Mg2+(aq) → Mg2+(aq)+ calmagite → Mg-calmagite + Mg2+(excess) → Mg-calmagite + Mg2+(excess) + EDTA4- → Mg-EDTA2- + calmagite
Pre-Lab Question:
1. A buffered solution containing lead ions, Pb2+, was titrated to the end point using 37.53mL of 0.10M EDTA. What is the mass of lead ions in solution?
Moles EDTA: (0.03753L)x(0.10M)=0.003753 mol Pb2+
Mass Pb2+: (0.003753 mol)x(207.2 g/mol)=0.778g
II. Procedure:
Only plastic laboratory equipment was used for this experiment as opposed to glass to avoid a chemical reaction destroying the glass. The student obtained a bottle of ~0.01M EDTA solution, a bottle of indicator, and a sample of the magnesium salt. Sample number of the magnesium salt was recorded. A 250mL Erlenmeyer flask with DI water thoroughly. The plastic 100mL buret was then rinsed with ~ 10-20mL EDTA. To ensure the entire inside of the buret was rinsed, it was inverted and twisted while holding the top closed. Washings were drained into the flask, which was subsequently rinsed with the EDTA washings. Excess EDTA was disposed of under the hood in a safe container. Concentration of the EDTA was provided by the laboratory professor as 0.01127M. Three ~0.15g samples of the magnesium salt were weighed and the mass was recorded. In the Erlenmeyer flask, the sample was dissolved in 20.0mL DI water, 10.0mL buffer, and two drops of indicator. The previously rinsed 100mL plastic buret was filled with standardized EDTA and clamped securely to the buret stand above the flask. Initial volume of EDTA in the buret was recorded. EDTA was added to the flask until the color of the solution changed from a red-pink to a very light blue, indicating the endpoint. The number of moles of Mg2+ in the magnesium sample was calculated. This procedure was repeated until three trials yielded accurate data. Mass of magnesium in each trial was calculated and recorded. Data as percent MgO for each trial was calculated and recorded. Calculations are below.
The buret was drained until empty, and returned to its proper location in the laboratory accordingly. Left over solution was discarded under the hood and the lab station was left clean after the laboratory professor signed the data.
III. Tabulated Experimental Data:
Concentration EDTA: 0.01127M
Molar ratio of magnesium ions (Mg2+) to EDTA: 1:1
Formula weight of MgO: 30.31
Average mass of Mg2+ in the magnesium sample:
Mg2+ as % MgO: 14.14%
IV. Sample Calculations:
For the calculation of Mg2+, since the molar ratio was 1:1, the following equations were used:
In both experiments, the average of the data from three trials was calculated.
The calculations of percent composition MgO from the best three trials is as follows:
Trial 1
Trial 2
Trial 3
V. Tabulated Results of Calculations:
VI. Discussion of Results:
The results of the experiment are precise and relatively accurate based on the actual amount of Mg2+ ions in the students’ individual sample, as provided by the laboratory instructor. Values are logical relative to the given information; the percent of Mg2+ ions should be small because a metal cation is smaller than the nonmetal anion since it has lost two electrons. The exact values were disclosed by the laboratory professor for purposes of calculating percent error and interpreting results.
VI. Conclusion:
Data calculated from the data of the three best trials yielded an average %Mg2+ of 14.14%. Given that the actual percentage in sample 54 was 14.7%, the average %Mg2+ yielded a percent error of 3.81%. Possible sources of error in this experiment were inexact measurements of any substance used in the experiment, magnesium sample sticking to the weigh boat therefore transferring a smaller amount than what was recorded, or over/under titrating. Data from the trials that the student chose not to use because of the large percent error they produced were accurate but not precise. Overall, the student accomplished the intent of the experiment. The reaction, summarized by:
M2+ + EDTA4- → M-EDTA2- in which the student observed the clear Mg2+ ion solution with calmagite indicator which was wine red at a pH of 9.3 turning to a light blue when the chelate was produced, product: Mg-EDTA2- + calmagite. The light blue color indicated that the last Mg2+ ion complexed with EDTA. Magnesium has an optimum pH range of 9-10, so the choice of indicator was logical and effective. The data from all three trials allowed the student to use stoichiometry to determine the moles of EDTA in the solution. As indicated in the purpose, the molar ratio of Mg2+ to EDTA is 1:1, so the moles of EDTA also indicated the moles of Mg2+.
The student became familiar with the process of chelating through active participation, performing the necessary steps for a reaction to take place in which a ligand (EDTA) was bonded to a metal ion producing the chelate. This type of titration is known as complexometric titration, because it involves the formation of a metal complex that can be quantitatively analyzed. In further experiments, it is recommended that the student conduct more trials for assurance that the calculated result will be as close as possible to the actual value.