A Laboratory Manual For Environmental Chemistry 9789389583588

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A Laboratory Manual for TM

Environmental Chemistry

A Laboratory Manual for

PART I – WATER ANALYSIS: Principles of Spectrometric Analysis / Sampling Techniques / Colour / Turbidity / Total Dissolved Solids / Conductance / pH Value / Acidity / Alkalinity / Dissolved Oxygen / Chemical Oxygen Demand / Biochemical Oxygen Demand / Dissolved Organic Carbon / Calcium Hardness of Water / Total Hardness of Water / Nitrate Nitrogen / Nitrate Nitrogen Brucine Method / Ammonia Nitrogen / Fluoride / Fluoride in Drinking Water (by Fluoride Electrode) / Chloride / Residual Chlorine / Sulphide / Sulphite / Sulphate / Phosphate / Iron / Chromium / Copper / Lead / Nickel / Cadmium / Mercury / Cyanides / Boron / Selenium / Arsenic / Suspended Solids / Determination of Chlorophyll-a / Determination of Productivity PART II – SOIL ANALYSIS: Collection of Soil Samples / pH Value / Soluble Salts / Identification of Calcium Carbonate, Nitrogen, Phosphorus and Potassium / Calcium Carbonate / Gypsum Requirement / Lime Requirement / Alkalinity / Organic Carbon / Total Nitrogen / Available Nitrogen / Ammonium Nitrogen / Nitrate Nitrogen / Nitrite Nitrogen / Available Phosphorus / Available Sulphur / Silica / Available Potassium / Sodium / Calcium and Magnesium / Micronutrient Elements / Pesticide Residues / Separation and Identification of Pesticide Residues from the Soil PART III – AIR ANALYSIS: Qualitative Analysis of CO, CO2, H2S, SO2, NO and NO2 / Carbon Monoxide / Carbon Dioxide / Hydrogen Sulphide / Sulphur Dioxide / Nitric Oxide and Nitrogen Dioxide / Particulates

Calcium in Egg Shells / Effect of pH on Germination of Seeds / Effect of Copper Salt on Germination of Seeds / Phosphate in Detergents / Benzoic Acid and Sorbic Acid in Soft Drinks / Sulphur Dioxide in Soft Drinks / Identification of Food Adulterants / Appendices / References / Index.

978-93-89583-58-8

` 315/9 789389 583588

Gopalan | Anand | Sugumar

PART IV – MISCELLANEOUS:

Environmental Chemistry

Contents

A Laboratory Manual for

The present book is meant for the students who opt for a course in “Environmental Chemistry” with laboratory work as a component of the course. Spread in 72 experiments, the analyses of soil, water and air have been described in a simple manner, so that most of these experiments can be conducted even by the beginners. The principles involved, preparation of the reagents, and the procedures are described for each experimental method. The authors hope that this manual would be useful in laboratories where soil, water and air are routinely tested.

Environmental Chemistry

Distributed by:

TM

A Laboratory Manual for Environmental Chemistry

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A Laboratory Manual for Environmental Chemistry

Gopalan Former Head of the Department of Chemistry,. Madras Christian College, Chennai

Amirtha Anand Reader in Chemistry, Maitreyi College, New Delhi

R. Wilfred Sugumar Head of the Department of Chemistry, Madras Christian College, Chennai

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©Copyright 2019 I.K. International Pvt. Ltd., New Delhi-110002. This book may not be duplicated in any way without the express written consent of the publisher, except in the form of brief excerpts or quotations for the purposes of review. The information contained herein is for the personal use of the reader and may not be incorporated in any commercial programs, other books, databases, or any kind of software without written consent of the publisher. Making copies of this book or any portion for any purpose other than your own is a violation of copyright laws. Limits of Liability/disclaimer of Warranty: The author and publisher have used their best efforts in preparing this book. The author make no representation or warranties with respect to the accuracy or completeness of the contents of this book, and specifically disclaim any implied warranties of merchantability or fitness of any particular purpose. There are no warranties which extend beyond the descriptions contained in this paragraph. No warranty may be created or extended by sales representatives or written sales materials. The accuracy and completeness of the information provided herein and the opinions stated herein are not guaranteed or warranted to produce any particulars results, and the advice and strategies contained herein may not be suitable for every individual. Neither Dreamtech Press nor author shall be liable for any loss of profit or any other commercial damages, including but not limited to special, incidental, consequential, or other damages. Trademarks: All brand names and product names used in this book are trademarks, registered trademarks, or trade names of their respective holders. Dreamtech Press is not associated with any product or vendor mentioned in this book. ISBN: 978-93-89583-58-8 EISBN: 978-93-90078-95-0

Preface The space available for human habitation is the space just above the surface of the earth. It is an extremely small fraction of the total volume of our universe. Therefore, pollution of air, water and soil within this space would endanger the very survival of mankind. Rapid industrialisation all over the world is mainly responsible for this undesirable ‘environmental pollution’. Due to increasing industrialisation, pollution control is a major concern and challenge for the people of 21st century. The importance of this subject has been realized as a result of which ‘Environmental Chemistry’ is being introduced as a subject in schools, colleges and universities. This book “A Laboratory Manual for Environmnetal Chemistry” has been written in order to cater to the needs of the students who opt for a course in “Environmental Chemistry” with laboratory work as a component of the course. In this manual, the the analyses of soil, water and air have been described in a simple mannner so that most of these experiments can be conducted even by the beginners in this subject. The principle involved, preparation of the reagents and the procedure are described for each experimental method. The authors sincerely hope that this manual would prove to be useful in laboratories where soil, water and air are routinely tested. The authors are thankful to Dr. (Mrs) Shanthi Kesavan, Reader in Chemisty, Maitreyi College for performing and standardizing some of the experiments and to I.K. International Publisher for bringing out this manual neatly. The authors dedicate this book to the memory of Prof. K.S.Nagarajan. Suggestion for further improvement of this book will be greatly appreciated.

R. Gopalan Amirtha Anand R. Wilfred Sugumar

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Contents

1

1. Principles of Spectrometric Analysis

1

PART ONE – WATER ANALYSIS 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30.

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Sampling Techniques Colour Turbidity Total Dissolved Solids Conductance pH Value Acidity Alkalinity Dissolved Oxygen Chemical Oxygen Demand Biochemical Oxygen Demand Dissolved Organic Carbon Calcium Hardness of Water Total Hardness of Water Nitrite Nitrogen Nitrate Nitrogen (Brucine Method) Ammonia Nitrogen Fluoride Fluoride in Drinking Water (By Fluoride Electrode) Chloride Residual Chlorine Sulphide Sulphite Sulphate Phosphate Iron Chromium Copper Lead

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7 10 12 14 15 17 21 23 26 29 32 35 36 38 41 44 46 49 51 53 55 57 59 61 63 65 67 69 71

viii

31. 32. 33. 34. 35. 36. 37. 38. 39. 40.

Contents

Nickel Cadmium Mercury Cyanides Boron Selenium Arsenic Suspended Solids Determination of Chlorophyll- a Determination of Productivity

73 75 77 79 81 83 85 87 88 89

PART TWO – SOIL ANALYSIS 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. 61. 62. 63.

Collection of Soil Samples pH Value Soluble Salts Identification of Calcium Carbonate, Nitrogen, Phosphorus and Potassium Calcium Carbonate Gypsum Requirement Lime Requirement Alkalinity Organic Carbon Total Nitrogen Available Nitrogen Ammonium Nitrogen Nitrate Nitrogen Nitrite Nitrogen Available Phosphorus Available Sulphur Silica Available Potassium Sodium Calcium and Magnesium Micronutrient Elements Pesticide Residues Separation and Identification of Pesticide Residues from the Soil

93 95 97 99 102 104 107 110 111 115 118 120 122 124 126 128 130 132 134 135 136 138 140

PART THRE – AIR ANALYSIS 64. Qualitative Analysis of CO, CO2, H2S, SO2, NO and NO2

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145

Contents

65. 66. 67. 68. 69. 70.

ix

Carbon Monoxide Carbon Dioxide Hydrogen Sulphide Sulphur Dioxide Nitric Oxide and Nitrogen Dioxide Particulates

147 148 150 152 154 156

PART FOUR – MISCELLANEOUS 71. 72. 73. 74. 75. 76. 77.

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Calcium in Egg Shells Effect of pH on Germination of Seeds Effect of Copper Salt on Germination of Seeds Phosphate in Detergents Benzoic Acid and Sorbic Acid in Soft Drinks Sulphur Dioxide in Soft Drinks Identification of Food Adulterants Appendices References

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159 161 163 165 167 169 170 175 191

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Principles of Spectrometric Analysis

Photometric methods provide a simple means for determining minute quantities of constituents. The variation in the intensity of colours of the system with variation in concentration of some component is the basis of this analysis. Generally, colour is developed due to the formation of a coloured compound when an appropriate reagent is added. Colorimetric method involves determining the concentration of a substance by measuring its relative absorption of light with respect to a known concentration of the same substance. Therefore, it is also regarded as absorption photometry in the visible region of light. The instruments normally used for this purpose are filter photometer and spectrophotometer. Filter photometer, which is also known as photoelectric colorimeter, is employed with light contained within a narrow range of wavelength (l) provided by passing white light through filters. Filters are materials in the form of plates of gelatin, glass, etc., and these transmit light of only a limited spectral region. In spectrophotometer, light of definite wavelength (l) extending to the ultraviolet region of the spectrum constitutes the source of light. Here, the photoelectric cells are used to measure the light transmitted from the solution. These analyses are governed by Lambert’s law and Beer’s law. The Lambert’s law states that when monochromatic light passes through a transparent medium, the rate of decrease in intensity of the incident light with the thickness of the medium, is proportional to the intensity of light. Beer’s law states that the intensity of a beam of monochromatic light decreases exponentially as the concentration of absorbing substance increases arithmetically. Combining both the laws, (I ) log 0 = Îbc ( It ) where Io and I t are the intensities of incident and transmitted lights, c is the concentration in mol L-1, b is the thickness of the medium or length of cell in cm and Î is a constant called the molar absorptivity. (I ) The ratio log 0 is known as absorbance (A) (or optical density) of the medium. ( It ) A = Îbc or A a c

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A Laboratory Manual for Environmental Chemistry

This is called the Beer-Lambert law and is the fundamental equation used in colorimetry. Photoelectric colorimeters or spectrophotometers are used for the measurement of concentrations of coloured solutions. These instruments have either a set of filters to be chosen visually or in-built filters. A number of companies manufacture these instruments and they provide the instruction manuals for use of their respective instruments. Determination of the concentration of a copper sulphate solution can be considered as an example in order to become more clear. Procedure: First fill the cell with pure solvent i.e., water in this case. Choose visually the filters of colour complimentary to the colour of copper sulphate solution. The table given below can be referred. Colour of the solution

Complimentary colour (colour of the filter)

Yellowish-green Yellow Orange Red Purple Violet Blue Greenish-blue Bluish-green

Violet Blue Greenish-blue Bluish-green Green Yellowish-green Yellow Orange Red

Wavelength of the colour transmitted (nm) 400-435 435-480 480-490 490-500 500-560 560-580 580-585 595-610 610-750

The possible filters are yellow, orange and red (in the wavelength range ±15 nm) . Insert one of these filters in the optical path and adjust the setting to get 100% transmittance or 0% absorbance for water. Now, discard the water from the cell and rinse the cell with the copper sulphate solution, the concentration of which is to be measured. Then fill the cell with this solution. Note down the transmittance/absorbance. Repeat these steps with each of the chosen filters adjusting each time with water for 100% transmittance. The filter for which the solution shows the least transmittance or the most absorbance (lmax) must be selected. Let us suppose it is orange filter. Use this filter for further work. Weigh accurately 0.2496 g CuSO4.5H2O and make upto 100 mL with distilled water in a standard flask. This is the stock solution. Prepare a series of solutions by pipetting 1, 2, 3, 4, 5 mL of the stock solution in separate 10 mL standard flasks and making upto the mark with distilled water. Shake the solutions well and determine the absorbance (or optical density) of each of the solutions. Plot a graph between absorbance and concentration. This is known as calibration curve. A straight line passing through the origin indicates that Beer-Lambert law is applicable. Now, measure the absorbance of the unknown solution and determine its concentration from the calibration curve. The concentrations may be in mol L–1 or in ppm (calculation is given in Appendix 1). In the instruments with in-built filters, the wavelengths can directly be selected and the above procedure repeated to find out the lmax. Let us suppose it is 600 nm. Select this value for further work.

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Principles of Spectrometric Analysis

3

ATOMIC SPECTROSCOPY Atomic spectroscopy is the principal tool for measuring metallic elements at parts per million and lower levels in environmental laboratories. This technique is automated with mechanical sample-charging devices that allow each instrument to turn out hundreds of analyses per day. In atomic spectroscopy, a liquid sample is aspirated through a tube into a flame that is hot enough to break the molecules into atoms. The concentration of an element in the flame is measured either by absorption or emission of radiation. In atomic absorption spectrophotometry (AAS), radiation of the correct frequency is passed through the flame and some of the incident light will be absorbed by the sample. The intensity of the transmitted radiation is measured to determine the concentration of the sample. For atomic emission spectrophotometry (AES) or flame emission spectrophotometry (FES), no external source of light is required. Radiation is emitted by atoms whose electrons have been promoted to excited states in the flame. The spectrum of gaseous atoms in a flame has extremely sharp lines with widths of 10-3 to 10-2 nm. Because the lines are so sharp, there is little overlap between the spectra of different elements in the same sample. The lack of overlap allows some instruments to measure over 60 elements in a sample simultaneously. The procedure for the preparation of standard solutions and drawing the calibration curve will be the same as in the case of colorimetric/spectrophotometric analysis. For operating the instrument, the instruction manual provided must be referred to. The Division of Analytical Chemistry of the International Union of Pure and Applied Chemistry (IUPAC) has made recommendations to standardize the terms used in spectrophotometry. According to these guidelines, single words are preferred. For example, wavelength instead of wave length and absorbance instead of optical density. The recommended terms, symbols and definitions are given in Table 1. Table 1 Name

Symbol

Spectrophotometry nomenclature Definition

Absorbance

A

log T

Absorptivity

a

= A/bca

Path length

b

Molar absorptivity

e

Internal cell length or sample length in cm = A/bcb

Transmittancec Wavelength unit

T nm mm

lt/lo 10-9 m 10-6 m

Name not recommended Optical density (OD), extinction, absorbency Absorbency index, extinction coefficient l or d Molar absorbency index Molar extinction coefficient Transmittancy, transmission mm (millimicron) m (micron)

a. The concentration is in g L–1 b. The concentration is in mol L–1 c. It/Io is the ratio of radiant power transmitted to radiant power incident.

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PART ONE

WATER ANALYSIS

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2 1

Sampling Techniques

Water sample for analysis should be collected in such a manner that the sample truly represents the water source or the main body of water. Tap water: The pipeline should be flushed before the sample is collected. Water obtained initially from a tap may contain contaminants from the pipeline, and also its temperature may be higher than that of the source. For analysis of natural or wastewater, the following principal types of sampling procedures are adopted: (1) Judgemental sampling is the result of a bias of the analyst and usually occurs when one collects the sample from a place where the concentration of the pollutants is thought to be high or low. This type of sampling is not usually representative of the entire site, but is useful in locating ‘Best case’ or ‘Worst case’ scenario of a pollutant source. (2) Systematic sampling usually involves dividing the site into equal sized areas and sampling each area. Creating a measured grid or regular pattern of sample sites is an easy way to set up a systematic sampling scheme. (3) Random sampling involves selecting sample sites with no particular pattern or reason. The choice of the sites is truly a random process. (4) Spot or grab samples are discrete portions of water samples taken at a given time. A series of grab samples, collected from different depths at a given site, reflect variation in constituents over a period of time. (5) Composite samples are essentially weighted series of grab samples, the volume of each being proportional to the rate of flow of water stream at the time and site of sample collection. Sample may be composite over a period of time e.g., 24 h. Such composite samples are useful in computing the material balance of a stream of water body over a period of time. Composite sampling makes it possible to sample a large area with fewer samples.

SAMPLE BLANKS Blank samples help to determine if there are extraneous sources of pollutants. For example, if the sample bottles are completely clean, some contamination might be introduced into the sample. The following are the different types of blanks: (1) Field blank is taken when there might be the possibility of air contamination of the sample during the sampling process. The field blank is a bottle of pure

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A Laboratory Manual for Environmental Chemistry

deionized water that is taken to the sample site and left open and exposed to the air for the same amount of time that the real samples are exposed to the air. The bottle is taken to the lab and analysed in the same manner as the real samples. (2) Trip blank is used to measure possible contamination from container or any preserving reagents added to the samples. The trip blank is a bottle of pure deionized water that is taken to the sample site but never opened, returned to the lab and analysed with other samples. (3) Equipment blank gives an indication of possible contamination from sampling equipment. It is taken by rinsing the equipment after it has been cleaned and analysing the rinsed solution for possible contamination.

GROUNDWATER The majority of groundwater samples are collected from wells drilled into the subsurface. While collecting samples from the wells, a process called well purging is carried out. Well purging is the withdrawal of a specified volume of groundwater prior to the collection of the sample. The aim of purging is to ensure that the water sample truly represents the properties and the conditions of the subsurface environment and not those of the well itself.

LAKE WATER In relatively stagnant bodies of water, such as lakes and lagoons, it is important to collect sample from depth, particularly where sharp temperature changes, or thermoclines, occur. If the temperature remains constant throughout the body, the sample should be collected from the depth of the body of water, between the surface and the bottom.

RIVER WATER Flowing bodies of water should be sampled across the entire width of the channel flow because water moves faster in the centre than at the edges. In cross-sectional sampling, flow velocities should be measured and recorded at each sampling location.

REAL TIME OR IN SITU ANALYSIS With the advent of portable field analytical kit, several water quality parameters can now be quantified on location and in real time. Parameters such as pH, salinity, dissolved oxygen and temperature can be measured directly by inserting the probe into water.

ANALYSIS IN A LABORATORY As a general rule, accurate and precise analysis for low-level pollutants in water can be done only in a laboratory environment using modern sample-preparation

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Sampling Techniques

9

methods and analytical equipment. The laboratory procedures minimize the risk of contamination during sample preparation and analysis. However, water analysis cannot be done quickly. Storage of samples in a controlled laboratory environment is necessary.

SAMPLE CONTAINERS Table 2

Major types of sample containers and methods of preservation

Analyte

Sample container

Preservation/Storage

Cations (Pb2+, Ni2+, Zn2+)

Plastic (HDPE) bottle

Add nitric acid (final pHE@EJO

!

PROCEDURE Light the candle and pour the turbid water sample slowly into the graduated tube placed over the candle. Watch the candle flame through the tube by keeping the eye directly above the tube. Stop the addition of the sample when the outline of the candle flame is no longer clearly visible. Read the turbidity value directly from the calibrated tube corresponding to the liquid level. A small amount of the highly turbid water will make the outline of the candle flame indiscernible. When the turbidity of water sample is very high, it needs to be first diluted accurately with distilled water and then the turbidity of the diluted solution is measured.

Note 1. Turbidity below the value of 25 units cannot be accurately measured with Jackson Candle Turbidimeter. It can be measured by nephelometric method. In the nephelometer, a light source illuminates the sample and one or more photoelectric detectors are used to indicate the intensity of scattered light at right angle to the path of the incident light. The values are expressed in nephelometric turbidity units (NTU). Knowledge of the turbidity variation in raw water supplies is of prime importance to the environmental engineer. Turbidity is used in conjuction with other information to determine whether a supply requires special treatment by chemical coagulation and filtration before it may be used for public water supply.

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5 1

Total Dissolved Solids

The total amount of dissolved chemical species in water is called total dissolved solids (TDS). It is a good general measurement of the concentration of ionic substances in water. In general, fresh water has less than 1500 mg L–1 of TDS and saline water has above 5000 mg L–1 of TDS.

PROCEDURE Place a clean 150 mL beaker in a drying oven at 103 – 105o C for one hour. Remove the beaker from the oven, place it in a desiccator until it cools and then weigh the beaker (w1 g). Filter 100 mL water sample through 0.45 mm filter paper into the beaker. Evaporate the sample on a water bath to reduce the volume of the sample to approximately 10 mL. Allow the beaker to cool and place it in the oven at 103 – 105o C for 1 – 11/2 hours. Keep the beaker in the desiccator and after cooling weigh it (w2 g). Repeat the process for concordant weight. TDS = [(w2 – w1)/100 ] ´ 106 ppm

Notes 1. If the water sample contains any solid which is volatile at about 110° C, then those would be lost on drying in an oven. However, only rarely such highly volatile solids are present in water. 2. The amount of volatile solids in water may be found out by heating the total solids obtained strongly in a muffle furnace (600° C) for 20 minutes and recording the weight loss. For this experiment, a silica or platinum crucible, instead of a china dish should be used. The loss in weight would correspond to volatile solids and the residue weight to non-volatile solids. 3. The measurement of salinity is a special case of measuring TDS in solutions with very high salt content, such as brackish water or sea water. Salinity is a dimensionless number that expresses the mass of dissolved salts in a given mass of solution. Mass of dissolved salt (mg ) Salinity = Mass of solution (g ) This ratio is also called parts per thousand or ppt. The symbol commonly used for salinity is ‰ (per thousand). Sea water has an average salinity of 35‰ (per thousand). 4. Groundwater has high TDS, low turbidity and less bacteria. Surface water has low TDS, high turbidity and more bacteria.

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6 1

Conductance

Pure water has a very high resistance to the flow of electricity because the only ions that are present are trace quantities of H3O+ and OH – arising from the autoionization of water. But natural water usually contains dissolved ionic salts and these ions provide a pathway for the flow of electrons through the solution. Therefore, it conducts electricity. The amount of current that flows is proportional to the concentration and the types of dissolved ions.

REAGENTS Standard potassium chloride solution (0.010 M): Dry KCl (AR) in an oven at 110°C for three hours and then cool it in a desiccator. Weigh exactly 0.7456 g of KCl (AR), dissolve it in conductivity water and make upto 100 mL. Pipette 10 mL of this solution into a 100 mL standard flask and make upto the mark with previously boiled and cooled conductivity water.

PROCEDURE Take about 50 mL of standard KCl solution in a dry container and immerse the container in a constant temperature water bath maintained at 25 °C. After 30 minutes, arrange a dip-type conductance cell inside the KCl solution and measure the resistance (R1). Remove the conductance cell, wash it with distilled water and rinse it with the sample water. Take the sample water in another beaker and arrange it in the constant temperature bath. Fix the conductance cell inside the sample water. After about 30 minutes, measure the resistance (R2).

CALCULATION Conductivity of the water sample = (in mS cm–1)

1411.8 ´ R1 R2

1411.8 = conductivity of 0.010 M KCl at 25° C R1 = resistance in ohms of the standard (0.010M) KCl R2 = resistance in ohms of the water sample

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Notes 1 The correct conductivity of the water sample is obtained only when it is measured soon after the collection of the sample, preferably in the field (site) itself. Laboratory fumes are likely to dissolve in the stored water sample and alter the conductivity. 2. The conductivity and the amount of dissolved salts are related in the following manner: Total concentration (mg L–1) = 0.64 ´ conductivity in mS cm–1. 3. Conductance can also be measured directly using conductivity meter. The conductivity cell is placed into the sample and the value is read from the display. If the temperature of the sample is not 25° C, apply correction using the equation c25 = ct (1 – 0.025Dt), where c25 is the conductance at 25° C, c t is the conductance measured at the sample temperature and Dt is the temperature difference. 4. The traditional unit used to measure electrical conductivity is mmho cm–1, which is equivalent to the SI unit deciSiemens per metre (dS m–1). Siemens is equal to ohm–1 (W–1) and is also known as mho. However, for smaller conductivities, microsiemens per centimetre (mS cm–1) is used. Very pure deionized water has a typical conductivity of 1mS cm–1, whereas rainwater can have conductivity in the range of 20 – 40 mS cm–1. 5. Water with a conductivity value greater than 2.25 dS m–1 is unsuitable for irrigation.

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7 1

pH Value

PRINCIPLE The pH of a sample water can be measured using a pH meter, which employs glass electrode. Different types of pH meter are commercially available— (a) battery-operated portable model, (b) line-operated laboratory model, (c) line-operated laboratory model with digital read-out, and (d) portable recording model. The pH meter should first be calibrated with a set of standard buffer solutions. Then the pH of the water sample can be determined. The glass electrode and the reference calomel electrode are used to determine the pH of a solution. Instead of these two separate electrodes, a combined electrode may be used. Ag/AgCl electrode may also be used as a reference. The pH meter works by using an electrical circuit that contains a glass electrode, an external reference electrode and a voltmeter that gives a measure of electrical potential in the circuit. The glass electrode consists of a current-carrying wire that dips into a solution of known pH. This solution is encapsulated in a thin glass membrane made of aluminosilicate glass (Fig. 2).

Internal reference solution (0.1M HCl) Thin glass membrane

Fig. 2

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Ag/AgCl electrode

The glass electrode used for measuring pH

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A Laboratory Manual for Environmental Chemistry

The molecular structure of glass is such that oxygen atoms comprise much of the surface of glass. These oxygen atoms have unshared pairs of electrons and hence carry a partial negative charge. When the glass electrode is dipped into an acidic sample, the protons attach themselves to the outside of the glass membrane. In order to maintain electrical neutrality, some of the protons that are stuck to the inner surface of glass membrane desorb from the surface. The consequence is a change in the pH of the inner solution which results in an electrical potential (Fig. 3).

O Reference solution inside the glass membrane

H+

–

O O

H+ –

OH OH+ Outer solution to be measured OH OH

O

Glass electrode placed in an acid solution (low pH)

+HO

Reference solution inside the glass membrane

–

O O

HO +HO

–

O

Outer solution to be measured

O

HO

Glass electrode placed in a basic solution (high pH)

Fig. 3 The surface of the glass electrode in acidic and basic solutions

An external reference electrode, Ag/AgCl is coupled to the glass electrode to complete the circuit and the voltmeter gives a measure of electrochemical potential. The line diagram of the cell is: Ag(s) | AgCl(s) | Cl – (aq) | H + (aq, outside) || H + (aq, inside), Cl– (aq) | AgCl(s)|Ag(s)

REAGENTS All the solutions should be prepared in previously boiled and cooled distilled water.

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pH Value

'

Standard buffer solutions Potassium tetroxalate: Dissolve 3.175 g KH2C4O8.2H2O in water and make upto 250 mL. Potassium hydrogen tartrate: Take 250 mL distilled water in a glass-stoppered bottle. Add excess KHC4H4O6 and shake the stoppered bottle vigorously for about 10 minutes to obtain a saturated solution. Filter and preserve the solution by adding 0.1 g thymol to it. Potassium hydrogen phthalate: Dissolve 2.552 g KHC8H4O4 in water and make upto 250 mL. Potassium dihydrogen phosphate: Dissolve 0.860 g KH2PO4 and 0.8875 g Na2HPO4 in water and make up the solution to 250 mL. Borax: Dissolve 0.9502 g Na2B4O7.10H2O in water and make upto 250 mL.

PROCEDURE Turn the pH meter on. Rinse the electrode with deionized water, shake it dry and place it in a beaker containing the first buffer solution at pH 7.00. Wait for the reading to stabilize, and then adjust the reading to pH 7.00. Rinse the electrode with deionized water, shake it dry and place it in a beaker containing the second buffer solution of known pH (Table 3). Allow the reading to stabilize and then adjust it to the appropriate pH for the buffer solution used. Check the calibration by measuring pH of the 7.00 standard. If the difference is more than 0.05 pH units, repeat standardization procedure. Take the sample water in a clean 100 mL beaker and immerse the electrode into it. Allow the reading to stabilize for about 2 minutes and read the display to obtain the pH of the sample. Table 3 pH of standard buffer solutions Buffer solution

0.05 M Saturated solution 0.05 M 0.025 M 0.01 M

Potassium tetroxalate Potassium hydrogen tartrate Potassium hydrogen phthalate Potassium dihydrogen phosphate Borax

Temperature (°C) 20

25

30

35

40

1.68 3.56

1.68 3.56

1.79 3.56

1.69 3.54

1.70 3.54

4.00

4.01

4.01

4.02

4.03

6.88

6.86

6.85

6.84

6.84

9.22

9.18

9.14

9.10

9.07

Notes 1. The calibration of the pH meter should be done with a buffer solution having pH closer to that of the sample water.

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2. The pH values are temperature-dependent. Therefore, the container having the sample may be maintained in a constant temperature water bath (25° C) when the pH is recorded. Similarly, the buffer also can be maintained in the water bath when the pH meter is calibrated. 3. As soon as the water sample is brought to the laboratory, its pH should be determined immediately. Delay in the measurement may involve change in the pH value of the sample because of absorption of laboratory fumes, loss of dissolved gases and also chemical changes. The accurate pH value of a water sample is obtained by determining this value at the time of sample collection on the spot itself using a battery-operated portable pH meter. 4. Immerse the electrodes in distilled water after use. A dry electrode requires several hours of soaking before it responds to H + correctly. 5. Do not use highly viscous liquids. 6. When [H]+ is very low and [Na+] is high, the electrode responds to Na+ as if Na+ were H+, and the apparent pH is lower than the true pH. This is called alkaline error or sodium error. 7. Solid-state pH sensors: There are pH sensors that do not depend on a fragile glass membrane. The field effect transistor is a tiny semiconductor device whose surface binds H+ from the medium in which the transistor is immersed. The higher the concentration of H + in the external medium, the more positively charged is the transistor’s surface. The surface charge regulates the flow of current through the transistor, which therefore behaves as a pH sensor.

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)?E@EJO

PRINCIPLE Natural water may contain both dissolved carbon dioxide and mineral acids. The mineral acids in water are measured by titration with sodium hydroxide to a pH of 4.5 using methyl orange indicator. This acidity is called methyl orange acidity. Both the carbon dioxide acidity and the mineral acid acidity can be determined by titration with sodium hydroxide using phenolphthalein indicator to a pH of 8.3. This total acidity is called phenolphthalein acidity. The presence of chlorine in water makes the methyl orange end point difficult to identify. This interference is avoided by the addition of sodium thiosulphate. The acidity is expressed in terms of CaCO3 mg mL–1 or mg L–1.

REAGENTS Standard sodium hydroxide solution (0.02 M): Prepare 0.2 M NaOH by dissolving 2 g NaOH in water and making up this solution to 250 mL. Pipette 25 mL of this solution and make it upto 250 mL. Standardize the diluted solution with 0.02 M oxalic acid. Standard potassium hydrogen phthalate solution (0.02 M): Dissolve 4.085 g-KHC8H4O 4 in water and make up the solution to 1000 mL. Sodium thiosulphate solution: Dissolve 2.5 g in water and make upto 100 mL. Methyl orange indicator Phenolphthalein indicator All the solutions are prepared in previously boiled and cooled distilled water.

PROCEDURE Methyl orange acidity: Pipette 100 mL water sample into a conical flask. Add a drop of thiosulphate solution. Add 2 drops of methyl orange indicator and titrate with standardized NaOH in the burette. The end point is the appearance of faint orange colour (characteristic of pH 4.5). Repeat the titration to obtain concordant titre values.

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CALCULATION 1000 mL of 1 M NaOH º

M CaCO3 2

= 50 g of CaCO3 Therefore, the methyl orange acidity of water sample (in mg L–1 of CaCO3) =

v2 × M × 50 ×1000 v1

v1 = volume of water sample in mL v2 = volume of NaOH consumed in mL M = molarity of NaOH solution The amount of dissolved carbon dioxide (mg L–1) Amount of CaCO3 (mg L–1 ) ´ 44 100 Phenolphthalein acidity (Total acidity): Pipette 100 mL water sample into a conical flask. Add a drop of thiosulphate solution. Add 2 drops of phenolphthalein indicator to it and titrate with standardized NaOH in the burette. The end point is the appearance of pale pink colour. (The colour is that of the pH 8.4 with NaHCO3 colour standard). Phenolphthalein acidity of water sample (in mg L–1 of CaCO3)

=

=

50 × M × v 2 ×1000 v1

v1 = volume of water sample in mL v2 = volume of NaOH consumed in the titration in mL M = molarity of NaOH solution

Notes 1. Stirring the analyte during the titration would result in the loss of dissolved carbon dioxide. Therefore, after a preliminary titration, a second titration may be carried out by adding a little less than the indicated amount of titrant to the water and then completing the titration quickly. 2. If the solution turns yellow on adding methyl orange, then methyl orange acidity is absent. 3. Where biological processes of treatment are used, the pH must be maintained within the range of 6.0 to 9.5. This criterion often requires adjustment of pH to favourable levels and calculation of the amount of chemicals needed is based upon acidity values.

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Alkalinity

PRINCIPLE In pure water, addition of even small quantities of a strong acid results in a drastic pH change. For example, if one drop of 0.1 N of HCl is added to 250 mL of pure water, the pH will change from 7.0 to 4.7. Most natural waters are protected from such drastic change by the presence of ionic compounds that react quickly with acid to buffer the system. This capacity of a solution to neutralize the acid is called acid neutralizing capacity (ANC) or total alkalinity. It is expressed in terms of CaCO3 (mg L–1). The alkalinity of the sample can be determined by titrating water with hydrochloric acid. If phenolphthalein is used as an indicator, the end point corresponds to complete neutralization of OH – ions and half neutralization of CO32–. This partial alkalinity is called phenolphthalein alkalinity. The reaction are: NaOH + HCl ® NaCl + H2O Na2CO3 + HCl ® NaHCO3 + NaCl During methyl orange end point, HCO3– is originally present and that obtained from CO32– is neutralized. NaHCO3 + HCl ® NaCl + H2O + CO2 (The total alkalinity corresponds to the amount of acid required to react with the hydroxide, carbonate and bicarbonate present in water.)

REAGENTS Distilled water: Boil and cool. Use this distilled water for preparing all the reagents. Standard hydrochloric acid solution (0.01 N): Standardize this with 0.01 NNa2CO3 solution. Standard sodium carbonate solution (0.01 N): Dissolve 0.53g Na2CO3 (AR) in water and make up the solution to 1000 mL. Sodium thiosulphate solution: Dissolve 2.5 g in water and make up the solution to 100 mL.

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Sodium bicarbonate: NaHCO3 (AR) Methyl orange indicator. Phenolphthalein indicator.

PROCEDURE Phenolphthalein alkalinity: Pipette 100 mL water sample into a conical flask. Add a drop of thiosulphate solution. Then add 2 drops of phenolphthalein indicator and titrate with the standardized hydrochloric acid in the burette. The end point is the disappearance of pale pink colour. Keep the contents of the conical flask stoppered for further titration to determine the total alkalinity.

CALCULATION Phenolphthalein alkalinity (in mg L–1 of CaCO3) =

50 ×N × V2 ×1000 V1

V1 = volume of water sample in mL V2 = volume of HCl used in mL N = normality of HCl Total alkalinity: To the solution obtained after determining the phenolphthalein alkalinity, add 2 drops of methyl orange indicator. Check the burette reading and continue the titration with the same hydrochloric acid in the burette. The end point is the colour change from yellow to red-orange (pH = 4.5).

CALCULATION Total alkalinity (in mg L–1 of CaCO3) =

50 ×N × V3 ×1000 V1

V1 = volume of water sample in mL V3 = volume of HCl used in mL (including the first titration)

Notes 1. The total alkalinity may also be determined using bromocresol green indicator. The end point in this case is the change of colour from blue to green at a pH of 4.5. 2. The total alkalinity can be determined potentiometrically also. 3. If pH of water sample is below 4.5, there is no need to determine alkalinity, since the sample will have no acid neutralizing capacity.

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4. The total alkalinity is also referred to as m value and the phenolphthalein alkalinity is called p value. A plot of p, m values and pH is used to determine the amount of calcium in drinking water (Tilliman’s curve). 5. The p and m values are used to determine the various types of alkalinities. i. If p = m, water sample contains only OH– alkalinity. ii. If p = m/2 or 2p = m, water sample contains only CO32– alkalinity. iii. If p = 0 ; m > 0, water sample contains only HCO3– alkalinity. iv. If p > m/2 or 2p > m, water sample contains OH– and CO 32– alkalinities. v. If p < m/2 or 2p < m, water sample contains CO32– and HCO3– alkalinities. 6. A water sample with high alkalinity is resistant to changes in pH. 7. There is a fundamental difference between the expression of acid-base properties in terms of pH and alkalinity. The pH can be considered to be an intensity factor, which measures the concentration of alkali or acid immediately available for reaction. In contrast, the alkalinity is a capacity factor, which is a measure of the ability of water sample to sustain reaction with added acid or base.

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Dissolved Oxygen

The amount of oxygen dissolved in water is called the dissolved oxygen (DO). It is expressed in milligrams per litre. The solubility of atmospheric oxygen in fresh water depends on the atmospheric temperature and pressure. When an aqueous solution is saturated with air at 1 atmospheric pressure and 25°C, the amount of oxygen present in the solution is 8.32 mg L–1. It may be in the range of 7–15 mg L–1. The solubility of oxygen is less in salt water (sea water) than in fresh water. In polluted water also, the solubility of oxygen decreases. DO is necessary for the life of fish and other aquatic organisms. When its concentration is less than 4 ppm, the water system is unsuitable for aquatic life. Oxygen is also needed to enable bacteria to oxidize organic matter present in water. Low concentration or absence of oxygen is an indicator of pollution in water. Its determination is therefore, essential.

PRINCIPLE The Winkler method is used to estimate DO. This method is based on the fact that the dissolved oxygen oxidizes Mn2+ to a higher oxidation state under acidic conditions; the oxidized manganese then liberates iodine from potassium iodide. The amount of iodine liberated is equivalent to the dissolved oxygen originally present. The iodine liberated is estimated by titration with sodium thiosulphate solution. 1 O2 ® MnO2 + H2O 2 MnO2 + 2I– + 4H + ® Mn2+ + I2 + 2H2O I2 + 2S2O32– ® 2I– + S4O62–

Mn2+ + 2OH– +

A sample of water collected in the field and brought to the laboratory is likely to undergo a change in its DO value, because of changes in temperature and also due to occurrence of biological reactions with time. Therefore, to obtain correct DO value, the water sample must be “fixed” immediately after collection in the field itself. This fixing is done by adding to the sample, solutions of KF, MnSO4, alkaline iodide-azide solution and concentrated sulphuric acid. After such a fixing, titration can be done at a convenient time later in the laboratory.

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REAGENTS Standard sodium thiosulphate solution (0.01 N): Dissolve 0.2482 g Na2S2O3.5H2O in water and make up the solution to 100 mL. Add to this 0.1 g Na2CO3 to preserve the solution. Manganese sulphate solution: Weigh about 48 g MnSO4. 4H2O and dissolve it in 100 mL water. Filter this solution if it is not clear. Potassium fluoride solution: Dissolve 40 g KF.2H2O in 100 mL water. Alkaline iodide-azide solution: Dissolve 50 g NaOH and 13.5 g NaI in distilled water and dilute to 1 litre. Dissolve 1 g NaN3 in 50 mL distilled water and add this solution with mixing to the alkaline iodide solution. Standard potassium dichromate solution (0.001 N): Dissolve 0.049 g K2Cr2O7 (AR) in distilled water and make up this solution to 1000 mL. Starch indicator: Prepare a fresh solution from starch powder. Sulphuric acid (AR)

PROCEDURE Standardization of thiosulphate: Pipette 20 mL standard K2Cr2O7 into a conical flash. Add 10 mL each of 4 N sulphuric acid and 5% sulphate KI. Titrate the liberated iodine with the thio in the burette. When the colour of the solution in the flask becomes straw yellow, add 2 mL starch solution and continue the titration. The end point is the disappearance of the blue colour and appearance of light green colour. Repeat the titration to obtain concordant titre values and calculate the strength of the solution. Pipette into a stoppered bottle 300 mL sample water. Add to this, quickly, 1 mL KF solution, 2 mL MnSO4 solution and 2 mL alkaline iodide-azide reagent. Stopper the bottle and shake the contents well without allowing air to enter. Then add 2 mL concentrated H2SO4, stopper and shake the bottle gently to dissolve the precipitate formed. Measure out 100 mL clear solution obtained above, into a conical flask and titrate the liberated iodine with the standardized thio in the burette. Add 2 mLstarch solution when the colour of the solution turns straw yellow. Continue the titration by adding the thio in drops till the blue colour disappears. Repeat the experiment to obtain concordant titre values.

CALCULATION 1000 mL of 1 N thiosulphate º 8 g oxygen \ Dissolved oxygen (mg L–1) =

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V2 × N ×1000 × 8 V1

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V1 = volume of water sample in mL V2 = titre value in mL N = normality of thiosulphate.

Notes 1. If the titre value is very low (low DO), a large volume of the sample water may be taken. 2. Fe3+, if present in water, will interfere with this estimation by liberating iodine from KI; this is masked by fluoride ions. 3. Similarly, nitrite ion in water can liberate iodine; this is prevented by the addition of NaN3 which destroys the nitrite ion: NaN3 + H + ® HN3 + Na+ HN3 + NO2– + H + ® N2 + N2O + H2O 4. The dissolved oxygen may be determined by polarography also. Oxygen in water is reduced, in two stages, at the mercury electrode:

2e– ® 2H2O2 O2 + 2H + ¾¾¾ 2e– ® 2H2O H2O2 + 2H + ¾¾¾ 10 mL sample is taken in a polarographic cell and 0.1 mL KCl (1M) and a small amount of mercury are added. Dropping mercury electrode with a head of 50 cm Hg is inserted and readings of current are measured at the potentials from –0.1 – to –1.5 V. The height of the wave i in the plot of current vs potential is proportional to the DO. 5. The polarographic method is not desirable for DO analysis in domestic / industrial wastewaters as Hg electrode gets poisoned by impurities. Hence, the membrane electrode method is used. Two metal electrodes, Ag and Pb are immersed in saturated KHCO3 solution separated from the test solution by polyethylene membrane (0.06 mm thick). This galvanic cell is plugged to a pH meter to give a direct reading of DO. The scale of 0–14 pH becomes 0–14 mg L–1 DO. 6. DO can be analysed using dissolved oxygen meter. DO meter uses oxygen-sensing membrane electrode. The electrode is shielded from the sample solution by a membrane that is selectively permeable only to oxygen. When oxygen diffuses through the membrane, transfer of an electron takes place between the electrode and the oxygen molecule. This causes a current to flow, which is measured. The amount of current is proportional to the amount of dissolved oxygen in the sample.

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Chemical Oxygen Demand

The chemical oxygen demand (COD) represents the amount of oxygen in milligrams required to oxidize all the organic pollutants present in water to carbon dioxide and water by a strong oxidizing agent. This oxidation can be carried out by acidified potassium dichromate and the reaction can be represented by the following general equation: CnHaOb + cCr2O72– + 8cH+ ® n CO2 +

a + 8c H2O + 2cCr3+ 2

2 a b n+ – 3 6 3 Correction for the presence of chloride: Chloride ion consumes dichromate by itself getting oxidized to chlorine. For each 1.0 mg L–1 of chloride ion, 0.23 mg of oxygen (COD) must be subtracted from the total COD measured for chloridecontaining water sample. Therefore, the chloride content of the sample must be determined with a separate aliquot of the water sample to make this correction in COD. The interference by chlorides can be avoided by the addition of mercuric sulphate to the sample. The mercuric ion combines with the chloride ion to form an inactive, weakly ionized HgCl2. No correction in the calculated COD value will be required if the chloride ions are “masked” using mercuric sulphate. Silver sulphate is used to catalyse the oxidation of straight chain aliphatic compounds, aromatic hydrocarbons and pyridine. where c =

REAGENTS Standard potassium dichromate solution (0.0125 N): Dissolve 0.613g K2Cr2O7 (AR) in distilled water and dilute to 1000 mL. Phenanthroline-ferrous sulphate ( ferroin) indicator: Dissolve 1.48g 1,10phenanthroline monohydrate and 0.7 g FeSO4.7H2O in 100 mL distilled water. Sulphuric acid (AR) Mercuric sulphate (AR) Silver sulphate (AR)

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Standard ferrous ammonium sulphate solution (0.0125 N): Dissolve 4.90 g FeSO4.(NH4)2(SO4).6H2O (AR) in a mixture of 200 mL distilled water and 15 mL concentrated H2SO4 taken in a beaker. Cool and make up this solution to 1 litre. Pipette 25 mL standard K2Cr2O7 into a conical flask. Add 100 mL distilled water and 25 mL concentrated H2SO4. Cool and add 0.5 mL phenanthrolineferrous sulphate (ferroin) indicator. Titrate this solution with the ferrous solution taken in the burette. The end point is the sharp change of blue to red colour. Repeat the titration to obtain concordant titre values and calculate the strength of ferrous ammonium sulphate solution.

PROCEDURE Pipette 50 mL water sample into a 500 mL conical flask. Pipette into this 25 mL standard dichromate solution. Add 50 mL concentrated sulphuric acid carefully with gentle mixing. Then add about 1 g Ag SO4 and 1 g HgSO4. Drop carefully 2 or 3 boiling chips into the flask and then attach a reflux condenser. Reflux the mixture for two hours on a heating mantle. Cool the flask and rinse the inside of the condenser with a small amount of distilled water allowing the washings to run into the flask. Then add 0.5 mL ferroin indicator and titrate the unreacted dichromate with standardized ferrous ammonium sulphate in the burette. The end point is the sharp change of colour from blue to reddish-brown. Simultaneously, conduct a duplicate with the same quantity of sample water. Perform a blank using 50 mL distilled water (instead of sample water), titrating in the same way and conducting all the procedures including refluxing.

CALCULATION COD (mg L–1) =

(V2 – V3 ) × N × 8 ×1000 V1

V1 = volume of water sample taken in mL V2 = blank titre value in mL V3 = titre value with the water sample in mL N = normality of the ferrous ammonium sulphate solution In this calculation, the oxidant is not directly involved because the amount of oxidant used is related to the amount of reducing agent involved in the titration.

Notes 1. Recommended as the standard method until 1965, the permanganate test has been replaced by the dichromate test described above. This test utilizes potassium permanganate (KMnO4) instead of dichromate as the oxidizing agent. The wastewater sample is boiled with a measured excess of permanganate in acid solution (H2SO4) for 30 min. The pink solution is cooled and a known excess of ammonium oxalate is added, the solution

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becoming colorless. Excess oxalate is then titrated with KMnO4 solution until the pink colour is restored. Oxalate used is calculated by difference and permanganate utilized is calculated from simple stoichiometry. 2. Instrumental COD methods are very fast and yield reproducible results. The Aquarator is designed to measure oxygen demand in the range of 10–300 mg L–1. Samples of higher concentration are handled by preliminary dilution of the sample. A 20-mL sample, homogenized if necessary, is injected by a syringe into the Precision Aquarator. The sample is swept through a platinum catalytic combustion furnace by a steam of dry CO2, which oxidizes the contaminants to CO and H2O. Water is stripped out in a drying tube and reaction products are then passed through a second platinum catalytic treatment. The CO concentration is measured by an integral non-dispersive infrared analyser senstized for carbon monoxide. The resultant reading is directly converted to COD by the use of a calibration chart.

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12 1Biochemical Oxygen Demand Biochemical oxygen demand (BOD) is defined as the amount of oxygen required by bacteria for stabilizing decomposable organic matter in water under aerobic conditions. BOD is calculated by measuring the oxygen consumed (decrease in dissolved oxygen) by bacterial and chemical action in a closed sample of water maintained at 20°C for five days. BOD is expressed in mg L–1.

a b 3 ö æ æa 3 ö CnHaObNb + ç n + – – c÷ O2 ¾® nCO2 + ç – c÷ H2O + cNH3 è è2 2 ø 4 2 4 ø –1 BOD values may range from 30 to 100 mg L depending on the quality of water. Clean water samples have low BOD values whereas polluted water samples have high BOD values.

PRINCIPLE An aliquot of the water sample is maintained in an incubator at 20°C for five days in a closed BOD bottle, without allowing air to enter. During this time, the water sample is assumed to be stabilized, i.e., the bacterial decomposition gets completed. Measuring the dissolved oxygen in the water sample before and after incubation would indicate the amount of oxygen used for stabilizing the water. A water sample with a low BOD can be straightaway used for the BOD determination. But a water sample with a high BOD must be diluted accurately and pretreated before the determination of BOD.

REAGENTS Phosphate buffer solution: Dissolve 8.5 g KH2PO4, 21.75 g K2HPO4, 33.4 Na2HPO4.7H2O and 1.7 g NH4Cl in 500 mL distilled water. Dilute this solution to one litre. Ferric chloride solution: Dissolve 0.25 g FeCl3-6H2O in water and dilute the solution to one litre. Calcium chloride solution: Dissolve 27.5 g CaCl2 in one litre water. Magnesium sulphate solution: Dissolve 22.5 g MgSO4.7H2O in one litre water.

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Water for dilution: Take 2 litres distilled water in a 3-litre bottle. Shake this partially filled bottle for about 10 minutes so that the distilled water gets saturated with atmospheric oxygen. Then add to this 2 mL phosphate buffer solution, 2 mL magnesium sulphate solution, 2 mL calcium chloride solution and 2 mL ferric chloride solution.

PROCEDURE BOD determination without dilution: If the water sample is fairly clean, it will have a low BOD value and can be used as such. Fill four BOD bottles (200 mL capacity) upto the brim and tightly stopper these. Make sure that there are no air bubbles inside the bottles. Use the water in two bottles to determine the dissolved oxygen immediately. Seed the other two bottles with 1 mL domestic wastewater and place in an incubator or a constant temperature water bath at 20° C. After five days, determine the dissolved oxygen in these two bottles. BOD (mg L–1) = DO(b) – DO(a) DO(b) = DO before incubation; DO(a) = DO after incubation. BOD determination after dilution: This method is used when the BOD value is high. Pipette 100 mL water sample into each of the two BOD bottles. Fill the bottles with the dilution water prepared and stopper the bottles tightly. Determine the DO of the diluted water using one of the bottles immediately. Incubate the other bottle for 5 days at 20° C and then determine the DO of the incubated water.

CALCULATION BOD (mg L–1) = DO(b) – DO(a) (without dilution) BOD (mg L–1) = [DO(b) – DO(a)] ´ f (with dilution) f = dilution factor

Notes 1. Commercially available BOD bottles are of known volumes, provided with cup seal and vinyl tubing fixed at the mouth of the bottle. 2. If the BOD of the water sample is likely to be greater than the DO level, the sample should be accurately diluted before estimation. 3. Sometimes, the sample water may be very low in bacterial content so that the water may not get stabilized in 5 days. Then the sample has to be seeded with a small, known volume of domestic wastewater. When this is done, a seed correction in the measured BOD value should be made. The value for correction is determined by measuring the BOD value of the wastewater used for seeding. The wastewater may be diluted and then its BOD value is determined.

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4. COD/BOD ratio indicates the ease with which a wastewater sample can be biologically degraded. If the ratio COD/BOD is less than 1.7, the wastewater can be easily and completely biodegraded. If the ratio is between 1.7 and 10, the wastewater can be degraded but not completely. If the value is closer to 1.7, the sample can be biodegraded slowly. If the ratio is more than10, the wastewater cannot be biodegraded. Any attempt to biodegrade such wastewaters would be futile. Thus, the ratio COD/BOD helps in assessing the biodegradability of wastewater.

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Dissolved Organic Carbon

PRINCIPLE Dissolved Organic Carbon (DOC) is defined as the concentration of carbon of all the dissolved organic compounds. It is analysed using a DOC analyser. The organic compounds are oxidized to CO2 and H2O and the amount of CO2 produced is determined by IR detector. y CxHyOz + nO2 ¾¾® xCO2 + H2O 2

PROCEDURE The wastewater sample has to be pretreated to remove dissolved inorganic carbon and suspended particle matter. Suspended particles are removed by filtration or centrifugation. The presence of bacteria may also contribute to the DOC value. A bacterium is represented as C5H7O2N based on the chemical analysis of cells. Using a cellulose acetate filter or membrane with a pore diameter of 0.45 mm, the bacteria may be removed from the sample. Inorganic carbon is removed by acidifying the sample with 15% HCl solution to maintain a pH of 3 – 4 and then passing pure oxygen or nitrogen through the sample. The sample is then oxidized in the furnace of the DOC analyser at 980°C or it is oxidized using potassium persulphate as the oxidizing agent and promoting the oxidation by the free radical OH produced from the exposure to UV light. Thus, all the dissolved organic carbon is converted into CO2, which is detected using infrared analyser. Alternatively, potassium persulphate and phosphoric acid may be added to the sample. Phosphoric acid is sparged with air or nitrogen to expel carbon dioxide formed from bicarbonate and carbonate in solution. After sparging, the sample is pumped into a chamber containing a UV lamp emitting radiation at 184 nm. The OH free radicals formed by the UV radiation are the active species, which bring about rapid oxidation of dissolved organic compounds. At the end of the oxidation, carbon dioxide is sparged from the system and measured with a gas chromatographic detector or IR detector.

Notes 1. TOC—Total organic carbon is estimated using the same instrument without filtering wastewater. 2. TOC : DOC + Suspended carbon.

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14 1 Calcium Hardness of Water Historically, ‘hardness’ is defined in terms of capacity of a water sample to precipitate soap. Soap is precipitated by most cations with multiple charges. In natural waters, the concentration of calcium and magnesium far exceeds that of any other metal ion. Thus, hardness has come to mean the total concentration of calcium and magnesium expressed in terms of calcium carbonate equivalent. The hardness of water due to dissolved calcium is referred to as calcium hardness.

PRINCIPLE The water sample is mixed with a buffer of pH 12 (1M NaOH) so that Mg2+ present in the water is precipitated as magnesium hydroxide. Then the calcium in water is titrated with EDTA using murexide indicator.

REAGENTS EDTA solution: Dissolve 3.72 g Na2EDTA.2H2O in water and make this upto one litre to give a 0.01 M solution. Standardize this solution with a standard Zn2+ (0.01M) solution. Pipette 20 mL standard Zn2+ solution into a conical flask, add 5 mL buffer and 3 – 4 drops of Eriochrome black T indicator. The end point is the colour change from wine red to blue. Sodium hydroxide solution: Dissolve 40 g pure NaOH in water and make upto one litre (1 M solution). Murexide: Mix 0.1 g murexide with 9 g NaCl(AR). Grind this mixture in a mortar with a clean pestle.

PROCEDURE Pipette 20 – 100 mL water sample into a clean conical flask. Add 10 mL sodium hydroxide solution (1M) and dilute the solution to 100 mL. Then add 1g murexide-sodium chloride solid and titrate the solution quickly with the EDTA taken in the burette. The end point is the change from pink to purple colour. The titration should be completed quickly to avoid the possible formation of calcium carbonate precipitate. Repeat the titration to obtain concordant titre values.

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CALCULATION Hardness is expressed in mg L–1 of as CaCO3. 1 mL of 0.01 M EDTA º 1.0 mg CaCO3 The calcium hardness =

V2 × M ×105 V1

V2 = titre value in mL M = molarity of EDTA solution V1 = volume of water sample in mL

Notes 1. If there is interference by Fe, Zn, Al, etc., the addition of masking agents, such as potassium cyanide and hydroxylamine hydrochloride is necessary. These two are added before the addition of sodium hydroxide. 2. Instead of murexide, hydroxynaphthol blue indicator can be used. The indicator solid is added to the sample and titrated to a blue-violet end point. 3. Alternatively, 5 – 10 drops of calconcarboxylic acid indicator (0.4% in methanol) can be used. The end point is the colour change from wine red to pale blue.

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Total Hardness of Water

PRINCIPLE The total hardness of water is measured as the sum of calcium and magnesium ions. This can be determined by titrating the water sample with the disodium salt of EDTA using a visual indicator. The water sample is buffered with NH4OH-NH4Cl to give a pH of 10. Eriochrome black T is used as an indicator. In this titration, actually, both Ca2+ and Mg2+ are titrated with the EDTA. If the water sample has a low magnesium content relative to its calcium content, a solution of Mg2+-EDTA is added before titration. The indicator dye forms a weak complex with calcium and magnesium. As EDTA is added, it begins to form complexes with calcium and magnesium ions in water. At the end point, the EDTA begins to take calcium and magnesium away from the dye complex. Hence, the colour change takes place. Na2(EDTA) + Ca2+ ® [Ca(EDTA)]2– + 2Na+ + 2H+

REAGENTS Standard Zn2+ solution (0.01M): Weigh 0.7185 g ZnSO4.7H2O (AR), dissolve it in a minimum amount of dilute H2SO4 and make up the solution to 100 mL (0.01M). EDTA solution (0.01M): Dissolve 3.72 g Na2EDTA.2H2O in one litre distilled water. Eriochrome black T: Dissolve 0.25 g solid in 100 mL methanol Buffer: Mix 70 g NH4Cl (AR) with 570 mL liquor ammonia and dilute the resulting solution to 950 mL. Mg2+-EDTA solution: Dissolve 3.120 g MgSO4.7H2O and 4.716 g disodium ethylenediaminetetraacetate dihydrate in 50 mL distilled water. This solution is then added to a buffer solution prepared by dissolving 70 g NH4Cl in 570 mL liquor ammonia diluted to 950 mL with distilled water. This buffer is used only if the magnesium content of the water sample is low. Potassium cyanide. Hydroxylamine hydrochloride.

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39

PROCEDURE Pipette a suitable volume of the water sample (20 – 100 mL, depending on the hardness) into a conical flask. (If the water is very hard, a 10 mL sample is sufficient for titration and this can be found out by a pilot titration). Add 5 mL buffer solution, 0.5 g potassium cyanide and about 1.0 g hydroxylamine hydrochloride. To this add 3 – 4 drops of Eriochrome black T indicator. The solution becomes wine red. Titrate the solution with EDTA taken in the burette. The end point is the disappearance of the last shade of red and formation of steel blue colour. Repeat the titration to obtain concordant titre values. Standaradize the EDTA solution with a standard Zn2+ (0.01M) solution (Refer: Experiment 14)

CALCULATION 1 mL of 0.01 M EDTA º 1.0 mg CaCO3 Total hardness =

V × M ×105 V2 × M ×1000 = 2 V1 0.01× V

V2 = titre value in mL M = molarity of EDTA solution V1 = volume of water sample in mL

Temporary and permanent hardness The ions that contribute to water hardness, namely Ca2+ and Mg2+ have their counter ions as carbonate, bicarbonate, hydroxide, chloride and sulphate. Based on the nature of the counter ions, hardness is classified into two types: 1. Temporary or carbonate hardness 2. Permanent or non-carbonate hardness

Temporary or carbonate hardness Temporary hardness is due to the presence of bicarbonates of calcium and magnesium. Temporary hardness is removed by boiling the water. Boiling results in the precipitation of CaCO3 or MgCO3, which can be filtered off.

Permanent or non-carbonate hardness Permanent hardness is caused by the presence of chlorides and sulphates of calcium and magnesium. Permanent hardness cannot be removed by boiling. Total hardness = Temporary hardness + Permanent hardness

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PROCEDURE Permanent hardness Take about 100 mL water sample in a 250 ml beaker. Boil it for about 30 minutes. Cool and filter it. Pipette 50 mL of this water into a clean conical flask. Add 5 mL buffer solution, 0.5 g potassium cyanide and about 1.0 g hydroxylamine hydrochloride. Then add 3 – 4 drops of Eriochrome black T indicator. Titrate the wine red solution against EDTA till the colour changes to steel blue.

CALCULATION 1 mL of 0.01M EDTA º 1.0 mg CaCO3 \

Permanent hardness =

V3 × M ×1000 V × M ×105 = 3 0.01× 50 50

V3 = titre value in mL M = molarity of the EDTA solution Calculate the temporary hardness of the water by subtracting the permanent hardness from the total hardness.

Notes 1. Metal ions, Fe(III), Zn(II), Cd(II), Al(III), etc., present in water interfere with the water hardness titrations. These are masked or inhibited by the addition of potassium cyanide and hydroxylamine hydrochloride, (NH2OH HCl). Mn(II), if present in water, can be masked by 1, 10-phenanthroline. 2. Phosphates do not interfere in the total hardness determination. However, they interfere in the calcium determination. The other common anions do not interfere in the determination of total hardness or temporary hardness. 3. An alternative method for determination of the concentration of calcium and magnesium in water is atomic absorption spectrophotometry. 4. The counter ions of calcium and magnesium contribute to the alkalinity of water. Thus, hardness and alkalinity are related water quality parameters.

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16 1

Nitrite Nitrogen

PRINCIPLE In this method, two organic reagents, sulfanilamide and N-(1-naphthyl)-ethylenediamine dihydrochloride are used to estimate the nitrite. Under acidic conditions, nitrite ion as nitrous acid reacts with the amino group of sulfanilamide to form a diazonium salt. This salt then combines with N-(1-naphthyl)-ethylenediamine dihydrochloride and forms a bright pinkish-red azo dye. SO2NH2

SO2NH2 –

+ NO2

2H+

N+

NH2

N

SO2NH2 + N

+ 2H2O

NHCH2CH2NH2 ◊ 2HClÆ

+

N NH2SO2

N

N

NHCH2CH2NH2 ◊ 2HCl

Red purple dye (lmax = 543 nm)

The intensity of the colour produced is directly proportional to the amount of nitrite nitrogen present in the water sample. Therefore, the nitrite nitrogen can be estimated by comparison of the colour intensity developed by the method of visual colour standard series method or by photometric measurement. The photometric measurement is preferred over the visual method because colour standards for visual comparison are not stable for a long time.

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REAGENTS Standard KMnO4 solution (0.025N): Dissolve 0.79 g solid in one litre water. Standard nitrite solution: Prepare a solution of nitrite by dissolving 1.232 g NaNO2 (AR) in deionized water and dilute to 1000 mL. Pipette 25 mL standard KMnO4, 1 mL H2SO4 (AR) and 25 mL NaNO2 solution into a stoppered flask. Gently shake the stoppered flask and after fifteen minutes add 1g KI. Shake the contents of the flask gently and titrate the liberated iodine with standard sodium thiosulphate (0.01N) solution using starch indicator. The end point is the first disappearance of the blue colour. Repeat the titration to obtain concordant titre values and calculate the mass of nitrite nitrogen per litre of the standard solution. (V × N1 ) – (V2 × N 2 ) × 7 ×1000 Nitrite nitrogen (mg L–1) = 1 V3 V1 = volume of standard KMnO4 solution used in mL N1 = normality of KMnO4 solution V2 = volume of standard sodium thiosulphate consumed in mL N2 = normality of sodium thiosulphate solution V3 = volume of NaNO2 taken for titration in mL

Colour-forming reagent Mix 1.0 g sulfanilamide, 0.1 g N-(1-naphthyl) - ethylenediamine dihydrochloride and 10 mL of 85 wt% phosphoric acid and dilute to 100 mL. Store in a dark bottle.

PROCEDURE Dilute the stock nitrite solution accurately 500 times in two or three steps, so that the final diluted solution will contain 0.005 mg L–1 nitrogen. Pipette into a series of Nessler tubes the following volumes of diluted standard nitrite solution: 1,2,3,4,5,6,8,10 mL. Add 30 mL distilled water and 1 mL colourforming reagent to each Nessler tube. Adjust the volume in each Nessler tube accurately to 50 mL and stir the solution. Pipette a suitable volume of the water sample to be analysed into another Nessler tube and develop the colour by adding exactly the same amount of the reagent as with the standard solutions. Compare the colour intensity of the test solution with the colour intensities of the standard solutions and fix the nitrite nitrogen concentration range for the test solution. Carry out another set of determination by using a series of standard solutions within the concentration range estimated initially to find out the exact concentration of the test solution. The matching of colours should be completed within fifteen minutes after the addition of reagent. Alternatively, the concentration can be determined by measuring the absorbance of solution at 543 nm using a spectrophotometer. A calibration curve is needed to get the concentration.

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Nitrite Nitrogen

43

The estimation is repeated with different volumes of water sample for the confirmation of result. When the water sample has a pH greater than or less than seven, it should be neutralized with 1N HCl or NaOH respectively, before adding the reagent into the Nessler tube.

Note 1. Nitrite can also be allowed to react with sulfanilic acid and HCl to form a diazonium salt. This is turn will react with 1-naphthylamine to form a pink dye (l max = 520 nm). (Refer: Experiment 54)

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Nitrate Nitrogen Brucine Method

PRINCIPLE Brucine is a naturally occurring complex organic compound (heptacyclic alkaloid). It reacts with nitrates under acidic conditons at an elevated temperature to produce a yellow colour. Such solutions obey the Beer’s law only at low nitrate nitrogen concentration of 0.1 to 1 mg L–1. The intensity of the colour developed is a function of both time and temperature, therefore, these two factors must be carefully fixed during estimation to obtain correct results. The presence of chloride in water does not interfere in this method.

REAGENTS Brucine reagent: Take 50 mL water and 3 mL conc. HCl in a beaker. Heat this just to boiling and add 1 g brucine sulphate and 0.1 g sulfanilic acid with stirring. Cool and dilute this solution to 100 mL. Sulphuric acid solution: Mix carefully 500 mL conc. H2SO4 with 100 mL distilled water. Standard nitrate solution: Dissolve 0.07218 g KNO3 in water and make upto 100 mL in a standard flask. Dilute this 10 times so that the diluted solution contains 0.01 mg mL–1 N.

PROCEDURE Prepare a series of 50 mL standard solutions of the nitrate from the stock standard nitrate solution to obtain a range of concentration, 0.01–1.00 mg L–1 N. Pipette 2 mL aliquots of the standard solutions into different dry 100 mL breakers provided with glass rods. To each beaker add 1 mL brucine-sulfanilic acid reagent and 10 mL sulphuric acid. Stir gently for about five minutes. Cover the beakers with watch glasses and keep them in the dark for 30 minutes. The absorbance of each

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Nitrate Nitrogen Brucine Method

45

solution is then measured at 410 nm. A graph is drawn by plotting the absorbance values against the nitrate nitrogen concentrations of the standard solutions. Treat a suitable volume (2-5mL) of the water sample with the same amount of reagents in an identical manner and read the absorbance of this solution. Then, using the absorbance-concentration plot, the nitrate nitrogen concentration of water sample can be found out. Repeat this experiment with a different volume of the water sample.

CALCULATION Concentration of nitrate nitrogen = concentration from graph ´ 14 mg L–1 (ppm)

Notes 1. The presence of Iron (II) or Iron (III) in concentration greater than 1 mg L–1 interferes with this estimation. Strong oxidants and reductants also interfere in this method. Pretreatment of the water sample with a dilute solution of sodium arsenite can be done to destroy the oxidants present. Chlorinated water has to be dechlorinated before the nitrate estimation. This is done by adding 2 mL NaAsO2 solution (0.025 N) for a residual chlorine content of 1 mg L–1 . 2. If the nitrate content of the water sample is very high, the yellow colour developed would be very dark. In such case, the water sample can be diluted with distilled water to bring the nitrate nitrogen concentration in the 0.1-1.0 mgL–1 range. 3. The nitrate can be reduced quantitatively to nitrite either by passing the solution through a column containing cadmium filings coated with copper or using zinc/sodium chloride mixture and then estimating as nitrite. (Refer: Experiments 16 and 54). 4. The nitrate can also be estimated using phenol-disulphonic acid. (Refer: Experiment 53). 5. UV spectrophotometric method: Measurement of absorbance at 220 nm enables rapid determination of nitrate. Since dissolved organic matter may also absorb at 220 nm and nitrate does not absorb at 275 nm, a second measurement is made at 275 nm. Acidification with 1N HCl prevents interference from hydroxide and carbonate. To 50 mL water sample add 1 mL of 1N HCl. Read the absorbance against distilled water set to zero absorbance. Take the absorbance readings at 220 nm (A) and 275 nm (B). Absorbance for sample = A - 2B.

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18 1

Ammonia Nitrogen

The nitrogen which occurs as ammonium ion or as NH3 is considered to be ammonia nitrogen. Thus, ammonia nitrogen in solution exists in the equilibrium NH4+ S NH3 + H+

PRINCIPLE The ammonia nitrogen present in a water sample can be estimated by direct Nesslerization. A known volume of the water sample is treated with Nessler’s reagent to produce a yellowish-brown colloidal dispersion. The intensity of this colour produced is directly proportional to the amount of NH3 originally present. The Nessler’s reagent is an alkaline solution of potassium mercuric iodide. It is represented as K2HgI4 or 2KI.HgI2. The reaction between NH3 and Nessler’s reagent may be represented by the following equation: 2K2HgI4 + NH3 + 3KOH ® I-Hg-O-Hg-NH2 + 7KI + 2H2O The colour developed with the water sample by Nessler’s reagent is then visually matched with the colours developed similarly with a series of standard ammonia solutions. If the water sample is turbid or coloured, it should be pretreated before Nesslerization. This is done by the addition of ZnSO4 and NaOH solutions.

REAGENTS Zinc sulphate solution: Dissolve 10 g ZnSO4.7H2O in 100 mL water. Sodium hydroxide solution: Dissolve 60 g NaOH in 100 mL water. Nessler’s reagent: Dissolve 10 g HgI2 and 7 g KI in about 25 mL water in a beaker. In another beaker dissolve 16 g NaOH in 50 mL water. Mix the solutions in the two beakers and dilute to 100 mL. If the solution thus obtained is not clear, quickly filter it using a Whatman Number 40 filter paper. Standard ammonia solutions: Prepare 100 mL standard solution of NH4Cl by dissolving 1 g NH4Cl (AR) in water. From this stock solution prepare several diluted ammonia solutions, such that the aliquots to be used in Nessler tubes would have the range of ammonia nitrogen concentration between 0.01and 0.06 mg L–1 (ppm). Chapter_18.p65

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Ammonia Nitrogen

47

PROCEDURE Pretreatment of coloured or turbid water sample: Take 500 mL water sample and add 5 mL ZnSO4 solution and 1.5 mL NaOH solution to it. Allow the floc to settle, which takes about fifteen minutes. Filter the water through a Whatman Number 40 filter paper. Collect the filtrate after discarding the first 100 mL. Add 3-4 drops of 1 M EDTA solution to it to prevent calcium hydroxide or magnesium hydroxide from getting precipitated and producting turbidity on adding Nessler’s reagent. Pipette 5 mL clear water sample into a Nessler tube. Make up the solution to 50 mL. Add 1 mL Nessler’s reagent and mix the solution thoroughly with a glass rod. Compare visually the colour intensity produced by the water sample with the colour intensities developed similarly with a series of standard ammonia solutions. After fixing the ammonia concentration range of the sample, conduct another set of experiments to fix the actual ammonia concentration of the sample. Photometric measurement: The concentration of the coloured product may be found photometrically in a 1 cm cell at the wavelength range 420 nm. Reliable photometric measurements of ammonia concentration are obtained upto 0.5 mg L–1 N. A photometric calibration curve may be constructed by measuring the absorbance of a series of standard solutions in this concentration range.

CALCULATION The NH4+ content of the water sample can be calculated by using the weight of NH4Cl employed for preparing the standard solution. Concentration of ammonia nitrogen = concentration from the graph x

14 18

mg L–1 (ppm)

Notes 1. Nessler’s reagent should be freshly prepared; old samples should not be used. 2. The same sample of Nessler’s reagent must be used for both the water sample and the standard ammonia solutions. 3. When the water is highly turbid or coloured, it should be distilled to separate ammonia from interfering substances. All the ammonia present in the turbid or coloured water will get distilled with the steam and absorbed in the condensate. The distillation is done after adding a phosphate buffer to the water sample to maintain the pH at 8.2. In the absence of the buffer during distillation, the pH decreases due to the release of H+, suppressing the formation of NH3: NH4+ S NH3 + H+

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4.

5. 6.

7.

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This equilibrium is displaced far to the right when the pH is above 7. Ammonia distilled may be absorbed in excess boric acid and back titrated with standard acid. The orange-brown reaction product tends to precipitate at higher concentrations. Then, a protective colloid (gelatin or gum arabic) may be added to stabilize the suspension formed. Calcium and magnesium interfere with this estimation; addition of tartrate solution masks these metals. Ammonia is also determined spectrophotometrically by reaction with phenol in the presence of hypochlorite. The absorbance of the blue product is measured at 625 nm. Ammonia can be determined using an ammonia-selective compound electrode. To the water sample, 1.0 mL NaOH is added to keep all ammonia in the NH3 form and the voltage is recorded to the nearest millivolt.

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19 1

Fluoride

PRINCIPLE Zirconium ion reacts with alizarin dye to form a coloured product (reddish “lake”). When fluoride is added to this reddish lake, the fluoride ion combines with zirconium ion to form a stable complex, ZrF62–. This reaction between the zirconium ion and fluoride ion decreases (bleaches) the colour intensity of the lake. Zr4+ + Alizarin ® Zr - Alizarin lake (reddish) Zr - Alizarin lake + 6F – ® Alizarin + ZrF62– (reddish) (yellow) Therefore, the bleaching action is proportional to the fluoride ion concentration. In this case, Beer’s law is valid in an inverse manner i.e., the absorbance of the bleached solution is inversely proportional to [F–]. Then, by visual colorimetry or using a photometer, the amount of fluoride in a water sample can be determined. The bleaching action of fluoride is slow and therefore comparisons are made one hour after mixing the lake and the fluoride. In addition, the colour intensity is temperature dependent. Therefore, in this determination, time and temperature are carefully maintained constant in constructing the absorbance-concentration calibration curve. Chlorine, if present in the water, interferes with the estimation. It must be eliminated by the addition of sodium thiosulphate.

REAGENTS Sodium thiosulphate solution: Dissolve 1 g of Na2S2O3.2H2O in water and dilute to 250 mL. Zirconium-alizarin reagent: Add 100 mL conc. HCl to 300 mL distilled water in a beaker. In another beaker, take 400 mL distilled water and add 30 mL conc. H2SO4 carefully to it. Cool both the diluted acids and mix them. Add 0.3 g ZrOCl2.8H2O to 50 mL distilled water. Dissolve separately 0.07 g Alizarin Red S in 50 mL distilled water. Add this solution with stirring to the

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zirconium solution. After 5 minutes, add the HCl-H2SO4 solution to the zirconium-alizarin solution and dilute the resulting solution to 1000 mL with distilled water. Mix this reagent thoroughly and use it after an hour. Store this reagent in a brown bottle. Standard fluoride solution: Dry NaF (AR) at 110° C for 2 hours and cool it in a desiccator. Dissolve 0.5525 g NaF in water and make up this solution to 250 mL. Pipette 25 mL stock fluoride solution and dilute it accurately to 250 mL.

PROCEDURE Prepare a series of standard NaF solutions from the diluted stock solution. Pipette a fixed aliquot from each of these standard solutions into five Nessler tubes. Add 1 drop of thio solution to each Nessler tube and then add 5 mL zirconium-alizarin reagent. Dilute with distilled water upto the mark and stir. In another Nessler tube, pipette a suitable volume of water sample and add a drop of thio and 5 mL zirconium-alizarin reagent. Dilute and stir the solution. After an hour, match the sample with the standards and determine the concentration range of the fluoride ion in the sample. Perform another set of experiments to fix the concentration of fluoride in the sample water.

Notes 1. The colorimetric method of determining fluoride is affected by the original colour and turbidity of water. To avoid such interference, the water sample should be distilled in borosilicate glass distillation apparatus before fluoride determination. Under acidic condition, fluoride is distilled out. 2. For spectrophotometric estimation, 4,5-dihydroxy-3-(p-sulpha-phenylazo)2,7-naphthalene-disulphonic acid trisodium salt (SPADNS dye) is used instead of alizarin for developing the “lake” with Zr. In this method, there is no waiting period for colour development, the reaction between F– and Zr-dye (bleaching) being instantaneous. SPADNS solution is prepared by dissolving 0.9 g of it in water and diluting it to 500 mL. Zirconyl acid reagent is prepared by dissolving 0.133 g ZrOCl2 . 8H2O in 25 mL water and adding 350 mL conc. HCl. The solution is diluted to 500 mL. Zirconyl acid-SPADNS reagent is prepared by mixing equal volumes of these two solutions. Fluoride bleaches this reagent and the absorbance of the bleached solution is read at 570 nm. 3. An optimum fluoride level should be maintained in drinking water, as it affects the teeth both at low and high concentrations.

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20 1

Fluoride in Drinking Water (By Fluoride Electrode)

The drinking water should contain approximately 1 mg L–1 (ppm) of fluoride for people to maintain good dental and skeletal health. When the fluoride level is low for a long time in drinking water, dental caries (tooth decay) occurs. When the fluoride level is high, dental fluorosis (mottled or discoloured enamel) results. Therefore, maintaining an optimal level of fluoride ion in drinking water is important. High fluoride leads to skeletal fluorosis.

Principle Ion-selective electrode: An ion-selective electrode is sensitive to a particular ion. The fluoride ion-selective electrode (Fig. 4) consists of a single crystal of europium (II)-doped lanthanum fluoride, LaF3. The response mechanism in this electrode is similar to the ion exchange type mechanism for the glass electrode. Fluoride ions, from the solution to be determined, migrate into the crystal (in the form of membrane). Conduction in the crystal phase occurs by a lattice defect mechanism. Mobile fluoride ions move into vacant crystal position. The size, shape and charge of the mobile ion are important in this conduction mechanism. With the LaF3 - Eu (II) defect crystal, the Ag/AgCl electrode fluoride ions are best suited for vacancy filling and thus the electrode becomes selective. When this electrode, along with a reference calomel electrode, is immersed in a fluoride solution, a KCl — NaF solution potential (E) is developed which is proportional to the logarithm of the activity of the fluoride ion in solution. E = 0.058 log [F–] + constant A calibration curve is drawn plotting measured potentials versus fluoride concentration. Then, the potential for the

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LaF3 — Eu (II) membrane

Fig. 4 Ion-selective electrode

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sample water is measured and the concentration of F– in the sample can be deduced from the calibration curve. The fluoride electrode responds only to hydrated F – but not to HF or FeF63–. If the pH of the medium is more than 9, there will be erroneous response by the electrode. In addition, the potential measured is proportional to the activity of the ion and not to the concentration. Therefore, the sample and the standards are brought to a constant ionic strength to render the electrode response reproducible. A buffer solution, called Total Ionic Strength Adjustment Buffer (TISAB) is used— as suggested by Frant and Ross—for obtaining reproducible results. TISAB is a mixture of an acetate buffer (to maintain pH), EDTA (to mask metal ions and thus release F – ) and sodium chloride (to maintain ionic strength).

REAGENTS TISAB reagent: 57 mL glacial acetic acid, 58 g NaCl and 3.3 g EDTA are taken in a beaker. After stirring the contents well, the beaker is cooled and a solution of NaOH (5 M) is added slowly with stirring until the pH of the solution is between 5.0 and 5.5. Dilute the resulting solution to 1000 mL. Standard NaF solutions: Weigh 1.05 g NaF (AR) and dissolve it in water. Make up this solution to 250 mL (0.1M). From this solution prepare four diluted fluoride solutions to obtain fluoride concentrations of 0.01, 0.001, 0.0001 and 0.00001 M respectively.

PROCEDURE To 10 mL of each standard F– solution, add 10 mL TISAB, mix gently and insert the fluoride electrode and calomel reference electrode into the solution. Wait for about three minutes and read the potential using a millivoltmeter correct to 0.1 mV. Plot the potential values against log [F–]. A straight line will be obtained. Then, measure the potential of 10 mL water sample after adding 10 mL TISAB to it. Read the concentration of fluoride ion in the water from the calibration curve.

Notes 1. The fluoride solutions are prepared in plastic beakers and flasks. The water sample should be collected only in plastic container. The potential measurements are also done in plastic beakers. 2. Fluoride contents of toothpaste, coffee and tea decoctions may also be determined by this method.

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21 1

Chloride

Chloride ion is essential to the electrolytic balance in our bodies. Since there is a continuous intake and excretion of chlorine from all animals, it is one of the most abundant anions found in wastewater and is a good marker ion for pollution sources. Chloride gives water a salty taste, detectable at a level of 250 ppm if the cation present is sodium. With calcium and magnesium as counter ions, this salty taste is not detectable until the chloride concentration reaches upto 1000 ppm. High chloride levels can pose a threat to crops and freshwater aquatic plants, which do not have mechanism to excrete excess salt. High levels of chloride in a water system increase the rate of corrosion of metallic pipes.

PRINCIPLE Mohr’s Method The chloride present in the water sample is titrated with silver nitrate solution. Cl- + Ag+ S AgCl (Ksp = 3 ´ 10-10) Potassium chromate is used as an indicator. At the end point, the concentration of chloride ion in solution approaches zero and the silver ion concentration increases to a level at which the solubility product of silver chromate is exceeded. Then, silver chromate is precipitated as a reddish-brown product. 2Ag+ + CrO42- S Ag2CrO4 (K sp = 5 ´ 10–12) The appearance of reddish-brown colour signals the end point. A small excess of Ag+ is required to produce a visible quantity of Ag2CrO4; therefore, an indicator correction has to be made in the titre value by performing blank titration. The pH of the water sample should be in the 7-8 range; when it is less, CrO42is converted to Cr2O72-; when it is more, Ag+ is precipitated as AgOH. For all the titrations, a definite volume of the indicator should be accurately pipetted into the analyte.

REAGENTS Distilled water: Redistilled water should be used for the preparation of all solutions.

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Standard sodium chloride solution (0.005 M): Dry NaCl (AR) at 140° C for an hour and cool it in a desiccator. Dissolve 0.2923 g of it in distilled water and make up this solution to 1000 mL. Silver nitrate solution (0.005 M): Dissolve 0.8496 g AgNO3 (AR) in distilled water and make up this solution to 1000 mL. Store this in a brown bottle. Standardize this solution using the standard NaCl solution and 1 mL K2CrO4 indicator. Make indicator correction in the titre value. Potassium chromate indicator: Dissolve 5g K2CrO4 in 100 mL distilled water.

PROCEDURE Pipette 100 mL water sample into a conical flask. Pipette 1 mL K2CrO4 indicator into it. Titrate the solution with the standardized silver nitrate solution taken in the burette. The end point is the appearance of reddish-brown colour. Perform a duplicate titration in an identical manner. Carry out a blank titration using 100 mL deionized, chloride-free water and 1 mL indicator. Subtract this titre value from that obtained for the water sample for indicator correction.

CALCULATION 1000 mL of 1M AgNO3 º 35.45g ClChloride in water sample (mg L–1 or ppm) 35.45 × V2 × M × 1000 V1 V 1 = volume of water sample in mL V 2 = corrected titre value for water sample titration in mL M = molarity of AgNO3

=

Notes 1. If the pH of water sample is less than 7, neutralize with 1M NaOH; if it is greater than 10, neutralize with 1M H2SO4. 2. Intense colour of water, sulphite, sulphide and cyanide, if present in water, interfere in this estimation. For decolorising the water, add 3 mL Al(OH)3 suspension to each 100 mL water and stir. Keep it aside for 30 minutes, filter and wash the floc. Collect the filtrate, including the washings, quantitatively for titration. Sulphite can be eliminated by the addition of 1 mL of 30% H2O2 to each 100 mL water sample, keeping pH at 7. Cyanide and sulphide can be eliminated by acidifying the sample with 1M H2SO4 to pH 4 and boiling for about 15 minutes. Then the cooled sample is neutralized with 1M NaOH. 3. If the chloride content of water sample is very low, 0.001 M AgNO3 can be used for titration. 4. The chloride content of water can be determined accurately by potentiometric titration using copper/copper sulphate as the reference electrode and a silver wire as the working electrode.

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Residual Chlorine

Drinking water is chlorinated with either free chlorine or hypochlorites for disinfection. Therefore, it contains residual chlorine, i.e., the excess not consumed by pollutants in water.

PRINCIPLE Starch-iodide method: This method depends upon the oxidizing power of free and combined chlorine in water to convert iodide ion to free iodine. This oxidation is represented as Cl2 + 2I- ® I2 + 2 ClI2 + starch ® blue colour The iodine released is titrated with a standard solution of sodium thiosulphate. The end point is the disappearance of blue colour. I2 + 2Na2S2O3 ® Na2S4O6 + 2NaI The amount of residual chlorine in water can then be calculated from the titre value.

REAGENTS Standard potassium dichromate solution (0.001M): Dissolve 0.147 g K2Cr2O7 (AR) in water and make up this solution to 500 mL. Sodium thiosulphate solution (0.005 M): Dissolve 1.245 g Na2S2O3.2H2O in previously boiled distilled water and make up this solution to 1000 mL. Standardize this solution by pipetting 25 mL of standard K2Cr2O7 solution, adding 10 mL of 10% KI solution and using starch indicator. Starch indicator: Prepare a fresh solution using soluble starch powder. Potassium iodide solution: Use AR sample and prepare a 10% solution. Acetic acid: Glacial

PROCEDURE Pipette 100 mL water sample into a conical flask. Add 5 mL glacial acetic acid and then add 5 mL KI solution. Gently swirl the flask. Titrate the liberated iodine

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with standardized thio in the burette. When the colour of the solution becomes straw-yellow, add 1 mL starch solution and continue the titration by adding thio in drops and with good stirring. The end point is the first disappearance of blue colour. Carry out a duplicate titration.

CALCULATION 1000 mL of 1M thiosulphate º 35.45g chlorine Available chlorine (in mg L–1 or ppm) =

35.45 ´ V2 ´ M ´ 1000 V1

V1 = volume of water sample in mL V2 = volume of thiosulphate in mL M = molarity of thiosulphate

Notes 1. If the water is highly coloured and turbid, the end point in the iodometric method is likely to be masked. 2. The iodometric method is convenient for water samples containing high chlorine concentrations (1 mg L–1). 3. Oxidants like manganic and ferric ions, if present in water, will interfere in this estimation. Under such circumstances, the titration is carried out at neutral pH. 4. The chlorine content of a water sample decreases with time. Therefore, the estimation should be completed soon after collecting the sample. 5. Orthotoludine is used as a colorimetric indicator for chlorine residuals. It is oxidized in acidic solution by chlorine to produce holoquinone—a yellowcoloured compound. Nitrites and oxidized forms of manganese oxidize orthotoludine to produce holoquinones, thus producing false indications of chlorine residuals. Hence, the acid orthotoludine procedure is no longer considered as a standard method for residual chlorine determination.

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Sulphide

PRINCIPLE The iodimetric method of determining sulphide is suitable for water samples containing more than 1 mg L–1 (ppm) sulphide. The water sample is treated with concentrated sulphuric acid and then carbon dioxide is blown through it. All the sulphide is carried away as hydrogen sulphide by the carrier gas. The outcoming gaseous mixture is then passed through a solution of zinc acetate to trap the hydrogen sulphide. The sulphide thus trapped is then estimated iodimetrically using starch indicator. H2S + I2 ® 2H+ + 2I- + S 2Na2S2O3 + I2 ® Na2S4O6 + 2NaI

REAGENTS Zinc acetate solution (0.1 M): Dissolve 11 g (CH3COO)2Zn.2H2O in water and dilute to 500 mL. Iodine solution (0.01 M): Dissolve 20 g KI in water and then add 0.1270 g I2 to this solution and stir thoroughly. Dilute this to 1 litre. Standardize this solution with 0.01 M Na2S2O3. Sulphuric acid: Concentrated (AR) Carbon dioxide gas Starch solution: Prepare a fresh solution Hydrochloric acid: Concentrated (AR)

PROCEDURE Fit a 2-holed rubber bung to a wide-mouth 500 mL bottle. Insert a bent glass tube (inlet) through one hole so that it reaches almost the bottom of the bottle. Through the other hole insert a bent delivery tube. Connect the outer end of the delivery tube to two 250mL conical flasks in series (Fig.5).

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Water sample + H2SO4 Zinc acetate solution

Fig.5

Apparatus for the estimation of sulphide

The whole assembly should be airtight. Pour 500 mL water sample into the wide-mouth bottle. Add 10 mL conc.H2SO4 to it. Pour 100 mL zinc acetate solution into each of the conical flasks and reassemble the apparatus airtight. Pass CO2 gas through the water sample for an hour. After this, combine the solutions in the two flasks quantitatively. To the combined solution, add a known excess of iodine solution (25 mL) and 10 mL conc. HCl. Stir the solution thoroughly. Titrate the unreacted iodine with standardized thio in the burette using starch indicator. The end point is the disappearance of blue colour. Conduct a blank experiment with 25 mL iodine solution in an identical way with 500 mL distilled water and determine the blank titre value.

CALCULATION 1000 mL of 1M iodine º 32 g S2– Total sulphide in water sample (mg L–1) =

(V2 – V1 ) × M × 32 ×1000 V3

V1 = volume of thiosulphate with water sample in mL V2 = volume of thiosulphate consumed in the blank titration in mL V3 = volume of water sample in mL M = molarity of thio thiosulphate

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Sulphite

PRINCIPLE Sulphite ion is a reductant. The water sample containing this ion is mixed with hydrochloric acid and KI solution. This mixture is then titrated with potassium iodate (KIO3). It first oxidizes sulphite and at the end of this oxidation, any excess KIO3 added will liberate iodine from KI. The end point, namely, the liberation of iodine, can be detected with starch.

REAGENTS Standard potassium iodate solution (0.0002 M): Dissolve 0.428 g dry KIO3 and 0.25 g NaHCO3 in distilled water. Make up this solution to 1000 mL. Dilute 100 mL of this to 1000 mL exactly. Potassium iodide solution: Dissolve 5 g KI in 100 mL water. Starch indicator: Prepare a fresh solution using soluble starch powder. Hydrochloric acid solution: Mix 100 mL conc. HCl with 100 mL distilled water.

PROCEDURE With a measuring cylinder, measure out 100 mL HCl into a conical flask. Pipette 100 mL water sample quickly into this, keeping the pipette tip below the liquid surface in the conical flask. Add 1 mL starch indicator and 5 mL KI solution. Titrate with standard KIO3 solution in the burette. The end point is the first appearance of a stable blue colour. Repeat the experiment with another 100 mL water sample to get concordant readings.

CALCULATION 1000 mL of 1M KIO3 º 80 g SO32-

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The amount of sulphite (in mg L–1 or ppm) =

80 × V2 × M ×1000 V1

V1 = volume of water sample in mL V2 = volume of potassium iodate used in mL M = molarity of potassium iodate

Note 1. The hydrochloric acid solution used should not liberate iodine from KI.

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Sulphate

PRINCIPLE The sulphate ion, present in a water sample, can be determined by turbidimetry. In the turbidimetric method, white light is passed through a finely divided suspension of the substance to be estimated. The light transmitted by the suspension is compared with the light transmitted by a suspension of known concentration and then the concentration of the sample is determined. The sulphate ions present in the water sample are converted to a suspension of barium sulphate using barium chloride reagent. Then, by turbidimetry, the concentration of barium sulphate is estimated. SO42– (aq) + Ba2+(aq) ® BaSO4(s)

REAGENTS Barium chloride solution: Dissolve 10 g BaCl2.2H2O (AR) in 100 mL water. NaCl-HCl reagent: Dissolve 24 g NaCl in 100 mL water containing 2 mL conc. HCl. Standard sulphate solution: Dissolve 0.3698 g Na2SO4 (AR) in water and make up the solution to 250 mL. This is a standard solution of 1000 mg L–1 (ppm) SO2– 4.

PROCEDURE Take five 50 mL numbered beakers and pipette 2,4,6,8,10 mL standard sulphate solution into these beakers. Add 10 mL NaCl-HCl solution to each and stir. Then add 10 mL barium chloride solution. Stir the mixture using a magnetic stirrer for 1 minute. At the end of the stirring period, quickly transfer the solution into the spectrophotometer cell. Read the absorbance of the suspension at 420 nm. Draw the absorbance vs sulphate concentration calibration line. Pipette 20 mL water sample into another 50 mL beaker and prepare the suspension in an identical manner. Measure its absorbance and read the

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concentration of SO2– 4 directly from the calibration line. Conduct a duplicate experiment with another 20 mL water sample.

Calculation ü Concentration of the sulphate ions in the ý = concentration from graph þ ´

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Phosphate

PRINCIPLE Phosphorus occurring as orthophosphate (HxPO4x–3) can be measured quantitatively by colorimetric method. Phosphate ion combines with ammonium molybdate under acidic condition to form a complex compound known as ammonium phosphomolybdate. PO43– + 12(NH4)2MoO4 + 24H+ ® (NH4)3PO4.12MoO3 + 21NH4+ + 12H2O The molybdenum content as ammonium phosphomolybdate is then reduced to a blue coloured solution called “molybdenum blue” or “heteropoly blue”. Either stannous chloride or ascorbic acid can be used as the reducing agent. (NH4)3PO4.12MoO3 + Sn2+ ® molybdenum blue + Sn4+ The intensity of blue colour, which is proportional to the amount of phosphate originally present, can be measured either by visual colorimetry or by spectrophotometry.

REAGENTS Sulphuric acid solution: Add 250 mL conc.H2SO4 to 500 mL distilled water slowly with cooling and stirring. Cool and add 5 mL conc. HNO3 to this and dilute the solution to 1 litre. Ammonium molybdate solution: Dissolve 6g ammonium molybdate in 50 mL distilled water in a beaker. In another beaker take 100 mL distilled water and add carefully 100 mL conc. H2SO 4 to it with cooling. To this cold acid, add ammonium molybdate solution and stir well. Stannous chloride solution: Add 2.5 g SnCl2.2H2O to 100 mL glycerol taken in a beaker. Heat the beaker in a water bath with constant stirring of its contents to get a clear solution.

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Phenolphthalein indicator Standard phosphate solution: Weigh 0.0439 g pure KH2PO4, dissolve it in water and make upto 100 mL. Pipette 10 mL of this solution into a standard flask and dilute to 100 mL. This gives a standard solution of 100 mg L–1(ppm) PO3– 4.

PROCEDURE Take five 25 mL, numbered, standard flasks. Pipette 1,2,4,8,10 mL diluted standard phosphate solution into these flasks. To each, add a drop of phenolphthalein indicator. If pink colour develops, add a few drops of sulphuric acid to discharge the pink colour. Then add 5 mL ammonium molybdate reagent and 0.5 mL stannous chloride reagent. Make upto the mark with distilled water, stopper and shake the solution well. Measure the absorbance of each of the solutions at 620 nm in a spectrophotometer. Plot the five absorbance values against the corresponding phosphate concentrations. A straight line would be obtained. Pipette 10 mL water sample into another 25 mL standard flask. Add the exact amounts of the reagents that were added to the standards, make upto the mark and measure the absorbance. Read the concentration of phosphate in the water sample from the linear plot corresponding to the absorbance of the sample.

Notes 1. Alternatively, ammonium phosphomolybdate may be treated with ammonium metavanadate to form a yellow vanadomolybdophosphoric acid complex. The intensity of absorbance of the complex at 420 nm is measured spectrophotometrically. It is proportional to the original concentration of orthophosphate in the sample. 2. The total phosphorus content is determined by performing an oxidative acid digestion to convert the various forms of phosphorus to orthophosphate, HxPO4x–3. For this purpose, the water sample (100 mL) is treated with 1 mL of 18 M sulphuric acid and 0.8 g ammonium persulphate and boiled gently for 1½ hours. 3. 1-Amino-2-naphthol-4-sulphonic acid can also be used as the reducing agent. It is prepared by dissolving 6 g NaHSO3 and 100 mg Na2SO3 in 40 mL water and then adding 100 mg 1-amino-2-naphthol-4-sulphonic acid. 4. Based on the total phosphorus content (nutrient), water bodies can be classified as oligotrophic or mesotrophic or eutrophic (vide appendix 2). 5. Phospate can also be estimated using sodium molybdate and hydrazine sulphate.(Refer: Experiment 55).

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Iron

Iron is present in soil mainly as insoluble ferric oxide and iron sulphide. In some places, it also occurs as ferrous carbonate, which is very slightly water-soluble. The carbon dioxide present in groundwater can react with ferrous carbonate forming soluble ferrous bicarbonate. FeCO3 + CO2 + H2O ® Fe(HCO3)2 This ferrous iron dissolved in water gets slowly oxidized to ferric iron under aerobic conditions. Therefore, natural water may contain both Fe2+ and Fe3+.

PRINCIPLE Ferric iron reacts with excess ammonium thiocyanate forming a blood-red complex ion [Fe(SCN)6]3– in acid solution. This coloured complex can then be estimated either by visual colorimetry or spectrophotometry. When the water sample contains both Fe2+ and Fe3+, only the ferric iron is estimated when thiocyanate is used as the reagent. The ferrous iron present in water sample can be oxidized to ferric iron by potassium permanganate and then the total iron estimated. The ferrous iron can be determined as the difference between total iron and ferric iron.

REAGENTS Ammonium thiocyanate solution (5M): Dissolve 380 g NH4CNS (AR) in water and make up the solution to one litre. Hydrochloric acid solution (5 M): Dilute 400 mL conc. HCl to one litre. Potassium permanganate solution (0.002 M): Dissolve 0.0316 g solid in 100 mL water. Standard ferric solution: Dissolve 0.8608 g Fe2(SO4)3 (NH4)2SO4 24H2O in 25 mL warm 1:1 H2SO4. Cool and make up this solution to 100 mL. This is a standard ferric solution of 1000 mg L–1 (ppm).

PROCEDURE Estimation of ferric iron: Pipette into six, numbered 50 mL standard flasks 1,2,4,6,8,10 mL standard Fe3+ solution. Add to each flask 1 mL dilute hydrochloric

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acid and 5 mL ammonium thiocyanate solution. Make upto the mark with distilled water and shake the flask. Measure the absorbance of each solution at 480 nm. Draw a linear plot by plotting the absorbance values against iron concentrations. Pipette 20 mL water sample into another 50 mL standard flask and add 5 mL dilute hydrochloric acid and 5 mL ammonium thiocyanate solution to it. Make upto the mark with distilled water and shake the flask. Measure the absorbance of this solution. Using this absorbance value, mark in the graph the Fe3+ concentration of the water sample. Determination of total iron: Pipette 20 mL water sample into a 50 mL standard flask. Add 5 mL dilute sulphuric acid and 1 drop potassium permanganate to it. If the colour is discharged, add one more drop of permanganate solution. Add 5 mL ammonium thiocyanate solution and make upto the mark with distilled water and shake the flask. Measure the absorbance of this solution and find out the Fe3+ concentration (total iron) from the graph. Subtracting Fe3+ concentration determined from this total iron concentration, the concentration of ferrous iron can be obtained.

Notes 1. If the iron concentration of water sample is very high, it can be diluted with distilled water suitably. 2. Only the minimum amount of potassium permanganate solution should be added for oxidizing ferrous iron. The minimum amount of permanganate can be found out separately by mixing a known amount of water sample with potassium permanganate solution in a beaker.

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Chromium

PRINCIPLE Hexavalent chromium (Cr2O72–) reacts with diphenylcarbazide (C6H5NHNH)2 CO under acidic condition to form a red-violet soluble complex. Traces of chromium can be estimated either by visual colorimetry or spectrophotometry using this reagent. Lower-valent Cr can be oxidized to Cr(VI) by KMnO4. Iron, copper and nickel, which may also be present in water, will interfere with this method. These metals should be removed from the sample by peroxidation treatment and filtration.

REAGENTS Sulphuric acid (1:1) solution: Add, with cooling, 50 mL conc. H2SO4 to 50 mL water. Diphenylcarbazide reagent: Dissolve 0.2 g reagent in 100 mL of 95% ethyl alcohol. Add 4N H2SO4 to this solution and mix well. Potassium permanganate solution (0.002 M): Refer experiment 27. Sodium azide: 0.5% solution in water. Standard K2Cr2O7 solution: Dissolve 0.3535 g K2Cr2O7 (AR) in water and make up this solution to 100 mL. Pipette 10 mL of this and dilute to 100 mL. Pipette 5 mL of the second solution and make it upto 250 mL. This diluted solution contains 1 mg L–1 (ppm) Chromium(VI).

PROCEDURE Pipette 2,4,6,8,10 mL final, diluted standard Cr(VI) solution into five Nessler tubes. Add to each tube, 5 mL 1:1 H2SO4 and 4 mL diphenylcarbazide solution. Dilute upto 50 mL mark with distilled water and stir. Pipette 50 mL water sample into a 100 mL beaker. Add 5 mL 1:1 H2SO4 to it and gently boil the solution. To the boiling solution, add KMnO4 solution drop by drop until a very faint pink colour persists (avoid adding excess KMnO4). Stop boiling the solution and add the sodium azide solution drop by drop just to discharge the pink

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colour of the solution. Cool and make up the solution to 100 mL in a standard flask. Pipette 25 mL of the made up solution into a Nessler tube and add 5 mL diphenylcarbazide reagent. Stir and add distilled water upto the 50 mL mark and stir. Compare the colour intensity with a series of standard solutions and find out the concentration range. Perform another set of experiments to find out the exact concentration of chromium. This estimation can be carried out spectrophotometrically also by reading the absorbance of the coloured solution at 540 nm.

Note 1. The colour comparison or absorbance measurement should be completed within 15 minutes after the addition of diphenylcarbazide.

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Copper

PRINCIPLE Divalent copper forms a coloured complex with cuprethol, whose concentration can be determined spectrophotometrically at 435 nm or visually using Nessler tubes.

REAGENTS Cuprethol reagent: Dissolve 2 g diethanolamine in 100 mL pure methyl alcohol. Add to it a solution of 1 mL carbon disulphide in 100 mL methyl alcohol. Store it in a tightly stoppered bottle. This is cuprethol reagent. Hydrochloric acid solution: Mix 50 mL conc. HCl with 50 mL distilled water. Sodium acetate solution: Dissolve 60 g sodium acetate trihydrate in 100 mL water. Warm, if necessary, to get a clear solution. Sodium pyrophosphate: Dissolve 3 g Na4P2O7.10H2O in water and make up the solution to 100 mL. Standard copper solution: Dissolve 3.929g CuSO4.5H2O in water and make upto 100 mL. Pipette 10 mL of this solution into another 100 mL standard flask and make upto the mark. This solution contains 1000 mg L–1 (ppm) copper.

PROCEDURE Pipette 2,4,6,8,10 mL diluted standard copper sulphate solution into five, numbered, 25 mL standard flasks. To each, add 0.5 mL HCl, 2 mL pyrophosphate solution, 5 mL sodium acetate solution (to maintain the pH at 5-6) and stir. After ten minutes, add 1 mL cuprethol reagent and make upto the mark with distilled water. Read the absorbance of the yellow solution at 435 nm. Construct the absorbance-concentration plot. Develop the colour with 10 mL water sample by adding exactly the same quantities of the reagents in another 25 mL flask. Make upto the mark with distilled water, shake and record the absorbance of this solution. Using the calibration graph, find the concentration of copper in the water sample.

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Notes 1. This estimation can also be done by visual colorimetry using Nessler tubes. 2. High concentrations of iron, lead, manganese, etc., interfere with this estimation. 3. The cuprethol reagent should be freshly prepared. It should not produce turbidity on adding to the dilute solution of cupric sulphate. 4. Neocuproine method: In neutral or weakly acidic solution, copper reacts with 2,9-dimethyl-l,10-phenanthroline to form a complex, which can be extracted with chloroform-methanol mixture. This yellow solution has maximum absorbance at 457 nm and hence copper is determined spectrophotometrically at 457 nm. 5. Sodium diethyldithiocarbamate (SDDC) forms a 1:2 copper-SDDC complex, which can be extracted into xylene. The absorbance of xylene extract is measured at 436 nm. 6. A solution containing Cu2+ ions can be directly aspirated into an airacetylene flame of an atomic absorption spectrophotometer (AAS) and measured at 325 nm. (Refer: Experiment 55)

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30 1

Lead

PRINCIPLE Dithizone or diphenylthiocarbazone forms highly coloured chalets with metals. Tetravalent lead forms a red chelate complex with dithizone (H2Dz), C6H5 – N = N – C – NH – NH – C6H5 || S The complex formed has the formula Pb(HDz)2 and is soluble in CCl4 solvent. The red complex can therefore be extracted with CCl4 and its concentration determined spectrophotometrically at 515 nm. Interfering metals (Cu, Cd, Zn, etc.) can be masked by cyanide ion.

REAGENTS NH4OH-Na2SO3-KCN Solution: Dissolve 3.5 g KCN and 12g Na2SO3 in water and add 400 mL concentrated ammonia solution to this solution. Dilute this mixture to one litre with water. Dithizone solution (0.005%): Dissolve 0.05 g dithizone in 1000 mL carbon tetrachloride. Keep the solution tightly stoppered. Standard lead solution: Dissolve 0.2496 g Pb(NO3)2 (AR) in distilled water and make up this solution to 250 mL. Pipette 10 mL of this and dilute accurately to 250 mL. From this dilute solution pipette 10 mL and dilute it accurately to 250 mL. The last solution contains 1 mg L–1 (ppm) lead.

PROCEDURE Pipette 10 mL final, diluted standard lead solution into a 100 mL standard flask. Add 40 mL distilled water, 25 mL NH4OH-Na2SO3-KCN solution and 25 mL dithizone solution. Stopper and shake the flask well. Draw out the clear CCl4 layer and measure the absorbance of this solution at 515 nm. Repeat the experiment with 20 mL, 30 mL, 40 mL and 50 mL standard lead solution, adding

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30 mL, 20 mL, 10 mL and 0 mL water respectively, and adding the same quantities of the reagents used earlier. Construct the calibration curve by plotting the absorbance values against lead concentrations. Pipette 10 mL water sample into a 250 mL separating funnel. Add 25 mL NH4OH-Na2SO3-KCN solution to it and then add 10 mL dithizone solution. Stopper the funnel and shake the contents of the funnel vigorously for about 2 minutes without spilling the liquid. Keep the funnel aside over a ring stand and after about 10 minutes separate the carbon tetrachloride layer into a clean, dry boiling tube. Measure the absorbance of this solution at 515 nm. Read the concentration of lead in the water sample from the calibration curve.

Notes 1. Potassium cyanide is highly poisonous; it should be kept in the fume hood and its solution should not be pipetted by mouth. 2. CCl4 is toxic and it should not be pipetted by mouth; a burette or a pipette with a bulb could be used. 3. Dithizone solution is not stable; therefore, only a fresh solution should be used. 4. Lead solution can be directly aspirated into air-acetylene flame of AAS and its concentration can be measured at 283.3 nm.

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Nickel

Nickel occurs naturally in trace amounts in most plant and animal tissues. This metal is extensively used in industries. Therefore, it is present in some water sources.

PRINCIPLE Ni(II) forms a scarlet red coloured chelate complex with dimethylglyoxime. This divalent insoluble complex is converted into a soluble, red, tetravalent complex in the presence of an oxidizing agent, e.g., bromine. By measuring the colour intensity of the solution containing this product, traces of nickel can be estimated. Iron interferes with this estimation, which is masked by the use of ammonium citrate.

REAGENTS Ammonium citrate solution (10%) : Dissolve 25g ammonium citrate in water and dilute to 250 mL. Bromine water: Saturated. Ammonium hydroxide solution: Mix 25 mL liquor ammonia with 75 mL water. Dimethylglyoxime solution (1% ): Dissolve 1g DMG in 100 mL methylated spirit. Standard nickel solution: Weigh 4.78 g NiSO4.7H2O (AR) and dissolve it in water. Make up this solution to 100 mL. Pipette 10 mL of this into another 100 mL flask and make upto the mark. This diluted solution contains 1000 mg L–1 (ppm) Nickel.

PROCEDURE Into five Nessler tubes, pipette 2,4,6,8,10 mL diluted standard nickel solution. To each tube, add the reagents in the following order: 5 mL ammonium citrate solution, 2 mL bromine water, 2 mL ammonia solution and 1 mL dimethylglyoxime solution. After adding each reagent, mix the solution thoroughly. Make up the solution to the 50 mL mark with distilled water and stir.

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In another Nessler tube, pipette 20 mL water sample and add the same amounts of the reagents as with the standard solutions. Make up the solution to 50 mL mark and stir. Compare the colour intensity of this solution with the colour intensities of the standard solutions and fix the approximate nickel concentration range. Carry out another set of experiments with suitable volumes of the standard solution and determine the concentration of nickel in the water sample. Repeat the experiment with another 20 mL sample of water.

Notes 1. Appreciable quantities of cobalt, copper and manganese, if present in water, interfere with this estimation. 2. If the coloured solution formed is turbid, it implies that the nickel concentration is high. Then, the solution has to be diluted suitably before estimation.

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Cadmium

PRINCIPLE Cadmium reacts with dithizone (diphenylthiocarbazone) to form pink to red coloured complex, which can be extracted with chloroform. The organic extract can be measured spectrophotometrically at 518 nm.

REAGENTS Sodium potassium tartrate solution: Dissolve 25 g KNaC4H 4O6.4H2O in 100 mL distilled water. NaOH-KCN solution: 40% NaOH + 1% KCN in distilled water. Hydroxylamine hydrochloride: 20% solution. Dithizone solution (0.01%): Dissolve 0.01 g in 100 mL chloroform. Standard cadmium sulphate solution: Dissolve 1.5 g CdSO4 (AR) in 100 mL water (stock).

PROCEDURE Dilute 1 mL stock solution of cadmium sulphate to 1000 mL. This contains 5 mg L–1 (ppm) cadmium. Pipette 20 mL water sample into a separating funnel and add 1 mL sodium potassium tartrate solution, 5 mL NaOH-KCN solution, 1 mL hydroxylamine hydrochloride and 15 mL dithizone. Stopper the funnel and shake for 1 minute. Transfer the chloroform layer into a beaker. Add 10 mL chloroform to the funnel, shake for 1 minute and separate the chloroform layer. Mix with chloroform extract in the beaker and measure the absorbance at 518 nm against the reagent blank. Repeat the procedure for a series of standard solutions and plot a graph of absorbance versus concentration. Determine the amount of cadmium in the water sample from the graph.

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Notes 1. Cadmium solution can be directly aspirated into an air-acetylene flame of AAS and the absorbance can be measured at 229 nm. 2. If the amount of cadmium lies in the range of 1-100 mg L–1 (ppm), polarographic methods can be employed.

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Mercury

PRINCIPLE Mercury reacts with dithizone (diphenylthiocarbazone) in acidic medium to produce an orange chelate. It is extracted with CCl4 and measured at 490 nm. C6H5N=NCSNHNHC6H5 + Hg ®

H5C6HN—N H5C6N

N

C

S

Hg

S

N

NHC6H5

N

NC6H5

C

REAGENTS Potassium permanganate Hydrogen peroxide Sulphuric acid solution (1M)

: 5% solution : 30% solution : Take 5.6 mL conc. acid and make upto one litre Dithizone solution : Dissolve 20 mg in 100 mL CCl4 Standard mercuric chloride solution : Dissolve 2.82 g HgCl2 in 100 mL water. Dilute 1 mL of this solution to 1000 mL. This contains HgCl2 15 mg L–1 (ppm) mercury.

PROCEDURE Place 100 mL water sample in a 500 mL distillation flask. Add 5-10 mL of 5% KMnO4 and reflux for four hours. Cool and add a few mL of 30% H2O2 to remove KMnO4. Boil to remove excess H2O2. Cool and add a few drops of 1 M H2SO4. Extract the metal twice with 10 mL dithizone solution. Combine the extracts and measure the absorbance at 490 nm against a reagent blank. Dilute the standard solution and prepare a series of standard solutions containing mercury in the range of 2-10 mg L–1 (ppm). Repeat the procedure for these

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solutions and construct a calibration graph. Determine the amount of mercury in the water sample from the calibration graph.

Notes 1. Mercury ion can be reduced to elemental mercury using SnCl2 and it can be swept by air into the absorption cell of flameless AAS. Mercury is then estimated by measuring absorbance at 253.7 nm. 2. The metals, copper, lead, nickel, cadmium and mercury can also be determined by anodic stripping voltammetry (ASV). Its limits of detection range to ppb and below.

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Cyanides

Free cyanide occurs in wastewaters as a result of discharge from metal refining and cleaning units, electroplating industries and coke ovens. Cyanide is very toxic and therefore utmost care must be taken while handling samples containing cyanides.

PRINCIPLE The sample solution is distilled in the presence of mineral acids to remove all cyanide and it is absorbed in sodium hydroxide. The distillation is carried out in the presence of magnesium chloride solution to remove the thiocyanate formed. acid distillation

® HCN Simple cyanides ¾¾¾¾¾¾ The cyanide content may be determined by titration method or spectrophotometrically or by ion selective method, depending on the concentration of cyanide ions present in the sample.

REAGENTS Magnesium chloride solution (5%): Dissolve 5 g magnesium chloride in 100 mL water. Sodium hydroxide solution (1M): Dissolve 4 g sodium hydroxide in 100 mL water. Sulphuric acid solution (9M): Mix 500 mL conc. acid with 500 mL water. p-dimethylamino-benzalrhodamine indicator: Dissolve 0.02 g reagent in 100 mL acetone . Silver nitrate solution (0.02M): Dissolve 1.7g AgNO3 in 500 mL distilled water.

PROCEDURE To 250 mL water sample taken in a distillation flask, add 50 mL sulphuric acid (9M) and 20 mL magnesium chloride solution (5%) and reflux for about an hour. Collect the HCN gas in 500 mL of 1M NaOH.

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To 20 mL distillate, add 0.5 mL p-dimethylamino-benzalrhodamine indicator and titrate with 0.02 M silver nirate solution till the colour changes from canary yellow to salmon blue. From the titre value, calculate the amount of cyanide in a litre of the water sample.

Notes 1. The cyanide in alkaline distillate can be converted into CNCl by reacting with chloramine T at pH less than 8. It is then treated with pyridinebarbituric acid and the resulting red solution can be spectrophotometrically determined at 578 nm. 2. The distillate can be treated with 0.5 mL each of conc. HCl and bromine water and then 0.5 mL arsenious acid solution. To this, 5 mL pyridine-pyrazolone reagent can be added and the cyanide content can be read at 620 nm. 3. Cyanide content can also be determined using ion selective electrode and measuring the potential of the distillate. 4. The type of analysis depends on the concentration of the cyanide ions. If the concentration exceeds 1 mg L–1, titration method is used. Spectrophotometric method is used if the amount is between 20 mg L–1 and 1 mg L–1. Ion selective electrode method is used if the concentration is between 0.05 and 10 mg L–1.

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35 1

Boron

Boron is an essential element for plant growth, but if the concentration of boron is more than 2 ppm in the irrigation water, it is harmful for plants. Drinking water should not contain more than 1 ppm boron.

PRINCIPLE Acidification and evaporation of boron with curcumin reagent gives a red coloured rosocyanin, which is soluble in alcohol. This can be determined colorimetrically.

REAGENTS Curcumin reagent: To 80 mL ethyl alcohol, add 40 mg curcumin and 5.0 mg oxalic acid. Mix well and then add 42 mL of hydrochloric acid and make upto 100 mL using ethyl alcohol. Sodium hydroxide solution (1N): Add 4 g sodium hydroxide to 100 mL distilled water and stir well. Hydrochloric acid solution: Transfer 10 mL concentrated hydrochloric acid into 110 mL water. Standard borax solution: Dissolve 10 mg borax (AR) in 100 mL water to get the stock solution containing 22 mg L–1 (ppm) boron.

PROCEDURE Take 1 mL water sample in an evaporating dish and add 2 mL sodium hydroxide solution. Evaporate on a water bath at 55° C to dryness. Ignite the contents at 550° C and cool. Add 2.5 mL hydrochloric acid and mix well. Centrifuge and pipette the supernatant liquid and make upto 25 mL. Take 1 mL of this solution in the dish and add 4 mL curcumin reagent and 4 mL hydrochloric acid. Gently heat on a water bath at 55° C for an hour. Add 10 mL ethyl alcohol and dissolve the residue. Measure the absorbance at 540 nm against a reagent blank. Prepare a series of standard solutions by appropriately diluting the stock borax solution and repeat the experiment with these solutions. Construct a calibration graph with absorbance versus concentration. From the graph, note the concentration of boron in the water sample corresponding to its absorbance.

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Note 1. Alternatively, 2 mL water sample can be mixed with 2 drops conc.HCl and 10 mL conc.H2SO4. After cooling, the solution can be mixed with carmine reagent and the absorbance can be measured at 585 nm.

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36 1

Selenium

Selenium is a suspected carcinogen and it can lead to hair loss. In the case of cattle, prolonged exposure may prove to be fatal.

PRINCIPLE Selenium is oxidized to selenate using acidified potassium permanganate and then it is reduced to selenite using HCl. It is then allowed to form a chelate with diaminobenzidine at pH 1.5. The chelate is extracted into toluene and the absorbance of the yellow extract is measured at 420 nm.

REAGENTS Potassium permanganate solution (0.02M): Dissolve 0.316 g solid in 100 mL water. Hydrochloric acid solution (0.1M): Add 2 mL conc. HCl to water and make upto 250 mL. Calcium chloride solution: Dissolve 30 g CaCl2.2H2O in one litre water. Sodium hydroxide solution (0.1M): Dissolve 0.4 g pellets in 100 mL distilled water. EDTA–sulphate reagent: Dissolve 100 g EDTA and 200 g sodium sulphate in one litre water and add conc. ammonia till the solution becomes clear. Diaminobenzidine reagent: Dissolve 100 mg 3, 3-diaminobenzidine hydrochloride in 10 mL water. Standard sodium selenate solution: Dissolve 10 mg sodium selenate (AR) in 100 mL water to obtain the stock solution. This contains 42 mg L–1 (ppm) selenium.

PROCEDURE To 100 mL water sample, add a drop of methyl orange and neutralize with 0.1M hydrochloric acid to orange-red end point. Add three drops of KMnO4 and 5 mL calcium chloride solution. Heat to boiling. Maintain the purple tint by adding KMnO4 intermittently. Reduce the volume to 25 mL.

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Add 5 mL sodium hydroxide solution (0.1M) and evaporate to dryness. Cool and add 5 mL conc. hydrochloric acid and 1 mL ammonium chloride solution (25%). Heat on a steam bath for ten minutes. Add 5 mL EDTA-sulphate reagent and add ammonia to clear the solution. Adjust the pH to 1.5. Add 1 mL diaminobenzidine solution and heat on a steam bath. Cool and add NH4OH to adjust the pH to 8.0. Transfer the contents into a separating funnel. Add 10 mL toluene and shake for 30 seconds. Remove the organic layer and measure the absorbance at 420 nm. Prepare a series of standard solutions by appropriately diluting the standard selenate solution and repeat the experiment with these solutions. Construct a calibration plot of absorbance with concentration. From this plot, find the concentration of selenium in the water sample corresponding to its absorbance.

Note 1. Selenium can also be determined using atomic absorption spectroscopy. Selenium is reduced to SeH2, which is directly aspirated into Ar-H2 flame. The measurement is done at 196 nm.

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37 1

Arsenic

Arsenic occurs in water as a result of mineral dissolution, industrial discharges and the application of insecticides containing arsenic. The usual level of arsenic in potable water is within 10 ppb. Chronic effects can appear from the accumulation of arsenic in the body at low intake levels for prolonged periods.

PRINCIPLE Arsenic is reduced to AsH3 by zinc in acid medium and AsH3 is led through a scrubber containing glass wool impregnated with lead acetate solution into a tube containing silver diethyldithiocarbamate dissolved in pyridine. Arsenic forms a red complex with this salt and the absorbance of the complex is measured at 535 nm.

REAGENTS Potassium iodide solution: Dissolve 15 g KI in 100 mL water. Stannous chloride solution: Digest tin bits in conc. HCl for about thirty minutes to obtain an approximately 4% SnCl2 solution. Zinc metal powder Lead acetate solution: Dissolve 10 g lead acetate in 100 mL water. Silver diethyldithiocarbamate reagent: Dissolve 0.05 g reagent in 100 mL pyridine. Concentrated HCl Standard sodium arsenite solution: Weigh 10 mg arsenious oxide in a beaker and dissolve it in a minimum amount of 10% sodium hydroxide solution. When the oxide is completely dissolved, transfer it quantitatively into a 100 mL standard flask and make upto the mark with water. This contains 76 mg L–1 (ppm) arsenic (highly poisonous).

PROCEDURE Take 25 mL water sample in a clean Gutzeit generator bottle and add 5 mL conc. HCl, 2 mL KI solution and 0.5 mL SnCl2. It will take 15 minutes for the complete reduction of As(V) to As(III).

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Add 3 g zinc powder to the solution and connect the generator to the scrubberabsorber assembly. It will take 30 minutes to complete the evolution of AsH3. Warm the generator to ensure the complete removal of AsH3. AsH3 passes through the scrubber containing glass wool impregnated with lead acetate and then through the absorber containing diethyldithiocarbamate reagent. Pour the solution from the absorber into 1 cm cell and measure the absorbance at 535 nm. Prepare a series of standard solutions by appropriately diluting the standard arsenite solution (warning: use the burette; do not pipette the solution) and repeat the experiment with these solutions. Draw the absobrance-concentration plot. From this plot, determine the concentration of arsenic in the water sample corresponding to its absorbance.

Note 1. AsH3 can be directly aspirated into Ar-H2 flame of atomic absorption spectrophotometer and the absorbance can be measured at 193.7 nm.

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38 1

Suspended Solids

The term “solids” comprise a wide variety of organic and inorganic materials in water. Suspended and dissolved particles in water are termed as total suspended solids (TSS) and total dissolved solids (TDS) .

PRINCIPLE A sample of water is evaporated as such to obtain the total solids present in it. The total dissolved solids is determined by filtering the sample before evaporation. The difference between total solids and total dissolved solids gives the total suspended solids present in the water sample.

PROCEDURE i. Accurately weigh a clean, dry 100 mL silica crucible and record its mass (w1). Transfer 100 mL unfiltered sample of water in the crucible. Evaporate the sample by placing it in an air oven at 105o C for an hour. Record the mass of the crucible after cooling it in the desiccator (w2). Total solids (mg L–1) =

(w1 – w 2 ) x10000 , where S is the mass of the S

sample. ii. Accurately weigh a clean, dry 100 mL silica crucible and record its weight (w1). Filter 100 mL water sample and transfer the filtered sample into the crucible. Evaporate the crucible in an air oven at 105o C for an hour. Cool in a desiccator and obtain the mass of the crucible (w2). Total dissolved solids (mg L–1) =

(w1 – w 2 ) x10000 , where S is the of the S

mass sample. iii. Total suspended solids can be obtained by subtracting the mass of total dissolved solids from that of total solids. Total suspended solids = Total solids –Total dissolved solids

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39 1

Determination of Chlorophyll- a

PRINCIPLE An estimate of the green plant matter in a water body can be obtained by measuring the chlorophyll-a content of the water body. The chlorophyll-a content is determined by the extraction of the pigment into acetone, followed by the measurement of its visible absorbance.

PROCEDURE Filter 100-500 mL water sample through a glass-fibre filter. Record the volume filtered. Roll the filter and place it in a small vial or centrifuge tube. Add 10 mL 90% acetone to the vial, stopper it and store in the dark at 4°C for 24 hours. Centrifuge the extract and measure the absorbance of the extract at 665 and 750 nm. Subtract the 750 nm value from 665 nm value. The 750 nm value serves to correct for any turbidity. Add a drop of 2 M HCl to 3 mL sample and measure the absorbance at 665 and 750 nm and subtract the values. The remaining figure is taken as the absorbance due to pigments other than chlorophyll-a.

CALCULATION Chl-a ( mg L–1) = 29 (A–A a ) ´

Vol. extract (mL) ´ 10–3 Vol. of sample (L)

where A is the corrected absorbance before acidification and A a is the corrected absorbance after acidification.

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40 1Determination of Productivity PRINCIPLE The production of green plant matter (productivity) occurs as a result of photosynthesis with a proportionate production of oxygen. CO2 + H2O ® CH2O + O2 So, the determination of oxygen produced in a water body will serve as a measure of productivity. However, respiration also occurs in any water body, resulting in the consumption of oxygen. A measure of this loss is needed to calculate productivity. The productivity is determined in situ by the measurement of dissolved oxygen change in sub-samples held in ‘light’ and ‘dark’ bottles.

PROCEDURE Take four BOD bottles and make two of them light-proof by wrapping them with aluminium foil. These bottles will act as light and dark bottles. Obtain a sample of water from the middle of the upper half of the water body and measure the dissolved oxygen (Oi). Fill a pair of light and dark bottles with this water sample. Suspend these bottles in the water body at depths corresponding to where the water sample was taken. Expose the bottles for several hours. The usual exposure period is from dawn to noon or noon to dusk. At the end of the exposure period, remove the bottles and measure the dissolved oxygen in each (Ol and Od for light and dark bottles respectively).

CALCULATION Respiration (R) = O2 decrease in dark bottle = Oi – Od Net photosynthesis (Pn ) = O2 increase in light bottle = Ol – Oi Gross photosynthesis (Pg ) = Respiration + net photosynthesis = (Oi – Od) + (Ol – Oi) = Ol – O d

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Daily Pg (mg O2 L–1 day–1) = Pg ´

Length of photoperiod per day ´ water depth (m) Length of exposure time

Daily Pg (mg O2 m–3 day–1) = Pg ´

Length of photoperiod per day ´ 103 ´ water depth (m) Length of exposure time

The factor 103 converts the volume concentration from mg L–1 to mg m–3 Daily R(mg O2 m–3 day–1) = R ´

24 ´ 103 ´ water depth(m) Exposure time (hr)

Daily Pn (mg O2 m–3 day–1) = Daily Pg – Daily R Daily Pn (mg C m–3 day–1) = Daily Pn (mg O2 m–3 day–1) ´

12 32

(based on CO2 + H2O ®CH2O + O2)

Note 1. The length of the photoperiod is roughly 12 hours. However, it is to be obtained from the meteorological department for the given day.

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PART TWO

SOIL ANALYSIS

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41 1

Collection of Soil Samples

Only a minute fraction of the large soil mass of the field is used for analysis in the laboratory. Therefore, it is important to have a truly representative sample of a particular area. If the field is uniform in colour, slope, crop growth and texture, it can be treated as a single sampling unit. If there is any variation in any of the above, separate samples must be collected from each area. During sampling, recently fertilized plots, spots near trees, compost piles and non-representative locations must be avoided. Salt-crusts on the surface of soils should be sampled separately. Once the area is selected, sample should be taken at various points. The points should be selected along a zigzag path in the field. Samples should not be taken from 60 cm border of the plot. If the land is fertilized and if the straight line path is chosen, then all the sampling points may fall in the fertilized band making the sample unrepresentative of the plot soil. For a real composite sample, all soil cores must have equal volume and weight. Any of the tools, like tube auger, screw-type auger, post-hole auger, spade, or knife can be used for digging the plot. Spade or tube auger is satisfactory for moist and soft soil. Screw-type auger is convenient for hard or dry soil, while post-hole auger is useful for wet area. Using a spade, a V-shaped cut is made first upto a depth of 15 cm (plough-layer) and a slice of about 2 cm thickness is removed. While taking a soil sample, only a soil volume is sampled (i.e., area ´ depth) and not mere area. Samples are collected from 3 or 4 well-distributed spots after scrapping the surface a little. These are collected in a suitable container—a trough or a bucket. Subject the sample collected to the following processes:

SIEVING Pass the bulk sample through a 6 cm (4 meshes per inch) sieve by rubbing with fingers. Discard the stones and gravels on the sieve.

DRYING Dry it in air at room temperature in shade (at about 25° C and the relative humidity of 20-60%). Some analyses require fresh moist samples from the field and some need oven-dried (100-110° C) samples.

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GRINDING Break up the bigger particles of soil by grinding lightly with a roller or with wooden pestle and mortar. Pass the heavy soils through a 2 mm sieve before they get dried completely.

MIXING Place the sieved, dried and ground sample on a piece of clean cloth. Grasp the opposite corners of the cloth and hold one corner down pulling the other across the sample. Repeat this in the reverse direction. Carry out the same procedure using the other two corners. Continue this for 6 or 7 times so that the mixing will be thorough.

CONING AND QUARTERING Partition the soil as follows. Keep the mixed soil at the centre of a clean paper. Make it into a cone. Flatten and divide it into two halves through the centre using a flat metal spatula. Then divide each half further into two halves. Separate the four portions. These are called ‘quarters’. Discard the two diagonally opposite ‘quarters’. Mix the other two thoroughly and repile. Continue this fixed procedure (cutting, randomly choosing opposite quarters, mixing and grinding, and repiling) until a sample of the size needed for replicate analysis in the laboratory is obtained. This technique minimizes bias in the sampling.

STORING Place the ‘quartered’ sample of soil in paper cartons using a polythene bag as an inner lining, or in an airtight glass jar with a screw cap. Label it properly with the location from where the soil is taken and the date of sampling.

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pH Value

PRINCIPLE The pH of the soil is defined by the following equation: Soil pH = –log aH+ , where aH+ is the activity of the hydrogen ion. The effective concentration of hydrogen ions includes H+ formed by the dissociation of soluble acids and soil particles. The soil pH will indicate the nature of the soil (i.e., acidic, alkaline or neutral). The growth of the crop suffers under strongly alkaline (very high pH) and strongly acidic (very low pH) conditions. If the pH is known, suitable measures can be taken for soil reclamation. Determination of the pH of a soil suspension by electrometry is a satisfactory method. The instrument used generally is a glass electrode pH meter. Digital pH meter can also be used. The pH determination is preferably done at the moisture saturation percentage because it is easy for routine wetting of different soils to equipotential moisture status; the moisture films are sufficiently thick to give good contact with glass electrode; this moisture content is the highest probable in the field; and small changes in dilution will produce only a very small change in pH.

REAGENTS Buffer solution of pH 4.0: Dissolve 5.105 g potassium hydrogen phthalate (KHC8H4O4) in distilled water and make upto 500 mL. Buffer solution of pH 9.2: Dissolve 9.54 g sodium tetraborate (borax) in distilled water and make upto 500 mL.

PROCEDURE Set the temperature knob of the instrument at 25º C since the pH scale 0 to 14 is valid only at 25º C. This is because the dissociation increases with temperature. Calibrate the pH scale with buffers of known pH. Take the soil sample in a 50 mL breaker to half its height (approximately 25 mL). Add distilled water in small amounts. Do not stir the soil and continue

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the addition of water until it just wets the entire mass. Add a few more drops very slowly till the surface glistens slightly. Stir the soil with a glass rod (avoid stirring until the soil becomes completely wet as otherwise a puddle mass will be formed through which the water moves very slowly) and then add a few drops of water. Make a ‘thin paste’ of the soil. Check this by ensuring that the paste barely flows together to close a hole left by the rod. If not, add a little more water or soil (in case free water stands on the surface) and remake the paste. The soil is now at ‘moisture (or water) saturation percentage’ and this is an equipotential moisture content for various soils. The surface of such a soil glistens and the soil is also said to be at the ‘flow point’ or ‘liquid limit’. Insert carefully the electrode (s) of the pH meter in this paste. Measure the pH when the reading is constant. Ensure the absence of water film around the glass electrode by moving it a little. Repeat the measurement 2 or 3 times to get consistent readings. Each time, wash the electrodes with distilled water and wipe them with a filter or tissue paper. Check the calibration frequently with buffers. Stir the soil suspension well whenever a reading is taken.

Notes 1. The electrodes should be kept in distilled water when not in use. 2. Terminology commonly used to describe the acid-base status of soils is as follows: pH less than 4 strongly acidic 4 to 5 moderately acidic 5 to 6 slightly acidic 6 to 8 neutral 8 to 9 slightly alkaline 9 to 10 moderately alkaline greater than 10 strongly alkaline

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43 1

Soluble Salts

PRINCIPLE Ions carry electricity. Hence, electrical conductivity of the soil-water system increases with the amount of soluble salts present in the soil. A linear relationship exists between the conductivity (specific conductance) of water-extract of the soil and the ionic strength of the salts, as found by analysis. Therefore, the concentration of soluble salts can be directly determined by measuring the electrical conductivity. This helps in ascertaining the salinity of the soil and hence in the reclamation of the soil. Conductance (C) is usually measured in millimhos (1 mho = 1000 millimhos). Conductivity (k) is the conductance of 1 cubic centimetre of an electrolyte solution at 25° C.This can be measured with a cell of known cell constant. The conductivity of the solution is conveniently expressed in millimhos cm–1, which is equivalent to the SI unit decisiemens per metre (dS m–1 ). Cell constant (K) is usually determined with a standard KCl solution. k KCl , where kKCl and C KCl are the conductivity (known) and the C KCl conductance of standard KCl respectively. Thus, for the soil solution,

K =

ksoil = K ´ Csoil

REAGENT Standard potassium chloride solution (0.05 M): Weigh accurately 0.3728 g dry KCl (AR) and dissolve it in conductivity (double-distilled) water. Make up the volume to 100 mL. The conductivity of this solution is 6.669 millimhos cm–1 at 25º C.

PROCEDURE For measuring the conductance, an electrical resistance bridge and a cell with Pt black-coated electrode or an AC salt bridge can be used. The instrument, which directly gives the conductance in millimhos cm–1 at 25º C, may also be used.

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Check and calibrate the instrument with standard KCl solution. Measure the conductance of this solution with the given cell. Knowing the conductivity and conductance, calculate the cell constant, K. Then, take 30 g soil sample in a 100 mL conical flask. Add 60 mL distilled water and shake well for about an hour. Allow it to stand for about 10 minutes and determine the conductance of the supernatant liquid after setting the temperature knob in the instrument at 25º C.

CALCULATION Conductance of 0.50 M KCl solution = CKCl Cell constant, K =

=

k KCl C KCl

6.669 cm–1 C KCl

Conductivity of soil solution, ksoil = K ´ Csoil millimhos cm–1 So, concentration of salts = 12.5 ´ ksoil ppm (mg L–1). This equation is based on the fact that a linear relationship exists between the conductivity of the water extract of a soil and the concentration of salts in the soil (the concentration of salts in the soil may be independently found by analysis but the procedure is elaborate). The proportionality constant was found to vary for different soils; however, for most of the soils, the value 12.5 is a close approximation. (Reference 10, p.243).

Note 1. Soils are classified based on their electrical conductivity (EC) of the saturated extract and exchangeable sodium percentage (ESP). Exchangeable sodium percentage is the fraction, expressed as a percentage of exchangeable sodium ions compared to total exchangeable ions. · Normal soils have EC values less than 4 ds m–1 and ESP values below 15%. · Saline soils have EC values greater than 4 ds m–1 with an ESP value less than 15%. · Sodic soils have EC values less than 4 ds m–1 but ESP greater than15%. · Saline-sodic soils have both high EC (>4 ds m–1) and high ESP (>15%).

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44 1Identification of Calcium Carbonate,

Nitrogen, Phosphorus and Potassium

PRINCIPLE Nitrogen is usually present in soil as ammonium, nitrate or nitrite ion. Phosphorus is present as orthophosphate. Potassium occurs as its silicate. Calcium carbonate is present mainly in the free state. The given soil sample is extracted with Morgan’s extracting solution, which is a mixture of sodium acetate trihydrate and acetic acid solutions. This soil extract is tested with the respective reagents for the presence of various ions. As the micronutrients (Zn, Cu, Fe and Mn) are present in very small quantities in soil, the qualitative detection is not successful for these elements.

REAGENTS Extracting solution: Dissolve 50 g sodium acetate trihydrate in about 250 mL distilled water. Add 15 mL glacial acetic acid and make upto 500 mL. Nessler’s reagent: Dissolve 12 g HgI2 and 8 g KI together in ammonia-free water and make upto 50 mL. To this add 50 mL sodium hydroxide (6M) solution. Allow it to stand overnight and then decant. =-naphthylamine solution: Dissolve 1.22 g solid in 250 mL of 20% (v/v) hydrochloric acid. Sulphanilic acid solution: Dissolve 1.5 g solid in 250 mL of 2% (v/v) hydrochloric acid. Ammonium molybdate solution: Dissolve 62.5 g ammonium nitrate in 60 mL water and add 90 mL nitric acid of specific gravity 1.42. Dissolve separately 6.25 g ammonium molybdate (AR) in 40 mL water. Add this slowly to the nitrate solution with constant stirring. Make up the volume to 250 mL with water. Place the solution at 60º C for 6 hours in a water bath. Then, allow the solution to stand overnight. Filter through a Whatman No. 42 filter paper, if necessary. This reagent is quite stable for at least 3-4 months.

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Stannous chloride solution: Prepare a stock solution by dissolving 20 g crystalline SnCl2 in 50 mL conc. HCl. Warm, if it does not dissolve in cold. Keep this solution in an amber coloured bottle. Just before use, dilute 1 mL stock solution to 130 mL with distilled water. To keep the stock solution for long, add a piece of tin metal (AR).

PROCEDURE Weigh accurately about 5 g soil sample and transfer it to a glass bottle or a conical flask carrying a lid. Add 50 mL extracting solution and shake well for about an hour. Filter through Whatman No. 1 filter paper. Analyse the filtrate (i.e., soil extract) as follows: Sl. No.

Experiment

1. To a portion of the extract, add a few drops of dilute HCl (1:1) Pass the gas through Ba(OH)2 solution 2. To another portion of the extract, add a few drops of Nessler’s reagent. 3. Treat a few mL extract with a few drops of dilute acetic acid. Then add a few drops each of sulphanilic acid and = -naphthylamine solutions. 4. Add a small quantity of dilute sulphuric acid to a few drops of the extract. Then add a few crystals of KI followed by 2 or 3 drops of fresh starch indicator. 5. Brown-ring test for nitrate (if NO2 is absent): To a part of the extract, add a little dilute sulphuric acid and freshly prepared ferrous sulphate solution. Then add, slowly and carefully, concentrated sulphuric acid in drops along the sides of the test tube. 6. If nitrite is present, eliminate it as below. Acidify a portion of the extract with dilute sulphuric acid and heat with solid NH4Cl or urea until no more gas is evolved. This treatment decomposes nitrite into nitrogen which volatilizes off.Test a small portion of this solution with dilute sulphuric acid, KI and starch to ensure the complete removal of

Observation

Inference

Effervescence takes place Milkiness appears Reddish-brown colouration or precipitate Red or pink colouration

Presence of carbonate (CaCO3 ) Presence of carbonate Presence of ammoniacal nitrogen Presence of nitrite nitrogen

Blue colour

Presence of nitrite nitrogen

A brown-ring at the Presence of nitrate junction of two nitrogen layers

(Contd.)

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Identification of Calcium Carbonate, Nitrogen, Phosphorus and Potassium Sl. No.

a. b.

7.

8.

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Experiment nitrite. Then divide the nitrite-free solution into two parts. Perform brown-ring test with one part of the solution. To another part, add a pinch of zinc dust and heat. Then add dilute sulphuric acid, KI and starch indicator. Now add zinc dust and heat in order to reduce nitrate to nitrite To a little of the extract, add a few mL each of conc. HNO3, ammonium molybdate and dilute SnCl2 solutions. To a small quality of extract, add saturated solutions of cobaltnitrate and sodiumnitrite followed by acetic acid.

101

Observation

101

Inference

Formation of a brown-ring No blue colour

Presence of nitrate nitrogen Indicates that nitrite has been removed completely.

Blue colour

Presence of nitrate nitrogen Presence of phosphorus

Blue colour

Yellow precipitate

Presence of potassium

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45 1

Calcium Carbonate

PRINCIPLE In calcium carbonate, the carbon is regarded as inorganic carbon. Based on the intensity of effervescence, the amount of CaCO3 in the soil can be qualitatively divided into (a) very high (b) high (c) medium (d) low, and (e) traces. Quantitative estimate can then be made for soil of high and very high CaCO3 content by the reaction between a known mass of sample and a known excess volume of standard hydrochloric acid. The unreacted acid is back-titrated with a standard alkali. From the quantity of acid consumed, the amount and hence the percentage of CaCO3 in the soil can be calculated. Since CaCO3 gives the measure of alkalinity of soil, its estimation is desirable in order to make the soil suitable for plant growth. CaCO3 + 2HCl ® CaCl2 + H2O + CO2 (from soil) HCl + NaOH ® NaCl + H2O (unreacted)

REAGENTS Hydrochloric acid: Prepare approximately 0.5 M solution by adding 40 mL conc. HCl in water and making upto one litre. Sodium hydroxide: Prepare approximately 0.5 M solution by dissolving 20g pellets in water and making upto one litre. Standardize using oxalic acid. Phenolphthalein indicator: Dissolve 2.5 g indicator solid in 125 mL rectified spirit. To this solution, add 125 mL distilled water and stir.

PROCEDURE Weigh accurately about 1 g soil sample and place it in a dry 250 mL conical flask provided with a lid. Add 30 mL HCl (~0.5M) and warm the contents. Keep the solution aside for 10 minutes. Then titrate the mixture with standard NaOH solution using phenolphthalein indicator. Carry out a duplicate under identical conditions.

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Calcium Carbonate

103

Perform a blank titration with 10 mL standard HCl (~0.5M) against NaOH solution (standard) using phenolphthalein indicator.

CALCULATION Mass of the soil taken 10 mL standard HCl 30 mL standard HCl where M is the molarity of NaOH. W g soil + 30 mL HCl Volume of HCl consumed by W g soil 1000 mL 1M NaOH

= Wg º V1 mL NaOH (M) º 3V1 mL NaOH (M) º V2 mL NaOH (M) = (3 V1 –V2) = V mL NaOH (M) º 36.45 g HCl

36.45 ´ V ´ M g HCl 1000 According to the reaction CaCO3 + 2HCl ® CaCl2 + H2O + CO2

VmL NaOH(M) º

2 ´ 36.45 g HCl º 100 g CaCO3 36.45 ´ V ´ M 100 ´ 36.45 ´ V ´ M g HCl º 1000 2 ´ 36.45 ´ 1000

=

VM g CaCO3 20

Hence, the percentage of CaCO3 in the soil sample = =

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V M ´ 100 20 ´ W 5 VM W

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46 1

Gypsum Requirement

PRINCIPLE The physical properties (such as permeability of soil to water) deteriorate in saline and alkaline soils. When such a soil, containing excessive percentage of sodium (i.e., degree of alkalisation), is shaken with saturated gypsum (CaSO4.2H2O) solution, calcium is exchanged for sodium. The loss of calcium from the saturated solution of gypsum is called the gypsum requirement of the soil. In this experiment, the soil is treated with a known excess of saturated gypsum solution and the unreacted amount of gypsum is determined by EDTA titration. Therefore, gypsum requirement is the equivalence of CaSO4.2H2O or sulphur that must be added to reclaim the soil by displacing the exchangeable sodium. CaSO4.2H2O + 2Na+ ® Na2SO4.2H2O + Ca2+ In this experiment, “normality” is used, since in soil agricultural science laboratories, the gypsum requirement is always estimated in terms of milliequivalents (meq) of gypsum.

REAGENTS Saturated gypsum solution: Shake 2.5 g CaSO4.2H2O (AR) with 500 mL distilled water for 15 minutes and filter the solution. Standard calcium chloride solution (0.01 N): Dissolve 0.25 g dry CaCO3 (AR) in a minimum quantity of 0.2 N HCl.Boil the solution gently to remove CO2. Make up the solution to 500 mL with distilled water. EDTA solution (approximately 0.01N.): Dissolve 0.465 g EDTA and 0.02 g MgCl 2 (AR) in distilled water and make up the solution to 250 mL. Determine the exact normality of this solution, by titrating it with standard calcium chloride solution. Buffer of pH 10: Dissolve 35 g NH4Cl (AR) in 285 mL ammonia (specific gravity 0.91) and make up the solution to 500 mL with distilled water. This is the NH4ClNH4OH buffer having a pH 10.

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Eriochrome black T indicator: Refer experiment 15. Mn, Ni, Co and Cu interfere during the titration. Therefore, use 4% hydroxylamine hydrochloride (to keep Mn in divalent form) and 0.2% sodium cyanide (to mask Cu, Co or Ni) along with the indicator. These are the masking reagents.

PROCEDURE Accurately weigh about 2g soil sample and shake it thoroughly with exactly 50 mL saturated gypsum solution in a clean 250 mL conical flask. After about half an hour, filter the solution through dry Whatman No. 1 filter paper. Discard the first few mL of the filtrate and collect the rest. Then pipette 10 mL filtrate in a conical flask, add 1 mL buffer, 2 drops of indicator and 2 drops each of the masking reagents and titrate with standard EDTA solution. The end point is a colour change from wine red to blue. Simultaneously, perform a blank titration with 10 mL saturated gypsum solution using the same amounts of buffer and indicator as with the test filtrate.

CALCULATION Mass of the soil = Wg Volume of EDTA solution for 10 mL CaCl2 solution = V1 mL Strength of standard CaCl2 solution = 0.01 N Strength of EDTA solution =

10 ´ 0.01 = ‘a’ N V1

Volume of EDTA for 10 mL gypsum solution = V2 mL Volume of EDTA for 10 mL test solution (filtrate) = V3 mL Volume of gypsum required for soil = (V2 – V3) mL of ‘a’ N EDTA Concentration of gypsum required for soil solution =

(V2 – V3 )a 10

= ‘b’ meq L–1 For 50 mL soil solution, the gypsum required = b ´

50 1000

= 0.05b meq Since 50 mL solution contains Wg soil, the amount of gypsum needed for Wg soil is also 0.05b meq.

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For 100 g soil, the amount of gypsumü 0.05 b ´ 100 ý = required þ W 5b meq W 1 meq gypsum per 100 g soil º 1.7 metric tons gypsum per acre (30 cm soil depth) Hence, the amount of gypsum required per acre soil =

=

5b ´ 1.7 metric tons W

5b ´ 1000 ´ 2.471 kg/hectare W (1 metric ton = 1000 kg and 1 hectare = 2.471 acre) =

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Lime Requirement

PRINCIPLE A slightly acidic condition of the soil is always favourable for crop growth but high acidity has an adverse effect. Such acidic soils therefore need the application of lime to reduce the acidity and to improve productivity. The quantity of lime required varies according to the nature of soil. Lime requirement of a soil is defined as the CaCO3 equivalence of exchangeable hydrogen. Liming materials include limestone, burned lime, dolomitic limestone, slaked lime, etc. All these contain CaCO3 as the chief constituent and small amounts of other constituents. Therefore, only the neutralizing equivalence of any liming material, i.e., “CaCO3 equivalent”, forms the important factor in the determination of “lime requirement” of the soil. As it is determined in terms of milliequivalents of exchangeable hydrogen, the concentrations are expressed in “normality”. In this experiment, the amount of exchangeable hydrogen in the soil is determined by pH measurement. The percentage of CaCO3 present in the liming material, which is termed as the percentage CaCO3 equivalence of that liming material is determined by titration technique. Using these two, the quantity of the liming material to be used for a particular soil can be estimated. CaCO3 + 2HCl ® CaCl2 + H2O + CO2 (in liming material) HCl + NaOH ® NaCl + H2O (excess)

REAGENTS Buffer mixture (for extracting hydrogen): Dissolve 4g p-nitrophenol, 20 g calcium acetate and 0.3 g magnesium oxide in 500 mL distilled water. Adjust the pH to 7.0 with HCl or MgO as needed. Standard hydrochloric acid solution (1N): Add 80 mL conc. HCl to water and make upto one litre Standard sodium hydroxide solution (1N): Dissolve 40g pellets in water and make upto one litre

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Phenolphthalein indicator Liming material: Limestone, burned lime or slaked lime.

PROCEDURE Part I : Weigh 5 g dry soil sample accurately and place it in a 100 mL dry beaker. Add to it 5 mL distilled water and 10 mL buffer solution. Stir the solution well and allow it to stand for 45 minutes. Measure the pH of the suspension using pH meter. If the pH is below 6.0, repeat the test with half the amount of soil. Part II : Select any liming material such as limestone, burned lime or slaked lime. Dry and grind it. Weigh accurately about 0.5 g liming material and place it in a 250 mL conical flask. Then add 20 mL HCl(1N). Swirl the flask well and heat to boiling. Keep it on a steam or water bath for 20 minutes to complete the reaction. Now add 50 mL distilled water and boil the solution for exactly 1 minute over a gentle flame. Cool to room temperature. Back-titrate the solution with 1N NaOH solution using phenolphthalein indicator. Perform a duplicate.

CALCULATION Part I : There is a perfect linearity between the pH and the milliequivalent (meq) of exchangeable hydrogen. Therefore, in the pH range 7.0-6.0, each 0.1 unit change in pH = 1 meq. of exchangeable hydrogen per 100 g soil. If the pH is below 6.0, 0.1 unit change in pH = 2 meq exchangeable hydrogen per 100 g soil (from soil science literature) Calculate the meq exchangeable hydrogen per 100 g soil corresponding to the pH measured. Let it be ‘a’. Part II : Mass of the liming material = Wg Strength of HCl = 1 N Strength of NaOH = 1 N Volume of NaOH for the neutralization of 20 mL HCl = 20 mL Volume of NaOH for neutralizing excess HCl (in test solution) = V mL Volume of HCl consumed by CaCO3 in liming material = (20 – V) mL According to the equation CaCO3 + 2HCl ® CaCl2 + H2O + CO2 1000 mL of 2N HCl º 100 g CaCO3 (20 – V) mL of 1N HCl º =

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(20 – V) ´ 1 ´ 100 1000 ´ 2 (20 – V) g CaCO3 20

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Lime Requirement

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So, the percentage of CaCO3 present in the liming material (or % CaCO3 equivalence) =

(20 – V) 100 ´ 20 W

=b 1 meq exchangeable hydrogen per 100 g soil = 1000 pounds CaCO3 per 2 million pounds soil (from soil science literature) ‘a’ meq per 100 g soil = 1000 ´ ‘a’ pounds CaCO3 per 2 million pounds soil Quantity of lime (pure CaCO3) required by 2 million pounds soil = 1000 x ‘a’ pounds ‘b’ pounds CaCO3 is present in 100 pounds liming material Hence, 1000 x ‘a’ pounds CaCO3 will be present in

100 ´ 1000 ´ ‘a’ pounds b

liming material So, quantity of liming material needed by 2 million pounds soil =

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100 ´ 1000 ´ ‘a’ pounds b

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48 1

Alkalinity

Alkalinity is an important parameter for plant growth. The soil should neither be too acidic nor be too alkaline. So, its determination gives information about the quality of the soil and its suitability for plant growth. Principle : Refer to experiment 9. Reagents : Refer to experiment 9.

PROCEDURE Weigh 20 g soil sample and extract it with 250 mL Morgan’s extracting solution (Experiment 44). Make up the solution to 500 mL with the extracting solution. Pipette 20 mL soil extract in a conical flask and proceed as given in experiment 9.

CALCULATION Refer to experiment 9.

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Organic Carbon

PRINCIPLE Carbon occurs in four forms in soil: (1) mineral forms like carbonates and bicarbonates, (2) elemental forms like coal, charcoal, graphite, etc., (3) resistant organic residues of animals, plants and micro-organisms; these are collectively termed as “humus”, and (4) freshly added organic residues subject to rapid decomposition in soils. These release their nutrient elements to the crop. Among the various forms, the fourth form is easily oxidizable and so is called oxidizable organic carbon. This organic carbon gives an indication of the fertility status of the soil. Organic matter can be determined by three methods: a. total organic matter determination by wet oxidation and weight loss. b. organic carbon determination as CO2 either by wet oxidation with chromic acid or by dry combustion. c. oxidizable organic matter determination. In soil fertility, the easily oxidizable organic matter is important. Presence of hydrated salts and carbonates in the soil influences the weight change and so the first method is susceptible to errors. In the second method, apart from oxidizable carbon, carbon in carbonate and carbon as graphite, charcoal or coal are included and so a correction must be made. The third method excludes all other forms of carbon and only the oxidizable organic carbon is involved. Therefore, this method has the advantage over the other methods. Estimation of the percentage of organic carbon can be done either by titration or by colorimetric method.

I. TITRATION METHOD In this experiment, the organic carbon is oxidized by potassium dichromate in presence of concentrated sulphuric acid. The heat of dilution of sulphuric acid is used for this purpose. The dichromate left unreacted is back-titrated with ferrous ammonium sulphate solution. Thus, the amount of dichromate required to oxidize the soil carbon is determined. The oxidation can be expressed as:

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4Cr6+ + 3C0 ® 3C4+ + 4Cr3+ Generally, the H and O contents of organic matter are not considered in stoichiometric equations. But the organic H requires the same amount of dichromic acid as organic C requires. For example, the reaction of CH4 with dichromic acid can be represented with organic carbon and organic hydrogen separately as below. 2H2Cr2O7 + 3C0 + 6H2SO4 ® 2Cr2 (SO4)3 + 3CO2 + 8H2O and 2H2Cr2O7 + 3(4H0) + 6H2SO4 ® 2Cr2(SO4)3 + 14H2O But due to the presence of organic oxygen, carbon in -COOH group does not need dichromate for its oxidation (i.e., RCOOH ® RH + CO2) and thus lowers the amount of dichromate. These opposite effects due to organic H and O balance each other and therefore no complication will arise.

REAGENTS Potassium dichromate solution (0.16 M): Dissolve 23.52 g K2Cr2O7 (AR) in water and make upto 500 mL. Ferrous ammonium sulphate solution (0.5 M): Mix 10 mL conc. H2SO4 with 490 mL water. Add to this mixture, 98 g hydrated crystalline ferrous ammonium sulphate. This is more stable than FeSO4. Diphenylamine indicator: Dissolve 0.25 g diphenylamine in a mixture of 10 mL water and 50 mL conc. H2SO4. Sulphuric acid : Concentrated, Specific gravity 1.84 Phosphoric acid: 1:1 with water. Solid sodium fluoride

PROCEDURE Take about 5 g soil sample and remove from it, the sulphides and chlorides, if present. For the removal of sulphide, add 2 M sulphuric acid and warm until no evolution of H2S is observed. Then wash the soil with distilled water. Dry it in an oven at 105-110° C. For removing the chloride, wash the soil with distilled water several times until no turbidity appears when a drop of wash-water is tested with AgNO3. Dry it again in the oven. Weigh accurately about 0.5 g sulphide-and chloride-free sample of the soil and place it in a dry 250 mL conical flask. Pipette 10 mL K2Cr2O7 solution (0.16 M). Swirl the conical flask gently to mix the reagent well with the soil. Now add carefully through the sides of the conical flask, 20 mL concentrated sulphuric acid in one lot and swirl again. Leave the flask as such for about 40 minutes to allow the oxidation to proceed. Then dilute the solution (to approximately 200 mL) with distilled water till it becomes light yellow. Add 5 mL syrupy phosphoric acid (1:1), 1 g solid NaF and 2 drops diphenylamine indicator. Titrate the contents

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113

with ferrous ammonium sulphate solution till the colour changes from blue-violet to green. Perform a duplicate under identical conditions. Simultaneously, run a blank without the soil. Pipette 10 mL dichromate solution into a clean conical flask. Add 20 mL dilute sulphuric acid, 5 mL phosphoric acid, 1 g NaF and 2 drops indicator. Titrate this mixture with ferrous ammonium sulphate solution till the end point (blue-violet to green) is reached. Repeat the titration to obtain concordant values. Phosphoric acid and sodium fluoride help to complex the Fe3+ ions produced by the oxidation of Fe 2+ which would, otherwise, cause premature end point.

CALCULATION Standardization of ferrous ammonium sulphate (FAS) solution: Titre value = V mL K2Cr2O7 (0.16 M) 10 mL FAS (M) º V mL K2Cr2O7 (0.16 M) 1000 mL 1M K2Cr2O7 = 1 mole

V ´ 0.16 moles 1000 1 mole K2Cr2O7 º 6 moles Fe2+

V mL 0.16 M K2Cr2O7 =

V ´ 0.16 6 ´ 0.16 ´ V moles K2Cr2O7 º 1000 1000 =

0.96 ´ V moles Fe2+ (in 10 mL FAS) 1000

0.96 ´ V 10 Determination of percentage of organic carbon:

Hence, molarity of FAS (M1) =

Mass of the soil sample = Wg Volume of K2Cr2O7 required for oxidation of organic carbon = ( Vb – Vs) mL FAS (M1) Vb = volume of FAS from blank titration Vs = volume of FAS from titration with soil Number of moles of K2Cr2O7 required for oxidation =

(Vb – Vs ) ´ M1 ´ 1 1000 ´ 6

According to the equation, 2 moles K2Cr2O7 º 36 g carbon

(Vb – Vs ) ´ M1 36 ´ (Vb – Vs ) M1 moles K2Cr2O7 º 1000 2 ´ 1000 ´ 6

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3 ´ (Vb – Vs ) M1 g carbon 1000 3 ´ (Vb – Vs ) ´ M1 ´ 100 Hence, percentage of organic carbon = 1000 ´ W =

=

3 ´ (Vb – Vs ) ´ M1 10 ´ W

II. COLORIMETRIC METHOD Principle In this method, the organic carbon present in the soil sample reduces K2Cr2O7 to green chromium sulphate in presence of sulphuric acid. The amount of chromium sulphate formed is a measure of the organic carbon in the soil. The intensity of the green colour and hence the amount of chromium sulphate are determined colorimetrically.

Reagents K2Cr2O7 (AR) solution: 0.16M Sulphuric acid : Concentrated Anhydrous sucrose (AR)

PROCEDURE Weigh anhydrous sucrose of 5, 10, 15, 20 and 25 mg accurately and place in separate 100 mL dry standard flasks. Add first 10 mL K2Cr2O7 solution and then 20 mL conc. H2SO4 to each of the flasks. Swirl the flasks gently and keep them aside for 40 minutes. Make the volumes upto the mark with distilled water and shake well. Adjust the colorimeter or the spectrophotometer reading (at 660 nm) to zero with the blank solution (without sucrose). Measure the absorbance values for the standard solutions. Draw a calibration curve by plotting the concentration of carbon in sucrose solutions (mg L–1) and the absorbance. Weigh accurately about 0.5 g soil and transfer it to a dry 100 mL standard flask. To this add 10 mL K2Cr2O7 solution , swirl for a while and then add 20 mL conc. H2SO4. Again swirl gently and keep aside for 40 minutes. Make the volume upto the mark with distilled water and shake well. Centrifuge the solution and read the absorbance of the centrifugate at 660 nm. Note the concentration of carbon (mg L–1) corresponding to the absorbance of the soil solution from the calibration curve.

CALCULATION Percentage of organic carbon in the given soil sample 100 where W is the mass of the soil taken. = (Concentration from graph) ´ W

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Total Nitrogen

PRINCIPLE The estimation of the amount of total nitogen in the soil is essential to evaluate the fertility of the soil. There are two methods for this purpose. (a) Duma’s method, and (b) Kjeldahl’s method. In the former, the soil nitrogen is converted into free nitrogen and the volume of gas produced is measured. In the latter procedure, nitrogen in the soil is first converted into ammonia and then to ammonium sulphate. But Kjeldahl’s procedure is preferred to Duma’s due to the following reasons: i. Appraratus needed is simpler and controlled conditions are not required. ii. Since the soil required to be used is very small in Duma’s method, there is a possibility of large error in sampling. In Kjeldahl’s method, a sample containing nitrogen is heated with conc. H2SO4. Nitrogen in the sample is converted to (NH4)2SO4. Ammonia is liberated by treating the resultant liquid with excess alkali. Ammonia thus produced is absorbed in excess standard acid. The excess acid remaining after neutralizing NH3 is back-titrated with standard alkali. From this, the amount of NH3 and hence nitrogen in the soil is determined. N (in soil)

+ H2SO4 ® (NH4)2SO 4 (conc.)

(1)

2(NH4)2SO4 + 4NaOH ® 2Na2SO4 + 4H2O + 4NH3

(2)

NH3 + H2O ® NH4OH

(3)

2NH4OH + H2SO4 ® (NH4)2SO4 + 2H2O

(4)

Excess H2SO4 + 2NaOH ® Na2SO4 + 2H2O (titration)

(5)

During H2SO4 digestion, nitrates are usually lost by volatilization of HNO3. Therefore, salicylic acid is added to the reaction mixture so that the nitrate is made to combine with salicylic acid in presence of conc. H2SO4. Also, thiosulphate is added to reduce any interfering mercuric salt, if present. A catalyst mixture of

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anhydrous sodium sulphate and copper sulphate is also used. Here, Na2SO4 raises the boiling temperature of H2SO4.

REAGENTS Sulphuric acid: Concentrated, Specific gravity 1.84. Salicylic acid Sodiumthiosulphate Sodium sulphate Anhydrous copper sulphate Sodium hydroxide solution. (40%) : Dissolve 40 g NaOH pellets in 100 mL distilled water. Sulphuric acid: 0.05 M solution Sodium hydroxide solution: 0.1M for titration Phenolphthalein indicator.

PROCEDURE Weigh accurately about 2 g soil sample and place it in a 500 mL Kjeldahl’s flask. Add 1 g salicylic acid to 25 mL concentrated sulphuric acid. Slowly add this solution to the soil and allow the mixture to stand for 40 minutes. Now add 5 g Na2S2O3 and leave the contents aside for another 40 minutes. Add to this mixture, the catalyst mixture containing 5 g Na2SO4 and 1 g CuSO4. Stopper the flask loosely. Heat it first gently to avoid frothing and then strongly. A brown colour is produced. When this colour disappears after about 35 minutes, remove the flask and digest over a hot water bath for about an hour. Cool and cautiously add 100 mL distilled water. Then transfer 100 mL of 40% NaOH to the flask and put into it a few porcelain pieces. Connect the flask immediately to the distillation unit. Keep exactly 100 mL of 0.05 M H2SO4 in the receiver. Collect the ammonia evolved in the H2SO4. Take care to ensure that NH3 bubbles slowly under the surface of H2SO4. Continue distillation for 1 hour. Test the distillate with litmus paper to confirm that no more NH3 passes over. Swirl the receiver well and pipette 10 mL solution from it. Titrate with 0.1 M NaOH solution using phenolphthalein indicator. Carry out the blank titration with 10 mL H2SO4 (0.05 M) alone. Perform a duplicate experiment under identical conditions.

CALCULATION Mass of the soil sample 10 mL 0.05 M H2SO4 (blank) 100 mL 0.05 M H2SO4 Molarity of H2SO4

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= Wg. º v1 mL M1 NaOH º 10 v1 mL M1 NaOH = V1 mL M1 NaOH = 0.05

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Total Nitrogen

Molarity of NaOH =

%

100 ´ 0.05 ´ 2 = M1 V1

From the titration with excess H2SO4, 10 mL solution º v2 mL M1 NaOH 100 mL solution º 10 v2 mL M1 NaOH = V2 mL M1 NaOH Volume of H2SO4 required for complete neutralization of NH3 = (V1 – V2) = V mL of M1 NaOH 1000 mL 1M NaOH º 49 g H2SO4

49 ´ V ´ M1 g H2SO4 1000 From the chemical equation given in the principle, V mL M1 NaOH º

98 g H2SO4 º 28 g nitrogen (from eq.4)

49 ´ V ´ M1 28 ´ 49 ´ V ´ M1 g H2SO4 º = 0.014 ´ V ´ M1 g nitrogen 1000 98 ´ 1000 The percentage of nitrogen in the soil sample =

0.014 ´ V ´ M1 ´ 100 W

Note 1. NH3 can be absorbed in boric acid forming ammonium borate. This is titrated against 0.01 M H2SO4 using methyl orange indicator and is expressed in terms of N2.

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Available Nitrogen

PRINCIPLE Most of the nitrogen in the soil exists in organic form and it becomes available to plants by biological transformations. Several methods are available for the estimation of nitrogen supplying capacity of soils. Some of these are: (1) the determination of organic matter or total nitrogen (2) the evaluation of rate of mineralisation or nitrification, and (3) the estimation of the amount of NH3 evolved by KMnO4 oxidation. Apart from these, hydrolyses by acids and bases are also utilised. For alkaline hydrolysis, NaOH, Ba(OH)2, Ca(OH)2, mixture of NaHCO3 and Na2 CO3 or MgO can be used. H2SO4 can be used for acid hydrolysis. Among these Ca(OH)2 gives the most satisfactory result. Therefore, in this experiment, the amount of available nitrogen in the soil is estimated by hydrolysis with Ca(OH)2. The soil is treated with Ca(OH)2 when the organic nitrogen present in it undergoes hydrolysis. As a result, NH3 is released. It is collected in a known excess of sulphuric acid solution. NH3 is neutralized completely and the unreacted H2SO4 is back-titrated with standard alkali. Hence, the quantity of NH3 produced and of nitrogen can be determined.

REAGENTS Calcium hydroxide (AR) Sulphuric acid: Accurately 0.05 M solution Sodium hydroxide for titration: Approximately 0.1M solution Phenolphthalein indicator

PROCEDURE Weigh accurately about 2 g soil sample and place it in a 500 mL Kjeldahl’s flask. Add 0.5 g Ca(OH)2, 250 mL distilled water and a few porcelain pieces. Connect the flask to the distillation set up. Place exactly 100 mL H2SO4 (0.05 M ) in the receiver. Use heating mantle for steady distillation. Carry out the hydrolysis for 45

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Available Nitrogen

'

minutes from the time of boiling. Remove the receiver, swirl well and pipette 10 mL solution into a clean and dry conical flask. Titrate the excess acid with approximately 0.1 M NaOH using phenolphthalein indicator. Titrate 10 mL sulphuric acid alone for blank value. Perform a duplicate under identical conditions.

CALCULATION Same as in the experiment ‘Total nitrogen’ (Experiment 50).

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Ammonium Nitrogen

PRINCIPLE Nessler’s reagent, which is an alkaline solution of HgI2 in potassium iodide, is used for the determination of ammonium ions. In this method, Nessler’s reagent is added to a solution of ammonium salt when a reaction occurs between the liberated ammonia and the reagent. The orange-brown product formed has the maximum absorption at 460 nm. By measuring the absorbance at 460 nm, the concentration of ammonium nitrogen is determined.

REAGENTS Extracting solution: Refer to experiment 44. Nessler’s reagent: Refer to experiment 44. Standard solution of ammonium sulphate: Weigh accurately 0.3667 g ammonium sulphate (AR) and dissolve it in one litre of 1:1 extracting solution. This is 100 mg L–1 NH4+ solution. Dilute 25 mL of this solution to 250 mL with 1:1 extracting solution to make a 10 mg L–1(ppm) NH4+ solution.

PROCEDURE Pipette a series of 0.5, 1.0, 1.,5, 2.0, 2.5 and 3.0 mL standard NH4+ solution (10 mg L–1ppm) into different 50 mL standard flasks. Add 1 mL Nessler’s reagent to each flask and dilute upto the mark with 1:1 extracting solution. Shake well and allow the solutions to stand for 10 minutes. After 10 minutes but before 20 minutes measure the absorbance of each solution at 460 nm. Draw a calibration curve with absorbance and concentration in mg L–1. Weigh accurately 5 g soil sample and shake it well with 25 mL extracting solution for 45 minutes. Filter the solution through Whatman No. 1 filter paper. Transfer 2 mL soil extract into a 50 mL standard flask, add 1 mL Nessler’s reagent and make up the volume to 50 mL with 1:1 extracting solution. Measure the absorbance at 460 nm after 10 minutes of mixing but before 20 minutes. Read the corresponding concentration from the calibration curve.

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CALCULATION Refer “Appendix 6”.

(concentration from graph) 50 25 ´ ´ 1000 2 5 –1 = a mg g

Ammonium content of the soil =

Amount of ammonium nitrogen = a ´

14 mg g–1 18

Note 1. If a precipitate or turbidity appears with the soil sample or if the absorbance goes beyond limit, appropriate dilution must be carried out.

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Nitrate Nitrogen

PRINCIPLE The nitrate in the soil is extracted first with Morgan’s extracting solution. The extract is then treated with phenol disulphonic acid in concentrated sulphuric acid. Nitrate is converted into nitric acid which is utilised for the nitration of phenol disulphonic acid. The product, a nitro compound, gives an intense yellow colour in presence of an alkali. The intensity of the colour is measured from which the concentration of nitrate can be evaluated. As there is a possibility of volatility loss of nitric acid on heating, the reaction must be carried out in cold.

REAGENTS Calcium carbonate powder Extracting solution: Refer to experiment 44. Phenol disulphonic acid: Transfer 12.5 g pure phenol to a dry beaker. Add 75 mL concentrated sulphuric acid (AR, nitrate-free) followed by 37.5 mL nitrate-free AR fuming sulphuric acid very carefully. Mix the contents well and keep the beaker on a hot water bath for at least 2 hours. Cool and store in an amber coloured bottle. Ammonia solution (1:1): Mix. 100 mL liquor ammonia (specific gravity 0.88) and 100 mL water. Standard nitrate solution : Dissolve 0.1805 g KNO3 (AR)in 250 mL extracting solution. This is a 100 mg L–1 (ppm) nitrate solution.

PROCEDURE Weigh accurately 5 g soil sample and shake it vigorously with 25 mL extracting solution for 45 minutes. Filter the solution through Whatman No. 1 filter paper. The filtrate is the soil extract. Take 2 mL soil extract in a porcelain dish and to this add 1 g CaCO3. Evaporate this solution to dryness on a steam bath. Cool and add 2.5 mL phenol disulphonic

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acid. Swirl the dish in order to mix the contents well. After 15 minutes, add 10 mL distilled water and transfer the solution to a 100 mL standard flask. Now add 1:1 ammonia solution in drops with constant swirling till a yellow colour is produced. Add 2.5 mL more and make the solution upto the mark with distilled water. After 10 minutes, measure the absorbance at 420 nm. Pipette 0.5, 1.0, 1.5, 2.5 and 3.0 mL aliquots of the standard nitrate solution into small beakers or porcelain dishes. Evaporate the solutions to dryness and proceed as with the soil extract to produce the yellow colour. Measure the absorbance of the solutions at 420 nm. Draw the calibration curve with absorbance and concentration in mg L–1 (ppm). Note from the calibration curve, the concentration of nitrate in the soil extract corresponding to its absorbance.

CALCULATION Refer to “Appendix 6” Nitrate content of the soil =

(concentration from graph) ´ 100 ´ 25 = a mg g –1 1000 ´ 2 ´ 5

Amount of nitrate nitrogen = a ´

14 mg g–1 62

Notes 1. Nitrate nitrogen may be estimated by Brucine method also. Refer to experiment 17. 2. It may also be reduced quantitatively and estimated as nitrite. Refer to experiments 16 and 54.

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Nitrite Nitrogen

PRINCIPLE Nitrite is usually determined by a method based on diazotisation reaction. In this experiment, sulphanilic acid is diazotised with nitrite ion in presence of an acid. The diazotised sulphanilic acid is then coupled with a-naphthylamine. A purplered azo dye is formed which absorbs maximum at 520 nm. By determining its absorbance at this wavelength, the concentration of nitrite in the soil can be evaluated.

REAGENTS

ü ý a-naphthylamine solution þ Extracting solution

Sulphanilic acid solution

Refer to experiment 44.

Standard potassium nitrite solution: Dissolve 0.1848 g KNO2 (AR) in one litre extracting solution. This is a standard nitrite solution of 100 mg L–1 (ppm). Prepare a 10 mg L–1 solution by diluting 10 mL solution to 100 mL with the extracting solution.

PROCEDURE Make a series of standard solutions of nitrite by pipetting 0.5, 1.0, 1.5, 2.0, 2.5 and 3 mL of 10 mg L–1 solution into various 50 mL standard flasks. Add to each of the flasks 1 mL sulphanilic acid and 1 mL a-naphthylamine solutions. Make up the solutions to 50 mL with the extracting solution. Shake well and keep the flasks aside. After 10 minutes, measure the absorbance of the purple coloured solutions at 520 nm. Draw a calibration curve with absorbance versus concentration in mg L–1. Shake thoroughly about 5 g (accurately weighed) soil sample with 25 mL extracting solution for about 45 minutes. Filter the solution through Whatman No. 1 filter paper and treat 2 mL of this extract exactly the same way as the standards to get a purple solution. Read the absorbance of this solution at 520 nm. Note the corresponding concentration of nitrite from the calibration curve.

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CALCULATION Refer to “Appendix 6” Nitrite content =

(concentration from graph) 50 25 ´ ´ = a mg g–1 1000 2 W

Amount of nitrite nitrogen = a ´

14 mg g–1 46

Note 1. Nitrite can also be estimated using sulphanilamide. Refer to experiment 16.

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55 1

Available Phosphorus

PRINCIPLE In soil, phosphorus occurs as orthophosphate. Although this is in various combinations, the entire amount would not be available to plants. Only a small fraction is available and this determines the P-fertility status of the soil. Many reagents can be used to extract phosphorus from the soil but sodium bicarbonate solution (0.5 M) of pH 8.5 is the most suitable for extracting reactive forms of P (phosphates of iron, aluminium and ammonium) in the soil. In this experiment, the soil extract is treated with acidic sodium molybdate solution when phosphate and molybdate ions condense to give phosphomolybdate or molybdophosphoric acid. This product on selective reduction with hydrazine sulphate produces an intense blue complex (uncertain composition). The blue complex exhibits a maximum absorption at 620 nm. Other oxidizing and reducing agents, and ions such as silicate, arsenate, copper, antimony and lead must be absent, as these interfere in P estimation. + + . PO3– 4 + 12Na 2 MoO4 + 24H ® Na3 PO4 12 MoO 3 + 21Na + 12H2O . . Na3PO4 12MoO3 + N2H4 H2SO4 ® Molybdenum blue.

REAGENTS Sodium molybdate solution : Dissolve 6.25 g sodium molybdate (AR) in 5 M sulphuric acid. Make up the volume to 250 mL with 5 M sulphuric acid. Hydrazine sulphate solution: Dissolve 0.75 g hydrazine sulphate (AR) in doubledistilled water and make upto 500 mL. Extraction solution: Dissolve 21 g NaHCO3 (phosphate-free) in 500 mL distilled water. This is 0.5 M solution. Adjust the pH to 8.5 using 10% NaOH solution. Standard phosphorus solution: Weigh accurately 0.1098 g dry potassium dihydrogen phosphate (KH2PO4) (AR), dissolve it in distilled water and make up the solution to 250 mL. This is a 100 mg L–1 (ppm) phosphorus solution.

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Available Phosphorus

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PROCEDURE Accurately weigh about 1 g soil sample and transfer it to a dry 100 mL conical flask. Add a little P-free activated carbon and exactly 25 mL extracting solution. Shake the contents of the flask well for about an hour and filter through dry Whatman No.1 filter paper. The filtrate is the soil extract. Pipette separately 0.5, 1, 2, 4 and 5 mL standard phosphorus (100 mg L–1) solution into a series of 25 mL standard flasks. To each flask, add 2.5 mL molybdate solution and 1 mL hydrazine sulphate solution. Make upto the mark with the extracting solution. Shake the contents thoroughly. Warm the solutions on a boiling water bath for exactly 10 minutes and cool . Measure the absorbance of these solutions at 620 nm. Draw the calibration curve with absorbance versus concentration of P in mg L–1 (ppm). Repeat this procedure with 2 mL aliquot of the soil extract. Using its absorbance, read the concentration of P in the soil extract from the calibration curve.

CALCULATION Refer to “Appendix 6” Mass of the soil = W g. Concentration of P in soil extract (from graph) = a mg L–1 Phosphorus available to the soil =

a 25 25 ´ ´ mg g–1 1000 2 W

Note 1. Phosphorus may also be estimated using ammonium molybdate and stannous chloride. Refer experiment 26.

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56 1

Available Sulphur

PRINCIPLE Sulphur is an essential nutrient for plant growth and some crops need such high amount of sulphur that it exceeds the amount of phosphorus needed. Much of the sulphur is present in organic form in the soil. Therefore, organic matter is an important source of available sulphur. The soil sample is fused with sodium carbonate and sodium nitrate. This converts all the sulphur of the soil into sulphate. The sulphate is then precipitated as barium sulphate and estimated gravimetrically.

REAGENTS Sodium carbonate (AR) Sodium nitrate (AR) Hydrochloric acid: Concentrated Barium chloride: 10% solution

PROCEDURE Take in a platinum crucible a known mass (1-5 g depending on the sulphur content) of the dry soil sample. Add to it, sodium carbonate (five times the mass of the soil) and 1 g sodium nitrate. Fuse this mixture in an electrothermal bunsen. This may take about 3 to 4 hours. Then add a small amount of water and boil gently for about 15 minutes. Cool and filter the solution. Collect the filtrate and washings quantitatively. Make up the solution to 250 mL. Pipette 50 mL of the made-up solution into a beaker and neutralize the unused sodium carbonate in it by carefully adding hydrochloric acid. Add a slight excess of hydrochloric acid followed by 25 mL of 10% BaCl2 solution. Test for completion of precipitation. Filter the precipitate of BaSO4 formed using a Whatman No. 42 filter paper and wash with water to free it from C1–. Dry the precipitate over an iron cone. Separate the dried precipitate as much as possible from the filter paper and collect

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Available Sulphur

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it on a glazed paper. Incinerate the filter paper carefully in a weighed silica crucible. Moisten it with a drop of conc. H2SO4. Heat it slowly until the fume formation stops. Then transfer the precipitate quantitatively from the glazed paper to the crucible and heat the crucible strongly for about 30 minutes. Cool it in a desiccator and weigh. Repeat heating, cooling and weighing until a constant mass is recorded.

CALCULATION Mass of soil taken = w1 g Mass of BaSO4 formed = w2 g Available sulphur in the soil =

w2 ´ 32 ´ 106 mg L–1 (ppm) w1 ´ 233.4

Note 1. The use of sintered crucible may be avoided as it becomes very difficult to remove the precipitate after the experiment and the porous disc gets damaged.

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57 1

Silica

PRINCIPLE Silica in the soil is determined usually by colorimetry. The soil is first fused with anhydrous sodium carbonate and the silica is brought into solution. This solution is then treated with a solution of ammonium molybdate in presence of an acid when an intense yellow complex, molybdosilicic acid is produced. This acid, (H2[SiMo12O40]), is reduced with a mixture of sodium hydrogen sulphite and 1-amino-2-naphthol-4-sulphonic acid to give molybdenum blue. This has uncertain composition. Similar colouration is produced by phosphate also and so phosphates are eliminated as ammonium magnesium phosphate precipitate in acetic acid. Alternatively, these can be masked by adding ammonium citrate. Freshly distilled (in an all-Pyrex apparatus) water, stored in polythene bottles is used.

REAGENTS Ammonium molybdate Solution: Dissolve 4 g AR ammonium molybdate in a little water. Add to it 4 mL concentrated sulphuric acid. Make up this solution to 50 mL. Tartaric acid solution: (10%) Dissolve 10 g substance in 100 mL water. This is to maintain the pH at 4.5-5.0, necessary for the reduction. Reducing agent: Prepare a solution of 5 g NaHSO3 in 30 mL water. Make another solution containing 0.4 g anhydrous NaHSO3 and 0.08 g 1-amino-2naphthol-4-sulphonic acid in 10 mL water. Mix the two solutions and make upto 50 mL. Standard solution of silica: Weigh accurately 0.107 g pure, dry silica and fuse with 1 g Na2CO3 (AR) in a platinum crucible. Cool the melt and dissolve in a small quantity of double-distilled water. Make up the solution to one litre. Keep this solution in a polythene bottle. This is a 100 mg L–1 (ppm) silica solution.

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Silica

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PROCEDURE Pipette the aliquots of 1, 2, 3, 4, 5 and 6 mL of 100 mg L–1 standard solution into a series of 50 mL standard flasks. To each flask, add 0.5 mL ammonium molybdate solution and mix. After 2 minutes, add 2.5 mL tartaric acid solution followed by 0.5 mL reducing agent. Dilute the solution upto the mark with doubledistilled water. Measure the absorbance of the solutions at 620 nm after 15 minutes. Draw the calibration curve with absorbance and concentration in mg L–1 . Dry the soil sample at 110° C for about an hour and weigh 0.2 g from it. Mix it with 2 g Na2CO3 (AR) in a platinum crucible and fuse till the melt is liquefied. Treat this liquid exactly the same way as the standards and measure the absorbance at 620 nm. From the calibration curve determine the concentration of silica in the given sample.

CALCULATION Mass of the soil

= 0.2 g

Silica content of the soil = (Concentration from graph) ´ =

Volume of solution Mass of soil

(Concentration from graph) 50 ´ mg g–1 1000 0.2

Notes 1. The pH should be 4.5 – 5.0. 2. Sodium molybdate may be used instead of ammonium molybdate. 3. Phosphate gives similar colouration and therefore must be removed by precipitation as ammonium magnesium phosphate in acetic acid solution or masked using ammonium citrate.

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58 1

Available Potassium

PRINCIPLE The term ‘available potassium’ incorporates both exchangeable and water soluble forms of the nutrient present in the soil which are available to the plants. This can be determined in the neutral normal ammonium acetate extract of the soil. The ammonium ion provides a sharp and rapid separation of potassium from the exchange complex while other cations bring about a gradual separation. Therefore, ammonium acetate is preferred to other extracting reagents. The estimation of potassium in the extract is usually carried out by flame photometry. Chemical methods may also be employed but are not found suitable for soil testing purpose. The non-exchangeable potassium also contributes to the potassium availability to crops. All forms of potassium can be extracted using strong inorganic acids such as H2SO4 (6N) and hot HCl (1N). This includes exchangeable, water soluble and non-exchangeable potassium. Then, total potassium is estimated by flame photometry. The difference between this value and the one obtained for exchangeable potassium gives the amount of non-exchangeable potassium in the soil . However, it was found that the results are not as quantitative as expected. The potassium content is first extracted from the soil using 1M ammonium acetate. This extract is aspirated into flame and then the reading is taken. Similarly, readings are taken for solutions of known potassium concentration and a calibration curve is drawn. The concentration of potassium in the soil extract is then read from the calibration curve. When a solution containing a metallic salt or a metallic compound is aspirated into a flame (acetylene/air, LPG, etc.), a vapour which contains metal atoms is formed. Some of these gaseous metal atoms are excited to higher energy levels by the energy provided by the flame. These levels are so high as to permit the emission of radiation characteristic of that metal, when the atoms return to the ground state. The intensity of radiation emitted will be directly proportional to the number of atoms in the flame, i.e., the concentration of the particular metal or element in the test solution. This is the basis of flame emission spectroscopy (FES) which was earlier referred to as flame photometry.

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REAGENTS Ammonium acetate solution (1 M): This is the extracting reagent. Prepare solutions of 2 M ammonium hydroxide and 2 M acetic acid. Mix these two solutions in the ratio of 1:1. Adjust the pH to 7.0 either with ammonia or with acetic acid. Standard solution of potassium: Dissolve 0.9545 g KCl (AR) in 500 mL extracting reagent in a standard flask; this is a 1000 mg L–1 K solution. Dilute suitably with extracting reagent to obtain solutions of 2, 4, 6, 8, 10 and 12 mg L–1 (ppm) K..

PROCEDURE Weigh accurately about 10 g soil sample, freshly taken from the field and add to it 50 mL extracting reagent in a conical flask provided with a stopper. Shake the mixture vigorously for about 20 minutes and filter it through a Whatman No. 42 filter paper. Set the flame photometer with proper sensitivity control, gas supply, etc. Then light the gas. Adjust the air supply to the required pressure. Regulate the gas supply (follow the instructions given in the manual of the instrument) and place the appropriate filter. Standardize the instrument with the extracting reagent for zero reading and with the standard solution of maximum concentration for reading 100. Then aspirate various solutions of known concentration (2,4,6,8,10 and 12 mg L–1 K) into the flame and note the flame photometer reading at each concentration. Plot the readings against concentration in mg L–1. This is the calibration curve. Check the reading periodically with standard solution for 100. Introduce the soil extract into the flame, note the reading and evaluate the concentration of K from the calibration curve.

CALCULATION Mass of the soil

= Wg.

Concentration of K in soil = Concentration from the graph ´ =

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Volume of extracting solution mg/L (ppm) Mass of soil

Concentration from graph 50 mg g-1 ´ 1000 W

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59 1

Sodium

PRINCIPLE The sodium content of a soil sample can also be determined by flame photometry, the principle of estimation being the same as that for determining the potassium content.

REAGENTS Ammonium acetate solution (1 M): Extracting solution. Refer to experiment 58 Standard solution of sodium: Dissolve 1.271 g NaCl (AR) in 500 mL extracting solution. This is a 1000 mg L–1 (ppm) sodium solution. To get a series of 2,4,6,8,10 and 12 mg L–1 solutions, dilute suitably the 1000 mg L–1 (ppm) solution using the extracting solution. PROCEDURE AND CALCULATION: Similar to those given for the determination of potassium.

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60 1

Calcium and Magnesium

PRINCIPLE The calcium and magnesium salts present in the soil (a known mass) are extracted with a solution of ammonium acetate. The extract is then titrated with EDTA using Eriochrome black T indicator. The extracting solution is 1 M ammonium acetate solution.

REAGENTS Refer to the experiments on “Hardness of water” (14 and 15)

PROCEDURE Extraction of soil—Refer to “Available Potassium”(Experiment 58) Determination of Ca and Mg - Refer to “Hardness of water” (Experiments 14 and 15)

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61 1

Micronutrient Elements

PRINCIPLE Elements such as Zn, Cu, Fe and Mn are present in the soil in very small amounts and therefore these are called ‘micronutrients’. However, these are essential to crop growth. The estimation of these elements in the soil is helpful in determining their deficiencies and compensating for the same. When a solution of a metallic salt or compound is sprayed into a flame, the solute molecules dissociate to give gaseous atoms, some of which may be excited to higher energy levels. But a larger number of gaseous metal atoms will remain in the ground or unexcited state. These atoms absorb radiant energy of their own resonance wavelength. So, when a light beam of resonance wavelength is passed through the flame containing the atoms under consideration, a part of the light will be absorbed. The absorption obeys the Beer-Lambert law and the absorbance is directly proportional to the number of ground state atoms in the flame i.e., the concentration of the element in the given solution. This is the basis of atomic absorption spectrophotometry (AAS). There may be interference by the radiation emitted from excited atoms of the same metal when they return to the ground state (the excitation occurs due to the absorption of incident light by some metal atoms). In AAS, a light source of specific pulse frequency is used. This is matched with an AC amplifier which will respond only to signals of that particular frequency. Therefore, emitted radiation will not be measured. The soil is first extracted with diethylenetriamine-pentaacetic acid (DTPA). All the four elements can be determined with the same extract using appropriate hollow cathode lamps in atomic absorption spectrophotometer.

REAGENTS Extracting solution: Take 0.9835 g DTPA and 0.73 g CaCl2.2H2O in a 500 mL standard flask. Add about 15 mL double-distilled water followed by 7.5 mL triethanolamine. Make up the volume to the mark with double-distilled water. Adjust the pH to 7.3 with 2 M hydrochloric acid.

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Micronutrient Elements

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Stock solution of Zn: Weigh accurately 0.25 g pure zinc metal and transfer it into a 250 mL standard flask. Then dissolve it in 2.5 mL concentrated hydrochloric acid. Make the solution upto the mark with extracting solution. This is a 1000 mg L–1 (ppm) solution. Dilute appropriately with extracting solution to get 100 mg L–1. Stock solution of Cu: Dissolve exactly 0.25 g pure copper turnings or wire in about 10 mL nitric acid (1:1). Make up the volume to 250 mL with the extracting solution. This is a 1000 mg L–1 Cu solution. Dilute this appropriately with extracting solution to get 100 mg L–1 (ppm) solution. Stock solution of Fe: Accurately weigh 0.25 g pure iron metal and dissolve in 1:1 nitric acid (about 10 mL) in a 250 mL standard flask. Make up the solution to 250 mL with extracting solution. Prepare a 100 mg L–1 (ppm) solution by diluting 10 mL of this solution to 100 mL with extracting solution. Stock solution of Mn: Dissolve 0.3958 g MnO2 (AR) in 10 mL nitric acid (1:1). Make up the solution to 250 mL with extracting solution. Dilute 10 mL of this solution to 100 mL with extracting solution. This is a 100 mg L–1 (ppm) Mn solution.

PROCEDURE Weigh accurately 50 g soil sample and take it in a 250 mL conical flask provided with a lid. Add 100 mL extracting solution and shake thoroughly for 2 hours. Filter the solution through Whatman No. 40 filter paper. Use 20 mL extract (filtrate) for each of the estimations of Zn, Cu, Fe and Mn. Use appropriate hollow cathode lamp in the atomic absorption spectrophotometer. Construct a separate calibration curve for each element as follows: Prepare the standard solutions of 1, 2, 3, 4, 5 and 6 mg L–1 of the metal by pipetting 0.5, 1,0, 1.5, 2.0, 2.5 and 3.0 mL stock solution (100 mg L–1), respectively in a series of 50 mL standard flasks. Make up the volume to the mark with extracting solution. Set the atomic absorption spectrophotometer with the suitable hollow cathode lamp and then take the absorbance readings for the standard solutions. The optimum absorbance range is 0.2-0.8. Plot a graph with absorbance and concentration of element in mg L–1. Using the soil extract, measure the absorbance of each element as before. From the corresponding calibration curve and the measured absorbance, determine the concentration of the particular element in the soil.

CALCULATION Same as for “Available Potassium” (Experiment 58).

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62 1

Pesticide Residues

Due to the extensive use of pesticides in agriculture, some residues are present in crops and food products. A part of each of these toxic materials is retained by the soil. The two important classes of pesticides are the organochlorine and organophosphorus types; the former class includes DDT(dichlorodiphenyltrichloroethane), aldrin, dieldrin, heptachlor and heptachlor epoxide. These are very stable and hence more persistent in the environment than the organophosphorus pesticides which are less stable. Parathion and malathion are two of the organophosphorus pesticides.

IDENTIFICATION Principle The organochlorine pesticides can be identified with standard solutions using chromatographic technique. Chromatography is a process for separation of different molecular mixtures. It depends on the redistribution of the molecules of the mixture between two or more phases. Thin layer chromatography is generally used for identification and column chromatography for separation of the components in a mixture. In both the cases, distribution takes place between a ‘stationary liquid’ phase sorbed on a support and a mobile phase in contact with it. Retardation factor or Rf values are used to distinguish between various components. Rf =

Distance between the starting line and centre of the final zone of solute (in cm) Distance between the starting line and solvent front (in cm)

REAGENTS Standard pesticides: Endrin, DDT, dieldrin, heptachlor, heptachlor epoxide and DDD standards. Dilute the standards to 100 ppm in benzene. Spray reagent: Dissolve 0.2 g silver nitrate in 2 mL distilled water. Add 20 mL 2-phenoxyethanol and make up the solution to 400 mL with acetone. Then add 2 drops of 30% hydrogen peroxide. Solvent: Mix 98 mL n-heptane with 2 mL acetone

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Pesticide Residues

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PROCEDURE Draw a thin line with a pencil at the edges about 2 cm from one of the silica gel-coated glass plate. Mark six small dots at equal distances on the line corresponding to the six pesticides. Mark the 10 cm level. By a microsyringe or micropipette, apply a 10 mL sample of the pesticide solution. Hold the pipette vertically to the plate and gently touch the spot on the plate indicated for that sample. Repeat this 3 or 4 times. Dry the spot with air. Similarly, apply all the standard pesticides. Then transfer the plate to a tall beaker or jar containing the developing solvent. Keep the plate upright and take care to see that the sample end is immersed to a depth of 0.5 cm in the solvent. Close the jar with the lid. When the solvent front has travelled 10 cm on the plate, remove the plate and dry it with air for 5 minutes. Then place it in the vented hood chamber and spray with the spraying reagent. Repeat this spraying 3 or 4 times at 1-minute intervals. Dry the plate in the hood for 10 minutes. Then place it under an ultraviolet light source for about 15 minutes. Observe the brown or grey spots. (Protect the eyes from radiation with goggles). Measure the distance travelled by the pesticides from the original spots. Record the results. Calculate the Rf value for each pesticide.

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63 1 Separation and Identification

of Pesticide Residues from the Soil

PRINCIPLE The pesticides retained by the soil undergo degradation and are leached into water. Presence of pesticides affects the quality of the soil thereby having an impact on human food chain. Therefore, it is essential to determine and monitor the pesticide residues in the soil sample. In this experiment, the pesticide residues are extracted from the soil by ‘column chromatography’. The residues are dissolved in hexane and identified by ‘thin layer chromatography’, as in the previous experiment.

REAGENTS Methylene chloride solution (1:1): Mix 100 mL methylene chloride with 100 mL petroleum ether. Methylene chloride solution (20%): Mix 200 mL methylene chloride with 800 mL petroleum ether. Florisil (60/100 mesh): 50 g—activated at 650° C Anhydrous sodium sulphate Hexane (AR)

PROCEDURE Reactivate the florisil at 140° C for 12 hours. Then add 2.5 mL distilled water and keep in an airtight container for 48 hours. Mount the chromatograph column vertically. Use a ring stand for this purpose. Place some glass-wool at the bottom of the column. Add 30 g florisil to the

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column and pack it tightly without any gap. Add anhydrous sodium sulphate to a height of 1.5 cm. Wash the florisil with 60 mL 1:1 methylene chloride-petroleum ether solution. Reject the washings. Weigh 10 g dry soil sample and mix it with 20 g florisil thoroughly. Transfer this mixture to the column. Take 780 mL methylene chloride-petroleum ether (20%) mixture in a separating funnel and add this to the column. Elute with a flow rate of 6 mL per minute. Separately collect the pesticides in the following sequence. DDT, DDD – 180 mL Heptachlor and its epoxide – 300 mL Dieldrin – 660 mL Endrin – 780 mL Evaporate the eluates in Erlenmeyer flasks to dryness on a water bath inside a vented hood and dissolve each of the residues in 6 mL hexane. Proceed as in the previous experiment for the identification of pesticides in the extracts. Measure the distance each pesticide has migrated from the original spot. Determine the Rf values. Compare the results with those of the standards.

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2)46604--

AIR ANALYSIS

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64 1 Qualitative Analysis of CO,

CO2, H2S, SO2, NO and NO2

Air may be polluted by dust, fumes, smoke, gas, mist, vapour, etc. These pollutants affect the human, plant and animal lives. Pollution can be natural, like volcanic explosions, or man-made such as those from the biological and chemical processes carried out by man. Carbon monoxide, carbon dioxide, hydrogen sulphide, sulphur dioxide, nitric oxide, nitrogen dioxide and particulates are the major pollutants in air.

PRINCIPLE First, air is collected in a sampler and then a required quantity is pumped each time through a suitable absorbent for a particular pollutant analysis. Specific reagents can be added to the absorbed solution of air and the presence of various gases can be detected. Sometimes, the reagent itself may act as an absorbent.

PROCEDURE The sampler for air and the disperser for the absorbent are shown in figures 6 and 7, respectively.

Fig.6 Sampler

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Fig.7

Disperser

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A Laboratory Manual for Environmental Chemistry

As in figure 6, fit a plastic bag inside a 1 litre glass jar. Suck on the bent tube 2 so that air will fill the bag. Connect the tube 1 with the disperser (Fig.7) carrying a glass frit as shown by the dotted lines. Take the absorbent for each test in this disperser. Analyse the air as follows: Experiment 1.

2.

3. 4. a. b.

c.

Chapter_64.p65

Keep in the disperser 2 mL cuprous chloride in conc. HCl. Blow gently through tube 2. Let in some sample air from the bag through a few mL lime water or Ba (OH)2 solution. Pass sample air through a solution of ferrous sulphate. Blow sample air through distilled water. Divide this solution into three parts. To one part add a few drops of dilute H2SO4 solution To the second part, add a few drops each of acetic acid and lead acetate solutions. To the third portion, add a small quantity of BaCl2 solution

146

Observation

Inference

White crystalline precipitate

Presence of CO

White precipitate or turbidity.

Presence of CO2

Dark brown colouration

Presence of NO

Dark brown colouration Black precipitate

Presence of NO2 Presence of H2S

White precipitate, Presence of SO2 soluble in conc. HCl

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65 1

Carbon Monoxide

PRINCIPLE CO is determined by non-dispersive infrared spectrometer (NDIR). In this technique, the radiation is not dispersed as in normal spectrometer where the measurement is carried out at fixed frequency or wave number. A gas gets heated when it absorbs IR radiation. This results in an increase of pressure if the volume remains constant. The change in pressure is directly proportional to the intensity of radiation. Unlike in a monochromatic radiation, here the gas absorbs at all its characteristic wavelengths. The detector consists of the same gas whose concentration is to be determined. CO strongly absorbs IR radiation at certain wavelengths. Therefore, when IR radiation is passed through CO absorbed solution, CO absorbs a certain amount of energy. This is related to the concentration of CO.

PROCEDURE Fill the reference cell with clean air and the sample cell with air containing CO. Both the cells are quite long i.e., 1 metre. Radiation from the source is split along two paths by means of a ‘chopper’. In one path is kept the reference cell and in the other the sample cell. The detector consists of two chambers filled with CO gas. The chambers are divided by a thin metal diaphragm. Depending on the amount of CO in the sample, there is a difference in absorption between the reference air and the sample air. This determines the absorption of IR radiation by CO gas in the two chambers of the detector. Due to the difference in absorption, pressure difference will arise in the two compartments. As a result, the diaphragm bulges towards one chamber. This movement of the diaphragm is detected and recorded. The output is a series of pulses, which is proportional to the concentration of CO in sample air.

Note 1. For estimation of CO, which has a concentration of 10 mg L–1 or less, gas chromatography is used. The sample is subjected to catalytic reduction by H2 over nickel at 360° C and measured with a flame ionization detector. CO2 + 4H2 ® CH4 + 2H2O

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66 1

Carbon Dioxide

PRINCIPLE CO2 is absorbed in a known excess volume of barium hydroxide solution. BaCO3 is precipitated. Then the excess of Ba(OH)2 is titrated with standard hydrochloric acid. A blank is run with Ba(OH)2 alone. Knowing the amount of unreacted Ba(OH)2, the concentration of CO2 in air can be evaluated. CO2 + excess Ba(OH)2 ® BaCO3 + H2O Ba(OH)2 unreacted + 2HCl ® BaCl2 + 2H2O

REAGENTS Barium hydroxide solution (approximately 0.005 M): Dissolve 1.8 g Ba(OH)2 . 8H2O (AR) in one litre hot, distilled water. Close the container and allow the solution to stand overnight. Barium carbonate, if present, will settle down. Filter and store the solution in an airtight bottle. Hydrochloric acid : 0.01 M solution Propan-1-ol (AR) Phenolphthalein indicator

PROCEDURE Place exactly 40 mL Ba(OH)2 solution in a wash bottle carrying a glass frit disperser (Fig.7). This itself acts as the absorbent. Add a few drops of propan-1-ol to lower the surface tension and to cause mild foaming. Pump 5 litres air through the absorbent at a rate of 0.5 litre per minute. The absorption may be slow. Take care to see that small air bubbles are formed and these have a long contact time with Ba(OH)2 solution. Then allow the precipitate to settle down. From the clear supernatant solution, pipette exactly a known quantity of the solution. Add a few drops of phenolphthalein indicator to the solution and titrate it with standard HCl. The end point is the disappearance of pink colour. To avoid mixing in any more air, have some cotton wool around the

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Carbon Dioxide

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tip of the burette or pass N2 over Ba(OH)2 solution. Carry out the blank titration with 10 mL Ba(OH)2 solution alone.

CALCULATION 40 mL Ba(OH)2 (blank) º a mL of 0.01 M HCl 40 mL Ba(OH)2 + 5 L absorbed air sample º b mL of 0.01 M HCl º (a–b) mL of 0.01M HCl Ba(OH)2 consumed by CO2 in air =

(a – b) × 0.01 moles HCl 1000

Ba(OH)2 + 2HCl ® BaCl2 + 2H2O 2 moles HCl º 1 mole Ba(OH)2

(a – b) × 0.01 (a – b) × 0.01 moles Ba(OH) 2 moles HCl º 2 ×1000 1000 = (a – b) ´ 5 ´ 10 –6 moles Ba(OH)2 Ba(OH)2 + CO2 ® BaCO3 + H2O 1 mole Ba(OH)2 º 44 g CO2 (a – b) ´ 5 ´ 10–6 moles Ba(OH)2 º 44 ´ (a – b) ´ 5 ´ 10–6 g CO2 Hence, the amount of CO2 in the air sample =

44 × (a – b) × 5 ×10 – 6 ×1000 5

= (a – b) ´ 44 ´ 10 –3 mg L–1

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67 1

Hydrogen Sulphide

PRINCIPLE An air sample containing H2S is absorbed in zinc acetate solution. It is then made to react with N, N-diethyl-p-phenylenediamine in presence of NaClO4 and ferric ammonium sulphate. The complex formed is extracted with chloroform and its absorbance measured. Using the standard sulphide solutions, a calibration curve is drawn.

REAGENTS Zinc acetate solution: Dissolve 0.25 g zinc acetate(AR) in 100 mL distilled water. Sodium perchlorate solution (1 M): Dissolve 6.123 g sodium perchlorate in 50 mL water. Sulphuric acid solution (50% W/V): Mix well 27.2 mL concentrated sulphuric acid (sp.gr.1.84) with 100 mL water. Sulphuric acid solution (1 M): Add slowly 5.3 mL concentrated sulphuric acid to 100 mL water. N, N-diethyl-p-phenylenediamine: Dissolve 2 g in 100 mL 50% sulphuric acid (W/V).

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Ferric ammonium sulphate solution: Dissolve 10 g solid in 100 mL of 1 M sulphuric acid. Standard sulphide solution: Weigh accurately 244 mg pure sodium sulphide crystals and dissolve it in one litre distilled water. This is a 100 mg L–1 (ppm) sulphide solution. Chloroform (AR)

PROCEDURE Into a set of five 25 mL standard flasks, pipette 0.5, 1.0, 1.5, 2.0 and 2.5 mL stock solution of Na2S. To each flask, add the following solutions: 2 mL NaClO4 ,1 mL N-N-diethyl-p-phenylenediamine and 1 mL ferric ammonium sulphate. Make up each to 25 mL with chloroform. After 10 minutes, but before 20 minutes, measure the absorbance of these solutions at 670 nm. Draw the calibration curve by plotting absorbance values against the concentration of sulphide in mg L–1 (ppm). Take, in a standard wash bottle fitted with glass disperser, 50 mL zinc acetate solution. This is the absorbent. Pump 500 litres air at the rate of 12 litres per minute through the abrosbent. After the pumping was stopped, add the following solutions in order. 2 mL NaClO4, 1 mL N,N-diethyl-p-phenylenediamine and 1 mL ferric ammonium sulphate. Shake the solution well and after 10 minutes, extract the solution 2 or 3 times, each time with 5 mL chloroform. Transfer the extracts into a 25 mL standard flask. Make up the volume to the mark with chloroform. Measure the absorbance in the same cell used for the standards at 670 nm. Determine the concentration of sulphide in the given sample of air from the calibration curve.

Note 1. In all absorbance measurements, blank solution should contain all the reagents except the substance under investigation.

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68 1

Sulphur Dioxide

PRINCIPLE SO2 is trapped in a solution of mercuric chloride and potassium chloride. The solution is then reacted with formaldehyde and pararosaniline hydrochloride. A violet red dye is formed. Its absorbance is measured. A calibration curve is plotted with the absorbances of standard sulphite solutions and their concentrations. [HgCl4]2– + SO2 + H2O ® HgCl2. SO32– + 2H+ + 2Cl– SO2 + H2O +HCHO ® HOCH2SO3H +

NH3

+

+

C

H3N

NH3 + 3 HOCH2SO3H

Cl +

NH3

HO3

S

CH2

NH

C

CH2SO3H

NHCH2SO3H + Cl– + 3H2O + 3H+

Red violet

REAGENTS Absorption solution: Weigh 3.4 g HgCl2 (AR) and 1.875g KCl (AR) and dissolve these together in distilled water. Make up this solution to 250 mL with water.

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Pararosaniline reagent: Dissolve 1 g pararosaniline in 100 mL distilled water. Mix 2 mL of this solution with 3 mL conc. HCl and dilute to 50 mL with water. Formaldehyde solution (0.2%) Dilute 2.5 mL commercial formaldehyde (40%) solution to 500 mL with distilled water. Standard solution of sulphite: Weigh accurately 0.1575 g sodium sulphite (AR), dissolve in a little distilled water and make upto one litre with water. This is a stock solution of 100 mg L–1 (ppm) sulphite.

PROCEDURE Pipette aliquots of 0.5, 1.0, 1.5, 2.0 and 2.5 mL standard sulphite solution into separate 25 mL standards flasks. Add 10 mL absorption solution, 1 mL pararosaniline reagent and 1 mL HCHO to each solution. Make upto the volume with water. Mix well and after 15 minutes, measure the absorbances at 560 nm. Draw the calibration curve by plotting absorbance values against concentrations of sulphite in mg L–1 (ppm). Place 10 mL absorption solution in a small impinger (disperser). Pump in 50 litres air at the rate of 1.5 litres per minute. When pumping is complete, add 1 mL each of pararosaniline reagent and HCHO solution. Shake the solution thoroughly. Transfer the solution to 25 mL standard flask. Wash the inner sides of the impinger with water and transfer the washings to the standard flask. Make upto 25 mL with water. Mix the solution thoroughly and after 15 minutes, measure the absorbance in the same cell as before at 560 nm. From the calibration curve, determine the concentration of sulphite corresponding to the absorbance of the sample.

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69 1

Nitric Oxide and Nitrogen Dioxide

PRINCIPLE A sample of air is absorbed in a mixture of N-(1-naphthyl)-ethylenediamine and sulphanilic acid. Diazotisation and coupling reactions take place. The absorbance of the purple azo-dye is measured. Another sample of the same volume of air is passed separately through the oxidizing agent, CrO3 in water. NO in air gets oxidized to NO2. Then it is pumped through the absorption reagent and all the NO2 is trapped. This NO2 includes the original as well as that from NO. The absorbance of the dye is then measured. Using the calibration curve, the concentration of NO2 alone from the first experiment and that of total NO and NO2 from the second experiment are evaluated. The concentration of NO is then found out from the difference between these two values. Equations: Refer experiment 16.

REAGENTS Absorbing reagent: Dissolve 2.5 g anhydrous sulphanilic acid in a mixture of 400 mL water and 70 mL glacial acetic acid. Warm, if necessary. Cool and add 10 mL N-(1-naphthyl)-ethylenediamine hydrochloride solution in water (0.1%). Make upto 500 mL with water. This reagent is stable for a long time if kept in a refrigerator. Oxidizing agent: Dissolve 17 g CrO3 in 100 mL water very cautiously. Soak 20 g finely ground fire-brick in this solution for 45 minutes. Dry the granules at 120° C. Place these granules in a tray over a saturated solution of sodium acetate in a closed container. Leave it overnight. Standard solution of sodium nitrite: Weigh accurately 177 mg NaNO2(AR) and dissolve it in water. Make up the solution to one litre with water. This stock solution contains 100 mg L–1(ppm) of nitrite.

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155

PROCEDURE Pump 2 litres air through 20 mL absorbing reagent in an impinger. Maintain the rate of flow at 0.5 litre per minute. After 15 minutes, measure the absorbance at 550 nm. Pump exactly the same volume of air though the oxidizing agent and then through the absorbing reagent. Adjust the rate to 0.04 litre per minute to effect the oxidation completely. After 15 minutes, note the absorbance of the solution at 550 nm. Pipette 0.5, 1.0, 1.5, 2.0 and 2.5 mL standard solution of nitrite into separate 25 mL standard flasks. To each of these, add 20 mL absorbing reagent and make upto the volume with water. Measure the absorbances at 550 nm. Plot the absorbance values against concentrations of nitrite in mg L–1 (ppm). From the calibration curve, determine the concentration of NO2 and (NO+NO2) in mg L–1 (ppm). Find out the difference between these two values, which gives the concentration of NO.

Note 1. Chemiluminescence method : NO2 is converted into NO by passing the air sample over a thermal converter. It is then allowed to react with ozone to produce electronically excited NO2* molecule. This molecule loses energy and returns to the ground state by emitting radiation in the range 600-3000 nm. This radiation is measured by a photomultiplier tube (Chemiluminescence detector). NO + O3 ® NO*2 + O2 NO2* ® NO2 + hv

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70 1

Particulates

PRINCIPLE Particulates are estimated using a high volume sampler (Fig.8). This consists of an electric motor and a pump. A gas meter is attached to this unit to indicate the volume of air pumped through the sampler. This is fitted with a filter made up of glass fibre and the outlet of the pump is connected to the filter. The sampler is capable of handling large volumes (about 2 ´ 10 6 L) of air. When air passes through the filter, particles of diameter in the range 0.1 – 100 micron are collected on the filter. The difference in the mass of the filter before and after the passage of air will give the amount of particulate matter in the air.

PROCEDURE Maintain the filter at 30° C and at the relative humidity 45 – 55% for 20 hours. Weigh it accurately. Place the filter in the sampler and pass 2x106 litres of air through it. Maintain the rate at 1500 L min–1. After the passage of air (it will take about 22 hours) remove the filter. Again leave it for 20 hours at 30° C and at 45 – 55% relative humidity. Weigh it again. Calculate the difference in the mass of the filter. Report the amount of particulates as mg L–1 (ppm) of air.

Filter Air inlet

Electric motor

Fig. 8

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Pump

High volume sampler

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PART FOUR

MISCELLANEOUS

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71 1

Calcium in Egg Shells

The strength of an egg shell depends upon the amount of calcium salt formation in the form of shell. The shell has to withstand the pressure exerted on it by the mother bird during incubation. The accumulation of DDT in the bird (through the food chain) leads to thinning of egg shells. DDT interferes with the calcium metabolism in the bird resulting in weaker or delicate egg shells. Such eggs break easily before they are hatched. The near extinction of falcons in USA is attributed to the effect of DDT on the egg shells of these birds.

PRINCIPLE The egg shell is ashed and the ash is dissolved in hydrochloric acid. The resulting solution containing Ca2+ is titrated with EDTA.

REAGENTS EDTA( 0.05M): Dissolve 18.6 g EDTA in distilled water and make up the solution to one litre Buffer solution: Mix 70g of NH4Cl (AR) with 570mL of liquor ammonia and dilute the resulting solution to one litre. Eriochrome black T indicator: Dissolve 0.1 g solid in a mixture of 7.5 mL triethanolamine and 2.5 mL ethanol.

PROCEDURE Break the eggs and discard the albumen and yolk. Wash the shells repeatedly with distilled water. Dry in a desiccator overnight. Grind the dry shells in a mortar with a pestle to obtain a fine powder. Weigh 0.2 g powder into a porcelain crucible. Heat the crucible in a muffle furnace at 700°C or in a Bunsen electrothermal heater until the conversion of shell to ash is complete (this may take 12-15 hours). Cool the crucible and add 3 mL distilled water and 1 mL of 12M HCl. Transfer the solution into a 100 mL conical flask. Add 10 mL buffer and 3 or 4 drops of indicator. Gently swirl the flask and titrate the solution with the EDTA in the burette. The end point is the change of wine red colour of the solution to blue colour.

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Repeat the experiment with another 0.2 g sample of the egg shell and check for concordancy.

CALCULATION 1000 mL 1 M EDTA º 40.08 g Ca Percentage Ca in the egg shell =

40.08 × M × V ×100 1000 × w

V = titre value in mL for the solution of egg shell w = mass of egg shell powder taken M = molarity of EDTA

Notes 1. This titre value actually corresponds to Ca and Mg present in the shell. However, the Mg content is negligible compared to the Ca content in shells. 2. The Ca content of the egg shell would vary with the type of bird (egg).

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72 1 Effect of pH on Germination

of Seeds

The pH of soil influences the germination of seeds. Fertile soils for growing crops have a pH of 6.5 – 7.0. However, some types of plant can grow at pH values outside these limits. The pH of the soil may affect the germination, rate of growth and other plant characteristics. Some seeds will not germinate when the soil pH is very low or very high.

REAGENTS Disodium hydrogen phosphate: Dissolve 28 g salt in water and dilute the solution to one litre. Citric acid: Dissolve 192 g solid in water and dilute to one litre. Mustard seeds: 50 g.

PROCEDURE Prepare a set of 8 buffer solutions in the pH range 4– 8 by mixing disodium hydrogen phosphate and citric acid solutions in 50 mL standard flasks as indicated below: Table 4

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Volume of Na2HPO4 (mL)

Volume of citric acid (mL)

pH

19.2 25.8 29.0 31.6 36.4 41.2 46.8 48.6

30.8 24.2 21.0 18.4 13.6 8.8 3.2 1.4

4.0 5.0 5.5 6.0 6.5 7.0 7.5 8.0

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Take 24 Petri dishes and arrange them in 8 groups, 3 in each group. Number the groups, 1 to 8. Place a circular filter paper in each of the dishes just fitting inside the dish. Add 10 mL of each buffer solution to each of the eight groups of Petri dishes so that in a group, 3 dishes have the same buffer solution. Place 50 mustard seeds on the surface of each filter paper in the Petri dish. No two seeds should be in contact with each other. Leave the dishes in a warm place. After seven days, count the number of seeds that have germinated in each dish. Every day check the Petri dishes and ensure that the filter papers are moist. If the papers are found dry, gently add distilled water without disturbing the seed (buffer solution, instead of distilled water, should not be added). Calculate the average percentage of germination for each group of 3 dishes. Plot these percentage values against pH values and find out the optimum pH range for the germination of mustard seeds.

CALCULATION The percentage of seed germination for each dish =

No. of seeds germinated ´ 100 Total No. of seeds

For each group, average percentage germination =

Sum of percentage germination for 3 dishes 3

Note 1. This experiment can be repeated for other types of seeds, e.g., paddy.

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73 1

Effect of Copper Salt on Germination of Seeds

Copper is essential for plant growth as a micronutrient. But excess copper is toxic for plants. If the copper content of the soil exceeds certain value, germination, plant growth and yields of plant products are severely affected.

REAGENTS Copper sulphate solutions: Prepare the following solutions using CuSO45H2O crystals. 0.4 g in 100 mL (1000 mg L–1 Cu) 0.5 g in 100 mL (1250 mg L–1 Cu) 0.6 g in 100 mL (1500 mg L–1 Cu) 0.7 g in 100 mL (1750 mg L–1 Cu) 0.8 g in 100 mL (2000 mg L–1 Cu) 1.0 g in 100 mL (2250 mg L–1 Cu) Mustard seeds: 50 g.

PROCEDURE Take 21 Petri dishes, label and arrange them in seven groups of three each. Place a circular filter paper just to fit inside each dish. Pipette into the Petri dishes of each group 10 mL of each of the copper sulphate solutions. The copper sulphate solution added should completely moisten the filter paper and there should also be a trace of surplus solution in the dish. In one group of dishes, pipette 10 mL distilled water. On each filter paper, place 50 mustard seeds such that no two seeds are in contact with each other. Leave the dishes in a warm place. Check the dishes everyday, and if they are dry, moisten them with the minimum amount of distilled water (copper sulphate solution should not be added instead of water). After seven days, count the seeds that have germinated in each dish. Calculate the average percentage of seed germination for each concentration of copper.

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Average percentage of seed germination =

No. of seeds germinated ´ 100 Total No. of seeds

Plot germination percentage values against Cu (II) concentrations. Write, in three or four sentences, the conclusions of this experiment.

Note 1. This experiment can be repeated with other types of seed. For each type of seed, determine the limit of copper concentration, above which no germination will occur. Similar experiments may be carried out to find out the poisoning limits by salts of other heavy metals such as mercury, lead, silver, zinc, cadmium, chromium, nickel and cobalt.

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74 1

Phosphate in Detergents

Most detergents contain some form of phosphate as a builder. The phosphate is usually present in the form of sodium tripolyphosphate (Na5P3O10).

PRINCIPLE Phosphate is determined spectrophotometrically by forming phosphovanadomolybdic acid. This can be used for the determination of phosphate without reducing it to heteropoly blue. (Refer: Expteriment 26). In this procedure, the polyphosphates in detergents are converted into HxPO4x–3 species by hydrolysis in acid medium. P3O105– + 2H2O ® 2HPO42– + H2PO4–

REAGENTS Standard phosphate stock solution (0.001M): Dissolve 0.136 g potassium dihydrogen phosphate (KH2PO4) in 1L of 0.1% H2SO4. Working standards of 0.001, 0.0002, 0.0004 and 0.0006 M solutions are prepared by the dilution of stock standard. Ammonium vanadomolybdate solution: Add 1.0 g ammonium vanadate (NH4VO3) to a mixture of 300 mL water and 200 mL HNO3. Add 40 g ammonium molybdate in 400 mL water. Mix these two solutions and dilute to one litre. Sulphuric acid solution (2M): Add 10 mL conc. H2SO4 to 100 mL water. Antifoaming agent: 30% silicone emulsion.

PROCEDURE Transfer 1 g detergent sample to a 100 mL round-bottom flask fitted with a reflux condenser. Add 1 or 2 drops of antifoaming agent and 25 mL of 2M H2SO4. Boil gently for 30 minutes. Cool and transfer to a 250 mL volumetric flask and make up. Transfer 2 mL of this solution to a test tube and add 5 mL ammonium vanadomolybdate solution. Mix and wait for colour development (4 to 8 minutes) and then measure the absorbance at 415 nm.

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Take 2 mL samples of diluted standard phosphate solutions, add 5 mL ammonium vanadomolybdate reagent to each and measure the absorbances at 415 nm. Prepare a calibration graph by plotting absorbance versus phosphate concentration. Determine the phosphate concentration of sample from the calibration curve.

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75 1 Benzoic Acid andSorbic Acid in Soft Drinks

Benzoic acid and its sodium salt are among the most widely used antibacterial agents in foods. Many berries (e.g. raspberries) contain appreciable amounts (0.05%) of benzoic acid as part of their natural composition. Sorbic acid (CH3CH=CH–CH=CH COOH) is present in some fruits and is a selective growth inhibitor for moulds, yeasts and bacteria. Both benzoic and sorbic acids are used as preservatives in soft drinks.

PRINCIPLE Benzoic or sorbic acid is determined by measurement of ultraviolet absorbance after extracting them with ether.

REAGENTS Ether (AR) Anhydrous sodium sulphate (AR) Hydrochloric acid solution(2M): Dissolve 19 mL conc HCl in 100 mL water. Sodium hydroxide solution(0.01M): Dissolve 0.1g NaOH pellets in 250 mL water.

PROCEDURE Place 25 mL degassed soft drink in a 100 mL separating funnel. Mix with a few drops of conc. HCl. Extract the mixture with three portions of ether, using volumes of 50, 25 and 25 mL. Combine the extracts and wash with 15 mL water. Dry the extract by swirling with anhydrous sodium sulphate. Evaporate the extract to dryness on a steam bath. Dissolve the residue in 100 mL of 0.01M NaOH. Record the uv spectrum of the solution from 190 to 270 nm.

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CALCULATION In dilute NaOH solution, benzoic acid has a maximum absorption at 200 nm and sorbic acid has a maximum absorption at 254 nm. The Î values are 130 M–1cm–1 and 2.48 ´ 104 M–1 cm–1 for benzoic and sorbic acids, respectively. Calculate the concentrations of these acids using the relationship A = Î bc where A is the absorbance, b is the cuvette length in cm , c is the concentration (M) and Î is the molar absorptivity.

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Sulphur Dioxide in Soft Drinks

Sulphur dioxide is used as a food preservative and is the most effective inhibitor of the deterioration of dried fruits and fruit juices.

PRINCIPLE Sulphur dioxide is determined by the redox reaction with iodine solution. SO2 + I2 + H2O ® SO3 + 2H + + 2I –

REAGENTS Sodium hydroxide solution(2M): Dissolve 8 g NaOH pellets in 100 mL water. Sulphuric acid solution (2M): Add carefully 10 mL conc. H2SO4 to 100 mL water. Iodine solution (0.001 M): Dissolve 2g KI in water and add 0.127 g I2 and stir thoroughly. Dilute to 1000 mL. Starch solution: Prepare a fresh solution using soluble starch powder.

PROCEDURE Place 100 mL degassed soft drink in a conical flask. Add about 5 mL of 2M NaOH and let it stand for about half an hour. Add about 1 mL fresh starch and 10 mL of 2M H2SO4 and titrate with 0.001M iodine solution until a permanent blue colour is obtained. Express sulphur dioxide concentration as mg L–1 of soft drink.

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77 1

Identification of Food Adulterants

INTRODUCTION TO FOOD ADULTERATION Food adulteration is the intentional addition of non-permitted foreign matter to food or the removal of useful ingredients from food. According to the Prevention of Food Adulteration Act of India 1954, it also includes the willful addition of those substances, which adversely affect the quality of food and their incidental contamination during the period of growth, processing, transportation and distribution. The International Codex Alimentarius Commission has suggested many internationally adopted food standards, including the provisions in respect of food hygiene, food additives, pesticide residues, contaminants, labeling, analysis and sampling. In India too, we have two national standards-the Prevention of Food Adulteration and the Fruit Products Order, both being amended from time to time. In spite of many rules and Acts, the adulteration of food is steadily increasing. This has an impact on public health especially in the poor and under-nourished population. For example, the non-permitted dyes are used for colouring, which are not absolutely safe. Most of them are carcinogenic in nature. They may cause pathological lesions in the vital organs like liver, spleen and kidney. Similarly, lead chromate or metanil yellow may cause anaemia, paralysis, abortion, mental retardation and brain damage in children. In order to prevent a large number of diseases that are disastrous to our health, it is desirable and advisable to detect various adulterants present in the food stufs.

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Food Article

Adulterant

Test

Observation and Inference

1. Mustard oil

Argemone oil

Treat sample of mustard oil with ferric chloride in the presence of HCl

Needle-shaped brown crystals show the presence of Argemone oil. Turbidity shows the presence of castor oil.

2. Edible oils (Groundnut & coconut oil)

Mineral oil or castor oil

Dissolve the sample in petroleum ether and acidify the solution with dilute HCl. To this add a few drops of 1% ammonium molybdate in 2M H2SO4 and shake vigorously

Turbidity shows the presence of castor oil.

3. Vegetable oil

Tea seed oil

Mix 0.8 mL acetic anhydride, 1.5 mL chloroform and 0.2 mL conc.H2SO4 in a test tube. Cool it in an ice bath for 2 minutes. Then add 6 to 7 drops of the given oil. If the liquid appears cloudy, add a few more drops of acetic anhydride to make it clear. Keep the test tube in the ice bath for 5 minutes. Add 10 mL dry ether to the above solution and shake. To a little amount of the ghee or butter in a test tube, add equal amount of conc. HCl and a little sugar and shake vigorously. Keep it

Development of green colour indicates the presence of tea seed oil.

Brown colour appears which changes to red within a minute and then fades away.

Confirms the presence of tea seed oil. Change in colour from blue to red shows the presence of vanaspati.

4. Ghee and butter

Vanaspati

5. Coffee powder

Chicory powder

Take a small amount of coffee powder in a test tube and add water.

Tamarind powder

Sprinkle the coffee powder Red colouration shows on a filter paper and add the presence of tamarind powder. 1% Na2CO3 solution

standing for 5 minutes. Red colour separation if the chicory content is high

(Contd.)

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Adulterant

Test

Observation and Inference

6. Tea

Saw dust and some colourants.

Sprinkle tea powder on a wet filter paper.

(a) Colour separation indicates non-permitted colour. (b) Pinkish Red colour indicates used-up tea leaves or outer coats of dhal.

7. Dhal

Tonic dyes (metanil yellow) and kesari dhal

To the sample, add 25 mL conc. HCl and keep in a water bath for 15 minutes.

Pale red colour shows non-permitted colour like metanil yellow.

8. Chilli powder

Brick powder

(a) Add water and shake.

Brick powder settles down showing contamination. Brickred flame colour. Presence of calcium salts in brick-powder.

(b) Perform flame test (introduce the paste of chilli powder & conc. HCl into flame) 9. Turmeric Powder

(a) Yellow aniline dyes

Shake a little of the turmeric powder (suspended in water) with alcohol.

The solution turns yellow immediately, if aniline yellow dye is present.

(b) Nonpermitted colourants like metanil yellow (c) Lead

To a little of the sample add conc. HCl

Red with no further change in colour shows the presence of metanil yellow.

chromate 10. Asafoetida

(a) Foreign resins

Powder a gram of asafoetida and shake it with alcohol. Filter the extract and to 5 mL of it add 10% ferric chloride. (b) Galbanum To a few mL of petroleum colophony ether extract, add 5 mL of resin. 5% aqueous copper acetate. Allow the layers to separate Burn on a nickel spatula

Olive green colour shows the presence of foreign resins.

A bluish green colour in the petroleum ether shows the presence of colophony resins.

Burns like camphor with a luminous flame shows pure sample. (Contd.)

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Adulterant

Test

173 Observation and Inference

11. Milk, Khoa and bura

Starch

Add a little water and boil for a few minutes. Cool and add iodine solution.

12. Pulses (green peas and dhal)

Colour Dye stuffs

Keep the sample immersed Colour separation in water for about half indicates adulteration. an hour and stir.

13. Bura

Washing soda

(a) Dip a red litmus paper in solution. (b) Add a few drops of dilute HCl

Blue colouration shows the presence of starch

Blue colour shows the presence of washing soda. Effervescence indicates the presence of adulterant.

14. Edible oils Karanjia (Mustard (pungam) oil oil or groundnut oil)

To two drops of the oil sample, add 2 ml SbCl3 in CHCl3 solution (20% W/V)

Appearance of canary yellow to orange yellow colour immediately shows the presence of karanjia oil.

15. Ghee

Rancid fat

To about 5 mL melted fat, add 5 mL 0.1% phloroglucinol in alcohol and 5 mL conc. HCl. Shake the contents well.

Deep pink or red colour indicates the presence of rancid fat.

16. Vinegar

Mineral acid

17. Honey

Invert sugar

Evaporate 2 mL vinegar sample in a crucible with excess of ammonia and heat. Dissolve the residue in water and add Nessler’s reagent. Dissolve 1g honey in 10 mL cold water. Transfer this solution into a small separating funnel and shake the solution with ether. Decant the ether layer and evaporate in a porcelain dish to half the original volume and add 1% solution of resorcinol in conc. HCl.

Formation of brown precipitate shows the presence of ammonium ion indicating the presence of mineral acid in vinegar. Pink colour indicates the presence of invert sugar in honey sample.

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APPENDICES

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Appendix 1

1

Water used for preparing solutions must be distilled water unless otherwise specified. Even if it is mentioned simply as “water”, distilled water must be used. For very dilute solutions, it is convenient to express concentration in terms of parts per million (ppm) or parts per billion (ppb). (ppm) =

Mass of solute ×106 Mass of solute ×109 ; ppb = Mass of solution Mass of solution

If solvent is water, the density of the solution is essentially 1 g mL–1. Therefore,

ppm =

mg solute mg solute = 6 L solution 10 mg water

1 ppm corresponds to 1 mg mL–1 or 1 mgL–1 or 10–4 wt % 1 ppb corresponds to 1 ng mL–1 or 1 mg mL–1 or 10 –7 wt %

How small? To get an idea of how small I part in a million (1 ppm) is, compare the following equivalents: – 2.5 cm in 25 km – one minute in 2 years – one paise in Rs.10000 – one drop in 50 L water – one mg in I L – one mg in 1 kg – one mg in 1g – one mg in 1 mL – one Hz in 1 MHz

Relationship between different units 1 M = mol L–1 = mol dm–3 = 103 mol m–3 L = dm3 = 10–3 m3 mL = cm3 = 10–6 m3 mL = 10–3 cm3

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1

Appendix 2

Classification of water bodies according to primary productivity, nutrient and chlorophyll-a content. Classification

Primary Productivity (mg C m2 day–1)

Oligotrophic Mesotrophic Eutrophic

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0 – 136 137 – 409 410 – 547

178

Total P (mg L–1) 0.03

Chlorophyll-a (mg L–1) 0.3 – 2.5 2.5 – 15 15 – 140

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Appendix 3

1

Guidelines of Bureau of Indian Standards (BIS) and Significance/Health Effects of Chemical Parameters- Water Quality Parameters – IS 10500:1991

Parameters

BIS Guideline value Significance/Health effects (maximum allowable)

Colour in Pt/Co scale 25 Turbidity 10 NTU Total dissolved solids 2000 mg/L pH

6.5 – 8.5

Alkalinity as ‘CaCO3’

600 mg/L

Hardness as ‘CaCO3’

600 mg/L

Calcium as ‘Ca’

200 mg/L

Magnesium as ‘Mg’

100 mg/L

Iron as ‘Fe’

1.0 mg/L

Manganese as ‘Mn’

0.3 mg/L

Aluminium as ‘Al’

0.2 mg/L

Copper as ‘Cu’

1.5 mg/L

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179

Consumer acceptance decreases Consumer acceptance decreases Undesirable taste; gastro intestinal irritations; corrosion or incrustation. Beyond this range, mucous membranes affected; bitter taste; corrosion; aquatic life affected. Taste becomes unpleasant, boiled rice is yellowish. Poor lathering with soap; deterioration of clothes; scale forming; incrustation in water supply structure; skin irritation; boiled meat/food becomes poor in quality Poor lathering and deterioration of clothes; incrustation in water supply structure; scale formation. Poor lathering and deterioration of clothes; with sulphate laxative; incrustation in water supply structure. Poor or sometimes bitter taste, colour and turbidity; staining of materials; iron bacteria causing slime. Poor taste, colour and turbidity; staining; black slime. Neurological disorders; Alzheimer’s disease (causes dementia). Astringent taste; discolouration and corrosion of pipes and utensils; liver damage; mucosal irritation, renal damage and depression; restricts growth of aquatic plant. (Contd.)

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180 Parameters

A Laboratory Manual for Environmental Chemistry BIS Guideline value Significance/Health effects (maximum allowable)

Zinc as ‘Zn’

15 mg/L

Ammonia as ‘NH3’ Nitrite as ‘NO2’

– –

Nitrate as ‘NO3’

100 mg/L

Sulphate as ‘SO4’

400 mg/L

Chloride as ‘Cl’ Fluoride as ‘F’

1000 mg/L 1.5 mg/L

Phosphate as ‘PO4’ Arsenic as ‘As’

– 0.05 mg/L

Mercury as ‘Hg’

0.001 mg/L

Cadmium as ‘Cd’

0.01 mg/L

Lead as ‘Pb’

0.05 mg/L

Chromium as ‘Cr’

0,05 mg/L

Cyanide as ‘CN’ Pesticide Anionic Detergent

0.05 mg/L 0.001 mg/L 1.0 mg/L

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180

Astringent taste; opalescence in water, gastro intestinal irritation; vomiting, dehydration, abdominal pain, nausea and dizziness. Indicates pollution; growth of algae. Forms nitrosoamines which are carcinogenic. Blue baby disease (methemoglobinemia); algal growth. Taste affected; laxative effect; gastro intestinal irritation. Taste affected; corrosive. Dental and skeletal fluorosis; nonskeletal manifestations. Algal growth. Toxic; bioaccumulation; central nervous system affected; carcinogenic. Highly toxic; causes ‘minamata’ disease (neurological impairment and renal disturbances); mutagenic. Highly toxic; causes ‘itai-itai’ disease (painful rheumatic condition); cardio vascular system affected; gastro intestinal upsets and hypertension. Causes plumbism (tiredness), lassitudes, abdominal discomfort, irritability, anaemia; bioaccumulation; impaired neurological and motor development, and damage to kidneys. Carcinogenic; ulcerations, respiratory problems and skin complaints. Highly toxic. Affects central nervous system. Undesirable froth, foaming

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Appendix 4

1 U.S.National Primary Drinking Water Standards for Inorganic and Microbiological Constituents of Water Contaminant

Maximum Contaminant Level (MCL)

Antimony (Sb) Asbestos Arsenic (As) Barium (Ba) Beryllium (Be) Cadmium (Cd) Chromium (Cr) Cyanide (CN–) Fluoride (F–) Mercury (Hg) Nickel (Ni) Nitrate (NO3– as N) Nitrite (NO2– as N) Total Nitrite/Nitrate Selenium (Se) Sulphate (SO2– 4 )

Inorganic 0.006 mg/L 7 MFL (million fibers per litre longer than 10 mm) 0.05 mg/L 2 mg/L 0.004 mg/L 0.005 mg/L 0.1 mg/L 0.2 mg/L 4 mg/L (secondary MCL of 2 mg/L triggers public notice) 0.002 mg/L 0.1 mg/L 10 mg/L 1 mg/L 10 mg/L 0.05 mg/L Regulation deferred because of relatively low health risks

Thallium (Tl) Lead (Pb)

0.002 mg/L 0.015 mg/L (No MCL set. This number represents the Action Level (AL), the level at which the authorities must do something to remove the contaminant.) 1.3 mg/L (No MCL set. This is the AL.) Microbiological < 1 per 100 mL 0 per 100 mL 0 per 100 mL 0 per 100 mL

Copper (Cu) Total coliform Fecal coliform Giardia lamblia Viruses

Source: United States Environmental Protection Agency.

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A Laboratory Manual for Environmental Chemistry U.S.National Secondary Drinking Water Standards

Contaminant

Maximum Contaminant Level (MCL)

–

Chloride (Cl ) Fluoride (F –) Sodium (Na+) Surfactants (soaps ü ý and detergents) þ Iron (Fe) Manganese (Mn) pH Sulphate (SO42–) Total dissolved ü ý solids (TDS) þ Zinc (Zn) Colour Corrosivity Odour

250 mg/L 2.0 mg/L 20 mg/L (health advisory) 0.5 mg/L 0.3 mg/L 0.05 mg/L 6.5 – 8.5 pH units 250 mg/L 500 mg/L 5.0 mg/L 15 colour units Noncorrosive 3 (threshold odour number)

Source: United States Environmental Protection Agency

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Appendix 5

1

Saturation Values of Dissolved Oxygen as a Function of Temperature Temperature (°C) 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35

DO (mg/L) 14.62 14.22 13.83 13.46 13.11 12.77 12.45 12.14 11.84 11.56 11.29 11.03 10.78 10.54 10.31 10.08 9.87 9.67 9.47 9.28 9.09 8.92 8.74 8.58 8.42 8.26 8.11 7.97 7.83 7.69 7.56 7.43 7.31 7.18 7.07 6.95 (Contd.)

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184

A Laboratory Manual for Environmental Chemistry Temperature (°C)

DO (mg/L)

36 37 38 39 40 41 42 43 44 45 46 47

6.84 6.73 6.62 6.52 6.41 6.31 6.21 6.12 6.02 5.93 5.84 5.74

Altitude Correction for Dissolved Oxygen Measurements Altitude (ft)

Altitude (m)

– 276 0 278 558 841 1126 1413 1703 1995 2290 2587 2887 3190 3496 3804 4115 4430 4747 5067 5391 5717 6047 6381 6717 7058 7401 7749 8100 8455

– 84 0 85 170 256 343 431 519 608 698 789 880 972 1066 1160 1254 1350 1447 1544 1643 1743 1843 1945 2047 2151 2256 2362 2469 2577

Correction Factora 1.01 1.0 0.99 0.98 0.97 0.96 0.95 0.94 0.93 0.92 0.91 0.90 0.89 0.88 0.87 0.86 0.85 0.84 0.83 0.82 0.81 0.80 0.79 0.78 0.77 0.76 0.75 0.74 0.73 (Contd.)

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Appendices Altitude (ft)

Altitude (m)

8815 9178 9545 9917 10293 10673 11058

2687 2797 2909 3023 3137 3253 3371

185 Correction Factora 0.72 0.71 0.70 0.69 0.68 0.67 0.66

a

Multiply the saturation value at sea level times this correction factor to obtain the saturation value at the specified altitude.

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Appendix 6

1

SOIL ANALYSIS Formula to be used in the colorimetric / spectrophotometric estimations

Amount of the particular ü = (concentration in soil extract) ´ (dilution) species present in the soil ýþ ´ (soil dilution) concentration from Volume of coloured Volume of extracting graph in mg –1 solution (mL) solution (mL) ´ ´ = 1000 volume of aliquot (mL) mass of soil (g) = a mg g–1

Note 1. Concentration from graph is in mg L–1. It is converted to mg mL–1 by dividing it by 1000.

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Appendix 7

1 NATIONAL AMBIENT AIR QUALITY STANDARDS (CPCB 1994) Pollutant

Time Weighted Average

Industrial Area

Sensitive Area

ANNUAL AVERAGE * 24 HOURS ** ANNUAL AVERAGE 24 HOURS

80 g/m3 120 g/m3 80 g/m3 120 g/m3

15 g/m3 30 g/m3 15 g/m 3 30 g/m3

ANNUAL AVERAGE 24 HOURS

360 g/m3 500 g/m3

70 g/m3 100 g/m3

RESPIRABLE ü ANNUAL AVERAGE PARTICULATE ïý 24 HOURS MATTER (RPM)ïþ LEAD ANNUAL AVERAGE 24 HOURS

120 g/m3 150 g/m3

50 g/m3 75 g/m3

1.0 g/m3 1.5 g/m3

0.50 g/m3 0.75 g/m3

SOx NOx SUSPENDED ü PARTICULATE ïý MATTER(SPM)ïþ

CARBON ü MONOXIDEý þ

8 HOURS 1 HOUR

5.0 g/m 3 10.0 g/m

3

1.0 g/m3 2.0 g/m3

* Annual Arithmetic mean of minimum 104 measurements in a year taken twice a week 24 hourly at uniform interval. ** 24 hourly/8 hourly values should be met 98% of the time in a year. However, 2% of the time, it may exceed but not on two consecutive days.

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Appendix 8

1 AIR QUALITY STANDARDS IN RESPECT OF NOISE NOISE POLLUTION (REGULATION AND CONTROL) RULES, 2000 Area code

Category of Area/Zone

Limits in dB(A) Leq* Day time

(A) (B) (C) (D)

Industrial area Commercial area Residential area Silence area

Night time

75 65 55 50

70 55 45 40

Note 1. Day time shall mean from 6.00 a.m. to 10.00 p.m. 2. Night time shall mean from 10.00 p.m. to 6.00 a.m. 3. Silence zone is defined as an area comprising not less than 100 metres around hospitals, educational institutions and courts. The silence zones which are declared as such by the Competent Authority. 4. Mixed categories of areas may be declared as one of the four abovementioned categories by the Competent Authority. * dB(A) Leq denotes the time weighted average of the level of sound in decibels on scale A which is relatable to human hearing. A “decibel” is a unit in which noise is measured. “A” in dB(A) Leq, denotes the frequency weighting in the measurement of noise and corresponds to frequency response characteristics of the human ear. Leq: It is an energy mean of the noise level over a specified period.

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Appendix 9

1

SATURATION INDEX OR LANGELIER INDEX This provides a method of calculating, from the methyl orange alkalinity, calcium hardness, total solids and water temperature, the pH at which the water would be just saturated with calcium bicarbonate. This is referred to as the saturation pH or pHs . The difference between the actual pH and saturation pH is called the saturation index or Langelier index. A negative Langelier index indicates that the water will tend to dissolve calcium carbonate and a positive index indicates that it will tend to deposit calcium carbonate. Data for rapid calculation of the Langelier index of waters based on the Langelier formula are given in the table: Total solids mg/L

A

50–300 400–1000

0.1 0.2

Temp °C 0–1 2–6 7–9 10–13 14–17 18–21 22–27 28–31 32–37 38–43 44–50 51–56 57–63 64–71 72–81

B 2.6 2.5 2.4 2.3 2.2 2.1 2.0 1.9 1.8 1.7 1.6 1.5 1.4 1.3 1.2

Calcium hardness mg/L CaCO3 10–11 12–13 14–17 18–22 23–27 28–34 35–43 44–55 56–69 70–87 88–110 111–138 139–174 175–220 230–270 280–340 350–430 440–550 560–690 700–870 880–1000

C

Methyl orange alkalinity

0.6 0.7 0.8 0.9 1.0 1.1 1.2 1.3 1.4 1.5 1.6 1.7 1.8 1.9 2.0 2.1 2.2 2.3 2.4 2.5 2.6

10–11 12–13 14–17 18–22 23–27 28–35 36–43 44–55 56–-69 70–88 89–110 111–139 140–176 177–229 230–270 280–359 360–449 450–559 560–699 700–889 890–1000

1. Obtain values of A,B,C and D from the table. 2. pHs = (9.3 + A + B) – (C + D) 3. Saturation index = pH + pHs.

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D 1.0 1.1 1.2 1.3 1.4 1.5 1.6 1.7 1.8 1.9 2.0 2.1 2.2 2.3 2.4 2.5 2.6 2.7 2.8 2.9 3.0

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References

1

1. Experiments on Land Pollution, D I Williams and D. Anglesea, Wayland Publishers Ltd., England, 1978. 2. Experiments on Air Pollution, D I Williams and D. Anglesea, Wayland Publishers Ltd., England, 1978. 3. Experiments on Water Pollution, D I Williams and D. Anglesea, Wayland Publishers Ltd., England, 1978. 4. Environmental Science : Laboratory manual, Maurice A Strobbe, The C V Mosby Company, Saint Louis, 1972. 5. Chemistry for Environmental Engineering, Chair N. Sawyer and Perry L McCarty, McGraw Hill Book Company, New York, 1978. 6. Environmental Protection, Chanlett, McGraw Hill Book Company, New York, 1975. 7. The Pollution Handbook, Richard Mabey, Penguin, 1978. 8. Experimental in Environmental Chemistry, P.D. Vowles and D.W. Connel, Pergamon Press, Oxford, 1980. 9. Environmental Chemical Analysis, Iain L Marr and Malcolm S Cresser, Blackie and Son Ltd., London, 1983. 10. Soil Chemical Analysis, M L Jackson, Prentice-Hall of India Pvt. Ltd., New Delhi, 1973. 11. Environmental Chemistry, S E Manahan, Willard Grant Press, Boston, 1983. 12. Vogel’s Textbook of Quantitative Inorganic Analysis (Revised Fourth Edition), J Bassett, R C Denney, G H Jeffery and J Mendham, Longman Group Ltd., London, 1978. 13. Manual of Soil Laboratory Testing - Vol.I, K.H. Head, Pentech Press, London, 1980. 14. Proc. Soil. Soc. Am., Fireman and Reeve, 13,494 (1948). 15. Standard methods of Chemical Analysis, F J Welcher (Part B), Robert E Krieger Publishing Co, USA, 1975. 16. Analytical Chemistry - Practice. John H. Kennedy, Harcourt Brace Jovanovich Publishers, New York, 1984. 17. Instrumental Analysis II Edn. Gary D. Christian an James E O’Reily Allyn and Bacon, Inc. Boston, 1986. 18. Fundamentals of Environmental Pollution, Krishnan Kannan, S. Chand and Company Ltd., New Delhi, 1997. 19. Environmental Chemistry, Colin Baird W.H. Freeman and Company, New York, 1995.

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Environmental Practicals

20. Pollution Science, Ian L. Pepper, Charles P Gerba, mark L. Brusseau, Academic Press, New York, 1996. 21. Marine Chemistry, Vol.I, II Edn. Dean F. Martin, Marcel Dekker Inc. New York, 1972. 22. A manual of Chemical and Biological Methods. Methods for Seawater Analysis, Timothy R. Parsons, Yoshiaki Maita and Cavol M. Lalli, Pergamon Press, New York, 1985. 23. Environmental Pollution Analysis, S.M. Khopkar, New Age International (P) Limited, New Delhi, 1995. 24. Environmental Chemistry, A.K. De, New Age International (P) Limited, New Delhi, 1996. 25. Introduction to wastewater treatment process, R.S.Ramalho, Academic Press, New York, 1977. 26. Industrial water pollution control, W. Wesley EcKenfelder, McGraw Hill Book Company, New York, 1989. 27. Standard Methods for the Examination of Water and Wastewater, A.D. Eaton, L.S. Clesceri and A.E. Greenberg, Eds. 19th Edn. American Public Health Association, Washington D.C. 1995. 28. The Chemistry of Water, Susan E. Kegley and Joy Andrews, University Science Books, California, 1998. 29. Exploring Chemical Analysis, Daniel C Harris, W.H. Freeman and Company, New York, 1997. 30. Contemporary Instrumental Analysis, Kenneth A. Rubinson and Judith F. Rubinson, Printice Hall International, Inc, New Jersey, 2000. 31. The Water Encyclopedia, F. Van der Leeden, F.L. Troise and D.K. Todd, Lewis Publishers, Boca Raton, FL, 1991. 32. Analytical Chemistry, Gary D. Christian, John Wiley & Sons, Inc. New York, 2001 33. Environmental Chemistry – a global perspective, Gary W. VanLoon and Stephen J. Duffy, Oxford University Press, Oxford, 2001.

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A Laboratory Manual for TM

Environmental Chemistry

A Laboratory Manual for

PART I – WATER ANALYSIS: Principles of Spectrometric Analysis / Sampling Techniques / Colour / Turbidity / Total Dissolved Solids / Conductance / pH Value / Acidity / Alkalinity / Dissolved Oxygen / Chemical Oxygen Demand / Biochemical Oxygen Demand / Dissolved Organic Carbon / Calcium Hardness of Water / Total Hardness of Water / Nitrate Nitrogen / Nitrate Nitrogen Brucine Method / Ammonia Nitrogen / Fluoride / Fluoride in Drinking Water (by Fluoride Electrode) / Chloride / Residual Chlorine / Sulphide / Sulphite / Sulphate / Phosphate / Iron / Chromium / Copper / Lead / Nickel / Cadmium / Mercury / Cyanides / Boron / Selenium / Arsenic / Suspended Solids / Determination of Chlorophyll-a / Determination of Productivity PART II – SOIL ANALYSIS: Collection of Soil Samples / pH Value / Soluble Salts / Identification of Calcium Carbonate, Nitrogen, Phosphorus and Potassium / Calcium Carbonate / Gypsum Requirement / Lime Requirement / Alkalinity / Organic Carbon / Total Nitrogen / Available Nitrogen / Ammonium Nitrogen / Nitrate Nitrogen / Nitrite Nitrogen / Available Phosphorus / Available Sulphur / Silica / Available Potassium / Sodium / Calcium and Magnesium / Micronutrient Elements / Pesticide Residues / Separation and Identification of Pesticide Residues from the Soil PART III – AIR ANALYSIS: Qualitative Analysis of CO, CO2, H2S, SO2, NO and NO2 / Carbon Monoxide / Carbon Dioxide / Hydrogen Sulphide / Sulphur Dioxide / Nitric Oxide and Nitrogen Dioxide / Particulates

Calcium in Egg Shells / Effect of pH on Germination of Seeds / Effect of Copper Salt on Germination of Seeds / Phosphate in Detergents / Benzoic Acid and Sorbic Acid in Soft Drinks / Sulphur Dioxide in Soft Drinks / Identification of Food Adulterants / Appendices / References / Index.

978-93-89583-58-8

9 789389 583588

Gopalan | Anand | Sugumar

PART IV – MISCELLANEOUS:

Environmental Chemistry

Contents

A Laboratory Manual for

The present book is meant for the students who opt for a course in “Environmental Chemistry” with laboratory work as a component of the course. Spread in 72 experiments, the analyses of soil, water and air have been described in a simple manner, so that most of these experiments can be conducted even by the beginners. The principles involved, preparation of the reagents, and the procedures are described for each experimental method. The authors hope that this manual would be useful in laboratories where soil, water and air are routinely tested.

Environmental Chemistry

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