Solved by a verified expert:Combining two ionic compounds in solution can cause precipitation of anew, insoluble, ionic compound. In this experiment you will use theprecipitation of cobalt ions to determine the concentration of a cobalt(II)chloride (CoCl2) solution. This analytic technique is called precipitatetitration.The solution of cobalt chloride is titrated by a solution of sodium hydroxide(NaOH). A double displacement reaction results in the formation of cobalthydroxide which precipitates out of the solution.The endpoint is visualized by adding a pH indicator. As long as there arecobalt ions in the solution, the sodium hydroxide is neutralized and the pHremains just slightly acidic. As soon as the cobalt ions are used up, excesssodium hydroxide makes the solution basic.The choice of the indicator is not a trivial one, since the cobalt(II) chloridesolution has a pinkish color to begin with. Phenolphthalein is not a goodchoice because it is already pink. Thymolphthalein is the only indicator inour stock room that will do the job. This indicator is colorless in acids andturns blue in the base pH range of 9.4 to 10.6.While you might think that we need to visualize the endpoint right at thechange from pH of 7 in order to obtain accurate results, the extra volumeof titrate required to raise the pH to 10 is less than one drop.The concentration of the cobalt(II) chloride solution is determined bycomparing the concentrations of the titrant and titrate. The equation for thechemical reaction between cobalt chloride and sodium hydroxide is:CoCl2 (aq) + 2NaOH (aq) → 2NaCl (aq) + Co(OH)2 (s)At the endpoint, exactly enough NaOH has been added to remove thecobalt ions from the solution in the form of the solid precipitate, Co(OH) 2.Notice, however, that 2 molecules of NaOH are required to react with onemolecule of CoCl2. Therefore, if we know the total moles of NaOH used inthe titration, then there will have been half as many moles of CoCl2 in thetitrated sample.This means that we must adjust the familiar formula for a titration toaccount for the ratio in which the CoCl2 and NaOH react:(Moles of CoCl2 ) = (Moles of NaOH) ÷ 2which can be expressed in terms of concentrations as:C1 × V1 = (C2 × V2) ÷ 2where:•••C1 is the concentration of the CoCl2V1 is the volume of the CoCl2 solution being titratedC2 is the concentration of the NaOH solutionV2 is the total volume of NaOH added up to the endpoint •Part 1: Coarse TitrationNOTE: The procedures described in this lab assume that you have alreadydone the Titration Tutorial and are familiar with the technique. If you havenot yet done the Titration Tutorial Lab, please do it now.Place a 150 mL Erlenmeyer Flask from the Containers shelf onto theworkbench.••Add 10 mL of Cobalt Chloride Solution (CoCl2) from the Materials shelf tothe Erlenmeyer Flask.••Dilute the solution by adding 10 mL of water. This dilution makes it easierto visualize the end point, but remember that the concentration of theCobalt Chloride Solution (CoCl2) relates the moles of Cobalt Chloride(CoCl2) to the original 10 mL.••Add 2 drops of Thymolphthalein to the Erlenmeyer flask .••Place a burette from the Containers shelf and place it on the workbench.••Fill the burette with 50 mL of 0.1 M Sodium Hydroxide (NaOH). Move themouse cursor over the Burette’s glass tube to display the volume of NaOHsolution and record it in your Lab Notes.••Move the Erlenmeyer flask onto the lower half of the burette to connectthem.••Perform a coarse titration, adding large increments of the NaOH solutionfrom the burette by pressing and holding the black knob at the bottom ofthe burette. Each time you add the solution, check the volume remaining inthe burette. As the Cobalt Chloride (CoCl2) in the Erlenmeyer Flask isused up in the reaction with NaOH, the pink color will disappear. At theendpoint of the titration, the solution suddenly turns blue.••Record the last burette volume at which the solution in the flask was stillpink as well as the volume at which the solution turned blue. They give youthe range for your fine titration.••1Clear your station by dragging your containers to the recycling bin beneaththe workbench.Part 2: Fine TitrationPrepare the tiration as before by repeating steps 1 through 7 in Part 1. 23Quickly add enough NaOH to just get into the range of the coarse titrationbut still have the solution in the Erlenmeyer flask appear pink. This is near,but not yet at, the titration’s end point and is the bigger of the two volumesyou recorded in Part 1.45Add NaOH one drop at a time. When a drop causes the solution in theErlenmeyer flask to turn blue. Record the start and end volumes of theNaOH solution in the burette in your lab notes: the volume of the burettewhen the reaction occurred, and the volume just before.67Repeat the fine titration two more times for accuracy, and record theresults in your lab notes.8Clear your station by dragging your containers to the recycling bin beneaththe workbench. (Remember to press Save Notes so you don’t lose yourcalculations.)Precipitation Titration of CobaltChlorideExperiment 21. Record your results for each of the 3 trials:aVolume of CoCl2 (mL)b Volume of NaOH added (mL)c Concentration of NaOH added2. Calculate the concentration of the CoCl2 solution (moles/L) using theformula developed in the Background. 3. What is the average concentration of the CoCl2 solution? To how manysignificant digits can this concentration be reported? (Consider the accuracyof the burette and the volume in one drop.)
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