Water Hardness Determination Using Complexometric EDTA Titration

Abstract

This laboratory experiment aimed to determine the water hardness of a mineral water sample using complexometric EDTA titration. The hardness of water is essential to understand its suitability for various applications. Complexometric titration with EDTA, a chelating agent, was employed to analyze a mineral water sample for calcium ions. The results revealed that the sample water had a significant concentration of calcium and magnesium ions, indicating it to be hard water.

Introduction

Water hardness is a measure of the amount of calcium and magnesium present in sample water.

These calcium and magnesium ions have the capacity to replace sodium or potassium ions and form sparingly soluble products or precipitates. Water hardness is involved in various aspects of industrial and biochemical processes. Large amounts of ppm CaCO3 in water can form precipitates when interacted with soap and form rings known as “scum” in several utensils and appliances. The formation of these “scum” in electrical appliances degrades its efficiency and will eventually reduce its life span.

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In addition, these can cause impairments on fabric as well, and damage water treatment plants and piping systems at a water hardness of 300 ppm CaCO3.

Calcium is necessary for aquatic animals such as fish. It serves an important role in bone formation, blood clotting, and metabolic processes of the fish and prevents the loss of important salts in the body which helps in the functioning of its vital organs such as the heart. Small amounts of calcium in water can be life-threatening to aquatic organisms like the fish.

Thus, determination of water hardness is important. One method of determining water hardness is through complexometric titration. In this process, a ligand is involved in the said titration.

Metal ion reacts with a particular ligand forming a complex and the equivalence point is determined by an indicator. The ligand used in the experiment is Ethylenediaminetetraacetic Acid (EDTA) with Eriochrome Black T indicator. EDTA is an efficient chelating agent and has an ability to bind with metal ions. Because of this, EDTA is also used in food preservation, an anti-coagulant in blood, and, when EDTA is combined with Fe(II), can even be used as an effective absorbent of harmful NO (nitric oxide). The purpose of this experiment is to determine the hardness of water through complexometric titration.

Methodology

Before commencing the experiment, we prepared the following solutions:

• 500 mL of 0.1000 M stock EDTA solution
• 250 mL of 0.0100 M standard EDTA solution
• 250 mL of 0.050 M standard CaCO3 solution
• 50 mL of 0.0050 M working standard CaCO3 solution
• 250 mL of 1.0 M NH3-NH4+ buffer solution (pH 10)

Ethylenediaminetetraacetic Acid (EDTA) was chosen as the titrant due to its ability to form stable chelates with metal ions, making it suitable for complexometric titration.

For the preparation of the 0.1000 M stock EDTA solution, we followed these steps:

1. Weighed 18.6 g of Na2H2EDTA2H2O and transferred it into a 400 mL beaker.
2. Added 200 mL of distilled water and 1.0 g of MgCl26H2O crystals into the beaker and mixed until the crystals dissolved.
3. Heated the solution for faster dissolution and added NaOH pellets to increase solubility.
4. Transferred the solution into a 500 mL volumetric flask and diluted it to the mark with distilled water.
5. Stored the solution in a dry and clean reagent bottle.

The 0.0100 M standard EDTA solution was prepared by diluting 25 mL of the 0.1000 M stock EDTA solution to the mark in a 250 mL volumetric flask. The 0.050 M standard CaCO3 solution was prepared by dissolving 1.2510 g of pure CaCO3 in a 250 mL beaker, followed by the addition of distilled water and drops of 6 M HCl until complete dissolution. After evaporation and cooling, the solution was transferred to a 250 mL volumetric flask and diluted to the mark.

The 0.0050 M working standard CaCO3 solution was prepared by diluting 5 mL of the 0.050 M standard CaCO3 solution in a 50 mL volumetric flask. For the NH3-NH4+ buffer solution of pH 10, 2.06 g of NH4Cl was dissolved in 14.3 mL of concentrated ammonia and diluted to the mark in a 250 mL volumetric flask.

Buffer solutions were used to maintain a constant pH during titration, preventing interference from other species. The specific pH of 10 was chosen to ensure optimal conditions for the experiment.

For the standardization of the 0.010 M EDTA Solution, we prepared three 250 mL Erlenmeyer flasks, each containing 10 mL of the 0.0050 M working standard CaCO3 solution. To each flask, we added 75 mL of distilled water, 3 mL of the NH3-NH4+ buffer solution, and 2-3 drops of Eriochrome Black T (EBT) indicator. EBT indicator was chosen for its suitability in magnesium titration.

The water sample from commercial mineral water Viva was analyzed by measuring 50 mL into each of three 250 mL Erlenmeyer flasks. We added 75 mL of distilled water, 3 mL of the NH3-NH4+ buffer solution, and 2-3 drops of EBT indicator to each flask. The solutions were titrated one at a time with the 0.010 M standard EDTA solution.

Results and Discussion

Complexometric titration with EDTA was employed in this experiment to determine water hardness. The reaction between the analyte (CaCO3 solution, water sample) and the titrant (EDTA) formed complexes involving calcium and magnesium ions, which allowed us to assess water hardness. EDTA, being a weak acid, typically forms a 1:1 stoichiometric ratio with metal ions, resulting in a single endpoint and accurate stoichiometric calculations.

Several factors can influence the titration with EDTA, including the presence of complex-forming ions, organic solvents, metal ion components, and pH. In this experiment, we maintained a constant pH of 10 using the NH3-NH4+ buffer solution to ensure optimal conditions. This pH range was essential for achieving accurate results and a clear color change with the EBT indicator.

The specific pH of 10 was chosen because it allowed EDTA to deprotonate just enough to bind with the metal ions effectively. Excessive buffer addition could lead to defective endpoints, and a pH of 12, for example, would result in different outcomes due to the formation of precipitates with magnesium and calcium.

In the analysis of Viva mineral water, it was found to contain 54 mg/L of calcium and 14 mg/L of magnesium. After calculating the total hardness of the sample in terms of ppm CaCO3, the experimental value was determined to be 139.5 ppm CaCO3, while the claimed total hardness on the label was 192.6 ppm CaCO3. This indicates that the calculated value from the experiment is lower than the labeled value but still falls within the category of hard water according to the water hardness scale.

Conclusion

The Complex solutions were formed by titration with the chelating agent EDTA. With the use of complexometric titration the total hardness of water sample was determined. It was found out that the water hardness of Viva mineral water is classified as “hard” in terms of calcium and magnesium ions content that was expressed in terms of ppm CaCO3. The claimed total hardness of Viva Company is larger than the experimental value meaning it has less metal ion content than expected. The results of the experiment can be improved with the addition of KCN. It might not be visible that the endpoint was violet but it would be safer to eliminate iron discrepancies in the results.

References

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