Complexometric Titration

Complexometric Titration: Complexometric titration has found an important place in the analysis of pharmaceuticals, especially for metals and their salts, certain anions and indirectly some drugs. 

The development of this technique is comparatively of recent origin. Though many types of substances (ligands) are used as a titrant in complexation phenomena, ethylene diamine tetra-acetic acid disodium salt, commonly known as EDTA (also known as sodium edetate) is most widely used. 

Complexometric titrations are those reactions in which, a simple metal ion is transformed into a complex ion by the addition of a reagent which is known as ligand. The complex formed is stable and water-soluble. 

Complexometric titrations are similar to acid-base titrations. In acid-base titrations, acid is that species that donates protons or accepts electrons. In complexometric titrations, metal ion accepts electrons and ligand donates them. Thus, in a ligand molecule, there is the presence of at least one lone pair of electrons through which co-ordinate linkage with metal ions takes place. The ligand molecule usually possesses oxygen, nitrogen or sulphur in one or more numbers in their structure. There is a particular number of ligand molecules associated with a metal ion (like 2, 4, 6 etc.) which is called coordinate linkage with metal ions takes place. The ligand molecule may have several sites (through which co-ordination linkage with metal ion occurs) present in them like unidentate (single site as in cyanide ion), bidentate (as in glycerin, oxalic acid); multidentate (many sites) such as in EDTA. 

Ethylenediamine tetra-acetic acid (EDTA) disodium salt is a very versatile complexing agent. The disodium salt is used in the preparation of the solution because of its greater water solubility. The important features of EDTA are: 

  1. Complexes formed are stable. 
  2. Complex formation is quantitative and instantaneous. 
  3. Complexes formed are water-soluble. 
  4. Complex formation occurs with most metals of the periodic table in a 1: 1 ratio.

Besides EDTA many other amino polycarboxylic acids have been used in certain specific determination because of certain advantages offered. However, EDTA remains as most popular titrant. 

In these reactions, acid is liberated and hence, buffers are employed to maintain the pH of the solution. Further, pH has a marked effect on the stability of complexes formed. 

Since the reaction of metal ion Mn+ with ligand EDTA involves the formation and production of acid and the reaction is reversible, the appropriate buffer is used, neutralized and removed, the complex will break. Complexes of alkaline earth metals are stable in an alkaline medium and decompose in neutral and acidic solutions. Some metals like aluminium, lead, mercury form complex under mildly acidic conditions while others like bismuth, iron, chromium, form and remain stable under distinct acidic pH. The monovalent ions like sodium, potassium and silver form weak complexes with no stability. 

The stability of the complex is governed by the law of mass action. Complex formation is an equilibrium process in which metal ion M reacts with a ligand molecule. 

M + nL = MLn 

Thus the stability constant Ks is governed by 

Ks = [MLn]/[M][L]

where in ml is a complex formed. As the magnitude of Ks increases, the stability of the complex is more and the concentration of unbound metal ion [M] and ligand [L] is less. 

The formation of the complex may be a stepwise reaction e.g. Cu2+ and ligand NH3. Since Cu2+ has co-ordination number four, it can react with four ammonia molecules one at a time. 

Complexes of colourless metal ions are themselves colourless and those of coloured ions are intensely coloured. 

The coloured complexes observed are endpoints in the visual titrimetric method and for these alternate method of analysis is adopted.   

Some definition: 

A ligand is an ion or molecule that binds to a central metal atom to form a complex (alternatively known as a coordination entity). 

Ligands are usually thought of as electron donors attracted to the metal at the centre of the complex. Metals are electron acceptors. 

Ligands may be neutral or negatively charged species with electron pairs available. Water is a common ligand. Here a water molecule is shown with its two lone pairs of electrons.

Ligands (Complexometric Titration)

An electron pair from the ligand, such as water, provides both of the electrons for the bond that is formed between itself and the central metal atom or ion. 

Here a single ligand L, which could be water, donates a pair of electrons to form a bond with a metal atom M. 

L → M 

Monodentate Ligands 

Monodentate ligands have only one atom capable of binding to a central metal atom or ion. H2O and NH3 are examples of neutral monodentate ligands. When H2O is a ligand, oxygen is the donor atom binding to the metal. When NH3 is a ligand, nitrogen is the donor atom binding to the metal. 

Examples of electrically charged monodentate ligands are halide ions, such as: F, Cl, Br, I, and cyano, CN

The overall charge on a complex is the arithmetic sum of the oxidation state of the metal in the centre plus the charge(s) brought to the complex by each ligand. 

For example, if complex forms between Fe2+ and six CN ligands, the complex will have a -4 charge, and the formula is written as [Fe(CN)6]4−

Polydentate Ligands

Polydentate Ligands (Complexometric Titration)

Example of a bidentate ligand-ethane-1, 2-diamine has two lone pairs of electrons.  A ligand molecule with more than one donor atom is called a polydentate ligand. These are given specific names, depending on how many donor atoms they contain. 

Bidentate Ligands 

Bidentate ligands have two atoms capable of binding to a central metal atom or ion.  Ethane-1, 2-diamine (shown in the image) is an example of a bidentate ligand. Both of the nitrogens in this molecule can act as electron donors, binding with a central metal atom or ion. 

Other examples of bidentate ligands are the acetylacetonate ion, known as (acac), and the oxalate ion (ox).

Bidentate Ligands (Complexometric Titration)

Example of a tridentate ligand. 1, 4, 7-triazaheptane (also known as diethylenetriamine) has three lone pairs of electrons.

Tridentate Ligands and Higher Polydentate Ligands 

Tridentate ligands have three atoms capable of binding to a central metal atom or ion.  Molecules with four donor atoms are called tetradentate ligands; five donor atoms, pentadentate; and six donor atoms hexadentate. 

A complex that contains a polydentate ligand is called a chelate.  The image shown below of EDTA binding to a central metal atom is an example of a chelate. Also, see crown compounds.

Tridentate Ligands and Higher Polydentate Ligands

An example of a hexadentate ligand: EDTA (also known as ethylene diamine tetraacetic acid).  The image shows a complex formed between EDTA and a central metal species. If the metal were Fe in oxidation state III, the complex would be [Fe(EDTA)]. 

Chelating agents are chemical compounds that react with metal ions to form a stable, water-soluble complex. They are also known as chelants, chelators, or sequestering agents. 

Chelating agents have a ring-like centre that forms at least two bands with the metal ion allowing it to be excreted. Chelating agents are usually organic compounds (a compound that contains carbon). 

Specific chelating agents bind iron, lead, or copper in the blood and can be used to treat excessively high levels of these metals. Chelating agents may also be used in the treatment of heavy metal poisoning.   


The technique involves titrating metal ions with a complexing agent or chelating agent (ligand) and is commonly referred to as complexometric titration. This method represents the analytical application of a complexation reaction. In this method, a simple ion is transformed into a complex ion and the equivalence point is determined by using metal indicators or electrometrically. Various other names such as chilometric titrations, chilometry, chilatometric titrations and EDTA titrations have been used to describe this method. All these terms refer to the same analytical method and they have resulted from the use of EDTA (Ethylene diamine tetraacetic acid) and other chilons. These chilons react with metal ions to form a special type of complex known as a chelate.

Metal ions in solution are always solvated, i.e. a definite number of solvent molecules (usually 2, 4 or 6) are firmly bound to the metal ion. However, these bound solvent molecules are replaced by other solvent molecules or ions during the formation of a metal complex or metal co-ordination compound. 

The molecules or ions which displace the solvent molecules are called Ligands. Ligands or complexing agents or chelating agents can be any electron donating entity, which can bind to the metal ion and produce a complex ion. An example of a complexation reaction between Cu(II) ion and four ammonium molecules in an aqueous solution may be  expressed by the following equation:

ammonium molecules in an aqueous solution

Bonding in Complexes 

The bonds are either ordinary covalent bonds in which the metal and the ligand contribute one electron each, or co-ordinate bonds in which both electrons are contributed by the ligand. Thus, the hexacyanoferrate ion may be considered to consist of three ordinary covalent bonds and three co-ordinate bonds, although in the complex the bonds are identical hybrid bonds that are directed towards the apices of a regular octahedron.

The hexacyanoferrate iron (III) ion (Complexometric Titration)
The hexacyanoferrate iron (III) ion 

The negative charge on the complex ion is equal to the total number of negative groups minus the valency of the metal ion. When neutral groups only are involved, the charge on the complex is positive and is equal to the metal ion, e.g. [Cu(NH3)4]2+

Werner’s Co-ordination Number 

Werner (1891) first noticed that for each atom there were an observed maximum number of small groups which can be accommodated around it. This number, which is called Werner’s co-ordination number, depends purely upon steric factors and is in no way related to the valency of the ion. Thus, although the valency shell of the elements of the third period is theoretically capable of expanding up to 18 electrons, and that of the fourth to 32 electrons, there is, in practice, a limit to the number of small groups which can be accommodated owing to limitations of space around the ion. For example, in the [BF4] ion,  the octet is completed and the maximum co-ordination number is reached, but in the  [AlF6]ion the outer shell contains 12 electrons and cannot expand to the maximum number of 18 electrons since the maximum co-ordination number has been reached. 

Within the limits imposed by Werner’s co-ordination number, there is a tendency for the metal to attain or approach inert gas structure, and this is probably the driving force for complex formation. 

Chelate Compound or Chelate

Complexes involving simple ligands, i.e., those forming only one bond are described as co-ordination compounds. A complex of a metal ion with 2 or more groups on a multidentate ligand is called a chelate or a chelate compound. There is no fundamental difference between co-ordination compound and a chelate compound except that in a chelate compound, the ring influence the stability of a compound. Thus, a chelate can be described as a heterocyclic ring structure in which a metal atom is a member of the ring. The stability of a chelate is usually much greater than that of a corresponding unidentate metal complex. 

Chelating agent 

Ligands having more than one electron-donating group are called chelating agents. The most effective complexing agent in ligands is amino and carboxylate ions. 

The solubility of metal chelates in water depends upon the presence of hydrophilic groups such as COOH, SO3H, NH2 and OH. When both acidic and basic groups are present, the complex will be soluble over a wide range of pH. When hydrophilic groups are absent, the solubilities of both the chelating agent and the metal chelate will be below, but they will be soluble in organic solvents. The term sequestering agent is generally applied to chelating agents that form water-soluble complexes with bi- or poly-valent metal ions. Thus, although the metals remain in solution, they fail to give normal ionic reactions. Ethylenediamine tetraacetic acid is a typical sequestering agent, whereas, dimethylglyoxime and salicylaldoxime are chelating agents, forming insoluble complexes.

Chelating agent 

As a sequestering agent, ethylene diamine tetra-acetic acid reacts with most polyvalent metal ions to form water-soluble complexes which cannot be extracted from aqueous solutions with organic solvents. Dimethylglyoxime and salicylaldoxime form complexes that are insoluble in water, but soluble in organic solvents; for example, nickel dimethylglyoxime has a sufficiently low solubility in water to be used as a basis for gravimetric assay. 

EDTA forms chelates with nearly all metal ions and this reaction are the basis for the general analytical method for these ions by titration with a standard EDTA solution. Such titrations are called complexometric or chilometric or EDTA titrations. 

Reagent EDTA 

Disodium salt of EDTA is a water-soluble chelating agent and is always preferred. It is non-hygroscopic and a very stable sequestering agent (Ligands that form water-soluble chelates are called sequestering agents). 

Some chelating agents form water-insoluble chelates with metal ions. e.g. oxine or 8-hydroxy quinoline.

Reagent EDTA 

EDTA and 8-hydroxy quinoline are important reagents used in analytical chemistry.  Sequestering agents are used to liberate or solubilize metal ions. The agents which form water-insoluble chelates are used to remove the metal ions from the solution by precipitation. 

EDTA has the widest general application in the analysis because of the following important properties: 

  • It has a low price.
  • The special structure of its anion has 6 ligand atoms.
  • It forms strainless five-membered rings.
  • Disodium EDTA is used as an M/20 solution. 

Purification of Disodium EDTA: 

Commercial samples of disodium EDTA may be purified for use as a primary standard by adding ethanol to a saturated aqueous solution until the first permanent precipitate appears; filter and add an equal volume of ethanol; filter the precipitated disodium EDTA, wash with acetone and ether, and dry to constant weight at 80°C, drying may require four days. The official material contains not less than 98% dihydrate. 

Preparation of M/20 Disodium EDTA: Dissolve 18.6 gm of disodium EDTA in water and make the volume up to 1000 ml and standardize the prepared solution.

Standardization of Disodium EDTA: Weigh accurately about 200 mg of CaCO3 in a titration flask. Add 50 ml of water and a minimum quantity of dil. HCl to dissolve CaCO3. Adjust the pH of the solution to 12 by adding NaOH. Add 300 mg of hydroxy naphthol blue indicator and titrate with the prepared M/20 disodium EDTA solution, until the solution is deep blue. 

The HCl solubilizes the CaCO3 by converting it to CaCl2. The NaOH makes the solution alkaline and maintains the pH at about 12 so that the Ca-EDTA complex would be stable and any Mg, which might be present as a contaminant, would not react. The coloured Ca-indicator complex gives up Ca to EDTA, liberating the free uncomplexed indicator, which is blue. 

Factors influencing EDTA reactions:

  • The nature and activity of the metal ion.
  • The pH at which the titration is carried out.
  • The presence of interfering ions such as CN, Citrate, Tartrate, F and other complex-forming agents.
  • Organic solvents also increase the stability of the complex. 

Nature and stability of metal complexes of Ethylenediaminetetra-acetic acid: 

Ethylenediaminetetra-acetic acid forms complexes with most cations in a 1: 1 ratio, irrespective of the valency of the ion:

Nature and stability of metal complexes of Ethylenediaminetetra-acetic acid

where M is a metal and [H2X]2− is the anion of the disodium salt (disodium EDTA) which is most frequently used. 

Effect of pH on Complex Formation 

Ethylenediamine tetra-acetic acid ionizes in four stages (pK1 = 2.0, pK2 = 2.67, pK3 = 6.16 and pK4 = 10.26) and, since the actual complexing species is Y4−, complexes will form more efficiently and be more stable in alkaline solution. If, however, the solubility product of the metal hydroxide is low, it may be precipitated if the hydroxyl ion concentration is increased too much. On the other hand, at lower pH values when the concentration of Y4− is lower, the stability constant of the complexes will not be so high. Complexes of most divalent metals are stable in ammonical solution. Those alkaline earth metals, such as copper, lead and nickel, are stable down to pH 3 and hence can be titrated selectively in the presence of alkaline earth metals. Trivalent metal complexes are usually still more firmly bound and stable in strongly acid solutions; for example, the cobalt(III) edetate complex is stable in concentrated hydrochloric acid. Although most complexes are stable over a fair range of pH, solutions are usually buffered 6 at a pH at which the complex is stable and at which the colour change of the indicator is most distinct. 

Colour of complexes: 

There is always a change in the absorption spectrum when complexes are formed and this forms the basis of many colourimetric assays. 

Stability of Complexes: The general equation for the formation of a 1:1 chelate complex, MX, is 

M + X → MX 

where M is the metal ion and X is the chelating agent. 

Stability constant (K) = [MX]/[M][X] 

An increase in temperature causes a slight increase in the ionization of the complex and a slight lowering of K. the presence of electrolytes having no ion in common with the complex decreases K, whilst the presence of ethanol increases K, probably due to the suppression of ionization. 

Principle: Many principles of acid-base titrations are used in complexometric titration. In a complexometric titration, the free metal ions disappear as they are changed into complex ions. In acid-base titrations, the endpoint is marked by a sudden change in pH. Similarly, in EDTA titration, if we plot the pM (negative log of metal ion concentration) v/s volume of titrant, we will find that at the endpoint, the pM rapidly increases. This sudden pM raise results from the removal of traces of metal ions from the solution by EDTA. 

Any method, which can determine this disappearance of free metal ions, can be used to detect the endpoint in complexometric titrations. An endpoint can be detected usually with an indicator or instrumentally by the potentiometric or conductometric (electrometric) method. 

Three factors are important in determining the magnitude of a break in the titration curve at an endpoint. 

1. The stability of the complex formed: The greater the stability constant for complex formed, the larger the charge in free metal concentration (pM) at the equivalent point and the more clear would be the endpoint. 

2. The number of steps involved in the complex formation: Fewer the number of steps required in the formation of a complex, the greater would be the break in the titration curve at the equivalent point and clear would be the endpoint. 

3. Effect of pH: During a complexometric titration, the pH must be constant by the use of a buffer solution. Control of pH is important since the H+ ion plays an important role in chelation. 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.

Thus, the stability of the metal complex is pH-dependent. The lower the pH of the solution, the lesser would be the stability of the complex (because more H+ ions are available to compete with the metal ions for ligand). Only metals that form very stable complexes can be titrated in an acidic solution, and metals forming weak complexes can only be effectively titrated in an alkaline solution. 

Methods of End Point Detection 

An endpoint in complexometric titration can be detected by the following two methods: 

1. Indicators method of Endpoint Detection:  The endpoint in complexometric titrations is shown using pM indicators. The concept of pM arises as follows: 

Therefore, if a solution is made such that [X] = [MX], pM = −pK (or pM = pK’, where K’ = dissociation constant). This means that, in a solution containing equal activities of the metal complex and free chelating agent, the concentration of metal ions will remain roughly constant and will be buffered in the same way as hydrogen ions in a pH buffer. Since, however, chelating agents are also bases; equilibrium in a metal-buffer solution is often greatly affected by a change in pH. In general, for chelating agents of the amino acid type (e.g., edetic acid and ammonia triacetic acid), it may be said that when [X] = [MX], pM increases with pH until about pH 10, when it attains a constant value. This pH is, therefore, usually chosen for carrying out titrations of metals with chelating agents in buffered solutions. 

The pM indicator is a dye that is capable of acting as a chelating agent to give a dyemetalcomplex. The latter is different in colour from the dye itself and also has low stability constant than the chelate-metal complex. The colour of the solution, therefore, remains that of the dye complex until the endpoint, when an equivalent amount of sodium EDTA has been added. As soon as there is the slightest excess of EDTA, the metal-dye complex decomposes to produce free dye; this is accomplished by a change in colour. 

Organic compounds form coloured chelates with ions in a pM range that is unique to the cation and the dye selected. To be useful, the dye-metal chelates usually will be visible at 10-6-10-7 M concentration. Many of these indicators also have typical properties of acid-base indicators and the colour changes are the result of the displacement of H+ by a metal ion. Metal indicators must comply with the following requirements:

The compound must be chemically stable throughout the titration. 

  • It should form a 1: 1 complex which must be weaker than the metal chelate complex. 
  • The Colour of the indicator and the metal complexed indicator must be sufficiently different. 
  • The Colour reaction should be selective for the metal being titrated. 
  • The indicator should not compete with the EDTA.

Mechanism of action of the indicator: Let the metal be denoted by M, the indicator by I and the chelated by EDTA. At the onset of the titration, the reaction medium contains the metal indicator complex (MI) and excess metal ions. When an EDTA titrant is added to the system, a competitive reaction takes place between the free metal ions and EDTA. Since the metal indicator complex (MI) is weaker than the metal-EDTA chelate, the EDTA which is being added during the titration is chelating the free metal ions in solution at the expense of the MI complex. Finally, at the endpoint, EDTA removes the last traces of the metal from the indicator and the indicator changes from its complexed colour to its metal-free colour. The overall reaction is given by:

Mechanism of action of the indicator

Table.1: Indicators used in complexometric titrations

Indicators used in complexometric titrations

2. Instrumental methods of Endpoint detection: Spectrophotometric detection: The change in absorption spectrum when a metal ion of a complexing agent is converted to the metal complex, or when one complex is converted to another can usually be detected more accurately and in a more dilute solution by spectrophotometric than by visual methods. Thus, in disodium EDTA titrations, an accurate endpoint can be obtained using 0.001 M solutions. In practice, an indicator giving a colour change in the visible region is generally employed, but coloured ions may be titrated without an indicator using spectrophotometric methods. Also, it is sometimes possible to use an endpoint in the ultraviolet region for ions and complexes which are colourless in the visible region. 

Amperometric titration: The effect of complex formation on the half-wave potential of an ion is to render it more negative. If the electrode potential is adjusted to a value between that of the half-wave potential of the free cation and that of the complex, and disodium EDTA solution is added slowly, the diffusion current will fall steadily until it equals the residual current, that is, until the last trace of free cation has been complex. This is the endpoint and the amount of standard disodium EDTA solution added is equivalent to the amount of metal present. 

Potentiometric titration: Since disodium EDTA reacts preferentially with the higher valency state of an ion, it will reduce the redox potential according to the equation, 


  • E = Eo + loge [Ox]/[Red] 
  • E = The potential of the electrode 
  • Eo = The standard electrode potential 
  • [Ox] = Activity of ions in the oxidized state 
  • [Red] = Activity of ions in the reduced state 

This method is of limited application owing to the lack of suitable indicator electrodes.  Iron (III) and copper (II), however, can be titrated in this way. Back titration of excess disodium EDTA with ferric chloride in acid solution is possible for some ions.   

Structure of PM indicator 

  1. Alizarin fluorine blue (alizarin complexone). 
  2. Murexide. 
  3. Xylenol orange. 
  4. Tiron (disodium, 1, 2-dihydroxy phenol-3, 5-disulphonate). 
  5. Mordant black II (Eriochrome black T, Solochrome black T). 
  6. Catechol violet. 
  7. Calcone.

Types of Complexometric Titrations 

1. Direct Titration: It is the simplest and the most convenient method used in chelometry. In this method, the standard chelon solution is added to the metal ion solution until the endpoint is detected. This method is analogous to simple acid-base titrations.  e.g. Calcium gluconate injection, calcium lactate tablets and compound sodium lactate injection for the assay of calcium chloride (CaCl2.6H2O). 


Slow complexation reaction: Interference due to the presence of other ions. 

1. Back Titration: In this method, excess of a standard EDTA solution is added to the metal solution, which is to be analyzed, and the excess is back titrated with a standard solution of a second metal ion. e.g. Determination of Mn. This metal cannot be directly titrated with EDTA because of precipitation of Mn(OH)2. An excess known volume of EDTA is added to an acidic solution of Mn salt and then ammonia buffer is used to adjust the pH to 10 and the excess EDTA remaining after chelation, is back titrated with a standard Zn solution kept in burette using Eriochrome Black T as an indicator. This method is analogous to the back titration method in acidimetry. e.g. ZnO. 

2. Replacement Titration: In this method, the metal which is to be analyzed, displaces quantitatively the metal from the complex. When direct or back titrations do not give sharp endpoints, the metal may be determined by the displacement of an equivalent amount of Mg or Zn from a less stable EDTA complex. 

Mn+2 + Mg EDTA-2 → Mg+2 + Mn EDTA-2 

Mn displaces Mg from Mn EDTA solution. The freed Mg metal is then directly titrated with a standard EDTA solution. In this method, the excess quantity of Mg EDTA chelate is added to the Mn solution. Mn quantitatively displaces Mg from Mg EDTA chelate. This displacement takes place because Mn forms a more stable complex with EDTA. By this method, Ca, Pb, Hg may be determined using the Eriochrome Black T indicator. 

3. Indirect Titration: This is also known as alkalimetric titration. It is used for the determination of ions such as anions, which do not react with EDTA chelate. Protons from disodium EDTA are displaced by heavy metals and titrated with sodium alkali. 

e.g. Barbiturates do not react with EDTA but are quantitatively precipitated from alkaline solution by mercuric ions as a 1: 1 complex. 

Some important elements which could be determined by complexometric titration are as follows: 

Direct Titration: Analysis of Cu, Mn, Ca, Ba, Br, Zn, Cd, Hg, Al, Thallium, Sn, Pb, Bi, Vanadium, Cr, Mo, Gallium, Fe, Co, Ni, and Pd. 

Indirect Titration: Analysis of Na, K, Ag, Au, As, C, N, P, S, Cl, Br, I and F.

Titration Selectivity, Masking and Demasking Agents 

EDTA is a very unselective reagent because it complexes with numerous doubly, triply and quadruply charged cations. When a solution containing two cations that complex with EDTA is titrated without the addition of a complex-forming indicator, and if a titration error of 0.1% is permissible, then the ratio of stability constants of EDTA complexes of the two metals M and N must be such that KM / KN ≥ 106 if N is not to interfere with the titration of M.  strictly, of course, the constants KM and KN considered in the above expression should be the apparent stability constants of the complexes. If the complex-forming indicators are used, then for a similar titration error KM / KN ≥ 108

The following procedures will help to increase the selectivity:

  • Use of masking and demasking agents.
  • pH control.
  • Use of selective metal indicators.
  • Classical separation.
  • Solvent extraction.
  • Removal of anions.
  • Kinetic masking.   

Use of Masking and Demasking Agents: 

Masking agents act either by precipitation or by the formation of complexes more stable than the interfering ion-EDTA complex. 

Masking by Precipitation: Many heavy metals e.g. Co, Cu and Pb, can be separated either in the form of insoluble sulphides using sodium sulphide or as insoluble complexes using thioacetamide. These are filtered, decomposed and titrated with disodium EDTA. Other common precipitating agents are sulphate for Pb and Ba, oxalate for Ca and Pb, fluoride for Ca, Mg and Pb, ferrocyanide for Zn and Cu, and 8-hydroxyquinoline for many heavy metals.  Thioglycerol (CH2SH.CHOH.CH2OH) is used to mask Cu by precipitation in the assay of lotions containing Cu and Zn. 

Masking by Complex formation: Masking agents form more stable complexes with the interfering metal ions. The most important aspect is that the masking agent must not form complexes with the metal ion under analysis. 

The different masking agents used are enlisted below: 

Ammonium fluoride will mask aluminium, iron and titanium by complex formation. 

Ascorbic acid is a convenient reducing agent for iron (III) which is then masked by complexing as the very stable hexacyanoferrate (II) complex. This latter is more stable and less intensely coloured than the hexacyanoferrate (III) complex.

All these complexes are stronger than the corresponding edetate complexes and are almost colourless. Cobalt, copper and nickel form intense yellowish-green complexes with the reagent under the above conditions. Cobalt and copper, but not nickel, are displaced from their edetate complexes by dimercaprol. 

Potassium cyanide reacts with silver, copper, mercury, iron, zinc, cadmium, cobalt and nickel ions to form complexes in an alkaline solution that are more stable than the corresponding edetate complexes, so that other ions, such as lead, magnesium, manganese and the alkaline earth metals can be determined in their presence. Of the metals in the first group mentioned, zinc and cadmium can be demasked from their cyanide complexes by aldehydes, such as formaldehyde or chloral hydrate (due to the preferential formation of a cyanohydrin), and selectively titrated. 

Demasking: It is the process in which the masked substance regains its ability to enter into a particular reaction. This enables a determination of a series of metal ions in one solution containing many cations. 

pH control method: The formation of a metal chelate is dependent on the pH of the reaction medium. In a weakly acid solution, the chelates of many metals are completely dissociated such as alkaline earth metals, whereas chelates of Bi, Fe3+ or Cr are readily formed at this pH. Thus, in an acidic solution, Bi can be effectively titrated with a chelating agent in the presence of alkaline earth metals. This method is based upon the differences in stability of the chelates formed between the metal ions and the chelating agent. 

Use of selective metal indicators: These indicators are the metal complexing agents which react with different metal ions under various conditions. Several selective metal indicators have been used and they are specific for a particular ion.   

Applications of Complexometric Titrations 

Complexometric titrations have been employed with success for the determination of various metals like Ca, Mg, Pb, Zn, Al, Fe, Mn, Cr etc. in different formulations that are official in I.P., and also for the determination of the Hardness of water. 

Determination of calcium in different formulations: Calcium can be determined in almost every formulation by EDTA-titrations. e.g. Five membered heterocyclic rings are formed with EDTA, which are stain-free, and thus highly stable.

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