Gravimetric Determination of Calcium as Calcium ... The concentration of calcium in a sample can be determined by gravimetric analysis. ... this answe...
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CHM 212 Experiment 5: Gravimetric Determination of Calcium as Calcium Oxalate A two week lab Purpose: To determine the amount of calcium in a sample by precipitating the ion as calcium oxalate. Background The concentration of calcium in a sample can be determined by gravimetric analysis. In this experiment an unknown Ca2+-containing sample will be analyzed by precipitating the Ca2+ using oxalate (C2O42-). In the presence of basic oxalate solution, Ca2+ forms an insoluble precipitate (Ksp = 1.7 x 10-9): Ca2+(aq) + C2O42-(aq) CaC2O4.H2O(s) The resulting precipitate is, however, soluble in the presence of acidic solution because the oxalate anion is a weak base. Large relatively pure crystals that are easily filtered will be obtained if the precipitation is carried out slowly. This can be done by dissolving Ca2+ and C2O42- in acidic solution and gradually raising the pH by thermal decomposition of urea: O H2N
C
NH2
+
3H2O
heat
CO2
+
+
2NH4
-
+ 2 OH
Prelab Question 1. Why shouldn’t you handle your sintered glass funnels with your hands (other than not wanting to burn yourself)? Procedure: (1) Dry three medium-porosity sintered-glass funnels for 1 – 2 h at 1100C. Cool them in a desiccator for 30 mins and weigh them. Repeat this procedure with 30-min heating periods until successive weighings agree to within 0.3 mg. Use paper, paper towel or tongs, not your fingers, to handle the funnels. NOTE: if you placed your funnels in the oven during the last lab period you will probably only need to reweigh once to achieve agreement. (2) Transfer exactly 25 mL of unknown to each of three 250 – 400-mL beakers and dilute each with ~ 75 mL of 0.1 M HCl. (3) To each of the above solutions add 5 drops of methyl red indicator solution. This indicator is red below pH 4.8 and yellow above pH 6.0.
(4) Add ~25 mL of ammonium oxalate solution to each beaker while stirring with a glass rod. Remove the glass rod and rinse it into the beaker. Add ~15 g of solid urea to each sample, cover with a watchglass, and boil gently for ~30 mins until the indicator turns yellow. (5) Filter each hot solution through a weighed funnel, using suction generated by the aspirator. Add ~3mL of ice-cold water to the beaker, and use a rubber policeman to help transfer the remaining solid to the funnel. Repeat this procedure with small portions of ice-cold water until all of the precipitate has been transferred. Finally, use two 10-mL portions of ice-cold water to rinse each beaker, and pour the washings over the precipitate. (6) Dry the precipitate, first with aspirator suction for 1 min, then in an oven at 1100C for 1 – 2 h. Bring each filter to constant mass. The product is somewhat hygroscopic, so only one filter at a time should be removed from the dessicator, and weights should be done rapidly. (7) Calculate the average molarity of Ca2+ in your unknown solution.
Report the standard
deviation. NOTE: You may not be able to complete this experiment in one lab period. Consult with your TA for an appropriate point to stop and appropriate sample storage if you are unable to complete the experiment. Postlab Questions (1)
You are performing what is ideally a very accurate determination of Ca2+ by weighing the amount of a precipitate. This determination assumes that your precipitate is pure: what are known problems that can affect the purity of a precipitate?
(2)
How would you determine whether or not your precipitate was pure?
(3)
What would you estimate the lowest possible amount of Ca2+ that you could accurately detect would be? You need to think about your entire experimental procedure here, including any approximations that you have made in the method. There will be multiple factors affecting this answer.
(4)
How would you accurately experimentally determine the lowest possible amount of Ca2+ that you could accurately detect?
