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b.) In the systematic analysis of cations II. and III. Although group cations precipitate as sulfides, precipitation reagents and environments are different? Explain why.
c.) III. Why do some precipitate as sulphides and others as hydroxides in a qualitative systematic analysis of group cations?
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- 1. How is complete precipitation assured in the Group III Cations? 2. In the confirmatory test of Magnesium ion, why is it necessary to heat the centrifugate in water bath for 5 minutes? 3. If Magnesium hydroxide is a white gelatinous precipitate, why is the confirmatory result indicated a blue precipitate?Inorganic Synthesis: Production of Alum from Coke Cans If acid is added slowly, a white precipitate appears and subsequently disappears. What is the white precipitate? Support your answer with appropriate chemical equations.You are assigned an unknown solution that contains Group III cations. To -1 mL of this solution was added 6 M NH3 the solution was agitated to mix well, and a reddish-brown precipitate with a gelatinous solid clinging to the inner walls of the test tube was observed. The solution was centrifuged and the supernatant was tested for completeness of precipitation by adding an additional drop of 6 M NH3. No cloudiness was observed as the drop of reagent diffused through the solution. The supernatant was then carefully decanted into a clean test tube, labeled (1st solution), and saved for further testing later. The precipitate remaining in the test tube was washed with a small amount of water, centrifuged, and the wash decanted and discarded. To the precipitate was added about 10 drops of 6 M NAOH plus ~1 mL H20 and the test tube was vigorously agitated. The resulting suspension was centrifuged and the supernatant liquid was transferred to another clean test tube and clearly labeled (2nd…
- 1. How quantity Qsp and Ksp can be used in predicting precipitate formation? 2. How precipitation titrations work? 3. How will you differentiate Mohr method, Volhard method and Fajans method?Sometimes it is not possible to indicate the end point of a titration.a) How can one proceed then and what is the name of the type of titration that can be performed? Briefly describe. An example in which this method can be used is in the determination of mercury, which forms strong complexes with EDTA, but for which there is no suitable indicator that can indicate the end point. b) You are given the task of determining the Hg2 + concentration in a sample solution? After adding an excess of EDTA, the sample solution is titrated with a magnesium solution. 20.00 ml of a 0.0452 M EDTA solution was added to 30.00 ml of sample solutionThe excess EDTA was determined by adding 0.0500 M Mg 2+ solution, consuming 4.37 ml to the end point.In a separate experiment, the lab partners I. Goofed and T. Klutz treat their unknown with HCl and obtain a precipitate. They treat their precipitate with hot water and decant the supernatant. On treatment of this solid with ammonia, it dissolves. On careful acidification with HNO3 a white precipitate forms. They report that their unknown contains silver. They receive only partial credit. Why?
- Why homogenously formed precipitates are better suited for analysis than a precipitate formed by direct addition of a precipitating agent?Calculate the equilibrium concentration of N2O4 in 1.00 L of a solution prepared from 0.129 mol of N2O4 with chloroform as the solvent. (Assume that the change in concentration of N2O4 is small enough to be neglected).Suppose that 0.323 g of an unknown sulfate salt is dissolved in 50 mL of water. The solution is acidified with 6 M HCl, heated, and an excess of aqueous BaCl2 is slowly added to the mixture resulting in the formation of a white precipitate. 1) Assuming that 0.433 g of precipitate is recovered calculate the percent by mass of SO42− in the unknown salt. 2) If it is assumed that the salt is an alkali sulfate determine the identity of the alkali cation.
- Why are many ionic precipitates washed with electrolyte solution instead of pure water?The NaOH titrant in this experiment was prepared to be approximately 0.1 M and then was standardized to determine its exact concentration. What possible reasons could there be for not simply weighing the solid NaOH, dissolving to a known volume and calculating its molarity?The solubility of Sr(OH)2 is 3.2 x 104 g/L a) Write a balanced equation for the solubility equilibrium b) Write the expression for the solubility product constant Ksp and calculate its value c) Calculate the pH of a saturated solution of Sr(OH) 2 d) 25 ml of 3 x 10-4 SOLUTION of SrS04 is added to 35 ml of 3.21 x 10-6 molar NaH solution. Does a precipitate form? Explain and show calculations to support your answer