QUANTITATIVE DETERMINATION OF TOTAL HARDNESS IN DRINKING WATER BY COMPLEXOMETRIC EDTA TITRATION JOHN JOSEPH CHUA DEPARTMENT OF FOOD SCIENCE AND NUTRITION, COLLEGE OF HOME ECONOMICS UNIVERSITY OF THE PHILIPPINES, DILIMAN, QUEZON CITY 1101, PHILIPPINES DATE SUBMITTED: 26 MARCH 2015 DATE PERFORMED: 6 APRIL 2015
ABSTRACT Complexometric Ethylenediaminetetraacetic Acid (EDTA) Titration is used to quantitatively determine the total hardness of Hidden Spring Drinking Water. Total Water Hardness, which is measured in ppm CaCO3 , is the intensity of the presence of calcium and magnesium ions in water that is detected by formation of complexes. Through the application of knowledge regarding complex formation and complex stability, the methodology of the titration for the determination of 2calcium and magnesium ion concentration is concisely discussed. EDTA, in the form of H 2 2Y at pH 10, is used as the complexing agent of titration. Through three trials of the complexometric titration, the computed total water hardness for the experiment is 215.1450363 ppm CaCO3 , whereas the total water hardness indicated on the bottle label is 206.5834385 ppm CaCO3. With respect to the total water hardness on the bottle label, the experimental value deviates by 4.144%. Overall, the experiment is a success due to its relative closeness to the value indicated on the bottle label.
INTRODUCTION
Water hardness is the traditional measure of the capacity of water to react with soap; hard water requiring considerably more soap to produce lather [1]. Predominance of metallic ions such as Calcium and Magnesium cations in hard water allows formation of insoluble mineral deposits which may reduce efficiency in industrial processes such as heating and in the cleansing action of soaps [2]. Table 1. Water Hardness Scale (3)
Water Hardness Soft Moderately Soft Moderately Hard Hard Very Hard
ppm CaCO3 0-20 20-60 61-120 121-180 >180
Water hardness is assigned with levels of ppm CaCO3, which are tabulated in table 1. Amount of CaCO3 in milligrams per liter is a
general term that indicates the total quantity of divalent salts present in a water sample. It does not specifically identify whether calcium, magnesium, or other divalent salt is the main cause of the hardness of water but it simplifies calculations due to the ~100 g/mol formula mass of CaCO3 [4]. Since hard water hinders efficiency in industrial processes, soft water is preferred. On the other hand, hard water is inclined to human consumption. Presence of the metallic ions Ca 2+ and Mg2+ can aid in the human body’s calcium and magnesium needs [3]. Although there are no convincing evidences that relate hardness of water to adverse effects, more research are needed to establish direct effects and relations regarding human consumption [1]. Complexometric titrations involve the application of complexation reactions. In this method of titration, a metal ion reacts with a suitable ligand to form a complex, and the equivalence point is determined by 1
an indicator or an appropriate instrumental method [5]. Through the addition of the titrant, identification of the amount of metal ions in the analyte can be quantitatively determined. It is already mentioned that total water hardness is characterized by presence of calcium and magnesium ions; therefore, complexometric titration is used to determine total hardness in drinking water. To determine the amount calcium and magnesium, which determines the total water hardness in the drinking water sample, Ethylenediaminetetraacetic acid or EDTA is used as the complexing agent or the titrant in the solution. EDTA is a widely used complexometric titrant. It also has six potential sites for bonding a metal ion: four carboxyl groups and two amino groups, wherein the latter has an unshared pair of electrons. With these potential sites, EDTA is said to be a hexadentate ligand, which also makes it a chelating agent. [4]
Figure 1. EDTA Structure [6]
The metal cations to be given a quantitative analysis have a charge of 2+, which indicates that to maximize the complexation of these metal cations, EDTA must be kept in its form as H 2Y2-, which is its halfneutralized form. To maintain this form of EDTA, a pH of at least 10 is required. Complex formation constant, denoted as K f , is an equilibrium constant which is the ratio of the concentration of products versus the concentration of the reactants raised to their appropriate stoichiometric factors. A large value for the K f indicates that the
formation of the complex studied is very probable. EDTA is used in the determination of total water hardness as a titrant because the Formation Constants (Kf ) of EDTA with the discussed metal cations are relatively high. Formation Constants for EDTA Complexes Cation K MY log K MY Ag+ 2.1 x 10 7 7.32 Mg2+ 4.9 x 10 8 8.69 2+ 10 Ca 5.0 x 10 10.70 2+ 8 Sr 4.3 x 10 8.63 Ba2+ 5.8 x 10 7 7.76 Mn2+ 6.2 x 10 13 13.79 Table 2. Formation Constants for EDTA Complexes
As tabulated in table 2, the Kf of Mg2+ with EDTA is 4.9x108 and 5.0x10 10 for Ca 2+. It has been previously stated that the species required for complexation of the said metal ions is H2Y2-, which is present at pH 10. Therefore, for the complexation to be favored, the system should be maintained at pH 10. Complexometric titration also deals with visible indicators that mark the end point of the titration. CaH2Y and MgH2Y are colorless complexes, but with the addition of Eriochrome Black T as an indicator, physical observations in the titration endpoint can be seen. Eriochrome Black T is used to generate an unambiguous color change from wine red to sky blue. [4] To observe this color change, complexometric titration involving EDTA develops a series of displacement reactions, which lie on which has higher formation constants. Mg2+ + ErioT → ErioT-Mg (1) According to (1), addition of Eriochrome Black T to the analyte forms a wine red complex, which confirms the presence of Mg2+. Ca2+ + H2Y2- → CaH2Y (2) Mg2+ + H2Y2- → MgH2Y (3) Ca2+ + MgH2Y2- → CaH2Y + Mg2+ (4)
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Addition of the titrant EDTA, which is in the form H2Y2-, will react with Ca2+ to form CaH2Y complex (2) which has a K f value of 5.0x1010; Mg2+ in the drinking water sample that did not complex with Eriochrome Black T will now complex with EDTA to form MgH2Y (3) which has a Kf value of 4.9x108. ErioT-Mg + H 2Y2- → MgH2Y + ErioT (5) At (5), the initial excess of EDTA after complexation with all the Ca 2+ and Mg2+ in the analyte will displace Eriochrome Black T in its complexation with Mg2+ during (1). The displacement reaction results in the free Eriochrome Black T, which produces the sky blue color. The presence of the said color marks the end point of titration, where no metal cations are to form complexes with EDTA. Therefore, the total water hardness, which is expressed in ppm CaCO3, can be calculated from the volume of titrant (EDTA) used and the molarity of the titrant acquired from standardization. To acquire quantitative data of the aforementioned variables, the solutions in table 3 are prepared. Solution EDTA Ca2+(aq) NH3-NH4+ Buffer Drinking Water Sample Eriochrome Black T
Volume 500 mL 100.0 mL 250.0 mL
Molarity 0.100 M 0.0500 M 1.0 M
150 mL X mL
Table 3. Solution Preparation
In the preparation of 0.100 M of EDTA, 1 g of MgCl2∙6H20 and NaOH pellets are added. Addition of the said species increases the solubility of NA2H2EDTA because of the diverse ion effect. The addition of MgCl 2 allows the complexation of EDTA with Mg 2+ which ensures a sharp endpoint of titration. In equation (3), Mg 2+ formed a complex with EDTA, resulting in the presence of MgH2Y complex in the titrant. Therefore, there are “free” EDTA and “Magnesiumbound” EDTA. Calcium ions in the drinking water sample react with the “free” EDTA
forming CaH2Y, while excess Ca2+ will displace Mg2+ in the MgH2Y complex, as seen in equation (4). The displacement is due to the higher K f value of CaH 2Y than MgH2Y, so the production of CaH 2Y is more probable to happen. Because of the hastened formation of CaH2Y complex, excess EDTA will be met faster, thus a sharper endpoint. In preparation of 0.0500 M Ca 2+ equation (6) is involved.
(aq),
the
CaCO3 + CO2 + H2O ↔ Ca 2+ + 2HCO3- (6) Since equation (6) is endothermic, heating allows dissolution of the solid reactant CaCO3, forming Ca 2+ completely. If CaCO3 produced insoluble solids in the solution, addition of HCl will dissolve the precipitate in the reaction (7). CaCO3 + 2HCl → CaCl 2 + H2O + CO2 (7) Addition of the NH3-NH4+ buffer ensures that upon the addition of the titrant there is minimal change in pH. Computed amounts of each reactant compose the buffer which maintains the pH of the system at pH 10. After preparation of the required solutions, a 10 mL of diluted aliquot of the Ca 2+(aq) solution is used for standardization of the EDTA titrant. 3 mL of the NH 3-NH4+ buffer is added, along with the addition of Eriochrome Black T Indicator. The same procedure has been done with the water sample. 50 mL of the drinking water is used for the complexometric titration. All of the titrations are done three times. RESULTS AND DISCUSSION
Complexometric EDTA Titration has been done to show the amount of Ca 2+ and Mg2+ in the drinking water sample, which is referred to as total water hardness once expressed in ppm CaCO3. In an overview, the complexometric titration involved solution preparation, EDTA Solution Standardization, and titration of the drinking water sample. The
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following data presented are the computed results from the methodologies. Trial mL EDTA used M EDTA Average M EDTA
1
2
3
5.9
5.8
6.2
0.0084
0.00852
0.007976
0.008294425 M
Table 4. Standardization of EDTA Solution
Three trials of the standardization of the EDTA solution correspond to the three computed molarities indicated on table 4. The mean value is used as the molarity of EDTA solution for further calculations. All the data underwent and passed the Grubbs Test. Theoretically, the molarity of the EDTA solution is 0.01 M. The computed value of 0.008294425 M deviated from the theoretical value by 17.05575%. The deviation may be caused by inaccurate weighing of the reagents, fast addition of titrant so endpoint was observed more immediately and the percent purity of the analytical reagent is taken to account. The titer of the solution is computed to make the calculations for ppm CaCO 3 easier. Using the average molarity of the standardized EDTA solution, the titer is computed and is expressed as mg CaCO 3 per mL of standardized solution. The computed titer is 0.829608109 mg CaCO 3/mL Standardized EDTA Solution. Trial Net Volume of EDTA Average Volume Standard Deviation Relative S.D Confidence Limits
1
2
3
12.9
13
13
12.966666666667 0.057735027 4.45257277 12.823333
13.11
Table 5. Volume of Titrant Used and Some Statistical Measures
The data on table 5 indicate the volumes of titrant used during the titration involving the drinking water sample. The volume of
EDTA used during the titration has a standard deviation of approximately 0.057735, which is relatively small. Using the computed titer, the total hardness of water can be known. The amount of CaCO3 computed from the titer is 10.75725181 mg. Therefore, the experimental ppm of CaCO 3 is 215.1450363 ppm. The drinking water sample is very hard, according to the boundaries given on table 1. Ions Ca2+ Mg2+ Total Water Hardness
ppm 25 35
ppm CaCO3 62.40648379 144.1769547
206.5834385 ppm CaCO 3
Table 6. Ions and Total Water Hardness Indicated on Bottle Label
The experimental value deviates from the theoretical value by 4.144%. The small deviation is possibly brought about by instrumental error, impurity of reagents and solvents used and imprecise color comparing. Instrumental errors include wrong usage of glassware and wrong weighing of reagents. Presence of calcium and magnesium ions in water that was used in dilution of all solutions will result in an increased computed total hardness. The same result will happen if there are varying degrees of sky blue that is used in determining the endpoint of the titration. If the pH of the system changes by a greater extent, the species of EDTA will change to a more protonated or more deprotonated form which will reduce the complexation of EDTA with the metal ions calcium and magnesium. Drastic changes in pH may be brought about by wrong preparation of the buffer. SUMMARY AND CONCLUSIONS
Through the complexometric EDTA titration of drinking water, the total hardness of Hidden Spring Drinking water is identified. The experimental total water hardness is 215.15 ppm CaCO 3 while the total water 4
hardness indicated on the label in 206.58 ppmCaCO3. The percent error with respect to the indicated total water hardness is 4.144%. Generally, the drinking water is found to be very hard, which is ideal for drinking. Because of this fact, the experiment can be considered a success because of relatively low standard deviations and closeness to the ideal water hardness of the computed ones.
REFERENCES
[1] WHO. (n.d). Hardness in Drinking-water: Background document for development of WHO Guidelines for Drinking-water Quality. Health criteria and other supporting information. Geneva: http://www.who.int/water_sanitation_heal th/dwq/chemicals/en/hardness.pdf [2] Casiday, C., Frey, R. (n.d). Water Hardness: Inorganic Reactions Experiment. Department of Chemistry, Washington University St. Louis. St Louis, MO: http://www.chemistry.wustl.edu/~edudev /LabTutorials/Water/FreshWater/hardnes s.html [3]Oram, B. (n.d). Hard Water Hardness Calcium Magnesium Water Corrosion Mineral Scale. Dallas, PA: http://www.waterresearch.net/index.php/watertreatment/tools/hard-water-hardness [4] University of Canterbury. (n.d). Determination of Total Calcium and Magnesium ion Concentration. College of Science, University of Canterbury. [5] Skoog, D., et.al., (2004). Fundamentals of Analytical Chemistry. Belmont, CA: Brooks/Cole-Thomson Learning.
[6] OpenStar CNX. (n.d). U.S Students in Open-Source Book Program. http://cnx.org/resources/60c9ece3d13106 99f7311bf16087ff7e/graphics4.jpg
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