Sciencemadness Discussion Board - NH3, H2SO4, and HNO3 from fertilizer

NH3, H2SO4, and HNO3 from fertilizer

Melgar - 25-2-2010 at 16:17

Ammonium sulfate (NH4)2-SO4 is a common industrial by-product that's widely used as an acidifying nitrogen fertilizer. I can get it for $6 for a 20-pound bag. Heating it beyond its melting point yields NH3 and Ammonium bisulfate, which is quite acidic. Heating ammonium bisulfate causes an equilibrium to be set up with NH3, H2O, and SO3 on one side and NH4HSO4 on the other side. Raising temperature and lowering pressure puts the equation more in favor of the gases.

Now, platinum will catalyze the oxidation of ammonia to NO2. So what I'm thinking is, once the gases have started forming, I can start slowly pumping in oxygen and run some current through a thin platinum wire I would have placed in the sealed flask prior to heating it. Once the wire gets hot enough to catalyze the oxidation of ammonia, the exothermic reaction should keep it that hot for some time. However, I'm a bit worried that the NO2 would form nitric acid, react with the ammonia to form ammonium nitrate, then immediately decompose. Of course, that may be a minor side reaction since NH3 would preferentially react with the sulfuric acid that would also be formed. Simultaneously distilling off the nitric acid somehow could also drive the reaction in the preferential direction. Anyway, the end product would be a roughly equimolar mixture of concentrated nitric and sulfuric acids, and we all know what we can do with that!

Also, when making ETN, is there any danger in neutralizing the acid with NH4OH instead of baking soda? In that case, the solution could be evaporated and the excess base would just evaporate too. The (NH4)2-SO4 would precipitate first, then the NH4NO3 could be dried later. This would have the obvious advantage of providing a source of ammonium nitrate, I just hope there isn't also a chance of it blowing up on me. I don't think so because it's in an aqueous solution, but that's what I thought about NI3, and I turned out to be quite wrong about that.

User - 25-2-2010 at 17:11

Catalytic from NH3 to NO2/x is a process that definitely needs controlled conditions.
If your planning to do this with acceptable yields in mind you should consider pressure/temp control and high a high efficiently absorption system. Don't underestimate the reactions you are proposing.
I'd say show some nice reports.

[Edited on 26-2-2010 by User]

per.y.ohlin - 25-2-2010 at 19:24

Ammonium nitrate is an extremely insensitive explosive. I don't know of any way to reliably detonate pure AN at room temperature, and ammonium nitrate based explosives require a large booster to detonate. If you were to neutralize with ammonia, the ammonium nitrate would not be a significant detonation hazard. I can't think of any possible side reactions when neutralizing with ammonia, so I believe it would be safe. As for a source of ammonium nitrate, ice packs seem easier.

Melgar - 25-2-2010 at 20:26

Quote: Originally posted by per.y.ohlin

Ammonium nitrate is an extremely insensitive explosive. I don't know of any way to reliably detonate pure AN at room temperature, and ammonium nitrate based explosives require a large booster to detonate. If you were to neutralize with ammonia, the ammonium nitrate would not be a significant detonation hazard. I can't think of any possible side reactions when neutralizing with ammonia, so I believe it would be safe. As for a source of ammonium nitrate, ice packs seem easier.

Well, the only reason I was concerned was because the AN would be combined with ETN for a brief period, and there's always a chance they might not separate completely. I agree, the chances that would cause problems are remote, but I just wanted to make sure. I was mostly interested in that question because by neutralizing with ammonia, you can dump in as much as you want and not have to worry about rinsing it off later like you do with baking soda, since it'll evaporate on its own. The fact that you get ammonium nitrate when you evaporate the liquid is a secondary benefit, but a benefit nonetheless.

Also, I spilled two beakers over because of all the bubbles from the baking soda, so I was looking for another base to use. I was going to use sodium hydroxide, but didn't want to use something that strong in case it might damage the ETN. Another nice thing about neutralizing with ammonia is that any nitric acid fumes are immediately visible in the vicinity when they combine with the ammonia vapor, so no more accidentally getting a whiff of that stuff. The microscopic crystals that form do set off smoke detectors though.

franklyn - 26-2-2010 at 03:37

And how is this magic to be performed ? Most of this supposed process is
not detailed in the brief outline. If this is the extent to which it has been
thought out , it is not even a half baked idea.
The one pot preparation for nitric acid is a nitrate salt boiled in sulfuric acid.
Nitric acid made from the Oxidation of Ammonia is a multiple step process
not amenable to a single reaction procedure as suggested by this lightbulb
reactor. Ammonium Sulfate gives up the first Ammonia molecule at 100 ºC
Nitric acid boils at 120 °C before Ammonium Bisulfate melts at 147 ºC, and is
completely dissociated before further decomposition of Ammonium Sulfate
occurs at over 280 ºC. Oxidation of ammonia by a mantle of heated platinum
gauze occurs around 850 ºC. The most important issue is not even addressed ,
how is the necessary oxygen to be supplied.

Nitric acid is made from the oxidation of Ammonia in three stages. The first
step is the oxidation of ammonia gas ( NH3 ) with air to form nitric oxide ( NO )
by passing through a screen of platinum-rhodium catalyst at high temperature
and pressure to achieve a high conversion efficiency. The nitric oxide is cooled
and further oxidized with more air to form nitrogen dioxide ( NO2 ). This is
absorbed in water forming nitric acid ( HNO3 ) and nitrous acid ( HONO ) which
itself is further oxidized to nitric acid. Temperature and conditons at each step
require the contents to be segregated in their own separate reaction chamber.
This is described in greater detail below along with references and a process
patent attached.

. . . 4 (NH4)2SO4 => 4 NH4HSO4 + 4 NH3
Ammonia oxidation
1 ) . 4 NH3 + 5 O2 => 6 H2O + 4 NO
Nitric Oxide oxidation
2 ) . .4 NO + 2 O2 => 4 NO2
Nitrogen Dioxide absorption
3 ) 4 NO2 + 2 H2O => 2 HNO3 + 2 HONO
. . . 2 HONO + O2 => 2 HNO3

The overall reaction is _ 4 NH3 + 8 O2 => 4 H2O + 4 HNO3

Twice the molar volume of oxygen to that of ammonia is needed.
The equal molar volume of water also produced without a means of drying
will limit concentration so it cannot exceed ~ 75 % at best.

_________________________________________________

References
http://www.backyardchem.com/nitric-acid.php
http://chemistry.slss.ie/resources/downloads/ch_cw_nitricand...
http://www.platinummetalsreview.com/pdf/pmr-v11-i1-002-009.p...

Production Process Nitric acid is commercially available in two forms: weak (50 to 70
percent nitric acid) and concentrated (greater than 95 percent nitric acid). Different
processes are required to produce these two forms of acid. For its many uses, weak
nitric acid is produced in far greater quantities than is the concentrated form. Virtually
all commercial production of weak nitric acid in the United States utilizes three common
steps: (1) catalytic oxidation of ammonia ( NH3 ) to nitric oxide ( NO ), (2) oxidation of
nitric oxide with air to nitrogen dioxide ( NO2 ), and (3) absorption of nitrogen dioxide in
water to produce "weak" nitric acid.

The first step of the acid production process involves oxidizing anhydrous ammonia over
a platinum-rhodium gauze catalyst to produce nitric oxide and water. The exothermic
reaction occurs as follows: 4 NH3 + 5 O2 -> 4 NO + 6 H2O + heat. This extremely rapid
reaction proceeds almost to completion, evolving 906 kilojoules per mole ( kJ/mole )
(859 British thermal units per mole [Btu/mole]) of heat. Typical ammonia conversion
efficiency ranges from 93 to 98 percent with good reactor design. Air is compressed,
filtered, and preheated by passing through a heat exchanger. The air is mixed with
vaporized anhydrous ammonia and passed to the converter. Since the explosive limit of
ammonia is approached at concentrations greater than 12 mole percent, plant operation
is normally maintained at 9.5 to 10.5 mole percent. In the converter, the ammonia-air
mixture is catalytically converted to nitric oxide and excess air. The most common
catalyst consists of 90 percent platinum and 10 percent rhodium gauze constructed
from squares of fine wire. Up to 5 percent palladium is used to reduce costs.

Operating temperature and pressure in the converter have been shown to have an
influence on ammonia conversion efficiency. Generally, reaction efficiency increases
with gauze 8 temperature. Oxidation temperatures typically range from 750 ºC to 900 ºC .
Higher catalyst temperatures increase reaction selectivity toward NO production, while
lower catalyst temperatures are more selective toward less useful nitrogen ( N2 ) and
nitrous oxide (N2O). The high-temperature advantage is offset by the increased loss
of the precious metal catalyst. Industrial experience has demonstrated and the industry
has generally accepted conversion efficiency values of 98 percent for atmospheric
pressure plants at 850 ºC and 96 percent for plants operating at 0.8 megaPascals
(MPa) (8 atmospheres [atm]) and 900 ºC. As mentioned earlier, the ammonia oxidation
reaction is highly exothermic. In a well-designed plant, the heat byproduct is usually
recovered and utilized for steam generation in a waste heat boiler. The steam can be
used for liquid ammonia evaporation and air preheat in addition to nonprocess plant
requirements. As higher temperatures are used, it becomes necessary to capture
platinum lost from the catalyst. About 300 milligrams per ton of nitric acid is lost this
way. Consequently, a platinum recovery unit is frequently installed on the cold side of
the waste heat boiler. The recovery unit, composed of ceramic-fiber filters, is capable
of capturing 50 to 75 percent of the platinum lost by Nitric Oxide oxidation.

The nitric oxide formed during the ammonia oxidation process is cooled in the cooler /
condenser apparatus, where it reacts noncatalytically with oxygen to form nitrogen
dioxide and its liquid dimer, dinitrogen tetroxide. The exothermic reaction, evolving 113
kJ/mole 4 (107 Btu/mole), proceeds as follows: 2 NO + O2 -> 2 NO2 + heat. This slow,
homogeneous reaction is highly temperature- and pressure-dependent. Lower
temperatures, below 38 ºC, and higher pressures, up to 800 kilopascals (kPa) (8 atm),
ensure maximum production of NO2 and minimum reaction time. Furthermore, lower
temperatures and higher pressures shift the reaction to the production of NO2 ,
preventing the reverse reaction ( dissociation to 2 NO and O2 ) from occurring.

The final step for producing weak nitric acid involves the absorption of NO2 and N2O4
in water to form nitric acid ( as NO2 is absorbed, it releases gaseous NO ). The rate
of this reaction is controlled by three steps: (1) the oxidation of nitrogen oxide to NO2
in the gas phase, (2) the physical diffusion of the reacting oxides from the gas phase
to the liquid phase, and (3) the chemical reaction in the liquid phase. The exothermic
reaction, evolving 135 kJ/mole (128 Btu/mole), proceeds as follows: 3 NO2 (g) + H2O
-> 2 HNO3 (aq) + NO(g) + heat , and also 3 N2O4 (liq) + 2 H2O => 4 HNO3 + 2 NO(g)
The absorption process takes place in a stainless steel tower containing numerous
layers of either bubble cap or sieve trays. The number of trays varies according to
pressure, acid strength, gas composition, and operating temperature. Nitrogen dioxide
gas from the cooler/condenser effluent is introduced at the bottom of the absorption
tower, while the liquid dinitrogen tetroxide enters at a point higher up the tower.
Deionized process water is added at the top, and the gas flows countercurrent to both
liquids. Oxidation occurs in the free space between the trays, while absorption takes
place in the trays. Because of the high order of the oxidation process in absorbers,
roughly one-half the volume of the absorber is required to absorb the final 3 percent
of nitrogen oxide gas concentration. Because lower temperatures are favorable for
maximum absorption, cooling coils are placed in the trays. Nitric acid in concentrations
of 55 to 65 percent is withdrawn at the bottom of the tower. Secondary air is used to
improve oxidation in the absorption tower and to bleach remaining nitrogen oxides from
the product acid. Absorption efficiency is further increased by utilizing high operating
pressure in the absorption process. High-pressure absorption improves efficiency and
increases the overall absorption rate. Absorber tail gas is reheated using recovered
process heat and expanded through a power recovery turbine. In a well designed plant,
the exhaust gas turbine can supply all the power needed for air compression with
excess steam available for export.

.

Attachment: Manuf of Nitric Acid from Ammonia US1872638.pdf (462kB)
This file has been downloaded 940 times

hissingnoise - 26-2-2010 at 04:16

Quote:

The one pot preparation for nitric acid is a nitrate salt boiled in sulfuric acid.

That's it in a nutshell!
For the amateur there's really no other practical route to strong HNO3.
Use K or NaNO3 - the ammonium salt produces a diluted HNO3 because the NH4 ion is oxidised by hot H2SO4.

Melgar - 26-2-2010 at 04:57

Quote:

[quote=172624&tid=13434&author=franklyn]
And how is this magic to be performed ? Most of this supposed process is
not detailed in the brief outline. If this is the extent to which it has been
thought out , it is not even a half baked idea.

The reaction can be carried out in ways other than exactly how nitric acid plants do it, you know. See:
http://www.chem.umn.edu/services/lecturedemo/info/catalysis_...

Quote:

The one pot preparation for nitric acid is a nitrate salt boiled in sulfuric acid.

Yeah, I know. I've done that reaction already.

Quote:

Nitric acid made from the Oxidation of Ammonia is a multiple step process
not amenable to a single reaction procedure as suggested by this lightbulb
reactor.

I just have to generate the NO. The rest takes care of itself, provided O2 and water are present.

Quote:

Ammonium Sulfate gives up the first Ammonia molecule at 100 ºC
Nitric acid boils at 120 °C before Ammonium Bisulfate melts at 147 ºC, and is
completely dissociated before further decomposition of Ammonium Sulfate
occurs at over 280 ºC. Oxidation of ammonia by a mantle of heated platinum
gauze occurs around 850 ºC.

Well, your HNO3 boiling point is wrong, you're thinking of the water azeotrope which doesn't apply thanks to the sulfuric acid present. Without water the boiling point is 83ºC.

Also, when ammonium bisulfate decomposes, did you ever think about what it decomposes into? Well, turns out it decomposes into SO3, NH3, and H2O. SO3, as we all know, reacts with water to make sulfuric acid, and NH3 reacts with sulfuric acid to make ammonium bisulfate.

Quote:

The most important issue is not even addressed ,
how is the necessary oxygen to be supplied.

How is that the most important issue? Oxygen is readily available in the form of air, oxygen tanks, or chlorine bleach mixed with hydrogen peroxide. I think it's a lot more important to make sure the NH3 doesn't become N2, at least for the most part.

Quote:

Nitric acid is made from the oxidation of Ammonia in three stages. The first
step is the oxidation of ammonia gas ( NH3 ) with air to form nitric oxide ( NO )
by passing through a screen of platinum-rhodium catalyst at high temperature
and pressure to achieve a high conversion efficiency. The nitric oxide is cooled
and further oxidized with more air to form nitrogen dioxide ( NO2 ). This is
absorbed in water forming nitric acid ( HNO3 ) and nitrous acid ( HONO ) which
itself is further oxidized to nitric acid. Temperature and conditons at each step
require the contents to be segregated in their own separate reaction chamber.
This is described in greater detail below along with references and a process
patent attached.

. . . 4 (NH4)2SO4 => 4 NH4HSO4 + 4 NH3
Ammonia oxidation
1 ) . 4 NH3 + 5 O2 => 6 H2O + 4 NO
Nitric Oxide oxidation
2 ) . .4 NO + 2 O2 => 4 NO2
Nitrogen Dioxide absorption
3 ) 4 NO2 + 2 H2O => 2 HNO3 + 2 HONO
. . . 2 HONO + O2 => 2 HNO3

The overall reaction is _ 4 NH3 + 8 O2 => 4 H2O + 4 HNO3

Twice the molar volume of oxygen to that of ammonia is needed.
The equal molar volume of water also produced without a means of drying
will limit concentration so it cannot exceed ~ 75 % at best.

Did you catch the part about simultaneously producing H2SO4? That's a most excellent drying agent, you know. And even though the industrial process is complicated, it's that way mainly for three reasons: to maximize efficiency, provide a continuous process, and avoid releasing nitrogen oxides into the environment. Efficiency is not that critical to me, I'm not producing enough to generate significant pollution, and I don't care about having a continuous process, so no need for all those complicated optimizations.

[Edited on 2/26/10 by Melgar]

User - 26-2-2010 at 05:05

I absolutely adore your optimism.

Quote:

And how is this magic to be performed ? Most of this supposed process is
not detailed in the brief outline. If this is the extent to which it has been
thought out , it is not even a half baked idea.

The reaction can be carried out in ways other than exactly how nitric acid plants do it, you know. See:
http://www.chem.umn.edu/services/lecturedemo/info/catalysis_...

This is a demo that is well known, it does not produce any significant amounts for sure.

[Edited on 26-2-2010 by User]

Melgar - 26-2-2010 at 05:06

Quote: Originally posted by hissingnoise

Quote:

The one pot preparation for nitric acid is a nitrate salt boiled in sulfuric acid.

Not quite. Ammonium nitrate decomposes at high temperatures, sulfuric acid or not. The NO3 ion is what oxidizes the NH4, which sucks because it also lowers your yields. Thanks for the advice though, I'm just throwing out an idea for potentially cheaper materials since both sulfuric acid and nitrate salts have been fairly expensive and/or hard to find lately. Considering I can buy a 20-pound bag of ammonium sulfate for less than the price of a liter of sulfuric acid, I think it could be a viable process.

hissingnoise - 26-2-2010 at 06:08

Quote:

Not quite. Ammonium nitrate decomposes at high temperatures, sulfuric acid or not. The NO3 ion is what oxidizes the NH4, which sucks because it also lowers your yields.

Just to be clear, NH4NO3 begins to decompose above ~174*C but when in excess, H2SO4 decomposes NH4NO3 above ~ 100*C.

Melgar - 26-2-2010 at 07:45

Quote: Originally posted by hissingnoise

Quote:

Not quite. Ammonium nitrate decomposes at high temperatures, sulfuric acid or not. The NO3 ion is what oxidizes the NH4, which sucks because it also lowers your yields.

Just to be clear, NH4NO3 begins to decompose above ~174*C but when in excess, H2SO4 decomposes NH4NO3 above ~ 100*C.

In other words H2SO4 catalyzes the decomposition of NH4NO3 at lower-than-usual temperatures? That's what I'd expect given H2SO4's propensity to push reactions in whatever direction favors the formation of water.

Quote:

This is a demo that is well known, it does not produce any significant amounts for sure.

Well, considering the platinum is releasing enough energy to stay red hot, and considering the reaction on platinum favors the production of NO, then either a significant amount of NO is being produced or that's a perpetual energy machine there. Either way it's a winner.

[Edited on 2/26/10 by Melgar]

quicksilver - 26-2-2010 at 08:37

NH4NO3 is still sold in 15.5-0-0 as "calcium nitrate". Why attempt such a thing if the material is freely available?

hissingnoise - 26-2-2010 at 09:08

Whether hot H2SO4 is acting as oxidiser or catalyst is a moot point in this case.
But if you have difficulty getting H2SO4 or Na/KNO3 then you have a real problem because the other "method" you've outlined will mean a hell of a lot of time and effort for very little product.
And prices for KNO3-H2SO4 are unlikely to go beyond anyone's means. . .

Melgar - 26-2-2010 at 11:12

Quote: Originally posted by quicksilver

NH4NO3 is still sold in 15.5-0-0 as "calcium nitrate". Why attempt such a thing if the material is freely available?

Quote:

Whether hot H2SO4 is acting as oxidiser or catalyst is a moot point in this case.
But if you have difficulty getting H2SO4 or Na/KNO3 then you have a real problem because the other "method" you've outlined will mean a hell of a lot of time and effort for very little product.
And prices for KNO3-H2SO4 are unlikely to go beyond anyone's means. . .

Honestly, it's largely just to see if I can do it, and what the yields are like if I can. Plus, if it goes really well then it wouldn't be much more work than making HNO3 from H2SO4 and KNO3, only instead of using H2SO4 I'd be getting it, as well as ammonia gas. But all in all, I think it's a cool exercise in theoretical chemistry, and I already have all the stuff except for a platinum/rhodium thermocouple wire that cost me $30. We'll see how it goes!

hissingnoise - 26-2-2010 at 12:07

Blind optimism - that's the spirit!

chief - 26-2-2010 at 12:20

There was an old patent here around after which NH4NO3 was the _preferred_ nitrate for heating with H2SO4, since everything melts nicely and the reaction is more complete thereby ... ; also the paper said the HNO3 were quite clean from Water ...

The H2SO4 would come from electrolysis of a sulfate ...

hissingnoise - 26-2-2010 at 13:02

My experience was that RFNA came over at the start but this gradually lightened as distillation progressed and near the end the distillate consisted of water, essentially.
I don't remember (I may have) if I used an excess of H2SO4.
An NH4NO3/H2SO4 mix would be preferred for nitration of say, erythritol but I won't use it for HNO3. . .

Quote:

The H2SO4 would come from electrolysis of a sulfate ...

That electrolysis needs an inert anode like Pt!
Can't you just use drain-cleaner?

[Edited on 26-2-2010 by hissingnoise]

franklyn - 26-2-2010 at 19:59

@ Melgar

You have not yet given a lucid description of this imagined process you refer to
apart from randomly citing unrelated reactions in no coherent sequence without
stipulating reaction constraints.
If heated in confinement , Ammonium Sulfate dissociates as a high pressure
supercritical gas of steam , Sulfur trioxide and Ammonia , which may dissociate into
nitrogen and hydrogen itself , reducing Sulfur trioxide into Sulfur dioxide and steam.
How to enter air / oxygen into this bomb and vent the vessel remains unexplained
a need apparently not understood by you. What becomes of the supposed ( NOx )
formed as it accumulates at high temperature and pressure since it surely must
decompose. But it gets worse , the catalytic oxidation process is necessarily anhydrous
and the Sulfur oxides are not going to just stand by and only observe. Feeding this
stew into a platinum wire mesh is not going to yield nitric oxide ( NO ) in any amount
that is not hot off the catalyst , since the tendency will be to form water and free
nitrogen. That will certainly put a stop to the pressure buildup - in an instant.
Ignoring that temperature and pressure are all that matters in thermal equilibrium ,
in just the precisely needed optimal amounts at each separate stage , is not a plan.

(NH4)2SO4 => 2 H2O + SO2 + N2 + 2 H2 , <- this is the likely result.

Imperial Chemical Industries in England , developed a high temperature process for
conversion of Ammonium bisulfate to produce Ammonia and Sulfur trioxide. Ammonia
is further oxidized under these oxidative conditions to nitrogen and water , which is
added to the Sulfur trioxide producing Sulfuric acid , the end product.

a similar process patent US20080056983 is covered here , the prior art and research
cited is what is important , a sampling is below _
http://www.freshpatents.com/Process-for-producing-ammonia-an...

U.S. Pat. No. 4,081,515 (Gruhier) carries out the decomposition of ammonium bisulfate
by heating it to 400 ºC. The heating is carried out without a catalyst present or with
copper, molybdenum, or tungsten catalyst present. The primary product is sulfur dioxide,
which is not desired in the recovery of ammonia and sulfuric acid.

U.S. Pat. No. 3,282,646 (Bonfield) discloses the production of ammonia and sulfur
dioxide from ammonium sulfate. The ammonium sulfate was heated to 250 ºC. to
drive off ammonia and produce ammonium bisulfate. The temperature was increased
to 450 ºC. and carbon monoxide, hydrogen sulfide, hydrogen, or nitrogen was bubbled
through the molten salt. The reaction produced ammonia and sulfur dioxide, the latter
being an undesirable byproduct.

Halstead (J. Appl. Chem., 1970, 20:129-132) describes the decomposition of ammonium
sulfate at 400 ºC. The first reaction converts ammonium sulfate to ammonium bisulfate
and ammonia. In a second reaction, the ammonium bisulfate dehydrates to form a water
molecule and ammonium pyrosulfate, (NH4)2S2O7. Further heating of the ammonium
pyrosulfate to form sulfur dioxide and nitrogen, which are not desired products from
ammonium sulfate, is also carried out.

Kiyoura and Urano (Ind. Eng. Chem. Process Des. Develop., 1970, 9(4):489-494) describes
the thermal decomposition of ammonium sulfate to ammonia, sulfur dioxide, sulfur trioxide
and other gases. Kiyoura states that the simple decomposition of ammonium bisulfate to
sulfuric acid and ammonia is not an adequate mechanism to describe the thermal
decomposition reaction.

Liske et al. (Journal of Hazardous Substance Research, 2000, 2:8.1-8.17) teach the
complete thermal decomposition of ammonium sulfate to several gaseous products.
This combustion process does not recover ammonia and sulfuric acid.

Dugger et al. (Ammonium Sulfate Decomposition, 1955, RMO-2036, United States
Atomic Energy Commission) describes the thermal decomposition of ammonium sulfate
in the presence of zinc oxide. The reaction products are ammonia and sulfur dioxide.

.

Melgar - 27-2-2010 at 11:38

Thanks for all the patent links, even though the tone of your responses indicates you aren't really trying to be helpful. At sub-plasma temperatures, I doubt I have to worry about a significant amount of ammonia disassociating into N2 and H2, though I probably do have to worry about NH3 + O2 = N2 + H2O.

I drew a picture of my proposed setup, which is basically just a modified vacuum distillation set, not a bomb. I admit I wasn't sure what it should look like at first, which is why I was soliciting advice, but I think I have an idea now. The vacuum distillation setup will be run in a way such that no liquids are carried over to the collection flask, just gases. And the water would collect the nitrates that form, probably in the form of nitric acid and ammonium nitrate. Potassium hydroxide could go in there too if someone wanted to make KNO3.

chemsetup.png - 90kB

hissingnoise - 27-2-2010 at 14:02

What can one say - it's certainly artistic. . .

quicksilver - 28-2-2010 at 12:28

@Melgar :
your O2 and vacuum may need some alteration or you'll just loose the O2 to no avail.

@ hissingnoise :
I've not seen a flask set up for electrodes. IS there such a commercial thing? I thought about a 3 neck and a stopper but the heat may screw even a polypropylene stopper up in short order. I suppose you could use a tweaked Titrator but that's serious money. I suppose if you kept the whole thing on a really small scale you could get away with a stopper....

[Edited on 28-2-2010 by quicksilver]

hissingnoise - 28-2-2010 at 13:05

Quote:

. I suppose if you kept the whole thing on a really small scale you could get away with a stopper....

The wire would have to sealed into the stopper to be dependable, and that would call for glassworking skills.
And dependable, in this case just means unlikely to leak.
For one thing, with the temperatures needed for the reaction, cold water in the condenser could be problematic. . .
Also, I'd say franklyn has it just about right!

[Edited on 28-2-2010 by hissingnoise]

densest - 28-2-2010 at 19:04

Quote:

The wire would have to sealed into the stopper to be dependable, and that would call for glassworking skills.
And dependable, in this case just means unlikely to leak.

I've successfully made mostly-tight seals for wires to flasks two ways, as long as the pressure differential isn't more than .1 bar or so. One is to grind grooves down the taper of a glass standard taper stopper, then run the wires in the grooves, wrapping them in plumber's teflon tape. The other way is to carefully put wires inside teflon tubing and heat the assembly until the teflon softens but does not decompose - it won't really melt. Fill out to a suitable size with larger diameter teflon tubing or tape. Compress each wire into a hole in a teflon stopper with a hollow tool until the plastic firmly solidifies. A zig-zag or partial loop in the wire helps prevent the wire from sliding which would enlarge the hole. Neither method will make a gastight seal since teflon will not adhere to the metal. It is sufficient to keep liquid inside if not pressurized or at depth, mostly.

I will assert that I think the proposed method would not have a good yield. If the (NH4)2SO4 is really really cheap or HNO3 is really unavailable, it might work. I wonder about miscellaneous reactions between NH3 and SOx where NH3 reduces the SO3 to SO2? Perhaps some V2O5 heated in the gas + oxygen stream might help? Or would it just confuse things?

One might try to separate the SO3 from the NH3:

nh42so4.gif - 3kB

chief - 1-3-2010 at 01:22

Quote: Originally posted by hissingnoise

That electrolysis needs an inert anode like Pt!
Can't you just use drain-cleaner?

[Edited on 26-2-2010 by hissingnoise]

No such drain-cleaner in germany ... ; anyhow Sulfates are everywhere, and cheaply ...
==> Carbon rods for welding should do it, or not ? After all they even stand the far more agressive chlorate-electrolyses-environment for a while ...

For distillation of HNO3 from Nitrate it also would be sufficient to boil down a H2SO4-solution with residual sulfate, as it doesn't harm the HNO3-generation ...

So the input for the process would be some sulfate-fertilizer and NH4NO3 ... ; only: What about the electrolysis of ammonium-sulfates ?
==> Might that become dangerous (reminding of the NH4Cl-electrolytical dangers) ?
==> Then the NH4NO3 would have to be converted with soda to NaNO3 + NH3 ...

Melgar - 1-3-2010 at 02:10

Hmm, thanks for the input.

Quicksilver - Obviously I'd only have one valve open at a time. The vacuum valve would probably only be open for a short time at the beginning, and maybe later to get rid of inert gases if they're building up.

Densest - you have a good point with the possible SO3 + NH3 = SO2 + N2 + H2O reaction. However, I am hoping that SO3 reacts preferentially with H2O and condenses out fairly quickly. I'm also counting on the fact that of the decomposition products, only NH3 is a gas at room temperature. I think I'll also rearrange my distillation setup so that the condensation in the condenser runs back to the reflux column. As for the wires, they are very thin; .005 inch diameter. They can run right through the silicone grease at any joint. And I only need to heat them once at the beginning. The oxidation reaction is quite exothermic, and this heat keeps the wire glowing for as long as there's ammonia and O2 present.

Chief - Actually, you'd only need ammonium sulfate fertilizer, which is (NH4)2SO4. I did make ammonium nitrate starting with just ammonia. The yield wasn't anything great, but this was uncontrolled in open air.

hissingnoise - 1-3-2010 at 03:46

Quote:

As for the wires, they are very thin; .005 inch diameter. .

-Melgar, at .005" the surface area will be so small you'd need a very long wire to ensure a working surface-area big enough to produce usable quantities of NO.
You might as well face it - you're grasping at straws. . .
-Chief, I've seen 96%H2SO4 D/C on eBay UK - eBay Deutschland might have it too.

[Edited on 1-3-2010 by hissingnoise]

watson.fawkes - 1-3-2010 at 05:04

Quote: Originally posted by hissingnoise

at .005" the surface area will be so small you'd need a very long wire to ensure a working surface-area big enough to produce usable quantities of NO.

I must agree; that surface area is mighty small. There are only two catalysts used in industry for catalytic oxidation of ammonia for nitric acid production: platinum and cobalt oxide. I did a review of patent literature last year on the subject, and most of it was about how to conserve and/or recover the platinum catalyst. Another good piece of it was how to increase surface area. My suggestion is to figure out how to use cobalt oxide on a small scale. It likely means a tube furnace with good temperature control and a packed-column catalyst bed. The desired reaction only happens within a certain temperature range, and you've got both external heating and an exothermic reaction to manage. Ceramic beads are used as catalyst supports ordinarily. Some creativity about how to make this (with smaller dimensions) or an equivalent would be a real contribution.

Melgar - 1-3-2010 at 11:20

Quote: Originally posted by watson.fawkes

Quote: Originally posted by hissingnoise

at .005" the surface area will be so small you'd need a very long wire to ensure a working surface-area big enough to produce usable quantities of NO.

Thank you very much for the suggestion. Unfortunately, it appears sulfur has a tendency to poison cobalt oxide catalysts. Looks like it's back to platinum then, although this time I'll try electroplating it onto titanium wire to increase the surface area. That should work, right?

hissingnoise - 1-3-2010 at 13:16

Attaboy! With determination like yours, everything's possible!

watson.fawkes - 1-3-2010 at 14:18

Quote: Originally posted by Melgar

Thank you very much for the suggestion. Unfortunately, it appears sulfur has a tendency to poison cobalt oxide catalysts.

So take the sulfur out. Use a phosphate-ammonium fertilizer. Or better yet, if you're after nitric acid, just by a tank of liquid ammonia at said fertilizer store.

Quote:

Looks like it's back to platinum then, although this time I'll try electroplating it onto titanium wire to increase the surface area. That should work, right?

If you can get a reliable method of doing that, please do publish how.

Melgar - 2-3-2010 at 15:15

Quote: Originally posted by watson.fawkes

Quote: Originally posted by Melgar

Thank you very much for the suggestion. Unfortunately, it appears sulfur has a tendency to poison cobalt oxide catalysts.

So take the sulfur out. Use a phosphate-ammonium fertilizer. Or better yet, if you're after nitric acid, just by a tank of liquid ammonia at said fertilizer store.

Yesterday I found out that my local hydroponics store actually has bags of unblended fertilizer that they use to make their own blends for customers, they just don't normally offer it like that for sale. Anyway, I got a pound of calcium nitrate fertilizer, then mixed it with H2SO4 drain opener and distilled the result, giving me 200 mL or so of fuming white nitric acid.
So, I'm not quite as interested in the ammonium sulfate route to nitric acid anymore, although it would still be helpful to have a means of turning a sulfate salt into sulfuric acid. Can you actually get tanks of anhydrous ammonia at ag stores? I thought they put a stop to that because meth cooks were using it.

Quote:

Looks like it's back to platinum then, although this time I'll try electroplating it onto titanium wire to increase the surface area. That should work, right?

If you can get a reliable method of doing that, please do publish how.

Heh, I take it it's easier said than done? I'm assuming the first step is dissolving the platinum in aqua regia, but that's proving to be hard enough by itself.

[Edited on 3/2/10 by Melgar]

watson.fawkes - 2-3-2010 at 15:30

Quote: Originally posted by Melgar

Can you actually get tanks of anhydrous ammonia at ag stores? I thought they put a stop to that because meth cooks were using it.

Well, then don't be a meth cook. And don't look like one or act like one either.

Melgar - 2-3-2010 at 20:29

Quote: Originally posted by watson.fawkes

Quote: Originally posted by Melgar

Can you actually get tanks of anhydrous ammonia at ag stores? I thought they put a stop to that because meth cooks were using it.

Well, then don't be a meth cook. And don't look like one or act like one either.

No, I mean, I think you have to prove you're a farmer to buy it now. But I don't think it's used by the farmers near where I live anyway. My dad's a farmer, and he mostly gets urea or a blend, although he has gotten ammonium nitrate before. Ammonium sulfate was interesting to me because it's really cheap due to it being an industrial byproduct, whereas the phosphates are fairly expensive. Plus, I need a better source of sulfuric acid that doesn't have all that crap that's in the drain opener. Heating the ammonium bisulfate to drive off ammonia seems like it'd work, since presumably any remaining sodium bisulfate would precipitate out at room temperature.

Also, I'm thinking of having a continuous stream of air blowing over the molten sodium bisulfate, through a condenser, then over the catalyst. The air would then be bubbled through a solution of potash (potassium carbonate) where the NO2 would form nitric acid, then potassium nitrate. So hopefully, that's three useful chemicals (NH3, KNO3, and H2SO4) from one not-so-useful one. Of course, I could always just use the ammonia for other stuff if I didn't need any nitrate salts at the time.

quicksilver - 3-3-2010 at 09:24

C.A.N. 15.5-5-5 has generally replaced ammonium nitrate in packaged purchases. I live in a large farm area. AN generally is available by open truck delivery to very few farmers who actually NEED AN. Berry and Tobacco benefit from straight AN bu "CAN" works for most all needs so that's what being sold. One provider is able to get AN but no one wants it anymore. We even have some bagged C.A.N. now but some places are doing the "Be Aware for America" thing from the Fertilizer Institute (YES - there is such a thing) and folks just don't like having their name on some list.
To have a delivery (generally) your purchase has to be in line with your planting acreage and coincide with time of year nitration (for most that Nov- Feb) Then you have to have it stored in a special area! NO ONE puts up with that crap. So no one orders AN when "CAN" works just fine.

hissingnoise - 3-3-2010 at 09:54

The CAN I know quicksilver, is AN prills with a calcium carbonate coating.
The AN is easily separated from it but your CAN might be a different nitrate fertiliser.
If it's the same, then it makes Govt. restrictions laughable in their pointlessness.

Melgar - 3-3-2010 at 22:14

Yeah, I mean, you can just mix the aforementioned ammonium sulfate with calcium nitrate in water, and calcium sulfate will precipitate out, leaving ammonium nitrate still dissolved.

Also, I found out that ozone will apparently oxidize ammonia pretty well. I have an ozone generator to get rid of the occasional bad smell, and I turned it on in a room where i had spilled ammonia, and the room gradually filled up with a white, smoke-like mist. Set off the smoke detector too.

not_important - 4-3-2010 at 05:20

Note that calcium sulfate is slightly soluble in water, and that ammonium salts enhance the solubility; this has been documented many times before on this board. For use in making HNO3 this solubility is likely not a problem, but it is if you are after NH4NO3 for other purposes.

The catalytic oxidation of ammonia is a delicate balance between low conversion rates and three competing reactions. The first is the cracking of NH3 to N2 and H2, the second is oxidation of NH3 to N2 and H2O, and the thirdis the reaction of nitrogen oxides with ammonia to give H2O and N2. That 3rd is used to reduce NOx emissions in combustion exhaust, first NH3 is added to react with the nitrogen oxides and the the stream flows through a slip catalyst where excess NH3 is oxidised to N2 and H2O.

Industrially this is made practical by using high surface area catalysts in thin layers with careful control of mix ratios, temperatures, flow rates, and quick quenching. Not easy to do on a small scale at home, while a demonstration may produce traces of NOx it is going to give disappointing yields.

The smoke you observed is in part ammonium nitrite - NH4NO2. When dissolved in water it tends to react with itself to produce N2 and H2O, reducing the yields bassed on NH3 or O3 consumed.

Someone here once posted a description of a set-up that used copper, strong aqueous NH3, and air, to produced a solution of Cu(NO3)2 or the corresponding tetra-ammonium complex.

Melgar - 4-3-2010 at 06:46

Quote: Originally posted by not_important

Note that calcium sulfate is slightly soluble in water, and that ammonium salts enhance the solubility; this has been documented many times before on this board. For use in making HNO3 this solubility is likely not a problem, but it is if you are after NH4NO3 for other purposes.

The catalytic oxidation of ammonia is a delicate balance between low conversion rates and three competing reactions. The first is the cracking of NH3 to N2 and H2, the second is oxidation of NH3 to N2 and H2O, and the thirdis the reaction of nitrogen oxides with ammonia to give H2O and N2. That 3rd is used to reduce NOx emissions in combustion exhaust, first NH3 is added to react with the nitrogen oxides and the the stream flows through a slip catalyst where excess NH3 is oxidised to N2 and H2O.

Industrially this is made practical by using high surface area catalysts in thin layers with careful control of mix ratios, temperatures, flow rates, and quick quenching. Not easy to do on a small scale at home, while a demonstration may produce traces of NOx it is going to give disappointing yields.

The smoke you observed is in part ammonium nitrite - NH4NO2. When dissolved in water it tends to react with itself to produce N2 and H2O, reducing the yields bassed on NH3 or O3 consumed.

Someone here once posted a description of a set-up that used copper, strong aqueous NH3, and air, to produced a solution of Cu(NO3)2 or the corresponding tetra-ammonium complex.

I read some papers on the subject, and it's really not as delicate as you make it seem. For one, the first reaction, NH3 to N2 + H2 is practically nonexistant, since it's not favored thermodynamically except maybe at really high temps. Kind of like how water doesn't tend to disassociate into hydrogen and oxygen. As for the oxidation to water and nitrogen, that happens mainly above 1400C with platinum catalysts. That's a pretty wide range, and keeping it in that range should be easy enough just by estimating the temperature due to color from black-body radiation.

Of course, you're right that getting the surface area high enough to convert most of the ammonia would be difficult. And now that I have a reliable source of calcium nitrate, my incentive to get that working isn't as great. And it is good to know that oxidizing with ozone won't help me because of the ammonium nitrite reaction.

As for purifying ammonium nitrate, couldn't it just be carefully heated to just over its melting point to get the calcium sulfate to precipitate out? I know mixtures of salts usually have different melting points than the salts by themselves, but the huge difference in melting points should minimize that effect, no?

not_important - 5-3-2010 at 04:42

Quote: Originally posted by Melgar

... For one, the first reaction, NH3 to N2 + H2 is practically nonexistant, since it's not favored thermodynamically except maybe at really high temps. Kind of like how water doesn't tend to disassociate into hydrogen and oxygen...

...

As for purifying ammonium nitrate, couldn't it just be carefully heated to just over its melting point to get the calcium sulfate to precipitate out? I know mixtures of salts usually have different melting points than the salts by themselves, but the huge difference in melting points should minimize that effect, no?

From S. Strelzoff's Technology & Manufacture of Ammonia
Percent NH3 in equilibrium with [N2+3H2] at 1 atmosphere pressure and various temperatures

T Cent. Percent NH3
350 0,838
400 0,405
450 0,214
500 0,1216
550 0,0736
600 0,0470

So even at 350C the equilibrium gives less than a percent NH3; catalysts simply speed up reaching equilibrium without changing the final ratios. Increased pressure favours higher NH3 percentages, as would be expected; at 800 atmospheres and 500 C the equilibrium is 51% NH3 which demonstrates why the H-B process is run at high pressures.

Melting ammonium nitrate is always a little dicey, especially if impure. Plus it's not straightforward to use solubility values for water as solvent to solubility values in a non-aqueous solvent such as liquid ammonium nitrate. If all you are after is a route to nitric acid the amount of calcium sulfate remaining in solution shouldn't be a problem. If you need pure NH4NO3 then using hot MeOH to dissolve the NH4NO3 is likely the best route (solubility given in other threads); warm anhydrous NH3 under pressure works well as a solvent too, but is hardly a home lab method.

Melgar - 5-3-2010 at 06:40

Quote: Originally posted by not_important

Quote: Originally posted by Melgar

Damn, you're right. I figured for sure that reaction couldn't be occurring, because that was never given in any of the products in the papers I read regarding catalytic oxidation of ammonia. But then I realized, no shit, that's because there's OXYGEN in the mix along with HYDROGEN at HIGH TEMPERATURES. Gee, I wonder where all the hydrogen went?

chief - 5-3-2010 at 06:49

@Melgar: You have a reliable source of calcium nitrate ?
==> Since CaSO4 is nearly insoluble, the following should/could work: Ca(NO3)2 + H2SO4 ==> 2 HNO3 +CaSO4 ;
the CaSO4 could be let settle down or pressed out, and ready would be the conc. HNO3 ...

[Edited on 5-3-2010 by chief]

[Edited on 5-3-2010 by chief]

quicksilver - 5-3-2010 at 07:08

If you live in the US you can't miss calcium nitrate.....especially in rural areas. They (The magic Fertilizer In$titute) found they could make MORE money with bagged CaNO3 but it has a great deal of AN in it - it's called 15.5-0-0 and it's all over the SouthWestern US. To get AN now you have to (in my area) order one acre at least (300lbs) and it's always delivered unbagged. The companies agreed NOT to bag it now. So you better damn well have a feed trough and a way to cover it. PLUS, you have to sign a chit saying "I am a good boy or girl and I won't do anything with this except burn the Hell out of my soil". If you even attempt to order outside of turn over season they laugh in your face and hang up.

Oh and 15.5-0-0 is totally legal and righteous. No one will flip out. You'll start seeing it all over eBay because it's like $27 a bag!!!! They put so much clay on it you can build a sand castle while you wait for your nitrate to clean up.

[Edited on 5-3-2010 by quicksilver]

Melgar - 5-3-2010 at 10:44

Quote: Originally posted by chief

@Melgar: You have a reliable source of calcium nitrate ?
==> Since CaSO4 is nearly insoluble, the following should/could work: Ca(NO3)2 + H2SO4 ==> 2 HNO3 +CaSO4 ;
the CaSO4 could be let settle down or pressed out, and ready would be the conc. HNO3 ...

Heh, yep, that was the first thing I did last week when I got home with my bag of calcium nitrate. Yielded a nice clear acid too; I was afraid I was distilling water until I took off the vacuum and white vapors started drifting out.

My dad used to be able to get ammonium nitrate, in those huge-ass fertilizer spreaders. Once I filled up a five gallon bucket with that stuff and stuck it in the shed, but then quite a long time later when I went back to get it, it was gone. My dad said he got rid of it... damn it.. Now he says they can't get it anymore. Well, presumably it's still possible to get it but it's so much paperwork that everyone just uses something else.

I'm still interested in ammonium sulfate, but now more as a source of ammonia and sulfuric acid. And it was cool to at least get a little bit of ammonium nitrate from that reaction.

quicksilver - 5-3-2010 at 13:57

There is some recent interesting history about AN. In New York there was this congressman named Shumer (D-NY) who tried to have it banned totally (The Calcium Nitrate is calcium-ammonium nitrate - a big mix of nitrates with a lot of clay) But the rural states flipped. To ban nitrates could actually destroy agri-business-states checker-boarded all over the US. So his "Ban" was shut down hard!
But the "Fertilizer Institute" got a great deal of pressure from the Fed to do something. So they put together a campaign called "Be Aware for America" and generally told folks to watch WHO was buying. well no one likes doing extra work so that fizzled. but since the prill towers don't care what drops down them, someone came up with putting come clay and mica and AN with CaNO3; the idea being that people with crazy ideas wouldn't play with water and filters and evaporation, etc. That works fine for some states but some have vast chemical plants and the AN was a by-product of their endeavors. All they needed was a tower. So ammonium nitrate is still being made at the hundred ton level in the US but a lot has shifted to Northern Europe.

Ammonium Nitrate is REALLY only needed by a few crops. Berry's benefit & tobacco, and in areas with a Fall with few leaves to mulch the soil with sharp season-crop demands. So realistically If you grow corn or perhaps cotton in sand soil conditions, you'd want AN - & you can still get it but you can't buy it bagged any longer. It is delivered only at the acre level (which is (about 300lbs). But C.A.N. does everything that AN would and they made it the same size so the rollers don't need bigger holes or any alteration in equipment.
In terms of quality Yara, the Norwegian company produces the best small bagged product.
http://www.yara.co.uk/about_us/history/corporate_history.htm...

Hydroponics growers knew about that for some time. We have an FFA & 4H club very close and our Fair has some serious contestants. There was a 100+ lbs hydroponically grown pumpkin last year: the product WORKS!

[Edited on 5-3-2010 by quicksilver]

grndpndr - 2-4-2010 at 21:33

Yara sent a 4 0z sample of thier solution grade calcium nitrate double salt its very nearly completely soluble
OT .01%.I dont think thats alot of clay to worry bout
but as I said its a dbl salt 5Ca(NO3)2-NH4NO3.-10h20
.If interested Im sure its available about anywhere "yara" Viking Ship Brand Calcinit ,'solution grade' 15.5-0-0 +19Ca.If Ebays to be trusted 35lbs shipped for $38 doesnt sound like gouging.Im not aware of how it will react attempting to synthesise AN
via ammonium sulfate or HNO3 using the double salt and H2SO4.But I intend to at least try distilling HNO3 with the sample im not lookin to make FRNA/WFNA the 68% azeotrope will do for most of my needs.Simplicitys sake Ill use dilute H2SO4.I really cant forsee any insurmountable problems but thats the reasoning for the post. Thanks.Maybe some of the clay
problems can be avoided sticking to solution grade calcinit.Ive heard the DBL salt calcium nitrate in question referred to as ammonium calcium nitrate to differentiate it from the clay laden CAN, different animals altogether according to MSDS/manufacturers.

[Edited on 3-4-2010 by grndpndr]