Experiment 5. Volumetric Determination of Calcium in TUMS® via EDTA... Techniques Weighing TCD-4,

Chem 2LA3
Complexometric Titration of Ca
Experiment 5. Volumetric Determination of Calcium in TUMS® via EDTA Complexation
Techniques
Weighing
Heating
Volumetric Analysis
Complexometric Titrations
TCD-4, 5
TCD-9
TCD-39
TCD-45
Scenario (continued from EXP-3). Further investigations have revealed that more than two dozen
batches of TUMS® products were mislabelled during manufacturing at GlaxoSmithKline (GSK) Inc.
whose dosage remains unknown. We have been hired by GSK to develop an alternative volumetric
method for unequivocal identification of the calcium dosage of large numbers of TUMS® tablets that is
accurate and inexpensive, yet rapid relative to classical gravimetric analyses. Volumetric titrations are
a widely used analytical method that uses a selective chemical reaction or process for the quantification
of a specific target analyte, based on the volume of titrant (i.e., calibrated reagent) required to complete
a reaction. In this context, you recall that ethylenediaminetetraacidic acid (EDTA) is a multidentate
chelating agent that forms stable complexes with various metal ions in solution, such as Ca2+. Your
objective is to quantify the amount of Ca2+ in the unknown TUMS® sample by using a volumetric
titration method based on a competitive metal-ligand complexation reaction, as well as determine the
weight-percent (wt %) of CaCO3 it contains. You will also compare its performance relative to your
previous gravimetric analyses in order to assess whether it offers any distinct advantages for calcium
determination in terms of analysis time, precision or accuracy.
OH
Na+
SO3-
Eriochrome Black T
O
EDTA
pK1 = 0.0 (CO2H)
pK4 = 2.69 (CO2H)
pK2 = 1.5 (CO2H)
pK5 = 6.13 (NH+)
+
HO
NH
O
HO
NO2
O
NH+
OH
N
pK3 = 2.0 (CO2H) pK6 = 10.37 (NH+)
In3-
HO
at pH < 10, EDTA exists as Y
4-
N
OH
O
Kf1
1. Ca
2+
3-
(aq) + In (aq)
CaIn- (aq)
Kf2
2. Ca
2+
4-
(aq) + Y (aq)
CaY2- (aq)
Since Kf1 << Kf2,the addition of EDTA titrant to
a sample containing CaIn- results in:
3. CaIn- (aq) + Y4- (aq)
Red
CaY2- (aq) + In3- (aq)
Blue
Suggested Procedure. The complexation agent EDTA is a polyprotic aminocarboxylic acid which can
form a chelate with up to six sites (i.e., hexadentate) with Ca2+, usually in a 1:1 stoichiometry as shown
in the chemical equilibria above. Prior to using EDTA as a titrant to determine the calcium content of
your antacid, you must first standardize the titrant against a standard solution made from pure calcium
carbonate with a known concentration. Since both free Ca2+ and the Ca2+-EDTA complex (CaY2-) are
colourless, you will need to use Eriochrome Black T (In3-) as a colorimetric metal indicator to determine
the endpoint of your titration. Under the alkaline conditions, the indicator will change from red to blue
when the endpoint is reached. Note that all the free Ca2+ is initially complexed with Eriochrome black T
prior to the start of the titration thus forming CaIn-, hence the colour of the solution appears red. As you
approach the equivalence point with the addition of EDTA to the sample solution, calcium will be
stripped from CaIn- by EDTA via competitive displacement, thereby generating the free uncomplexed
indicator (In3-, blue colour) which signals the titration endpoint. Note that the thermodynamic stability of
the CaIn- complex is much weaker than the hexadentate EDTA complex CaY2- (i.e., Kf1 << Kf2), which is
a key feature to this direct complexometric titration method.
AN-3
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Chem 2LA3
Complexometric Titration of Ca
Pre-Lab
In the notebook:
1. Write out the chemical equilibria involved in this complexometric titration method for calcium
determination, as well as note the colour changes expected during the titration.
2. Prepare data tables along the lines of Table 1 and 2 below.
3. What is the purpose of adding Eriochrome Black T to the sample solution prior to starting the
titration? Sketch the chemical structures of the two complexation agents used in the titration.
4. What is the concentration of CaCO3 you are making in Part (a), step 1? Given that you will be
provided with ≈ 0.01 M EDTA, how much titrant do you expect to use to titrate a 25 mL aliquot?
5. Explain why the stability of most EDTA-metal complexes is lower under acidic conditions. Why
is a pH 10 buffer solution used for the experiment?
6. If your gravimetric analysis was accurate, you “already know” the wt% CaCO3 in your TUMS
sample. Use this value to calculate an estimated CaCO3 concentration in the solution you are to
make up in Part (b) below, and the anticipated volume of 0.01 M EDTA titrant you will require to
reach the equivalence point?
Directions
You will be provided with stock solutions of EDTA and HCl, as well as NH4Cl buffer at pH 10. Examples
of colour changes expected for the titration involving Eriochrome black T will also be shown.
(a) Standardization
1. Weigh out ~0.30 g of pure CaCO3. Record the exact mass and transfer it quantitatively to a 250
mL beaker.
2. Dissolve the CaCO3 by adding ~25 mL distilled water and slowly add 0.5 mL drops of 12 M HCl.
Swirl to dissolve. Add more HCl if needed.
3. Quantitatively transfer the solution into a clean 250 mL volumetric flask and dilute to the mark
with distilled water. Calculate the moles of CaCO3 and molarity of your standard solution.
4. Rinse well a clean buret with EDTA stock solution at least three times. Fill with EDTA solution
and check for air bubbles. Record initial volume on buret.
5. Transfer 25.00 mL of the CaCO3 standardization solution to a clean 250 mL Erlenmeyer flask
using a volumetric pipette. Add 5 mL of the pH 10 buffer and 4-5 drops of Eriochrome black T
indicator. Record observations on colour of solution. Place solutions on hot plate to warm (do
not boil).
6. Perform a quick trial titration and add EDTA titrant to the sample while swirling the flask to mix
until the endpoint is reached, as indicated by a colour transition from red/purple Æ blue. Record
the final volume and net volume of titrant used.
7. Once you have an approximate titrant volume required to reach the endpoint, repeat the titration
again more carefully/efficiently such that you have three consistent measurements. Please note
that as you reach the endpoint, the reaction can be slow, so add titrant dropwise and mix well
for ≈ 3-5 s prior to adding further titrant until the colour changes from red Æ purple Æ clear blue.
The first consistent blue colour change represents your titration endpoint. Keep the solution
warm near the endpoint by periodically warming solution on the hot plate. The addition of more
Eriochrome black T may be needed near the endpoint.
8. Calculate the average molarity of the EDTA solution from your replicate measurements to the
appropriate number of significant figures Æ report the average and standard deviation.
Table 1. Titrant Standardization Data Table
Replicate 1
Mass of CaCO3 (g)
Replicate 2
Replicate 3
Amount of Ca2+ in solution (moles)
Initial buret reading (mL)
AN-3
Page 2 of 4
Chem 2LA3
Complexometric Titration of Ca
Final buret reading (mL)
Net volume of EDTA used (mL)
[EDTA] (M)
(b) Complexometric Titration of the TUMS®
1. Weigh out ~0.35 g of dried crushed TUMS® stored in a dessicator. Record the exact mass and
transfer it quantitatively to a 250 mL beaker.
2. Dissolve the antacid by adding ~ 25 mL distilled water and 0.5 mL of 12 M HCl. Swirl to
dissolve. Add more HCl if needed. Quantitatively transfer into a clean 250 mL volumetric flask
and dilute to the mark with distilled water.
3. Transfer 25 mL of the dilute sample solution to a clean 250 mL Erlenmeyer flask using a
volumetric pipette. Add 5 mL of the pH 10 buffer and 4-5 drops of Eriochrome black T indicator.
Place solutions on hot plate to warm (do not boil).
4. Perform a quick trial titration with EDTA until the endpoint is reached. Record the final volume
and net volume of titrant used. Use this volume to plan out your subsequent titrations.
5. Repeat and carefully complete three replicate titrations that provide you consistent results. Keep
the solution warm near the endpoint by periodically warming solution on the hot plate. The
addition of more Eriochrome black T may be needed near the endpoint.
6. Calculate the mass of Ca2+ and percentage of CaCO3 in each sample and report the mean and
error at the 95% confidence level.
Table 2. Antacid Sample Data Table
Replicate 1
Replicate 2
Replicate 3
Mass of TUMS® (g)
Initial buret reading (mL)
Final buret reading (mL)
Net volume EDTA used (mL)
Amount of EDTA used (mol)
Amount of Ca2+ in sample (mol)
Mass of Ca2+ in sample (mg)
% CaCO3 in sample
Mass of Ca2+ per tablet* (mg)
*
Average dried weight of a single TUMS tablet: 1.291 g
Waste Disposal
Ensure to all aqueous solutions are disposed in the aqueous base waste container and solids in the
Solid Waste Container in the fumehood.
Discussion/Conclusions
Report the average mass of Ca2+ per tablet and average %CaCO3 content in the mislabelled TUMS®
sample, including the error at the 95% confidence level. Which one of the following three possible
dosages of TUMS® is the unlabelled unknown sample you analyzed? Calculate the % error (accuracy)
of your result – comment? Is the volumetric titration based on EDTA complexation faster, more precise
and more accurate relative to your previous gravimetric analysis? Provide evidence to justify your
conclusion!
Regular Strength: 200 mg Ca2+
Extra-strength: 300 mg of Ca2+
AN-3
Ultra-strength: 400 mg of Ca2+
Page 3 of 4
Chem 2LA3
Complexometric Titration of Ca
Reference
1. D. C. Harris. Exploring Chemical Analysis, 4th Edition: Ch. 6, 13, W.H. Freeman and Company:
New York, NY, 2009.
2. http://www.tums.ca/faq_contents.aspx
EXPERIMENT 5. REPORT SUMMARY: COMPLEXOMETRIC TITRATION OF CALCIUM
Date:
Student Name:
Summary Data Table
Replicate 1
Replicate 2
Replicate 3
Mean + 95% C.L.
Standard [EDTA] (M)
Mass Ca2+ in sample (mg)
% CaCO3 in sample
Mass of Ca2+ per tablet* (mg)
*
Average dried weight of a single TUMS tablet: 1.291 g
Dosage of unlabelled TUMS®?
% Error:
Are your gravimetric and volumetric results statistically equivalent at 95% confidence level in
terms of reported mass of Ca2+ per tablet?
AN-3
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