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Unit 18 Electroanalytical Methods

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Unit 18 Electroanalytical Methods 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 Meas urement 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, 5 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 p H = 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 = −log10 [H ] ..… (18.1) 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 Kw is the ionic product constant of water, its value is 1 × 10 at 25 C. Taking logarithm of both sides of equilibrium of equation 18.2 and substituting 'p' for negative logarithm we get 0 pH + pOH = pKw = 14, at 25 C . .… (18.3) For pure water [H+] = [OH−] = 1×10−7 (at 250C), which gives the pH value of pure water equal to 7 at this temperature. 6 For an acidic solution [H+] > [OH−] and pH is below 7, whereas for a basic solution Electroanalytical [OH−] > [H+] and pH is above 7. Methods 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 ind icators method which also known as colorimetric method for the measurement of the pH. 18 .2.3 Colo rimetric 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 Measu rement 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 absorbanc e 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) Colour A Colour B Our eyes can generally detect only one colour The dissociation constant K of indicator HIn is if the ratio of the concentration of the + - two colour forms is [H] [In] K = ....... (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) (pH) = (pH) ± ( )us( ) ..… (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)u = cell potential of glass electrode and reference electrode system with unknown solution (Ecell)s = cell potential of glass electrode and reference electrode system with standard buffer In order to understand this operational definition of pH, we will take up general principles of potentiometry. 8 18.3.1 Principle of Potentiometry Electroanalytical Methods Potentiometry deals with the measurement of difference in potential between two electrodes which have been combined to form an electrochemical cell.
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