How to test for and remove lead(II) sulfate from a sample of 1 M sulfuric acid?

If you had a clear, colourless sample of dilute sulfuric acid and suspected lead contamination, probably lead sulfate, how would you test for this? It's solubility is low in sulfuric acid solutions, and lower near freezing, so I thought cooling might produce a precipitate, but I'm not sure what level if any this would be able to detect. Raising the concentration by open boiling or distilling would, from what I can gather, produce a filterable precipitate, but that is not something I want to test.

Would there be any easy way to remove any detected contamination? By easy, I mainly mean not distilling or boiling and filtering conc. sulfuric. Ideally, I would be looking for a method that does not involve raising the concentration at all. Chemical conversion to something even less soluble? Electrolysis?

EDTA can be used to combat lead intoxication. Would there be a way of using it to remove lead compounds?

• PbSO4 is slightly soluble in concentrated sulfuric acid. But by diluting it to 30% H2SO4, PbSO4 is not soluble any more. Its solubility is 43 mg per liter. Is this concentration acceptable for you ? – Maurice Apr 11 '20 at 15:40
• Thanks for your answer. I based my comment about solubility on an old paper, D. Craig & G. Vinal 1939, which shows solubility of about 6.5mg/l in 1M sulfuric and actually decreasing at higher and slightly lower concentrations. I don't really know what a reasonable limit is though. I gather it's a serious accumulative toxin, so I'm really looking to get it as low as possible with my limited apparatus. – eisd Apr 11 '20 at 17:02
• Quick and dirty method: Take a few mL of your sulfuric acid and neutralize it with ordinary household ammonia solution. Use litmus paper to get approximate neutrality: red means acidic and blue means basic. Then test the neutralized solution with lead testing strips. Hardware store often have them. If anything precipitates during the neutralization step, then that, in itself, indicates that the 1 M sulfuric acid was contaminated with whatever. – Ed V Apr 11 '20 at 19:04
• What I suggested has nothing to do with possible purification of the acid, assuming it needed it. If it did have lead it it, I would take it to a recycling facility that can handle used lead-acid battery acid and then buy new battery acid. – Ed V Apr 11 '20 at 19:17
• That sounds inexpensive and worth a try. Thanks. There seem to be quite sensitive test kits available for applications worth the cost. I'm not going to be drinking this, but it has a workspace exposure limit, according to UK EH40 WEL, of 15ppb 8hrs or 45ppb 15min, so the potential based just on solubility that a clear solution still has 43mg/l or even 6.5mg/l seems way too much. – eisd Apr 12 '20 at 12:55

To answer your question on how to test for Lead sulfate, I refer you to this paper, to quote:

In 1929 Fischer [1] introduced a sensitive reagent for small quantities of lead which has found interesting applications in widely diverse fields. Diphenylthiocarbazone, or dithizone as it is commonly called...

And further:

We have made no study of this reagent, except to find means of using it in our problem of measuring the solubility of lead sulfate. Dissolved in chloroform, dithizone imparts a noticeable green tint to the solution, provided the solution is not too concentrated. The lead-dithizone complex, however, has a bright cherry-red color.

Also of import;

Dithizone is known to react with metals other than lead. The precautions taken to avoid interference by other metals are discussed on page 58. These precautions, together with the spectroscopic analysis of the lead salt and the redistillation of water and acid, make it highly improbable that the results were affected by other metals. Confirmation of this is found in the analysis of solutions of known lead content.

If Pb is present, dilution appears to be the best path due to solubility.

[EDIT] Here is an theoretical path to the removal of heavy metals that does not require dilution. I am suggesting UV light therapy employing TiO2 photocatalyst to act on cold H2SO4 containing heavy metal impurities. Heavy metals are expected to precipitate.

With respect to the mechanics of the photolysis, I referred, for example, to this source. Per cited Equation [R1]:

$$\ce{ hv -> e- + h+ }$$

Per Equation [R3], expected reaction with an electron hole:

$$\ce{H2O + h+ -> .OH + H+ }$$

$$\ce{H2SO4 = H+ + HSO4- }$$

$$\ce{H+ + e- = .H }$$

where the hydrogen radical behaves as a e-/H+ pair, see Hydrometallurgy 2008: Proceedings of the Sixth International Symposium, p. 818, a commercial reductive leaching equation, to quote:

$$\ce{ PbS + 2 •H = Pb + H2S }$$ (5)

In the current case:

$$\ce{ .H + PbSO4 -> Pb (s) + HSO4- }$$

However, the associated transient sulfate radical is also created:

$$\ce{ .OH + HSO4- -> H2O + .SO4- }$$

$$\ce{ .SO4- + .SO4- -> S2O8(2-) }$$

So, the sulfate ions is not alone. Possibly more advantageous, further pump air/oxygen into the solution during the UV light treatment. Then per Equation [R3], expect the superoxide radical anion formation:

$$\ce{O2 + e- = .O2- }$$

which becomes, in acidic conditions, the $$\ce{.HO2}$$ radical, which may similarly react with Lead sulfate as follows:

$$\ce{ .HO2 + PbSO4 -> PbO2 (s) + HSO4- }$$

or, just oxidize any liberated Pb.

I would even suggest you try this process out by creating a test solution (with at least added Lead) and heat in H2SO4 to form a heavy metal impurity and then apply the UV light/TiO2 (and also added oxygen) therapy.

• Seems to be marginally more obtainable than Potassium Rhodizonate and reasonably safe to store and use, so this is definitely a helpful answer to the testing part of the question. If I can test before and after, that gives me a good basis for trying out purification methods, but I still need some ideas for that. – eisd Apr 11 '20 at 20:33