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CHRIS25
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254 mL of 70% HNO3 is 252 g of HNO3
593 mL of 30% HNO3 is 209 g HNO3.
How did you do that?
Also, since it has been wisely put forward that I should put in excess copper, then it seems to my logic that if I had 10ml of HNO3 and kept adding
copper until there was no more reaction then I would have the same thing essentially, copper nitrate? And all this maths becomes null and void?
[Edited on 29-9-2012 by CHRIS25]
[Edited on 29-9-2012 by CHRIS25]
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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Vargouille
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The density of a solution of 70% HNO3 solution is 1.42 g/mL at 25C. By multiplying it by the ratio (70 g HNO3/ 100 g sln), I received the answer of
0.994 g HNO3/ mL sln. Multiplying this by 254 mL gives 252 g HNO3. Previously, I calculated the amount of HNO3 in a mL of 30% solution (0.35289 g
HNO3/ mL sln) and simply multiplied it by 593 mL.
You could do that, of course, but doing the maths gives you an idea of how much of the reagents you need, so you can have everything ready before the
start of the reaction. One problem with your process is that once most of the HNO3 is depleted in the reactions, the final reactions are less visible.
The solution, of course, is to simply leave it to react for a long time, but as from this video, the reaction can take an exceedingly long time to go to completion.
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CHRIS25
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Hi thanks. Ok That is straightforward. By the way Im putting all this into a PDF. A lot of work here to note. So, I watched that video, two days?
But that was 100ml of his unknown phosphoric and nitric mixture. Still I have time. Will perform the reaction and re-evaluate my other concoptions
now that I have this new maths to be applying. But when you say excess copper needs to be added, what kind of percentage of the stoichemetry amount
would you suggest?
Also, if you don't mind, I notice that Copper trihydrate is Cu(NO3)2.3H2O. Since my reaction theoretically ends with 2H2O, what will I actually be
getting since there is no dihydrate that I have found. There is a monohydrate though, I suspect that since H2O can not exist, (its formula being
2H2O) then this is the monohydrate? Remember I am not a chemist?
[Edited on 29-9-2012 by CHRIS25]
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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Vargouille
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I would suggest a 10% stoichiometric excess. It should be enough. As for the hydration states, I wouldn't presume that copper nitrate doesn't exist as
a dihydrate, only that that hydration state hasn't been reported. However, the 2H2O formed from the reaction is unrelated to the hydration state of
the crystallized copper nitrate, especially considering that the reaction takes place with water as a solvent. There are a few factors, including
ambient temperature, and humidity, that impact the hydration state of crystallized salts. I can not honestly say that I am sure of the hydration state
of the salt you will produce, because of those reasons, and due to the hygroscopy of copper nitrate, the tendency to absorb water from the atmosphere.
My best guess is that it will be either in the trihydrate or hexahydrate form, both of which form blue rhombic crystals.
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CHRIS25
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So the excess copper will, as was suggested earlier by someone, ensure that all of the Nitric acid part be used up. Leaving me with copper nitrate in
water, though probably slightly acidic? Which I can then gently boil away and place the remaining wet solids into a dessicator bag. Is this a
suitable method to proceed? I am not after the anhydrous, this seems impossible for someone like me and actually I don't need it since my end product
will be re-dissolved again anyway. I have read that attempting to dehydrate any of the copper 2 nitrate hydration states produces CuO. But the action
of boiling away until just the visible liquid is gone does not even come close to de-hydrating does it?
[Edited on 30-9-2012 by CHRIS25]
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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Vargouille
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"Gently boiling" is still far above the decomposition temperature of copper nitrate, 80C. Applying heat above that temperature, even if the copper
nitrate is in solution, will decompose it. Also, apparently, the hexahydrate decomposes at 26.4C, so it may be unlikely that you form that hydration
state, depending on ambient temperature. I haven't yet found what the decomposition constitutes, mind you, but that's something that you can search
up. In any case, the best solution is after filtering to remove the copper, to put it into a dessicator bag to dry, perhaps for weeks. If you keep the
temperature low, around 50C, you should be able to remove a large amount of excess water, but I would prefer the "slow boat to China" route.
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CHRIS25
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Right. I have been searching on google for the various decomposition rates at different temperatures: Frustrating when all you turn up is the
obvious chit chat, and videos showing you what happens. Nothing around really that I can find listing various temperatures. Anyway this requires a
sophisticated search for those tech documents - never know how to search for those, requires someone who knows exactly how to type in the correct
words in the correct order - trouble with google is that unless you have one of those thick manuals that deliver you weeks of reading on how to target
your query (phew - no time for that), people like me will always miss about 80% of what really is out there. I think I'll take a speed boat at first,
gently heat keeping it below 40 degrees with my hot plate. I think that should do it, I saw how much heat it took to decompose that copper nitrate
in a test tube, so 40 degrees or below should get rid of the water and keep the Cu(NHO3)2 in one piece. I can't wait weeks, I will probably put half
away for weeks, and work on the other half. Many appreciations for your constant time on this, I have actually learned far more than just the copper
nitrate question, which actually is what I want to do, I can never settle for how answers, I need to understand the answer otherwise the answer has no
purpose for me and I can not use it.
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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