Log in

View Full Version : concentrating nitric?


S. Toppholzer
June 5th, 2002, 03:48 PM
There is the commonly available 75% nitric acid - could it be concentrated like sulfuric acid?

10fingers
June 5th, 2002, 04:04 PM
Once nitric acid reaches a concentration of 68% it cannot be further concentrated by heating to drive off water. It forms an azetrope at this level. To further concentrate it must be distilled from conc. H2SO4. There are other ways to do it but this is probably the easiest.

Einstein
July 29th, 2002, 11:57 AM
If I have gotten it right, you have to add strong (98% should do right?) H2SO4 to your nitric acid, then distillate the nitric, boil remaining H2SO4 until you have 98% again and do this whole thing couple of times. This should give you 90+% HNO3, IF I´m correct.

PYRO500
July 29th, 2002, 02:26 PM
Yes that will work.

megalomania
July 30th, 2002, 02:39 AM
It should actually give you 100% nitric acid if you use strong enough sulfuric acid to begin with.

knowledgehungry
October 10th, 2002, 04:37 PM
in terms of concentrating hno3 from h20 might it be possible to use electrolylosis to break down the h20 in the solution of nitric??
i know that even if it was theorytically possible i know that it would most likely be too dufficult and risky to do but i was just wondering if it has been tried or if there is no way that it would work. Is it likely that the hno3 would decompose too?
Just for you all to think about...

kingspaz
October 10th, 2002, 07:02 PM
knowledgehungry, i know you mean well but by saying 'i know that it would most likely be too dufficult and risky' it makes you look silly.
HNO<sub>3(aq)</sub> + H<sub>2</sub>O<sub>(l)</sub> &lt;=&gt; H<sub>3</sub>O<sup>+</sup><sub>(aq)</sub> + NO<sub>3</sub><sup>-</sup><sub>(aq)</sub>

this shows that solutions of nitric acid are ionised. so when you put current through it the NO<sub>3</sub><sup>-</sup> ion will go to the cathode and the H<sub>3</sub>O<sup>+</sup> will go to the anode so not only the water will be destroyed. good thoguht though.
also look in forum matters under new features to see how to use the proper format for writing chemicals instead of h2o and the like.

<small>[ October 10, 2002, 06:04 PM: Message edited by: kingspaz ]</small>

Marvin
October 11th, 2002, 01:58 AM
You can electrolyse nitric acid to make it more concentrated, but its more complex than simply sticking 2 electrodes in and away you go. The method you are hoping it will work by is flawed for reasons other than have been mentioned already. H3O+ being positive moves towards the cathode where it is reduced, which normally results in the production of nascent hydrogen. Every process however, that would normally produce hydrogen in acids fails in conc nitric acid oweing to its agressive oxidising action with the production of NO2 gas. NO3- moves towards the anode, where its oxidised. The solution eventually ends up giving half an oxygen atom per electron regardless of the actual species oxidised by the anode, any NO3 produced ending up back as NO3-. The failiar of this to liberate hydrogen effectivly makes simple electrolytic concentration impossible. A more complicated cell with seperate anolyte and catholyte compartments can be used, the NO2 produced at the cathode is fed into the anolyte to be reoxidised and concentration effectivly proceeds by ion migration rather than actual electrolysis of the anolyte at the expense of the catholyte. You dilute as much acid as you concentrate, and the anolyte needs replenishing to keep the process going. Over 99% nitric acid can be achieved eventually....

NO2 is extremely toxic. Fear it.

Richy
October 12th, 2002, 07:56 AM
how concentrated is the nitric acid from the H2SO4 + KNO3 reaction?
i am using concentrated battery acid (boiled at 100c till no more steam then boiled at 130c for 20 mins) and 98% potassium nitrate. the final HNO3 product is orangy and fuming.

vulture
October 12th, 2002, 01:46 PM
Theoretically the concentration is 99%, but everytime your HNO<sub>3</sub> sees air, it will decrease in concentration because it will evaporate and decompose partially, that's why it's called fuming. If you leave your HNO<sub>3</sub> exposed to air it will autodilute to 68%. Use as soon as possible.
You'll notice that alot of nitration recipes use a nitrate salt with concentrated H<sub>2</sub>SO<sub>4</sub> because then almost 99% HNO<sub>3</sub> will be created in the reaction vessel which is often superior to earlier prepared HNO<sub>3</sub>.

Marvin
October 12th, 2002, 03:41 PM
The decomposition of nitric acid has virtually nothing to do with air. Pure nitric acid decomposes into water, oxygen gas and NO2, which colours the solution brown. Its accelerated by heat and light, as usual. Dont assume that decomposition will stop at the azeotropic point of the acid, thats just when the vapour pressure of the liquids is in the same proportion of the constitution of the acid at boiling point. Although I say 'virtually nothing', the eventual equilibrium depends on the concentration of oxygen (its partial pressure), and the amount of NO2 in solution, as well as the amount of water.

Acid mixtures prepaired by dissolving nitrates are never 'superior' to equivalent mixtures of acids. The only time such mixtures are used industrially is when nitric acid cannot be obtained. While nitration mixtures are generally expressed as a mixture of absolute nitric, sulphuric and water content, its accepted that they will almost never be made in any scale by mixing pure nitric and sulphuric acids. The industrial cost of pure nitric acid is too great for this to make sense, instead they are usually made by mixing the calculated amounts of azeotropic nitric acid, oleum and if needed water to result in an identical mixture than if pure nitric and sulphuric had been used. The only reason anyone would see 'alot' of directions using nitrate salts, would be if they restricted themselves to fringe literature, while some of these methods may work, they arnt accepted in reliable texts. Anyone who is serios about this sort of chemistry should buy the pyrex bits for a still, quickfit works well, and vacuum distillation under it results in less decomposition, as other members have also reported. Anyone that isnt serios, shouldnt be doing any of this at all.

One of the big problems with dissolving nitrate salts in sulphuric acid, is a rather complex equilibrium results that keeps quantities of nitrate and sulphate and hydrogensulphate undissolved. Its the specific application of Le Chatelier's principle usually taught as 'The common ion effect'. The stronger the concentration of nitric acid in solution, the lower the solubility of the alkali nitrate in it. The other problems are more obvios.

There are several good methods of making pure nitric acid without distilling from NO2, but owing to its toxicity, none of these are suitable for home production. If you add liquid, or concentrated NO2 gas to water a rather violent reaction is produced where the liquid apears to 'boil'. NO2 reacts with the water to produce nitric and nitrous acids as in the expected dilute system, but in the concentrated acid, the nitrous acid decomposes releasing NO gas. Its the reoxidation of the NO to NO2 in the gas phase that plays the dominant role in the oxidation. Industrially the correct amounts of liquid NO2 and water with oxygen gas at 50atmospheres is sealed into containers and eventually produces over 99% nitric acid solution.

The reason conc nitric acid fumes in air is exactly the same reason as HCl gas fumes in air. They are good enough dehydrating agents that they cause the water vapour in the air to precipitate as acid solution, the solution having a lower vapour pressure than the ambient concentration of water. Nitric acid is only decomposing if it turns brown.

Richy
October 13th, 2002, 02:26 AM
excellent, thanks guys. i have had it stored in an airtight container for 3 days now. ill use it in some NG today. wish me and my fingers luck.

vulture
October 13th, 2002, 06:42 AM
</font><blockquote><font size="1" face="Verdana, Arial, Helvetica">quote:</font><hr /><font size="2" face="Verdana, Arial, Helvetica"> The decomposition of nitric acid has virtually nothing to do with air. </font><hr /></blockquote><font size="2" face="Verdana, Arial, Helvetica">What I meant is that in an open container the nitric acid will decompose quickly because the NO2 gas can escape. In a closed vessel the partial NO2 pressure builds up and stops the equilibrium from moving further to the right. Also, unless you work in a dry room it's going to attract water from the air.
Hmm, I wonder if adding conc. H<sub>2</sub>SO<sub>4</sub> or conc. H<sub>3</sub>PO<sub>4</sub> to the nitric for storage would improve storage life. Also I wonder If one could drain the H<sub>2</sub>SO<sub>4</sub>/water or H<sub>3</sub>PO<sub>4</sub>/water layer or does this mix with the HNO<sub>3</sub>?

</font><blockquote><font size="1" face="Verdana, Arial, Helvetica">quote:</font><hr /><font size="2" face="Verdana, Arial, Helvetica"> Acid mixtures prepaired by dissolving nitrates are never 'superior' to equivalent mixtures of acids. </font><hr /></blockquote><font size="2" face="Verdana, Arial, Helvetica">From the perspective of the home experimenter they are. If you use enough H<sub>2</sub>SO<sub>4</sub>, you'll get better nitrating than with 70% nitric acid.
Ofcourse 70% nitric acid also nitrates better with an excess of sulfuric acid, but then why go through the hassle of preparing the nitric acid?

</font><blockquote><font size="1" face="Verdana, Arial, Helvetica">quote:</font><hr /><font size="2" face="Verdana, Arial, Helvetica"> Nitric acid is only decomposing if it turns brown. </font><hr /></blockquote><font size="2" face="Verdana, Arial, Helvetica">Green (yes green! Had it once myself after distilling), Yellow and colors between yellow-brown are signs of decomposition.
Deep red HNO<sub>3</sub> isn't necessarily a bad thing as long as there isn't much water in it.

</font><blockquote><font size="1" face="Verdana, Arial, Helvetica">quote:</font><hr /><font size="2" face="Verdana, Arial, Helvetica"> Dont assume that decomposition will stop at the azeotropic point of the acid </font><hr /></blockquote><font size="2" face="Verdana, Arial, Helvetica">At the azeotropic point the decomposition will decrease significantly because dissociated acid won't decompose that easily. Also at the azeotrope water and nitric acid will decompose at the same rate, so the solution will stay at 68%, but the volume will decrease.

</font><blockquote><font size="1" face="Verdana, Arial, Helvetica">quote:</font><hr /><font size="2" face="Verdana, Arial, Helvetica"> NO2 is extremely toxic. Fear it.
</font><hr /></blockquote><font size="2" face="Verdana, Arial, Helvetica">Very true! It's a silent killer. I once breathed a significant amount of NO<sub>2</sub>/HNO<sub>3</sub> gas when opening a distilling setup. Besides a scorched feeling in my throat and nose, I felt no other effects for half an hour. But then I got a terrible headache, high hearthbeat, a feeling of anxiety and stress and severe nausea. This lasted an entire night. <img border="0" title="" alt="[Frown]" src="frown.gif" />

<small>[ October 13, 2002, 05:54 AM: Message edited by: vulture ]</small>

aster
December 23rd, 2002, 01:37 AM
i am making my mercury fulminate following the mega recipe, when the mercury react with the nitric acid (68%)they released the nitrogen dioxide gas, and the product coloured green, (settle down in the lower layer of the container/glass). if i use more volume of acid than recommended, and closed the open end of the beaker or erlenmeyer by placed a mica sheet, then will the nitrrogen dioxide gas that trapped in will react with the water in the nitric acid (68%)and increasing the concentration of the nitric acid left?(so i can get more concentrated acid) then i can simply separate them carefuly by decant the upper layer, am i right? is there any danger to use this left nitric acid that i assumed became more concentrated than before for syntesis another recipe that require more concentrated acid?

aster
December 30th, 2002, 04:23 AM
humm .......i realize something wrong with my idea above, so from the reaction of the mercury and HNO3 the left acid will be more dilute than concentrated, because when NO2 gas formed, then the H2O will left and dilute the remaining acid, maybe using separated glass, the one for reacting mercury with HNO3 and the other just fill with HNO3 (and keep closed and wrapped dark), so the stream of the NO2 gas can react with the water in the glass fill only with HNO3 and increasing the concentration. being closed and wrapped dark so the HNO3 won't decompose upon influence of light, ...any one ever try this? i am wonder how much this (idea) can concentrated the acid ?

Marvin
December 30th, 2002, 09:18 PM
</font><blockquote><font size="1" face="Verdana, Arial, Helvetica">quote:</font><hr /><font size="2" face="Verdana, Arial, Helvetica"> What I meant is that in an open container the nitric acid will decompose quickly because the NO2 gas can escape. In a closed vessel the partial NO2 pressure builds up and stops the equilibrium from moving further to the right. Also, unless you work in a dry room it's going to attract water from the air. </font><hr /></blockquote><font size="2" face="Verdana, Arial, Helvetica">In fact exposure to air does play a part which is why I said 'virtually', but not becuase of loss of NO2 which is very soluable in the conc acid, loss of NO2 in normal storage will be much less than loss of the acid itself by simple evaporation even in an open container. In an ordinary partially sealed container, ie effectivly sealed, but pressure can vent, the atmosphere will push out all the air, ending up as mostly oxygen. Anyone who stores nitric in an open container deserves everything they get, and it will be a lot more serious than the nitric absorbing water. The other main product apart from the diluting water is oxygen, and this is the real problem becuase of its low solubility and the resultent high pressure of oxygen that forms if you try to prevent its escape. Making containers resistant to nitric acid/NOx at several 10's of atmospheres is certainly possible industrially, but for home chemists is not recommended. What matter for home storage, are kinetic factors, mainly heat and light, rather than equilibrium conditions that are never really aproached and impractical to alter.

</font><blockquote><font size="1" face="Verdana, Arial, Helvetica">quote:</font><hr /><font size="2" face="Verdana, Arial, Helvetica"> Hmm, I wonder if adding conc. H2SO4 or conc. H3PO4 to the nitric for storage would improve storage life. Also I wonder If one could drain the H2SO4/water or H3PO4/water layer or does this mix with the HNO3? </font><hr /></blockquote><font size="2" face="Verdana, Arial, Helvetica">The decomposition turns acid into NOx/O2/H2O, so adding a strong acid, or a dehydrating agent will tip the balence in the favour of more decomposition, not less, with or without a means for removing the water continually. The more agressive particualy the dehydrating action, will increase the speed as well as affect the equilibrium result. Both phosphoric and sulphuric are miscable with nitric acid or we wouldnt be using mixed acids for nitrations. Nitric acid and phosphorous oxide forms mixed layers of mainly nitrogen pentoxide and some form of phosphoric acid, but I havnt seen a way to exploit this without using P2O5. Maybe HPO3 could be made to dehydrate the acid and form a decantable pyrophosphoric layer, but I suspect this wouldnt work well.

Nitrate salts in a nitrating mixture produce hydrogen sulphate, reducing acidity, wasting acid, the ppt causes problems of its own but for some reason the mixture behaves much worse than that of pure nitric acid in the same total volume. Ive tried nitrating mixtures with mixed acids, and nitrate salts and for want of a better explanation the salts seem to kill the dehydrating properties of the sulphuric acid. Adding additional sulphuric acid counters this, but at the same time dilutes the nitric in the mixture and the best comprimise performs much less well than ordinary mixed acid and at considerable extra waste of sulphuric. Some people may be desperate enough to try this but the results are never superior to distilling the acid yourself and saving on resources. Ive used 70% nitric much less, simply becuase I cant buy it here but the dinitrations on an unactivated benzene ring with 70% nitric/sulphuric went much better with a relativly low excess of nitric acid than any of my experiments with sulphuric/nitrate mixtures did. As a secondary point mixed acid can be recycled a number of ways if its been prepared from acids rather than salts (cf the TNT synth in Davis for the most elegent), Id go as far as to suggest recycling nitrate salt mixtures is completly unfeasable without using the sort of distillation that would have better been used to prepare the nitric acid in the first place.

Ive had greenish acid before and I confess I assumed it was an external contaminent like copper. Looking at it logically it could be an affect of dissolved N2O3 which is blue, and the yellow/red colour of dissolved NO2 resulting in a green hue, but I'm speculating. Brown nitric acid is a real problem as its a much more agressive oxidising agent than nitric acid is, and this is certainly one factor in runaway reactions. Getting rid of the NO2 often involves effectivly adding water, eg use of urea. Storing brown nitric acid is only an advantage over long periods of time with pressurised oxygen so I see no reason not to get rid of as much of the NO2 as possible prior to cool storage.

</font><blockquote><font size="1" face="Verdana, Arial, Helvetica">quote:</font><hr /><font size="2" face="Verdana, Arial, Helvetica"> At the azeotropic point the decomposition will decrease significantly because dissociated acid won't decompose that easily. Also at the azeotrope water and nitric acid will decompose at the same rate, so the solution will stay at 68%, but the volume will decrease. </font><hr /></blockquote><font size="2" face="Verdana, Arial, Helvetica">I dont follow your reasoning here, there is no sudden point at which the acid becomes dissociated, only a smooth curve that has nothing to do with the Az point at all. The Az point is simply the point at which the partial pressure of the gas is in the same ratio as the liquid its above. Theres a tendency to ascribe a 'molecular formula' to this, particulaly in older textbooks, but the points change with temperature and pressure and often bear little relation to the mixtures that crystalise out on freezing for example. There is however a gut feeling that tends to creep up when thiking about these ideas, that the molecules 'prefer' to be in these ratios and that the composition will tend to form the 'prefered' mixture. I'm not accusing you of this, just suggesting this might be a gut impulse that is misleading the logic. If there was a massive great Az discontinuety in Water/peroxide mixtures would this affect its decomposition? Rather more obviosly not, and Id expect nitric acid decomposition to depend smoothly on the temperature and the concentration of the acid and each product. Pure nitric acid, isnt pure nitric acid when its liquid, its a mixture of nitric acid, water and nitrogen pentoxide. If its the nitrogen pentoxide that is the major kinetic factor in decomposition, then the concentration of water might have nothing directly to do with it at all.

aster, Whats missing is oxygen, the oxygen converted to water by the dissolving metal. Pure nitric acid has been made industrially, and probably still is, by mixing the right proportion of NO2, water, and oxygen at 50 atmospheres and results in 99%+ acid. NO2 is very nasty however, and making nitric, or increasing the concentration of nitric acid this way isnt worth the health problems in my opinion. If you think you can setup a draft cabinet good enough to avoid health problems, and pure oxygen, preferably at pressure safely without anything corroding, or splatting nitric acid everywhere, then I'm sure small amounts of higher concentration acid can be produced, slowly.

NO2 has an unpaired electron, which is what makes it so reactive, paramagnetic, and brown. It destroys lung tissue very fast, so much so it makes your body think it has a lung infection, its your bodies reponse to this that makes you feel ill, or die. The effects of which vary between 'mild infection' symptoms all the way to fatal puminery edema as much as several days after the exposure, when of course no infection was ever actually present. There is also cumulative long term damage which can be symptomless apart from the impaired lung function.

If you choose to play with NO2, you will need much more than luck, and much more information that anyone would be able to find on the internet or should expect to find at a forum.

aster
December 31st, 2002, 08:34 AM
ok, your explanation is very helpful, thanks. another idea is making more concentrating nitric acid by adding few H2SO4 98% then using calcium nitrate to convert the H2SO4 into HNO3 and Calcium sulfate that will precipitate upon forming, what is the plus minus of using this method? it will be more safe than disolving NO2 gas into HNO3 solution i think...

jarrod
January 2nd, 2003, 06:37 AM
A bit off topic… :rolleyes:
But I have a really good site with pics and detailed info on distillation, but its in GERMAN. If any one out there can translate this for me I would be grateful. I have already tried the free translators and they get general idea but are shotty overall.

In den Destillationskolben links füllt man 25 ml konz. Schwefelsäure. Vorsichtig, in Portionen und unter guter Kühlung mit Eiswasser schüttet man 30 g gut getrocknetes, fein gepulvertes Kaliumnitrat durch einen Trichter hinein. Dazu gibt man noch ein wenig Silbernitrat (AgNO3). Bei dieser Prozedur bilden sich bereits stechend riechende Dämpfe, nicht einatmen! Nach etwa einstündigem Stehen im verschlossenen Kolben wird die Mischung vorsichtig erwärmt; die zuerst übergehenden Tropfen (erkennbar an der gelb-braunen Farbe) werden gesondert aufgefangen.
Der als Vorlage dienende Rundkolben muß (am besten durch Einstellen in ein Becherglas mit kaltem Wasser) ebenfalls gut gekühlt werden. Bei dem Versuch destilliert fast reine Salpetersäure über (im Temperaturbereich von 83 bis 85 °C bei Atmosphärendruck). Sobald sich größere Mengen brauner Stickoxide bilden und sich das abtropfende Destillat bräunlich verfärbt, wird die Reaktion abgebrochen.
Sind präparativ größere Mengen reiner Salpetersäure darzustellen, so arbeitet man auf jeden Fall mit einer Schliffapparatur und destilliert wiederholt im Wasserstrahl-Vakuum (Siedekapillare verwenden). Der Siedepunkt der Säure liegt bei einem Druck von 26 mbar nur noch zwischen 36 und 38 °C!. Hier reicht zur Erwärmung des Destillierkolbens ein Wasserbad aus

<a href="http://www.chemikalienlexikon.de/cheminfo/1248-lex.htm" target="_blank">the site</a>

thanx
:confused:

vulture
January 2nd, 2003, 08:41 AM
</font><blockquote><font size="1" face="Verdana, Arial, Helvetica">quote:</font><hr /><font size="2" face="Verdana, Arial, Helvetica"> I dont follow your reasoning here, there is no sudden point at which the acid becomes dissociated </font><hr /></blockquote><font size="2" face="Verdana, Arial, Helvetica">That is not what I said. Nitric acid of 68% is fully dissociated into H<sup>+</sup> and NO<sub>3</sub><sup>-</sup>. Above that concentration, there will be HNO<sub>3</sub> molecules present which can decompose easily into NO<sub>2</sub>, water and oxygen.
Dissociated acid decomposes much slower.

Rhadon
January 2nd, 2003, 08:53 PM
jarrod: The procedure you pasted here is a well-known one. The only thing that could be new to some persons might be the addition of small amounts of silver nitrate. Yet, I'll do a very crude translation:

</font><blockquote><font size="1" face="Verdana, Arial, Helvetica">quote:</font><hr /><font size="2" face="Verdana, Arial, Helvetica">
The distillation flask which is shown on the left side is filled with 25 ml of conc. H2SO4. Then there are, cautiously and in small portions, added 30 g of well dried and finely ground KNO3 by means of a funnel. Then there is added a small amount of AgNO3. At this point there will be evolved vapour with pungent odour. Do not breathe them! After standing for about an hour in the closed flask, the mixture is heated cautiously. The first drops of the distillate (which can be recognized by their yellow-brownish colour) are collected a separate vessel.
The round bottom flask which is used for collecting the distillate has to be cooled well, e.g. by placing it in a beaker which is filled with cold water. In the experiment the nitric acid that distills is almost pure (@83 - 85 °C). As soon as larger amounts of brown NO2 are formed and the distillate has a brownish colour, the reaction is stopped.
Are larger amounts of HNO3 to be prepared, one should use a ground distillation apparatus and distill multiple times under vacuum in any case. The boiling point of the acid is about 36 - 38 °C @26 mbar, so a water bath is sufficient for heating.
</font><hr /></blockquote><font size="2" face="Verdana, Arial, Helvetica">

Marvin
January 9th, 2003, 09:12 PM
</font><blockquote><font size="1" face="Verdana, Arial, Helvetica">quote:</font><hr /><font size="2" face="Verdana, Arial, Helvetica">
That is not what I said. Nitric acid of 68% is fully dissociated into H+ and NO3-
</font><hr /></blockquote><font size="2" face="Verdana, Arial, Helvetica">A finite point in the nitric acid equilibrium at which you could say one species was not present would qualify as 'Sudden' in physical chemistry. An acid is is only assumed to be completely dissociated at infinite dilution. Its also not clear to me why you chose the azeotropic point to be the point at which HNO3 stopped existing.

</font><blockquote><font size="1" face="Verdana, Arial, Helvetica">quote:</font><hr /><font size="2" face="Verdana, Arial, Helvetica">
Dissociated acid decomposes much slower.
</font><hr /></blockquote><font size="2" face="Verdana, Arial, Helvetica">H+(aq) cant decompose, and if nitrate decomposed detectably (From a chemical point of view) then we'd expect a RT solution of sodium nitrate which will also generate free nitrate ions to visibly evolve oxygen.

</font><blockquote><font size="1" face="Verdana, Arial, Helvetica">quote:</font><hr /><font size="2" face="Verdana, Arial, Helvetica">
Above that concentration, there will be HNO3 molecules present which can decompose easily into NO2, water and oxygen.
</font><hr /></blockquote><font size="2" face="Verdana, Arial, Helvetica">You are making the assumption that its HNO3 molecules that decompose, something I thought Id made clear in my last post may not be the case and probably isnt. It seems much more likly that its the nitrogen pentoxide that decomposes, certainly in the high concentration acid and this is supported by the observation that solid pure nitric acid does not seem to decompose at all of its own accord.