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ANALYSIS OF REAL SAMPLE (Chem3118) LAB MANUAL
Abdu Hussen Ali Email: abdelmelik9@gmail.com
JULY 2021
TULUAWLIYA, ETHIOPIA

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ANALYSIS OF REAL SAMPLE, LABORATORY MANUAL Chem3118
Preface and Acknowledgments
Analysis of Real Sample is dealing with sampling, preservation and preparation of samples for the
determination of the major, trace elements, inorganic compounds (speciation) and organic
compounds; biological samples; food and beverages; water and waste water samples; soils and
related samples. It’s also concerned with developing the tools used to examine these properties.
Thus, it is important that students of chemistry do experiments in the Lab to more fully understand
the theories they study. The manual helps students understand the timing and situations for the
various techniques. Each experiment is presented with concise objectives, a comprehensive list of
techniques, and detailed lab introduction, step-by-step procedures and discussion questions at the
end of each lab. It is also important that you carefully prepared for each experiment by reading the
related text material before coming to the lab. This way you can maximize the laboratory
experience. I encourage you to discuss ideas for improvements or suggestions for new experiments
with me.
This manual was completed with support of chemistry staff members of Mekdela Amba
University. Therefore, I am thankful for their efforts.

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TABLE OF CONTENT
Content Page
preface and Acknowledgments 1
Table of Content 3
Course Policies and Information 4
Learning Outcomes 4
Laboratory Precautions 5
Laboratory Maintenance 6
Some Laboratory Equipments 8
Experiment-1:Spectrophotometric determination of iron in meat through derivation with
ferrozine 12
Experiment-2: Determination of vitamin c content in fruit juice 16
Experiment-3: Spectroscopic analysis of caffeine and benzoic acid in soft drink 22
Experiment-4: Determination of acid content of vinegar 28
Experiment-5: Determination of fluoride ion using an ion selective electrode 34
Experiment-6:Analysis of turbidity, colour, ph, and alkalinity of water 40
Experiment-7:Determination of permanent hardness due to Ca2+
and Mg2+
in tap water by
EDTA method 47
Experiment-8:Determination of chemical oxygen demand (COD) of wastewater using open
reflux method 52
Experiment-9: Soil sample collection and preparation for heavy metal analysis 59
Experiment-10: Determination of manganese in soil sample using FAAS 66
Experiment-11: Determination of soil organic matter by walkley and black method 70
Experiment-12:Determination of extraction efficiency of some organic solvents using soxhlet
extractor 75
Experiment-13:Determination of the lipid content of snack food using soxhlet extraction
method 80
REFERENCES 85

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Course Policies and Information
In this laboratory, you will be working as a team with two or three persons in each group. During
the first lab period your instructor will assign you to a group. You will then introduce yourself to
group members and get to know other members of your group. Then your instructor will read you
important safety rules. In this meeting of the laboratory, you will also be given your first lab
experiment and the rest of this lab period you will work on a plan of action for the first experiment.
Each student in the group must have a lab Notebook and bring it to the lab every week. You should
keep a good notebook with all the calculations and the results because your instructor will grade
your Lab notebooks at the end of each experiment. Finally, each member of your group has to
write a 5 or 6-page lab report after completing the experiment. This is going to be a individual
report therefore, even if your results are the same. The reports you write must be your own work.
If your instructor finds out that your report is exactly the same with another member of your group,
you will not receive any credit for that report and he/she may consider it as cheating.
Learning Outcomes
At the end of this course students should be able to:
x Select appropriate sampling and preservation of a particular real sample
x Identify preparation methods for analysis of metals by different methods
x Perform experiments on water, soil and air
x Familiarize the students with the techniques of sampling, storage, and analysis of
real samples.
Format of the Lab Report
You should prepare your lab reports by handwriting. They should include tables and illustrations
where necessary. Typically, a lab report should contain the following sections: title page,
introduction, experimental section, results and discussion, Conclusion and references. Your
title page should be a separate page including the title of the project which might simply be the
name of the experiment, your name, name of the course and the date the report is due.

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LABORATORY PRECAUTIONS
A. Inside the Laboratory
1. Do not eat, drink beverages or chew gum in the laboratory. Do not use laboratory glassware
as containers for food or beverages
2. Wear safety goggles and aprons
3. Always keep the working area clean and orderly
4. Know the locations and operating procedures of all safety equipment.
5. Notify the instructor immediately of any unsafe condition you observe
B. Handling Chemicals
1. All chemicals in the laboratory are to be considered dangerous. Do not touch, taste or smell
any chemical unless specifically instructed to do so
2. Check the label on chemical bottles twice before removing any of the contents.
3. Never return unused chemicals to their original containers.
4. Acid must be handled with extreme care. ALWAYS ADD ACID SLOWLY TO WATER.
5. Handle flammable hazardous liquids over a pan to contain spills. Never dispense flammable
liquids anywhere near an open flame or source of heat.
C. Handling glassware and Equipment
1. Always lubricate glassware (tubing, thistle tubes, thermometers, etc.) before attempting to
insert it in a stopper.
2. When removing an electrical plug from its socket, grasp the plug, not the electrical cord.
Keep your hands dry when working with electricity.
3. Do not immerse hot glassware in cold water, it may shatter.
4. Report damage electrical equipment immediately.
D. Heating Substances
1. TURN OFF THE GAS AT GAS OUTLET VALVE after using.
2. Never leave a lit burner unattended. Never leave anything that is being heated or is visibly
reacting unattended.
3. Use tongs or heat-protective gloves when holding or touching heated apparatus.

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LABORATORY MAINTENANCE
1. Make sure your Laboratory space(s) is cleaned. Also clean all the equipment’s and returned
them to their assigned positions. Failure to do so will result to a zero grade for the experiment.
NO exceptions please.
2. All glassware must be cleaned before it is put away.
3. Use sponges to clean bench tops and wiping of non-hazardous materials.
4. Laboratory instructors are the ONLY one allowed to clean up corrosive or toxic materials
5. Sweep up broken glassware with a broom and collect with the dust pan and then place in the
special container provided for glasses.
6. No debris of any type should be left in the sink. Put all debris in allocated containers
7. Make sure all drawers are properly closed and locked when necessary.
GENERAL INFORMATION
1. Dispense organic solvents, strong acids and bases and other volatile solvents in the fume hoods.
2. No fee will be collected for broken equipment or glassware. Each broken glassware and
equipment’s will be replaced with two of similar type by the culprit.
LABORATORY TECHNIQUES
1. Use proper utensils such as crucible tongs to hold or move hot items.
2. Make sure there are no flammable materials near you when lightning a burner
3. Add boiling chips to liquids before heating them up. This will help to prevent bumping or boil
over.
4. Place test tubes in a slanting position away from yourself and others when heating liquids. Heat
liquids at the surface of the liquid.
5. Do not heat up a closed system
6. Heat all substances that emit noxious fumes under the hood
7. Use funnel to transfer liquids into a narrow neck container
8. Use a bulb or pump to pipette a liquid. Never use your mouth
9. Avoid smelling anything unless instructed to do so. While sniffing, gently waffle the material
towards your nose when allowed to do so

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10. Do not return excess reagent to its original container
11. Do not experiment with the chemicals in the laboratory except those that your are scheduled
to do.
12. Do not use your pipette or spatula to remove samples from the stock container. Use the one
provided by the laboratory technologist
13. Correctly label test tubes or other containers indicating their contents
14. Strong acids and bases should be added to water and not vice versa.
EMERGENCIES AND FIRES
1. Laboratory instructors are in-charge of all emergencies. Follow instructions as directed
2. All laboratory users should learn how to locate the following materials: safety shower, eyewash,
blankets, fire extinguishers, first aid kit, fire alarm
3. Laboratory users should notify the laboratory instructors of any fire.
4. Turn off all gas jets if it is the source of the fire
5. All laboratory users should learn how to use the fire extinguisher
ACCIDENTS AND INJURIES
1. The chemistry department does not treat injuries or illness. Any injury or illness will be referred
to the University of Mekdela Amba.
2. It is the responsibility of the laboratory instructor(s) on duty to prevent further injury by taking
the appropriate action after the incident. Arrangement should be made to immediately transport
the victim to the Medical Center. If the injury is minor and the student can walk to the Medical
Center, such student should be accompanied by another person to the Medical Center.
3. An accident report form must be filled at all times even when the victim declines Medical
treatment.

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SPECIAL WASTE
1. The laboratory will provide label containers for hazardous waste. Read the label very well and
dispose the waste appropriately.
2. At no time should organic or toxic wastes such as mercury, lead, chromium be dumped down
the drain.
3. Ask when in doubt about proper disposal of waste.
SOME LABORATORY EQUIPMENTS

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EXPERIMENT-1
Spectrophotometric Determination of Iron in Meat through Derivation with
Ferrozine
Objectives: To determine the iron content of food samples using derived UV-Vis
Spectrophotometry
Introduction
Determination of the concentration distribution of soluble reactive species is key to understanding
biogeo-chemical processes in natural settings. Iron is one of the most reactive elements in aquatic
environments, and its cycling is coupled to that of the major biogeo-chemical elements (C, O, S
and P) and trace elements such as heavy metals. It is present in the hydrosphere under two
oxidation states, II and III, which are thermodynamically stable under anoxic and oxic conditions,
respectively.
Chromogens are chemicals that react with compounds of interest and form colored products that
can be quantified using spectroscopy. Several chromogens that selectively react with minerals are
available. In this case, ferrozine is used to measure ferrous iron in an ashed food sample. The meat
samples are first ashed to dissociate the iron bound to proteins, and thus ash residue is solubilized
in dilute HCl. The acid is necessary to keep the mineral in solution. Ferrozine complexes with
ferrous iron but not with ferric iron. Prior to the reaction with ferrozine, the solubilized ash is first
treated with ascorbic acid to reduce all forms of ferric iron to the ferrous form. This step is
necessary with ashed samples as this procedure would be expected to reduce all the iron present
in the meat.
Spectroscopic analysis is based on the change in the intensity of the colour of a solution with
variations in concentration. These methods represent the simplest form of absorption analysis. The
human eye is also used to compare the colour of the sample solution with a set of standards until
a match is found. An increase in sensitivity and accuracy results when a spectrophotometer is used
to measure the absorbace. Basically, it measures the fraction of an incident beam of light which is
transmitted by a sample at a particular wavelength. Iron oxides are dissolved in hot, diluted

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ascorbic acid. Ascorbic acid rapidly reduces Fe(III) to Fe(II) in acidic solution. Ferrozine forms a
stable and colored complex with Fe(II) in the pH range 4-10, and this makes a sensitive
determination of iron possible by means of spectrophotometry.
Principle of Method
Ferrous iron in extracts or ashed samples reacts with ferrozine reagent to form a stable colored
product which is measured spectrophotometrically at 562 nm. Iron is quantified by converting
absorbance to concentration using a standard curve.
Materials and Instruments
Meat sample, 16 Test tubes (18 × 150 mm) porcelain crucible, volumetric flask (250 mL),
pipettes (10 mL,25 mL), muffle furnace, hot plate, spectrophotometer, analytical balance
Chemicals and reagents
¾ Ferrozine, ascorbic acid, ammonium acetate and ferric stock solution
¾ Ferrozine reagent (0.493 g of ferrozine in water and dilute to liter in a volumetric flask
¾ Ascorbic acid (0.02% in 0.2 N HCl, made fresh daily)
¾ Ammonium acetate (30% w/v)
¾ Iron stock solution (10 μg iron/mL)
¾ Solutions of 0.1 N and 0.2 N HCl
PROCEDURE
Ashing
1. Place 5 g sample into the crucible and weigh accurately. Make a triplicate of measurement
2. Heat on the hot plate until the sample is well charred and has stopped smoking.
3. Ash in muffle furnace at 550 Ԩ until the ash is white.

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Iron Measurement
1. Prepare working standards of 10, 8, 6, 4 and 2μg iron/mL from a stock solution of 1 μg
iron/mL using serial dilution. In addition, prepare a blank solution.
2. Dissolve ash in small amount of 1 N HCl, and dilute to 50 mL in volumetric flask with
0.1 N HCl.
3. Put a triplicate of 0.50 mL of appropriately diluted samples and standards into 10 mL test
tubes.
4. Add 1.25 mL ascorbic acid (0.02 % in 0.2 N HCl, made fresh daily). Mix the solution
thoroughly and let it set for 10 minutes.
5. Add 2.0 mL 30 % ammonium acetate and mix the solution well (pH needs to be 3 for
color development)
6. Add 1.250 mL ferrozine (1mM in water). Mix the solution and let set in dark for 15
minutes
7. Group the contents of the two standard water blanks and use this to zero the
spectrophotometer at 562 nm (single beam instrument) or place in the reference position
(dual beam instrument).
8. Take your readings three times
Data and Calculations
I. Calculation of percentage of ash
Calculating the percentage of ash using
% Ash =
ࢃ૛ିࢃ૚
ࢃ࢙
‫כ‬ ૚૙૙
Where, W1: weight of crucibles
W2: weight of crucibles with ash sample
Ws: weight of sample

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II. Absorbances of Standards and sample
(μg iron/mL) Trial- 1 Trial-2 Average
Unknown
2
4
6
8
10
Absorbances of replicate (Rep) Samples
Rep 1 = ___________________
Rep 2 = ___________________
Rep 3 = ___________________
Calculation of total iron in sample:
1. Plot the standard curve and determine the content of iron (μg iron/mL) in the dissolved ash
solution.
2. Calculate the iron (μg iron/g) in the sample using
ࣆࢍ࢏࢘࢕࢔
࢓ࡸ࢕ࢌ࢙࢕࢒࢛࢚࢏࢕࢔
ൈ૞૙࢓ࡸࢇ࢙ࢎ࢙࢕࢒࢛࢚࢏࢕࢔
࢓ࢇ࢙࢙࢕ࢌ࢓ࢋࢇ࢚࢚ࢇ࢑ࢋ࢔
=ൈ
ࣆࢍ࢏࢘࢕࢔
ࢍ࢓ࢋࢇ࢚
Where x stands for mass of iron per dry mass in the sample
Questions
1. What is the purpose on adding ascorbic acid in the solution?
2. Why the sample are ashed before analysis?
3. Which chemical are used to form stable colour products to be spectoro-photometrically
determined?

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EXPERIMENT-2
Determination of Vitamin C content in Fruit Juice
Objective: To determine the content of vitamin c in fruit juice
Introduction
Vitamin C, known chemically as ascorbic acid, is an important component of a healthy diet. The
history of Vitamin C revolves around the history of the human disease scurvy, probably the first
human illness to be recognized as a deficiency disease. Its symptoms include exhaustion, massive
hemorrhaging of flesh and gums, general weakness and diarrhea. Resultant death was common.
Scurvy is a disease unique to guinea pigs, various primates, and humans. All other animal species
have an enzyme which catalyzes the oxidation of L- gluconactone to L-ascorbic acid, allowing
them to synthesize Vitamin C in amounts adequate for metabolic needs.
L-Ascorbic Acid -- Vitamin C
The RDA (Recommended Daily Allowance) for Vitamin C put forward by the Food and Nutrition
Board of the National Research Council is 60 mg/day for adults. It is recommended that pregnant
women consume an additional 20 mg/day. Lactating women are encouraged to take an additional
40 mg/day in order to assure an adequate supply of Vitamin C in breast milk. Medical research
shows that 10 mg/day of Vitamin C will prevent scurvy in adults. There has been much controversy
over speculation that Vitamin C intake should be much higher than the RDA for the prevention of
colds and flu. Vitamin C and the Common Cold, that humans should be consuming around 500
mg of Vitamin C a day (considered by many doctors to be an excessive amount) to help ward off
the common cold and prevent cancer.
Vitamin C is a six carbon chain, closely related chemically to glucose. It is a simple, inexpensive,
four-step process for synthesizing ascorbic acid from glucose. This method has been used for

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commercial synthesis of Vitamin C. Vitamin C occurs naturally primarily in fresh fruits and
vegetables. Vitamin C content of some foodstuffs are described below:
Vitamin-C (mg/100g) Foods
100 – 350 Chili peppers, sweet peppers, parsley, and turnip greens
25 – 100 Citrus juices (oranges, lemons, etc.), tomato juice, mustard greens,
spinach, brussels sprouts
10 – 25 Green beans and peas, sweet corn, asparagus, pineapple,
cranberries, cucumbers, lettuce
10 Eggs, milk, carrots, beets, cooked meat
Vitamin C is a water-soluble, antioxidant vitamin. It is important in forming collagen, a protein
that gives structure to bones, cartilages, muscles, and blood vessels. Vitamin C also aids in the
absorption of iron, and helps maintain capillaries, bones, and teeth. It is the most common
electroactive biological compound and one of the most ubiquitous vitamins ever discovered. Rich
sources include blackcurrant, citrus fruit, leafy vegetables, tomatoes, green and red peppers.
Ascorbic acid is known for its reductive properties. Hence, it is used on a large scale as antioxidant
in food and drinks. Due to its content variation caused by the thermal lability, vitamin C represents
an important quality indicator that contributes to the antioxidant properties of food.
Traditional methods for ascorbic acid assessment involve titration with an oxidant solution:
dichlorophenol indophenol (DCPIP), potassium iodate or bromate. Chromatographic methods,
particularly HPLC with electrochemical detection, has turned out to be a selective and sensitive
method for ascorbic acid assessment in foodstuffs and biological fluids. Fluorimetric methods and
UV-VIS absorbance-based determinations were also used for ascorbic acid estimation. However,
the determination of vitamin C concentration in a solution by a redox titration using iodine is
simple and best method rather than the above method. Vitamin C, more properly called ascorbic
acid, is an essential antioxidant needed by the human body. As the iodine is added during the
titration, the ascorbic acid is oxidised to dehydroascorbic acid, while the iodine is reduced to iodide
ions.
ascorbic acid + I2 → 2 I− + dehydroascorbic acid

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Due to this reaction, the iodine formed is immediately reduced to iodide as long as there is any
ascorbic acid present. Once all the ascorbic acid has been oxidised, the excess iodine is free to
react with the starch indicator, forming the blue-black starch-iodine complex. This is the endpoint
of the titration. The method is suitable for use with vitamin C tablets, fresh or packaged fruit juices
and solid fruits and vegetables.
The amount of ascorbic acid in a fruit juice sample will be determined by titrating a weighed
amount of the sample with iodine. The iodine will immediately react with the ascorbic acid until
all of the ascorbic acid has been exhausted. The next drop of iodine cannot be reduced to iodide
(I-
) and, thus, reacts with the starch causing the solution to turn blue-black. Thus, the amount of
iodine necessary to bring about the color change is an indicator of the amount of ascorbic acid
present in the sample. In a titration procedure a solution of unknown analyte concentration is mixed
with a solution with a known concentration of a compound that reacts with the analyte. (The
analyte is the compound being analyzed; in this experiment it is ascorbic acid.) Measuring the
amount of known solution required to just completely use up the analyte allows the calculation of
the concentration of analyte in the unknown solution. (The known solution is called the titrant.
In this experiment, the titrant is the iodine solution.) Usually a burette is used to measure the
amount of the known solution required.
Material and Instrument
¾ burette and stand
¾ volumetric flask
¾ pipette
¾ measuring cylinders
¾ conical flasks
Chemical and Reagent
¾ 2% Lugol’s iodine solution, acetic acid
¾ commercial fruit juice
¾ Starch solution
¾ Sodium thiosulfate

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Standardization of Iodine
A solution of iodine can be standardised by titration against a known concentration of sodium
thiosulfate according to the following equation.
I2(aq) + 2S2O3
2–
(aq) → 2I–
(aq) + S4O6
2–
(aq)
A solution of sodium thiosulfate has already been prepared for you. The exact concentration is
written on the label. You are going to use this solution to determine the exact concentration of the
I2 solution used for determining the concentration of vitamin C in your fruit.
1. Collect about 200 mL of a solution of I2 in a clean dry stoppered 250 mL conical flask. Prepare
the burette for titration and as shown by your demonstrator, by washing with water and then
three times with a small amount of the I2 solution. Remember to restopper the iodine solution
in your 250 mL flask.
2. In a clean and dry 250 mL conical flask collect 120 mL of the sodium thiosulfate solution.
3. Pipette a 25.00 mL aliquot of the sodium thiosulfate solution into a clean 250 mL conical flask.
Add a half a Ni spoonful of Vitex reagent and 10 drops of 0.2 M acetic acid. Titrate with the I2
solution until you get a permanent colour change for at least 30 seconds. The endpoint is a light
blue colour. Record your initial and final volumes of the titration.
4. Repeat at least twice.
Vitamin C Determination
Your demonstrator will allocate you a fresh juice OR a preserved juice.
Fresh fruit juice preparation
Begin with step (6) if you have been allocated a preserved juice.
1. Weigh 2 clean and dry Petri dishes on the top loading balance. Record their masses.
2. In duplicate, accurately weigh about 5 g of a fruit using a top loading balance. Record the exact
mass.
3. For each sample, cut the fruit into small pieces, place it in the microcloth and squeeze the fruit
juice into a clean 250 mL conical flask via a funnel.
4. Rinse the cloth twice with an additional 5 - 10 mL of deionized water. Squeeze the water through
the cloth and allow the filtrate to mix with that from step (3).

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5. Use a spatula to scrape the remaining pulp onto the pre-weighed Petri dish and reweigh. Record
the mass. Add half a Ni spoonful of Vitex and 10 drops of 0.2 M acetic acid. Continue with step
(10).
Preserved fruit juice
6. Weigh a 10 mL measuring cylinder on the top loading balance and record its mass.
7. Measure 5 mL of fruit juice in the pre-weighed 10 mL measuring cylinder and reweigh it.
8. Place the fruit juice into a clean 250 mL conical flask and dilute with 20 mL of deionised water.
If necessary, filter off any pulp using the microcloth and rinse with deionised water into a clean
250 mL conical flask. Add a spatulaful of Vitex and 10 drops of 0.2 M acetic acid.
9. Repeat steps (6) - (7) to obtain a duplicate sample.
10. Fill the burette with the standardised I2 solution and titrate your fruit juice until a permanent
pale blue colour persists for at least 30 seconds. Record your initial and final volumes of I2.
Data and Calculation
Part 1: Standardization of I2
Titration 1 Titration 2
Initial volume (mL)
Final volume (mL)
Titre (mL) Average=
Concentration of thiosulfate ion = 2.040 × 10–3
M (from bottle)
¾ Calculate the concentration of your I2 solution?

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Part 2: Vitamin C Determination
Sample 1 Smple 2
Mass of juice
Initial volume (mL)
Final volume (mL)
Titre (mL)
[I2] = ------------ (from Part 1)
Reaction: I2(aq) + ascorbic acid → 2I–
(aq) + dehydroascorbic acid
Therefore 1 mol of I2 reacts with 1 mol of ascorbic acid
Molar mass of ascorbic acid (C6H8O6) = 176.12 g mol–1
¾ Calculate the mass of vitamin C in mg per g of fruit or mg per mL of fruit juice for your
assigned sample?
Content of Vitamin C in fruit juice =
‫ܖܗܑܜ܉ܚܜܑܜܕܗܚ܎܌܍ܖܑܕܚ܍ܜ܍܌܋ܖܑܕ܉ܜܑܞ܎ܗܛܛ܉ܕ‬
‫ܖ܍ܓ܉ܜ܍ܔܘܕ܉ܛ܍܋ܑܝܒ܎ܗܛܛ܉ܕ‬
Questions
1. Why are I2 is added to each of our flasks during titrating in this experiment? What is the
function?
2. Why were the fruit juices centrifuged and filtered?
3. Using your average milligrams of Vitamin C per gram or milliliter of product as the
correct value, determine the percent error in the manufacturer or text’s claim (show
calculations)?
4. What can you conclude about the labeling of this preserved juice or reference value? How
do your account for any discrepancies? Does the manufacturer or reference overstate or
understate the amount of Vitamin C in the product? If so, why might they do this? Explain
below.

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EXPERIMENT-3
Spectroscopic analysis of caffeine and benzoic acid in soft drink
Objective: To analysis caffeine and benzoic acid in soft drink
Introduction
Soft drinks come under the category of junk food products. This is because of their nutritional
value is less and fat, sugar, salt and calories contents are high. Soft drinks are manufactured from
carbonated water, sugar cane syrup, caffeine and extract of kola nut and coca leaves. A soft drink
is a beverage that contains carbonate water, sweetener, Flavoring agents, Preservatives like salts
of benzoic acid, Caffeine, Coloring agents etc. Caffeine, is an alkaloid (C8
H10
O2
N4
.H2
O) found in
coffee, tea, cacao, and some other plants. Caffeine is added to soft drinks as a flavoring agent, and
from dietary sources is the most frequently and widely consumed central-nervous-system stimulant
today. Caffeine is a non-polar organic compound, which goes by many names including caffeine,
guaranine, mateina etc. The molecular weight of caffeine is 196.19 g/mol, with a melting point of
238 Ԩ and a sublimation point at 178 Ԩ. The chemical name for caffeine is 1, 3, 7-
trimethylxanthine or 1, 3, 7-trimethyl-2, 6-dioxopurine.
The caffeine drug increases the blood pressure, stimulates the central nervous system, promotes
urine formation and stimulated the action of the heart and lungs. Caffeine is used in treating
migraine because it constricts the dilated blood vessels and thereby reduces the pain. It also
increases the potency of analgesics such as aspirin, and it can somewhat relieve asthma attacks by
widening the bronchial airways. Caffeine produces increased mental alertness and reduces fatigue
and increases the heart rate slightly. It is relatively nontoxic, but clearly has addictive potentials.
Other withdrawal symptoms in heavy users include fatigue and difficulty in concentration. Over
use can lead to insomnia, gastrointestinal disturbances and hypertension. Caffeine is also useful
in pain medication. Butalbital, sedative barbiturate drug, is typically combined with other
ingredients such as aspirin, acetaminophen, caffeine or codeine, to create a medication used to
relieve headache or muscle pain in the neck and shoulders. It works by decreasing the activity of
the central nervous system. Propoxyphene is also a drug containing caffeine and other ingredients
such as aspirin and acetaminophen. This drug is used to treat mild to moderate pain. It is a narcotic

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that works by depressing the activity of the central nervous system to create an anesthetic or pain
killing effect.
Benzoic acid (C6
H5
COOH) is the simplest of carboxylic acids of the aromatic series. It is used as
a food preservative. Food preservation is the prevention of chemicals decomposition and the
development of harmful bacterial in foods. Generally affected by the sterilization of the food (that
is by the destruction of bacterial in it) is by heating in sealed vessels or making the conditions
unfavorable for the development of bacterial. Yeast, a unicellular micro-organism producing
zymase, converts sugars (hexose) into alcohol and carbon dioxide. This is because benzoic acid
inhibits the growth of yeast and moulds. Benzoic acid is also a white crystal sparingly soluble in
cold water, moderately soluble in hot water, having melting point of 122.4 Ԩ and sublimes if
rapidly heated. High amount of benzoic acid (added as preservative) is harmful for liver and it
disturbs carbohydrate metabolism which may lead to accumulation of fat causing obesity and
impairment of liver also affects removal of toxic waste materials from body which leads to several
metabolic disorders. Thus consumption of soft drinks having large quantity of benzoic acid causes
severe health hazards.
Both caffeine and benzoic acid are aromatic in nature and absorb ultraviolet radiation. So for the
estimation of Caffeine and benzoic acid spectrophotometric method has been used.

24 | P a g e
By: Abdu H. (MSc)
Material and instrument
¾ UV/Visible spectrophotometer
¾ Chemical balance
¾ Measuring cylinder
¾ Magnetic stirrer
¾ Filter
¾ Beakers
¾ Thermometer
¾ Separatory funnel
¾ Funnel
¾ Erlenmeyer flask
Chemical and Reagents
¾ Benzoic acid standard
¾ Caffeine standard
¾ Sodium carbonate
¾ Methylene chloride
¾ Magnesium sulphate
¾ Potassium chlorate
¾ Hydrochloric acid
¾ Ammonia solution
Sampling
¾ The following soft drinks are taken from a supermarket in: Fanta, Pepsi, Miranda and coca
cola.
Preparation of a stock solution
1. A 0.01 g of benzoic acid standard and caffeine standard were weighed separately and dissolved
into two different 100 mL volumetric flasks and topped with distilled water to the 100 mL mark.
2. Prepare 5 ppm, 10 ppm, 15 ppm, 20 ppm, and 25 ppm from the stock solution.
3. Pipette 0.5 mL of stock solution into 50 mL volumetric flask and diluted with distilled water to
the 50 mL mark.

25 | P a g e
By: Abdu H. (MSc)
4. Pipette 0.75 mL of stock solution into 50 mL volumetric flask and diluted with water to the 50
mL mark.
5. Pipette 1mL stock solution into 50 mL volumetric flask and diluted with distilled water to the
50 mL mark.
6. Pipette 1.25 mL of stock solution into 50 mL volumetric flask and was topped with distilled
water to the 50 mL mark.
Extraction of caffeine
1. A brand of soft drink was recorded and 150 mL was carefully measured into a conical flask.
2. To this, 2.0 g of sodium carbonate was added to it. The mixture was then tested with a red litmus
paper and the litmus registered a blue color.
3. A 50 mL of methylene chloride was added to the mixture in the conical flask and the flask
swirled gently for at least 5 minutes and poured into a separating funnel and allowed to settle
for about 5 -10 minutes.
4. The organic layer was drained into a clean 250 mL conical flask.
5. A fresh 50 mL sample of methylene chloride was added to the mixture in the separating funnel
and the flask stoppered.
6. The funnel was gently inverted a few times to allow the remaining caffeine to be extracted into
the methylene chloride layer.
7. The lower layer was then separated and combined with the first extract.
8. The total was treated with 5g anhydrous magnesium sulphate to remove water.
9. The methylene chloride was filtered through a cotton pad into a 250 mL Erlenmeyer flask.
10. The extract in the Erlenmeyer flask was placed on a water bath to evaporate methylene
chloride. A small amount of the precipitate was placed on a watch glass and mixed with 2 -3
drops of concentrated hydrochloric acid.
11. Few crystals of potassium chlorate were mixed well with a glass rod and mixture evaporated
to dryness on a water bath.
12. The watch glass was cooled and moistens with a drop of 2M ammonia solution.
13. Residue turned purple which was an indication of the presence of caffeine.
14. The rest of the precipitate was diluted with methylene chloride and taken to the UV for the
absorbance to be taken.

26 | P a g e
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Isolation of Benzoic Acid
1. 150 mL of soft drink was poured into a conical flask and acidified with 2 drops of dilute
hydrochloric acid.
2. 50 mL of methylene chloride was added and the flask swirled gently for at least 5 minutes.
3. The mixture was transferred into 250 mL separating funnel and allowed to settle for about 5 -
10 minutes.
4. The organic layer was drained into a beaker and allowed to evaporate on a water bath, leaving
a residue of benzoic acid.
5. The residue was diluted with methylene chloride and sent to the UV for the absorbance to be
taken.
Part 1: Standard caffeine and benzoic acid solution
Standard concetration
of caffeine (ppm)
UV absorbace Standard concetration
of benzoic acid (ppm)
UV absorbace
5 5
10 10
15 15
20 20
25 25

27 | P a g e
By: Abdu H. (MSc)
Part 2: Caffeine and benzoic acid in soft drink
Name of soft drink Average absorbace (benzoic
acid)
Average absorbace
(Caffeine)
Mirinda
Fanta
Pepsi
Coca cola
¾ Calculate the concentration of benzoic acid and caffeine in each soft drink?
Questions
1. What is the difference between caffeine and benzoic acid?
2. What is the purpose of caffeine and benzoic acid found in soft drink?
3. Calculate and compare the amount of benzoic acid and caffeine in each soft drink?

28 | P a g e
By: Abdu H. (MSc)
EXPERIMENT-4
Determination of acid content of vinegar
Objective: To determine the acetic acid content of vinegar
Introduction
Vinegar is an acidic liquid, which is made both naturally and synthetically, from the oxidation of
ethanol, CH3CH2OH, in an alcohol-containing liquid such as wine, fermented fruit juice (e.g.
cider) or beer. It has been used since ancient times as an important cooking ingredient, e.g. in salad
dressings and on fish and chips. The key chemical component of vinegar is acetic acid, CH3COOH
(systematic name: ethanoic acid). The trivial name, acetic acid, is derived directly from the word
for vinegar, which, for example, in Italian is called aceto. The word vinegar itself derives originally
from the Latin vinum aegrum, meaning “feeble wine”. In fact, when a wine has “gone off” and has
acquired a sour taste, this is due to the oxidation of the ethanol in the wine to acetic acid. (The
“corking” of wine, i.e. tainting of the wine by compounds transferred from or through the cork, is
due to a totally different chemical process.) The acetic acid content of vinegar can vary widely,
but for table vinegar it typically ranges from 4 to 8 % v/v.
To determine the amount of acetic acid in vinegar (typically 4-8% by mass) we will use titration.
Titration is a common analytical method used to measure the amounts of compounds in solution.
The glassware you will be using is called a buret. The buret holds one of the reactants, called the
titrant, and conveniently adds it into a reaction vessel which contains the second reactant. The
titrant in this experiment will be a sodium hydroxide (NaOH) solution. In this experiment we titrate
acetic acid with sodium hydroxide (a strong base). The reaction of acetic acid with sodium
hydroxide is shown below:
HC2H3O2 (aq) + NaOH (aq) → NaC2H3O2(aq) + H2O (l)
acetic acid sodium acetate
This equation is an acid-base reaction; also know as a neutralization reaction. The acetic acid
(HC2H3O2) found in the vinegar will react with the NaOH until all of the acetic acid is neutralized.
When an acid, such as acetic acid reacts with a base like NaOH, the products are a salt (NaC2H3O2,

29 | P a g e
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sodium acetate) and water (H2O). If you know the concentration of the sodium hydroxide solution
and the volume that you need to add to the acid, then you can figure out how much acetic acid is
in the vinegar.
In an acid-base titration, the point at which both acid and base have been completely consumed
can be detected and is known as the “equivalence point”. The amount of one reactant (the analyte)
can be calculated from the known concentration and the volume of reactant in a standard solution
(the titrant) using the balanced chemical equation. The end point in this experiment will be detected
with an acid/base indicator. An acid/base indictor is a coloured substance with two or more
different colors depending on the value of the pH of the solution. The standard solution may be
prepared in two ways – the direct or indirect method. In the direct method, a precisely weighed
quantity of the pure solute (primary standard) is dissolved and diluted to a known volume in a
volumetric flask. The concentration of the standard solution is then calculated from the known
mass of the solute and the known volume of the solution. If the solute used to prepare the standard
solution is pure and the solution is stable (does not decompose), then the compound is referred to
as a primary standard.
However, often it is not possible to obtain the solute in sufficiently pure form to be suitable as a
primary standard. For example, NaOH(s) reacts with gases (H2O and CO2) in the air which means
that NaOH (s) is not pure enough to be used as a primary standard. In this case the standard solution
is prepared by an indirect method. A solution is prepared at approximately the desired
concentration and it is then standardized against another primary standard to determine its exact
concentration.
¾ 50 mL beaker
¾ 100 mL volumetric flask (with cap)
¾ 50 mL burette
¾ 250 mL Erlenmeyer flask
¾ 25 mL pipette
¾ Titration apparatus

30 | P a g e
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¾ KHC8H4O4
¾ phenolphthalein indicator solution
¾ Distilled water
¾ Sample of commercial grade vinegar (may be obtained from home)
Preparation of the Potassium Hydrogen Phthalate Standard
For the titration of the vinegar in this experiment the following specific reaction will be used to
calculate the acetic acid content of the vinegar sample:
HC2H3O2 (aq) + NaOH (aq) → H2O (l) + NaC2H3O2 (aq)
1. Weigh precisely (on an electronic balance) ~ 1.5 grams of pure potassium hydrogen phthalate
(KHC8H4O4) into a 50 mL beaker.
2. Dissolve the acid in ~50 mL of water and transfer carefully into a 100 mL volumetric flask.
3. Rinse the beaker several times with small portions of water to ensure quantitative transfer.
4. Make the volume up to the mark with distilled water, cap the flask and mix thoroughly.
This solution will be used to standardize a solution of sodium hydroxide.
Standardization of the Sodium Hydroxide Solution
Potassium hydrogen phthalate, the primary standard, reacts with sodium hydroxide as shown
below:
1. Select a clean 50 mL burette, rinse it with a small portion of the sodium hydroxide solution, and
fill it to just below the zero mark.
2. Read and record the initial volume to the nearest 0.01 mL.
3. Rinse a 250 mL Erlenmeyer flask with distilled water to make sure it is clean.

31 | P a g e
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4. Pipette exactly 25.00 mL of the potassium hydrogen phthalate solution into the flask.
5. Add three drops of phenolphthalein indicator and titrate with the sodium hydroxide solution
until a permanent colour change is first detected. The palest pink colour denotes the end point
of this reaction. The colour should persist throughout the entire solution when swirled for at
least 10 seconds. The colour will fade slowly on standing.
6. Repeat the titration on a second 25.00 mL aliquot of the primary standard solution. If the
titration volumes do not agree within ±0.1 mL continue to repeat the titration.
7. Report your two best titrations.
The Determination of Acetic Acid in Vinegar
The acetic acid (CH3COOH) concentration in commercial vinegar may be easily determined by
titrating a suitable sample of the vinegar with the standardized sodium hydroxide solution.
1. Pipette exactly 10.00 mL of the commercial vinegar sample into a 250 mL Erlenmeyer flask
and add ~5 mL of distilled water.
2. Using three drops of phenolphthalein indicator, titrate the acetic acid with the standard base to
a pale pink equivalence point. Record the burette readings.
3. Repeat the titration at least once more using a fresh aliquot of vinegar. Results should agree
within ±0.2 mL or additional titrations are required.
4. Report your two best titrations.

32 | P a g e
By: Abdu H. (MSc)
Part A: Standardization of sodium hydroxide solution
Titration 1 2
Mass of beaker (g)
Mass of beaker + KHP (g)
Mass of KHP (g)
Volume of NaOH to neutralize the
KHP solution (mL)
Part B: Molarity of acetic acid and percent of vinegar
No Volume of NaOH used (mL)
Titration 1
Titration 2
Calculations
1. Determine the number of moles of sodium hydroxide required to titrate the vinegar for each
titration from the known molarity and the titration volume (V = V2 - V1) of sodium hydroxide.
Be sure that the volume of the sodium hydroxide has been converted from milliliters to liters (1
L = 1000 mL).
MNaOH =
௠௢௟௘௦ே௔ைு
௏௢௟௨௠௘ே௔ைு
so moles NaOH= (VNaOH) x (MNaOH)
2. The moles of acetic acid are equal to the moles of sodium hydroxide at the equivalence point.
The equivalence point is close to the endpoint so we can use the endpoint value. The endpoint
is when the phenolphthalein changes color.
moles acetic acid = moles sodium hydroxide
3. Determine the mass of acetic acid present in each titration from the molar mass (sometimes
called molecular weight) of acetic acid and moles of acetic acid.
Molar mass =
࢓ࢇ࢙࢙
࢓࢕࢒ࢋ࢙
so mass of acetic acid = molar mass x moles
4. The mass of the vinegar for each titration is found from the measurements using the

33 | P a g e
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analytical balance (m2 - m1).
5. The percentage mass of acetic acid in the vinegar is found from the mass of acetic acid and the
mass of vinegar.
% mass =
࢓ࢇ࢙࢙࢕ࢌࢇࢉࢋ࢚࢏ࢉࢇࢉ࢏ࢊ
࢓ࢇ࢙࢙࢕ࢌ࢜࢏࢔ࢋࢍࢇ࢘
x 100
Questions
1. How accurately does the 50 mL of water used to dissolve the KHP in the standardization
of the NaOH solution need to be measured? Explain.
2. When transferring the KHP in Part A, if some of the KHP missed the opening to the
Erlenmeyer flask and fell onto the weighing pan and stayed there, how would the calculated
molarity of the NaOH solution compare to the actual value (i.e., is the calculated
concentration more, less or the same as the actual value)? Explain.
3. During the titration of KHP in part A of this experiment, you obtain a dark pink endpoint
(instead of a pale pink endpoint). Will this result in the calculated molarity of the NaOH
solution being higher, lower or the same as the actual molarity? Explain your answer
4. How does obtaining a dark pink endpoint (instead of a pale pink endpoint) in the titration
in Part B, affect the calculated mass % of acetic acid in vinegar compared to the actual
value (what it should be)? That is, is the calculated mass % greater than, less than or equal
to the actual value? Explain.

34 | P a g e
By: Abdu H. (MSc)
EXPERIMENT-5
Determination of Fluoride Ion Using an Ion Selective Electrode
Objective: To determine the F- ion concentration in toothpaste
Introduction
The primary purpose of brushing the teeth with dentifrice is to clean the accessible tooth surface
of dental plaque, stains and food debris. Fluoride (F-) is an important anion present in various
environments, clinical and food samples. In many countries, fluoride is purposely added to the
water supply (water fluoridation) as sodium fluoride (NaF) and to toothpastes in 0.1%
concentration as sodium monofluorophosphate, Tin difluoride or sodium fluoride {Na2 3POF,
SnF2, NaF. In topical fluoride agents are the main dental products
used in caries prevention. Though a small amount of fluoride is beneficial, and has been used to
treat osteoporosis, fluoride causes mottled teeth and bone damage at about 5mg L-1
when it is
present in water.
Studies have shown that bone cancer in male children and uterine cancer deaths are linked to water
fluoridation due to fluoride’s gradual build up in the bones thereby causing adverse changes to the
bone structure. Recent independent research has shown that fluoride build up in the brain of
animals when exposed to moderate levels of fluoride. Two new epidemiological studies have also
confirmed fluorides’ neurotoxic effects on the brain, as children exposed to higher levels of
fluoride had lower IQs., showed that rats drinking 1ppm fluoride (NaF) in water had histologic
lessions in their brain similar to Alzheimer’s disease and dementia. Fluoride has also been reported
to cause birth defects and perinatal deaths, impaired immune system, acute adverse reactions,
severe skeletal fluorosis at high levels, osteo- arthritis, acute poisoning and contributes to the
development of repetitive stress injury.
The determination of fluoride concentration in the various samples requires very sensitive
methods. Electro analysis, spectra analysis, chromatography and miscellaneous methods with
various adaptations have been employed in the analysis of fluoride. In many recent applications,
ion-selective electrode (ISE) methods are replacing existing time consuming and expensive

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analytical methods with resulting increases in efficiency and simplicity of measurement. They are
cost-effective, and sufficiently sensitive, selective, accurate and precise. The fluoride selective
electrode is a solid-state type electrode consisting of a lanthanum fluoride crystal sealed over the
end of an inert plastic tube which contains an internal electrode and filling solution usually, 0.1M
NaCl and 0.1M NaF. A potential arises because of the difference in fluoride activity on either side
of the crystal. The ionic strength and the pH of sample and standard solutions should be matched
when determining F- concentration using F-
ISE.
Ion-selective electrode (ISE) is a type of membrane electrode incorporates a special ion-sensitive
membrane which may be glass, a crystalline inorganic material or an organic ion-exchanger. The
membrane interacts specifically with the ion of choice, in our case fluoride, allowing the electrical
potential of the half cell to be controlled predominantly by the F-
concentration. The potential of
the ISE is measured against a suitable reference electrode using an electrometer or pH meter. The
electrode potential is related to the logarithm of the concentration of the measured ion by the Nernst
equation. If the measurements are made with very little current flowing in the cell, the reference
electrode potentials are fixed, and if the sample solution is essentially the same matrix for all
measurements the junction potentials are also unchanged. Then the measured cell potential can be
expressed as ‫ܧ‬௠௘௔௦ ൌ ‫ܭ‬ െ ͲǤͲͷͻͳ͸ ݈‫݃݋‬
௔೔೚೙೔೙೙೐ೝ
௔೔೚೙೚ೠ೟೐ೝ
Where K is a constant and ‘’a’’ is the activity of the analyte ion. The ISE filling solution contains
a large concentration (activity) of the analyte ion and is essentially unchanged during operation of
the electrode (aion inner is fixed). Thus, at 25 Ԩ ‫ܧ‬௠௘௔௦ ൌ ‫ܭ‬ ൅ ͲǤͲͷͻͳ͸ ݈‫݃݋‬ ܽ௜௢௡௢௨௧௘௥
For fluoride ion solutions at 25o
C and constant ionic strength,
]
log[
05916
.
0
F
K
Emeas
Thus, for an ideal fluoride ISE, the cell potential is linearly related to the logarithm of the fluoride
ion con centration and should increase 59.16 mV for every 10-fold decrease in the [F-
]. When the
ionic strength of all standards and samples is constant, the response of a real fluoride ISE is
described by a similar relationship
]
log[
)
05916
.
0
(
F
K
Emeas E
where β is the electromotive efficiency and typically has a value very close to unity ( 0.98)

36 | P a g e
By: Abdu H. (MSc)
Direct Potentiometric Measurement
To check if the electrode is working properly, you will measure the cell potential of three fluoride
standards prepared in a Total Ionic Strength Adjustment Buffer (TISAB). The TISAB contains an
acetic acid/acetate buffer that fixes the pH of the solution at about 5. At this pH the formation of
HF is negligible and the concentration of OH-
, the only other anion that the electrode responds to
is insignificant. It also contains NaCl to establish a high and constant ionic strength, and a
complexing agent that removes cations that could interfere by forming complexes with fluoride.
From a linear least-squares fit to a plot of Emeas versus log [F-
] you can obtain the slope [S =
β(0.05916)]. Typically S equals 56 ± 2 mV.
Method of Standard Addition
The method of variable volume standard addition will be used to determine the fluoride content of
an unknown solution. In this approach, a solution containing fluoride will be mixed with the
TISAB and the potential will be measured. Then successive amounts of a fluoride standard
solution will be added and the potential will be measured after each addition. The following
describes how the unknown fluoride concentration can be obtained from these measurements. The
measured potential (E) can be represented by C
S
K
E log

Where K is a constant;
S: is the slope of the calibration curve and equals β (0.05916) and
C: is the analyte ion (F-
) concentration
The equation can be rearranged to give
S
K
E
C

10
The analyte ion concentration after any addition of the standard is given by
std
o
std
std
o
o
V
V
V
C
V
C
C

Where C0 is the analyte concentration before any standard is added;
V0 : is the volume of the solution before any standard is added;
Cstd: is the concentration of the standard solution;
Vstd: is the volume of standard solution that is added.

37 | P a g e
By: Abdu H. (MSc)
Substituting this expression for C in the previous equation gives
S
K
S
E
std
o
std
std
o
o
V
V
V
C
V
C
/
/
10
10

This equation can be rearranged to give
std
std
S
K
o
o
S
K
std
o
S
E
V
C
V
C
V
V /
/
/
10
10
)
(
10

A plot of 10E
/S
(Vo + Vstd) versus CstdVstd will give a linear plot with an x-intercept (y = 0) equal
to the negative of the amount (μg) of analyte in the solution before addition of the standard. The
analyte concentration (μg/mL) in the original unknown solution (Cunk) can then be determined by
dividing by the volume of the unknown fluoride solution (Cunk).
Procedure
Preparation of Fluoride Standard Solutions
By serial dilution of the 1000 μg/mL fluoride standard solution, prepare 50 mL each of 200, 20
and 2 μg/mL fluoride standards in 50-mL volumetric flasks. After thorough mixing, transfer
each diluted standard solution to a labeled plastic reagent bottle for storage. Calculate the
concentration of each diluted standard using the exact concentration of the stock solution. If
you do not have fluoride standard solution in your lab, you can prepare it from solid NaF dried
at 100 Ԩ for hour .
Calibration of Electrode
1. Carefully pipette 25.0 mL of the most dilute fluoride standard into a 50-mL volumetric flask
and dilute to the mark with the TISAB. Stopper the flask and thoroughly mix the solution.
2. Transfer this solution to a 100 mL plastic beaker. Place the beaker on a stirring plate, add a
magnetic stirring bar and begin stirring at a constant rate.
3. Connect the fluoride ISE to a pH meter and set the meter to the mV mode. Rinse the electrode
with deionized water and blot dry.
4. Lower the electrode into the standard solution and when the reading is stable record the mV
value.

38 | P a g e
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5. Repeat steps 1-4 for each of the remaining fluoride standards.
6. Estimate of the slope (S) from the difference in the mV readings for each factor of 10 of increases
in the fluoride ion concentration. If your value is outside the expected range, consult your lab
instructor.
Analysis of Unknown
1. Accurately weigh about 0.2 g of toothpaste into a 100 mL beaker. Add 10 mL of 1N KCl and about
40 mL of water to the beaker.
2. Boil the mixture gently for 3-5 minutes, breaking up the toothpaste with a stirring rod if necessary.
3. Cool the solution, quantitatively transfer the liquid to a 100 mL volumetric flask and dilute to
volume with KCl
4. Prepare a 500 μg/mL fluoride standard by pipeting 5.0 mL of the 1000 μg/mL fluoride standard
solution into a 10 mL volumetric flask and diluting to the mark with the TISAB.
5. Carefully pipette 50.0 mL of prepared toothpaste which contains the TISAB at the same
concentration as used for the standard calibrations into a 100 mL plastic beaker. Place the beaker
on a stirring plate, add a magnetic stirring bar and begin stirring at a constant rate.
6. Rinse the ISE with deionized water and blot dry.
7. Lower the electrode into the unknown solution; when the reading is stable record the mV value.
8. Pipet 1.0 mL of the 500 μg/mL F-
standard solution into the unknown solution and record the mV
value when the reading is stable.
9. Make three additional 1.0 mL additions of the standard solution and record the mV reading after
each addition as before.
10. When finished, rinse the ISE with deionized water and place it in the storage container.
Calculations
a. Determination of Calibration Slope
1. Using EXCEL, plot the mV reading for the diluted fluoride standards versus the log of the
actual fluoride ion concentration.
2. Fit the data points with a linear least-squares line and from the equation for the line obtain the
slope (S).

39 | P a g e
By: Abdu H. (MSc)
b. Determination of Unknown Concentration by Standard Addition
1. Using the slope determined in ‘’a’’, plot 10E/S
(V0+Vstd) versus CstdVstd. Remember to include
the initial reading with no added standard.
2. Fit the data points with a linear least-squares line and obtain the equation for the line.
3. Use the equation for the line to determine the x-intercept and from this calculate the fluoride ion
concentration in the unknown solution. Report the fluoride ion concentration (μg/mL) in the
unknown solution.
Questions
1. Why is the calibration plotted in log concentration?
2. Explain how to determine the concentration of fluoride by ion selective electrode?
3. What is the importance of TISAB in this analysis? Comment on your results and the
technique in general?

40 | P a g e
By: Abdu H. (MSc)
EXPERIMENT-6
Analysis of Turbidity, Colour, pH, and Alkalinity of Water
Objective: To perform alkalinity, pH, turbidity and colour analysis on a given set of water
samples
Introduction
Turbidity is caused by suspended materials which absorb and scatter light. These colloidal and
finely dispersed turbidity-causing materials do not settle under quiescent conditions and are
difficult to remove by sedimentation. Turbidity is a key parameter in water supply engineering,
because turbidity will both cause water to be aesthetically unpleasant and cause problems in water
treatment processes, such as filtration and disinfection. Turbidity is also often used as indicative
evidence of the possibility of bacteria being present.
Turbidity measurements performed using proprietary nephelometric instruments are expressed as
Nephelometric Turbidity Units (NTU). The nephelometric apparatus is designed to measure
forward scattering of light at 90
o
to the path of an incandescent light beam. Suspended particles
present in a water sample reflect a portion of the incident light off the particle surface. The light
reflected at 90
o
is measured by a photoelectric detector and is compared against light reflected by
a reference standard.
Many surface waters are coloured, due primarily to decomposition of organics, metallic salts or
coloured clays. This colour is considered as apparent colour as it is seen in the presence of
suspended matter, whereas true colour is derived only from dissolved inorganic and organic
matters. Samples can be centrifuged and/or filtered to remove turbidity in order to measure true
colour. Waters which obtain their colour from natural organic matter usually pose no health
hazard. However, because of the yellowish brown appearance of such waters, the consumers may
not find the water aesthetically acceptable. Consumers of highly coloured but already properly
treated water may not believe the water is in fact properly treated. Many processing industries

41 | P a g e
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require low coloured water. PUB requires drinking water to meet the highest desirable World
Health Organisation (WHO) standards of less than 5 colour units.
One standard colour unit is defined as a 1 mg/L concentration of platinum in the form of potassium
chloroplatinate K
2
PtCl
6
. Measurement of colour is by comparison of the sample with standard
colour solutions using a spectrophotometer. A straight line calibration curve is initially developed
by plotting absorbance versus platinum-cobalt colour standard. In practice, the absorbance of a
sample is determined and corresponding concentration is read off the calibrated curve. When
measuring true colour, pre-treatment has to be carried out to remove turbidity. Unfortunately,
either centrifugation or filtration has some effect on true colour. So when reporting the true colour
value, specify the details of the pre-treatment method and its operating conditions. Likewise, the
colour value of water is extremely pH dependent, too, and invariably increases as the pH of the
water is raised. When reporting a colour value, specify the pH at which colour is determined.
pH is a way of expressing the hydrogen-ion concentration of a solution. As acids and bases in
solution dissociate to yield hydrogen ions [H
+
] and hydroxyl ions [OH
-
] respectively, pH is used
to indicate the intensity of the acidic or alkaline condition of a solution. Alkalinity is a measure
of the acid-neutralizing capacity of dissolved substances in water and equals the amount of strong
acid required to lower the solution from initial pH to about 4.5. Many materials may contribute to
the alkalinity of water. For most practical purposes, it is due primarily to presence of salts of weak
acids (mainly bicarbonate and carbonate) and hydroxide (at high pH).
pH and alkalinity are key water quality parameters in environmental engineering practice. In the
water supply and treatment fields, these parameters have great influence on the chemical
coagulation, disinfection and softening processes, and corrosion control for water distribution pipe
networks. Effective chemical coagulation of water, for instance, occurs only within a specific pH
range. Chemicals used for coagulation release, as a by-product of their reactions with water to
form insoluble hydroxide precipitates, hydrogen ions (acid-causing). If unchecked, these hydrogen
ions could lower the pH of the water sufficiently to render the coagulants ineffective. The presence
of sufficient amount of alkalinity in the water can react and remove the hydrogen ions released by
the coagulants, thus buffering the water in the pH range where the coagulant can be effective. In

42 | P a g e
By: Abdu H. (MSc)
pure water, water molecules dissociate into equal amounts of hydrogen and hydroxyl ions (10
-7
moles/L). From the law of mass action, it can be shown that, for pure water at about 25°C:
[OH-
] [H+
] = Kw = 10 -14
The pH value of a solution has been defined to be the negative log of the hydrogen ion
concentration:
pH = log [H+
]
The pH scale runs from 0 to 14, with pH 7 representing neutrality. Acid conditions increase as pH
values decrease, and alkaline (base) conditions increase as the pH values increase. Measurement
of the hydrogen ion concentration is made by pH meters via a glass electrode and a calomel
reference electrode. The alkalinity of water is its quantitative capacity to neutralize acids. The three
major forms of alkalinity ranked in order of their association with high pH values are
(1) hydroxide alkalinity, [OH
-
],
(2) carbonate alkalinity, [CO
3
2-
], and
(3) bicarbonate alkalinity [HCO
3
-
]
Their ability to react with H
+
ends at pH 4.5 when both have turned into carbonic acid (H
2
CO
3
). In
nature, bicarbonates are the major form of alkalinity because they result from the reactions of CO2
on calcium and magnesium rocks. Some CaCO3
(up to about 20 mg/L) may also go into solution
as Ca
2+
and CO
3
2-
ions. For all practical purposes, alkalinity due to other sources in natural waters
may be ignored.
Alkalinity of waters is measured by means of titration with a standard solution of a strong acid
(usually H
2
SO
4
) to designated pHs, and is reported in terms of equivalent CaCO
3
. Alkalinity
depends on the end-point pH or indicator used. Either titration curve technique or colour indicators
can be used for the determination. The alkalinity measurement is based on the titration curve for a
hydroxide-carbonate-bicarbonate mixture, as shown in figure below.

43 | P a g e
By: Abdu H. (MSc)
Titration curve for a hydroxide-carbonate-bicarbonate mixture
For samples whose initial pH is above 8.3, the titration is made in two steps. In the first step, the
titration is conducted until the phenolphthalein indicator end-point is reached (i.e. pH of about 8.3)
with a colour change from pink to colourless. During this first phase, the acid added to the sample
reacts with [OH
-
] alkalinity, if present, and [CO
3
2-
] alkalinity as follows:
OH-
+ H+
→ H2O
CO3
2-
+ H+
→ HCO3
-
In the second phase, the titration is continued until the methyl orange indicator end-point is reached
with a colour change from yellow to red (i.e. pH of about 4.5). During this phase, the addition of
acid changes the HCO
3
-
ions, initially present as well as those produced by Reaction (1-4), into
carbonic acid;
HCO3
-
+ H+
→ H2CO3

44 | P a g e
By: Abdu H. (MSc)
In the above titration, the result of first step is known as phenolphthalein alkalinity and the
overall titration is known as total alkalinity (the amount of acid required to react with all the
hydroxide, carbonates and bicarbonates in the sample) respectively. When the pH of a sample is
less than 8.3, a single titration is made to the methyl orange end point. Also for routine work, it is
common that only the total alkalinity is determined. It is possible to determine the various
components of alkalinity (i.e. hydroxide, carbonate and bicarbonate fractions) from a combination
of titration, pH measurements and chemical equilibrium equations. An understanding of the
buffering capacity of alkalinity can be derived from an evaluation of figure above. At the inflection
points of pH 8.3 (phenolphthalein alkalinity) and pH 4.5 (total alkalinity), the carbonate system
will react with a considerable pH change when only a small fraction of titrant is added. However,
at the points where only half of the initial carbonate has been converted to bicarbonate and only
half of the resultant bicarbonate has been converted to carbonic acid, considerably more titrant is
required to effect a pH change. It is during these conditions that the buffering capacity is exhibited.
Quantifying the alkalinity to the inflection points is a measure of this buffering capacity.
¾ Turbidimeter
¾ UV visible spectroscopy
¾ Beaker
¾ pH meter
¾ Magnetic stirrer
¾ Erlenmeyer flask
¾ Methyl orange indicator
¾ Sulphuric acid
Sampling
Surface water and Tap water collect from around your campus.

45 | P a g e
By: Abdu H. (MSc)
Determination of Turbidity
1. Select the operating range at ”AUTO” mode of the Turbidimeter.
2. Fill a clean sample cell to the mark with the test sample and place it in the cell holder. The
sample cell must be clean, dry and free of fingerprints. Wipe the outside of the cell with a lens
tissue and align the dot on the sample cell with the raised mark on the spill ring around the cell
holder opening. Be sure the cell is kept down completely and held in place by the spring clip.
Cover the sample with the light shield.
3. The digital readout is in Nephelometric Turbidity Units (NTU).
Determination of Colour
1. Place the cell containing the blank (distilled water) in the sample compartment with the transparent
sides facing the light source of UV visible spectroscopy. Close the sample compartment lid.
(Note: Do not touch the transparent sides of the sample cell and keep it clean).
2. Press “AUTOZERO” key to set the zero absorbance.
3. Discard the distilled water and place the cell containing the sample in the measuring position. Close
the sample compartment lid.
4. Press “START” key to measure colour.
5. Record the “Conc” as colour units for the sample
Determination of pH
1. Calibrate the pH meter according to instructions supplied by the Lab instructor.
2. Pour sample into a clean beaker.
3. Rinse the probe thoroughly with distilled water to prevent any carry-over. Switch to pH mode.
4. Immerse the probe in the sample.
5. Establish equilibrium between probe and sample by stirring to insure homogeneity. Gently drop
a stirring bar into the sample and place the beaker on a magnetic stirrer. Start the magnetic stirrer
and adjust the speed to give thorough but gentle mixing.
6. Read and record the pH.
7. Rinse the electrode thoroughly with distilled water.
8. When not in use, the electrode should be replaced in the beaker containing water.

46 | P a g e
By: Abdu H. (MSc)
Determination of Alkalinity (Total Alkalinity)
1. For each sample, place 100 mL of sample in an Erlenmeyer flask.
2. Add 3 drops of methyl orange indicator solution to the flask.
3. Titrate sample with 0.02 N H2
SO4
(sulphuric acid), constantly swirling the flask content above
a white surface until just after the colour of the flask content change from yellow to red.
4. Record the volume of titrant used.
5. Calculate Total Alkalinity as follows:
Total alkalinity as mg/L CaCO3 =
‫ۯ‬ൈ‫ۼ‬ൈ૚૙૙૙ൈ૞૙
‫܍ܔܘܕ܉ܛ܎ܗۺܕ‬
where: A = volume of 0.02 N H
2
SO
4
used for methyl-orange end point.
N = Normality of H
2
SO
4
, 0.02 N.
Sample Turbidity
(NTU)
Apparent
colour
True colour pH Alkalinity
(mg/L CaCO3)
Tap water
Surface
water
Calculate the total alkalinity of the tap and surface water sample?
Questions
1. Why you rinse the electrode in the glass beaker?
2. From your results, are there any observable relationships between turbidity and apparent
colour and between turbidity and true colour?
3. What form of alkalinity would you expect to predominate in Tap and Surface waters?

47 | P a g e
By: Abdu H. (MSc)
EXPERIMENT-7
Determination of Permanent Hardness due to Ca2+
and Mg2+
in Tap Water
by EDTA Method
Objective: To determine the level of permanent hardness of tap water by EDTA Method
Introduction
Hard waters are generally considered to be those waters that require considerable amounts of soap
to produce foam and that also produce scale in water pipes, heaters, boilers and other units in which
the temperature of water is increased. Hard water is appropriate for human consumption similar to
that as soft waters, however it produces adverse actions with soap and thus their use for cleaning
purposes is unsatisfactory and thus their removal from water is required. Hardness of waters varies
from place to place. In general, surface waters are softer than ground waters. Waters are commonly
classified based on degree of hardness:
Classification of hardness types
Hardness (mg/L) Degree of hardness
0-75 Soft
75-100 Moderately hard
150-300 Hard
300 Very hard
Hardness:
Hardness is caused by polyvalent metallic cations, though the divalent cations, such as calcium
and magnesium cations are usually the predominant cause of hardness. In addition, hardness is
also caused by Ca2+
and Mg2+
ions. For example, when hard water is heated, Ca2+
ions react with
bicarbonate (HCO3
-
) ions to form insoluble calcium carbonate (CaCO3) (Eq. 1). This precipitate,
known as scale, coats the vessels in which the water is heated, producing the mineral deposits on
your cooking dishes. Equation 2 presents magnesium hardness.
Ca2+
(aq) + 2HCO3
-
(aq) → CaCO3(s) +H2O +CO2 1a
Mg2+
(aq) + 2OH-
(aq) → Mg(OH)2 (s) 1b

48 | P a g e
By: Abdu H. (MSc)
Total hardness is defined as the sum of the calcium and magnesium concentrations, both expressed
as calcium carbonate in mg/L. When hardness (numerically) is greater than the sum of carbonate
and bicarbonate alkalinity, amount of hardness equivalent to the total alkalinity is called
“Carbonate hardness”.
Carbonate hardness (mg/L) = Alkalinity (2a)
When alkalinity Total hardness:
Carbonate hardness (mg/L) = Total hardness (2b)
The amount of hardness in excess of this is called “Non-carbonate hardness (NCH)”. These
are associated with sulfate chloride, and nitrate ions.
Temporary hardness is due to the presence of bicarbonates of calcium and magnesium ions. It can
be easily removed by boiling. When water is boiled, temporary hardness producing substances
(bicarbonates) are precipitated as insoluble carbonates or hydroxides. This precipitate can be
removed by filtration. However, Permanent hardness is due to the presence of chlorides and
sulphates of calcium and magnesium ions. This type of hardness cannot be removed by boiling.
The filtrate obtained contains permanent hardness producing substances and is estimated against
EDTA using EBT indicator.
The estimation of hardness is based on complexometric titration, is used to find the total calcium
and magnesium content of milk, sea water and various solid materials. It can also be used to
determine the total hardness of fresh water provided the solutions used are diluted. The combined
concentration of calcium and magnesium ions is considered to be the measure of water hardness.
The method uses a very large molecule called EDTA which forms a complex with calcium and
magnesium ions. EDTA is short for ethylenediaminetetraacetic acid. A blue dye called Eriochrome
Black T (ErioT) is used as the indicator. This blue dye also forms a complex with the calcium and
magnesium ions, changing colour from blue to pink in the process. The dye–metal ion complex is
less stable than the EDTA–metal ion complex. For the titration, the sample solution containing the
calcium and magnesium ions is reacted with an excess of EDTA. The indicator is added and
remains blue as all the Ca2+
and Mg2+
ions present are complexed with the EDTA.

49 | P a g e
By: Abdu H. (MSc)
A back titration is carried out using a solution of magnesium chloride. This forms a complex with
the excess EDTA molecules until the end-point, when all the excess EDTA has been complexed.
The remaining magnesium ions of the magnesium chloride solution then start to complex with
ErioT indicator, immediately changing its colour from blue to pink. Estimation of hardness by
EDTA method is based on the principle that EDTA forms complexes with hardness causing metal
ions in water. The complexes are stable within pH range of 8 to 10. Thus, to maintain the pH range
buffer solution (NH4Cl and NH4OH mixture) is used. Eriochrome Black-T (EBT) indicator is used
to indicate the completion of complexation reaction.
ቈ
‫ܽܥ‬ାଶ
‫݃ܯ‬ାଶ቉ ൅ ‫ܶܤܧ‬ ՜ ൤
‫ܶܤܧܽܥ‬
‫݃ܯ‬
൨ ܿ‫ݔ݈݁݌݉݋‬
unstable complex
When this solution is titrated against EDTA, it replaces the indicator from the indicator complex.
When all the hardness causing ions are complexed by EDTA, the indicator is set free and end point
is marked by color change from purple red to blue. The total hardness is thus determined.
Ca EBT Ca EDTA
complex EDTA EBT
Mg Mg
ª º ª º
o
« » « »
¬ ¼ ¬ ¼
The formed complex is blue in color. The temporary hardness is removed by boiling and then
precipitate formed is removed by filtration and the permanent hardness in filtrate is determined
by titration with EDTA.
Temporary hardness = total hardness - permanent hardness
The most common multivalent metal ions in natural waters are Ca2+
and Mg2+.
In this experiment,
we will find the total concentration of metal ions that can react with EDTA and we will assume
that this equals the concentration of Ca2+
and Mg2+.
In a second experiment, Ca2+
is analyzed
separately after precipitating Mg (OH)2 with strong base NaOH.
¾ Desiccator
¾ Volumetric flask
¾ Burette
¾ Pipette

50 | P a g e
By: Abdu H. (MSc)
Reagents Required
¾ EDTA: Na2H2EDTA. 2H2O
¾ Buffer (pH 10): Add 142 mL of 28 % aqueous NH3 to 17.5 g of NH4Cl and dilute to 250
mL with distilled water.
¾ Eriochrome black T indicator: Dissolve 0.2 g of the solid indicator in 15 mL of
triethanolamine plus 5 mL of absolute ethanol. (Alternatively, Calmagite could be used by
dissolving 0.05 g in 100 mL of water. The color changes are the same for both indicators)
¾ Hydroxynaphthol blue indicator
¾ 50 % (w/w) NaOH: Dissolve 100 g of NaOH in 100 g of H2O in a 250-mL plastic bottle.
Store tightly capped. When you remove solution with a pipette, try not to disturb the solid
Na2CO3 precipitate.
¾ Unknowns: Collect water from streams or lakes. To minimize bacterial growth, plastic jugs
should be filled to the top and tightly sealed. Refrigeration is recommended
Procedure
1. Dry Na2H2EDTA .2H2O at 80 Ԩ for 1 hour and cool in the desiccator. Accurately weigh
out ~0.6 g and dissolve it with heating in 400 mL of water in a 500-mL volumetric flask.
Cool to room temperature, dilute to the mark, and mix well.
2. Pipet a 50.00-mL sample of tap water into a 250-mL flask.
3. To each sample, add 3 mL of pH 10 buffer and 6 drops of Eriochrome black T indicator
4. Titrate with EDTA from a 50-mL burette and note when the color changes from wine red
to blue
5. Repeat the titration with three samples to find an accurate value of the total Ca2+
+ Mg2+
concentration.
6. Perform a blank titration with 50 mL of distilled water and subtract the value of the
blank from each result.
¾ Let V1 mL volume of EDTA consumed during titration and let V be the volume of tap
water taken. Thus, total hardness of the sample = 1 1000
V x
V
ppm of CaCO3 equivalent
¾ For the determination of permanent hardness due to Ca2+
, pipette out the same volume of
unknown sample as previous into clean flasks

51 | P a g e
By: Abdu H. (MSc)
7. Add 30 drops of 50% (w/v) NaOH to each solution and swirl for 2 minutes to precipitate
Mg(OH)2 (which may not be visible). Add~0.1 g of solid hydroxynaphthol blue to each flask.
This indicator is used because it remains blue at higher pH than does Eriochrome black T.
8. Collect the filtrate into volumetric flask and titrate in the same way as above. After reaching
the blue end point, allow the sample to stand for 5 min with occasional swirling so that any
Ca(OH)2 precipitate may redissolve. Then titrate back to the blue end point if the blue color
turns to red upon standing.
9. Perform a blank titration with 50 mL of distilled water. Calculate the permanent hardness of
water due to calcium as =
௏ൈேൈହ଴ൈଵ଴଴଴଴
௩௢௟௨௠௘௢௙௦௔௠௣௟௘௧௔௞௘௡
CaCO3 equivalent.
Where V = volume of EDTA consumed during the titration
N= normality of EDTA
Permanent hardness due to Mg+2
= Total hardness - permanent hardness due to Ca+2
N.B. In this experiment temporary hardness is assumed to be negligent
Calculation
Calculate the total and permanent hardness of the water sample in ppm of CaCO3?
Questions
1. Why is hardness of water expressed in terms of calcium carbonate equivalent?
2. Mention the disadvantages of hard water for industrial purpose.
3. Why is the colour of solution wine red before titration and blue colour at the end of
titration?
4. State the salts responsible for temporary and permanent hardness of water?
5. Why is ammonium hydroxide-ammonium chloride buffer added during the determination
of hardness of water?

52 | P a g e
By: Abdu H. (MSc)
EXPERIMENT-8
Determination of Chemical Oxygen Demand (COD) of Wastewater Using
Open Reflux Method
Objective: To determine the COD of Wastewater
Introduction
Water pollution and its impacts on the environment are serious issues for present world. To limit
the water pollution and improve the water quality, advanced wastewater treatment technologies
are invented. These technologies are implemented by removing physical, chemical and biological
contaminants from wastewater and producing an environmentally safe fluid waste stream (treated
effluent) and a solid waste (treated sludge). It may then even be possible to reuse sewage effluent
for drinking water with the help of more advanced technologies. If untreated wastewater
containing contamination enters into the surface and ground water resources, it leads to a serious
environmental and human health risk. To minimize the potential risks from untreated wastewater
entering freshwater resources, industrial wastewater plants go through a water quality assessment
by monitoring some parameters. Water quality professionals assess water quality by measuring the
concentrations of these parameters and comparing with their standards. Some of the unique
analytical parameters of the water pollution control industry are biochemical oxygen demand,
chemical oxygen demand, taste, odor, color, chlorine demand, hardness, alkalinity and
biodegradability tests. Finding excessive levels of one or more of these parameters can serve as an
early warning of potential pollution problems. One of these parameters are COD and BOD that
indicate the amount of organic pollution and water degradation.
COD is defined as the amount of oxygen equivalents consumed in oxidizing the organic
compounds of samples by strong oxidizing agents such as dichromate or permanganate. It is
expressed in milligrams per liter (mg/L) that indicates the mass of oxygen consumed per liter of
solution. The higher the chemical oxygen demand, the higher the amount of pollution in the water
sample. COD is considered one of the most important quality control parameters of an effluent in
wastewater treatment facility. COD values are used to monitor wastewaters before (influent) and
after (effluent) treatment, and, therefore, their reliability is important to protect the environment

53 | P a g e
By: Abdu H. (MSc)
and to guarantee the economical sustainability of the treatment facility. COD measurements are
commonly made on samples of wastewater treatment facility or of natural waters contaminated by
domestic and industrial wastes. COD is measured as a standardized laboratory assay in which a
closed water sample is incubated with a strong chemical oxidant under specific conditions of
temperature and for a particular time. A commonly used oxidant in COD assays is potassium
dichromate (K2Cr2O7) which is used in combination with boiling sulfuric acid (H2SO4).
Chemical Oxygen Demand (COD) is rapidly measured parameters as a means of measuring
organic strength for streams and polluted water bodies. The test can be related empirically to BOD,
organic carbon or organic matter in samples from a specific source taking into account its
limitations. The test is useful in studying performance evaluation of wastewater treatment plants
and monitoring relatively polluted water bodies. COD determination has advantage over BOD
determination. COD results can be obtained in 3-4 hrs as compared to 3-5 days required for BOD
test. Further, the test is relatively easy, precise, and is unaffected by interferences as in the BOD
test. The intrinsic limitation of the test lies in its inability to differentiate between the biologically
oxidizable and biologically inert material and to find out the system rate constant of aerobic
biological stabilization.
The open reflux method is suitable for a wide range of wastes where a large sample size is
preferred. The closed reflux methods are more economical in the use of metallic salt reagents and
generate smaller quantities of hazardous waste, but require homogenization of samples containing
suspended solids to obtain reproducible results. The dichromate reflux method is preferred over
procedures using other oxidants (e.g. potassium permanganate) because of its superior oxidizing
ability, applicability to a wide variety of samples and ease of manipulation. Oxidation of most
organic compounds is up to 95-100% of the theoretical value. The organic matter gets oxidized
completely by potassium dichromate (K2
Cr2
O7
) with silver sulphate as catalyst in the presence of
concentrated H
2
SO
4
to produce CO
2
and H
2
O. The excess K
2
Cr
2
O
7
remaining after the reaction is
titrated with ferrous ammonium sulphate [Fe (NH
4
)
2
(SO
4
)
2
]. The dichromate consumed gives the
oxygen (O
2
) required for oxidation of the organic matter. The chemical reactions involved in the
method are as under:

54 | P a g e
By: Abdu H. (MSc)
2K2
Cr
2
O
7
+ 8 H
2
SO
4
→2 K
2
SO
4
+ 2Cr
2
(SO
4
)3 + 8 H
2
O + 3O2
C
6
H
12
O
6
+ 6O
2
→ 6CO
2
+ 6H
2
O
Cr
2
O7
- 2
+ 6Fe
+2
+ 14H
+
→ 6Fe
+3
+ 2Cr
3+
+ 7H
2
O
Interferences
Oxidation of most organic compounds is 95 to 100% of the theoretical value. Pyridine and related
compounds resist oxidation and volatile organic compounds will react in proportion to their contact
with the oxidant. Straight-chain aliphatic compounds are oxidized more effectively in the presence
of a silver sulfate catalyst. The most common interferent is the chloride ion. Chloride reacts with
silver ion to precipitate silver chloride, and thus inhibits the catalytic activity of silver. Bromide,
iodide, and any other reagent that inactivates the silver ion can interfere similarly. Such
interferences are negative in that they tend to restrict the oxidizing action of the dichromate ion
itself. However, under the rigorous digestion procedures for COD analyses, chloride, bromide, or
iodide can react with dichromate to produce the elemental form of the halogen and the chromic
ion. Results then are in error on the high side. The difficulties caused by the presence of the
chloride can be overcome largely, though not completely, by complexing with mercuric sulfate
(HgSO4) before the refluxing procedure. Although 1 g HgSO4 is specified for 50 mL sample, a
lesser amount may be used where sample chloride concentration is known to be less than 2000
mg/L, as long as a 10:1 weight ratio of HgSO4: Cl-
is maintained. Do not use the test for samples
containing more than 2000 mg Cl-
/L.
Halide interferences may be removed by precipitation with silver ion and filtration before
digestion. This approach may introduce substantial errors due to the occlusion and carry down of
COD matter from heterogenous samples. Ammonia and its derivatives, in the waste or generated
from nitrogen-containing organic matter, are not oxidized. However, elemental chlorine reacts
with these compounds. Hence, corrections for chloride interferences are difficult. Nitrite (NO2
-
)
exerts a COD of 1.1 mg O2/mg NO2
-
-N. Because concentrations of NO2
-
in waters rarely exceed 1
or 2 mg NO2
-
-N/L, the interference is considered insignificant and usually is ignored. To eliminate
a significant interference due to NO2
-
, add 10 mg sulfamic acid for each mg NO2
-
-N present in the

55 | P a g e
By: Abdu H. (MSc)
sample volume used; add the same amount of sulfamic acid to the reflux vessel containing the
distilled water blank.
Reduced inorganic species such as ferrous iron, sulfide, manganous manganese, etc., are oxidized
quantitatively under the test conditions. For samples containing significant levels of these species,
stoichiometric oxidation can be assumed from known initial concentration of the interfering
species and corrections can be made to the COD value obtained.
Apparatus and equipment
¾ 250 or 500 mL Erlenmeyer flask with standard (24/40) tapered glass joints
¾ Friedrich’s reflux condenser (12 inch) with standard (24/40) tapered glass joints
¾ Electric hot plate
¾ Volumetric pipettes (10, 25, and 50 mL capacity)
¾ Burette, 50 mL with 0.1 mL accuracy
¾ Analytical balance, accuracy 0.001g
¾ Volumetric flasks (1000 mL capacity)
¾ Boiling beads
¾ Magnetic stirrer and stirring bars.
Reagents and standards
a. Standard potassium dichromate solution, 0.25 N (0.04167 M): Dissolve 12.259g K
2
Cr
2
O
7
dried at
103 °C for 24 h in distilled water and dilute to 1000 mL. Add about 120 mg sulphamic acids to take
care of 6 mg/L NO2-N.
b. Sulphuric acid reagent: Add 10 g of Ag
2
SO
4
to 1000 mL concentrated H
2
SO
4
and let stand for one
to two days for complete dissolution.
c. Standard ferrous ammonium sulphate approx. 0.25 N (0.25 M): Dissolve 98 g Fe (NH4
)2
(SO4
)2
.6H2
O
in about 400 mL distilled water. Add 20 mL concentrated H
2
SO
4
and dilute to 1000 mL.
d. Ferroin indicator: Dissolve 1.485 g 1, 10-phenanthroline monohydrate and 695 mg FeSO
4
.7H
2
O in
distilled water and dilute to 100 mL.

56 | P a g e
By: Abdu H. (MSc)
e. Mercuric sulphates: HgSO4
crystals analytical grade
f. Potassium hydrogen phthalate (KHP) Standard: Dissolve 425 mg lightly crushed dried potassium
hydrogen phthalate (HOOC. C6
H4
.COOK) in distilled water and dilute to 1000 mL. This solution
has a theoretical COD of 500 μg O2
/mL. This solution is stable when refrigerated, up to 3 months
in the absence of visible biological growth.
Sample collection, preservation and Sample preparation
Preferably collect wastewater/tap water in glass bottles. Remove settleable solids by sedimentation
or decantation. If there is delay between collection and analysis, preserve sample by acidification
to pH≤2 using concentrated H2
SO4
. Samples can be preserved for maximum 7 days. All samples
high in solids should be blended for 2 minutes at high speed and stirred when an aliquot is taken
for analysis. Select the appropriate volume of sample based on expected COD range, e.g. for COD
range of 50-500 mg/L take 25-50 mL of sample. Sample volume less than 25 mL should not be
pipetted directly, but serially diluted and then a portion of the diluted sample taken. Dilution factor
should be incorporated in calculations.
a) 500 mL of sample diluted to 1000 mL = 0.5 mL sample/mL of diluent, 50 mL = 25 mL of
sample.
b) 100 mL of sample diluted to 1,000 mL = 0.1 mL sample/mL diluent, 50 mL of diluent =
5 mL of sample
Calibration
Since the procedure involves chemical of organic matter by potassium dichromate as oxidizing
agent which is a primary standard, calibration is not applicable. For standardization of ferrous
ammonium sulphate, dilute 10 mL standard K
2
Cr
2
O
7
to about 100 mL. Add 10 mL concentration
of H
2
SO
4
and allow it to cool. Titrate with ferrous ammonium sulphate (FAS) to be standardized
using 2-3 drops of ferroin indicator. Calculate normally
Normality of FAS =
required
FAS
of
mL
CrO
K
of
mL 7
2
The deterioration of FAS can be decreased if it is stored in a dark bottle.

57 | P a g e
By: Abdu H. (MSc)
Procedure
1. Place 0.4 g HgSO4
in a 250 mL reflux sample
2. Add 20 mL sample or an aliquot of sample diluted to 20 mL with distilled water. Mix
well
3. Add clean pumic stones or glass beads.
4. Add 10 mL 0.25 N (0.04167M) K
2
Cr2
O7
solution and mix.
5. Add slowly 30 mL concentrated H
2
SO
4
containing Ag
2
SO
4
mixing thoroughly. This slow
addition along with swirling prevents fatty acids to escape due to generation of high
temperature. Alternatively attach flask to condenser with water flowing and then add
H
2
SO
4
slowly through condenser to avoid escape of volatile organic substance due to
generation of heat
6. Mix well. If the color turns green, either take fresh sample with lesser aliquot or add
more potassium dichromate and acid.
7. Connect the flask to condenser. Mix the contents before heating. Improper mixing will
result in bumping and blow out of flask content.
8. Reflux for a minimum of 2 hours. Cool and then wash down condenser with distilled
water.
9. Disconnect reflux condenser and dilute the mixture to about twice its volume with distilled
water. Cool to room temperature and titrate excess K2
Cr2
O7
with0.1M FAS using 2-3 drops
of ferroin indicator. The sharp color change from blue green to reddish brown indicates
end-point or completion of the titration. After a small time, gap, the blue-green color may
reappear. Use the same quantity of ferroin indicator for all titrations.
10. Reflux blank in the same manner using distilled water instead of sample. Alternate
procedure for low COD samples less than 50 mg/L: Follow similar procedure with two
exceptions (use standard 0.025 N (0.004167 M) K2
Cr2
O7
and titrate with standardize 0.025
M FAS. The sample volume should be 5 mL. Exercise extreme care with this procedure
because even a trace of organic matter on the glassware or from the atmosphere may cause
gross errors. Compute amount of HgSO4
to be added based on chloride concentrations.
Carry blank reagent through the same procedure.

58 | P a g e
By: Abdu H. (MSc)
Calculations
COD as mg/L =
ሺࢇെ࢈ሻൈࡺൈૡ૙૙૙
ࢂ࢕࢒࢛࢓ࢋ࢕ࢌ࢙ࢇ࢓࢖࢒ࢋሺ࢓ࡸሻ
Where a = Volume of FAS used for blank
b = Volume FAS used for sample
N = normality of FAS
8000 = Milieq. Wt. of O
2
x 1000
Questions
1. What is the significant of COD?
2. What is the colour change at the end point in the determination of COD?
3. Why is the blank titre value higher than sample titre value?
4. Why dil. H2SO4 is used to dissolve FAS crystals while preparing standard solution?

59 | P a g e
By: Abdu H. (MSc)
EXPERIMENT-9
Soil Sample Collection and Preparation for Heavy Metal Analysis
Objectives: To Familiarize Students with Soil Sample Collection and Preparation for Further
Analysis
Introduction
THE SOIL SYSTEM
Soil is defined as “the unconsolidated mineral material on the immediate surface of the earth that
has been subjected to and influenced by genetic and environmental factors. The true soil
component can also be defined as all mineral and naturally occurring organic materials with a
particle size less than 2 mm. The physical and chemical characteristics of the soil system influence
the transformation, retention, and movement of pollutants through the soil. Clay content, organic
matter content, texture, permeability, pH and cation exchange capacity will influence the rate of
migration and form of the chemical found in leachate migrating from the waste. Elevated
concentrations of heavy metals in soils are of potential long term environmental and health
concerns because of their persistence and cumulative tendency in the environment, and their
associated toxicity to biological organisms. These factors must be considered by the investigator
when designing a soil sampling plan. Furthermore, restricted use of contaminated lands and the
costs of soil remediation also pose liabilities and financial burdens on landowners and other
stakeholders. As a consequence, environmental assessment of lands with respect to heavy metal
contamination, and identification its environmental and health implications have become
increasingly important in environmental research. For a reliable and cost-effective investigation of
heavy metal contamination of soils, a well-planned sampling strategy, appropriate selection of
analytical methods, and careful interpretation of results are of vital importance. The soils are
contaminated with heavy metals, analysis of the heavy metal concentrations of the soils will be
adequate, and sampling of the soils will be relatively simple. However, if knowledge of the spatial
distribution of heavy metals in soils is also sought, a systematic sampling approach will be
required.

60 | P a g e
By: Abdu H. (MSc)
Sampling Strategy
Sampling is the process of obtaining representative sample which reliably represents the
population under question both in composition and size. The sampling process must ensure that
the items chosen are representative of the bulk of material or population. Sampling is inherent to
any research program in science because the measurement of properties of the total population is
impossible or difficult for any realistic study. However, it is clear that the larger the sample size
the more closely your sample data will match the entire population. The goal of sampling is thus
to produce a sample that is representative of the target population. The following sampling
techniques are commonly in soil sampling.
RANDOM SAMPLING
The basis of most sampling plans in environmental sampling is the concept of random or
probabilistic selection of the sample to be collected and the subsample that is to be analyzed. The
random sampling strategy is the simplest methods, where soil samples are collected randomly and
stochastically independently across the site of interest. It can be used as a quick sampling program
of a pilot study. In random sampling of a site, each sample point within the site must have an equal
probability of being selected. The same can be said for the selection of particles within a sample.
Each and every particle within the sample must have an equal chance of being selected. Each
particle that is not in the sample should have a zero probability of being selected. Properly
designed sampling plans based upon the laws of probability provide a means of making decisions
that have a sound basis and are not likely to be biased. On contrary, there are nonrandom samples
collected for a particular reason. They are based solely on the choice of collector. Such samples
are called purposive samples.
A major disadvantage of this sampling strategy is that soil samples may not represent the whole
study site. Therefore, this sampling strategy is usually employed in relatively homogenous sites
and applicable to investigations where the major objective is to determine whether heavy metal
concentrations of the soils are elevated above background and/or legislative standards.

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