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UNIT 18 ELECTROANALYTICAL METHODS
Structure
18.1 Introduction
Objectives
18.2 pH Metry
Definition of pH
Measurement of pH
Colourimetirc Measurement of pH
18.3 Electrometric Measurement of pH
Principle of Potentiometry
Electrodes
Measurement of pH using pH Meter
pH of Water and Waste Water
Acid Rains and pH
pH of Soils
18.4 Ion Selective Electrodes
18.5 Counductometry
Some Basic Concepts of Conductometry
18.6 The Measurement of Conductance
The Wheatstone Bridge Principle
Measurement of Conductance of a Solution
Experimental Measurement
18.7 Application of Conductometry
18.8 Summary
18.9 Terminal Questions
18.10 Answers
18.1 INTRODUCTION
Electroanalytical methods find applications in all branches of Chemistry, industries,
engineering and a number of other technologies. The possibility of the determination
of low level of pollutants has prompted the use of these methods in environmental
studies.
An electroanalytical method can be defined as one, in which the electrical response of
a chemical system or sample is measured. These methods can be classified into a
number of types characterized by measuring the electrical response in terms of
different electrical quantities such as: potential, current, quantity of current, resistance
and voltage etc. and bear the corresponding names as potentiometry, amperometry,
coulometry, conductometry and voltammetry etc. During the past few years, there has
been sudden increase in interest in electroanalytical techniques. This is partially
attributed to the development in instrumentation and partially due to the heavy
demands by environmental scientists for the determination of a large number of heavy
metal, organic and inorganic substances present in water and soil samples.
In this unit we will study how to measure pH of water and soil samples using pH
metry. We will also discuss the potentiometric measurement of concentration of ions
selectively with the help of ion selective electrodes. Then we will discuss
conductometry.
Objectives
After studying this unit, you will be able to:
• define pH,
• define electrode potential,
• describe the use of some electrodes,
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Instrumental Methods • measure the pH of a solution,
of Analysis • define conductivity,
• measure the conductivity of a solution,
• apply the concept of pH metry, ion selective potentionmetry and conductivity for
water and soil analyses
18.2 pH METRY
There is a widespread usage of electrochemical methods in general and of
potentiometric determination of pH and concentration of several ions in particular.
Measurement of pH is one of the most important and widely used test in water
analysis. For natural water treatment as well as for waste water treatment a large
number of reactions e.g. coagulation, disinfection, water softening, acid base
neutralisation etc. are all pH dependent. Most of chemical laboratories are equipped
with pH meters. Modernization of potentiometry by the development of ion selective
electrodes has increased the interest in the study of environmental samples.
The principle of potentiometry is applied to measure the potential difference in terms
of pH unit on pH scale by suitably modifying the common voltmeter to high input
impedance mV meter and such pH measurement can be termed as pH metry instead of
potentiometry. In pH metry, pH meter is used to measure the pH. Before going in
further details of potentiometric measurement of pH, let us know the basic concept of
pH.
18.2.1 Definition of pH
The hydrogen ion concentration plays an important role in many areas of chemistry
A rigorous definition of and its determination and control is of great practical value in the study of
pH would obviously environment.
involve activities,
accordingly: The shorthand notation of hydrogen ion concentration is given in terms of pH for
pH = −log a +
a H 'puissance de hydrogen'.
The pH value, originally formulated in 1909 by S.P. Sorensen, is defined as the
negative logarithm of hydrogen ion concentration:
pH = −log [H+] ..… (18.1)
10
where [ ] represents equilibrium concentration and logarithm is taken to the base 10.
In practice 'p' preceding a variable is used to express the negative logarithm of that
variable. Likewise, pOH is to designate the negative logarithm of hydroxyl ion
concentration.
+
−
In aqueous solutions the product of [H ] and [OH ] is always a constant at a particular
temperature. Thus,
+ −
KW =[H ] [OH] ..... (18.2)
−14 0
where K is the ionic product constant of water, its value is 1 × 10 at 25 C.
w
Taking logarithm of both sides of equilibrium of equation 18.2 and substituting 'p' for
negative logarithm we get
pH + pOH = pK = 14, at 250C . .… (18.3)
w
+ 7 0
− −
For pure water [H ] = [OH] = 1×10 (at 25 C), which gives the pH value of pure
water equal to 7 at this temperature.
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+ − Electroanalytical
For an acidic solution [H ] > [OH] and pH is below 7, whereas for a basic solution
− + Methods
[OH] > [H ] and pH is above 7.
Neutral
Acidic Range Basic Range
0 7 14
pH Scale
SAQ 1
+
Find the concentration of H ions of a solution for which pH value is 4.5.
…………………………………………………………………………………………
…………………………………………………………………………………………
………………………………………………………………………………………...
18.2.2 Measurement of pH
The pH of a solution is commonly found by the use of either an indicator or a pH
meter. Because of their accuracy and speed, pH meters have superseded the older
indicator method in many applications. However, the indicator method remain in use
because it is simple and convenient specially for field work in pollution analysis. In
next part of this section we are taking a brief description of indicators method which
also known as colorimetric method for the measurement of the pH.
18.2.3 Colorimetric Measurement of pH
For the approximate and rapid estimation of pH and in studies in non-aqueous media,
it is convenient to make use of coloured indicators. Colorimetric measurement can be
carried out visually or photometrically.
Visual Measurement of pH
The use of coloured indicators for the visual measurement of pH is well known. The
approximate pH of a solution can be determined by comparing its reaction with
different indicators or on papers impregnated with the indicator solution. In this
method the colour change is observed in a particular pH range. The chief advantage is
the low cost and also the method is suitable for routine pH measurement. A very
common example is litmus which is red below pH 5 and blue above pH 8. The colour
changes from red to blue when pH changes from 5 to 8. To find colour changes in a
wide range of pH, the mixtures of indicators, the so called universal indicators are to
be used. For example, the Kolthoff universal indicator is a mixture of five indicators
and gives a conspicuous colour change within unit pH values. The colours at different
pH values are given in Table 18.1.
Table 18.1: Variation of colour of Kolthoff Universal Indicator with change
in pH.
pH 1 2 3 4 5
Colour R R-P R-O O Y-O
pH 6 7 8 9 10
Colour L-Y Y-G G G-B V
Abbreviations: R=Red, P=Pink, O=Orange, Y=Yellow, LY=Light Yellow,
G=Green, B=Blue & V=Violet 7
Instrumental Methods Photometric Measurement of pH
of Analysis
The visual method for pH measurement using indicators has low accuracy due to
difficulties of light intensity estimation. The accuracy can be increased by
instrumental means using a colorimeter or a spectrophotometer to measure the
absorbance at a particular wavelength. Indicators are considered to behave as weak
acids or weak bases and the degree of dissociation of indicator substance depends on
hydrogen ion concentration in solution. Consider, e.g. an indicator acid, HIn, which
dissociates as
+ −
HIn H + In ..… (18.4)
Our eyes can generally Colour A Colour B
detect only one colour
if the ratio of the The dissociation constant K of indicator HIn is
concentration of the + -
two colour forms is K = [H] [In] ....... (18.5).....18.6
10:1. Only the colour [HIn]
of the more
concentrated form is −
+ [In]
seen. logK =+log[H]log...... (18.6)
[HIn]
−
[In]
orpH=+pK log...... (18.7)
[HIn]
−
Indicator colours are indicated by the In and HIn concentration ratio which depends
on degree of dissociation and hence the pH can be indicated by the intensity of either
colour A or colour B with the assumption that the Beer's law is obeyed. To get
satisfactory results by photometric measurement, it is necessary to keep the indicator
concentration as small as possible. The principle of photometric measurement is
discussed in detail in Unit 19 of this course. In next section, we will take up the
principle of pH metry.
18.3 ELECTROMETRIC MEASUREMENT OF pH
The electrometric method of pH determination is based on the measurement of
potential of a pH cell, whereby the potential of a hydrogen sensitive electrode is
directly proportional to pH, and pH is defined in an operational manner on a
potentionmetric scale.
The pH meter is calibrated potentiometrically with an indicator electrode (glass) and a
reference electrode using a standard buffer. The operational pH is defined as:
EE−
( ) ( )
(cell)(cell)
us
(pH) = (pH) ± ..… (18.8)
u s 0.0591
where
(pH)u = potentiometrically measured pH of the sample (unknown solution)
(pH)s = assigned pH of the standard buffer used for calibration
(E ) = cell potential of glass electrode and reference electrode system with
cell u
unknown solution
(E ) = cell potential of glass electrode and reference electrode system with
cell s
standard buffer
In order to understand this operational definition of pH, we will take up general
principles of potentiometry.
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