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Titrimetric Analysis

Quantitative chemical analysis carried out
determining the volume of a solution
accurately known concentration which
required to react quantitatively with
measured volume of the substance to
determined.
by
of
is
a
be
Classification

Neutralisation Reactions

Complex Formation Reactions

Redox Reactions

Precipitation Reactions
Basics

Equivalence and end points

Standards
Basics

Equivalence and end points
 Precise
and accurate titrations require the
reproducible determination of the end point which
either corresponds to the stoichiometric point of the
reaction or bears a fixed and measurable relation to
it.
Basics

Equivalence and end points
 Monitor
a property of the titrand which is
removed at the end point.
 Monitor a property which is readily
observed when excess titrant has been
added.
 Two main methods
 Coloured
indicators
 Electrochemical techniques.
Basics

Colour Change Indicators
 Common
to a wide variety of titrations.
 In
general terms a visual indicator is a compound that
changes from one colour to another as its chemical form
changes.
= InB + nX where X may be H+, Mn+ or e-, and the
colour is sensitive to the presence of H+, Mn+, oxidants
or reductants.
 InA
Basics

An indicator constant is defined as:
KIn = [InB][X]n / [InA]
[X]n = KIn ([InB] / [InA])
npX = pKIn + log10([InB] / [InA])
pH = pKa + log10 ([InB] / [InA])
Basics

Potentiometric Measurements

Measuring the change in potential
during the titration.
 Acid-base
titrations.
 Precipitation
 Redox
titrations.
titrations.
Basics

Monitor the change of Ecell during the
course of a titration where the indicator
electrode responds to one of the reactants or
the products.

A plot of Ecell against the volume of titrant is
obtained.

Precision of better than 0.2%.
Basics
Basics
 The
Nernst Equation
aA + bB + …+ ne- = xX + yY + ...
x[Y]y...
[X]
RT
0
E=E ln
a[B]b...
[A]
nF
Basics

RT/F ln 10 = 0.059158 V thus:
E = E0 - (0.059 V/n) log10
 And
[X]x[Y]y...
[A]a[B]b...
E = E0 at unity concentrations
Basics

Conductimetric Indication
 The
electrical conductance of a solution is a measure
of its current carrying capacity and is determined by
its total ionic strength.
 It
is a non-specific property.
 Conductance
is defined as the reciprocal of resistance
(Siemans, -1).
Basics

A conductance cell consists of two
platinum electrodes of large surface area.

5-10 V at 50 -10,000 Hz is applied.

Control of temperature is essential.
Basics

Acid-base titrations especially at trace levels.

Relative precision better than 1% at all levels.

Rate of change of conductance as a function of added
titrant used to determine the equivalence point.

High concentrations of other electrolytes can
interfere.
Basics
Basics

Standards
 Certain
chemicals which are used in defined
concentrations as reference materials.
 Primary
standards.
 Secondary
standards.
Basics

Primary Standards
 Available
in pure form, stable and easily dried
to a constant known composition.
 Stable
 High
in air.
molecular weight.
 Readily
soluble.
 Undergoes
stoichiometric and rapid reactions.
Basics
 Acid-base
 Na2CO3,
reactions.
Na2B4O7, KH(C8H4O4), HCl (cbpt.)
 Complex
formation reactions.
 AgNO3, NaCl
 Precipitation
reactions.
 AgNO3, KCl
 Redox
reactions.
 K2Cr2O7, Na2C2O4, I2
Basics

Secondary Standards
A
substance that can be used for
standardisations, and whose concentration of
active substance has been determined by
comparison to a primary standard.
Classification

Neutralisation Reactions

Complex Formation Reactions

Redox Reactions

Precipitation Reactions
Neutralisation Titrations

The neutralisation reactions between acids
and bases used in chemical analysis.

These reactions involve the combination of
hydrogen and hydroxide ions to form water.
Neutralisation Titrations

For any actual titration the correct end point
will be characterised by a definite value of
the hydrogen ion concentration.

This value will depend upon the nature of the
acid and the base, the concentration of the
solution and the nature of the indicator.
Neutralisation Titrations

A large number of substances called neutralisation indicators change colour according to the
hydrogen ion concentration of the solution.

The end point can also be determined
electrochemically by either potentiometric or
conductimetric methods.
Theory of Indicator Behaviour

An acid/base indicator is a weak organic acid or a
weak organic base whose undissociated form
differs in colour from its conjugate base or
conjugate acid form.

The behaviour of an acid type indicator is
described by the equilibrium;
Theory of Indicator Behaviour
In- + H3O+

HIn + H2O

The behaviour of an base type indicator
is described by the equilibrium;

In + H2O
InH+ + OH-
Theory of Indicator Behaviour

The equilibrium constant takes the form:
[H3O+][In-]
= Ka
[HIn]

Rearranging:
-]
[HIn
[H3O+] = Ka
[In-]
Theory of Indicator Behaviour

pH (acid colour) = -log(Ka . 10) = pKa +1

pH (base colour) = -log(Ka / 10) = pKa -1

Therefore; indicator range = pKa ± 1
Theory of Indicator Behaviour

The human eye is not very sensitive to colour
change in a solution containing In- and HIn.

Especially when the ratio [In-] / [HIn] is greater
than 10 or less than 0.1.

Hence the colour change is only rapid within the
limited concentration ratio of 10 to 0.1.
Theory of Indicator Behaviour
Neutralisation Titrations

Strong acids and bases

Weak acids

Weak bases

Polyfunctional acids

Applications
Neutralisation Titrations
Strong acids and bases.
 When both reagent and analyte are strong
electrolytes, the neutralisation reaction can
be described by the equation:

H3O+ + OH-
2H2O
Neutralisation Titrations

The H3O+ concentration in aqueous solution
comprises of two components.
 The
reaction of the solute with water.
 The
dissociation of water.

[H3O+] = CHCl + [OH-] = CHCl

[OH-] = CNaOH + [H3O+] = CNaOH
Neutralisation Titrations

Using these assumptions you can calculate the
pH of a titration solution directly from
stoichiometric calculations and therefore
simulate the titration curves.

This is useful in determining the correct
indicator for a new titration.
Neutralisation Titrations
Neutralisation Titrations

Examples:
 HCl,
HNO3
 NaOH,
KOH, Na2CO3
 Standards:anhydrous
boiling HCl.
Na2CO3 and constant
Neutralisation Titrations

Weak acids and bases

Examples
 Ethanoic
acid
 Sodium cyanide

Four types of calculation are required to derive
a titration curve for a weak acid or base.
Neutralisation Titrations
 Solution
contains only weak acid. pH is calculated from
the concentration and the dissociation constant.
 After
additions of the titrant the solution behaves as a
buffer. The pH of each buffer can be calculated from
there analytical concentrations.
 At
the equivilence point only salt is present and the pH is
calculated from the concentration of this product.
 Beyond
the equivilence point the pH is governed largely
by the concentration of the excess titrant.
Neutralisation Titrations

Effect of Concentration

Effect of reaction completeness

Indicator choice; Feasibility of
titration
Neutralisation Titrations
Neutralisation Titrations

Polyfunctional acids and bases

Typified by more than one
dissociation reaction.
Neutralisation Titrations

Phosphoric acid

Yield multiple end points in a titration.
Neutralisation Titrations
Neutralisation Titrations

Sulphuric Acid

Unusual because one proton behaves as a strong
acid and the other as a weak acid (K2 = 1.20 x
10-2).
Neutralisation Titrations

Applications:
 Determination
of the concentration of
analytes which are either acid or bases.
 Determination of analytes which can be
converted to acids or bases.
Complexometric Titrations

Titrations between cations and complex
forming reagents.

The most useful of these complexing agents
are organic compounds with several electron
donor groups that can form multiple
covalent bonds with metal ions.
Complexometric Titrations

Most metal ions react with electron-pair
donors to form coordination compounds or
complex ions.

The donor species, or LIGAND, must have
at least one pair of unshared electrons
available.
Complexometric Titrations

Inorganic Ligands
 Water
 Ammonia
 Halides

Organic Ligands
 Cyanide
 Acetate
Complexometric Titrations

The number of bonds a cation forms
with an electron donor is called the
COORDINATION NUMBER.

Typical values are 2, 4 and 6.

The species formed as a result of
coordination can be electrically positive,
neutral or negative.
Complexometric Titrations

Complexometric methods have been
around for more than a century.

Rapid expansion in the 1940’s based on
a class of coordination compounds
called CHELATES.
Complexometric Titrations

A chelate is produced when a metal ion
coordinates to two or more donor groups
within a single ligand.

For example the copper complex of glycine.
Complexometric Titrations
NH2
Cu2+ + 2 H
C
C
OH
H
O
C
O
O
O
C
Cu
C N
H2 H2
O
+ 2H +
N C
H2 H2
Complexometric Titrations

A ligand with a single donor group is
called unidentate.

Glycine is bidentate.

Tri, tetra, penta and hexadentate
chelating agents are also known.
Complexometric Titrations

Multidentate ligands have two advantages
over unidentate ligands.

They react more completely with cations to
provide a sharper endpoint.

The reaction is a single step process.
Complexometric Titrations

Tertiary amines that also contain carboxylic acid
groups form remarkably stable chelates with many
metal ions.

Ethylenediaminetetraacetic Acid EDTA
HOOCCH2
CH 2CO O H
N
HOOCCH2
CH2
CH2
N
CH 2CO O H
Complexometric Titrations

EDTA can complex a large number of metal ions.

Approximately 40 cations can be determined by
direct titration.

EDTA is usually used as the disodium salt,
Na2H2EDTA
H2EDTA2- + M2+ [M(EDTA)]2- + 2H+
Complexometric Titrations

Because EDTA complexes most cations, the
reagent might appear at first glance to be totally
lacking in selectivity.

However, great control can be acheived by pH
regulation and the selection of suitable indicators.
Complexometric Titrations

Indicators are generally complexing agents which
undergo a colour change when bonded to a metal
ion.
H2EDTA2- + [M(Ind)] [M(EDTA)]2- + Ind2- + 2H+
Complexometric Titrations

Typical indicators are:
 Murexide
 Solochrome
black
 Calmagite
 Bromopyrogallol
 Xylenol
orange
red
Complexometric Titrations

Typical applications:
 Determination
of cations
 Hardness of water
Redox Titrations
Basics
 Potassium Permanganate
 Potassium Dichromate
 Cerium IV
 Iodine

Redox Titrations

Basics
 Electrode
 Indicators
Potentials
Redox Titrations

Electrode Potentials
 Derived
from Nernst equation.
 Calculations of cell potentials leads to
theoretical titration curves.
 EOX = ERED = Esystem
Redox Titrations
Redox Titrations

Indicators
 Potentiometric
 Coloured
 EIn
indicators
= EOX = ERED = Esystem
 Specific:
Starch
 Oxidation
/ Reduction Indicators
Redox Titrations
Redox Titrations

Potassium Permanganate
MnO4- + 8H+ + 5e-  Mn2+ + 4H2O

Standardisation
 Sodium

oxalate or arsenic (III) oxide
Many Analyses
Redox Titrations

Hydrogen Peroxide:


2MnO4- + 5H2O2 + 6H+  2Mn2+ + 5O2 + 8H2O
Nitrites:

2MnO4- + 5NO2- + 6H+  2Mn2+ + 5NO3- + 3H2O
Redox Titrations

Persulphates:
 Add
an excess of iron (II)
 S2O82- + 2Fe2+ + 2H+  2Fe3+ + 2HSO4 The
excess iron (II) is determined by back titration against
standardised permangenate.
 MnO4- + 8H+ + 5Fe2+  Mn2+ + 5Fe2+ + 4H2O
Redox Titrations

Potassium Dichromate
CrO72- + 14H+ + 6e-  2Cr3+ + 7H2O

Standardisation
 Against
metallic iron
 1 mole K2CrO7 = 6 moles Fe
Redox Titrations

Iron (II):
CrO72- + 14H+ + 6Fe2+  2Cr3+ + 6Fe3+ 7H2O
 Indicators
include N-phenylanthranilic acid and sodium
diphenylamine sulphonate.

Chlorates:
 Reduced
 The
with an excess of iron (II)
excess iron (II) is determined by back titration
against standardised dichromate.
Redox Titrations

Iodine

Iodometric Titrations
I2 + 2e- 2I-

Standardisation
 Standardised
sodium thiosulphate or
arsenic (III) oxide

Many Analyses
Redox Titrations

Hydrogen Peroxide:


Thiosulphates:


H2O2 + 2I- + 2H+ I2 + 2H2O
2S2O32- + I2  S4O62- + 2I-
Hydroxyl Groups:

2OH- + I2  IO- + H2O + 2I-
Redox Titrations

Others:
 Copper
 Dissolved
oxygen
 Chlorine
 Arsenic
(V)
 Sulphides
 etc..........
Redox Titrations

Cerium (IV) Sulphate

Very strong oxidising agent (1.43V)
Ce4+ + e- Ce3+

Standardisation
 Sodium

oxalate or arsenic (III) oxide
Many Analyses
Precipitation Titrations

Titrations between analytes and reagents
resulting in the formation of a precipitate.

The most useful of these precipitating
reagents is silver nitrate.

Titrimetric methods based upon the use
of silver nitrate are sometimes called
Argentometric titrations.
Precipitation Titrations

Used for the determination of many
anions including:
 halides
 divalent
anions
 mercaptans
 certain fatty acids
Precipitation Titrations

Precipitation titrations are based on
the SOLUBILITY PRODUCT of the
salt, KSP.

The smaller KSP, the less soluble the
silver salt and the easier it is to
determine the endpoint
Precipitation Titrations

Endpoint determination is by
coloured indicators (usually back
titrations) or turbidity methods.

The most accurate is the
VOLHARD METHOD.
Precipitation Titrations

VOLHARD METHOD

A back titration of thiocyanate ions
against the excess silver ions using an
iron (II) salt as the indicator.
Precipitation Titrations
Ag+ + SCN-
AgSCN
Fe3+ + SCN-
FeSCN2+
Blood Red