Synthesis of Potassium Nitrate KNO3

[Setharier]

Hopefully this hasn't been yet discussed around. I just feel I have responsibility of bringing this into knowledge due to the fact EU actually plans to limit or even partially ban commercial sale of KNO3, NH4NO3, HNO3, H2SO4, hexamine and various other general "terrograde" chemicals due to threat of terrorism. The way to obtain KNO3 from commercial dozen-fertilizers is explained in this conversation:

>> I have found a couple of methods for synthesizing
>> Potassium Nitrate from common chemicals. This method
>> relies on Calcium Nitrate (fertilizer) as the precursor
>> for making KNO3.
>>
>> The reaction is pretty simple:
>>
>> Ca(NO3)2 + K2SO4 --> 2KNO3 + CaSO4
>>
>> The molecular weight of these compounds is:
>>
>> Calcium Nitrate = Ca(NO3)2 = 164.0879
>> Calcium Sulphate = CaSO4 = 136.1416
>> Potassium Nitrate = KNO3 = 101.1032
>> Potassium Sulphate = K2SO4 = 174.2602
>>
>> The right percentage of chemicals required to balance the
>> reaction from a stoichiometricly is:
>>
>> Calcium Nitrate ................... 48.5
>> Potassium Sulphate ................ 51.5
>>
>> Which results in:
>>
>> Potassium Nitrate ................. 59.8
>> Calcium Sulphate .................. 40.2
>>
>> Assuming you start with 485 grams of Calcium Nitrate and
>> 515 grams of Potassium Sulphate, your result is approximately
>> 598 grams of Potassium Nitrate and 402 grams of Calcium
>> Sulphate. Ideally, you would do this by dissolving the
>> chemicals in distilled water and boiling to facilitate the
>> ion exchange.
>>
>> The Calcium Sulphate resultant is technically insoluble in
>> water (and becomes even more insoluble when the temperature
>> is increased). By adding distilled water to the solution you
>> can separate the Potassium Nitrate from the Calcium Sulphate
>> very easily. The solubility of Potassium Nitrate in water is
>> as follows:
>>
>> 13 g/100 ml (0 °C)
>> 32 g/100 ml (20 °C)
>> 64 g/100 ml (40 °C)
>> 110 g/100 ml (60 °C)
>> 169 g/100 ml (80 °C)
>> 246 g/100 ml (100 °C)
>>
>> For a temperature of 20C and a result of 598 grams of KNO3,
>> you would need approximately 1869 m/l of water to dissolve
>> the potassium nitrate (598g / 32 = 1869 m/l), plus a little
>> bit more for evaporation. If you heat the solution slightly,
>> you should get Potassium Nitrate dissolved in water, and
>> Calcium Sulphate should precipitate out of the solution.
>>
>> Simply siphon off the clear solution (this is the "supernatant").
>> You may even be able to filter the Calcium Sulphate from the
>> solution by pouring the solution through filter paper. Boil the
>> supernatant until you start to see the Potassium Nitrate appear
>> as surface tension on the top of the hot solution. Remove the
>> container from the heat, cool the solution to 0 C, and let the
>> KNO3 crystals precipitate out of the solution.
>>
>> You should be able to get 477 grams of pure KNO3 (calculated
>> as (1 - (13g / 246g)) x 598g KNO3 = 566g AN), with 32 grams
>> left in the solution. You can repeat this process on the
>> remaining solution to get another 30 grams of KNO3.

[Jome skanish]

This requires the fertilizer-grade chemicals to be pure, and water free, which they usually isn't.

Ca(NO3)2 for example is mostly encountered as the tetrahydrate, Ca(NO3)2*4H2O. Here in Sweden we have two brands of useful non-conspicuous fertilizers, one consists of AN and potassium sulfate, and one consists of mainly calcium nitrate tetrahydrate with some 5% of the nitrogen coming from NH4 (as AN)

Most people in here should be able to figure out the ratios to form KNO3 from these based on the stated NPK-value.

In "our" case, we needed to first calculate the amount of ammonium ions in the fert. grade calcium nitrate, and then these would each grab a -NO3 rendering it inert as the AN won't precipitate when the supernatant is boiled down. So, here it was "moles of nitrate / kg - moles of ammonia / kg = necessary amount of K2SO4, which in turn was calculated after the K-value of fert. II since K forms very few insoluble salts, on the contrary to Ca.

What I am trying to say is, making a specific "how-to" writeup will not be very useful for anyone (...) if the principles behind are not understood.

EDIT: Tetrahydrate, it was! And the crap is insoluble in ethanol, just as K2SO4, meaning we have an easy way of purifying our KNO3.

[Hinckleyforpresident]

We're dissolving the nitrate in water... It certainly doesn't matter if it's the dihydrate, or if it's anhydrous. All that you need to know is how much more the molar weight of the hydrates is in order to get the proper ratios.

[Lewis]

This procedure may be useful to rocketry enthusiasts or black powder makers, but I don't see how KNO3 is any more useful in HE production than Ca(NO3)2. (other than the hydrate issue, which I assume can be resolved by heating the calcium salt carefully, but I could be wrong)

Regardless, it seems a method for creating either ammonium or sodium nitrates would be best for nitrations or distillation of HNO3. Thankfully, I don't have to worry about this sort of thing in my location. A simple five minute drive to the nearest farm and garden location got me 25kg of prilled NH4NO3, with not so much as a questioning look. (it's winter, too)

I suppose there are advantages to living in the breadbasket of the world.

[Jome skanish]

Hinckley: You are assuming that the agricultural chemicals are pure, or can be easily purified. That is not always the case, as with "my" calcium nitrate which contained both insoluble crap and ammonium nitrate. I once tried to boil it down after dissolving and filtrating (the AN was supposed to stay in solution) and was treated to a rather impressive demonstration of over saturation and rapid crystallization once the solution was stirred.

If AN is desired it could be had if the calcium nitrate is reacted with ammonium sulfate, precipitating gypsum. AN in stores? Never take anything for granted...

[Hinckleyforpresident]

If it came down to it, you could separate out the nitrate by fractionally distilling the fertilizer with sulfuric acid and only collecting the HNO3. Unfortunately, one would need a large excess of sulfuric acid just to keep the calcium bisulfate mostly in solution. This process would require the fertilizer to be anhydrous, which could be done by throwing a cookie sheet of fertilizer into the oven at 200*F.

If a nitrate salt was required, it could be made by neutralizing the acid with the appropriate base.