In the Laboratory
The Softening of Hard Water and Complexometric Titrations
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An Undergraduate Experiment Helena Ceretti, Enrique A. Hughes, and Anita Zalts* Instituto de Ciencias, Área Química; Universidad Nacional de General Sarmiento Roca 850, (1663) San Miguel, Prov. Bs.As., Argentina; *
[email protected] Throughout history, the quality and quantity of water available to humans have been vital factors in determining their welfare. The acceptable quality is determined by the water’s intended use; for example, water with a medium salt content may be used for irrigating crops, but not for drinking. The available water often doesn’t fulfill the minimum requirements and has to be treated in some way so as to render it acceptable. The type and degree of water treatment are strongly dependent upon the water’s source and intended use. Calcium is the cation found at the highest concentration in most freshwater systems; it occurs along with magnesium and sometimes with iron(II). These ions, which are generally present as bicarbonates or sulfates, account for water hardness and produce an insoluble “curd” by reaction with soap (1). Although ions that cause water hardness do not form insoluble products with detergents, they do adversely affect detergent performance. Another problem caused by hard water is the formation of mineral deposits. This solid product coats the surfaces of hotwater systems, clogging pipes and reducing heating efficiency. Dissolved salts such as calcium and magnesium bicarbonates and sulfates can be especially damaging in boiler feedwater (2). Clearly, the determination of water hardness and its removal (water softening) are essential for many uses of water. One widely used method for determining the calcium and magnesium ion content of water is complexometric titration with EDTA (ethylenediaminetetraacetic acid) using Eriochrome Black T or Calmagite as visual end-point indicators (3–8). Water softening may be achieved by various techniques. On a large scale, the lime–soda process is preferred. Calcium may be removed from water very efficiently by the addition of orthophosphate. Another alternative is the use of ion exchange, the reversible transfer of ions between an aqueous solution and a solid material capable of bonding ions. This last is an effective and economical technique widely used for domestic water treatment that does not require the removal of all ionic solutes, but just those cations responsible for water hardness. A number of materials have ion-exchanging properties. For this purpose, aluminum silicate minerals or zeolites can be used (8). The discovery in the mid-1930s of synthetic ionexchange resins composed of organic polymers with attached functional groups marked the beginning of modern ionexchange technology. The undergraduate laboratory experiment described here introduces students to (i) complexometric water hardness titration, (ii) characteristics of an ion exchanger, and (iii) uses of an ion exchanger for water softening. The experiment consists of the following stages Demonstration of ion-exchange properties of special resins. A mini-ion-exchange column is made with 10 mL of a slurry 1420
of a cationic resin (e.g., Dowex 50 W-X8) and a disposable syringe. A small portion of CuSO4 solution is added and eluted. The following observations can be made: the top portion of the column turns green; the eluate is colorless and tests positive for acid and sulfate ions. This should lead to the deduction that the Cu2+ ions are retained, displacing an equivalent amount of H+ ions. The anions (in this case sulfate) pass through the column unaffected. Analysis of hard water by complexometric titration with EDTA. Complex formation and equilibrium and the influence of pH and concentration can be discussed at this stage. A sample of hard water (tap water; if necessary an artificial sample can be prepared easily) is titrated with standardized EDTA solution, using Eriochrome Black T or Calmagite as an indicator. Ion-exchange softening of the water sample. The disadvantages of hard water are presented to introduce the need for softening. This can lead to the subject of water purification in general. A second portion of the water sample is passed through the mini-column; the eluate and washings are collected, tested with litmus, and titrated with standardized EDTA solution (this volume is practically zero, indicating total removal of Ca2+ and Mg2+ ions). Alternatively, the eluted protons may be titrated. Requirements The entire experiment can be carried out in one lab period of two hours. Only the usual lab equipment is needed (disposable syringes, flasks, burets, pipets, etc.). No special safety considerations are necessary (9). All reagents are either inexpensive or reusable, and all are readily available. Conclusion This experiment is easy to perform. It gives students hands-on experience in complexometric and acid–base titrations and in the use, reuse, and regeneration of ion-exchange resins. The following concepts are applied and can be discussed: quantitative analysis; titrations; indicators; equilibria: acid– base reactions, complex reactions, and ion-exchange in resins; water purity; and water hardness and softening. Note W A handout for students, notes for the instructor, and two illustrative figures are available as supplementary materials for this article on JCE Online at http://jchemed.chem.wisc.edu/Journal/issues/1999/Oct/ abs1420.html.
Literature Cited 1. Manahan, S. E. Environmental Chemistry, 6th ed.; Lewis: Boca Raton, FL, 1994; pp 61–62.
Journal of Chemical Education • Vol. 76 No. 10 October 1999 • JChemEd.chem.wisc.edu
In the Laboratory 2. Manahan, S. E. Op. cit.; pp 235–239. 3. Yappert, M. C.; DuPré, D. B. J. Chem. Educ. 1997, 74, 1422. 4. Harris, D. C. Exploring Chemical Analysis, 1st ed.; Freeman: New York, 1996; Chapter 15. 5. Harris, D. C. Quantitative Chemical Analysis, 4th ed.; Freeman: New York, 1995; pp 793–794. 6. American Public Health Association, American Water Works Association, and Water Pollution Control Federation, Joint Edito-
rial Board. Standard Methods for the Examination of Water and Wastewater, 15th ed.; American Public Health Association: Denver, 1980; Method 2340, pp 185–191. 7. Soriano, D. S.; Draeger, J. A. J. Chem. Educ. 1993, 70, 414. 8. Smoot, A. L.; Lindquist, D. A. J. Chem. Educ. 1997, 74, 569– 570. 9. Material Safety Data Sheets. http://www.phys.ksu.edu/~tipping/ msds.html (accessed Jul 1999).
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