QUANTITATIVE DETERMINATION OF TOTAL HARDNESS IN DRINKING WATER BY COMPLEXOMETRIC EDTA TITRATION

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QUANTITATIVE DETERMINATION OF TOTAL HARDNESS IN DRINKING WATER BY COMPLEXOMETRIC EDTA TITRATION UNIVERSITY OF THE PHILIPPINES, DILIMAN, QUEZON CITY 1101, PHILIPPINES ______________________________________________________________________________ ABSTRACT The determination of water hardness is a useful tool to provide a measure for quality of water within households and for industrial uses. In this experiment complexometric EDTA titration was used to determine the mineral content of Hidden Spring Mineral Water. A 0.0500 M EDTA solution was standardised with 99.5% pure calcium carbonate. Using the ammonia-ammonium buffer the environment of the reaction was kept alkaline at pH = 10 and the indicator used was Eriochrome Black T or EBT. The analysis of the water sample involved the the titration of 50 mL of the mineralised water mixed with buffer solution and indicator solution with EDTA solution. The change of color from red wine (color of EBT bound to metal ion) to clear blue (color of free EBT ion) indicated the endpoint. The method yielded the hardness to be 219.63 ppm CaCO3 which classifies the water as very hard. The total cation content computed from the content given on the waters label was 206.87 ppm. The relative standard deviation was computed and it resulted to 0 ppm and confidence interval at 95% confidence level obtained was 217.7 ± 0. During the experiment intrinsic errors of the method and typical errors like human errors may have occurred.

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INTRODUCTION

2 Complexometric titration is considered to be one of the best ways to determine water hardness.

METHODOLOGY In this experiment the total hardness of Hidden Spring Mineral Water was determined quantitatively using complexometric titration with ethylenediaminetetraacetic acid (EDTA) as the titrant. The large molecule known as EDTA forms a complex with calcium and magnesium. A blue dye known as Eriochrome Black T (EBT) was used as an indicator. EBT also forms a complex with calcium and magnesium ions which causes the

Minerals that are dissolved in groundwater produce metal ions like manganese, calcium, strontium, aluminum, barium, zinc, magnesium and iron. These metal ions are considered to be good for the body, which is why mineralised water contains these ions. Calcium and magnesium are the most abundant ions in mineralised water because they have become commercialised. The ions present in water give it a “hardness”. The determination of water hardness is a useful test to measure the quality of mineral water for households and industrial uses. The hardness of water a concern because hard water increases soap consumption, leaving a soapy sink, causing burn out of electrodes or discolouring plumbing fixtures and utensils. [1] The total hardness of water is determined by the sum of the calcium and magnesium salts dissolved in water. This is expressed as milligrams per litre (mg/L) of water, or for more practical reasons parts per million (ppm) of CaCO3. The hardness of water can be determined by the table below.

change of the solutions color from red wine to clear blue endpoint. [3]

Figure 1. Ethylenediaminetetraacetic acid (EDTA) Figure 2. Eriochrome Black T (EBT)

Table 1. The water hardness scale. [2] Water Hardness

ppm CaC

Soft

0-20

Moderately Soft

21-60

Moderately Hard

61-120

Hard

121-180

Very Hard

> 180

Water hardness is usually determined in one of two ways. One of these methods would be complexometric titration, also known as chelatometry. Complexometric titration is a form of volumetric analysis where the formation of a colored complex is used to indicate the end point of a titration. This titration depends on the combination of ions to form a soluble complex of ions like magnesium, calcium, iron, zinc and lead. These titrations are useful for determining the mixture of different metal ions in a solution. [1]

3 this experiment NaOH pellets, facilitates dissolution of acid form of EDTA. A carbonate error can cause discrepancies in the pH reading. The addition of HCl while dissolving CaCO3 during solution preparation is important for all reactions because the metal ions and EDTA are pH dependent. The divalent ions, solutions must be kept basic for the reaction to complete. [4] Ligands are mostly basic and they bind to H+ ions throughout a wide range of pH. Some of the H+ ions displace from the ligands by the metal during chelate formation. The buffer was used to hold the pH constant. As observed, the standard was allowed to react in a basic medium by the addition of the basic buffer of pH 10. In this experiment, NH3-NH4Cl buffer was used. A buffer was added for the pH of the whole reaction to remain constant. A constant pH is needed in the titration process since the EDTA and EBT have polyprotic properties, making them unstable. At different pH levels the hydronium ion concentration may interfere with the complexation of the EDTA with the calcium and magnesium ions. Also, the effective stability constant of EDTA varies with the pH because it depends on its degree of ionization. [4] EBT was used as an indicator in this experiment. Most complexometric titrations use indicators that are chelating agents and have metal complexes that have different color from the reagents themselves. The same principle on the ionization and dissociation of EDTA also applies to EBT since it is also polyprotic. [5] At the beginning of the titration of EDTA to 10mL 0.0050 M working standard CaCO3 with 3mL buffer and 2-3 drops of EBT, the solution becomes wine red. This is due to the complexation of the Mg 2+ ions with the indicator. Upon titrating with EDTA, the color gradually changed from wine red to clear blue, indicating that the endpoint was reached. The endpoint occurs when Ca2+ ions complexes with EDTA, and at the same time the Mg 2+ complexes with it. The Mg-EBT complex breaks as illustrated in the equation.

MgIn- + Y4- + H+ ⟶ MgY2- + Hln2-

For the standardisation of 0.01M EDTA solution three 250mL erlenmeyer flasks were set and 10mL of 0.0050M working standard, CaCO 3 solution, was measured in to each. Then 75 mL of distilled water was added into the 10mL working standard. Three mL of the buffer solution NH3-NH4 Cl with pH = 10, followed by 2-3 drops of EBT indicator was added to the solution. Immediately after the addition of the buffer and indicator the titration with the standard EDTA solution followed. The solution was titrated until the clear blue endpoint was reached. For the analysis of the Hidden Spring Mineral Water 50mL of the water sample was measured into three separate 250mL erlenmeyer flasks. Then three mL of the buffer and 2-3 drops of the indicator was added to the sample immediately followed by titration with standard EDTA solution until the clear blue endpoint was observed.

RESULTS AND DISCUSSION During the preparation of the EDTA solution MgCl 2 x 6H2O crystals were added to the dissolved EDTA salt. When determining the hardness of water it is determining the concentration of calcium carbonate. Calcium ion determination via EDTA is also known as displacement titration because EDTA binds with calcium ions very slowly. By addition of another metal, like magnesium in the form of salt it could bind to EDTA quicker forming a Mg-EDTA complex which is less stable than CaEDTA complex. The Mg-EDTA has a higher formation constant giving it a higher tendency to form complexes than Ca-EDTA. Then the free Ca 2+ ions in the solution replace the Mg forming CaEDTA complex and resulting in free Mg 2+ ions which can be titrated easily. The amount of liberated Mg 2+ is equivalent to the original concentration of Ca 2+. Ca2+ + EDTA4- ⟶ CaEDTA

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Kf = 5.0x1010 Kf = 4.9x108

Mg2+ + EDTA4- ⟶ MgEDTA There is a greater energy for the complex to form since the Ca-EDTA has a lower Kf, formation constant making it is less stable. Adding more Mg 2+ ions will make the endpoint sharper. EDTA is considerably insoluble in water, and only dissolves when pH is neutralized at 8. Addition of base, in

4 indicator used and determine whether it is suitable for the type of titration. In this experiment EBT was used since it selectively binds with the ions. Neglecting to add a buffer can also cause error in the concentration of the standardised EDTA. The reaction is very sensitive to pH change, so it is important to note that proper selection of kind of buffer based on each pH range, in this experiment it is at 10.

SUMMARY AND CONCLUSIONS The standardized EDTA resulted to a 0.0084 M after titrating it thrice with primary standard CaCO 3. Hidden Spring Mineral Water was used as the sample and has an initial Ca2+ content of 25 ppm and Mg2+ content with 35 ppm. The water sample was titrated with standardized EDTA and EBT was used as an indicator. The selection of the indicator was based on the range of its pH, 8 to 10. NH3-NH4Cl buffer with pH at 10 was used to maintain the pH of the solution because EDTA is pH dependent. After three trials of titration, the computed hardness of water were 219.36 ppm, 214.30 ppm, 219.36 ppm and making the average hardness 217.67 ppm. The values were checked using the Grubbs-test to eliminate outliers. In this experiment the outlier was 214.30 ppm making the new average 219.36 ppm. After computing the relative standard deviation obtained was 0 ppm. The calculated confidence interval at 95% confidence level was 217.7 ppm to 217.7 ppm. Based from the labels ion content the total ion content was computed to be 206.8665931 ppm which was smaller than the value obtained through titration. The difference in the value may be caused by either over titration or even inaccurate solution preparation. These factors may have affected parameters such as the concentration of standardised EDTA and total hardness based from the titration data. The two calculated hardness have showed the level of hardness of Hidden Spring Miner Water which is very hard. The hardness of the water shows that

wine red

clear blue

The computed average hardness from the three trials is 217.67382700 ppm (see Appendix C for calculations). According to the water hardness scale in Table 1 the water sample is very hard.

Table 2. Mineral Water. Trial

Analysis of Hidden Spring

1

Water Sample

50mL

Final volume EDTA

13mL

Initial volume EDTA

0mL

Net volume EDTA

13mL

Total Hardness

219.36

The Grubbs test was first conducted to determine which values will be accepted, and only the values from trial 1 and trial 3 were accepted from the calculations making the new average 219.63 ppm. The standard deviation computed was 0 ppm and the relative standard deviation was also 0 ppt. The confidence limit was also computed with confidence interval at 95% and the resulting interval was 217.7 ± 0. From the Hidden Spring’s given ion content of Ca 2+ and Mg2+ the total ion content could be computed. This is done adding the total CaCO 3 ppm from CaCO3 and MgCO3. The computed value was 206.87 ppm (see Appendix for calculations). Based on Table 1, the calculated ion content classifies the water as very hard. The values obtained and calculated may be effected by certain errors made throughout the experiment. Discrepancies may be caused by several factors. Over titration of the sample analysis is a common source of error. The delayed change of color of the solution could be a possible cause for over titration. It is important to take note of the

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5 pertinent ions such as Ca and Mg , are present and are acceptable for human consumption. 2+

[3] N.p.. Web. 29 Jan 2014.

2+

REFERENCES

[4] N.p.. Web. 29 Jan 2014. .

[1] Spellman, F. The Science of Water: Concepts and Applications, Second Edition. 2000. Boca Raton, Florida: CRC Press.

[5] Chang, R. 2007. Chemistry: Ninth Edition. New York: McGraw Hill. 549-553.

[2] Institute of Chemistry, University of thePhilippines. Analytical Chemistry: Laboratory Manual. 2013.