Download Complexometric Titrations

Complexometric Titrations
Introduction

The simple salt consists of two radicals, e. g. (acidic and basic
radicals), for example:

FeCl3, CoSO4, Cu(NO3)2, …etc.

Fe3+, Co2+, Cu2+, ... Basic radical

Cl-, SO42-, NO3-, ... Acidic radical

The complex salt consists of more than two radicals, such as:

[Fe(H3N)6]Cl3, [Co(NH3)4]SO4, [Cu(NH3)4] …etc.

The chemical species [Fe(HN3)6]Cl3, [Co(NH3)4]SO4, [Cu(NH3)4],
[Co(NH3)4Cl2]Cl, [Co(NH3)5Cl]Cl2 are called metal complexes or
coordination compounds.

The formation of a coordination compound from a metal ion is called
complexation.

The part [Co(NH3)6]3+ in [Co(NH3)6]Cl3 molecule is called complex
species or complex ion.

The formula of the complex species is written within bracket [ ].

The six NH3 groups are called the ligands.

The ligand can be defined as any molecule or ion that has at least one
electron pair (Lewis base) which can be donated to a metal ion (Lewis
acid).

Transition elements exhibit two types of valencies, namely primary
valence and secondary valence.

The primary valence is also known as ionisable valence and secondary
valence is otherwise known as nonionisable valence.

Anions can satisfy primary valence whereas anions or neutral
molecules can satisfy secondary valence.

In modern terms, the primary valence corresponds to the oxidation
number and the secondary valence corresponds to the coordination
number.
Cl
Cl
H3N
Co
Example: [Co(NH3)3Cl3]
H3N
Cl
NH3
Types of ligands

Monodentate ligands
When the ligand has only one donor atom, it is said to be monodentate.
It may be a neutral molecule such as

or a negatively charged ion such as


Chelation

When the ligand has two or more donor atoms, the terms bidentate,
tridentate, tetradentate …etc., generally multidentate frequently used.

The cyclic compound form when the multidentate ligand attached to
the central metal atom or ion by more than one coordinating atom, is
called chelate.

The process is known as a chelation.
Bidentate Ligands
O
O
C
CH2 CH2
H2N
C
O
*
2-
*
O
*
oxalate ion
*N
*N
Donor Atoms
*
Ethylenediamine (en)
CH
CH
C
CH
HC
C
C
HC
C
CH
CH
NH2
CH
ortho-phenanthroline
tris(ethylenediamine)cobalt(III) complex, [Co(en)3]3+
Polydentate Ligands
O
O
C
O
*
N
O
*
CH
*
C
O
CH2 C
CH2
*
CH2 CH2 N
* CH C
2
2
EDTA
Donor Atoms
O
O
O
*
Polydentate ligand and chelation
EOS
The Chelate
[Pt(en)2]
The square planar
complex of Pt2+ and
ethylenediamine has two
ligands with a total of
four points of attachment.
Ethylenediamine
has two nitrogen
atoms that can
donate electrons to a
single metal ion.

Complex species are of different types; cationic, neutral and anionic:

[Pt(NH3)4]2+
Cationic

[Pt(NH3)4Cl2]
Neutral

[PtCl4]2-
Anionic

The molecules (or ions) of complexes have characteristic stereo shapes
such as tetrahedral, square planar, octahedral, …etc., depending on the
number of ligands bonded to the central metal.
Examples of complexes
Cationic complex ions
Neutral complex
Anionic
complex ions
Significant Definitions

Complex ion: Species where transition metal ion is surrounded by a
certain number of ligands.

Coordination number: The total number of monodentate ligands
attached to the central metal ion in the complex molecule.

For example, in the complex ion [Cu(NH3)4]2+, as four monodentate
ligands are attached to the central metal ion, the coordination number is
four.

The coordination number can be also known as the number of sigma
bonds between the ligands and the central metal ion.

Coordination numbers from two to nine are known in complexes.
Complex ion

Complex ion is formed by the union of a simple ion with either other
ions of opposite charge or with neutral molecules as shown by the
following examples.

AgCN + CN−
[(AgCN)2]− (complex ion)

AgCl + 2NH3
[Ag(NH3)2]Cl (complex molecule)

CuCl2 + 4NH3
[Cu(NH3)4]Cl2 (complex molecule)

Complexation process: is a 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: Mm+(H2O)n + L
[Mm+(H2O)(n-1)L] + H2O
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:

M+L
ML;
K1 = [ML]/[M][L]

ML + L
ML(n−l) + L
ML2;
MLn;
K2 = [ML2]/[ML][L]
Kn = [MLn]/[ML(n−l)] [L]


The equilibrium constants K1, K2, ..., Kn are referred to as stepwise
stability constants or the overall stepwise formation constant K;

K = K1 × K2 ×..., Kn

Knowledge of stability constant values is of considerable importance in
analytical chemistry, since they provide information about the
concentrations of the various complexes formed by a metal in specified
equilibrium mixtures; this is valuable in the study of complexometry,
and of various analytical separation procedures such as solvent
extraction, ion exchange, and chromatography.
Complexones

Ethylenediaminetetra-acetic acid (EDTA)

The formula (I) is preferred to (II), since it has been shown from
measurements of the dissociation constants that two hydrogen atoms
are probably held in the form of zwitterions.
I
II
The formula of EDTA (I) and the zwitterions (II).

The values of pK are respectively pK1 = 2.0, pK2 = 2.7, pK3 = 6.2, and
pK4 = 10.3 at 20 ºC; these values suggest that it behaves as a
dicarboxylic acid with two strongly acidic 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.

To simplify the following discussion EDTA is assigned the formula
H4Y: the disodium salt is therefore Na2H2Y and affords the complexforming ion H2Y2− in aqueous solution; it reacts with all metals in a 1:1
ratio.

The reactions with cations, e.g. M2+, may be written as:

M2+ + H2Y2−
MY2− + 2H+
Stability Constant of Metal EDTA Complexes

The stability of a complex is characterized by the stability constant (or
formation constant) K:


Mn+ + Y4−
(MY)(n−4)+
K = [(MY)(n−4)+]/[Mn+][ Y4−]
Metal ion Indicators

The success of an EDTA titration depends upon the precise
determination of the end point. 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.

(b) The color reaction should be specific or at least selective.

(c) The metal-indicator complex must possess appreciable stability, but
the metal-indicator complex must, however, be less stable than the
metal-EDTA complex.

(d) The color contrast between the free indicator and the metalindicator complex should be such as to be readily observed.
 (e)
The indicator must be very sensitive to metal ions.
 (f)
The above requirements must be fulfilled within the pH
range at which the titration is performed.
 The
indicators which form complexes with specific metal
cations can serve as 1:1 complexes.
 The
molar ratio metal: indicator = 1:1 are common, but 1:2complexes and 2: l-complexes also occur.
 The
metal ion indicators, like EDTA itself, are chelating agents.
Theory of the visual use of metal ion
indicators.

Discussion will be confined to the more common 1: l-complexes.

The use of a metal ion indicator in an EDTA titration may be written as:
M-In + EDTA → M-EDTA + In

This reaction will proceed if the metal-indicator complex (M-In) is less
stable than the metal-EDTA complex (M-EDTA).

The former dissociates to a limited extent, and during the titration the
free metal ions are progressively complexed by the EDTA until
ultimately the metal is displaced from the complex M-In to leave the
free indicator (In).

The indicator color change is affected by the hydrogen ion
concentration of the solution.

Thus Eriochrome black T, who may be written as H2In−, exhibits the
following acid-base behavior:
H2InRed

pH
HIn2-
5.3 - 7.3
Blue
pH
10.5 - 12.5
In3Yellow - orange
In the pH range 7 − 11, in which the dye itself exhibits a blue color,
many metal ions form red complexes; these colors are extremely
sensitive, as is shown, for example, by the fact that 10−6 − 10−7 molar
solutions of magnesium ion give a distinct red color with the indicator.
Eriochrome black T
Murexide
Types of EDTA Titrations


A. Direct titration.
The solution containing the metal ion to be determined is buffered to the
desired pH (e.g. to pH = 10 with NH4Cl/NH4OH, i.e. the ammoniacal
buffer) and titrated directly with the standard EDTA solution.
Mn+ + In
Color 1
[M-In] + EDTA
[M-In]
Color 2
[M-EDTA] + In
Colorless
Color 1
Mn+ + In + 10 ml of
buffer solution pH = 10
B. Back-titration.

Many metals cannot, for various reasons, be titrated directly; thus they
may precipitate from the solution in the pH range necessary for the
titration, or they may form inert complexes, or a suitable metal
indicator is not available.

In such cases an excess of standard EDTA solution is added, the
resulting solution is buffered to the desired pH, and the excess of the
EDTA is back-titrated with a standard metal ion solution.

A solution of zinc chloride or sulphate or of magnesium chloride or
sulphate is often used for this purpose.

The end point is detected with the aid of the metal indicator which
responds to the zinc or magnesium ions introduced in the back-titration.
Zn2+ or Mg2+
EDTA(excess) + Mn+ + In +10 ml
of buffer solution pH = 10
Mn+ + EDTA(excess) → [M-EDTA]
EDTA + Zn2+ or Mg2+ → [Zn or Mg – EDTA]
Zn2+ or Mg2+ + In → [M-In]
C. Replacement or substitution titration.

Substitution titrations may be used for metal ions that do not react (or
react unsatisfactorily) with a metal indicator, or for metal ions which
form EDTA complexes that are more stable than those of other metals
such as magnesium and calcium.

The metal cation Mn+ to be determined may be treated with the
magnesium complex of EDTA, when the following reaction occurs:

Mn+ + MgY2−
(MY)(n−4)+ + Mg2+

The amount of magnesium ion set free is equivalent to the cation
present and can be titrated with a standard solution of EDTA and a
suitable metal indicator.

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.
D. Alkalimetric titration.

When a solution of disodium EDTA, Na2H2Y, is added to a
solution containing metallic ions, complexes are formed with
the liberation of two equivalents of hydrogen ion:


Mn+ + H2Y2−
(MY)(n−4)+ + 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.
E. Miscellaneous methods.

Exchange reactions between the tetracyano-nickelate(II) ion
[Ni(CN)4]2− and the element to be determined, whereby nickel ions are
set free, have a limited application.

Thus silver and gold, which themselves cannot be titrated
complexmetrically, can be determined in this way.


[Ni(CN)4]2−+ 2Ag+
2[Ag(CN)2]− + Ni2+
the equivalent amount of nickel thereby set free is determined by rapid
titration with EDTA using an appropriate indicator (murexide, bromo
pyrogallol red).