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

J. LINGAO1 and M. FRANCISCO2

1NATIONAL INSTITUTE OF MOLECULAR BIOLOGY AND BIOTECHNOLOGY, COLLEGE OF SCIENCE2NATIONAL INSTITUTE OF MOLECULAR BIOLOGY AND BIOTECHNOLOGY, COLLEGE OF SCIENCE UNIVERSITY OF THE PHILIPPINES, DILIMAN, QUEZON CITY 1101, PHILIPPINESDATE SUBMITTED:DATE PERFORMED:

ABSTRACT

Complexometric EDTA titration was used to determine the mineral content of Viva Mineralized Water. The EDTA solution of 0.050M was first standardized with 99,5% pure calcium carbonate. The environment of the reaction was alkaline, with pH 10 provided by the ammonia-ammonium buffer solution. The metallochrome Eriochrome Black T was used the indicator for this experiment. The same procedures were applied to the analysis of the 50mL aliquot sample. The color changes were expected to be from a wine red (color of EBT bound to a metal ion) to a clear blue solution (color of free EBT ion). The method yielded a hardness of 141.8045 ppm CaCO3 for the group, and 143.598 ppm for the collective data which differ by 26.338% and 25.406%, respectively, with the 192 ppm claim of Viva.

INTRODUCTION

One method for testing the quality of water is total hardness. Total hardness is defined as the sum of Ca2+ and Mg2+ hardness, in mg/L or commonly expressed as ppm CaCO3. This means thatregardless of the amount of the various componentsthat make up hardness, they can be related to aspecific amount of calcium carbonate. That is,hardness is expressed in mg/L or ppm of CaCO

3Water is a good solvent and dissolves minerals that come into contact with it. Calcium is dissolved in water as it passes over and through limestone (CaCO3) deposits. Magnesium is dissolved as water passes over and through dolomite (CaMg(CO3)2) and other magnesium bearing formations. Since it is impossible to separate the two cations when in solution, hardness is expressed as a combination of Ca2+ and Mg2+. Other soluble, divalent, metallic ions also contribute to water hardness; however, their levels are much less compared to Ca2+ and Mg2+ and are, thus, not included in measurements. Hardness can be expressed in grains per gallon, milligrams per liter (mg/L), or parts per million (ppm).

Water consumers are concerned about the hardness of the water they use. Hard water consumes more soap and detergents for washing, leaves scums on sinks, causes burnouts quickly, and creates scaling in industrial equipment.

Minerals dissolved in groundwater produce metal ions such as calcium, magnesium, manganese and iron [1] which are good for the body. Thus, mineralized water or water containing these ions, usually calcium and magnesium, was commercialized. Water containing such ions makes it “hard”. The total hardness of water is defined by the content of calcium and magnesium combined (calcium having the greatest amount present in water compared to the other minerals [2]) in milligrams per liter of water or commonly expressed as ppm CaCO3. Moreover, it is expressed as a combination of the two ions because there is no way to separate the two cations when both present in a solution. In this experiment, a technique called complexometric EDTA titration was used to determine the ppm CaCO3 of Viva Mineralized Water. Since drinking water is analyzed, it is expected that it contains more minerals, thus, hard water. Hard water is not as beneficial if used for a different purpose other than as drinking water. Calcium in hard water, when added with soap, breaks the surface tension in the bubbles so they cannot form, therefore, saponification is hindered; whereas, soft water cannot break the soap film due to lack of calcium carbonate.

In the experiment, total hardness in drinking water is quantitatively determined by using a technique called complexometric EDTA titration. Complexometric titration (sometimes chelatometry) 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.

EDTA, ethylenediaminetetraacetic acid, has four carboxyl groups and two amine groups that can act as electron pair donors, or Lewis bases. The ability of EDTA to potentially donate its six lone pairs of electrons for the formation of coordinate covalent bonds to metal cations makes EDTA a hexadentate ligand. However, in practice EDTA is usually only partially ionized, and thus forms fewer than six coordinate covalent bonds with metal cations. Disodium EDTA is commonly used to standardize aqueous solutions of transition metal cations. Disodium EDTA (often written as Na2H2Y) only forms four coordinate covalent bonds to metal cations at pH values ≤ 12. In this pH range, the amine groups remain protonated and thus unable to donate electrons to the formation of coordinate covalent bonds. Note that the shorthand form Na4-xHxY can be used to represent any species of EDTA, with x designating the number of acidic protons bonded to the EDTA molecule. EDTA forms an octahedral complex with most 2+ metal cations, M2+, in aqueous solution. The main reason that EDTA is used so extensively in the standardization of metal cation solutions is that the formation constant for most metal cation-EDTA complexes is very high, meaning that the equilibrium for the reaction:

lies far to the right. Carrying out the reaction in a basic buffer solution removes H+ as it is formed, which also favors the formation of the EDTA-metal cation complex reaction product. For most purposes it can be considered that the formation of the metal cationEDTA complex goes to completion, and this is chiefly why EDTA is used in titrations / standardizations of this type. it is almost always necessary to use a complexometric indicator to determine when the end point has been reached. Common indicators are organic dyes such as Fast Sulphon Black, Eriochrome Black T. Color change shows that the indicator has been displaced (usually by EDTA) from the metal cations in solution when the endpoint has been reached. Thus, the free indicator (rather than the metal complex) serves as the endpoint indicator.

This hardness test is often executed through complexometric EDTA titration. Complexation reaction involves reactions of complex ions or undissociated neutral molecules in a solution. The most important requirement is that, the complex is extremely soluble [3]. These complexes or coordination compounds are formed when metal ions react with electron-pair donors. These donor species, also called ligands, must have at least one pair of unshared electrons available for covalent bond formation [4]. A chelate is produced when a metal ion coordinates with two or more donor groups of a single ligand to form a five- or six-member heterocyclic ring [4]. In this experiment, EDTA is used as the titrant. It is a chelating agent; capturing the metal ions contained in the water, binding them to itself very tightly and therefore softening the water [2]. It can form several bonds to a metal ion, therefore increasing the stability of the reaction and making it proceed into completion. This property of EDTA makes it a multi-dentate or polydentate ligand. Multi-dentate ligand has many dents or claws or areas of attachment where a metal ion(s) to bind to it. Polydentate complexing agents are useful as titrants since they react in a 1:1 mole ratio, which produces only one endpoint. EDTA's chelating property is also used in eye drops and other ophthalmic solutions not only to prevent precipitation but also to prevent infection by gram-negative bacteria like Pseudomonas that have calcium and magnesium in their outer cell walls. This chelating ability of EDTA also increases the shelf life of perfumes [5]. In addition to that, it is also frequently used in soaps and detergents because it forms complexes with calcium and magnesium ions. These ions which are in hard water are bound to the EDTA and cannot interfere with the cleaning action of the soap or detergent. EDTA is also used in foods. Certain enzymes are responsible for food spoilage. EDTA is used to remove metal ions from these enzymes. It is used to promote color retention in dried bananas, beans, chick peas, canned clams, pecan pie filling, frozen potatoes and canned shrimp. It is used to improve flavor retention in canned carbonated beverages, beer, salad dressings, mayonnaise, margarine, and sauces. It inhibits rancidity in salad dressings, mayonnaise, sauces and salad spreads [2]. 

 Fig. 1 Structure of EDTA

The reactions in this experiment are:2H2Y2- + Ca2+ + Mg2+ ⟶ CaY2- + MgY2- + 4H+H2Y2- + MgIn- ⟶ MgY2- + HIn2- + H+ METHODOLOGY A 500mL 0.050 M stock EDTA solution was prepared by dissolving an appropriate amount of salt by distilled water. Then, MgCl2•6H2O was added to increase the solubility of the titrant due to diverse ion effect. But more importantly, because the CaIn- is not very stable the presence of Mg2+ in the solution ensures a sharp endpoint through the complexation of the EBT indicator with the magnesium ions present [6]. Then, if the crystals still do not dissolve because the dissolution of EDTA is very slow, NaOH may be added because it converts EDTA into a more soluble salt form and makes the solution basic therefore allowing the dissolution of EDTA which is at pH 8. 

Then, a 250 mL of 0.0050 M standard EDTA solution was prepared from 0.050M stock by diluting 25mL aliquot in a volumetric flask. In the preparation of the CaCO3 solution, 1.2484 grams of the pure CaCO3 was dissolved in a beaker. Then, a few drops of HCl were added slowly until all of the solids were dissolved. The addition of the acid should be slow so as to prevent the production of carbon dioxide and so, loss of material occurs. When dissolution was completed, the solution was then transferred to a volumetric flask and diluted to mark. The 50mL working standard of 0.0050M was then prepared from the stock solution by diluting 5mL of aliquot into 50mL in a volumetric flask. Next, 250mL of pH 10 ammonia-ammonium chloride buffer was made by mixing 142 mL of concentrated NH3 with 17.5 g NH4Cl. The pH was adjusted by the help of a pH meter or pH paper. The 50 mL of 0.1% (w/v) EBT in ethanol was prepared by dissolving 5g of EBT in 50mL of ethanol. In the standardization of the 0.01M EDTA solution, 10 mL of the 0.0050 M working standard CaCO3 solution was pipetted into a 250 mL E. flask then, 75 mL of distilled water was added. 3 mL of the buffer was added along with 15 drops of EBT. The solution was titrated from a wine red solution until a clear blue endpoint was obtained. This was done three times. In the analysis of the water sample, 50 mL of Viva Mineralized Water was pipetted into a 250ML E. flask then added with 3 mL of the buffer solution and 10 drops of EBT. The solution was again titrated from a wine red solution until a clear blue endpoint was obtained. RESULTS AND DISCUSSION As observed, the standard was allowed to react in a basic medium by the addition of the basic buffer of pH 10. A buffer was added so that the pH while the whole reaction occurs is constant. A constant pH is needed in the titration process since the EDTA and EBT have polyprotic properties, therefore unstable; and only a single endpoint is needed to be observed (EDTA can be protonated up to six while EBT is usually up to three[7]).  At a different pH, the increase in the concentration of hydronium ion may interfere with the complexation of the EDTA with the calcium and magnesium ions. Moreover, the effective or conditional stability constant of EDTA varies with pH because of its dependence on its degree of ionization, . In this case, the of the unprotonated form of EDTA in pH 10 is 0.35. It means that α αthere is a higher chance that the EDTA will be dissociated and produce its lower forms and bind or displace the ions in the reaction [3]. 

 Fig. 2 Values of 4 for EDTA at different pHα

However, if the pH is increased, the sharpness of the endpoint is also increased; therefore, the endpoint will be too sharp for a feasible titration. Also, the use of too much buffer would make the system very resistant to pH change, therefore, the endpoint will not be as sharp as predicted or worse, the endpoint will not be observed. 

Another reason for using an environment of pH 10 where the calcium and magnesium ions are being analyzed is that it satisfies one of the optimum conditions for complexometric titration of these two metal ions: the minimum solution pH. For calcium, the minimum solution pH for it to complexate with EDTA is 7.3 while for magnesium, it is 10. Therefore, the minimum solution pH of the environment where the reaction occurs is at pH 10[3]. At this pH, Ca-EDTA is stable and any magnesium present will not interfere with the reaction [8].  Fig. 3 Optimum conditions for complexometric titrations of selected elementsIn this experiment, EBT (eriochrome black T) was used as the indicator. Most complexometric titrations are performed with indicators which are themselves chelating agents and whose metal complexes have a different color from the reagents. Indicators having this special property are called metallochromic indicators. EBT is a tribasic weak acid. Since it also is polyprotic, the same principle on the ionization and dissociation of EDTA applies. 

At the start of the titration of EDTA to 10mL 0.0050M working standard CaCO3 with 3mL buffer and 15 drops of EBT, the solution is wine red. This is actually due to the complexation of the Ca2+ ions with the indicator. Ca2+ + HIn2- CaHIn (Kf = 2.5x105)As the reaction reaches the equivalence point, the calcium in the standard reacts with the EDTA added. The magnesium ions that were bound originally with the EDTA were then displaced by the calcium. Since the formation constant, Kf, of calcium-EBT is much lower than the Kf value of the Ca-EDTA complex, the Ca-EDTA formation dominates and the displaced magnesium ions complexate with the free indicator originally bound to the calcium. Moreover, the lower the formation constant, the less stable the complex is because there is greater energy needed for the complex to form, causing it to become unstable.Ca2+ + H2Y2- CaH2Y (Kf = 5.0x1010) Ca2+ + MgH2Y CaH2Y + Mg2+ Mg2+ + CaHIn MgHIn + Ca2+

Near the endpoint, assuming that all the calcium ions were consumed, the free magnesium ions present react with EDTA until the endpoint is reached where the magnesium ions bound with EBT were reacted with the EDTA, displacing the EBT and is signaled by the blue color of the solution. The blue color is due to the free EBT ions in the solution.Mg2+ + H2Y2- MgH2YMgHIn + H2Y2- MgH2Y + HIn2-

In the sample analysis of 50mL of Viva, 3mL of buffer and 10 drops of the indicator was added before the titration. As the indicator was added, the solution turns wine red. This is because of the magnesium ions present in the water sample react readily with the indicator added.The magnesium ions bind with the EBT and not the calcium ions since the Kf of Mg-EBT is higher than Ca-EBT.At the start of the titration, more Mg-EBT ions form as the Mg2+ from the Mg-EDTA reacts with the free EBT ions in the solution; therefore making the color of the solution turn into a darker shade of red.Mg2+ + HIn2- MgHIn (Kf = 1.0x107)Before the equivalence point, the EDTA added reacts and consumes the calcium in the water.Ca2+ + H2Y2- CaH2Y (Kf = 5.0x1010)Ca2+ + MgH2Y CaH2Y + Mg2+

As the endpoint is approached, all the calcium is consumed; then the present magnesium until the free magnesium ions reacted with EDTA.

Mg2++H2Y2-MgH2Y (Kf=4.9x108)

At the endpoint, the magnesium ions that were attached to EBT originally were reacted with EDTA to form Mg-EDTA complex and leave the EBT as a free ion and therefore gives off the blue color of the solution.

MgHIn + H2Y2- MgH2Y + HIn2-In the standardization process, the following volumes needed to titrate the calcium

carbonate standard were obtained:Table1. Volumes of titrant used in the standardization of 0.0050 M EDTA solution | Trial 1 | Trial 2 | Trial 3 | Vol. working std CaCO3, mL | 10 | 10 | 10 | Final vol. mL | 17 | 34.3 | 31.1 | Initial vol. mL | 0 | 17 | 14.1 | Net vol. mL | 17 | 17.3 | 17 |

These volumes were used to calculate the molarity and titer of the EDTA solution. The working equations are found in the Appendix.Table2. Molarity and titer of EDTA solution | Trial 1 | Trial 2 | Trial 3 | Molarity, M | 0.00292 | 0.00287 | 0.00292 | Titer, mg/L | 0.292272 | 0.287204 | 0.292272 |

Using these values, the total hardness of water, given by ppm CaCO3, and some statistical parameters were determined.Table3. Total Hardness and some statistical parameters | Trial 1 | Trial 2 | Trial 3 | Total Hardness, ppm | 141.8045 | 145.2915 | 141.8045 | Confidence Limits | 141.8045 + 0 | Range | 0 | Relative Standard Deviation | 0 |It can be observed that the relative standard deviation is zero. This is because 145. 2915, through Q-test, was eliminated since it was detected as an outlier in the data set.The results were then compared to the results obtained by the other groups. The following values were obtained:Table4. Collective data of the groups and calculated measures of central tendency Grp. No. | Trial 1 | Trial 2 | Trial 3 | Average | 1 | 146.1488 | 143.2832 | 138.6981 | 142. 71 | 2 | 144.085 | 142.9136 | 142.3279 | 143.1088 | 3 | 141.8045 | 145.2915 | 141.8045 | 141.8045 | 4 | 145. 807 | 149.73 | 147.7685 | 147.7685 | 5 | 148.1345 | 146.2672 | 149.3793 | 147.927 | 6&9 | 143.2135 | 144.3826 | 144.3826 | 144.3826 | 7 | 140.0737 | 141.8173 | 141.8045 | 142.9669 | 8 | 140.234 | 138.5238 | 140.234 | 140.234 | Mean | | | | 143.598 | StDev | | | | 1.447971 | RSD | | | | 10.0835 |

Note that the values in bold formatting are detected outliers and therefore were eliminated in the calculation of average, mean, standard deviation and relative standard deviation.As observed, the calculated values of ppm CaCO3 per group indicate that the total hardness of Viva Mineralized Water is approximately 144 ppm. The relative standard deviation (10) means that the experiment is very precise. And the calculated average ppm of the group (141.8045) is only 1.25% far from the collective average (143.598).On the other hand, the calculated result of the group is 26.338% far from the claimed 192 ppm of Viva while the collective average differs from the claimed ppm by 25.406%.CONCLUSION

The total hardness determined by the group and the obtained collective results of the whole class is lower from the claimed hardness of Viva. The relative standard deviation is outstandingly low and therefore signifies that the experimentation process is precisely done. Thus, the results obtained in this experiment are considered valid. As a conclusion, it can be said that Viva is hard (120-180ppm) as opposed to its claim of being very hard (180ppm and above).

REFERENCES[1] Waterfilters.net. “Understanding Hard Water and Water Softening”. <February 8, 3013>[2] Determination of Water Hardness By Complexometric Titration Class Notes.html <February 8, 2013>[3] Khopkar, S. M. Basic Concepts of Analytical Chemistry. New Age Internationals, 2004; pp. 63-76[4] Skoog, D. A.; West, D. M.; Holler, F. J.; Crouch, S. R. Fundamentals of Analytical Chemistry, 8th ed.; : David Harris, 2004; pp. 563-565[5] J. Roger Hart; J. Chem. Educ., 1984, 61 (12), p 1060.[6] Institute of Chemistry, University of the Philippines. General Chemistry II: Laboratory Manual. June 2011.[7] Sinex, Scott A. EDTA - A Molecule with a Complex Story. 1 August 2007. 25 January 2012 <http://www.chm.bris.ac.uk/motm/edta/edtah.htm>.[8] Dr. Husain, A. Theoretical Basis of Analysis: Complexometric Titrations.

1. What is water hardness and why is it expressed as ppm CaCO3?Hard water is a mixture of calcium and magnesium, together with bicarbonate, sulphate, chloride, etc.Water hardness is an aesthetic quality of water, and is caused mostly by the minerals calcium and magnesium, but is classified or measured based on the level of concentration of calcium carbonate. How hardness is classified is based on the following scale:

Table 1. The water hardness scaleWater Hardness ppm CaCO3Soft 0 – 20Moderately soft 20 – 60Moderately hard 61 – 120Hard 121 – 180Very hard > 180

A hard water with 300 mg/l as CaCO3 would precipitate 300 milligrams of calcium carbonate per litre.

The hardness is not entirely due to calcium – magnesium is usually present to some extent and other multivalent cations. When the hardness is expressed as CaCO3, it is calculated as if the magnesium, etc. were there as calcium. It also includes the bicarbonate ions such as chloride, sulphate and nitrate.

The hardness of water is expressed in terms of ppm ( equalent of CaCO3) because the molicular weight of calcium carbonate is 100. it is easy to calculate. this is one of the main reason for expressing the hardness of water in ppm.also calcium carbonate is insoluble in water there four it is easy to calculate its amount in water.

2. How does complexometric titration work?IUPAC has defined'complexometric titration' as a titration based onreaction of a metal ion with a ligand to form a solublecomplex and in which one of the two reactants isused a titrant.[3

3. Why was EDTA used as complexing agent/-titrant?The ligands act as a "complexing agent" for the metal cation. Multidentate ligands, also called chelating agents, provide more than one pair of electrons for complexation, and therefore can form mutliple coordinate covalent bonds between the ligand and the metal cation. The resulting chelates are extremely stable; having very large formation constants (Kf).EDTA is the most important chelating agent in analytical chemistry. The tetrabasic form of this acid forms complexes with virtually all metal ions. In this form EDTA is a hexadentate ligand; each of the acid oxygens and each of the amine nitrogens can donate one electron pair to the metal. The metal ion is usually held in a one-to-one complex with the EDTA molecule, 6 EDTA to metal coordinate covalent bonds having formed.

4. Why were MgCl2 · 6H2O crystals and NaOH pellets added in the preparation of EDTA solution?

MgCl2•6H2O was added to increase the solubility of the titrant due to diverse ion effect. But more importantly, because the CaIn- is not very stable the presence of Mg2+ in the solution ensures a sharp endpoint through the complexation of the EBT indicator with the magnesium ions present [6]. Then, if the crystals still do not dissolve because the dissolution of EDTA is very slow, NaOH may be added because it converts EDTA into a more soluble salt form and makes the solution basic therefore allowing the dissolution of EDTA which is at pH 8.

pon preparation of stock EDTA solution, MgCl26H2Ocrystals were added to the dissolved salt EDTA. Thefact that Mg-EDTA has a higher formation constanthence a higher tendency to form complexes than Ca-EDTA, Mg2+ from EDTA solution can easily displaceCa2+ions and form Mg-EDTA complexSince Ca-EDTA has a lower Kf , formation constant,Ca-EDTA is less stable because there is a greaterenergy for the complex to form, so adding more Mg2+ ions will make the endpoint sharper. On the otherhand, EDTA is essentially insoluble in water, and willonly dissolve when pH is neutralized to 8. Addition ofbase, in this experiment NaOH pellets, facilitatesdissolution of acid form of EDTA

5. Why was HCl added in the preparation of CaCO3 solution?Carbonate error can cause discrepancy in pH readingso adding HCl while dissolving CaCO3 during thesolution preparation is important for all reactionsbetween metal ions and EDTA are pH dependent,and for divalent ions, solutions must be kept basic(and buffered) for the reaction to go to completion [6].Most ligands are basic and bind to H+ ions throughouta wide range of pH. Some of these H+ ions arefrequently displaced from the ligands (chelatingagents) by the metal during chelate formation,sobuffer was used to hold the pH constant. In thisexperiment, 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 [7]

6. What are the pertinent chemical equations involved during titration? Explain how the color of the solution turns to clear blue using the stability of complexes involved in the titration.

Viva! Mineralized water was titrated with thestandardized EDTA. The starting color sample uponaddition of EBT indicator was wine red due to theexistence of Mg2+ ions (3), and upon titrating withEDTA, the color gradually changes from wine red toclear blue, indication that the endpoint is reached,which occurs when Ca2+ ions complexes with EDTA(1), the same time the Mg2+ complexes with it (2)sequentially and the Mg-EBT complex breaks asillustrated in the equation

In the sample analysis of 50mL of Viva, 3mL of buffer and 10 drops of the indicator was added before the titration. As the indicator was added, the solution turns wine red. This is because of the magnesium ions present in the water sample react readily with the indicator added.

The magnesium ions bind with the EBT and not the calcium ions since the K f of Mg-EBT is higher than Ca-EBT.

At the start of the titration, more Mg-EBT ions form as the Mg2+ from the Mg-EDTA reacts with the free EBT ions in the solution; therefore making the color of the solution turn into a darker shade of red.

Mg2+ + HIn2- à MgHIn (Kf = 1.0x107)

Before the equivalence point, the EDTA added reacts and consumes the calcium in the water.

Ca2+ + H2Y2- à CaH2Y (Kf = 5.0x1010)Ca2+ + MgH2Y à CaH2Y + Mg2+

As the endpoint is approached, all the calcium is consumed; then the present magnesium until the free magnesium ions reacted with EDTA.

Mg2++H2Y2-àMgH2Y (Kf=4.9x108)

At the endpoint, the magnesium ions that were attached to EBT originally were reacted with EDTA to form Mg-EDTA complex and leave the EBT as a free ion and therefore gives off the blue

color of the solution. 7. What is the importance of maintaining the pH at 10 and choosing NH3-NH4Cl as the buffer?A buffer was added so that the pH while the whole reaction occurs is constant. A constant pH is needed in the titration process since the EDTA and EBT have polyprotic properties, therefore unstable; and only a single endpoint is needed to be observed (EDTA can be protonated up to six while EBT is usually up to three[7]).  At a different pH, the increase in the concentration of hydronium ion may interfere with the complexation of the EDTA with the calcium and magnesium ions. Moreover, the effective or conditional stability constant of EDTA varies with pH because of its dependence on its degree of ionization, α. In this case, the α of the unprotonated form of EDTA in pH 10 is 0.35. It means that there is a higher chance that the EDTA will be dissociated and produce its lower forms and bind or displace the ions in the reaction [3].

However, if the pH is increased, the sharpness of the endpoint is also increased; therefore, the endpoint will be too sharp for a feasible titration. Also, the use of too much buffer would make the system very resistant to pH change, therefore, the endpoint will not be as sharp as predicted or worse, the endpoint will not be observed. 

Another reason for using an environment of pH 10 where the calcium and magnesium ions are being analyzed is that it satisfies one of the optimum conditions for complexometric titration of these two metal ions: the minimum solution pH. For calcium, the minimum solution pH for it to complexate with EDTA is 7.3 while for magnesium, it is 10. Therefore, the minimum solution pH of the environment where the reaction occurs is at pH 10[3]. At this pH, Ca-EDTA is stable and any magnesium present will not interfere with the reaction [8].

8. What are the possible sources of errors and their effect on the calculated parameters? Rationalize.