Complexation titration. Determination of weight of salts Ca2+, Mg2+, Zn2+, Cu2+, Bi3+ in drugs, definition of hardness of water.
Complexation titration is used for definition of many metal cations. In particular, direct complexation titration to define drugs of magnesium, calcium, zinc, lead, bismuth, and a method of back titration – aluminium. Sometimes in complexymetry it is possible to define not only cation of metals, but also anion (phosphates, sulphates, oxalates, chromates, etc.), applying displacement and back titration in a combination with a precipitation method.
Complexation titration is pharmacopoeia’s method. This method express, exact also is used in the analysis of substances and drugs. In particular, this method is used in the analysis of such pharmaceutical preparations, as a alumag (contains aluminium, magnesium); magnesium sulphate; calcium gluconate, calcium lactate, calcium chloride; oxide, sulphate zinc; the basic nitrate and subcytrate colloidal bismuth; Xeroformium (contains bismuth).
These titrations are based on complexation reactions.
Most often used reagent is EDTA – EthyleneDiamineTetraAcetic acid. There are also other similar chelating agents (EGTA, CDTA and so on) used. In some of other methods Ag+ is used as a titrant for determining cyanides and Hg2+ as a titrant in Cl- determination.
Changing property of the solution is usually the concentration of the complexed substance, although in some cases it can be much more convenient to express results in terms of titrant concentration. As its concentration changes by many orders of magnitude, and is almost always smaller than 1, we use negative logarithmic scale, similar to that used in pH definition.
In the case of determination of metals detection of the endpoint is mainly based on substances that change color when creating complexes with determined metals. One of these indicators is eriochrome black T, substance that in pH between 7 and 11 is blue when free, and black when forms a complex with metal, other examples are pyrocatechin violet and murexide. It is important that formation constant for these complexes is low enough, so that titrant reacts with complexed ions first.
The earliest titrimetric applications involving metal–ligand complexation were the determinations of cyanide and chloride using, respectively, Ag+ and Hg2+ as titrants. Both methods were developed by Justus Liebig (1803–1873) in the 1850s. The use of monodentate ligand, such as Cl– and CN–, however, limited the utility of complexation titrations to those metals that formed only a single stable complex, such as Ag(CN)2– and HgCl2. Other potential metal–ligand complexes, such as CdI42-, were not analytically useful because the stepwise formation of a series of metal–ligand complexes (CdI+, CdI2, CdI3–, and CdI42–) resulted in a poorly defined end point.
The utility of complexation titrations improved following the introduction by Schwarzenbach, in 1945, of aminocarboxylic acids as multidentate ligands capable of forming stable 1:1 complexes with metal ions. The most widely used of these new ligands was ethylenediaminetetraacetic acid, EDTA, which forms strong 1:1 complexes with many metal ions. The first use of EDTA as a titrant occurred in 1946, when Schwarzenbach introduced metallochromic dyes as visual indicators for signalling the end point of a complexation titration.
Classification of complexation titration:
§ Mercurimetry – titrant is solution of Hg(NO3)2
§ Fluoridometry – titrant is solution of NaF
§ Cyanidometry – titrant is solution of KCN
§ Complexonometry – titrant is solution of sodium edetate
Requirements to reactions in complexation titration
§ reactions between the titrant and analyte must be stoichiometricaly, quantitatively
§ formation constant of complex should be more than b ³ 108
§ reaction of complex compound formation should be quickly
§ there should be a possibility of fixing of an equivalence point or the end point
§ in the conditions of titration carrying out competing reactions should not proceed
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 same analytical method and they have resulted from the use of EDTA (Ethylene diamine tetra acetic acid) and other chilons. These chilons react with metal ions to form a special type of complex known as 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 has the ability to 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:
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 which have been shown to be directed towards the apices of a regular octahedron.
The hexacyanoferrate iron (III) ion
The negative charge on the complex ion is equal to the total number of the negative groups minus the valency of the metal ion. Wheeutral 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-ordinatioumber, depends purely upon steric factors and is io 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-ordinatioumber 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-ordinatioumber 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.
Classification of Ligands
1. Unidentate Ligands: Ligands that are bound to metal ion only at one place are called unidentate ligands (one toothed). NH3, for example, is a unidentate ligand capable of complexing with cupric ions. Halide ions, cyanide ions and NH3 are common examples of unidentate ligands. The formation of complex Cu (NH3)42+ proceeds in the following steps:
Considering the overall reaction:
2. Bidentate and Multidentate Ligands: Many ligands are known that contain more than one group, capable of binding with metal ions. Such ligands are known as multidentate ligands or chelating agents. They include bidentate ligands (2 donar atoms), tridentate ligands (3 donar atoms), quadridentate ligands, etc.
Thus, ethylene diamine is an example of bidentate ligand. H2N-CH2-CH2-NH2
Ethylene diamine tetra acetic acid (EDTA) is an example of multidentate ligand.
Mercurimetry
§ Titrant: secondary standard solution of Hg(NO3)2
§ Standardization: on primary standard solution of sodium chloride NaCl:
Hg(NO3)2 + 2 NaCl = HgCl2 + 2 NaNO3
!!!The main lack of mercurimetry – high toxicity of Mercury compounds.
Indicators:
1. Solution of sodium pentacianonitrozoferrate (ІІІ) (sodium nitroprusside) Na2[Fe(CN)5NO], which forms with Hg2+-ions insoluble white salt:
Na2[Fe(CN)5NO] + Hg(NO3)2 = 2NaNO3 + Hg[Fe(CN)5NO]¯
Mercurimetric determination of chloridic with sodium nitroprusside
§ Mercurimetric determination of iodide is based on such reaction (without indicator):
until e.p.
Hg2+ + 4I- = [HgI4]2-
In the end point we observe appearance of red precipitate:
[HgI4]2- + Hg2+ = HgI2¯
red precipitate
Mercurimetric determination of iodide
§ Determination of thiocyanide SCN- – ions is based on reaction:
Hg2+ + 2SCN- = Hg(SCN)2
§ As indicators we use solution of Fe (III) salts
§ To the end point:
Fe3+ + 3SCN- = Fe(SCN)3
we observe red colour of solution
§ In the end point:
2Fe(SCN)3 + 3Hg(NO3)2 = 3Hg(SCN)2+2Fe(NO3)3.
red colour of solution disappears
Determination of mercury (ІІ) salts
§ Titrant – solution of potassium thiocyanide KSCN
§ Indicator – ions of Fe3+.
§ To the end point:
Hg2+ + 2SCN- = Hg(SCN)2
In the end point appears red colour of iron (ІІІ) thiocyanide solution:
§ In the end point:
Fe3+ + 3SCN- = Fe(SCN)3
Fluoridometry
§ Defined ions: aluminium Al3+, zirconium ZrIV, thorium ThIV, calcium Ca2+.
§ We use following reactions for determination:
Al3+ + 6F- = [AlF6]3-
ZrIV + 6F- = [ZrF6]2-
ThIV + 6F- = [ThF6]2-
§ Ions of Са2+ we determine by back titration:
Ca2+ + 2F- (excess) = CaF2¯
6F- (residue) + Al3+ = [AlF6]3-
Indicators:
§ acid-base indicators (Methyl orange), as solutions have acidic reaction of medium:
Al3+ + HOH = AlOH2+ + H+
§ Alizarine sulfate (forms complex with zirconium ZrIV – red-violet colour, with thorium ThIV– violet)
§ Fluoridometry serves for definition of rather high contents of substances (0,2-0,5 mol/L). The relative error can reach 1-3 %.
Fluoridometric determination of aluminium with Methyl orange
Fluoridometric determination of aluminium with Alizarine (ammoniac medium)
Cyanidometry
§ Titrant: secondary standard solution of potassium cyanide KCN
§ Standardization: on standard solution of AgNO3:
Ag+ + 2CN- = [Ag(CN)2]-
[Ag(CN)2]- + Ag+ = 2AgCN¯
in the end point appears turbidity
§ Defined substances: heavy metals, which form stable cyanidic complexes of specific structure, for example, [Ni(CN)4]2- , [Co(CN)4]3-, [Zn(CN)4]2-.
§ The reaction of methods:
[Ni(NH3)4]2+ + 4CN- = [Ni(CN)4]2- + 4NH3
§ Medium: ammoniac
§ Indicator: suspension of silver iodide AgI.
In the end point: dissolves the precipitate of silver iodide AgI
AgI¯ + 2CN- = [Ag(CN)2]- + I-
turbidity transparent solution
Potassium cyanide is strong toxin!
Chelate Compound or Chelate
Complexes involving simple ligands, i.e., those forming only one bond are described as co-ordination compound. 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, ring influence the stability of compound. Thus, a chelate can be described as a heterocyclic ring structure in which a metal atom is a member of ring. The stability of a chelate is usually much greater than that of corresponding unidentate metal complex.
Chelating agent
Ligands having more than one electron donating groups are called chelating agents. The most effective complexing agent in ligands are amino and carboxylate ions. All the multidentate ligands important in analytical chemistry contain the structure component as follows:
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 low, 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. Ethylenediaminetetra-acetic acid is a typical sequestering agent, whereas, dimethylglyoxime and salicylaldoxime are chelating agents, forming insoluble complexes.
As a sequestering agent, ethylenediaminetetra-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 which 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 is the basis for 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 which form water soluble chelates are called sequestering agents).
There are cheating agents that form water insoluble chelates with metal ions. E.g. – oxine or 8-hydroxy quinoline.
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 solution by precipitation.
EDTA has the widest general application in analyses because of the following important properties:
– It has low price.
– The special structure of its anion which has 6 ligand atoms.
– It forms strainless five-membered rings.
Disodium EDTA is used as 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
Preparation of M/20 Disodium EDTA: Dissolve
Standardization of Disodium EDTA: Weigh accurately about 200 mg of CaCO3 in a titration flask. Add 50 ml of water and minimum quantity of dil. HCl to dissolve CaCO3. Adjust the pH of the solution to 12 by adding NaOH. Add 300 mg of hydroxyl naphthol blue indicator and titrate with the prepared M/20 disodium EDTA solution, until the solution is deep blue in colour.
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 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 complex.
Nature and stability of metal complexes of Ethylenediaminetetra-acetic acid: Ethylenediaminetetra-acetic acid forms complexes with complexes with most cations in a 1:1 ratio, irrespective of the valency of the ion:
where M is a metal and [H2X]2- is the anion of the disodium salt (disodium EDTA) which is most frequently used. The structures of these complexes with di-, tri- and tetravalent metals contain three, four and five rings respectively:
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 of the 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 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 colorimetric assays.
Stability of Complexes: The general equation for the formation of a 1:1 chelate complex, MX, is
where M is the metal ion and X the chelating agent.
where [ ] represents activities. 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.
Chemistry and Properties of EDTA
Ethylenediaminetetraacetic acid, or EDTA, is an aminocarboxylic acid. The structure of EDTA is shown
EDTA, which is a Lewis acid, has six binding sites (the four carboxylate groups and the two amino groups), providing six pairs of electrons. The resulting metal–ligand complex, in which EDTA forms a cage-like structure around the metal ion, is very stable.
The actual number of coordination sites depends on the size of the metal ion; however, all metal–EDTA complexes have a 1:1 stoichiometry.
Metal–EDTA Formation Constants. To illustrate the formation of a metal–EDTA complex consider the reaction between Cd2+ and EDTA
CdCl2 + H4Y ® CdH2Y + 2HCl
where H4Y is a shorthand notation for the chemical form of EDTA. The formation constant for this reaction
quite large, suggesting that the reaction’s equilibrium position lies far to the right.
EDTA Is a Weak Acid. Besides its properties as a ligand, EDTA is also a weak acid. The fully protonated form of EDTA, H6Y2+, is a hexaprotic weak acid with successive pKa values of
pKa1 = 0.0; pKa2 = 1.5; pKa3 = 2.0; pKa4 = 2.68; pKa5 = 6.11; pKa6 = 10.17.
The first four values are for the carboxyl protons, and the remaining two values are for the ammonium protons. A ladder diagram for EDTA is shown
The species Y4– becomes the predominate form of EDTA at pH levels greater than 10.17. It is only for pH levels greater than 12 that Y4– becomes the only significant form of EDTA.
Conditional Metal–Ligand Formation Constants. Recognizing EDTA’s acid–base properties is important. The formation constant for CdY2– assumes that EDTA is present as Y4–. If we restrict the pH to levels greater than 12, then equation provides an adequate description of the formation of CdY2–. For pH levels less than 12, however, Kf overestimates the stability of the CdY2– complex.
At any pH a mass balance requires that the total concentration of unbound EDTA equal the combined concentrations of each of its forms.
CEDTA = [H6Y2+] + [H5Y+] + [H4Y] + [H3Y–] + [H2Y2–] + [HY3–] + [Y4–]
To correct the formation constant for EDTA’s acid–base properties, we must account for the fraction, , of EDTA present as Y4–.
If we fix the pH using a buffer, then is a constant. Combining with Kf gives
where Kf´ is a conditional formation constant* whose value depends on the pH. As shown in Table 9.13 for CdY2–, the conditional formation constant becomes smaller, and the complex becomes less stable at lower pH levels.
EDTA Must Compete with Other Ligands. To maintain a constant pH, we must add a buffering agent. If one of the buffer’s components forms a metal–ligand complex with Cd2+, then EDTA must compete with the ligand for Cd2+. For example, an ammonia buffer (NH4Cl/NH3OH) includes the ligand NH3, which forms several stable Cd2+–NH3 complexes. EDTA forms a stronger complex with Cd2+ and will displace NH3. The presence of NH3, however, decreases the stability of the Cd2+–EDTA complex.
We can account for the effect of an auxiliary complexing agent**, such as NH3, in the same way we accounted for the effect of pH. Before adding EDTA, a mass balance on Cd2+ requires that the total concentration of Cd2+, CCd, be
CCd = [Cd2+] + [Cd(NH3)2+] + [Cd(NH3)22+] + [Cd(NH3)32+] + [Cd(NH3)42+]
* Conditional formation constant is the equilibrium formation constant for a metal–ligand complex for a specific setof solution conditions, such as pH.
** Auxiliary complexing agent is a second ligand in a complexation titration that initially binds with the analyte but is displaced by the titrant.
Principle of Complexometric Titration
Many principles of acid-base titrations are used in complexometric titration. In complexometric titration, the free metal ions disappear as they are changed into complex ions. In acid-base titrations, the end point is marked by sudden change in pH. Similarly, in EDTA titration, if we plot pM (negative log of metal ion concentration) v/s volume of titrant, we will find that at the end point, the pM rapidly increases (Fig. 1). This sudden pM raise results from removal of traces of metal ions from solution by EDTA.
Any method, which can determine this disappearance of free metal ions, can be used to detect end point in complexometric titrations. End point can be detected usually with an indicator or instrumentally by potentiometric or conductometric (electrometric) method.
There are three factors that are important in determining the magnitude of break in titration curve at end point.
1. The stability of complex formed: The greater the stability constant for complex formed, larger the charge in free metal concentration (pM) at equivalent point and more clear would be the end point.
2. The number of steps involved in complex formation: Fewer the number of steps required in the formation of complex, greater would be the break in titration curve at equivalent point and clear would be the end point.
3. Effect of pH: During a complexometric titration, the pH must be constant by 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.
Equation below shows complexation between metal ion and H+ ion for ligand:
Thus, stability of metal complex is pH dependent. Lower the pH of the solution, lesser would be the stability of 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 acidic solution, and metals forming weak complexes can only be effectively titrated in alkaline solution.
Complexometric EDTA Titration Curves
The complexometric EDTA titration curve shows the change in pM, where M is the metal ion, as a function of the volume of EDTA.
Calculating the Titration Curve. As an example, let’s calculate the titration curve for 50.0 mL of 5.00 ´10–3 M Cd2+ with
Since the titration is carried out at a pH of 10, some of the EDTA is present in forms other than Y4–. In addition, the presence of NH3 means that the EDTA must compete for the Cd2+. To evaluate the titration curve, therefore, we must use the appropriate conditional formation constant
Kf˝ = ´ ´ Kf = (0.35)(0.0881)(2.9 ´ 1016) = 8.9 ´ 1014
Because Kf˝ is so large, we treat the titration reaction as though it proceeds to completion.
The first task in calculating the titration curve is to determine the volume of EDTA needed to reach the equivalence point. At the equivalence point we know that
Moles EDTA = Moles Cd2+
or
MEDTAVEDTA = MCdVCd
Solving for the volume of EDTA
shows us that 25.0 mL of EDTA is needed to reach the equivalence point.
Before the equivalence point, Cd2+ is in excess, and pCd is determined by the concentration of free Cd2+ remaining in solution. Not all the untitrated Cd2+ is free (some is complexed with NH3), so we will have to account for the presence of NH3. For example, after adding 5.0 mL of EDTA, the total concentration of Cd2+ is
To calculate the concentration of free Cd2+ we use equation
[Cd2+] = ´CCd = (0.0881)(3.64 ´10–3 M) = 3.21 ´10–4 M
Thus, pCd is
pCd = –log[Cd2+] = –log(3.21 ´10–4) = 3.49
At the equivalence point, all the Cd2+ initially present is now present as CdY2–. The concentration of Cd2+, therefore, is determined by the dissociation of the CdY2– complex. To find pCd we must first calculate the concentration of the complex.
Letting the variable x represent the concentration of Cd2+ due to the dissociation of the CdY2– complex, we have
Once again, to find the [Cd2+] we must account for the presence of NH3; thus
[Cd2+] = aCd2+ ´CCd = (0.0881)(1.93 ´10–9 M) = 1.70 ´10–10 M
giving pCd as 9.77.
After the equivalence point, EDTA is in excess, and the concentration of Cd2+ is determined by the dissociation of the CdY2– complex. Examining the equation for the complex’s conditional formation constant, we see that to calculate CCd we must first calculate [CdY2–] and CEDTA. After adding 30.0 mL of EDTA, these concentrations are
Substituting these concentrations into equation and solving for CCd gives
Thus,
[Cd2+] = ´CCd = (0.0881)(5.60 ´10–15 M) = 4.93 ´10–16 M
and pCd is 15.31. Figure and Table show additional results for this titration.
Complexometric titration curve for 50.0 mL of 5.00 ´10–3 M Cd2+ with
Data for Titration of 5.00 ´ 10–3 M Cd2+ with
at a pH of 10.0 and in the Presence of
|
Volume of EDTA |
(mL) pCd |
|
0.00 |
3.36 |
|
5.00 |
3.49 |
|
10.00 |
3.66 |
|
15.00 |
3.87 |
|
20.00 |
4.20 |
|
23.00 |
4.62 |
|
25.00 |
9.77 |
|
27.00 |
14.91 |
|
30.00 |
15.31 |
|
35.00 |
15.61 |
|
40.00 |
15.78 |
|
45.00 |
15.91 |
|
50.00 |
16.01 |
Methods of End Point Detection
End point in complexometric titration can be detected by the following two methods:
1. Indicators: The end point in complexometric titrations is shown by means of pM indicators. The concept of pM arises as follows:
If K is the stability constant,
then, [M] = [MX]/[X]K
or log [M] = log [MX]/[X] – log K
and pM = log [X]/[MX] – pK
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 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 which is capable of acting as a chelating agent to give a dye-metal complex. The latter is different in colour from the dye itself and also has a low stability constant than the chelate-metal complex. The colour of the solution, therefore, remains that of the dye complex until the end point, 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.
Over 200 organic compounds form colored 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–
– Compound must be chemically stable throughout the titration.
– It should form 1:1 complex which must be weaker than the metal chelate complex.
– Colour of the indicator and the metal complexed indicator must be sufficiently different.
– Colour reaction should be selective for the metal being titrated.
– The indicator should not compete with the EDTA.
Mechanism of action of indicator: Let the metal be denoted by M, indicator by I and chelate by EDTA. At the onset of the titration, the reaction medium contains the metal-indicator complex (MI) and excess of metal ion. When 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 course of the titration is chelating the free metal ions in solution at the expense of the MI complex. Finally, at the end point, 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:
Structures of some important indicators used in complexometric titrations are given in Fig. 2. Many compounds have been used as indicators (Table-1), like:
– Triphenyl methane dyes
– Phthalein and substituted phthaleins
– Azo dyes
– Phenolic compounds
2. Instrumental methods of End point 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 more dilute solution by spectrophotometric than by visual methods. Thus, in disodium EDTA titrations an accurate end point can be obtained using 0.001M 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 end point 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 complexed. This is the end point 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 = E0 + loge [Ox]/[Red]
where, E = the potential of the electrode
E0 = the standard electrode potential
[Ox] = activity of the ions in the oxidized state
[Red] = activity of the 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.
– High frequency titrator: This method is particularly suitable for dilute solutions, in some cases with concentrations as low as 0.0002M. The ions may be titrated directly in buffered solution or excess reagent can be added to the unbuffered solution and the liberated protons titrated with standard alkali. Since buffer solution and other extraneous electrolytes reduce the sensitivity of the titration, their concentration must be kept to a minimum.
Selecting and Evaluating the End Point
The equivalence point of a complexation titration occurs when stoichiometrically equivalent amounts of analyte and titrant have reacted. For titrations involving metal ions and EDTA, the equivalence point occurs when CM and CEDTA are equal and may be located visually by looking for the titration curve’s inflection point.
As with acid–base titrations, the equivalence point of a complexation titration estimated by an experimental end point. A variety of methods have been used to find the end point, including visual indicators and sensors that respond to a change in the solution conditions. Typical examples of sensors include
1) recording a potentiometric titration curve using an ion-selective electrode (analogous to measuring pH with a pH electrode),
2) monitoring the temperature of the titration mixture,
3) and monitoring the absorbance of electromagnetic radiation by the titration mixture.
Finding the End Point with a Visual Indicator. Most indicators for complexation titrations are organic dyes that form stable complexes with metal ions. These dyes are known as metallochromic indicators. To function as an indicator for an EDTA titration, the metal–indicator complex must possess a colour different from that of the uncomplexed indicator. Furthermore, the formation constant for the metal–indicator complex must be less favourable than that for the metal–EDTA complex. The complex are often intensely coloured and are discernible to the eye at concentrations in the range at 10–6 to 10–7 M.
The indicator, Inm–, is added to the solution of analyte, forming a coloured metal–indicator complex, MInn-m. As EDTA is added, it reacts first with the free analyte, and then displaces the analyte from the metal–indicator complex, affecting a change in the solution’s colour. The accuracy of the end point depends on the strength of the metal–indicator complex relative to that of the metal–EDTA complex. If the metal–indicator complex is too strong, the colour change occurs after the equivalence point. If the metal–indicator complex is too weak, however, the end point is signalled before reaching the equivalence point.
Eriochrome Black T is a typical metal-ion indicator that is used in the titration of several common cations. Eriochrome Black T forms red complexes with more than two twenty metal ions, but the formation constant of only a few are appropriate for end-point detection. Except, Eriochrome Black T behaves as an acid-base indicator as well as metal ion indicator:
The metal complexes of Eriochrome Black T are generally red. Until the equivalence point in a titration, the indicator complexes the excess metal ion, so the solution is red. When EDTA becomes present in slight excess, the solution turns blue as a consequence of the reaction:
MIn– + HY3– « HIn2– + MY2– (Yn– – EDTA ions)
red blue
A limitation of Eriochrome Black T is that its solutions decompose slowly with standing.
Most metallochromic indicators also are weak acids or bases. The conditional formation constant for the metal–indicator complex, therefore, depends on the solution’s pH. This provides some control over the indicator’s titration error. The apparent strength of a metal–indicator complex can be adjusted by controlling the pH at which the titration is carried out. Unfortunately, because they also are acid–base indicators, the colour of the uncomplexed indicator changes with pH. For example, calmagite, which we may represent as H3In, undergoes a change in colour from the red of H2In– to the blue of HIn2– at a pH of approximately 8.1, and from the blue of HIn2– to the red-orange of In3– at a pH of approximately 12.4. Since the colour of calmagite’s metal–indicator complexes are red, it is only useful as a metallochromic indicator in the pH range of 9–11, at which almost all the indicator is present as HIn2-.
A partial list of metallochromic indicators, and the metal ions and pH conditions for which they are useful, is given in Table. Even when a suitable indicator does not exist, it is often possible to conduct an EDTA titration by introducing a small amount of a secondary metal–EDTA complex, provided that the secondary metal ion forms a stronger complex with the indicator and a weaker complex with EDTA than the analyte. For example, calmagite can be used in the determination of Ca2+ if a small amount of Mg2+–EDTA is added to the solution containing the analyte. The Mg2+ is displaced from the EDTA by Ca2+, freeing the Mg2+ to form the red Mg2+–indicator complex. After all the Ca2+ has been titrated, Mg2+ is displaced from the Mg2+–indicator complex by EDTA, signaling the end point by the presence of the uncomplexed indicator’s blue form.
Selected Metallochromic Indicators
|
Indicator |
Useful pH Range |
Useful for |
|
Calmagite |
9–11 |
Ba, Ca, Mg, Zn |
|
Eriochrome Black T |
7.5–10.5 |
Ba, Ca, Mg, Zn |
|
Eriochrome Blue Black R |
8–12 |
Ca, Mg, Zn, Cu |
|
Murexide |
6–13 |
Ca, Ni, Cu |
|
PAN |
2–11 |
Cd, Cu, Zn |
|
Salicylic acid |
2–3 |
Fe |
Quantitative Applications
With a few exceptions, most quantitative applications of complexation titrimetry have been replaced by other analytical methods.
Selection and Standardization of Titrants. EDTA is a versatile titrant that can be used for the analysis of virtually all metal ions. Although EDTA is the most commonly employed titrant for complexation titrations involving metal ions, it cannot be used for the direct analysis of anions or neutral ligands. In the latter case, standard solutions of Ag+ or Hg2+ are used as the titrant.
Solutions of EDTA are prepared from the soluble disodium salt, Na2H2Y ´ 2H2O. Concentrations can be determined directly from the known mass of EDTA; however, for more accurate work, standardization is accomplished by titrating against a solution made from the primary standard CaCO3. Solutions of Ag+ and Hg2+ are prepared from AgNO3 and Hg(NO3)2, both of which are secondary standards. Standardization is accomplished by titrating against a solution prepared from primary standard grade NaCl.
Inorganic Analysis. Complexation titrimetry continues to be listed as a standard method for the determination of hardness, Ca2+, CN–, and Cl– in water and wastewater analysis. The determination of Ca2+ is complicated by the presence of Mg2+, which also reacts with EDTA. To prevent an interference from Mg2+, the pH is adjusted to 12–13, precipitating any Mg2+ as Mg(OH)2. Titrating with EDTA using murexide or Eriochrome Blue Black R as a visual indicator gives the concentration of Ca2+.
Titration with inorganic complexing agents. Complexometric titrations with inorganic reagents are among the oldest volumetric methods.
Cyanide is determined at concentrations greater than 1 ppm by making the sample alkaline with NaOH and titrating with a standard solution of AgNO3, forming the soluble Ag(CN)2– complex. The end point is determined using p-dimethylaminobenzalrhodamine as a visual indicator, with the solution turning from yellow to a salmon colour in the presence of excess Ag+.
Now sometime is used the titration of halide ions with mercury(II) ions, called mercurimetry. Chloride is determined by titrating with Hg(NO3)2, forming soluble HgCl2:
Hg(NO3)3 + 2NaCl = HgCl2 + 2NaNO3
The sample is acidified to within the pH range of 2.3–3.8 where diphenylcarbazone,
which forms a coloured complex with excess Hg2+, serves as the visual indicator, or sodium nitroprousside Na3[FeNO(CN)5]. Xylene cyanol FF is added as a pH indicator to ensure that the pH is within the desired range. The initial solution is a greenish blue, and the titration is carried out to a purple end point.
Types of Complexometric Titrations
Complexometric titrations are of 4 types:
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 end point 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).
Limitations -slow complexation reaction
-Interference due to presence of other ions
2. 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 of 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 blackT as indicator. This method is analogous to back titration method in acidimetry. e.g.- ZnO
3. 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 end points, the metal may be determined by the displacement of an equivalent amount of Mg or Zn from a less stable EDTA complex.
Mn displaces Mg from Mn EDTA solution. The freed Mg metal is then directly titrated with a standard EDTA solution. In this method, excess quantity of Mg EDTA chelate is added to 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 Eriochrome blackT indicator.
4. 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 a heavy metal and titrated with sodium alkali.
solution by mercuric ions as 1:1 complex.
Method: Barbiturate to be analyzed is taken in a flask and heated with excess of mercury in alkaline solution. When precipitated Hg-barbiturate complex is formed, it is filtered and dissolved in excess of standard EDTA solution. The unreacted EDTA solution is then back titrated with a standard Zn solution.
Some important elements which could be determined by complexometric titration are as follows:
i) 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.
ii) 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 which 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 the stability constants of the 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
Quantitative Calculations. The stoichiometry of complexation reactions is given by the conservation of electron pairs between the ligand, which is an electron-pair donor, and the metal, which is an electron-pair acceptor; thus
This is simplified for titrations involving EDTA where the stoichiometry is always 1:1 regardless of how many electron pairs are involved in the formation of the metal–ligand complex.
Definition of the calcium chloride in 10 % calcium chloride solution
10,00 ml of a investigated solution place in a measured flask spaciousness of 100,0 ml and dilute volume of this solution by distilled water to a label, mix. 10.00 ml of the received solution place in a conic flask, add 15 ml of distilled water, 5 ml of an ammonia buffer solution,
To start of titration:
CaCl2 + H2Ind + 2NH3 = CaInd + 2NH4Cl
To end-point:
CaCl2 + Na2H2Y + 2NH3 = CaNa2Y + 2NH4Cl
or the ionic equation
Ca2+ + H2Y2- + 2NH3 = CaY2- + 2NH4+
In the end-point:
CaInd + Na2H2Y + 2NH3 = CaNa2Y + 2NH4+ + Ind2-
Or the ionic equation
CaInd + H2Y2- + 2NH3 = CaY2- + 2NH4+ + Ind2-
Calculate mass of calcium chloride CaCl2×2H2O in 1 ml of an investigated solution:
In 1 ml of a preparation should be 0,097-
Definition of the calcium gluconate in tablets of calcium gluconate 0,5
Nearby
To start of titration:
C12H22CaO14 + H2Ind + 2NH3 = CaInd + (NH4)2C12H22O14
To end-point:
C12H22CaO14+ Na2H2Y + 2NH3 = CaNa2Y + (NH4)2C12H22O14
or the ionic equation
Ca2+ + H2Y2- + 2NH3 = CaY2- + 2NH4+
In the end-point:
CaInd + Na2H2Y + 2NH3 = CaNa2Y + 2NH4+ + Ind2-
Or the ionic equation
CaInd + H2Y2- + 2NH3 = CaY2- + 2NH4+ + Ind2-
Calculate mass of calcium gluconate C12H22CaO14·H2O in 1 tablet of an investigated tablet:
where b – average mass of the one of tablets, which calculate from the weight of the 10 tablets.
Content of calcium gluconate С12Н22СаО14×Н2О in one tablet should be 0,475-
Definition of the total water’s hardness
100,0 ml of the investigated sample of water place in a conic flask spaciousness of 250 ml, add 5 ml of an ammoniac buffer solution,
To start of titration:
Ca2+ + H2Ind + 2NH3 = CaInd + 2NH4+
Mg2+ + H2Ind + 2NH3 = MgInd + 2NH4+
To end-point:
Ca2+ + H2Y2- + 2NH3 = CaY2- + 2NH4+
Mg2+ + H2Y2- + 2NH3 = MgY2- + 2NH4+
In the end-point:
CaInd + H2Y2- + 2NH3 = CaY2- + 2NH4+ + Ind2-
MgInd + H2Y2- + 2NH3 = MgY2- + 2NH4+ + Ind2-
Total hardness of water (Н (Н2О)) express total quantity of calcium and magnesium cations in mmol into one litre of water:
Use of masking and demasking agents:
Masking agents act either by precipitation or by formation of complexes more stable than the interfering ion-EDTA complex.
a) 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-hydroxy quinoline for many heavy metals. Thioglycerol (CH2SH.CHOH.CH2OH) is used to mask Cu by precipitation in the assay of lotions containing Cu and Zn.
b) 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.
– Dimercaprol (2,3-Dimercaptopropanol); (CH2SH.CHSH.CH2OH). Cations of mercury, cadmium, zinc, arsenic, tin, lead and bismuth react with dimercaprol in weakly acidic solution to form precipitates which are soluble in alkaline solution.
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 alkaline solution which 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.
– Potassium iodide is used to mask the mercury(II) ion (HgI4)2- and is specific as for mercury. It can be used in the assay of mercury(II) chloride.
– Tiron (disodium catechol-3,5-disulphonate) will mask aluminium and titanium as colourless complexes. Iron forms highly coloured complexes and is best masked as its hexacyanoferrate(II) complex.
– Triethanolamine [N(CH2.CH2.OH)3] forms a colourless complex with aluminium, a yellow complex with iron(III), the colour of which is almost discharged by adding sodium hydroxide solution, and a green manganese(III) complex which oxidizes mordant black II. For these reasons, if murexide is used in the presence of iron and manganese it is best to amsk them with triethanolamine; similarly, mordant black II can be used in the presence of triethanolamine-aluminium complex.
pH control Method: The formation of a metal chelate is dependent on the pH of the reaction medium. In 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 acidic solution, Bi can be effectively titratedwith 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.
Classical separation: These may be applied if they are not tedious; thus the following precipitates may not be used for separations in which, after being re-dissolved, the cations can be determined complexometrically: CaC2O4, nickel dimethylglyoximate, Mg(NH4)PO4, 6H2O, and CuSCN.
Solvent extraction: This is occasionally of value. Thus, Zinc can be separated from copper and lead by adding excess of ammonium thiocyanate solution and extracting the resulting zinc thiocyanate with 4-methylpentan-2-one (isobutyl methyl ketone); the extract is diluted with water and the zinc content determined with EDTA solution.
Removal of Anions: Anions, such as orthophosphate, which can interfere in complexometric titrations, may be removed using ion exchange resins.
Kinetic masking: This is a special case in which a metal ion does not effectively enter into the complexation reaction because of its kinetic inertness. Thus the slow reaction of chromium (III) with EDTA makes it possible to titrate other metal ions which react rapidly, without interference from Cr (III); this is illustrated by the determination of iron (III) and chromium (III) in a mixture.
Applications of Complexometric Titrations
Complexometric titrations have been employed with success for 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 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.
Assay of CaCO3: Accurately weighed amount of CaCO3 is dissolved in water and then acidified with HCl. A mixture of naphthol green and murexide is then added and titrated with EDTA, kept in burette.
1ml of M/20 disodium EDTA ≡
Calcium Lactate tablets: 20 tablets are finely powdered and an accurately weighed amount of the powder, representing about 0.5gm of calcium lactate, is transferred to a crucible, ignited until free from carbon and then cooled. 10 ml water is added and the residue is dissolved by adding dropwise dil. HCl solution.this solution is then transferred to a container, diluted to 150 ml with water and the assay is completed as is given under general procedure.
1ml of M/20 disod. EDTA ≡ 0.01542gm of Ca lactate
Calcium Lactate injection: Measure out a suitable volume of the injection, equivalent to about 0.5gm of Ca lactate. Transfer to the titration flask and proceed as given under general procedure.
Calcium Gluconate: An accurately weighed quantity (0.8gm) is dissolved in water (150ml) containing dil HCl (5ml). to the acidified solution is added, solution of NaOH (15ml), murexide indicator (4mg), solution of naphthol green (3ml). The reaction mixture is titrated with M/20 disod. EDTA until the solution is deep blue in colour.
1ml of M/20 disod. EDTA ≡ 0.02242gm of Ca gluconate
Calcium Gluconate injection: An accurately measured volume of the injection, equivalent to 0.8gm of Calcium gluconate is taken in a titration flask and proceeded as above.
Calcium Gluconate tablet: 20 Tablets are finely powdered. An accurately weighed amount of the powder, equivalent to 0.8gm of Calcium gluconate is transferred to a crucible and proceeded as described under Calcium lactate tablets.
Determination of Magnesium: Dissolve an accurately weighed sample (75mg) of Mg in sufficient water to make 100ml. pipette out 50ml of this solution in a titration flask, add 50ml water, 5ml of NH3 buffer solution and a few drops of eriochrome blackT as indicator. Titrate it to a deep blue colour.
Each ml of M/20 disodium EDTA ≡
This method could be used for the assay of Mg stearate and Mg sulphate.
Determination of Hardness of Water: Water hardness due to Ca and Mg is expressed as the amount of Ca and Mg ions in ppm. Actually, the hardness is due to both Ca and Mg salts but he two are determined together in the titration. The total Ca and Mg is titrated with standard EDTA solution using eriochrome blackT as indicator.
Method: Disodium salt of EDTA has the general formula: Na2H2Y2.H2O, where Y is the tetravalent anion of EDTA. When Ca is titrated with H2Y2-, a very stable complex is formed.
Mg2+ forms a similar complex which is far less stable than the Ca complex. When a sample containing Ca and Mg ions is titrated with a solution of EDTA, the Ca2+ are first complexed as CaY2-. As more reagent is added, all the Ca2+ is combined as complex. Mg ion forms MgY2-. The desired end point if the titration is the point at which all the Ca and Mg ions of the solution have combined with the complexing agent.
Titration is a process in which a standard reagent (titrant) is added to a solution of an analyte until the reaction between the analyte and reagent is judged to be complete.
Titration
http://www.youtube.com/watch?v=9DkB82xLvNE
Titration can be:
1) direct titration – titrant add to an analyte solution and react with determined substrance;
2) back-titration – is a process in which the excess of a standard solution used to react with an analyte is determined by titration with a second standard solution. Back-titrations are required when the reagent is slow or when the standard solution lacks stability. For example:
CaCO3 + HCl = CaCl2 + H2O + CO2
surplus (titrant 1)
HCl + NaOH = NaCl + H2O
residue titrant 2
3) substitute-titration – is a process in which a standard solution used to react with an additional (substitute) substance, amount of which is equivalent an analyte amount. Substitute-titrations are required when the analytes are unstable substance or when is impossible to indicate the equivalent (end) point in direct reaction. For example:
CrCl2 + FeCl3 = CrCl3 + FeCl2
analyte substitute
5FeCl2 + KMnO4 + HCl = 5FeCl3 + KCl + MnCl2 + 4H2O
Equivalence point is the point where sufficient titrant has been added to be stoichiometrically equivalent to the amount of analyte. The equivalence point of a titration is a theoretical point that caot be determined experimentally, but can be determined experimentally the end point.
End point is the point in a titration when a physical change that is associated with the condition of chemical equivalence occurs.
We can estimate its position by observing some physical changes with various indicating techniques:
a) without any special means. The visible changes occur in titrated solution – change of titrant or analyte colour, turbidity arise, precipitation formation;
b) with internal indicator using. The special chemical substances called indicators are added to the analyte solution. Typical indicator changes include the appearance or disappearance of a colour, a change in colour, or the appearance or disappearance of turbidity;
c) with instruments. This instruments respond to certain properties of the solution that change in a characteristic way during the titration.
The difference in volume between the equivalence point and the end point is the titration error.
A standard solution (or titrant) is a reagent of exactly known concentration that is used in a titrimetric analysis. Standard solutions are the main participants in all titrimetric methods of analysis. The titrant solutions must be of known composition and concentration. Ideally, we would like to start with a primary standard material.
Primary standard is an highly purified compound that serves as the reference materials for a titrimetric method of analysis. Important requirements for a primary standard are:
1. High purity.
2. Stability toward air.
3. Absence of hydrate water so that the composition of the solid does not change with variations in relative humidity.
4. Ready availability at modest cost.
5. Reasonable solubility in the titration medium.
6. Reasonable large molar mass so that the relative error associated with weighing the standard is minimised.
A secondary standard is compound whose purity has been established by chemical analysis and serves as the reference material to a titrimetric method of analysis.
stopcock buret for standartization
Gay-Lussac burette Buret with bottle of standard solution
Mohr burette Buret with rubber shutter
Microburet:
а) Shilov air-powered Buret; б) stopcock buret
The concentration of the standard solutions can be established by two basic methods:
1. Direct method – a carefully weighed quantity of a primary standard is dissolved in a suitable solvent and diluted to an exactly known volume in a volumetric flask. A made solution is referred to as a primary standard solution (titrant).
Volumetric flask — for preparing liquids with volumes of high precision. It is a flask with an approximately pear-shaped body and a long neck with a circumferential fill line.
2. Standardisation – concentration of a volumetric solution (titrant) is detrmined by using to titrate
1) a weighed quantity of a primary standard,
where:
CN and V are concentration and volume of secondary standard solution
m and Em are mass and equivalent weight of primary standard
2) “standard titrimetric substance” (primary standard),
More often in an ampoule contains 0,1 mol (0,1 equivalents) of substances, it is necessary for preparation of 0,1 mol/L solution.
3) a measured volume of another standard solution.
where:
CN2 and V2 are concentration and volume of secondary standard solution
CN1 and V1 are concentration and volume of primary standard solution
A titrant that is standardised against a secondary standard or against another standard solution is referred to as a secondary standard solution (titrant).
Equivalence Points and End Points
For a titration to be accurate we must add a stoichiometrically equivalent amount of titrant to a solution containing the analyte. We call this stoichiometric mixture the equivalence point. Unlike precipitation gravimetry, where the precipitant is added in excess, determining the exact volume of titrant needed to reach the equivalence point is essential. The product of the equivalence point volume, Veq, and the titrant’s concentration, CT, gives the moles of titrant reacting with the analyte.
Knowing the stoichiometry of the titration reaction(s), we can calculate the moles of analyte.
Unfortunately, in most titrations we usually have no obvious indication that the equivalence point has been reached. Instead, we stop adding titrant when we reach an end point of our choosing. Often this end point is indicated by a change in the color of a substance added to the solution containing the analyte. Such substances are known as indicators. The difference between the end point volume and the equivalence point volume is a determinate method error, often called the titration error. If the end point and equivalence point volumes coincide closely, then the titration error is insignificant and can be safely ignored. Clearly, selecting an appropriate end point is critical if a titrimetric method is to give accurate results.
Units of concentration of standard solutions
The concentration of standard solutions (titrants) are generally expressed in units of either molarity (CM, or M) or normality (CN, or N).
Molarity (M) – is the number of moles of a material per liter of solution.
Normality (N) – is the number of species equivalents per liter of solution.
Sometime is used also one unite of concentration – titer (T). Titer established the relationship between volume of titrant and amount of analysed substance present. The most commonly titer is in units of mg analysed substance per ml of titrant. This system was developed to assist in doing routine calculations. It reduces the amount of time and training for technicians.
Equivalents law
Titrimetry is based on equivalents law:
or number of analyte equivalent present = number of standard reagent added,
or one equivalent of one material will react exactly with one equivalent of another
The weight of one equivalent of a compound depends on reference to a chemical reaction in which that compound is a participant. Similarly, the normality of a solution caever be specified without knowledge about how the solution will be used. Equivalent value is based on the type of reaction and the reactants:
1. One equivalent weight of a substance participating in a neutralisation reaction is that amount of substance that either react with or supplied one mol of hydrogen ions in that reaction.
2. One equivalent weight of a participant in an oxidation-reduction reaction is that amount that directly or indirectly produces or consumer one mol of electrons.
3. The equivalent weight of a participant in a precipitation or a complex-formation reaction is that weight which or provides one mole of the univalent reacting cation.
Volume as a Signal
Almost any chemical reaction can serve as a titrimetric method provided that three conditions are met. The first condition is that all reactions involving the titrant and analyte must be of known stoichiometry. If this is not the case, then the moles of titrant used in reaching the end point cannot tell us how much analyte is in our sample. Second, the titration reaction must occur rapidly. If we add titrant at a rate that is faster than the reaction’s rate, then the end point will exceed the equivalence point by a significant amount. Finally, a suitable method must be available for determining the end point with an acceptable level of accuracy. These are significant limitations and, for this reason, several titration strategies are commonly used. A simple example of a titration is an analysis for Ag+ using thiocyanate, SCN–, as a titrant.
This reaction occurs quickly and is of known stoichiometry. A titrant of SCN– is easily prepared using KSCN. To indicate the titration’s end point we add a small amount of Fe3+ to the solution containing the analyte. The formation of the redcolored Fe(SCN)2+ complex signals the end point. This is an example of a direct titration since the titrant reacts with the analyte.
If the titration reaction is too slow, a suitable indicator is not available, or there is no useful direct titration reaction, then an indirect analysis may be possible. Suppose you wish to determine the concentration of formaldehyde, H2CO, in an aqueous solution. The oxidation of H2CO by I3–
is a useful reaction, except that it is too slow for a direct titration. If we add a known amount of I3–, such that it is in excess, we can allow the reaction to go to completion.
The I3– remaining can then be titrated with thiosulfate, S2O32–.
This type of titration is called a back titration.
Calcium ion plays an important role in many aqueous environmental systems. A useful direct analysis takes advantage of its reaction with the ligand ethylenediaminetetraacetic acid (EDTA), which we will represent as Y4–.
Unfortunately, it often happens that there is no suitable indicator for this direct titration. Reacting Ca2+ with an excess of the Mg2+–EDTA complex
releases an equivalent amount of Mg2+. Titrating the released Mg2+ with EDTA
gives a suitable end point. The amount of Mg2+ titrated provides an indirect measure of the amount of Ca2+ in the original sample. Since the analyte displaces a species that is then titrated, we call this a displacement titration.
When a suitable reaction involving the analyte does not exist it may be possible to generate a species that is easily titrated. For example, the sulfur content of coal can be determined by using a combustion reaction to convert sulfur to sulfur dioxide.
Passing the SO2 through an aqueous solution of hydrogen peroxide, H2O2,
produces sulfuric acid, which we can titrate with NaOH,
providing an indirect determination of sulfur.
Calculations in titrimetric method of analysis
|
T = |
N = |
m =
|
|
mx(is) = |
mx(al) = |
ax = |
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T – titer (g/ml); |
N – normality (number of equivalents/l); |
|
Nt – nomality of used titrant (N); |
Vt – volume of used titrant (ml); |
|
m – mass of substance (g); |
meq – mass of one equivalent (g); |
|
mx(al) – amount of analyte, determined as aliquot of sample (g); |
mx(is) – amount of analyte, determined as individual sample (g); |
|
meqx – mass of one equivalent of analyte (g); |
W – dilution of analyte sample (ml); |
|
Vs – aliquot of sample solution (ml); |
px – mass of sample (g); |
|
ax – percentage of substance in sample (%) |
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Indicators of Titrimetry Methods
Indicators are the chemical compounds, which give some external effect attached to concentrations of reactive species according to equivalence point. This external effect can be accompanied by change, appearance or disappearance of colouring, and formation of slightly soluble compounds (precipitate formation).
On appliance technique indicators are external and internal.
Internal indicators are introduced into titrated solution. An end point install on changes of colour of analysed mixture.
The external indicators are used when internal indicators using is impossible. Reaction with external indicators runs out of analysed mixture. Some drops of analysed solution put on peace of filter paper, impregnated with indicator, or mix with drop of indicator solution on porcelain plate.
For effect the reactions appearance indicators are reversible and unreversible.
Reversible indicators – changes the colour can be repeated many times as changes the system state.
Unreversible indicators – colour changes ones with destruction of indicator molecule. The unreversible indicators are less comfortable and thinly use.
