QUANTITATIVE DETERMINATION OF TOTAL HARDNESS IN DRINKING WATER BY COMPLEXOMETRIC EDTA TITRATION S.M. BLANCAFLOR1 N.A.S. VERACRUZ2 1INSTITUTE
OF BIOLOGY, COLLEGE OF SCIENCE INSTITUTE OF GEOLOGICAL SCIENCES, COLLEGE OF SCIENCE UNIVERSITY OF THE PHILIPPINES, DILIMAN, QUEZON CITY 1101, PHILIPPINES DATE SUBMITTED: 22 AUGUST 2011 DATE PERFORMED: 17 AUGUST 2011 2NATIONAL
ABSTRACT The experiment aims to determine the total hardness in drinking water through complexometric EDTA titration. A volume of commercial mineral water sample, added with NH3-NH4+ buffer solution followed by EBT indicator, was titrated with 0.01 M standardized EDTA solution. The procedure was repeated three times to obtain accurate experimental results. The data obtained were used to solve for total hardness, measured in ppm CaCO 3, and to determine whether it is soft, moderately soft or hard, or hard. Pooled standard deviation was performed from individual data sets to determine the preciseness of volumes of titrant and total hardness in water. Results show that hardness of the water sample is relatively hard, with support from various statistical parameters used to test the data.
INTRODUCTION One common analysis for water, whether it is for domestic or industrial use, is total hardness. Hardness is the total concentration of alkaline earth ions, which are mainly Ca2+ and Mg2+, in water. It is commonly expressed as the equivalent number of milligrams of CaCO3 per liter (ppm CaCO3). Water whose hardness is less than 60 mg CaCO3 per liter is considered to be “soft”. If the hardness is above 270 mg/L, the water is considered to be “hard”. [1] Determination of water hardness is a useful test to measure quality of water for household and industrial uses. Hard water is not a health hazard. Indeed, drinking hard water contributes to a small amount of calcium and magnesium toward the total human dietary needs of these ions. It also causes soap scum, clogged pipes and boilers. [2] A volume of commercial mineral water sample, added with NH3-NH4+ buffer solution followed by eriochrome black T (EBT) indicator, was titrated with standardized ethylenediaminetetraacetic acid (EDTA) solution of 0.01 M. Procedure was repeated three times.
Data obtained were subjected to pooled standard deviation. This helps in assessing the preciseness of the data. [4] Total hardness was calculated to determine the total concentration of Ca2+ and Mg2+ in water. The purpose of this experiment is to be able to apply the concept of complexometric titration in the determination of total hardness in drinking water. RESULTS AND DISCUSSION In the preparation of 500 mL 0.1 M stock EDTA solution, 18.6 g Na2H2EDTA•2H2O were weighed. Distilled water (200 mL) and 1.0 g magnesium chloride hexahydrate (MgCl2•6H2O) crystals were then added and mixed. The solution was heated to dissolve crystals, since dissolution of EDTA may be slow, and was then cooled to room temperature. It was then diluted to 500 mL with distilled water. Standard EDTA solution of 0.01 M (250 mL) was prepared by applying concepts in preparation of diluted solutions.
EDTA, a chelating agent, was used to determine permanent hardness of water sample. It is used to capture metal ions, which causes water to soften. Metal ions are not removed from water, but were rather tightly bonded. EDTA is also frequently used in soaps and detergents because it forms complexes with Ca2+ and Mg2+ ions. These ions, which are in hard water, are bound to EDTA and cannot interfere with the cleaning action of soap or detergent. [2] Mg(II) ions from MgCl2•6H2O form chelates that are not as stable as those produced when EDTA reacts to other ions present. Mg2+ ions, initially bounded to EDTA, are displaced by Ca2+ upon titration. They then combine with EBT indicator, producing a wine-red color, and only recombine with EDTA once all Ca2+ ions are complexed. When titrating with EBT, conversion of MgIn- to HIn2- becomes a more appropriate marker of endpoint as compared to the conversion of CaIn- to HIn2-, which happens before equivalence. [1][3] 2H2Y2- + Ca2+ + Mg2+ CaY2- + MgY2- + 4H+ H2Y2- + MgIn- MgY2 + HIn2- + H+ On the other hand, addition of NaOH decreases turbidity of the solution by ensuring complete dissociation of crystals. It also serves to neutralize H+ ions released from EDTA. In the preparation of 250 mL 0.01 M standard CaCO3 solution, 1.25 g pure CaCO3 were weighed. Distilled water (20 mL) and 6 M HCl dropwise were then added to dissolve CaCO3 completely. Solution was evaporated until about 10 mL was left. It was then diluted to 250 mL with distilled water. Working standard CaCO3 solution of 0.005 M (50 mL) was prepared by applying concepts in preparation of diluted solutions. Likewise, 250 mL 1 M NH3-NH4+ buffer solution of pH 10 from ammonium chloride crystals (NH4Cl) and concentrated ammonia (NH4OH), and 50 mL 0.1 % (w/v) EBT in ethanol were prepared. EBT marks the endpoint of titration and indicates presence of metal ions. Prior to attaining equivalence point, the indicator binds with ions in excess to form wine-red complexes. Once EDTA displaces all EBT bound to metal ions, solution turns blue, marking the endpoint. [1][2]
Calculating EDTA and metal ion content is tedious as the concentrations are pH dependent. However, if pH was kept constant, interference of other cations are prevented, and computations are made possible. In the experiment, an enough endpoint for calcium titrations requires pH 8 or higher. Likewise, pH 7 or above is necessary in order for blue color to be visible when EBT is used as metal ion indicator. Thus, solution must be kept basic, so as to be suitable for cation analysis. [3] As pH must be kept consistent throughout the experiment, a buffer is used. Buffer prevents any unexpected changes on how the indicator behaves. Furthermore, EDTA used is slightly acidic, since it releases H+. Conducive pH level at which experiment must be carried out should be at 8 or greater, and NH3-NH4+ buffer solution serves to meet this requirement. However, buffer should not be in excess as it may react with analyte to form Ca(OH)2 or Mg(OH)2, rendering a poor endpoint. In EDTA standardization, working standard CaCO3 solution (10 mL) was pipetted into each of three 250mL Erlenmeyer flasks, and 75 mL distilled water was added to each of the flasks. Buffer solution (3 mL) was added followed by 20-30 drops of EBT indicator. Solution was then titrated with standard EDTA solution. Procedure was repeated for two more trials. Table 1. Data obtained for standardization of EDTA solution Trial 1 2 3 Volume of working Ca(II) 10 10 10 standard, mL Volume of EDTA, mL Final volume, 7.0 6.9 6.8 mL Initial volume, 0 0 0 mL Net volume, 7.0 6.9 6.8 mL As buffer solution and EBT were added, solution was immediately titrated with standard EDTA solution until
color changes from wine-red to blue. From table 1, obtained net volumes of titrant were relatively precise. If color at the endpoint is violet instead of blue, the solution contains iron. This can be avoided in titration of succeeding samples by adding small amount of KCN once buffer is added. KCN contains CN- ions that can be employed as masking agent when iron is present. This prevents EDTA from reacting with iron present and eliminates interference of iron in Ca2+ and Mg2- analysis. In water sample analysis, 50 mL of commercial mineral water was measured into each of three 250mL Erlenmeyer flasks. Similar to the method of EDTA standardization, 3 mL buffer solution followed by 2030 drops of EBT indicator were added. Solution was titrated with standard EDTA solution. This was repeated for two more trials. Table 2. Data obtained for water sample analysis Trial 1 2 3 Volume of water 50 50 50 sample, mL Volume of EDTA, mL Final volume, 8.8 18.5 30 mL Initial volume, 0 8.8 18.5 mL Net volume, 8.8 9.7 11.5 mL Akin to standardization of EDTA, same amounts of buffer and EBT were added, and solution was immediately titrated with standard EDTA solution until endpoint. From table 2, obtained net volumes of titrant are not that defined. Titration error, e.g. failure to determine the difference between equivalence point and endpoint, could be one reason. Table 3. Calculated values of molarity of titrant, titer of analyte, and total hardness of water sample Total Trial Molarity Titer Hardness 1 7.14 x 10-3 0.715 125.84
2 3 Average RSD Confidence interval
7.24 x 10-3 7.35 x 10-3 7.24 x 10-3 n/a
0.725 0.736 0.725 n/a
n/a
n/a
140.65 169.28 145.26 152 90.44, 200.08
From table 3, calculated total hardness of commercial water sample is not relatively precise, mainly due to the obtained data from table 2. Nevertheless, mean total hardness is 145.26, which is considered hard. A mineral concentration of 120-180 ppm is considered hard. From the label of water sample, total hardness is said to be 68 ppm, with only Ca2+ and Mg2+ ions considered. Other dissolved ions, such as bicarbonates, can contribute to water hardness. The data obtained from table 1, 2 and 3, were subjected to pooled standard deviation to more accurately determine the standard deviation of the analysis. [4] Table 4. Pooled standard deviation of individual data sets from table 1, 2 and 3 Pooled standard deviaiton Standardization of 0.669 EDTA solution Water sample analysis 0.550 Total hardness 16.07 Calculated spooled of both volumes of titrant from EDTA standardization and water sample analysis are reasonably small, thus, data obtained by different groups were precise. However, spooled of total hardness of water sample shows that ppm CaCO3 of different groups was somehow dispersed. Even so, most of the calculated ppm CaCO3 was in the range 120-180 ppm, which is considered hard and a property of most commercial mineral water samples. Possible sources of errors might include human errors (transferring and pipetting liquids, parallax error), weighing errors, and titration errors (difference between equivalence point and endpoint). In terms of instrumental errors, a faulty or improperly calibrated pH meter would have led to the experiment being conducted at an incorrect pH level. Addition of too much EBT, NaOH and buffer will greatly affect
experimental results. All the mentioned errors might decrease or increase volume of titrant used.
REFERENCES [1] Harris, Daniel C. Quantitative Chemical Analysis, Seventh Edition. New York: W.H. Freeman and Company, 2007.
SUMMARY AND CONCLUSIONS The experiment has shown that commercial drinking water is relatively hard, with high concentrations of Ca2+ and Mg2+. A small standard pooled deviation means that precise data from different groups were obtained. Other ions, aside from Ca2+ and Mg2+, could also be present in the water sample, which may contribute to total hardness.
[2] Skoog, Douglas A, F. James Holler, and Donald M West. Fundamentals of Analytical Chemistry, Eighth Edition. New York: Saunders College Publishing, 1998. [3] Spurlock, Deborah. Determination of Water Hardness By Complexometric Titration Class Notes. Indiana University Southeast. March 14, 2011. http://homepages.ius.edu/DSPURLOC/c121/week13. htm (accessed August 20, 2011).
As a whole, the experiment was a success. The main objective, which was to relate complexometric titration in determining water hardness, was push through. Calculated values of total hardness were in relevance with the theoretical data, which states that drinking water is hard.
[4] Walpole, Ronald E. Introduction to Statistics, Third Edition. New York: Macmillan Publishers, 1968.
APPENDIX Working equations and sample calculations 𝑚𝑚𝑜𝑙 𝐶𝑎𝐶𝑂3 × % 𝑝𝑢𝑟𝑖𝑡𝑦 0.050 × 99.98 = = 0.00714 𝑀 100 100 𝑚𝐿 𝐸𝐷𝑇𝐴 7.0 𝑚𝑔 𝑇 = 𝑀𝐸𝐷𝑇𝐴 × 𝐹𝑊 𝐶𝑎𝐶𝑂3 = 0.00714 × 100.09 = 0.715 𝑚𝐿 𝑇 × 𝑚𝐿 𝐸𝐷𝑇𝐴 0.715 × 8.8 𝑝𝑝𝑚 𝐶𝑎𝐶𝑂3 = = = 125.84 𝑝𝑝𝑚 𝐶𝑎𝐶𝑂3 𝐿 𝑠𝑎𝑚𝑝𝑙𝑒 0.05 𝑠 22.17 𝑅𝑆𝐷 = × 1000 𝑝𝑝𝑡 = × 1000 = 152 𝑝𝑝𝑡 𝑋 145.26 𝑡𝑠 4.30 (22.17) 𝐶𝑜𝑛𝑓𝑖𝑑𝑒𝑛𝑐𝑒 𝑙𝑖𝑚𝑖𝑡𝑠 = 𝑋 ± = 145.26 ± = (90.44, 200.08) 𝑛 3 𝑀𝐸𝐷𝑇𝐴 =
𝑃𝑜𝑜𝑙𝑒𝑑 𝑠𝑡𝑎𝑛𝑑𝑎𝑟𝑑 𝑑𝑒𝑣𝑖𝑎𝑡𝑖𝑜𝑛𝑇𝑜𝑡𝑎𝑙
𝑎𝑟𝑑𝑛𝑒𝑠𝑠
=
𝑁
𝑁1 𝑖=1 (𝑥𝑖
3 – 𝑋1 )2 + … + 𝑘=1 (𝑥𝑘 – 𝑋3 )2 = 16.07 𝑁1 + 𝑁2 + 𝑁3 + … − 𝑁𝑠
Raw data
Nathan & Pia Standard EDTA H2O Analysis Elmer & Nikki Standard EDTA H2O Analysis
Trial 1
Trial 2
Trial 3
7.0 8.8
6.9 9.7
6.8 11.5
8.7 9.4
8.3 9.9
8.0 9.6
Marc & Janel Standard EDTA H2O Analysis Hector & Jess Standard EDTA H2O Analysis Lads & Felice Standard EDTA H2O Analysis Arah, Lorraine & Marjorie Standard EDTA H2O Analysis Reymar & Johanne Standard EDTA H2O Analysis Morph & Donne Standard EDTA H2O Analysis Morgan & Katya Standard EDTA H2O Analysis
5.8 9.3
5.5 9.2
5.6 9.5
5.2 9.7
5.3 9.9
5.3 9.5
7.2 9.5
6.8 9.3
7.4 10.0
6.8 10.4
6.4 10.6
9.9 10.9
4.5 9.1
4.7 9.3
4.4 9.0
8.8 9.9
8.5 9.0
8.6 8.8
6.5 10.0
6.9 10.3
6.5 10.1
