Materials Included In Kit

Additional Materials Required

Safety Precautions

The pH 10 buffer solution is mildly toxic by ingestion, a severe body tissue irritant and produces harmful vapors; avoid breathing the fumes. Wear chemical splash goggles, a chemical-resistant apron and chemical-resistant gloves. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. Final products may be disposed of according to Flinn Disposal Method #26b.

Teacher Tips

• Typically, the titration of calcium and magnesium by EDTA is carried out in the presence of Erochrome Black-T, but endpoints tend to not be as sharp. Experience in developing this procedure has shown that calmagite gives sharp, easily identified endpoints. Calmagite solution is blue in basic solution. However, when calmagite forms a complex with calcium or magnesium, the solution is pink in color.

• In a basic solution, the acidic hydrogens, that is the hydrogens attached to oxygen, are pulled away from the EDTA molecule. That leaves each of these four oxygen atoms with three nonbonding pairs electrons and a negative charge. The two nitrogen atoms each have a nonbonding pair of electrons. The EDTA^4– (as shown in Figure 1) can bond to calcium or magnesium ions in six places: the four singly bonded oxygen atoms and the two nitrogen atoms, each donating a pair of electrons to the metal ion. The resulting Ca–EDTA and Mg–EDTA complexes both have a minus two charge.
• During the titration the following processes occur:
  1. Tap water containing Ca²⁺ and Mg²⁺ is placed in a beaker. Some pH 10 buffer is then added. The buffer maintains a basic pH so that the singly bonded oxygen atoms in EDTA are not bonded to hydrogen atoms and can bond to the metal ions.
  2. A small amount of calmagite solution is added. The calmagite immediately bonds with some of the free metal ions forming a red complex.
  3. EDTA is then added dropwise until all of the metal ions have been pulled away from the calmagite. This is seen when the calmagite indicator solution changes from red (metal ions attached to the calmagite) to blue (no metal ions attached to calmagite). This signals the end of the titration.

• Water with a range of 0 to 60 parts per million is considered soft, 60 to 120 parts per million—medium hard, 120 to 180 parts per million—hard, and 180 parts per million and above is considered very hard.
• Smaller beakers or plastic cups may also be used if you are short on glassware.
• If there is some question on the purity of your distilled water, you may want to run a trial test on your own, since significant amounts of dissolved minerals will affect student results.
• In order to set up an accurate control (beaker 4), it is imperative that students rinse both the graduated cylinder and their beakers with distilled or deionized water.
• Presence of the relatively small amounts of copper, iron, and/or aluminum ions in solution may result in interference problems. These may be masked effectively or removed from solution by the addition of stronger complexing agents. Potassium tartrate will mask aluminum, while the pH 10 buffer will precipitate iron as the hydroxide. Copper requires either fluoride ion or cyanide ion (which are very toxic and not commonly used in high school labs). If you have hard water due to the presence of bicarbonates, no masking is needed, since the solutions do not sit long enough before titration for interference to occur.
• Very soft water (<50 ppm) may require a change in the concentration of EDTA solution. Try a 0.005 M solution of EDTA if this is the case.

Further Extensions

1. If your kitchen (or other) faucet has a purification filter, try analyzing both the filtered and the unfiltered water. Discuss the results.
2. Repeat the experiment as described in the procedure, but use 2 mL of water and 1 mL of buffer, instead of 5 mL and 2 mL, respectively. Compare both the hardness values obtained and the consistency of data and discuss the results.
3. Try substituting volumes of 10 mL for the water and 4 mL of buffer in the original procedure. Compare the hardness values obtained and the consistency of data. Discuss the results.

Sample Data

Average value of calcium in Beakers 1, 2 and 3 (in ppm) ___188___

Answers to Questions

1. Convert the number of drops of EDTA used in each titration to milliliters of EDTA used.
   

   (Drops of EDTA)/(Drops of water per mL) = mL of EDTA used
   

2. Convert the milliliters of EDTA used in each titration to liters of EDTA used. Record these values in the data table.
   

   (mL of EDTA)/(1000 mL/L) = Liters of EDTA used
   

3. Calculate the number of moles of EDTA used in each titration, using the volumes calculated in the previous question and the molar concentration of the EDTA, 0.010 M. Record these values in the data table.
   

   (Liters of EDTA used) x 0.010 M EDTA = moles of EDTA used
   

4. Given that one mole of EDTA reacts with one mole of calcium, calculate the number of grams of calcium that were present in each of 5.0 mL samples. (Hint: Use the molar mass of calcium.) Record these values in the data table.
   

   (moles of EDTA used) x 40.1 g calcium/mol = grams of calcium per sample
   

5. Taking the density of water to be 1.0 g/mL (or 1000g/L), a calcium ion concentration of 1.0 g/L would correspond to 1000 ppm. Convert the results from Question 4 from grams of calcium per 5.0 mL sample to grams of calcium in 1.00 L. Then multiply this value by 1,000,000 to get parts per million of calcium. Record each individual sample values and the average value of parts per million calcium in the data table.
   

   (grams of calcium per sample)/5.00 mL = grams calcium in 1.00 mL
   (grams of calcium in 1.00 mL) x (1 x 10⁶) = parts per million calcium
   

6. Identify two major sources of experimental error in this analysis and suggest ways in which the procedure could be modified to minimize those errors.
   

   The biggest problems center around the impossibility of getting quantitative transfer of liquids from the graduated cylinder into the reaction vessel. A secondary difficulty involves the tendency of Beral pipets to give an overflow of drops. Another is finding a consistent endpoint.

References

Special thanks to John G. Little, St. Mary’s High School, Stockton, CA, for providing the idea and instructions for this activity.

Dick, J. G. Analytical Chemistry; McGraw-Hill: New York, 1973.

Lindstrom, F., and Diehl, H., Analytical Chemistry 1960, 32, 1123.

Pierce, W. C. Haenisch, E. L., and Sawyer, D. T., Quantitative Analysis; John Wiley and Sons: New York, 1958.

Scwartzenbach, G., and Flaschka, H., (tran. by Irving, H.M.N.H.), Complexometric Titrations; Methuen and Co. Ltd.; London, 1969.

Welcher, F. J., The Analytical Uses of Ethylenediamine Tetraacetic Acid; D.Van Nostrand Co.: Princeton, NJ, 1958.

Introduction

In this activity, the hardness of ordinary tap water will be measured using a titrametric method.

Concepts

• Water hardness

• Chelates

Background

If tap water is boiled for a long period of time, it begins to leave a film on the walls of its container. This film-like substance can be observed in pots or pans at home and also on beakers and flasks in the laboratory. Most of these whitish deposits are actually residues of sulfate and carbonate salts of calcium. Magnesium compounds are also present in these film deposits. When the quantity of these deposits becomes high, the water is described as being “hard.” Water becomes harder as the number of calcium and magnesium compounds increases. Ordinary soaps cannot dissolve in hard water and therefore leave a grayish film on containers and clothes. Synthetic detergents have the same functions as soap, but because they are soluble in hard water, a gray film does not result. This is the reason why synthetic detergents are used to clean clothing rather than standard soap.

Both calcium and magnesium ions have a strong attraction for a large organic molecule called ethylenediaminetetraacetic acid, commonly called EDTA (see Figure 1). The metal ions and EDTA molecules combine to form a large complex ion known as a chelate. A chelate is formed when metal ions are attached to coordinate links of two or more nonmetal atoms in the same molecule.

In this activity, EDTA is added to a water sample in known amounts until all of the magnesium and calcium in the tap water have been chelated away by the complex. An indicator is added to help clearly see the actual point where this occurs. Due to the fact that this reaction only works in a moderately basic solution, a pH 10 buffer solution is also added to the water sample. In addition, there are generally far more calcium ions than magnesium ions in water samples, so the final results will be reported in parts per million of calcium.

Materials

Safety Precautions

The pH 10 buffer solution is mildly toxic by ingestion, a severe body tissue irritant and produces harmful vapors; avoid breathing the fumes. Wear chemical splash goggles, chemical-resistant apron and chemical-resistant gloves. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Procedure

1. Label four, 50-mL beakers, 1–4. Beaker 4 will be used as the control.
2. Rinse a 10-mL graduated cylinder and Beakers 1–3 with tap water and shake to remove most of the remaining water. Do not dry the inside of the beakers with any type of towel.
3. Measure exactly 5.0 mL of tap water into the first three beakers, labeled 1–3.
4. Rinse the 10-mL graduated cylinder and Beaker 4 with distilled or deionized water and shake to remove the excess water. Add exactly 5.0 mL of distilled water from the graduated cylinder to the control beaker (4).
5. Using the 10-mL graduated cylinder and a thin stem Beral-type pipet, add 2.0 mL of the pH 10 buffer and 5 drops of calmagite indicator to each of the four beakers. Swirl gently to mix.
6. Using a microtip Beral-type pipet, add 1 drop of the 0.010 M EDTA solution to beaker 4 and swirl to mix. This should produce a clear blue color. If the blue coloration does not result, continue adding EDTA dropwise, swirling after each addition, until the clear blue color remains. This beaker is the control beaker. Set aside for future comparison.
7. Add 10 drops of EDTA to the mixture in Beaker 1 and swirl to mix. If the indicator does not change from red to blue, add 5 more drops of EDTA and swirl again. Continue using 5-drop increments, while recording the number of drops used, until a blue color results that does not change back to red. To find the correct endpoint, compare the sample to the control (beaker 4). They should be the same color. Make sure that the solution in beaker 1 is a pure blue color with no trace of purple or red. Record the final number of drops used in the data table.
8. Titrate the samples in beakers 2 and 33 following the same procedure presented in step 7. When near the expected endpoint, add only a few drops at a time and swirl. This will provide a more precise value.
9. Calibrate your micropipet to determine how many drops are in 1.0 mL. To do this, place about 2 mL of water in the 10-mL graduated cylinder. Note the actual volume to the nearest 0.01 mL. Using the same micropipet used for the EDTA titration, add, while counting, more drops until exactly 1.0 mL of water has been added. Record the number of drops needed to increase the volume by 1.0 mL in the data table. Place the drops directly into the water; do not let the drops of water hit the sides of the graduated cylinder because they may not entirely reach the water in the graduated cylinder.
10. See instructor for proper disposal procedures.

Student Worksheet PDF

10274_Student1.pdf