26 Jun The endpoint is visualized by adding a pH indicato
Question
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:
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•
•
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.
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•
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 .
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Place a burette from the Containers shelf and place it on the workbench.
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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.
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•
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.
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•
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 don’t lose your
calculations.)
Precipitation Titration of Cobalt
Chloride
Experiment 2
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.)
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