Complexometric Titrations
ANALYTICAL CHEMISTRY
FOR PHARMACY STUDENTS
COMPLEXOMETRY
Second Year
ALAA KHEDR, PH.D.
PHARMACEUTICAL CHEMISTRY DEPARTMENT
FACULTY OF PHARMACY
KING ABDULAZIZ UNIVERSITY
JEDDAH, KSA
0
Complexometric Titrations
COMPLEXOMETRIC
TITRATIONS
ANALYTICAL IMPORTANCE OF COMPLEXES:
Coordination complexes are neutral or ionic compounds that involve the formation of at
least one coordinate covalent bond between the metal ion and a complexing agent.
Coordination complexes play an important role in many analytical processes as shown by the
following examples:
- Many coordination compounds are insoluble in water and form the basis of rather specific
precipitation reactions in gravimetry.
- Complexing agents that form coordination complexes with certain specific metal ions may
often be used as masking agents to prevent a metal ion from undergoing an undesired
reaction that would result in an interference.
- Because coordination compounds are often highly colored, they are frequently used in
many colorimetric and photometric procedures.
- Furthermore, the solubility of many neutral coordination compounds in organic solvents
allows the extraction of metal ions from aqueous solution into the organic phase and thus
forms the basis for some analytical separations.
COMPLEXATION:
The processes of complex-ion formation can be described by the general term
complexation. A complexation reaction with a metal ion involves the replacement of one or
more of the coordinated solvent molecules by other nucleophilic groups. The groups bound to
the central ion are called ligands and in aqueous solution the reaction can be represented by
the equation:
M(H2O)n + L = M(H2O)(n-1) L + H2O
Here the ligand (L) can be either a neutral molecule of a charged ion, and successive
replacement of water molecules by other ligand groups can occur until the complex ML n is
1
Complexometric Titrations
formed; n is the coordination number of the metal ion and represents the maximum number
of monodentate ligands that can be bound to it.
Ligands may be conveniently classified on the basis of the number of points of
attachment to the metal ion. Thus simple ligands, such as halide ions or the molecules H 2O
or NH3, are monodentate, i.e. the ligand is bound to the metal ion at only one point by the
donation of a lone pair of electrons to the metal. When, however, the ligand molecule or ion
has two atoms, each of which has a lone pair of electrons, then the molecule has two donor
atoms and it may be possible to form two coordinate bonds with the same metal ion; such a
ligand is said to be bidentate and may be exemplified by consideration of the tris(ethylenediamine) cobat(III) complex, [Co(en)3]3+. In this six-coordinate complex of cobalt(III), each of
the bidentate ethylenediamine molecules is bound to the metal ion through the ion pair
electrons of the two nitrogen atoms. This results in the formation of three five-membered
rings, each including the metal ion; the process of ring formation is called chelation and like
this complex is called a chelate.
Multidentate ligands contain more than two coordinating atoms per molecule, e.g. 1,2diaminoethanetetra-acetic acid (ethylenediaminetetra-acetic acid, EDTA), which has two
donor nitrogen atoms and four donor oxygen atoms in the molecule, can be hexadentate.
3+
H2
N
H2
N
Co
N
H2
N
H2
N
H2
N
H2
O
O
CH2 C OH
HO C CH2
N CH2 CH2 N
CH2 C OH
HO C CH2
O
O
Tris(ethylenediamine)Cobalt(III) complex
EDTA
A ligand involved in complexation can be either charged (acidic or anionic functional
groups) or uncharged groups (Table 1).
2
Complexometric Titrations
In the foregoing it has been assumed that the complex species does not contain more
than one metal ion, but under appropriate conditions a binuclear complex, i.e. one containing
two metal ions, or even a polynuclear complex, containing more than two metal ions may be
formed. Thus, interaction between Zn2+ and Cl- ions may result in the formation of binuclear
complexes, e.g. [Zn2Cl6]2-, in addition to simple species such as ZnCl3- and ZnCl42-. The
formation of bi- and polynuclear complexes will clearly be favored by a high concentration of
the metal ion; if the latter is present as a trace constituent of a solution, polynuclear
complexes are unlikely to be formed.
Table 1. Typical functional groups involved in chelation
Acidic or Anionic Groups
Carboxylic acid
or
Oximes
O
O
C O
C OH
NO
NOH
or
Hydroxyles
R O R OH
or
Phenols
Sulfonic acids
or
or
O
O
S O
S OH
O
O
R S
Amines
Carbonyl
Thiocarbonyl
Nitroso
OH
O
or
Mercaptyl
Uncharged Groups
Cyclic amine
H
,
N
,
N
NH2
O
C
S
C
NO
N
R SH
STABILITY OF COMPLEXES:
The thermodynamic stability of a species is a measure of the extent to which this
species will be formed from other species under certain conditions, provided that the system
is allowed to reach equilibrium. Consider a metal ion M in solution together with a
monodentate ligand L, then the system may be described by the following stepwise
equilibria, in which, for convenience, coordinated water molecules are not shown:
M + L
ML + L
ML(n-1) + L
3
ML;
K1 = [ML] / [M] [L]
ML2;
MLn;
K2 = [ML2] / [M] [L]
Kn = [MLn] / [M(n-1)] [L]
Complexometric Titrations
The equilibrium constants K1, K2, ...., Kn are referred to as stepwise stability constants.
An alternative way of expressing the equilibrium is as follows:
M + L
ML;
1 = [ML] / [M] [L]
ML + 2L
ML2;
2 = [ML2] / [M] [L]2
ML + n L
MLn;
n = [MLn] / [M] [L]n
The equilibrium constants 1, 2, ...., n are called the overall stability constants and are
related to the stepwise stability constants by the general expression.
n = K1 x K2 x ... Kn
In the above equilibrium it has been assumed that no insoluble products are formed nor any
polynuclear species.
A knowledge of stability constant values is of considerable importance in analytical
chemistry, since they provide information about the concentration of the various complexes
formed by a metal in specific equilibrium mixtures; this is highly valuable in the study of
compleximetry, and of various analytical separation procedures such as solvent extraction,
ion exchange, and chromatography. The stability of a complex will obviously be related to (a)
the complexing ability of the metal ion involved, and (b) characteristics of the ligand.
CHELATE EFFECT:
The process of chelation highly affects the stability of the formed complexes. Thus,
multidentate ligands usually form stronger metal complexes than do similar monodentate
ligands. For example, the reaction of Cd2+ with two molecules of ethylenediamine has a much
larger equilibrium constant than its reaction with four molecules of methylamine.
4
Complexometric Titrations
2+
Cd
+ 2 H NCH CH NH
2
2
2
2
+ 4 CH NH
3
2
methylamine
K = 2 x 1010
N
H
N
H
H3C
2+
H
N
Cd
ethylenediamine
Cd
2+
H
N
NH2 CH3
NH2
2+
Cd
H3C
NH2
K = 3 x 106
NH2 CH3
Another example is given by the complexes formed by the nickel(II) ion with: (a) the
monodentate NH3 molecule, (b) the bidentate ethylenediamine (1,2-diaminoethane), and (c)
the hexadentate ligand 'penten' {H2N.CH2.CH2.)2N.CH2.CH2.N(CH2.CH2.NH2)2} show an overall
stability constant value for the ammonia complex (a) of 3.1x10 8, which is increased by a
factor of about 1010 for the complex of ligand (b) and is approximately ten times greater still
for the third complex (c).
The chelate effect, then, is the observation that multidentate ligands form more stable
metal complexes than do similar monodentate ligands. The chelate effect is most
pronounced for ligands such as EDTA or DCTA (trans-diaminocyclohexane-tetracetic acid),
which can occupy all six coordination sites about a metal ion.
COMPLEXONES:
The formation of a single complex species rather than the stepwise production of such
species will clearly simplify complexometric titrations and facilitate the detection of end
points. Schwarzenbach realized that the acetate ion is able to form acetate complexes of low
stability with nearly all polyvalent cations, and that if this property could be reinforced by the
chelate effect, then much stronger complexes would be formed by most metal cations. He
found that the aminopolycarboxylic acids are excellent complexing agents; the most
important of these is 1,2-diamino-ethanetetraacetic acid (ethylene-diaminetetraacetic acid,
H4Y). The formula (complexone III) is preferred to (complexone II), since it has been shown
5
Complexometric Titrations
from measurements of the dissociation constants that two hydrogen atoms are probably held
in the form of zwitterions. The value of pK are respectively pK1= 2.0, pK2= 2.7, pK3= 6.2 and
pK4= 10.3 at 20°C; these values suggest that it behaves as a dicarboxylic acid with two
strongly acetic groups and that there are two ammonium protons of which the first ionizes in
the pH region of about 6.3 and the second at a pH of about 11.5.
Various trivial names are used for ethylenediaminetetra-acetic acid and its sodium
salts, and these include Trilon B, Complexone III, Sequestrene, Versene, and Chelaton 3; the
disodium salt is most widely employed in titrametric analysis. To avoid the constant use of
the long name, the abbreviation EDTA is utilized for the disodium salt.
Other complexing agents (complexones) which are sometimes used include: (a)
nitrilotriacetic acid (NITA or NTA or Complexone I; this has pK1= 1.9, pK2= 2.5 and pK3= 9.7),
(b) trans-1,3-diaminocyclohexane-N,N,N',N'-tetra-acetic acid: this should presumably be
formulated as a zwitterion structure like (I); the abbreviated name is CDTA, DCyTA, DCTA or
Complexone IV, (c) 2,2'-ethylenedioxybis {ethyliminodi(acetic acid)} also known as ethylene
glycol-bis (2-aminoethyl ether) N,N,N',N'-tetra-acetic acid (EGTA or Complexone V), and (d)
triethylenetetramine-N,N,N',N",N'",N'"-hexa-acetic acid (TTHA or Complexone VI).
6
Complexometric Titrations
CH2
COOH
CH2 COO
+
HN
CH2
HOOC
HOOC
HOOC
-
OOC
CH2
CH2
N C C N
H2 H2
CH2
CH2
CH2
+
+
HN C C NH
H2 H2
CH2
CH2
CH2COOH
H2
C H N
H2C
C
CH2COOH
C
CH2COOH
C H N
H2
CH2COOH
N
COOH
Complexone II
CH2
H2C
Complexone I
COOH
COOH
COO
Complexone III
COOH
Complexone IV
CH2COO
CH2COOH
(CH2)2
O
(CH2)2
Complexone V
O
(CH2)2
N
CH2COOH
CH2COO
CH2COO
HN
CH2COOH
+
HN CH2COO
+
(CH2)2
(CH2)2
(CH)2
+
HN
+
HN
Complexone VI
CH2COO
CH2COOH
CH2COO
7
-
-
Complexometric Titrations
CDTA often forms stronger metal complexes than does EDTA and thus finds
applications in analysis, but the metal complexes are formed rather more slowly than with
EDTA so that the end-point of the titration tends to be drawn out with the former reagent.
EGTA finds analytical application mainly in the determination of calcium in a mixture of
calcium and magnesium and is probably superior to EDTA in the calcium / magnesium
water-hardness titration.
However, EDTA has the widest general application in analysis because of its powerful
complexing action and commercial availability. The spatial structure of its anion, which has
six donor atoms, enables it to satisfy the coordination number of six frequently encountered
among the metal ions and to form strainless five-membered rings on chelation. The resulting
complexes have similar structures but different from one another in the charge they carry
(Fig. 1).
2CO
CH2
O
CO
CH2
N
O
CH2
M
O
N
CO
CH2
CH2
CH2
O
CO
Fig.1 : Structure of a divalent metal-EDTA chelates.
To simplify the following discussion EDTA is assigned the formula H 4Y: the dissolution
salt is therefore Na2H2Y and affords the complex-forming ion H2Y2- in aqueous solution; it
reacts with all metals in 1:1 ratio. The reactions with cations, e.g. M2+, may be written as:
M2+ + H2Y2-
8
MY2- + 2H+ .............................................................. (1)
Complexometric Titrations
For other cations, the reactions may be expressed as:
M3+ + H2Y2-
MY- + 2H+ ............................................................... (2)
M4+ + H2Y2-
MY + 2H+ ................................................................ (3)
or
Mn+ + H2Y2-
(MY)(n-4)+ + 2H+ ................................................. (4)
One mole of the complex-forming H2Y2- reacts in all cases with one mole of the metal ion and
in each case, also, two moles of hydrogen ion are formed. It is apparent from equation (4)
that the dissociation of the complex will be governed by the pH of the solution; lowering the
pH will decrease stability of the metal-EDTA complex. The more stable the complex, the lower
the pH at which an EDTA titration of the metal ion in question may be carried out. Table (2)
indicates minimum pH values for the existence of EDTA complexes of some selected metals.
Table 2: Stability with respect to pH of some metal-EDTA complexes.
Minimum pH at
which complexes
exist
Selected metals
1-3
Zr4+; Hf4+; Th4+; Bi3+; Fe3+
4-6
Pb2+; Cu2+; Zn2+; Co2+; Ni2+, Mn2+; Fe2+; Al3+; Cd2+; Sn2+
8-10
Ca2+; Sr2+; Ba2+; Mg2+
It is thus seen that, in general, EDTA complexes with metal ions of the charge number 2
are stable in alkaline or slightly acidic solution, whilst complexes with ions of charge numbers
3 or 4 may exist in solutions of much higher acidity.
STABILITY CONSTANTS OF EDTA COMPLEXES:
The stability of a complex is characterized by the stability constant (or formation
constant) K:
Mn+ + Y4-
(MY)(n-4)+ ........................................................................... (5)
K = [(MY)(n-4)+] / [Mn+] [Y4-] ...................................................................................... (6)
Some values for the stability constants (expressed as log K) of metal-EDTA complexes
are collected in Table 3:
9
Complexometric Titrations
Table 3: Stability constants (as log K) of metal-EDTA complexes.
Mg2+
8.7
Zn2+
16.7
La3+
15.7
Ca2+
10.7
Cd2+
16.6
Lu3+
20.0
Sr2+
8.6
Hg2+
21.9
Sc3+
23.1
Ba2+
7.8
Pb2+
18.0
Ga3+
20.5
Mn2+
13.8
Al3+
16.3
In3+
24.9
Fe2+
14.3
Fe3+
25.1
Th4+
23.2
Co2+
16.3
Y3+
18.2
Ag+
7.3
Ni2+
18.6
Cr3+
24.0
Li+
2.8
Cu2+
18.8
Ce3+
15.9
Na+
1.7
THE "APPARENT" OR "CONDITIONAL" STABILITY CONSTANT:
In equation (6) only the fully ionized form of EDTA, i.e. the ion Y4-, has been taken into
account, but at low pH values the species HY3-, H2Y2-, H3Y- and even undissociated H4Y may
well be present; in other words, only a part of the EDTA uncombined with metal may be
present as Y4-. Further, in equation (6) the metal ion Mn+ is assumed to be uncomplexes, i.e.
in aqueous solution it is simply present as the hydrated ion. If, however, the solution also
contains substances other than EDTA which can complex with the metal ion, then the whole
of this ion uncombined with EDTA may no longer be present as the simple hydrated ion. Thus,
in practice, the stability of metal-EDTA complexes may be altered (a) by variation in pH and
(b) by the presence of other complexing agents. The stability constant of the EDTA complex
will then be different from the value recorded for a specified pH in pure aqueous solution; the
value recorded for the new conditions is termed the 'apparent' or 'conditional' stability
constant. It is clearly necessary to examine the effect of these two factors in some detail.
(A) PH EFFECT:
The apparent stability constant at a given pH may be calculated from the ratio K/,
where
 is the ratio of the total uncombined EDTA (in all forms) to the form Y4-. Thus KH, the
apparent stability constant for the metal-EDTA complex at a given pH, can be calculated from
the expression:
10
Complexometric Titrations
KH = K/
or:
log KH = log K - log  ............................................................................................. (7)
Figure (2) shows the effect of pH on the apparent stability constant values K' for metalEDTA chelates.
It is clear from this figure that the calcium chelate is too weak to be titrated in acid
solution, while mercury chelate is strong enough to be titrated in that medium.
At pH 13, all KH values are equal to Kf (formation or stability constant) because  = 1
and EDTA is completely dissociated to Y4-.
The value of log  is small at high pH values, and it therefore follows that the larger
values of log KH are found with increasing pH. However, by increasing the pH of the solution
the tendency to form slightly soluble metallic hydroxides is enhanced owing to the reaction:
(MY)(n-4)+
+ nOH-
M(OH)n
+ Y4-
The extent of hydrolysis of (MY)(n-4)+ depends upon the characteristics of the metal ion,
and is largely controlled by the solubility product of the metallic hydroxide and, of course, the
stability constant of the complex. Thus iron(III) is precipitated as hydroxide (K sol= 1x10-36) in
basic solution, but nickel(II), for which the relevant solubility product is 6.5x10-18, remains
complexed. Clearly the use of excess EDTA will tend to reduce the effect of hydrolysis in basic
solutions. It follows that for each metal ion there exists an optimum pH which will give rise to
a maximum value for the apparent stability constant.
(B) THE EFFECT OF OTHER COMPLEXING AGENTS (Z):
If another complexing agent (say NH3) is also present in the solution, then in equation
(6) [Mn+] will be reduced owing to complexation of the metal ions with ammonia molecules. It
is convenient to indicate this reduction in effective concentration by introducing a factor ,
defined as the ratio of the sum of the concentrations of all forms of the metal ion not
complexes with EDTA to the concentration of the simple (hydrated) ion. The apparent stability
11
Complexometric Titrations
constant of the metal EDTA complex, taking into account the effects of both pH and the
presence of other complexing agents, is then given by:
KHZ = K / .
log KHZ = log K - log  - log 
EDTA TITRATION CURVES:
If, in the titration of a strong acid, pH is plotted against the volume of the solution of the
strong base added, a point of inflexion occurs at the equivalence point. Similarly, in the EDTA
titration, if pM (negative logarithm of the 'free' metal ion concentration: pM = - log [Mn+]) is
plotted against the volume of EDTA solution added, a point of inflexion occurs at the
equivalence point; in some instances this sudden increase may exceed 10 pM units. The
general shape of titration curves obtained by titrating 10.0 mL of a 0.01 M solution of a
metal ion M with a 0.01 M EDTA solution is shown in Figure 3. The apparent stability
constants of various metal-EDTA complexes are indicated at the extreme right of the curve. It
is evident that the greater the stability constant, the sharper is the end point provided the pH
is maintained constant.
pM16
1016
Complex stability
constant value
14
12
1010
10
8
106
6
104
4
12
2
4
6
8
10 12
0.01M EDTA, mL
Fig.3: EDTA titration curves and the effect of Kst (complex stability constant).
12
Complexometric Titrations
Titration curves for 0.1 M Ca2+ versus 0.1 M EDTA at pH 7 and pH 10 are shown in
Figure 4. This figure indicates the effect of pH on apparent stability constants and
correspondingly on the shape of the titration curve.
pCa
mL EDTA
Fig.4: Titration curves of Ca2+ with EDTA at pH 7 and pH 10.
METAL ION INDICATORS (METALLOCHROMIC INDICATORS):
Several methods can be used to detect the end point in EDTA titrations. The most
common technique is to use a metal ion indicator. A metal ion indicator is a compound
whose color changes when it binds to a metal ion. Several common indicators are shown in
Table 4. For an indicator to be useful, it must bind metal less strongly than EDTA does.
A typical analysis is illustrated by the titration of Mg2+ with EDTA using Eriochrome
black T (Erio T) as the indicator. We can write the reaction as follows:
MgIn
(red)
+
EDTA 
(colorless)
MgEDTA
+
(colorless)
In
(blue)
A small amount of indicator (In) is added to the Mg2+ to form a red complex. As EDTA is
added, it reacts first with free, colorless Mg2+ and then with the small amount of red MgIn
13
Complexometric Titrations
complex. (The EDTA must therefore bind to Mg2+ better than the indicator binds to Mg2+). The
change from the red of MgIn to the blue of unbound In signals the endpoint of the titration.
Most metal ion indicators are also acid-base indicators. Some pKA values are listed in
Table 4. Because the color of free indicator is pH-dependent, most indicators can be used
only in certain pH ranges. For example, xylenol orange (pronounced zyl-een-ol) changes from
yellow to red when it binds to a metal ion at pH 5.5. This is an easy color change to observe.
At pH 7.5 the change is from violet to red and rather difficult to see. A spectrophotometer can
be used to measure an indicator color change, but it is more convenient if we can see it.
Common metal ion indicators, other than those mentioned in Table 4, are Eriochrome
Blue-Black B, Patton and Reeders' indicator, Calcon, Calichrome, Fast Sulpone Black F,
Catechol violet, Bromopyrogallol Red, Methylthymol Blue, Warianine Blue and Zincon.
Some metal ion indicators are unstable. Solutions of azo indicators (compounds with N=N- bonds) deteriorate rapidly and should probably to prepared each week. Murexide
solution should be prepared fresh each day.
INDICATOR BLOCKING:
For an indicator to be useful in the titration of a metal with EDTA, the indicator must
give up its metal ion to the EDTA. If a metal does not freely dissociate from an indicator, the
metal is said to block the indicator. Eriochrome black T is blocked by Cu 2+, Ni2+, Co2+, Cr3+,
Fe3+, and Al3+. It cannot be used as an indicator for the direct titration of any of these metals.
Eriochrome black T can be used for a back titration, however, for example, excess standard
EDTA can be added to Cu2+. Then indicator is added and the excess EDTA is back-titrated with
Mg2+.
The success of an EDTA titration depends upon the precise determination of the end
point. The most common procedure utilizes metal ions indicators. The requisites of a metal
ion indicator for use in the visual detection of end points include:
(a) The color reaction must be such that before the end point, when nearly all the metal ion
is complexed with EDTA, the solution is strongly colored.
14
Complexometric Titrations
(b) The color reaction should be specific or at least selective.
(c) The metal-indicator complex must possess sufficient stability, otherwise, because of
dissociation, a sharp color change is not obtained. The metal-indicator complex must,
however, be less stable that the metal-EDTA complex to ensure that, at the end point,
EDTA removes metal ions from the metal indicator-complex. The change in equilibrium
from the metal-indicator complex to the metal-EDTA complex should be sharp and rapid.
(d) The color contrast between the free indicator and the metal-indicator complex should be
such as to be readily observed.
(e) The indicator must be very sensitive to metal ions (i.e. to pM) so that the color change
occurs as near to the equivalence point as possible.
(f) The above requirements must be fulfilled within the pH range at which the titration is
performed.
TYPES OF EDTA TITRATIONS:
I. Direct titration:
In a direct titration, analyte is titrated with standard EDTA. The analyte is buffered to an
appropriate pH at which the conditional formation constant for the metal-EDTA complex is
large enough to produce a sharp end point. Since most metal ion indicators are also acidbase indicators, they have different colors at different values of pH. An appropriate pH must
be one at which the free indicator has a distinctly different color from the metal-indicator
complex.
In many titrations an auxiliary complexing agent, such as ammonia, tartarate, citrate, or
triethanolamine, is employed to prevent the metal ion from precipitating in the absence of
EDTA. For example, the direct titration of Pb2+ is carried out in ammonia buffer at pH 10 in
the presence of tartarate, which complexes the metal ion and does not allow Pb(OH)2 to
precipitate. The lead-tartarate complex must be less stable than the lead-EDTA complex, or
the titration would not be feasible.
15
Complexometric Titrations
II. Back titration:
In a back titration a known excess of EDTA is added to the analyte. The excess EDTA is
then titrated with a standard solution of a second metal ion. A back titration is necessary if
the analyte precipitates in the absence of EDTA, if it reacts too slowly with EDTA under
titration conditions, or if it blocks the indicator. The metal ion used in the back titration
should not displace the analyte metal ion for its EDTA complex. Examples for using the back
titration technique are:
(a) Preventing precipitation:
Al3+ precipitates as Al(OH)3 at pH 7 in the absence of EDTA. An acidic solution of Al3+
can be treated with excess EDTA, adjusted to pH 7-8 with sodium acetate, and boiled to
ensure complete complexation of the ion. The Al3+-EDTA complex is stable in solution at this
pH. The solution is then cooled, eriochrome black T indicator is added; and back titration with
standard Zn2+ is performed.
(b) Preventing blocking of the indicator:
An indicator is said to be blocked when it forms a metal ion complex whose stability
constant is greater than that of the metal-EDTA complex. In this case no color change can be
observed at the end of the titration. Blocking also occurs when the metal-indicator complex
dissociates so slowly that the titration with EDTA cannot be carried out in a reasonable time.
For example, Ni2+ forms a slowly dissociating complex with the pyridylazonaphthol (PAN). It is
therefore not feasible to add indicator to the free metal ion. Excess EDTA can be added to the
Ni2+ and a back titration with Cu2+ is performed.
III. Displacement titration:
For metal ions that do not have a satisfactory indicator, a displacement titration may be
feasible. In this procedure the analyte usually is treated with excess Mg(EDTA)2- chelate to
displace Mg2+, which is later titrated with standard EDTA.
Mn+ + MgY2-  MYn-4 + Mg2+ ........................................................... (8)
Hg2+ is determined in this manner. The formation constant of Hg(EDTA)2- must be greater
than the formation constant for Mg(EDTA)2-, or else Reaction (8) will not work.
16
Complexometric Titrations
An interesting application is the titration of calcium. In the direct titration of calcium
ions, solochrome black gives a poor end point; if magnesium is present, it is displaced from
its EDTA complex by calcium and an improved end point results.
An another example of the displacement titration; there is no suitable indicator for Ag+.
However, Ag+ will displace Ni2+ from the tetracyanonickelate ion:
2Ag+ + Ni(CN)42-
 2Ag(CN)2- + Ni2+ .............................................. (9)
The liberated Ni2+ can then be titrated with EDTA to find out how much Ag+ was added.
IV. Alkalimetric titration:
When a solution of disodium ethylenediaminetetra-acetate, Na2H2Y, is added to a
solution containing metallic ions, complexes are formed with the liberation of two
equivalents of hydrogen ions:
Mn+
+ H2Y2-
(MY)(n-2)+ + 2H+
The hydrogen ions thus set free can be titrated with a standard solution of sodium
hydroxide using an acid-base indicator or a potentiometric end-point. Alternatively, an iodateiodide mixture is added as well as the EDTA solution and the liberated iodine is titrated with a
standard thiosulphate solution.
The solution of the metal to be determined must be accurately neutralized before
titration; this is often a difficult matter on account of the hydrolysis of many salts and
constitutes a weak feature of alkalimetric titration.
V. Indirect titrations (determination of anions):
Anions that form precipitates with certain metal ions may be analyzed with EDTA by
indirect titration. For example, sulfate can be analyzed by precipitation with excess Ba2+ at pH
1. The BaSO4 precipitate is filtered and washed. Boiling the precipitate with excess EDTA at
pH 10 brings the Ba2+ back into solution as Ba(EDTA)2-. The excess EDTA is back-titrated with
Mg2+.
17
Complexometric Titrations
Alternatively, an anion may be precipitated with excess metal ion. The precipitate is
filtered and washed, and the excess metal ion in the filtrate is titrated with EDTA.
Phosphate may be titrated by forming MgNH4PO4 and titrating the Mg2+ in the
precipitate.
Cyanide is titrated by adding an excess of Ni2+ and titrating back the Ni2+ not used up in
the formation of the tetracyano-nickeltate (Ni(CN)4)2-.
Halogens may be determined by precipitating as AgX which is then dissolved in a
solution of (Ni(CN)4)2- producing an equivalent amount of free Ni2+ ions which can be (DIRECT)
titrated with complexone (III). Fluorides may be determined by precipitation as calcium
fluoride and then determining the calcium content in the precipitate after separating it by
filtration. Anions such as CO32-, CrO42-, S2- and SO32- can also be determined by indirect
titration with EDAT.
VI. Masking:
A masking agent is a reagent that protects, without physical separation, some
component of the analyte from reaction with EDTA. Masking can be performed by the
following methods:
1- Masking agents:
For example, Al3+ reacts with F- to form the very stable complex AlF63-. The Mg2+ in a
mixture of Mg2+ and Al3+ can be titrated by first masking the Al3+ with F-, leaving only the
Mg2+ to react with EDTA.
Cyanide is a common masking agent that forms complexes with Cd2+, Zn2+, Hg2+, Co2+,
Cu2+, Ag+, Ni2+, Pd2+, Pt2+, Fe2+, and Fe3+, but not with Mg2+, Ca2+, Mn2+, or Pb2+. If cyanide is
first added to a solution containing Cd2+ and Pb2+, only the Pb2+ is then able to react with
EDTA. Fluoride can mask Al3+, Fe3+, Ti4+ and Be2+. Triethanolamine masks Al3+, Fe3+ and
Mn2+, and 2,3-dimercaptopropanol masks Bi3+, Cd2+, Cu2+, Hg2+, and Pb2+.
2- Kinetic masking:
18
Complexometric Titrations
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.
3- Masking by adjustment of the oxidation state of the element:
This procedure is of limited application, but it has been used, for example, in removing
the interference of Fe3+ with other tri- and tetravalent ions by reducing it to the ferrous state
by ascorbic acid. Cu2+ can be reduced to Cu+ by acidic solution of hydroxylamine or ascorbic
acid. Also, Hg2+ can be reduced to metallic mercury and Cr3+ can be oxidized with alkaline
peroxide to the chromate ion which does not react with EDTA.
Demasking:
Refers to the release of a metal ion from a masking agent. Cyanide complexes can be
demasked by treatment with formaldehyde in acetic acid medium or chlorohydrate.
OH
M (CN ) nmm  mH 2 CO  mH   mH 2 C
+ Mn+
CN
[Zn(CN)4]2- + 4 HCHO + 4H+  4 HCH(OH)CN
+ Zn2+
EDTA SELECTIVITY:
EDTA is very unselective chelating agent because it reacts with numerous bi-, tri-,
tetravalent metals. But, the selectivity can be highly increased by:
(a) Masking and demasking:
This allows individual components of complex mixtures of metal ions to be analyzed by
EDTA titrations. The use of masking and selective demasking agents permits the successive
titration of many metals.
Thus, a solution containing Mg, Zn and Cu can be titrated as follows:
19
Complexometric Titrations
(i) Add excess of standard EDTA and back-titrate with standard Mg solution using
Solochrome Black (Eriochrome Black T) as indicator. This gives the sum of all the metals
present.
(ii) Treat an aliquot portion with excess of KCN and titrate as before. This gives Mg only.
(iii) Add excess of chloral hydrate (or of formaldehyde-acetic acid solution, 3:1) to the titrated
solution in order to liberate the Zn from the cyanide complex, and titrate untie the
indicator turns blue. This gives the Zn only. The Cu content may then be found by
difference.
(b) Suitable control of the pH of the solution:
This, of course, makes use of the different stabilities of metal-EDTA complexes. Thus
bismuth and thorium can be titrated in an acidic solution (pH= 2) with xylenol orange or
methylthymol blue as indicator and most divalent cations do not interfere. A mixture of
bismuth and lead ions can be successfully titrated by first titrating the bismuth at pH 2 with
xylenol orange as indicator, and then adding hexamine to raise the pH to about 5, and
titrating the lead.
(c) Classical separation:
These may be applied if they are not tedious; thus the following precipitates may be
used for separations in which, after being re-dissolved. The cations can be determined
complexometrically: CaC2O4, nickel dimethylglycoximate, Mg(NH4)PO4.6H2O and CuSCN.
(d) 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 (isobutylmethylketone); the extract is diluted with
water and the zinc content determined with EDTA solution.
(e) Choice of indicator:
20
Complexometric Titrations
The indicator chosen should be one for which the formation of the metal-indicator
complex is sufficiently rapid to permit establishment of the end point without undue waiting
and should preferably be reversible.
(f) Removal of anions:
Anions, such as orthophosphate, which can interfere in complexometric titrations may
be removed using ion exchange resins.
TITRATIONS INVOLVING UNIDENTATE LIGANDS:
Because of the stepwise formation of successive complexes, unidentate ligands are
only rarely suitable for the titration of metal ions. However, there are a few examples of
important titrations based upon such ligands, and we shall consider briefly the two bestknown cases.
1- Cyanometric titrations:
The cyanide ion forms complexes with many metal ions of which those of silver, nickel
and mercury are the most useful in estimating these ions. The Liebeg's method for the
determination of the cyanide ion has been described in connection with the argentometric
titrations. It has also been mentioned that the cyanide complexes formed with nickel, cobalt,
cupric and zinc ions have been utilized in the determination of these ions, where an excess of
the cyanide ion is added to an ammoniacal solution of the metal ion, and the excess or
unreacted cyanide is titrated back with silver nitrate.
The addition of a silver nitrate solution to a solution containing cyanide ion causes the
formation of the soluble dicyanoargentate(I) ion, Ag(CN)2. After sufficient silver has been
added to react with all the cyanide by this reaction, the first excess will react to form
insoluble silver dicyanorgentate(I), Ag[Ag(CN)2], which is often referred to and written as silver
cyanide, AgCN. The ionic equations for the reaction are:
21
Complexometric Titrations
Ag+ + 2CNAg+ + Ag(CN)2-
Ag(CN)2Ag[Ag(CN)2]
As a consequence, the appearance of a permanent precipitate indicates the
quantitative completion of the first reaction and can be taken as the equivalence point.
The end point is not sharp, owing to formation of solid silver dicyanoargentate(I) caused
by local excess concentrations in the zone of mixing of the titrant and the solution. This
precipitate is slow to re-dissolve in the equivalence point region. The end point can be
sharpened by the addition of ammonia and iodide ion because the ammonia will prevent the
precipitation of silver dicyanoargentate(I) but does not prevent the precipitation of silver
iodide (Deniges modification of Liebeg’s method).
In the determination of copper by cyanometric method, the reaction goes according to
the following equations:
Cu2+ + 4NH3 = [Cu(NH3)4]2+
[Cu(NH3)4]2+ + 4CN- + H2O = [Cu(NH3)]2- + CNO- + 2NH4+ + 2NH3
Here, it can be noticed that the cyanide ion has displaced the ammonia molecule, a fact
which indicates that copper-cyanide complex is more stable than the copper-ammonia one.
Also, one cyanide group has been oxidized at the expense of the copper which is converted to
the cuprous stage, Cu+.
As the cyanide complex is colorless and the copper-ammine is deep blue, the end point
of the reaction is marked by the disappearance of the blue color. However, the use of metal
indicators has made the detection of the end point quite easy by the sharp change of the
color. In the copper instance, murexide is used and during the titration, the deep blue color
gradually fades to pale yellow, the, at the end point, it changes into pink-violet even in very
dilute solutions. The same indicator can be used in the estimation of nickel by the
cyanometric method.
2- Mercurimetric estimations:
22
Complexometric Titrations
Many soluble mercuric salts are very slightly ionized e.g. chloride, thiocyanate, and
cyanide; while other salts are highly ionized, e.g. nitrate and perchlorate. This property is
made use of in mercurimetric determinations where a standard mercuric nitrate or
perchlorate solution is added to the halide and the end point is detected by including another
compound which forms with the excess mercury a color or turbidity.
The choice of a suitable indicator is governed by the instability constant of the complex
as well as by the solubility product of the mercury salt. The dissociation constants of mercury
complexes decreases in the following order: Cl < Br < CNS < I < CN. After the equivalence
point of the titration is reached, the first excess titrant reacts with the indicator, e.g.
nitroprusside to give a turbidity of mercuric nitroprusside in presence of little nitric acid.
Hg2+ + [Fe(CN)5NO]2-  Hg[Fe(CN5NO]
Diphenylcarbazone and diphenylcarbazide can also be used as indicators; they form
with Hg2+ oxidation products with an intense blue-violet color at pH 2.
C6H5
C6H5NH-NH
++
C O + ½ Hg
C6H5N=N
Hg
2
N-NH
C O
+ H+
N=N
C6H5
In the Volhard process for the determination of halides and thiocyanate, the presence of
precipitates which are sometimes colored, hinders to a certain extent the sharpness of the
color change of the ferric ion indicator. This difficulty has been solved by replacing the silver
nitrate by the highly ionizable mercury perchlorate Hg(ClO4)2. The procedure is to add an
excess of standard mercuric perchlorate solution to the bromide solution with nitric,
perchloric or sulphuric acids. The excess mercury is then titrated back with a standard
thiocyanate using ferric alum as indicator. Chlorides cannot be determined by this method
since the complex it forms with the mercuric ion is less stable than the thiocyanate one. In
this case direct titration of the chloride with mercuric nitrate or perchlorate is useful using
either sodium nitroprusside, or diphenyl carbazone as indicators.
23
Complexometric Titrations
The method is suitable for the determination of even very small amounts of chloride in
acid medium, and the end point is sharp. Metallic ions which give insoluble nitroprussides,
e.g. Cu2+ interfere.
The mercury(II) ion-chloride system is unusual in that the last two of the successive
complexes in the formation of HgCl42- are of much lesser stability than the first two, as shown
by the following successive formation constants:
Hg2+ + Cl-
HgCl-
[ HgCl  ]
K1=
= 5.5 x 106
[ Hg 2  ][Cl  ]
HgCl- + Cl-
HgCl2
K2=
HgCl2 + Cl-
HgCl3- K3=
HgCl3- + Cl-
HgCl42- K4=
[ HgCl 2 ]
= 3.0 x 106
[ HgCl  ][Cl  ]

[ HgCl 3 ]
= 7.1
[ HgCl 2 ][Cl  ]
2
[ HgCl 4 ]

[ HgCl 3 ][Cl  ]
= 10
Thus, in the titration of a chloride solution with an ionized mercury salt such as
mercury(II) nitrate or perchlorate, there is a sudden drop in pHg (pHg= - log [Hg2+]) when the
formation of HgCl2 is essentially complete.
Iodide ion forms with mercuric ion a very stable complex (pK= 30.2). The reaction can
be used to estimate iodides or mercuric salts by titrating the iodide with a mercuric ion when
a soluble colorless complex is first formed. Near the end point, excess mercuric ions
decompose the complex and produce red mercuric iodide, thus;
4I- + Hg2+ = HgI42HgI42- + Hg2+ = 2 HgI2 (red ppt.)
The end point can be detected by diphenylcarbazone indicator if the titration is carried
out in aqueous alcoholic medium (1:9).
The thiocyanate is determined by titration with standard mercuric nitrate solution, using
ferric alum as indicator. The titration reaction is:
2SCN- + Hg2+ = Hg(SCN)2
24
Complexometric Titrations
An excess of the titrant beyond the stoichiometric point, will react with the red ferric
thiocyanate complex resulting in the disappearance of its color.
2FESCN2+ + Hg2+ = 2 Fe3+ + Hg(SCN)2
The reaction goes quantitatively to the right because the mercuric thiocyanate is less
ionized than the ferric thiocyanate complex.
Cyanide is determined, indirectly, by adding a known volume in excess of standard
mercuric nitrate solution,
2CN- + Hg2+ = Hg(CN)2
The excess mercuric nitrate is back-titrated with standard thiocyanate solution using
ferric alum as indicator, in a manner analogous to the Volhard method. In this case, no
precipitate is formed, and the mercuric cyanide will not interfere with the back-titration
because mercuric cyanide is less ionized than mercuric thiocyanate.
REFERENCES
1- Vogel's Textbook of Quantitative Chemical Analysis, 5th Edn. ELBS edn., London (1991).
2- Gary D. Christian, Analytical Chemistry, Jon Willey & Sons, Inc., New York (1994).
3- Daniel C. Harris, Quantitative Chemical Analysis, W.H. Freeman and Company, San
Franciso (1982).
4- Charles T. Kenner and Kenneth W. Busch, Quantitative Analysis, Macmillan Publishing
Co., Inc., New York (1979).
QUESTIONS
Write shortly on the following:
Ligand-Coordination number-Chelation-Binuclear complex-Masking.
Show by equations:
Stepwise stability constant of the complex.
Cyanometric determination of Cu2+.
Relation between K and KHZ.
Masking and selective demasking of Zn2+.
25
Complexometric Titrations
Show by drawing:
Effect of apparent stability constant and of pH on EDTA titration curves.
Spatial structure of Ca-EDTA chelate.
Chemical structure of a metallochromic indicator.
Show by examples:
Types of ligands - Chelate effect - EDTA replacement titration.
Explain why:
- EDTA has the widest general applications in compleximetry.
- Al3+ when determined by EDTA; back titration should be applied.
- Hg2+ can be determined successfully by titration which Cl-.
Explain how:
EDTA selectivity can be increased.
Anions can be determined by EDTA.
A mixture of Cu, Zn and Mg cations can be analyzed.
26