During complexation reactions the concentration of the analyte ion (for example, a metal ion) changes most rapidly at the end point. As noted in Topic C6, the most widely used complexing agent is ethylenediaminetetracetic acid or EDTA, and Table 1gives a selection of metal EDTA formation constants.
Using the values of a4 given in Topic C6, Table 1, we may calculate the practical, or conditional, formation constant at a particular pH
K’MY = KMYa4
From the data in the tables, it can be calculated that magnesium could be titrated at pH 10, but not at low pH. This has already been discussed in Topic C6.
Complexation titrations
Key Notes
The techniques of titrimetry, using both visual and potentiometric end point detection, are used to measure species, particularly metal ions, in a wide range of samples.
Reactions producing an insoluble product are valuable analytical tools for the titrimetric determination of halide and other anions.
Oxidation and reduction titrations may be used to measure many species, especially metals in high or low valency states, iodine and iodides, and easily oxidized organic compounds.
Related topics Complexation, solubility and redox equilibria (C6) Redox titrations
Complexation titrations
Precipitation titrations
Table 1. Metal-EDTA formation constants at 25∞C
Cation KMY log (KMY)
Ag+ 2.0¥ 107 7.3
Mg2+ 4.9¥ 108 8.7
Ca2+ 5.0¥ 1010 10.7
Fe2+ 2.1¥ 1014 14.3
Fe3+ 1.0¥ 1025 25.1
Zn2+ 3.2¥ 1016 16.5
Cd2+ 2.9¥ 1016 16.5
V3+ 8.0¥ 1025 25.9
It is possible to titrate two cationic species in a solution by performing the titration at different pH values. However, if a solution of high pH must be used, this might cause precipitation of metal hydroxides or other insoluble species. In order to prevent this, secondary complexing agentscan be added to retain the metal ion in solution. Ammonium chloride and triethylamine are typical reagents for this purpose.
Zinc ions, which might otherwise form insoluble Zn(OH)2at pH of 10, may be converted to soluble zinc amine complexes. These are less stable than the EDTA complex and the zinc may then be reacted quantitatively.
Secondary complexing agents may also act as masking agents. Examples of this are the use of cyanide ions to form stronger complexes with heavy metal ions so that magnesium can be titrated, or masking Fe2+and Mn2+ using hydroxylamine in water hardness determinations.
Standard solutions of EDTA may be prepared from the dry disodium salt (Na2H2Y, RMM 336 or the dihydrate, RMM 372), by dissolving a known amount in water free of heavy metals. Alternatively, the solution may be standardized by a standard magnesium salt solution.
In complex and precipitation titrations, as in others, the end point corre- sponds to a rapid change in the concentration of species. This may be detected by instrumental methods, particularly potentiometry (see Topic C3) and by visual indicators discussed below. Using suitable indicators, or potentiometric measurements, it is possible to detect two or more end points.
Complexometric indicatorsbehave in a similar way to titrating complexing agents such as EDTA. They generally change color with pH, but one species,for example HIn2-, will react with excess metal ions Mn+:
HIn2-+Mn+=MIn(n-3)++H+
blue red
Some selected indicators for complexometric titrations are given in Table 2.
In order to use potentiometric methodsto study complexometric titrations, an electrode specific to the metal ion may be used (see Topics C3 and C5), for example, a copper ISE to follow the reaction of copper with EDTA.
Alternatively, a ‘J’-shaped electrode with a small mercury pool may be used together with a small amount of added Hg-EDTA complex. This acts in a similar way to the Class 2 electrodes, where the complexes determine the concentration of ions in contact with the mercury pool:
Mn++HY3-=MY(n-4)++H+ Hg2++HY3-=HgY2-+H+
Hg2++2e-=Hg
C7 – Titrimetry II: complexation, precipitation and redox titrations 91
Table 2. Indicators for complexometric titrations
Indicator Free color Complex color* Metal ions
Eriochrome black T Blue Red Ba, Cd, Ca, Pb, Zn
Pyrocatechol violet Yellow Blue Al, Bi, Cd, Co, Cu
Fe, Mg, Mn, Ni, Zn
Xylenol orange Yellow Red Bio, Cd, Pb, Th, Zn
Calcon carboxylic acid Blue Red Ca, Cd, Mg, Mn, Zn
*varies with metal and pH
The major application of complexation titrations is for the determination of the concentrations or amounts of metallic elements in water, food and other industrial samples.
Example
The calcium and magnesium ions in hard water may be determined. The solution is adjusted to pH 12 with NaOH, when Mg(OH)2is precipitated. The calcium is then titrated with EDTA using calcon carboxylic acid as indicator. Both calcium and magnesium are then determined in by titrating a sample with EDTA at pH 10 using eriochrome black T, and finding the magnesium by difference.
For precipitation reactions, the change in the concentration of either ion forming the precipitate may be considered. Since the changes often involve many orders of magnitude of concentration, it is again convenient to use the pX notation. For example, for the reaction of silver ions with chloride to form an insoluble silver chloride precipitate
Ag++Cl-=AgCl (s) the concentration may be expressed as:
pAg = -log (a(Ag+)) ~ -log (c (Ag+))
Figure 1 shows the pAg values in the titration of sodium chloride by silver nitrate as a function of the volume of silver nitrate added. This figure shows that, before the end point pAg is very high (that is, the concentration of silver ions is small) and changes little, because there is still an excess of chloride and the silver is almost completely removed as precipitate. After the end point, there is an excess of silver ions, the concentration increases and pAg decreases. In the region around the end point, where the amounts are nearly equal, the change in pAg with volume added is very large. If a mixture of iodide and chloride ions is titrated (dashed line), the iodide, which is less soluble, precipitates first and pAg is even higher than for chloride. Then the chloride precipitates. Both end points can be found.
As noted in Table 1of Topic A5, silver nitrate, sodium chloride and potassium chloride are primary standards for silver halide precipitation reactions. Other Precipitation
titrations
14 12 10 8 6 4 2 0
0 5 10 15 20 25 30 35
Volume AgNO3 added
pAg
(a) (b)
Fig. 1. Silver-halide titrations. (a) Chloride alone (solid line); (b) iodide plus chloride (dashed line).
precipitation titrations (e.g., barium with sulfate, zinc with ferrocyanide) are less commonly performed.
Indicators for silver-halide precipitation titrations are of two types. The first react specifically when an excess of titrant becomes present immediately after the end point -for example, if a small amount of potassium chromate is added, it will react with excess silver ions to produce deep red silver chromate in neutral solutions (Mohr’s method). In acid solutions, the silver is titrated with potassium thiocyanate (KCNS) solution (Volhard’s method). Iron (III) ammo- nium sulfate solution is added and reacts with an excesss of thiocyanate to produce a deep red iron thiocyanate species.
Adsorption indicatorssuch as fluorescein adsorb onto the precipitate when excess silver ions are present and the precipitate takes on a pinkish color.
As with other indicators, the change of color is detectable by eye over a range:
log (c(ion)) = ±1
A selection of indicators for precipitation titrations is given in Table 3.
Table 3. Indicators for precipitation titrations
Precipitation Color 1 Color 2 Ions detected
CrO42- Yellow Deep red Ag+
Fe3+ Light brown Deep red CNS-
Fluorescein Green-yellow Pink Ag+
It is considerably easier to titrate mixed chloride and bromide in solution by potentiometry. The use of a silver ion selective electrode, or even a silver wire, together with a double junction reference electrode, since the chloride ions from a calomel electrode would react, allows the determination of the silver ion concentration. Other precipitation titrations may be followed using suitable ion selective electrodes.
The major applications of precipitation titrations involve the determination of halides with silver, or the reverse, or the determination of silver in acid solutions with thiocyanate.
Example
An insecticide containing chlorine was digested in nitric acid to convert the chlorine to soluble chloride. Silver nitrate was added in excess, and the excess titrated with potassium thiocyanate by Volhard’s method. It is important to know whether all the chlorine is converted to chloride.
Redox titrations Oxidation-reduction or redoxtitrations are used for determining metals with two well-defined oxidation states,and indirect methods for the determination of organic compounds.
For redox reactionsthe concentrations both of the oxidized species, Ox, and of the reduced species, Red, will change simultaneously. Considering a cell with a redox electrode and a reference electrode:
SCE ||a (Ox), a(Red) |Pt the cell emf is given by:
E =En+(RT/F) ln (a (Ox)/a (Red)) -ESCE
C7 – Titrimetry II: complexation, precipitation and redox titrations 93
Therefore, as the reaction proceeds during titration, the ratio of the concentra- tions will change and the emf will alter. The potentiometric titration curve will resemble those described in Topic C5. A summary list of redox reagents is given in Table 3of Topic C6.
For standard solutions, sodium oxalate and iron(II) ammonium sulfate and potassium iodate are suitable, but potassium permanganate and iodine solu- tions decompose on standing and must be standardized before use.
The indicators for redox reactions are reagents whose oxidized and reduced forms differ in color:
In(Ox) (color 3) +ne-=In(Red) (color 4) An example of this is 1,10-phenanthroline iron (II)
[Fe (C12H8 N2)3]3++e-=[Fe (C12H8N2)3]2+
oxidized form, pale blue =reduced form, deep red
In several cases, the indicator reaction additionally involves hydrogen ions, so the change is pH dependent. Table 4 lists commonly used redox indicators.
Table 4. Indicators for redox titrations
Redox Oxidized Reduced EIn/V Solution
color color
1,10-phenanthroline iron(II) complex Pale blue Red 1.11 1 M H2SO4
Diphenylamine Violet Colorless 0.76 Dilute acid
Methylene blue Blue Colorless 0.53 1M acid
Phenosafranine Red Colorless 0.28 1M acid
For a redox indicator where one electron is involved, at 25∞C, the color change takes place at electrode potentials in the range
E =EIn ±0.059
One further useful indicator employed in redox titrations involving iodine is starch, or more synthetic equivalent materials. The starch forms a blue-black complex with iodine, which is rendered colorless when all the iodine has been removed.
The applications of redox titrations include the determination of metals, with two well-defined oxidation states, which are present in metallurgical samples and ores. In order to dissolve the material, it may be necessary to use oxidative conditions, for example, concentrated nitric acid. This will convert the majority of the ions into their higher oxidation state, and in order to titrate them they must first be reduced quantitatively. This may be done by passing the acidified solution through a Jones reductor, which contains a zinc-mercury amalgam.
The effluent may then be titrated using a suitable oxidant. Some organic compounds, such as phenols, may be determined by bromination with a bromate/bromide mixture, followed by back titration of the excess using thiosulfate.
Section C – Analytical reactions in solution