Randy Holmes-Farley
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Vanadium pentoxide- need help from the big brains
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salty joe
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I tried to dissolve 1/2 g vanadium pentoxide in 250 mL water (half strength). I shook the daylights out of it but there were still some solids. so I added some calcium hydroxide and now I have over 1/4" of light orange fluffy sediment. No way would the amount of calcium hydroxide I added equal that volume. Also, the water is now almost clear.You're welcome.
Let us know if you hit any snags.
The good news is I still have 99.5 g of vanadium pentoxide.
Randy Holmes-Farley
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I'm not sure I understand what you are saying, but perhaps you added too much and have excess calcium hydroxide?
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In general, I would not add excess solids of anything since trace elements can bind to the solid surfaces. It happens, for example, to copper with excess calcium hydroxide:
reefs.com
To start, I made a solution of 1.0 ppm copper as copper sulfate (ACS reagent grade) in deionized water (700 mL). It had no detectable color by eye. When this sample was analyzed by ICP, it had an easily quantified set of emission peaks for copper, and this emission intensity was used as a standard (1.0 ppm).
To this solution I added 2.4 grams of calcium hydroxide (ACS Reagent Grade) and stirred it on a magnetic stirrer for 1 hour. This amount of lime is far more than necessary to saturate this solution (which would take just over 1 g of Ca(OH)2). The solution then settled for 24 hours, and a slightly cloudy sample was removed. This sample was analyzed and found to contain 340 ppb copper before filtration and 133 ppb copper after filtration through a 0.45mm polypropylene syringe filter. The value of 133 ppm copper represents a droop of 87% in the copper concentration.
The original solution was allowed to continue settling for 6 days total, and another sample was withdrawn. This sample was found to contain 160 ppm copper before filtration and 125 ppm copper after filtration through a 0.45mm polypropylene syringe filter.
Additional calcium hydroxide was added to the original solution (2 g/600 mL), the solution was stirred for an hour, and the solution was allowed to settle for 24 hours. This sample was found to contain 33 ppb copper before filtration and 25 ppb copper after filtration through a 0.45mm polypropylene syringe filter. The value of 25 ppb copper represents a drop of 80% from the 125 ppm copper solution just before this second lime addition, and a total drop of 97.5% from the initial copper sulfate solution.
As a control, I tested a solution of calcium hydroxide (2 g/50 mL of deionized water) without any copper sulfate added. In analyzing this sample, I could detect no real signal from copper, either before or after filtering. This result implies that the “natural” copper concentration in these samples is below about 10 ppb.
The first result might be explained simply by errors in the literature values or on my part. The Dyer et al article shows that these calculations can be off and should be used as a guide rather than a strict value.17 But the second drop is really only consistent with an entirely different phenomenon: that of surface absorption or coprecipitation. The binding of metals to the surface of inorganic materials, such as calcium carbonate, iron oxide or hydroxide, alumina, and clay (e.g., kaolin) is well known.33-44 In this case, both calcium carbonate and calcium hydroxide are present in these solutions, and these minerals provide surfaces to which copper and other metals may absorb strongly. Additionally, as CaCO3 is precipitated in such mixtures, copper and other metals are likely to get incorporated into the growing crystals, further lowering the solution metal concentration.
Consequently, the purification of limewater may not be entirely driven by pH, but also by the presence of “clean” mineral surfaces to which impurities such as copper will adhere.
Chemistry And The Aquarium: Metals In Limewater
This article expands on the basis of the claim that limewater may be able to deliver lower quantities of some metals compared to calcium reactors and it also provides some experimental data that demonstrates this claim.
reefs.com
Precipitation of Copper from Limewater: Experimental Results
The fact that a number of aquarists have reported blue precipitates in their limewater residue suggested to me that there may be something more to this story. Sure, they could have been using especially impure lime, or had especially impure source water. Nevertheless, it was easy enough for me to run some experiments to see what was really happening. In all of the experiments to follow, I monitored copper concentrations using a modern analytical lab technique: Inductively Coupled Plasma (ICP) using atomic emission detection. I used two different emission peaks (324.754 and 327.395 nm).To start, I made a solution of 1.0 ppm copper as copper sulfate (ACS reagent grade) in deionized water (700 mL). It had no detectable color by eye. When this sample was analyzed by ICP, it had an easily quantified set of emission peaks for copper, and this emission intensity was used as a standard (1.0 ppm).
To this solution I added 2.4 grams of calcium hydroxide (ACS Reagent Grade) and stirred it on a magnetic stirrer for 1 hour. This amount of lime is far more than necessary to saturate this solution (which would take just over 1 g of Ca(OH)2). The solution then settled for 24 hours, and a slightly cloudy sample was removed. This sample was analyzed and found to contain 340 ppb copper before filtration and 133 ppb copper after filtration through a 0.45mm polypropylene syringe filter. The value of 133 ppm copper represents a droop of 87% in the copper concentration.
The original solution was allowed to continue settling for 6 days total, and another sample was withdrawn. This sample was found to contain 160 ppm copper before filtration and 125 ppm copper after filtration through a 0.45mm polypropylene syringe filter.
Additional calcium hydroxide was added to the original solution (2 g/600 mL), the solution was stirred for an hour, and the solution was allowed to settle for 24 hours. This sample was found to contain 33 ppb copper before filtration and 25 ppb copper after filtration through a 0.45mm polypropylene syringe filter. The value of 25 ppb copper represents a drop of 80% from the 125 ppm copper solution just before this second lime addition, and a total drop of 97.5% from the initial copper sulfate solution.
As a control, I tested a solution of calcium hydroxide (2 g/50 mL of deionized water) without any copper sulfate added. In analyzing this sample, I could detect no real signal from copper, either before or after filtering. This result implies that the “natural” copper concentration in these samples is below about 10 ppb.
Precipitation of Copper from Limewater: Interpretation
Clearly, the first addition of excess calcium hydroxide dropped the copper concentration from 1 ppm to 133 ppb even though 133 ppb is below the solubility of copper at ANY pH according to literature values.17 Likewise, the addition of a second portion of solid calcium hydroxide to the solution, which should not have impacted the pH at all as there was already substantial excess solid calcium hydroxide before the second addition of solid lime, dropped the concentration from 125 ppb to 25 ppb.The first result might be explained simply by errors in the literature values or on my part. The Dyer et al article shows that these calculations can be off and should be used as a guide rather than a strict value.17 But the second drop is really only consistent with an entirely different phenomenon: that of surface absorption or coprecipitation. The binding of metals to the surface of inorganic materials, such as calcium carbonate, iron oxide or hydroxide, alumina, and clay (e.g., kaolin) is well known.33-44 In this case, both calcium carbonate and calcium hydroxide are present in these solutions, and these minerals provide surfaces to which copper and other metals may absorb strongly. Additionally, as CaCO3 is precipitated in such mixtures, copper and other metals are likely to get incorporated into the growing crystals, further lowering the solution metal concentration.
Consequently, the purification of limewater may not be entirely driven by pH, but also by the presence of “clean” mineral surfaces to which impurities such as copper will adhere.
salty joe
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My orange vanadium solution cleared up just like the blue copper solution did from the linked article.
Since I don't have any sodium hydroxide, maybe I'll see if I can get 1/4 g vanadium pentoxide to dissolve in 250 mL water.
Will temperature affect solubility?
Would kalkwasser likely not be an issue assuming no solids are in it?
Since I don't have any sodium hydroxide, maybe I'll see if I can get 1/4 g vanadium pentoxide to dissolve in 250 mL water.
Will temperature affect solubility?
Would kalkwasser likely not be an issue assuming no solids are in it?
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Randy Holmes-Farley
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Color is a poor indicator of vanadium because different forms of it have different colors.
I'd also suggest doing this differently than adding excess.
I am not 100% sure that calcium hydroxide is a good idea, since there are many forms of vanadate that may be present.
Sodium hydroxide is cheap and easy to get, but in the absence of it, here's what I'd do:
Take some clear kalkwasser with no solids on the bottom.
Add a small amount of the vanadium pentoxide and see if it dissolves. If it does, then one can make a dosing solution that way.
salty joe
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I decided to use sodium hydroxide after all. Can I add small amounts of dry sodium hydroxide until the vanadium pentoxide dissolves?
I'd like 1/2g vanadium pentoxide dissolved in 250mL water to be my working solution. If I only added just enough sodium hydroxide to dissolve the vanadium do you think I could store it in a glass bottle?
I'd like 1/2g vanadium pentoxide dissolved in 250mL water to be my working solution. If I only added just enough sodium hydroxide to dissolve the vanadium do you think I could store it in a glass bottle?
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Randy Holmes-Farley
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Randy Holmes-Farley
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You’re welcome. Let us know how it turns out.
salty joe
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It turned out beautiful! A nice looking yellowish with not a hint of cloudiness.
I used the tiny scoop from a Salifert test kit, I only needed 7 scoops of lye to dissolve 1/2g vanadium pentoxide in 250 mL water.
Thanks a ton!
I used the tiny scoop from a Salifert test kit, I only needed 7 scoops of lye to dissolve 1/2g vanadium pentoxide in 250 mL water.
Thanks a ton!
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salty joe
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So my vanadium mix turned clear! No precipitate or anything. It's stored in a glass bottle with a tight fitting lid and kept at room temp in a dark cabinet. Seems super strange, I assume it still has the same potency.
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You mean colorless, when it was colored before? The color may come from other ions that slowly changed oxidation state to colorless.
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Ik
Ok, probably an impurity redox reaction, but maybe someone involving the vanadium.
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I think that is likely.
