Download CHE 1400 Lab Manual - Al Akhawayn University

CHE 1401
School of Science & Engineering
LABORATORY MANUAL FOR
CHEMISTRY AND THE ENVIRONMENT
Last Update: 7 July 2015
Last update: June 2011
1
CHE 1400
Name: ________________________
Section: ________________________
LABORATORY MANUAL FOR
CHEMISTRY AND THE ENVIRONMENT
Last Update: 7 July 2015
Last update: June 2011
1
CHE 1400
TABLE OF CONTENTS
Laboratory introduction
Common laboratory equipment
iii
viii
Experiment 1:
Physical Properties of a Compound
1
Experiment 2:
Atoms and Light
7
Experiment 3:
Solids in Cigarette Smoke
11
Experiment 4:
Chemical Reactions
17
Experiment 5:
Solubility
23
Experiment 6:
Acids, Bases, Buffers and pH
26
Experiment 7:
Acid Neutralization by Antacid
38
Experiment 8:
Oxidation and Reduction
42
Experiment 9:
Clarification of Water
46
Experiment 10:
Saponification
51
Experiment 11:
Water Analysis (field trip)
56
Appendix I
Appendix II
62
64
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Laboratory introduction
Laboratory safety, chemicals, equipment and techniques
Laboratory safety
Common sense precaution and a proper understanding of the techniques and
chemicals being used make a chemical laboratory no more dangerous than parts of the
home, such as the kitchen or garage.
1. Wear safety glasses. Safety glasses must be worn at all times when working in the
laboratory. It’s particularly important to avoid the use of contact lenses since these
lenses can result in more serious injuries if a toxic chemical is splashed into the
eye.
2. Wear suitable clothing. Clothing acts as a protection against spilled chemicals or a
burning liquid. Clothing which exposes large areas of bare skin can be a major
hazard; open toed shoes or sandals increase the possibility of injury to the foot. A
lab coat serves to clothes and skin, and should be worn in the laboratory at all
times. Hair extending below the shoulder blades should also be secured to avoid
damage to the hair itself, as well as to reduce the possibility of spilling a chemical.
3. Watch out for broken and hot glassware. Cuts and burns are the two most
common types of laboratory injury. Do not use any glassware that is broken,
chipped or badly cracked. If you have any questions on using a piece of glassware,
check with the instructor. Hot glass looks no different than a piece of cold glass.
Most burns result from trying to handle a piece of hot glass or other hot material.
Be aware of any glassware or other items that have been heated on a hot plate or
in a Bunsen burner flame.
4. Locate and be familiar with safety equipment. Know where to find and how to use
safety and first aid equipment, such as fire extinguishers, eye wash stations, safety
showers, etc. Remember that even if you know the location of safety equipment, it
is of very little use unless you know how and when to use it.
5. Be careful with the chemicals. Consider all chemicals as potentially dangerous
until you know otherwise. Avoid rubbing your eyes if there is any possibility that
your hands may be contaminated with a chemical. Never eat or drink anything
while working in the laboratory.
6. Read all labels. Read the label on any reagent bottle before you use it. Be sure to
read the complete name of the chemical since many chemical names sound and
look very similar. Also be sure to read the concentration of the reagent since a 5.0
M solution of sulfuric acid is very different than of a 0.005 M solution of the same
acid.
7. Dispose of all chemicals properly. Discard any chemical in the proper container.
Mixing of certain chemicals can start a fire or produce other hazards. Always
dispose of unused chemicals; never return a chemical to the original container.
8. Wash any chemical off of your skin. If corrosive chemicals get on your skin or in
your eyes, flush the affected area immediately with a large volume of clean water.
Be sure you notify the laboratory instructor as soon as possible.
9. Smell any chemical cautiously. Never taste a chemical. If you need to smell a
chemical do not smell the source directly. Instead use a cupped hand to bring a
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small amount of the vapor to your nose.
10. Use the fume hood when appropriate. Any chemical, which is irritant, dangerous
or which has an unpleasant odor, should always be used in a properly operating
fume hood.
11. Perform only authorized experiments. Perform only the experiments noted in the
laboratory manual or those given to you by your laboratory instructor.
12. Add concentrated acids and bases to water. Always use the dilute acids and bases
provided to you. If you are instructed to dilute a strong acid or base, pour the acid
or base into the water. Remember that diluting many acids or bases can generate
large amount of heat.
13. Clean up your workstation. Be sure to empty all reagents into the appropriate
waste container, wash out any used glassware and wipe off the bench top. Be sure
the gas, water and all electrical equipment are turned off.
14. In case of accident. Notify your instructor immediately if either you or another
student had any type of accident.
15. Proper use of equipment. Misuse can lead to injury. Please do not hesitate to ask
questions if you are unsure about the use of equipment.
Chemicals
Laboratory reagents shared by several students are stored in the fume hoods. Special
reagents and unknowns will be issued by the laboratory instructor.
1. Aqueous solutions (water based) can, in most cases, be poured down the sink
followed by a large volume of water.
2. Concentrated acids, concentrated bases, and organic solvents (halogenated and
non- halogenated) should not be poured down the sink. Acids and bases should be
first neutralized, then disposed of in the sink with a large volume of water. Water
soluble organic acids should be diluted with water, then poured down the sink.
Water-insoluble organic acids should be disposed in the containers provided.
3. Dispose of solids in the labeled waste containers: never dispose solids in the sink.
4. Do not place stoppers from the bottles on the bench top. Stoppers should be held
in the hand until they can be returned to the original container.
5. Never return unused reagents to a reagent container. Dispose of the reagent in an
appropriate manner.
6. Return reagent bottles to the work area shelf (reagents should never be taken out
of the fume hoods). Do not take the reagent bottles to your work bench area where
they will be unavailable for the other students.
7. Use only the amounts (approximately) of the reagent called for in the laboratory
manual, or indicated by the laboratory instructor.
Equipment
Each student will have the use of a set equipment at an assigned work place. It should
be remembered that each set of equipment is used by other students during other
laboratory sessions.
1. Keep both the work area and the laboratory equipment clean and in good working
condition. Do not leave dirty equipment laying around after the laboratory session
2. Do not borrow equipment from other desks. If you need extra equipment or are
having problems with broken or malfunctioning equipment, obtain the additional
equipment or replacement equipment from the stock room.
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Laboratory techniques
The laboratory principles that you learn in lecture and from the textbook are the result
of numerous laboratory experiments. Years of observations, experimentation,
interpretation and predication are necessary before a principle becomes well enough
established to be included in an introductory chemistry course.
Most laboratory work is made up of the same basic techniques. In this laboratory
session you will study the basic rules, equipment and techniques used in a chemical
laboratory. These include:
- handling and transferring solid and liquid chemicals
- using graduated cylinders, pipettes, burettes and volumetric flasks
- using a laboratory balance
- evaporating liquids
- titrating a solution
- separating a liquid and a solid
Laboratory equipment
From the list below and the pictures given on the following page, identify the
common laboratory equipment at each work station and give a brief indication of
what each piece of equipment is used for. Some of the equipment shown in the picture
may not be at the work station.
Number
Item
Use
Common laboratory equipment
1-2
graduated cylinder
________________________________________________
3
beakers
________________________________________________
4
stirring rods
________________________________________________
5
wash bottle
________________________________________________
6
funnel
________________________________________________
7-8
Erlenmeyer flasks
________________________________________________
9-11
test tubes
________________________________________________
12-13
test tubes racks
________________________________________________
14
glass plate
________________________________________________
15
wire gauze
________________________________________________
16
crucible tongs
________________________________________________
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Number
Item
Use
17
spatulas
________________________________________________
18
litmus paper
________________________________________________
19
watch glasses
________________________________________________
20
evaporating dish
________________________________________________
21
dropping pipettes
________________________________________________
22
test tube holder
________________________________________________
23-24
test tube brushes
________________________________________________
Special laboratory equipment
1
reagent bottle
________________________________________________
2
condenser
________________________________________________
3
500-mL Erlenmeyer
________________________________________________
4
1000-mL beaker
________________________________________________
5
petri dish
________________________________________________
6
Buchner funnel
________________________________________________
7
Buchner flask
________________________________________________
8
volumetric flask
________________________________________________
9
500-mL Florence flask
________________________________________________
o
10
110 C thermometer
________________________________________________
11
100-mL graduated cylinder
________________________________________________
12
50-mL buret
________________________________________________
13
glass tubing
________________________________________________
14
U-tube
________________________________________________
15
porous cup
________________________________________________
16
crucible & cover
________________________________________________
17
mortar & pestle
________________________________________________
18
glass bottle
________________________________________________
19
pipettes
________________________________________________
20
ring stand
________________________________________________
21
buret clamp
________________________________________________
22
double buret clamp
________________________________________________
Number
23
Item
Bunsen burner
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24
burette brush
________________________________________________
25
clay triangle
________________________________________________
26
rubber stoppers
________________________________________________
27
wire loop
________________________________________________
28
pneumatic trough
________________________________________________
29
rubber pipettes bulb
________________________________________________
30
iron support ring
________________________________________________
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COMMON LABORATORY EQUIPMENT
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EXPERIMENT 1
Physical properties of a compound
OBJECTIVES
1. To identify a compound based on its physical properties
2. To learn how to properly assemble and use a simple boiling point apparatus
Relates to chapter 1 of “Chemistry for changing times, 13th Ed.”.
INTRODUCTION
Gold is a yellow solid, salt is a white crystal, lead is heavy, chlorine is a greenish-yellow
gas and water is a clear colorless liquid. These are all physical properties of chemical
substances. Physical properties can be used to identify a chemical substance. The
common physical properties include color, odor, density, solubility, crystal structure,
melting point and boiling point. Additional physical properties such as conductivity,
malleability, etc. can also be determined.
1. Solubility
The solubility of a substance is most accurately defined as the maximum mass (expressed
in grams) of the test substance that dissolves in a known mass (usually 100 grams) of
another substance at a given temperature. The test substance is referred to as the solute
and the substance in which the test substance is being dissolved is the solvent. A given
chemical has different solubilities in different solvents, depending on the similarities in
molecular composition of the two substances. Usually, “like dissolves like“ and therefore
highly polar compounds like sodium chloride (table salts) are soluble in high polar
solvents like water, but not in low polar solvents like oils. Similarly, low polar substances
like butter dissolve in low polar solvents (oils) but not in high polar ones (water).
The degree of solubility can also be expressed in a rough way by the terms « soluble »,
« slightly soluble » or « insoluble ». A soluble solid dissolves quickly and easily in the
solvent soluble liquid mixes quickly and easily with the solvent. An insoluble substance
will not dissolve or mix with the solvent at all. A slightly soluble substance will dissolve
or mix to a limited degree.
2. Density
Density is the mass per unit volume. A substance with a high density has a large mass in
a small volume. While we often say that lead is « heavy », what we really mean is that
lead has a high density. Density is commonly given in terms of grams per milliliter
(g/mL), although other units of mass and volume can be used.
3. Boiling point
When a liquid is gradually heated, there is a point at which the temperature of the liquid
no longer increases, but bubbles of vapor (gas) form spontaneously, and continue to do so
until the entire volume of liquid has been converted to a gas. This constant temperature is
called the boiling point of the liquid. The boiling is dependent on both the liquid being
heated and the atmospheric pressure on the liquid. Boiling points are normally given at
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normal atmospheric pressure.
EXPERIMENTAL PROCEDURE
Note down the number of the experimental unknowns on your laboratory report sheet.
1. Solubility
a. Solubility in water
Into a 25-mm test tube add about 1 mL of water. Add about 1 mL of the liquid unknown.
Shake the test tube several times and note if there is any change in the appearance of the
mixture. If the two liquids mix (even if shaking is required) to give a single mixture
without a meniscus, then that unknown is soluble in water.
If the unknown liquid does not mix with water and you observe two distinct layers with a
meniscus between them (even after shaking), then that unknown is insoluble in water.
b. Solubility in ethanol
Into a 25 mm test tube add about 1 mL of ethanol. Add about 1 mL of the liquid
unknown. Shake the test tube several times and note if there is any change in the
appearance of the mixture. If the two liquids mix (even if shaking is required) to give a
single mixture without a meniscus, then that unknown is soluble in ethanol.
The meniscus is the slightly curved surface at the top of a column of liquid.
If the unknown liquid does not mix with ethanol and you observe two distinct layers with
a meniscus between them (even after shaking), then that unknown is insoluble in ethanol.
c. Solubility in cyclohexane
Into a 25-mm test tube add about 1 mL of ethanol. Add about 1 mL of the liquid
unknown. Shake the test tube several times and note if there is any change in the
appearance of the mixture. If the two liquids mix (even if shaking is required) to give a
single mixture without meniscus, then that unknown is soluble in cyclohexane.
If the unknown liquid does not mix with ethanol and you observe two distinct layers with
a meniscus between them (even after shaking), then that unknown is insoluble in
cyclohexane
Record your results for the solubility of the unknown in water, ethanol and cyclohexane
on the laboratory report worksheet.
2. Density
Place a 25-mm test tube into a 100 mL beaker and weigh the assembly to the nearest
0.1 g on a top loading balance (be sure that the balance is zeroed).
Record the weight of the beaker/test tube assembly on the laboratory report worksheet.
Pipette 1 mL of the unknown into the 25-mL test tube and reweigh.
Use a rubber pipetting bulb to draw the liquid into the pipet. DO NOT pipet any liquid by
mouth!
Make sure that you pipet the correct volume of liquid.
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If you don’t know how to use a pipette, please check with the laboratory instructor.
Record the weight of the beaker/test tube assembly containing the unknown on the
laboratory report worksheet.
Be sure that you use the same balance each time, otherwise you may get slightly different
values due to differences in the balances.
Calculate the mass of the unknown by subtracting the mass of the [beaker/test tube
assembly] from that of the [beaker/test tube assembly plus unknown].
If you wish, you may tare the balance with the test tube/beaker assembly and then get the
mass of the unknown directly.
If you use this method, be sure to get the mass of the unknown immediately after you
have tared the balance. Otherwise, another student may tare the balance to their
glassware.
Calculate the densities of the unknown by dividing the mass of the unknown by the
volume of that unknown.
Complete a second determination of the density and use the average of your two
experimental results.
3. Boiling point
3.1. Assemble the boiling point apparatus
Assemble the boiling point apparatus as shown in Figure 3.1. Place 1 to 2 mL of your
unknown liquid into a 75-mm test tube. Next to the thermometer bulb secure a 10 cm
capillary tube in an inverted position (open end down) using a rubber band. Place the
thermometer bulb and capillary tube into the liquid. Place the apparatus into a water bath.
Figure 3.1 Apparatus for determining the boiling point of a liquid.
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3.2. Measure the boiling point
Make sure that the water bath contains 5-8 « boiling chips ». Slowly heat the water in the
water bath. You should notice a bubble of gas escaping from the end of the tube every
once in a while. As the temperature of the water increases, the rate at which the bubbles
are formed should increase. When a rapid and continuous stream of bubbles escapes from
the capillary tube, discontinue heating by either shutting off the hot plate, removing the
test tube from the water bath, or removing the water bath and test tube from the hot plate.
Caution! The water and the glassware may be quite hot.
As soon as the bubbles from the capillary tube STOP, take a note of the temperature.
Record the temperature of the water bath as the boiling point of the liquid unknown.
To get as accurate a reading as possible, watch the capillary tube closely and record the
temperature immediately.
3.3. Repeat the measurement
Determine the boiling point of the unknown a second time. The capillary tube and all
liquid must be removed. Reinsert the capillary tube before heating is resumed. Also, the
water bath should be cooled by removing some of the hot water and adding some cool
water along with some fresh boiling stones.
4. Identification of the unknown compound
Based on your observation, experimental results, calculations and the information given
in Table 1, determine the identity of the liquid unknown. If you cannot make a
reasonable determination based on all of your data, make a determination on two of the
three factors (solubility, density and boiling point) and repeat the experiment for the third
factor. For example, if you observe that your unknown is soluble in all three solvents and
has a boiling point of 54 °C (indicating it is acetone) but the density was calculated as
0.95 g/mL (instead of 0.79 g/mL), then carefully repeat the density experiment to see if
you get the same value.
Explain any problems you had in either the procedures or in the identification of the
unknown.
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Experiment 2
Physical properties of a compound
Name(s)
Date
Laboratory Instructor
Unknown n°_
REPORT SHEET
1-Solubility
Solvent
Soluble
Insoluble
Water
Ethanol
Cyclohexane
2-Density
Trial 1
Trial 2
Mass of test tube and beaker (g)
___________
___________
Mass of test tube, beaker & unknown (g)
___________
___________
Mass of unknown liquid (g)
___________
___________
Volume of unknown liquid (mL)
___________
___________
Density of unknown liquid (g/mL)
___________
___________
Average density of unknown liquid
___________ g/mL
3-Boiling point
Observed b.p. 1: ______ °C
Observed b.p. 2: _____ °C
Average b.p. : ________ °C
4.Identification of the unknown compound
Unknown compound: ________________________________________
Explanations and comments:
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TABLE 1- Physical Properties of Some Common Laboratory Chemicals
Compound
Density
Melting point
Boiling point
(g/cm3)
(oC)
(oC)
Water
Ethanol
Acetone
0.79
-95
56
s
s
Acetamide
1.00
82.3
221
s
s
-
-
Acetanilide
1.22
114
304
-
s
s
s
Anthracene
1.28
216
-
-
-
s
s
Benzamide
1.08
132
290
s
s
-
s
Benzoic acid
1.07
122
249
-
s
s
s
Benzoin
1.31
137
344
-
s
s
-
2-Butanone
0.81
-86
80
s
s
s
s
Cyclohexane
0.79
6.5
81
-
s
s
Cyclohexene
0.81
-104
83
-
s
s
s
Ethanol
0.79
-117
79
s
s
s
Ethyl acetate
0.90
-84
77
s
s
s
s
Heptane
0.68
-91
98
-
s
s
s
n-Hexane
0.66
-95
69
-
s
-
-
Methanol
0.79
-94
65
s
s
s
s
Naphthalene
0.96
80.5
218
-
s
s
s
1-Propanol
0.80
-127
97
s
s
s
s
2-Propanol
0.79
-90
82
s
s
s
s
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Solubility
Acetone
Cyclohexane
s
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EXPERIMENT 2
ATOMS AND LIGHT
Light is not just light
OBJECTIVES
1. To observe the colors of light emitted from excited atoms and relate these wavelengths
to the energy changes in the electrons within the atom.
2. To learn to use a diffraction grating to separate the emission spectrum into separate
wavelengths of light.
Relates to chapter 3 of “Chemistry for changing times, 13th Ed.”.
BACKGROUND
Atoms are made up of positive nuclei surrounded by negative electrons. These electrons
have different amounts of energy as a rule, but the important fact is that electrons cannot
possess just any amount of energy but only certain amounts of energy. How do we know
that? This knowledge comes from a study of the electromagnetic spectrum in the x-ray,
ultraviolet, and visible portions of the spectrum. In today’s experiment, a detailed
observation of the light coming from hot atoms, will lead you to the same conclusion.
When an atom is heated, electrons absorb energy in definite amounts and as they cool,
they emit that extra energy which we see as a particular color of light. Electrons in
different kind of atoms absorb and then emit different amounts of energy, which create
different spectra. The spectra or combination of colors, observed from a discharge tube,
can be used to identify what element is glowing. This was the way helium was discovered
in the sun before it was known on earth. Without a diffraction grating the one most
prominent color is seen. Atoms with many electrons, like tungsten, give off so many
colors of light we observe a complete spectrum, i.e., white light. Through a diffraction
grating, we would observe a rainbow of the complete spectrum from tungsten.
A diffraction grating is like a prism in that it will separate the different wavelengths of
light. It is interesting that, when electrons behave as waves as well as particles (just as
light does), properties of waves lead to the same those electrons around nuclei can have
only certain energies.
Some instruments, such as the atomic absorption spectrometer make use of the fact that
every element has a particular set of energies, which are absorbed when its electrons are
excited. These particular energies, or wavelengths, of light can be used to detect the
presence of a particular element.
Other instruments identify elements by the wavelengths of light, which are re-emitted
when the atoms are excited. When a solution containing the atoms of an element is heated
in a flame, we can use these energy signatures to detect which elements are present in the
solution.
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WASTE AND THE ENVIRONMENT
Only concentrated hydrochloric acid poses a hazard. Concentrated acid can damage
plumbing if not neutralized or diluted. Once used in the flame test, it should be flushed
down the drain with lots of water.
PROCEDURE
NOTE: the room needs to be darkened for best results in this experiment
1. Hold a diffraction grating up to your eye and view the candle light through it. Record
your observations.
2. Using the diffraction grating, view each of the following light sources when
demonstrated by the instructor:
a- candle
b- helium light bulb
c- neon light bulb
d- argon light bulb
e- krypton light bulb
Record your observation(s) for each light source.
3. Without the grating, observe the colors as you place salts of the following metals into
the flame of a Bunsen burner. Make sure to wash the spatula with concentrated
hydrochloric acid between each salt. Record these observations.
a- Lithium
b- Sodium
c- Barium
d- Strontium
e- Potassium
f- Copper
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Experiment 2
Atoms and light
Name(s)
Date
Laboratory Instructor
Unknown n°_
REPORT SHEET
I – Describe the view of each of the following through a diffraction grating:
a- Candle
________________
b- Argon
________________
c- Krypton
________________
d- Neon
________________
e- Helium
________________
II – Describe the order of salts of these metals in a flame:
a- Lithium
________________
b- Sodium
________________
c- Barium
________________
d- Strontium ________________
e- Potassium ________________
f- Copper
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III – Questions
1- Could the colors of the flame be used to identify different metal ions in a mixture?
2- Why is a yellow shirt yellow?
3- What can we learn from a study of light coming from different stars?
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EXPERIMENT 3
SOLIDS IN CIGARETTE SMOKE
Smoke gets in your eyes……….. and that’s not all!
OBJECTIVES
1. To use the process of mass difference to make quantitative measurements of the solids
produced by a cigarette.
2. To determine whether the smoker or a person exposed to second-hand smoke captures
the largest amount of smoke solid.
3. To make a quantitative determination of the segment of a cigarette which delivers the
smallest mass of solid particles per mass of cigarette consumed.
Relates to chapter 13 and 19 of “Chemistry for changing times, 13th Ed.”.
BACKGROUND
Cigarette smoke is generated by burning tobacco leaves. The smoke produced in this
combustion reaction has at least two major components, which are harmful to the human
body.
Carbon monoxide (CO), one of the major components of cigarette smoke, is an odorless,
invisible, but toxic gas. Carbon monoxide attaches to hemoglobin (making carboxyhemoglobin) and prevents hemoglobin (as oxy-hemoglobin) from carrying oxygen to the
rest of the body. The affinity of hemoglobin for carbon monoxide is 200 times greater
than for oxygen. A cigarette smoker often has a carboxyhemoglobin (COHb) level of 5%
instead of the normal 0.5%. An increased level of COHb can cause increased stress on
the heart, impaired time discrimination, and high blood pressure.
Another harmful component in cigarette smoke is small solid particles. These particles
may be fly ash, organic tar, or mineral dust. Some of these solids have been found to be
carcinogenic or cancer causing. The size of these particles is a few microns (10-6 m) or
less. When solid particles enter the lung, they can irritate and destroy alveoli (tiny sacs
where gases are exchanged) and cause emphysema.
The Surgeon General of the United States has published research indicating smoking is
linked to lung cancer. A smoker may have 10 times the chance of lung cancer as that of a
non-smoker. There is also research indicating that cigarette smoke not only causes
damage in the lungs but also inhibits the reactions of the body to repair itself. An
increased chance of heart trouble has also been linked to smoking.
Not all of the smoke from a cigarette goes through the cigarette to the smoker, and a large
part of the components of the smoke is exhaled by the smoker into the air. This causes an
atmosphere of smoke for other people to breathe. Does this cause illness in non-smokers
exposed to smoke? This is a hotly debated question and the research that has been done is
not conclusive. However, a study in Japan indicated that lung cancer in women increases
as the amount of smoke in the home increases. Another study concluded that chronic
exposure to tobacco smoke at work is deleterious to the non-smoker and significantly
reduces functioning of small airways. The most frightening research is on the effect of
second-hand smoke on children. Babies of women who smoke during pregnancy tend to
weigh less and develop more slowly than those of non-smoking mothers. The children of
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parents who smoke have been found to be more prone to bronchial ailments.
WASTE AND THE ENVIRONMENT
The damage to the environment that occurs during this experiment is to the air.
PROCEDURE
PART I: DETERMINING SOLIDS IN CIGARETTE SMOKES
1. Each pair of students will use the pre-assembled setup, shown in Figure 4.1, aimed at
simulating the consumption of a cigarette and determining the quantities of solids
released within the smokes.
This setup comprises:
- a piston-powered vacuum pump (P = 0.85 atm; air flow = 38 L/min)
- vacuum rubber tubing
- one T hose connector
- 3 filter flasks named “Cigarette flask”, “Funnel flask” and “Pump guard”
- 2 pieces of 6 μm filter paper (Whatman grade 3)
- one piece of 2.5 μm filter paper (Whatman grade 5)
- 3 one-hole cork stoppers
- 2 funnels
- one stand
- 3 clamps
- one cigarette.
Note: The purpose of the filter flask named “pump guard” is to protect the
(expensive) vacuum pump by preventing any residual fine solid from
getting into it.
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Figure 3.1: Setup used to simulate the consumption of a cigarette.
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2. Remove the two pieces of filter paper from the apparatus (one from the cigarette
flask, the other from the funnel flask) and weigh them.
3. Replace the two weighed pieces of filter paper in the apparatus.
4. Weigh a non-filter tip cigarette.
5. One pair of students inserts the cigarette into the funnel connected to the cigarette
flask, as shown in Figure 3.1.
6. The other pair of students simply connects the second funnel into the funnel flask.
This will serve to collect room smoke.
7. Clamp the cigarette funnel pointing up and the second-hand smoke funnel about a
half-inch above the cigarette funnel. The funnels are to be far enough apart to allow
air in but close enough together to catch all the smoke.
8. Make sure the pressure release valve of the pump is fully opened. Then, turn on the
vacuum pump and light the cigarette.
9. If needed, adjust the pressure release valve so that the cigarette burns slowly, over a
period of about one minute.
10. When the cigarette has burned down to the last centimeter, turn off the vacuum pump
and put out the cigarette.
11. Remove the filter paper from both suction flasks and weigh each paper.
12. Weigh the unconsumed part of the cigarette. Do not weigh the ash. Subtract this
weight from the initial weight of the cigarette.
13. Record all weights on the report page.
14. Calculate the number of milligrams of solids collected per gram of cigarette
consumed for both smoker and second-hand smoker.
15. Repeat steps 2-14 for a filter-tip cigarette using new pieces of filter paper.
16. Subtract the initial weight of the filter paper from the final weight to obtain the
weight of solids collected.
For each cigarette complete steps 16-18.
17. Multiply the mass of the solids collected by 1000 and divide by the weight of
cigarette consumed to obtain milligrams of solids per gram consumed.
masssolid
collected
masscigarette
( g )  1000
consumed
(g)

________________ mg solid collected
________________ g cigarette
18. Divide the solids per cigarette of the cigarette flask (smoker smoke) by the solids per
cigarette of the funnel flask (room smoke) to get the amount of solids in smoker
smoke compared to solids in room smoke.
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CHE 1400
PART II: SOLIDS FROM DIFFERENT PARTS OF A CIGARETTE
1. Using a pencil or marking pen, mark off the length of the tobacco part of a filter
cigarette in thirds.
2. Weigh two pieces of filter paper. Place one piece in each flask.
3. Place the marked filter-tip cigarette into the funnel of the cigarette flask. Turn on the
vacuum pump. Light the cigarette and as quickly as possible position the funnel over
the cigarette.
4. Allow the cigarette to burn until one-third of the tobacco portion has been consumed.
Extinguish the cigarette with a drop of water from an eyedropper as you turn off the
pump.
5. Weigh both pieces of filter paper and record those weights on the report page.
6. Replace the 2 pieces of filter paper, turn on the pump, and relight the cigarette. Allow
the second third of the cigarette to burn. When the second third is consumed, turn off
the pump and extinguish the cigarette. Weigh the filter paper and record the weight
on the report page.
7. Repeat step 6 for the last third of the cigarette.
8. Subtract the initial filter paper weight from the weight after one-third was consumed
to find the number of grams of solids in the first third of the cigarette.
9. Subtract the weight after one-third consumed from the weight after two-third was
consumed to find the grams of solids in the second third of the cigarette.
10. Subtract the weight after two-thirds consumed from the weight after last third was
consumed to find the grams of solids from the last third.
11. Divide the grams of solids from the smoker smoke by the grams of solids from the
room smoke for each third of the cigarette.
Throw all solid residues in the trash. Be sure the cigarettes are totally extinguished.
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CHE 1400
Experiment 3
Solids in smoke
Name(s)
Date
Laboratory Instructor
REPORT SHEET
PART I - SOLIDS IN CIGARETTE SMOKE
filter cigarette
non-filter cigarette
Cigarette
flask
Funnel
flask
Cigarette
flask
Funnel
flask
Initial filter paper weight (g)
__________
__________
__________
__________
Final filter paper weight (g)
__________
__________
__________
__________
Weight of solids collected (g)
__________
__________
__________
__________
[=smoker solids]
[=room solids]
[=smoker solids]
[=room solids]
Initial cigarette weight (g)
_________
__________
Final cigarette weight (g)
_________
_________
Weight of cigarette consumed (g)
_________
_________
masssolid
collected
masscigarette
( g )  1000
consumed
(g)

______ mg/g
smoker solids

room solids
______ mg/g
_________
______ mg/g
______ mg/g
_________
1. Which of the cigarettes in part I produced the most solids in smoker smoke and in
room smoke? Can you explain the reason for your results?
2. Which cigarette produced the lowest ratio
smoker solids
?
room solids
3. What conclusions can be drawn from the data about being in a room with a smoker?
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CHE 1400
PART II: EFFECTIVENESS OF A CIGARETTE FILTER
filter cigarette
Cigarette flask
Funnel flask
[= smoker smoke]
[= room smoke]
Initial filter paper weight (g)
___________
___________
Final filter paper weight after ⅓ of
cigarette consumed (g)
___________
___________
Final filter paper weight after ⅔ of
cigarette consumed (g)
___________
___________
Final filter paper weight after a last third of
cigarette consumed (g)
___________
___________
msolids from the first third (g)
___________
___________
msolids from the second third (g)
___________
___________
msolids from the last third (g)
___________
___________
smoker smoke
for the first third =
room smoke
___________
smoker smoke
for the second third =
room smoke
___________
smoker smoke
for the last third =
room smoke
___________
4. Which portion of the cigarette in part II yielded the most solids in smoker smoke?
Explain your result.
5. Does the portion of the cigarette affect the amount of solids in the room smoke?
If so, why?
6. Does the filter protect a smoker from solids? Does a filter protect someone nearby?
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CHE 1400
EXPERIMENT 4
CHEMICAL REACTIONS
OBJECTIVES
1. To learn to observe and record events of a laboratory experiment and to draw
conclusions from these observations.
2. To gain practice in writing and balancing chemical equations.
Relates to chapter 5 of “Chemistry for changing times, 13th Ed.”.
INTRODUCTION
We live in a world where chemical reactions are continually taking place around and
within us. Life itself is dependent on a very large number of highly complex chemical
reactions. These include photosynthesis, which is the production of sugars by plants, and
respiration, which is the burning of sugars by animals to produce energy. Chemical
reactions are used throughout many industries to produce such products as synthetic
fibers, drug, cosmetics, plastics, certain type of food detergents, fertilizers, metal alloys
and many others. Numerous chemical reactions occur as part of the natural world such as
the rusting of iron, tarnishing of silver, weathering of rocks, etc.
A chemical reaction is a process during which a chemical change occurs. In the course of
a chemical reaction, the starting material(s) undergo a chemical change, to form new
chemical entities (products) that have different chemical and physical properties.
Chemical reactions can be described by chemical equations. By convention, the reactants
are written to the left and the products to the right of an arrow, therefore indicating the
direction of the reaction. As an example, the reaction between the element sodium, Na,
and chlorine, Cl2, (reactants) to form sodium chloride (product) can be written as:
Na + Cl2
NaCl
(not balanced)
This equation is not balanced, because there are 2 chlorine atoms on the left side of the
equation, but only 1 on the right side. This indicates that an atom of chlorine has been
destroyed, which is a direct violation of the law of conservation of matter. According to
this law, atoms can neither be created nor destroyed in a chemical reaction. If there are 2
chlorine atoms at the start of the reaction, then there must be 2 at the completion of the
reaction.
We cannot balance a chemical equation by changing the formulas of either the reactants
or the products. This would be a violation of the law of definite proportions, which states
that compounds are always found with the same proportion or ratio of each element. In
the above example sodium chloride is always made up of one atom of sodium and one
atom of chlorine.
If we wrote the formula for sodium chloride as NaCl2, this would indicate that sodium
chloride was composed of 1 atom of sodium and 2 atoms of chlorine.
We can, however, balance a chemical equation by altering the number of formula units in
the equation. Each formula unit of sodium chloride contains 1 atom of chlorine. Since
there are 2 atoms of chlorine in the reactants, this would mean that 2 formula units of
sodium chloride are produced:
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CHE 1400
Na + Cl2
2 NaCl
(not balanced)
Now, however, there are 2 atoms of sodium on the right, and only one on the left. This
would mean that 2 atoms of sodium are needed to carry out this reaction.
2 Na + Cl2
2 NaCl
(balanced)
As written, note that there are 2 atoms of sodium and 2 atoms of chlorine on both sides of
the equation.
Chemical equations sometimes indicate the physical state of the reactants and products.
The most common notation is (s) for a solid, (l) for liquid and (g) for gas. The above
equation could be written as:
2 Na (s) + Cl2 (g)
2 NaCl (s)
In this experiment, where the emphasis is on learning how to write and balance chemical
equations, these symbols will not be used.
Chemical changes are usually accompanied by an observable physical change. This
includes changes in color, the formation of an insoluble precipitate, the evolution of a
gas, or a change in temperature. Other, more subtle changes are possible.
The above reaction is an example of a synthetic reaction in which several reactants are
put together to form a new compound. The opposite situation is a decomposition reaction
in which a compound breaks down into elements or simpler compounds; a common
example of this type of reaction is the decomposition of potassium chlorate (KClO3) into
potassium chloride (KCl) and oxygen (O2)
2 KClO3 + heat
2 KCl + 3 O2
Another common type of reaction is the exchange reaction, where the cations and anions
of two compounds are « exchanged ». For example, sodium bromide (NaBr) and lead
nitrate (Pb(NO3)2) react to form sodium nitrate (NaNO3) and lead bromide (PbBr2)
2 NaBr + Pb(NO3)2
2 NaNO3 + PbBr2
In this reaction the PbBr2 product is a precipitate which can be noted either by
underlining, with the use of an arrow pointing downwards, or by indicating the product is
a solid.
2 NaBr + Pb(NO3)2
2 NaNO3 + PbBr2 (s)
How do we know that the observed precipitate is PbBr2 and not NaNO3? A simple test
would show that NaNO3 is very water-soluble and could not be a precipitate; which by
definition is NOT water-soluble. The only other possible precipitate is the PbBr2.
Although such tests are simple, it would require a considerable amount of time for you to
make all the tests necessary to identify the precipitates in this experiment. Therefore, the
two general solubility rules given below should be helpful.
All compounds containing the cations Na+, K+ or NH4+, are water-soluble and therefore
cannot be precipitates.
All compounds containing the anion NO3- are water-soluble and therefore cannot be
precipitates.
For sections 1 and 2 of this experiment, you will mix chemicals in aqueous solution
together, in small test tubes. In some cases a reaction will occur. This will be indicated by
a change in the appearance (including temperature) of the mixture. In other cases there
will be no evidence of a reaction. You need to make careful observations in each case,
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CHE 1400
and record your observations on the data sheet. In some cases you will find that more
than one reactant will react with a given reactant, in which cases you should identify the
common component among the reactants.
In sections 3 and 4 you will test certain chemicals with the indicators methyl orange and
phenolphthalein. Carefully note any color changes during this part of the experiment and
try to determine the chemical characteristics common to acids, and chemical
characteristics common to bases. Table 3.1 indicates the colors given by each indicator.
Table 3.1 Colors given by phenolphthalein and methyl orange
Indicator
Phenolphthalein
Methyl orange
Color when acidic
colorless
red
Color when basic
pink
yellow
EXPERIMENTAL PROCEDURE
1. Reaction with silver nitrate
Into eight (8) 25-mm test tubes add approximately 0.5 mL of each of the 8 solutions
listed below:
1) dilute hydrochloric acid HCl
2) barium nitrate Ba(NO3)2
3) sodium hydroxide NaOH
4) sodium chloride NaCl
5) hydrogen nitrate (nitric acid) HNO3
6) barium hydroxide Ba(OH)2
7) barium chloride BaCl2
8) potassium hydroxide KOH
Then add about 0.5 mL of silver nitrate, AgNO3, to each of the tubes. Carefully observe
to see if there is any precipitate.
Record your observations in the space provided on the report sheet. If no observable
reaction occurs, then report “NR” (No Reaction).
2. Reactions with sodium sulfate
Into eight (8) 25-mm test tubes add approximately 0.5 mL of each of the 8 solutions
listed below:
1) dilute hydrochloric acid HCl
2) barium nitrate Ba(NO3)2
3) sodium hydroxide NaOH
4) sodium chloride NaCl
5) hydrogen nitrate (nitric acid) HNO3
6) barium hydroxide Ba(OH)2
7) barium chloride BaCl2
8) potassium hydroxide KOH
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CHE 1400
Then add about 0.5 mL of sodium sulfate, Na2SO4, to each of the tubes. Carefully
observe to see if there is any precipitate.
Record your observations in the space provided on the report sheet. If no observable
reaction occurs, then report “NR” (No Reaction).
3. Reactions with Phenolphthalein Indicator
Into eight (8) 25-mm test tubes add approximately 0.5 mL of each of the 8 solutions
listed below:
1) dilute hydrochloric acid HCl
2) barium nitrate Ba(NO3)2
3) sodium hydroxide NaOH
4) sodium chloride NaCl
5) hydrogen nitrate (nitric acid) HNO3
6) barium hydroxide Ba(OH)2
7) barium chloride BaCl2
8) potassium hydroxide KOH
Add one drop of phenolphthalein solution to each test tube and mix the solution. Note
any change in the color of the solution.
Record your observations in the space provided. No equation is needed.
4.Reactions with Methyl Orange Indicator
Into eight (8) 25-mm test tubes add approximately 0.5 mL of each of the 8 solutions
listed below:
1) dilute hydrochloric acid HCl
2) barium nitrate Ba(NO3)2
3) sodium hydroxide NaOH
4) sodium chloride NaCl
5) hydrogen nitrate (nitric acid) HNO3
6) barium hydroxide Ba(OH)2
7) barium chloride BaCl2
8) potassium hydroxide KOH
Add one drop of methyl orange solution to each test tube and mix the solution. Note any
change in the color of the solution.
Record your observations in the space provided. No equation is needed.
Note: The presence of dissolved CO2 in the water can give confusing results. If needed,
add ONE DROP of the NaOH solution to any questionable test tube(s) and see if this
alters your conclusions.
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CHE 1400
Experiment 4
Chemical Reactions
Name(s)
Date
Laboratory Instructor
REPORT SHEET
-A
Reagents
Equations
Hydrogen chloride + silver nitrate
Barium nitrate + silver nitrate
Sodium hydroxide + silver nitrate
Sodium chloride + silver nitrate
Hydrogen nitrate + silver nitrate
Barium hydroxide + silver nitrate
Barium chloride + silver nitrate
Potassium hydroxide + silver nitrate
What ions are common to all the solutions that react with AgNO3?
-B
Reagents
Equations
Hydrogen chloride + sodium sulfate
Barium nitrate + sodium sulfate
Sodium hydroxide + sodium sulfate
Sodium chloride + sodium sulfate
Hydrogen nitrate + sodium sulfate
Barium hydroxide + sodium sulfate
Barium chloride + sodium sulfate
Potassium hydroxide + sodium sulfate
Which ion is common to all the solutions that react with Na2SO4?
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CHE 1400
-C
Reagents
Color observed
Hydrogen chloride + phenolphthalein
Barium nitrate + phenolphthalein
Sodium hydroxide + phenolphthalein
Sodium chloride + phenolphthalein
Hydrogen nitrate + phenolphthalein
Barium hydroxide + phenolphthalein
Barium chloride + phenolphthalein
Potassium hydroxide + phenolphthalein
Which ion is common to all the solutions that interact with phenolphthalein?
-D
Reagents
Color observed
Hydrogen chloride + methyl orange
Barium nitrate + methyl orange
Sodium hydroxide + methyl orange
Sodium chloride + methyl orange
Hydrogen nitrate + methyl orange
Barium hydroxide + methyl orange
Barium chloride + methyl orange
Potassium hydroxide + methyl orange
What ions are common to all the solutions that interact with methyl orange?
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CHE 1400
EXPERIMENT 5
SOLUBILITY
Vinegar and oil separate
OBJECTIVES
1. To understand and predict which liquids will be miscible with each other.
2. To measure mass and volume and then calculate density for several liquids.
3. To observe soap as both an emulsifying agent and as a cleaning agent.
Relates to chapter 6 of “Chemistry for Changing Times, 13th Ed”.
BACKGROUND
Molecules which are similar to each other in shape do usually dissolve in each other, that
is, they are very miscible. Water is a polar molecule, which is usually immiscible with
non-polar oils. Vinegar and oil dressing is an example of the denser water solution of
vinegar forming a separate layer under the less dense oil layer.
To blend these into one liquid, called an emulsion, requires an emulsifying agent. The
emulsifier works by having one end soluble in water and the other end soluble in oil. The
water and oil molecules are thus held together. Eggs are used as an emulsifying agent in
mayonnaise. In an old jar of mayonnaise often a top layer of oil can be observed because
some of the oil has separated.
Immiscibility can be mimicked by two liquids if one is very viscous or thick.
Viscosity is the ability to flow. A very viscous liquid will flow very slowly and thus will
inhibit mixing and miscibility. Syrup and water will be used as an example in this
investigation. The syrup and water can be mixed by putting energy into the system. The
energy is added by stirring vigorously.
Liquids that are miscible may dissolve in each other so well that the new volume is less
than the sum of the volumes. The molecules fit inside each other and require less space
together than by themselves.
Another aspect of solubility is that one set of intermolecular attractions may be stronger
than another. In this investigation, the attraction of oil for cloth will be stronger than the
attraction of gum cloth. Thus the oil will loosen the gum/cloth attraction and allow the
gum to be removed.
Soap actually works much the same way in removing soil from cloth or other materials.
The soil/clothe attraction is broken by stronger attractions of soap/soil. Soap acts as an
emulsifying agent. Soap has a long organic end which is soluble in oil or grease and an
ionic end which is soluble in water. Small balls of oil are surrounded by soap molecules
with the organic end of the soap in the oil. The ionic end is then left sticking out to attract
water. The entire ball is called a micelle. Watch the soap commercials on TV, which
show dirt leaving a dish or cloth. The dirt is rolled into a ball.
WASTE AND THE ENVIRONEMENT
The solutions used in this experiment are not toxic.
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CHE 1400
PROCEDURE
Part I
1.
Weigh a clean and dry 10-mL graduated cylinder.
Fill the cylinder with exactly 10 mL of syrup and reweigh.
Pour the syrup in a 50-mL beaker and set it aside.
2.
Wash the cylinder, dry the outside, fill with exactly 10 mL of water and reweigh.
Pour the water very carefully down the side of the beaker so that the water does not
mix with the syrup. Set the beaker aside.
3.
Dry the cylinder inside and out. Fill the cylinder with exactly 10 mL of cooking oil
and reweigh. Pour the oil carefully down the side of the beaker. When you finish you
should have three separate layers.
4.
With a stirring rod mix the three layers well and set the beaker on the lab bench.
Two layers will reform.
5.
With an eyedropper remove enough of only the top layer to put exactly 5 mL into the
10-mL graduated cylinder. Weigh the cylinder and liquid. Record the masses in the
report sheet. Clean and dry the cylinder.
6.
With an eyedropper reach through the solution and remove enough of only the
bottom layer to put exactly 5 mL into the 10-mL graduated cylinder. Weigh the
cylinder and liquid. Record the masses on the report sheet.
7.
Add 1 mL of liquid soap to the beaker. Stir well and set the beaker on the lab bench.
Part II
1.
Place exactly 5 mL of water in a 10-mL graduated cylinder. In a second 10-mL
cylinder place exactly 5 mL of alcohol. Pour the two liquids together in the water
cylinder. Read the new volume.
2.
Chew a piece of gum to remove most of the sweet taste. Place the gum in a cloth
square and squeeze to stick the gum to the cloth. Pour a small amount of oil and rub
it into the cloth, loosening the gum from the cloth. When the gum is removed, wash
the oil out with soap.
Pour the solutions down the drain with plenty of water.
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CHE 1400
Experiment 5
Solubility
Name(s)
Date
Laboratory Instructor
REPORT SHEET
Part I
Syrup
Water
Oil
Weight of cylinder and liquid (g)
__________
__________
__________
Minus weight of cylinder (g)
__________
__________
__________
Weight of liquid (g)
__________
__________
__________
__________
__________
__________
Density of liquid 
weight of liquid

10 mL
Which liquid is the densest?
Which liquid is on the bottom of the beaker?
Which two layers mix in procedure 4?
Top layer
Bottom layer
Weight of cylinder and liquid (g)
__________
__________
Minus weight of cylinder (g)
__________
__________
Weight of liquid (g)
__________
__________
__________
__________
weight of liquid

5 mL
Is the density of the mixture an average of the original two layers?
Density of liquid 
Does the soap act as an emulsifying agent?
Why would forming an emulsion be important when washing greasy dishes?
Part II
Volume of water: 5 mL.
Volume of alcohol: 5mL.
Measured volume of water (5 mL) + alcohol (5 mL): __________ mL
Is the volume different, why?
What is in oil that allows it to remove the gum?
Why would using oil be better to remove gum from hair than scissors?
Would peanut butter work similar to the vegetable oil to remove gum?
Just like egg, soap would allow the vinegar and oil in a dressing to remain mixed, but
would it be good to eat?
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CHE 1400
EXPERIMENT 6
ACIDS, BASES, BUFFERS and pH
OBJECTIVES
1. To gain an understanding of the relationship between pH and the concentration of
acids and bases.
2. To observe the difference in properties between strong acids and weak acids.
3. To learn how a buffer is prepared and to observe how it functions.
Relates to chapter 7 of “Chemistry for changing times, 13th Ed.”.
INTRODUCTION
Acids and bases
We often encounter acidic solutions in our daily life. Many foods owe their characteristic
tastes to the presence of specific acids. For example, citrus fruits contain citric acid, soft
drinks often contain phosphoric acid and vinegar is little more than a solution of dilute
acetic acid. Other acidic foods include tomatoes, coffee, apples and cabbage. Generally,
the acids in food give them a noticeably sour taste. Our stomachs then secrete
hydrochloric acid; which is needed to help to digest food.
Commonly encountered bases include many cleaning agents such as soaps, ammonia, and
lye (sodium hydroxide). A few foods such as egg whites are basic in nature, and many
plants contain a special group of bases called alkaloids. Morphine, caffeine, nicotine and
cocaine are all examples of plant alkaloids. Many commonly occurring minerals, such as
limestone, magnesite, dolomite and soda ash are basic in nature.
There are several definitions of both acids and bases, but the Arrhenius definition is one
of the most useful. According to this definition, acids are compounds that can act as
proton donors (forming hydronium ions, H3O+), and bases are compounds that act as
proton acceptors (forming hydroxide ions, OH-).
Strength of acids and bases
When any acid reacts with water it releases the hydronium ion, H3O+, along with the
corresponding anion. A strong acid is a compound that dissociates (breaks up) completely
into ions when added to water. Examples of strong acids are HCl, HF, HNO3, H2SO4 and
H3PO4. Since most of these acids do not contain any carbon they are often referred to as
mineral acids.
HCl + H2O
H3O+ + ClH2SO4 + H2O
H3O+ + HSO4-
These reactions go totally to the right, meaning that no associated broken up HCl or
H2SO4 exists in the solution. Therefore, one mole of HCl in water releases one mole of
Hydronium ions along with one mole of chloride anions. One mole of H2SO4 in water
releases one mole of hydronium ions along with one mole of HSO4- anions.
Similarly strong base completely dissociates to form the hydroxide ion, OH- and a
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CHE 1400
corresponding anion:
Na+ + OH-
NaOH
CaOH+ + OH-
CaO + H2O
In contrast a weak acid, such as acetic acid CH3COOH or formic acid
HCOOH, dissociate only to a small extent
H3O+ + CH3COO-
CH3COOH + H2O
H3O+ + HCOO-
HCOOH + H2O
The double arrows in the equations for a weak acid denote that the reactions proceed in
both directions (at different rates) at the same time. The longer arrow pointing to the left
indicates that the majority of the molecules are in the associated form.
Because nearly all of these acids contain at least one atom of carbon, they are often
referred to as organic acids. The double arrow indicates that a given molecule, of acetic
acid for instance, may give up an atom of hydrogen (H) to form a proton (H+), then later
regain hydrogen from a passing hydronium. Although a given molecule of acetic acid
may release and gain a hydronium ion repeatedly over time, the percentage of acetic acid
molecules in the associated (CH3COOH) and dissociated (H3O+ + CH3COO-) states
remain constant. One mole of a weak acid dissolved in water will generate less than one
mole of hydronium ions.
The amount of a weak acid that dissociates is defined by the equilibrium constant
expression. This value is a property of a weak acid and relates the ratio of the moles of
dissociated and associated acid present in the solution. For acetic acid (CH3COOH) this
ratio is:
H O  CH 3COO 
[1]
Ka  3
 1.8  10 5
CH 3COOH 



As a first approximation it can be assured that the only source of hydronium ions is the
acid and therefore (H3O+) = (CH3CCO-). It can also be assumed that only a small
percentage of those acids are dissociated and therefore the value of (CH3COOH) does not
change. The above equation can then be rearranged to:
Ka 
H O 
 2

3
CH 3COOH 
H O  

3
Ka  CH 3COOH 
[2]
Similarly for a weak base, such as ammonia, only a small number of the ammonia
molecules dissociate into the ammonium and hydroxide ions.
NH3 + H2O
NH4+ + OHThe degree of dissociation can also be expressed through an equilibrium constant
expression:

NH 4 OH 
Kb 
 1.8  10 5
[3]
NH 3 
Note that, even though ammonium hydroxide is a base and acetic acid is an acid, they

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
28
CHE 1400
have the same degree of dissociation, as shown by the values of Ka and Kb.
The strength of an acidic solution can be expressed by the concentration of the
hydronium ion (H3O+) which is defined as the number of mole of H3O+ per liter of a
solution. In pure water, a small number of water molecules dissociate into both the
hydronium and hydroxide ions:
to the extent that H 3O    OH    1.0 107
Since the concentration of the two ions are numerically equal, the solution is considered
to be neutral, neither acidic nor basic. Due to the nature of water, small concentrations of
both H3O+ and OH- are always present. Numerically, this is expressed as the ion product
of water and is defined as follows:
[4]
H 3O   OH   1.0 1014
H2O
H3O+ + OH-

 

The (H3O+) from a strong acid, e.g HCl, is the same as the concentration of the acid
(HCl). For example, in a 0.01 M solution of HCl, H 3O    0.01 M . For a
1.0 x10-5 M solution of HNO3, H 3O    1.0  10 5 M .
Similarly, for a strong base, the (OH-) is the same as the base concentration. For example,
in a 0.01 M solution of KOH, the OH   0.01 M . Since the product of (H3O+) and
(OH-) ions is a constant, the concentration of the hydronium ion can be easily calculated.


H O  OH   1.0 10
10
1.0 10
H O   1.0OH


0.01


14
3
14


3
14
 1.0 1012
Due to the large numerical range possible for the hydronium ion it is more convenient to
express this value in terms of the power of the hydronium ion concentration, more
commonly referred to as the pH. The pH is defined as:

pH   log H 3O 

[5]
Since in pure water, H 3O    1.0  10 7 , then for a neutral solution of pure water the pH
is 7.0. As the hydronium ion concentration increases, the value of the pH numerically
decreases. As the hydronium ion concentration decreases, the value of the pH
numerically increases. It should be remembered that the pH scale is a logarithmic one. A
pH of 6.0 is ten times more acidic than water, a pH of 5.0 is 100 times more acidic and a
pH of 4.0 is 1000 times more acidic.
Buffer solutions
Buffer solutions are solutions that tend to resist changes in pH. A typical buffer is a
solution of a weak acid (such as CH3COOH) and the salt of a weak acid (such as
CH3COO- Na+). If a small amount of base is added to the buffer solution, the base reacts
with the weak acid to form more of the salt and water:
CH3COOH + NaOH
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CH3COO-Na+ + H2O
29
CHE 1400
If some acid is added to the buffer solution it reacts with the anion of the salt (which
behaves as a base, i.e. a proton acceptor) to reform the weak acid and an additional salt:
CH3COO-Na+ + HCl
CH3COOH + NaCl
Since in either case the added base or acid is neutralized, there is very little change in the
pH of the solution. The addition of acid or base to a buffered solution does change the pH
of that solution, but only to a small extent. If a sufficiently large amount of base (or acid)
is added so that all of the CH3COOH (or CH3COO-Na+) is consumed, then that capacity
of the buffering solution has been exceeded, and it will lose its buffering ability.
EXPERIMENTAL PROCEDURE
1 - pH values of common household materials
Solutions of the items listed below, which are commonly found in the household, have
different pH values.
Ammonia
Dishwashing detergent
Table salt
Vinegar
Lemon juice
Household cleaner
Baking soda
Oven cleaner
Using a disposable pipette wet a small strip of indicator paper with one of the solutions.
Determine the pH of the solution by comparing the color of the wet indicator paper with
the color chart provided.
Record the pH of the solution on the report sheet. Indicate if the solution is acidic (A),
neutral (N), or basic (B).
Continue with the remaining 7 solutions.
2- pH values of acids, bases and buffers
2.1- HCl, a strong acid
2.1.1
Put about 4 mL of 0.1 M HCl into a small test tube.
Place a glass rod into the 0.1 M HCl solution, and then touch the tip of the rod on a small
strip of the pH indicator paper. Determine the pH of the solution by comparing it to the
color chart provided.
Record the calculated (H3O+), assuming that the HCl completely dissociates.
From the above (H3O+), calculate the expected pH from equation [5].
Record the observed pH.
2.1.2
Rinse a 10-mL graduated cylinder thoroughly with distilled water and then dilute 1.0 mL
of sample of the 0.1 M HCl up to 10 mL with distilled water. Pour about 6 mL into a
small test tube and save the remaining 4 mL.
Wet a small strip of the indicator with the diluted solution from section 2.1.2.
Determine the pH of the solution by comparing it to the color chart provided.
Record the concentration of the HCl.
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CHE 1400
Record the calculated (H3O+), assuming that the HCl completely disassociates.
From the above (H3O+), calculate the expected pH from equation [5].
Record the observed pH.
2.1.3
Rinse the 10-mL graduated cylinder thoroughly with distilled water and then dilute a 1.0
mL sample of the HCl solution (saved from section 2.1.2) to 10 mL of distilled water and
transfer to a small test tubes.
Wet a small strip of the indicator with the diluted solution from section 2.1.3.
Determine the pH of the solution by comparing it to the color chart provided.
Record the concentration of the HCl
Record the expected (H3O+), assuming that the HCl completely disassociates.
From the above (H3O+), calculate the expected pH from equation [5].
Record the observed pH.
2.2- NaOH , a strong base
2.2.1
Put about 4 mL of the 0.001 M NaOH solution into a small test tube
Wet a small strip of the indicator with the 0.001 M NaOH solution. Determine the
pH of the solution by comparing it to the color chart provided.
Record the calculated (OH-), assuming that the NaOH completely disassociates.
Calculate the (H3O+) from equation [4].
From the above (H3O+), calculate the expected pH from equation [5].
Record the observed pH.
2.2.2
Rinse a 10-mL graduated cylinder thoroughly with distilled water and then dilute a 1.0mL sample of the 0.001 M NaOH up to 10 mL with distilled water.
Pour about 6 mL into a small test tube and save the remaining 4 mL.
Wet a small strip of the indicator with the diluted from section 2.2.2 solution.
Determine the pH of the solution by comparing it to the color chart provided.
Record the (NaOH)
Record the calculated (OH-), assuming that the (NaOH) completely disassociates.
Calculate the (H3O+) from equation [4].
From the above (H3O+), calculate the expected pH from equation [5].
Record the observed pH.
2.2.3
Rinse the 10-mL graduated cylinder thoroughly with distilled water and then dilute a mL
sample of the NaOH solution (saved from section 2.2.2) up to 10 mL with distilled water
and transfer to a small test tube.
Wet a small strip of the indicator with the diluted solution from section 2.2.3.
Determine the pH of the solution by comparing it to the color chart provided.
Record the (NaOH).
Record the calculated (OH-), assuming that the NaOH completely disassociates.
Calculate the (H3O+) from equation [4].
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CHE 1400
From the above (H3O+), calculate the expected pH from equation [5].
Record the observed pH.
2.3. CH3COOH , a weak acid
2.3.1
Put about 4 mL of the 0.1 M CH3COOH into a small test tube.
Wet a small strip of the indicator with 0.1 M CH3COOH solution. Determine the pH of
the solution by comparing it to the color chart provided.
Calculate the expected (H3O+) from equation [2].
From the above (H3O+), calculate the expected pH from equation [5].
Record the observed pH.
2.3.2
Rinse a 10-mL graduated cylinder thoroughly with distilled water, and then dilute a
1.0 mL sample of the 0.1 M CH3COOH with distilled water up to 10 mL. Pour about
6 mL into a small test tube and save the remaining 4 mL.
Wet a small strip of the indicator with the diluted solution from section 2.3.2.
Determine the pH of the solution by comparing it to the color chart provided.
Record the concentration of the CH3COOH
Calculate the expected (H3O+) from equation [2].
From the above (H3O+), calculate the expected pH from equation [5].
Record the observed pH.
2.3.3
Rinse a 10-mL graduated cylinder thoroughly with distilled water and then dilute a
1.0 mL sample of the CH3COOH solution (saved from section 2.3.2) to 10 mL with
distilled water and transfer to a small test tube.
Wet a small strip of the indicator with the diluted solution from section 2.3.3.
Determine the pH of the solution by comparing it to the color chart provided.
Record the concentration of the CH3COOH.
Calculate the expected (H3O+) from equation [2].
From the above (H3O+), calculate the expected pH from equation [5].
Record the observed pH.
2.4. CH3COOH and CH3COO-Na+, a buffer solution
2.4.1
Mix 10 mL of 0.2 M CH3COOH with 10 mL of 0.2 M CH3COO-Na+. This will give
20 mL of a solution containing 0.1 M CH3COOH and 0.1 M CH3COO-Na+.
Wet a small strip of the indicator with the CH3COOH/CH3COO-Na+ solution.
Determine the pH of the solution by comparing it to the color chart provided.
Record the observed pH of the CH3COOH/CH3COO-Na+ solution.
2.4.2
Add one drop of 0.1 M HCl to the CH3COOH/CH3COO-Na+ solution.
Wet a small strip of the indicator with the CH3COOH/CH3COO-Na+ solution.
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CHE 1400
Determine the pH of the solution by comparing it to the color chart provided.
Record the observed pH of the CH3COOH/CH3CCO-Na+ solution.
2.4.3
Add nine more drops of 0.1 M HCl to the CH3COOH/CH3COO-Na+ solution in
section 2.4.2.
Wet a small strip of the indicator with the CH3COOH/CH3COO-Na+ solution.
Determine the pH of the solution by comparing it to the color chart provided.
Record the observed pH of the CH3COOH/CH3COO-Na+ solution.
2.4.4
Add one drop of 0.1 M NaOH to a fresh sample of the CH3COOH/CH3COO-Na+ solution.
Wet a small strip of the indicator with the CH3COOH/CH3COO-Na+ solution.
Determine the pH of the solution by comparing it to the color chart provided.
Record the observed pH of the CH3COOH/CH3COO-Na+ solution.
2.4.5
Add nine more drops of 0.1 M NaOH to the CH3COOH/CH3COO-Na+ solution.
Determine the pH of the solution by comparing it to the color chart provided.
Record the observed pH of the CH3COOH/CH3COO-Na+ solution.
2.5 Water, an unbuffered solution.
2.5.1 Wash a test tube thoroughly and add about 4 mL of distilled water.
Wet a small strip of the indicator with the water. Determine the pH of the water by
comparing it to the color chart provided.
Record the observed pH of the water.
2.5.2 Add one drop of 0.1M HCl to the water.
Wet a small strip of the indicator with the water. Determine the pH of the water/acid
solution by comparing it to the color chart provided.
Record the observed pH of the water/acid solution.
2.5.3 Add nine more drops of 0.1 M HCl to the water/acid solution in section 2.5.2.
Wet a small strip of the indicator with the water/acid solution. Determine the pH of the
solution by comparing it to the color chart provided.
Record the observed pH of the water/acid solution.
2.5.4 Add one drop of 0.1 M NaOH to fresh sample of distilled water.
Wet a small strip of the indicator with the water/base solution. Determine the pH of the
solution by comparing it to the color chart provided.
Record the observed pH of the water: base solution.
2.5.5 Add nine more drops of 0.1 M NaOH to the water/base solution in section 2.5.4.
Wet a small strip of the indicator with the water/base solution. Determine the pH of the
solution by comparing it to the color chart provided.
Record the observed pH of the water/base solution.
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CHE 1400
Experiment 6
Acids, bases, buffers and pH
Name(s)
Date
Laboratory Instructor
REPORT SHEET
1. pH of common household materials.
Indicate “A” (Acidic), “N” (Neutral) or “B” (Basic).
1. Ammonia solution
pH:
______________
2. Lemon juice
pH:
______________
3. Baking soda solution
pH:
______________
4. Dishwashing detergent
pH:
______________
5. Table salt solution
pH:
______________
6. Oven cleaner
pH:
______________
7. Vinegar
pH:
______________
8. Household cleaner solution pH:
______________
2. pH values of acids, bases and buffers
2.1 HCl , a strong acid
2.1.1
2.1.2
2.1.3
(HCl)
=
0.1 M
Calculated (H3O+) =
______________
Expected pH =
______________
Observed pH =
______________
(HCl)
=
0.01 M
Calculated (H3O+) =
______________
Expected pH =
______________
Observed pH =
______________
(HCl)
=
+
0.001 M
Calculated (H3O ) =
______________
Expected pH =
______________
Observed pH =
______________
Are the observed pH values consistent with the expected ones?
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CHE 1400
2.2. NaOH, a strong base
2.2.1
2.2.2
2.2.3
(NaOH) =
0.1 M
Calculated (OH-) =
______________
Calculated (H3O+) =
______________
Expected pH =
______________
Observed pH =
______________
(NaOH) =
0.01 M
Calculated (OH-) =
______________
Calculated (H3O+) =
______________
Expected pH =
______________
Observed pH =
______________
(NaOH) =
0.001 M
-
______________
+
Calculated (H3O ) =
______________
Expected pH =
______________
Observed pH =
______________
Calculated (OH ) =
How much do the pH values vary with each tenfold dilution of the base? Is this change
expected?
2.3. CH3COOH , a weak acid
2.3.1
Lab Manual
(CH3COOH) =
0.1 M
Calculated (H3O+) =
______________
Expected pH =
______________
Observed pH =
______________
35
CHE 1400
2.3.1.1 Are the observed pH values consistent with the expected ones?
2.3.1.2 How does the pH of this weak acid compare to the pH of a strong acid of the same
molarity?
2.3.2
(CH3COOH) =
+
2.3.2
0.01 M
Calculated (H3O ) =
______________
Expected pH =
______________
Observed pH =
______________
(CH3COOH) =
0.001 M
Calculated (H3O+) =
______________
Expected pH =
______________
Observed pH =
______________
2.3.3.1 How do the changes in pH of a weak acid diluted 100 fold compare to the change
in the pH of a strong acid diluted 100 fold?
2.4 CH3COOH and CH3COO-Na+, a buffer solution
2.4.1 Initial condition
observed pH = ______________
2.4.2 After adding 1 drop of HCl
observed pH = ______________
2.4.3 After adding 10 drops of HCl
observed pH = ______________
2.4.4 After adding 1 drop of NaOH
observed pH = ______________
2.4.5 After adding 10 drops of NaOH
observed pH = ______________
What is the observed change in the pH of the buffer solution after the addition of:
1 drop of acid?
____________________________________
10 drops of acid?
____________________________________
1 drop of base?
____________________________________
10 drops of base?
____________________________________
Does a solution of acetic acid and sodium acetate act as a buffer? Explain.
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CHE 1400
2.5 Water, an unbuffered solution
2.5.1 Initial condition
observed pH = ______________
2.5.2 After adding 1 drop of HCl
observed pH = ______________
2.5.3 After adding 10 drops of HCl
observed pH = ______________
2.5.4 After adding 1 drop of NaOH
observed pH = ______________
2.5.5 After adding 10 drops of NaOH
observed pH = ______________
What is the observed change in the pH of the buffer solution after the addition of:
1 drop of acid?
____________________________________
10 drops of acid?
____________________________________
1 drop of base?
____________________________________
10 drops of base?
____________________________________
Does a solution of water act as a buffer? Explain
Compare the effects of acids and bases on a buffer solution and unbuffered water.
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CHE 1400
EXPERIMENT 7
ACID NEUTRALIZATION BY ANTACID
How to stop heart-burn
OBJECTIVES
1. To achieve neutralization in a titration reaction.
2. To observe the effects of decreasing hydrogen ion concentration on the indicator
phenolphthalein.
3. To calculate the amount of acid absorbed per gram of antacid based on data gathered
from titrations.
Relates to chapter 7 of chemistry for changing times, 13th Ed.
BACKGROUND
Stomach acid is a combination of gastric juices and an acid very similar to hydrochloric
acid. Sometimes eating rich food or experiencing stress may cause more than the usual
amount of stomach acid to be produced. This causes discomfort to the person.
Commercial antacids are primarily composed of basic or alkaline compounds and binders
to hold the tablet together, but sometimes fillers are also added. The basic compounds
react with or neutralize the excess stomach acid, which causes acid stomach or « heart
burn ». Some of the basic compounds are hydroxides such as magnesium hydroxide,
Mg(OH)2, or aluminum hydroxide, Al(OH)3. The hydroxide ion is released when the
hydroxide dissolves in water. The reaction between the hydroxide ion and the acidic
hydrogen ion reduces the amount of acid and relieves the discomfort.
Mg(OH)2 (s)
H2O
Mg2+ + 2 OH-
OH- + H+
H2O
The hydrogen ion doesn’t actually exist by itself. It is always combined with some
molecule. Usually that molecule is water and will form ions such as H3O+ or H5O2+.
Because we don’t know the exact form of the ion, it is easier to just write hydrogen ion as
H+ or H+ (aq).
Other basic compounds used, as antacids are carbonates such as calcium carbonate,
CaCO3, and sodium carbonate, Na2CO3. These react with the hydrogen ion to form
carbonic acid, which quickly dissociates into water and carbon dioxide. The carbon
dioxide is a gas and may cause belching which also helps to relieve stomach distress.
CO32- + 2 H+
H2CO3
H2O + CO2
This investigation involves doing a back titration. To an excess amount of acid, the
antacid will be added. The amount of base necessary to neutralize the solution is equal to
the excess acid in the solution. The titration of an excess of reagent added is a back
titration. Thus, the less base required to reach neutrality, the more acid was absorbed by
the antacid. The reaction is:
Lab Manual
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CHE 1400
H+
from
acid
+
OH-
H2O
from
antacid
Phenolphthalein is a compound that is colorless in its acidic form. When it loses a
hydrogen ion to get its basic form, it gives a pink color. Since it will lose its hydrogen ion
to other bases, it is used as an indicator.
Which antacid neutralizes the most acid? It is time for you to find out.
WASTE AND ENVIRONMENT
Concentrated acids and bases can damage plumbing if not neutralized or diluted.
Pour the solutions down the drain with plenty of water.
PROCEDURE
1. In a clean Erlenmeyer flask, weigh 1.00 g of liquid antacid named “Maalox 1”.
2. Then, add 50 mL of 0.36 M hydrochloric acid solution to the flask. The quantity of
acid added is in excess. Add 2-3 drops of phenolphthalein and a magnetic stirring bar.
3. Place the Erlenmeyer flask on the top of a magnetic stirrer and start stirring the
solution for 5 sec.
1. Fill the burette with a 0.1M NaOH solution. Start adding NaOH to the Erlenmeyer
flask, quickly in the range of 0-14 mL, and dropwise afterwards.
A pink color can be noted and will fade quickly with stirring. Stop adding when one
drop causes a pink color that does not fade within 30 seconds: the end point is
reached. Record the volume of NaOH added.
2. Repeat steps 1-3 for the second liquid antacid named “Maalox 2”.
Flush all solutions down the drain with plenty of water.
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CHE 1400
Experiment 7
Acid neutralization by antacid
Name(s)
Date
Laboratory Instructor
REPORT SHEET
Brand name of antacid
Mass of antacid (g)
Vbase added (mL)
nacid added  C ACID  V ACID added  0.36 M  50 mL 
Antacid 1
Antacid 2
“Maalox 1”
“Maalox 2”
1.00 g
1.00 g
__________ mL
__________ mL
____________ mmol
nbase added  C BASE  VBASE added 
________ mmol
________ mmol
nacid neutralized = nacid added - nbase added =
________ mmol
________ mmol
________ mmol/g
________ mmol/g
nacid neutralized per gram of antacid
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CHE 1400
IV. QUESTIONS
1. Read the antacid label to find the basic compounds and write the balanced chemical
equations showing the reaction with the hydrogen ion
Basic compounds:
______________________________________________
Equation 1:
.
Equation 2:
.
2. What happens to the pH of your stomach if you take more antacid than necessary to
neutralise the acid?
3. Which antacid has more neutralizing power per gram?
4. Which antacid appears to be best at relieving acid stomach?
5. In recent commercials some antacids brag that they contain calcium. If calcium is
used by the body to form bones and teeth, why does it matter that it is in an antacid?
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CHE 1400
EXPERIMENT 8
OXIDATION AND REDUCTION
Those travelling electrons
OBJECTIVES
1. To observe several redox reactions and note the changes of substances from their
elemental form to their ionic form and vice-versa.
2. To create a voltaic cell and measure the current it produces.
Relates to chapter 8 of “Chemistry for changing times, 13th Ed.”.
BACKGROUND
Oxidation-reduction reactions are reactions in which electrons are transferred from one
element to another. Many everyday reactions are based on oxidation-reduction (redox)
reactions. Even rust is produced by a redox reaction:
4 Fe + 3 O2
2 Fe2O3
(rust)
One common misconception is that water causes rust. Although objects in water will rust
more quickly, water only provides the medium for the ions to travel more quickly.
Copper ions and aluminium metal
Redox reactions occur because one element has a stronger attraction for electrons than
another element. As an example, the copper 2+ ion will take electrons from aluminum
metal, producing copper metal and aluminum ions.
3 Cu2+ + 2 Al (s)
2 Al3+ + 3 Cu (s)
Due to the presence of a thin and dense layer of aluminium oxide, Al2O3, on the surface
of any piece of aluminium exposed to the oxygen of the air, Cu2+ ions cannot react
directly with Al (s) as described in the equation above.
The solution to this problem is to add Cl- ions to the solution by adding a few drops of
NaCl. Chloride ions react easily with aluminium to form AlCl3. In presence of Cl- ions,
the reaction is:
3 CuCl2 (aq) + 2 Al (s)
Lab Manual
2 AlCl3 (aq) + 3 Cu (s)
42
CHE 1400
Iodine and zinc metal
Another redox reaction occurs between zinc and iodine in solution.
Zn (s)
+
Zn2+ + 2 I-
I2
(in alcohol)
purple
colorless
Chlorine will react with the iodide ion.
2 I- + Cl2
H+
2 Cl- + I2
colorless
purple
Voltaic cell
Batteries are based on redox reactions, which are
arranged to cause the electrons to flow through
an external circuit. Different metals can be used
to set up a redox cell, which will produce a
voltage between the two metals. This
arrangement is called a voltaic cell.
A simple voltaic cell can be produced by two
metal strips (eg copper and zinc) with different
affinities for electrons which are connected by an
electron path. Paper towels, which have been
soaked into sulfuric acid, H2SO4, can serve as an
ionic bridge. The free hydrogen ions will be
attracted by the copper strip (cathode) where they
will be reduced into hydrogen gas, while free
sulfate ions will remain in solution to balance the
zinc ions released by the zinc strip (anode).
Finally, a voltmeter connects the two metal strips
to allow a path for the flow of electrons.
The voltage read is the attraction of one metal for
the electrons compared to the other metal’s
attraction.
WASTE AND THE ENVIRONEMENT
 The solutions in this investigation are not toxic.
 Wet zinc dust in air can burst into flames. Placing the zinc solutions or damp zinc
solids on a metal pad protects from fire damage and sets up conditions for an oxidation
reduction reaction which will form zinc oxide.
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CHE 1400
 Zinc oxide has a low toxic hazard rating.
PROCEDURE
1. a) Place a small amount of zinc powder in a small beaker. Cover the metal with
tincture of iodine solution and set it aside.
b) Wait a few minutes. If the brown color of the solution over the zinc has faded,
pour the liquid into a second beaker leaving the zinc powder behind. Add several
drops of bleach to the solution. Then, add a few drops of 1 M acetic acid. Observe
the change of color.
2.
a) Place 25 mL of 1 M copper sulfate (CuSO4) in a 50-mL beaker. Roll a «4 by 4»
piece of aluminum foil into a roll. Place the aluminium roll in the beaker and add 5
drops of NaCl solution.
b) Place 5 mL of 1 M copper sulfate in a second 50-mL beaker and set it aside.
c) Wait a few minutes. Remove the aluminium foil from the beaker of copper sulfate.
The resulting orange/black solid is copper metal. Observe the difference in the
aluminum foil. Compare the color of the two beakers of copper sulfate.
3.
Make a voltaic cell using a strip of zinc metals and a strip of copper metal. Cut a
piece of paper towel to about the size of a strip. Wet these pieces with 1 M sulfuric
acid H2SO4. Place one strip on the bottom, then the piece of wet paper towel and the
other strip on top. Call the instructor to measure the voltage produced by this voltaic
cell.
The aluminum foil may be thrown away in the trash.
The zinc solutions from steps 1 and 5 should be placed on a metal pan to dry.
The zinc oxide formed can then be buried in a landfill.
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CHE 1400
Experiment 8
Oxidation and reduction
Name(s)
Date
Laboratory Instructor
REPORT SHEET
I. COPPER AND ALUMINIUM
1. What does the difference in color in the copper solution indicate?
2. What happened to the aluminium foil?
3. In what form is the copper at the end of the reaction?
4. Could copper be recycled by collecting it on aluminium foil?
II. ZINC, IODINE AND CHLORINE
1. In what form is the zinc (Zn(s) or Zn2+(aq)) when the purple color is gone from the
solution?
2. Does bleach change the zinc ions back to the metal?
3. In what form is the chlorine (Cl2(aq) or Cl-(aq)) when the purple color appears in the
solution?
III. Cu/Zn/H2SO4 VOLTAIC CELL
voltage:
________________V
1. Would other metal strips work?
2. Would other acidic solutions like HCl work in the voltaic cell instead of H2SO4 ?
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CHE 1400
EXPERIMENT 9
CLARIFICATION OF WATER
Is it clean enough to drink?
OBJECTIVES
1. To understand the process of cleaning a particulate matter as one of several steps
necessary to make it drinkable.
2. To produce an alternate usable compound from used aluminum cans.
Relates to chapter 12 and 14 of “Chemistry for changing times, 13th Ed.”.
BACKGROUND
Cleaning water enough for it to be usable is a major task for most of the world. As we
increase our population it becomes a larger problem. A primary method to clean water is
to use a flocculent, a compound that produces a precipitate that will settle to the bottom,
thus trapping suspended solids as it settles. This compound increases the speed at which
suspended particles will settle and traps some particles, which would not otherwise settle.
One example of this type of compound is aluminum hydroxide, a gel-like solid. The
aluminum hydroxide will increase the amount of solids which settle out, and the speed at
which they settle out, but the time period is often days or weeks. If we could make the
aluminum hydroxide from aluminum cans, we also reduce litter.
Making aluminium hydroxide from aluminium cans is not usually done as it is too
expensive and making new aluminium cans is a better use for the old aluminium cans.
Another method of cleaning water is to pass it over activated charcoal. This is considered
an advanced or tertiary method. Activated charcoal works by adsorbing the heavy
molecules in the water. The molecules are caught on the irregular surface of the activated
charcoal. After cleaning the water, the charcoal can be regenerated by heating it to
500-1000 °C with steam or carbon dioxide.
In this investigation, the aluminium hydroxide is compared to water that settles without
the precipitate and to the original dirty water. The dirty water is produced so that it is a
visible “dirty” suspension. The settling is speeded up by centrifuging so that the length of
time fits into the lab time.
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CHE 1400
METHOD
Aluminum can be dissolved by the action of a strong base (steps [1]-[2]) to afford the
aluminium compound NaAl(OH)4. Next, this compound is converted into aluminium
sulphate, Al2(SO4)3, by the addition of sulfuric acid, H2SO4 (step [3]). The second goal of
sulfuric acid is to dissolve any remaining aluminium hydroxide, Al(OH)3, formed in
step [1]. Finally, aluminium ions present in the resulting solution are reacted with sodium
hydrogen carbonate (also called sodium bicarbonate, a cheap base) to produce a large
precipitate of aluminium hydroxide, Al(OH)3 (step [4]).
unreacted part removed by filtration
2 Al (s) + 6 H2O (l)
[1]
2 Al(OH)3 (s) + 3 H2 (g)
[2] + 2 NaOH (aq)
2 NaAl(OH)4 (aq)
[3] + 4 H2SO4 (aq)
Al2(SO4)3 (aq) + Na2SO4 (aq) + 8 H2O (l)
[4] + 6 NaHCO3 (s)
+ dirty water
2 Al(OH)3 (s) + 3 Na2SO4 (aq) + 6 CO2 (g)
+ dirt
+ dirt-free water
removed by centrifugation
Notes:
- In step [3], it is not necessary to add a large amount of sulfuric acid.
- In step [4], the filtrate from the aluminum sulfate solution contains aluminum ions in an
acidic solution. By adding sodium hydrogen carbonate until the solution is basic,
hydroxide ions are produced. These ions will precipitate with aluminum ions as
aluminum hydroxide (Al(OH)3), a gel like precipitate which will trap large particles.
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CHE 1400
WASTE AND THE ENVIRONMENT
 Acidic and basic solutions can damage plumbing unless neutralized or diluted.
 Aluminum ions are note classed as toxic but it is not good to put metal ions into the
water system.
PROCEDURE
1- Cut two 1 cm x 3 cm strips of aluminium.
2- Place the strips in an evaporating dish with 15 mL of 6 M sodium hydroxide (NaOH).
Make sure the strips are completely covered.
3- Warm the dish with a hotplate in a fume hood. Observe fizzing. Heat gently to
prevent boiling. Stir gently to prevent foaming over.
4- Continue until most of the fizzing stops (about 5 min.). Cool the solution.
5- Add 20 mL of 3 M sulfuric acid (H2SO4) and stir well to dissolve as much remaining
aluminum hydroxide (Al(OH)3) as possible.
6- Filter the aluminum sulfate solution by suction filtration using the Buchner funnel
suction apparatus as shown in Figure 9.1.
Wet the filter paper slightly with distilled water to seat the filter before adding the
solution.
Figure 9.1 Buchner funnel with safety flask
7-
Add 10 mL of the filtrate to 10 mL of dirty water in a 100-mL beaker. Add solid
sodium hydrogen carbonate (NaHCO3) slowly while mixing until the solution is
basic to pH paper (pH  7). Stir well and pour some of the slurry into a small test
tube and mark the tube “1”.
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CHE 1400
8-
Although the solids will settle over a period of time, we can speed the process by
centrifuging. Place about the same volume of dirty water in a second test tube
marked “2”. Balance the centrifuge by placing the test tubes on opposite sides.
Centrifuge for two minutes.
9-
Place the same volume of dirty water in a third test tube marked “3”. Compare the
three test tubes visually. Then, compare the 3 test tubes in a spectrophotometer set at
490 nm. Record the absorbance.
 The solutions can be flushed down the drain with a lot of water.
 The solids containing aluminum should be collected to be buried in a toxic waste
site.
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CHE 1400
Experiment 9
Clarification of water
Name(s)
Date
Laboratory Instructor
REPORT SHEET
I.
ALUMINUM HYDROXIDE
Compare the volume of aluminum hydroxide obtained to the original volume of
aluminum metal.
II.
USE OF ALUMINUM HYDROXIDE
CLARITY
OF
WATER
III.
Absorbance of test tube with dirty water
______________
Absorbance of test tube with centrifuged dirty water
______________
Absorbance of test tube with clarified water
______________
QUESTIONS
1.
What were the bubbles which formed when aluminum reacted with alkali?
2.
Is it easy to see aluminum hydroxide floating in water? Explain.
3.
What were the bubbles which formed when sodium hydrogen carbonate was
added to the sulfuric acid solution?
4.
What happens to the colorful decorations on the aluminum can when you
dissolve it?
5.
Is the clear water you have prepared now safe to drink? Why or why not?
6.
Why isn’t a centrifuge used to clean the city water?
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CHE 1400
EXPERIMENT 10
SAPONIFICATION
from butter to lye soap
OBJECTIVES
- To understand the process of saponification.
- To produce small amounts of soap from butter and sodium hydroxide.
- To calculate the average molar mass of soap produced from butter.
Relates to chapter 9 of “Chemistry for changing times, 13th Ed.”.
BACKGROUND
In the times of our grandparents and great grandparents, a wide range of cleaning
products was not available. In fact the main cleaning product was lye soap. It was used
for a large number of cleaning jobs: bathing, hair washing, clothes washing, and dish
washing. It was usually formed into a bar but it could be cut into chips to make it more
soluble for dishwashing or clothes washing.
Lye soap was made by the reaction of a fat or oil with lye (sodium hydroxide, also called
caustic soda). Although many soaps are now made with coconut oil, hog lard was
probably more commonly used in the days of our grandparents. Palm oil or whale oil can
also be used. Some restaurants have even used the grease collected in the kitchen to make
soap to wash the dishes. Potassium hydroxide can be used instead of sodium hydroxide.
The soap making process is called saponification. The reaction is: a fat plus sodium
hydroxide produces soap plus glycerol. Before the reaction, the oils and fats are tri-esters
(three ester groups). The reaction breaks the ester group apart so that a carboxylate ion
(RCOO-) and an alcohol (R-OH) are formed. By removing the molecules from water the
carboxylate ion combines with the sodium ion to form a solid that is called soap.
After the reaction is complete, the soap is separated from the glycerol by “salting out”.
O
C O
O
R C O CH2
O
R2 C O CH +
O
R3 C O CH2
1
A fat
O
C O Carboxylate
Ester
O
R C ONa
O
R2 C ONa +
O
R3 C ONa
1
3 NaOH
Soap
HO CH2
HO CH
HO CH2
Glycerol
The solution is mixed with a concentrated sodium chloride solution. This electrolyte
causes the dispersed soap to coagulate. The soap is then washed several times with
concentrated sodium chloride to remove the excess lye. R is a straight hydrocarbon chain
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CHE 1400
of 12-20 carbon atoms. The three R groups (R1, R2, R3) are often different from each
other. In your butter sample, there are three main R groups; linoleate, oleate, and
palmitate. Each produces one of the three most abundant soap molecules derived from
butter, sodium linoleate (approx. 50%), sodium oleate (approx. 30%), and sodium
palmitate (approx. 20%).
O
CH3(CH2)CH CHCH2CH CH(CH2)7CONa
O
CH3(CH2)7CH CH(CH2)7CONa
Sodium Linoleate
Sodium Oleate
O
CH3(CH2)14CONa
Sodium Palmitate
WASTE AND ENVIRONMENT
 The compounds in this investigation are not toxic.
 Acidic and basic solution can harm plumbing if not neutralized or diluted.
PROCEDURE
1. Boil a mixture of 10.0 g of butter and 15 mL of 20% sodium hydroxide (NaOH) in a
100-mL or 150-mL beaker until all of the water is evaporated. Boil the mixture as
strongly as possible in order to re move the water as quickly as possible.
Take the following precautions to prevent injury caused by the spattering of the soap
solution. Vigorous stirring with a long rod (at least 8 in.) will prevent spattering and
frothing if the boiling is not too rapid. The long rod is used to keep your hand far from
the beaker. Keep your head back and bellow the opening of the beaker. When it is
necessary to rest from stirring, remove the burner. And continue to stir until the
boiling ceases. Resume stirring before replacing the burner in one hand so that you can
remove it intermittently to avoid burning the soap. One purpose of the stirring is to
prevent the soap from forming a solid at the bottom of the beaker. A solid layer on the
bottom will often char.
2. When all of the water has apparently been removed let the mixture cool slightly. If a
waxy solid forms, the process is complete. Let the mixture cool. If a syrup liquid
results, the reaction is not complete and heating must be resumed.
3. Pour 20 mL of the concentrated sodium chloride solution into the beaker containing
the soap. Break up the soap into small pieces with the stirring rod so that all of the
soap is washed free of glycerol and sodium hydroxide.
4. Filter the soap solution with a piece of filter paper in a funnel. Fold a piece of filter
paper in half and half again (see Figure 10.1); tear off the outer corner of the fold.
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CHE 1400
Open the paper up between the first and second sheet, and place the paper in the
funnel. Pour the soap solution through the filter paper into a beaker.
Figure 10.1 Folding filter paper for gravity filtration
5. Scrape the soap back into the beaker and repeat the washing with each of the other
two 20 mL portions of concentrated sodium chloride.
6. Work the soap on a piece of dry filter paper or paper toweling to remove the last part
of the wash water. Weigh the soap.
7. Show the soap to your instructor for approval.
8. Add a piece of soap about the size of a pea to 5 mL of water in a test tube and with
your thumb sealing the top of the test tube, shake vigorously. Use pH paper to
determine the pH of the solution.
9. Wash your hand with a small piece of your soap.
10. Place 5 mL of distilled water in a test tube. Place 4 mL of tap water and 1 mL of
magnesium salt solution into another test tube. Put a small piece of soap into each test
tube and shake vigorously. Now compare the lathering in both tubes.
Look for a gelatinous precipitate in each test tube.
 The solids can be thrown in the trash.
 The solutions can be flushed down the drain with plenty of water.
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CHE 1400
Experiment 10
Saponification
Name(s)
Date
Laboratory Instructor
REPORT SHEET
I. SOAP WEIGHT:
________________ g
Instructor’s approval of soap: ________________
II. QUESTIONS
1. Explain why you washed the soap with salt solution rather than water.
2. What was the pH of soap in water?
3. How do you explain this pH?
4. Describe any observations about washing your hands with a small piece of the soap.
5. In step 10, which aqueous solution allows the soap to lather more freely?
6. Saponification of your butter sample produces mainly 3 soaps: sodium linoleate,
sodium oleate and sodium palmitate. Sodium linoleate results from saponification of
the ester of the fatty acid named linoleic acid. What are the chemical names of the
other 2 fatty acids?
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CHE 1400
7. The three most abundant soap molecules produced from your butter sample are
sodium linoleate, often written as C18H31O2Na; sodium oleate, written as
C18H33O2Na; and sodium palmitate, written as C16H31O2Na. Calculate the average
molecular weight of your butter sample.
Soap component
Formula
Portion of the
component in soap
Molar mass
(g/mol)
Molar mass x
Percentage
(g/mol)
sodium linoleate
C18H31O2Na
50%
_________
_________
sodium oleate
C18H33O2Na
30%
_________
_________
sodium palmitate
C16H31O2Na
20%
_________
_________
Average Molar mass of soap molecule =
_________ g/mol
The butter would be 3 soap molecules as fatty acids combined on a 3-carbon backbone.
To find the average molecular weight of butter, multiply the average soap by 3, subtract
69 for the 3 sodium atoms, which are removed, and add 41 for the
3-carbon backbone.
The 3–carbon backbone is glycerol without the 3 OH groups.
M butter  3  M soap  M Na   M CH2 CH CH2
 3  M soap  23  41
 3  M soap   28
M butter  _____________________ g/mol
O
R C O CH2
O
R C O CH +
O
R C O CH2
3 NaOH
Butter
O
R1 CONa
O
R2 CONa
O
R3 CONa
Soap
8. Calculate the expected yield of soap
nbutter 
nsoap
3

msoap  3  10.0 g 
msoap
mbutter

M butteer 3  M soap
 msoap  3  mbutter 
M soap
M butter
___________ g / mol
 ____________ g
___________ g / mol
9. How does your actual weight compare to the calculated yield? Does your soap still
have water in it?
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HO CH2
+
HO CH
HO CH2
Glycerol
CHE 1400
EXPERIMENT 11
Water analysis
OBJECTIVE
To learn how a water sample is collected and stored.
Perform some chemical analyses on a water sample.
Relates to chapter 14 of “Chemistry for changing times, 13th Ed.”.
APPARATUS AND CHEMICALS
conductivitimeter
150-mL beakers
100-mL, graduate cylinders
Buffer solutions
Vacuum filtration apparatus
Analytical balance
Hot plates
pH meter with electrodes
100 mL, 25-mL pipets
Weighing papers
Filter paper
INTRODUCTION
Water is the universal solvent and many contaminants (impurities) are easily dissolved
upon its contact. They may give water a bad taste, color, odor, or cloudy appearance
(turbidity), and cause hardness, corrosiveness, or staining. They can also damage growing
plants and transmit disease. At low levels, impurities generally are not harmful in water.
Removing all contaminants would be extremely expensive and in nearly all cases would
not provide greater protection of health. At high level (waste water) many of these
impurities are treated and removed or rendered harmless. Chemists are concerned with
the purity of water but regulatory agencies are concerned with setting standards to protect
the environment and public health. One mean of establishing and assuring the purity and
safety of water is to meet standards for various contaminants found in water. In this
project you be able to get some practice on how water is collected and stored and to
perform some routine chemical analysis.
A. Sample Collection and Storage
Site data should be recorded for all sampling locations. The information generally
required includes time, date, grid references of site, weather, temperature, method of
collection and information about any local activities that might influence the results.
Some of these data may only be applicable for certain classes of water samples.
A number of sampling devices are available for taking water samples from small ponds
and from different depths in large stratified lakes. The simplest system uses a weighted
bottle which is suspended at the required depth. The stopper is then removed by a sharp
pull on a separate line.
Water samples are especially subject to alteration in chemical composition due to
microbiological activity and chemical reactions. Heavily polluted waters can undergo
changes in composition within an hour of collection, and most natural waters are affected
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CHE 1400
to some degree. Some tests (particularly pH and dissolved gases) should, if possible, be
carried out in the field.
Glass sample containers are frequently recommended for storage since polythene vessels
can be porous to gaseous constituents and have been found to absorb phosphorus.
In general, to minimize possible bias of results caused by any changes occurring during
storage it is important to:
1 - Analyze the samples as soon as possible. Ensure that solution collectors at sites are
emptied regularly.
2 - Fill containers to exclude air.
3 - Keep sample cool, but do not freeze.
Physical preservation methods
Fine filtration
If the interest is only in the dissolved fraction then fine filtration, which removes many of
the microorganisms, can be applied. It will also remove fine mineral matter and any
traces of turbidity which could affect a later analytical stage.
An alternative approach is centrifuging which can be used to separate various particle
sizes.
Temperature reduction
Although some microbial activity appears to continue even at 0 °C the rapid cooling of
samples after collection is generally to be recommended.
Preliminary and general tests
For reasons given in the previous section tests on waters should be made as soon as
possible after sampling and in some cases in the field. This particularly applies to labile
and gaseous constituents.
Odor, turbidity, and color
Odor can serve as a guide to gross pollution of water. For example, characteristic odors
are associated with chlorination plants, untreated sewage and chemical industry effluents.
Color in water may be a true color due to dissolved material or an apparent color when
suspended material is present. The latter is quite common in natural waters, seen for
example when algal blooms impart a greenish tinge.
Turbidity may be used as estimate of undissolved substances in the sample. It is generally
measured by visual comparison with standards or photometrically, using a neophelometer
or spectrophotometer. Turbidity and color control light penetration in lake which in turn
affect phytoplankton population.
SOLID
 Total suspended solids (TSS)
The finer suspended matter in natural waters is usually of an organic nature
representing colloidal matter, which has been flocculated under the influence of
bacteria and protozoa. Inorganic suspended matter is chiefly restricted to siliceous
material resulting from the erosion of mineral soils.
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CHE 1400
 Total dissolved solids (TDS)
It is often convenient to determine the dissolved solids in the filtrate remaining from
the TSS determination.
 Total organic matter (TOM)
TOM is all organic matter that can be found in a given water sample
Alkalinity (and acidity)
The alkalinity of water is its capacity to neutralize a strong acid, and the values obtained
will depend on the pH of the titration end-point. In practice it is the bicarbonate,
carbonates and hydroxides in solution that largely determine the alkalinity although there
are minor contributions from silicates and other anions.
Total alkalinity is determined by titration to the equivalent point of carbonic acid which
occurs between pH 4.2 and 5.4 depending on the carbon dioxide content of water.
It has long been the practice in water analysis to determine solids as dissolved, suspended
and organic.
Conductivity
Conductivity is a property of water governed by the total ionic content. Although it is
non-specific and varies with the proportion of species presents, it is often measured,
because of its value in characterising waters. It expresses the resistance of 1 cm cube of
water to the passage of a current, usually at 25°C (specific resistance).
PROCEDURE
1. TSS determination
Filtrate 100 mL of your sample and then determine the weight of the solid in the filter
paper
2. TDS determination
Evaporate the filtrate (liquid) to a small volume (from 100 to 50 mL)
Transfer to a weighed 100 mL beaker for evaporation. Dry at 105°C to a constant
weight. Cool and weigh. Express the result in mg/L.
3. TOM determination
Transfer to a small pre-weighed evaporating beaker 50mL of your sample. Evaporate
to dryness at constant temperature and weigh beaker plus contents.
Ash in a muffle furnace, leaving at 500 °C for 1 hour. The loss in weight of the residue
gives the TOM in the sample. The method is only approximate and estimates of total
organic carbon are preferable when organic contents are low.
4. Alkalinity (and acidity)
Measure pH of water with a pH meter
If pH> 8.3 add 3-4 drops of phenolphthalein indicator and titrate against 0.01 M HCl.
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CHE 1400
5. Conductivity
Add the unfiltered sample into two beakers and bring to the required temperature
(preferably 25 °C) by immersion in a water bath. Immerse the electrodes in each
beaker that contain water samples. Record the conductivity in the second tube (having
used the first as a rinse).
Check the sample temperature just after immersion of the electrode.
Source:
A.P. Rowland & H.M. Grimshaw, in Chemical Analysis of Ecological Materials (S.T. Allen,
Editor) 2nd Edition, Oxford, UK: Blackwell, 1989, p. 62.
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CHE 1400
Experiment 11
Water Analysis (I) – Field Trip
Name(s)
Date
Laboratory Instructor
REPORT SHEET 1/2
Observations (in the field)
Water temperature
(oC)
pH
Conductivity γ
(µmhos/cm or µS/cm)
Amount of plants, or living
things you observe
Think about it and ANSWER these questions in the field
Areas with large amount of water insects and underwater plants are warmer because these living
things produce heat. Compare your results with other groups that collected water from areas with
and without plants/insects, is there a temperature difference?
Plants need CO2 from the water as food. When CO2 dissolves in water, an acid is produced, if
CO2 is removed from water by the plants. What would happen to the pH in areas where there is
plenty of plants? Does this agree with your observations of pH?
The conductivity is a measure of the amount of ions in water. Plants need CO 32-, NO3- and PO43ions and sunlight to produce food. Given the conductivity measurements that you have made on
the water that you took and combining it with the pH measurements, do you think that there can
be life in the area where you took the water?
The solubility of ionic compounds in water depends on the temperature. Higher temperatures
dissolve more of the compound to produce more ions. Does your measurement of conductivity
and temperature respectively, when compared to that of other groups, agree with this
observation?
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CHE 1400
Experiment 11
Water Analysis (II)
Name(s)
Date
Laboratory Instructor
REPORT SHEET 2/2
(Water analysis will be performed at the AUI Chemistry laboratory)
Part 1 - SOLID
Mass of total suspended solids (TSS) in mg
________________
Mass of total dissolved solids (TDS) in mg
________________
Mass of total organic matter (TOM) in mg
________________
Alkalinity (and acidity) pH
________________
If pH is larger than 8.3 continue otherwise skip to conductivity measurements
Part 2 – Determination of Alkalinity.
Concentration of standard HCl solution: 0.01mol/L
Trial 1
Trial 2
0
0
1
Buret reading, initial (mL)
2
Buret reading,final (mL)
2
Buret reading, initial (mL)
3
[HCl] (mol/L)
4
Volume of HCL used (mL)
_____________
_____________
5
Amount of HCL added (mol)
_____________
_____________
6
Amount of OH- in satd solution (mol)
_____________
_____________
7
Volume of sample solution (mL)
8
[OH-], equilibrium (mol/L)
Part 3 - Conductivity γ
Lab Manual
_____________
0.01
25.0
_____________
_____________
0.01
25.0
_____________
γ = _____________ µmhos/cm (or µS/cm)
61
APPENDIX I
CHE 1400
Common polyatomic ions
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CHE 1400
Common ions
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63
APPENDIX II
Lab Manual
CHE 1400
64