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Combining two ionic compounds in solution can cause precipitation of a new, insoluble, ionic compound




Combining two ionic compounds in solution can cause precipitation of a;new, insoluble, ionic compound. In this experiment you will use the;precipitation of cobalt ions to determine the concentration of a cobalt(II);chloride (CoCl2) solution. This analytic technique is called precipitate;titration.;The solution of cobalt chloride is titrated by a solution of sodium hydroxide;(NaOH). A double displacement reaction results in the formation of cobalt;hydroxide which precipitates out of the solution.;The endpoint is visualized by adding a pH indicator. As long as there are;cobalt ions in the solution, the sodium hydroxide is neutralized and the pH;remains just slightly acidic. As soon as the cobalt ions are used up, excess;sodium hydroxide makes the solution basic.;The choice of the indicator is not a trivial one, since the cobalt(II) chloride;solution has a pinkish color to begin with. Phenolphthalein is not a good;choice because it is already pink. Thymolphthalein is the only indicator in;our stock room that will do the job. This indicator is colorless in acids and;turns 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 the;change from pH of 7 in order to obtain accurate results, the extra volume;of titrate required to raise the pH to 10 is less than one drop.;The concentration of the cobalt(II) chloride solution is determined by;comparing the concentrations of the titrant and titrate. The equation for the;chemical 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 the;cobalt 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 one;molecule of CoCl2. Therefore, if we know the total moles of NaOH used in;the titration, then there will have been half as many moles of CoCl2 in the;titrated sample.;This means that we must adjust the familiar formula for a titration to;account for the ratio in which the CoCl2 and NaOH react;(Moles of CoCl2) = (Moles of NaOH) 2;which can be expressed in terms of concentrations as;C1 V1 = (C2 V2) 2;where;C1 is the concentration of the CoCl2;V1 is the volume of the CoCl2 solution being titrated;C2 is the concentration of the NaOH solution;V2 is the total volume of NaOH added up to the endpoint;Part 1: Coarse Titration;NOTE: The procedures described in this lab assume that you have already;done the Titration Tutorial and are familiar with the technique. If you have;not yet done the Titration Tutorial Lab, please do it now.;Place a 150 mL Erlenmeyer Flask from the Containers shelf onto the;workbench.;Add 10 mL of Cobalt Chloride Solution (CoCl2) from the Materials shelf to;the Erlenmeyer Flask.;Dilute the solution by adding 10 mL of water. This dilution makes it easier;to visualize the end point, but remember that the concentration of the;Cobalt 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 the;mouse cursor over the Burette's glass tube to display the volume of NaOH;solution and record it in your Lab Notes.;Move the Erlenmeyer flask onto the lower half of the burette to connect;them.;Perform a coarse titration, adding large increments of the NaOH solution;from the burette by pressing and holding the black knob at the bottom of;the burette. Each time you add the solution, check the volume remaining in;the burette. As the Cobalt Chloride (CoCl2) in the Erlenmeyer Flask is;used up in the reaction with NaOH, the pink color will disappear. At the;endpoint of the titration, the solution suddenly turns blue.;Record the last burette volume at which the solution in the flask was still;pink as well as the volume at which the solution turned blue. They give you;the range for your fine titration.;1;Clear your station by dragging your containers to the recycling bin beneath;the workbench.;Part 2: Fine Titration;Prepare the tiration as before by repeating steps 1 through 7 in Part 1.;2;3;Quickly add enough NaOH to just get into the range of the coarse titration;but 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 volumes;you recorded in Part 1.;4;5;Add NaOH one drop at a time. When a drop causes the solution in the;Erlenmeyer flask to turn blue. Record the start and end volumes of the;NaOH solution in the burette in your lab notes: the volume of the burette;when the reaction occurred, and the volume just before.;6;7;Repeat the fine titration two more times for accuracy, and record the;results in your lab notes.;8;Clear your station by dragging your containers to the recycling bin beneath;the workbench. (Remember to press Save Notes so you dont lose your;calculations.);PrecipitationTitrationofCobalt;Chloride;Experiment2;1. Record your results for each of the 3 trials;a;Volume of CoCl2 (mL);b Volume of NaOH added (mL);c Concentration of NaOH added;2. Calculate the concentration of the CoCl2 solution (moles/L) using the;formula developed in the Background.;3. What is the average concentration of the CoCl2 solution? To how many;significant digits can this concentration be reported? (Consider the accuracy;of the burette and the volume in one drop.)


Paper#15539 | Written in 18-Jul-2015

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