TITRIMETRY / Overview 105
Certified Reference Materials
In order to validate speciation methods more effectively, a number of certified reference materials
(CRMs) have been produced to allow laboratories to
measure the accuracy of their techniques. A range of
sediment and biota materials are available such as
the sediment materials PACS-2 from the National
Research Council of Canada, NIES-12 from the
National Institute for Environmental Studies, Japan
and the CRM-462 from Community Bureau of Reference (BCR), EU. These materials have been rigorously homogeneity and stability tested and the levels
of organotins have been certified by a range of techniques utilizing either ‘definitive’ methods or multiple
independent methods.
See also: Atomic Absorption Spectrometry: Interferences
and Background Correction. Atomic Emission Spectrometry: Principles and Instrumentation; Interferences and
Background Correction; Flame Photometry; Inductively Coupled Plasma; Microwave-Induced Plasma. Atomic Mass
Spectrometry: Inductively Coupled Plasma; Laser Microprobe. Countercurrent Chromatography: Solvent Extraction with a Helical Column. Derivatization of Analytes.
Elemental Speciation: Overview; Practicalities and Instrumentation. Extraction: Solvent Extraction Principles; Solvent Extraction: Multistage Countercurrent Distribution;
Microwave-Assisted Solvent Extraction; Pressurized Fluid
Extraction; Solid-Phase Extraction; Solid-Phase Microextraction. Gas Chromatography: Overview. Isotope Dilution Analysis. Liquid Chromatography: Overview.
Further Reading
Abalos M, Bayona J-M, Comañó R, et al. (1997) Analytical procedures for the determination of organotin
compounds in sediment and biota: A critical review.
Journal of Chromatography A 788: 1–49.
Alonso JIG, Encinar JR, Rodrı́guez-González P, and
Sanz-Medel A (2002) Determination of butyltin compounds in environmental samples by isotope-dilution
GC–ICP-MS. Analytical and Bioanalytical Chemistry
373: 432–440.
Davies AG (1997) Organotin Chemistry. Weinheim: VCH.
Ebdon L, Hill SJ, and Rivas C (1998) Organotin compounds in solid waste: A review of their properties and
determination using high-performance liquid chromatography. Trends in Analytical Chemistry 17:
277–288.
Ebdon L, Pitts L, Cornelis R et al. (eds.) (2001). Trace
Element Speciation for Environment, Food and Health.
Cambridge: Royal Society of Chemistry.
Leroy MJF, Quevauviller P, Donard OFX, and Astruc M
(1998) Determination of tin species in environmental samples. Pure and Applied Chemistry 70:
2051–2064.
Rajendran RB, Tao H, Nakazato T, and Miyazaki A (2000)
A quantitative extraction method for the determination
of trace amounts of both butyl- and phenyltin compounds in sediments by gas chromatography-inductively
coupled plasma mass spectrometry. The Analyst 125:
1757–1763.
Rodrı́guez-González P, Encinar JR, Alonso JIG, and SanzMedel A (2003) Isotope dilution analysis as a definitive
tool for the speciation of organotin compounds. The
Analyst 128: 447–452.
Smedes F, de Jong AS, and Davies IM (2000) Determination of (mono-, di- and) tributyltin in sediments. Analytical methods. Journal of Environmental Monitoring 2:
541–549.
Suzuki T, Kondo K, Uchiyama M, and Murayama M
(1999) Chemical species of organotin compounds in sediment at a marina. Journal of Agricultural and Food
Chemistry 47: 3886–3894.
TITRIMETRY
Contents
Overview
Potentiometric
Photometric
Overview
A Townshend, University of Hull, Hull, UK
& 2005, Elsevier Ltd. All Rights Reserved.
Introduction
A titration is defined as ‘the process of determining
the quantity of a substance A by adding measured
increments of substance B, the titrant, with which it
reacts until exact chemical equivalence is achieved
106 TITRIMETRY / Overview
(the equivalence point)’. The titrant is usually added
as a standardized solution, but electrolytic generation, as in coulometric titrations, is also possible.
The achievement of the equivalence point is indicated in one of two ways. The first is visually, for
example, by the addition of an indicator, which
changes color, or by fluorescence or a similar property at or close to the equivalence point. The change
indicates the endpoint of the titration, and the indicator should be chosen so that the endpoint is as
close as possible to the equivalence point. The second
method is to measure a physical property of the solution being titrated (e.g., conductivity, pH, absorbance) and identify the equivalence point by
processing the signals obtained. The latter approach
is particularly useful when titrations are automated.
The amount of titrant added is usually measured by
volume (by dispensing the solution from a burette),
and in this case, titrimetry is an example of volumetric
analysis. Occasionally, the titrant is measured by
weight (especially if greater accuracy is required) or by
amount of electricity (as in coulometric titrations).
Titrimetry is one of the oldest analytical techniques,
originating in the middle of the eighteenth century as a
rapid means of quality control of industrial processes,
such as acid manufacture. Since that time, the equipment has been refined, the procedures have been automated, and the number of chemical reactions utilized
greatly increased, but the basic principles are unchanged. Its continued popularity stems from the simplicity of equipment and execution, wide applicability,
and high accuracy and precision (greater than most
instrumental techniques), all of which make it particularly applicable to the determination of major and
minor components of samples. Skilled titrimetric analysis should give results with a precision of o0.2% at
the 1 10 2 mol l 1 level.
Titrimetry may be classified with respect to the
types of reaction that are involved. The major
reactions are acid–base reactions (hence acid–base
titrimetry), redox reactions (redox titrimetry), complexing reactions (compleximetric titrimetry), and
precipitation reactions (precipitation titrimetry),
which will be discussed in more detail below.
Titrimetry may also be classified by the nature of
the endpoint measurement. The use of electrical
measurements gives rise to potentiometric, amperometric, and coulometric titrations. Measurement of
heat changes is used in thermometric titrimetry, and
of absorbance in photometric and turbidimetric titrations. Radiometric titrations measure changes in
radioactivity during the titration. All of these techniques are dealt with in other articles in this Encyclopedia. This article discusses only those titrations
that use visual indicators.
General Manual Titrimetric Technique
The simple equipment used for a typical titration is
shown in Figure 1. A sample is measured into the
flask from a pipette, or by weighing. The accuracy
and precision of manual titrations using visual indicators is critically dependent on the use of correct
experimental technique by the analyst. As in most
analytical procedures, precise measurement of the
amount of sample (or sample aliquot) is necessary,
but it is most important in titrimetry if an accuracy of
o0.2% is to be achieved. Thus, proper use of the
pipette, burette, and balance, and a careful sample
preparation procedure is crucial. Measurement
devices of high quality, such as class A pipettes and
burettes, should be used.
When carrying out a titration the burette, holding
the titrant solution, is clamped vertically just above
the flask. Generally, the flask containing the solution
being titrated should be placed on a white tile. Where
the color change is somewhat difficult to detect, a
reference solution should be used for comparison
(this is a solution held in a flask similar to that being
used for the titration, has the same volume, contains
the same indicator, and has been adjusted to the
endpoint). Titrant is added from the burette to the
solution in the flask, which is then swirled by hand.
As the endpoint approaches (which will be signaled
by a transient color change in the portion of the
solution where the titrant is added), the titrant
should be added dropwise. When very close to the
endpoint, fractions of drops can be added by touching the tip of the burette with a partly formed drop
on the inside of the flask, and washing down with
water from a wash bottle. The detection of the endpoint, when using a 50 or 10 cm3 burette, should be
possible to within 0.02 cm3. The volume of titrant
run out of the burette is used to calculate the concentration of analyte in the titrated solution.
Standards
Titrimetric analysis depends upon the availability of
solutions of accurately known concentration
for use as titrants. Such standard solutions may be
prepared and themselves standardized by titration
with solutions prepared from materials of guaranteed
purity and composition, preferably solutions of
chemicals known as primary standards. These materials have the following properties, as listed by
Dodge:
1. They are easily obtained in an analytically pure
state.
TITRIMETRY / Overview 107
D 20°C
ml
Maker (here: Brand)
0
Country of manufacture
1
Trademark of Brand
2
Rated volume
49
BRAND
Class: all instruments having
the international symbol 'A'
are suitable for official
certification.
'S' stands for rapid delivery
50
25
25 ml
D 20°C
ml
AS
0.03 ml
Ex=15s
20°C
Tolerance
Calibration (here Ex = deliver)
and waiting time (here 15 s)
Calibration temperature
(A)
(B)
(C)
0
1
2
3
4
(D)
(E)
Figure 1 Equipment used for a manual titration. (A) Transfer pipette; (B) burette; (C) information provided on volumetric glassware;
(D) conical or Erlenmeyer flask (normally 250 or 100 cm3); (E) burette reader. (Reproduced with permission from Belcher et al. (1970)
and Rudolf Brand and Co. Working with Volumetric Instruments. Wertheim: Brand.)
2. They should be unalterable in air at ambient or
moderately high temperatures.
3. They should have a high equivalent weight,
thus decreasing the effect of small weighing
errors.
4. They should be readily soluble under the conditions of the analysis, thus allowing immediate titration in the cold.
5. On titration, no interfering product should be
present.
108 TITRIMETRY / Overview
Table 1 Some primary titrimetric standards
Compound
Name
Formula
Anhydrous sodium carbonate
Sodium borate (borax) (recrystallized)
Sulfamic acid
Potassium hydrogenphthalate
Potassium hydrogenbiiodate
Silver nitrate
Sodium oxalate
Arsenic(III) oxide
Potassium dichromate
Ammonium hexanitrocerate(IV)
Potassium iodate
Potassium bromide
Calcium carbonate
Zinc oxide
Ni, Zn, Cu metals
Anhydrous disodium EDTA
Na2CO3
Na2B4O7.10H2O
NH2SO3H
KHC8H4O4
KH(IO3)2
AgNO3
Na2C2O4
As2O3
K2Cr2O7
(NH4)2Ce(NO3)6
KIO3
KBrO3
CaCO3
ZnO
Ni, Zn, Cu
C10HI4N2O8Na2
Type
Type of titration
Weak base
Weak base
Strong acid
Weak acid
Strong acid
Acid–base
Acid–base
Acid–base
Acid–base
Acid–base
Argentimetric
Redox
Redox
Redox
Redox
Redox
Redox
Compleximetric
Compleximetric
Compleximetric
Compleximetric
Reductant
Reductant
Oxidant
Oxidant
Oxidant
Oxidant
Source of calcium ions
Source of zinc ions
Source of metal ions
Complexing agent
6. They should be colorless, before and after titration, to avoid interference with indicators.
Equivalence
point
Primary standards for particular titrations are given
in Table 1. There are relatively few compounds that
satisfy all these conditions.
It is also possible to buy concentrated standard
solutions, which, after accurate dilution, can be used
for titrimetry.
A
−1
12
A = 0.1 mol
10
B = 0.01 mol
−1
C = 0.001 mol
B
−1
C
pH
8
6
Acid–Base Titrations
In acid–base titrations, an acid is determined by titration with a base, or vice versa. The essential reaction is between H þ and OH , giving water. The pH
at the endpoint depends on the dissociation constants
of the reactants and products. Thus, the titration of
a strong, i.e., completely dissociated, acid with a
strong, almost completely dissociated, base reaches
equivalence at pH 7.0. A typical example is the titration of hydrochloric acid with sodium hydroxide
solution. If a weaker, i.e., less dissociated base, such
as ammonia, is used, the equivalence pH is o7.0; the
weaker the base, the lower the equivalence pH. Likewise, if a weak, i.e., less dissociated, acid is titrated
with a strong base, the equivalence pH is 47.0, the
pH increasing with increasing weakness of the acid.
For the titration of a weak acid with a weak base, the
equivalence pH depends on the relative dissociation
constants of the acid and base, but the limited pH
change means that the equivalence point is not as
sharp as in a strong acid or strong base titration. The
change of pH during the course of such titrations is
illustrated in Figures 2 and 3.
4
C
B
A
2
0
5
Titrant (ml)
10
Figure 2 Titration of 10 ml of hydrochloric acid of various concentrations with sodium hydroxide solution of the same concentration. (Redrawn from Belcher et al. 1970.)
The extent of the pH change also depends on the
concentration of the analyte and titrant. Figure 2
shows how the pH change decreases with decreasing
concentration of HCl and NaOH. The indicators for
such titrations are chosen to change color very close
to the pH at equivalence, and are described in detail
in a separate article.
Some acids are polybasic, i.e., they give rise to
more than one hydrogen ion. Phosphoric acid, for
TITRIMETRY / Overview 109
a
12
12
10
10
b
8
pH
pH
8
6
6
4
4
2
2
0
5
Titrant (ml)
0
10
25
50
75
NaOH (ml)
Figure 3 Titration of 10 ml of 0.1 mol l 1 acetic acid with (A)
0.1 mol l 1 sodium hydroxide; (B) 0.1 mol l 1 ammonia. (Redrawn
from Belcher et al. 1970.)
Figure 4 Titration of 25 ml of 0.1 mol l 1 H3PO4 with 0.1 mol l 1
NaOH. (Redrawn from Belcher et al. 1970.)
Table 2 Examples of acid–base titrimetric analyses
Analyte
Titrant
Indicator
Conditions
CaCO3
H3PO4
HCl
NaOH
NH4þ
Boric acid
HCl
NaOH
Methyl red or phenolphthalein
Methyl orange
Thymolphthalein
Methyl red
Phenolphthalein
–
First endpoint
Second endpoint
Distillation from NaOH solution
Binding with mannitol or sorbitol to increase acidity
example, produces three hydrogen ions:
þ
H3 PO4 "H2 PO
4 þH
½I
2
þ
H2 PO
4 "HPO4 þ H
½II
3
þ
HPO2
4 "PO4 þ H
½III
The change in pH on titration with NaOH is shown
in Figure 4. The first dissociation (reaction [I]) occurs
most easily, and titration with NaOH gives an
equivalence point at pH 4. Further titration, of the
second, more strongly bound hydrogen ion (reaction
[II]), also gives rise to an equivalence point, at pH 9.
The third hydrogen ion does not give rise to a sharp
endpoint. Thus, phosphoric acid may be determined
by titration to the first or second endpoints.
Some examples of acid–base titrimetric analyses
are given in Table 2.
Precipitation Titrations
These are titrations in which the analyte and titrant
react to form a precipitate. The only common titrant
used is silver nitrate (argentimetric titrations), and
its use is mainly restricted to the determination of
chloride, bromide, iodide, cyanide, and thiocyanate,
although in principle any species that is precipitated
by silver ions could be determined. Direct titrations
involve the use of potassium chromate (Mohr’s
method) or fluorescein derivatives as indicator. An
indirect (back-titration) procedure is also popular, in
which excess of precipitant (Ag þ ) is added to the
sample, and the unreacted Ag þ titrated with
thiocyanate ions (Volhard’s method), using iron(III)
as indicator. The mechanisms of the indicator reactions are described in another article.
Redox Titrations
In these titrations, a reducing agent is titrated with an
oxidizing agent, or vice versa. The common oxidizing
titrants are potassium permanganate (KMnO4),
potassium dichromate (K2Cr2O7), cerium(IV) sulfate
(Ce(SO4)2), iodine (I2), potassium iodate (KIO3), and
potassium bromate (KBrO3), all of which are solids,
and sodium hypochlorite (NaClO), which is available as a solution. The most important reducing titrants are iron(II) salts, often ammonium iron(II)
sulfate ((NH4)2Fe(SO4)2 6H2O), sodium thiosulfate
(Na2S2O3 H2O), and arsenic(III) oxide (As2O3).
110 TITRIMETRY / Overview
The stoichiometry of the reaction between one of
these titrants and a particular analyte is established
by combining the appropriate half-reactions. For the
titrants above, the half-reactions are as follows:
given by the Nernst equation:
Oxidants
where n is the number of electrons involved in the
above half-reaction, and [ox] and [red] are the concentrations (or better, activities) of the oxidized and
reduced forms of the species, respectively. E0 is a
constant known as the standard potential, which is
the idealized potential when [ox] ¼ [red]. A similar
parameter, but measured under actual experimental
conditions, is known as the formal potential (EF),
and is more immediately useful for application in
redox titrimetry. Some potentials for the above
half-reactions are given in parentheses in the above
equations.
The change in oxidation potential during the titration of a reductant (iron(II)) by an oxidant (cerium(IV)) is shown in Figure 5. The final oxidation
potential increases with the strength of the oxidant
used. Sometimes, it is possible to make a simple calculation of the equivalence potential (EEP) as follows.
For the reaction that can be written as in [XVII]:
þ
2þ
MnO
þ 4H2 O
4 þ 8H þ 5e -Mn
ðacidic conditions; E0 ¼ 1:50 VÞ
½IV
þ
MnO
4 þ 4H þ 3e -MnO2 þ 2H2 O
ðneutral conditionsÞ
½V
2
MnO
4 þ e -MnO4
ðalkaline conditionsÞ
½VI
3
þ
Cr2 O2
7 þ 14H þ 6e -2Cr þ 7H2 O
ðEF ¼ 1:0 V; 1 mol l1 HClÞ
½VII
Ce4þ þe -Ce3þ
ðEF ¼ 1:44 V; 1 mol l1 H2 SO4 Þ
I2 þ 2e -2I
ðE0 ¼ 0:54 VÞ
þ
IO
3 þ 6H þ 6e -I þ 3H2 O
½VIII
E ¼ E0 þ
0:059
½ox
log
n
½red
½1
Fe2þ þ Ce4þ "Fe3þ þ Ce3þ
½IX
½XVII
then
½X
EF ðCe4þ =Ce3þ Þ þ EF ðFe3þ =Fe2þ Þ
2
1:44 þ 0:68
¼
2
¼ 1:06 V
EEP ¼
þ
BrO
3 þ 6H þ 6e -Br þ 3H2 O
ðE0 ¼ 1:51 VÞ
ClO þ 2Hþ þ 2e -Cl þ H2 O
½XI
½XII
Reductants
½2
1.50
Fe2þ -Fe3þ þ e
ðEF ¼ 0:68 VÞ
Excess Ce(IV)
1.40
½XIII
1.30
AsðIIIÞ-AsðVÞ þ 2e
ðE0 ¼ 0:09 VÞ
½XIV
ðEF ¼ 0:58 V; 1 mol l1 HClÞ
1.10
1.00
½XV
0.90
For example, iron(II) can be determined by titration
with dichromate, so combination of the appropriate
half-reactions [VII] and [XIII], so as to achieve a charge
and mass balance, gives the overall reaction [XVI]:
2þ
Cr2 O2
þ 14Hþ -2Cr3 þ 6Fe3þ þ 7H2 O
7 þ 6Fe
1.20
E (V )
2
2S2 O2
3 -S4 O6 þ 2e
½XVI
The driving force for each half-reaction is measured
by its oxidation potential, E, measured in V, which is
0.80
Excess Fe(II)
0.70
0
25
Titrant (ml)
50
Figure 5 Titration of 25 ml of 0.1 mol l 1 iron(II) with 0.1 mol l 1
cerium(IV). (Redrawn from Vogel’s Textbook of Quantitative
Inorganic Analysis, 4th edn (1978). London: Longman.)
TITRIMETRY / Overview 111
For more complex systems, however, especially those
involving oxoanions, such simple calculations are
not valid.
Indicators for redox titrations will be chosen to
change color reversibly by oxidation or reduction at
a potential as close as possible to the equivalence
potential (starch indication for iodine is an exception). This aspect is described in detail in another
article.
Some analytes may be determined by titration with
an oxidant, after their reduction. There are several
ways of carrying out such reductions. One commonly used reductant is tin(II) chloride, in which the excess of tin(II) is destroyed by addition of mercury(II)
chloride:
SnðIIÞ þ 2HgðIIÞ þ 2Cl -SnðIVÞ þ Hg2 Cl2
Redox titrations are still widely used. Table 4
summarizes some applications of redox titrations.
Compleximetric Titrations
Compleximetric titrations are used mainly to determine metal ions by use of complex-forming reactions. Although in theory many complexing agents
(cyanide, thiocyanate, fluoride, 1,2-diaminoethane,
etc.) could be used for this purpose, in practice the
titrants are almost always compounds having the
iminodiacetic acid functional group:
CH2COOH
N
½XVIII
The mercury(I) chloride is unaffected by oxidants
during the subsequent titration. Sulfite (or SO2) and
hydrogen sulfide are alternative reductants. Metals
may also be used. Small pieces of metal (zinc – a
Jones reductor, silver – a Walden reductor) are used
to fill a column, through which the analyte solution is
passed. The effluent is titrated with oxidant. A comparison of the reduction products of the two reductor
columns is given in Table 3.
Table 3 Comparison of the Jones and Walden reductors
Walden (HCl)
Jones (H2SO4)
Fe(III)-Fe(II)
Ti(IV) not reduced
Cr(III) not reduced
V(V)-V(IV)
Mo(VI)-Mo(V)
Cu2 þ -Cu(I)
Fe(III)-Fe(II)
Ti(IV)-Ti(III)
Cr(III)-Cr(II)
V(V)-V(II)
Mo(VI)-Mo(III)
Cu2 þ -Cu(0)
CH2COOH
and by far the most popular of these is ethylenediaminetetraacetic acid (EDTA):
CH2COOH
HOOCCH2
N
CH2
CH2
N
HOOCCH2
CH2COOH
(H4L)
This ligand fulfills many of the qualities of a good
compleximetric titrant. It forms complexes with
most metal ions (those with the alkali metals are
too weak to be useful); all the complexes have exact
1:1 stoichiometry, because the ligand is hexadentate,
and can therefore occupy up to six coordination positions on the metal ions, as shown in Figure 6. The
reaction with most metal ions is rapid (Cr3 þ is a
well-known exception), and the complexes are water
Table 4 Examples of redox titrations
Analyte
Titrant
Indicator
Iron(II)
Iron(III)
H2O2
Iron
Ethanol
Oxalate
Nitrite
Copper(II)
Acids
Available chlorine
Antimony(III)
Magnesium(II)
KMnO4
KMnO4
KMnO4
K2Cr2O7
K2Cr2O7
Ce(SO4)2
Ce(SO4)2
Na2S2O3
Na2S2O3
Na2S2O3
I2
Na2S2O3
Self-indicating
Self-indicating
Self-indicating
Diphenylamine sulfonic acid
N-Phenylanthranilic acid
Nitroferroin
Ferroin
Starch
Starch
Starch
Starch
Starch
Ascorbic acid
Ammonia
KIO3
NaClO
Self-indicating (I2/CCl4)
Bromothymol blue
Condition
SnCl2 reduction
In iron ore, SnCl2 reduction
Add excess oxidant, heat, back-titrate with iron(II)
Add excess oxidant, heat, back-titrate with iron(II)
Add excess oxidant, back-titrate with iron(II)
Iodide oxidized to iodine
þ
5I þ IO
3 þ 6H -3I2 þ 3H2 O
Oxidation of I to I2
Precipitate Mg2 þ with 8-quinolinol, add excess KBrO3/KBr to
brominate precipitate, determine excess KBrO3 by oxidation
of I -I2
Add Br and excess NaClO
112 TITRIMETRY / Overview
Table 5 Stability constants of some EDTA complexes and
optimum pH for titration of the metal ions with EDTA
2−
CO
Metal ion
O
CH2
CO
2þ
Mg
Ca2 þ
Ba2 þ
Mn2 þ
Zn2 þ
Ni2 þ
Co2 þ
Fe2 þ
Fe3 þ
Cu2 þ
Hg2 þ
Pb2 þ
Bi3 þ
CH2
O
N
CH2
M
CH2
N
O
CO
CH2
CH2
O
log k1
Optimum pH for titration
8.7
10.7
7.8
13.8
16.5
18.6
16.3
14.3
25.1
18.8
21.8
18.0
27.9
10
7.5
12–13
5.5
4
3
4
5
1
3
5.0–5.5a
4
1–3
The behavior of Hg2 þ is somewhat anomalous because of
hydroxocomplex formation. Values obtained from Pribil (1982).
a
CO
Figure 6 Structure of chelate of EDTA anion (L4 ) with a metal
ion M2 þ . (Redrawn from Vogel’s Textbook of Quantitative
Inorganic Analysis, 4th edn (1978). London: Longman.)
CH2COOH
N
CH2COOH
soluble and colorless (unless the metal ion itself is
colored).
The reaction between a typical metal ion and
EDTA (H4L) can be written as
nþ
M
þ H4 L"ML
ð4nÞ
þ
þ 4H
½XIX
that is, as a competition between the metal ion and
hydrogen ions for binding with L2 . The stability of
binding of Mnþ with L2 is measured by its stability
constant k1, which is the equilibrium constant for the
reaction:
Mnþ þ L4 "MLð4nÞ
k1 ¼
½MLð4nÞ ½Mnþ ½L4 ½XX
½3
where [ ] denotes concentrations (better, activities).
Some typical stability constants are given in Table 5.
Metals forming weaker complexes, therefore, require
less acidic (i.e., higher pH) conditions for complex
formation. The optimum pH values for the titration
of a number of metal ions are included in Table 5.
Metals forming stronger complexes can be titrated at
lower pH values, at which the weaker complexing
metals do not react, thus selective titration of, for
example, bismuth can be carried out in the presence
of lead at pH 1–2. The reaction of metal ions with
EDTA (reaction (XIX)) generates H þ . Thus, to
prevent a pH change during the titration, the solution must be adequately buffered.
N(CH2COOH)3
CH2COOH
N
CH2COOH
DCTA
NTA
Figure 7 Two alternative titrants to EDTA.
Adjustment of pH often will not give sufficient
selectivity and, of course, is inappropriate if a weakercomplexing metal (MW) has to be titrated in the
presence of a more strongly complexing metal (MS).
In such circumstances it is possible to ‘mask’ MS by
adding another complexing agent that complexes
much more strongly with MS than with MW, so that
MW but not MS will react with EDTA. The use of
such masking agents is widespread in compleximetric
titrimetry. A typical example is the use of triethanolamine to mask iron(III) when calcium is titrated with
EDTA in alkaline solution.
A wide range of visual indicators is available for
compleximetric titrations. These generally function
by forming a colored complex with the metal ion
being titrated, which causes a color change when the
metal ion is removed from the complex by reaction
with EDTA and releases the free ligand. These indicators are described in detail in another article.
Other EDTA-type compounds are sometimes used
as titrants, including 1,2-diaminocyclohexaneN,N,N0 ,N0 -tetraacetic acid, which generally forms
stronger complexes than EDTA, and nitrilotriacetic
acid, which generally forms weaker complexes than
EDTA (Figure 7).
TITRIMETRY / Overview 113
Nonaqueous Titrimetry
Most titrations are carried out in aqueous solution,
including all those described above. In some circumstances, however, it is advantageous to use other
solvents, especially organic solvents. Such nonaqueous titrations are normally used for acid–base reactions, but redox reactions may also be applicable.
The Karl–Fischer titration of water, in particular,
is based upon redox reactions in a nonaqueous
medium.
The ionization of a molecule HB in a solvent S is
influenced by the solvation of the ions:
HB þ nS"HSþ þ BS
½XXI
The ease of dissociation to form HS þ (solvated H þ )
increases with increasing basicity of the solvent, i.e.,
N2
N2
with increasing binding strength between H þ and
the solvent. Thus, an acid that is very weak in aqueous solution will be stronger in a more basic solvent
such as pyridine or dimethylformamide, and will give
a bigger ‘pH’ change on titration. Phenols, for example, which are too weak acids to be titrated in
aqueous solution, can be titrated in pyridine solution
with tetrabutylammonium hydroxide in benzene–
methanol (9:1, v/v) as titrant, and thymolphthalein in
methanol as indicator.
Similarly, bases that are very weak in aqueous
solution (e.g., amines) show increased basicity in
solvents of greater acidity, such as anhydrous acetic
acid. Perchloric acid in acetic acid may be used
as the titrant, with crystal violet in acetic acid as
indicator.
Because many of the solvents used are aggressive,
volatile, and obnoxious, nonaqueous titrations are
normally carried out in a closed environment, which
also minimizes the ingress of moisture (Figure 8). It is
essential to ensure that all apparatus used is dry and,
especially for titrations in basic solvents, a stream of
nitrogen is used to prevent access of carbon dioxide
to the solution being titrated. The titrant is stored in
a reservoir connected directly to the burette.
Compounds that may be determined by nonaqueous titrimetry include amines, amino acids, phenols,
and Schiff’s bases. Carbonyl compounds (by oxidation and titration of the released H þ ) can also be
determined. Such titrations are especially useful in
the pharmaceutical industry.
See also: Indicators: Acid–Base; Redox; Complexometric, Adsorption, and Luminescence Indicators. pH.
Quality Assurance: Internal Standards. Water Determination.
Further Reading
Figure 8 Apparatus for visual titration with nonaqueous
solvents. Automatic burette protected with guard tubes; a guard
tube and blow-bulb are attached to the bottom outlet of the
burette. The nitrogen flow can be omitted for titrations in acidic
solvents. (Redrawn from Belcher et al. 1970.)
Belcher R, Nutten AJ, and Macdonald AMG (1970) Quantitative Inorganic Analysis, 3rd edn. London: Butterworths.
Kolthoff IM and Belcher R (1957) Volumetric Analysis
(Redox Titrations), vol. III. New York: Interscience.
Kolthoff IM and Stenger VA (1947) Volumetric Analysis,
vols. I and II. New York: Interscience.
Kucharsky J and Safarik L (1965) Titrations in Non-Aqueous Solvents. Amsterdam: Elsevier.
Pribil R (1982) Applied Complexometry. Oxford: Pergamon.
Schwarzenbach G and Flaschka H (1969) Complexometric
Titrations, 2nd English edition. London: Methuen.
West TS (1969) Complexometry with EDTA and Related
Reagents. Poole: BDH Chemicals.