Soil Testing

Certificate Course in Agricultural Biotechnology

BIOFERTILIZER TECHNOLOGY

Institute team:

Dr. B.K.Datta
Dr. R.Datta
Dr. S.K.Das
Dr. S.K.Si
Mr. S. Sahoo
Mr. D. Biswas
Mr. S. Giri
Mr. T. Nayek

Expert Consultants:

Dr. P.K. Singh, IRAI, New Delhi
Dr. O.P. Rupela, ICRISAT, Hyderabad
Dr. R.K.Basak, BCKV, Mohanpur, W.B
Dr. D.J. Bagyaraj, Univ. of Agric. Sciences, Bangalore
Dr. R. Kale, Univ. of Agric. Sciences, Bangalore
Dr. Sunil Pabbi, IARI, New Delhi
Dr. Aloke Adholeya, TERI, New Delhi

Vivekananda Institute of Biotechnology

First published in July, 2004

by Vivekananda Institute of Biotechnology
Nimpith, South 24 Parganas, West Bengal, India

Designed by Tellywallah, Calcutta
Printed by Swapna Printing Works Private Limited, Calcutta

All rights reserved.

c Vivekananda Institute of Biotechnology, 2004

This manual is sold subject to the condition that it shall not, by way of trade or otherwise,
be lent, resold, hired out or otherwise circulated without the publisher's prior consent
in any form of binding or cover other than that in which it is published.

Foreword

Fifty years back a monk, Swami Buddhanandji, was deeply inspired by Swami
Vivekananda’s ideals. He started a different kind of journey of life and a seed
of development was sown at Nimpith. Late Swami Buddhanandaji established
this Ashram.

Vivekananda Institute of Biotechnology is a branch of that tree. The
Institute initiated its activities in 1991. Its aim is to develop an advanced,
functional, research as well as a resource center for the people of the Sunderbans.
One of the chosen fields is Agricultural Biotechnology. For the last few years the
Institute has been implementing programmes on entrepreneurship development,
in this field, in rural areas. The present manual is one of such activity to support
this programme, which is the need of the hour.

The Department of Science and Technology, Govt of India and United
Nations Development Programme have come forward benevolently in bringing
out this manual. It is a result of the combined efforts of scientists at our Institute
and of other National Institutes and Universities. Senior scientists like Dr.
P.K.Singh, Dr. O.P.Rupela, Dr. Radha Kale, Dr. Ranjan Basak, Dr. Sunil Pabbi,
Dr. D.J. Bhagyaraj and Dr. Alok Adholeya have enriched this effort by their
valuable guidance and by sharing their experience . Shri Gautam Bose, Shri Amit
Kumar and Shri Pradip Nair of Tellywallah have worked hard and extensively to
make it in this present form.

We hope that this great effort will be used by the rural agro-biotechnologists,
whose services, we believe, will bring a new dawn to rural India.

Nimpith Swami Sadananda
July 2004 Chairman

Preface

In today’s world, technology is moving very fast, in certain sectors at a faster
pace. Biotechnology is such an area. The rural India which is depending on
agriculture for its day to day life provides an immense market for new technologies.
The only condition is the proper training and marketing.

The present manual is the outcome of the project ‘Technical Human
Resource Development – Vocational Training For Employment Generation’
supported by UNDP & DST, Govt. of India. The objective of the programme is
to develop human resources through competency based training in innovative
areas for the production and service sectors in new, high technology areas based
on market needs.

Agricultural biotechnology is the area in which we have worked on under
this project. This manual is first of its kind in this series. It deals with Biofertilizer
Technology which has six modules - Soil testing and fertilizer recommendation,
Production and Application of Blue-green algae, Azolla, Microbial inoculants,
Vesicular Arbuscular Mycorrhiza and Vermicompost.

The target are the rural youth, who have passed their 10th std. It is
designed and presented in such a way that a complicated subject like soil testing
or microbial inoculant production becomes an easily adaptable skill, a demystification
of technology indeed. All the techniques mentioned here are of world standard
but no doubt many other options can also be opted for, for instance, in the
section of microbial inoculant production only the use of vessel is mentioned in
this manual though the use of other types of fermentors or shakers are also
possible.

This effort is the result of hard work of a team of our Institute, the same
was complemented by experts of other organizations, which are of world repute.
Mr. Khudiram Sardar, Mr. Diwakar Haldar, , Mr.Tarun Das, Mr. Deepankar Haldar,
Mr.Tapan Haldar and Mr. Shubankar Malik have helped while filming the experiments
for this manual.

Shri Gautam Bose , Shri Amit Kumar, Shri Pradip Nair and Shri Partha
Bhattacharya of Tellywallah gave their best to make this dream a reality by
filming, designing, adding inputs and finally printing the manual.

Nimpith Dr. B.K. Datta
July 2004 Principal Scientist, VIB

Using this manual

This manual has been written to complement classroom lectures pertaining to
the Biofertilizer Technology Section of the “Certificate Course in Agricultural
Biotechnology” taught at the Vivekananda Institute of Biotechnology — VIB.

However, it may also be read by anyone with school level knowledge of
Science who is interested in setting up a soil testing lab or a microbial inoculant
production facility.

We have departed from traditional styles of writing Course material.
Instead of dry and forbidding lists of procedures and equations, the manual tries
to expose the student to the various aspects of the three disciplines that this
course straddles — Microbiology, Agricultural Science and Chemistry. Of course,
one cannot entirely avoid equations and procedures! But we felt it is possible
to present these topics in a friendlier manner.

Our approach to making the subject “friendlier” and also maintaining
sufficient scientific detail was two-pronged.

We’ve split the material being discussed into two parts — quite literally.
The odd-numbered pages in this manual contain descriptions of experiments
and processes in a step-by-step manner, devoid of detailed explanations — rather
like a traditional lab manual which expects students to follow the steps described
bothering to evoke an interest in the topic at hand.

The objective of the VIB course was the opposite. As scientists we are amazed
and intrigued by our respective fields of study. We wanted the reader to feel this
amazement as well; to undestand that science is an exciting subject!

To help us out, we enlisted the help of Shubham, an inquisitive,
imaginary friend, who can’t stop asking questions. Shubham can
be found loitering around on the even-numbered pages — asking
questions about the text on the facing page. The even-numbered
pages also contain supplementary information about the topic at
hand. Definitions, little bits of trivia, brief forays into the history
of science, suggestions for further reading, tips on how to simplify
a process and so on may also be found on even numbered pages.

The course material was meant to enable an interested student set up
his own Soil Testing Centre and an Inoculant Production Centre. Which is why
the material in this manual is presented in a “modular” fashion — typical of
industrial processes. Therefore, operations like autoclaving, working in a Laminar
Flow Cabinet and using a Fermentor have been dealt with in a separate section.
The step-by-step description concentrates on following the process being discussed
segment-wise.

The line at the bottom of each odd-numbered page is a “station map”.
This helps the students to see the “whole railway line” and on which station he
stands — this is typical of any industrial process where one goes from
A (A test tube of bacteria) to B (application of the inoculant in the field) and the
various stages that must be crossed to do so.

Using this manual

In sections where extra explanation on the left hand pages was unnecessary,
the step-by-step description continues.

Addressing the reader directly (”You could then....” “We might look at
this from another angle....”) is something most textbooks never do but we felt
there was no reason why the manual should not attempt to make the reader
feel as if he was in a classroom while reading it.

The manual is accompanied by an interactive CD-ROM. This contains all
the course material viewable in a non-linear fashion. This enables a student
to quickly refer to topics without having to flip through the manual. Further, all
the “modules” that are part of the production process, were filmed during the
making of this manual and the footage, with narration, is available on the CD.
This could be used in the classroom too, during, say, the first time the topic
is brought up for discussion by the instructor.

The glossary in this manual is available in a searchable format on the
CD-ROM.

The CD-ROM contains two computer programmes written specifically for Soil
Testing Centres. Chemicalc and X-Base are a scientific formula calculator and
a Soil Database, respectively. X-Base allows Soil Testing Centres to share their
data over the Internet with no additional software or hardware requirements
apart from a very rudimentary computer system and a telephone line.

A few of the pages also contain material that is not immediately relevant
to the course (eg., a brief description of the gene cascade in R.Meliloti that leads
to the production of nitrogenase, may be found in the pages that describe
Rhizobia). These portions are italicised. These sections are rare and are only
found, if at all, in the introductory part of the section. These could, if nothing
else, be food for thought for an inquisitive reader.

Biotechnology is the fastest growing scientific field around the world. As
scientists around the world learn more about the amazing internal workings of
living things, literature — this manual included — needs to be updated with a
regularity, which, to a scientist, is more exciting than monotonous. The authors
would greatly appreciate any feedback about this manual.

VIB hopes that this initiative will help fuel the next agricultural revolution
in our country — one powered not just by fertilizers and technology but also by
a more aware and knowledgable farmer who understands the science behind
the word 'biotechnology'...

VIB team

Nimpith
July 2004

Soil Testing - Collection and preparation of soil samples

Composite soil samples, packed for lab. analysis 1-1

Introduction..

A bit about the Soil Collection Process

The process of soil testing begins in the field — where collection of samples
is done. This needs to be done methodically — to ensure that the samples
taken from the field represent the soil characteristics of the entire field.
The process stated here is simple and easy to follow — though it does seem
elaborate when you first read about it!
Locating areas to take samples...

Start at the bottom -left corner of the field and walk
along the path indicated by the arrows. Along the
way, mark areas (with a piece of wood — or anything
easy to locate) from where you will be taking soil
samples. Avoid marking prohibited areas (see the
next page).

The marked areas are shown in Red in the figure.

Notice, that the red dots along the path are not
necessarily on the path itself.
That’s perfect — because this is a random sample.

Apparatus required

Plastic bucket, spade and wooden stakes or markers.

Why is it necessary to go through such an elaborate process?

It’s quite simple — ideally you’d want to test all the soil in your field
— but that would take a lot of time. So, it becomes necessary to take
samples from different places in the field.

But how do you ensure that the samples taken to the testing centre
represent, more or less, the soil in the field? The answer is the reason
behind the elaborate process.

By taking samples from randomly selected spots, you ensure that
your cumulative soil sample represents the soil in most areas of the field. The
zig zag walk is just to establish a technique of selecting random spots.

1-2 Collection and Preparation of soil samples

Collecting soil samples

The locations from where soil samples
are taken may be marked by wooden
stakes as shown. The stakes need not
be numbered.

The locations are chosen by walking
along a zig-zag path starting at any
corner of the field. These are marked by
white spots in the picture.

A typical location is shown in the next
picture.

Each collection location needs to be
cleared of vegetation and dust. This is
done by scraping away a very thin layer
of soil.

Collection Grinding Partitioning Storage

Drying Mixing Sieving

Collection and Preparation of soil samples 1-3

Prohibited samples...

Prohibited Samples?

It’s just another thing to do with the statistics (see page
2.1). Your soil samples need to represent the average soil
characteristic of the field. However, every field has areas
in it that tend to distort the average characteristic because
the soil there has properties very different from the rest
of the soil.

(It’s called deviation from the mean, by the way)

So, you need to ignore these areas when collecting samples.

These include

— Areas near gates, farmways, buildings etc. and areas on crop hills and in rows.
— Areas where organic/chemical manure is or was kept.
— Areas which are permanently in the shade.

Why must the pit be 6” (15 cm) deep?

6” is the depth of tillage i.e., the depth up to which the root system of the crop
penetrates. Since soil testing is carried out to determine the availability of
nutrients to crops, samples i.e., the furrow slices, are 6” long.

Why should a PVC bucket be used?

An iron bucket may have rust, which might contaminate the sample. Presence
of iron would skew the results of the organic carbon test as we shall see later.
Jute or nylon bags which may have been used to store fertilizer etc. must not
be used too.

PVC (or any other plastic) would not contaminate the soil. Besides, they are easy
to clean.



Remove a wedge-shaped lump of soil
from the cleared sampling location and
discard it.

The resulting pit should be about 6 inches
(15 cm) deep.

From the two larger surfaces of the pit,
remove a half inch thick slice of soil —
called the furrow-slice. Thus, from each
sampling location on the field, two furrow-
slices are obtained.

Carry these in a PVC bucket.





After collection — what now?

You now have a bucketful of furrow-slices. However, it would be time-consuming
to test all the soil in each of the furrow-slices separately and then average the
results. Ideally, you would want to test a relatively small sample of soil which,
nevertheless, represented the soil present in the entire field.

This is achieved by thoroughly mixing the soil. Further, to simplify the lab
experiments, the samples are ground and sieved... We’ll discuss this issue in the
next section...



After all the samples are collected, they
could be taken directly to the laboratory
— if one is close enough. Else, the
samples may be prepared on location
itself, and a composite soil sample may
be sent for testing.

The samples are now air-dried in the
shade.




Drying the soil samples

Why are the samples dried?

In dried soil, any reversible chemical reactions that usually take place in it are
in equilibrium.

However, drying does change the chemical constituents of soils. Ferrous iron
is oxidised to ferric iron, exchangeable potassium content increases or decreases
depending upon the soil and hydrogen ion activity changes to some extent.

Therefore, concentration of ferrous iron (if required) should be determined with
a field-moist soil sample. Also, concentration of exchangeable potassium and
soil pH may also be determined without drying the sample.

Dried soil is easier to grind.

For reasons mentioned earlier, metallic apparatus must not be used for grinding
as particles might break off and contaminate the soil.

Therefore, a wooden mortar and pestle must be used.


Grinding the soil samples

Spread a clean polythene sheet on the
ground and place a wooden mortar on
it. Transfer the dried soil samples to the
mortar.

With a wooden pestle, grind the soil to
break down any aggregates.

After grinding, transfer the soil to the
polythene sheet and spread evenly across
the surface.




Mixing

Why is mixing important? After all, we just ground the samples — that
should have mixed everything quite well.

Mixing is necessary to ensure that the composite soil sample (all the soil samples
taken from different areas in the field) represents the field’s soil composition as
closely as possible even in small quantities.

For example, to determine the amount of phosphorus, as little as 2.5 g of soil
is used in the experiment. You would have collected nearly 3 kg of soil from the
field of which only about 500g of soil is sent to the soil-testing centre.

Therefore, unless properly mixed, it is likely that soil from some parts of the
field might not reach the testing centre.

Yes, during grinding, mixing does take place — but it is random and might not
be enough. Besides, like the collection stage, the mixing stage described here
is to establish a process that minimises chances of statistical error.


Mixing

Lift one end of the sheet and fold it till
the soil collects in the centre. Repeat the
process with the diagonally opposite
corner as shown in the picture... and
then with the other two corners...

This will cause the soil to collect in the
centre of the sheet. This process is called
Mixing and must be repeated 5 times.

Cone the soil and flatten the top.




Partitioning the soil sample

Partitioning?

At this stage, you have about 3kg of a composite soil sample.
This would be a rather unwieldy package to transport to the
testing labs. Besides, the lab will need only about 300g of soil
for all the tests.

Here is where the purpose behind the monotonous mixing process
becomes apparent. Because the sample is mixed, statistically,
the average chemical composition of the field is represented by
surprisingly small amounts of soil — as little as a few grams!

So, you don’t need to send in all the soil you’ve collected painstakingly — you
could send in 500g which is relatively easier to transport and store.

There is a catch though — selecting the final soil sample must also be done at
random . This is fulfilled by the next Stage — Partitioning.

This involves halving the mass of the soil sample in successive stages till it is
about 500g. At each stage, a random portion of soil is selected.

The method described here is one of many that may be used to quarter the
soil sample. There are others, such as : The Riffle Technique and The Paper
Quartering Technique.

However, the method described needs no extra apparatus and is very easy.
Hence its use in this manual.



Divide the soil into 4 equal portions, as
shown.

Discard any two diagonally opposite
portions.

This process is called Partitioning.





Continue partitioning the sample till about
500g of soil remains.

The amount of soil retained depends
upon the number of experiments that
the Soil Chemist intends to conduct. 500
g is enough if all tests are to be carried
out.

Transfer all the soil to a sieve.




Sieving the soil sample

What mesh size is to be used?

A fine sieve (80 mesh) is used for determination of oxidisable
organic carbon and the elements i.e., Nitrogen, Phosphorus and
Potassium.
A coarse sieve (20 mesh) is used for determination of soil pH
and salinity.

The entire volume of the partitioned sample should pass through
the sieve. Soil aggregates that are too large to be sieved should
be ground in a mortar and sieved again.


Sieving the soil sample

Sieve the soil.

Preparation is complete. The soil particles
have been ground, mixed, partitioned
and sieved.

Before the soil is transported to the lab,
it must be stored in polythene bags.




Packing the soil sample


Packing the soil sample

Seal the open end of the bag with thread.

Label the bag. The information required
by the laboratory for testing is shown in
the Information Sheet (see Page 9-28)

The samples are now ready for transport.
The picture shows 3 polythene bags —
these are samples from adjoining fields
all headed to the testing centre.




Soil testing - Determination of soil pH

pH electrodes dipped into a soil-water suspension for pH measurement 1-21

Introduction

A bit about pH...

pH is the quantitative measure of acidity or alkalinity of liquid
solutions. A solution with a pH value less than 7 is considered
acidic and a solution with a pH more than 7 is considered alkaline.
pH 7 is considered neutral.
Soil pH between 6 and 8 is safe for most crops. If the tested
sample has a pH value outside this “safe range”, steps must be
taken to artificially correct the problem.

The acidity of a solution is directly proportional to its hydrogen
ion concentration.

The term pH is derived from p representing the German word potenz, ‘power’,
+ H, the symbol for hydrogen.

pH meters are extremely sensitive instruments. They consist of one (or two)
glass electrodes connected to a digital display. The pH of a solution is displayed
when the electrodes of the meter are dipped in it.

Soil water suspension : A suspension, as opposed to a solution, is a
heterogeneous mixture, i.e., its constituents may be separated by physical
means. The mixture of soil in water is therefore a suspension, not a solution.

Apparatus and reagents required
Buffer tablets of pH 4.0 and 7.0, a top loading balance, a 100mL beaker, a
wash bottle, a glass rod and a pH meter

Determination of Lime Requirement : An acidic soil is treated with Lime to
increase its pH. Recommendation of liming is also done after a pH test. The
difference being the addition of an extra ingredient to the soil.

Take 5g of soil instead of 20g as shown here. Add 5mL of distilled water and
then add 10mL of SMP Extractant Buffer.

Proceed with the pH exactly as shown in the following pages. When you obtain
the pH value, refer to the Lime Recommendation Table on page 9-28.

1-22 Determination of soil pH

Preparing a soil suspension

Weigh out 20g of soil.

Transfer the soil to a 100mL beaker.

Add 50ml of distilled water. This creates
a soil-water suspension.

The soil : water ratio for conducting this
test should be 1 : 1.25

Stir the suspension occasionally for about
half an hour or shake in a shaker for 5
minutes.

Preparing the soil suspension Measuring the pH of the sample

Calibrating the pH meter

Determination of soil pH 1-23

Using the pH meter...

Calibration? Why is it necessary?

Think about this — how does the meter know a solution’s pH? It doesn’t. It’s
just programmed to display different pH values depending upon the voltage
across its electrodes.

The electrodes, though, are sensitive to a whole lot of other things — like
temperature for instance. So, even though a change in room temperature will
not change the pH of a solution, it will cause the electrodes to report the pH
incorrectly.

And this is true of any measuring instrument. We calibrate by measuring a known
amount and then re-programming the meter to display that amount — this is
sometimes as simple as pushing a switch.

In this experiment, we calibrate the meter with 2
buffer solutions — with pH values of 7.0 and 4.0.
See page 1-26 for a description of buffer solutions.

First, we dip the
electrodes in the pH
7.0 buffer. While we
were conducting the
experiment, the
meter read 7.2. This
was because it was set to measure correctly at a
slightly lower room temperature. So, the adjustment
knob was turned till the display read 7.0.

Between readings, wash the electrodes with distilled
water and wipe them dry with a piece of clean
tissue paper.

Repeat the process with a buffer solution of pH 4.0. The volumetric flask in the
third picture contains the pH buffer solution.

The meter might need a few minutes to “warm
up”. The time varies from model to model and you
should check the literature that came with your
meter. Generally, “warming up” takes a few minutes.
Also, it takes a few seconds for the display to
stabilise after you’ve dipped the electrodes in a
solution. So, wait a while before noting down a
pH reading.


Calibration

Calibrate the pH meter with any two
known buffer solutions.

Just prior to taking any readings, stir the
soil-water suspension with a glass rod.

Dip the electrodes of the pH meter into
the suspension and take a reading.




A buffer story

A bit about Buffer Solutions

We know about pH and how it describes the acidity or alkalinity of a solution.
Now, why is a solution acidic? Or alkaline? Modern definitions of acidity refer to
the ability of the compounds in a solution to accept or release electrons — the
Lewis Concept.

But historically, an acid was a compound that, in solution, could release hydrogen
ions into the solution, and a base was a compound which could accept hydrogen
ions.

Since HCl dissociates into H+ and Cl- ions, in solution, it is an acid.

There are situations when we want the pH of a solution to remain constant —
irrespective of change in the concentrations of its acidic or alkaline constituents.
This is done by adding an acid-base pair to the solution that acts as a reservoir,
or buffer. What this reservoir does is suppress or increase the dissociation of
other compounds depending upon the pH of the solution. A commonly used
buffer solution is the NH4Cl - NH4OH pair. These are readily soluble chemicals
and keep each other’s dissociated concentrations in check — in line with the
solubility product principle.

The NH4Cl - NH4OH pair is an alkaline buffer and maintains the pH of the solution
at around 8.5.

You can read more about Buffer Solutions in books on Physical Chemistry.

A thin “chemical film” is deposited on the electrodes each time a pH measurement
is taken. Unless removed, this film causes the electrodes to report inaccurate
values. Therefore, between readings, wash the electrodes of the pH meter with
a stream of distilled water and then wipe them dry with tissue paper.

When not in use keep the electrodes dipped in distilled water.


Meter readings...

The pH of the sample tested is 7.79. The
display on most pH meters takes about
a minute to stabilise.




Soil Testing - Determination of salinity

Conductivity readings are taken from the supernatant liquid.. 1-29

Introduction

A bit about salinity...

The determination of the quantity of water-soluble salts is
of special importance for arid, semi-arid as well as coastal
areas. It helps in taking reclamation measures as well as in
the selection of crops which differ in their tolerance to salts.

While we could measure the concentrations of the salts by
chemical analysis, it would be time-consuming, expensive
— and entirely unnecessary. That’s because, we don’t need
to identify all the salts lurking about — we’re only interested
in the water-soluble ones. The concentration of all of these salts taken together
is what matters to plants.

So, we take a more practical approach to measure soil salinity — we add distilled
water into the soil and stir it till the soluble salts get dissolved. Then, we measure
the electrical conductivity of the water.

And how does that tell us anything about salinity? Indirectly, it does — because
in solution, ions are the carriers of electric charge and therefore, the electrical
conductivity of a solution is directly proportional to its soluble salt concentration.

The conductivity of a soil sample is measured with the help of a conductivity
meter and is expressed in mmhos/cm. or, in SI units, in dS/m.

You don’t even need to calculate the concentrations for the purposes of
recommending fertilizers. The recommendation is based upon the conductivity
measurement itself. ( See the recommendation tables). Most soil testing labs
mention “Electrical Conductivity” or just “E.C.” in their reports.


0.01N KCl solution, a 100mL beaker and a conductivity meter.

Using a conductivity meter...

An electrical conductivity meter is very similar to a pH meter. It also consists
of an electrode connected to a digital display. Measurements are made by dipping
the electrode in the solution being tested.

The precautions to be observed while using this instrument are the same as
those with a pH meter (see page 1-24 ).

1-30 Determination of salinity

Calibrating the conductivity meter

This is a typical Conductivity Meter. Like
the pH meter, its electrode must be kept
immersed in distilled water when not in
use.

Calibrate the meter. This is done with
distilled water and a 0.1N Potassium
Chloride solution.

Distilled water should display 100 in the
digital panel.

Then, dip the electrode in a 0.1N KCl
solution.

Calibration of the Conductivity meter Measuring conductivity

Supernatant Liquid

Determination of salinity 1-31

A bit about supernatant liquid

What is Supernatant liquid?

The liquid that floats above a suspension after it
has been allowed to stand for a while is termed
“Supernatant”. When measuring conductivity, the
conductivity cell should remain in the Supernatant
liquid and not touch the soil below as shown.

So, why are we measuring the conductivity of the
supernatant liquid? After all, during the pH
experiment we had specified that the soil should
be in suspension while the reading was being taken.

The supernatant liquid is a solution. The salts present in the soil dissolve in water
and dissociate into ions, which are charged particles. The concentration of soluble
salts in the soil may, therefore, be calculated from the conductivity of the
supernatant liquid.

Soil - water ratio? Why is that important?

Soil to water ratio should be 1:2. The ratio influences the amount of salts in the
extract. Some laboratories use different ratios while conducting this test. Either
way, the ratio must (and is) always mentioned in a soil analysis report.


Conductivity readings...

Adjust the cell constant knob till the
meter displays 14.1 m.mhos/cm.

The soil water suspension from the pH
experiment is allowed to stand till a clear
supernatant liquid is obtained.
After setting the range switch to
maximum, the electrode is dipped in the
supernatant liquid.

Reduce the range setting on the meter
one at a time till the most appropriate
setting is found.

In this case, the conductance of the soil
sample is 1.20 m.mhos/cm.

Calibration of the Conductivity meter Measuring conductivity

Supernatant Liquid

Determination of salinity 1-33

Soil Testing - Determination of available organic Carbon

Diphenylamine indicator being added drop by drop 1-35

Introduction

A bit about Oxidisable organic carbon...

Decomposed plants and microbial residues are the constituents of
organic matter. The percentage of oxidisable organic matter can
be determined by multiplying its percentage of organic carbon by
1.724.

Oxidisable organic carbon consists of partly decomposed residues
of plants, animals and microorganisms. This constitutes most of
the usable carbon present in the soil.

The other forms of carbon which are present, but not useful as a source of
nutrients, include inorganic carbon (such as carbonates), elemental carbon (such
as coal and graphite) and completely decomposed organic carbon.

For areas known to have very low organic matter content take 2g of soil in the
conical flask, for peat soils, take 0.05g and for areas known to have about
1-2% of organic carbon content, take 0.5g of soil.


1N potassium dichromate solution, 0.5N ferrous ammonium sulphate solution,
diphenylamine indicator, concentrated sulphuric acid and 85% orthophosphoric
acid solution.

500mL conical flask, titration setup (50mL burette, chromyl chloride solution
to clean the burette and titration stand) 10mL bulb-type pipette, chemical
balance, 1000mL volumetric flask and two watch glasses.

Why are two conical flasks used?

This experiment is based upon the Walkley and Black method according
to which soil is digested with chromic acid resulting in the oxidation of
its organic content.

The excess chromic acid is determined by titration with a standard ferrous
ammonium sulphate solution.

After titration, in the case of the soil sample, the amount of titrant consumed
is obtained. The amount of titrant consumed in a blank titration (without soil)
could be calculated stoichiometrically. But this would require accurate weighing
of all the reagents involved in the reaction.

Therefore, it is much simpler to perform a blank titration to obtain the required
figures i.e. the volume of ferrous iron solution consumed.

1-36 Determination of oxidisable organic carbon

Oxidising the carbon in the soil sample

Take two 500mL conical flasks.

Add 1g of soil to one of the flasks.

Add 10mL of K2Cr2O7 to each of the
flasks with a pipette.

Oxidizing the carbon in the soil Titration of the soil sample suspension

Blank Titration Calculations

Determination of oxidisable organic carbon 1-37

Precautions...

Concentrated sulphuric acid is a very corrosive chemical. It fumes in contact
with moisture. Observe the following precautions when using sulphuric acid :

The acid must be poured into the beaker along a glass rod or along its inner
walls.

DO NOT use a pipette to measure out the acid - if any of the acid gets into
your mouth, there might not be enough of it left to talk about the experience!

This step should ideally be carried out in an Exhaust Cabinet because the fumes
are extremely corrosive as well. Don’t try to smell the fumes however tempting
it might seem!

What happens during the half hour?

The oxidisable organic carbon in soil is oxidised by potassium dichromate

3C + 2K2Cr2O7 + 8H2SO4 = 3CO2 + 8H2O + 2K2SO4 + 2Cr2 (SO4)3

Potassium dichromate is converted to potassium sulphate and chromium sulphate.
Cr6+ is reduced to Cr3+. The colour of the oxidised form of chromium(Cr6+) is
yellow (or orange) and that of it’s reduced form (Cr 3+ ) is green.

The volume of K2Cr2O7 solution added to the soil should be large enough so that
only a small fraction of it is reduced - which is indicated by yellow (or orange)
c o l o u r o f t h e r e a c t i o n m e d i u m a f t e r c o m p l e t i o n o f ox i d a t i o n .

This occurs over a period of 30 minutes.

Because sulphuric acid fumes, reagents
might get deposited on the watch glass.
Therefore, when adding water, if you see
flecks of chemicals deposited on the watch
glass, rinse them and allow the water to
drip into the conical flasks.


Oxidising the carbon in the soil sample

With a measuring cylinder, add 20 mL.
of concentrated sulphuric acid to each
of the flasks.

Cover the flasks with watch-glasses and
allow them to stand for about half an
hour.

Then, add about 200 mL of distilled
water to each of the flasks.




Titration

A bit about titration...

Titration is a process by which the amount of an oxidisable or reducible substance
in solution is determined by measuring the volume of a standard reagent required
to react with it.

The burette used must be cleaned with chromyl chloride prior to titration. Dirty
burettes are the most common cause of errors.

Carry out the blank titration first. This will give you a general idea about the
volume of titrant, i.e., Ferrous ammonium sulphate that will be consumed. With
this value in mind, the titration of the soil sample usually takes less time.

Observe the colours that the solution assumes during the process. In the first
phase, the solution is a dark burgundy. After a while, it turns violet. This indicates,
approximately, the mid point of the titration. Local action is also observed at
this point. The remainder of the titration needs to be carried out carefully, i.e.,
by agitating the contents of the flask after every 2 drops.

The end-point is indicated by a sudden change of colour of the solution to viridian,
or dark green.

During the titration of the soil sample, all the indicative colours are more cloudy
than those observed during the blank titration. This is due to suspended soil
particles.

Why is Orthophosphoric acid used?

Orthophosphoric acid, H3PO4, is added so that the colour change at
end point is clearly defined.

Diphenylamine should be added just prior to titration. This is to avoid the
potassium dichromate from oxidising the indicator instead of the organic
content of the soil sample being tested.
During titration, a small amount of diphenylamine is oxidised, however, the
error is negligible.


Blank Titration

Add 10mL of orthophosphoric acid to
each of the flasks.

Just prior to titration, add about 10 drops
of Diphenylamine indicator to the flask.
The solution turns a dark burgundy.

The blank titration is done first.

Titration is done with a 0.5N Ferrous
ammonium sulphate solution.




Local action... And a few calculations

What is local action?

Local action is a phenomenon observed midway during titration. At
this stage, even though the titration is not complete, a faint, localised
“end point” may be observed in the solution where the titrant drops
fall.
To observe local action during this experiment allow a drop of titrant,
i.e., Ferrous ammonium sulphate, to drop on the solution without
agitating the flask as is normally done during titration.The solution
in the immediate vicinity of the drop turns green momentarily.

A few Calculations...

The CD-ROM has a Chemical calculator that does all the work for you but
since we’ve set out to understand the science behind Agricultural Biotechnology,
let’s dive headfirst into yet another bout with theory — and learn a bit about
Stoichiometry.

The Appendix contains an article that explains why you need to add and divide
all these numbers...

The percentage of oxidisable organic carbon (%OC) in the soil sample is given
by

% O.C. = [VK x (1– VS/VB) x SK x 0.3] / W

where

Vk = Volume of Potassium dichromate solution

VS = Volume of Ferrous iron solution consumed in titration with soil

VB = Volume of Ferrous iron solution consumed in blank titration

Sk = Strength of Potassium dichromate solution

W = Weight of soil


Indicative colours during titration...

Midway through the titration, the colour
of the solution turns to clear purple.
At this stage, local action may be
observed.

End point is indicated by a sudden change
of colour to viridian, or dark green.

Titrate the soil sample as well.

Notice that all the indicative colours with
the soil sample are cloudy.

The picture shows the titrated soil sample
suspension at end point.




Soil Testing - Determination of available Nitrogen

Ammonia bubbling up the neck of a Kjeldahl flask 1-45

A bit about the Experiment

A bit about the Experiment

Plants generally take up nitrogen as nitrate under aerobic conditions.
In anaerobic situations, some crops, such as rice, can take up
nitrogen as ammonium ions. Most of the nitrogen present in soil
is present in complex compounds. This is considered as a potential
reserve source and, as such, it may be measured to assess the
nitrogen-supplying capacity of the soil.

Soil testing centres do not usually conduct a separate test for
determining the quantity of available nitrogen in a soil sample brought to them
for testing. Instead, they calculate this quantity directly from the quantity of
oxidisable organic carbon.

And how exactly is that possible? The ratio of the amount of oxidisable organic
carbon is proportional to the amount of nitrogen in a given area. The ratio is
unique to each region. These ratios have been tabulated. In Nimpith, where VIB
is located, the ratio is 1 : 5. With this value, we need only perform the organic
carbon test to determine the quatities of both nitrogen and oxidisable organic
carbon.


Boric acid solution, Mixed indicator, 0.32% potassium permanganate solution,
2.5% sodium hydroxide solution, liquid paraffin.

Kjeldahl flask(s), distillation setup, titration setup, 250mL conical flask and
a few glass beads.

Glass beads and liquid paraffin

These are used to reduce frothing and the formation of bubbles in the solution
when the flask is heated.

The bubbles may carry soil into the delivery tube and deposit them in the conical
flask connected to the other end of the tube. The presence of soil makes it hard
to detect the end point when we titrate the contents of the conical flask. More
on this topic later...

1-46 Determination of Available Nitrogen

Extracting the Nitrogen as Ammonia

Take 20g of soil in a Kjeldahl Flask.

Add 20mL of distilled water.

Coat a few glass beads in liquid paraffin
and put them in the flask.

Extracting the Nitrogen as Ammonia Calculations

Titration of the condensate

Determination of Available Nitrogen 1-47


A bit about Kjeldahl and the flask he invented...

A Danish chemist called J.G.C.T. Kjeldahl came up with the brilliant idea of
estimating nitrogen concentrations in organic substances by distilling it out as
ammonia — which can be easily assayed. For boiling the organic substance, he
made a round-bottomed glass flask with a long neck. A special heat-resistant
glass is used which does not crack when heated to high temperatures and is
expposed to relatively cooler liquids at the same time.

The entire assembly is called a Kjeldahl setup or unit and the flask also bears
its inventor’s name.



Then add 100mL each of 0.32%
potassium permanganate and 2.5%
sodium hydroxide solutions.

Heat the flask to about 80oC on an electric
heater.

Ammonia is evolved. The gas escapes
into the delivery tube attached to the
Kjeldahl flask.





Why do we use mixed indicator?

In this test the pH changes at two distinct points. The first is when
the ammonia is absorbed by the boric acid and the solution changes
from bright pink to green. The second occurs during the titration
of the solution with sulphuric acid. The solution then changes back
to pink.

These changes occur at different pH values and a single indicator
is not sufficient since indicators exhibit a colour shift only in a
small pH range. Thus, we need two indicators which will show us both these
changes. Hence a mixed indicator — which is a mixture of Methyl Red, Bromocresol
green and Ethanol — is used in this experiment.



The evolved gases condense and are
collected in a conical flask containing
Boric Acid solution and Mixed Indicator.

The Ammonia is absorbed by the acid —
indicated by a change in colour of the
solution to green. Continue boiling the
contents of the Kjeldahl flask till about
100mL of distillate is collected.

Titrate the distillate with 0.02N sulphuric
acid.





My end point is brown!

You did not pour in enough paraffin. Or, perhaps, you didn't use
enough glass beads. These ingredients are added to reduce the
surface tension of the solution in the Kjeldahl flasks. This greatly
reduces bubble formation...

The bubbles often carry small amounts of soil and deposit it in
the conical flask. This “muddies” the indicative colours during
titration and hence the end point appears brownish...

A few calculations

Substitute the observed volume, V, of sulphuric acid consumed in the following
equation to calculate the amount of available nitrogen (in kg per hectare) of
the soil sample —

V X 31.36 kg/Ha


Titration of the condensate... and Calculations

Local action is observed distinctly during
titration.

End-point is indicated by a change in
colour from green to a brownish-pink.

Note the value of sulphuric acid
consumed.

Carry out a blank titration — with the
contents of the conical flask
corresponding to the Kjeldahl flask
without soil.




Titration of the condensate... and Calculations


Soil Testing - Determination of available Potassium

Flame view - the orange-red colour indicates the presence of potassium 1-55

A bit about Potassium

A bit about potassium

In soil, potassium may be found in four compound forms - Water
soluble, Exchangeable, Fixed and Lattice-bound. Of these, plants
are interested only in the first two since they cannot assimilate
potassium when it is present in the last two types of compounds.
Potassium is the most abundant meta-cation in plant cells. Oddly
though, soil humus furnishes very little potassium during
decomposition. Also, it occurs in plants only as a mobile, soluble
ion, K+, rather than as an integral part of any specific compound -
but, it is known to affect important aspects of a plants life such as
cell division, formation of carbohydrates, translocation of sugars and resistance
of the plant to certain diseases. Over 60 enzyme actions are known to require
potassium for activation.

Which is why it forms the “K” part of the NPK trio - the three major important
elements that plants require for proper growth. Incidentally, the “K” comes from
“Kalium” which is what potassium used to be called.

And a bit about the experiment

In the next experiment, when testing for phosphorus, we will learn about a
technique called curve fitting. This experiment also uses the same principle but
the curve fitting itself is done electronically by the machine itself.

So what machine are we talking about? It’s called a Flame Photometer. It consists
of two parts — the gas compressor (that’s the first picture on the right) and the
Aspirator/Measurement unit (the second picture). The principle on which this
gizmo operates is that every element, when burnt in a flame, emits energy in
a set series of wavelengths. Simply put, each element burns with a different
colour. Further, the intensity of colour is directly proportional to the concentration
of the element. So, by measuring the intensity of colour of flame aspirated with
the sample, and comparing it with a known set of colour intensities, the photometer
c a n d e t e r m i n e t h e c o n c e n t ra t i o n o f p o t a s s i u m i n t h e s a m p l e .

The gas compressor regulates the flow of LPG to the Photometer. The gas burns
with a nearly colourless (or faint blue) flame. The soil sample extract is then
sucked into the flame in minute quantites. This causes the water to vapourise
instantly and the compounds in it burn in the flame. The colour of the flame
changes depending upon the elements present in the extract. This is detected
by electronic sensors which calculate the intensity of the colour.

Potassium burns with an orange-yellow flame.

Apparatus and reagents list is on page 1-60

1-56 Determination of available Potassium

Calibrating the photometer

Set the Compressor to supply gas at a
pressure of 0.45kg/cm2.

Ignite the flame and then calibrate the
photometer with solutions whose
potassium concentrations are known.

The calibration must be done with 4
standard solutions.

Calibrating the Flame Photometer Measuring the concentration

Extracting the Potassium from the sample Calculations

Determination of available Potassium 1-57

Calibrating a Flame Photometer

Enter “Calibration” mode We’re working with pretty The smart meter now
using the control panel. high concentrations of realises that it needs to
The buttons you have to Potassium here. plot a Standard Curve. So
press will vary for different you need to tell it how
models. Some meters are designed many standard samples
for micro-analysis — like you’re going to use to plot
Usually, you’ll see a the one here. In our case, the curve. We’ll use 4
numbered list of options. we need to tell it to expect samples. Most meters are
In this case, we press “5” Potassium in high good enough to accurately
on the numeric keypad to concentrations. “fit a curve” with this
enter Calibration mode. number of samples.

It now wants to know the Sampling time! Aspirate The flame colour changes
concentrations, in ppm, of each of the stock solutions immediately to a bright
the standard solutions one by one — in the order orange-yellow. The
we’re going to use. Here, in which you keyed them change in intensity of the
we use solutions with 100, into the meter. We typed colour will be barely
75, 50 and 25 ppm in the concentrations in noticable to the naked
concentrations. descending order (100,75, eye. The meter however,
50,25), so, we’ll have to can distinguish each
a s p i ra t e t h e 1 0 0 p p m colour precisely.
solution first.


Calibrating a Flame Photometer

After each sample is The flame becomes After aspirating distilled
aspirated, the meter colourless when distilled water, repeat with the next
demands a “washing” with water is aspirated. It might Standard sample. Note
distilled water — just like also be a faint blue colour. that the display in the
the E.C. meter and the pH The picture here has been picture reads STD4, or
m e t e r. N o t i c e h o w deliberately modified to Standard Sample No. 4.
different measuring exaggerate the colour of
apparatus all have the the flame — so that you This was taken when we’d
same operating principles. can easily compare the already aspirated the first
It’s really quite simple — flame colours with and three samples. Now, we
no magic, just science! without the sample. aspirate the 4th sample.

Between each sample, the Calibration is over. Now, Shubham, it seems, has
meter will ask you to the meter is ready to test no question to ask on this
aspirate distilled water. the soil sample extract - topic and wants to get on
After all the 4 samples are we don’t know the with the experiment!
aspirated, the meter concentration of Potassium
sounds a satisfied beep in this. The meter will
and tells you happily that analyse the colour of the
calibration is over. flame, plot its density on
the standard curve that it’s
drawn for itself and tell us
t h e c o n c e n t ra t i o n o f
Potassium.




Extracting the potassium from the soil sample

Ammonium Acetate? Why is this used?

Like in the Phosphorus experiment, we need to find a way to
extract the Exchangeable Potassium from the soil sample. That’s
what Ammonium acetate is used for.

The ratio of soil : Ammonium acetate should be 1:5. That is, if
you used 5.0g of soil, take 25mL of ammonium acetate.

Ammonium acetate dissociates to yield ammonium ions

CH3COONH4- CH3COO- + NH4+
The NH4+ ions replace the K+ ions, held on exchange sites of soil colloids. As a
result, K+ ions are released into solution. Perfect for our purpose!

The chemical equation above has two arrows pointing in both directions. This
indicates a reversible reaction — i.e. one that occurs simultaneously in both
directions. However, each reversible reaction has an equilibrium point at which
the rates of both the forward and backward reactions remain constant.

Chemistry can be a really exciting subject! You can find out more about reversible
equations in any textbook on Physical Chemistry. See the Appendix for a list.


1N ammonium acetate solution of pH 7.0, 1000ppm potassium solution of
pH 7.0.

10mL pipette, 150mL conical flask with a rubber stopper, 50mL volumetric
flask, 100mL measuring cylinder, a funnel, Whatman no. 42 filter paper and
a flame photometer. This experiment is shown using a direct read-out electronic
flame photmeter.



Take 5g of soil in a 150mL conical flask.

Add 25mL of 1N Ammonium Acetate.
The pH of the solution should be 7.0.

Cork the flask and agitate its contents
for about 30 minutes. This can be done
with a mechanical shaker.





Whatman No. 42... Whatman No. 1... Who is this man called
What?

It’s a brand name. “Whatman”, the company, makes filter paper
— and a lot of other paper products used in Chemical analysis.
The paper is graded according to its relative porosity. Hence,
No. 1, No. 42 etc.

An interesting feature of these papers is that they are “ashless”.
This means that you can burn them and they do not leave behind
a residue. This property is useful in a lot of experiments — such
as gravimetric analysis, in which the filter paper is burned after it is used for
filtration thereby leaving all the precipitate behind for weighing. Neat!

Also, whoever started up the company was probably called “Whatman”... If that
helps at all....



Filter the suspension through Whatman
No. 42 paper.

The filtrate is used for determining
Potassium concentration. Transfer the
filtrate to beakers for use with the Flame
Photometer.

Aspirate the filtrate (the soil sample
extract).




Calculations...

The results and some calculations....

Easy as pie! The smart Flame Photometer tells you the concentration of
Potassium in the extract after politely asking you to wait for a while. In our
case, the concentration of Potassium was 22.3.

The K 2 O content of the soil is calculated using this formula -
K20 (in kg/hectare)= [ 2 X CK ppm X Ve ] / Ws
where

CK ppm = Concentration of the Potassium in ppm obtained from the
Photometer
Ve = Volume of Ammonium Acetate used

Ws = Weight of soil taken, in grams

You can use Chemi-Calc, the calculator on the CD-ROM, to do the calculations.
Or, if you want to show off a bit as well (like the Photometer did), use this trick

Multiply the CK ppm amount by 10!
Remember that the ratio of soil to extractant used should be 1 : 5. Which means
that if you followed the steps, you would have taken 5g of soil and 25mL of
Ammonium acetate. Those values then cancel out to give 5 in the numerator
part of the equation. Multiply that by 2(also in the numerator) and you get
[ CK ppm X 10]

But beware, the trick works ONLY if you measured out the Ammonium acetate
and the soil carefully to at least a couple of decimal places. So 25.01mL and
5.02g of soil is fine. But 25.5 mL and 5.2g of soil means your experiment
will be approximately OK but you cannot show off with the calculation trick!

The formula gives you the amount of K2O present in the soil (in kg per hectare).
To find the amount of Potassium multiply the result from the calculation above
by 0.83.

However, this is not necessary for recommendation since we are interested in
the K2O amount — Why? Because that’s the compound that Fertilizers Companies
refer to! They aren’t very smart, are they?


Measuring the concentration of Potassium in the soil sample

Presence of potassium is indicated if the
flame changes to a yellow-orange colour.

The minute colour differences between
the colours emitted by different elements
are not distinguishable by the naked eye.
The flame view is provided primarily for
adjusting the stability of the flame and
verifying that the nozzles of the aspirator
are not contaminated by residues from
previous experiments.

The concentration of Potassium, in ppm,
is displayed on the screen after a few
seconds.

Most soil testing labs are not equipped with direct-display Photometers like the
one used here. Older equipment requires the user to plot a Standard Curve
manually on Graph Paper. If such equipment is used, the process described in
Section 7 (the procedure used to determine the amount of Phosphorus) is
applicable here as well.




Soil Testing - Determination of available Phosphorus

The blue colour indicates presence of phosphorus 1-67

Preparing a Standard Curve

A bit about available Phosphorus...

Phosphorus occurs in soil in both organic and inorganic forms, most of which is
not easily available to plants. A portion of the total Phosphorus is absorbed by
plants during their growth in the form of H2PO4= . This is what we refer to as
available phosphorus.
What is a Standard Curve and what is it’s use?

Simply put, it is a graphical means of determining unknowns that are
variables of a linear equation. The catch is, that this is an equation of the
form y = n x, where n might vary randomly as x varies. So how is that
linear, you may ask. It is, approximately, because if we define n as n + e, then
we find that e is a very small positive or negative number.

The easier method to solve the problem, is to plot, on graph paper, a few (x,y)
pairs and draw ONE line — the Standard Curve — connecting as many points
as possible. If e was a relatively large number then we would have no option but
to resort to esoteric mathematical tools like regression analysis because then
the equation would cease to be approximately linear. But because e is a small
number, we would find that most points are either on — or very close to — the
straight line drawn and points are scattered almost equally on either side of the
line.

Then, to find the value of y for any given x (or vice versa) all you need to do is
find the corresponding point on the line. Thus, we’ve found the solution to a
linear equation by graphical means. Let’s call it the Graph Technique. Sounds
difficult? It isn’t. As an exercise, try plotting the following value pairs on a sheet
of graph paper

(x,y) = (0 , 0) , (1.1 , 1) , (4.9 , 5) , ( 6.2 , 6) , (7.3 , 7) and (11 , 11).

Draw a line that joins the points - you will find that the points are scattered to
either side of the straight line joining (0 , 0) and (11 , 11).

Using this line, find out the value of y when x is 8. You get y = 8. Now, the
actual value of y might not be 8 exactly, but the graph shows that it would be
pretty close to 8, if not exactly 8. In most cases, as in our present experiment,
the small error is negligible.

Olsen's extractant or Bray and Kurtz No. 1 extractant, Standard 100ppm
phosphate solution, ammonium molybdate reagent (containing antimony
potassium tartrate and ascorbic acid), 2,4-dinitrophenol, P-free charcoal.
Eight 25mL volumetric flasks, 10mL graduated pipette, 10mL measuring
cylinder, a funnel, graph paper, Whatman No. 42 filter paper and a colorimeter.

1-68 Determination of available Phosphorus


Take eight 25 mL volumetric flasks and
add 1mL, 2mL, 3mL, 4mL, 5mL and
10mL of 2 ppm Phosphate solution to
six of them.

Leave two flasks blank.

Add 5mL of a 2ppm Standard Phosphate
solution to one of the empty flasks.

Add 5mL of Olsen's Reagent to the same
flask.

Preparing a Standard Curve Measuring the concentration of Phosphorus

Extracting Phosphorus from the soil sample Calculations

Determination of available Phosphorus 1-69


A step-by-step description of the process...
First we need to plot a Standard Curve. In this test, it is plotted to determine,
approximately, the relationship between the intensity of colour and the needle
deflection of the Colorimeter (the reading on the scale) — the assumption being
that the relationship would be linear. It is linear, by the way... Which makes this
an ideal candidate for using the Graph Technique described on page no. 1-68.

Now, we need a few stock solutions of Phosphorus to plot the Standard Curve.
We take 6 such solutions. Then, we need to find out how much the Colorimeter
would read when the solutions are placed in it.

At this stage two problems crop up —

(1) The Phosphorus in the soil is present primarily as a Phosphate of Calcium,
Iron and Aluminium — which are all colourless. So the Phosphorus needs to be
extracted AND

(2) the new compound, in solution, must have a colour whose intensity varies
linearly with respect to its concentration.

The extraction part is done by adding either Olsen’s extractant which is a 0.5M
NaHCO3 solution or by Bray and Kurtz No.1 Extractant.
Ammonium molybdate reagent or "Reagent B" (which is a cocktail of Ammonium
Molybdate, Tartarate and ascorbic acid) is used to obtain a heteropoly complex
called phosphomolybdic acid, which is reduced, partially, by the ascorbic to give
a blue coloured solution. The amount of the complex produced is directly
proportional to the Phosphorus concentration. Problem solved!

The two flasks to which we did NOT add any stock solution are used to simplify
a technical hitch (see page 1-78) when Olsen's extractant is used.

The Flask with Olsen's reagent described in the facing page will be used to carry
out a mini-titration to determine the amount of acid required to lower the
solution's pH to 3.

The other flask will be used as a "blank". This lets us know if any of the chemicals
we are using contain Phosphorus as an impurity. Shubham will be lurking on
these pages to elaborate as we discuss the steps of the experiment!



Add two drops of 2,4-dinitrophenol
indicator to the flask.

The solution turns yellow.

Then, with a graduated pipette or
dropper, add 2.5M Sulphuric acid to the
flask till the yellow colour is discharged.

This indicates that the pH of the solution
is 3.

Note down the volume of acid consumed.
In our test, 0.4mL of acid was required
to lower the pH when 5mL of Olsen's
reagent was added.

This flask is not needed any more and
may be removed from the work area.

This leaves 7 flasks — 6 with standard
phosphate solutions and one empty flask
for the blank test.





More on the pH value

The value obtained in the previous step — 0.4mL of acid —
pertained to 5mL of Olsen's extractant. So, if we use, say, 50mL
of the reagent, we need to add 4mL of 2.5M sulphuric acid to
adjust the pH of the solution.

Since we've standardised the procedure (always using 5mL of
filtrate and so on) the value obtained earlier — 0.4mL — is used
throughout.



We now have seven flasks left.

To each of the flasks, add 5mL of Olsen's
reagent.

Add exactly 0.4mL of 2.5M Sulphuric
Acid to each of the flasks to adjust the
pH of the solution to 3.

Add approximately 10mL of distilled
water to each of the flasks.

Then, add 4mL of Reagent B — the
Ammonium Molybdate reagent mixture
- to each of the flasks.

The picture shows the reagent being
pipetted into the flask containing 10mL
of the Standard Phosphate solution.




Using a colorimeter

Using the Colorimeter

Colorimeters are usually analogue — measurements are indicated by the deflection
of a needle over a semicircular scale — and they look very ancient in the
laboratory, surrounded by digital equipment. However, they are very precise
instruments as well!

Colorimeters have a light-proof slot where the sample to be tested is inserted.
If the sample is opaque, a full-scale deflection is observed, whereas a transparent
sample does not cause any deflection of the needle.

Calibration

The zero and full scale deflections are set by
“measuring” a sample of distilled water and
a black, opaque cylinder (that comes as a
standard accessory with the meter and is
designed specifically for this purpose.)

Turn the adjustment knob to set the needle
to the zero when calibrating with distilled
water. With the opaque cylinder in the metering
slot, turn the adjustment knob so that the
needle deflects all the way and stops in front
of the “infinity” mark.

More on the Standard Curve...

After you’ve measured the optical densities of all the six solutions,
plot them on a sheet of graph paper.

The concentrations of Phosphorus in each of the flasks, in parts
per million, go on the x-axis, while the corresponding optical
densities go on the y-axis.

Draw a straight line that connects most of the dots. That’s your
Standard Curve! By the way, you might be wondering why it’s called the Standard
Curve when it is a straight line. Well, it would be a curve if the equation were
not linear. For instance, if the relationship between y and x were quadratic
( y = nx2 + c) then the graph would be a curve. Besides, (and this might sound
silly) mathematically, a straight line is also a curve!



Add distilled water to each of the flasks
till the volume of solution in them is
exactly 25mL

The 25mL value is to be maintained for
all solutions. The value determines the
concentration of the phosphate solution.

Remember that ppm is a measure of
concentration.

Allow the flasks to stand for 10 minutes.
The solutions assume a blue tinge.

Measure the optical densities of all the
solutions using a colorimeter.

The blank sample should register a "zero"
deflection while the sample with 10mL
of the standard phosphate solution should
register the highest optical density.




Phosphorus free or not...

On to the extraction
We have the Standard Curve and now we proceed to extracting
the Phosphorus from the soil sample.

This is the part that you will be doing more often since a Standard
Curve only needs to be prepared once a day.

Remember to note down the volume of acid used to adjust the
pH of Olsen's Reagent. In our case, the volume is 0.4mL.

Olsen’s Method

This method is used to extract Phosphorus in soils with pH above 6.

The Bray and Kurtz Method is employed to extract Phosphorus from soils with
pH below 6. The reagent used for extraction, in this case is a solution of 0.03
N Ammonium fluoride in 0.025 N HCl.

Olsen's reagent is added at 1 : 20 ratio to the soil. Thus, when 2.5g of soil are
taken for extraction, 50mL of Olsen's Reagent is to be used.

Bray and Kurtz Reagent is added at a 1 : 10 ratio.

Usually you will find yourself using Olsen’s Method more often. The steps followed
for employing both the methods are identical except for the ratio mentioned
above.

Preparation of the Standard Curve should also be done with the same reagent
used for extraction.


Extracting Phosphorus from the sample

Take 2.5g of soil in a 150mL conical flask.

Add about 0.5g of Phosphorus-free
charcoal to the conical flask.

Prior pH tests on the soil sample indicate
that its pH is 7.2. Hence, Olsen's Reagent
is used as an extractant.

Add 50mL of Olsen's Reagent to the
flask.




Extracting Phosphorus from the sample

2,4-dinitrophenol and ascorbic acid

2,4-dinitrophenol is used as an indicator. The test is to be carried out at a pH
of 3.0. Therefore, we add 2.5M sulphuric acid, drop by drop, till the yellow colour
of the 2,4-dinitrophenol is discharged indicating that the pH of the solution is
exactly 3.0.

This causes a conflict with Reagent B — the ammonium molybdate mixed with
antimony potassium tartarate and ascorbic acid. Ascorbic acid cannot be used
if we use 2,4-dinitrophenol.

Some laboratories use stannous chloride in conjunction with pure ammonium
molybdate. The problem with this method is that the blue complex formed is
very unstable and the colour disappears in a few minutes. In a soil testing centre,
where you could be testing 100 samples everyday, you cannot use Stannous
Chloride.

So we have a bit of problem! The solution is to carry out a kind of “mini-titration”
to determine exactly how much 2.5M sulphuric acid is required to bring the pH
of the solution down to 3.0. We do this by adding 5mL of Olsen’s Extractant to
a 10mL volumetric flask and then adding two drops of 2,4-dinitrophenol. The
solution assumes a yellow colour. Then, we add 2.5M Sulphuric acid to the flask
drop by drop through a graduated pipette, till the yellow colour is discharged.
This gives us the amount of acid required to correct the pH for 5mL of extractant.

Note that one does not need to lower the pH when using Bray and Kurtz Extractant.
After addition of this reagent, we directly add 4mL of Reagent B and then top
up with distilled water to 20mL.


Soil Testing

Recomendados

Recomendados

Más contenido relacionado

Destacado

Destacado (20)

Similar a Soil Testing

Similar a Soil Testing (20)

Último

Último (20)

Soil Testing