Quantitative Determination of Total Hardness In Drinking Water Essay

Custom Student Mr. Teacher ENG 1001-04 12 March 2016

Quantitative Determination of Total Hardness In Drinking Water

Abstract

This experiment is about the determination of water hardness through the use of complexometric EDTA titration. Determination of water hardness is important to find out the most suitable water hardness under particular circumstances. This was conducted for the purpose of applying the concept of complexometric titration using an efficient chelating agent, EDTA. Sample mineral water was analyzed using standard EDTA with EBT as indicator, and calcium ions present in the solution were calculated to determine the hardness of the water sample. At the end of the experiment, the results indicated that the mineral sample water has large amounts of calcium and magnesium ions—an implication of a hard water sample.

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. 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 the actual experiment, solutions of 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, and 250 mL of 1.0 M NH3-NH4+ buffer solution were prepared quantitatively. In this experiment, the titrant used was Ethylenediaminetetraacetic Acid (EDTA), a polydentate with six bonding sites. Polydentates aid in obtaining sharper endpoints since they react more completely with cations. Likewise, reaction with polydentates only involves a single step process compared to using monodentates as titrants which involves at least two intermediate species. Among polydentates, EDTA was chosen as the titrant since it is versatile and forms most sufficiently stable chelates because of its several complexing sites which gives rise to a cage-like structure isolating the cations from solvent molecules.

For the preparation of 500 mL of 0.1000 M stock EDTA solution, 18.6 g of Na2H2EDTA2H2O was weighed to the nearest 0.1 mg and was transferred into a 400 mL beaker. 200 mL of distilled water and 1.0 g MgCl26H2O crystals were added into the beaker and mixed until the crystals were dissolved. MgCl26H2O was added to obtain a sharper endpoint since CaIn- complex ion is less stable and endpoint will come earlier than actual. The solution was heated for faster dissolution and NaOH pellets were added to the turbid solution to produce salt EDTA making the pH of the solution higher and increasing the solubility of the EDTA. Into a 500 mL volumetric flask, the solution was transferred and was diluted to mark with distilled water. The solution was stored in a dry and clean reagent bottle. The 250 mL of 0.0100 M standard EDTA solution was prepared by getting 25 mL from 0.1000 M stock EDTA solution and diluting it to mark with distilled water in a 250 mL volumetric flask. For the preparation of 250 mL of 0.050 M standard CaCO3 solution, 1.2510 g of pure CaCO3 was weighed to the nearest 0.1mg into a 250 mL beaker and 20 mL distilled water was added.

Drops of 6 M HCl were added until the CaCO3 was completely dissolved. The beaker was covered using a watch glass and was put over a hot plate. The solution was evaporated until an amount of 10 mL was left. After cooling the solution, the washings were collected by rinsing the watch glass into the beaker using distilled water. 20 mL more distilled water was added into the solution and it was transferred into a 250-mL volumetric flask. The solution was diluted to mark and was stored in a plastic polyethylene bottle since glass bottle can leach and ions from it will contaminate the solution. The 50 mL of 0.0050 M working standard CaCO3 solution was prepared by dilution of 5 mL 0.050 M standard CaCO3 into a 50-mL volumetric flask. For NH3-NH4+ buffer solution of pH 10, 2.06 g of NH4Cl was dissolved in 14.3 mL of concentrated ammonia and was diluted to mark in a 250-mL volumetric flask. Buffer solution was used since buffers are resistant to pH changes[13].

Maintaining the pH is important in preventing interference of other species during titration since different chelates form at a particular pH.[14] For the standardization of 0.01 M EDTA Solution, 10 mL each of 0.0050 M working standard CaCO3 solution was transferred into each of the three 250-mL Erlenmeyer flask using a pipette. Then, into each flask, 75 mL of distilled water was added followed by 3 mL of the NH3-NH4+ buffer solution and 2-3 drops of Eriochrome Black T (EBT) indicator. Although use of EBT indicator is unsatisfactory in calcium, it is ideal to use in magnesium titration[15], and since MgCl2 was put earlier, the number of calcium ions can be determine using EBT indicator[16]. One at a time, the solutions were titrated with the 0.010 M standard EDTA solution. Water sample was analyzed by measuring 50 mL of commercial mineral water Viva into each of the three 250-mL Erlenmeyer flask. Then, into each flask, 75 mL of distilled water was added followed by 3 mL of the NH3-NH4+ buffer solution and 2-3 drops of EBT indicator. One at a time, the solutions were titrated with the 0.010 M standard EDTA solution. RESULTS AND DISCUSSION

Complexometric titration was used in the experiment since the reaction between the aqueous solutions of the analyte (CaCO3 solution, water sample) and titrant (EDTA) forms a complex. Which involves a coordination center composed of Ca2+ and Mg2+ and the chelating agent EDTA. EDTA, a weak acid, commonly forms 1:1 stochiometric ratio when it reacts to form soluble complexes with metal ions, this means that a single endpoint would be observed. Most of the time EDTA reacts with metals regardless of their charges. These would all correlate to a sharp endpoint in titration and a smooth calculation in stoichiometry. Titration with EDTA is affected by several factors such as the existence of complex forming ions and of organic solvents that affects the stability of the complex, the metal ion components, and the pH wherein the titration was performed. The pH range for optimal indications using EBT indicator and for better results in titration using EDTA method is from 8-10.

Lower pH would form a colorless complex with EDTA while a high pH makes it hard to distinguish using the metal indicator In the experiment the pH was kept constant at 10, this was possible with the presence of the buffer solution of NH3 –NH4Cl. It has a buffer capacity that satisfies the optimal pH range. Buffer solutions resist pH change that might be caused by other cations and the weak acid titrant, EDTA. Also, the indicator EBT would behave as it should be if there are no fluctuations in the pH. The specific pH was essential because at the pH of 10 EDTA would deprotonate just enough to bind with the metals involved. If too much buffer was added to the solution, the titration would yield defective endpoints. For example the pH was at 12, the solution would be too basic that it might form precipitates with magnesium and calcium which in turn would cause different results. The endpoint of the solution in the first trial was blue so we opt not to put KCN in the solution. KCN bonds with iron so that iron would not affect the color change of the indicator. If iron is present in the sample it would affect the color endpoint and turn to violet instead of blue. Chemical equations that express the reaction in the titration can be shown in figure 1.

Figure 1. Chemical equations involved in the titration.
In the sample analysis of Viva mineral water, it contained 54mgCa/L and 14mgMg/L. After computing for the total hardness of the sample using ppm CaCO3 it was found out that the claimed total hardness was 192.6 ppm CaCO3 while the computed average ppm CaCO3 from the experiment was 139.5 ppm CaCO3this means that the calculated value from the experiment is less than the calculated total hardness of Viva mineral water according to the indicated value in the label but still in the range of hard according to the water hardness scale in table 1. Table 1 Water Hardness Scale

The unit ppm CaCO3 was used because water is mostly composed of calcium and magnesium ions. Both of these ions can be expressed in terms of CaCO3 One possible source of error is the human error from differentiating color change of the indicator EBT. The solution might have turned violet but not observed making the titrant endpoint wrong because of the presence of iron. Other possible sources of error are excess buffer solution that will increase pH, calibration error of pH meter, wrong volume reading, and over titration.

SUMMARY AND 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

[1] Carillo, K.J.D., Ballesteros, J.I., et al. Analytical Chemistry Laboratory Manual, 2009 edition, UP Chemistry Alumni Foundation, 2009, p. 67

[2] Skoog, D.A., West, D.M., et al., Introduction to Analytical Chemistry, 8th edition, Cengage Learning Asia Pte Ltd., 2012, p. 403

[3]Hardwater,http://water.me.vccs.edu/concepts/hardwater.html

[4] Wurts, W.A., Understanding Water Hardness, http://www.ca.uky.edu/wkrec/Hardness.htm

[5] Skoog, D.A., West, D.M., et al., Introduction to Analytical Chemistry, 8th edition, Cengage Learning Asia Pte Ltd., 2012, p. 372

[6] Skoog, D.A., West, D.M., et al., Introduction to Analytical Chemistry, 8th edition, Cengage Learning Asia Pte Ltd., 2012, p. 386

[7] Ethylenediaminetetraacetic acid disodium salt dehydrate,http://www.sigmaaldrich.com/etc/medialib/docs/Sigma/Product_Information_Sheet/e5134pis.Par.0001.File.tmp/e5134pis.pdf

[8] Liu, N. et. al., Evaluation of Nitric Oxide Removal from Simulated Flue Gas by Fe(II)EDTA/Fe(II)citrate Mixed Absorbents, http://pubs.acs.org/doi/abs/10.1021/ef300538x?prevSearch=Uses%2Bof%2BEDTA&searchHistoryKey=

[9] Skoog, D.A., West, D.M., et al., Introduction to Analytical Chemistry, 8th edition, Cengage Learning Asia Pte Ltd., 2012, p. 372

[10] Skoog, D.A., West, D.M., et al., Introduction to Analytical Chemistry, 8th edition, Cengage Learning Asia Pte Ltd., 2012, p. 384

[11] Carillo, K.J.D., Ballesteros, J.I., et al. Analytical Chemistry Laboratory Manual, 2009 edition, UP Chemistry Alumni Foundation, 2009, p. 69

[12] Ethylenediaminetetraacetic acid disodium salt dehydrate,http://www.sigmaaldrich.com/etc/medialib/docs/Sigma/Product_Information_Sheet/e5134pis.Par.0001.File.tmp/e5134pis.pdf

[13] Whitten, K.[et. Al.], Chemistry.8th ed., Thomas Higher Education. USA. 2007, p. 742

[14] Skoog, D.A., West, D.M., et al., Introduction to Analytical Chemistry, 8th edition, Cengage Learning Asia Pte Ltd., 2012, p. 401

[15] Skoog, D.A., West, D.M., et al., Introduction to Analytical Chemistry, 8th edition, Cengage Learning Asia Pte Ltd., 2012, p. 399

[16] Skoog, D.A., West, D.M., et al., Introduction to Analytical Chemistry, 8th edition, Cengage Learning Asia Pte Ltd., 2012, p. 400

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