In the Laboratory
Quantitative Analysis of Sulfate in Water by Indirect EDTA Titration
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Deirdre Belle-Oudry Department of Chemistry, University of Arizona, Tucson, AZ 85721;
[email protected]
Measurement of common anions in water samples is typically carried out using the standard EPA method of ion chromatography. However, for many teaching laboratories that are not fortunate enough to be equipped with such instrumentation, alternative wet chemical methods for quantifying anion content are required. In addition to providing a low-cost alternative technique for sulfate analysis, the experiment described herein is an excellent tool to illustrate the notion that in some instances simple wet chemical analyses are equally if not more valuable than their instrumental counterparts. The titration of metal ions (e.g., Mg2+ and Ca2+) with chelating agents such as EDTA is an important concept included in many general chemistry laboratory curricula. Anionic species can also be indirectly measured by EDTA titration. This type of titration involves precipitating the anion of interest with an excess of a metal cation forming an insoluble salt, then backtitrating the excess metal ions with EDTA. In this specific example, excess BaCl2 is added to a water sample to precipitate the sulfate ion as BaSO4. The excess Ba2+ is then back-titrated with a standardized EDTA solution (1–3). The hardness of the water sample must be accounted for as well. Students in our program measure water hardness in their general chemistry lab. The inclusion of this more advanced EDTA experiment in our upper-level analytical lab course is an example of how one can integrate lab activities and concepts at different levels in the curriculum. This twist on the standard direct EDTA titration is most suitable for a second- or third-year analytical chemistry laboratory course. It can be used to introduce or to reinforce concepts such as chelation, standardization, and water hardness, as well as to teach more advanced concepts such as indirect and back-titrations. Several different types of samples may be analyzed using this method, including groundwater, soils (4), seawater (5, 6), drinking water, and beverages (7). We opted to use “mock” drinking water containing well-known amounts of sulfate, calcium, and magnesium ions, since (i) a strong emphasis in our analytical chemistry lab course is placed on the accuracy of the students’ results, so it is important to know exactly what amount is in the sample to assess their accuracy, (ii) we sought an experiment of environmental relevance that would appeal to students in a nonmajors course as well as to chemistry majors, and (iii) it builds on students’ previous exposure to EDTA titrations in general chemistry. The concentrations of sulfate and metals in the mock water samples are within the range of 1–2 mM (~100–200 ppm SO42−). This is consistent with typical sulfate levels in drinking water. While sulfate is not considered a primary drinking water contaminant, it may affect the taste and odor of water. The National Secondary Drinking Water Regulations therefore recommend that the maximum level of sulfate in drinking water be limited to < 250 ppm (8). Experimental Overview The experiment consists of four parts and is carried out during two 3-hour lab periods. In Part 1, the mock water sample (prepared in advance by the teaching staff ) is pipetted
volumetrically into glass bottles. A few drops of concentrated HCl are added to each bottle to create an acidic environment conducive to the precipitation of BaSO4. An excess of 0.0100 M BaCl2 solution is then added to each bottle. The bottles are covered with aluminum foil, boiled, and stirred for about an hour while the student completes Part 2 of the experiment. The bottles are removed from the hot plate and, once cool, capped and set aside for two days until the next lab meeting. According to Sijderius (1), the precipitation requires 24 hours to come to completion. Therefore the experiment works best when carried out over two lab periods. In Part 2 of the experiment the student prepares and standardizes a ~0.005 M EDTA solution in ultrapure water. A standardized 0.0500 M ZnSO4 solution is first diluted 1:10; this serves as their primary standard. Each student prepares an indicator solution by dissolving a few mg of Eriochrome Black T (EBT) in ~20 mL of ultrapure water. EBT was chosen as the indicator for this experiment since it gives a clear end point when used with the metal ions of interest (Ba2+, Ca2+, and Mg2+) (9). The pH of the solution is first adjusted to a suitable range (~9.5–10) by adding pH 10 buffer solution. This step is crucial, as the EBT indicator transition is best observed in this range (9). A few drops of indicator are added to the EDTA just before titration with the standard ZnSO4 solution. The end point is observed when the solution turns from pink to pure blue. Note that this is a gradual change, and the solution first turns purple before turning blue. The hardness of the water sample is measured in Part 3. Salts containing Mg2+ and Ca2+ were added to mimic a real water sample and demonstrate that these interfering ions must be accounted for in back-titration of the Ba2+. A portion of the water sample is added to an Erlenmeyer flask and the pH is adjusted with pH 10 buffer solution. The solution is then titrated with the standardized EDTA solution. The solution turns from blue to fuchsia at the end point. Finally, in Part 4, the excess, unprecipitated Ba2+ is found by titrating the samples from Part 1 with the standardized EDTA solution. The procedure is nearly identical to that in Part 3, with a bit more pH 10 buffer added to counteract the effect of the acid added in Part 1. Note that it is not necessary to filter the BaSO4 precipitate before this final titration. By measuring the total amount of metal in the solution of the precipitated samples and subtracting out the amount of metal due to water hardness measured in Part 3, the amount of excess Ba2+ is found. The concentration of SO42− in the sample can easily be calculated:
SO 4 2
total amount added Ba2 VSO
4
amount excess Ba2
2
Hazards Barium chloride is toxic and can irritate the skin, eyes, and respiratory system (10). Having a solution prepared in advance
© Division of Chemical Education • www.JCE.DivCHED.org • Vol. 85 No. 9 September 2008 • Journal of Chemical Education
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In the Laboratory Table 1. Typical Student Results from EDTA Experiment Student
Hardness/(mmol L–1)
Sulfate Concentration/ mM
Measured
True
Measured
True
A
2.06 ± 0.02
2.00
1.19 ± 0.02
1.20
B
2.05 ± 0.01
2.00
1.37 ± 0.01
1.40
C
2.06 ± 0.01
2.00
1.23 ± 0.01
1.20
D
2.13 ± 0.03
2.00
1.23 ± 0.03
1.20
E
1.99 ± 0.02
2.00
1.48 ± 0.02
1.40
Note: Uncertainties were calculated at the 95% confidence level.
helps to minimize inhalation risk. However, students should avoid getting this solution on their skin. Concentrated hydrochloric acid and the pH 10 ammonium buffer are corrosive and may cause burns. Zinc sulfate is environmentally harmful. All chemicals from this experiment should be considered harmful and disposed of as hazardous waste. Results and Discussion Typical results from this experiment are fairly accurate. For a recent class of 24 students, the average percent error was about 2.56% for sulfate and 3.43% for hardness. The method also proved fairly precise, with typical percent relative standard deviations of less than 1.5%. Table 1 shows representative results obtained by five students in the class. For each student, the results shown represent an average of 3–5 trials. Comparisons with other methods for sulfate determination, such as ion chromatography and gravimetric analysis of BaSO4 precipitate, are useful for illustrating the advantages and disadvantages of each method. Gravimetric analysis (EPA Method 375.3) is comparable in accuracy (1.9% error) but is typically not as precise (4.7% RSD) as this EDTA method (11). Recent measurements in our laboratory of these same water samples using a Dionex DX-300 ion chromatograph (IC) showed IC to be more accurate (< 1% error). However, the IC analysis was much more time consuming, requiring 7 hours of lab time to obtain enough data for a statistically meaningful result. Additionally, neither gravimetry nor IC provides the water hardness information that this EDTA titration gives. Overall, the experiment described herein achieves the objective of providing a fast, precise, and fairly accurate method that teaches students a number of important analytical concepts. This experiment can be implemented in any quantitative analysis course or even in an honors general chemistry course. It is best done over the course of two lab periods, since the BaSO4 precipitation takes about 24 hours to complete (1). Students can usually complete the first three parts of the experiment within one three-hour lab period and finish the final titration of the excess Ba2+ during the second period. To shorten the time required, one could prepare a standardized EDTA solution in advance. Another experimental variation would be to substitute the measurement of the water hardness with removal of the interfering cations by passing the sample through a cat-
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ion exchange resin (12) before precipitation. One could also complex the metal ions with masking agents before titration with EDTA. Real drinking water samples could be substituted for the lab-prepared sample for courses in which accuracy is not as crucial. Conclusions The indirect titration of sulfate is a valuable exercise for teaching a number of important concepts to students in the analytical chemistry lab, from the utility of wet chemical techniques to water quality. It requires inexpensive and common chemicals and equipment. The procedure can be varied in several ways to fit the goals of a particular course and student skill levels. Literature Cited 1. Sijderius, R. Analytica Chimica Acta 1954, 11, 28–32. 2. Munger, J. R.; Nippler, R. W.; Ingols, R. S. Anal. Chem. 1950, 22, 1455. 3. Kowalczyk, G. S.; Simpson, C. P. An Indirect Determination of Sulfate by Back-Titration of Barium with EDTA. In Abstracts, 37th Middle Atlantic Regional Meeting of the American Chemical Society; American Chemical Society: Washington, DC, 2005. 4. Purokoski, P.; Lakanen, E. Acta Agr. Scand. 1959, 9, 355–360. 5. Howarth, R. W. Limnol. Oceanogr. 1978, 23, 1066–1069. 6. Page, J. O.; Spurlock, W. W. Analytica Chimica Acta 1965, 32 (6), 593–595. 7. Krasnova, N. S.; Gamarnik, B. A. Zhurnal Analiticheskoi Khimii 1989, 44 (5), 948–950. 8. U.S. EPA Drinking Water Contaminants. http://www.epa.gov/ safewater/mcl.html#inorganic (accessed May 2008). 9. Ueno, K. J. Chem. Educ. 1965, 42, 432–433. 10. Sigma-Aldrich Home Page. http://www.sigmaaldrich.com (accessed May 2008). 11. Sulfate by Gravimetric Determination (Method 375.3). In Methods for the Analysis of Water and Wastes (MCAWW), EPA/600/4-79/020; U. S. E. P. A. Office of Research and Development: Washington, DC, 1983. 12. MacKellar, W. J.; Wiederanders, R. S.; Tallman, D. E. Anal. Chem. 1978, 50 (1), 160–163.
Supporting JCE Online Material
http://www.jce.divched.org/Journal/Issues/2008/Sep/abs1269.html Abstract and keywords Full text (PDF) Links to cited URLs and JCE articles Supplement Student handouts
Instructor notes including answers to the questions in the student handout
JCE Featured Molecules for September 2008 (see p 1296 for details) Structures of EDTA and Eriochrome Black T discussed in this article are available in fully manipulable Jmol format in the JCE Digital Library at http://www.JCE.DivCHED.org/JCEWWW/ Features/MonthlyMolecules/2008/Sep/.
Journal of Chemical Education • Vol. 85 No. 9 September 2008 • www.JCE.DivCHED.org • © Division of Chemical Education
