Sciencemadness Discussion Board - Making nitric acid

Making nitric acid

Silver lord - 6-12-2013 at 03:25

Need feedback from you experts:

I need (weak) nitric acid for refining silver.

My Idea is:

Mix Calcium nitrate with Sodium hydrogen sulphate in distilled water:

Ca(NO3)2+NaHSO4 ->CaSO4+HNO3+NaNO3

Let it settle and pour off HNO3 + NaNO3

Add Sulphuric acid to the mixture:

2HNO3 + 2NaNO3 + H2SO4(35%) -> Na2SO4 + 4HNO3

Freeze until settled and then pour off Nitric acid (~50% ?)

The tricky part I think, is to know how much sulphuric acid I need.

Will it work? Do I have to add heat somewhere?

Pulverulescent - 6-12-2013 at 03:43

Quote:

The tricky part I think, is to know how much sulphuric acid I need.

Can't your pharmacist get you the dilute HNO₃ you need?

Silver lord - 6-12-2013 at 03:49

In Sweden nitric acid is highly regulated.

You have to be a registred company with a defined need to purchase...

Pulverulescent - 6-12-2013 at 04:30

'Sorry, thought you were in the UK . . .

Omit the sodium salt?
Dissolve the calcium salt in water, add (stoichiometric?) quantity of H₂SO₄, cool and decant HNO₃ solution from precipitated CaSO₄.
The obtained impure acid can be distilled if needed . . .

[Edited on 6-12-2013 by Pulverulescent]

Silver lord - 6-12-2013 at 05:08

Wanted to do it in two steps to save sulphuric acid.

Sodium hydrogen sulphate is easier to obtain for me.

But I am not sure it works...

TheChemiKid - 6-12-2013 at 05:51

Check out this video.
If you cannot perform a distillation, then just dissolve the bisulfate in a minimum amount of water, then add to the nitrate. You can cool down the final solution to push out as much sulfate as possible.

[Edited on 6-12-2013 by TheChemiKid]

metalresearcher - 6-12-2013 at 13:56

But now the question arises : how can one get Ca(NO3)2 ? Is it available in house / garden stuff?

Silver lord - 7-12-2013 at 04:40

Quote: Originally posted by Pulverulescent

'Sorry, thought you were in the UK . . .

Omit the sodium salt?
Dissolve the calcium salt in water, add (stoichiometric?) quantity of H₂SO₄, cool and decant HNO₃ solution from precipitated CaSO₄.
The obtained impure acid can be distilled if needed . . .

[Edited on 6-12-2013 by Pulverulescent]

Thanks for answer.

CaSO4 is almost insoluble in water. NaNO3 on the other hand has very good solubiliy.
Some of the salt will probably stay with the calcium sulfate, but most I think will be in solution with the nitric acid -or?

Yes I can add sulphuric acid directly to the calcium nitrate, but the idea was to save sulphuric acid...

To get the first reaction running, do I need to add heat?

Silver lord - 7-12-2013 at 04:45

Quote: Originally posted by metalresearcher

But now the question arises : how can one get Ca(NO3)2 ? Is it available in house / garden stuff?

In Sweden calcium nitrate is available as fertilizer.

bfesser - 7-12-2013 at 06:15

Search before posting and don't ask for spoon feeding. This has been answered on SM and elsewhere on the internet countless times.

To start you off, here's a link to the last time this was discussed on SM—only two weeks ago:
<a href="viewthread.php?tid=27492">Is there a way to dissolve silver without HNO3?</a>

MrHomeScientist - 9-12-2013 at 10:38

Not to encourage spoonfeeding, but I felt like doing some stoichiometry with the two proposed reaction paths. These were:

1) Ca(NO₃)₂ + NaHSO₄ --> CaSO₄ + HNO₃ + NaNO₃
2HNO₃ + 2NaNO₃ + H₂SO₄(35%) --> Na₂SO₄ + 4HNO₃

2) Ca(NO₃)₂ + H₂SO₄ --> CaSO₄ + 2HNO₃

It turns out you would indeed save sulfuric acid with (1), using only half as much as in (2). This comes at the cost of buying sodium bifulfate, as well as the extra time executing the 2 step reaction and chilling to crystallize sodium sulfate. Method (2) has the advantage of being very simple, requiring only a filtration step, as well as yielding highly pure nitric acid. Calcium sulfate is very insoluble so there will be negligible contamination as a result (not counting impurities in your starting materials of course). By contrast, the sodium sulfate at the end of (1) is still somewhat soluble at 0 C (5g / 100g water), so you will almost certainly have some sodium and sulfate ion contamination. It's up to you to judge whether this impacts your end use, and whether the additional cost of bisulfate offsets the savings in sulfuric acid.

Also, the concentration of your final product nitric acid is wholly dependant on the concentrations of your other solutions - use as little water as possible dissolving each reactant for highest final concentration.

vmelkon - 10-12-2013 at 18:06

I don't know if it occurred to anyone but if you want to do
Ca(NO3)2+NaHSO4 in water, then you aren't forced to keep the HNO3 + NaNO3 solution. You can distill it and collect the HNO3 vapors.

Another way would be to dry distill Ca(NO3)2+NaHSO4 and collect the HNO3 vapors.

Nitric acid for the 10,000th time.

Varmint - 11-12-2013 at 07:32

All I want is nitric acid to create sivler nitrate.

Many are familiar with the balance equation: 2KNO3 + H2SO4 = 2HNO3 + K2SO4.

So, lets say I take 2 Mol KNO3 (202.2064G) and 1 mol H2SO4 (53.3038mL). Should be 100% stoichiometric, and the specific post reaction masses should be: 126.02G HNO3 and 174.259G of K2SO4. Correct?

Since my sulfuric is 98+% concentration, and theoretically consumed completely in the reaction, this should leave how much water?

The reason for the question is simple, K2SO4 is fairly poorly soluable at 20C (24.1G/100mL), so shouldn't there be a huge mass of K2SO4 at the bottom of the flask?

Lets assume the H2O component in the equation (only source of water is from the sulfuric) just so happened to look like pure water for this exercise. In 53mL of water, we can only expect 12.773G of K2SO4 to be in solution (with about 161.5G at the bottom of the flask). Right? (I know its wrong, thus I need help.)

Taking it further, I could cool the contents to 0, and leave only 7.4G in solution, the rest having crashed out.

Cooling even further to near the nitric freezing point, presumably we could get virtually all of the sulfate to drop out, leaving a fairly clean nitric.

This does not work of course, so, what part of the big picture am I not accounting for?

Thanks,

DAS

Galinstan - 11-12-2013 at 07:40

I think it is because when you mix your sulphuric acid and potassium nitrate your don't form nitric acid and potassium sulphate your get an ion soup so the sulphate remains in solution as sulphate ions with K+ and H+ and NO3- it doesn't just form potassium sulphate and crash out as nice a this would be

Varmint - 11-12-2013 at 07:52

OK, that is as I understand it (I'm trying to play dumb enough to get all the details).

So, now that we have this soup, is there anything we can do to cause anything to partially crash out, or, is distillation the singular way out?

And, what if I were to add my silver to the soup, would I get silver nitrate in some porportion with a sulfated complex?

Seriously, all I want is a nice 10 or 20G of silver nitrate, and I don't want to pay what the manufacturers want since they still seem to think silver is at its all time high.

thanks

Dariusrussell - 11-12-2013 at 07:53

Well one issue off the bat that I see, is you don't want to run this stoichiometrically. If you want to yield all the HNO3, then you will need H2SO4 in excess. Also yes there will be a massive cake of K2SO4 in your flask.
Also during distillation, there should be very,very little dissolved solid in the water, if there is any at all. For making silver nitrate, you don't need R/WFNA, 70% would suffice.
In which case performing a second distillation adding some distilled (Or better yet Deionized) water, you would be left with significantly less contaminants. I would be more worried about how pure your silver is.

Ahhh whoops, I thought you were going to distill it.

[Edited on 12-11-13 by Dariusrussell]

Varmint - 11-12-2013 at 08:50

.999 fine Ag, so purity should be good enough.

And why won't stoich suffice? I can understand 1-2% to offset the water in the sulfuric, but I ran the stoich taking into account the mol at the correct density.

Thanks

[Edited on 11-12-2013 by Varmint]

hyfalcon - 11-12-2013 at 09:03

Calcium nitrate would be a better starting nitrate. Then the calcium sulphate would crash out. Make a god awful mesh unless you had enough water in the solution to allow the calcium sulphate to settle.

Moderator Warning

bfesser - 11-12-2013 at 09:10

Varmint, the last nitric acid/silver nitrate topic hadn't even made it off the 1-day 'Today's Posts' list. Your subject line also shows that you're aware of the existing topics but that you defiantly chosen to ignore them. I suggest that you read through all of them in detail before posting again. Blatant requests for spoon-feeding are unacceptable. There's more than enough information on this subject all over the internet and this website if you put in more minimal effort. In addition, posting in the way you have is just plain rude—akin to trolling or spamming. I expect to see some references included in posts of any future questions, to show that you've at least tried to answer them on your own; if not, the thread will be closed without further warning and subsequent new topics deleted.

[edit] Read the SM Guidelines (link in sig.) before posting anything more.

[Edited on 11.12.13 by bfesser]

Varmint - 11-12-2013 at 09:15

hyfalcon:

Describe "god awful mess" if you would.

As is obvious, I don't need greater than 70%, so I could easily afford to add so distilled water, in fact I shouldn't need 70% to create the silver nitrate.

So, can you help work me through the definition of mess? Interesting to see calcium sulfate solubility first rises then falls across the span 0-100C, but still never goes past 0.265G/100mL. So that means huge shitloads of water, placing final concentration way down in the noise.

Interesting indeed.

Thanks for the info, and again, expand on your comments if you would.

DAS

Varmint - 11-12-2013 at 09:27

Bfesser, you are such an arrogant prick, it is unbelievable. Take some of your own medicine asshole, and read all the nitric acid threads, and every one of them is trying to get above 70 for nitrations for things that go boom. it is as though what I'm asking for then is below the radar of common interest, and is blatantly neglected when taking the sulfuric/nitrate path. It is automatically assumed high concentration is the goal. In my case, it is not.

Is there a reason you have to be such a snotty prick? You should get a job man, stop living of food stamps and make something of yourself other than the asshole you relish making yourself here.

So delete everything. do it. you so badly want to be in control of something since you lost control of your life. Go ahead, I'm certain you will pound your chest with glee. but you will still be the biggest asshole on the internet, bar none.

DAS

bfesser - 11-12-2013 at 09:30

<a href="viewthread.php?tid=27851">Nitric Acid and Silver Nitrate</a>

[edit] As I've pointed out in previous posts—which you would know, had you bothered to read them— I am well aware that highly concentrated nitric acid is undesirable for dissolution of silver. I absolutely do not want to discourage discussion on this topic; merely limit the number of new topics being added to the mess. This is a message board, not IRC; replying to existing threads is preferred. I'm very disappointed with having to close this topic, but your last post certainly warrants it. [closed]

[Edited on 11.12.13 by bfesser]