Hallo to all,
even searching, I haven't found nothing about a thread talking of ammonium salts, so I've decided to create this one.
Ammonium Bicarbonate
NH4HCO3
Mol. 79,06 g/mol
Density 1,586 g/cm3
Sol. 17,4g in 100ml H2O at 20°C
M.P 106°C
B.P Decomposition
Ammonium Bifluoride
(NH4)HF2
Mol. 57,04 g/mol
Density 1,5 g/cm3
Sol. 63g in 100ml H2O at 20°C
M.P 126°C
B.P 230°C - decomposition
Ammonium Chloride
NH4Cl
Mol. 59,49 g/mol
Density 1,527 g/cm3
Sol. 29,7g in 100ml H2O at 20°C
M.P 338°C
B.P 520°C - sublimation
Ammonium Fluoride
NH4F
Mol. 37,04 g/mol
Density 1,01 g/cm3
Sol. 82g in 100ml H2O at 20°C
Ammonium Sulphate
(NH4)2SO4
Mol. 132,14 g/mol
Density 1,77 g/cm3
Sol. 75,5g in 100ml H2O at 20°C
M.P 273°C - decomposition
Just some infos. about common ammonium compound...
How can I prepare ammonium oleate?
Does anyone have idea about an exotic/interesting ammonium salt?
Thanks at all for help! Ammonium fans!The_Davster - 8-10-2005 at 13:40
Ammonium perchlorate, chlorate permanganate, are all interesting in their own little way.
Or a rather exotic ammonium compound I prepared recently: diethylammonium chloride.
Ammonium oleate I imagine could be prepared by the action of Ammonium (bi)carbonate or hydroxide on oleic acid. Look up their solubilities. If oleic
acid is insoluble in water it is easy, just shake it with a solution of NH3 or NH4CO3 until it dissolves, then evaporate to get crystals of ammonium
oleate.kazaa81 - 8-10-2005 at 14:03
Thank you very much, rogue chemist!
The salts you've mentioned are quite interesting, and also the synthesis of NH4 oleate will be helpful to me. I want to prepare oleic acid from
olive oil and NaOH (cold saponification process, no heat required), then adding 90% H2SO4.
Returning on ammonium, ferric ammonium citrate is fluorescent if I remember right.
(How to diethylammonium Cl? NH4Cl + C2H5OH? You could send more details via U2U).
Thanks all for help!
[Edited on 8-10-2005 by kazaa81]
[Edited on 9-10-2005 by kazaa81]
Some more data about ammonium salts...
kazaa81 - 9-10-2005 at 12:24
There's some more infos about ammonium salts.....
Ammonium Aluminum Sulphate
(NH4)Al(SO4)2 * 12 H2O
Mol. 453,33 g/mol
Density 1,64 g/cm3
Sol. 15g in 100ml H2O at 20°C
M.P 93,5°C
B.P 120°C -10 H2O
Ammonium Carbonate
(NH4)2CO3
Mol. 96,09 g/mol
Density -
Sol. 100g in 100ml H2O at 15°C
M.P 58°C - decomposition
B.P -previously decomposition-
Ammonium Chlorostannate
(NH4)2SnCl6
Mol. 367,52 g/mol
Sol. 33g in 100ml H2O at 20°C
Density 2,4 g/cm3
M.P decomposition
B.P -previously decomposition-
Ammonium Chromiumsulphate
(NH4)Cr(SO4)2 * 12H2O
Mol. 478,38 g/mol
Sol. 20g in 100ml H2O at 20°C
Density 1,72 g/cm3
M.P -
B.P 95°C -9 H2O
Ammonium Ironsulphate (ferrous)
(NH4)2SO4 * FeSO4 * 6 H2O
Mol. 392,16 g/mol
Sol. 27g in 100ml H2O at 20°C
Density 1,864 g/cm3
M.P decomposition
B.P -previously decomposition-
Ammonium Ironsulphate (ferric)
(NH4)Fe(SO4)2 * 12 H2O
Mol. 428,21 g/mol
Sol. 123g in 100ml H2O at 20°C
Density 1,71 g/cm3
M.P 230°C -12 H2O
B.P -
Ammonium Nitrate
NH4NO3
Mol. 80,05 g/mol
Sol. 118g in 100ml H2O at 0°C
Density 1,725 g/cm3
M.P 170°C
B.P 210°C - decomposition
I hope there aren't mistakes....
Oleic Acid [cis-9-octadecenoic]
CH3-(CH2)7-CH=CH-(CH2)7-COOH
C18H34O2
Mol. 282,47 g/mol
Density 0,895 at 18°C
Sol. not sol. in H2O, much soluble in ether, benzine, acetone,
high sol. in alcohol, chloroform
M.P 16°C
B.P 286°C
Please post any idea about ammonium and related!
Ammonium fans!
[Edited on 9-10-2005 by kazaa81]kazaa81 - 15-10-2005 at 07:00
Other informations about ammonium salts....
Ammonium Acetate
(NH4)CH3COO
Mol. 77,08 g/mol
Density 1,17 g/cm3
Sol. 148 g in 100ml H2O at 4°C
M.P 114°C
B.P 273°C -decomposition
Ammonium Cerium (IV) Nitrate
(NH4)2Ce(NO3)6
Mol. 548,23 g/mol
Density - g/cm3
Sol. 141 g in 100ml H2O at 25°C
M.P 107°C
B.P -
Ammonium Dichromate
(NH4)2Cr2O7
Mol. 252,07 g/mol
Density 2,15 g/cm3
Sol. 36 g in 100ml H2O at 20°C
M.P 180°C -explosive decomposition
B.P
Ammonium Bismuth Citrate
NH4(BiOH)C6H5O7
Mol. 433,12 g/mol
Density - g/cm3
Sol. - g in 100ml H2O at 20°C
M.P -
B.P -
Ammonium Citrate
NH3 * (C6H8O7)2
Mol. 226,19 g/mol
Density 1,48 g/cm3
Sol. - g in 100ml H2O at 20°C
M.P -
B.P -
Ammonium Chromate
(NH4)2CrO4
Mol. 152,07 g/mol
Density - g/cm3
Sol. 34 g in 100ml H2O at 20°C
M.P -
B.P -
Ammonium Oleate
(NH4)C18H3302
Mol. 299 g/mol
Density g/cm3
Sol. g in 100ml H2O at 20°C
M.P
B.Pchloric1 - 15-10-2005 at 15:03
Dont kinow where you are from kazaa but in my area they seel huge bags of ammonium sulfate at garden suplly centers. I merely dissolve as much as
possible in boiling water, filter pout crap, and let the solution cool until I get white needle like crystals. The sulfate is pretty versatile. It
gives the ammonium chlorate with potassium chlorate and can precipitate barium from its solutions leaving the ammonium salt in solution. Ammonium
carboionate can be made by heating CaCO3 with Ammonium Sulfate also.12AX7 - 15-10-2005 at 22:23
Quote:
Originally posted by chloric1
It gives the ammonium chlorate with potassium chlorate
...Isn't NH4ClO3 more soluble than KClO3 or if you typoed, KClO4?
Timjimmyboy - 15-10-2005 at 22:53
ammonium carbonate - bubble carbon dioxide into ammonia solution
ammonium chloride - bubble chlorine into ammonia solution
and so on and so forth....Pyrovus - 16-10-2005 at 03:38
Quote:
Originally posted by jimmyboy
ammonium chloride - bubble chlorine into ammonia solution
I'm hoping you mean bubble hydrogen choride into ammonia. Chlorine will give NCl3 woelen - 16-10-2005 at 05:25
Quote:
Originally posted by jimmyboy
ammonium carbonate - bubble carbon dioxide into ammonia solution
ammonium chloride - bubble chlorine into ammonia solution
and so on and so forth....
Please don't do the latter without thinking! You'll end up getting chloramine, a very dangerous and carcinogenic gas. If lots of chlorine
are bubbled through such a solution, you may even get NCl3, but most likely long before that point is reached, you'll have left the room, due to
the unbearable stench of NH2Cl.
NH2Cl definitely is not similar to NH4Cl!!chloric1 - 16-10-2005 at 07:03
Quote:
Originally posted by 12AX7
Quote:
Originally posted by chloric1
It gives the ammonium chlorate with potassium chlorate
...Isn't NH4ClO3 more soluble than KClO3 or if you typoed, KClO4?
Tim
Yes tim you are right ammonium chlorate is more soluble. That is exactly my point. Potassium chlorate is relativley slightly soluble in the cold but
in hot water the dissolution in quite consdierable. When the proper molar proportions of ammonium sulfate and potassium chlorate are mixed and enough
hot water is added to dissolve, all four ions are mixed. It is therefore good to concentrate VERY carefully until a slurry is formed and a large
volume of alcohol is added to precipitate the potassium sulfate as the ammonium chlorate is easily alchohol soluble. This is from the "Handbook
of Preparative Inorganic Chemistry" by Brauer. It is the initial setup for barium chlorate preparation.
I have not physically tried this in fear of potential expolision hazards but I think I could slightly modify the process to make it marginally safer.
There is a US patent from September 2003 about making alkaline chlorate from ammonium chlorate that is obtained from sodium chlorate, ammonia, and
carbon dioxide. this would precipitate baking soda and leave ammonium chlorate in solution. Being sodium chlorate is not always easy to come by I
might do the more dnagerous ammonium sulfate process and add a little aqua ammonia to retard acid production. Sulfates + Chlorates BAD!!kazaa81 - 16-10-2005 at 07:43
Does anyone have a ferric ammonium citrate preparation?
It's an interesting compound that is fluorescent (glow).
Any other exotic/interesting ammonium compound?
I've heard which the reaction between a barium salt and an ammonium one is so endotermic that the water in which the 2 salts are dissolved will
frozen!
Ammonium fans!Darkblade48 - 16-10-2005 at 10:47
Quote:
Originally posted by kazaa81
Does anyone have a ferric ammonium citrate preparation?
CHAPTER I--FOOD AND DRUG ADMINISTRATION, DEPARTMENT OF HEALTH AND HUMAN
SERVICES (CONTINUED)
PART 184--DIRECT FOOD SUBSTANCES AFFIRMED AS GENERALLY RECOGNIZED AS SAFE--Table of Contents
Subpart B--Listing of Specific Substances Affirmed as GRAS
Sec. 184.1296 Ferric ammonium citrate.
(a) Ferric ammonium citrate (iron (III) ammonium citrate) is
prepared by the reaction of ferric hydroxide with citric acid, followed
by treatment with ammonium hydroxide, evaporating, and drying. The
resulting product occurs in two forms depending on the stoichiometry of
the initial reactants.
(1) Ferric ammonium citrate (iron (III) ammonium citrate, CAS Reg.
No. 1332-98-5) is a complex salt of undetermined structure composed of
16.5 to 18.5 percent iron, approximately 9 percent ammonia, and 65
percent citric acid and occurs as reddish brown or garnet red scales or
granules or as a brownish-yellowish powder.
(2) Ferric ammonium citrate (iron (III) ammonium citrate, CAS Reg.
No. 1333-00-2) is a complex salt of undetermined structure composed of
14.5 to 16 percent iron, approximately 7.5 percent ammonia, and 75
percent citric acid and occurs as thin transparent green scales, as
granules, as a powder, or as transparent green crystals.
(b) The ingredients meet the specifications of the Food Chemicals
Codex, 3d Ed. (1981), pp. 116-117 (Ferric ammonium citrate, brown) and
p. 117 (Ferric ammonium citrate, green), which is incorporated by
reference. Copies are available from the National Academy Press, 2101
Constitution Ave. NW., Washington, DC 20418, or available for inspection
at the Office of the Federal Register, 800 North Capitol Street, NW.,
suite 700, Washington, DC 20408.
[[Page 503]]
(c) In accordance with Sec. 184.1(b)(1), the ingredients are used in
food as nutrient supplements as defined in Sec. 170.3(o)(20) of this
chapter, with no limitation other than current good manufacturing
practice. The ingredients may also be used in infant formula in
accordance with section 412(g) of the Federal Food, Drug, and Cosmetic
Act (the act) (21 U.S.C. 350a(g)) or with regulations promulgated under
section 412(a)(2) of the act (21 U.S.C. 350a(a)(2)).
(d) Prior sanctions for these ingredients different from the uses
established in this section do not exist or have been waived.
[53 FR 16864, May 12, 1988]kazaa81 - 16-10-2005 at 11:28
Thank you, darkblade!
Maybe this will help a little about information, not in preparation...the fact which it's structure isn't yet well known is interesting!
Ferric Ammonium Citrate (also called: ammonium ferric citrate, iron ammonium citrate)
Fe(NH4)3(C6H5O7)2
CAS # 1185-57-5
Ammonium Bromide
NH4Br
CAS # 12124-97-9
Mol. 98,98 g/mol
Density 2,43 g/cm3
M.P -
B.P 425°C - sublimationDarkblade48 - 16-10-2005 at 14:20
Quote:
Originally posted by kazaa81
Thank you, darkblade!
Maybe this will help a little about information, not in preparation...the fact which it's structure isn't yet well known is interesting!
(a) Ferric ammonium citrate (iron (III) ammonium citrate) is
prepared by the reaction of ferric hydroxide with citric acid, followed
by treatment with ammonium hydroxide, evaporating, and drying. The
resulting product occurs in two forms depending on the stoichiometry of
the initial reactants.
I tried that method once (took some iron, put it as both cathode and anode in a solution of NaCl and proceeded to run 9V through it to obtain the gray
iron hydroxide (some of it oxidized at the water surface to Fe2O3, but that's besides the point).
Upon throwing in some citric acid (notice the "precise" measurements) and pouring in ~ 100 mL of ammonium hydroxide (just your normal
regular kind, nothing concentrated), I got a dark brown liquid.
I didn't bother recrystallizing it though, but the procedure seems to work.kazaa81 - 17-10-2005 at 12:31
How can FeCl3 be (brown-red) prepared?
If one dissolve Fe in HCl, the yield is FeCl2 (green).
FeCl2 (green) + sodium hypochlorite ----> FeCl3?
Thanks all for help!neutrino - 17-10-2005 at 13:49
Simply bubbling air through the mix should work. I know that this works in oxidizing copper to copper II.12AX7 - 17-10-2005 at 13:56
Idunno, the FeCl2 solution I've had sitting out for weeks has a rusty skin on it but damned if the solution ain't turning from green to
brown.
If you don't want anything else in solution, TCCA would be a good idea.
TimDarkblade48 - 17-10-2005 at 15:45
Quote:
Originally posted by kazaa81
How can FeCl3 be (brown-red) prepared?
If one dissolve Fe in HCl, the yield is FeCl2 (green).
FeCl2 (green) + sodium hypochlorite ----> FeCl3?
Thanks all for help!
This might work, since the addition of the hypochlorite to the acid will liberate Cl2 gas, and for sure, using Cl2 gas will easily make the FeCl2
--> FeCl3...but then there is the issue of safety, and how effective this actually is.kazaa81 - 18-10-2005 at 12:25
I will try it....
If it works, everything would go fine!
Has anyone experience about this or ideas to post?
Ammonium fans!denatured - 19-10-2005 at 05:16
Hello
Kazza81:
Could you add the uses of each ammonium compounds for those who don't know?kazaa81 - 20-10-2005 at 11:50
I will add them as soon as possible...
I have recently discovered a thing:
in Italy press is PARTIALLY free, as Mongolia or other 3rd world natons. Now, say if this is possible this for a nation like Italy....... thank you,
b3rlusconi!
[Edited on 22-10-2005 by kazaa81]kazaa81 - 22-10-2005 at 08:32
Most of ammonium compounds are used as fertilizer...
Ammonium Dihydrogen Phosphate
(NH4)H2PO4
Mol. 115,03 g/mol
Density 1,80 g/cm3
Sol. 37g in 100ml H2O at 20°C
M.P 190°C
B.P -bio2 - 23-10-2005 at 03:09
...(NH4)H2PO4 ......
The mono H form of this orthophosphate is quite interesting..
(NH4)2HPO4
CRC says it decomposes 155deg and soluble in water 57.5g/100g at 10deg.
Got over a kilo of the fine dust like yellow powder out of a fire extinguisher that has only monoammonium phosphate listed as an ingredient. There
must be some other component added to this stuff because this is what was observed in an attempt to make some phosphoric acid.
First off this stuff won't dissolve in water to any appreciable extent even at the boiling point. It's so fine and floats so I ground it
under hot water stirred it for hours with a 1000rpm magnet and no luck dissolving any noticable amount in the water.
OK so now I try acid and finally determine nope concentrated Sulfuric or HCl ain't working either. This was after trying different dilutions.
Being frustrated by this time I decided to just decompose the shit with some heat so into a flask immersed in a 200deg oil bath it goes.
Finally got the oil to the smoke point trying to melt/decompose this wonderful substance and nothing happened.
In desperation (lol) I place some fresh few grams in a stainless steel beaker and put over the propane burner. After 10minutes as hot as the damn
burner would get it finally it sort of gells and gets viscous but no ammonia evolution. I even melted some NaOH with it some time later and after the
cool goo dried
a couple days it still didn;t dissolve in water.
Strange stuff indeed. No wonder its sprayed on fires.gm137 - 29-4-2007 at 04:52
I've had similar problems with MAP; the powder used for extinguishers has a metallic sterate added to repel water and prevent caking - which also
prevents is dissolving in anything aqueous.Pyrovus - 1-5-2007 at 03:23
On the subject of exotic ammonium salts, I wonder if it might be possible to create ammonium hydride, NH5. Ordinarily you would expect the hydride
anion to act as a base, and deprotonate the ammonium ion to give ammonia gas and hydrogen. However, if you could trap the ammonium ion inside a crown
ether or something similar, you could prevent it from interacting with the hydride ions and render the ammonium hydride stable. It's basically the
same idea as was used to prepare stable alkali metal anion salts - trap the alkali metal cation in a crown ether to prevent it reacting with the
anion. And on the note of alkali metal anions, ammonium natride would be pretty nice .
[Edited on 1-5-2007 by Pyrovus]Matchheads - 11-5-2007 at 16:24
what ammonium salt and what reagent is it going to be if I want to pour the reagent onto a bucket of the amm salt and get huge clouds of ammonia?JohnWW - 11-5-2007 at 18:56
The NH4+ cation is smaller than, and has less available orbitals than, the Na+, K+, and other alkali metal cations which have been successfully
chelated in hexadentate crown-ether ligands, primarily for the purpose rendering them soluble in aprotic organic solvents. In fact, NH4+ has no spare
2s and 2p orbitals available with which to accept electron pairs from crown-ether molecules, unline Na+, K+, and even Li+. Even if it could somehow,
steric factors would allow only a tetradentate ligand to chelate it; but there is no known case of any second-row element, after filling its 2s and 2p
orbitals with 8 electrons, then engaging in further bond-formation by utilizing empty 3s and 3p (and 3d) orbitals. By contrast, third-row elements can
expand beyond 8 valence electrons (up to 12, with more than 6 ligands being sterically precluded) by utilizing empty 3d orbitals.
However, it may just be possible to make He, Ne, and possibly Ar atoms accept electrons to become anions which are isoelectronic with monatomic Li,
Na, and K, but even more reactive, if they could be stabilized somehow. Stabilization with a hexadentate crown-ether ligand in the case of Ne- and
Ar-, and tetradentate crown-ether ligand in the case of He-, might just do it. If, as Pyrovus says, crown ether-chelated chelated alkali metal anions
have been obtained by inducing them to accept electrons, it just might be possible with inert gas atoms, although the spin-pairing of the two valence
electrons in an alkali metal anion may make them more stable than an inert-gas anion..
Even if NH5, as the salt NH4+H-, could be somehow obtained by mixing a NH4+ salt and an alkali metal hydride salt (spontaneously flammable in air) in
a non-oxidizing polar aprotic solvent, it is likely to be a fairly ephemeral species, because of the driving-force of the heat of formation of H2 from
H+ and H-.
It may be for the same reason that all attempts to prepare NF5, as the salt NF4+F-, have been unsuccessful, in spite of the existance as very stable
compounds of NF4BF4, NF4PF6, NF4SbF6, and NF4ClO4,. However, even these latter compounds are not easy to prepare, as they require either the addition
of F+ to NF3, or loss of an electron by NF3 with subsequent bonding by an F atom. For example, NF4SbF6 can be prepared by reacting NF3 with the salt
KrFSbF6, one of the few isolatable Kr compounds (obtained by reaction of KrF2 with SbF5); the KrF+ cation decomposes to Kr and F+, then the latter
reacts with NF3. ClF6+ has been obtained in the same manner. Methods that have been used to try to prepare NF5 have been primarily heating NF3 with F2
under high pressure. Similarly, although the OF3+ cation, containing tetravalent oxygen, is theoretically thermodynamically stable, all attempts to
prepare it, as the salt of either F- or BF4- or SbF6- or other anions, have been unsuccessful, even though of course H3O+ exists in all aqueous
solutions of acids.
.