1
$\begingroup$

I was just trying to make a stock solution of some Co(NO3)2 . One in H2O, another in Methanol. Both don't completely dissolve. They're almost dissolved and appear near translucent but when shining a laser pointer through, I can see the path of light (indicating it's a colloid).

I'm just making a 1M solution in 50 mL test tubes.

Anyone have any experience with this? Shouldn't this salt be completely soluble? Perhaps i'm dissolving a lot within this 50 mL volume and it's forcing some particles to slightly aggregate?

Thank you!

P.S: The one in H2O I made was in early December. Every week or so a bit of precipitate collects at the top and when I shake it, it dissipates. Sort of like a sediment but on top instead of bottom. I assume this is just aggregated Co(NO3)2 and i'm just redissolving it. The one in Methanol, I just prepared today but it's exhibiting similar properties and I haven't observed this precipitate in the methanol yet given it's been less than a day.

Improve this question
$\endgroup$

1 Answer 1

Reset to default
2
$\begingroup$

A first thought would be impurities in $\ce{Co(NO3)2}$. There are varying grades of purity of laboratory chemicals, ranging from technical (just good enough for some crude uses) to American Chemical Society (ACS)... but even ACS grade reagents might have 5% impurities! 99.999% ("five nines") pure $\ce{Co(NO3)2}$ is available, if the use justifies the price.

Another possibility is some contaminant in the labware (or even in the lab air) is reacting with the $\ce{Co(NO3)2}$. For example, if a glass vessel were to be cleaned in an alkaline detergent, and not sufficiently washed (perhaps using an acid rinse), adherent alkali might precipitate an insoluble cobalt compound. You could test for this by rinsing all labware used to measure and to hold that stock solution with a small amount of nitric acid, and then rinsing again with highest purity water available. Use that highest quality water to make the stock solution.

Also consider the quality of glassware -- use good grade borosilicate, not soda-lime glassware.

Improve this answer
$\endgroup$
3
  • $\begingroup$ i do use alkali to clean. i really hope it's not leaving some residue causing premature precipitation. wouldn't the pH need to be at a certain point though for the Co 2+ to precipitate as a Co(OH)2, or atleast have the process be kinetically/thermodynamically favourable? i feel like the tiny amount of residue left won't be enough to swing the pH that drastically, as i do rinse out the glassware thoroughly with water afterwards. $\endgroup$
    – Ahmer Imam
    Commented Jan 10 at 2:04
  • $\begingroup$ One uses the experimental method to verify a hypothesis. Make up some more solution, following the suggested procedure for acid rinse and water rinse. $\endgroup$
    – DrMoishe Pippik
    Commented Jan 10 at 3:06
  • $\begingroup$ For water solution, does the behaviour differ if kept acidic by addition of small amount of nitric acid? Co(II) salts should be stable, but if there are Fe salt traces, it can be hydrated Fe(III) oxide. $\endgroup$
    – Poutnik
    Commented Jan 10 at 4:22

Not the answer you're looking for? Browse other questions tagged or ask your own question.