?Experiment 4: Qualitative Analysis of Cations and Precipitation Reactions Essay

Experiment 4: Qualitative Analysis of Cations and Precipitation Reactions

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The overall goal of experiment four was to determine the identity of unknown cations presented to the student. But in order to know the identity of these unknowns, in part 1, Ag+, Pb+, and Hg22+ were presented to the student in aqueous solutions and then precipitated through experimentation. In part 2, the same procedure was enacted to determine which substances precipitated through qualitative analysis. Solubility rules were also a major theme as solubility is important in determining whether a reaction will produce a precipitate. Starting out the experiment, HCl was added to the solution in the test-tube in order to form a reaction between the HCl, Ag, Pb, and Hg cations. The products of this addition of HCl were PbCl2, AgCl, and Hg2Cl2.

After the HCl was added, the solution turned a milky white color. It was important to not add too much HCl because excess HCl would have caused an aqueous complex of PbCl2 and AgCl to form instead of the desired solid PbCl2. The mixture was then centrifuged in order to let the solid particles of the three ions to fall to the bottom. Another drop of HCl was added to test if the reaction had been completed. If the solution were to turn milky white again then it would’ve signaled an incomplete reaction between the cations and HCl. The next objective was to separate the lead(II) ion from the mixture. The test tube with all three solids precipitated at the bottom was heated, which allowed the PbCl2 to dissolve. The supernatant fluid containing the Pb(II) ions was then separated from the solid mercury(I) ions and solid silver ions decanting, adding a drop of acetic acid to the supernatant fluid, and also adding two drops of K2CrO4 to the test tube. A milky yellow mixture was observed as this indicated the presence of the lead (II) ion. This insoluble yellow precipitate was the insoluble compound of PbCrO4.

In part 2 of the experiment, after the supernatant liquid was separated and the drops of acetic acid and K2CrO4 were added, the solution did not turn a milky white indicating the absence of lead (II) ions in the unknown. The next step was to display the mercury (I) ions present. NH3 was added to the test-tube with the saved white and grey precipitate. The mixture was then centrifuged and a grey precipitate was found at the bottom of the test-tube. This grey participate signalizes the presence of mercury (I) ion in the mixture. The compound that shows the presence of mercury (I) ion is HgNH2Cl. The unknown sample contained mercury (I) ions because of the observed precipitation of the grey substance to the bottom of the test tube after the addition of NH3 and centrifuging. Last but not least, the presence of silver ions had to be tested. One ml of nitric acid was added to the supernatant solution from the previous step in order to neutralize it and produce AgCl salt.

Once the nitric acid was added, the solution turned a milky white color and white precipitate was observed floating around in solution. In the unknown, silver ions were present because of the observed white precipitate formed when adding nitric acid. More than one ml had to be added in order for the precipitate to be observed, so approximately 25 drops were added as opposed to 20(20 drops equals approximately one ml). Although the experiment faced a couple minor problems, it went smoothly overall. One of these difficulties was a problem with one of the centrifuges.

I had set the timer to 3 minutes but it proceeded for a minute and then stopped. I had to in turn go and reset the timer. This technical difficulty was not a personal fault, just an equipment problem. Another problem faced was the rapid evaporation of the boiling water in the 250 ml beaker used to heat the initial test-tube. For part one of the experiment, there was sufficient water to heat the test-tube, but as part 2 came around, the water level had dropped sufficiently and the test-tube wasn’t able to be heated to its fullest potential. This could have affected the outcome of the visibility of lead(II) ions.

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