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Author: Subject: OTC Phosphoric Acid
S.C. Wack
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[*] posted on 27-1-2005 at 13:13


Organic solvent comes to mind as something to try. Phosphoric acid/phosphates have to get fairly hot before any harm comes to glassware, no need to worry in this case. It is actually recommended as a heating bath in Vogel. The damage to glass (looks like frosted/ground glass) might have something to do with the composition.
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Magpie
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[*] posted on 27-1-2005 at 15:58


Thanks Mr Wack - that was something I had thought of also but was concerned that the "cleaners and degreasers" might just put some of this organic solvent into the aqueous phase. But I think it deserves a try on a test tube scale, of course. I have toluene and hexane (Coleman lantern fuel) as possible solvents. I think I'll try the hexane first. Any traces of this left in the phosphoric acid should evaporate off when I concentrate it. Also it might be smart to just leave the dye in the acid to start with as an indicator to see if it follows the hexane or the water.



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[*] posted on 27-1-2005 at 16:30


It may be worth a try to react all the phosphoric acid with a calcium base or CaCl2 and filter to remove all the phosphoric acid. Then boil/distill the remaining mixture to see what else you can find.

Something else you might want to check. Does if foam or sud on shaking. I seem to remember some phosphoric acid cleaner we had around here having soap in it. This was probably 10 years ago so I may be mistaken. I wasn't the chem nerd I am now so I didn't do a full analysis.
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[*] posted on 27-1-2005 at 22:42


Yes Mumbles it foamed when I poured it into the 250 mL grad cylinder. It's loaded with soaps/detergents.

I just finished the extraction with Coleman lantern fuel. BTW as a correction this solvent is not strictly hexane but is a mixture of C5-C9 hydrocarbons according to the Coleman MSDS. I think this is more properly called "petroleum naptha.

I started out with 10 mL of the OTC phosphoric acid cleaner. I had to extract it 5 times with 5 mL of solvent before the foam disappeared. The dye stayed mostly with the aqueous (lower) layer but slowly disappeared. The foam of course was on top and was discarded with each solvent wash. The final yield was 6 mL of aqueous at a sp gr of 1.10. The original cleaner sp gr was 1.14. So, I conclude that this procedure is probably successful for cleaning up the phosphoric acid but it is labor intensive and the yield is low (~50%).

[Edited on 30-1-2005 by Magpie]




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[*] posted on 28-1-2005 at 11:27


So I could use a phosphate salt and hydrochloric acid to manufacture my own phosphoric acid? The process would go like this:

Na3PO4 + 3HCl ---> H3PO4 + 3NaCl

Where can I get sodium phosphate? and will sodium chloride precipitate out of my phosphoric acid? Also, will I be able to vacuum distill in order to obtain high a concentration of phosphoric acid without destroying my glassware?




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JohnWW
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[*] posted on 28-1-2005 at 11:51


That reaction is the opposite to what actually happens.
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[*] posted on 28-1-2005 at 12:36


Sodium phosphate can be acquired at almost every hardware store. It is sold as TSP, trisodium phosphate, it is used as a cleaner.



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S.C. Wack
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[*] posted on 28-1-2005 at 13:15


Quote:

That reaction is the opposite to what actually happens.


It takes some heat to get HCl, HI, HNO3, or SO3 from the corresponding salts. Excess HCl will prevent problems in any case. I've noticed the smell of Cl2 rather than HCl with NaCl/H3PO4, oddly enough. No reason why a cheap plastic vacuum dessicator/H2SO4 can't be used here, AFAIK. But yeah, there is a problem getting all of the NaCl to precipitate.

And so you may want CaSO4 as a byproduct instead.

Looking at the patent literature, they are using alcohols, mixed with a little aromatic or aliphatic alkane or ethers, to extract H3PO4 from the leftover salts.

The hardware store TSP, who knows what all else is in it in small amounts, but since it is sold in a leaking cardboard box, you can be sure that it contains some DSP and Na2CO3 from atmospheric CO2.
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[*] posted on 28-1-2005 at 17:12


Quote:

I like it for cleaning glassware, I guess that makes me an environmental terrorist.

[Edited on 29-11-2004 by S.C. Wack]


S.C Wack I have heard fromat least two sources that the phosphate pollution was exaggerated much like the cfc's eating all the ozone. It was a public panic move to get more expensive products to market. Think about it. Now the only refrigerants are tetrafluoroethane and after 15 years they are just now coming down in cost. The phosphates used in cleaners in teh 1960's where replaced by PETROLEUM derivatives such as EDTA , ethanolamines, ethoxy surfactants etc. Which of course cost more. You can thank Dow chemical, the pocketed politians, and the scare media for that.




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Magpie
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[*] posted on 29-1-2005 at 19:27


Tom Haggen wrote:


Quote:

Na3PO4 + 3HCl ---> H3PO4 + 3NaCl


And JohnWW responded:

"That reaction is the opposite of what actually happens"

JohnWW I don't understand why this would be true. HCl is a strong acid and should be completely dissociated. H3PO4 on the other hand is a weak acid with K1 = 7.5 x 10^-3. So it is basically undissociated. Please explain.




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[*] posted on 30-1-2005 at 13:30


Because HCl is much more volatile than H3PO4. Any HCl formed in an exothermic reaction, or in a mixture subject to heat, would be liable to be lost as vapor, tilting the equilibrium.
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[*] posted on 10-2-2005 at 02:01


Add H2SO4 to Na3SO4.
Na2SO4 will precipitate out at room or lower temps. Filter off Na2SO4 preferrably using vacuum filtration.

3H2SO4 + 2Na3PO4 --> 2H3PO4 + 3Na2SO4

H2SO4 Ka = very large
NaHSO4 Ka = 1.3 E-2

H3PO4 Ka = 7.5 E-3
NaH2PO4 Ka = 6.2 E-8
Na2HPO4 Ka = 4.8 E-13

The aqueous ionization constant for the monoprotic Hydrogen Sulfate ion is greater than that of even the triprotic Phosphoric acid;

Equillibrium shifts toward Phosphoric acid.

This plus the fact that H2SO4 is relatively non-volatile (please everybody don't start a flame-war over the volatility of H2SO4) means that the problem mentioned in the last response (dealing with HCl) is alleviated.

Depending upon the conc. of H2SO4, Na3PO4 may need to be dissolved in respectively varying amounts of water first.

Upon vacuum evaporation of product, the phosporic acid becomes concentrated and upon cooling the Na2SO4 precipitates even more. Again, vacuum filter off.

:oTC H3PO4




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[*] posted on 12-5-2005 at 16:44
phos.........try this


farm supply has the same thing for $5.00 a gallon.....two ways to release a hydro-bond...boiling....bad idea.....when white phos has less that 33% h2o....it blows....but you can release bond by using metals.....magnesium????????
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[*] posted on 8-1-2006 at 10:17


I made some phosphoric acid a few weeks ago using Rooto sulfuric acid and pottery grade (Ca)3(PO4)2. After vacuum filtering off the voluminous quantity of CaSO4*2H2O I estimated the 270 mL of filtrate to be H3PO4 at 11wt%. Since I made this without benefit of a lab procedure I'm sort of following the industrial process outlined in Shreeve's Chemical Process Industries. The dilute phosphoric is concentrated by vacuum evaporation. So I set up my vacuum distillation glassware and proceeded to concentrate it at 27"Hg vacuum. But violent bumping caused me to abort.

After making my own ebulliator tube and buying another thermometer adapter I was ready to try it again. This time things went well and my 270 mL acid was boiled down to about 50 mL. Final pot temperature was 55C.

The concentrate is cloudy, and fine white material has settled on the bottom of the bottle. I'm not sure what this is but suspect it may be CaSO4, or could it even be H3PO4?

I'm trying to come up with a good filter medium. I'm afraid to use paper as I don't know if it would disintegrate. I'm thinking of trying some nylon stocking. Any suggestions?

Edit: see picture below



[Edited on 8-1-2006 by Magpie]
EbC: Picture size!

[Edited on 26-1-2007 by chemoleo]

concentrated.jpg - 19kB




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[*] posted on 8-1-2006 at 13:34


Here's a picture of the vacuum distillation setup for the phosphoric acid concentration:

EbC: Picture size!

[Edited on 26-1-2007 by chemoleo]

vacuum.jpg - 67kB




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[*] posted on 8-1-2006 at 17:48


What was the final concentration of your acid?

You mentioned a vacuum of 27”. Did you mean 27mm?
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[*] posted on 8-1-2006 at 18:36


Neutrino I realize this is an importance piece of information, i.e., the concentration. I will find out tommorrow when I titrate it against some standard NaOH.

No it is 27"Hg vacuum. You can see this on the 2nd picture. 30"Hg is a perfect vacuum. You may be thinking of absolute pressure which in this case would be 3"Hg or 76mmHg absolute.




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[*] posted on 8-1-2006 at 19:40


Magpie, I tried to make some sulfuric acid via the wet process sometime last year, I got the details here. However I didn't get good results, I used concentrated sulfuric acid and calcium phosphate that was like dust. When half of the calcium phosphate was added I had a solid mass so more water was added. The whole procedure was a mess and in the end I only got a 10% yield or so, where did all the acid go, did it just stick in the calcium sulfate and never see the light of day, and it was difficult to filter, you had the right idea with using dillute solutions, when its all boiled down are you going to assay the purity in some way? I'm eager to see how things go.



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[*] posted on 9-1-2006 at 16:03


I titrated some of the supernate from my H3PO4 today. I titrated it to the 1st endpoint using 0.188N NaOH and phenol red indicator(pH at color change: 6.8-8.2). The result was 11.0 molar, or ~ 70wt%.

Bromic I see from your web site what the settled material is. Thanks for that information. I will filter it out of my acid.

I had to deal with a lot of gypsum also. That is the main reason I decided to add a 6" Buchner funnel to my lab equipment. I also used diatomaceous earth which helps a lot with snotty precipitates.

My reaction charge was 104 g (Ca)3(PO4)2, 64 mL of 86% sulfuric acid, and 18 mL H2O. I ended up adding another 500 mL H20 just to make the slurry filterable . Then vacuum filtered to remove the gypsum. This was tedious but not unpleasant. Next time having my 6" (15cm) Buchner, which is 4 times larger than my 7cm Buchner, will really help. After all the vacuum filtering I ended up with just 270 mL of filtrate. This was then vacuum evaporated as described above.

Yield = [(~50mL)(0.70)(1.526g/mL)/98.08]/[(2moles/mole)(104 g/310.18)]x100% = 81%

Edit:

Phenol red isn't the best indicator for this titration so I ran it again this time using methyl orange (pH at color change: 3.0-4.4). This titration gave a more realistic result of 8.4 molar, or ~58wt% H3PO4.

(These are my only two indicators as they were readily available at my local pool supplier.)

Revised yield = [(~50mL)(0.58)(1.46g/mL)/98.08]/[(2moles/mole)(104 g/310.18)]x100% = 64%

[Edited on 10-1-2006 by Magpie]




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[*] posted on 9-1-2006 at 16:20


US 4390509 Process for manufacturing ammonium phosphate utilizing an oxalic acid acidulating process
oxalic acid + phosphate --> phosphoric acid
http://v3.espacenet.com/origdoc?DB=EPODOC&IDX=US4390509&...


US 3645682 A PROCESS FOR PRODUCING PHOSPHORIC ACID BY THE USE OF ION EXCHANGE RESINS
http://v3.espacenet.com/origdoc?DB=EPODOC&IDX=US3645682&...

OTC oxalic acid
http://www.dap.com/product_details.aspx?product_id=324
http://www.dap.com/docs/tech/00079208.pdf


Ion exchange resins are used in water softening.
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[*] posted on 14-12-2010 at 02:05


In the bones there is ossein and calcium phosphates. I read on a lot of places that HCl can dissolve calcium phosphates from the bone, but not ossein. If it dissolves them then phoshoric acid is left free. We can filter osein and if we used more bones than HCl, we will have solution of phosphoric acid and CaCl2 :D

Edit:

From wikipedia:

Preparation of hydrogen halides
Phosphoric acid reacts with halides to form the corresponding hydrogen halide gas (steamy fumes are observed on warming the reaction mixture). This is a common practice for the laboratory preparation of hydrogen halides.
NaCl(s) + H3PO4(l) → NaH2PO4(s) + HCl(g)
NaBr(s) + H3PO4(l) → NaH2PO4(s) + HBr(g)
NaI(s) + H3PO4(l) → NaH2PO4(s) + HI(g)


If we would use more bones than HCl, we would just filter excess of bones and ossein, then wait until some phosphoric acid reacts with cacl2 to hydrogen chloride and calcium phosphate. What would be left is phosphoric acid I think with calcium phosphate that can be filtered.

[Edited on 14-12-2010 by Random]

[Edited on 14-12-2010 by Random]
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[*] posted on 29-12-2012 at 19:25


Some time ago I made about 50mL of ~70% phosphoric acid using pottery grade Ca3(PO4)2 and Rooto con H2SO4. I have used up my supply so am making another batch. The reaction is:

Ca3(PO4)2 + 3H2SO4 + 6H2O --> 2H3PO4 + 3CaSO4*2H2O

This is consistent with the "Dorr strong acid" industrial process for making H3PO4, and the reaction seems to proceed satisfactorily for my home chemistry preparations. However, in reviewing this I have come across something that doesn't make sense, ie,:

The Ksp for CaSO4*2H2O = 2.4 10^-5 @25°C
The Ksp for Ca3(PO4)2 = 1 x 10^-25 @25°C

Given these two values for the solubility products, how is it possible for the above reaction to proceed to apparent completion?
----------------------------------
Edit: In case anyone is interested, the answer is: in an acidic solution the PO4--- ion is in very low concentration due to the formation of the weak acids H3PO4, H2PO4- , and HPO4--. SO4-- forms an acid also, ie, HSO4-, but it is not anywhere near as weak.

[Edited on 31-12-2012 by Magpie]




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[*] posted on 30-12-2012 at 06:55


An OTC alternative is Dairyland milkstone remover and acid rinse, available at Tractor Supply. It's 42% phosphoric acid, and while it still contains surfactants and dye, it is supposedly non-foaming. http://www.tractorsupply.com/cattle-handling/dairyland-milks...

Anyway, I have no good suggestions on how to remove the cleaners and degreasers. Seems like you'd need to know what they are.




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[*] posted on 30-12-2012 at 07:38


Almost 2 years after my last post, I found a passing remark by Ditte [Compt. rend. 90, 1163 (1880)] on sodium phosphates and HCl. He speaks of saturation with HCl gas; I didn't saturate.

"Je ferai remarquer seulement que l'action de l'acide chlorhydrique sur le phosphate de soude permet de préparer l'acide phosphorique avec une très grande facilité. Il suffit de diriger, dans une solution de phosphate de soude, un courant d'acide chlorhydrique, de manière à saturer la liqueur: tout le sel marin se précipite; le liquide clair, décanté et distillé, dégage de l'acide chlorhydrique qui peut servir à une opération nouvelle, et le résidu dans l'appareil distillatoire consiste en acide phosphorique sirupeux pur."





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[*] posted on 30-12-2012 at 07:56


WOW, that is interesting to know!

Now I will be able to make my own phosphoric acid with Trisodium phosphate sold at the store as TSP.

I should read more often french text....




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