Complexometric titration

Complexometric titration

1. Complexometric Titration Metal ion (M+) + EDTA (HY) Metal- EDTA complex (MY) + H + M +n + H 2 Y -2 MY n-4 + 2H + Na 4-x H x Y can be used to represent any species of EDTA, with x designating the number of acidic protons bonded to the EDTA molecule. Due to low solubility of EDTA in water, its disodium dihydrate EDTA salt i.e. Na 2 H 2 Y.2H 2 O is used. Complex is formed by the reaction of metal ion (M n+ ) with either an anion e.g.[Ag(CN) 2 ] - or neutral molecule e.g.[Ag(NH 3 ) 2 ] + The metal ion is known as Central metal atom. The anion or neutral molecule is known as Ligand (L)

2. It is a volumetric analysis in which formation of a colored complex is used to indicate the end-point of the titration. It is used for the determination of a mixture of different metal ions present in the solution. Complexometric titration\ Chelatometry is based on essentially Lewis acid-base reactions in which electron pair is donated from one chemical to another. Disodium EDTA (Na 2 H 2 Y ⋅ 2H 2 O: better soluble in water) is most common chelating agent used for Complexometric titration because it creates water soluble stable metal complex in alkaline medium. (except the  alkali metals  which form too weak complexes) All the complexes have exact 1:1  stoichiometry   (regardless of the charge of the cation ). They react stoichiometrically and can be used to quantitatively determine the metal ions in the sample by titration. The titration needs to be carried out at alkaline medium to consume H+ which produced during reaction, so that the reaction goes to the right and stability of the complex formed (MY) is increased. The  versatility , sensitivity, and general convenience of  complexometric titrations are dependent on the correct choice of indicators for endpoint detection. In deprotonated form of EDTA , six binding sites —four negatively charged carboxylate groups and two tertiary amino groups—that can donate six pairs of electrons to a metal ion to form coordinate covalent bond. [EDTA = Hexadentate ligand ] Some terms Ligand : Compound having at least one pair of unshared electrons available for bond formation Chelate : It is a complex formed between the ligand containing two or more donor groups and metal to form ring structure. (heterocyclic rings or chelate rings). Chelates are usually insoluble in water but soluble in organic solvent. Chelating agents: Organic molecules/ ligands containing two or more donor groups which combine with metal to form complex having ring structure. Most chelating agents consist of N or O. Sequestering agent: Ligands which form water soluble chelates e.g. EDTA.

3.   The resulting metal–ligand complex, in which EDTA forms a cage-like structure around the metal ion, is very stable. Central metal atom= acts as Lewis acid (electron acceptor) Ligand= acts as Lewis base (electron donor) Coordinate bond (dative)= The bond formed between central metal atom (ion) (acceptor) and the Ligand (donor) Dative bond is similar to covalent bond(formed of two electrons) Buti n dative bond the electrons pair are donated from one atom to the other. The atom gives electron pair is known as donor, while the atom accept electron pair is known as acceptor. Coordination number =The no. of coordinate bonds formed to a metal ion by its ligands

4. In general, the fully ionized form of EDTA will form a 1:1 complex with all metal ions in solution. Thus, we have : Ligands are any combination of anions that can donate an electron pair to form coordinate covalent bonds such as EDTA. Some examples of polyvalent metal ions are given below : Bivalent Metal ions : Ca 2+ , Mg 2+ , Zn 2+ , Trivalent Metal ions : Fe 3+ , Al 3+ , Cr 3+ , Tetravalent Metal ions : Sn 4+ ,Ce 4+ , Cr 4+ , Pt 4+ . There are two main classes of ligands : Unidentate ligands : Ligands that are bound to a metal ion at one place are called as unidentate ligands . Example: NH 3 Bidentate or multidentate ligands : These ligands contain more than one group which is capable of binding with the metal ions. Example: ethylene diamine , EDTA

5. CHELATING AGENTS A chelating agent is a chemical compound that can bind with a metal ion and prevent that metal from undergoing other chemical reactions. These chelating agents are composed of atoms with lone pairs. These lone pairs can be donated to a metal ion. The donation of one lone electron pair to the metal atom forms a  coordinate covalent bond . The number of coordinate covalent bonds present in a coordination complex is called the coordination number. EDTA is a common example of a chelating agent. It is a multidentate ligand . This means it can bind with the metal ion via several atoms by forming coordinate covalent bonds. There are some other chelating agents that are Bidentate . They form only two coordinate covalent bonds. Chelating agents are very useful in isolating heavy metals to remove them from drinking water , to deactivate metal ions that can cause precipitation, to limit available metal ion content, etc.

6. Effect of pH on complex formation: During the formation of stable, 1:1 water soluble, colorless complex there will be the release of two H + ions, as the number (n) of M-EDTA complex increases in the solution, there will be the removal of 2n[H + ] ions, acidity increases in the solution. So we need to maintain the p H  of the solution with suitable basic buffer in order to have a M-EDTA complex. Hence complexometric titrations are generally carried out in basic p H . Therefore, the stability of the metal complex is pH dependent . The lower the pH of the solution the more hydrogen ions are available to compete with the metal ion for the ligand, and the equilibrium shifts to the left , causing a decrease in the stability of the complex . This equation illustrates the competition between the metal ion and the hydrogen ion for the ligand: M 2+ + H 2 Y 2- = MY 2- + 2H +

7. STABILITY OF COMPLEXES Generally, the formation of a 1 : 1 chelate complex (MY n − 4 )may be designated by the following equation : where, M = Metal ion, and X = Chelating ion. Hence, the stability constant, K f , may be expressed as : K f =[MY n-4 ]/[M n+ ][Y 4- ] The equilibrium constant for the reaction of a metal with a ligand is called the formation constant, K f , or the stability constant. The formation constant K f = [MY n − 4 ] /[M n+ ][Y 4− ] describes the reaction between Y 4− and a metal ion. Generally, complex ions with polydentate ligands have much higher formation constants than those with monodentate ligands . This additional stability is known as the  chelation effect

8. Stability of Metal-Ligand Complexes The stability of complexes is influenced by a number of factors related to the ligand and metal ions. Nature of the metal ion: Small ions with high charges lead to stronger complexes. Nature of the ligand: The ligands forming chelates impart extra stability ( chelon effect). For example the complex of nickel with a multidentate ligand is more stable than the one formed with ammonia. Basicity of the ligand: Greater basicity of the ligand results in greater stability of the complex.

9. 4. Size of chelate ring: The formation of five- or six - membered rings provides the maximum stability. 5. Number of metal chelate rings: The stability of the complex is directly related to the number of chelate rings formed between the ligand and metal ion. Greater the number of such rings, greater is the stability. 6. Steric effects: These also play an important role in the stability of the complexes.

10. Requirement for Metal ion indicator The colour must be sufficiently intense, so that a minimum amount of indicator can be used. The colour contrast between the indicator and Metal-indicator complex should be readily observable. The Metal-indicator complex should possess sufficient stability to ensure a sharp colour change, however it should be less stable than Metal-EDTA complex. The change in equilibrium from metal-indicator complex to the Metal-EDTA complex should be sharp and rapid.

11. 5. The colour reaction of the indicator should be selective. 6. The indicator must be very sensitive to metal ions so that the colour change occurs at near the equivalence point. 7. The indicator must be stable in the titration medium. 8. The indicator must be stable on storage also. 9. All the above requirements must be fulfilled in the pH range in which the proposed titration is to carried out . 10 . It should be commercially available in adequate purity.

12. Types of complexometric titration: Direct titration Back titration Replacement titration Alkalimetric titration Indirect titration 1. Direct titration: It is the simplest and most convenient method in which the standard solution of EDTA is slowly added to the metal ion solution till the end point is reached. The solution containing metal ion is buffered to a desired pH and titrated directly with the standard EDTA solution. A blank titration may be performed by omitting the sample to check the presence of impurities in reagent. Ca 2+ , Mg 2+ , Zn 2+ ions are determined by this method. Metal ion + EDTA metal- EDTA complex

13. 2. Back titration A measured amount of EDTA is added in excess to react with the metal sample to be examined. The resulting solution will contain unreacted EDTA which is then back titrated with another metal ion standard solution in the presence of indicator. ZnCl 2 , ZnSO 4 , MgCl 2 , MgSO 4 is used as standard metal ion solution. Al +3 , Co +2 , Pb +2 , Mn +2 , Hg +2 , and Ni +2 can be determine by using Back titration method. Metal ion 1 + EDTA in excess metal- EDTA complex 1 + unreacted standard EDTA standard metal ion 2 metal- EDTA complex 2

14. 3. Replacement titration When direct or back titrations do not give sharp endpoints or when there is no suitable indicator for the analyte the metal may be determined by this method. In this method , determination of metal ion is performed by displacing magnesium or zinc ions from EDTA complex with equivalent amount of calcium ion and the liberated magnesium or zinc ions are then titrated with standard EDTA solution. To a calcium salt solution, ammonia-ammonium buffer is added. To this, standard known volume of Mg-EDTA solution is added. In this reaction, stable Ca-EDTA complex is form and Mg ions are liberated which may be titrated with a standard EDTA solution. Cadmium , lead and mercury can also be determined by this method. Ca 2+ + Mg-EDTA Ca- EDTA + Mg 2+ Stable complex Mg 2+ + EDTA Mg-EDTA

15. 4. Alkalimetric titration In an unbuffered solution, when a solution of disodium edetate (Na 2 H 2 Y) is added to the metal ion solution, complex is formed with the liberation of two equivalents of H + ions. The free H + ion then is titrated with a standard solution of alkali , NaOH using acid-base indicator or potentiometric method of detecting end point. Only metals forming EDTA complex of high stability constant can be determined by this method.

16. 5. Indirect titration This method is used to determine the ions such as Halides, phospates , and sulphates that do not form complex with EDTA. In the determination of sulphate ion, SO 4 -2 ion solution is treated with excess of standard solution of Barium ion. The formed precipitate of BaSO 4 is filtered off and unreacted Barium ions present in filtrate is titrated with EDTA. In this way, indirect determination of the amount of sulphate ion present in the sample solution is possible. Sulphate ion + excess barium ion barium sulphate ppt + (Known amount) unreacted barium ion barium- EDTA COMPLEX Filter out ppt and Titrate barium ion With EDTA

17. End point detection in complexometric titration The equivalence point of a complexation titration occurs when we react equivalent amounts of the titrand and titrant . A variety of methods are available for locating the end point, including indicators and sensors that respond to a change in the solution conditions. a. Metallochromic indicator Most indicators for complexation titrations are organic dyes —known as metallochromic indicators—that form stable complexes with metal ions. The indicator, In m– , is added to the titrand’s solution where it forms a stable complex with the metal ion, MIn n– . As we add EDTA it reacts first with free metal ions, and then displaces the indicator from MIn n– . If MIn n– and In m– have different colors , then the change in color signals the end point.

18. At the onset of the titration the buffered reaction medium contains the metal-indicator complex (MI) that has formed, and excess metal ions. When EDTA, the titrant , is added to the system a competitive reaction takes place between the free metal ions and the EDTA. Since the metal-indicator complex is weaker than the metal-EDTA chelate (K, > K MI ), the EDTA which is being added to the system during the course of the titration is progressively chelating the free metal ions in solution at the expense of the metal-indicator complex. Finally, at the equivalence point, EDTA removes the last traces of metal from the indicator, and the indicator changes from its complexed color to its metal-free color.

19. Eriochrome Black T (EBT) Indicator Eriochrome Black T becomes wine red in colour when bound with metal ions while remaining blue in colour when free from metal ions EDTA is colourless whether it’s bound to metal ions or not Eriochrome Black T binds with metal ions loosely while EDTA binds with metal ions strongly. When all metal ions are bound to EDTA, indicator EBT remains free in the sample and the solution turns blue. 

20. b. Spectrophotometric detection Limitation of metallochromic indicator is that if the solution is already colored, it will be difficult to observe the indicator’s color change at end point. For example, when titrating Cu 2+ with EDTA, ammonia is used to adjust the titrand’s pH. The intensely colored Cu(NH 3 ) 4 2+ complex unclear the indicator’s color, making an accurate determination of the end point difficult. If at least one species in a complexation titration absorbs electromagnetic radiation, then we can identify the end point by monitoring the titrand’s absorbance at a carefully selected wavelength.

21. For example, we can identify the end point for a titration of Cu 2+ with EDTA in the presence of NH 3 by monitoring the titrand’s absorbance at a wavelength of 745 nm, where the Cu(NH 3 ) 4 2+ complex absorbs strongly. At the beginning of the titration the absorbance is at a maximum. As we add EDTA, however, the reaction decreases the concentration of Cu(NH 3 ) 4 2+ and decreases the absorbance until we reach the equivalence point ,owing to the color change from an intense deep blue of the copper( ll )- tetraamine complex to the less intense light-blue color of the copper( ll )-EDTA chelate . The end point is detected when the absorbance does not change upon the addition of more titrant . absorbance Volume of EDTA

22. c. Amperometric titration An amperometric method is used for end-point detection of complexometric titrations of metal ions in unbuffered solutions with disodium ethylenediaminetetraacetate (Na 2 H 2 Y). A platinum disc microelectrode is used as a sensor for monitoring the hydrogen ions delivered during the titration. Synthetic solutions containing Ca 2+ and Mg 2+ ions, each alone or in mixture, are examined Before the equivalence point, the overall current is mainly related to the concentration of hydrogen ions, which increases up to the point where all Ca 2+ ions have been titrated. After the equivalence point, the overall current is instead related to the concentrations of both H + and H 2 Y 2- The intersection point of the two straight lines should be assumed as the end-point of the titration.

23. d. Potentiometric method End point can be detected by measuring the change of potential at an indicator electrode, immersed in the solution being analyzed, as the titrant is being added. The end point of the titration is indicated by a large and rapid change of potential. The equivalence point is indicated by a marked change in the titration curve, when potential is plotted versus volume of titrant added to the solution. The mercury electrode system can be used to detect the end point in the titration of almost all metal ions. When the electrode is immersed in a solution containing a mercury(II) salt the following equilibrium is established:

24. The potential of the electrode is given by Nernst equation: When the potential of the system is measured, the standard potential E is known, and the concentration of free mercury(II) ion is calculated by: E= E -(0.0591/2) log[Hg 2+ ] In the titration procedure using mercury electrode system, the electrode is immersed in the solution containing the metal ion to be titrated and a small quantity of a mercury(II)-EDTA chetate (usually a few drops of a 0.01 M solution),

25. Mercury(II) is displaced from the mercury(II)-EDTA chelate by the metal ion (that is, a divalent ion M 2+ ), and the following exchange equilibrium results: From this equilibrium it can be seen that the concentration of M 2+ , which is in excess of the Hg-EDTA 2- at the start of the titration, determines the amount of free mercury(II) in solution ;thus the potential of the mercury electrode depends upon the ratio [M 2+ ]/[M-EDTA 2- ]. During the early part of the titration with EDTA, the ratio [M 2+ ]/[M-EDTA 2- ] changes only slightly; however, in the vicinity of the end point this ratio undergoes its largest changes and the potential of the mercury electrode changes accordingly. As in all potentiometric titrations, the equivalence point is indicated by a sharp inflection in the titration curve.

26. TITRATION SELECTIVITY AND MASKING REAGENTS Chelons are very unselective titrants because stable chelates are formed with a great number of metal ions. When several metal ions are present in a solution, the titration data usually give the total metal content. Thus, it is necessary to find practical methods which increase the selectivity of the titrant . Several methods have been designed to increase titrant selectively : pH adjustment of the reaction medium, the use of selective metal indicators, the use of selective precipitants, the addition of masking agents to the reaction medium.

27. 1. pH adjustment of the reaction medium The formation of a metal chelate is dependent on the pH of the reaction medium. In weakly acidic solution the chelates of many metals, such as the alkaline earths, are completely dissociated, whereas chelates of bismuth, iron(II) or chromium readily form. Thus, in acidic solution, bismuth can be effectively titrated with a chelon in the presence of the alkaline earth metals. This method is based upon the differences in stability of the chelates formed between the metal ions and the chelon . 2. Use of selective metal indicators These indicators are metal complexing agents which preferentially react with different metal ions under various conditions. Pyrocatechol Violet s a well-known example of selective metal indicator. This substance preferentially reacts with bismuth in acidic solutions (pH 2 to 3). When bismuth and zinc are present in a mixture, the bismuth is directly titrated with chelon in an acidic solution using Pyrocatechol Violet as indicator.

28. At the completion of this titration, an ammonia buffer is added to adjust the pH to 10 and the system is again titrated to the Pyrocatechol Violet end point with the chelon solution. The volume of titrant employed in the second part of the titration is used to calculate the zinc concentration. 3. Use of selective precipitant Some of the more common precipitating agents which separate out interfering metal ions, or groups of metal ions are oxalate for removal of calcium ions, sulfide for the heavy metals, fluoride for aluminum, and hydroxide for some transition metals, iron, and magnesium. 4. The addition of a masking agent to the reaction medium These substances are usually auxiliary chelating or complexing compounds, which selectively react with the interfering metal ions forming more stable complexes than does the chelon .

29. Therefore, these interfering ions are said to be masked and are not titrated by the chelon . Since masking a cation in solution is an equilibrium process, it is possible to reverse this reaction by the addition of suitable demasking agents; thereby, a, series of metal ions can be titrated in one solution. One example of the use of masking and demasking reagents in chelometry is the analysis of a mixture of three metals: copper, cadmium, and calcium. Direct titration of the mixture with EDTA gives the sum of the three metals. Copper and cadmium may be masked with the addition of cyanide to the solution, leaving only the calcium ion titratable . When formaldehyde or chloral hydrate is added to the cyanide-containing mixture, only the cadmium is demasked , and the EDTA titrates the sum of the calcium and cadmium. In this manner the concentrations of all three ions are determined by three individual titrations.

30. The reaction sequences involved in the various steps of this titration procedure are represented as: Step 1: All three metals are titrated: Step 2: Only calcium titrated (Cu and Cd masked by cyanide):

31. Step 3 : Calcium and cadmium titrated (Cu masked):

32. Common Masking Agents Masking agent Metals masked Cyanide Nickel, cobalt, zinc, cadmium, copper, iron(II) Iodide Mercury Ascorbic acid Iron(III), copper Triethanotamine BAL (2,3-dimercapto-1 - propan -l- ol ) Iron( IlI ), aluminum, manganese, Tin, antimony, cadmium, mercury, zinc , Tiron (disodium catcchol-3,5-disulfonate) Aluminum, iron(III)

33. PREPARATION OF BUFFER SOLUTIONS: 1. Ammonia Buffer Dissolve 54 g of ammonium chloride in 100 ml of distilled water, add 350 ml of ammonium hydroxide solution, and dilute to 1000 ml with distilled water. 2 . Acetate Buffer Dissolve 136 g of sodium acetate trihydrate in 500 ml of water; add 29 ml of glacial acetic acid and sufficient water to make 1 liter. Preparation of 0.1 M EDTA solution To prepare a 0.1 M EDTA solution, dissolve 37.22 g of the commercially available disodium EDTA dihydrate in water. When solution is complete, dilute to 1 liter. Standardization of 0.1M EDTA solution using calcium carbonate Dissolve an accurately weighed quantity (200 to 400 mg) of calcium carbonate in 100 ml of water with the aid of a minimum amount of hydrochloric acid Adjust the pH of the solution to about 12 with a freshly prepared 0.1 N sodium hydroxide solution. Add 40 mg of murexide powder and 3 ml of Naphthol Green B indicator. Titrate to a deep-blue color with EDTA solution. Calculate the molarity by the formula: M EDTA = W CaCO3 / ( MW CaCO3 * Vol EDTA ) where W CaCO3 is the weight of calcium carbonate sample, MW CaCO3 the molecular weight of calcium carbonate (100.1), and Vol EDTA is the volume, of EDTA solution required for the titration. The EDTA solution should be standardized against the metal ion for which it is to be used as titrant and under the conditions of the actual titration. For example, if zinc is to be determined, the EDTA solution should be standardized by titrating a zinc solution prepared from pure zinc.

34. 1. DETERMINATION OF ALUMINUM (Back Titration) Dissolve an accurately weighed sample of the aluminum-containing substance, in water with the aid of a few drops of hydrochloric acid, if necessary. Dilute the solution to 100 ml with water. To 25 ml of this solution, add 25 ml of the standardized 0.1 M EDTA solution and 25 ml of water. Heat the solution on a water bath for 30 min, add 10 ml of acetate buffer, and titrate with lead nitrate or zinc sulfate solution using a few drops of Xylenol Orange as indicator. At the end point, the indicator changes from a lemon-yellow color to red. The volume of 0.1 M EDTA solution consumed by the aluminum is found by subtracting the volume of lead nitrate or zinc sulfate from the total volume of EDTA added to the system. Each milliliter of 0.1 M EDTA is equivalent to 2.698 mg of aluminum.

35. 2. DETERMINATION OF CALCIUM Calcium ions may be determined by four various titrimetric procedures a. Direct Titration Dissolve an accurately weighted sample of the calcium salt, equivalent to about 100 mg of calcium, in 50 ml of water with the aid of a few drops of - hydrochloric acid, if necessary. When solution is complete, dilute the sample with water to volume in a 100 ml volumetric flask and mix thoroughly. Accurately pipette 50 ml of the sample-solution into a suitable titration vessel containing 50 ml of water, adjust the pH of the solution to 12 with a freshly prepared 1.0 N sodium hydroxide solution, add 300mg Hydroxy Naphthol Blue indicator, and titrate rapidly with 0.1 M EDTA until the color changes to a deep blue. Each milliliter of the titrant is equivalent to 4.008 mg of calcium.

36. b. Direct Titration In the Presence of a Magnesium Salt Prepare the sample solution as described under the direct titration procedure. Pipette 50 ml of the solution into titration vessel containing 50 ml of water; add by means of a pipette 2 ml of 0.1 M magnesium sulfate solution, 5 ml of ammonia buffer, and a few drops of Eriochrome Black T indicator. Titrate with 0.1 M EDTA solution to a deep-blue color. The volume of 0.1 M EDTA consumed by the calcium is calculated by substracting the volume of 0.1 M magnesium sulfate, in milliliters, from the total volume of 0.1 M EDTA required to obtain the deep-blue color at the end point. c. BackTitration Prepare a sample solution as previously described. 50 ml of the solution into a suitable vessel, add by means of a pipette 50 ml of 0.1M EDTA solution, add 5 ml of ammonia buffer, and a few drops of Eriochrorne Black T indicator. Titrate with 0.1 M magnesium sulfate solution until the color changes from blue to wine-red. The difference in volume between the 0.1 M EDTA solution and the magnesium sulfate solution is the volume of 0.1 M EDTA solution consumed by the calcium. Each milliliter of this solution is equivalent to 4.008 mg of calcium.

37. d. Replacement Titration A solution of Mg-EDTA, equivalent to 50 mG / mI of the chelate is prepared by dissolving 19.90 g of MgSO 4 .7H 2 O and 30.1 g of Na 2 –EDTA.2H 2 O in water and diluting to 1 liter. Prepare the sample solution as previously described. Pipette 5O ml of the sample solution into-a titration vessel containing 50 ml of water; add 25 ml of Mg-EDTA solution, 5 ml of ammonia buffer, and a few drops of Eriochrome Black T indicator. Titrate to a deep blue with 0.1 M EDTA solution. Each milliliter of titrant is equivalent to 4.008 mg of calcium.

38. 3. DETERMINATION OF MAGNESIUM Dissolve an accurately weighed sample, equivalent to about 75 mg of mangesium . in sufficient water to make 100 ml. Pipette 50 ml of this solution into a suitable vessel containing 50 ml of water; add 5 ml of ammonia buffer, and a few drops of Eriochrome Black T indicator, Titrate to a deep blue color with 0.1 M EDTA solution. Each milliliter of titrant is equivalenL to 2.432 mg of magnesium. 4. DETERMINATION OF LEAD Dissolve an accurately weighed sample, equivalent to about 250 mg of lead, in sufficient water to make 100 ml. Pipette 50 ml of the sample solution into a suitable vessel, add 5 ml of acetate buffer, a few drops of Xylenol Orange indicator, and titrate with 0 I M EDTA solution to a purple-yellow end point. Each milliliter of titrant is equivalent to 20.72 mg of lead . This procedure is similar to the BP assay for lead acetate. 5. DETERMINATION OF ZINC Dissolve an accurately weighed sample of zinc salt, equivalent to about 200 mg of zinc, in sufficient water to make 100 ml. Pipette 50 ml of the sample solution into a vessel containing 50 ml of water, and add 5 ml of acetate buffer and a few drops of Xylenol Orange indicator. Titrate to a purple-yellow end point with 0.1 M ETA solution. Each milliliter of the titrant is equivalent to 6.538 mg of zinc.

39. Assay of calcium gluconate Preparation of 0.05 M EDTA Dissolve 18.6 g of Disodium EDTA dihydrate in sufficient water to produce 1000 ml. Standardization of 0.05 M EDTA Weigh accurately about 0.1 g of calcium carbonate and transfer it to a conical flask . Dissolve in 3 ml of HCl and 10 ml of purified water. Boil for 10 min, cool and dilute to 50 ml with water. Add few drops of Erichrome Black T indicator and titrate with prepared EDTA till end point. Note down volume of EDTA solution consumed during titration. Repeat procedure two more times. Calculate the molarity by: Molarity = (W CaCO3 * 0.05)/ (V EDTA *5.0045) Each ml of 0.05 M EDTA is equivalent to 5.0045 mg calcium carbonate Assay of Calcium gluconate Weigh 0.5 g of calcium gluconate and dissolve it in 20 ml of hot water containing 2 ml of hydrochloric acid. Dilute it to 100 ml of distilled water. From it, pipette out 50 ml of solution and add 10 ml of ammonium buffer solution. Titrate the solution against 0.05M EDTA using mordant black II indicator till color changes from pink to blue. Each mL of 0.05 M disodium edetate is equivalent to 22.42 mg of Calcium gluconate .