Posted by Sarah on Sunday, February 12, 2012 at 1:06pm.
What about using an electrochemical set up to show that AgCl is slightly soluble. You would measure the (Ag^+) in the cell.
One difficulty you will have at very low amounts of NaCl and AgNO3, the precipate is so scarce you cannot see it. So my question is, in your test for a precipate, what is your test?
thanks for your suggestion. but maybe is it possible to use the titration method to determine whether or not the precipate would for form, for example, adding the large concentration of AgNO3 and NaCl which would cause the precipitate to form and then use those values obtained from the titration to work out the Ksp value.However, if the ionic concentration is equal to or more than the Ksp then the precipitate would form. And, for the small concentration, i would add the small volumes of NaCl and AgNO3, doing the same method and calculating the Ksp, where if the ionic concentration is less than the Ksp, then no precipitate would form.Would that method be alright or not effective in finding out if precipate would form or not?
And the method you mentioned about using electrochemical set-up to show that AgCl is slightly soluble,can you explain in detail what the experiment would consist of? Thanks!!
With regard to the electrochemical set up, I was proposing that you prepare AgCl in sufficient quantity, add it to a fresh sample of water, let it form a saturated solution, filter, then measure the (Ag^+) in the filtrate. That would show that you have Ag^+ in solution. That would show that "insoluble" salts do, in fact, dissolve to a small extent. That would also get around Bob Pursley's legitimate concern that very small amounts of AgCl ppts are not visible to the naked eye. The electrochemical method is one way in which Ksp values are determined in the first place.
With regard to your proposal, I don't think it gets around the concern Bob Pursley has about "seeing" the very small amounts of AgCl formed.
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