General Chemistry Lab Manual

WOLLEGA UNIVERSITY
COLLEGE OF NATURAL AND COMPUTATIONAL SCIENCE
DEPARTMENT OF CHEMISTRY
GENERAL CHEMISTRY MANUAL (CHEM 1012)
Prepared by:
1. Kemal Mohammed (BSc)
2. Mosisa Dugasa (BSc)
Edited by: Chala Boru (MSc)
June, 2021
Nekemte, Ethiopia

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Contents
To The Student ............................................................................................................................... iii
Laboratory Report Format.............................................................................................................. iii
Introduction to Laboratory Safety Rules ..........................................................................................v
Table 1.1 .the following cases must be notified firstly to the laboratory instructor........................ix
Experiment 1: Preparation of solutions and concentration calculation ............................................1
Experiment 2: Mass and Volume Measurements to Define Density................................................5
Experiment 3: Bunsen burner.........................................................................................................10
Experiment 4: Physical and Chemical Changes.............................................................................13
Experiment 5: Diffusion of gases: Determination of Graham’s rate laws .......................................2
Experiment 6: Acid-base reactions: use of acid-base indicators ......................................................6
Experiment 7: Determination of solubility of salts: Investigating the solubility of ionic and
covalent compounds .......................................................................................................................11
Experiment 8: Simple and Fractional Distillation ..........................................................................16
Experiment 9: Separation of mixtures: Extraction; Distinguishing compounds and mixtures;
Separation of a mixture using a magnet; Recrystallization; and Filtration ....................................22
Experiment 10: Instrumental analysis; Colorimetric Determination of Acetaminophen...............28
Experiment 11: Investigating the heat involved in a chemical reaction (Calorimetry):
Investigating endothermic reaction; Investigating exothermic reaction; and Effect of temperature
on reaction rate ...............................................................................................................................32
REFERENCES...............................................................................................................................35

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To The Student
To work in the laboratory most efficiently, you should read the experimental procedures in
advance and understand in detail before you started the laboratory works. After you did the
experiment, try to answer the review questions at the end of each experiment. These questions
will help you to understand the experiment in advance.
Chemistry is an experimental science. Therefore, chemists perform basic experimental analysis.
Qualitative analysis is to determine the nature of processes, which are often unanticipated and
sometimes unpredictable. Quantitative analysis is to determine the amount of a measurable
change in mass, volume, or temperature etc.
The objectives to perform laboratory works:
 To develop the skills necessary to obtain and evaluate a reliable original result.
 To record your results for future use.
 To be able to draw conclusions regarding your results (with the aid of some coaching and
reading in the beginning).
 To learn to communicate your results critically and knowledgeably.
By attentively reading over the experimental procedures in detail, and carefully following
directions you can safely able to accomplish the objectives in the laboratory.
Laboratory Report Format
A laboratory report is a written composition of the results of the experiment. A laboratory reports
explain what you did in experiment, what you learned and what the result meant. It should be
written precisely and clearly, using good grammar and punctuation. The laboratory report formats
differ depending up on the type of experiment that you do but the general format includes the
following main points.
Cover page
The cover page should contain the name of the University, the name of the college and
department, Experiment number, the students’ name and name of any laboratory partners, the
instructor’s name and the date.
The title of the experiment
The title of the experiment reflects the nature (factual content) of the work. The title says what
you did. It should be brief and describe the main point of the experiment or investigation.
Purpose or objective of experiments
Objective state the purpose of the experiment and outlines the significant goals intended to be
learned from each experiments.

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Objective should be specific and has to be written using smart words. The objective should
answer the questions, what will the student know after completing the experiment and what will
student be able to do after completing the experiment?
Theory/introduction
Theory explains the background of the experiment briefly. It is usually written in few lines. The
object of theory is to give the reader a sense what you are doing, why you are doing it in the way
you are doing the experiment, and what the results you have determined. It should not directly
copy from the laboratory manual.
Apparatus and chemicals
List of all chemicals and apparatus you used for doing the experiment. Sometimes it is necessary
to provide the specification of the chemicals and apparatus used.
Procedure
A procedure refers to details of steps followed, in chronological order, while doing the
experiment. A procedure has to be written using passive voice. It is very important that you read
and understand the procedure before you start working.
Data and observation
Enter data into the notebook as the work is being performed. This means that loose pieces of
paper used for the intermediate recordings are prohibited.
Write the chemical reaction that takes place, if available.
Show the mathematical formulas utilized for all calculations and also a sample calculation.
Construct data tables whenever useful and appropriate. Numerical data obtained from your
procedure usually is presented as a table. Data encompasses what you recorded when you
conducted the experiment. It’s just the facts, not any interpretation of what they mean.
Both numerical data and important observation should be recorded.
Experimental data should be presented, with the correct units, neatly and succinctly in tabular
and/or graphical form rather than verbally.
Data should be presented in chart or graph form. All charts and graphs must have a title. Columns
and rows of charts should be labeled. Axes of graphs should be labeled and units given in the
label. Any data manipulation, such as equations or calculations, should be included and explained
briefly.

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Results and discussion
Interpretation of data and comparisons with the literature are presented in this section. Describe in
words what the data means. This is where you interpret the data and determine whether or not a
hypothesis was accepted. This is also where you would discuss any mistakes you might have
made while conducting the investigation. You may also wish to describe ways the study might
have been improved. This is the section where the results are explained, and a student can show
the instructor that he or she has a thorough understanding of the concept of the experiment and
the results obtained.
This is probable the most important section of the lab report. It is here that you report your
experimental data, analyze and interpret your results, and draw specific conclusions.
Discussion of results is a concise interpretation of the results.
Conclusion
This section summarizes the pertinent concepts discussed in the results and discussion section.
The conclusion should restate the result of the experiment. Conclusion should be brief, as it refers
back to the objectives and considers how and to what degree they have been met. Review the
purpose of the experiment and summarize the implication of the experiments.
Introduction to Laboratory Safety Rules
The chemistry laboratory may be considered as a place of discovery and learning. However, by
very nature of laboratory work, it can be a place of danger if precautions are not taken. Therefore
it is necessary to take care of your own health and safety and that of others working in the
laboratory. The responsibility for laboratory safety rules with each and every student in the
laboratory. You must use common sense and work carefully to avoid chemical spills, broken
glassware, and fires. This ensures not only your own safety, but that of your laboratory mates.
Knowing the level of hazards of each chemical you are using, make you ready the care you can
do during laboratory session. Doing that, you will not expose to any harmful chemicals during
your laboratory work.
The compounds you are doing with may have hazardous properties associated with them. So it is
important to follow the safety rules outlined in the manual such as: safety goggles for eye
protection are recommended and laboratory coats are to be worn by the students always when
they are working in the chemistry laboratory. Always don’t forget to wash your hands thoroughly
when you are leaving the laboratory.

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The location and use of the safety equipment in laboratory will be reminded by your instructor
the first day of your laboratory class and you should become familiar with the proper use of the
safety equipment location and their use. Examples shower, eye-wash fountain, fire blanket and
fire extinguisher.
1. Report any accidents which occur immediately to the laboratory supervisor.
Safety rules
The laboratory can be but is not necessarily a dangerous place. When intelligent precautions and
a proper understanding of techniques are employed, the laboratory is no more dangerous than any
other classroom. Most of the precautions are just common-sense practices. These include the
following:
1. Know what you have to do before entering the laboratory. Read the experiment carefully
before starting the laboratory works.
2. Do not engage in games in the laboratory. Failure to follow this rule will result in immediate
dismissal from the laboratory and subsequent conduct action.
3. Eating, drinking, and smoking are strictly prohibited in the laboratory at all times
4. Know where to find and how to use safety and first-aid equipment.
5. Wear approved eye protection when required while in the laboratory. Your safety eye
protection may be slightly different from that shown, but it must include shatterproof lenses and
side shields to provide protection from splashes.
6Consider all chemicals to be hazardous unless you are instructed otherwise. Dispose of
chemicals as instructed by your instructor. Follow the explicit instructions given in the
experiments.
7. Do not pipet solutions by mouth. Rubber pipet bulbs are provided at each laboratory station.
8If chemicals come into contact with your skin or eyes, wash immediately with copious amounts
of water and then consult your laboratory instructor.
9. Never taste anything. Never directly smell the source of any vapor or gas; instead by means of
your cupped hand, bring a small sample to your nose (see figure below). Chemicals are not to be
used to obtain a "high" or clear your sinuses.

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Figure 1.1: Wafting vapors towards one’s nose
1. Perform in the hood any reactions involving skin-irritating or dangerous chemicals
and/or ill-smelling chemicals. A typical fume exhaust hood is shown below.
Figure 1.2: Fume hood found in the laboratory
Exhaust hoods have fans to exhaust fumes out of the hood and away from the user. The hood
should be used when noxious, hazardous, and flammable materials are being studied. It also has a
shatterproof glass window, which may be used as a shield to protect you from minor explosions.
Reagents that evolve toxic fumes are stored in the hood. Return these reagents to the hood after
their use.
2. Do not perform any unauthorized experiments.
3. Clean up all broken glassware immediately.
4. Always pour acids into water, not water into acid, because the heat of solution will cause
the water to boil and the acid to spatter.
5. Never point a test tube that you are heating at yourself or your neighbor. It may erupt
like a geyser.
Figure 1.3: Beware of spattering

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1. Avoid rubbing your eyes unless you know that your hands are clean.
2. Notify the instructor immediately in case of an accident.
3. Many common reagents, for example, alcohols, acetone, and especially ether, are highly
flammable. Do not use them anywhere near open flames.
4. Observe all special precautions mentioned in experiments.
5. When finished with your Bunsen burner for a given portion of an experiment, turn it
off.
6. Exercise good housekeeping practices in the laboratory. Be sure that the laboratory
benches remain free of disorder during the experiment. In the event of a spill, clean the
area immediately and be sure to use a wet sponge to wipe off the work station at the end
of the laboratory session.
7. Learn the location of fire protection devices. In the unlikely event that a large chemical
fire occurs, a powder extinguisher and a CO2 extinguisher are available in the laboratory.
Figure 1.4: Powder and CO2extinguisher
In order to activate the extinguisher, you must pull the metal safety ring from the
handle and then depress the handle. Direct the output of the extinguisher at the base of
the flames. The carbon dioxide smothers the flames and cools the flammable material
quickly. If you use the fire extinguisher, be sure to return the extinguisher in at the stockroom so
that it can be refilled immediately. If the carbon dioxide extinguisher does not extinguish the fire,
evacuate the laboratory immediately and call the security. One of the most frightening and
potentially most serious accidents is the ignition of one’s clothing.
Therefore, certain types of clothing are hazardous in the laboratory and must not be worn. Since
sleeves are most likely to come closest to flames, any clothing that has bulky or loose sleeves
should not worn in the laboratory. Ideally, students should wear laboratory coats with tightly
fitting sleeves. Long hair also presents a hazard and must be tied back. If a student's clothing or
hair catches fire his or her neighbors should take prompt action to prevent severe burns. Most
laboratories have a water shower for such emergencies. A typical laboratory emergency water
shower has the following appearance.

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Figure 1.5: A safety shower
In case someone's clothing or hair is on fire, immediately lead the person to the shower and pull
the metal ring. Safety showers generally dump 151 to 190 liters of water, which should extinguish
the flames. These showers cannot be shut off once the metal ring has been pulled. Therefore, the
shower cannot be demonstrated.
Table 1.1 .the following cases must be notified firstly to the laboratory instructor.
Burn: expose the burned area to
the tap water (5-10 min.), apply
first aid.
Cut / Injury: Wash with water and apply first aid.
Fainting: Provide fresh air. Lay
down and put the head lower than
the body.
Fire: (Notify the assistant immediately) Put the Bunsen
burner off. Use shower in case of hair and clothes caches
fire. Use the fire extinguisher when necessary.
Bleeding: compress on the wound,
keep the wound above the heart
level and get medical help.
Chemical Spill: clean in a manner appropriate to the
chemical. Aqueous solutions can be removed with water.
Information your assistant.
Acid Burns: Use NaHCO3
solution
Base Burns: Use Boric acid or
Acetic acid solution
Chemicals Spilled in the eye: The Eye is washed
immediately with plenty of water for at least 15
minutes (use the eye-wash shower rooms) Get medical
help.

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Experiment 1: Preparation of solutions and concentration calculation
Objective: To practice the preparation of solutions of known concentration from
a solid and by dilution from a stock solution
Theory: Preparation of a solution is an essential skill in the study of chemistry. The solutions
which are prepared are often used in determining quantitative relationships in chemical reactions.
Many of the reactions of qualitative and quantitative chemical analyses take place in solutions.
Analytical chemistry deals with solution measurements and concentrations, from which we
calculate mass and vice- versa. Thus, we prepare solutions of known concentration for calibration
of instruments response or to titrate sample solutions. Laboratory experiments and different types
of research often require preparation of chemical solutions in their procedure.
Solution is a uniform homogeneous mixture of two or more substances. The individual
substances may be present in varying amounts. The relative amount of a given solution
component is known as its concentration. Often, though not always, a solution contains one
component with a concentration that is significantly greater than that of all other components.
This component is called the solvent and may be viewed as the medium in which the other
components are dispersed, or dissolved. It is a substance which does the dissolving (typically a
liquid, such as water or alcohol). Solutions in which water is the solvent are, of course, very
common on our planet. A solution in which water is the solvent is called an aqueous solution.
A solute is other component of a solution that is typically present at a much lower concentration
than the solvent. Solute is a substance which is dissolved, or has gone into solution (typically a
solid).Its concentrations are often described with qualitative terms such as dilute (of relatively
low concentration) and concentrated (of relatively high concentration).Concentrations may be
quantitatively assessed using a wide variety of measurement units, each convenient for particular
applications.
Preparing a solution of known concentration is perhaps the most common activity in any
analytical lab. The method for measuring out the solute and solvent depend on the desired
concentration units, and how exact the solution’s concentration needs to be known. Pipets and
volumetric flasks are used when a solution’s concentration must be exact; graduated cylinders,
beakers, and reagent bottles suffice when concentrations need only be approximate.
Two methods for preparing solutions are described in this section.

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A. Preparing stock solutions
Stock solution is a solution of known concentration from which other solutions are prepared. It is
prepared by weighing out an appropriate portion of a pure solid or by measuring out an
appropriate volume of a pure liquid and diluting to a known volume. Exactly how this is done
depends on the required concentration units. For example, to prepare a solution with a desired
molarity you would weigh out an appropriate mass of the reagent, dissolve it in a portion of
solvent, and bring to the desired volume. To prepare a solution where the solute’s concentration is
given as a volume percent, you would measure out an appropriate volume of solute and add
sufficient solvent to obtain the desired total volume.
Example 1.1
Prepare 250 mL of 0.1 M sodium hydroxide.
(Molar mass of NaOH= 40 g/mol)
Mass of NaOH= 40 g/mol x 0.1 mol/L x 0.25 L= 1 g NaOH
Dissolve 1g of NaOH in about 100 mL of distilled water, and then add more water until final
volume is 250 ml.
B. Preparing solutions by dilution
Dilution is the process whereby the concentration of a solution is lessened by the addition of
solvent. A known volume of the stock solution is transferred to a new container and brought to a
new volume. Solutions with small concentrations are often prepared by diluting a more
concentrated stock solution. Since the total amount of solute is the same before and after dilution,
C1× V1 = C2 × V2
WhereC1 is the concentration of the stock solution, V1 is the volume of the stock solution being
diluted, C2 is the concentration of the dilute solution, and V2 is the volume of the dilute solution.
Again, the type of glassware used to measure V1 and V2 depends on how exact the solution’s
concentration must be known
Prepare 100 mL of 1.0 M hydrochloric acid from Concentrated (12.1 M) hydrochloric acid.
M1V1= M2V2
(12.1 M) (V1) = (1.0 M) (100 mL)
V1= 8.26 mL Conc. HCl
Add 8.26 mL of concentrated HCl to about 50 mL of distilled water, stir, and then add water up
to 100ml

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Apparatus: Volumetric flask, Stirrer, Beaker Measuring cylinder & Analytical balance
Chemicals: Sodium chloride (NaCl) and Distilled Water
Procedure
I. If starting with a solid,
Use the following procedure:
Procedure 1
Prepare 0.5 M NaCl in 250 mL volumetric flask
1. Determine the mass in grams of one mole of solute, the molar mass, MMs.
2. Decide volume of solution required, in liters, V.
3. Decide molarities of solution required, M.
4. Calculate grams of solute (gs) required using equation , mass (g) = molar mass (g/mol) x
molarity (mol/L) x volume (L)
II. If starting with a solution or liquid reagent
Use the following procedure:
Procedure 2
Prepare 0.25 M NaCl from the solution you have prepared in procedure (1) in 100 mL volumetric
flask
1. When diluting more concentrated solutions, decide what volume (V2) and molarity (M2)
the final solution should be. Volume can be expressed in liters or milliliters
2. Determine molarity (M1) of starting, more concentrated solution.
3. Calculate volume of starting solution (V1) required using equation
Note: V1 must be in the same units as V2. Using dilution law: M1V1= M2V2

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Review Questions
1) Describe how you would prepare the following three solutions:
(a) 500 mL of approximately 0.20 M NaOH using solid NaOH
(b) 1 L of 150.0 ppm Cu2+
using Cu metal
(c) 2 L of 4% v/v acetic acid using concentrated glacial acetic acid
2) A laboratory procedure calls for 250 mL of an approximately 0.10 M solution of NH3.
Describe how to prepare this solution using a stock solution of concentrated NH3 (14.8 M).
3) A sample of an ore was analyzed for Cu2+
as follows. A 1.25-g sample of the ore was dissolved
in acid and diluted to volume in a 250 mL volumetric flask. A 20 mL portion of the resulting
solution was transferred by pipet to a 50 mL volumetric flask and diluted to volume. An analysis
showed that the concentration of Cu2+
in the final solution was 4.62 ppm. What is the weight
percent of Cu in the original ore?

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Experiment 2: Mass and Volume Measurements to Define Density
Objective: To measure volume and mass, to evaluate precision of the measurements, and to use
the data to calculate the density.
Theory: Mass and Volume Measurements depend on careful observation and the use of good
laboratory techniques. In this experiment you will become familiar with some basic operations
that will help you throughout the course. Your success in future experiments will depend upon
your mastering these fundamental operations. Because every measurement made in the
laboratory is really an approximation, it is important that the numbers you record reflect the
uncertainty of the device you are using to make the measurement. The assumption is made in this
course that an uncertainty of at least one unit exists in the last digit.
10.01 mL could actually be 10.00 mL or 10.02 mL
The error of a measurement is defined as the difference between the measured and the true value.
This is often expressed as percent (%) error, which is calculated as follows:
% 𝐸𝑟𝑟𝑜𝑟 =
Measured Value − True Value
True Value
× 100
In chemical measurements we try to eliminate errors. Errors may be divided into two broad types,
systematic and random. Systematic errors occur regularly and predictably because of faulty
methods or defective instruments, or even because of incorrect assumptions (for example, a
reagent bottle that has a missing or incorrect label). Random errors are more difficult to define.
An example is a weighing error due to air currents near the balance. Line current fluctuations for
electronic instruments also lead to random errors. Random errors can make the measured quantity
either too large or too small and are governed by chance. Systematic errors always affect the
measured quantity in the same direction.
Accuracy is the closeness of agreement between a measured value and the true value. True values
can never be obtained by measurement. However, we accept values obtained by skilled workers
using the best instruments as true values for purposes of calculation or for judging our own
results. Precision describes the reproducibility of our results. A series of measurements with
values that are very close to one another is a sign of good precision. It is important to understand,
though, that good precision does not guarantee accuracy.

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Mass
The directions in this experiment are written for use with a digital pan balance. These balances
are very accurate, giving a mass measurement to the thousandth place (0.001 g). This is
considered the accuracy of the balance. Even though this is a digital instrument, the balance is
making an estimate at the thousandths place. All measurements recorded in your data sheet
should reflect this degree of precision. When using the analytical balance please be gentle with
these sensitive instruments and use the following procedures:
1. Before placing your item to be massed on the balance check the display to make sure it is
reading zero. If not press the tare bar once to zero out the balance.
2. Carefully place the item on the pan.
3. Allow sufficient time for the mass to be measured.
4. Record the mass, remove the item.
5. Never place chemicals directly on the balance pan; first place them on a weighing boat or in a
container and then place the boat or container on the balance pan.
6. Clean up any materials you spill on or around the balance.
Never make any adjustments to the balance. If it seems out of order tell to your instructor
Volume
The graduated cylinder is the device you will be using in lab. to measure liquid volumes.
Graduated cylinders are tall, cylindrical vessels with graduations scribed along the side of the
cylinder. Since volumes are measured in these cylinders by measuring the height of liquid, it is
crucial that the cylinder have a uniform diameter along its entire height, which is the case with
the glass cylinders you are using in lab.
Beakers and flasks are marked to indicate only approximate volumes and should not be used to
measure volumes during experiments.
When measuring volume in graduated cylinder, read the point on the graduated scale that
coincides with the bottom of the curved surface, called the meniscus, of the liquid. Volumes
measured in the graduated cylinder should be estimated and recorded to the correct number of
significant figures.
The volume of a solid object can also be measured using the graduated cylinder by water
displacement. Water displacement is filling a graduated cylinder with enough water to completely
cover the solid object, when it is placed into the cylinder. Be sure to place a rubber stopper in the
bottom of the graduated cylinder to prevent the solid object from cracking the bottom of the
cylinder. The difference between the measured volume of water before and after adding the solid
object is the volume of that object.

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Density
The density of material shows the denseness of that material in a specific given area. A material’s
density is defined as its mass per unit volume. Density is essentially a measurement of how
tightly matter is packed together. It is a unique physical property for a particular object. The
principle of density was discovered by the Greek scientist Archimedes. It is easy to calculate
density if you know the formula and understand the related units the symbol ρ represents density
or it can also be represented by the letter D.
Density =
Mass
Volume
𝝆 =
𝒎
𝒗
Where, m and v are density, mass, and volume respectively. For most solids, the density is given
in g/cm3
.
Apparatus: Analytical Balance, different size beakers, graduate cylinder, Erlenmeyer flask,
Rubber stopper Test tube & Metric ruler.
Chemicals: Sodium chloride and Metal sample
Procedures
A. Mass
2. Describe the precision of the analytical balance.
3. Be sure to zero out the balance and use the same balance throughout your massing. Place a
clean, dry 150 mL beaker on the pan of the balance, read and record the mass of the beaker.
4. Do the same with the 125 mL Erlenmeyer flask.
5. Do the same with the rubber stopper.
6. Go to your instructor’s lab station and get one of the metal samples and do the same.
7. Do the same with a weighing boat.
8. Carefully add approximately 2.5 grams of NaCl to the weighing boat while it is still on the
pan in the balance. Read and record the mass of the weighing boat with the NaCl. Save this
for part B.
9. Mass the second weighing boat.

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10. Leave the second weighing boat on the balance and use the tare bar to zero the balance. Now
add approximately 2.5 grams of NaCl to the boat and read and record the mass of only the
NaCl. Save this for Part B.
B. Volume
1. Fill a small test tube to the brim with water. Transfer the water to a graduated cylinder,
measure, and record the volume of water.
2. Fill a large test tube to the brim with water. Transfer the water to a graduated cylinder,
measure and record the volume of water.
3. Fill a 125 mL Erlenmeyer flask to the brim with water, measure and record the volume of
water.
4. Measure 5.0 mL of water and pour into the small test tube. Use the metric ruler to measure
and record the height in centimeters. In the future you will often find it convenient to
estimate this volume simply by observing the height of the liquid in the test tube.
5. Use the rubber stopper you massed in Part A and determine the volume of this solid object as
described in the introductory reading of this lab.
6. Place the rubber stopper in the bottom of the 50 mL graduated cylinder to prevent breakage.
Take the metal sample you massed in Part A and determine the volume of this solid object.
C. Density
The definition of density is the mass of an object divided by the volume of the object.
𝐷𝑒𝑛𝑠𝑖𝑡𝑦 =
𝑚𝑎𝑠𝑠
𝑣𝑜𝑙𝑢𝑚𝑒
1. Use your data in this experiment to determine the density of the rubber stopper. . __ g /mL
2. Use your data in this experiment to determine the density of the metal sample. _______ g / mL

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Report sheet
Read the appropriate section of the lab introduction or procedure before starting each section.
A. Mass
1. Describe the precision of the analytical balance ___________________ g
2. Mass of a clean, dry 150 mL beaker ___________________ g
3. Mass of a 125 mL Erlenmeyer flask___________________ g
4. Mass of the rubber stopper___________________ g
5. Mass of one of the metal sample___________________ g
6. Mass of a weighing boat___________________ g
7. Mass of the weighing boat with the first sample of NaCl___________________ g
8. Mass of the NaCl ___________________ g
9. Mass of only the second sample of NaCl placed in the tarred weighing boat ________g
B. Volume
1. Volume of a small test tube___________ mL
2. Volume of a large test tube___________ mL
3. Volume of a 125 mL Erlenmeyer flask___________ mL
4. Height in centimeters of 5.0 mL of water in a small test tube ___________ cm
5. Volume of the rubber stopper you massed in part A___________ mL
6. Volume of the metal sample used in part A___________ mL

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Experiment 3: Bunsen burner
Objective: To learns how to light and adjust a burner flame and to locate the hottest part of the
flame.
Theory: Often a chemist needs heating sources to heat materials. The Bunsen burner is a
convenient source of heat in the laboratory. Although there are several varieties, their principle of
operation is the same and is similar to that of the common gas stove. The Bunsen burner requires
gas and air, which it mixes in various proportions. The amount of air and gas mixed in the
chamber is varied by use of the collar. The relative proportions of gas and air determine the
temperature of the flame.
Figure 3.1: Different parts of the flame and basic reaction in the flame
Apparatus: Spark lighter, Bunsen burner, support stand, w/ ring wire, screen gloves, &
thermometer
Chemicals: iron wire, Aluminum wire & copper wire
Part 1: Parts of a Bunsen burner!
1. Use one of the Bunsen burners from the lab. (In the space below)
2. Identify and label its parts using the guide sheet.
3. Give the purpose of each part of the Bunsen burner.
5. Continue this until the water boils (100 degrees Celsius). Write down the amount of time it
took to boil.
6. Write data on board and Obtain data from all groups, make sure you have data from all heights.

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Figure 3.2 .Schematic of Bunsen burner at different height to heat sample
Examine your burner and locate the gas and air flow adjustments (valves). Determine how each
valve operates before connecting the burner to the gas outlet. Close both valves; connect a rubber
hose to the gas outlet on the burner and the desk; then open the desk valve about two-thirds of the
way. Strike a match or use a gas lighter.
Hold the lighted match to the side and just below the top of the barrel of the burner while
gradually opening the gas valve on the burner to obtain a flame about 7 or 10 cm high. Gradually
open and adjust the air valve until you obtain a pale blue flame with an inner cone.
Table 3.1: Flame temperatures can be observed using the melting points of metals.
Adjust the burner to a non-luminous flame to measure the temperatures in the various regions of
the flame. Use crucible tongs to hold 2 cm strips of iron wire, copper wire, and aluminum wire in
the various regions of the flame. The melting point of iron is 1535 °C, which of copper is 1083
°C and that of aluminum is 660 °C. On the Report sheet, record the estimated temperature of the
flame in the regions designated in Figure 3.3.
Figure 3.3. Regions of the flame for temperature measurement
Metal Melting point (oC)
Iron (Fe) 1535
Copper (Cu) 1083
Aluminium (Al) 660

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REPORT SHEET
A. Bunsen Burner
Indicate the approximate temperature of the following regions of the flame
(see Figure 3.3)
a. Region : Top of the outer cone__________ °C
b. Region : Center of the outer cone __________ °C
c. Region : Top of the inner cone __________ °C
d. Region : Center of the inner cone __________ °C

13
Experiment 4: Physical and Chemical Changes
Objective: To distinguish chemical change and physical change
Theory: One of the basic areas of interest for chemists is the study of the regrouping of atoms to
form new substances. In order to determine if such a chemical change has occurred, there
should be a change in the properties of the reactants that can be observed. The reaction, for
example, of two colorless solutions to produce a mixture of two new colorless solutions
could be quite difficult for us to observe. It would be much easier to follow the course of a
reaction if one of the following occurred:
1. An unexpected color change occurred during the reaction.
2. One of the new materials was a gas that was insoluble in the solution and escaped
to the atmosphere as bubbles.
3. One of the new materials was a precipitate that settled out of solution.
4. A characteristic odor (gas) either appeared or disappeared.
Other changes that only involve changes in form or appearance are called physical
changes. These do not produce new substances but only change the physical properties of the
material; for instance, when wheat is ground to make flour or when water is frozen to make ice.
Simply mixing two substances to form mixture, such as the mixing of sand and salt, is another
example of a physical change.
Changes in temperature often accompany both chemical and physical changes. A temperature
change, therefore, only indicates that there has been a change, but one must investigate
the reaction further to determine whether the change was chemical or physical.
BEFORE LAB: Write the formula and state notation for each reactant so you are able to find
the reagents.
REPORT: For the lab report, in the conclusion column indicate whether a chemical or
physical change has occurred. If a gas forms, identify the gas by name or formula; refer to your
pre study for a list of possible gases. Where there is a space for the reaction, write the
balanced, molecular equation including state notations, (s), (l), (g) or (aq).

14
For physical changes, the reaction could actually produce no new products or it could be
simple dissolving of a solid.
(For simply dissolving a solid, write a net ionic equation.) You are strongly encouraged to
refer to other labs, your text, other texts, the internet and your notes for help with the more
complex reactions.
WASTE: drain – down the drain with water, HM – heavy metal waste container.
Chemicals: copper metal, concentrated nitric acid, solid potassium chloride, solid iodine,
crystals of iodine ,hexane , zinc metal , dilute nitric acid, nickel (II) sulfate solution , sodium
carbonate solution, copper(II) sulfate solution , aqueous hydrochloric acid, copper metal , silver
nitrate solution, solid calcium chloride, iron(III) chloride, potassium thiocyanate, calcium
carbide, potassium permanganate solution & hydrogen peroxide solution
Apparatus: fume hood, measuring cylinder, analytical balance, Bunsen burner, dropper,
crucible & test tube

1
Procedure
Experiments Observations Conclusi
on
Waste
1. IN THE HOOD, add a small
piece of copper metal to 2 mL of
concentrated nitric acid.
Change:
Gas:
HM
REACTION:
2. Heat a few crystals of solid
potassium chloride in a dry test
tube over a Bunsen burner.
Change: cool then drain with water
3. Observe the color of solid
iodine. Place a couple crystals of
iodine into 1 mL hexane (C6H14).
Change: Organic
REACTION:
4. Mix 1 mL of nickel (II) sulfates
solution with 1 mL of sodium
carbonate solution.
Change: HM
REACTION:
5. Add 1 mL of copper (II) sulfate
solution to 1 mL of aqueous
hydrochloric acid.
Change: HM
REACTION:
6. Add a small piece of copper
metal to 2 mL of silver nitrate
solution. Observe immediately and
after approximately 15 minutes
Change: HM
REACTION:
7. Record the temperature of 5 mL
of water. Add a dime sized amount
of solid calcium chloride to the
Change: Drain

2
water. Swirl. Record the
temperature after swirling.
REACTION:
8. Add 1 mL of dilute hydrochloric
acid solution to 1 mL of silver
nitrate solution.
Change: HM
REACTION:
9. Mix 3 mL of iron (III) chloride
with 6 drops of potassium
thiocyanate.
Change: HM
REACTION:
10. Place 5 mL of dilute
hydrochloric acid in a test tube.
Place a thermometer in the acid
and record the temperature. Add 5
mL of dilute potassium hydroxide
solution. Record the temperature of
the mixture after adding the
sodium hydroxide.
Change: Drain
REACTION:
11. Mix 1 mL of potassium
permanganate solution with 1 mL
dilutes hydrochloric acid and 1 mL
of 3 % hydrogen peroxide solution.
Place a glowing splint into the
mouth of the test tube
Change:
Gas:
HM
REACTION: (Find the reaction between hydrogen peroxide and potassium permanganate in an acidic solution. You may write either the
net ionic equation or the molecular equation.

1
Physical and Chemical Changes
Pre study
1. Classify the following as chemical (C) or physical (P) changes.
A. Two clear, colorless solutions are mixed and a yellow solid forms.___________
B. Helium boils at 4.22 K.___________
2. Write the following underlined reaction in terms of a balanced equation. Write each reactant
and product (you have to determine the products) as a formula, including state notations. All
solutions are aqueous. A light green solution of iron (II) nitrate is mixed with a solution of
sodium carbonate resulting in the formation of a yellowish precipitate and a solution.
3. Look up the reaction between solid copper and the oxidizing agent nitric acid that produces
nitrogen monoxide gas (NO, aka nitric oxide). Write the balanced equation
4. Look up the properties of the following gases & Fill in the following table.
Gas Supports
Combustion*
Burns
(Flammable)
Odor Color
Yes/No (None or if any, describe it)
a. acetylene (C2H2)
b. carbon dioxide (CO2)
c. hydrogen (H2)
d. nitrogen (N2)
e. nitrogen dioxide (NO2) No No
f. oxygen (O2)
g. sulfur dioxide (SO2) No No
*If a gas supports combustion, it must be present for other substances to burn. Supports
combustion does not mean flammable.

2
Experiment 5: Diffusion of gases: Determination of Graham’s rate laws
Objective: To compare the rates of diffusion of HCl and NH3 gases and to determine
graham’s rate law
Theory; in this experiment, the relative rates of diffusion of two gases will be determined.
Rates of diffusion yield information that can lead to calculation of the molecular weights of
gases. Gases consist of particles that are in constant rapid motion. This motion causes gases to
travel across space (diffuse) and completely mix with each other. It is this diffusion that
eventually causes one to notice smells such as perfume, fish, ammonia, etc.
In this experiment, the rates of diffusion of two gases, ammonia (NH3) and hydrogen chloride
(HCl) will be investigated. These gases are convenient to use for such an experiment because,
when they meet and react, they form a white smoke consisting of ammonium chloride
(NH4Cl):
NH3 (g) + HCl (g) →NH4Cl(s)
Therefore, if ammonia gas and hydrogen chloride gas are released simultaneously a opposite
ends of a glass tube, a white ring of smoke will form at the location where they meet.
This experiment will demonstrate rates of diffusion, a property of gases investigated by Thomas
Graham. In 1829, he proposed his law of diffusion which states that the rate of diffusion of a gas
is inversely proportional to the square root of its density:
R 
1
√𝑑
However, since the ideal gas law indicates that the density of a gas and its molecular weight are
proportional, we can write:
R 
1
√𝑀𝑊
If the rates of diffusion of two gases are compared, this yields the following equations:
𝑅1
𝑅2
=
√𝑀𝑊2
√𝑀𝑊1
Thus, if the rates of diffusion of two gases are known and the molecular weight of one of them is
known, the molecular weight of the other gas can be calculated:
𝑀𝑊2 =
𝑅1
2
𝑅2
2 𝑀𝑊1

3
In this experiment, the distance each gas travels will be measured as well as the time it takes for
them to meet and react (D = distance, t = time):
𝑅𝑁𝐻3
𝑅𝐻𝐶𝑙
=
𝐷𝑁𝐻3
𝑡
=
√𝑀𝑊𝐻𝐶𝑙
√𝑀𝑊𝑁𝐻3
=
𝐷𝐻𝐶𝑙
𝑡
This can be rearranged to:
𝑀𝑊𝐻𝐶𝑙 =
𝑅2
𝑁𝐻3
𝑅2
𝐻𝐶𝑙
𝑀𝑊𝑁𝐻3
Chemicals: concentrated NH3 solution, concentrated HCl solution, tap water
Apparatus: gas diffusion apparatus, cotton
PROCEDURE
Wear your safety glasses while doing this experiment. You will carry out the experiment with a
partner .Obtain a gas diffusion apparatus from the stockroom, consisting of a glass tube and two
corks which have holes drilled in them and a stopwatch. Clamp the glass tube horizontally on a
ring stand (See Figure 5.1). Loosely place a wad of cotton into each cork and in the hood saturate
one wad of cotton with drops of concentrated NH3 solution and the other wad of cotton with
drops of concentrated HCl solution. Be sure to keep the corks far apart to avoid a premature
reaction.
Figure 5.1: Gas Diffusion Apparatus
Once the corks are ready, return to your lab bench and insert the corks simultaneously into the
opposite ends of the glass tube. One partner will begin to measure the time. Now, very carefully
observe the glass tube and at the time of the first appearance of a white ring of ammonium
chloride, record the time and mark the location of the ring on the tube. Measure and record the
distance each gas traveled.
To repeat the experiment, clean the glass tube with tap water. Rinse with deionized water and
then with alcohol. Dry the tube completely, and then clamp the tube horizontally, once again.
Discard the cotton wads in the waste beaker in the hood and then place clean cotton in the corks
and saturate as before, then repeat the experiment

4
Pre-Laboratory Assignment
1. A gas of unknown molecular weight is found to diffuse at a rate of 0.19 cm/s compared to
0.59 cm/s for helium gas. Calculate the molecular weight of the unknown gas.
2. Which of all of the possible gases should have the highest rate of diffusion?
3. A balloon inflated with helium gas will rise to the ceiling of a room. However, after a certain
period of time, the balloon will descend to the floor. Why does the balloon eventually descend?
4. If the cork containing the concentrated NH3 solution were inserted into the glass tube several
seconds after the insertion of the cork containing HCl, what effect would this have on the
calculated value of the molecular weight of HCl?
Report Sheet
Section Name
TRIAL
First Second Third
Time of insertion of stoppers
Time of appearance of smoke
Total time elapsed (seconds)
Distance travelled by NH3 (cm)
Distance travelled by HCl (cm)
Rate of diffusion of NH3 (cm/s)
Rate of diffusion of HCl (cm/s)
MW of NH3 17.0 17.0 17.0
Calculated MW of HCl
Mean MW of HCl
True value of the MW of HCl 36.5
Percent error of mean value
Show your calculations below:

5
Post-Laboratory Questions
1. How many times faster will CH4 gas diffuse compared to C4H8 gas?
2. If CH4 gas and C4H8 gas are released simultaneously at the left and right ends respectively of
a 50.0 cm long glass tube, at what distance from the left end of the tube will they meet?
3. Methane gas, CH4, diffuses 2.3 times faster than an unknown gas at the same temperature and
pressure. What is the molecular weight of the unknown gas?
4. Does the first appearance of the white smoke indicate the first contact of the NH3 and HCl
molecules? Explain your answer. If the answer is “no”, how will this affect the calculated value
of the molecular weight of HCl?

6
Experiment 6: Acid-base reactions: use of acid-base indicators
Objective: To Practice acid base reaction by using acid base indicators
Theory: This experiment demonstrates the analytical technique, titration. In a titration, a
solution is delivered from a burette until it completely consumes another solution in a flask.
You will chemically react an acid with base, until the endpoint of the reaction is shown by the
indicator phenolphthalein. You will first determine the concentration of a base solution
(standardization) and then of an unknown acid solution using previously standardized base. The
reaction of a strong acid and a strong base goes quickly: i.e.
HCl (aq) +NaOH (aq) NaCl (aq) +H2O (l)
Acid base salt water
This reaction is often called neutralization.
Part A: To standardize the NaOH solution, we will react it with potassium hydrogen phthalate,
KHC8H4O4 Because of its complex formula; this compound is often called “KHP”. The molar
mass of KHP is 204.2 g/mol.
KHC8H4O4 (aq) +NaOH (aq) KNaC8H4O4 (aq) +H2O (l)
Starting with a known mass of KHP then recording the volume of NaOH needed to reach the
endpoint, we can calculate the molarity of the base .We want the titrated solution to be a very
pale pink, not bright rosy red, at the endpoint.
Part B: Once the concentration of the NaOH solution in the burette is known, we can
determine the concentration of an HCl solution by titrating it with the NaOH solution from part
A. Phenolphthalein is used as the indicator by adding it to the HCl solution.
Chemicals: Solid KHC8H4O4 (“KHP”), Deionized water, Phenolphthalein solution, Approximately
0.2M NaOH solution & Unknown HCl solution
Apparatus: Weighing vial, scoop or spatula, 125-mL Erlenmeyer flasks or 2 flasks, and a 250 mL
beaker, Wash bottle with deionized water, 50-mL burette, burette holder and ring stand, Plastic
funnel & 500-mL Florence flask.

7
Procedure:
Part A: Standardization of an unknown base
1. Obtain and rinse burette with deionized water from your wash bottle.
CAUTION: do not stick the end of the burette under the faucet! It leaks all over the floor. Be
sure to open the valve at the bottom and let some water rinse out the stopcock. If you think the
tip is clogged, ask your instructor to check it.
2. Practice reading the meniscus in your burette while cleaning it with deionized water. Read your
burette by estimating between the 0.1-mL marks. In other words, your recorded volume
measurements include an uncertain digit at 0.01 mL. If the meniscus is right on a mark, record
the second decimal place as a zero. Hold the clean burette in the burette clamp that is attached
to the ring stand.
3. Obtain a 500 mL Florence flask, clean it well with deionized water, and then pour
approximately 16-17 mL of 6 M NaOH. Ask your instructor which bottle of sodium
hydroxide your need to use. Dilute this solution with deionized water to just below the neck
of the flask. Do not worry about exact measurements; you will use the technique of
standardization to determine the exact concentration. Mix the solution well. Do not discard
this solution until the entire experiment is complete.
4. When your burette is clean, rinse it 2 times with small portions (about 15- 20 mL) of your
NaOH, discarding the rinses. Then clamp the burette in on to a burette clamp. Close the valve
at the bottom, place a plastic funnel in the top opening and carefully pour base solution into
the burette until the solution level is near the 0.00 mL mark). Make sure there are no air
bubbles trapped in the tip of the burette. Record the initial base volume reading for this trial.
(0.00 or slightly below is OK)
5. Weigh three samples of KHP (about 1.00 to 1.10g into three separate 125-mL
Erlenmeyer flasks and/or beakers. Label the flasks and record you weights.
6. Dissolve the KHP crystals in your flasks in about 30 mL of deionized water if some KHP sticks
to the sidewalls of the flask; wash it down with deionized water from your wash bottle. If the
KHP doesn’t dissolve in a short time, you may gently warm the solution in a hot tap water bath.
Add 2-3 drops of phenolphthalein solution from the dropper bottle on the shelf.
7. Place the flask under the burette. A piece of white paper under the flask makes it easier to see
the pale pink color at the endpoint. Open the valve and allow base to flow from the burette into
the flask. Swirl continually to mix the solutions.

8
As you get close to the endpoint, the solution will begin to show pink color that goes away
when you mix. Slow the rate of base addition to one drop at a time, If you splash the solution
up onto the sidewalls of the flask, spray a stream of water from your wash bottle over the
inside of the flask. The extra water will not affect your results. When one drop of base
changes the solution from colorless to pale pink, close the burette valve, rinse down the flask
one last time, and make sure that the pink color lasts for at least 30 seconds. If so, record the
final burette volume reading.
8. Discard the titrated solution into the sink. Do at least three successful titrations that achieve
a pale pink color of the indicator. If color is bright rosy red, you have overshot the endpoint
and cannot include that trial in your calculations. If you have time after completing the
standardization of the base, you may continue right into Part B, or keep the NaOH solution,
tightly stoppered, in your locker or in the cabinets in the back of the room, until the next lab.
9. At the end of the lab period, drain and discard the leftover base solution from the burette and
rinse the burette with several portions of water. Return the burette.
Part B: Determination of an unknown acid
1. Prepare a burette, filled with standardized NaOH solution, and three clean 125-mL
Erlenmeyer flasks as you did earlier.
2. Your instructor will assign you one of the several unknown HCl solutions for your
experiment. Be sure to record the unknown code in your notebook and on your lab report.
Pour about 100 mL of your assigned unknown HCl solution into a 150-mLbeaker.
3. Use a volumetric pipette to measure out 25.00 mL of your unknown acid solution.
(Instructor will demonstrate the use of the bulb and pipette).
4. Place the acid solution in an Erlenmeyer flask. Add 2-3 drops of indicator. (Remember,
the phenolphthalein changes from colorless to pink at the endpoint.)
Place the flask under the burette, record the initial burette reading, and then add base from the
burette until you reach the endpoint, as before. Record the final burette reading, discard the sample
and repeat until you have at least three successful trials. Return the cleaned burette

9
This value will be used in calculations for part B of the experiment
Part B:
Data Table: (Be sure to record the correct number of significant figures.)
HCl solution Unknown Code___________
Trial 1 Trial 2 Trial 3
Volume of HCl solution used ________mL _________mL ____mL
Initial burette reading of NaOH ________mL _________mL _____mL
Final burette reading of NaOH ________mL _________mL _____mL
Calculations: Show a complete calculation from Trial 1 for each of the following, *then
complete the work for the remaining calculations in your notebook.
Volume of NaOH solution used ________mL _________mL _____mL
Convert the mL values to liters ________L _________L _____L
Moles of NaOH used in each sample (Use the molarity of NaOH from part A): ____mol _____mol ____mol
Moles of HCl reacting ________mol _________mol _____mol
(Refer to the first equation, in the introduction)
Molarity of HCl solution M M M
Average molarity (experimental value) of the unknown HCl solution __________M
(To be filled in by the instructor)
Actual molarity of the unknown HCl M
Percent error %

10
REVIEW QUESTIONS
Name ___________
Section
A student weighed out 1.54 g of KHP. How many moles is this? How many moles of NaOH will
react with this KHP sample?
1. The equation for the reaction of KHP and NaOH shows a 1:1 ratio for the two reactants.
Experimentally, how do you know that the base solution that you delivered from the burette
had just as many moles as were in the KHP sample that you weighed out in the beginning of
the experiment?
2. What might happen to your calculated NaOH molarity if you use tap water rather than
deionized (purified) water to dissolve the KHP crystals or to rinse down the walls of the flask
during the titration? (Tap water contains some calcium bicarbonate.)
3. Write and balance the equation for a neutralization of a sulfuric acid solution of unknown
concentration by sodium hydroxide. Calculate the molarity of an unknown sulfuric acid
solution if a 25.0-mL sample of the acid solution consumes 27.2 mL of 0.138 M NaOH
solution in a titration.

11
Experiment 7: Determination of solubility of salts: Investigating the solubility of
ionic and covalent compounds
Objective: To investigate solubility of salt in water, some ionic compounds and covalent
compounds
Theory
The solubility of a solute is the amount of solute dissolved in a given amount of a solvent at
equilibrium at specified conditions. The usual units used to express solubility are gram of solute per
100 grams of solvent at a specified temperature. Solubility of different substances usually varies
with temperature. A solution is said to be saturated if there is undissolved solute in equilibrium
with the solution. If a solution contains more solute that can dissolve at a given condition it is called
a supersaturated solution, and if less solute dissolves in the solution than it can dissolve at a given
temperature it is said to be unsaturated.
General Solubility Rules for Inorganic Compounds
Nitrates (NO3
–): All nitrates are soluble.
Acetates (C2H3O2
–): All acetates are soluble; silver acetate is moderately soluble.
Bromides (Br–) Chlorides (Cl–) and Iodides (I–): Most are soluble except for salts containing silver,
lead, and mercury.
Sulfates (SO4
2–): All sulfates are soluble except barium and lead. Silver, mercury (I) and calcium
are slightly soluble.
Hydrogen sulfates (HSO4
–): The hydrogen sulfates (aka bisulfates) are more soluble than the
sulfates.
Carbonates (CO3
2–), phosphates (PO4
3–), chromates (CrO4
2–), silicates (SiO4
2–): All carbonates,
phosphates, chromates, and silicates are insoluble, except those of sodium, potassium, and
ammonium. An exception is MgCrO4, which is soluble.
Hydroxides (OH–): All hydroxides (except lithium, sodium, potassium, cesium, rubidium, and
ammonium) are insoluble; Ba (OH) 2, Ca (OH) 2 and Sr (OH) 2 are slightly soluble.
Sulfides (S2–): All sulfides (except sodium, potassium, ammonium, magnesium, calcium and
barium) are insoluble. Aluminum and chromium sulfides are hydrolyzed and precipitate
as hydroxides.
Sodium (Na+), potassium (K+), ammonium (NH4
+): All sodium, potassium, and ammonium salts
are soluble. (Except some transition metal compounds)

12
Silver (Ag+): All silver salts are insoluble. Exceptions: AgNO3and AgClO4; AgC2H3O2and Ag2SO4
are moderately soluble.
A summary of the general properties of ionic compounds:
1. Ionic compounds do not contain molecules. They are aggregates of positive ions and
negative ions. In the solid state, each ion is surrounded by ions of the opposite charge,
producing an orderly array of ions called crystal.
2. At room temperature ionic compounds are hard and rigid crystalline solids. This is due to the
existence of strong electrostatic forces of attraction between the ions.
3. Ionic compounds have relatively high melting and boiling points. This is due to the presence
of strong electrostatic forces between the ions. These forces can be overcome only by applying
very large amounts of energy.
4. Ionic compounds can conduct electric currents when molten or in aqueous solution. This is
due to the presence of mobile ions in molten state or in solution. However, ionic compounds do
not conduct electricity in the solid state.
5. Ionic compounds are soluble in polar solvents such as water. They are insoluble in non-polar
solvents such as benzene.
Chemicals: distilled water & saturated solution of Sodium chloride solution
Apparatus: Evaporating dish, Bunsen burner, Analytical Balance, wire gauze, thermometer,
glass rod, Beaker & Crucible tongs.
Procedure 1:
1. Take a dry evaporating dish and weigh accurately.
2. Record the temperature of a saturated solution of Sodium chloride
3. Add about 10 to15 mL of the saturated Sodium chloride solution into the evaporating dish and
take the mass of the solution and the evaporating dish accurately.
4. Take a beaker nearly filled with water and boil the water with a burner flame.
5. With crucible tongs carefully put the evaporating dish with solution on the beaker. Do not heat
the evaporating dish directly on the flame.
6. Continuously stir the solution in the evaporating dish till a fine powder of Sodium chloride salt
remains.

13
Make sure that there is an opening between the mouth of the beaker and the bottom of the
evaporating dish. Never heat the beaker to dryness. If you have a small volume of water remaining
in the beaker carefully add more water.
7. Cool the evaporating dish; dry the moisture on its outer bottom with a piece of paper.
8. Weigh the evaporating dish with the salt accurately and put the salt in a container prepared for
this purpose.
i. Solubility of salts in the water
Name(s) ________________________ Laboratory Instructor___________________
Section__________________________ date_________________________________
Format sheet
Mass of empty evaporating dish (X) ______________________g
Temperature of the saturated Sodium chloride (T) ________ o
C
Mass evaporating dish + solution (y) _____________g
Mass evaporating dish + residue (n) _____________g
Mass of solution alone (y-x) ____________________g
Mass of solute alone (n-x) ______________________g
Mass of water lost by evaporating (y-x) – (n-x) = (y-n) __________________g
Solubility of the salt in 100grams of water at T o
C
mass of solute×100
mass of water lost by evaporation
=______________g

14
Error calculation
(In standard tables look for the solubility of saturated Sodium chloride at above given temperature
(T), and compare with your result above)
% error =
practical−theoretical
theoretical
× 100
Possible sources of error _________________
Procedure 2:
Chemicals: NaCl, CuCl2, ethanol, hexane, distilled water and benzene.,
Apparatus: Test –tube, beaker & dropper.
1. Place about 1 g each of sodium chloride (NaCl) and copper (II) chloride (CuCl2) in separate test
tubes. Add about 5 mL of water (polar solvent) and shake well.
2. Repeat experiment 1 using the following solvents instead of water. Ethanol (polar solvent),
hexane and benzene (non-polar solvents). These solvents are highly flammable and should be kept
away from flames.
Prepare a table as shown below and fill in the results of the solubility tests.
Substances Ethanol Hexane benzene water
NaCl (s)
CuCl2(s)
II. General properties of covalent compounds
1. Covalent compounds are generally liquids or gases at ordinary temperature. For example: water
and ethyl alcohol are liquids. Hydrogen chloride, methane and carbon dioxide are gases. Same
covalent compounds are solids (e.g. sugar)
2. As compared to ionic compounds, covalent compounds have relatively lower melting points and
boiling points.
3. They do not conduct electric current when molten or in aqueous solution, because they consist of
molecules rather than of ions.
4. Covalent compounds are insoluble in polar solvents such as water. They are soluble in non-polar
solvents such as benzene and carbon tetrachloride

15
Procedure 3:
Apparatus: Test-tubes, test tube rack
Chemicals: Naphthalene, graphite, iodine, ethanol, hexane and benzene.
1. Arrange 12 test tubes in three sets (A, B, C) of 4 test tubes each.
To each test tube of set A, add 1 g of naphthalene.
To each test tube of set B add 1 g of graphite and to each test tube of set C add 1 g of iodine.
2. Add about 10 mL of each the following solvents to the four test tubes of each set separately and
shake well. Water, ethanol, hexane and benzene.
Caution: Ethanol, hexane and benzene are all highly flammable. Observe and record whether the
solids are very soluble, slightly soluble or insoluble.
Solubility
water Ethanol hexane benzene
Naphthalene
Graphite
Iodine

16
Experiment 8: Simple and Fractional Distillation
Objective: 1. to purify 2-Propanol using simple distillation method
2. To purify and separate ethanol from Ethanol/Water Mixture
Theory
The use of distillation to separate the components in a mixture is based on the principle that the
boiling liquid in the vial and the vapor produced have a different composition. When a solution of a
liquid containing a dissolved solid is heated to the boiling point of the liquid, the vapor will have a
higher concentration of the liquid (i.e., the more volatile component of the solution / the component
with the lower boiling point). The vapor rises up the glassware where it cools and condenses. When
the vapor condenses, it is called the distillate. When enough distillate has collected in the Hickman
still head, it can easily be removed using a pipet. The distillate will then be enriched in the more
volatile component of the solution. Distillation can also be performed on a solution of multiple
liquids. In this case, the liquid with the lowest boiling point will be the most enriched in the
distillate, regardless of whether it is the major component of the solution, because it is the one that is
the most volatile. Distillation can also be utilized to manage natural resources. It is an indispensable
technique for obtaining drinking water from seawater. Distillation is the oldest and still most widely
used technology for desalination (removal of salt from saltwater). In the petroleum industry, oil
refineries use distillation to transform crude oil into fuels and chemical feed stocks. Distillation is
also employed by the alcohol and brewing industry to increase the alcohol content of fermented
products. In this experiment, simple distillation will be used to separate an organic liquid from an
organic solid. Simple distillation involves a single cycle of vaporization and condensation. Simple
distillation is used to purify liquids that contain either nonvolatile impurities, such as salts, or very
small amounts of higher- or lower boiling liquids. Simple distillation is not a practical method for
separating compounds with similar boiling points. In order to separate liquid mixtures where the
components have similar boiling points and/or are present in comparable amounts, fractional
distillation must be employed. In fractional distillation, insulated fractionating columns permit
multiple cycles of vaporization and condensation in a single operation. The column consists of
closely spaced packing material or “plates.” The vapor condenses on multiple surfaces in the
fractionating column and the resulting liquid revaporizes. At each stage in the series of vapor–liquid
equilibrium, the vapor becomes more enriched in the more volatile (lower-boiling) component.

17
Given a sufficient number of “plates,” the mixture will distill in fractions. Each fraction consisting
of only a single pure substance
Fractional distillation can be considered a series of simple distillations but rather than having many
condensers and receivers, the evaporation/condensation cycles take place in a single distillation
column. Distillation columns are high surface area tubes that allow multiple
vaporization/condensation cycles to occur at once. The added surface area in the column can come
from glass protrusions like in Vigreux columns, the ones we will be using, or from a packing
material such as glass beads or steel wool put in a hollow tube which is known as a Hemple column.
In fractional distillation, vapor rises up the column and condenses on the packing, then re-
evaporates, rises further up the column and condenses again. This evaporation/condensation
continues up the column and with each cycle the vapor becomes more and more pure in the lower
boiling component. In this way, fractional distillation accomplishes in one apparatus what would
require several simple distillation setups. The more vaporization/condensation cycles that the
mixture goes through, the better the separation. Although more surface area in the column leads to
better separation, it also makes the distillation process slower. Additionally, surfaces in the column
will be covered with condensation, and liquid that sticks to the surfaces will not be distilled. The
liquid condensed on the surfaces in the column is known as hold-up. When using fractional
distillation chemists try to strike a balance between the quality of the separation and the speed of the
distillation. During fractional distillation there is a temperature gradient over the length of the
column. The boiling point of a mixture varies with the composition of the mixture; the higher the
percentage of the low boiling component, the lower the boiling point will be. Thus as the mixture
travels up the column and becomes purer, the temperature of the column decreases. When the
distillate finally reaches the still head, the thermometer reading begins to rise. The thermometer
reads the boiling point of the vapor that is condensing on it. (Boiling point and condensation point
are the same temperature.) If the thermometer reads at the boiling point of the lower boiling
compound, the distilled liquid is pure. When the temperature rises above the boiling point of the low
boiling compound that means more and more of the high boiling component is starting to distill
over. To ensure a temperature gradient over the column, the still pot should be heated slowly. Most
of the vapor from the boiling still pot should condense and drip back down into the still pot, a
process known as refluxing.

18
Figure 8.1: Simple distillation Setup
Figure8. 2: Fractional distillation setup

19
apparatus : Fractional distillation setup, Simple distillation setup , 50-mL round bottom flask as the
distillation pot, 100-mL round bottom flask as the receiver, test tube, boiling stones to keep the
solution boiling smoothly!, Beaker, funnel, graduated cylinder, thermometer
Chemicals: 2-propanol, dye, water & ethanol
Procedure
Part A: Simple Distillation of 2-propanol
Safety: 2-propanol is a highly flammable liquid and a severe eye irritant no flames will be
allowed in lab while it is in use. As for every experiment, goggles must be worn, even though
you may not actually be working the chemicals, if there is anyone using 2-propanol in the lab.
You will use a 50-mL round bottom flask as the distillation pot, and a 100-mL round bottom
flask as the receiver. As a standard rule, anytime you are boiling an organic compound, you will
always include a few (not a handful!) boiling stones to keep the solution boiling smoothly!
Using a beaker (not a graduated cylinder) obtain about 30-40 mL of "impure" 2-propanol (this
sample of 2-propanol has had a small amount of a soluble, non-volatile dye added to it as an
impurity). Add the 2-propanol to the distillation pot (never pour anything through a ground-
glass opening without using a funnel), add the boiling stones, and begin the distillation
(remember to turn on the cooling water before you turn on the heat). Collect your distillate in a
pre-weighed graduated cylinder (10-mL or 25-mL or 50-mL cylinder).
Once the temperature starts to rise above room temperature, you should start to record the
thermometer reading every minute. You should start to record the temperature at any time up to
when the solution starts to boil. Record every minute the temperature you read. Continue to do
your distillation until you have collected about 20 mL of distillate. Be sure that the distillation
pot never goes dry (never let a heated flask go dry!).
You must plot your data by hand using graph paper, or you can use Excel or another graphing
program for a graph for inclusion in your lab notebook and written report.
Measure the volume of the distillate collected. Using a pre-weighed graduated cylinder (10-mL
or 25-mL or 50-mL cylinder), you should determine the density of the collected distillate. You
will also determine the refractive index of your distilled liquid. Your instructor will describe
how a refractive index is determined.

20
Dispose of your liquid, and any liquid remaining in the distillation pot, in the liquid wastes. Be
sure to make certain that no boiling stones are deposited into the liquid waste. Put the boiling
stones in the solid waste container.
Part B: Fractional Distillation of an Ethanol/Water Mixture
Safety: Ethanol is a flammable liquid and an irritant; avoid contact and inhalation -- wear
gloves while handling it. No flames will be allowed in lab while ethanol is in use. Goggles must
be worn whenever anyone is using chemicals.
Set up a fractional distillation apparatus as demonstrated by your instructor. Use glass beads to
pack the fractionating column (your instructor will demonstrate how to pack the column). Try
adding some glass beads directly to your fractionating column. If the glass beads stay in the
column, there is no problem, but if any beads go through, try adding a larger amount of glass
bead, and their packing inside the fractionating column should allow them to stay in place. Do
not ever use glass wool or anything besides beads in the fractionating columns.
To do this part of the experiment, you will use a 100-mL round bottom flask as the distillation
pot. You will need a number of receivers; it is best to use test tubes. Measure into one test tube
about 4 mL of water. Use this sample to know how much liquid you need to obtain about 4 mL
of distillate during this part of the experiment. Continue collecting 4-mL samples until you have
collected about 30 mL of distillate. Determine the refractive index of each collected sample, as
well as determining the refractive index for pure ethanol and pure water.
 Obtain about 50.0 mL of the 50% (v/v) ethanol/water mixture, and pour into the distillation
pot.
 Add a few boiling stones.
 Turn on the heating mantel to obtain a steady boiling mixture.
 Monitor time and temperature during the entire distillation process
Start recording the temperature as soon as your sample begins to boil. Record the temperature every
30 sec. Collect your distillate into test tubes. You should collect about 4 mL in each test tube, but it
is not necessary to measure each tube. As a comparison, add about 4 mL of water into a test tube.
Collect about the same amount of liquid into each of the tubes during the distillation process.
Continue recording the temperature until you stop collecting your samples. Collect about 30 mL of
distillate.

21
You must plot your data by hand using graph paper. You should have two plateaus, one for the
boiling point of the ethanol and the other for the boiling point of the water. Your graph for your lab
report and for your notebook must show these two plateaus.
After the distillation has finished, you will have a good determination for the boiling point of
ethanol. Determine the refractive index for each of your samples.
Do not ever throw any glass beads away. At an expense of about 25 cents ($0.25) per glass
bead, they are very expensive. Keep your glass beads in your fractionating column (add some
tissue to the top to prevent spillage) until the next lab period. Never throw away any glass
beads.

22
Experiment 9: Separation of mixtures: Extraction; Distinguishing compounds
and mixtures; Separation of a mixture using a magnet; Recrystallization; and
Filtration
Objective: To become familiar with the methods of separating substances from one another using
Distinguishing compounds and mixtures ,extraction, magnet, recrystallizations and Filtration.
Theory
Materials that are not uniform in composition are said to be impure or heterogeneous and are called
mixtures. Most of the materials we encounter in everyday life, such as cement, wood, and soil, are
mixtures. When two or more substances that do not react chemically are combined a mixture results.
Mixtures are characterized by two fundamental properties: First, each of the substances in the
mixture retains its chemical integrity; second, mixtures are separable into these components by
physical means. If one of the substances in a mixture is preponderant-that is, if its amount far
exceeds the amounts of the other substances in the mixture-then we usually call this
mixture an impure substance and speak of the other substances in the mixture as impurities.
The preparation of compounds usually involves their separation or isolation from reactants or other
impurities. Thus the separation of mixtures into their components and the purification of impure
substances are frequently encountered problems. You are probably aware of everyday problems of
this sort. For example, our drinking water usually begins as a mixture of silt, sand, dissolved salts,
and water. Since water is by far the largest component in this mixture, we usually call this impure
water. How do we purify it? The separation of the components of mixtures is based upon the fact
that each component has different physical properties.
The components of mixtures are always pure substances; either compounds or elements and each
pure substance possess a unique set of properties. The properties of every sample of a pure
substance are identical under the same conditions of temperature and pressure. This means that once
we have determined that a sample of sodium chloride, NaCl, is water soluble and a sample of silicon
dioxide (sand), SiO2, is not, we realize that all samples of sodium chloride are water-soluble and all
samples of silicon dioxide are not. Likewise, every crystal of a pure substance melts at a specific
temperature, and at a given pressure, every pure substance boils at a specific temperature.

23
Although there are numerous physical properties that can be used to identify a particular substance,
we will be concerned in this experiment merely with the separation of the components and not with
their identification. The methods we will use for the separation depend upon differences in physical
properties, and they include the following:
Filtration: This is the process of separating a solid from a liquid by means of a porous substance, a
filter, which allows the liquid to pass through but not the solid Common filter materials are papers,
layers of charcoal, and sand. Silt and sand can be removed from our drinking water by this process.
Extraction: This is the separation of a substance from a mixture by preferentially
dissolving that substance in a suitable solvent. This process is used to separate a soluble compound
from an insoluble compound.
Recrystallization (or Crystallization) is a technique used to purify solids. This procedure relies on
the fact that solubility increases as temperature increases (you can dissolve more sugar in hot water
than in cold water). As a hot, saturated solution cools, it becomes supersaturated and the solute
precipitates (crystallizes) out. In a recrystallization procedure, an impure (crude) solid is dissolved
in a hot solvent. As this solution is cooled, the pure product crystallizes out and the impurities stay
dissolved.
Magnetic separation; uses a magnet to pull out magnetic particles (such as iron filings);
Apparatus: filter paper,50 mL & 100 mL beaker ,funnel ,Ring Stand with ring ,electronic balance,
stir rod, evaporating dish, measuring cylinder, Buchner funnel, hot plate
Chemicals: sand, salt, distilled water, NaCl, NaCl-SiO2 mixture, benzoic acid
Procedure 1.
Separation of Salt and Sand Using Filtration Process
1. Using electronic balance, measure out a .5.1 gram sample of salt. Enter into data table
2. Using an electronic balance, measure out 1-2 gram sample of sand. Enter into data table
3. Place salt and sand in a 100 mL beaker.
4. Using a 25 ml graduated cylinder measure out 20 mL of water. Add the 20 mL of water to the
mixture. Using a glass stir rod, mix the solution until all of the salt has dissolved.
Mass of sand
Mass of salt
Mass of filter paper

24
5. Set up filtration by using filter paper and a funnel. Place the 50 mL beaker onto the ring stand.
Separate the sand from the mixture. The sand will be left on the filter paper, leaving a mixture of
salt and water in your 50 mL beaker.
Procedure 2.
Extraction of NaCl.
Weigh clean and dry, evaporating dish. Add between 5 and 7 mL of distilled water to the NaCl-
SiO2 mixture and stir gently for 5 minutes. Carefully decant the liquid from the first evaporating
dish into the second evaporating dish, leaving the solid behind. It is not crucial that all the liquid
be transferred at this point. Add 5-7 mL more of distilled water to the first dish and gently stir
for 5 minutes. Decant the liquid into the second evaporating dish. Repeat this process a third
time with 5-7 mL more of distilled water. This process effectively extracts the NaCl (now in
evaporating dish #2) from the SiO2 (in evaporating dish #1). Both components, however, are
impure; the water mixed with each of them must be removed.
Procedure 3
Purification of Benzoic Acid by Recrystallization
Safety: Benzoic acid is a severe irritant and a sensitizer (exposure to sensitizers does not cause
cancer, but can make you more susceptible to those substances, which do cause cancer), and is
therefore classified as a harmful solid. You may wish to wear gloves while handling it. Be sure
to wash your gloves and hands after handling it.
Before you begin the re-crystallization of benzoic acid, you should have determined its
solubility. If you did not find this information for benzoic acid, its solubility is 0.34 g per 100
mL of cold water. Obtain about 1.0 g of "impure" benzoic acid (this sample of benzoic acid has
a small amount of sodium chloride added to it). What kind of container should you use for the
solid? (Guideline 1--however, we will be using only about 15-20 mL of solvent, so use a 50-mL
or a 125-mL beaker for this re-crystallization). Heat about 50 mL of DI water in a 150-mL
beaker. Add about 15 mL of the heated water to the "impure" benzoic acid (in your beaker), and
place the benzoic acid/hot water mixture on the hot plate. Add more hot water to the benzoic
acid, as needed, until the benzoic acid has completely dissolved (usually at boiling conditions).
If the solid does not fully dissolve within about 5 minutes, using the initial 15-mL sample of hot

25
water, add more hot water in 5-mL increments. Once the solid has completely dissolved, add an
additional 2-5 mL of hot water to keep it dissolved. Remove the container from the hot plate
(and turn off the hot plate).
Based on the solubility of benzoic acid in water, you can estimate your recovery. For example, if
0.34 g of benzoic acid dissolves in 100 mL of cold water, then if you started with 1.0 g of benzoic
acid, the maximum you could recover by crystallization would be about 0.66 g if you used 100 mL
of water. If you used 50 mL of water, then only about 0.17 g would stay dissolved, and you would
recover a maximum of about 0.83 g of benzoic acid. How much water did you actually use? How
much benzoic acid should you recover?
Let the benzoic acid solution cool by placing on the bench top. After the mixture, with some crystals
present, has cooled to room temperature place the beaker in an ice bath to enhance crystallization
and crystal recovery, since most chemicals are less soluble at cooler temperatures. Never place the
beaker directly in an ice bath from the hot plate. Let crystals for normally by sitting on the bench. If
you cool the supersaturated mixture too soon, before you allow it to cool to room temperature, you
may actually trap impurities in the solid material. Letting nature for crystals naturally is much for
efficient and practical. Using vacuum filtration you should collect your crystals. Use a small
Buchner funnel placed on top of a 250-mL vacuum filter flask. The vacuum assembly consists of
your vacuum flask with Buchner funnel connected to a vacuum trap which is inserted into a
vacuum-trap-bottle which is then connected to the vacuum line. After pouring your crystalline
mixture into the Buchner funnel, wash your beaker with DI water and collect this additional
crystalline material in your funnel. Wash the solid material with a little DI water to removed filtrate
material and any soluble impurities. Let the vacuum run for an addition 5 minutes or so before
turning off the vacuum and collecting your crystals.
You will store your crystals until the next lab period in one of the drying ovens. Be certain that you
label an evaporating dish or small beaker with the required identifying information prior to drying.
During the next lab period, you will recover your material from the drying oven, weigh it and
determine the melting point of both the "impure" benzoic acid and your re-crystallized benzoic acid.

26
Procedure 4
Physical separation of Iron from Mixture of Iron and sulphur powder by using Magnet
1. Prepare a mixture containing iron powder and sulfur powder in the ratio 7:4 by mass. Do this by
weighing out 7 g of iron powder and 4 g of finely powdered sulfur onto separate pieces of filter
paper (or use weighing boats). Mix the two powders by pouring repeatedly from one piece of paper
to the other until a homogeneous mixture (by appearance) is obtained.
2. Note the appearance of the pure elements and the mixture. Demonstrate that iron can be separated
from the mixture by physical means. Do this by wrapping the end of a small bar magnet in a paper
tissue or cling film, and dipping it into a teaspoon sized heap of the mixture on a watch glass. The
iron will be attracted, but the sulfur remains on the watch glass.
3. Place about 2 g of the mixture into a borosilicate test tube.
4. Insert a plug of mineral wool (mineral fiber) into the mouth of the test tube. Clamp the test tube as
shown in the diagram.
5. Heat the powder mixture at the base of the test tube gently at first and then more strongly (use a blue
flame throughout). Heat until an orange glow is seen inside the test tube. Immediately stop heating.
Let the students see that the glow continues and moves steadily through the mixture.
6. Allow the test tube to cool down. At this point the students could carry out their own small scale
version of the reaction.
7. Once cool, it is possible to break open the test tube to show the appearance of the product, iron
(II) sulfide. The test tube can be broken open using a pestle and mortar. It is advisable to wear
protective gloves.

27
8. It may be possible to show that the product, iron (II) sulfide is non-magnetic. However, this is not
always successful. It has been suggested that using a very weak magnet is advisable

28
Experiment 10: Instrumental analysis; Colorimetric Determination of
Acetaminophen
Objective: To determine Acetaminophen content by using colorimetric method
Theory: When white light is passed through a colored substance, some of the light is absorbed. A
solution containing hydrated copper (II) ions, for example, looks pale blue because the solution
absorbs light from the red end of the spectrum. The remaining wavelengths in the light combine in
the eye and brain to give the appearance of cyan (pale blue). Some colorless substances also absorb
light, but in the ultra-violet region. Since we can't see UV light, we don't notice this absorption.
Different substances absorb different wavelengths of light. This unique property of all compounds
can be used to identify the substance. The presence of particular metal ions, for example, or of
particular functional groups in organic compounds determine the way particular compounds absorb
the different wavelengths of light, thus resulting in these compounds tending towards particular
colors. The amount of absorption also depends on the concentration of the substance if it is in
aqueous solution. Measurement of the amount of absorption can be used to find concentrations of
very dilute solutions. An absorption spectrometer measures the way that the light absorbed by a
compound varies across the UV and visible spectrum
In this practical, the concentration of the commonly used medicinal drug acetaminophen was
determined in a solution whereby it was reacted with iron (III) before the solution was analyzed
using UV-visible absorption spectrometry. In the reaction, acetaminophen reduces Fe3+
to iron (II) –
Fe2+
. Acetaminophen reacts with iron (III) in a stoichiometric ratio of 2:1 (i.e. two moles of Fe3+
oxidize one mole of acetaminophen.
This reaction can be used in order to analyses acetaminophen colorimetrically. The resulting Fe2+
ions react with the potassium hexacyanoferrate (III) solution to form the intense deep blue-colored
complex known as Prussian blue. By measuring the intensity of the blue color imparted to the
resulting solution due to the formation of this complex, the concentration of acetaminophen can be
precisely determined.

29
OH
HN CH3
O
Figure 10.1: Molecule of acetaminophen (also known as paracetimol or APAP)
Chemicals: acetaminophen tablet, deionized water, iron (III) chloride solution, potassium
hexacyanoferrate (III) solution, 5 moldm-3
hydrochloric acid
Apparatus: Uv-Visible spectrophotometer, Analytical balance, volumetric flask (6), beaker
Procedure
0.100 g of acetaminophen were accurately weighed in a beaker and dissolved in distilled water. The
newly-formed solution was then transferred quantitatively to a 1 dm3
volumetric flask and made up
to the measured mark with distilled water. The solution was then accurately diluted to make a 0.01
gdm-3
stock solution of acetaminophen by diluting 25 cm3
to 250 cm3
.
Once the stock solution had been prepared, a number of measured volumes of the 0.01 g dm-
3
acetaminophen stock solution along with distilled water were added to a series of 50 cm3
volumetric flasks labeled A-F as shown below.
Table 10.1: Measured volumes of in 50 cm3
volumetric flasks
Volumetric Flask Volume of acetaminophen stock
solution (cm3
)
Volume of distilled water
(cm3
)
A 10 0
B 8 2
C 6 4
D 4 6
E 2 8
F 1 9
2 cm3
of 0.02 moldm-3
iron (III) chloride solution and 4 cm3
of 0.002 moldm-3
potassium
hexacyanoferrate (III) solution were then added to each flask before being left for 10 minutes. After
10 minutes, 1 cm3
of 5 moldm-3
hydrochloric acid were added before making up to the mark with
distilled water. The absorbance of the solution was then measured after 20 minutes at 700 nm, using
distilled water to zero the instrument.

30
About 0.1 g of the powdered tablet was weighed accurately in a 250 cm3
beaker before being
dissolved in distilled water. This solution was transferred quantitatively into a 1 dm3
volumetric flask
before being made up to the mark with distilled water. 25 cm3
of this solution were then pipetted into
a 250 cm3
volumetric flask and made up to the volume with distilled water. 10 cm3
of this solution
was measured into a 50 cm3
volumetric flask before adding 2cm3
of 0.02 moldm3
aqueous iron (III)
chloride and 4 cm3
of 0.002 moldm-3
aqueous potassium hexacyanoferrate (III), leaving to stand for
10 minutes, adding 1 cm3
of 5 moldm-3
hydrochloric acid and making up to the mark with distilled
water. Once 20 minutes had passed, the absorbance was measured at 700 nm, once again using
distilled water in order to zero the instrument.
Precautions
 It was ensured that both solutions which were made up of acetaminophen and the crushed tablet
were shaken well and always transferred accurately and quantitatively.
 Both the acetaminophen and the crushed tablet solutions were thoroughly filtered before using in
order to remove any solid excipients from the solution.
 The curettes were handled from the opaque sides so as to avoid any undesired effect on the
samples from which the absorbance was achieved.
 All volumes were read at eye-level in order to maximize the avoidance of parallax errors.
 All waiting times specified in the procedure once the addition of iron (III) chloride, potassium
hexacyanoferrate (III) and hydrochloric acid reactants to the volumetric flasks had been made
were strictly adhered to in order to ensure the completion of the reactions taking place and more
importantly that the color of the solution obtained was in fact the final color, since this was the
only physical quantity which the experiment depended on.
The pharmaceutical tablet was crushed thoroughly in order to attain a homogenous powder. The
solution in which the powdered tablet was dissolved was also filtered in order to remove any solid
excipient from the table itself.
The cuvette used when obtaining the absorbance’s using the UV-visible absorption spectrometer
were handled from the opaque sides so as to avoid getting dirt on their clear sides which would have
affected the final results

31
Results, Data Analysis and Calculations
The colorimetric analysis of flasks A-F obtained by means of the experiment is shown below.
Flask Absorbance Reading
of acetaminophen
Absorbance Reading
of Crushed Tablet
Mean Absorbance
A
B
C
D
E
F

32
Experiment 11: Investigating the heat involved in a chemical reaction
(Calorimetry): Investigating endothermic reaction; Investigating exothermic
reaction; and Effect of temperature on reaction rate
Objective: To investigate the exothermic/endothermic nature of the process when ammonium
nitrate is dissolved in water, to investigate the exothermic/endothermic nature of the reaction
between sulphuric acid and sugar.
Theory
Exothermic Reaction
A chemical reaction that releases heat energy to the surroundings is known as an exothermic
reaction. During an exothermic process, heat is given out from the system to its surroundings and
this heat energy is written on the right side of the equation as shown below.
Reactants → Products + Heat
For example, the burning of carbon with oxygen produces carbon dioxide and heat is released
during the reaction. Thus, the reaction is exothermic and written as:
C + O2→CO2 + Heat
Endothermic Reaction
A chemical reaction which absorbs heat energy from the surroundings is known as an endothermic
reaction. During an endothermic process, heat flows into the system from its surroundings and the
heat is written on the left side of the equation.
Reactants + Heat → products
For example, the reaction between carbon and sulphur to form carbon disulphide is an endothermic
reaction because heat is absorbed in the reaction.
C + 2S + Heat → CS2 the amount of heat energy liberated or absorbed by a chemical reaction is
called heat of reaction or change in enthalpy for the reaction. It is symbolized as ∆H. Its unit is
expressed in kilojoules per mol(
𝐾𝐽
𝑚𝑜𝑙
). The change in enthalpy (∆H) is the difference between the
energy of the products and the energy of the reactants. ∆H = H p – Hr; where H p is the heat content
(energy) of the product, Hr is the heat content (energy) of the reactant

33
Apparatus: thermometer, cork, beaker, reagent bottle & stirrer.
Chemicals: ammonium nitrate, water, Concentrated H2SO4 and sugar.
Procedure1:
1. Take 100 mL of water in a beaker and record its temperature.
2. Dissolve 15 g of solid ammonium nitrate (NH4NO3) in the 100 mL of water.
3. Touch the outer surface of the beaker and record the temperature of the solution with the help of a
thermometer
Observations and analysis:
1. Does the beaker feel hot or cold when you touch it?
2. Is the temperature increased or decreased after the addition of NH4NO3?
3. What do you conclude from this experiment?
Procedure 2:
1. Take small amount of sugar in a beaker.
2. Add a little concentrated sulphuric acid to the sugar.
3. Touch the outer surface of the beaker and record your observation
Observations and analysis
1. Does the beaker feel hot or cold when you touch it?
2. Did you see any steam in the beaker?
3. What is the color of the product formed?
4. Write a balanced chemical equation.
5. What can you conclude from the experiment? [Caution-When mixing concentrated acid and
water, always adds the acid to the water; never add water to concentrated acid.]
Procedure 3:
Effect of Temperature on Reaction Rate
Objective: To study the effect of temperature on the rate of reaction between sodium thiosulphate
and hydrochloric acid.
Theory: Temperature usually has a major effect on the rate of reaction. Molecules at higher
temperatures have more thermal energy. Generally, an increase in the temperature of a reaction
mixture increases the rate of reaction of chemical reactions.
This is because as the temperature of the reaction mixture raises, the average kinetic energy of the
reacting particles increases. So, they collide more frequently and with greater energy.

34
The effect of temperature on rate of reaction can be experienced in our daily life. For example, a the
food is kept in refrigerator to slow down the rate of decomposition of food; and b during heart
surgery, the body of patient is cooled to slowdown the rates of biological reactions. Temperature not
only affects the rate of reaction but can even change the course of a reaction. For example,
At 200°C, NH4NO3(s) → N2O (g) + 2H2O (g)
Apparatus: 100 mL beakers, test tubes, thermometers, white paper, pencil.
Chemicals: 0.5 M dilute HCl solution, 0.1 M Na2S2O3 solution, pieces of ice.
Procedure 4:
1. Take 25 mL of 0.1 M Na2S2O3 solution in a test tube and 25 mL of 0.5 M HCl solution in
another test tube.
2. Prepare 3 such sets and maintain them at different temperatures.
Set (i) at 0°C [by keeping them in an ice bath].
Set (ii) at room temperature.
Set (iii) at 40 °C (by heating the two solutions in a water bath).
3. Put a cross signs on a white cardboard and place a clean dry 100 mL beaker above it.
4. Now, pour the contents of set (i) in the beaker and start a stopwatch immediately.
5. Carefully stir the mixture with thermometer and record the time taken for the cross to
disappear.
6. Repeat steps 3, 4, and 5 with set 2 and set 3 respectively.
7. Tabulate your results as temperature in °C versus time in minutes.
Observations and analysis:
a. What was the appearance of the mixture at the start of the reaction, and at the end of the
reaction? Explain the changes using the equation for the reaction.
b. Plot the graph of time (minutes) on the horizontal axis against rate on the vertical axis
c. Under which condition of temperature does the cross take?
i. the shortest time to disappear, and
ii. The longest time to disappear.
d. Draw a conclusion about the relationship between the average reaction rate and
temperature.

35
REFERENCES
1. Silberberg, M. Principles of General Chemistry: Student Solutions Manual. Publisher: MGH,
2006
2. Jo A.Beran. Laboratory Manual for Principles of General Chemistry, Edition [8 ed.], Publisher:
Wiley, 2007.
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