GENERAL CHEMISTRY
PHS 1015
Laboratory Manual
INDEPENDENCE COMMUNITY COLLEGE
Department of Chemistry
2014-2015
Table of Contents
Introduction: Laboratory Policies ...............................................................................................2
Experiment 1: Measurement & Significant Figures .................................................................3
Experiment 2: Chemical & Physical Properties ......................................................................23
Experiment 3: Identification Unknown Ion (Flame Test) .......................................................31
Experiment 4: Endothermic & Exothermic Chemical Reactions .........................................40
Experiment 5: Conductivity ......................................................................................................44
Experiment 6: Lewis Dot Structure & Molecular Models .....................................................49
Experiment 7: Stoichiometry .....................................................................................................62
Experiment 8: Alum Synthesis .................................................................................................66
Experiment 9: Alum Analysis ...................................................................................................72
Experiment 10: Identification Unknown Solution (Solubility) ...............................................79
Experiment 11: Acid-Base Titration (Handout) .....................................................................90
Experiment 12: Household Acids & Bases ..............................................................................91
Experiment 13: Boyles Law & Guy-Lussac ................................................................................
Boyle’s Law .......................................................................................................................96
Guy-Lussac ......................................................................................................................100
Experiment 14: Molar Volume (Handout) ............................................................................105
Experiment 15: Heat of Fusion ...............................................................................................106
Experiment 16: Evaporation & Intermolecular Interactions ..............................................110
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Introduction:
Laboratory Policies:
Laboratory part of science course are integral in the understanding of science. Science is about
exploring, observing, and interpreting experiments and chemistry is no different. 40% of the
grade for this course is based on the laboratory part of the class. There are 16 experiments to be
completed. The schedule of the experiments is shown in the syllabus and rarely changes. Be
prepared to conduct any experiment scheduled. There are no make-up days for missed
laboratory experiments. If you are going to miss a laboratory experiment, you can plan to
come into another class conducting the same experiment. You need to email the professor if
you are going to miss an experiment.
Proper laboratory attire is required, if you do not have proper laboratory attire you will be asked
to leave until you have proper attire. If you do not return to class and do not come to another
class time to make up the experiment, your will receive a zero (0) for the laboratory experiment
(both results and question sections)
Grading:
The laboratory experiments will be broken into two parts. You will receive a maximum of 15
points for the results part of the lab and a maximum of 10 points for the questions. Some
experiments have pre- and post- laboratory question and some have only post- laboratory
question. The only parts due are the results and any questions (only parts you write on). The
introduction and procedure portion of the laboratory book are for you to keep. These sections
will come in handy when you do the laboratory practical. The laboratory experiments are due
the next class date, late experiments will not be accepted without a valid excuse. An email in
advance is required to turn anything in late.
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Experiment 1: Measurement & Significant Figures
Purpose:
What are some of the techniques used to make scientific measurements in the laboratory? How
does one determine the uncertainty in any measurement? What is meant by “significant
figures”?
Introduction:
Conducting experiments or investigations are an important activity in all science classes. These
experiments or investigations require measurements to be recorded and analyzed. Every
measurement contains some amount of uncertainty due to the accuracy (or inaccuracy) in the
device or technique used to make the measurement. Some of the instruments or techniques used
to make measurements have large uncertainties and some have small uncertainties. Recognizing
how to record the measurements and determining the uncertainty of each measurement is your
goal for this experiment.
Measuring Mass:
The measurement of mass is conducted using a balance, because historically the mass of the
sample was determined by balancing it against a standard or known mass. Figure 1 & 2 are
examples of different scales or balances used to measure the mass of an object or sample. Figure
2 contains examples of the digital balances used in many chemistry courses across the country.
Figure 1. Triple beam Balance.
Figure 2. Example of spring-balance & electronic
digital balances.
Here are some guidelines for operating the electronic balance to successfully obtain a mass of
any sample:
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1. Be certain the balance has been “zeroed” (meaning the digital readout is “0”) before you
place any sample on the balance.
2. NEVER weigh chemicals directly on the balance pan. A suitable container or weighing
paper should always be used. Often, you will weigh the sample container empty (to get
an initial weight) and then weigh the contain containing the sample or chemical. Subtract
the two measurements to determine the mass of the sample or chemical.
3. Be certain that air currents are not effecting your measurements. Someone walking past
your balance or bumping the table that the balances are sitting on can effect a
measurement. When making measurements to 0.001 g, one should always close the door
or lid to the balance.
4. NEVER measure the mass of hot or warm objects, since the temperature difference will
change the density of the air surrounding the balance and give an inaccurate
measurement.
5. Record your measurement, to the proper number of significant figures. The balances
used measure to the 0.001 g or the 0.01 g.
6. Once your measurements have been made, close the door to the balance (if necessary),
and be sure the balance registers zero. Be sure to clean the balance if any chemical have
been spilled.
7. Be gentle with the balances, these are sensitive, delicate instruments and like an person,
responds best when treated properly.
When recording a mass, the digit furthest to the right of the number recorded is the uncertainty of
the measurement. For example, a mass of a penny was measured to be 1.57 g. The uncertainty
of the measurement is 1.57 ± 0.01 g. The uncertainty recognizes the + 0.01 g or -0.01 g in the
measurement. The uncertainty or last significant figure is always recognized as the last digit of
the measurement. The number of significant figures is defined as the number of digits recorded
in the measurement. If you recorded 1.57 g as the mass of the penny, then the number of
significant figures is 3 significant figures.
Measuring Volume:
Volume is another common measurement made in the laboratory. The most common type of
sample to measure the volume is a sample in the liquid state of matter. There are many different
devices that can be used to measure the volume of a sample, depending on the accuracy required
in your measurement. Typically volume is measured in mL but the SI unit for volume is L. It is
helpful to know:
1 mL = 1 cc = 1 cm3
A. Beakers & Erlenmeyer Flasks:
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Using a beakers and Erlenmeyer Flasks to measure volume is the least accurate of all the devices
used to measure volume. The measurement can be off by as much as 10%. For a 250 mL
beaker, the measurement could be ± 25 mL in error. Figure 3 is an example of breakers and
Figure 4 is an example of Erlenmeyer flasks.
Figure 3. 400 mL Breaker
Figure 4. Various sizes of Erlenmeyer Flasks
B. Graduated Cylinder:
Using a graduated cylinder to measure volume somewhat more precise, the error is less than 1%.
Most of the volume measurements conducted in Chemistry are conducted with graduated
cylinders. Examples of graduated cylinders are shown in Figure 5.
Figure 5. Various sizes of Graduated Cylinders.
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Depending on the size of the graduated cylinder, the uncertainty can be from ± 0.1 mL to ± 0.01
mL.
C. Pipets & Volumetric Flasks:
Pipets and volumetric flasks are calibrated to contain specific volumes of liquid. When filled to
the proper level, the volume contained is accurate to ± 0.01 mL. When using the pipet, proper
technique must be utilized for accurate measurements. Examples of volumetric flasks are shown
in Figure 6 and an example of a pipet is shown in Figure 7. The following steps make for the
proper use of a pipet:
1. The pipet must be clean and dried. When the pipet is draining, there should be no drops
left on the pipet walls.
2. Always use a pipet bulb or filler, NO mouth pipetting (pulling liquid into the pipet by
sucking on the pipet like a straw).
3. Equilibrate the pipet by drawing some of the liquid into the pipet, remove the bulb or
filler. Tilt the pipet so it is parallel to the floor and coat the inside wall of the pipet with
the solution. Allow the solution to drain out the tip of the pipet.
4. Use the pipet bulb or filler to draw the liquid into the pipet. Continue to draw the liquid
into the pipet until the meniscus is 1 – 2 cm above the calibration mark. Remove the
pipet bulb or filler and place a finger or thumb over the opening of the stem. Let the
liquid flow out of the pipet, into a waste beaker until the level of the meniscus is touching
the top of the calibration mark. Place the tip of the pipet into the receiving flask, with the
tip of the pipet touching the wall of the flask. The pipet needs to be straight up and down
so the receiving flask will be tilted. Allow the liquid to drain from the pipet while the tip
continues to touch the wall of the flask. Once the liquid has drained, separate the pipet
from the receiving flask. There is a small amount of liquid in the tip of the pipet, DO
NOT blow this out into the receiving flask. When properly used, the pipet is designed to
deliver the exact volume of liquid (± 0.01 mL) into the receiving flask.
Figure 6. Various sizes of Volumetric Flasks.
Figure 7. A Volumetric Pipet.
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D. Burets:
A buret (Figure 8) is used to add precisely measured volumes of liquid. This is different from
using a pipet to deliver the precisely measured volume of liquid in that unlike with a pipet the
volume of liquid can vary depending on the amount required. When using a pipet the volume of
liquid is dependent on the volume of the pipet, the buret has a valve that can be opened and
closed (Figure 9) to allow the user to add the liquid until a visual change has been achieved.
Figure 8. A Buret.
Figure 9. Buret valve (stopcock) open & closed.
As when using the pipet, working with the buret requires a particular technique to determine the
volume of the sample to the ± 0.01 mL. The following procedure should be helpful:
1. Equilibrate the pipet using the same procedure to equilibrate the pipet. Be sure to open
the stopcock to allow some of the solution to drain through the stopcock and tip. When
the solution has been removed from the buret, add enough solution to fill the buret past
the 0.00 mL mark.
2. Open the stopcock and tap the end with the stopcock to shake loose and air bubbles in the
stopcock/tip of the buret. Close the stopcock and determine the initial volume. If the
meniscus is still above the 0.00 mL mark, drain enough of the solution so the meniscus is
touching the top of the 0.00 mL mark. If the meniscus is below the 0.00 mL mark, read
the volume as the initial volume. Be sure to view the meniscus at eye level as shown in
Figure 10. An example of an initial and final volume measurement is shown in Figure
11.
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Figure 10. Proper method to read the meniscus of the buret.
Figure 11. The initial reading and final reading of a measurement using a buret.
3. Notice in the measurement with the buret, the hundredth place is an estimate. For the
initial reading, starting above the meniscus, the measurement is between the 9 and 10. So
this means the measurement will be starting with 9 mL. Counting the graduations
between 9 and 10, one views the meniscus to be between the 6th and 7th mark. This leads
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to initial measurement to start 9.6 mL. The meniscus is closer to the 0.6 mark than the
0.7 mark, estimating the meniscus one can say the meniscus is about 0.03 and therefore
the initial measurement of the volume is 9.63 mL. Read the final reading to the
hundredth place and verify the final reading to be 24.16 mL. To determine the volume
dispensed, subtract the initial reading from the final reading to get the volume used.
This laboratory experiment correlates with the section on measurements & mathematics in your
lecture book. Students are expected to read the introduction and appropriate lecture material to
answer the Pre-Lab Questions.
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Pre-Lab Questions:
1. What is the theoretical value for the density of water?
__________________________________
2. In one of the experiments students are asked to determine the density of 3 solid objects.
Describe the method you can use to determine the volume of the solid. You might want
to discuss this with your laboratory partners.
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
______________________________________________________
3. When using a buret to measure the volume of a liquid, the measurement should be made
to the nearest ________ mL.
4. The quantity of mass of 5 pennies was measured to be 12.54876 g. Which number in the
measurement has an uncertainty?
__________________________________________________
5. How many significant figures are contained in the measurement of the 5 pennies?
___________
6. What is the average mass of each penny? Be sure to consider the uncertainty in the
measurement and the number of significant figures.
______________________________________________
7. What is the uncertainty of each of the following measurements and how many significant
figures are contained in each of the measurements?
a. 840.3 mL ± _______mL
_______ Significant Figures
b. 0.0486 kg ± ______kg
_______ Significant Figures
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c. 4.184 L ± ______L
d. 5.000 mm ± ______mm
e. 760 mm Hg ± ______ mm Hg
_______ Significant Figures
_______ Significant Figures
_______ Significant Figures
8. Round each of the measurements in problem 4 to contain the number of significant
figures indicated in the parentheses.
a. (3) ______ mL
b. (2) ______ kg
c. (1) ______ L
d. (3) ______ mm
e. (2) ______ mm Hg
8. Carry out the following calculations and write the result using the correct number of
significant figures. All the numbers should be treated as measured quantities.
a. 0.847 x 2.10 = ______
b. 79.06/0.21 = ______
c. 17.2 x 1.000 = ______
d. 1.002/84.01 = ______
e. 0.748 + 12.02 + 0.0001 = ______
f. 12.54 – (12.874 + 12.0) = ______
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Procedure:
A. Distance measurements in cm & mm.
Using the ruler in your lab drawer, measure the length and width of shapes A, B, C, & D found
on page 10. Record the measurements in centimeters (cm) and millimeters (mm) on the Data &
Observations Sheet. Calculate the area and parameter of the shapes using the correct number of
significant figures.
B. Mass measurements.
Measure the mass of the three objects you receive from your instructor. Press the tare button to
zero the balance and place a container to hold the sample on the balance pan. Record the mass of
the empty sample container on the Data & Observations Sheet. Using the crucible tongs, place
the sample into the sample container. Record the mass of the sample and sample container.
Calculate the mass of the sample by subtracting the mass of the empty sample container from the
mass of the sample container with the sample. The difference is the mass of the sample. Repeat
for the remaining samples. Record the mass using the balance measuring to ± 0.01 g and ± 0.001
g on the Data Sheet.
C. Volume measurements.
Measure a volume of water using a 100 mL beaker, 50 mL graduated cylinder, and the 10 mL
graduated cylinder. Weigh the empty 100 mL beaker, 50 mL graduated cylinder, and the 10 mL
graduated cylinder and record the weight of each on the Data & Observations Sheet. Using a
Beral pipet, add 100 drops of water to each container and measure the volume. Remember each
of the containers will have a different uncertainty. Weigh the containers with the water in them;
determine the mass of the water in each of the containers. The density of an object or substance
is the mass of the object or sample, and the volume that object or sample occupies. Based on the
data collected, determine the density of water. Each of the containers used to determine the
volume will have a different volume with different uncertainty or significant figures. Determine
the percent error using the formula:
Theoretical value – Actual value x 100% = Percent Error
Theoretical value
The theoretical value for the density of water is 1.000 g/mL.
D. Determine the Density of Solids.
Using the mass measurements from Part B, determine the density of the 3 objects and using
Table 1, determine the unknown substances. You will need to determine the volume of the
object as accurately as possible so to get the most number of significant figures for your
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calculations. Notice the number of significant figures used in Table 1 and plan your experiment
accordingly.
Table 1. Density of Possible Unknown Solids
Element
Chemical Symbol
Density (g/mL)
Element
Chemical Symbol
Density (g/mL)
Zinc
Zn
7.140
Copper
Cu
8.960
Tin
Sn
7.265
Iron
Fe
7.874
Lead
Pb
11.35
Aluminum
Al
2.702
Nickel
Ni
8.908
Chromium
Cr
7.190
Silver
Ag
10.49
Gold
Au
19.32
E. Density Determination of Unknown Liquid.
Record the unknown identification code of the liquid unknown. Mark 3-50 mL beakers #1, #2,
#3 and measure the mass of each empty beaker. Using the buret, determine the volume of the
unknown liquids. Determine the mass of the each of the volumes of unknown liquid samples.
Using Table 2, determine the unknown composition of the liquid.
Table 2. Density of Possible Unknown Liquids
Unknown Solution
Density (g/mL)
Unknown Solution
Density (g/mL)
Saturated NaCl
1.200
90% Isopropyl Alcohol
0.812
Saturated Sucrose
1.343
70% Isopropyl Alcohol
0.845
Vinegar
1.101
50% Isopropyl Alcohol
0.898
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Print this page!
B
A
C
D
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Data & Observations:
A. Distance measurements in cm & mm.
Table 3. Measurement of shapes in cm.
Shape
Measured Length (cm)
Measured Width (cm)
A
B
C
D
Table 4. Parameter of shapes in cm.
Shape
Calculated Parameter (unrounded) (cm)
Calculated Parameter (rounded)
(cm)
A
B
C
D
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Table 5. Area of shapes in cm2.
Shape
Calculated Area (unrounded) (cm2)
Calculated Area (rounded) (cm2)
A
B
C
D
B. Mass of Object Measurements.
Identification code of Object 1 __________
Identification code of Object 2 __________
Identification code of Object 3 __________
Mass of Objects using 0.01 g Electronic Balance
Mass of Object 1 and container __________ g
Mass of container empty _______________ g
Mass of Object 1 _____________________ g
Mass of Object 2 and container __________ g
Mass of container empty _______________ g
Mass of Object 2 _____________________ g
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Mass of Object 3 and container __________ g
Mass of container empty _______________ g
Mass of Object 3 _____________________ g
Mass of Objects using 0.001 g Electronic Balance
Mass of Object 1 and container __________ g
Mass of container empty _______________ g
Mass of Object 1 _____________________ g
Mass of Object 2 and container __________ g
Mass of container empty _______________ g
Mass of Object 2 _____________________ g
Mass of Object 3 and container __________ g
Mass of container empty _______________ g
Mass of Object 3 _____________________ g
C. Volume Measurements.
Mass & Volume Measurements, Density Determination using 100 mL
Beaker
Mass of 100 mL beaker with water __________ g
Mass of 100 mL beaker empty ______________ g
Mass of water in 100 mL beaker ____________ g
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Volume Measurement using 100 mL beaker __________ mL
Density of water using 100 mL beaker ______________g/mL
Percent Error of Density of water using 100 mL beaker __________ %
Mass & Volume Measurements, Density Determination using 50 mL
Graduated Cylinder
Mass of 50 mL Graduated Cylinder with water __________ g
Mass of 50 mL Graduated Cylinder empty ______________ g
Mass of water in 50 mL Graduated Cylinder ____________ g
Volume Measurement using 50 mL Graduated Cylinder __________ mL
Density of water using 50 mL Graduated Cylinder ______________g/mL
Percent Error of Density of water using 50 mL Graduated Cylinder __________ %
Mass & Volume Measurements, Density Determination using 10 mL
Graduated Cylinder
Mass of 10 mL Graduated Cylinder with water __________ g
Mass of 10 mL Graduated Cylinder empty ______________ g
Mass of water in 10 mL Graduated Cylinder ____________ g
Volume Measurement using 10 mL Graduated Cylinder __________ mL
Density of water using 10 mL Graduated Cylinder ______________g/mL
Percent Error of Density of water using 10 mL Graduated Cylinder __________ %
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
D. Determine the Density of Objects.
Density of Object 1
Identification code of Object 1 ___________
Mass of Object 1 __________ g
Volume of Object 1 __________ mL
Density of Object 1 __________ g/mL
Element Object 1 __________
Chemical Symbol Object 1 __________
Density of Object 2
Identification code of Object 2 ___________
Mass of Object 2 __________ g
Volume of Object 2 __________ mL
Density of Object 2 __________ g/mL
Element Object 2 __________
Chemical Symbol Object 2 __________
Density of Object 3
Identification code of Object 3 ___________
Mass of Object 3 __________ g
Volume of Object 3 __________ mL
Density of Object 3 __________ g/mL
Element Object 3 __________
Chemical Symbol Object 3 __________
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E. Density of Unknown Liquid
Density of Unknown Sample 1
Unknown Identification Code Sample _______
Mass of Sample 1 __________ g
Volume of Sample 1 __________ mL
Density of Sample 1 __________ g/mL
Composition of Sample 1 __________
Density of Unknown Sample 2
Unknown Identification Code Sample _______
Mass of Sample 2 __________ g
Volume of Sample 2 __________ mL
Density of Sample 2 __________ g/mL
Composition of Sample 2 __________
Density of Unknown Sample 3
Unknown Identification Code Sample _______
Mass of Sample 3 __________ g
Volume of Sample 3 __________ mL
Density of Sample 3 __________ g/mL
Composition of Sample 3 __________
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Post-Lab Questions:
1. Oil (0.784 g/mL) and vinegar (1.101 g/mL) are immiscible. If you vigorously shake a
bottle of Italian salad dressing (primarily oil and vinegar with herbs and seasoning) and
let the bottle sit on the counter for a few minutes, what will happen to the Italian
dressing? Explain.
2. If you mistakenly did not tare the balance before you measured the mass of the graduated
cylinder with water, would your calculation of the density of the water be greater or less
than the actual density of water? You notice the balance has a reading of 0.03 g.
Explain.
3. Were the volume measurements made with the 10 mL graduated cylinder more, less, or
equally precise as the volume measurements made with the 50 mL graduated cylinder?
Explain.
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4. The density of water at 0 oC is 0.9999 g/mL, and the density of ice at 0 OC is 0.9168
g/mL. Based upon these densities, does ice float or sink in water? Explain.
5. Does water contract or expand when it freezes? Explain.
6. Suppose you have a 400 mL beaker with ice filled to the 250 mL mark (assume no air
pockets). What is the volume of liquid water formed when the ice melts? Show your
work.
7. The density of cod liver oil is 0.92 g mL-1, what is the volume occupied by a 1 lb bottle of
cod liver oil? Show your work.
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Experiment 2: Chemical & Physical Properties
Purpose:
Can a student differentiate between chemical change/properties and physical change/properties?
Introduction:
The properties of matter can be divided into two classes, physical properties and chemical
properties. Physical properties are those properties that can be observed or measure without
any change to the chemical nature of the substance. The best way to think about a physical
property is that a physical property describes the substance. The color, density, viscosity,
solubility, melting point, and boiling point are all physical properties of a substance. Chemical
properties describe a change in the chemical composition of the substance under specific
conditions.
Physical change is a change in the physical nature of the substance. Evaporation is the change
of a substance from the liquid state of matter to the gas state of matter. The substance did not
change in the chemical nature or composition, only in its state of matter.
Chemical change is a change in the chemical composition of the substance. Fermentation is the
process yeast use to convert sugar (glucose) in to ethyl alcohol, carbon dioxide, water, and
energy. Evidence of a chemical change include: color change, gas production, formation of a
precipitate (insoluble product), and release or absorption of heat. When heat is released from a
chemical reaction, the reaction is termed exothermic (giving off heat) and when the chemical
reaction absorbs heat, the chemical reaction is termed endothermic (absorbing heat).
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Pre-Lab Questions:
1. What are the three states of matter?
a. Describe each state of matter on the microscopic level.
b. Describe each state of matter on the marcoscopic level.
2. Identify each of the following as a chemical change/property or physical change/property.
Explain your choice.
a. an iron nail rusting
b. a test tube shattering when dropped on the floor
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c. paper shredded in a paper shredder
d. a silver fork losing its luster while sitting in the drawer
e. heated copper metal leaving a black residue when heated
f. sugar becoming a gel on a humid day
3. What is an observation? Give an example of an observation.
4. What is inference? Give an example of inference.
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Procedure:
A. Investigating properties.
Obtain the group of the five unknown coded, sealed vials, DO NOT open the vials. Based on the
appearance (physical properties) of the material in each vial and using the CRC Handbook of
Chemistry and Physics, determine which vial contains the carbon, iron, lead, sulfur, and
aluminum.
B. Chemical and physical properties of a substance.
Obtain the vial containing the white, crystalline unknowns. Based on the appearance (physical
properties) of the material in each vial and using the CRC Handbook of Chemistry and Physics,
determine which vial contains the sodium chloride (table salt) and sucrose (table sugar).
Determine if either or both the substances are soluble in water (dissolve in water). Place one of
the white, crystalline unknowns into an evaporating dish and the other white crystalline unknown
into another evaporating dish. Heat each evaporating dish with a Bunsen burner for 5 minutes.
Record your observations and inferences. Hint: You might want to fill out the data table 2 with
the physical properties from the CRC Handbook of Chemistry and Physics.
C. Chemical and physical properties of a substance (Part 2).
Obtain the vial containing the next four known substances. Record the physical appearance of
each substance. Record the solubility in water (dissolve in water) of each of the substances.
Place the paraffin, sulfur, malonic acid, and copper(II) sulfate pentahydrate into separate
evaporating dishes, record the appearance before and after heating.
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Data & Observations:
A. Investigating Properties.
CRC Handbook of Chemistry and Physics Edition: ________
Table 1. Physical Properties & Identification of Five Unknowns.
Unknown ID Code
1
2
3
4
5
Name of Unknown
Physical State
Molar Mass (g/mol)
Density (g/mL)
Melting Point (oC)
Boiling Point (oC)
Solubility in water at 0 oC, g/100
cc
Element Name
Chemical Symbol
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
B. Chemical and physical properties of a substance.
Table 2. Physical Properties & Identification of Table Salt & Table Sugar.
Unknown ID Code
1
2
Physical State
Molar Mass (g/mol)
Density (g/mL)
Melting Point (oC)
Boiling Point (oC)
Solubility in water at 0 oC, g/100 cc
Appearance before heating
Appearance after heating
Chemical Name
Chemical Symbol
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
C. Chemical and physical properties of a substance (Part 2).
Substance
Physical Appearance
Soluble in
water?
(Y/N)
Appearance
before heating
Appearance
after cooling
Paraffin
Sulfur
Malonic Acid
Copper(II)
Sulfate
pentahydrate
1. Description of the odor of heated malonic acid:
2. Observation of heated copper(II) sulfate pentahydrate before heating, after cooling, after
adding a few drops of water on the sample:
3. What can be inferred about the copper(II) sulfate:
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Post-Lab Questions:
1. What is another observation or physical property that can be made in the identification of
table salt & table sugar part of the lab? What would the inference be of this observation?
2. One warm summer evening, a group of students decide to go camping. For dinner they
cut up beef, potatoes, and onions into cubes, wrap everything into packets using
aluminum foil, and place the packets on the campfire. After 40 minutes, the contents of
the aluminum foil packets were opened revealing a brown, well-cooked meal. For desert,
the students ate crispy, golden brown, roasted marshmallows with graham crackers and
chocolate. One of the students boiled water for hot cocoa. During the evening, the
temperature dropped and the grass became damp with the evening dew.
a. List the physical changes described in the story.
b. List the chemical changes described in the story.
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Experiment 3: Identification Unknown Ion (Flame Test)
Purpose:
Using a flame test, can a student identify unknown ions?
Introduction:
Understanding the Bohr model of the atom is the basis for understanding the ability for one to
identify unknown ions using a technique call a flame test. Each element emits unique colored
bands combined to give a unique color to that element. In the Bohr model (Figure 1), the
electrons exist in energy levels around the nucleus. When an electron absorbs energy, the
electron moves to a higher energy level (further away from the nucleus). This situation can only
occur for a limited time and the electron must return to a lower energy level. When the electron
returns to a lower energy, the difference in the energies is released as a photon.
Humans visualize the energy in a photon as visible light in the electromagnetic spectrum. The
colored light seen corresponds to an energy difference between the electron in a higher energy
level and the electron transitioning to a lower energy level (Figure 2). The colored light seen
also corresponds to a specific wavelength. A wavelength (Figure 3) is defined as the distance
from the peak to peak or trough to trough. Typical units for wavelength is the unit used to
measure distance, nanometer, or another unit is the angstrom (Å). The wavelength is related to
the number of wave’s passing over a point, in an amount of time. This is called the frequency,
hertz, or cycles per second. The relationship between wavelength and frequency () is shown
in equation 1. When the electron transitions from a higher energy level to a lower energy level, a
photon is released. The photon corresponds to a specific amount of energy and related to a
specific wavelength () by the equation 2.
 = c/
Equation 1
c = the speed of light (3.00 x 108 m/s)
 = frequency
 = wavelength
E = h
Equation 2
E = Energy
h = Planck’s constant (6.626 x 10-34 J.s)
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Figure 1. Bohr Model of the atom. n = 1, n = 2, n = 3, are the different energy levels.
Figure 2. Electron transition from higher energy level to lower energy level.
32 | P a g e
Figure 3. Determining a wavelength measurement.
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Pre-Lab Questions:
1. Pretend your lab partner skipped class and you need to teach them how to light the
Bunsen burner. Write out the procedure to light the Bunsen burner as you would explain
the processes to your lab partner.
2. Using equations 1 & 2, rearrange equation 2 to solve for wavelength and substitute the
rearranged equation 2 into equation 1. The new equation should show the relationship
between the energy of the photon and the frequency of the photon.
3. What color flame indicates a “cool” flame?
4. What color flame indicates a “hot” flame?
34 | P a g e
5. Where is the flame the hottest?
6. What part of the Bunsen burner controls the gas flow?
7. What part of the Bunsen burner controls the air supply?
8. What is the method a student should use when the characteristic color of an ion is similar
to another ion?
35 | P a g e
Procedure:
A. Flame test of known ions.
Adjust the flame on the Bunsen burner to a very hot flame. Dip one end of a cotton swab into a
solution containing a known ion. Place the tip of the cotton swab with the solution on it into the
hot portion of the flame. You may need to repeat the flame test a couple of times to determine
the color of the flame. Record the corresponding ion and the characteristic color of the flame
emitted when the tip of the cotton swab was placed into the flame. The one side of the cotton
swab should only be used to one solution. Do not dip the cotton swab into another solution! Use
one cotton swab for each known solution. When you have completed the flame test for the
known ions, place the cotton swabs into a beaker with some water to make sure you extinguish
any hot embers on the cotton swab.
B. Flame test of unknown ions.
Each group is assigned two unknown solutions. Repeat the procedure to conduct the flame test
of the known ion solutions. Record the unknown identification codes, the corresponding
characteristic flame color, and the identity of the unknown ion. If you are uncertain of the
identity of the unknown ion, conduct the flame test of the suspected known ion solution and the
unknown ion solution at the same time. Sometimes the differences in the characteristic colors
are very slight and the only way to correctly identify the unknown ion solution is to conduct the
flame test side by side.
36 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Data & Observations:
A. Flame test of known ion solutions.
Table 1. Known ions and Characteristic color.
Chemical
Symbol
Element
Description of color
Na
K
Li
Sr
Ba
Cu
Table 2. Unknown ions identity and Characteristic color.
Unknown ID
Code
Description of Color
Unknown Identity
37 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Post-Lab Questions:
1. Name the colors of light in the electromagnetic spectrum from highest energy.
______
______
______
______
______
______
______
2. Red light has a longer wavelength than blue light.
c. Which of these colors has the higher frequency?
d. Which of these colors has the higher energy?
3. What are some of inaccuracies that may be involved when using the flame tests for
identification purposes? Explain.
38 | P a g e
4. The alkali metals cesium & rubidium were can be identified based on their characteristic
flame colors. Cesium was named after the sky and rubidium was named after the gem
color. What characteristic colors of light do you think are emitted when these metals are
heated in a flame?
5. When a glass rod is heated, the characteristic color of flame is a bright yellow flame.
Based on the flame test experiment, what metal ion do you think is predominately found
in glass?
6. What experiment would you conduct to verify the characteristic flame color was due to
theh metal cation and not the anion? Explain.
39 | P a g e
Experiment 4: Endothermic & Exothermic Chemical Reactions
Purpose:
Many chemical reactions give off energy. Chemical reactions that release energy are called
exothermic reactions. Some chemical reactions absorb energy and are called endothermic
reactions. You will study one exothermic and one endothermic reaction in this experiment.
In Part I, you will study the reaction between citric acid solution and baking soda. An equation
for the reaction is:
H3C6H5O7(aq) + 3 NaHCO3(s) 
 3 CO2(g) + 3 H2O(aq) + Na3C6H5O7(aq)
In Part II, you will study the reaction between magnesium metal and hydrochloric acid. An
equation for this reaction is:
Mg(s) + 2 HCl(aq) 
 H2(g) + MgCl2(aq)
Another objective of this experiment is for you to become familiar with using LabQuest. In this
experiment, you will use the program to collect and display data as a graph or list, to examine
your experimental data values on a graph, and to print graphs and data lists.
In this experiment, you will study one exothermic and one endothermic reaction.
Figure 1
Procedure:
1. Obtain and wear goggles.
Part I Citric Acid plus Baking Soda
2. Connect the Temperature Probe to LabQuest.
40 | P a g e
3. Place a Styrofoam cup into a 250 mL beaker as shown in Figure 1. Measure out 30 mL of
citric acid solution into the Styrofoam cup. Place the Temperature Probe into the citric acid
solution. Note: The Temperature Probe must be in the citric acid solution for at least 30
seconds before doing Step 5.
4. Weigh out 10.0 g of solid baking soda on a piece of weighing paper.
5. You are now ready to begin collecting data.
a. Record the temperature of the citric acid solution.
b. Add the baking soda to the citric acid solution. Gently stir the solution with the
Temperature Probe to ensure good mixing.
c. Temperature readings (in °C) can also be monitored in a display box.
d. Record the temperature of the solution when the temperature of the solution stabilizes.
6.
Dispose of the reaction products as directed by your instructor. Rinse the Temperature
Probe.
Part II Hydrochloric Acid Plus Magnesium
7. Measure out 30 mL of HCl solution into the Styrofoam cup. Place the Temperature Probe
into the HCl solution. Note: The Temperature Probe must be in the HCl solution for at least
30 seconds before doing Step 9.
8. Obtain a piece of magnesium metal from your instructor.
9. You are now ready to begin collecting data.
a. Record the temperature of the citric acid solution.
b. Add the magnesium strip to the HCl solution. Gently stir the solution with the
Temperature Probe to ensure good mixing.
c. Temperature readings (in °C) can also be monitored in a display box.
d. Record the temperature of the solution when the temperature of the solution stabilizes.
Caution: Do not breathe the vapors.
10. Dispose of the reaction products as directed by your instructor. Rinse the Temperature Probe.
41 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
DATA TABLE
Part I
Part II
Final temperature, t2
°C
°C
Initial temperature, t1
°C
°C
Temperature change, t
°C
°C
42 | P a g e
PROCESSING THE DATA
1. Calculate the temperature change, t, for each reaction by subtracting the initial temperature,
t1, from the final temperature, t2 (t = t2 – t1).
2. Tell which reaction is exothermic. Explain.
3. Which reaction had a negative t value? Is the reaction endothermic or exothermic? Explain.
4. For any chemical reaction, describe three ways you could tell a chemical reaction was taking
place.
5. Which reaction took place at a greater rate? Explain your answer.
43 | P a g e
Experiment 5: Conductivity
Purpose:
Students investigate conductivity of aqueous solution using a number of dissolved solutes. What
are the criteria for a solution’s ability to conduct electricity? What types of compounds, when
dissolved in water, conduct electricity? Why are some compounds better at conducting
electricity than others?
Introduction:
Electricity is the flow of negative charged particles. The ability for electricity to flow in an
aqueous solution is measured by the conductivity of the solution. An aqueous solution that
conducts electricity must have ions or be an ionic compound. When this type of compound
dissolves in water, the ions separate into cations and anions. This is what allows the negative
charged particles of electricity flow. Solutions that conduct electricity are called electrolytic
solutions. Covalent compounds do not form ions and therefore cannot conduct electricity.
These form non-electrolytic solutions and do not conduct electricity.
The periodic table can be used to predict whether a compound is ionic or covalent and how
much the compound will conduct electricity based on the location, of the elements, on the
periodic table. This principle is called the “trends of the periodic table”. This experiment
investigates this principle and the role of the periodic table in predicting the physical and
chemical properties of the compounds.
44 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Pre-Lab Questions:
1. Define an atom.
2. Define an ion.
3. Define a polyatomic ion.
4. Do all compounds contain ions? Explain. (What other types of compounds are there?)
5. Define electrolytic solutions.
6. Define non-electrolytic solutions.
45 | P a g e
Procedure:
A. Taking Measurements with the Conductivity Probe.
You should not need to calibrate the conductivity probe. Connect the conductivity probe to the
LabQuest data collection device. To make a measurement, into 400 mL beaker (waste beaker),
rinse the tip of the conductivity probe using the squirt bottle of distilled water and blot dry the tip
of the probe using a Kimwipe (lab tissue) to remove any water droplets. Insert the tip of the
conductivity probe into the sample to be tested. Important: Be sure the electrode surface is
completely submerged in the liquid. While gently swirling the probe, wait for the reading to
stabilize, this should take about 5 to 10 seconds. Rinse the end of the probe with the distilled
water over the waste beaker before taking another measurement of storing the probe.
Obtain a set of vial containing 0.01 M aqueous solutions for each group to test. Using the
procedure above to take measurements using the conductivity probe, record the conductivity of
each of the aqueous solutions.
46 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Data & Observations:
A. Taking Measurements with the Conductivity Probe.
Table 1. Conductivity of Aqueous Solutions.
0.01 M Solution
Conductivity (mg/L)
0.01 M Solution
Tap Water
NaNO3
Distilled Water
KNO3
LiCl
Mg(NO3)2
NaCl
Ca(NO3)2
KCl
Sr(NO3)2
Sucrose
Ba(NO3)2
Glucose
MgCl2
Na2SO4
CaCl2
Na3PO4
SrCl2
K3PO4
BaCl2
Al(NO3)3
NiCl2
H2O2
CuCl2
CH3OH
ZnCl2
C2H5OH
MgSO4
C3H7OH
CaSO4
Conductivity (mg/L)
47 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
1. Which solutions were electrolytic?
2. Which solutions were non-electrolytic?
3. Using the periodic table, locate the elements of the compounds for the solutions that were
electrolytic. Do you see a pattern for electrolytic solutions? Explain.
4. Using the periodic table, locate the elements of the compounds for the solutions that were
non-electrolytic. Do you see a pattern for non-electrolytic solutions? Explain.
5. For the electrolytic solutions, describe if there is a pattern seen in the numerical readings?
Explain the reason for these patterns.
48 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Post-Lab Questions:
1. How does the ionic and covalent bond differ?
2. Choose one of the electrolytic solutions and draw a view, at the molecular level, of the
particles in the solution.
3. Choose one of the non-electrolytic solutions and draw a view, at the molecular level, of
the particles in the solution.
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Experiment 6: Lewis Dot Structure & Molecular Models
Purpose:
Students investigate how to write Lewis Dot Structures and convert these structures into
Electronic & Molecular geometric shapes. What is the significance of the electronic &
molecular geometry shape on the physical & chemical properties of the compound?
Introduction:
Whether the chemical bonding be ionic or covalent in nature, the electron is the critical
component to the bonding theory. The outermost electrons or valance electrons are the electrons
on an atom that are chemically active and lead the element or compound having its physical &
chemical properties. The Lewis Dot representation of the atom, molecule or ion depicts each of
the valance electrons. Covalent bonds are bonds between atoms that share electrons. The ability
to represent a molecule or polyatomic ion using the Lewis Dot Structure allow one to determine
the special orientation of the electron cloud or electron density. The spatial orientation of the
all the electron clouds around a central atom lead to the ability to determine the electronic and
molecular geometries. The method used to determine the geometric shape of the molecules or
polyatomic ions is called VSEPR (valance shell electron pair repulsion) model. The ability to
represent the molecule or polyatomic ion using the Lewis Dot Structure will help determine
molecules that have resonance and distinguish whether a molecule is polar or non-polar.
50 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Pre-Lab Questions:
1. What is meant by the “octet rule”?
2. What is meant by “electronegativity”?
3. Does a negative ion have more or less electrons than the neutral atom? Explain.
4. Does a positive ion have more or less electrons that the neutral atom? Explain.
5. What is the acronym “VSEPR” mean?
6. What is meant by “resonance”?
51 | P a g e
Procedure:
Complete the tables to show the Lewis Structures and Electronic & Molecular Geometry for the
molecules and polyatomic ions, using the information from the textbook and the information
given below.
A. Writing the Lewis Dot representation for a monoatomic ion or atom.
Determine the number of valance electrons for the monoatomic ion or atom. Write the chemical
symbol for the monoatomic ion or atom. Imagine a box around the chemical symbol; place a dot
for each valance electron around the chemical symbol. A box has four sides, so the first four
electrons are unpaired, if there are more than four electrons, the electrons start getting paired.
B. Writing the Lewis Dot representation for a compound or polyatomic
ion.
Determine the number of valance electrons for each of the atoms in the compound or polyatomic
ion. The central atom tends to be the least electronegative atom but the central atom is NEVER
hydrogen. The remaining atoms are placed around the central atom to form a skeleton referred
to as a spider. Some general rules for forming the bonding and non-bonding pairs of electrons.
1. Valance electrons are the only electrons allowed to form bonds between atoms.
2. Atoms should initially be paired so one electron from each atom is shared to form a
single bond between the two atoms.
3. Check for compliance with the octet rule.
4. Oxygen atoms will not bond together except in rare cases (O2, O3, peroxides)
5. Electrons should be paired, no individual electrons.
6. If the octet rule is not fulfilled, try forming multiple bonds.
7. If there are free electrons or the octet rule is still being violated, try forming molecules
with resonance structures.
C. Determining the Electronic & Molecular Geometry of a Compound or
Ion
Determine the electronic and molecular geometry of each of the molecules or polyatomic ions.
1. Determine the correct Lewis Dot Structure.
2. Determine the number of regions containing bonding electrons and non-bonding
electrons (lone pair electrons).
3. Based on the number of total regions of electrons, determine the electronic geometry.
4. Based on the number of regions of bonding electrons and non-bonding electrons,
determine the molecular geometry.
5. Using the molecular models build the molecules or polyatomic ions.
52 | P a g e
6. Table 1 contains the information for electronic and molecular geometry, and bond angles.
Table 1. Table of Molecular Geometric Information.
Regions
of
electronic
density
Electronic
Geometry
2
Linear
3
4
5
6
Trigonal Planer
Tetrahedral
Trigonal
Bipyramidal
Octahedral
Bond
Angle
Number of
Bonding
regions
Number of
Non-bonding
(lone pair)
Regions
Molecular
Geometry
180
2
0
Linear
1
1
Linear
3
0
Trigonal Planer
2
1
Bent
4
0
3
1
2
2
5
0
4
1
3
2
2
3
6
0
5
1
4
2
3
3
2
4
120
109.5
120, 90
90
Tetrahedral
Trigonal
Pyramidal
Bent
Trigonal
Bipyramidal
See-Saw
“T”- Shaped
Linear
Octahedral
Square
Pydramidal
Square Planer
“T” – Shaped
Linear
53 | P a g e
D. Determining if a Molecule is Polar or Non-Polar.
Once the geometry of the molecule has been determined, the polarity of the molecule can be
determined using the following method.
1. Use the steps to determine the electronic and molecular geometries.
2. Determine if the molecule is symmetric or non-symmetric.
a. Symmetric molecule: Non-polar molecule
b. Non-Symmetric molecule: Polar molecule
54 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Data & Observations:
A. Writing the Lewis Dot Structure for monoatomic ion or atoms.
Table 2. Lewis Dot Structure for atom or monoatomic ion.
Element
Chemical
Symbol
Number of
Valance
Electrons
Lewis Dot Structure
Potassium
Arsenic
Magnesium
Sulfur
Bromine
Boron
Sodium ion
55 | P a g e
Iodide ion
Oxygen ion
Rubidium ion
Calcium ion
B. Molecular Models.
Group 1
Group 2
Group 3
Group 4
Group 5
Group 6
SiCl4
CHBr3
CH2Cl2
CF4
SiH4
NH4+
PCl3
ICl2+
ClO3-
OF2
AsBr3
SF2
NO3-
NO2-
SO2
COH2
O3
COF2
SO42-
PO43-
ClO4-
CH3-
BrO4-
NH3
O2
N2
CO2
CS2
CN-
CSO
HOOH
BF3
BBr3
CH3OH
CO32-
BCl3
56 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
_______________ (Assigned molecule/ion #1)
_______________ (Assigned molecule/ion #2)
1. Number of valance electrons
__________
1.
Number of valance electrons
__________
2. Lewis Dot Structure:
2. Lewis Dot Structure:
3. Electronic Geometry:
3. Electronic Geometry:
4. Molecular Geometry:
4. Molecular Geometry:
5. Bond Angles:
5. Bond Angles:
6. Molecular Model ______________
(Instructors Initials)
6. Molecular Model ______________
(Instructors Initials)
7. Polar or Non-Polar molecules: (Explain)
7. Polar or Non-Polar molecules: (Explain)
57 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
_______________ (Assigned molecule/ion #3)
_______________ (Assigned molecule/ion #4)
1. Number of valance electrons
__________
1.
Number of valance electrons
__________
2. Lewis Dot Structure:
2. Lewis Dot Structure:
3. Electronic Geometry:
3. Electronic Geometry:
4. Molecular Geometry:
4. Molecular Geometry:
5. Bond Angles:
5. Bond Angles:
6. Molecular Model ______________
(Instructors Initials)
6. Molecular Model ______________
(Instructors Initials)
7. Polar or Non-Polar molecules: (Explain)
7. Polar or Non-Polar molecules: (Explain)
58 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
_______________ (Assigned molecule/ion #5)
_______________ (Assigned molecule/ion #6)
1. Number of valance electrons
__________
1.
Number of valance electrons
__________
2. Lewis Dot Structure:
2. Lewis Dot Structure:
3. Electronic Geometry:
3. Electronic Geometry:
4. Molecular Geometry:
4. Molecular Geometry:
5. Bond Angles:
5. Bond Angles:
6. Molecular Model ______________
(Instructors Initials)
6. Molecular Model ______________
(Instructors Initials)
7. Polar or Non-Polar molecules: (Explain)
7. Polar or Non-Polar molecules: (Explain)
59 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Answer the following questions using the molecules/ions assigned to your group.
1. Which molecules are polar?
2. Which molecules are non-polar?
3. Which molecules/ions violate the “octet rule”? Explain.
4. Is it necessary to determine the polarity of an ion? Explain
60 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Post-Lab Questions:
1. Explain the factors to determine the molecular geometry of a molecule or polyatomic
ion?
2. Explain how the molecular geometry is different from the electronic geometry.
3. What is the key difference between a polar molecular and a non-polar molecule?
61 | P a g e
Experiment 7: Stoichiometry
Purpose:
In this activity you and your lab partners will analyze several sets of data and solve the problems
based upon the information provided. The reactants will be graham crackers (Gc), chocolate
(Ch), and marshmallows (M) and the final product will be a S’more (Gc2Ch2M) which can be
abbreviated as Sm.
The formula for the reaction is as follows:
2Gc + 2Ch + 1M  Gc2Ch2M (can be abbreviated as Sm)
1 mole Gc = 1 square of a graham cracker = 1.94 g
1 mole Ch = 1 rectangle of chocolate = 2.51 g
1 mole M = 1 marshmallow = 0.53 g
Procedure:
1. Use the balanced equation provided above to answer the data analysis questions.
2. Determine if your answers are correct by comparing your responses with your
classmates. Redo any problems that are not.
3. After washing your hands, obtain the necessary supplies to make your S’mores.
4. Use safety precautions when lighting the Bunsen Burners.
5. Always wear goggles when using Bunsen Burners.
6. ENJOY! 
7. Make additional S’mores depending upon supplies.
8. Pay close attention to clean up so that you can partake in Cooking Chemistry Labs in the
future :)
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Data Analysis:
1. How many grams are found in 2 moles of a graham cracker?
2. How many grams are found in 6 moles of chocolate?
3. How many grams are found in 10 moles of marshmallows?
4. How many moles are in 5.82 g of Gc?
5. How many moles are in 5.02 g of Ch?
6. How many moles are in 5.30 g of M?
7. What is the molar mass of a S'more?
8. How many atoms will be found in two moles of chocolate?
9. How many molecules will be found in one mole of a S'more?
63 | P a g e
10. What is the mole ratio between chocolate and graham crackers? Marshmallows and
graham crackers? Marshmallows and S'mores?
11. If you have three marshmallows, how many S'mores can you make?
12. If you have 5 squares of graham crackers, how many complete S'mores can you make?
13. Given 5 moles of Gc and excess of the other reactants, how many S'mores can you make?
How does this question relate to question #12?
14. If 10 moles Ch reacts with M and Gc, how many moles of S'mores will be produced?
15. If 12 moles of Sm were produced, how many moles of Ch were required to make this
product?
16. If you wanted to make enough S'mores for a class of 37 students, how many moles of
Ch, Gc, and M would you need?
17. Given 21.6 g of M, how many moles of S’mores can you make?
18. Given 36.7 g of Gc, how many grams of M will you need?
64 | P a g e
19. Given 8 moles of Gc, how many grams of M will be needed?
20. To produce 97.8 g of Sm, how many moles of Gc will you need?
21. Based on the amount of Gc, Ch, and M your group received, which reactant would be the
limiting reagent (reactant)?
22. How many moles of S’mores did you produce?
23. How many grams of S’mores did you produce?
24. What were the reactants in excess?
25. How many moles of each excess reactant were left?
26. How many grams of each excess reactant were left?
65 | P a g e
Experiment 8: Alum Synthesis
Purpose:
The term alum is a general family name for a crystalline substance composed of cations with 1+
and 3+ charges. In this experiment, you will synthesize a type of alum called potassium
aluminum sulfate dodecahydrate, KAl(SO4)2•12H2O. You will synthesize this compound by
placing the appropriate ions in one container in aqueous solution and then evaporate the water to
form the alum crystals.
2 Al(s) + 2 KOH(aq) + 6 H2O(l)  2 KAl(OH)4(aq) + 3 H2(g)
This particular compound has been chosen because it is relatively simple to prepare a pure
sample. The process of synthesizing this compound is interesting in that it involves both
chemical and physical reactions. Chemically, aluminum is oxidized from aluminum foil to
prepare the Al3+ ions. Physically, as the solution that contains the mixture of ions evaporates,
crystals will form which contain six waters of hydration bonded to the aluminum ion and six
waters bonded to the potassium ion.
Aluminum is considered a reactive metal, but because its surface is usually protected by a thin
film of aluminum oxide, it reacts slowly with acids. It does, however, dissolve quickly in basic
solutions. Excess hydroxide ion converts the aluminum to the tetrahydroxoaluminate (Al(OH)3)
precipitates. Continued addition of acid causes the hydroxide ions to be completely neutralized,
and the aluminum exists in solution as the hydrated ion [Al(H2O)6]3+. Aluminum hydroxide is
considered to be an amphoteric hydroxide because it dissolves in both acids and bases.
In this experiment, you will
Synthesize a sample of potassium aluminum sulfate dodecahydrate (alum).
Observe and record the process of synthesizing a compound.
Calculate the percent yield of your synthesis.
Procedure:
1. Obtain and wear goggles.
2. Obtain a piece of aluminum foil and measure its mass. For best results, you should have
about 1.00 g of aluminum. Tear the foil into small pieces and place the pieces in a 250 mL
beaker.
3. Set up a Büchner funnel and filter flask so that you are ready to filter the reaction mixture
that will be produced in Step 4.
4. Conduct the first part of the synthesis. CAUTION: Potassium hydroxide solution is caustic.
Avoid spilling it on your skin or clothing.
a. Use a graduated cylinder to measure out 25 mL of 3 M KOH solution.
66 | P a g e
b. Slowly add the KOH solution to the beaker of aluminum pieces. Notice that the reaction is
exothermic. Allow the reaction to proceed until all of the foil is dissolved.
c. Carefully pour the reaction mixture through your Büchner funnel and filter flask setup,
and rinse the filter paper with a small amount of distilled water. Note: The reaction
mixture contains three ions: K+, [Al(OH)4–], and excess OH–.
d. Rinse the beaker with distilled water, and pour the filtered liquid back into the beaker.
5. Allow the solution to cool to near room temperature. If you are pressed for time, you may
cover the beaker with plastic wrap or Parafilm, and store the liquid overnight.
6. Clean the Büchner funnel and filter flask, and prepare it for more filtering that you may need
to do in Step 7 or Step 10.
7. Complete the synthesis.
a. Use a graduated cylinder to measure out 35 mL of 3 M H2SO4 solution. CAUTION: The
reaction mixture must be cooled to room temperature before proceeding. Handle the
H2SO4 solution with care. It can cause painful burns if it comes in contact with the skin.
b. After the reaction mixture has cooled, slowly add the sulfuric acid solution to the beaker
of liquid. Stir the mixture constantly. The reaction is strongly exothermic, so be careful as
you stir the mixture. Note that aluminum hydroxide will precipitate initially. It will
dissolve as more sulfuric acid is added.
c. If there is some solid remaining in the beaker after the 35 mL of sulfuric acid has been
added, pour the mixture through the Büchner funnel and filter flask to separate the
undissolved solid from the mixture.
8. Gently boil your mixture until you have about 50 mL of liquid in the beaker.
9. Cool the beaker of solution. Choose one of the two methods listed below.
a. Allow the solution to cool overnight. In most cases, this gradual cooling forms a good
crop of alum crystals.
b. Prepare an ice bath for the 250 mL beaker. Place your beaker of solution, uncovered, in
the ice bath. Do not move the ice bath or the beaker. After about fifteen minutes, crystals
of alum will appear in the beaker. If there are no crystals after fifteen minutes, scrape the
bottom of the beaker with a glass stirring rod to create a rough spot for crystal growth.
You may also heat the solution to evaporate more water and cool the solution again.
10. Collect your alum crystals by pouring them onto the Büchner funnel and filter-flask setup.
Use vacuum filtration to wash the crystals on the filter paper with 50 mL of an aqueous
ethanol solution (50%). The crystals will not dissolve in this solution.
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11. Remove the filter and crystals from the Büchner funnel and allow the crystals to dry at room
temperature. Measure and record the mass of your sample of alum. Store the crystals for
further analysis.
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
PRE-LAB QUESTIONS
1. What are the two cautions you need to take when conducting the synthesis of alum experiment?
2. Fill in the blanks for the physical properties of alum:
a. Molecular Weight: _____________g/mol
b. Melting Point: _________________ oC
3. A student conducting the synthesis of alum experiment started with 0.9156 g of aluminum foil.
a. What is the theoretical yield of alum? (Show your work)
b. The student isolated 8.3181 g of alum. What was the percent yield the student obtained? (Show
your work)
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
DATA TABLE
Mass of Aluminum Foil used: ___________________g
Moles of Aluminum Foil used: __________________mol
Mole Ratio of Aluminum Foil to Alum: ___________ratio
Moles of Alum theoretically obtained: ____________mol
Theoretical Mass of Alum: _____________________g
Actual Mass of Alum obtained: _________________g
DATA ANALYSIS
1. Determine the theoretical yield of the alum. Use the aluminum foil as the limiting reagent
and presume that the foil was pure aluminum.
2. Calculate the percent yield of your alum crystals.
3. Discuss the factors that affected the percent yield.
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4. Write the balanced equations for the following: (a) aluminum and potassium hydroxide,
yielding [Al(OH)4] – and hydrogen gas; (b) hydrogen ions and [Al(OH)4] –, yielding
aluminum hydroxide; (c) aluminum hydroxide and hydrogen ions, yielding [Al(H2O)6]3+; and
(d) the formation of alum from potassium ions, sulfate ions, [Al(H2O)6]3+, and water.
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Experiment 9: Alum Analysis
Purpose:
After a compound has been synthesized, tests should be carried out to verify that the compound
formed is indeed the compound desired. There are a number of tests that can be performed to
verify that the compound is the one desired. In the previous experiment, you prepared alum
crystals, KAl(SO4)2•12H2O. In this experiment, you will conduct a series of tests to determine if
your crystals are really alum.
The first test is to find the melting temperature of the compound and compare this value with the
accepted (published) value for alum (92.5°C). The second test determines the water of hydration
present in the alum crystals. The third test is a chemical test to determine the percent sulfate in
your sample of alum.
In this experiment, you will
Determine the melting temperature of a sample of alum.
Determine the water of hydration of a sample of alum.
Procedure:
Part I: Determine the Melting Temperature
1. Obtain and wear goggles.
2. Connect the Temperature Probe to LabQuest.
3. Take a piece of weighing paper folded in half and using a micro-spatula remove a small
amount of alum from the sample you prepared in the past experiment. Place the alum into
the fold of the weighing paper, fold the weighing paper over, and using the edge of the
micro-spatula pulverize the alum sample. Use the micro-spatula to pile the alum in the
weighing paper. Push the open end of a capillary tube into the pile of the alum powder. Pack
alum into the capillary tube to a depth of about 0.5 cm by tapping the tube lightly on the table
top.
4. Use a rubber band to fasten the capillary tube to the Temperature Probe. The tip of the tube
should be even with the tip of the probe. Use a utility clamp to connect the Temperature
Probe to a ring stand (see Figure 1).
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Figure 1
5. Prepare a water bath to be heated by a hot plate.
6. Monitor the temperature readings on the Main screen. Immerse the capillary tube and
Temperature Probe in the water bath. Warm the alum sample at a gradual rate so that you can
precisely determine the melting temperature. The white powder will become clear when it is
melting. Observe the temperature readings and record the precise melting temperature when
the substance is completely clear.
7. Conduct a second test with a new sample of alum in a new capillary tube.
Part II Determine the Water of Hydration
8. Heat a crucible with cover over a burner flame until it is red hot. Allow the crucible to cool,
and then measure the total mass of crucible and cover. Handle the crucible with tongs or
forceps to avoid getting fingerprints on it.
9. Place about 2 g of your alum crystals in the crucible, and then measure the mass of the
crucible, cover, and alum. Record this measurement in the data table.
10. Set up a ring, ring stand and triangle over a lab burner. Use tongs or forceps to set the
crucible at an angle on the triangle and place the cover loosely on the crucible. Use a lab
burner to very gently heat the crucible of alum until you can see no vapor escaping from the
crucible. It is important that the vapor does not carry any alum with it. After the vapor is
gone, heat the crucible more strongly for five minutes, and then cool the crucible.
11. Measure and record the mass of crucible, cover, and alum after drying. This would be the
mass after the first heating.
12. Reheat the crucible and alum sample for five additional minutes. Cool and measure the mass
of the crucible again. This would be the mass after the second heating. If the two masses are
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the same (or very nearly so), the test is done. If not, repeat the heating and weighing until a
constant mass is obtained.
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
PRE-LAB QUESTIONS
1. Write the definition/explanation of each of the following terms:
a. Hydrated Salt
b. Anhydrous Salt
c. Hygroscopic substance
d. Desiccant substance
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2. A student was asked to identify an unknown hydrate by following the following the
procedure in this experiment to determine the percent water of alum. A 2.752 g sample
of the unknown sample is heated and weighted after cooling to constant weight of 1.941
g. The unknown is believed to be one of the following compounds: LiNO3.3H2O,
Ca(NO3)2.4H2O, or Sr(NO3)2.4H2O.
a. Calculate the percent water in the unknown sample.
b. Calculate the percent water in the three known samples.
c. What is the identity of the unknown sample? Explain your reasoning.
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
DATA TABLE
Part I Melting Temperature Test Results
Trial 1
Trial 2
Trial 1
Trial 2
Melting Temperature (°C)
Part II Water of Hydration Test Results
Mass of crucible and cover (g)
Mass of crucible, cover, and alum before heating (g)
Mass of Alum before heating (g)
Mass of crucible, cover, and alum after 1st heating (g)
Mass of anhydrous Alum after 1st heating (g)
Mass of crucible, cover, and alum after 2nd heating (g)
Mass of anhydrous Alum after 2nd heating (g)
Mass of crucible, cover, and alum after final heating (g)
Mass of anhydrous Alum after final heating (g)
Average mass of anhydrous Alum (g)
Moles of anhydrous Alum (mol)
Average mass of water lost (g)
Moles of water (mol)
Ratio of moles Alum : moles water
Chemical Formula of Alum
77 | P a g e
DATA ANALYSIS
1. Is your sample alum? Use the results of the three tests to support your answer. Discuss the
accuracy of your tests and possible sources of experimental error.
2. Suggest other tests that could be conducted to verify the composition of your alum.
3. If the melting temperature test was the only test that you conducted, how confident would
you be in the identification of your sample? Explain.
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Experiment 10: Identification Unknown Solution (Solubility)
Purpose:
Students investigate the procedure to identify unknown solutions bases on the observation of
chemical changes. What are the identities of the each of the unknowns?
Introduction:
Chemical change occurs when molecules of two substances are allowed to come in contact with
each other with enough energy for change to occur with the molecules. One can observe if a
chemical reaction occurred if there is an observable change. Observable changes are:
1.
2.
3.
4.
Precipitate forms (solid formed when two solutions are mixed)
Color change
Bubbles form (gases are released)
Heat generated
Observing one of these indicates a chemical reaction has occurred between two molecules.
These signs of a chemical reaction are especially evident when dealing with ionic compounds.
Ionic compounds are made up of ions, positively or negatively charged atoms or molecules.
Positively charged atoms or molecules are called cations, and negatively charged atoms or
molecules are called anions. The charged molecules are termed polyatomic ions. The
combination of cations and anions form the neutral ionic compound. Many ionic compounds
dissolve in water and when the ionic compound dissolves in the water, the ions separate into the
cations and anions. These are free to interact with other ions; cations react with other anions and
anions reacting with other cations. Many of the interactions of the cation and anion do not
generate one of the signs of a chemical reaction. But if there is a sign of a chemical reaction,
knowing the initial ionic compounds can be used to determine unknown solutions. Even the
observable signs of a chemical reaction can help determine unknown solutions:
1. Precipitation reactions produce a solid when two solutions are mixed together. The
solutions are ionic solutions and the combination of the specific cation and anion result in
the precipitate formation. The rules for precipitation can be used to determine whether
two solutions will produce a precipitate.
2. Color changes tend to involve Transition metal cations.
3. Bubbles forming tend to be single replacement/displacement reactions (redox). Or CO2
generated reaction involving carbonate or bicarbonate.
4. Heat generating reactions tend to be acids reacting bases, either the acid or base needs to
be strong.
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The rules for solubility (soluble means dissolved, insoluble means does not dissolve) outline
when two solutions are mixed together, whether the resulting mixture will have a precipitate
formed. These are the rules for solubility:
1. All soluble when cation is alkali metal or ammonium ion.
2. All soluble when anion is nitrate, acetate, perchlorate.
3. All soluble when anion is halide.
a. Except when cation is silver ion, lead(II) ion, mercury(II) ion.
4. All soluble when anion is sulfate.
a. Except when the cation is calcium, barium,
5. All insoluble when anion is hydroxide.
a. Except when the cation is alkali metal or alkaline earth metal.
6. All insoluble when anion is sulfide, carbonate, phosphate.
a. Except when cation is alkali metal.
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Pre-Lab Questions:
1. When a gaseous product is formed, will one see 1-large bubble or many small bubbles
form? (Hint: Think of a carbonated beverage.) Explain.
2. Define a precipitate.
3. Match the terms with the abbreviations, place the number of the abbreviation in the space
provided of the matching term:
_____ a. solution
1. unk
_____ b. precipitate
2. aq
_____ c. reaction
3. rxn
_____ d. aqueous
4. ppt
_____ e. unknown
5. soln
81 | P a g e
4. Complete the chemical equation for each of the following. Be sure to include the state of
each compound or if it is in solution. If a chemical reaction is observed, write the net
ionic equation.
BaCl2(aq) + KIO4(aq) 
Sr(NO3)2(aq) + Na2SO4(aq) 
Pb(NO3)2(aq) + KOH(aq) 
CuSO4(aq) + Na2CO3(aq) 
HCl(aq) + NaOH(aq) 
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Procedure:
Students will conduct the procedure for this experiment twice, once with a known set of
solutions and once with the unknown set of solutions. The object of this experiment is to
determine the unknown solutions.
A. Studying the reactions of the known set of solutions.
Students will obtain the six known solutions and a 96-well plate. Using Table 1 in the data &
observation section, construct a similar matrix using the 96-well plate. Choose one of the
solutions to be added to row 1, this will be solution 1. Write the name of the compound for
solution 1 in the space provided on Table 1. Add one drop of the solution to each of the first five
wells, in the first row, of the 96-well plate. Choose one of the solutions to be added to row 2,
this will be solution 2. Write the name of the compound for solution 2 in the space provided on
Table 1. Add one drop of the solution to each of the first four wells, in the second row, of the
96-well plate. Choose another solution for row 3; write the name of that solution in the space
provided on Table 1. Add one drop of the solution to each of the first three wells of the 96-well
plate. Choose a fourth solution to be added to row 4. Write the name of the solution on row 4 of
Table 1. Add a drop of solution 4 to the first 2 wells of the 96-well plate in row 4. Choose a
fifth solution to be added to row 5. Write the name of the solution on row 5 of Table 1. Add a
drop if solution 5 to the first well of the 96-well plate in row 5. The name of the sixth solution
will be written into the first column of Table 1. Add one drop of solution 6 to the first five wells
of column 1. Take solution 5 and write the name of solution 5 in the space provided in column
2. Add one drop of solution 5 to the first four wells of column 2. Take solution 4 and write the
name of solution 4 in the space provided in column 3. Add one drop of solution 4 to the first
three wells of column 3. Take solution 3 and write the name of solution 3 in the space provided
in column 4. Add one drop of solution 3 to the first two wells of column 4. Take solution 2 and
write the name of solution 2 in the space provided in column 5. Add one drop of solution 2 to
the first well of column 5. Observe each well for a sign of a reaction. If no observable sign of a
reaction, write N.R. on Table 1 where the two solutions intersect on the table. If there is a color
change, write the color of the solution. If a precipitate forms, write “ppt” and the color of the
precipitate. If bubbles form, write “bubbles”, and if heat was generated, write “heat”.
B. Studying the reactions of the unknown set of solutions.
Students will obtain a set of six unknown solutions and another 96-well plate. Using Table 2 in
the data & observation section, repeat the above procedure. In the space provided, write the
unknown number. Based on the results of the reactions of the unknowns and comparing those
data with the data in Table 1, the student will determine each of the unknowns and report those
results.
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Data & Observations:
A. Studying the reactions of the known set of solutions.
Table 1. Matrix of reactions of known solutions.
Solution 6
Solution 5
Solution 4
___________ ___________ ___________
Solution 3
Solution 2
___________ ___________
Solution 1
__________
Solution 2
__________
Solution 3
__________
Solution 4
___________
Solution 5
___________
84 | P a g e
B. Studying the reactions of the unknown set of solutions.
Table 2. Matrix of reactions of unknown solutions.
Unknown 6
Unknown 5
Unknown 4
___________ ___________ ___________
Unknown 3
Unknown 2
___________ ___________
Unknown 1
__________
Unknown 2
__________
Unknown 3
__________
Unknown 4
___________
Unknown 5
___________
85 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
C. Identify the Unknowns.
Unknown 1: ____________________
Experimental reasoning for your identification:
Unknown 2: ____________________
Experimental reasoning for your identification:
Unknown 3: ____________________
Experimental reasoning for your identification:
Unknown 4: ____________________
Experimental reasoning for your identification:
Unknown 5: ____________________
Experimental reasoning for your identification:
Unknown 6: ____________________
Experimental reasoning for your identification:
86 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
1. Write the molecular equation for each of the reactions that produced a precipitate.
2. Write the net ionic equation for each of the reactions that produced a precipitate.
87 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
Post-Lab Questions:
1. A student was given the following five known solutions; NiCl2, Pb(NO3)2, NaOH,
Na2C2O4, and CuSO4. After conducting an experiment mixing the solutions in pairs and
recording their observations in Table 3.
Table 3. Precipitate formation by mixing the pairs of known solutions.
CuSO4
Na2C2O4
NaOH
Pb(NO3)2
NiCl2
N.R.
N.R.
grenn ppt
N.R.
Pb(NO3)2
white ppt
white ppt
white ppt
NaOH
blue ppt
N.R.
Na2C2O4
white ppt
The student was then given the same five solutions as unknowns labeled Solutions A-E. After
conducted the same mixing experiment as with the known solutions, the student recorded their
observations in Table 4.
88 | P a g e
Table 4. Precipitate formation by mixing the pairs of unknown solutions.
Solution E
Solution D
Solution C
Solution B
Solution A
blue ppt
white ppt
N.R.
green ppt
Solution B
N.R.
N.R.
N.R.
Solution C
white ppt
white ppt
Solution D
white ppt
Identify the unknowns:
Solution A: ____________
Solution B: ____________
Solution C: ____________
Solution D: ____________
Solution E: ____________
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Experiment 11: Acid-Base Titration
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Experiment 12: Household Acids & Bases
Purpose:
Many common household solutions contain acids and bases. Acid-base indicators, such as litmus
and red cabbage juice, turn different colors in acidic and basic solutions. They can, therefore, be
used to show if a solution is acidic or basic. An acid turns blue litmus paper red, and a base turns
red litmus paper blue. The acidity of a solution can be expressed using the pH scale. Acidic
solutions have pH values less than 7, basic solutions have pH values greater than 7, and neutral
solutions have a pH value equal to 7.
In this experiment, you will use litmus and a pH Sensor to determine the pH values of household
substances. After adding red cabbage juice to the same substances, you will determine the
different red cabbage juice indicator colors over the entire pH range.
In this experiment, you will

Use litmus paper and a pH Sensor to determine the pH values of household substances.

Add cabbage juice to the same substances and determine different red cabbage juice
indicator colors over the entire pH range.
Figure 1
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Procedure:
1. Obtain and wear goggles. CAUTION: Do not eat or drink in the laboratory.
Part I Litmus Tests
2. Label 7 test tubes with the numbers 1–7 and place them in a test tube rack.
3. Measure 3 mL of vinegar into test tube #1. Refer to the data table and fill each of the test
tubes 2–7 to about the same level with its respective solution. CAUTION: Ammonia solution
is toxic. Its liquid and vapor are extremely irritating, especially to eyes. Drain cleaner
solution is corrosive. Handle these solutions with care. Do not allow the solutions to contact
your skin or clothing. Wear goggles at all times. Notify your teacher immediately in the event
of an accident.
4. Use a stirring rod to transfer one drop of vinegar to a small piece of blue litmus paper on a
paper towel. Transfer one drop to a piece of red litmus paper on a paper towel. Record the
results. Clean and dry the stirring rod each time.
5. Test solutions 2–7 using the same procedure. Be sure to clean and dry the stirring rod each
time.
Part II Red Cabbage Juice Indicator
6. After you have finished the Part I litmus tests, add 3 mL of red cabbage juice indicator to
each of the 7 test tubes. Record your observations. Dispose of the test-tube contents as
directed by your teacher.
Part III pH Tests
7. Prepare the pH Sensor for data collection.
a. Connect the pH Sensor to LabQuest and choose New from the File menu. If you have an
older sensor that does not auto-ID, manually set up the sensor.
b. Remove the pH Sensor from the sensor storage solution bottle by unscrewing the lid.
Carefully remove the bottle, leaving the cap on the sensor body.
c. Rinse the tip of the sensor with distilled water and place the sensor tip into a beaker
containing sensor soaking solution. Use a utility clamp to fasten the pH Sensor to a ring
stand, as shown in Figure 1.
8. Raise the pH Sensor from the sensor soaking solution and set the solution aside. Use a wash
bottle filled with distilled water to thoroughly rinse the pH Sensor. Catch the rinse water in a
250 mL beaker.
9. Obtain one of the 7 solutions in the small container supplied by your teacher. Raise the
solution to the pH Sensor and swirl the solution about the sensor. When the pH reading
stabilizes, record the pH value.
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10. Prepare the pH Sensor for reuse.
a. Rinse it with distilled water from a wash bottle.
b. Place the sensor into the sensor soaking solution and swirl the solution about the sensor
briefly.
c. Rinse with distilled water again.
11. Determine the pH of the other solutions using the Step 9 procedure. You must clean the
pH Sensor between tests using the Step 10 procedure.
12. When you are finished, rinse the sensor with distilled water and return it to the sensor
soaking solution.
93 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
PROCESSING THE DATA
1. Which of the household solutions tested are acids? How can you tell?
2. Which of the solutions are bases? How can you tell?
3. What color(s) is red cabbage juice indicator in acids? In bases?
4. Can red cabbage juice indicator be used to determine the strength of acids and bases?
Explain.
5. List advantages and disadvantages of litmus and red cabbage juice indicators.
94 | P a g e
DATA TABLE
Test
Tube
Solution
1
vinegar
2
ammonia
3
lemon juice
4
soft drink
5
drain cleaner
6
detergent
7
baking soda
8
antacid
Blue
Litmus
Red
Litmus
Red Cabbage
Juice
pH
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Experiment 13: Boyles Law & Guy-Lussac
Boyles Law
Purpose:
The primary objective of this experiment is to determine the relationship between the pressure
and volume of a confined gas. The gas we use will be air, and it will be confined in a syringe
connected to a Gas Pressure Sensor (see Figure 1). When the volume of the syringe is changed
by moving the piston, a change occurs in the pressure exerted by the confined gas. This pressure
change will be monitored using a Gas Pressure Sensor. It is assumed that temperature will be
constant throughout the experiment. Pressure and volume data pairs will be collected during this
experiment and then analyzed. From the data and graph, you should be able to determine what
kind of mathematical relationship exists between the pressure and volume of the confined gas.
Historically, this relationship was first established by Robert Boyle in 1662 and has since been
known as Boyle’s law.
Figure 1
Procedure:
1. Prepare the Gas Pressure Sensor and an air sample for data collection.
a. Connect the Gas Pressure Sensor to LabQuest. If you have an older sensor that does not
auto-ID, manually set up the sensor.
b. With the 20 mL syringe disconnected from the Gas Pressure Sensor, move the piston of
the syringe until the front edge of the inside black ring (indicated by the arrow in Figure 1)
is positioned at the 10.0 mL mark.
c. Attach the 20 mL syringe to the valve of the Gas Pressure Sensor.
2. You are now ready to collect pressure and volume data. It is easiest if one person takes care
of the gas syringe and another enters volumes.
a. Move the piston so the front edge of the inside black ring (see Figure 2) is positioned at
the 5.0 mL line on the syringe. Hold the piston firmly in this position until the pressure
value displayed on the screen stabilizes.
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Figure 2
b. Continue this procedure using syringe volumes of 10.0, 12.5, 15.0, 17.5, and 20.0 mL.
97 | P a g e
Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
DATA AND CALCULATIONS
Volume
(mL)
Pressure
(kPa)
Constant, k
(P / V or P • V)
5.00
10.0
12.5
15.0
17.5
20.0
PROCESSING THE DATA
1. Based on your data, what would you expect the pressure to be if the volume of the syringe
was increased to 40.0 mL. Explain or show work to support your answer.
2. Based on your data, what would you expect the pressure to be if the volume of the syringe
was decreased to 2.5 mL.
3. What experimental factors are assumed to be constant in this experiment?
98 | P a g e
4. One way to determine if a relationship is inverse or direct is to find a proportionality
constant, k, from the data. If this relationship is directly proportional, k = P/V. If it is inverse
proportionsal, k = P•V. Choose one of these formulas based on your knowledge of Boyle’s
Law, and calculate k for the seven Pressure-Volume data pairs in your data table. Based on
your answers to in the Data and Calculations Table, does the relationship hold true? Explain
your answer.
5. Graph the pressure vs volume relationship or the pressure vs inverse volume (1/volume)
relationship using Excel. Using the graph, determine if the relationship between pressure and
volume is a direct or inverse proportionality. Since the graph showing proportionality should
be a linear graph, you must have the trendline, r2 value, and the equation for the line on the
graph. Your graph should also have a title, and axis’s labeled including units.
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Experiment 13: Boyles Law & Guy-Lussac
Guy-Lussac Law
Purpose:
Gases are made up of molecules that are in constant motion and exert pressure when they collide
with the walls of their container. The velocity and the number of collisions of these molecules
are affected when the temperature of the gas increases or decreases. In this experiment, you will
study the relationship between the temperature of a gas sample and the pressure it exerts. Using
the apparatus shown in Figure 1, you will place an Erlenmeyer flask containing an air sample in
water baths of varying temperature. Pressure will be monitored with a Pressure Sensor and
temperature will be monitored using a Temperature Probe. The volume of the gas sample and the
number of molecules it contains will be kept constant. Pressure and temperature data pairs will
be collected during the experiment and then analyzed. From the data and graph, you will
determine what kind of mathematical relationship exists between the pressure and absolute
temperature of a confined gas. You may also do the extension exercise and use your data to find
a value for absolute zero on the Celsius temperature scale.
OBJECTIVES
In this experiment, you will

Study the relationship between the temperature of a gas sample and the pressure it exerts.

Determine from the data and graph, the mathematical relationship between pressure and
absolute temperature of a confined gas.

Find a value for absolute zero on the Celsius temperature scale.
Figure 1
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PROCEDURE
1. Obtain and wear goggles.
2. Prepare a boiling-water bath. Put about 800 mL of hot tap water into a l L beaker and place it
on a hot plate. Turn the hot plate to a high setting.
3. Prepare an ice-water bath. Put about 700 mL of cold tap water into a second 1 L beaker and
add ice.
4. Put about 800 mL of room-temperature water into a third 1 L beaker.
5. Put about 800 mL of hot tap water into a fourth 1 L beaker.
6. Prepare the Temperature Probe and Gas Pressure Sensor for data collection.
a. Connect the Gas Pressure Sensor to Channel 1 of LabQuest and the
Temperature Probe to Channel 2. Choose New from the File menu.
If you have older sensors that do not auto-ID, manually set up the
sensors.
Figure 2
b. Obtain a rubber-stopper assembly with a piece of heavy-wall plastic
tubing connected to one of its two valves. Attach the connector at
the free end of the plastic tubing to the open stem of the Gas Pressure Sensor with a
clockwise turn. Leave its two-way valve on the rubber stopper open (lined up with the
valve stem as shown in Figure 2) until Step 6d.
c. Insert the rubber-stopper assembly into a 125 mL Erlenmeyer flask. Important: Twist the
stopper into the neck of the flask to ensure a tight fit.
Figure 3
d. Close the 2-way valve above the rubber stopper—do this by turning the valve handle so it
is perpendicular with the valve stem itself (as shown in Figure 3). The air sample to be
studied is now confined in the flask.
9. Start recording data. Pressure readings (in kPa) and temperature readings (in °C) and are
displayed on the screen.
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10. Collect pressure vs. temperature data for your gas sample.
a. Place the flask into the ice-water bath. Make sure the entire flask is covered (see
Figure 3).
b. Place the Temperature Probe into the ice-water bath.
c. When the temperature and pressure readings have both stabilized, record the temperature
and pressure readings.
11. Repeat Step 10 using the room-temperature bath.
12. Repeat Step 10 using the hot-water bath.
13. Use a ring stand and utility clamp to suspend the Temperature Probe in the boiling-water
bath. CAUTION: Do not burn yourself or the probe wires with the hot plate. To keep from
burning your hand, hold the tubing of the flask using a glove or a cloth. After the
Temperature Probe has been in the boiling water for a few seconds, place the flask into the
boiling-water bath and repeat Step 10.
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
PROCESSING THE DATA
1. In order to perform this experiment, what two experimental factors were kept constant?
2. Based on the data and graph that you obtained for this experiment, express in words the
relationship between gas pressure and temperature.
3. Write an equation to express the relationship between pressure and temperature (K). Use
the symbols P, T, and k.
4. One way to determine if a relationship is inverse or direct is to find a proportionality
constant, k, from the data. If this relationship is direct, k = P/T. If it is inverse, k = P•T.
Based on your answer to Question 3, choose one of these formulas and calculate k for the
four ordered pairs in your data table (divide or multiply the P and T values). Show the
answer in the fourth column of the Data and Calculations table. How “constant” were
your values?
5. Graph the pressure vs volume relationship or the pressure vs inverse volume (1/volume)
relationship using Excel. Using the graph, determine if the relationship between pressure
and volume is a direct or inverse proportionality. Since the graph showing
proportionality should be a linear graph, you must have the trendline, r2 value, and the
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equation for the line on the graph. Your graph should also have a title, and axis’s labeled
including units.
6. According to this experiment, what should happen to the pressure of a gas if the Kelvin
temperature is doubled? Check this assumption by finding the pressure at –73°C (200 K)
and at 127°C (400 K) on your graph of pressure versus temperature. How do these two
pressure values compare?
DATA AND CALCULATIONS
Pressure
(kPa)
Temperature
(°C)
Temperature
(K)
Constant, k
(P / T or P•T)
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Experiment 14: Molar Volume
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Experiment 15: Heat of Fusion
Purpose:
Melting and freezing behavior are among the characteristic properties that give a pure substance
its unique identity. As energy is added, pure solid water (ice) at 0°C changes to liquid water
at 0°C.
In this experiment, you will determine the energy (in joules) required to melt one gram of ice.
You will then determine the molar heat of fusion for ice (in kJ/mol). Excess ice will be added to
warm water, at a known temperature, in a Styrofoam cup. The warm water will be cooled down
to a temperature near 0°C by the ice. The energy required to melt the ice is removed from the
warm water as it cools.
To calculate the heat that flows from the water, you can use the relationship
q = Cp•m•t
where q stands for heat flow, Cp is specific heat capacity, m is mass in grams, and t is the
change in temperature. For water, Cp is 4.18 J/g°C.
In this experiment, you will

Determine the energy (in Joules) required to melt one gram of ice.

Determine the molar heat of fusion for ice (in kJ/mol).
Procedure:
Part I Freezing
1. Connect the Temperature Probe to LabQuest and choose New from the File menu. If you
have an older sensor that does not auto-ID, manually set up the sensor.
2. Use a utility clamp to clamp the Temperature Probe on a ring stand as shown in Figure 1.
3. Place a Styrofoam cup into a 400 mL beaker as shown in Figure 1.
4. Use a 100 mL graduated cylinder to obtain 100.0 mL of water at about 60°C from your
teacher.
5. Obtain 7 or 8 large ice cubes.
6. Lower the Temperature Probe into the warm water (to about 1 cm from the bottom).
7. Record the temperature reading, in °C, is displayed to the right of the graph. You will need to
wait until the temperature reaches a maximum (it will take a few seconds for the cold probe
to reach the temperature of the warm water). This maximum will determine the initial
temperature, t1, of the water. As soon as this maximum temperature is reached, fill the
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Styrofoam cup with ice cubes. Shake excess water from the ice cubes before adding them (or
dry with a paper towel). Record the maximum temperature, t1, in your data table.
Figure 1
8. Use a stirring rod to stir the mixture as the temperature approaches 0°C. Important: As the
ice melts, add more large ice cubes to keep the mixture full of ice!
9. When the temperature reaches about 4°C, quickly remove the unmelted ice (using tongs).
Continue stirring until the temperature reaches a minimum (and begins to rise). This
minimum temperature is the final temperature, t2, of the water. Record t2 in your data table.
10. Use the 100 mL graduated cylinder to measure the volume of water remaining in the
Styrofoam cup to the nearest 0.1 mL. Record this as V2.
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
PROCESSING THE DATA
1. Use the equation t = t1 – t2 to determine t, the change in water temperature.
2. Subtract to determine the volume of ice that was melted (V2 –V1).
3. Find the mass of ice melted using the volume of melt (use 1.00 g/mL as the density of water).
4. Use the equation given in the introduction of this experiment to calculate the energy (in joules)
released by the 100 g of liquid water as it cooled through t.
5. Now use the results obtained above to determine the heat of fusion, the energy required to melt
one gram of ice (in J/g H2O).
6. Use your answer to Step 5 and the molar mass of water to calculate the molar heat of fusion for
ice (in kJ/mol H2O).
7. Find the percent error for the molar heat of fusion value in Step 6. The accepted value for molar
heat of fusion is 6.01 kJ/mol.
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DATA AND CALCULATIONS
Initial water temperature, t1
°C
Final water temperature, t2
°C
Change in water temperature, t
°C
Final water volume, V2
mL
Initial water volume, V1
mL
Volume of melt
mL
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Experiment 16: Evaporation & Intermolecular Interactions
Purpose:
In this experiment, Temperature Probes are placed in various liquids. Evaporation occurs when
the probe is removed from the liquid’s container. This evaporation is an endothermic process that
results in a temperature decrease. The magnitude of a temperature decrease is, like viscosity and
boiling temperature, related to the strength of intermolecular forces of attraction. In this
experiment, you will study temperature changes caused by the evaporation of several liquids and
relate the temperature changes to the strength of intermolecular forces of attraction. You will use
the results to predict, and then measure, the temperature change for several other liquids.
You will encounter two types of organic compounds in this experiment—alkanes and alcohols.
The two alkanes are pentane, C5H12, and hexane, C6H14. In addition to carbon and hydrogen
atoms, alcohols also contain the -OH functional group. Methanol, CH3OH, and ethanol,
C2H5OH, are two of the alcohols that we will use in this experiment. You will examine the
molecular structure of alkanes and alcohols for the presence and relative strength of two
intermolecular forces—hydrogen bonding and dispersion forces.
In this experiment, you will

Study temperature changes caused by the evaporation of several liquids.

Relate the temperature changes to the strength of intermolecular forces of attraction.
Figure 1
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Pre-lab exercise:
Prior to doing the experiment, complete the Pre-Lab table. The name and formula are given for
each compound. Draw a structural formula for a molecule of each compound. Then determine
the molecular weight of each of the molecules. Dispersion forces exist between any two
molecules, and generally increase as the molecular weight of the molecule increases. Next,
examine each molecule for the presence of hydrogen bonding. Before hydrogen bonding can
occur, a hydrogen atom must be bonded directly to an N, O, or F atom within the molecule. Tell
whether or not each molecule has hydrogen-bonding capability.
Procedure:
1. Obtain and wear goggles! CAUTION: The compounds used in this experiment are
flammable and poisonous. Avoid inhaling their vapors. Avoid contacting them with your skin
or clothing. Be sure there are no open flames in the lab during this experiment. Notify your
teacher immediately if an accident occurs.
2. Connect the Temperature Probes to LabQuest. If you have older sensors that do not auto-ID,
manually set up the sensors.
3. Wrap Probe 1 and Probe 2 with square pieces of filter paper secured by small rubber bands as
shown in Figure 1. Roll the filter paper around the probe tip in the shape of a cylinder. Hint:
First slip the rubber band on the probe, wrap the paper around the probe, and then finally slip
the rubber band over the paper. The paper should be even with the probe end.
4. Stand Probe 1 in the ethanol container and Probe 2 in the 1-propanol container. Make sure
the containers do not tip over.
5. After the probes have been in the liquids for at least 30 seconds, start data collection. Monitor
the temperature for 15 seconds to establish the initial temperature of each liquid. Then
simultaneously remove the probes from the liquids and tape them so the probe tips extend
5 cm over the edge of the table top as shown in Figure 1.
6. Monitor the temperature of both probes to determine when each probe maintains a constant
minimum temperature. Record your minimum temperature value for each probe (t2).
7. For each liquid, subtract the minimum temperature from the maximum temperature to
determine t, the temperature change during evaporation.
8. Repeating Steps 3–7 using 1-butanol with Probe 1 and n-pentane with Probe 2.
9. Repeating Steps 3–7, using methanol with Probe 1 and n-hexane with Probe 2.
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
PROCESSING THE DATA
1. Two of the liquids, n-pentane and 1-butanol, had nearly the same molecular weights, but
significantly different t values. Explain the difference in t values of these substances,
based on their intermolecular forces.
2. Which of the alcohols studied has the strongest intermolecular forces of attraction? The
weakest intermolecular forces? Explain using the results of this experiment.
3. Which of the alkanes studied has the stronger intermolecular forces of attraction? The weaker
intermolecular forces? Explain using the results of this experiment.
4. Plot a graph of t values of the four alcohols versus their respective molecular weights. Plot
molecular weight on the horizontal axis and t on the vertical axis. Be sure to label your
axis, have title, and since the graph is a linear graph you should have an equation for the line
and r2 value.
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PRE-LAB
Substance
Formula
Structural Formulas
Molecular
Weight
Hydrogen
Bonding
(yes or no)
ethanol
C2H5OH
1-propanol
C3H7OH
1-butanol
C4H9OH
n-pentane
C5H12
methanol
CH3OH
n-hexane
C6H14
DATA TABLE
Substance
t1 (oC)
t2 (oC)
t (t1 – t2) (oC)
ethanol
1-propanol
1-butanol
n-pentane
methanol
n-hexane
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