If very small concentrations of NaCl and AgNO3 are mixed, no precipitate forms. However, if large concentrations are mixed, a white precipitate forms.
Can someone tell me a method for testing the hypothesis that "insoluble" salts would dissolve to some extent in water and are slightly soluble. If the Ksp is low,the solid is not soluble and if high, it is soluble and if it the product of the ion concentration does the exceed the Ksp value, no precipate would form.So that's the hypothesis.So, can someone give me a experiment to test this hypothesis.
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.
To test the hypothesis that "insoluble" salts may dissolve to some extent in water and are slightly soluble, you can perform an experiment to determine the solubility of a salt and compare it to its Ksp value. Here is a step-by-step guide:
Materials needed:
1. NaCl (sodium chloride)
2. AgNO3 (silver nitrate)
3. Distilled water
4. Measuring spoons
5. Test tubes or small glass containers
6. Stirring rod or glass rod
7. Balance
8. Pipette or dropper
9. pH indicator (optional)
Procedure:
1. Begin by measuring out small quantities of NaCl and AgNO3. Start with concentrations that are known to be soluble and gradually increase the concentration until a precipitate forms. For example, you can start with 1 gram of NaCl and 1 gram of AgNO3 separately.
2. Label two test tubes or glass containers as Control and Experimental.
3. Add a small amount (around 10 mL) of distilled water to both the Control and Experimental containers. Dissolve the NaCl in the Control container and the AgNO3 in the Experimental container. Stir the solutions with a glass rod until the salts are fully dissolved.
4. Observe the solutions. If both solutions appear clear, it indicates that the salts are soluble in water at these concentrations. If a white precipitate forms in the Experimental container but not in the Control container, it suggests that the AgNO3 has a low solubility compared to NaCl.
5. If the solutions are clear and no precipitate forms, you can repeat steps 3 and 4, gradually increasing the concentration of the salts until a precipitate forms in the Experimental container.
6. Once a precipitate forms, record the concentrations of the NaCl and AgNO3 solutions at which the precipitate appears. These concentrations are the approximate solubility of each salt.
7. Calculate the solubility product constant (Ksp) for both NaCl and AgNO3 using the concentrations obtained in step 6. The Ksp is calculated by multiplying the concentrations of the ions in the solution. For example, for NaCl (NaCl -> Na+ + Cl-), if the concentration of Na+ is x and the concentration of Cl- is y, then Ksp = x*y.
8. Compare the calculated Ksp values with the literature values for NaCl and AgNO3. If the calculated Ksp values for both salts are lower than the literature values, it indicates that they are slightly soluble.
Note: It is always a good practice to repeat the experiment multiple times to ensure accuracy and consistency of the results. Also, using a pH indicator can help determine if any other reactions are occurring during the mixing process.
Remember, this experiment is designed to test the hypothesis that insoluble salts are slightly soluble and the Ksp value can indicate their solubility.