Statement of Goals
The goal of the group project was to calculate the percent calcium in a natural calcium tablet from Prime Eastern Pharmaceuticals in order to compare it with the reported amount of 800 mg of calcium per tablet to determine the validity of complaints regarding a reduced amount of mineral in the multivitamin. The quantitative determination of the analyte was accomplished through three experiments: complexometric titrimetry, flame atomic absorption analysis, and potentiometry using an Ion Selective Electrode. The complexometric titrimetry involved titrating the unknown calcium solution made from the calcium tablet with standardized EDTA. Quantitative determination of the analyte calcium was possible due to the stable and quick formation of the metal-EDTA complex with a 1:1 reaction stoichiometry. Flame atomic absorption analysis is useful in determining the amount of analyte in an unknown because the measure absorbances of standard solutions can be used to plot a calibration curve that can be used to determine the concentration of the metal in an unknown solution. Potentiometry is useful for determination of analyte when an ion sensitive electrode is used to find the amount of ion in standard solutions.

Week 1: Introduction and Background
The first test for calcium in the multivitamin was complexometric titrimetry, a titration of a solution of calcium tablet with a standardized solution of EDTA. EDTA forms a stable complex with calcium metal with a reaction stoichiometry of 1:1, so an EDTA titration can be used to quantitatively determine the metal ion concentration of calcium in a calcium tablet solution. The formation of calcium-EDTA complex is pH-dependent and needs to be buffered to pH so that formation can be quantitatively determined. A Calmagite indicator is used to determine the endpoint of the titration. A colored complex is formed when the indicator reacts with the excess calcium ion and then undergoes a color change when reacted with EDTA.
The reaction of Ca2+ with EDTA is given by: Ca2+ +H2Y2- ↔ CaY2- + 2H+
Calcium Determination with EDTA Procedure The standardized solution of EDTA was prepared by weighing a 0.9995 gram sample of EDTA salt by difference into a 250 mL volumetric flask. The solution was dissolved with about 200 mL of warm deionized water, left to cool to room temperature and then diluted to the mark with deionized water when all the EDTA salt had dissolved completely. The calcium solution was prepared by weighing 5 Natural Calcium tablets using a triple beam balance and then grinding the tablets into a fine powder using a mortar and pestle. A 0.1569 gram sample of the calcium powder was weighed into a 150 mL beaker. The solid was dissolved by adding 6 M HCl drop wise to the solid in the beaker and 50 mL of deionized water was added to the solid. The solution was brought to a boil to expel excess carbon dioxide. The solution was allowed to cool to room temperature and then neutralized to a pH of 6 by adding 1 M ammonium hydroxide drop wise. The solutions was gravimetrically filtered into a 250 mL volumetric flask and diluted to the mark with deionized water and mixed thoroughly. A 50.00 mL volumetric pipette was used to transfer 50.00 mL of the Natural Calcium solution into three Erlenmeyer flasks. A fourth Erlenmeyer flask was filled with 100 mL of deionized water and labeled as a blank. 15 mL of a buffer solution of pH 10, 12 drops of Calmagite, and 8 drops of methyl red indicator was added to all the solutions. Each of the solutions was titrated with the prepared EDTA solution from a wine-red color to a blue color at the end point.

Week 2: Introduction and Background Flame atomic absorption spectrometry is a technique used to measure the concentration of a specific metal present in a sample. The basic principle behind this technique is metal atoms emit specific wavelengths that can be detected allowing the determination of the concentration of metal present. Our unknown was analyzed using flame atomic absorption for its calcium concentration. To analyze a sample through flame atomic absorption, a liquid sample that contains the unknown metal in its ground state, is exposed to a lamp that contains a calcium cathode, which causes the calcium atoms to absorb energy as they move to a higher state, and then release energy at a specific wavelength as they drop back to the ground state. Released waves are captured by a detector, which relates to the amount of calcium in the unknown sample. Standards are used to construct a calibration curve. This plot allows us to determine the concentration of calcium in the unknown dietary tablet, after we measure its absorbance using flame atomic absorption spectrometry. With the completion of this study we will be able to determine the calcium concentration of a calcium substance and compare it to the label. As well as find the best way to measure the calcium content.
Flame Atomic Absorption Analysis Procedure
Equipment Used: Digital scale 25 mL volumetric flask 250 mL beaker 100 mL volumetric flask Pipet Filter paper Mortar & Pestle
Concentration of Calcium in a Natural Calcium tablet was calculated by using flame atomic absorption. The absorbances of four known calcium concentration standards were experimentally determined along with the absorbance of the unknown calcium solution. A calibration curve was generated and a line of best fit was determined from the four standard solutions. Based on the line of best fit, the concentration of the unknown calcium solution can also be calculated. To make the standard solutions, 0.1 grams of crushed calcium tablet were dissolved in 6 M HCl and 50ml of water. It was then boiled to remove any carbon dioxide formed and filtered into a 100ml volumetric flask, diluted to the mark with DI water and mixed. 0.1g of calcium carbonate was then weighed into a 100ml volumetric flask, diluted to the mark and mixed thoroughly. Standard solutions were prepared by pipetting 1 mL, 2 mL, 3 mL, and 4 mL of the 0.2M calcium carbonate solution into 25ml volumetric flasks and diluting to the mark with DI water. The calcium tablet solution was diluted twice by pipetting 5 mL of the solution into a 25 mL flask. Two of the diluted calcium tablet solutions were made. The solutions were then handed to the teaching assistant to measure their absorbances with an atomic absorption spectrophotometer.
References
M.W. Rowe, M. Hyman, A.E. Miller, A.C. Javier, E. Soriaga. Quantitative Analysis Laboratory Manual; Department of Chemistry, Texas A&M University: College Station.

Week 3: Introduction Potentiometric determination with an ion-selective electrode is a very useful method for determining the concentration of a certain element in a mixture or solution. In this experiment, a Calcium ion-selective electrode is used to detect the amount of Calcium in different standard solutions as well as an unknown solution. To prepare the unknown solution, the calcium tablet being investigated was ground to a fine powder, dissolved and neutralized before being analyzed with the electrode.

Potentiometric Determination Potentiometric determination uses an ion-selective electrode that detects the potential of the given solution and relates it to the concentration of the ion using the Nernst equation:
E=E^'+2.303 RT/nF log⁡(M) The determination works best for very dilute solutions and this is why many dilutions are performed for the standard and the unknown solutions.

Potentiometric Determination Procedure Weigh 1.001g of Calcium Carbonate into a 100-mL beaker. Dissolve with 6 M HCl until solution stops fizzing. Add 50 mL of DI water and boil off excess CO2, and dilute to the mark in 100- mL volumetric flask. Mix thoroughly. Use a 10.0-mL volumetric pipette to transfer 10.00 mL of the 0.1M Calcium Carbonate Solution. Dilute to the mark with DI water and mix well. Prepare four Calcium Carbonate calibration standards in 100-mL volumetric flasks according to the following table.

Solution Volume (mL) of 0.01 M Calcium Carbonate Volume (mL) of 1.0 M KCl
1 1.00 10.00
2 2.00 10.00
3 5.00 10.00
4 10.00 10.00

Dilute each solution to the mark with deionized water and mix thoroughly. Crush Calcium tablet to powder using mortar and pestle Weigh 0.30 g of Calcium tablet powder into a 100-mL beaker. Dissolve6 M HCl until solution stops fizzing. Add 50 mL of DI water and boil off excess CO2, neutralize, filter into volumetric flask and dilute to the mark. Mix thoroughly. Use a 2.0-mL volumetric pipette to transfer 2.0mL of unknown calcium tablet solution into a 100.0-mL volumetric flask. Add 10.0mL of KCl solution. Then dilute to the mark and mix vigorously. Transfer 60.0 mL of calibration standards and unknown solution to five 100.0 mL beakers. Use Ion-Selective Meter to measure potential in millivolts of each of the five solutions. (See Lab manual for instructions on use) Dispose of waste chemicals in proper containers.

Equipment: Glassware: 5 100-mL beaker 7 100-mL volumetric flasks 1 2.0-mL volumetric pipette 1 10.0-mL volumetric pipette 1 500.0-mL waste beaker 1 50.0-mL beaker Chemicals: 2 Calcium tablets 2g of Calcium carbonate 1.0 M KCl 6 M HCl DI Water Materials and Instruments: Kimwipes Ion-selective meter with calcium selective electrodes Mortar and pestle PC with Microsoft EXCEL

Safety Considerations: Handle chemicals with care, avoid skin and eye contact. Do not ingest chemicals.

Results and Discussion
Titration by EDTA The following table shows the volume EDTA used to fully titrate the unknown solution to the equivalence point and the respective grams of calcium and % calcium calculated from the results.
Trial 1 2 3 4
Volume(mL) 24.7 24.8 24.7 24.99
Mass Ca (g) 0.010633 0.010676 0.010633 0.0107576
% Ca 33.88375 34.02093 33.88375 34.281576 The mass of calcium in each sample was calculated as follows volume of EDTA (L)× molarity of EDTA (moles/L) × molar mass of calcium (g/mole)=grams calcium per sample Sample calculation
24.7 mL EDTA × (1 L)/(1000 mL) ×0.01074 (moles EDTA)/(L EDTA) × (1 mole Ca)/(1 mole EDTA) × (40.08 g Ca)/(mole Ca) =0.010633 g Ca per sample The percent calcium was then calculated by dividing the g Ca per sample by the total weight of the sample. An average of all four trials was found to be 34.075% Ca which was used for later calculations. This was used to find the average mass of Ca per tablet by converting the percentage back into a ratio and multiplying it by the average mass of a tablet (1.08052g) and multiplying by a thousand to make it mg which yielded a mass of 614.0227 mg Ca per tablet. The calculation for this mass is shown below
(34.075% Ca per sample)/(100%) ×1.08052 g tablet × (1000 mg)/g=614.0227 mg Ca per tablet

The Grubb’s Test was also performed for the %Ca found from each trial and it was found that all trials were to be accounted for because no calculated G values were higher than the G critical value. The calculations are summarized in the table below
G calc 0.713149533 0.018286 0.71315 1.408013
G crit 1.463 1.463 1.463 1.463
Gcalc>Gcrit? no no no no The standard deviation for the trials was 0.18755 which shows that the spread was very small so the test can be considered fairly accurate. However, since this analysis involves an eye test to discern when the sample is at the equivalence point since a color change must be observed there is always a slight amount of error simply due to the nature of the experiment.
Flame Atomic Absorption The following graph shows the calibration curve from the standards that was used to analyze the unknown samples. The equation A=0.0021C+0.0046 was found A being Absorbance and C concentration in parts per million. The two found absorbances for each unknown sample was fit to find the resulting concentration and %Ca per sample. The following table summarizes the findings.
Solution 1 2 concentration (ppm) 18.93333333 17.99111
Absorbance 0.0321 0.0326 concentration curve (ppm) 13.0952381 13.33333
Concentration (M/L) 0.000326727 0.000333 moles Ca 0.000816819 0.000832 g Ca per sample 0.03 0.03
% Ca 30.74 32.94 mg Ca per tablet 554.86 594.54

The concentration was converted from parts per million to molarity from the following sample calculation
13.095 (mg Ca)/(kg water) × (1000 kg water)/(1 m^3 water) × (1 m^3 water)/(1000 L water) × (1 g Ca)/(1000 kg Ca) × (1 mole Ca)/(40.08 g Ca)=0.000326 M Ca This was then converted to moles of Ca per sample by multiplying the concentration by the dilution factor. This was then converted to grams of Ca per sample by multiplying by the molar mass of Calcium and each percentage was found by dividing by the original amount of unknown sample used per trial and multiplying by 100%. The same steps were taken to find the mg Ca per tablet as the last experiment. Average mg Ca per tablet was found to be 574.7 mg Ca per sample with a standard deviation of 28.055. A much larger standard deviation is expected since there were only two samples and they varied greatly. Since they were far from each other it is safe to assume that there was error from some source in this analysis. The likely culprit is from the dilution step of the samples. Since they were diluted multiple times and it impossible to perfectly mix a sample an uneven amount of sample was likely in each sample. Since the concentrations were so small a very small deviation at that scale also leads to a much bigger difference when scaled up to the mg Ca per tablet. So, a small amount of error in the dilution could lead to the observed difference in percentage.
Potentiometric Determination of Calcium The following table and graph summarizes the potentials of the different trials and unknown found and the resulting calibration curve that the unknown was compared to. Measured Potentials (mV)
CaCO3 standards Trial 1 Trial 2 Average
1 -17.6 -18.9 -18.25
2 -9.6 -11.6 -10.6
3 1.1 0.4 0.75
4 8.8 8.6 8.7 unknown sample 3.2 2.2 2.7

The equation P=27.165C+89.945 was determined to fit the data P being the measured potential in mV and C being the log of concentration in molarity. Using this equation the molarity was found for the average unknown potential and the resulting percent calcium and mg of calcium per tablet were calculated summarized in the following tables.
Line of best fit:
Solution Sample potential (mV) log conc. concentration unknown 2.7 3.21167 0.000614229

Final Calculations
0.003071147 moles Ca
0.123091564 g Ca
33.63157484 % Ca
607.0566523 mg Ca

The moles of calcium were calculated by the following sample calculation
0.00061 (moles Ca)/(L water) ×(100 mL)/(2 mL) ×0.1 L=0.003 moles Ca
The concentration was essentially multiplied by the dilution factor and then multiplied by the original amount of water it was dissolved in. The remaining calculations were performed exactly as in the previous experiments. Like the previous experiment the dilution was the likely culprit for any error but since only one sample was made this error wouldn’t show up. The measuring potentiometer could also possibly have error from any contaminates from previous trials.
Conclusions
This table summarizes the average mg Ca found for each experiment
Experiment titration by EDTA atomic absorption Potentiometric mg Ca 614.03 574.70 607.06
Average 598.60
Standard deviation 20.9864