laboratory grade NaH2PO4 - The Planted Tank Forum
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post #1 of 16 (permalink) Old 01-17-2014, 03:01 AM Thread Starter
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laboratory grade NaH2PO4

So I just got some from a chemist and need to get about 85 gallons of water from 0ppm of phosphates to 5ppm.

Its in dry form and I have no idea what I'm doing lol.

Can anyone help?

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post #2 of 16 (permalink) Old 01-17-2014, 03:46 AM
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Hi Jarod,

You need to convert gallon to litre, ie: 85g = 322L
5ppm phosphates means 5mg phosphate per litre of water, which equals: 6.3mg NaH₂PO₄ per litre of water, or 322 x 6.3mg = 2029mg, or ≈ 2gram for the whole tank.

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post #3 of 16 (permalink) Old 01-17-2014, 04:11 AM Thread Starter
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Originally Posted by Greystoke View Post
Hi Jarod,

You need to convert gallon to litre, ie: 85g = 322L
5ppm phosphates means 5mg phosphate per litre of water, which equals: 6.3mg NaH₂PO₄ per litre of water, or 322 x 6.3mg = 2029mg, or ≈ 2gram for the whole tank.
Hmmm, too bad I know longer remember what 2 grams looks like.. ehehehehhe.

I sprinkled a little bit in the tank and I'll test phosphates in the morning. I have to play it safe because I have no idea how to measure grams without a scale.

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post #4 of 16 (permalink) Old 01-17-2014, 06:46 AM
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Hehe you should have had your chemist friend do the math for ya real quick.

NaH2PO4 has a molar mass of 120 g/mol
PO4 has a molar mass of 95 g/mol
Ratio of PO4 to the other elements is 79% PO4 to 21% NaH2
This ratio is necessary because ions split in aqueous solution
To reach 5 mg/L of NaH2PO4 in 322L of water

5 mg/L * 322 L = 1610mg NaH2PO4
That is only 79% of the PO4 you need so...
1610mg * 1.21 = 1948mg or about 2 grams of NaH2PO4 needed

Curious: when you add this to water does the H2(g) bubble out?

Last edited by Positron; 01-17-2014 at 06:50 AM. Reason: sadf
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post #5 of 16 (permalink) Old 01-17-2014, 06:55 AM
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post #6 of 16 (permalink) Old 01-17-2014, 10:53 AM
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Originally Posted by Positron View Post
Curious: when you add this to water does the H2(g) bubble out?
No. Hydrogen gas is not formed from the dissolution of sodium phosphate.

Anthony


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post #7 of 16 (permalink) Old 01-17-2014, 02:23 PM
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If you don't have a scale than you should turn the monosodium phosphate (NaH₂PO₄) into a saturated solution.
At room temperature 85g of the salt will dissolve in 100mL of water.
Just get a beaker of warm water and keep dissolving phosphate in it until no more will dissolve.

At room temperature the saturated solution has a density of 1.4 gr/mL. Therefore it will contain ≈ 650mg of NaH₂PO₄ per ml of liquid.
To get 2g, you need to measure 3mL of solution with a syringe.

Anything else I can help you with?

Regards
Cor
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post #8 of 16 (permalink) Old 01-18-2014, 12:31 AM Thread Starter
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Quote:
Originally Posted by Greystoke View Post
If you don't have a scale than you should turn the monosodium phosphate (NaH₂PO₄) into a saturated solution.
At room temperature 85g of the salt will dissolve in 100mL of water.
Just get a beaker of warm water and keep dissolving phosphate in it until no more will dissolve.

At room temperature the saturated solution has a density of 1.4 gr/mL. Therefore it will contain ≈ 650mg of NaH₂PO₄ per ml of liquid.
To get 2g, you need to measure 3mL of solution with a syringe.

Anything else I can help you with?
Ah ha, yes the saturated solution. Thanks its been a veryyy long time since I've had basic science classes lol

Thanks to everyone else too. I appreciate it and if its worth anything, I think I'm beggining to understand.

Jarod
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post #9 of 16 (permalink) Old 01-18-2014, 01:43 AM
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So I just got some from a chemist and need to get about 85 gallons of water from 0ppm of phosphates to 5ppm.

Its in dry form and I have no idea what I'm doing lol.

Can anyone help?
Hi anastasisariel,

1/2 teaspoon will raise 85 gallons of water from 0.0 ppm to 4.68 ppm.

Roy_________
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post #10 of 16 (permalink) Old 01-18-2014, 01:46 AM Thread Starter
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Hi anastasisariel,

1/2 teaspoon will raise 85 gallons of water from 0.0 ppm to 4.68 ppm.
Thanks for doing my homework sir!

Jarod
I'm in the Saint Louis, MO area and always looking to connect with other planted tank enthusiasts.

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post #11 of 16 (permalink) Old 01-18-2014, 01:51 AM
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Hi Jarod,

As an ex-St. Louis resident and MASI member it was my pleasure to help you out!

Roy_________
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post #12 of 16 (permalink) Old 01-18-2014, 02:38 AM Thread Starter
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Hi Jarod,

As an ex-St. Louis resident and MASI member it was my pleasure to help you out!
Just went to my first meeting last night.. was amazing!

Jarod
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post #13 of 16 (permalink) Old 01-19-2014, 04:29 PM
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If you don't have a scale than you should turn the monosodium phosphate (NaH₂PO₄) into a saturated solution.
At room temperature 85g of the salt will dissolve in 100mL of water.
Just get a beaker of warm water and keep dissolving phosphate in it until no more will dissolve.

At room temperature the saturated solution has a density of 1.4 gr/mL. Therefore it will contain ≈ 650mg of NaH₂PO₄ per ml of liquid.
To get 2g, you need to measure 3mL of solution with a syringe.

Anything else I can help you with?
Excellent advise! I've never thought to use the solubility of fertilizers to create a stock concentration. Looked up the solubility of the common fertilizers below.

At 20 degrees celcius

KNO3 470mg/ml
KH2PO4 226mg/ml
K2SO4 111mg/ml
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post #14 of 16 (permalink) Old 01-20-2014, 03:40 AM
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Good luck!

If you're going to try this method, do it with ordinary kitchen salt first. Keep the water warm when mixing the salt. Let it cool down a while to see if the solution will leave solid crystals.
You will soon find an easy way to make saturated solutions.

Let's exchange info so we can agree on concentration levels. (Sofar, I can only agree with the K₂SO₄ level @ 118mg/mL @25C)

Regards
Cor
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post #15 of 16 (permalink) Old 01-20-2014, 05:17 AM
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I was just going to suggest doing what Greystoke did. Super saturate by bringing the water to almost boiling and then letting stuff precip out for half a day. Then you know you have a fully saturated solution at the listed concentrations!

I'm only a physics major or else I would have suggested doing this a few days ago

EDIT: I remember reading somewhere in my general chem class (years ago) that sometimes it's possible for a solution to stay supersaturated under the right conditions. Like putting an object into the beaker of a supersaturated solution will quickly make things suddenly fall out. Weird, but perhaps this wouldn't be the best method after all?
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