metal from ebay in excess hardware store 31.4% HCl for like a week as it was a solid chunk and reacted very slowly. I boiled it
down somewhat before I ran out of metal and got a good crop of crystals.
The color of the solution is highly chloride and temperature dependent and is not an accurate gauge of anything. It's interesting to watch though. I
believe it was blue while hot and pink when cold. This was years ago.
[Edited on 30-9-2019 by UC235]
|
|
|
woelen
Super Administrator
Posts: 8014
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Quote: Originally posted by hodges  | Thanks for the suggestions. I ordered the cobalt metal.
Per woelen's advice, will not try going through the hydroxide state. Will plan on drying it as unionised suggested by placing in an airtight
container that has some NaOH in it to absorb the water and remaining HCl. Then, once it is completely dry (and free from HCl), will add water and
evaporate in open air to recrystalize it so I get the standard .6H2O. I can stop evaporation at the point right before it starts turning from pink to
blue if the humidity is low. |
I would not recrystallize it from water alone. I would add a tiny amount of acid. Add a single drop of conc. HCl to your solution in order to avoid
formation of basic salts and keeping the solution clear. Most of that HCl will evaporate away and will not get into the crystalline solid. Just be
sure that you do not allow all of the solution to evaporate to dryness. The last 10% of liquid should be decanted from your crystals and the crystals
should be dried between filter paper and then left in a container with NaOH to get rid of the last little amount of water.
The decanted liquid can be kept as solution, but it will be a relatively impure solution of CoCl2 (may contain other metals if your starting material
has some impurities and may contain some left over HCl from your drop of acid).
|
|
|
hodges
National Hazard
Posts: 525
Registered: 17-12-2003
Location: Midwest
Member Is Offline
|
|
| Quote: Originally posted by UC235 |
The color of the solution is highly chloride and temperature dependent and is not an accurate gauge of anything. It's interesting to watch though. I
believe it was blue while hot and pink when cold.
|
Yes. This is one of the experiments I wish to do with the cobalt chloride.
https://edu.rsc.org/resources/the-equilibrium-between-two-co...
The reaction [Co(H2O)6]2+(aq) + 4Cl–(aq) → [CoCl4]2–(aq) + 6H2O(l) is endothermic. Therefore, in accordance with Le Chatelier’s principle,
when the temperature is raised, the position of the equilibrium will move to the right, forming more of the blue complex ion at the expense of the
pink species.
As well as when there are excess Cl- ions. I would expect it to be blue when first formed (assuming use of concentrated HCl).
|
|
|
Arthur Dent
National Hazard
Posts: 553
Registered: 22-10-2010
Member Is Offline
Mood: entropic
|
|
I acquired 500g of Cobalt Carbonate from a pottery shop a while ago, the purpose of that purchase was to produce some Cobalt Chloride. I don't know if
there are major impurities, but the color's right.
I might try to make some this weekend, I have about 250 ml of HCl left so i'll try it out.
CoCO3 + 2 HCl(aq) → CoCl2(aq) + CO2
I imagine the reaction will be energetic so i won't need the hotplate to get the acid hot (thankfully) 
I don't have the equipment to heat it up to its anhydrous form so i'll settle for the hexahydrate.
I'll follow the recommendations above to dry out the liquor using some Calcium Chloride dessicant.
--- Art is making something out of nothing and selling it. - Frank Zappa ---
|
|
|
unionised
International Hazard
Posts: 5126
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
Dissolutions of carbonates in acid are often endothermic.
|
|
|
hodges
National Hazard
Posts: 525
Registered: 17-12-2003
Location: Midwest
Member Is Offline
|
|
Possibly endothermic overall, but I would still think heating will not be required to get a carbonate to react with a strong acid.
|
|
|
fusso
International Hazard
Posts: 1922
Registered: 23-6-2017
Location: 4 ∥ universes ahead of you
Member Is Offline
|
|
Isn't dissolving carbonates in acid exothermic?
|
|
|
hodges
National Hazard
Posts: 525
Registered: 17-12-2003
Location: Midwest
Member Is Offline
|
|
Baking soda reacting with vinegar gets cold. I'm guessing that has to do with with a gas being released (thus doing "work" on its surroundings).
Kind of like letting the air out of a tire causes a drop in temperature. My hunch is that if the gas were not allowed to escape (i.e. due to high
pressure that kept it in solution), it might no longer be endothermic.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
