Nitric Acid manufacture

[stanfield]

Like many guys on this site, I produce my own nitric acid by distilling sodium nitrate and sulfuric acid, the equation is : 3NaNO3+2H2SO4 ==> 3HNO3+NaSO4+NaHSO4
So, if I have 500g of NaNO3, how much sulfuric acid do I need ? (I have 95 % H2SO4)
I think i'm wrong somewhere in my calculations...

thanx...

[Anthony]

RAM of 1 molecule of HNO3 would be 1x1 1x14 3x16 = 63

For one molecule of HNO3 you'd need one molecule of NaNO3 to supply the NO3 and 0.5 molecules of H2SO4 to supply the H.

So you should need twice as many molecules of NaNO3 to H2SO4.

RAM of NaNO3 = 1x23 1x14 3x16 = 85
RAM of H2SO4 = 2x1 1x32 4x16 = 98

85/63 = 1.35
98/63 = 1.55

So, you need 1.74 times as much NaNO3 as H2SO4. so for 500gm you'd need 500gm/1.74 = 287gm H2SO4 and you'd get 371gm of HNO3

[Lagen]

stanfield: Actually there should be Na2SO4 but I see what you mean. Given that the equation is correct, and given your numbers, you would need 405g of the acid, the theoretical yield would be 391g HNO3 containing 20g water, or roughly 95% nitric. (371g anhydrous if you manage. But you won't. This reaction proceeds at higher temps and there will be lots of HNO3 decomposition.)
Anthony: I think a very high temperature would be necessary to convert NaNO3 to Na2SO4 completely (lots of decomposition again). I lean more towards stanfield's equation, or maybe only NaHSO4 would be produced, depending on the conditions. In that case (at lower temp., 50-60°C, vacuum):
NaNO3 + H2SO4 -> HNO3 + NaHSO4 and 607g H2SO4 -> 371g HNO3 anhydrous again, or 401g with all the H2O (only 92.5% this way, but more likely to succeed).

[This message has been edited by Lagen (edited July 20, 2001).]

[stanfield]

Anthony VS Lagen :
Two guys
Two answers
Who win ?
http://theforum.virtualave.net/ubb/smilies/smile.gif

more seriously, i found 405g like anthony...
but Lagen's calculation seems right too...

The temperature at the end of my Vigreux Column is around 70-80°C so, I dont think Water can pass since, it boils at 100°C !

Any reply is more than appreciated

[Lagen]

stanfield: 405g of which acid? Of the distillate? Or did you mean the 405g in my post? I'm confused. Anthony's reaction assumes more complete (in fact perfect) conversion and hence you would need much less H2SO4 than 400g.

Generally the amount of distillate would be around 400g too, it doesn't depend much on which of the reactions will happen. (Only the H2SO4 consumption and resulting purity.) In reality you shouldn't expect more than 85% from the first distillation, and should redistill with H2SO4. Some water will pass too, as the b.p. of H2O and HNO3 are too close together and they will not boil completely separately. It's similar to distilling alcohol, you don't get 100% at 80°C, but not even the 96% (azeotrope) from the first distillation. Some H2O also results from HNO3 decomposition (if the red fumes are there then H2O will be too).

Somebody with greater experience, please give us your comments...

[This message has been edited by Lagen (edited July 22, 2001).]

[Anthony]

Dude, go with Lagen's answer! I haven't taken into account the various inefficiencies that will occur.

It does raise a question of my own though, Lagen you say that considerably more H2SO4 is needed than I worked out. When doing nitrations with metalNO3/H2SO4 I have been using the amount of H2SO4 worked out as above. This is obviously a problem, how much extra H2SO4 do you think should be used to counter the effects of the inefficiencies?

[Lagen]

Well I'm not an organics chemist, just a chemistry hobbyist, but these are my guesses:

With plain nitric acid the nitration would look something like this:
R-H + HNO3 -> R-NO2 + H2O
H2SO4 would be added to take up the water. With a "mixed" acid it would look something like this:
(I know this is gross oversimplification.)
R-H + MNO3 + H2SO4 -> R-NO2 + M+ + HSO4- + H2O
R-H + MNO3 -> R-NO2 + M+ + OH-
The metal produced is basic so HSO4- -> SO4(2-)
Can be thought of as: MOH + MHSO4 -> M2SO4 + H2O
------------------------------------------------------------------------------
Overall imaginary equation: 2 R-H + 2 MNO3 + H2SO4 -> 2 R-NO2 + M2SO4 + 2 H2O

During distillation the preferred reaction is MNO3 + H2SO4 -> HNO3 + MHSO4. The "excess" H2SO4 is there to allow all the HNO3 to distill over by this mechanism. The other process, 2 MNO3 + H2SO4 -> 2 HNO3 + M2SO4, where you have just half the amount of H2SO4, proceeds according to MNO3 + H2SO4 -> HNO3 + MHSO4 at first. At some medium temperature (like 100°C I'd say) you end up with a MNO3 + MHSO4 mix (50/50). When this is pushed harder, like with 300°C or so, the MHSO4 acts as an acid: MNO3 + MHSO4 -> HNO3 + M2SO4. So strictly speaking, the excess H2SO4 is not there just because of inefficiency, but mainly because the "perfect" two stage process cannot proceed. I'd say that during the nitrations you have quite a different situation, you have couple of other things around, not just KNO3 + KHSO4 (at least it's in solution!), so the reactions can make it to the M(+)2SO4(2-)...

So I think, wherever you can do the nitration with plain nitric acid, just use metalNO3 + H2SO4 in stoichiometric proportion. Use an excess H2SO4 to the same extent as it is used in the nitric/sulfuric mix. But! I am sure that one cannot change the composition of the nitrating mix and expect to get exactly the same results. One example you all are familiar with: the nitration of PE. Using a mixed acid yields a product of poor quality... Nitrating hexamine to RDX with KNO3+H2SO4 sounds ridiculous, right? Sometimes it can work fine, sometimes it may work backwards, sometimes you won't get anything at all. The background mechanisms of nitration are pretty damn complicated and one cannot make a generalising statement, without having a particular reaction in mind.

For example, I found this info on the effects of increasing H2SO4 concentration:

In solution, nitric acid exists in the form of either NO3(-)H+ or N(O3)(-)H(+). Concentrated HNO3 on the other hand has this structure: NO2.OH or NOH.O2 (in equilibrium). This is called pseudo-acid, and has the form of a nitric ester. The pseudo-acid is capable of nitration and esterification, while the dissociated acid is not. Addition of water increases concentration of dissociated acid (NO3(-)H+). Adding H2SO4 increases the concentration of pseudo-acid. Sulfuric acid is stronger than nitric and passes its proton to the nitric acid: NO2.OH + H2SO4 -> NO(OH)2(+) "nitracidium ion" + HSO4(-), or NO2.OH + 2 H2SO4 -> N(OH)3(2+) "hydronitracidium ion" + 2 HSO4(-). So IMHO a nitrating mix composed of KNO3+H3SO4+HNO3 could behave differently under circumstances...

In the KNO3+H2SO4 mix, the dissociation would look like this: KNO3 + 2 H2SO4 -> NO2(+) + 2 HSO4(-) + K+ + H2O. As you can see, there is only a nitronium ion produced and no nitracidium or hydronitracidium, so again, this may or may not behave like genuine HNO3... I'm sorry to disappoint you, but I'd say stick to the standard, tested procedures and always consider the substitution of a MNO3+H2SO4 mix an experiment!

[This message has been edited by Lagen (edited July 22, 2001).]

[stanfield]

according to you that in my vigreux column, the temperature is ~80°C, what will be the concentration of my nitric acid (approximately) ? Do I need to re-distillate it for HE synthesis ?

thanx !

[Lagen]

That depends on what nitration you have in mind! If it requires something like 98%, you cannot expect that from a single step distillation. But there is a number of reactions where you use only 65-70% HNO3.

The concentration depends mostly on the strength of the sulfuric acid (you said yours is 95%), how much the HNO3 is decomposed, and whether Na2SO4 is formed. The maximum theoretical concentration from the numbers you gave would be 95%, but only if not too much nitrogen oxides appear! If you use a vacuum, you could approach that concentration, and only if you do not use more than the 405g H2SO4 and if the amount of HNO3 produced is around 400g. The other extreme: Without a vacuum, with 600+g H2SO4 95% it could be something like 85%. There are too many combinations and I don't know the details!

My guess is this: NaNO3+H2SO4->HNO3+NaHSO4. You aren't using vacuum and 80°C is a little bit too hot. You would use 600g H2SO4 and the conc. would be lower (like 85-90%).

[Anthony]

Thankyou for taking the time to post that information in detail Lagen. I will do some calculations with various known to work MNO3/H2SO4 nitration procedures and see what amount (if any) of excess H2SO4 they use, to see if there's a general figure.

[stanfield]

hey, Lagen,
do you think I can tell you the weight of my acid (example : 50 mL) and you, you could give me an approximate concentration ?
(for information : my acid is fuming and yellow, not red !)

thanx...

[Lagen]

Sure, no problem, I've got density tables here in my spreadsheet all ready to go!

What I would recommend to you, make an improvised hydrometer (I believe PHILOU posted the instructions some time ago) or buy one with a 1.0-2.0 range. The car battery kind you can get at gas stations etc. will not work, it's only 1.1-1.3.

I could email a decent excel table for HNO3 if you're interested (or post it here, but it's BIG).

[stanfield]

I weighed my acid :
I only have a 0.1 g electronic scale and the value oscillates between 75.8 and 75.7 g for 50 mL. so, what's the concentration ?

and could you send me your density table ?

thanx for all !

[Demolition]

It's easy to find the density.Just divide the weight of the acid by how many ml's there are.In your case
75.8 divided by 50 = 1.51
I'm not to sure but I think the concentration would be around 98-100%.
Did you distill it under vacuum?
Demolition

[stanfield]

I'll post a picture of my distillation apparatus soon...

see ya !

[John456]

Stanfield, what color is your acid? Mine is always a cloudy dark yellow even when its in the 98-100% range. It fumes thick white.

[stanfield]

my acid is yellow but no dark... it fumes thick white too...

here is my distillation apparatus :
http://www.geocities.com/st4nfield/distillation.jpg

see ya !

[stanfield]

damn... when I try to view my page it says 404 http://theforum.virtualave.net/ubb/smilies/frown.gif
I registered to geocities 5 min ago !!!

and you ?

[Lagen]

Stanfield, pretty good job!

But! These measurements can be very tricky! The difference in density of 100% and 90% HNO3 is only 2.6%! Which means the volume must be measured very precisely! A 1% error (0.5ml in your case) would thus lead to a 5% error in calculated concentration!

That's why I stick to the hydrometer - it affords better precision. Another good point is that you can do the measurement even on a small sample, with no decrease in precision. If you measure the volume (& weight) you need a large sample to do it precisely enough. The improvised kind would have another good point, you can make the scale to read in % directly which altogether saves a lot work if you do these measurements frequently. I am sending the density table to you (1000 rows). For everybody - here's a condensed form of it:<CENTER>
<TABLE BORDER=0 CELLSPACING=0 CELLPADDING=0>
<CAPTION ALIGN=TOP></CAPTION><TR><td width="55">w[%]</td><td width="120">d[g/ml]</td><td width="55">w[%]</td><td width="120">d[g/ml]</td><td width="55">w[%]</td><td width="120">d[g/ml]</td></tr><TR><td>1</td><td>1.0036</td><td>42</td><td>1.2591</td><td>84</td><td>1.4655</td></tr><TR><td>2</td><td>1.0091</td><td>44
</td><td>1.2719</td><td>85
</td><td>1.4689</td></tr><TR><td>4</td><td>1.0201</td><td>46</td><td>1.2847</td><td>86</td><td>1.4716</td></tr><TR><td>6</td><td>1.0312</td><td>48</td><td>1.2975</td><td>87
</td><td>1.4745</td></tr><TR><td>8</td><td>1.0427</td><td>50</td><td>1.3100</td><td>88</td><td>1.4773</td></tr><TR><td>10</td><td>1.0543</td><td>52</td><td>1.3219</td><td>89
</td><td>1.4796</td></tr><TR><td>12</td><td>1.0661</td><td>54</td><td>1.3336</td><td>90</td><td>1.4826</td></tr><TR><td>14</td><td>1.0781</td><td>56</td><td>1.3449</td><td>91
</td><td>1.4842</td></tr><TR><td>16</td><td>1.0903</td><td>58</td><td>1.3560</td><td>92</td><td>1.4873</td></tr><TR><td>18</td><td>1.1026</td><td>60</td><td>1.3667</td><td>93
</td><td>1.4886</td></tr><TR><td>20</td><td>1.1150</td><td>62</td><td>1.3769</td><td>94</td><td>1.4912</td></tr><TR><td>22</td><td>1.1276</td><td>64</td><td>1.3866</td><td>95
</td><td>1.4932</td></tr><TR><td>24</td><td>1.1404</td><td>66</td><td>1.3959</td><td>96</td><td>1.4952</td></tr><TR><td>26</td><td>1.1534</td><td>68</td><td>1.4048</td><td>96.5</td ><td>1.4972</td></tr><TR><td>28
</td><td>1.1666</td><td>70</td><td>1.4134</td><td>97</td><td>1.4988</td></tr><TR><td>30</td><td>1.1800</td><td>72</td><td>1.4218</td><td>97.5</td><td>1.5005</td></tr><TR><td>32
</td><td>1.1934</td><td>74</td><td>1.4298</td><td>98</td><td>1.5008</td></tr><TR><td>34</td><td>1.2071</td><td>76</td><td>1.4375</td><td>98.5</td><td>1.5044</td></tr><TR><td>36
</td><td>1.2205</td><td>78</td><td>1.4450</td><td>99</td><td>1.5066</td></tr><TR><td>38</td><td>1.2335</td><td>80</td><td>1.4521</td><td>99.5
</td><td>1.5091</td></tr><TR><td>40</td><td>1.2463</td><td>82</td><td>1.4589</td><td>100</td><td>1.5129</td></tr></table></CENTER>

[This message has been edited by Lagen (edited July 25, 2001).]

[Anthony]

stanfield, geoshities won't let you link directly to a file, you need to put the picture on a HTM or HTML page and then link to that.

[stanfield]

you'll have to type manualy the URL in your browser to acess my picture...

I did a search for "hydrometer" but nothing was found... I would like more info on this stuff plz... and how to make one if possible.

thanx for all !

[Lagen]

The instructions for the improvised hydrometer were posted here by PHILOU Zrealone on April 20 on the "No nitric acid" http://theforum.virtualave.net/ubb/smilies/smile.gif thread. It's on page 2.

For a professional hydrometer you would most probably have to go to a labware supplier. There are other kinds sold elsewhere (like the ones for car batteries, winemaking etc.) but they're not appropriate (do not range to 1.5+).

[Mick]

yeah, geoshits sucks

for those of you who wish to view the picture, just right click the link and select "Save Target As.."

and then save to your desktop or something and view it there

[stanfield]

Lagen, are you sure that no hydrometer (for car battery) on the market range to 1.5+ ? I will try to phone some companies to ask them about hydrometer...

If you found one in your contry, Lagen, you could buy and send it to me then, i'll re-buy it to you...sorry for my bad englsih http://theforum.virtualave.net/ubb/smilies/smile.gif

finally, if you find a labware supplier who can sell in France and who has (good)hydrometer, tell me too...

see ya !

[Lagen]

Now I have the guts of a car battery hydrometer in front of me, it only goes to 1.3. But obviously, one could cut the glass tube open, insert some more lead shot and re-seal it. All that one would need then, would be some good calibration solution (of a high and exactly known density).

The prices here are as follows:
w/out thermometer they all cost cca 8$:
1.0-2.0
1.4-1.6
w/thermometer they cost about twice as much:
1.45-1.50
1.40-1.50
1.50-1.60
I will look for some more sources if you're interested, I'm not sure if any would deliver to France, but surely you can get it close to your home too.

[nbk2000]

Hydgrometer is simple:

A large straw has one end heat sealed. Enough BBs to keep it floating upright about half way in the water are dropped in and a little wax dropped in and melted to hold the BBs in place.

Then calibrate. First in water at (I believe) 3*C (maximum density = 1). Then in another pure liquid of known denisty. The difference between the 2 is evenly divided and extrapolated to both ends of the scale. If you have a clear straw, you can mark the scale on a piece of paper that you then have sealed inside the straw.

Cost = $0

------------------
"The knowledge that they fear is a weapon to be used against them"

Go here (http://members.nbci.com/angelo_444/dload.html) to download the NBK2000 website PDF.

Go here (http://briefcase.yahoo.com/nbk2k) to download the NBK2000 videos.

[Chaos A.D]

ok, I havent tried it but from my knowlege of chem this process for making nitric acid wont work...

When you mix H2SO4 and CaNO3 the Ca precipitates out the SO4. you're left with a solution of NO3- and H+.

Nitric acid is a strong acid so the equation:
HNO3 --> H+ + NO3-
is not reversable.
That means the ions wont come together to form HNO3.

To make nitric acid, industry uses the Haber process... That uses NH3 and O2.

[Lagen]

http://theforum.virtualave.net/ubb/smilies/confused.gif To my best knowledge the Ca(NO3)2+H2SO4 process fails due to the CaSO4 being insoluble (even more so because of the SO4(2-) common ion) and forming a layer on the Ca(NO3)2 preventing it from being reacted. But to make dilute HNO3 w/ some Ca(NO3)2 in it, or to concentrate HNO3, this works pretty well. Please look here (http://theforum.virtualave.net/ubb/Forum1/HTML/000452.html). I have no idea what you meant by the equation thing. Please read in my earlier post here on the ions occuring in HNO3 soln. Why should the ions "combine"? There is no such thing as a HNO3 compound, except at 0 Kelvin. Even 100% HNO3 undergoes autoprotolysis, so some ions form. The H2SO4+KNO3 or H2SO4+NaNO3 rxn followed by distillation works well too and has been discussed to death.

[This message has been edited by Lagen (edited August 06, 2001).]

[BrAiNFeVeR]

Why would anyone want to make nitric acid ??? Isn't it less expensive to just buy it ? (less then 5€ /liter)
Making it yourself could only make sense if you could not buy it (which country do you live in ? UK ??)

Anyway, on making nitric acid you can find tons of information on the net (as it would be a nice experiment) I easily could give a pile of addresses, but hey, if you're too lazy too look for yourself ;-)

------------------
"Mess with me, and you'll end up with a .44 under your chin and your brains on the ceiling"

[Lagen]

I'm getting it for 1.5€/ltr. 65%. However I seriously doubt that you can buy any more than 90% HNO3 anywhere. Even that would have 2 bad points about it: for HE synthesis it would need distillation anyway, plus it would be a BIT more expensive than OTC H2SO4+fertilizer followed by distillation.

[Chaos A.D]

ok, just cause u got NO3- and H+ ions in the same solution it does not mean u have Nitric acid. If that was true you could make nitric acid out of any other acid.
What if you mix acetic acid(vinegar) and any nitrate...

CH3COOH + NaNO3 --> CH3COO- + Na+ + H+ + NO3-

Here there is H+ and NO3- in the same solution. This is not nitric acid. Nirctic acid is HNO3, not H+ and NO3-. Thats nitric after it has reacted with something.

Chemicaly it is not possible to produce nitric acid this was. Why would they(industries) make nitric using the costly Haber process if it was just this easy.

[FadeToBlackened]

HNO3 forms H+ and NO3- in solution.

[Lagen]

I hate to repeat myself, but if you really think it's impossible, I recommend you to make a dilute solution of calcium nitrate, and add (a calculated amount of) sulfuric acid to it. Filter out the precipitate and do some tests on the liquid, to prove it's not nitric acid. You will be surprised. It will colour your skin permanently yellow. What's that? Nitric acid. If you still don't believe me, then I can't help you any more. You just lost an easy way of making HNO3 (dilute, of course). Yes the Haber process is cheap to do on large scale, because there aren't any byproducts to throw away.

H+ + H2O tends to form H3O+ (oxonium cation).

[Lagen]

I have included a link to PHILOU's explanation of how this mechanism can be used for concentrating nitric acid - or producing more nitric acid in a nitric acid solution. As you just insisted it's impossible, I'm including a copy here for your kind perusal. Sorry to everyone else for the crappy post. <font face="Verdana, Arial" size="2">You can used this process with a little modification to produce nitric acid more concentrated than the original one without need of a distillator:
If you have concentrated H2SO4 more than your HNO3- and only if you have it!
-Mix hot medium conc or conc HNO3 and saturate it with Ca(NO3)2 by strong stirring.Heat it long to be sure it is saturated. Cool it until ambiant temp some Ca(NO3)2 cristals will come out proof of the saturation, filter them. In the clear solution can now be inserted under stirring the conc H2SO4 ( it has to be in minor quantity as compare to Ca(NO3)2 dissolved).
Then in your solution you will have a cloud of CaSO4 pure HNO3 more concentrated and some Ca(NO3)2 leftover.</font>

[Chaos A.D]

ok fine, if you actually did it and it works. I havent tried it but Im just going by what I learned it chem.

[Lagen]

Double replacement with precipitation will be one of the most common schemes you will encounter in "preparative" chemistry. In this one you can practically get just "dilute" HNO3 because the mass of the precipitate is comparatively large and it will take up a lot of space in the reaction mix. If you try to make more concentrated HNO3, you will end up with an almost solid mass of CaSO4 from which the HNO3 is nearly impossible to separate. Most precipitated stuff has an initial density of 0.1-0.05, you can figure out what that means in this case...