UPD Chem 26.1 - Formal Report for Experiment 7

QUANTITATIVE DETERMINATION OF TOTAL HARDNESS IN DRINKING WATER BY COMPLEXOMETRIC EDTA TITRATION A. P. AMBA1 and N. J. C. TANGGAAN2 1,2NATIONAL

INSTITUTE OF GEOLOGICAL SCIENCES, COLLEGE OF SCIENCE UNIVERSITY OF THE PHILIPPINES, DILIMAN, QUEZON CITY 1101, PHILIPPINES DATE SUBMITTED: 7 JULY 2015 DATE PERFORMED: 3 JULY 2015

ABSTRACT Water is very vital in sustaining life because no organism is independent of water for survival. Humans need water in a daily basis. About an average of 3 liters is recommended for daily water intake. Drinking water which comes from below the ground dissolves calcium and magnesium from rocks and the soil which makes the water hard. Water hardness is simply the amount of calcium and magnesium dissolved in water. Hard water can be beneficial since it contains minerals that the body needs but it also has disadvantages such as clogging of pipes when it builds up. The experiment aims to determine the total water hardness of a water sample by complexometric titration. It involves the formation of a complex through the reaction of the titrant and the analyte. Water samples of commercialized mineral water, specifically Hidden Spring, were titrated with EDTA using Eriochrome Black T indicator. Endpoint is indicated with a color change from wine red to clear blue. Through experimentation, it was found out that the average total hardness of the mineral water sample was 187.29 ppm with a percent error of 9.33% based from the theoretical value of the total water hardness that is 206.57 ppm and a relative standard deviation of 12.67ppt. The calculated total hardness states that the mineral water sample is hard which means that it has a high calcium and magnesium content.

INTRODUCTION When astronomers search for life outside the earth, they would first try look at the presence of water because when there is water, there might, just possibly, be life. Water is vital in sustaining life. This is because there isn’t an organism in this world that does not depend on water. There are also other uses of water in our lives. We grow plants using water, keep livestock using water, and drink water. Water is even a part of us. The human body contains about 70 percent water. [1] A certain amount of water intake is necessary for humans to survive or at least be healthy. Eleven to fifteen glasses or approximately 2.7-3.7 liters is recommended

for daily intake of water. Drinking water is primarily obtained from surface water or groundwater which is then transported to homes and other facilities to provide wider and easier access to consumers. [2] However, as the water is collected from below, it picks up minerals from the rocks and soil such as magnesium and calcium which then contributes to water hardness. The simple definition of water hardness is the amount of dissolved calcium and magnesium in the water. The more calcium and magnesium dissolved in the water, the higher the level of water hardness. One can observe it by washing hands or the dishes. With hard water, one requires more soap or detergent

to clean things because soap reacts with the calcium forming “soap scum”. But hard water can be beneficial to. Humans need minerals to stay healthy and hard drinking water generally contributes a small but significant amount toward total calcium and magnesium dietary needs. [3] Some studies even showed that hard water can lessen heart illnesses and provide calcium ions for bone growth. [4] But it also brings setbacks especially in residential settings. Minerals an build up in pipes and block the water transport, causes surfaces to be hazy as it leaves mineral deposits behind, and produces soap curd which reduces cleansing action of soaps and irritates the skin. [5] Water hardness is expressed in ppm CaCO3 although the hardness is not entirely because of the calcium. In the experiment, a certain type of titration called complexometric titration was used. Complexometric titration is a form of volumetric analysis in which the formation of a colored complex is used to indicate the end point of a titration. Complexometric titrations are particularly useful for the determination of a mixture of different metal ions in solution. An indicator capable of producing an unambiguous color change is usually used to detect the end-point of the titration. In theory, any complexation reaction can be used as a volumetric technique provided that the reaction reaches the reaction rapidly after portion of titrant is added, interfering situations do not arise, and a complexometric indicator capable of locating equivalence point with fair accuracy is available. [6] Ethylenediaminetetraacetic acid (EDTA) was used in the experiment. In a complexometric titration, the analyte and the titrant react to form a coordination complex [7]. Coordination compounds are complexes consist of a transition metal as the central atom bound to ions called ligands [8]. A ligand could be any ion with a pair of nonbonding electron. Ligands are classified as monodentate, bidentate, or polydentate depending on the number of electron pairs

they have that could attach to the metal ion [8]. Bidentate or polydentate ligands can form two or more bonds, respectively, to the metal ion. Molecules that form multiple bonds to the metal ion are called chelating compounds. This chelate effect accounts for the higher stability complex of the bidentate or polydentate ligands than the monodentate ones. In the experiment, complexometric titration was used to analyze the water sample using ethylenediaminetetraaceticacid (EDTA) as the titrant. EDTA is a hexadentate ligand which means that it could form six bonds with the metals in the water sample such as the calcium or magnesium ions. This reaction could then produce a coordination compound. EDTA could give different numbers of hydrogen depending on the pH required. Calcium and magnesium react weakly with EDTA; hence, the titration must be done in basic solution. The experiment was then set to achieve a constant pH of 10 for alkalinity. To maintain this pH, EDTA should be fully deprotonated to the form Y4-.Y is an even shorter notation used by analytical chemists for the EDTA. Upon titration, EDTA, in this form, reacts with both the calcium and magnesium ions in the water sample as expressed in equations (1) and (2). (1) (2) At pH equal to 10, the Eirochrome Black T (EBT) indicator was used and the endpoint of the titration process was exhibited by the color change from wine red to a clear blue. At the endpoint, the ppm of CaCO3was measured which consequently measured the total hardness of the water sample. METHODOLOGY Firstly, the solutions that were needed for the experiment were prepared, namely 500.0 mL of 0.1000 M stock EDTA solution, 100.0 mL of

0.0500 M stock Ca2+ solution, and 250.0 mL of 1.0 M NH3-NH4+ with pH 10 buffer solution obtained from NH4Cl solids and NH3 stock solution. Then, 250.0 mL of 0.0100 M working EDTA solution was prepared from the stock EDTA solution and 50.0 mL of 0.0050 M working standard Ca2+ solution from 0.0500 M stock solution. For the standardization of the EDTA solution, 10.0 mL of the working calcium solution was pipetted into each of three 250-mL Erlenmeyer flasks and 75 mL of distilled water was added to each flasks. 3 mL of buffer solution was added, which was followed with 2 drops of EBT indicator. The flask was swirled. The initial volume of the titrant was recorded. The calcium solution was titrated until endpoint was reached, which was a color change from wine red to clear blue. The final volume was recorded and the procedure was repeated for the other two flasks. For the analysis of the total hardness in drinking water, a sample of drinking water (in the case of this experiment, the brand was Hidden Spring) was procured. Then, using a 50-mL volumetric pipette, 50.0 mL of the water sample was measured and poured into a 250-mL Erlenmeyer flask. The titration steps in the standardization procedure were followed, and this was performed in triplicates. RESULTS AND DISCUSSION During the preparation of the stock EDTA solution, MgCl2∙6H2O crystals were added. These crystals were added in order to achieve a sharper endpoint. The EBT indicator forms a complex with the magnesium ions, which results to the wine red color in the solution. The endpoint of direct calcium titration is not that sharp compared to that of magnesium’s [9]. NaOH pellets were also added in order to make the dissolution of EDTA faster. These pellets convert EDTA into a more soluble salt, therefore increasing the solubility of the EDTA solution. Meanwhile, concentrated HCl

was added to the preparation of the stock Ca2+ solution while it was heated to make the dissolution of CaCO3 solids faster. The pH of the analyte needs to be maintained at pH 10. In order for EDTA to react with metal ions, the hydrogens that are attached to the carboxylate group in the molecule must be removed. So this means that the solution must be basic in order for the metal ions to bond with EDTA [10]. This is why the pH needs to be maintained at a basic level. A NH3-NH4+ buffer solution was chosen as the buffer because this is the solution that can be easily prepared at pH 10. The indicator used in the titration of the analyte was Eriochrome Black T, or EBT. It is widely known for its use in determining the water hardness in a water sample, since it bonds with magnesium and calcium, which then produces a wine red solution [11]. For the chemical equations that will be mentioned, EDTA will be denoted with the letter Y while the indicator EBT will be denoted with In. For the standardization of the EDTA solution, the primary standard that was used was calcium carbonate. Since it only contains calcium ions and no magnesium ions, the EBT indicator will then form a complex with the calcium ions and forms a wine red solution, which is shown in the equation below (3): Ca2+ + HIn- → CaIn- + H+ (3) It is noted that not all calcium ions form a complex with the indicator. The free calcium ions then form a complex with EDTA during titration, shown in the equation below(4): Ca2+ + Y4- → CaY2- (4) When all the calcium ions are consumed, the calcium in Ca-EBT will then form a complex with EDTA to form Ca-EDTA, since Ca-EDTA has a higher formation constant (Kf = 5.0 x 1010) than Ca-EBT. The solution then returns to its acidic form since it is titrated with EDTA

and then forms a clear blue solution (endpoint), as shown in the equation below (5):

interval of 187.29 ± 5.95 ppm at 95% confidence level. This was also classified as very hard water.

CaIn- + Y4- + H3O+ → CaY2- + HIn2- + H2O (5) (wine red) (clear blue)

The percent error of the experimental value was found out to be 9.33% from the theoretical value. Possible sources of error could be overtitration of the analyte, since the time it takes for EDTA to react is high, so the titration process should be done slower. The other ions that will also bond with EDTA are also not accounted in the experiment, which could be a source of error.

In the analysis of water sample, EBT will form mostly with magnesium ions than it does with calcium ions, since Mg-EBT (log Kf = 7.0) has a higher formation constant than Ca-EBT (log Kf = 5.4)[12]. Upon titration of the sample, EDTA forms a complex with the free calcium, shown in equation (4) and magnesium ions, which is shown in the equation below (6): Mg2+

+

Y4-

→

MgY2-

(6)

When all the free metal ions have formed a complex ion with EDTA, the magnesium in Mg-EBT will then react with EDTA, since CaEDTA and Mg-EDTA have a higher formation constant than Mg-EBT. The solution then returns to its acidic form since it is titrated with EDTA which then results to a clear blue solution (endpoint), which is shown in the equation below (7): CaIn- + Y4- + H3O+ → CaY2- + HIn2- + H2O (5) (wine red) (clear blue) The molarity of the standardized EDTA solution was found out to be 7.84 x 10-3 M and a titer of 0.7847 mg CaCO3 per mL of EDTA. Through these values, the calcium carbonate content of the water sample can then be obtained. The theoretical ppm of CaCO3 or the total cation content was obtained based from given values from the bottle of the water sample. It was found out that the theoretical value was 206.57 ppm CaCO3, which was classified to be very hard water. The concentration of CaCO3 obtained from the titration of the sample was found out to be 187.29 ppm, with a relative standard deviation of 12.67 ppt and a confidence

SUMMARY AND CONCLUSIONS Determination of water hardness is important for awareness. Water hardness can be quantitatively determined by complexometric titration. The volume of EDTA utilized in titration was used to determine the amount CaCO3 in the sample because of a complex formation reaction that occurs. The result of the experiment showed that the water sample, having an average CaCO3 concentration of 187.29 ppm which slightly differs with the theoretical value of 206.57 ppm, is hard. This value needs to be considered if the water sample is to be used for industrial purposes. Errors in the calculated result could be attributed to the possible miscalculation of the EDTA concentration that could have affected the calculated CaCO3 concentration. Also, there could have been errors in titration especially in volume measurements that made the calculated 187.29 ppm of the calcium carbonate lesser than the theoretical value. Also, the tolerance of the glassware and instruments used may be a possible source of errors in the experiment. Although the calculated and theoretical value hardness proved that the water sample is hard, it is highly recommended to conduct more experiments like to test hardness which gives more accurate results. The

standardization of EDTA and the titration of the sample itself should be done carefully to achieve the most accurate concentration leading to a more accurate result.

[7] "Complexometric Titration." , EDTA Titration. N.p., n.d. Web. 06 July 2015.

REFERENCES

[8] "Transition Metals." Coordination Complexes and Ligands. N.p., n.d. Web. 06 July 2015.

[1] Atteberry, Jonathan. "Why Is Water Vital to Life?" HowStuffWorks. HowStuffWorks.com, n.d. Web. 06 July 2015. [2] Water Quality Association. (May 2, 2000). Retrieved 8 May 2014 from: http://www.wqa.org/sitelogic.cfm?ID=477 [3] Water Quality Association. (May 2, 2000). Retrieved 8 May 2014 from: http://www.wqa.org/sitelogic.cfm?ID=477 [4] Lenntech. "Water Treatment Solutions." Calcium (Ca) and Water. N.p., n.d. Web. 06 July 2015. [5] Forbes, Rhomylly. "The Effects of Drinking Hard Water." EHow. Demand Media, 12 Dec. 2010. Web. 06 July 2015. [6] Schwarzenbach, Gerold, and Harry Irving. Complexometric Titration. London: Methuen and, 1960. Web.

[9] Kundu, P.C. Complexometric Titration of the Sum of Calcium and Magnesium Using Aluminon as Metal Indicator. Chemical Laboratory, Presidency College, West Bengal, India. 1962. Retrieved 5 July 2015. [10] Southway, C. APCH231: Chemical Analysis Complexometric Titrations EDTA. ND. Retrieved 5 July 2015. [11] Young, A., Sweet, T. Complexes of Eriochrome Black T with Calcium and Magnesium. Ohio State University. 1954. Retrieved 6 July 2015. [12] Hulanicki, A., Glab, S. Complexometric Indicators: Characteristics and Applications. International Union of Pure and Applied Chemistry. 1983. Retrieved 7 July 2015.

APPENDIX Sample Calculations: Molarity of Working Solution of CaCO3:

Standardization of EDTA solution: Trial 1:

Trial 2:

Trial 3:

Average M EDTA:

Titer in mg CaCO3 per mL EDTA:

ppm CaCO3 from ppm Ca:

ppm CaCO3 from ppm Mg:

Total Water Hardness:

Sample Analysis: Trial 1:

Trial 2:

Trial 3:

Average ppm CaCO3

Relative standard deviation (RSD):

Confidence Interval:

Percent Error: