---
title: lab
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author: [Nazmul Hossan](https://www.docswell.com/user/4459855)
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description: lab by Nazmul Hossan
published: November 15, 26
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---
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Soil and Environmental Chemistry
Laboratory Report – Illustrated Study Edition
The eight experiments at a glance: what each one measures, how, and what we found.
How to read this: This edition keeps the text, data, calculations and results of the original report and
adds apparatus pictures, procedure flow-charts and concept diagrams beside each experiment so
the work can be understood visually. All drawings are schematic illustrations (not photographs). The
Discussion sections are shortened to their key points.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Table of Contents
Table of Contents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2
Experiment 1: Determination of Soil pH by Using pH Meter
. . . . . . . . . . . . . . . . . . . . .3
Experiment 2: Determination of the Electrical Conductivity (EC) of Soil . . . . . . . . . . . . . . . 5
Experiment 3: Lime Requirement of Soil Sample . . . . . . . . . . . . . . . . . . . . . . . . . . . 7
Experiment 4: Determination of Organic Carbon of Soil Sample
. . . . . . . . . . . . . . . . . . 10
Experiment 5: Determination of Total Nitrogen in Soil Sample . . . . . . . . . . . . . . . . . . . 13
Experiment 6: Determination of Calcium (Ca²⁺) and Magnesium (Mg²⁺) by EDTA Titration . . . . . 16
Experiment 7: Determination of Chloride (Cl⁻) in Soil Sample
. . . . . . . . . . . . . . . . . . . 20
Experiment 8: Determination of Potassium (K⁺) in Soil by Flame Photometer
. . . . . . . . . . . 23
Summary: What the Results Tell Us . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 27
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Experiment 1: Determination of Soil pH by Using pH
Meter
Objective
To determine the active acidity or alkalinity (pH) of a given soil sample in order to assess its suitability for
specific crops and to determine the necessary soil amendments.
Principle
pH is technically defined as the negative base-10 logarithm of the hydrogen ion concentration expressed in
gram equivalents per litre. The pH of soil water is a crucial parameter that indicates the soil's acidity or
alkalinity. It is measured on a scale of 0 to 14, where pH 7 is neutral, values below 7 are acidic and values
above 7 are alkaline. Monitoring soil pH is essential for assessing its suitability for various biological and
agronomic processes.
pH = −log₁₀[H⁺]
H₂O ⇌ H⁺ + OH⁻ ; K = [H⁺][OH⁻] = 10⁻¹⁴
A neutral solution has a hydrogen ion concentration of 10⁻⁷, i.e. pH = 7.
Fig. 1.1: The pH scale. Each unit is a ten-fold change in H⁺; our soil (6.64) is almost neutral.
Apparatus
Fig. 1.2: Apparatus used (schematic).
Procedure
1. The soil sample was collected from the laboratory.
2. The sample was ground and passed through a 2 mm sieve.
3. 10 g of soil was weighed into a 100 mL beaker using a weighing balance, and 25 mL of distilled water was
added.
4. The suspension was stirred several times over the next 30 minutes.
5. The pH meter was switched on and allowed to stabilise.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
6. The electrode was placed into the soil suspension, and the reading was recorded once the display had
stabilised.
Fig. 1.3: Procedure at a glance.
Fig. 1.4: Setting up the measurement.
Observation
The stabilised reading obtained from the pH meter for the given soil sample was: pH = 6.64
Result
Result: The pH of the given soil sample was found to be 6.64. This value falls in the slightly acidic range.
At this pH, most essential plant nutrients remain reasonably available, although prolonged cultivation
without liming may gradually lower the pH further and reduce the availability of nutrients such as
phosphorus and molybdenum.
Discussion
On commonly used classification scales, 6.64 sits at the boundary between “slightly acidic” (6.1–6.5) and
“neutral” (6.6–7.3), close to the optimum range for most crops. This is consistent with the small lime
requirement found in Experiment 3. Because the method relies on a single stabilised reading, the main error
sources are incomplete equilibration of the suspension (reduced by the 30-minute stirring) and drift in
electrode calibration, which should be checked periodically with standard buffers.
In simple words: The pH meter “feels” how many hydrogen ions are in the soil-water mix. More H⁺
means more acidic. Our soil is almost neutral, which suits most crops.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Experiment 2: Determination of the Electrical
Conductivity (EC) of Soil
Objective
To determine the electrical conductivity of a given soil sample in order to estimate the concentration of total
soluble salts and assess the potential salinity hazard for crop production.
Principle
Electrical conductivity (EC) is a measure of a solution's ability to conduct an electric current, which depends
on the concentration of dissolved ions such as salts and minerals. It is typically expressed in microsiemens
per centimetre (μS/cm). Pure water is a very poor conductor, whereas water containing dissolved salts
conducts current approximately in proportion to the amount of salt present. Electrical conductance, the
reciprocal of resistance, is therefore used for salinity appraisal.
Example: EC = 0.000521 mhos/cm = 0.521 mS/cm = 521 μS/cm
Fig. 2.1: More dissolved ions carry more current, so the meter reads a higher EC.
Apparatus
Fig. 2.2: Apparatus used (schematic).
Procedure
1. 10 g of soil was weighed using a weighing balance.
2. The soil sample was poured into a 100 mL beaker.
3. 50 mL of distilled water was added to the sample.
4. The suspension was stirred for 30 minutes with a glass rod.
5. After 30 minutes, the EC meter electrode was placed in the beaker.
6. Once the reading had stabilised, the EC value was recorded.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Fig. 2.3: Procedure at a glance.
Observation
The stabilised reading obtained from the EC meter was: EC = 125 μS/cm
Result
Result: The EC value of the soil sample was found to be 125 μS/cm. This indicates a low soluble-salt
content, which is generally favourable for plant growth. The soil is non-saline and is unlikely to cause any
salinity stress to most crops.
Fig. 2.4: Where our soil sits on the salinity scale.
Discussion
In deciSiemens per metre, 125 μS/cm is about 0.125 dS/m, well below the threshold of roughly 2 dS/m used
to define the lower boundary of “slightly saline” soils, so salinity is not expected to limit crop establishment
or yield. Precautions: mix the suspension thoroughly for the full period and calibrate the EC probe against a
standard KCl solution, since electrode fouling or drift can bias the reading.
In simple words: Salts dissolve into charged particles (ions) that carry electricity. Higher conductivity
means saltier soil. At 125 μS/cm our soil is very low in salts.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Experiment 3: Lime Requirement of Soil Sample
Objective
To determine the amount of agricultural lime required to neutralise soil acidity and raise the pH of a given
soil sample to a target optimum level for crop production.
Principle
Plant growth is adversely affected by low soil pH. To bring the pH back towards neutrality, lime is added to
the soil. Liming may be carried out using CaO, Ca(OH)₂, CaCO₃ (calcite) or CaCO₃·MgCO₃ (dolomite). Soil pH
is the most important criterion determining the lime requirement, although texture, structure,
organic-matter content, and the type and amount of clay also influence it. Lime neutralises soil acidity by
reacting with H⁺ ions in the soil solution, forming water and carbon dioxide and thereby raising the pH:
CaCO₃ + 2H⁺ → Ca²⁺ + H₂O + CO₂
Fig. 3.1: Calcium from the lime takes the place of H⁺ on the soil surface, and H⁺ + OH⁻ forms water.
Apparatus and Reagent
Fig. 3.2: Apparatus and reagent (Ca(OH)₂ solution, 0.01 N).
Procedure
1. 10 g of soil was taken, 25 mL of distilled water was added, and the mixture was shaken for 30 minutes.
2. The pH of the resulting suspension was measured using the pH meter; this was recorded as the “0 drop”
reading.
3. One drop of Ca(OH)₂ solution was added, the suspension was mixed, and the pH was recorded.
4. The addition of a single drop was repeated, with the pH recorded after each drop, up to a total of ten
drops.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Fig. 3.3: Procedure at a glance.
Fig. 3.4: Adding Ca(OH)₂ one drop at a time while watching the pH.
Data Table
Drops of Ca(OH)₂
0
1
2
3
4
5
6
7
8
9
10
pH
6.94
7.04
7.13
7.32
7.48
7.58
7.64
7.67
8.04
8.13
8.19
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Fig. 3.5: Change in soil pH with successive additions of Ca(OH)₂ (determination of lime requirement).
Calculation
One drop of Ca(OH)₂ is taken as ≈ 0.05 mL (standard drop volume).
Total volume of Ca(OH)₂ added (10 drops) = 10 × 0.05 = 0.5 mL
Since 10 g of soil was used, the scaling factor to express the result per kg of soil is:
Scaling factor = 1000 / 10 = 100
Lime requirement (g CaCO₃ per kg soil) = [ (V × N × 50) / Weight of soil (g) ] × Scaling factor
Here V = 0.5 mL, N = 0.01 N, weight of soil = 10 g, scaling factor = 100 (50 is the equivalent-weight
conversion factor for CaCO₃):
Lime requirement = [ (0.5 × 0.01 × 50) / 10 ] × 100 = 2.5 g CaCO₃ per kg of soil
Result
Result: The lime requirement of the given soil sample was found to be 2.5 g of CaCO₃ per kg of soil.
The soil pH rose steadily from 6.94 to 8.19 with successive additions of Ca(OH)₂, with a comparatively
sharper rise beyond the 7th drop, indicating the buffering capacity of the soil against the applied base.
Discussion
The curve rises gradually up to about the 7th drop and then more steeply: early base is consumed
neutralising readily exchangeable acidity, after which small additions raise the pH more. Because the target
pH for most crops is close to neutral (6.5–7.0) and the initial pH (6.94) was already near it, only a modest
lime application was calculated. Main errors: the assumed fixed drop volume (0.05 mL) and electrode drift;
both can be reduced by recalibrating the meter and using a calibrated micro-burette.
In simple words: The soil's acidity is like a bank account of H⁺ ions. We add a known amount of
lime-water and watch the pH climb. How much we had to add tells us how much lime a whole kilogram of
soil would need.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Experiment 4: Determination of Organic Carbon of Soil
Sample
Walkley and Black Wet Oxidation Method
Objective
To determine the percentage of organic carbon present in a soil sample by the Walkley and Black
wet-oxidation method.
Principle
Soil organic carbon is oxidised by potassium dichromate (K₂Cr₂O₇) in the presence of concentrated sulfuric
acid (H₂SO₄); the heat generated on dilution of the acid assists the oxidation. Organic-carbon compounds act
as strong reducing agents, and this reducing capacity is exploited to determine organic carbon by digestion
with a standard oxidising agent, chromic acid:
4Cr⁶⁺ + 3C → 4Cr³⁺ + 3C⁴⁺
K₂Cr₂O₇ + 3C + 6H₂SO₄ → Cr₂(SO₄)₃ + 3CO₂ + 8H₂O + H₂Cr₂O₇ (unreduced)
The excess (unreduced) chromic acid remaining after oxidation is back-titrated with standard ferrous sulfate
solution:
K₂Cr₂O₇ + 6FeSO₄ + 7H₂SO₄ → K₂SO₄ + Cr₂(SO₄)₃ + 3Fe₂(SO₄)₃ + 7H₂O
and the amount of carbon oxidised is calculated from the quantity of chromic acid reduced.
Fig. 4.1: The story of the experiment: soil carbon uses up part of the dichromate; FeSO₄ measures the rest.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Apparatus and Reagents
Fig. 4.2: Apparatus (top row) and reagents (bottom rows). Standard 1 N K₂Cr₂O₇ = 49.03 g/L; FeSO₄ = 278 g FeSO₄·7H₂O +
15 mL conc. H₂SO₄ per litre.
Procedure
1. 2 g of air-dried soil (passed through a 2 mm sieve) was transferred to a 500 mL clean, dry conical flask.
2. 10 mL of standard 1 N K₂Cr₂O₇ solution was added using a pipette.
3. 10 mL of concentrated H₂SO₄ was added with a measuring cylinder and mixed thoroughly.
4. The flask was allowed to cool on a heat-resistant sheet with occasional shaking for 30 minutes.
5. After cooling, about 150 mL of distilled water, 10 mL of phosphoric acid and 0.2 g of sodium fluoride were
added.
6. 3 mL (about 60 drops) of diphenylamine indicator was added, producing a deep violet colour.
7. The excess chromic acid was titrated with standard ferrous sulfate solution until the colour changed
sharply to deep bottle-green (end point).
8. The titration was repeated three times, and a blank was run in the same manner, omitting the soil, to
standardise the FeSO₄ solution.
Fig. 4.3: Procedure at a glance.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Data Table
Observation
Initial burette
reading (mL)
Final burette
reading (mL)
Difference
(mL)
Mean (mL)
Soil sample (T) – 1
10
25
15.0
14.65
Soil sample (T) – 2
25
39.3
14.3
Blank (B) – 1
0
20.6
20.6
Blank (B) – 2
20.6
40.9
20.3
20.45
Fig. 4.4: Blank vs soil titre. The soil needs less FeSO₄ because soil carbon already used up some dichromate.
Standardisation of FeSO₄ Solution
Using S₁V₁ = S₂V₂ with S₁ = 1 N, V₁ = 10 mL and V₂ = mean blank volume = 20.45 mL:
S₂ = (S₁ × V₁) / V₂ = (1 × 10) / 20.45 = 0.489 N
Calculation
1000 mL of 1 N K₂Cr₂O₇ is equivalent to 3 g of carbon (equivalent weight of C = 12/4 = 3), so 1 mL
corresponds to 0.003 g C. Only about 77 % of soil carbon is oxidised under these conditions, so the factor
0.003 is multiplied by 1.3:
% Organic carbon = [ (B − T) × f × 0.003 × 1.3 × 100 ] / W
B = 20.45, T = 14.65, f = 0.489 (strength of FeSO₄), W = 2 g:
% Organic carbon = [ (20.45 − 14.65) × 0.489 × 0.003 × 1.3 × 100 ] / 2 = 0.55 %
Result
Result: The organic carbon content of the given soil sample was found to be 0.55 %. Every 100 g of soil
contains approximately 0.55 g of organic carbon, indicating a moderate level of organic matter.
Discussion
Multiplying by the Van Bemmelen factor of 1.724 gives soil organic matter of about 0.55 × 1.724 ≈ 0.95 %,
on the lower side for agricultural soils, suggesting compost or crop-residue incorporation. The duplicate titres
for soil (15.0 and 14.3 mL) and blank (20.6 and 20.3 mL) are reasonably close. Because Walkley–Black
oxidises only about 77 % of the carbon (hence the 1.3 factor), the result is an estimate; incomplete mixing
and subjective judgement of the end-point colour are the main error sources.
In simple words: Think of the dichromate as a fixed amount of “carbon-eating” chemical. Soil carbon
eats part of it; we then measure the leftover with FeSO₄. The less is left, the more carbon the soil had.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Experiment 5: Determination of Total Nitrogen in Soil
Sample
Kjeldahl Method
Objective
To determine the percentage of total nitrogen present in a given soil sample by the Kjeldahl method.
Principle
The Kjeldahl procedure involves three steps. Digestion: organic nitrogen is converted into ammonium
sulfate by heating with concentrated sulfuric acid and a catalyst. Distillation: the digest is made alkaline
with sodium hydroxide, liberating ammonia gas. Titration: the ammonia is absorbed in boric acid and
titrated against a standard acid.
Soil–N + H₂SO₄ —(heat, 360–410 °C)→ (NH₄)₂SO₄
(NH₄)₂SO₄ + 2NaOH → NH₃↑ + H₂O + Na₂SO₄
NH₃ + H₃BO₃ → NH₄H₂BO₃ (ammonium borate)
NH₄H₂BO₃ + H₂SO₄ → (NH₄)₂SO₄ + H₃BO₃
The digestion temperature is critical: below 360 °C the release of ammonia is slow and incomplete, while
above 410 °C some ammonia is lost from the mixture.
Fig. 5.1: The three Kjeldahl stages: digestion, distillation, titration.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Apparatus and Reagents
Fig. 5.2: Apparatus (top row) and reagents (bottom row). Digestion mixture: CuSO₄·5H₂O, selenium powder and anhydrous
Na₂SO₄.
Procedure
1. 5 g of finely powdered soil was placed in a 500 mL Kjeldahl flask, 20 mL of distilled water was added, and
the mixture was left for 20 minutes.
2. 25 mL of concentrated H₂SO₄ was added and mixed, and the flask was heated over a low flame for 15
minutes.
3. When white fumes of H₂SO₄ appeared, the flask was removed and 2–3 g of catalyst mixture was added.
4. Digestion was continued with slowly rising temperature for two hours or more until the liquid became
clear (light bluish colour).
5. The digest was cooled, diluted with 100 mL distilled water, and made up to volume in a 250 mL
volumetric flask.
6. An aliquot of the extract was distilled with 40 % NaOH in the micro-Kjeldahl apparatus, and the liberated
ammonia was collected in 10 mL of 4 % boric acid containing mixed indicator.
7. The distillate was titrated against standard H₂SO₄ until the colour changed to pink (end point).
8. A blank determination was carried out simultaneously using all reagents but no soil.
Fig. 5.3: Procedure at a glance.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Data Table
Observation
Initial burette
reading (mL)
Final burette
reading (mL)
Difference
(mL)
Average (mL)
Soil sample (T) – 1
1.1
3.1
2.0
1.8
Soil sample (T) – 2
3.1
4.7
1.6
Blank (B) – 1
6.7
8.2
1.5
Blank (B) – 2
8.2
9.1
0.9
1.2
T = 1.8 mL, B = 1.2 mL, f (normality factor of standard acid) = 0.018, w (weight of soil) = 5 g, volume of
extract used = 10 mL.
Fig. 5.4: The extra acid needed by the soil sample over the blank corresponds to the soil nitrogen.
Calculation
1000 mL of 1 N H₂SO₄ is equivalent to 14 g of N, so 1 mL of 1 N H₂SO₄ corresponds to 0.014 g N.
% N in soil = [ (T − B) × f × 0.014 × 250 × 100 ] / (w × volume of extract used)
% N in soil = [ (1.8 − 1.2) × 0.018 × 0.014 × 250 × 100 ] / (5 × 10) = 0.0756 %
Result
Result: The total nitrogen content of the given soil sample was found to be 0.0756 %, equivalent to
0.756 g N per kg of soil.
Discussion
0.756 g N/kg is in the low range typical of many cultivated tropical and sub-tropical soils. Comparing with the
organic carbon from Experiment 4 gives C : N ≈ 0.55 / 0.0756 ≈ 7.3 : 1, narrower than the 10–12 : 1 usually
reported for surface soils; this may indicate rapid mineralisation or simply combined analytical uncertainty
and would be worth repeating. The blank duplicates (1.5 and 0.9 mL) show a wider spread than the soil
duplicates (2.0 and 1.6 mL), suggesting the blank end point was harder to judge. Error sources: ammonia
loss if the delivery tube is not sealed below the boric acid surface, incomplete digestion, and indicator
subjectivity.
In simple words: Digestion turns all the soil's nitrogen into ammonium. Adding strong alkali turns it into
ammonia gas, which we catch in boric acid and then measure with acid. The amount of acid used is
proportional to the nitrogen.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Experiment 6: Determination of Calcium (Ca²⁺) and
Magnesium (Mg²⁺) by EDTA Titration
Objective
To determine the concentration of exchangeable Ca²⁺ and Mg²⁺ ions in a given soil sample by
complexometric titration with EDTA.
Principle
Calcium and magnesium in the soil are extracted using 1 N ammonium acetate (NH₄OAc) solution and
determined by titration with ethylenediaminetetraacetic acid (EDTA), which forms stable, soluble 1 : 1
complexes with both ions:
Ca²⁺ + EDTA⁴⁻ → [Ca–EDTA]²⁻
Mg²⁺ + EDTA⁴⁻ → [Mg–EDTA]²⁻
For calcium alone, murexide indicator is used at about pH 11–12; the colour changes from pink to light blue
at the end point (magnesium is masked as Mg(OH)₂ precipitate). For combined calcium and magnesium,
Eriochrome Black-T (EBT) is used at about pH 10. Magnesium is then obtained by subtracting the calcium
value from the combined (Ca + Mg) value.
Fig. 6.1: How EDTA grabs a metal ion, and the two titrations used.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Apparatus and Reagents
Fig. 6.2: Apparatus (top) and reagents (bottom). Also used: pH 10 buffer (NH₄Cl + NH₄OH), hydroxylamine hydrochloride,
potassium ferrocyanide, triethanolamine.
Procedure
A. Extraction of Ca and Mg
1. 10 g of air-dried soil was taken in a 250 mL conical flask and 100 mL of neutral 1 N NH₄OAc solution was
added.
2. The flask was shaken thoroughly and allowed to stand overnight, then filtered through Whatman No.
40/42 paper to obtain the soil extract.
B. Determination of calcium
1. 2 mL of soil extract was taken in a porcelain cup and 5 mL of 10 % NaOH solution was added to
precipitate Mg(OH)₂.
2. About 0.02 g of murexide indicator was added, and the solution was titrated with 0.01 N EDTA until the
colour changed from pink to light blue (end point).
3. A blank was run using distilled water in place of soil extract, and the EDTA was standardised against a
standard calcium solution.
C. Determination of calcium + magnesium
1. 2 mL of soil extract was taken in a porcelain cup with 25 mL distilled water, 10–15 mL of pH 10 buffer,
and a few drops each of potassium ferrocyanide, hydroxylamine hydrochloride and TEA solutions.
2. The solution was warmed gently, cooled, and 3 drops of Eriochrome Black-T indicator were added.
3. The solution was titrated with EDTA until the colour changed at the end point, and a corresponding blank
was run.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Fig. 6.3: Procedure at a glance.
Observations
Calcium determination
Observation
Initial (mL)
Final (mL)
Difference (mL)
Mean (mL)
Soil sample – 1
0
5.2
5.2
5.2
Soil sample – 2
5.2
10.4
5.2
Blank
33.5
34.5
1.0
1.0
Corrected Ca reading = 5.2 − 1.0 = 4.2 mL
Calcium + magnesium determination
Observation
Initial (mL)
Final (mL)
Difference (mL)
Mean (mL)
Soil sample – 1
0
8.2
8.2
8.1
Soil sample – 2
8.2
16.2
8.0
Blank
0
1.0
1.0
1.0
Corrected (Ca + Mg) reading = 8.1 − 1.0 = 7.1 mL; Mg reading = 7.1 − 4.2 = 2.9 mL
Fig. 6.4: The titre of the Ca + Mg titration is split into its calcium and magnesium parts.
Calculation
Calcium: 1 mL of 1 N EDTA ≡ 0.02 g Ca.
% Ca = (0.02 × mL EDTA × N × total extract volume × 100) / (mL extract used × weight of soil)
With EDTA = 4.2 mL, N = 0.0048, extract used = 2 mL, soil = 10 g, total extract = 100 mL:
% Ca = (0.02 × 4.2 × 0.0048 × 100 × 100) / (2 × 10) = 0.20 % (= 2000 ppm)
Magnesium: 1 mL of 1 N EDTA ≡ 0.012 g Mg; with EDTA = 2.9 mL:
% Mg = (0.012 × 2.9 × 0.0048 × 100 × 100) / (2 × 10) = 0.0835 % (= 835 ppm)
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Result
Nutrient
Percentage (%)
Concentration (ppm)
Calcium (Ca)
0.20
2000
Magnesium (Mg)
0.0835
835
Result: The available calcium and magnesium contents were found to be 0.20 % (2000 ppm) and
0.0835 % (835 ppm), respectively, indicating an adequate base status for most crops.
Discussion
The Ca : Mg ratio is about 2000 / 835 ≈ 2.4 : 1. A ratio of about 3–6 : 1 is often considered favourable, so a
lower ratio may indicate a comparatively high share of magnesium, which can occasionally interfere with
potassium and calcium uptake in sensitive crops. The duplicate titres agree well (5.2 and 5.2 mL; 8.2 and 8.0
mL). Error sources: subjective colour end points, incomplete masking of interfering ions and loss of extract
during filtration; titrate slowly near the end point and use fresh indicators.
In simple words: EDTA is a molecular “claw” that grabs calcium and magnesium one-to-one. When all
the metal is grabbed, the indicator changes colour. By doing one titration for Ca alone and one for Ca +
Mg, we get Mg by subtraction.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Experiment 7: Determination of Chloride (Cl⁻) in Soil
Sample
Mohr's Argentometric Titration Method
Objective
To determine the amount of chloride present in a soil sample by extracting soluble chloride with distilled
water and titrating it against standard silver nitrate (AgNO₃), using potassium chromate (K₂CrO₄) as the
indicator.
Principle
Chloride ions in the soil extract react with silver ions from AgNO₃ to form a white precipitate of silver
chloride:
Ag⁺ + Cl⁻ → AgCl↓
Once all the chloride has precipitated, the first excess of silver ions reacts with the chromate ions of the
indicator to form brick-red silver chromate, marking the end point:
2Ag⁺ + CrO₄²⁻ → Ag₂CrO₄↓ (brick-red)
Fig. 7.1: The colour story of Mohr's titration: yellow, white cloudiness, then the first brick-red tint.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Apparatus and Reagents
Fig. 7.2: Apparatus and reagents (0.01 N AgNO₃; 5 % K₂CrO₄ indicator).
Procedure
Preparation of soil extract
1. 10 g of air-dried, sieved soil was weighed into a 250 mL conical flask, and 50 mL of distilled water was
added (soil : water = 1 : 5).
2. The mixture was shaken for 30 minutes, allowed to settle, and filtered through Whatman No. 42 filter
paper to obtain the clear soil extract.
Titration
1. For the blank, 10 mL of distilled water was pipetted into a conical flask, 1 mL of 5 % K₂CrO₄ was added,
and the mixture was titrated with AgNO₃ until a permanent faint brick-red colour appeared.
2. For the sample, 10 mL of soil extract was pipetted into a clean flask, 1 mL of indicator was added
(solution turned yellow), and the mixture was titrated with 0.01 N AgNO₃ with continuous swirling until
the first permanent brick-red/reddish-brown colour appeared.
3. The titration was repeated to obtain concordant readings.
Fig. 7.3: Procedure at a glance.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Data Table
Observation
Initial (mL)
Final (mL)
Difference (mL)
Mean (mL)
Soil sample – 1
0
3.5
3.5
3.5
Soil sample – 2
3.5
7.0
3.5
Blank
7.0
8.8
1.8
1.8
Fig. 7.4: Sample vs blank titre; the difference is the silver used by chloride.
Calculation
Weight of soil = 10 g, total extract = 50 mL, extract titrated Vₛ = 10 mL, N(AgNO₃) = 0.01 N, equivalent
weight of Cl⁻ = 35.45 g/eq, V = 3.5 mL, B = 1.8 mL.
Cl⁻ (mg/L) = [ (V − B) × N × 35.45 × 1000 ] / Vₛ = [ (3.5 − 1.8) × 0.01 × 35.45 × 1000 ] / 10 = 60.265
mg/L
Accounting for the soil-to-water extraction ratio (10 g soil in 50 mL water):
Cl⁻ (mg/kg) = 60.265 × (0.05 / 0.01) = 301.325 mg/kg
Result
Result: The available chloride (Cl⁻) content of the soil sample was found to be 301.325 mg/kg (≈ 301
ppm), a low-to-moderate level with no apparent chloride toxicity; this level is not normally harmful to
most crops.
Discussion
The two replicate titrations gave identical titres (3.5 mL each), indicating good reproducibility. Chloride
toxicity in sensitive crops is generally associated with levels well above 350–700 mg/kg (thresholds vary with
crop, soil and irrigation water), so about 301 mg/kg is unlikely to be harmful, though it should be monitored
where chloride-sensitive crops or saline irrigation water are involved. Main errors: over-titration past the end
point (the pale brick-red is subtle against the yellow) and interference from other anions; titrate slowly with
swirling and compare with the blank colour.
In simple words: Silver grabs chloride first and makes a white cloud. Only when all the chloride is gone
does silver react with the yellow chromate indicator and turn the solution red-brown. The red tint tells us
we are done.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Experiment 8: Determination of Potassium (K⁺) in Soil by
Flame Photometer
Objective
To determine the amount of available potassium (K⁺) present in a soil sample using a flame photometer.
Principle
When a solution containing potassium ions is introduced into the flame of a flame photometer, the
potassium atoms are excited by the heat. As the excited atoms return to the ground state they emit light of
a characteristic wavelength, for potassium about 766 nm. The intensity of the emitted light is proportional to
the potassium concentration within the working range of the instrument, which is found by comparing the
extract with standard potassium solutions.
Fig. 8.1: Why a flame can “read” potassium: excited atoms give off light of a fixed wavelength.
Fig. 8.2: Inside a flame photometer (block diagram).
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Apparatus and Reagents
Fig. 8.3: Apparatus and reagents. Standards: 1.907 g dried KCl per litre ≈ 1000 mg K/L stock; working standards 0, 5, 10,
15, 20 mg K/L.
Procedure
A. Preparation of soil extract
1. 5 g of air-dried soil was weighed into a clean conical flask, and 25 mL of 1 N NH₄OAc solution was added.
2. The mixture was shaken for about 30 minutes, then filtered to obtain the clear soil extract.
B. Flame photometer measurement
1. The flame photometer was switched on and allowed to stabilise, and the K filter/wavelength (≈ 766 nm)
was selected.
2. Distilled water was aspirated to zero the instrument.
3. The potassium standards (5, 10, 15, 20 mg/L) and a blank were aspirated in turn, and their readings
recorded.
4. The soil extract was aspirated and its reading recorded.
5. The potassium concentration of the extract was determined from the standard calibration curve after
blank correction.
Fig. 8.4: Procedure at a glance.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Observation
No.
Solution
K concentration (mg/L)
Observed reading
Corrected reading
1
Standard 1
5
0.4
0.1
2
Standard 2
10
1.0
0.7
3
Standard 3
15
1.4
1.1
4
Standard 4
20
2.1
1.8
5
Blank
0
0.3
0.0
6
Soil extract
Unknown
2.1
1.8
Corrected reading = Observed reading − Blank reading (0.3). The raw readings still contain the instrument's
background signal, so only the blank-corrected values are used for quantification.
Fig. 8.5: Blank-corrected calibration curve with the soil extract marked (raw readings shown in grey for comparison).
Blank Correction
Corrected soil reading = Soil reading − Blank reading = 2.1 − 0.3 = 1.8
The corrected soil-extract reading (1.8) coincides exactly with the corrected reading of the 20 mg/L standard
(2.1 − 0.3 = 1.8), so the extract falls precisely on the calibration curve at that standard. Hence C = 20 mg
K/L.
Calculation
W = 5 g; V (extracting solution) = 25 mL; C = 20 mg/L; dilution factor D = 1.
Available K (mg/kg) = (C × V × D) / W = (20 × 25 × 1) / 5 = 100 mg K/kg soil
Result
Result: The available potassium content of the given soil sample was found to be 100 mg K/kg soil.
Discussion
Using blank-corrected readings for both the standards and the extract keeps the curve and the sample on
the same basis; here the corrected soil reading equals the corrected 20 mg/L standard, so no interpolation is
needed. 100 mg/kg is in the low-to-medium range (below roughly 100 mg/kg is typically low, 100–200
medium, above 200 high; thresholds vary by crop and local guidelines), so potassium fertilisation should be
considered for potassium-demanding crops. Error sources: incomplete extraction, glassware contamination
and matrix/ionisation interference; use matrix-matched standards and replicate aspirations.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
In simple words: Potassium glows with its own special light in a flame, and the brighter the glow the
more potassium there is. We compare the glow of the soil extract with solutions of known strength to
read off the amount.
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Soil and Environmental Chemistry – Laboratory Report (Illustrated Edition)
Summary: What the Results Tell Us
Traffic-light summary built from the interpretations given in the report's Results and Discussion sections.
How the experiments connect
Soil reaction and salts (1–3): pH 6.64, EC 125 μS/cm and a lime requirement of only 2.5 g CaCO₃/kg all
point the same way: a near-neutral, non-saline soil that needs little correction.
Organic matter and nitrogen (4–5): organic carbon (0.55 %) and total N (0.0756 %) give a C : N ratio of
about 7.3 : 1; both values are on the low side, so organic amendments and nitrogen management are the
main needs.
Ions and nutrients (6–8): Ca is adequate, the Ca : Mg ratio (≈ 2.4 : 1) is a little below the favourable
range, chloride is not a concern, and potassium (100 mg/kg) is at the low-to-medium boundary.
The classification thresholds quoted here are those given in the original report; exact limits vary by crop and
local guidelines.
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