Student showed me a protocol they got online from thermofischer for DNA precipitation (https://www.thermofisher.com/uk/en/home/references/protocols/nucleic-acid-purification-and-analysis/dna-extraction-protocols/phenol-chloroform-extraction.html)
and for the actual precipitation in 100% etOH it states to precipitate overnight at -20C, (that's fine, that's what i do for the best yield) but beside this it states for an accelerated protocol store the etOH sample mixture at -70C for 1 hour.... It was my understanding that for a quicker precipitation you incubate at room tempature and that leaving it at colder temperature slows the precipitation so this would have the opposite effect.
It doesn't make much difference to me as i've always down O/N -20 precipiations but I'd like to clear this up.
I have a solution of hydrochloric acid and bleach with dissolved gold flakes from pc scrap, forming chloroauric acid. I am now attempting to precipitate gold powder from solution. Most sources say sodium meta bisulphite is common use and effective. My solution however, seems not to react with the sodium. I have tried using to different sources of sodium, with a response from neither. The sodium simply gathers at the bottom. Do any of you know what I might be doing wrong, or perhaps know of another way to precipitate gold from chloroauric acid?
To make matters worse, after having to pay for the seeds for the crops that didn't grow and the scientist to figure out why my crops wouldn't grow, I can only afford a low-end DRO, and I know I won't have enough left to pay the free-market arbitration firm to find in my favor when we do figure out who to sue. My private security firm also demands upfront payment before they're willing to shoot people about things, and they're not exactly inexpensive.
> You wish to separate silver ions from calcium ions in a solution that is 0.10 M Ag+ and 0.10 M Ca2+. Both form insoluble precipitates with sulfite. Calculate the [H+] and pH at which the solution should be buffered in order to precipitate the maximum amount of one of the ions when the solution is made with 0.10 M H2SO3. Is the separation quantitative?
> (Ksp = 6.8 x 10-8 for CaSO3, Ksp = 1.49 x 10-12 for Ag2SO3)
Im really lost with this solubility material.
I understand what each does, I’m just having trouble telling which kind it is by simply looking at a chemical equation. In many cases, one reaction seems to match a few reaction types. Examples would be great.
We were studying this reaction: KCl (aq) + C4H6O6 (aq) (tartaric acid) = KC4H5O6 (s) + HCl (aq)
After that I asked the teacher about it, but she didn't give it much thought. I wonder: could a weak acid be forced to dissociate completely, to maintain the equilibrium, if we keep precipitating its conjugated base like in the above reaction and, because of that, make it as dissociated as a strong acid?
Hi all, my question is basically as the title says. I'm writing a lab report on a video I had to watch for uni where supposedly hexane was reacted with sulfuric acid however from what I gather theoretically the two shouldn't react at all? I've looked around to see what info I could come up with and watched several other videos of this all showing no reaction, I'm just wondering what it could be that would cause this to happen? Any ideas would be greatly appreciated.
*Salting-out
I understand that the process precipitates proteins by interecting with their charges and making them "less soluble", but how is DNA "protected" from this?
Full text of the question:
If your unknown solution's initial calcium concentration was 0.0003M, how much of the initial calcium ion would be removed from solution by the addition of 25.0mL of 0.01M sodium oxalate? Ksp of (CaOx) is 2.7 x 10^-9 . Assume the pH of the solution is 3.0.
(Also assume that the initial solution volume was 25.0ml, per the professor)
I don't know why this is giving me such a hard time. What's throwing me off is the pH of 3 I think. At that pH, wouldn't hydrogen oxalate form until the Ka Qc reached 10^-4.4 (pKa1 of oxalic acid is 1.46, pka2 is 4.4)? At which point any leftover oxalate could react with the calcium? What am I missing or overthinking here?
It is a follow up question to a lab we did where an unknown calcium solution was precipitated as CaC2O4, in which case we had an overall acidic solution that we added urea to, and then slowly boiled to thermally decompose the urea and release OH^- . As the pH of the solution rose, H2C2O4 was neutralized and freed up the oxalate ion to bond with calcium and then precipitate. For some reason putting it into raw numbers though is giving me hell.
Any help would be appreciated!
Is there any research on this? Any theories? All I coukd find was something about precipitation of salival proteins due to reactions with wine? Anyone know anything more than this? Is this what it is? Why does it happen?
