QUANTITATIVE DETERMINATION OF TOTAL HARDNESS IN DRINKING WATER BY COMPLEXOMETRIC EDTA TITRATIONJ. OLIVEROS1

1 INSTITUE OF BIOLOGY, COLLEGE OF SCIENCEUNIVERSITY OF THE PHILIPPINES, DILIMAN, QUEZON CITY 1101, PHILIPPINESDATE SUBMITTED: 27 FEBRUARY 2015DATE PERFORMED: FEBRUARY 2015

ABSTRACTThe purpose of the study was to determine the total hardness of a specific brand of drinking water using the concepts learned about the complexometric EDTA titration. The hardness of water is defined in terms of its content of calcium and magnesium ions, and by the complexometric EDTA titration, which is the process of metal ion titration until an end product of complexes is obtained, hardness can be determined. There total hardness of water was identified to categorize the use a specific type of water —whether it is used domestically or industrially. The determination of water hardness is also a useful test that provides a measure of quality of water for households, particularly for water potability. The working standard CaCO3 solution was added with buffer solution and Eirochrome Black T (EBT) indicator for the standardization of the EDTA titrant solution. The analysis of water sample involved the titration of a measured volume of mineralized water of the brand Hidden Spring (plus buffer solution and EBT indicator) with the EDTA solution. A color change from wine red to clear blue indicates the titration endpoint. The computed total hardness was 248.7 ppm (average of the three trials), which is classified as very hard water. This proves that Hidden Spring is a reliable source of mineral water. However, possible errors during the experiment may have caused by the intrinsic error of the method and typical random errors including human errors.

INTRODUCTION

Water hardness is important to fish culture and is a commonly reported aspect of water quality.  It is a measure of the quantity of divalent ions such as calcium, magnesium and/or iron in water.  There are many different divalent salts; however, calcium and magnesium are the most common sources of water hardness [1].

Hardness is traditionally measured by chemical titration.  The hardness of a water sample is reported in milligrams per liter (same as parts per million, ppm) as calcium carbonate (mg/l CaCO3).  Calcium carbonate hardness is a general term that indicates the total quantity of divalent salts present and does not specifically identify whether calcium, magnesium and/or some other divalent salt is causing water hardness. 

Hardness can be a mixture of divalent salts.  In theory, it is possible to have water with high hardness that contains no calcium [1].

The determination of water hardness is a useful test that provides a measure of quality of water for households and industrial uses. Originally, water hardness was defined as the measure of the capacity of the water to precipitate soap. Hard water is not a health hazard. People regularly take calcium supplements. Drinking hard water contributes a small amount of calcium and magnesium toward the total human dietary needs of calcium and magnesium. The National Academy of Science states that consuming extremely hard water could be a major contributor of calcium and magnesium to the diet [2].

Table 1. Water Hardness Scale

Water Hardness ppm CaCO3

Soft 0 – 20

Moderately Soft 20 – 60

Moderately Hard 61 – 120

Hard 121 – 180

Very Hard > 180

The ions involved in water hardness, i.e. Ca2+(aq) and Mg2+(aq), can be determined by titration with a chelating agent, ethylenediaminetetraacetic acid (EDTA), usually in the form of disodium salt (H2Y2-). The titration reaction is:

Eriochrome Black T is commonly used as indicator for the above titration. At pH 10, Ca2+(aq) ion first complexes with the indicator as CaIn+(aq), which is wine red. As the stronger ligand EDTA is added, the CaIn+(aq) complex is replaced by the CaY2-(aq) complex which is blue. The end point of titration is indicated by a sharp color change from wine red to blue.

Titration using Eriochrome Black T as indicator determines total hardness due to Ca2+(aq) and Mg2+(aq) ions. Hardness due to Ca2+(aq) ion is determined by a separate titration at a higher pH, by adding NaOH solution to precipitate Mg(OH)2(s), using hydroxynaphthol blue as indicator [3].

EDTA also forms a 1:1 molar ratio with metal ions in a complex. This results in increased stability for the metal-ion complex formed. For this reason, EDTA is primarily used as the titrant when performing complexometric titrations [8]. Figure 1 below shows the chemical structure of EDTA.

Figure 1. Chemical Structure of EDTA [9]

METHODOLOGY

The experiment started by preparing all the needed solutions for the study. The solutions prepared were 500 mL 0.1000 M stock Ethylenediaminetetraacetic acid (EDTA) solution, 100 mL 0.0500 M stock Ca2+ solution, and a 250 mL 1.0 M NH3-NH4

+ pH 10 buffer solution. To avoid confusion, the aforementioned solutions will be named solutions A, B, and C respectively.

Solution A was prepared by adding approximately 18.61 grams of Na2H2EDTA * H2O to the nearest 0.1 milligrams in a 400 mL beaker and adding 200 mL of distilled water and 1.0 gram MgCL2 * H2O upon the dissolution of the first solid added. Then the entire solution was quantitatively transferred in a 500 mL volumetric flask, filling to mark with distilled water and mixing.

Next, solution B was prepared by adding approximately 0.791 grams of pure CaCO3 (with a purity of 99.97%) to a 250 mL beaker and adding 40 mL distilled water.

HCl was then added to the solution as it was being heated and at the same time gently stirred.

This was done in order to dissolve the CaCO3

solids. Upon dissolution, the entire solution was transferred to a 100 mL volumetric flask, diluted to mark with distilled water and mixed.

Finally, the solution C was prepared by mixing the specified values for NH3 and NH4Cl (in this case, 14.072 mL of NH3 and 2.06 grams of NH4Cl into a volumetric flask, diluted to the mark and mixed. The pH was tested using a strip of pH paper and adjusted by either adding more acid or base to attain a pH value of 10.

Additional solutions were also prepared; 250.0 ml .0100 M working EDTA solution from .1 M EDTA and 50.0 ml .0050 M working standard Ca2+ from the .0500 M Ca2+ solution. These solutions will be labeled as Solutions A1 and B1

respectively. The solution A1 will be prepared by getting 25 ml from the .1000 M EDTA and diluting and storing it to a specified 250 ml volumetric flask. The solution B1 was prepared by getting 5 ml of the .0500 M Ca2+ solution diluting and storing it to a specified 50 ml volumetric flask.

Standardization of the .01 M EDTA solution started by getting three sets of 10 mL aliquots from the .0050M CaCO3 solution. 75 mL of distilled water was also added to each of the flasks. Three mL of the buffer solution was then added to the solution with 2-3 drops of the EBT indicator before titration.

The volumes (initial and final) of the titrant were then recorded. The solution was titrated until a change in color from wine red to a clear blue was observed. This part of the experiment was done within 3 trials.

After doing the standardization, a similar process was done for the analysis of the Hidden Spring water sample. However, instead of 10 mL, 3 sets of 50 mL water samples were added into 3 separate 250 mL Erlenmeyer flasks before

adding 3 mL of the buffer solution then titrating with the EDTA solution using the EBT as indicator. The color change observed for both was from wine red to dark blue.

RESULTS AND DISCUSSION

During the preparation of stock EDTA solution, MgCl2 * 6H2O crystals were added to the dissolved salt EDTA. It was added to the solution to increase its solubility via diverse ion effect.

The fact that Mg-EDTA has a higher formation constant hence a higher tendency to form complexes than Ca-EDTA, Mg2+ from EDTA solution can easily displace Ca2+ ions and form Mg-EDTA complex

Since Ca-EDTA has a lower Kf, formation constant, Ca-EDTA is less stable because there is a greater energy for the complex to form, so adding more Mg2+ ions will make the endpoint sharper. On the other hand, EDTA is essentially insoluble in water, and will only dissolve when pH is neutralized to 8. Addition of base, in this experiment NaOH pellets, facilitates dissolution of acid form of EDTA.

Carbonate errors can cause discrepancies in pH readings so adding HCl while dissolving CaCO3

during the solution preparation is important for all reactions between metal ions and EDTA that are pH dependent, and for divalent ions, solutions must be kept basic (and buffered) for the reaction to go to completion [5]. Most ligands are basic and bind to H+ ions throughout a wide range of pH. Some of these H+ ions are frequently displaced from the ligands (chelating agents) by the metal during chelate formation, so a buffer was used to hold the pH constant. In this experiment, NH3-NH4Cl buffer was used since EBT indicator only works when the pH is at 8 to 10, whereas the buffer has a pH equal to 10 [4].

2 4 2 10Ca + EDTA CaEDTA 5.0 10fK x

2 4 2 8Mg + EDTA MgEDTA 4.9 10fK x

Meanwhile NaOH was present to neutralize the carboxyl groups present in the EDTA crystals and as such increases its solubility, therefore ensuring its complete dissociation into the solution [5]

Equations (1), (2), and (3) below display the reaction process the various reactants and ions undergo in order to display a color change from wine red to dark blue as well as to display a sharper endpoint for the titration process.

H2Y2– + Ca2+ → CaY2– + 2H+ (1)H2Y2– + Mg2+ → MgY2– + 2H+ (2)

H2Y2– + Mg-In → MgY2– + HIn2–+ H+ (3)where Mg-In = wine red solutionand HIn2–+ = dark blue endpoint

Note how EDTA reacts to both calcium (equation (1)) and magnesium ions (equation (2)) before the magnesium ions bind to the EBT indicator (prior to titration) then forming a metal-ion complex post-titration (equation (3)). Figure 2 below displays the structure of a metal cation-EDTA complex.

Figure 2. Chemical Structure of a Metal (M)-EDTA Complexation [7]

The concentration of the two most common hard water ions, Ca2+ and Mg2+ was then determined by titrating the water sample with a known chelating agent, EDTA. It was standardized with CaCO3 thrice and the resulting average concentration of EDTA is 0.0095 M.

Table 2. Standardization of EDTA solution

Trial 1 2 3

Volume of workingCa(II) standard, mL

10 10 10

Net volume EDTA, mL

8 8.5 8.3

M EDTA .0098 .0093 .0095Average M EDTA .0095

Titer .9566187928

The net volume of EDTA used was obtained experimentally through the titration process while the MEDTA as well as the titer were both calculated from equations (4), (5), (6) and (7).

M stock Ca2+¿=

gCaCO3 ( 100.09g /mol ) (%purity )Lsolution

¿ (4)

Mworking standard=(M stock Ca) (V stock Ca)V working standard

(5)

MEDTA=(Mwor k ingstandard ) (V working standard )

V EDTA (6)

T=M (FW ) (7)

During standardization, purple color appeared upon addition of EBT indicator compared to the expected wine red color of the solution. This may happen since there are no Mg2+ ions in the solution, and the indicator only exists as wine-red complex due to the magnesium ions [6].

On the other hand, the water sample (Hidden Spring) was titrated with the standardized EDTA. The starting color sample upon addition of EBT indicator was wine red due to the existence of Mg2+ ions and upon titrating with EDTA, the color gradually changes from wine red to clear blue, indication that the endpoint is reached, which occurs when Ca2+ ions complexes with EDTA, the same time the Mg2+ complexes with it sequentially and the Mg-EBT complex breaks as illustrated in the equation.

2+ 2-Mg + HIn MgIn +H

- 4- + 2- 2-3 2MgIn +Y +H O MgY +HIn +H O

( - ) ( )red wine clear blue

Table 3. Hidden Spring Water Sample Analysis

Trial 1 2 3Volume of

water sample, mL

50 50 50

Net Volume of EDTA, mL

13 13 13

Total Hardness, ppm CaCO3

248.7 248.7 248.7

Average 248.7RSD 0 ppt

Confidence Interval

248.7 +/- 0

The values above were calculated using the formulas listed below in equations (8) and (9). While the formulas involving the RSD and confidence interval are listed down as equations (10) and (11).

MCa2+¿=

MEDTRA (V EDTA)Vwater sample

¿ (8)

ppmCaCO3=MCa2 +¿( 1molCaCO3

1molCa )( 100.09gmol

CaCO 3)(1000mg

1g)¿

(9)

RSD= sxmean

(1000 ppt) (10)

CI=mean±t(sx )√n

(11)

where t is a given value dependent on n-1 at 95% confidence level and n is the number of given data

When comparing the obtained values above with the theoretical values obtained from the labels indicated in the water bottle, It can be concluded that errors were still present in the experiment, however, the objective of the experiment was still achieved, that is to determine the total hardness of a water sample.

Theoretically, the water sample contained 25 ppm and 35 ppm of both Ca2+ and Mg2+ ions

respectively; and the calculated water hardness (CaCO3 ppm) based on these theoretical values are 62.44 ppm (from ppm Ca) and 144.10 ppm (from ppm Mg) for a total hardness of 206.54 ppm. Experimentally, the total hardness of water was 248.70 ppm, which according to Table 1, falls under the vey hard category. This fully supports the theoretical category of the Hidden Spring water sample to fall to the very hard category.

As mentioned, certain errors could have been committed while performing the experiment that may have affected the overall result, primarily the ppm value of calcium carbonate. Such errors committed that could have affected the ppm value of CaCO3 are overtitration (increase in VEDTA = increase in ppm CaCO3), the negligence to allow for EDTA to pass through the burette pre-titration (increase in ppm CaCO3 due to a most likely diluted EDTA titrant), too much EBT indicator added (indeterminate ppm CaCO3 due to a vague shift in color), too little EBT indicator added (increase ppm CaCO3 as the time for a color shift will be delayed), actual pH of the solution is less than 10 (decrease ppm CaCO3

due to the possible precipitation of metal hydroxides.) The addition of indicator before the 3 mL buffer solution to the analyte may also affect the overall process as some of the indicator may react with the solution before the buffer can work for an ideal set-up.

SUMMARY AND CONCLUSIONS

Complexometric titration was utilized to determine the hardness of the 50 mL Hidden Spring water sample. Water hardness is basically a set of standards used to determine whether or not a water sample is suitable for domestic or industrial use. Generally, very hard water is mostly suitable for drinking and not for industrial use for hard water can; clog pipes due to formation of precipitates called “scales /mineral deposits” when heated,

reduce the life of clothes due to the slimy film/yellow coating the “scales” form in clothes, and it can further deteriorate the performance of appliances due to excess mineral deposits. .

EDTA was used as the titrant because of its efficiency and stability in forming metal-ion complexes, specifically with the Mg2+ and Ca2+

ions found in the water sample. Because of various disassociations and formations among the reactants and ions involved as expressed in equations (1),(2), and (3), a color change from wine red to dark blue may be observed to determine the MEDTA as well as the ppm CaCO3. Various errors such as overtitration, addition of indicator before buffer (for the analyte), addition of too much or too little indicator, have influenced the accuracy of the data.

Lastly, after determining the total hardness of the water sample, and after comparing the results from the theoretical values, it can be concluded that the Hidden Spring water sample is classified with a total hardness that falls to the category of “very hard” which signifies its high-mineral content and potability as a type of drinking water.

REFERENCES

[1] Wurts, William. "WATER HARDNESS -- CALCIUM &  MAGNESIUM." WATER HARDNESS -- CALCIUM &  MAGNESIUM. Kentucky State University CEP at UK Research and Education Center, 1 Jan. 2008. Web. 1 Apr. 2015.

[2] Spurlock, D. "Determination of Water Hardness By Complexometric

Titration Class Notes." Determination of Water Hardness By Complexometric Titration Class Notes. Indian University Southeast, 31 Mar. 2014. Web. 1 Apr. 2015.

[3] "Hardness of Water by EDTA Titration." Edb.kedcity.net. Web. 1 Apr. 2015. <http://cd1.edb.hkedcity.net/cd/science/chemistry/s67chem/pdf/sOL_6_Water_hardness.pdf>.

[4] Andersen, Robert A. Algal Culturing Techniques. Burlington, Mass.: Elsevier/Academic, 2005. 54-55. Print.

5] Beynon, R.J., and J.S. Easterby. "Determination of Water Total Hardness by Complexometric Titration." Titrations.info. 6 Sept. 2009. Web. 1 Apr. 2015.

[6] KIRUTHIGA, B. "Complexometric Titration." Srmuniv.ac.in. DEPT OF PHARMACEUTICAL CHEMISTRY SRMUNIV. Web. 1 Apr. 2015. <http://www.srmuniv.ac.in/downloads/Complexometric_Titration.pdf>.

[7] Metal-edta.png. Digital image. Http://figures.boundless.com/12142/full/metal-edta.png. Web.

[8] "Buffers & Henderson-Hasselbal." Mhchem.org. Web. 1 Apr. 2015.

[9] Graphics 9. Digital image. Http://cnx.org/resources/0418e3a849c5083ecc35ba9e8c9249b4/graphics9.jpg. Web.

APPENDIX A - TABLES

Table 1. The water hardness scale.

Table 2. Standardization of EDTA solution [Complete version]

Table 3. Hidden Spring Water Sample Analysis

APPENDIX B – CALCULATIONS

Calculations

A. Standardization of EDTA Solution

AVERAGEM=¿¿

B. Analysis of Viva! Water Sample

Trial 1 2 3

Working Ca(II) standard

10 mL 10 mL 10 mL

Net Vol of EDTA 8 mL 8.5 mL 8.3 mL

M of EDTA