Wednesday, July 26, 2023

Seed BOD calculation

Mastering BOD Analysis: The Essentials of Seeding and Estimation

Biochemical Oxygen Demand (BOD) is a cornerstone parameter in water and wastewater analysis. It indicates how much dissolved oxygen microorganisms require to biologically degrade organic matter present in a sample. Accurate BOD results are essential for assessing pollution levels, evaluating treatment efficiency, and protecting receiving waters.

However, reliable BOD testing depends on two often‑misunderstood steps: proper seeding and realistic BOD estimation before dilution. This blog walks through the why, when, and how of seeding, explains practical BOD estimation techniques, and shows you how to select the right sample volume to avoid failed tests.

Why Seeding Matters in BOD Analysis

BOD testing is fundamentally a biological process. Microorganisms consume biodegradable organic matter and, in doing so, deplete dissolved oxygen. If sufficient and active microbes are not present, oxygen consumption will be low—even when organic pollution is high—leading to falsely low BOD values.

When Seeding Is Required

Seeding is necessary when the sample does not contain enough viable microorganisms, such as:

Industrial wastewater (especially toxic or disinfected streams)
Treated effluents after chlorination or advanced treatment
Samples stored for extended periods

In these cases, a microbial seed (commonly from domestic wastewater, activated sludge, or commercial seed preparations) is added to ensure biological oxidation can proceed.

When Seeding Is Not Required

Seeding is usually unnecessary for samples already rich in microorganisms, including:

Raw sewage
River or stream water receiving wastewater discharges

These samples naturally contain adequate bacterial populations to carry out biodegradation.

Proper seeding ensures the BOD test simulates natural conditions and produces representative, defensible results.

Estimating BOD Before Testing

One of the most common causes of BOD test failure is incorrect sample dilution. If too much sample is used, dissolved oxygen may drop to zero before the test ends. If too little is used, the oxygen depletion may be too small to measure accurately.

To avoid this, always estimate BOD before setting up dilutions.

Using COD as a BOD Estimation Tool

Chemical Oxygen Demand (COD) is a fast, chemical measurement of oxidizable material and is often used as a guide for BOD estimation.

A widely used rule of thumb:

Estimated BOD ≈ 70% of COD

This accounts for the fact that not all oxidizable material measured by COD is biologically degradable.

This estimate allows you to select dilutions that will result in an oxygen depletion of 2–7 mg/L, the optimal range for valid BOD results.

Calculating the Correct Sample Volume

Standard BOD bottles hold 300 mL of diluted sample. The following formula is used to calculate the sample volume required per liter of dilution water:

Sample volume (mL/L) = (X ÷ Estimated BOD) × 1000

Where X is the target oxygen depletion (mg/L).

Choosing X Values

Two dilutions (recommended): X = 2.5 and 4.0 mg/L
Single dilution: X = 3.0 or 3.5 mg/L

Always round to a practical volume for pipetting. For samples with very high BOD, perform a pre‑dilution with distilled water before preparing final BOD bottles.

Worked Example

Given: COD = 400 mg/L

Step 1: Estimate BOD

Estimated BOD = 400 × 0.7 = 280 mg/L

Step 2: Calculate sample volumes

X = 2.5 mg/L:
- (2.5 ÷ 280) × 1000 = 8.9 mL/L (≈ 9 mL/L)
- For a 300 mL bottle: 9 × (300 ÷ 1000) = 2.7 mL
X = 4.0 mg/L:
- (4.0 ÷ 280) × 1000 = 14.3 mL/L (≈ 14 mL/L)
- For a 300 mL bottle: 14 × (300 ÷ 1000) = 4.2 mL

Add the calculated sample volume to each BOD bottle, fill with aerated dilution water containing nutrients and seed (if required), and incubate at 20°C for 5 days. Measure dissolved oxygen before and after incubation to determine BOD.

Practical Tips and Common Pitfalls

Use high‑quality dilution water: Free from chlorine, toxic metals, and inhibitors
Seed controls are essential: Always run seed blanks to correct for seed oxygen uptake
Watch for toxicity: No oxygen depletion may indicate inhibitory substances
Avoid overloading: Zero DO invalidates the test
Always estimate first: Skipping estimation is the fastest way to fail a BOD test

Final Thoughts

Seeding and BOD estimation are not optional extras—they are fundamental to producing valid and meaningful BOD data. By estimating BOD from COD, selecting proper dilutions, and applying seeding only when necessary, you improve accuracy, save time, and reduce repeat testing.

Accurate BOD measurements support better environmental decisions, effective wastewater treatment, and healthier water bodies.

If you’ve applied these methods in your lab, share your experience or questions. For deeper dives, explore related topics such as COD testing methods, seed correction calculations, or microbial inhibition screening.

Happy testing—and keep your oxygen in the sweet spot!

Thursday, September 29, 2022

End point of colour change in Total Hardness.

Understanding End Point Colour Change in Titration (Water Testing)

In titrimetric analysis, the end point is the stage at which an indicator shows a visible colour change, signalling that the chemical reaction is complete. Correct identification of the end point is critical in water and wastewater testing because even a slight error in colour interpretation can lead to inaccurate analytical results. This article explains the concept of end point colour change and highlights common end points observed in routine water analysis.

What Is an End Point in Titration?

The end point is the point during titration at which the indicator changes colour permanently, indicating that the required amount of titrant has reacted with the analyte. Although the end point is close to the equivalence point, it is identified visually using indicators and therefore depends on proper observation and experience.

Correct recognition of the end point ensures:

Accurate test results
Good repeatability
Compliance with standard methods

Importance of End Point Colour Change

In environmental laboratories, most routine analyses such as hardness, alkalinity, chloride, and calcium determination rely on visual indicators. Misjudging the end point colour may result in:

Over‑titration or under‑titration
Incorrect calculation of concentration
Poor quality control results

Therefore, understanding the correct end point colour is essential for laboratory analysts.

Common End Point Colour Changes in Water Testing

1. Total Hardness

In total hardness determination using EDTA titration with Eriochrome Black T (EBT) indicator, the colour changes from wine red to clear blue. The appearance of a stable blue colour indicates that all calcium and magnesium ions have reacted with EDTA.

2. Calcium Hardness

For calcium hardness titration, the indicator commonly used produces a colour change from pink to purple or blue, depending on the method. The end point is confirmed when the colour change remains stable for at least 30 seconds.

3. Total Alkalinity

During alkalinity determination, two end points may be observed depending on the indicator:

Phenolphthalein alkalinity: pink to colourless
Total alkalinity: yellow to orange (with methyl orange or bromocresol green indicator)

Each colour change corresponds to neutralization of specific alkaline components in water.

4. Chloride

In chloride determination by argentometric method, the end point is observed as a colour change from yellow to reddish‑brown due to the formation of silver chromate after all chloride ions have precipitated.

5. Other Titrimetric Tests

Other routine titrations such as acidity, residual chlorine, and sulphide determination also rely on distinct end point colour changes defined by standard methods. Analysts must strictly follow method‑specified indicators and observation conditions.

Tips for Accurate End Point Detection

Use freshly prepared indicators
Perform titration under proper lighting conditions
Swirl the flask continuously during titration
Add titrant dropwise near the end point
Confirm the colour change is permanent

Conclusion

Understanding and correctly identifying end point colour changes is a fundamental skill in water and wastewater testing laboratories. Proper training, practice, and adherence to standard methods help ensure accurate titrimetric analysis and reliable test results. Consistent observation of end point colours improves analytical precision and supports effective laboratory quality control.

Let's see how many of know the end point of colour change of four parameters of water testing.

Total Hardness
Calcium Hardness
Total alkalinity
Chloride

All four parameters end point colour change is given in below picture.

From left to right

Calcium Hardness

Chloride

Total Alkalinity

Total Hardness

Now you have to name the colour write in comment box my dear Environmentalists

Thursday, September 15, 2022

TOC Analysis by titration method

TOC Analysis by Titration: A Simple Guide

Total Organic Carbon (TOC) analysis is a critical parameter in environmental monitoring, water quality assessment, pharmaceuticals, and many industrial processes. While modern TOC analyzers often rely on combustion or UV–persulfate oxidation, TOC analysis by titration remains an important classical approach—especially for educational labs, method validation, and low-resource settings.
This blog breaks down the concept, principle, procedure, advantages, and limitations of TOC analysis by titration in a clear and practical way.

What is Total Organic Carbon (TOC)?

Total Organic Carbon represents the amount of carbon bound in organic compounds present in a sample. It is commonly used as an indirect indicator of organic pollution in water and wastewater systems.

TOC typically includes:

Dissolved organic carbon (DOC)
Particulate organic carbon (POC)

In titration-based methods, TOC is usually determined by oxidizing organic matter and quantifying the carbon indirectly.

Principle of TOC Analysis by Titration

The titrimetric method for TOC analysis is based on three key steps:

Oxidation of organic carbon in the sample using a strong oxidizing agent (commonly potassium dichromate in acidic conditions).
Conversion of organic carbon to carbon dioxide (CO₂) during oxidation.
Back-titration of the excess oxidizing agent with a standard reducing agent (such as ferrous ammonium sulfate).

The amount of oxidant consumed is proportional to the organic carbon content in the sample.

This approach is closely related to the Chemical Oxygen Demand (COD) method and is sometimes referred to as a wet-chemical TOC estimation.

Reagents Commonly Used

Potassium dichromate (K₂Cr₂O₇)
Concentrated sulfuric acid (H₂SO₄)
Ferrous ammonium sulfate (FAS)
Ferroin indicator
Distilled or deionized water

Step-by-Step Procedure (Overview)

Measure a known volume of the water sample into a reflux flask.
Add a measured excess of potassium dichromate solution.
Carefully add concentrated sulfuric acid to initiate oxidation.
Reflux the mixture for a fixed time to ensure complete oxidation.
Cool the solution after refluxing.
Titrate the remaining dichromate with standard ferrous ammonium sulfate using ferroin as an indicator.
Perform a blank determination using distilled water.

Calculation of TOC

The TOC concentration is calculated based on the difference between the blank and sample titration values.

In simplified terms:

TOC ∝ (Dichromate consumed by organic matter)

The result is typically expressed as mg/L of carbon (C).

Calculation of TOC (APHA Style)

According to APHA Standard Methods for the Examination of Water and Wastewater, TOC estimation by wet chemical oxidation is calculated based on the amount of dichromate reduced during reflux and subsequent titration.

APHA Formula


	TOC (mg/L as C) = {(A - B) x N x 3000}/{V}

Where:

A = mL of ferrous ammonium sulfate (FAS) used for blank
B = mL of ferrous ammonium sulfate (FAS) used for sample
N = Normality of FAS
V = Volume of sample taken (mL)
3000 = Conversion factor (based on equivalent weight of carbon × 1000)

Example Calculation (APHA Format)

Data:

Sample volume (V) = 50 mL
Normality of FAS (N) = 0.1 N
Blank titration (A) = 24.0 mL
Sample titration (B) = 16.0 mL

Calculation:


	TOC = {(24.0 - 16.0) x 0.1 x 3000}/{50}


	TOC. = {2400}{50} = 48 mg/L as C

Reporting of Results (APHA Recommendation)

Results should be reported as mg/L Total Organic Carbon (as C)
Report to the nearest whole number for routine analysis
Include blank correction and reagent normality in the report

Advantages of TOC Analysis by Titration

Simple and cost-effective
Does not require sophisticated instrumentation
Suitable for academic and training laboratories
Useful for cross-checking instrumental TOC results

Limitations

Time-consuming compared to automated TOC analyzers
Uses hazardous chemicals (chromium compounds, strong acids)
Lower sensitivity for very low TOC levels
Interference from inorganic reducing substances

Applications

Water and wastewater analysis
Environmental monitoring
Teaching analytical chemistry concepts
Method development and comparison studies

Titration vs Instrumental TOC Methods

Aspect	Titration Method	Instrumental TOC
Cost	Low	High
Accuracy	Moderate	High
Speed	Slow	Fast
Automation	Manual	Fully automated

APHA Method Reference

This method aligns conceptually with APHA Standard Methods for the Examination of Water and Wastewater, latest edition, under:

Method 5310 – Total Organic Carbon (TOC)
Wet chemical oxidation approach (classical/reference technique)

Note: While APHA primarily recommends instrumental methods for routine TOC analysis, wet oxidation followed by titration is acceptable for instructional purposes, method comparison, and laboratories without TOC analyzers.

Practical Record Format (APHA Style)

Aim: To determine Total Organic Carbon (TOC) in a water sample by wet oxidation followed by titration.

Principle: Organic carbon present in the sample is oxidized by potassium dichromate in acidic medium. The excess dichromate is titrated with ferrous ammonium sulfate. The amount of dichromate consumed is proportional to the organic carbon content.

Reagents:

Potassium dichromate solution
Concentrated sulfuric acid
Ferrous ammonium sulfate (FAS)
Ferroin indicator

Procedure:

Take a measured volume of the sample in a reflux flask.
Add a known excess of potassium dichromate.
Add sulfuric acid carefully and reflux for a fixed time.
Cool and titrate the excess dichromate with FAS using ferroin indicator.
Perform a blank determination.

Calculation:


	ext{TOC (mg/L as C)} = rac{(A - B) 	imes N 	imes 3000}{V}

Result: Total Organic Carbon of the given sample = ______ mg/L as C.

Precautions:

Handle sulfuric acid and dichromate with care.
Ensure proper reflux time for complete oxidation.
Always run a reagent blank.

Final Thoughts

Although modern laboratories increasingly rely on automated TOC analyzers, TOC analysis by titration still holds educational and practical value. Understanding this classical APHA-aligned method helps analysts build strong fundamentals in water quality analysis and analytical chemistry.

Thursday, July 14, 2022

Analysis of Extractable organic ( Oil & Grease) in Hazardous Waste

Oil and grease are critical parameters routinely analyzed in water, wastewater, and hazardous solid waste. These substances originate from petroleum products, lubricants, fats, oils, waxes, and industrial residues. When present in high concentrations, oil and grease can clog treatment systems, interfere with biological processes, contaminate soil and groundwater, and pose serious environmental risks.

This blog explains the principle, materials, and step‑by‑step procedure for estimating extractable organic matter (oil & grease) in hazardous waste using solvent extraction, presented in a clear and plagiarism‑free format.

Why Oil & Grease Analysis Matters

Monitoring oil and grease is essential because:

It affects wastewater treatment efficiency
It can inhibit microbial activity in biological systems
It contributes to soil and groundwater contamination
Regulatory agencies set discharge limits for compliance

For wastewater-specific methods, oil and grease are often analyzed using liquid–liquid extraction. In solid and hazardous waste, solvent extraction using hexane is widely applied.

Principle of the Method

Oil and grease are separated from the sample by solvent extraction using n‑hexane. Hexane selectively dissolves non‑polar organic compounds such as oils, fats, and greases.

Because hexane has a boiling point of approximately 40–60°C, evaporation during analysis is performed at a temperature slightly above this range to ensure complete solvent removal without decomposing the extracted organics.

The final mass increase of the extraction flask corresponds to the amount of extractable organic matter present in the sample.

Apparatus and Materials

Analytical balance
Vacuum pump
Extraction thimble (filter paper)
Glass wool or small glass beads
Beakers and conical flasks
Pipettes
Porcelain mortar and pestle
Extraction flask
Water bath
Desiccator
pH indicator paper

Reagents

Concentrated hydrochloric acid (HCl)
Anhydrous sodium sulfate
n‑Hexane (extraction solvent)

Analysis Procedure

Weigh 20 ± 0.5 g of wet sludge with a known dry‑weight fraction and place it in a 150 mL beaker.
Acidify the sample to pH 2 using approximately 0.3 mL concentrated HCl.
Add 25 g of magnesium sulfate hydrate (MgSO₄·xH₂O) and stir until a smooth paste is formed.
Spread the paste along the sides of the beaker to aid evaporation. Allow it to stand for 15–30 minutes until solidified.
Transfer the dried material to a porcelain mortar and grind into a fine powder.
Place the powder into a paper extraction thimble.
Wipe the beaker and mortar with solvent‑moistened filter paper and add the wipes to the thimble.
Fill the thimble with glass wool or glass beads to ensure proper drainage.
Mix 10 g of solid sample (with known dry‑weight fraction) thoroughly with 10 g anhydrous sodium sulfate and place into the extraction thimble.
Perform solvent extraction at a rate of approximately 20 cycles per hour for 4 hours.
Filter the extract using grease‑free cotton into a pre‑weighed boiling flask, wearing gloves to prevent contamination.
Rinse the cotton and flask with fresh solvent.
Attach the flask to a distillation setup and evaporate hexane by immersing the flask in a 70°C water bath. Collect solvent for reuse.
When the distillation head reaches 50°C or the flask appears dry, remove the setup.
Sweep the flask with air for 15 seconds using a vacuum source to remove solvent vapors.
Wipe the flask exterior, cool it in a desiccator for 30 minutes, and weigh.

A solvent blank must be analyzed with each batch for quality control.

Calculation

Extractable Organics (%)

Where:

W₁ = Initial weight of empty boiling flask
W₂ = Final weight of flask after extraction

Results are reported as percentage of total dry solids.

Key Points for Accurate Results

Ensure complete drying before weighing
Avoid fingerprints or grease contamination
Use high‑purity hexane
Always run method blanks
Maintain consistent extraction cycles and temperature

Final Thoughts

The analysis of extractable organics (oil & grease) in hazardous waste is a vital component of environmental monitoring and waste management. When performed carefully, this method provides reliable data for regulatory compliance, treatment design, and pollution control.

Understanding oil and grease behavior also complements other parameters such as BOD, COD, and TOC, offering a complete picture of organic pollution.

If you would like a detailed post on BOD analysis in wastewater or simplified infographics for this method, feel free to ask.

Saturday, May 8, 2021

Measurement of Nitrite in Water and Wastewater

How to Measure Nitrite in Water and Wastewater: Methods, Health Risks, and Tips

Meta Description: Learn how to accurately measure nitrite in water and wastewater using the Griess method. Understand its environmental impact, health risks, and practical testing tips.

Introduction

Nitrite (NO₂⁻) might be a minor component in water chemistry, but it plays a critical role in water safety, environmental monitoring, and wastewater treatment. Elevated nitrite levels can be toxic to humans and aquatic life and may indicate issues in water treatment processes.

Measuring nitrite accurately is essential for public health, regulatory compliance, and ecosystem protection. This article explains how nitrite forms, why it matters, and the most effective methods for testing it in water and wastewater.

What is Nitrite and How Does it Form?

Nitrite is an intermediate in the nitrogen cycle, formed during the microbial conversion of ammonia to nitrate:

Ammonia → Nitrite → Nitrate

In water and wastewater:

Produced by ammonia-oxidizing bacteria during nitrification
Usually short-lived in healthy systems
Accumulation signals biological imbalance or incomplete treatment

Monitoring nitrite levels helps detect potential treatment failures and prevent environmental contamination.

Why Measuring Nitrite is Important

Nitrite is not just another chemical—it’s an indicator of water quality and safety. Key reasons to measure nitrite include:

Health Risks: Can cause methemoglobinemia (“blue baby syndrome”) in infants by interfering with oxygen transport in blood.
Aquatic Toxicity: Toxic to fish and other aquatic organisms, even at low concentrations.
Wastewater Monitoring: Indicates incomplete nitrification or oxygen deficiencies in treatment systems.
Regulatory Compliance: Safe drinking water standards limit nitrite to ≤ 1 mg/L (WHO & EPA).

Accurate nitrite measurement ensures safe water, efficient treatment, and environmental protection.

How to Measure Nitrite in Water and Wastewater

Several analytical techniques exist, but the colorimetric Griess method is the most widely used due to its accuracy, simplicity, and cost-effectiveness.

The Colorimetric (Griess) Method

Principle: Nitrite reacts with sulfanilamide in acidic conditions to form a diazonium salt. This intermediate reacts with N-(1-naphthyl)ethylenediamine (NED), producing a pink/red azo dye. The intensity of the color is directly proportional to the nitrite concentration.

Equipment & Reagents Needed:

Spectrophotometer (wavelength: 543 nm)
Sulfanilamide reagent
NED reagent
Phosphoric acid
Sodium nitrite (for standard solutions)

Detection Range: 0.001–1.0 mg/L

Step-by-Step Procedure

Take a measured sample of water or wastewater.
Add sulfanilamide reagent and mix thoroughly.
Add NED reagent and allow ~10 minutes for color development.
Measure absorbance using a spectrophotometer.
Determine nitrite concentration using a calibration curve.

Conversion to Nitrogen:


\text{mg/L as N} = \text{mg/L as NO₂⁻} \times \frac{14}{46}

Alternative Nitrite Testing Methods

Other methods are available for specialized applications:

Ion Chromatography: High precision, simultaneous detection of multiple ions.
Electrochemical Sensors: Real-time monitoring in treatment plants.
UV Spectrophotometry: Suitable for clear water but prone to interference.

Applications of Nitrite Measurement

Nitrite testing is critical in many areas:

Drinking Water Safety: Ensures regulatory compliance and safe consumption.
Wastewater Treatment: Helps optimize nitrification and denitrification processes.
Environmental Monitoring: Detects nitrogen pollution in rivers, lakes, and groundwater.
Research: Provides insight into microbial nitrogen cycling.

Tips for Accurate Nitrite Measurement

Neutralize residual chlorine before analysis.
Filter turbid or colored samples to reduce interference.
Use freshly prepared reagents and calibration standards.
Regularly calibrate your spectrophotometer.
Follow standard methods (APHA, ISO, EPA) for reliable results.

FAQs About Nitrite in Water

Q1: What is a safe level of nitrite in drinking water?

WHO and EPA recommend ≤ 1 mg/L as NO₂⁻.

Q2: Can nitrite turn into nitrate in water?

Yes, nitrite is oxidized to nitrate by nitrite-oxidizing bacteria in natural and treatment systems.

Q3: How fast does the Griess method work?

Color develops in about 10 minutes, making it ideal for routine laboratory testing.

Q4: Does nitrite affect aquatic life?

Even low concentrations can be toxic, especially to fish and sensitive invertebrates.

Conclusion

Measuring nitrite in water and wastewater is essential for public health, environmental safety, and wastewater management. The Griess colorimetric method is the most widely used technique due to its reliability, sensitivity, and ease of use. Accurate nitrite monitoring helps detect water treatment issues, prevent ecological damage, and ensure safe drinking water for communities world

Friday, May 7, 2021

Measurement of Silica by Molybdosilicate Method water and waste water samples.

Silica Analysis in Water and Wastewater: APHA Method

Silica (SiO₂) is a natural component of water, originating from the weathering of silicate minerals in rocks, soil, and sand. While generally not harmful to human health, high silica concentrations can cause scaling in boilers, fouling of membranes, and operational inefficiencies in water treatment systems. Reliable measurement of silica is essential for industrial water systems, wastewater reuse, and reverse osmosis (RO) processes.

The APHA molybdate blue method is a standardized and widely used procedure for silica determination, particularly for reactive silica.

Forms of Silica in Water

Reactive (Dissolved) Silica
- Mainly monosilicic acid (H₄SiO₄)
- Directly measurable by APHA methods
Polymeric or Colloidal Silica
- Forms from condensation of dissolved silica
- Reacts slowly and may require digestion
Particulate Silica
- Suspended solids (sand, silt, clay)
- Usually removed before analysis

Importance of Silica Analysis

Monitoring silica is critical for:

Preventing scale formation in boilers, cooling towers, and heat exchangers
Protecting RO membranes and industrial equipment
Optimizing demineralization and water reuse processes
Ensuring efficiency in wastewater treatment and zero liquid discharge systems
Extending equipment lifespan and reducing maintenance costs

Typical Silica Concentrations

Water Source	Silica Concentration (mg/L as SiO₂)
Surface water	1–30
Groundwater	10–100
Industrial wastewater	Highly variable
RO permeate / high-purity	< 1

APHA Method for Silica Determination

The Molybdate Blue Colorimetric Method measures reactive silica:

Silica reacts with ammonium molybdate in acidic conditions to form silicomolybdic acid
Reduction of this complex produces a blue color
The intensity of the blue color is proportional to silica concentration and is measured spectrophotometrically at ~815 nm

This method is widely used due to its simplicity, sensitivity, and cost-effectiveness.

Laboratory SOP (APHA Method)

Purpose

To determine reactive silica in water and wastewater samples using a standardized colorimetric procedure.

Scope

Applicable to drinking water, surface water, groundwater, industrial wastewater, and RO permeate.

Apparatus and Equipment

Spectrophotometer or colorimeter (~815 nm)
Plastic or polyethylene sample bottles
Volumetric flasks, pipettes, and test tubes

Reagents

Ammonium molybdate reagent
Acid reagent (e.g., sulfuric acid)
Reducing reagent (e.g., ascorbic acid solution)
Silica stock solution
Deionized, silica-free water

Sample Collection

Collect in plastic bottles (avoid glass)
Filter samples if particulate silica is not required
Analyze promptly at room temperature

Calibration Procedure

Prepare silica standards (0, 5, 10, 20, 30, 50 mg/L as SiO₂)
Add reagents to standards and blanks under identical conditions
Allow color development according to the method
Measure absorbance at ~815 nm against a reagent blank
Plot absorbance vs. silica concentration to create the calibration curve

Calculating the Slope (Calibration Factor)

Use the equation of the line from the calibration curve: Absorbance = m × [SiO₂] + c
m is the slope, which represents the change in absorbance per unit concentration
c is the y-intercept (blank absorbance)

For example, if the calibration curve equation is:

Absorbance = 0.0085 × [SiO₂] + 0.005

Then the slope m = 0.0085 Abs/mg/L, which is used in sample calculations.

Sample Analysis Procedure

Pipette a measured volume of sample into a clean reaction vessel
Add ammonium molybdate reagent under acidic conditions and mix
Add reducing reagent and allow full color development
Measure absorbance against the reagent blank
Determine silica concentration using the calibration curve and slope

Quality Control Measures

Include a reagent blank in each batch
Analyze duplicate samples to assess precision
Verify calibration with a mid-range standard
Recalibrate when instrument drift or new reagent batches occur

Example Calculation

Given:

Sample absorbance = 0.420
Calibration curve equation: Absorbance = 0.0085 × [SiO₂] + 0.005

Step 1: Solve for SiO₂ concentration

[SiO₂] = (0.420 − 0.005) ÷ 0.0085 ≈ 48.8 mg/L

Step 2: Apply dilution factor if used

Final Result: Reactive silica concentration = 48.8 mg/L as SiO₂

Reporting Guidelines

Report results in mg/L SiO₂
Specify that the analysis measures reactive silica
Include details of any filtration or dilution performed

Conclusion

The APHA molybdate blue method provides a reliable, sensitive, and standardized approach for reactive silica measurement in water and wastewater. Accurate silica analysis is essential for preventing scaling, protecting membranes, optimizing treatment processes, and ensuring sustainable water management. Using the calibration curve slope in calculations ensures consistent and reproducible results, which are critical for both industrial and municipal water systems.

Wednesday, April 28, 2021

Measuring Sulfur Dioxide (SO₂) in Ambient Air

Measuring Sulfur Dioxide (SO₂) in Ambient Air: A Practical Laboratory Guide.

Monitoring sulfur dioxide (SO₂) in ambient air is a key component of air‑quality assessment and public health protection. SO₂ is a major atmospheric pollutant generated primarily from fossil‑fuel combustion, power plants, refineries, and other industrial activities. Prolonged exposure can harm human health, damage vegetation, and contribute to acid rain formation.

This blog presents a practical, laboratory‑based guide to measuring ambient SO₂ using the widely accepted para‑rosaniline colorimetric method, explaining the principle, reagents, procedures, and calculations in a clear and user‑friendly manner.

Why Measure Ambient SO₂?

Accurate measurement of sulfur dioxide is essential because:
SO₂ irritates the respirat hiory system and aggravates asthma
It damages crops, forests, and building materials
It contributes to acid rain and secondary particulate formation
Regulatory agencies require routine monitoring for compliance
Reliable laboratory analysis supports environmental decision‑making and pollution‑control strategies

Principle of the Method

Ambient air is drawn through an absorbing solution of potassium tetrachloromercurate (TCM). Sulfur dioxide reacts with TCM to form a stable dichlorosulphitomercurate complex, which is resistant to oxidation by oxygen, ozone, and nitrogen oxides. This stability allows samples to be stored prior to analysis without significant SO₂ loss.

For analysis, the complex reacts with para rosaniline in the presence of formaldehyde, forming a colored compound. The color intensity is directly proportional to the amount of SO₂ present and is measured spectrophotometrically at 560 nm.

Roles and Responsibilities

Laboratory Chemist: Conducts sampling, analysis, and calculations
Technical Manager: Reviews analytical procedures and results
Quality Manager: Ensures SOP implementation and quality control

Key Reagents Used

The following reagents are critical for accurate SO₂ analysis:

Distilled water (free from oxidizing agents)
Potassium tetrachloromercurate (0.04 M) – absorbing solution
Sulphamic acid (0.6%) – removes nitrogen oxide interference
Formaldehyde (0.2%) – supports color development
Para rosaniline dye – produces measurable color
Iodine and sodium thiosulphate solutions – for standardization
Standard sulphite solution – used to prepare calibration standards

All reagents must be freshly prepared or stored under specified conditions to maintain analytical accuracy.

Preparation of Standards and Calibration

Standard Sulphite Solution

A sulphite solution is prepared using sodium sulphite or sodium metabisulphite and standardized by iodine–thiosulphate titration. This step determines the exact SO₂ concentration in the standard solution.

Working Sulphite–TCM Solution

A measured volume of the standardized sulphite solution is diluted and mixed with TCM. This working solution is stable for up to 30 days when stored under refrigeration and is used for preparing calibration standards.

Calibration Curve

Different volumes of the working sulphite–TCM solution are added to volumetric flasks to prepare standards containing known amounts of SO₂. After reagent addition and color development, absorbance is measured at 560 nm.

A straight‑line plot of absorbance versus SO₂ mass (µg) confirms proper calibration. The slope of this line is used to calculate the calibration factor (B).

Sample Analysis Procedure

Prepare a reagent blank, control, and sample solutions
Add sulphamic acid to remove nitrite interference
Add formaldehyde followed by para rosaniline
Allow color to develop for 30 minutes
Measure absorbance between 30–60 minutes at 560 nm using a 1 cm cuvette
Use distilled water, not the reagent blank, as the spectrophotometer reference

Strict temperature control is essential, as color intensity is temperature‑dependent.

Handling High Absorbance Samples

If absorbance lies between 1.0 and 2.0, dilute the sample 1:1 with reagent blank
Highly concentrated samples may require dilution up to six times
Always apply the correct dilution factor (D) during calculations

Calculations

SO₂ Concentration in Air

SO₂ concentration is calculated using:

SO₂ (µg/m³) = (SA × B × D) ÷ V₁

Where:

SA = Sample absorbance
B = Calibration factor
D = Dilution factor
V₁ = Volume of air sampled at STP (m³)

Conversion to ppm

The calculated mass concentration can be converted to ppm using standard gas‑law relationships.

Quality Control and Good Laboratory Practice

Analyze control samples with known SO₂ concentrations
Recalibrate if reagent blank absorbance deviates significantly
Clean cuvettes immediately after use
Maintain consistent temperature during calibration and analysis

Final Thoughts

The para rosaniline method remains a dependable and sensitive technique for measuring sulfur dioxide in ambient air. When performed with careful reagent preparation, calibration, and quality control, it provides accurate and reproducible results essential for air‑quality monitoring and regulatory compliance.

Consistent application of this method helps laboratories contribute reliable data toward protecting public health and the environment.

Friday, April 23, 2021

Methods phosphate analysis in water and wastewater

Phosphate Analysis in Water and Wastewater

Introduction

Phosphates are essential nutrients for plant growth, but when present in excess in water and wastewater, they become a major environmental concern. Elevated phosphate levels are one of the primary causes of eutrophication, leading to algal blooms, oxygen depletion, and degradation of aquatic ecosystems. Because of these impacts, phosphate analysis is a critical component of water quality monitoring and wastewater treatment operations.

This blog explores why phosphate analysis matters, common forms of phosphates found in water, analytical methods used for measurement, and their significance in environmental management.

What Are Phosphates?

Phosphates are chemical compounds containing phosphorus combined with oxygen, commonly found as:

Orthophosphates – the simplest and most reactive form
Condensed phosphates – polyphosphates and metaphosphates
Organic phosphates – phosphorus bound to organic molecules

In water and wastewater analysis, results are often reported as mg/L of PO₄³⁻ or as phosphorus (P).

Sources of Phosphates in Water and Wastewater

Phosphates enter water bodies from both natural and anthropogenic sources:

Natural Sources

Weathering of phosphate-containing rocks
Decomposition of organic matter

Anthropogenic Sources

Domestic sewage and human waste
Detergents and cleaning agents
Agricultural runoff containing fertilizers
Industrial effluents (food processing, fertilizer, chemical industries)

Why Is Phosphate Analysis Important?

Phosphate monitoring is essential for several reasons:

Prevention of eutrophication: Excess phosphates promote uncontrolled algal growth.
Protection of aquatic life: Algal decay reduces dissolved oxygen, harming fish and other organisms.
Regulatory compliance: Environmental agencies set strict discharge limits for phosphates.
Process control: Wastewater treatment plants rely on phosphate measurements to optimize biological and chemical removal processes.

Analytical Methods for Phosphate Determination

Several methods are used for phosphate analysis depending on accuracy requirements, sample type, and available instrumentation.

1. Colorimetric (Spectrophotometric) Method

This is the most widely used method for phosphate analysis.

Based on the reaction of phosphate with ammonium molybdate under acidic conditions
Forms a blue-colored complex (molybdenum blue)
Intensity of color is measured using a spectrophotometer

Advantages:

Simple and cost-effective
Suitable for routine laboratory analysis

Limitations:

Interference from silica, arsenate, or turbidity if not properly controlled

2. Ascorbic Acid Method

A refined colorimetric method commonly recommended by standard methods.

Produces a stable blue color
High sensitivity and reproducibility

This method is widely used in environmental laboratories for both water and wastewater samples.

3. Ion Chromatography

Separates phosphate ions from other anions
Quantification based on conductivity or UV detection

Advantages:

High precision and selectivity
Capable of multi-ion analysis

Limitations:

High equipment and maintenance cost

4. Inductively Coupled Plasma (ICP) Techniques

Measures phosphorus directly as an element
Suitable for trace-level analysis

Advantages:

Very high sensitivity
Minimal chemical interference

Limitations:

Expensive instrumentation
Requires skilled operation

Sample Collection and Preservation

Proper sampling is critical for accurate phosphate analysis:

Use clean, phosphate-free containers
Analyze samples as soon as possible
Refrigerate samples at 4°C if analysis is delayed
Acid digestion may be required to convert all forms of phosphorus into orthophosphate for total phosphate analysis

Phosphate Removal and Control in Wastewater

Phosphate analysis supports treatment strategies such as:

Biological phosphorus removal (EBPR)
Chemical precipitation using alum, ferric chloride, or lime
Tertiary treatment and filtration

Accurate monitoring ensures effective removal and regulatory compliance.

Regulatory Standards

Many environmental authorities specify maximum allowable phosphate or phosphorus concentrations in effluents and surface waters. Typical discharge limits range from 0.1 to 1.0 mg/L as phosphorus, depending on local regulations and receiving water sensitivity.

Conclusion

Phosphate analysis plays a vital role in protecting water quality and maintaining ecological balance. By understanding phosphate sources, applying appropriate analytical methods, and interpreting results correctly, water and wastewater professionals can effectively control nutrient pollution and meet environmental standards.

Regular monitoring, combined with efficient treatment processes, is key to reducing phosphate-related environmental impacts and ensuring sustainable water management.

Friday, March 26, 2021

Measurement of Acidity in water and wastewater

Understanding Acidity in Water: A Practical Lab Guide to Accurate Measurement

Water quality sits at the heart of environmental health, public safety, and countless industrial processes. While parameters like pH and turbidity often steal the spotlight, acidity is another critical factor that deserves attention. Elevated acidity can corrode pipelines, disrupt aquatic ecosystems, and compromise drinking water safety.

In this post, we’ll walk through a standard laboratory method for measuring acidity in water and wastewater samples. This isn’t just textbook chemistry—it’s a practical, widely used approach that helps protect water resources and ensures regulatory compliance.

Why Measuring Acidity Matters

Acidity testing determines the concentration of acidic substances present in a water sample. These acids may exist as free hydrogen ions or as compounds that release hydrogen ions when dissolved or hydrolyzed.

By neutralizing these acidic components with a standard alkaline solution, laboratories can quantify acidity and express it as milligrams per liter (mg/L) of calcium carbonate (CaCO₃). This standardized unit allows results to be compared across different samples, locations, and regulatory frameworks.

Acidity measurements are essential for:

Monitoring environmental pollution
Designing and evaluating wastewater treatment processes
Ensuring compliance with environmental and industrial discharge standards

Roles and Responsibilities in the Laboratory

Accurate acidity measurement is a team effort. In a typical laboratory setup:

Laboratory Chemist: Conducts the titration and records observations.
Technical Manager: Reviews data for accuracy, consistency, and technical validity.
Quality Manager: Ensures procedures follow approved standards and quality protocols.

Each role contributes to producing reliable, defensible results.

The Science Behind Acidity Testing

Acidity is determined using acid–base titration, one of the most fundamental techniques in analytical chemistry. The acidic components in the sample react with a standardized base—commonly sodium hydroxide (NaOH).

As the base is added, it neutralizes the acid. The endpoint of the reaction is detected using a color indicator, signaling that all acidic components have been neutralized.

Equipment and Glassware Required

This method relies on standard laboratory glassware:

Conical (Erlenmeyer) flasks
Burette for accurate titrant delivery
Pipettes or measuring cylinders for sample handling

No advanced instrumentation is required, making this method accessible and cost-effective.

Reagent Preparation

High-quality reagents are critical for accurate results. Proper preparation and storage are essential.

1. 0.02 N Sodium Hydroxide (NaOH)

Dilute 200 mL of 0.1 N NaOH to 1 liter with distilled water.
This solution serves as the titrant.

2. Phenolphthalein Indicator

Dissolve 80 mg of phenolphthalein in 100 mL of 95% ethanol.
Isopropyl alcohol or methanol may be used as alternatives.
The indicator turns pink under basic conditions, marking the titration endpoint.

3. 0.05 N Potassium Hydrogen Phthalate (KHP)

Dry approximately 15 g of KHP at 120°C for two hours.
Cool in a desiccator.
Accurately weigh about 10 g and dilute to 1 liter with distilled water.
This primary standard is used to standardize the NaOH solution.

Step-by-Step Acidity Testing Procedure

Sample Preparation
Measure 50–100 mL of the water or wastewater sample into a clean conical flask. The volume may be adjusted depending on the expected acidity.
Indicator Addition
Add 2–3 drops of phenolphthalein indicator to the sample.
Titration
Slowly titrate with 0.02 N NaOH from the burette while gently swirling the flask.
Endpoint Detection
Continue titration until a faint pink color appears and persists for at least 30 seconds.
Recording Results
Record the burette reading corresponding to the volume of NaOH used.
Standardization
Standardize the NaOH solution by titrating 10 mL of the 0.05 N KHP solution using the same procedure.

Calculation of Acidity

Acidity is calculated using the following formula:

Acidity (mg/L as CaCO₃) = (A × N × 50 × 1000) / V

Where:

A = Volume of NaOH used (mL)
N = Normality of NaOH
V = Volume of sample (mL)
50 = Equivalent weight of CaCO₃

This calculation converts titration data into a standardized and meaningful result.

Practical Tips and Common Pitfalls

Rinse all glassware with distilled water before use.
Always standardize the NaOH solution before analysis.
Highly colored or turbid samples may require modified techniques or indicators.
Wear appropriate personal protective equipment (gloves and safety goggles).
Store reagents properly to avoid degradation or contamination.

Final Thoughts

Acidity testing is a simple yet powerful tool in water and wastewater analysis. By following this standardized titration method, laboratories can generate accurate, reliable data that support environmental protection, public health, and industrial compliance.

Whether you’re a student learning analytical chemistry or a professional working in environmental monitoring, mastering acidity measurement is a valuable skill that directly contributes to safer and more sustainable water systems.

Have questions or hands-on experiences to share? Join the conversation and stay tuned for more practical lab insights!

Sunday, March 21, 2021

Biological Oxygen Demand BOD Testing

Biological Oxygen Demand (BOD): A Complete Laboratory Guide

Biological Oxygen Demand (BOD) is one of the most widely used parameters for evaluating organic pollution in water and wastewater. It reflects the amount of oxygen required by microorganisms to biologically decompose organic matter under aerobic conditions. This blog presents a clear, original, and laboratory-oriented explanation of the BOD test procedure, including reagent preparation, dilution techniques, incubation, and calculation

Purpose of BOD Analysis

The purpose of this procedure is to describe the laboratory method for measuring Biological Oxygen Demand (BOD) in water and wastewater samples.

Scope

This method is applicable to environmental laboratories involved in the analysis of:

Surface water
Groundwater
Treated and untreated wastewater
Industrial effluents

where BOD determination is required for monitoring, treatment efficiency, or regulatory compliance.

Roles and Responsibilities

Laboratory Chemist: Performs sample preparation, dilution, incubation, DO measurement, and calculation of BOD.
Technical Manager: Reviews analytical activities and validates results.
Quality Manager: Ensures SOP implementation and adherence to quality requirements.

Principle of the BOD Test

Biological Oxygen Demand is defined as the quantity of dissolved oxygen consumed by microorganisms while stabilizing biodegradable organic matter present in a water or wastewater sample under aerobic conditions. The reduction in dissolved oxygen over a fixed incubation period reflects the BOD of the sample.

Instruments and Equipment

BOD bottles (300 mL capacity)
BOD incubator maintained at 27 ± 1°C
Measuring cylinders and volumetric flasks
DO titration setup (as per Winkler method)

Reagents and Their Preparation

1. Phosphate Buffer Solution

Reagents required:

Potassium dihydrogen phosphate (KH₂PO₄): 8.5 g
Dipotassium hydrogen phosphate (K₂HPO₄): 21.75 g
Disodium hydrogen phosphate heptahydrate (Na₂HPO₄·7H₂O): 33.4 g
Ammonium chloride (NH₄Cl): 1.7 g
Distilled water: up to 1000 mL

Preparation: Dissolve all salts in distilled water, make up to 1000 mL, and adjust the pH to 7.2.

2. Magnesium Sulphate Solution

Reagents required:

Magnesium sulphate heptahydrate (MgSO₄·7H₂O): 82.5 g
Distilled water: up to 1000 mL

Preparation: Dissolve MgSO₄·7H₂O in distilled water and dilute to 1000 mL.

3. Calcium Chloride Solution

Reagents required:

Calcium chloride (CaCl₂): 27.5 g
Distilled water: up to 1000 mL

Preparation: Dissolve CaCl₂ in distilled water and dilute to 1000 mL.

4. Ferric Chloride Solution

Reagents required:

Ferric chloride hexahydrate (FeCl₃·6H₂O): 0.25 g
Distilled water: up to 1000 mL

Preparation: Dissolve FeCl₃·6H₂O in distilled water and dilute to 1000 mL.

5. Sodium Thiosulfate Solution (0.025 N)

Reagents required:

Sodium thiosulfate (Na₂S₂O₃): 6.205 g
Distilled water: up to 1000 mL

Preparation: Dissolve sodium thiosulfate in distilled water and dilute to 1000 mL.

Test Method

Preparation of Dilution Water

Aerate the required quantity of distilled water by bubbling compressed air for 1–2 days to achieve DO saturation.
Add 1 mL each of phosphate buffer, magnesium sulphate, calcium chloride, and ferric chloride solutions per liter of dilution water.
Mix thoroughly.
For samples lacking sufficient microbial population, add seed—generally 2 mL of settled sewage per liter of dilution water.

Sample Preparation and Pretreatment

Adjust sample pH to approximately 7.0 if highly acidic or alkaline.
Ensure the sample is free from residual chlorine. If chlorine is present, remove it using sodium thiosulfate.

Removal of Residual Chlorine

Take 50 mL of sample and acidify with 10 mL of 1+1 acetic acid.
Add approximately 1 g potassium iodide (KI).
Titrate with sodium thiosulfate using starch as an indicator.
Calculate the amount of sodium thiosulfate required per milliliter of sample and treat the BOD sample accordingly.

If the sample has unusually high DO (above 9 mg/L), reduce it by gentle aeration or agitation.

Sample Dilution

Prepare multiple dilutions to achieve:

At least 2 mg/L DO depletion
Residual DO not less than 1 mg/L after incubation
Approximately 50% DO depletion

Dilution is prepared by siphoning seeded dilution water, adding the required volume of sample, and making up to volume with dilution water.

Suggested Dilutions and BOD Ranges

% Dilution	Expected BOD (mg/L)
0.01	20,000 – 70,000
0.02	10,000 – 35,000
0.05	4,000 – 14,000
0.1	2,000 – 7,000
0.2	1,000 – 3,500
0.5	400 – 1,900
1	200 – 700
2	100 – 350
5	40 – 140
10	20 – 70
20	10 – 35
50	Up to 14
100	1 – 7

Incubation and DO Measurement

Fill labeled BOD bottles with prepared dilutions and stopper immediately.
Measure initial DO (D₀) in one bottle.
Incubate three bottles at 27°C for 3 days with a proper water seal.
Prepare blank bottles using dilution water only.
Determine DO for samples and blanks on day 0 and after 3 days using the Winkler method.

Calculation of BOD

Let:

D₀ = DO of sample on day 0 (mg/L)
D₁ = DO of sample after 3 days (mg/L)
C₀ = DO of blank on day 0 (mg/L)
C₁ = DO of blank after 3 days (mg/L)


BOD (mg/L) = {(D₀ - D₁) - (C₀ - C₁)}/{Decimal fraction of sample used}}

If the sample is seeded, determine the BOD contribution of the seed separately and apply the appropriate correction.

Final Remarks

BOD testing provides critical insight into the biodegradable organic load of water and wastewater. Accurate dilution, proper incubation, and careful DO measurement are essential for reliable results. When followed correctly, this method remains a cornerstone of environmental water quality assessment.

🌱 Healthy microbes tell the true story of water quality.

Monday, March 15, 2021

Iron (Fe) Analysis

Mastering Iron Detection in Water: The Phenanthroline Colorimetric Method

Iron is a common element in water sources and can cause aesthetic and operational issues, including rusty stains and metallic taste. Accurate measurement of iron in water and wastewater is essential for environmental monitoring, industrial processes, and safe water supply. The Phenanthroline colorimetric method is a reliable, sensitive, and easy-to-use technique for quantifying iron concentrations.

Why Measure Iron?

High iron levels in water are more than just cosmetic:

Cause orange-brown stains on plumbing, laundry, and utensils
Affect taste and color in food and beverages
Influence industrial processes and equipment longevity
Impact environmental quality for agriculture and ecosystems

Monitoring iron ensures compliance with regulatory limits and helps maintain water quality.

Principle of the Phenanthroline Method

This method measures iron by forming a colored complex:

Iron in the sample is converted to the ferrous form (Fe²⁺) using acid and a reducing agent like hydroxylamine.
At pH 3.2–3.3, 1,10-phenanthroline binds Fe²⁺, forming an orange-red tris-phenanthroline complex.
The color intensity is proportional to iron concentration (Beer's Law) and can be measured spectrophotometrically at 510 nm.

The method is robust, with stable complexes across pH 3–9 and rapid color development between pH 2.9–3.5.

Roles in the Lab

Lab Chemist: Handles sample preparation, treatment, and measurement
Technical Manager: Reviews procedures and ensures accuracy
Quality Manager: Oversees SOP compliance and quality assurance

This method applies to water and wastewater samples where precise iron measurement is required.

Equipment Needed

Spectrophotometer (set at 510 nm)
Conical flasks, volumetric flasks, pipettes
Glass beads (for boiling)

Reagents

Hydroxylamine Solution: 10 g hydroxylamine hydrochloride in 100 mL distilled water
Ammonium Acetate Buffer: 250 g ammonium acetate in 150 mL water + 700 mL glacial acetic acid
Sodium Acetate Solution: 200 g in 800 mL water
Phenanthroline Solution: 100 mg 1,10-phenanthroline monohydrate in 100 mL water at 80°C
0.1M Potassium Permanganate: 0.316 g KMnO₄ in 100 mL water
Stock Iron Solution: 1.404 g ferrous ammonium sulfate + 20 mL concentrated H₂SO₄ in 1 L water
Standard Iron Solutions: Dilutions from stock solution (e.g., 10 µg/mL or 1 µg/mL)

Safety: Handle acids and reagents with gloves and fume hood protection.

Procedure

Sample Preparation: Filter cloudy samples into clean flasks.
Calibration Curve: Prepare iron standards (1–5 mg/L Fe). Zero spectrophotometer with blank and measure absorbance.
Sample Treatment: Pipette 50 mL of sample into a 100 mL volumetric flask. Add 2 mL HCl, 1 mL hydroxylamine, and a few glass beads. Boil until reduced to 15–20 mL.
Cooling and Color Development: Cool, add 10 mL ammonium acetate buffer and 4 mL phenanthroline solution. Dilute to 100 mL with water and wait 10 minutes.
Measurement: Measure absorbance at 510 nm. Dilute samples if readings exceed the standard range.
Blank Check: Run a blank with distilled water instead of sample.

Calculation

Fe (mg/L) = (µg Fe in 100 mL final solution) ÷ mL of sample

This formula gives the iron concentration directly in the sample.

Why Use This Method?

The Phenanthroline method is preferred because it is:

Sensitive and accurate
Easy to perform with minimal equipment
Applicable to a wide range of water and wastewater samples
Provides reliable results that comply with regulatory standards (APHA, ISO)

Monitoring iron ensures safe water, prevents corrosion, and improves aesthetic quality. High levels can be treated using filtration or oxidation methods.

Pro Tip: Always follow lab safety protocols and verify your results against official standards.

Mastering this method allows scientists, students, and environmental technicians to measure iron accurately and maintain water quality effectively.

Tuesday, February 2, 2021

Ammonical Nitrogen Testing in waste water

Monitoring Ammonical Nitrogen in Water and Wastewater

Monitoring ammonical nitrogen (NH₃‑N) in water and wastewater is essential for evaluating pollution levels, treatment efficiency, and compliance with environmental regulations. Elevated ammonia concentrations can be toxic to aquatic organisms and often indicate contamination from sewage, industrial discharges, or agricultural runoff.

This article presents a clear, laboratory‑based Standard Operating Procedure (SOP) for determining ammonical nitrogen using the distillation and titrimetric method, a widely accepted and reliable analytical technique.

Why Measure Ammonical Nitrogen?

Ammonical nitrogen represents ammonia and ammonium compounds present in water. High levels may:

Indicate contamination from domestic or industrial wastewater
Cause toxicity to fish and other aquatic life
Interfere with drinking water treatment processes
Signal incomplete biological treatment in wastewater plants

Accurate measurement of ammonical nitrogen is therefore critical for environmental monitoring, regulatory compliance, and treatment process control.

Objective of the Method

The objective of this method is to determine the concentration of ammonical nitrogen (NH₃‑N) in water and wastewater samples using a standardized laboratory procedure that ensures accuracy, reliability, and repeatability of results.

Principle of Analysis

The method is based on the alkaline distillation of ammonia from the sample. The sample is buffered to a pH of approximately 9.5 using a borate buffer, which minimizes the decomposition of cyanates and organic nitrogen compounds.

Under alkaline conditions, ammonia is liberated and distilled into a receiving solution of boric acid. The absorbed ammonia is then determined by titration with standard sulfuric acid using a mixed indicator. The volume of acid consumed is directly proportional to the ammonical nitrogen content of the sample.

Apparatus and Equipment

The following laboratory equipment is required:

Pipettes
Conical flasks
Nitrogen distillation assembly
Heating mantle

Reagents

All reagents should be of analytical reagent grade.

Sodium Tetraborate (0.025 M): Dissolve 9.5 g Na₂B₄O₇·10H₂O in 1 L distilled water
Borate Buffer: Mix 500 mL of 0.025 M sodium tetraborate with 88 mL of 0.1 N NaOH
Sodium Hydroxide (6 N): Dissolve 240 g NaOH in 1 L distilled water
Mixed Indicator Solution: Methyl red and methylene blue dissolved in ethanol or propanol
Indicating Boric Acid Solution: Dissolve 20 g H₃BO₃ and add 10 mL mixed indicator; dilute to 1 L
Sulfuric Acid (0.1 N and 0.02 N): Prepare and standardize as required
Sodium Carbonate (0.05 N): Dissolve 2.5 g in 1 L distilled water

Test Method

Distillation Procedure

Take 250 mL of dechlorinated sample or dilute the sample to 250 mL.
Adjust the pH to approximately 7.0, if required.
Add 25 mL borate buffer and adjust the pH to 9.5 using 6 N NaOH.
Assemble the distillation unit and begin distillation at a rate of 6–10 mL per minute.
Collect the distillate in a 500 mL conical flask containing 50 mL indicating boric acid solution.
Perform a reagent blank under identical conditions and apply blank correction.

Titration

Titrate the collected distillate with 0.02 N sulfuric acid until a pale lavender endpoint is observed.

Calculation

Ammonical Nitrogen (mg/L) is calculated using the formula:

Ammonical Nitrogen (mg/L) = ((A − B) × N × 14 × 1000) / Sample Volume (mL)

Where:

A = Volume of sulfuric acid used for the sample (mL)
B = Volume of sulfuric acid used for the blank (mL)
N = Normality of sulfuric acid
14 = Atomic weight of nitrogen

Example Calculation

Given:

A = 12.5 mL
B = 0.5 mL
N = 0.02 N
Sample volume = 250 mL

Step 1: Blank correction 12.5 − 0.5 = 12.0 mL

Step 2: Multiply by normality 12.0 × 0.02 = 0.24

Step 3: Multiply by atomic weight of nitrogen 0.24 × 14 = 3.36

Step 4: Convert to mg/L 3.36 × 1000 / 250 = 13.44 mg/L

Conclusion

The distillation–titrimetric method is a reliable and widely accepted approach for determining ammonical nitrogen in water and wastewater. Strict adherence to this SOP ensures accurate results, supports regulatory compliance, and aids in effective water quality assessment. Consistent monitoring of ammonical nitrogen plays a vital role in protecting aquatic ecosystems and ensuring sustainable water management practices.

Wednesday, December 23, 2020

Total Hardness Analysis in Water

Total Hardness Determination (EDTA Titrimetric Method)

1. Purpose

This Standard Operating Procedure (SOP) describes a validated method for the determination of Total Hardness in water and wastewater samples using the EDTA titrimetric method. The procedure is designed to produce accurate, precise, and reproducible results suitable for routine laboratory analysis and regulatory monitoring.

2. Scope

This method applies to the analysis of drinking water, surface water, groundwater, and wastewater samples in which calcium (Ca²⁺) and magnesium (Mg²⁺) are the primary contributors to hardness.

3. Principle of the Method

Total hardness in water is caused mainly by dissolved calcium and magnesium salts. The determination is based on complexometric titration using Ethylenediaminetetraacetic Acid (EDTA) as the titrant.

At a controlled pH of 10.0 ± 0.1, calcium and magnesium ions react with Eriochrome Black T (EBT) indicator to form a wine‑red colored complex. During titration, EDTA preferentially complexes with calcium and magnesium ions. Once all metal ions are bound by EDTA, the indicator is released, producing a distinct color change from wine‑red to sky blue, which signifies the endpoint.

4. Responsibilities

Laboratory Analyst: Perform the analysis in accordance with this SOP and accurately record all observations and results.
Laboratory Supervisor: Ensure availability, calibration, and validation of equipment, reagents, and standards.

5. Apparatus and Equipment

The following equipment is required:

Calibrated pH meter
Burette (50 mL capacity)
Conical flasks (100 mL)
Volumetric flasks
Measuring cylinders
Analytical balance
Standard laboratory glassware

6. Reagents and Chemicals

All reagents shall be of Analytical Reagent (AR) grade.

6.1 Distilled Water

Used for reagent preparation and dilution.

6.2 Buffer Solution (pH 10)

Dissolve 16.9 g ammonium chloride (NH₄Cl) in 143 mL ammonium hydroxide, and dilute to 250 mL with distilled water.

6.3 Eriochrome Black T Indicator

Thoroughly mix 0.5 g Eriochrome Black T with 100 g sodium chloride (NaCl).

6.4 Inhibitor Solution (If Required)

For samples containing interfering ions, dissolve 4.5 g hydroxylamine hydrochloride in 100 mL of 95% ethyl alcohol or isopropyl alcohol.

6.5 Standard EDTA Solution (0.01 M)

Dissolve 3.723 g EDTA in distilled water and dilute to 1000 mL. Standardize the solution against 0.01 M zinc sulfate solution.

6.6 Standard Zinc Sulfate Solution (0.01 M)

Dissolve 2.8754 g ZnSO₄ in distilled water and make up to 1000 mL.

6.7 Standard Calcium Solution (1 mL = 1 mg as CaCO₃)

Dry analytical‑grade calcium carbonate at 180°C for 1 hour. Accurately weigh 1.000 g, dissolve using minimal concentrated hydrochloric acid, boil briefly, cool, add methyl red indicator, neutralize to an orange endpoint using 3N ammonium hydroxide, and dilute to 1000 mL with distilled water.

7. Standardization of EDTA Solution

Pipette 50 mL of standard calcium solution into a conical flask.
Add 1 mL buffer solution.
Add 1–2 drops of Eriochrome Black T indicator.
Titrate with EDTA solution until the color changes from purple to sky blue.
Record the burette reading and calculate the exact normality of the EDTA solution.

8. Sample Analysis Procedure

Transfer 50 mL of the water sample (or a suitable aliquot) into a conical flask.
If interference is anticipated, add 1 mL hydroxylamine hydrochloride solution.
Add 1–2 mL buffer solution to adjust the pH to 10.0–10.1.
Add 2–3 drops of Eriochrome Black T indicator; the solution will turn wine‑red.
Titrate with standardized EDTA solution, stirring rapidly initially and slowly near the endpoint.
Note the endpoint indicated by a color change from wine‑red to sky blue.

9. Calculation

Total Hardness (as CaCO₃), mg/L


Total Hardness = {A x N x 1000 x 100}/{V}

Where:

A = Volume of EDTA used (mL)
N = Normality of EDTA solution
V = Volume of sample taken (mL)

10. Process Flow (Summary)

Measure 50 mL of sample into a conical flask
Add buffer solution
Add Eriochrome Black T indicator (wine‑red color)
Titrate with 0.01 N EDTA solution
Observe endpoint color change to sky blue
Record titration volume
Calculate total hardness in mg/L

11. Precautions

Maintain pH strictly at 10 to ensure accurate endpoint detection.
Use freshly standardized EDTA solution.
Ensure all glassware is free from metal contamination.

12. Conclusion

The EDTA titrimetric method is a well‑established and reliable technique for determining total hardness in water and wastewater samples. When performed under controlled conditions, it provides precise and reproducible results essential for water quality assessment and effective treatment process control.