Sciencemadness Discussion Board

OTC Phosphoric Acid

tom haggen - 28-11-2004 at 23:52

Man, I can't believe how big of a slacker I am. With how much school work I need to do there is no reason why I should be posting messages right now. But here I go, Someone had mentioned that you can obtain Phosphoric acid OTC in another thread so I thought I would do some investigating. I went down to my local Hydroponics store where I get my 35% H2O2 by the gallon, and looked around. The product they had called PH Down contained a mixture of citric acid, phosphoric acid, and either ammonium bisulphate or ammonium bisulphite. I'm wondering If I will be able to separate the phosphoric acid by fractional distillation? I know that citric acid decomposes at a certain temp so I'm wondering how hard this will be. Any suggestions? I would also like to isolate the ammonium bisulphite cause I might be able to use that for titrations? I might be wrong about that last part but if anyone knows, let me know so I can hit two birds with one stone. Look out finals here I come!

Tacho - 29-11-2004 at 03:58

Phosphoric acid is sold OTC to prevent rust on steel. Look at your hardware store.

The brand name "Naval Jelly" comes to my mind. I think its an US brand.

Also serch for "rust converter".

Dont bother

chloric1 - 29-11-2004 at 04:01

Its not worth it. The temperature your citric acid would decompose the phosphoric acid would etch your glass! Thats right! Modest phosphoric acid attacks silicates at high temps. ONe method I ponder is adding muriatic acid to saturated TSP to precipitate NaCl. Sure you would still need to boil excess HCL but the etching of glass would be
negligable until the phosphoric acid is concentrated.


Better still would be vacuum evaporation of the HCL so lower temperautures could be used!

[Edited on 11/29/2004 by chloric1]

Organikum - 29-11-2004 at 04:54

Is TSP
- tri-sodium-phosphate
or
- triple-super-phosphate

or this this the same anyways?

confused
/ORG

vulture - 29-11-2004 at 06:08

Why is it that phosphoric acid is used for removing rust from steel? Is it because it passivates iron even when dilute?

The_Davster - 29-11-2004 at 06:08

It is trisodium phosphate, Na3PO4.

Source: The bag

Mendeleev - 29-11-2004 at 08:32

The rust remover is only 25%, but certain pH down brands contain 75% with nothing else added. It all depends on the brand, because some don't even have phosphoric acid, they use dilute nitric acid. You just have to look around. Oh, off-topic but interesting, Wink rust remover is 3% HF.

Mr. Wizard - 29-11-2004 at 09:18

Quote:
"Is TSP
- tri-sodium-phosphate
or
- triple-super-phosphate

or this this the same anyways?"
TSP is Tri Sodium Phosphate, except when you buy it in boxes labeled "TSP" , then it is most likely a cheaper non Phosphate detergent (Sodium Carbonate or Sodium Sesqui-carbonate, or silicate) that the makers have put in a box with a Trade Marked name of "TSP". This little bit of deceptive advertising takes advantage of the ignorance of most people who buy the product to clean walls before painting them. Checking the ingredients on the box will guide you. If the label spells out the words Tri Sodium Phosphate, it is most likely a pure product, if it just uses the initials "TSP", it isn't. Tri Sodium Phosphate has long grain like crystals and shouldn't foam when treated with a mild acid such as vinegar.

tom haggen - 29-11-2004 at 11:10

Well the stuff at my local hydroponics store is quite nasty indeed. Not only is it mixed with citric acid and ammonium bisulfite, it is contaminated with some nasty orange dye. So I don't really think I understood any of the answers given other than hot phosphoric acid will etch your glassware. That is a big fucking no no. Would you be able to separate the phosphoric acid in a vacuum distillation setup? Or would this too destroy your glass. On a side note just how many solvents out there will destroy your glass? I know NaOH will, and apparently Phosphoric acid will, any thing else come to anyone’s mind? Being some one that has worked with metal for a few years now I can honestly say I have never seen phosphoric acid used for removing rust. The only acid I have seen on the shop floor for these types of uses is hydrochloric acid. Now I've seen Nitric acid used in the metallurgical lab used for etching clean surfaces in order to observe crystalline lattice structures of metal. But no, I haven't seen phosphoric acid used in the metal shop.

[Edited on 29-11-2004 by tom haggen]

BromicAcid - 29-11-2004 at 11:19

Phosphoric acid doesn't have a boiling point that I'm aware of. It decomposes to pyrophosphoric acid and other polymeric solids, maybe at very high temps these will distill (>800C) but I'm not certain.

tom haggen - 29-11-2004 at 11:28

When I read the msds on phosphoric acid it stated that the bp of phosphoric acid was about 158C. But this was assuming you had anhydrous phosphoric acid. Separating the phosphoric acid is starting to sound like a problem that can't be solved,:(

BromicAcid - 29-11-2004 at 11:56

According to Hawley's Condensed Chemical Dictionary regarding the boiling point of phosphoric acid:
Quote:
... the 100% acid is in the form of crystals, d 1.834 (18C), mp 42.35 C, loses 1/2 H2O at 213C (to form pyrophosphoric acid)...

Regarding pyrophosphoric acid:
Quote:
A viscous, syrupy liquid that tends to solidify on long standing at room temperature. When diluted with water it is rapidly converted into orthophosphoric acid. Mp 54C. Soluble in water.
Derivation: By heating phosphoric acid at 250 - 260C. Further heating produces metaphosphoric acid.

Regarding metaphosphoric acid:
Quote:
Derivation: By heating orthophosphoric acid to redness

No boiling point in sight from my end, sorry. Although by the end of everything most everything else, the citric acid, ammonium bisulfide, and dye should all have volatized out, your resulting metaphosphoric acid could then be re-hydrolyzed. I bought some phosphoric acid once for cleaning stainless steel from a cooking supply place, it had detergents in it and when I tried to heat to get rid of everything it foamed over and the bubbles caused my glassware to crack.

S.C. Wack - 29-11-2004 at 12:00

Fluorides, hydroxides and phosphoric are the only glass etchers AFAIK. I have seen hot and concentrated phosphoric etch some of my glass in some long experiments - it looks nothing like what the hydroxides do, btw.

Phosphates are not hard to find around here. The Ca phosphates from the nursery were grey and smelled bad - that did not go well.

Home Depot sells TSP, in a box that says TSP, and is not crystalline. Yet it really is trisodium phosphate. Several other chains only sell "TSP substitute"s due to the phosphate thing.

The local chemical salvage store has barrels of the mono- and tri- Na phosphates, cheap - no one wants them.

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

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

Polverone - 29-11-2004 at 12:10

I've heated and concentrated phosphoric acid up to the metaphosphoric acid stage a number of times, and never noticed any attack on glassware. If there was any attack, it was much slower and less noticeable than that of HF or molten hydroxides.

Mendeleev - 29-11-2004 at 13:38

What about concentrated NaOH solution, will that do anything?

neutrino - 29-11-2004 at 14:44

Polverone: What type of glass is this? Phosphoric acid attacks Pyrex, so you must have some pretty good glass there.

tom haggen - 29-11-2004 at 14:48

What links phosphates to terrorist activity if you don't mind me asking?

EDIT: This topic sounds a little more advanced than I had first anticipated. I guess I will need to learn a few more things about the chemistry of phosphoric acid before I decide to continue this endeavor.


[Edited on 29-11-2004 by tom haggen]

[Edited on 29-11-2004 by tom haggen]

HNO3 - 29-11-2004 at 15:15

At Wal-Mart I can get phosphoric acid in the hardware section. It is sold as Concrete etchant. Never having needed phosphoric acid, I have never gotten any, so I can't tell you much about it, but the 1 gal. (3.78 L) bottle says its phosphoric acid based.

skippy - 29-11-2004 at 15:16

Polverone,

Was the metaphosphoric acid for elemental phosphorus production? If so, I'd love to hear about such activities :)

S.C. Wack - 29-11-2004 at 15:21

Quote:

What links phosphates to terrorist activity if you don't mind me asking?

Look on the labels of various cleansers since the 70's and you will probably see "contains NO phosphates". There was a bit of change when it became widely known that soluble phosphates change bodies of water, or rather what lives (but mostly dies) in them. Certain plants love the phosphate and cause some sort of oxygen problem IIRC. I'm sure it's in books and on Google.

Given the sewer - er - river that sewage is dumped into around here (on its way to the Mississippi), often raw, I have no qualms about adding an ounce of phosphate to it. It's too late.

budullewraagh - 29-11-2004 at 16:06

phosphorus is a shady element. i have a book about it called "the thirteenth element: a sordid tale of murder, fire and phosphorus"

interesting book. i think the main reason phosphorus compounds aren't too available is the fact that it isn't too hard to make certian organophosphates out of them...(ex: tabun, sarin)

Polverone - 30-11-2004 at 03:05

Yes, the metaphosphate production has mostly taken place in the course of attempts at phosphorus production. The most recent time, I was using an ordinary borosilicate test tube and simply wished to heat the acid as much as I could. The acid was heated until it turned to a solid, and it did not melt again even at the melting point of the glass. I had some hopes that it might have been dehydrated nearly to P2O5 and that I would get a strong reaction with water, but there was no such luck. I didn't notice the glass being attacked, though of course the final deformed tube resembled the starting item very little and I actually broke the glass to free the solid acid. I have also concentrated phosphoric acid (though not to that degree) in ordinary soda lime glass and not noticed any attack on the glass after it was cleaned out. What is the chemical mechanism of phosphoric acid attack on glass? I've seen enough references about it to trust that it occurs, but I've had enough experiences to say that it isn't nearly as aggressive as HF or molten alkali hydroxides.

Blackout - 3-12-2004 at 14:36

I've found a nitric/phosphoric acid mix at my local hardware store...

How can we synth. phosphoric acid from elemental P?

budullewraagh - 3-12-2004 at 14:55

depends on your allotrope. if it's white or red you can easily oxidize it. oxidize until you get the anhydride and then add to water

Magpie - 27-1-2005 at 12:23

I'd been wanting some phosphoric acid for some time (just to have it) and knew that Home Depot had a "Phosphoric Acid Cleaner" made by Aqua-Mix. When I got it home and poured some into a graduated cylinder (to check the sp. gr. by hydrometer) I noticed it was pink! Then I read the label a little more carefully and found out it contains "cleaners and degreasers." I added a little food grade activated carbon to a few mL. This eliminated the pink dye but now I'm wondering what other shit is in there and how I might remove it. I checked the MSDS on Aqua-Mix's website but it only lists phosphoric acid at 23% by volume (which corresponded to my sp. gr. reading btw).

Eventually I want to concentrate some of this acid up to 88% by evaporation. After reading this thread it seems I may want to do this at a reduced pressure to save my glassware.

If anyone has any suggestions about how I might remove the "cleaners and degreasers" I would apprecitate hearing them.

S.C. Wack - 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.

Magpie - 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.

Mumbles - 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.

Magpie - 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]

tom haggen - 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?

JohnWW - 28-1-2005 at 11:51

That reaction is the opposite to what actually happens.

The_Davster - 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.

S.C. Wack - 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.

chloric1 - 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.

Magpie - 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.

JohnWW - 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.

Protium - 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

phos.........try this

chemistkat - 12-5-2005 at 16:44

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????????

Magpie - 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

Magpie - 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

neutrino - 8-1-2006 at 17:48

What was the final concentration of your acid?

You mentioned a vacuum of 27”. Did you mean 27mm?

Magpie - 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.

BromicAcid - 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.

Magpie - 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]

Bill - 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.

Random - 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]

Magpie - 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]

bbartlog - 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.

S.C. Wack - 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."

plante1999 - 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....

blogfast25 - 30-12-2012 at 08:03

Plante, you really need to see a doctor, if you haven't already: there are so many treatments out there you wouldn't believe it. I've just been treated for something I've been suffering with for weeks (not chemistry related), believing there was no cure for it but time itself, and just two treatments have already made a world of difference.

Ooopsie: wrong thread!

[Edited on 30-12-2012 by blogfast25]

plante1999 - 30-12-2012 at 08:17

Quote: Originally posted by blogfast25  
Plante, you really need to see a doctor, if you haven't already: there are so many treatments out there you wouldn't believe it. I've just been treated for something I've been suffering with for weeks (not chemistry related), believing there was no cure for it but time itself, and just two treatments have already made a world of difference.

Ooopsie: wrong thread!

[Edited on 30-12-2012 by blogfast25]


I will see what I can do...

PS:Use the delete button or edit the content of the post!

Fantasma4500 - 2-1-2013 at 10:54

i remember watching this video from myst32YT (http://www.youtube.com/watch?v=S6GpMQ98u8I) where hes going for getting the nitric acid, but i also remember one of the other things that was in...
phosphoric acid, and water..
if you look at the video at 0:55 you can see the contents..
as this is described as a OTC source of nitric acid, then i guess you can get both nitric acid AND phosphoric acid as a nice OTC source of both... right?
i dont live in usa, nor do i like chemicals that etches glass :x
so i guess i wont be making concentrated phosphoric acid, or purifying it.

just thought it might be helpful, perhaps this might even be able to be bought other places in the world? i doubt usa is the only place they use nitric acid for cleaning concrete??

btw have anybody made a sticky thread yet where OTC sources of 'interesting' chemicals are posted?

Scr0t - 2-1-2013 at 11:34

Quote: Originally posted by Antiswat  

just thought it might be helpful, perhaps this might even be able to be bought other places in the world? i doubt usa is the only place they use nitric acid for cleaning concrete??


Over here a similar mixture H2O/H3PO4/HNO3 is sold as "Acid Descaler", It's used by dairy farmers to remove milkstone and associated protein deposits.

Some contain H2SO4 instead of H3PO4 and sometimes they contain surfactants.

The one I use is "Biocell, Best acid descaler", about 52% /w total acid and about ~33% /w H3PO4 85%.

I usually just evaporate the H2O/HNO3 into the atmosphere from a Pyrex dish and when Temp. reaches 180°C throw in some activated carbon pellets to reduce any residual HNO3, if any (depends on how quickly it's been evaporated).

No etching of the Pyrex has been observed for temperatures up to 220°C.

elementcollector1 - 14-1-2014 at 20:44

I recently added 40g of sulfuric acid to about 67g of sodium orthophosphate, as per stoichiometry to make sodium dihydrogen phosphate. However, what I did not expect was a gigantic amount of gas produced. This may be boiling water from the heat of the sulfuric acid and the hydration of the phosphate, but the odd thing is that it kept bubbling long after it should have been 'extinguished', when I added about 250 mL of water to the flask in an ice bath. The temperature is now well below room temperature, and yet it is still bubbling. Curiously, there is a yellow deposit on the side of the flask.

What is this gas? Why is there so much of it? It was certainly acidic from the fumes I smelled (at first), but now it doesn't smell like anything at all.

Did I lose all of my phosphoric acid/sodium dihydrogen phosphate?

Edit: There is now a milky white bilayer that suggests a precipitate. The odd yellow spot has disappeared. This was filtered off, and found to be translucent white crystals of a small and easily filterable nature. Given the original appearance of the "TSP" as a white, fine powder, and the nature of the reaction (it was initially the dry powder and concentrated sulfuric acid), it is unlikely that this is unreacted sodium orthophosphate. It could be sodium sulfate - I will have to check with a soluble calcium salt when I get time. It is actually unlikely to be the target sodium dihydrogen phosphate, as this has a solubility of 60g/100mL water at 0 degrees C. All of that should be in the supernatant liquid, as well as (likely) some of the sulfate. I will check the density of the precipitate tomorrow, first comparing with that of trisodium phosphate, then sodium dihydrogen phosphate and sodium sulfate. I will also check the density of the supernatant, to see if it corresponds with the density of 40g sulfuric acid in about 200 mL water.

I know a reaction happened, but given the aforementioned release of copious quantities of gas, I am confused as to what it is. Could there be stabilizers containing carbonates in my TSP?

In addition, a significant portion of the TSP (visual estimate: ~20g) went unreacted. This appears to be covered in a protective, transparent layer of *something* about a millimeter thick, and tons of small bubbles are formed for a brief period of time when this is shaken with the filtered supernatant liquid.

I think that next time, I will just distill the phosphoric acid and partially neutralize with NaOH...



[Edited on 1-15-2014 by elementcollector1]

S.C. Wack - 15-1-2014 at 19:55

Quote: Originally posted by elementcollector1  
What is this gas?


Well geez you're the chemist...freaking analyze it. Or look back to what was said 9 years ago.

BTW I suspect that the strong etching I got was in fact from HF, from OTC grade phosphate, or it was the glass or both.

blogfast25 - 16-1-2014 at 05:50

Quote: Originally posted by S.C. Wack  
[Or look back to what was said 9 years ago.



So what was said 9 years ago?

IrC - 16-1-2014 at 14:04

Quote: Originally posted by Antiswat  
btw have anybody made a sticky thread yet where OTC sources of 'interesting' chemicals are posted?


http://www.sciencemadness.org/talk/viewthread.php?tid=3254

posted on 1-19-2005 at 07:00 AM

Would not this qualify? Somewhere in 7 pages surely something 'interesting' must be in there?

blargish - 17-1-2014 at 07:29

I was browsing through Canadian Tire the other day and stumbled upon something in the pool section called Aquarius Flush. The MSDS says that it's 30 - 60% phosphoric acid.

Anyone else have experience with something similar?