CHEMISTRY LABS
1. Buttered Cats (Scientific Method) .........................................................................................................2
2. Density/Precision in Measurements
a. Honors Chemistry ......................................................................................................................4
b. Academic Chemistry..................................................................................................................5
3. %Water in a Hydrate ..............................................................................................................................7
4. Separation of a Mixture .........................................................................................................................9
5. Conservation of Mass
a. Honors Chemistry ....................................................................................................................10
i. Conservation of Mass Write-Up Guidelines ................................................................11
ii. Annotated Sample Formal Lab Report ........................................................................12
iii. Example of Student Report ..........................................................................................19
b. Academic Chemistry................................................................................................................24
6. Radioactive M&M’s (Nuclear Reactions)
a. Honors Chemistry ....................................................................................................................26
b. Academic Chemistry................................................................................................................30
7. “What’s in the Box?” (Indirect Observations)
a. Honors Chemistry ....................................................................................................................34
b. Academic Chemistry................................................................................................................35
8. Flame Tests
a. Honors Chemistry ....................................................................................................................37
b. Academic Chemistry................................................................................................................39
9. Periodic Poster Lab (Periodicity)
a. Honors Chemistry ....................................................................................................................41
b. Academic Chemistry................................................................................................................42
10. Jelly Bean Lab (Lewis structures/molecular shapes)
a. Honors Chemistry ....................................................................................................................51
b. Academic Chemistry: “Feeling Dotty” Prelab .......................................................................52
“From Dots to Jelly Beans” lab sheet .................................................54
11. Gold Pennies (Alloys/metallic bonding) ..............................................................................................55
12. VSEPR Balloons (Hybrid orbitals) ......................................................................................................56
13. Water’s Wicked Ways .........................................................................................................................57
14. Hydrogen Balloons (Activity series) ...................................................................................................58
15. Embarrassment of Reese’s (Molar volume) ........................................................................................59
16. Solubility Lab (Double-replacement reactions) ...................................................................................60
17. Electrolyte Lab (written for Physical Science) ....................................................................................62
18. Gas Laws .............................................................................................. Error! Bookmark not defined.
19. Boiling Point Elevation ........................................................................................................................65
20. Ice Cream Lab ......................................................................................................................................66
21. Determination of Kf for Half & Half (AP)...........................................................................................67
22. Equilibrium Lab ...................................................................................................................................68
23. %Acidity of Vinegar ............................................................................................................................70
24. Identification of Unknown Metal.........................................................................................................73
25. Heat of Formation of MgO (advanced level) .......................................................................................76
26. Heat of Neutralization ..........................................................................................................................78
27. Iodine Clock (dry lab)
a. Lab sheet ..................................................................................................................................80
b. Lab Data ...................................................................................................................................82
BUTTERED CATS
Background: Recently, some very clever Southeast Raleigh High School students discovered that a
Mythbusters myth had been erroneously ‘busted.’ Specifically, they discovered that toast would always
fall buttered side down if the butter/toast mass ratio were high enough. While investigating the buttered
toast phenomenon, one team of students decided that it would be interesting to also investigate the myth
that cats always land on their feet. To their surprise, they determined that while a cat will always land
on its feet if dropped from a specific height, the cat’s ability to land on its feet decreases if the height is
changed significantly.
Objective: Examine the plotted data and propose an experiment that could be used to answer ONE of
the following questions:
1. Does the butter mass ratio apply to other objects?
2. Can the height at which cats will always land on their feet be changed?
Available Materials
1.0 kg of butter
one 4.8 kg cat
four 40 g pieces of toast
one meter stick
one 2 meter tall ladder
four meters of string
one pair of safety scissors
a box of bandages
Procedure
1. On a separate sheet of paper, propose the question your group would like to address (the hypothesis).
You must combine the ‘cat effect’ with the ‘butter effect’ in some manner.
2. Write a step-by-step procedure that describes how you would conduct your experiment. Be sure to
include a control in your experiment.
3. Include a section entitled “Analysis.” Describe in writing or graphically the outcome of your
experiment assuming that your hypothesis proves to be correct.
4. Answer the questions below on your lab sheet.
Questions (Write out the questions and use complete sentences for your answers)
1. What is a control?
2. What was the control in your experiment?
3. What is a constant?
4. What will be kept constant in your experiment?
5. What is a variable?
6. What is an independent variable?
7. What was the independent variable in your experiment?
8. What is a dependent variable?
9. What is the dependent variable in your experiment?
10. What is necessary for your hypothesis to be accepted?
2
% time toast lands buttered side down
Effect of Dropping Buttered Toast
100
80
60
40
20
0
0.0
1.0
2.0
3.0
4.0
5.0
Amount of Butter (g)
Used (g)
*40 g piece of toast
used
6.0
7.0
Percent of successful cat landings
Percent of successful cat landings vs. dropped height
100
75
50
25
0
0.00
0.50
1.00
1.50
2.00
Dropped height (m)
3
PRECISION vs. ACCURACY
Data Collection
Each person in the lab group should do the following:
Teaspoon station: (Record weight to 2 decimal places, volume to 0 decimal places)
1. Weigh an empty Dixie cup and record the weight in your data table.
2. Fill a teaspoon (volume = 7 mL) with water and add to the Dixie cup.
3. Reweigh the cup and the water and record the total weight in your data table.
Graduated cylinder station: (Record weight to 2 decimal places, volume to 1 place)
1. Weigh an empty 10-mL graduated cylinder and record the weight in a group data
table.
2. Add about 5 mL of water to the graduated cylinder and then record the actual volume
of water.
3. Reweigh the graduated cylinder containing the water and record the total weight in
your data table.
Buret station: (Record weight and volume to 2 decimal places)
1. Weigh an empty Dixie cup.
2. After recording the initial buret volume, deliver about 5 mL of water to the cup and
record the final volume. Calculate the actual volume of water delivered.
3. Weigh the cup containing the water and record in your data table.
Data Analysis
1. Calculate the density of water for each person at each station (should have 12 density
values for a 4 person group).
2. Calculate the group’s average water density for each station (should have 3 average
values).
Answer the questions below on your group’s lab sheet.
Analysis Questions (Please answer in complete sentences)
1. Which station had the greatest variation in individual density measurements? Why do
think this was so?
2. Which station do you think resulted in a density value closest to the ‘true’ value?
Why?
3. Do you think that taking the average of several numbers adequately corrects for
variations in experimental data? Why or why not?
4
PRECISION vs. ACCURACY
Data Collection
Each person in the lab group should do the following:
Teaspoon station: (Record weight to 2 decimal places)
4. Weigh an empty Dixie cup and record the weight in your data table.
5. Fill a teaspoon (volume = 7 mL) with water and add the water to the Dixie cup.
6. Reweigh the cup and the water and record the total weight in your data table.
Graduated cylinder station: (Record weight to 2 decimal places, volume to 1 place)
4. Weigh an empty 10-mL graduated cylinder and record the weight in the group data
table.
5. Add about 5 mL of water to the graduated cylinder and then record the actual volume
of water (e.g. 5.2 mL) in the data table.
6. Reweigh the graduated cylinder containing the water and record the total weight in
your data table.
Buret station: (Record weight and volume to 2 decimal places)
4. After recording the initial buret volume in the data table, deliver about 5 mL of
water into one of the weighed Dixie cups and record the final volume. Calculate the
actual volume of water delivered and enter that value into the data table.
5. Weigh the cup containing the water and record in your data table.
Data Analysis
3. Calculate the density of water for each person at each station (should have 9 density
values for a 3 person group).
4. Calculate the group’s average water density for each station (should have 3 average
values).
Teaspoon (all weights should have 2 decimal places!)
Cup#
Wt Empty Cup
(g)
Wt Cup + H2O
(g)
Wt H2O (g)
Vol. H2O
(mL)
1
7
2
7
3
7
Average of density values using a teaspoon:
Density H2O
(g/mL)
g/mL
5
Graduated Cylinder (weights should have 2 decimal places; volumes should have 1 decimal place!)
Wt. of empty graduated cylinder:
Trial#
Wt Cylinder + H2O
(g)
Wt H2O
Vol. H2O
Density H2O
(g)
(mL)
(g/mL)
1
2
3
Average of density values using a graduated cylinder:
g/mL
Buret (weights AND volumes should have two decimal places!)
Cup
#
Wt
Empty
Cup (g)
Wt Cup +
H2O (g)
Wt H2O
(g)
Initial
Volume
Final
Volume
Vol. H2O
(mL)
Density H2O
(g/mL)
1
2
3
Average of density values using buret:
g/mL
Analysis Questions (Please answer in complete sentences)
4. Which station had the greatest variation in individual density measurements? Why do
think this was so?
5. Which station do you think resulted in a density value closest to the true value? Why?
6. Do you think that taking the average of several numbers adequately corrects for
variations in experimental data? Why or why not?
6
PERCENT OF WATER IN COPPER SULFATE PENTAHYDRATE
Materials
Goggles, goggles, goggles!!
Porcelain crucible and lid
1 chunk of copper sulfate pentahydrate (CuSO4·5H2O)
Water dropper
Bunsen burner
Ring stand with clay triangle
Crucible tongs
Balance
Procedure
1. Weigh the crucible and its lid and record in the data table.
2. Add the chunk of CuSO4·5H2O to the crucible and reweigh. Record the weight in the data
table.
3. Carefully place the crucible, copper sulfate and lid in the clay triangle. Light the Bunsen
burner and heat the copper sulfate until it is completely white (lift the lid from time to
time to check its progress).
4. Answer the questions 1 and 2 in part B while it is heating.
5. Once the copper sulfate no longer appears to be changing, turn off the Bunsen burner
and allow the crucible to cool while sitting in the clay triangle. Answer questions 3 and 4
in part B.
6. After the liner and the solid are cool, reweigh them and record the weight on your
group’s lab sheet.
7. AFTER you have the cooled weight of the dried solid, use dropper to add water just until
the original color has returned. Pay close attention to what is happening when you add
the drops of water, and record any observations.
8. Reweigh the crucible, lid and the solid and record on your lab sheet.
Calculations: Record your answers with the correct number of significant figures and
UNITS.
Wt of empty crucible + lid:
Wt of crucible, lid & CuSO4·5H2O:
Initial wt of CuSO4·5H2O:
Wt of crucible, lid & anhydrous copper sulfate:
Wt of anhydrous copper sulfate:
Wt of water lost by heating:
Wt of crucible, lid & rewetted solid:
Wt of rewetted solid:
Calculate the %H2O in your sample of CuSO4·5H2O.
Given that the %H2O in CuSO4·5H2O is 36.08%, calculate your percent error.
7
Part B: Questions
1. Describe the physical properties of the copper sulfate pentahydrate.
2. If you listen carefully while the copper sulfate is heating, you’ll hear a sizzling sound.
What do you think was causing this sound?
3. Describe the physical properties of the anhydrous copper sulfate.
4. Is the formation of the anhydrous copper sulfate a physical change or a chemical change?
Justify your answer. (There’s no wrong answer to this one, but you DO need to provide a
justification for your choice to get credit for this question.)
5. In addition to returning to its original color, what else did you observe happening when
you added the drops of water to the solid in step 7? Why do you think it did that?
6. How did the two weights of the CuSO4•5 H2O (step 2 weight vs. step 8 weight) compare?
Would you expect them to be the very similar or very different?
7. What are some sources of error that would have contributed to your %error? (Saying that
‘We screwed up” is not a valid source of error.)
8
Separation of a Mixture: An Inquiry Lab
(for groups of 2 or 3 students)
Objective: Separate a mixture of sand and iron filings, and than calculate the % mass of the two
substances in the mixture.
Materials:
Sand and iron filings mixture
2 numbered Dixie cups
2 Scoopulas
2 styrofoam plates
2 magnets
Powder funnel
Balance
Procedure:
1. Weigh each Dixie cup and record the weights in the data table below.
2. Obtain a sample of the sand and iron mixture from your teacher, place it in one of the weighed
Dixie cups and weigh the cup and the sample. Calculate the weight of just the mixture sample.
3. Pour the mixture out on one of the plates, and use the available materials to separate the sand
from the iron. (You have to figure out how to do this.)
DO NOT PUT THE MAGNETS DIRECTLY INTO THE MIXTURE!!!
There will be a 5 point deduction from your lab for doing so.
4. Collect the sand in one of the weighed Dixie cups and the iron in the other, and weigh the two
cups.
5. Calculate the %mass as follows:
% mass sand 
wt of sand
x 100
total sample wt
% mass iron 
wt of iron
x 100
total sample wt
Data:
Wt empty (g)
Wt + sample (g)
Wt + sand (g)
Wt + iron (g)
Cup #1
Cup #2
Calculations:
Sample wt:
g
Sand wt:
g
%mass sand:
%
%mass iron:
%
Iron wt:
g
9
Conservation of Mass
Objective: Use the reaction between acetic acid and sodium bicarbonate to investigate the law of
conservation of mass.
Materials
Baking soda
Vinegar
Balloon
100 mL graduated cylinder
Balance
Metric ruler
Tea candle
Procedure
Weigh the graduated cylinder empty and with about 30 mL of vinegar (a 5% solution of acetic acid) in
it. Weigh the balloon and then carefully add one scoop (~2 to 3 g) of baking soda (sodium bicarbonate,
NaHCO3) to the balloon. Wipe off any NaHCO3 that may have gotten on the outside of the balloon, and
then reweigh the balloon. Being careful not to let any of the NaHCO3 fall into the vinegar, stretch the
mouth of the balloon around the mouth of the graduated cylinder so that they form a good seal. Lift up
the end of the balloon so the NaHCO3 falls into the vinegar. Be sure to keep the balloon sealed on to the
graduated cylinder. Measure the approximate diameter of the balloon and obtain the combined weight
of the graduated cylinder, balloon and their contents once the reaction stops. After obtaining the
combined weight, light the candle, carefully (but quickly) remove the balloon from the graduated
cylinder and ‘pour’ the gas that was generated from the graduated cylinder over the candle flame and
observe what happens. Reweigh the graduated cylinder and balloon.
Discussion Questions
1. Was the total weight of the reactants different from the total weight of the products?
2. What happened to the total weight of the products after the gas was allowed to escape?
3. What happened to the candle flame? What does this tell us about what kind of gas was
generated?
10
Conservation of Mass
Formal Lab Report Guidelines
First page: Title page with names of lab partners
Body of the write-up contains the following sections
Introduction: Contains whatever background information is necessary to allow the reader to
understand the objective of the experiment. The introduction section concludes with an explicit
statement of the experimental objective.
Experimental: Describe what you did to collect your data in prose format (not step by step).
Include sufficient detail so that someone could repeat the experiment if they wished.
Results and Discussion: The first part of this section is where you present your data. For the
Conservation of Mass lab, you should have a data table that looks something like:
Trial
NaHCO3
Vinegar
Wt. 1
Wt. 2
Wt. 3
Wt. of gas Vol. of gas
#
(g)
(g)
(g)
(g)
(g)
(g)
(cm3)
1
2
3
Wt. 1: calculated total weight of balloon, NaHCO3, vinegar and graduated cylinder before mixing.
Wt. 2: total weight after mixing with balloon still attached.
Wt. 3: total weight after releasing gas.
Note that you have to calculate the weight and volume of the gas produced. (Recall that volume of a
sphere is calculated using: Vsphere  4  r 3 )
3
Following the data table is a discussion of your collected data. In addition to specifically addressing the
discussion questions listed at the end of the lab, you should also comment upon any other relationships
you saw in the data. For example, you might want to look to see if the weight of the gas depended upon
the amount of NaHCO3 you had or the amount of vinegar you used. Also, please discuss any possible
sources of error that may have existed and what possible impact these potential errors could have had on
your data.
Conclusion: The conclusion should restate the experimental objective and provide key numerical data
to support whatever conclusion you were able to draw.
The last page should have whatever extra information is needed to complete your report. This is where
you put your sample calculations and cite any sources you may have referenced within the body of your
report.
11
Optimization of Chocolate Chip Ratio
In Toll House Cookies
Group name: The Dough Boys
Lab partners: B. Silly
N.O. Kidding
I.M. Goofy
O. Really
Introduction
Chocolate chip cookies were first prepared in 1937 by
Mrs. Ruth Wakefield in Whitman, Massachusetts using the
recipe shown in Appendix A.1 Since that time, the development
of the ideal cookie recipe has become a multimillion dollar
enterprise. Not only are 7 billion chocolate chip cookies sold
every year, but half of the cookies prepared at home are also
The
Introduction
The
Introduction
should
provide
should
provide
enough
general
enough
general
background
to to
background
indicate
to to
thethe
indicate
reader
why
thisthis
reader
why
research
was
research was
worth
doing
in in
worth
doing
thethe
first
place.
first place.
Appropriate
Appropriate
references
should
references
should
be be
cited
as as
well.
cited
well.
chocolate chip cookies.
Many cookie consumers prefer that their cookies not only
contain a high percentage of chocolate chips, but that they also
be soft and chewy rather than crumbly.
Unfortunately, the
addition of extra chips can cause deterioration of the cookie
dough matrix leading to ready disintegration of the cookie upon
handling. Recently, the “cookie elastometer” shown in Figure 1
was developed by Dr. D. Chip and the American Institute for
Better Cookies. The device consists of 3 knife-edges attached to
a platform.
The elasticity of the cookie is determined by
measuring the indentations left
by the knife-edges after allowing
the device to rest on the cookie
After that, there
should
a there
Afterbe
that,
section
should be a
describing
section WHAT
question
or WHAT
describing
problem
needs
question
or to
be problem
addressed
and to
needs
HOW
it has beenand
be addressed
addressed
HOW it in
hasthe
been
past.
addressed in the
past.
The last section
states specifically
WHAT you are
planning
The lasttosection
investigate
–
states specifically
effectively
stating
WHAT you
are
theplanning
experimental
to
objective.
investigate –
effectively stating
the experimental
objective.
knife edges
for 30 seconds. Dr. Chip has also
measured
the
elasticity
of
cookie
Figure 1: Elastometer
13
standard materials in order to allow for the quantification of the
actual cookie elasticity.
In order to prevent rampant chocolate chip cookie
disintegration, the appropriate ratio of chips to cookie dough
must be determined for a particular cookie recipe. Additionally,
Statement of the
experimental
problem to be
solved.
since consumption of chocolate chip cookies often also involves
serving the cookies on plates, the dynamic stability of the
prepared cookies to stresses induced by dropping and stacking
must also be determined. The effect of increasing the percentage
of chocolate chips in the “official” Toll House recipe for
chocolate chip cookies on both cookie texture and stability will
Ah ha! The
purpose at last!
therefore be examined in this experiment.
Experimental
Three different batches of the cookie dough described in
Appendix A were prepared using 1 cup, 2 cups and 3 cups of
semi-sweet chocolate chips, respectively. Ten dough balls from
each batch were weighed prior to and after baking.
After
The Experimental
Section is just a
general
description of
what you did –
not a step by step
recipe.
calibrating the elastometer with Wrigley’s Spearmint gum (E =
1.0 for a 1 kg platform), the elasticity of five of the
cookies from each batch was determined using the cookie
elastometer.
The dynamic stability of the prepared
cookies was measured by dropping five of the cookies
14
from each batch from a height of 2 feet. Following each test, the
number of cookie pieces with dimensions larger than 1 cm were
counted.
Results and Discussion
The model 104B Elastometer was calibrated with a stick of
Wrigley’s Spearmint gum as described in the literature. The
2
following calibration data were obtained:
Gum indentation: 1.80 mm
Note that single
entries are used
for unique pieces
of data, while
tables are used
for sets of related
data points.
Platform weight: 1.22 kg
Conversion factor = k = 1.22 kg/1.80 mm = 0.678 kg/mm
Table 1 summarizes the experimental data collected during the course of the experiment.
Table 1: Weight loss, static and dynamic stability measurements (in standardized
squishiness units, ssu) for batches 1-3
Batch 1
Batch 2
Batch 3
Cookie Weight
#
loss (g)
E
(ssu)
#
Pieces
Weight
loss (g)
E
(ssu)
#
Pieces
Weight
loss (g)
E
(ssu)
#
Pieces
1
2
3
4
2.1715
1.8296
2.4172
2.3361
0.229
0.291
0.276
0.218
2
1
1
2
3.1339
2.2892
2.2721
2.5619
0.890
0.557
0.765
0.878
2
2
1
3
1.5751
2.3465
2.2216
1.8563
0.434
1.241
0.833
1.252
5
3
5
4
5
1.9585
0.294
3
2.1345
0.665
2
1.7148
0.775
3
Since the softest cookies should retain the most moisture, the average weight loss and
standard deviation (shown in parentheses) was calculated. Sample calculations for Table
2 are shown in Appendix B.
15
Table 2: Average weight loss, static and dynamic stability values
Batch
Avg. Wt. (g)
Avg. E (ssu)
Avg. #Pieces
1
2.1425 (0.2478)
0.262 (0.036)
1.8 (0.8)
2
2.4783 (0.3979)
0.751 (0.142)
2.0 (0.7)
3
1.9428 (0.3299)
0.907 (0.345)
4.0 (1.0)
Examination of the data indicates that the use of 3 cups of
chocolate chips had a detrimental effect upon the dynamic
elasticity of the cookie since this batch had the greatest incidence
of breakage. The low water loss for this batch can be understood
by considering that the chocolate chips are present in a much
larger ratio and thus the cookie dough, which contains the
majority of the water, is not making a large of a contribution.
The static elasticity measurement for this batch seems somewhat
anomalous, but is perhaps not a definitive result due to the
relatively large standard deviation associated with this particular
measurement. It would appear that Batch 2 provides the softest
cookies that tare not overly crumbly from disruption of the
cookie dough matrix.
The much larger static elasticity
coefficient of 0.762 ssu, in particular, makes this recipe for
cookies highly attractive.
A good
discussion of
your data
involves
critiquing what
each piece of
information tells
you and how
much merit there
is in that
information.
Calculating a
standard
deviation along
with an average
gives you a basis
of comparison for
your data. If
something is not
quite as expected,
say why that
could be so (and
“We screwed up”
doesn’t cut it!).
Use your data to
logically reach
some conclusion
about your
experimental
objective.
16
Conclusion
The optimal chocolate chip ratio in Toll House cookies was
investigated.
It was found that using 2 cups of chips in a
standard batch of cookies resulted in cookies with a high static
elasticity coefficient of 0.762 ssu and minimal cookie breakage
during the dynamic test (average of 2.0). An unexpected result
was the high moisture loss during cooking, the highest loss of all
three batches investigated.
Determination of the actual
chip/dough ratio rather than just a net weight would help indicate
where this anomalous loss is occurring.
The conclusion
should contain
numerical results
that describe how
your experimental
objective was
answered. (I know
it seems
redundant, but
necessary since
some readers will
skip the
Experimental
Section in order to
save time. If they
find your results
provocative
enough, they’ll go
back and read the
rest.
Peer Evaluation
N.O. Kidding: 5
I.M. Goofy: 4
O. Really: 5
Be sure to include
an evaluation of
your lab partners
with a 5 indicating
that they made a
significant
contribution to the
investigation and a
1 indicating that
they did not
contribute to the
investigation at all.
17
APPENDIX A: Toll House Cookie Recipe
1 cup of softened butter
2 eggs
1 teaspoon baking soda
¾ cup packed brown sugar
1 teaspoon vanilla extract
2 ¼ cups all purpose flour
2 cups semi-sweet chocolate chips
Cream sugars and butter. Stir in remaining ingredients
except for chips. Fold in chips. Drop dough by rounded
teaspoonfuls onto ungreased cookie sheet. Bake at 375°F
for 9 – 11 minutes.
APPENDIX B: Sample Calculations
Neat, hand -written
calculations are fine
in most cases.
References:
1. Ida Luvacookie, J. of Chocolate, 1987, 45, 111.
2. C. Chip, Advanced Chocolate Chip Cookie Quantification
(1996: Plenum), pp. 7-17.
Be sure to include
any references that
were cited in the
report.
18
Conservation of Mass
Group Name: The Magnificent Seven
B. Silly
Lab Partners
N.O. Kidding
I.M. Goofy
O. Really
19
Introduction
The Law of Conservation of Mass was created by the French Chemist Antoine
Lavoisier, often recognized as the father of modern chemistry, in 1789. The law
states that mass is neither created nor destroyed in any chemical reaction. This
means that the mass of the products of a chemical reaction must be equivalent to
the mass of the reactants. This principle can only be used in classical physics, for
it does not apply in special relativity.
Acetic acid (CH3COOH), also known as ethanoic acid, is the most
important of the carboxylic acids. It is called vinegar when the acid is natural
carbohydrates are oxidized and fermented and creates a dilute solution of acetic
acid. Acetic acids can be found in bananas, apple juice, beef, apricot, blue
cheese, and blueberries. It is the most organic acid and used to manufacture
plastics, insecticides, and a large range of chemical products. Sodium
bicarbonate (NaHCO3), a salt, is a widely used chemical compound that is also
known as baking soda. It is used in cooking/baking and for medical needs. The
reaction between acetic acid and sodium bicarbonate will be examined in the
experiment to investigate the law of conservation of mass.
Experimental
The experiment was completed twice to gather sufficient data and to establish a
mean. To begin preparations, the graduated cylinder alone was weighed and
recorded. Then the graduated cylinder was weighed again but this time with 30
mL of vinegar inside, which has a five percent solution of acetic acid. After, the
balloon was weighed by itself, and then about two to three grams of baking soda
was added to the balloon, which was weighed after that. To proceed to the
experiment, the balloon was carefully and firmly attached to the rim of the
graduated cylinder, without letting any of the baking soda enter the graduated
cylinder and react with the vinegar before the actual experiment was to be
conducted. Once the balloon was on and sealed and the group was ready to
begin, the balloon was lifted to allow the baking soda to drop down into the
graduated cylinder to reach the vinegar. The balloons seal upon the graduated
20
cylinder was held to make sure it stayed on as the baking soda and vinegar
reacted. The reaction was observed and when the two solutions seemed to be
done reacting, the diameter of the balloon was measured and recorded. The
balloon, graduated cylinder, and its contents, still together, were then weighed
and recorded. After, the balloon was taken off the graduated cylinder. The
balloon was pinched at the end so that none of the gases inside of the balloon
could escape. Then, a tea candle was lit, and the balloon's pinched end was
opened to allow the gases hit the flame. Once all of the gases were released, the
graduated cylinder with the solution and the emptied balloon were weighed
together.
Results and Discussion
Using the triple beam balance scale, the following objects were weighed:
First Trial:
Graduated Cylinder: 41.60 g
Graduated Cylinder with Vinegar: 70.84 g
Balloon: 2.97 g
Balloon with Baking Soda: 4.48 g
Diameter of Balloon: 7.5 cm
Graduated Cylinder and Balloon with new Solution: 75.27 g
Graduated Cylinder and Balloon after Gas is Released: 74.79 g
Second Trial:
Graduated Cylinder: 44.98 g
Graduated Cylinder with Vinegar: 71.51 g
Balloon: 3.02 g
Balloon with Baking Soda: 6.01 g
Diameter of Balloon: 9.5 cm
Graduated Cylinder and Balloon with new Solution: 76.69 g
Graduated Cylinder and Balloon after Gas is Released: 76.40 g
Table 1:
Trial
Number
NaHCO3
(g)
Vinegar
(g)
Wt. 1
(g)
Wt. 2
(g)
Wt. 3
(g)
Weight of
gas (g)
Volume of
gas (cm3)
1
1.51
29.24
75.32
75.29
74.29
1.00
221
2
2.99
26.53
76.77
76.64
76.40
0.24
449
Weight 1: Calculated total weight of balloon, baking soda, vinegar and graduated
cylinder before reaction
Weight 2: Total weight after mixing with balloon still attached
Weight 3: Total weight after releasing gas.
21
Once the experiment was completed, it can be seen that the weight of the gas
was almost exactly the same weight as the two solutions - vinegar and baking
soda- put into the experiment, exhibiting the law of conservation of mass. The
weight of the gas affected the total weight of the product created when mixing
baking soda and vinegar. The weight went down when the gas was released,
showing that it was a part of the product and thus necessarily because of the law
of conservation of mass, which states mass can not be created or destroyed. The
flame was affected when the gas was introduced to it; the flame blew out at the
point the gas was released from the balloon. This probably means that the gas
takes in oxygen or rejects oxygen, thus diminishing the flame, or takes in heat,
which also could have been the reason the flame went out. The weight of the gas
depended greatly on how much baking soda we added to the experiment. We
conducted the experiment twice to get different outcomes, but they were still
quite close. Because we did the experiment twice, the amounts and
measurements were different, almost by a large margin as the baking soda for
trial 1 weighed 1.51 while the baking soda for trial 2 weighed 2.99. The amount of
vinegar added were very close to being the same, and throughout the experiment
the scale was off continuously, and the measurements given were most likely not
the exact measurements, due to both human and mechanical error.
Conclusion
The reaction between acetic acid and sodium bicarbonate was evaluated in this
experiment. It was done so to witness the Law of Conservation of Mass. From
the experiment, it can be concluded that the Law of Conservation of Mass is true
and applies to all chemical reactions. When the baking soda and the vinegar
reacted, they created a new solution that was almost exactly the same weight as
the vinegar and baking soda before they were introduced to each other. Human
error caused the weights to be off by 0.1 in trial 1 and 0.88 in trial 2, but the
difference is too small to say that the Law of Conservation of Mass is incorrect.
Whatever the weight of the chemicals going into a reaction must have same
weight as the substance produced after the reaction was created.
22
Peer Evaluation
N.O. Kidding: 5
I.M. Goofy: 4
O. Really:
5
References: "acetic acid." Encyclopedia Britannica. 2008. Encyclopedia Britannica
Online. 20 Jan. 2008 <http://www.britannica.com/eb/article-9003505>. "bicarbonate
of soda." Encyclopedia Britannica. 2008. Encyclopedia Britannica Online. 21 Jan.
2008 <http://www.britannica.com/eb/article-9357220>.
Sample Calculations: (Trial 1)
Wt. of NaHCO3: 4.48 g - 2.97 g = 1.51 g
Wt of vinegar: 70.84 g - 41.60 g = 29.24 g
Wt 1: 4.48 g + 70.84 g = 75.32 g
Wt of gas: Wt2 - Wt3 = 75.29 g - 74.29 g = 1.00 g
Volume of gas = 4/3 π r3= 4/3 π (3.75)3= 221 cm
3
23
CONSERVATION OF MASS
OBJECTIVE: Use the reaction between acetic acid and sodium bicarbonate to investigate the law of
conservation of mass.
MATERIALS
Baking soda
Vinegar
Balloon
100 mL graduated cylinder
Powder funnel
Balance
Metric ruler
Tea candle
PROCEDURE
Weigh the graduated cylinder empty (1) and with about 30 mL of vinegar (a 5% solution of acetic acid)
in it (3). Weigh the balloon (2) and then carefully add one scoop of baking soda (sodium bicarbonate,
NaHCO3) to the balloon. Wipe off any NaHCO3 that may have gotten on the outside of the balloon, and
then reweigh the balloon (4). Being careful not to let any of the NaHCO3 fall into the vinegar, stretch the
mouth of the balloon around the mouth of the graduated cylinder so that they form a good seal. Lift up
the end of the balloon so the NaHCO3 falls into the vinegar. Be sure to keep the balloon sealed on to the
graduated cylinder. Measure the approximate diameter of the balloon in cm (5) and obtain the combined
weight of the graduated cylinder, balloon and their contents once the reaction stops (8) with the balloon
still attached to the graduated cylinder!!! After obtaining the combined weight, light the candle,
carefully (but quickly) remove the balloon from the graduated cylinder and ‘pour’ the gas that was
generated from the graduated cylinder over the candle flame and observe what happens. Reweigh the
graduated cylinder and balloon (9).
DATA
1. Wt empty graduated cylinder:
2. Wt empty balloon:
3. Wt graduated cylinder + vinegar:
4. Wt. balloon + NaHCO3:
5. Diameter of balloon:
6. Volume of balloon:
7. Calculated total weight (#3 + #4):
8. Measured total weight after reaction stops with balloon still attached:
9. Measured total weight after removing balloon:
10. Weight of the gas in balloon:
24
QUESTIONS
4. Was there a big difference between the calculated total weight (7) and the measured total weight
before the balloon was removed (8)? Why do you think this was so?
5. Was there a big difference between the measured total weight before the balloon was removed (8)
and the total weight after the balloon was removed (9)? Why do you think this was so?
6. Calculate the volume of the gas in the balloon using the equation below and enter the result into your
data table (6).
π (diameter in cm) 3
volume 
6
7. Calculate the volume of the gas using the mass of the gas and its density, which is 1.98 x 10-3 g/mL,
and show your work below.
8. Why is there such a large difference between the two calculated volumes?
9. What happened to the candle flame? What does this tell us about what kind of gas was generated?
25
M & M ISOTOPES LAB
Part I: Weighted Proportions
A. Introduction:
In this lab you will study the distribution of the naturally occurring isotopes of the rare element,
Candium, symbol Mm, Candium is found occurring in M & M 's. The isotope found in each M &
M is indicated by the color of the M & M. Below is a list of all the known isotopes of Candium
along with the corresponding M & M color.
M&M Color
Brown
Blue
Yellow
Red
Orange
Tan
Green
Isotopic Symbol
Mm-64
Mm-65
Mm-66
Mm-67
Mm-68
Mm-69
Mm-70
Mass Number
64
65
66
67
68
69
70
BEFORE YOU BEGIN

SAFETY AND WASTE DISPOSAL
1. Safety goggles are optional in this lab.
2. Do not eat any of the isotopes until instructed to do so!
3. Dispose of all wastes as directed by the instructor.
MATERIALS
1 package of M & M 's per group, calculator or spreadsheet program.
B. Procedure
1. Open your package of M & M 's. Separate them by color. Count how many of each color and
enter your data into the table on the next page.
2. Calculate the percent distribution of each of the isotopes.
3. Calculate the weighted proportion of each isotope as follows:
# M& M ' s of 1 color
X Mass Number
Total # of M& M ' s
Repeat for each isotope.
4. To calculate the atomic weight, add up the weight proportions of all the naturally occurring
isotopes.
26
DATA
M&M Color
Mass
Brown
Isotopic
Symbol
Mm-64
Blue
Mm-65
65
Yellow
Mm-66
66
Red
Mm-67
67
Orange
Mm-68
68
Tan
Mm-69
69
Green
Mm-70
70
# in
package
% Distribution
Weighted
Proportion
64
Total in Package
Atomic Weight of Candium in amu
QUESTIONS
1. How does your atomic weight for Candium compare to that of other groups? Why would
there be any differences?
2. Does any single isotope of Candium have a weight equal to that of the atomic weight of the
element Candium? Explain.
3. Did you find any Mm-69 isotopes (tan M & M 's)? Why not? What type of isotopes do the
tan M&M’s represent in this model?
27
Part II: Radioactive Decay of Candium
Question: How long will it take for a radioactive isotope to completely transmutate into a
new element?
Materials
Bag of M&M’s
graph paper or spreadsheet program
Procedure
1. Place all of the Candium atoms from part I back into their bag.
2. Hold the bag shut and gently shake for 10 seconds.
3. Gently pour out candy.
4. Count the number of pieces with the print side up. These atoms have "decayed".
5. Return only the pieces with the print side down to the bag. Reseal the bag.
6. Consume the "decayed" atoms.
7. Gently shake the sealed bag for 10 seconds.
8. Continue shaking, counting, and consuming until all the atoms have decayed (or you
only have one atom left…the model breaks down at that point).
9. Graph the number of undecayed atoms (on the y-axis) vs. ln(time) (on the x-axis).
Data and Observations
Half-life
Total Time
0
0
# of Undecayed Atoms
# of Decayed Atoms
0
1
2
3
4
5
6
7
8
28
Questions
1. What is transmutation?
2. Define half-life.
3. Determine the equation for the best-fit straight line through the data points on your
graph. Use this equation to CALCULATE the half-life of Candium. Show your work
below.
4. At the end of 2 half-lives, calculate the portion (in decimal form) of atoms that had
not decayed. Show your work below.
5. If you had plotted #undecayed atoms vs. time, what would your curve have looked
like? What kind of decay is this?
29
M & M ISOTOPES LAB
Part I: Weighted Proportions
A. Introduction:
In this lab you will study the distribution of the naturally occurring isotopes of the rare element,
Candium, symbol Mm, Candium is found occurring in M & M 's. The isotope found in each M &
M is indicated by the color of the M & M. Below is a list of all the known isotopes of Candium
along with the corresponding M & M color.
M&M Color
Brown
Blue
Yellow
Red
Orange
Tan
Green
Isotopic Symbol
Mm-64
Mm-65
Mm-66
Mm-67
Mm-68
Mm-69
Mm-70
Mass Number
64
65
66
67
68
69
70
BEFORE YOU BEGIN
1. Do not eat any of the isotopes until instructed to do so!
2. Dispose of all wastes as directed by the instructor.
MATERIALS
1 package of M & M 's per group and a calculator
B. Procedure
1. Open your package of M & M 's. Separate them by color. Count how many of each color and
enter your data into the table on the next page.
2. Calculate the percent distribution of each of the isotopes.
%Distribution 
# M & M' s of 1 color
x 100
Total # of M & M' s
3. Calculate the weighted proportion of each isotope as follows:
Weighted proportion 
# M & M' s of 1 color
x Mass Number for that color
Total # of M & M' s
Repeat for each isotope.
4. To calculate the atomic weight, add up the weight proportions of all the naturally occurring
isotopes.
30
DATA
M&M Color
Mass#
Brown
Isotopic
Symbol
Mm-64
Blue
Mm-65
65
Yellow
Mm-66
66
Red
Mm-67
67
Orange
Mm-68
68
Tan
Mm-69
69
Green
Mm-70
70
# in
package
% Distribution Weighted
Proportion
64
Total in Package
Atomic Weight of Candium in amu
QUESTIONS
1. How does your atomic weight for Candium compare to that of other groups? Why would
there be any differences?
2. Does any single isotope of Candium have a weight exactly equal to that of the atomic
weight of the element Candium? Explain.
3. Did you find any Mm-69 isotopes (tan M & M 's)? Why not? What type of isotopes do the
tan M&M’s represent in this model?
31
Part II: Radioactive Decay of Candium
Question: How long will it take for a radioactive isotope to completely transmutate into a new
element?
Materials: Bag of M&M’s and graph paper
Procedure
10. Place all of the Candium atoms from part I back into their bag.
11. Hold the bag shut and gently shake for 10 seconds.
12. Gently pour out candy.
13. Count the number of pieces that are “M” side up. These atoms have "decayed".
14. Return only the pieces that were print side down to the bag. Reseal the bag.
15. Consume the "decayed" atoms so they don’t get mixed back in with the undecayed atoms.
16. Gently shake the sealed bag for 10 seconds.
17. Continue shaking, counting, and consuming until all the atoms have decayed (or you only
have one atom left…the model breaks down at that point).
Data and Observations
Half-life
Total Time
0
0
1
10
2
20
3
30
4
40
5
50
6
60
# of Undecayed Atoms
# of Decayed Atoms
0
Questions
6. What is transmutation?
7. Define half-life.
32
Decay of Candium Atoms
#Undecayed
Candium atoms
8. Graph the number of undecayed atoms (y-axis) vs. the
total time (x-axis). Be sure to give your graph a title and to
label the axes. Draw a single, smooth line (don’t just connect
the dots) that best represents the trend in the data. An
example of what your graph should look like is shown at right.
Time (s)
9. According to your graph, what is the half-life of Candium?
10. At the end of 2 half-lives, calculate the portion (in decimal form) of atoms that had not
decayed. Show your work below.
# undecayed
Portion of undecayed atoms 
Total# at start
33
What’s in the Box?
Introduction
How can scientists explain what they have not actually seen? For example, if no light escapes from a
black hole, how can astronomers locate it? How can scientists characterize things too small to be seen
with the naked eye, like atoms and molecules? You will use indirect observations to determine the
contents of a toaster-pastry box universe that cannot be seen.
Procedure
1. Your teacher will provide you with a toaster-pastry box universe that contains 2 to 3 different
types of objects. Your task is to determine what is in the box.
2. Your team can do any test on the box universe as long as it doesn’t damage it. DO NOT
ATTEMPT TO OPEN THE BOX.
3. Determine the density of the box by weighing it on a balance to the closest 0.01 g and measuring
its dimensions (remember volume of a box is length x width x height) to the closest 0.01 cm.
Record your data in a table like the one below ON A SEPARATE SHEET OF PAPER (one per
lab group).
Box #
Box Weight
(g)
Volume
(cm3)
Density
(g/cm3)
Possible contents
Indirect observations
1
2
3
4
5
6
7
8
4. Record each type of item that you theorize to be in the box in the “Possible contents” column and
the indirect observation that allowed you to detect it in the “Indirect observations” column of
your data table.
5. Complete your data table by entering the collected data from the other 7 groups in the class.
Questions
Write your answers on the paper that has your data table. Be sure to indicate which box was yours.
1. What tests done by your team were most helpful in determining the contents of your universe?
2. All of the boxes contained 2 to 3 different types of objects in various quantities (eg. One of
object A and a handful of object B). Compare the data that you collected to that of the other
groups. Do you agree with all of their conclusions about the box contents? Did observations
made by other groups change any of your conclusions? Please explain both responses (ie. Don’t
answer this question with just a yes or a no).
3. What other tests might you perform to find out more information about what is in the box?
34
What’s in the Box?
Introduction
How can scientists explain what they have not actually seen? For example, if no light escapes from a
black hole, how can astronomers locate it? How can scientists characterize things too small to be seen
with the naked eye, like atoms and molecules? You will use indirect observations to determine the
contents of a toaster-pastry box universe that cannot be seen.
Procedure
6. Your teacher will provide you with a toaster-pastry box universe that contains 2 to 3 different types
of objects. Your task is to determine what is in the box.
7. Your team can do any test on the box universe as long as it doesn’t damage it. DO NOT ATTEMPT
TO OPEN THE BOX.
8. Determine the density of the box by weighing it on a balance to the closest 0.01 g and measuring its
dimensions (remember volume of a box is length x width x height) to the closest 0.01 cm.
9. Record each type of item that you theorize to be in the box in the “Possible contents” column and the
indirect observation(s) that allowed you to detect it in the “Indirect observations” column of your
data table.
10. Complete your data table by entering the collected data from the other 7 groups in the class.
Data
Box length:
cm
Box weight:
g
Box
#
Density
(g/cm3)
Box width:
Possible contents
cm
Box height:
cm
Box volume:
cm3
Indirect observations
1
2
3
4
5
6
7
8
35
Questions
4. What was your box #?
5. What tests done by your team were most helpful in determining the contents of your universe?
6. All of the boxes contained 2 to 3 different types of objects in various quantities (e.g. One of
object A and a handful of object B). Compare the data that you collected to that of the other
groups. Did observations made by other groups change your conclusions about what was in your
box? Explain. (i.e. Don’t answer this question with just a “yes” or a “no”).
36
FLAME TESTS
Just as a fingerprint is unique to each person, the color of light emitted by metals heated in
a flame is unique to each metal. When atoms of elements are heated to high temperatures,
the electrons absorb quanta of energy (become excited) and move to a higher energy level.
When the electrons no longer have sufficient energy to remain in the higher energy level,
they fall back down to a lower energy level and emit excess energy as a photon of light. The
energy (and the color) of the photons emitted, E, is equal to the difference in energy
between the two energy levels and can be found from the following expression:
E = hc/
Where h (Planck’s constant) is 6.63 x 10-34 J·s, c (the speed of light) is 2.99 x 108 m/s, and 
is the wavelength of the photon emitted in meters. While energy is inversely proportional to
the wavelength of a photon, energy is directly proportional to the frequency, , of the
photon:
E=h
Where  is in units of cycles/s or Hz.
OBJECTIVES
 Test various metal salt solutions in a hot flame and observe the characteristic color
given off by each excited atom
 Identify the metal ion present in two unknown metal salt solutions.
 Use a spectrometer to measure the two most prominent emission lines for both the
known and unknown metal salt compounds.
SAFETY: This lab is an all-time favorite among students, but the following safety rules
MUST be followed for everyone’s sake.
 NO NAKED EYEBALLS!
 NEVER LEAVE A LIT BUNSEN BURNER UNATTENDED!!!



ABSOLUTELY NO HORSEPLAY!!! I’LL THROW YOU OUT OF THE ROOM AND GIVE YOU
A ZERO FOR THE LAB IF YOU START FOOLING AROUND AT ALL!
TIE BACK LONG HAIR, ROLL UP LOOSE SLEEVES AND TIE BACK ANYTHING THAT
POSES A POTENTIAL FIRE HAZARD.
Be careful not to touch anyone or anything with the hot nichrome wire.
PROCEDURE
1. Being careful to keep track of the location of each known metal salt compound, put a
SMALL amount (covering just the TIP of the scoopula!!!) of each compound into a
cleaned 12 well tray.
2. Add 10-15 drops of distilled water to each well. The idea here is to make a watery
paste…not to dissolve the compound.
37
3. After lighting your Bunsen burner, clean your nichrome wire by rinsing the loop in
dilute HCl, followed by distilled water, put the loop into the flame then cool by
rinsing in a second beaker of distilled water. Repeat several times until there is very
little color emitted when the loop is first placed into the flame.
4. After cooling your wire in the second beaker of distilled water, dip the cleaned wire
into your first known compound solution/paste. Record the emission color and the
identity of the metal ion present in your notebook. Being careful not to lean too
close to the flame (this will take some serious teamwork), use the spectrometer to
measure the wavelengths of the two most prominent emission lines and record in your
notebook.
5. Repeat step 4 for the remaining known and unknown compounds being careful to
clean the wire as in step 3 between each compound (you don’t want to be
contaminating your flame color with previous compounds!).
DATA
1. Calculate the energy of each emission line you observed (remember to change nm
into meters first).
2. Organize your observations, measurements and calculated energies into an
appropriate data table.
Questions for Discussion
1. What caused the colors that you observed during this lab? Why were the colors different?
2. Compare the position in the periodic table of each element examined in this experiment
to the energy of the photons that are emitted. Do you observe any trends?
3. What were the metal ions in your two unknowns? Justify your answer.
4. What were some possible sources of error that could have complicated proper
identification of your unknown compounds?
Peer evaluation
Be sure to include an evaluation of how helpful each of your lab group members were in
completing this lab in your formal lab report using the following scale:
Score
5
4
3
Fantastic
Good
OK
2
Helped a little
1
They were in
my group??
38
FLAME TESTS
Just as a fingerprint is unique to each person, the color of light emitted by metals heated in
a flame is unique to each metal. When atoms of elements are heated to high temperatures,
the electrons absorb quanta of energy (become excited) and move to a higher energy level.
When the electrons no longer have sufficient energy to remain in the higher energy level,
they fall back down to a lower energy level and emit excess energy as a photon of light. The
energy (and the color) of the photons emitted, E, is equal to the difference in energy
between the two energy levels and can be found from the following expression:
E = hc/
Eq. 1
Where h (Planck’s constant) is 6.63 x 10-34 J·s, c (the speed of light) is 2.99 x 108 m/s, and 
is the wavelength of the photon emitted in meters.
OBJECTIVES
 Test various metal salt solutions in a hot flame and observe the characteristic color
given off by each excited atom
 Identify the metal ion present in two unknown metal salt solutions.
 Use a spectrometer to observe the emission lines in a mercury gas discharge tube (a
fluorescent lamp).
SAFETY: This lab is an all-time favorite among students, but the following safety rules
MUST be followed for everyone’s sake.
 NO NAKED EYEBALLS!
 NEVER LEAVE A LIT BUNSEN BURNER UNATTENDED!!!



ABSOLUTELY NO HORSEPLAY!!! I’LL THROW YOU OUT OF THE ROOM AND GIVE YOU
A ZERO FOR THE LAB IF YOU START FOOLING AROUND AT ALL!
TIE BACK LONG HAIR, ROLL UP LOOSE SLEEVES AND TIE BACK ANYTHING THAT
POSES A POTENTIAL FIRE HAZARD.
Be careful not to touch anyone or anything with the hot nichrome wire.
PROCEDURE
6. Being careful to keep track of the location of each known metal salt solution, put a
SMALL amount (enough to fill a well ¾-full) of each solution into a cleaned 12 well tray.
7. After lighting your Bunsen burner, clean your nichrome wire by rinsing the loop in dilute
HCl, followed by distilled water, put the loop into the flame then cool by rinsing in a
second beaker of distilled water. Repeat several times until there is very little color
emitted when the loop is first placed into the flame.
8. After cooling your wire in the second beaker of distilled water, dip the cleaned wire into
your first known solution. Record the emission color next to the corresponding metal ion
in your data table.
39
9. Repeat step 8 for the remaining known and the two unknown compounds being careful to
clean the wire as in step 7 between each compound (you don’t want to be contaminating
your flame color with previous compounds!).
10. Look through the spectrometer at a fluorescent lamp, and record the colors of the emission
lines and their measured wavelengths from the scale in the spectrometer.
DATA
Flame Colors
Li+1
Na+1
K+1
Ca+2
Sr+2
Ba+2
Cu+2
U#1
U#2
Line colors and wavelengths from spectroscope:
QUESTIONS
5. What caused the colors that you observed during this lab? Why were the colors different?
6. What were the metal ions in your two unknowns? How do you know?
7. What were two possible sources of error that could have complicated proper
identification of your unknown compounds? (“We screwed up” is not a valid response.)
8. Calculate the energy of the purple line you observed using the spectroscope. (If the
purple line fell on the 4.5 on the spectroscope’s scale, the wavelength of the line was 4.5
x 10-9 m.)
40
Honors Chemistry
P
Er
I
O Dy K
P Os Te Re
Objective: To assemble a periodic table that illustrates the trends in atomic and ionic radii,
electronegativity and ionization energy values of elements 1 through 86.
Note: Neatness counts in this assignment and points will be deducted for unnecessarily messy projects.
However, bonus points can be earned for creativity or unusually attractive projects.
Directions:
1. Use the periodic table with atomic and ionic radii to identify the element symbol for each paper
circle.
2. Neatly write the electron dot structure for the atom or ion in the center of the circle. If the circle
represents an ion, include the appropriate oxidation number. (See examples below)
167
Li
68
Li+1
3. Once you have all of the atoms and ions labeled, carefully cut out each circle and arrange onto your
diagram of the periodic table on the piece of poster board. (You might want to wait to draw the lines
separating out the blocks on the periodic table until after you have an idea of the space
required…notice that you’re also going to be adding values for the ionization energy and
electronegativity to the blocks.)
4. Neatly glue each atom and ion into place then label the ionization energy in one color and
electronegativity in another color. (Data on next page)
5. Include a key for the different colors involved and be sure to give your periodic table a title.
6. Write your names on the back of the periodic table and submit.
41
Chemistry
P
Er
I
O Dy K
P Os Te Re
Objective: To assemble a periodic table that illustrates the trends in atomic and ionic radii,
electronegativity and ionization energy values of elements 1 through 86.
Note: Neatness counts in this assignment and points will be deducted for unnecessarily messy projects.
However, bonus points can be earned for creativity or unusually attractive projects.
Directions:
7. Use the periodic table with atomic and ionic radii to identify the element symbol for each paper
circle.
8. For the atoms (blue paper): write the element symbol and its electron dot structure in their circle.
9. For the ions (yellow paper): write the element symbol and the oxidation number for the ion in the
circle. (See examples below)
167
Li
68
Li+1
10. Once you have all of the atoms and ions labeled, carefully cut out each circle and arrange onto your
diagram of the periodic table on the piece of poster board. (You might want to wait to draw the lines
separating out the blocks on the periodic table until after you have an idea of the space
required…notice that you’re also going to be adding values for the ionization energy and
electronegativity to the blocks.)
11. Neatly glue each atom and ion into place then label the ionization energy in one color and
electronegativity in another color. (Data on next page)
12. Include a key for the different colors involved and be sure to give your periodic table a title.
13. Write your names on the back of the periodic table and submit.
42
Add the Following data to your periodic table. Be sure to use one color ink/colored pencil for the
ionization energies and a different one for the electronegativity values.
Element
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
1st I.E. (eV)
13.6
24.5
5.4
9.3
8.3
11.3
14.5
13.6
17.4
21.6
5.1
7.6
6.0
8.2
10.5
10.4
13.0
15.8
4.3
6.1
6.6
6.8
6.7
6.8
7.4
7.9
7.9
7.6
7.7
9.4
6.0
7.9
9.8
9.8
11.8
14.0
E-neg
2.20
---0.98
1.57
2.04
2.55
3.04
3.44
3.98
---0.93
1.31
1.61
1.90
2.19
2.58
3.16
---0.82
1.00
1.36
1.54
1.63
1.66
1.55
1.83
1.88
1.91
1.90
1.65
1.81
2.01
2.18
2.55
2.96
2.96
Element
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
Cs
Ba
La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
1st I.E. (eV)
4.2
5.7
6.2
6.6
6.8
7.1
7.3
7.4
7.5
8.3
7.8
9.0
5.8
7.3
8.6
9.0
10.5
12.1
3.9
5.2
5.6
6.8
7.5
7.9
7.8
8.4
9.0
9.0
9.2
10.4
6.1
7.4
7.3
8.4
9.3
10.7
E-neg
0.82
0.82
0.95
1.22
1.33
1.6
2.16
1.9
2.2
2.28
2.2
1.93
1.69
1.78
1.96
2.1
2.05
2.1
2.66
2.6
0.79
1.3
1.27
1.3
1.5
2.36
1.9
2.2
2.2
2.28
2.54
2
1.62
2.33
2.02
2
43
H
2
3
4
5
53 Atomic diameter
12
Li
Be
167 112
68 31 Ionic diameter
Na Mg
190 145
95 65
K
Ca Sc
Ti
V
243 194 184
176
171
133 99 81(+3) 68(+4) 52(+5)
Rb Sr
Y
Zr
Nb
265 219 212
206
198
148 113
Cs Ba La
Hf
Ta
298 253 217
208
200
169 135
6
7
8
9
10
11
12
13
14
15
B
87
C
67
Al
118
Si
111
N
56
171
P
98
212
As
114
222
Sb
133
Cr
166
52(+6)
Mo
190
Mn
161
46(+7)
Tc
183
Fe
156
64(+3)
Ru
178
Co
152
74(+2)
Rh
173
Ni
149
72(+2)
Pd
169
Cu
145
72(+2)
Ag
165
Zn
142
74(+2)
Cd
161
Ga
136
62(+3)
In
156
Ge
125
53(+4)
Sn
145
W
193
Re
188
Os
185
Ir
180
Pt
177
Au
174
Hg
171
Tl
156
Pb
154
16
O
48
140
S
88
184
Se
103
198
Te
123
221
Bi Po
143 135
17
He
31
F
42
136
Cl
79
181
Br
94
196
I
115
216
At
127
227
Ne
38
Ar
71
Kr
88
Xe
108
Rn
120
Ionic
227
221
216
46
62
135
212
72
140
184
181
198
72
45
Atomic
136
184
125
217
212
176
206
208
88
265
190
167
103
46
Atomic
98
31
145
161
243
166
219
38
156
149
194
171
53
118
42
112
298
253
111
145
48
47
Atomic
174
165
136
171
161
114
154
145
156
56
156
71
94
135
133
123
143
142
87
67
115
88
152
120
127
108
79
48
Atomic
198
200
193
180
169
183
190
188
177
185
178
173
49
Ionic
31
65
95
52
12
99
171
222
68
133
148
81
68
52
113
169
136
53
64
74
196
74
50
From Dots to Jelly Beans!
Materials: 7 three-inch balloons, minimum of 70 multi-colored jelly beans (cheap, stale ones work
best), round toothpicks and masking tape.
Safety: Be careful when inserting toothpicks to avoid jamming the toothpick under your fingernail or
into the side of your finger.
Directions:
EVERYONE IN THE GROUP MUST COMPLETE THEIR OWN DOT STRUCTURE
TABLES!!!
1. In a table, neatly draw the electron dot structure for the following atoms:
Hydrogen, oxygen, nitrogen, boron, fluorine, sulfur, phosphorus, aluminum, chlorine, carbon
2. In a second table, make the Lewis dot structures for the following compounds:
H2, F2, HF, HCl, H2O, H2S, NH3, BF3, PCl3, BH3, AlCl3, H2O2, CH4, CCl4, CCl2F2
Bonus compounds: O2, N2, C2H4, C2H2
3. Inflate your balloons so that they are all approximately the same size. Make two pairs by tying
the ends of two balloons together. Make one triple by tying the ends of the remaining three
balloons together.
4. Once your Lewis dot structures have been approved by your teacher, use different colors of
jellybeans for each type of atom and the toothpicks for the bonds between the atoms to build
your jellybean molecules. Use half of a toothpick to show the location of any lone pairs on the
molecules. REMEMBER: the electrons in the bonds repel each other so the bonds and lone pairs
need to be positioned so that they’re as far away from each other as possible. (If you are using
marshmallows instead of jelly beans, use a magic marker to color code the marshmallows for the
different types of atoms.)
5. Use your balloons to help you build the correct 3D model. The knot at the center represents the
central atom while the inflated balloons represent the electrons in a bond or in a lone pair.
6. Label each one of your jelly bean molecules with the chemical formula (e.g. H2O) using a folded
over piece of masking tape.
7. Be sure to put all lab group members’ names on a piece of paper that will be submitted with the
jelly bean models.
Feeling
Dotty
Add up the number of valence electrons for each compound and write the number on the
line in the box. Then, draw the Lewis dot structure for each compound in the box. (The
central atom is underlined.)
H2
#VE:_____
F2
#VE:_____
HF
#VE:_____
BH3
#VE:_____
BF3
#VE:_____
H2O
#VE:_____
H2S
#VE:_____
PH3
#VE:_____
52
PF3
#VE:_____
NH3
#VE:_____
CH4
#VE:_____
CF4
#VE:_____
CH2F2
#VE:_____
O2
#VE:_____
N2
#VE:_____
HCN
#VE:_____
Bonus: SF4
#VE:_____
Bonus: C2H4
#VE:_____
53
From
Dots to Jelly
Beans!
Materials: Minimum of 70 multi-colored jelly beans (cheap, stale ones work best), round
toothpicks and masking tape.
Safety: Be careful when inserting toothpicks to avoid jamming the toothpick under your
fingernail or into the side of your finger.
Directions:
8. Everyone in the group must complete their own copy of the “Feeling Dotty” Lewis Dot
Structures pre-lab worksheet.
9. Once your Lewis dot structures have been approved by your teacher, use different
colors of jellybeans for each type of atom and a full-length toothpick for each bond
between the atoms to build your jellybean molecules. Use half of a toothpick to show
the location of any lone pairs on the molecules. REMEMBER: the electrons in the
bonds and in the lone pairs repel each other and need to be positioned so that they’re
as far away from each other as possible. Use your balloon models to help you get the
best 3D geometry.
10. Label each one of your jelly bean molecules with the chemical formula (e.g. H 2S)
using a folded over piece of masking tape.
11. Be sure to put all lab group members’ names on the tray holding your jelly bean
models.
12. Make sure everyone submits their “Feeling Dotty” pre-lab worksheet to the “In-Box”
and your tray in the designated area.
54
Preparation of "Gold" Pennies
Bring 2 pennies with you to lab (the cleaner and shinier the better).
PRELAB QUESTIONS: Answer all Prelab and Procedure Check questions on your lab group’s paper (one per
group) BEFORE starting the lab.
1. What is the name given to an alloy made from the two metals found in a penny?
2. Pennies minted prior to 1982 are made from an alloy composed of approximately 98%
copper and 2% zinc.
a. Why might a penny be made of an alloy instead of pure copper?
b. If one of these pennies were examined microscopically, it would appear uniform
throughout. Is this a heterogeneous alloy or a homogeneous one?
c. Would the alloy in the penny be classified as an interstitial alloy or a solid solution?
d. Sketch a picture of what the penny alloy would look like on the atomic level.
3. Pennies minted after 1982 are composed of a zinc core surrounded by a copper shell.
a. Would this be considered an alloy? Explain.
b. Why do you think a copper shell is used?
Procedure Check Questions
A. In order to make the silver-colored penny, the penny is placed in a beaker containing what solution
and what solid material?
B. What will you use to heat the silver colored penny?
C. How should you heat the silver penny (strongly or gently)?
PROCEDURE: Pair up with someone. While each person should prepare their own pennies, the two
people can both use the same beaker containing the zinc and zinc chloride solution.
1. Place ~1 g of granular zinc in a 150 mL beaker on a hot plate.
2. Add enough 1 M zinc chloride solution to cover the zinc metal and heat to a gentle boil.
3. Add a penny and simmer until the penny becomes silvery. Stir it occasionally.
4. Use tongs to remove the penny from the evaporating dish and rinse it with water (dropping it in a
beaker of clean tap water works well).
5. Using a pair of forceps (tweezers) GENTLY heat the penny on a hot plate until it turns gold. Flip
the penny back and forth while heating. You can ruin your
penny by overheating it.
6. Drop the penny in the beaker of rinse water. Remove the penny from the beaker and dry.
POSTLAB QUESTIONS: Answer on your lab group’s lab sheet.
1. What metal made the penny appear silvery?
2. What made the penny appear gold in color?
3. If the penny is heated to strongly it changes back to a copper color. Suggest a reason for this.
4. If you had actually changed your penny into gold, how much would it be worth? You may use
any means you wish to find the selling price of gold.
55
VSEPR Balloons
Objective: To construct three dimensional balloon models of various molecular
geometries.
Materials: Multicolored 3” balloons, Sharpie marker, string, index cards for labels.
Directions:
1. Draw the Lewis dot structures for the following 13 molecules: (see tutorial on my website
Boxall.srhs.net for help)
Note: For full credit on this part of the lab, I want evidence that each person in the group
to draw the structures for at least three of the molecules…i.e. I expect to see ~4 different
handwriting styles and the name of the person that completed each structure.
BeH2, BH3, CH4, NH3, H2O, PCl5, SF6, NO2-1, SF4, ClF3, I3-1, IF5, XeF4
2. Label each dot structure with the compound name and the appropriate AXE notation.
(This portion is worth ~ 30% of your lab grade)
3. Next to your Lewis dot structures, identify the molecular geometry (the geometry of
just the atoms…no lone pairs) of your molecules using one of the following terms: (2
terms are repeated, all terms are used, worth 13% of lab grade)
Bent, linear, octahedral, see-saw, square planar, square pyramid,
T-shape, tetrahedral, trigonal-planar, trigonal-pyramidal,
trigonal-bipyramidal
4. Balloon models of the molecules can be made by tying either two or three balloons
together. The knot at the center represents the central atom and each balloon
represents either a terminal atom (all the same color) or a lone pair (a different color
than the terminal atoms).
5. Combine the dyad and triad sets of balloons together to make 3-dimensional models of
the molecules being sure to place the lone pairs (you might want to blow those balloons
up a bit bigger) where their electron repulsion from other lone pairs or bonds is
minimized. For example, I3 ends up being a linear molecule with the three lone pairs in
equatorial positions in order to maximize the separation between them.
6. Draw two dots on the lone pair balloons to represent the two electrons. Use a piece of the
index card and the string to label your molecule with its formula.
7. When completed, hang your balloon models up on the hooks near your lab station and turn
in your Lewis dot structures (make sure everyone’s name is on the group lab sheet and on
the balloon models).
56
Water’s Wick-ed Ways: WATER DISCOVERY LAB
1. Float a paperclip in a beaker of cold water. What happens if you add one drop of soap to the water?
2. Pour water from one beaker to another using a piece of string. Experiment with different pouring
angles and distances.
3. How many drops of water can you fit onto a penny? Dry the penny off and then see how many
drops of rubbing alcohol you can fit onto the penny.
Questions
1. Soaps consist of a long hydrophobic tail and a polar, hydrophilic head group. How might the
molecular structure of the soap affect the surface tension of the water in the beaker?
2. Why was it necessary to wet the string before trying to pour the water from beaker to beaker in
#2?
3. The stick structure of rubbing alcohol (isopropanol) is shown below. Why would the expected
result for #3 be that the number of drops of water is greater than the number of drops of
isopropanol?
H
H O H
H C
C C
H
H H H
57
The Activity Series:
Hydrogen Balloons
Background: The dissolution of metals in an acid is an example of a single replacement reaction as
shown below:
M(s) + 2 HX(aq)  MX2(aq) + H2(g)
where M is a metal and HX is an acid. In order for a metal to replace hydrogen in the compound, the
metal must be above hydrogen in the activity series. The more active the metal is (the higher up on the
activity series), the faster the reaction will progress, and the faster the hydrogen gas will be generated.
Objective: To experimentally determine the relative positions of four metals on the activity series with
respect to the acidic proton in hydrochloric acid.
Materials:
3 M HCl
Graduated cylinder
Copper wool
Iron filings
Mossy zinc
Tin shot
Porcelain twelve well plate
Florence flask
12” Balloon
2” piece of string
Procedure:
1. Put a small amount of each metal into separate wells on the porcelain plate.
2. At your lab station, add enough HCl to cover the metals.
3. After you have determined the order of activity (most bubbles to least), list them in the correct
order below and come tell me which metal has the greatest activity.
4. When I have approved your selection, put 5 g of the metal into the Florence flask and add 100
mL of the acid to the flask. Stretch the mouth of the balloon over the mouth of the flask to collect
the hydrogen gas. Carefully swirl the flask (don’t get any solution up into the balloon) to keep
the reaction as vigorous as possible.
5. When the balloon is no longer expanding (or the diameter is approximately 6”), pinch the neck
of the balloon so the hydrogen doesn’t escape, remove the balloon and tie it off.
6. Tie a piece of string to the balloon and take it over to the ignition station.
7. Dangle the balloon from a ring on a ringstand, and use the candle taped to a meter stick to ignite.
8. Boom! 
Activity Series:
________
________
Most Active
________
________
Least Active
58
An Embarrassment of Reese’s
Goal: To determine the molar volume of three different materials.
Materials:
Unpopped popcorn
Uncooked long grain rice
Small candies (Reese’s pieces, plain M&Ms, Skittles, etc.)
Metric rulers
Balance
Selection of graduated cylinders
Directions:
One mole of water (H2O) weighs 18.0 g. The density of water is 1.0 g/cm 3, so
one mole of water would have a molar volume of 18.0 cm3. This volume can be
pictured as a cube with sides of 2.6 cm each and a footprint area (the area that
the cube is resting on) of 6.9 cm2, OR an 18.0 cm tall column of water with a
1.0 cm2 footprint OR any other shape in between that has a total volume of
18.0 cm3. Your objective is to determine the molar volume of the popcorn, rice
and candy pieces (please note what kind of candy you used), and to express that
volume in terms of the most reasonable* footprint area from the list given
below and the corresponding thickness. Please record all measurements neatly
in the data section below (and on the back of this paper, if necessary), and
remember to use the appropriate number of significant figures.
Footprint areas
Football field: 5.35 x 103 m2
V1
=
V2
Wake County: 2.21 x 109 m2
North Carolina: 1.40 x 1011 m2
Contiguous 48 States: 9.63 x 1012 m2
* Please explain why your chosen footprint area is reasonable.
Data
59
SOLUBILITY LAB
Prelab: In the introduction for your lab report, include a definition of solubility, lattice energy and a
brief description of how the size and charge of the ions in an ionic compound affect both
solubility and lattice energy.
Goals: (1) To systematically combine ten aqueous solutions and record any chemical changes.
(2) To write net ionic equations for all of the observed chemical reactions.
(3) To formulate some general solubility rules based on experimental observations.
Safety: Several of these compounds are toxic. Please be sure to wash your hands at the conclusion of
the lab and ABSOLUTELY NO NAKED EYEBALLS PERMITTED!!! If you can’t be
bothered to keep your goggles on, I can’t be bothered to give you a grade for doing the lab.
Procedure:
1. Clean a twelve-well tray and rinse well (soap will react with some of the solutions).
2. Clean the plastic pipets with several changes of tap water.
3. Go to Reagent Central and get a sample of each solution (fill the wells about ¾ full). BE SURE
TO WRITE DOWN WHAT SOLUTION YOU’RE PUTTING IN EACH WELL!!!
4. Binary mix the ten solutions (e.g. 1 drop of solution A and 1 drop of solution B) using the
reaction matrix chart as a guide. Use a clean stirring rod to mix the drops if necessary. If you
think you’ve mixed up your pipets, rinse them out before using.
Caution: Drop the drops – do not touch the solution on the plastic with the tip of
the pipet or you will contaminate everything!
5. Devise an accurate code to describe any precipitates that you see. (e.g. ‘bird dropping white’ not
just ‘white precipitate’) One of the reactions will make bubbles.
6. When you and your lab partners are satisfied that all of the ‘No Reaction’ spots really don’t have
a reaction (and there are some), rinse the solutions in the 12 well tray down the drain, clean off
the sheet protector and rinse out the pipets with water. (DO NOT throw the pipets away!)
Data
1. Draw a picture of the reaction matrix chart in your lab notebook. Include a key to your code for
the reaction observations.
2. For every precipitate forming reaction that you observed, write a balanced net ionic chemical
equation.
Discussion
In your discussion section, use your observed reactions to formulate some general solubility rules. For
example, we’re starting with sodium salts of a lot of compounds so it would appear that compounds with
sodium cations have a high solubility. Look at which cations and anions produced the most
precipitates…this will help guide you in formulating your solubility rules. DO NOT include solubility
rules for compounds that you did not observe. For example, don’t include the fact that mercury(I) and
silver form insoluble chlorides because you personally did not observe this fact.
Conclusion
In your conclusion section, include a list of your solubility rules.
60
Materials:
Need about 50 mL per 30 students of each of the following in wash bottles
0.2 M Na3PO4
8.2g /250 mL
1.0 M NaOH
1.0g / 250 mL
0.2 M Na2CO3
5.3g / 250 mL
0.2 M NaI
8.3g KI/ 250 mL
0.2 M Pb(NO3)2
16.6g / 250 mL
1.0 M HCl
21 mL of 12M in 250 mL
0.1 M FeCl3
6.8g FeCl3•6H2O/ 250 mL
0.2 M CuSO4
12.5g CuSO4•5H2O / 250 mL
0.1 M Co(NO3)2
7.3g Co(NO3)2•6H2O/ 250 mL
0.1 M Ba(NO3)2
6.5g / 250 mL
(10) 12 well trays
(10) glass stirring rods
(10) beakers for rinse water
(10) sheet protectors with reaction matrix
(30) disposable pipettes (3 per lab group)
The lab groups should have no more than 3 people per group in order to ensure maximal student
participation.
61
Group Member Names:
ELECTROLYTE SOLUTIONS
Background
Metals are not the only substances that conduct electricity. Electrolytes also conduct electricity and are
essential to the proper functioning of our bodies. Strong electrolytes are compounds that will completely
ionize in water and will conduct electricity easily. Weak electrolytes do not completely ionize in water
and are limited in how much electricity they can conduct. Nonelectrolytes do not ionize at all and will
not conduct any electricity.
In this lab, you are going to compare the conductivity of solutions of salt, sugar and Gatorade to evaluate
whether these substances are electrolytes or nonelectrolytes.
1. Using the plastic weigh boat labeled “SALT,” weigh out 2.5 g of table salt.
2. Put the weighed portion of salt into a beaker and add 50 mL of tap water using a graduated
cylinder.
3. Stir the solution with a stirring rod until all of the salt has dissolved.
4. Put both ends of the conductivity tester into the salt water and record the brightness of the light
bulb (Salt BB) in your data table. DO NOT ALLOW THE TWO WIRES TO TOUCH
EACH OTHER!!!! I’m going to take 5 points off of the lab grade for any group that I see
doing that.
5. Weigh out another 2.5 g of table salt as before and add to the beaker containing your solution
(you’re making the salt water saltier).
6. Put both ends of the conductivity tester into the salt water and record the brightness of the
lightbulb using the saltier water in your data table.
7. Add 50 mL of water to your salt water solution and test with the conductivity tester. Record the
bulb brightness.
8. Rinse out your beaker 5 times with water and then repeat steps 1-7 with sugar instead of salt.
Use the plastic weigh boat labeled “SUGAR.”
9. Rinse out your beaker and get a sample of Gatorade.
10. Observe the brightness of the bulb on the conductivity tester with Gatorade and record in the data
table.
11. Calculate the m/v% concentration of the salt and sugar solutions.
12. Answer the questions on the back of this sheet.
Sample
2.5 g + 50 mL H2O
Salt BB*
Salt m/v%
Sugar BB*
Sugar m/v%
Gatorade
5 g + 50 mL H2O
5 g + 100 mL H2O
No water added
62
Questions
1. Which solution produced the brightest bulb? What was its m/v% concentration?
2. Is salt water an electrolyte or a nonelectrolyte? How can you tell?
3. Is sugar water an electrolyte or a nonelectrolyte? How can you tell?
4. Gatorade commercials tell us that it contains electrolytes. Do you think this is true? If so, estimate
the m/v% concentration of the electrolyte in Gatorade.
63
Having a Gas with the Gas Laws
1. Put about 1 cm (width of a finger) of water in an aluminum soda can. Heat the can on
the hotplate until the water is boiling vigorously. Take the beaker tongs and quickly put
the can into the bucket of water. There are a couple of ways to do this…only one will
result in a satisfying sound. (i.e. getting the can to float is not the effect that you’re
after). Record your observations in the space below and then explain your
observations/results in terms of gas pressure, volume and temperature.
2. While you’re waiting for the water in #1 to heat up, get a Cartesian diver containing food
coloring from the table in the back of the room. Fill your 2L bottle most of the way up
with tap water and drop the diver into the bottle. Put the cap on the 2L bottle and then
squeeze the bottle until the diver begins to drop (if it’s too difficult to get it to drop, try
adding some more water to the 2L bottle). Observe what happens inside the diver and at
the open end of the diver. Record your observations in the space below and then explain
your observations/results in terms of gas pressure, volume, Pascal’s principle1, and
buoyancy (Archimedes’ Principle2).
3. AFTER you have completed and cleaned up from parts 1 and 2, get a piece of dry ice
from me and carefully put it into a film canister. Put the top on the film canister and
observe what happens. DO NOT HANDLE THE DRY ICE WITH YOUR HANDS!!! The
temperature of the dry ice (-67C) will ‘burn’ your fingers!! Record your observations in
the space below and then explain your observations in terms of gas pressure, volume
and number of moles of gas.
Pascal’s Principle: Pressure applied to an enclosed fluid will be distributed evenly among all surfaces that the
fluid contacts.
2
Archimedes’ Principle: a body immersed in a fluid is buoyed up by a force equal to the weight of the fluid
displaced by the object.
1
64
Some Like It Hot!!
Boiling Point Elevation
Objective: To determine the boiling point elevation constant for water.
Procedure
1. Weigh an empty 250 mL beaker, then fill to the 100 mL mark on the beaker and reweigh.
Calculate the weight of the water in the beaker in kg.
2. Put the beaker on the hot plate and heat to boiling.
3. While the water is heating, weigh out 10 g* of NH4NO3 into each of the three weighing boats at
your station. Make sure you know which one is which.
4. Once the water is boiling, measure the boiling temperature with the thermometer. Do not allow
the tip of the thermometer to touch the bottom of the beaker.
5. Add the contents of weighing boat #1 to the beaker (the water should stop boiling) and stir with
the stirring rod until the ammonium nitrate dissolves. Add water to the beaker if the level of the
water has dropped below the 100 mL mark.
6. Measure the boiling point of the solution once it begins to boil again.
7. Repeat steps 5 and 6 for the contents of weighing boats 2 and 3.
8. Once you have your fourth boiling point, unplug the hotplate and clean up your lab station.
*I’ll deduct 5 points from your lab score if any of your weights are exactly 10.00g.
Data and Calculations
Mass of water in kg:
Weighing
Boat
Weight of
NH4NO3 (g)
Boiling point of water:
Moles of
NH4NO3 (mol)
Soln. molality
(mol/kg)
Boiling point of
solution (C)
T
(C)
#1
#2
#3
1. Use Excel to make a scatter plot with T on the y-axis vs. molality on the x-axis. Be sure to
include the data point from the ‘pure’ water. Be sure to give your plot a title and to label the
axes.
2. Right-click on the data points and choose the “Add Trendline” option, then select a linear fit on
the top page and “Display Equation” and “R-squared value” on the Options page.
3. Print out your plot and staple it to this lab sheet.
Questions
1. If the equation for calculating the boiling point elevation of a solution is T=i K b m ( i is the
number of particles that each solute particle breaks up into, and m is the molality of the
solution), what is the boiling point elevation constant, Kb?
2. If i for another solute with the same molar mass was 4, what would your fourth boiling point have
been?
65
Freezing Point Depression of Water
Aka…Keeping It Cool with the Ice Cream Lab
DIRECTIONS
1. Dissolve 2 tablespoons of sugar in 1 cup of milk, then add no more than a half
teaspoonful (one capful) of vanilla or other extract for flavor.
2. Pour milk mixture into the small zip-close sandwich bag and close securely.
3. Half fill a 1 gallon zip-close freezer bag with ice cubes, then add about 1/2 cup
salt.
4. Place the sandwich bag into the freezer bag and securely close the freezer bag.
Gently massage the bags together to keep the inner bag mixing without heating it
with your hands (you do NOT want the inner bag to open).
5. Once the ice cream has frozen (or you’ve gotten fed up with waiting for it to
freeze), clean one corner of the inner bag off and then cut the corner. This helps
prevent salt from getting into your ice cream.
6. Squeeze your ice cream into a cup or a bowl and enjoy. 
QUESTIONS
1.
How does adding the salt help make the ice cream?
2.
Why must you continuously mix the ice cream as it freezes?
3.
List any problems you encountered as you made the ice cream.
4.
What would you change if you were to repeat this activity?
66
Determination of Kf for Half & Half
Colligative properties of solutions such as freezing point depression, boiling point elevation and osmotic
pressure depend only upon the number of solute particles that are present and are independent of the
chemical identity of the solute. The amount that the freezing point of a solvent will be depressed
(lowered), Tf, can be calculated using:
Tf = Kfm
where Kf is the freezing point depression constant for that particular solvent and m is the molality of the
solution. (Recall molality = moles of solute particles/kg solvent)
Objective: Determine the Kf for Half & Half (HH) using sugar as the solute
Procedure: (Make very careful measurements and record EVERYTHING! I want to publish this
eventually in the Journal of Chemical Education)
1. Set up an “ice” bath using isopropyl alcohol in a large styrofoam cup and pieces of dry ice.
2. Carefully weigh an empty, large test tube. Fill it halfway with HH and reweigh.
3. Clamp the test tube so that the HH is below the level of the alcohol, insert the mixing loop and
monitor the temperature with a digital thermometer while mixing continuously. (Will probably need
to record the temperature ~2 minutes…I’m expecting the cooling curve to be a bit squashy)
4. Once the HH has frozen, remove the test tube, extract the thermometer and start with another sample
(Could be preparing this one while the first one is freezing.) After weighing the test tube with the
HH, add a small amount of sugar and reweigh. If the weight difference is less than 0.2 g, add some
more sugar.
5. Repeat step 3. While the HH/sugar mixture is freezing, calculate the molality of the solution.
6. Repeat with another HH/sugar solution that has a different molality. Don’t add so much sugar that it
doesn’t all dissolve…that will skew the results.
Calculations
1. Calculate the molality of each solution.
2. For the neat HH and each HH/sugar solution, determine the freezing point. If the freezing point of
the HH is not evident from your temperature measurements, plot the Temperature vs. time data so
you can see where the plateau is.
3. Calculate Tf for each HH/sugar solution.
4. As a class, plot your Tf vs. molality. Do a linear regression analysis (best fit line) on the data and
determine the equation for the line.
5. The slope of the line will be Kf.
67
EQUILIBRIUM AND LE CHÂTELIER'S PRINCIPLE
Goals
 To observe the effect on equilibrium of adding or removing products and reactants.
 To predict the direction of shift in an equilibrium upon the change in concentration of one of
the components.
Background
Le Châtelier’s Principle states that:
If a stress is applied to a system at equilibrium, the system will respond by shifting in the
direction that reduces the stress and reach a new equilibrium condition.
In this experiment, the equilibrium between iron(III) thiocyanate, FeSCN+2, and it’s ions will be
investigated. This is a particularly easy equilibrium system to study because the dissociated ions (Fe +3
and SCN-1) are essentially colorless while the undissociated iron(III) thiocyanate is a dark red.
Fe3+ +SCN1– ⇋ FeSCN2+
Colorless ⇋ red
Lab Procedure
1. Place a piece of scrap paper under the ceramic spot plate so you can label four of the wells 1 — 4.
In each of these four wells, place 2 drops of 0.05 M NaSCN (sodium thiocyanate), 2 drops of 0.01
M Fe(NO3)3 solution, and 3 drops of water. Make sure you have taken the correct
concentrations of each solution. Mix each with a stirring rod; all of the solutions should appear
red.
2. Individually fill three of the wells at the bottom of the plate ¾ full with 0.10 M Fe(NO 3)3, 0.05 M
NaSCN and 1.0 M NaNO3 and take your plate back to your lab station.
3. Add 2 more drops of water to well #1. This well will serve as your color comparison for the following
experiments.
4. Add 2 drops of the 0.10 M Fe(NO3)3 to Well #2 and record your observations.
5. Add 2 drops of 0.05 M NaSCN to Well #3 and record your observations.
6. Add 1 drops of 1.0 M NaNO3 to Well #4 and record your observations.
7. After answering the questions below, rinse all of your lab equipment and leave your lab station
ready for the next class to use.
Question 1. When Fe(NO3)3 was added to the system,
a. Which ion in the equilibrium system caused the "stress"?
b. Which way did the equilibrium shift?
c. What happened to the concentration of SCN1–?
d. What happened to the concentration of FeSCN2+?
Question 2. When NaSCN was added to the system,
a. Which ion in the equilibrium system caused the "stress"?
b. Which way did the equilibrium shift?
c. What happened to the concentration of Fe3+?
d. What happened to the concentration of FeSCN2+?
68
Question 3. When you added NaNO3, did anything happen? Can you explain this result?
ADDITIONAL EQUILIBRIUM QUESTIONS
For each of the following, indicate the direction the equilibrium would shift AND what would happen to
the concentrations of each substance in equilibrium.
1. The following equilibrium may be established with carbon dioxide and steam.
CO (g) + H2O (g) ⇋ CO2 (g) + H2 (g) + heat
What would be the effect of each of the following on the equilibrium and concentrations?
a.) The addition of more H2O?
b.) The removal of some H2?
c.) Raising the temperature?
d.) Increasing the pressure?
e.) Addition of a catalyst?
2. What would be the effect of each of the following on the equilibrium involving the synthesis of
methanol?
CO (g) + 2 H2 (g) ⇋ CH3OH (g)
a.) The removal of CH3OH?
b.) An increase in pressure?
c.) Lowering the concentration of H2?
d.) The addition of a catalyst?
3. A small percentage of nitrogen gas and oxygen gas in the air combine at high temperatures found
in automobile engines to produce NO gas, which is an air pollutant.
N2 (g) + O2 (g) + heat ⇋ 2 NO (g)
a.) Higher engine temperatures are used to minimize carbon monoxide production. What effect does
higher engine temperatures have on the production of NO gas? Why?
b.) What effect would high pressures have on the production of NO gas? Why?
69
TITRATION LAB
PROCEDURE: Everyone in the group must do a titration!!! I will deduct 5pts from the lab grade
for identical titration results within a lab group.
1. Weigh clean, dry and empty 125 mL Erlenmeyer flask, then add ~25 mL vinegar (HAc) to the
flask. Reweigh and record the weight of just the HAc in the table below.
2. Calculate the volume of the vinegar used given that the density of the HAc solution is 1.01 g/mL.
3. Add 2 drops of phenolphthalein to the HAc in the flask, measure your initial pH and record below.
4. Carefully refill your buret and adjust the volume so that you are starting at 0.00 mL (the topmost
line).
5. Add approximately 10 mL of the sodium hydroxide solution to the flask and swirl until the pink
color disappears. Measure and record the pH of the solution (2nd pH)
6. Add approximately 5 mL of the sodium hydroxide solution to the flask and swirl until the pink
color disappears. Record the pH of this solution (3rd pH).
7. Slowly add the sodium hydroxide to the HAc solution, swirling continuously until a faint pink color
persists. Record the final volume of the buret and the pH of the solution (4th pH).
8. Add 1 mL of NaOH, and measure the pH again (5th pH).
9. Repeat step 10 and record (6th pH).
10. Calculate the molarity of the HAc (MHAc) using the equation: MNaOHVNaOH = MHAcVHAc
11. Calculate the mass% of HAc from the molarity.
12. Given that the mass% of HAc in household vinegar is 5.25%, calculate your percent error.
DATA:
1. What is the Molarity of the NaOH solution? (calc. from info on reagent bottle) __________
2. Complete the table below
Name of Experimenter 
Mass of vinegar used (g)
Calculated VHAc (mL)
Initial pH
2nd pH
3rd pH
4th pH
5th pH
6th pH
Final buret volume (mL)
VNaOH (mL)
MHAc (M)
Mass% HAc
%error
Group average mass%
70
Use the pH vs. volume NaOH data to sketch your titration curve on the graph below. Label the buffer
zone and the equivalence point.
Titration of Acetic Acid with NaOH
pH
14
13
12
11
10
9
8
7
6
5
4
3
2
1
0
0
10
20
30
Total Volume NaOH Added (mL)
40
50
QUESTIONS:
1. Write out the acid-base reaction that occurred when you added NaOH to the acetic acid,
HC2H3O2.
2. Looking at your titration curve, is acetic acid a strong acid or a weak acid? How can you tell?
3. If acetic acid is a weak acid, use your titration curve to estimate its Ka. (Remember that the pH
at ½ Veq is equal to the pKa of a weak acid.)
4. What are some possible sources of error (Remember, ‘we screwed up’ isn’t a valid source of
error.) in this experiment?
5. How did the group’s average mass% compare to the ‘true value’ (low or high)?
71
6. If all three values were higher than the ‘true value,’ what could cause this systematic error? Or if
all three were lower, what could cause that kind of systematic error?
Bonus Experiment: Determination of m/m% malic acid (H2C4H4O5) in Warhead candies.
1. Weigh a Warhead candy, place in a rinsed Erlenmeyer flask and add about 25 mL of water.
2. Stir with a stirring rod until the candy has completely dissolved, then add two drops of
phenolphthalein indicator.
3. Titrate as you did with the acetic acid, being careful to add the NaOH solution in smaller
increments as the persistence of the pink color increases.
4. Record your final volume of NaOH, and complete the calculations below.
5. If you think that different colors have different acidities (are more sour) AND you have time,
repeat for a different color of Warhead.
Calculations
Molarity of NaOH:
Volume NaOH: ___________________
Moles malic acid = (VNaOH)(MNaOH)/2 =
Grams malic acid: ________________
m/m% malic acid in Warhead candy:
Color of candy: __________________
Question: What did you think of this experiment? Did you learn more about titrations, acids and
bases by doing this experiment than you did by titrating the vinegar?
72
Identification of an Unknown Metal
Introduction:
Calorimetry is the measure of heat flow into or out of a system. The heat flow is measured in a device
called a calorimeter. An ideal calorimeter would insulate the substance in the calorimeter so well that
NO HEAT would be lost to the surroundings. The calorimeter you will use in this experiment (two
nested Styrofoam coffee cups) is far from an ideal one, but we will “assume” that no heat flows in or out
of the calorimeter.
When heat flows into or out of a substance the temperature of the substance usually changes. This
change can be used to monitor the flow of heat energy. In order to determine the exact amount of heat
(q) that flows, we need to know the temperature change in the substance (ΔT), the mass of the substance
(m), and the specific heat capacity (Cp) of the substance.
Specific heat capacity is defined as the amount of heat (in joules) required to raise the temperature of
one gram of the substance by 1C. The mathematical relationship relating the three quantities above is:
q = mCpΔT
In this experiment you will measure the specific heat capacity of several metals by placing the hot metal
in a weighed amount of cold water. Heat will flow from the metal to the water until they reach the same
final temperature. The amount of heat absorbed by the water, qwater, can be calculated using the equation
of
qwater = mwaterCwaterΔTwater
The amount of heat absorbed by the water is the amount of heat given off by the metal. The heat change
for the two processes is the same, but with opposite signs.
- qmetal = qwater
The amount of heat the metal absorbed can be used to calculate Cmetal using
qmetal = mmetalCmetalΔTmetal
(Note: The negative sign in the equation accounts for the fact that metal decreases in temperature and
the water increases in temperature; i.e., an exchange of heat.)
The value Cwater = 4.184 J/gC. After performing the experiment, the only unknown will be Cmetal. This
will be calculated from the other experimental data.
The densities of the metal samples and their molar masses, as calculated using the law of Dulong and
Petit (shown below), will also be determined to aid in identification of your sample of an unknown
metal.
Law of Dulong and Petit:
24.9 J mol-1 C-1 = Cmetal(Molar Mass)
NOTES: Goggles are necessary, and long hair must be tied back. Record all data to the correct decimal
place in your lab handbook, and submit your attached lab sheet upon completion of the lab. NEVER stir
anything with a thermometer, and never rest one on the bottom of the beaker.
PROCEDURE – Density
To determine the density of your metal sample, you need to know the mass (in grams), and the volume
(in mL = cc = cm3). Determine the volume as precisely as possible by measuring water displacement.
73
PROCEDURE - Specific Heat
To determine the initial (high) temperature of your metal sample, suspend it in a beaker of boiling water
and keep it there until boiling has proceeded steadily for about two minutes. Record the temperature of
the boiling water with a thermometer making sure it does not touch the bottom of the beaker. This is the
same as the initial metal temperature.
While the water/metal mixture is boiling, record the mass of the empty calorimeter (2 nested coffee
cups). Next, add just enough water to cover your piece of metal (estimate), and determine the combined
mass.
Record the temperature of the cold water and of the metal (the boiling water) just before combining
them. Carefully immerse the metal sample in the cold water so that none splashes out, and record the
final temperature of the mixture. It should change quickly at first, then level off, then cool back down
slowly. Record the level part.
NOTE: Measure the density and specific heat at least twice (three times if there is a large difference
between the two trials). Make sure everything is dry for each trial.
DATA ANALYSIS
You are to:
 Determine the density of both metals
 Determine the specific heat of both metals
 Calculate the molar mass of both metals
 Calculate the %error in your measured values for the density, specific heat and molar mass of
copper
Remembering that the “law” of Dulong and Petit is only accurate to within 10% or so (high or low),
locate a region of the periodic table that likely contains your unknown element. Also recall that many
elements are not reasonable laboratory samples (i.e., sodium and magnesium are too reactive in water,
neon is a gas, and silver is too expensive to use). With this information, identify several metals that your
unknown sample may be. Consult your textbook, the Handbook of Chemistry and Physics, or any other
reference available for tables of the physical property you measured. Give your best estimate of the
unknown's identity, and then calculate the %error in your measurements assuming that you’ve chosen
the correct metal. (Ideally, the %error values should be comparable to what you found with the copper
sample.)
74
Data: Identification of an Unknown Metal
COPPER
Mass (g)
Trial 1
Trial 2
Trial 3
Trial 1
Trial 2
Trial 3
Volume (mL)
Density (g/mL)
%error Density
Wt cold water (g)
Twater (C)
TCu (C)
Cp(Cu) (J/gC)
%error Cp(Cu)
Calc’d molar mass
%error molar mass
UNKNOWN METAL #_____
Mass (g)
Volume (mL)
Density (g/mL)
Wt cold water (g)
Twater (C)
TCu (C)
Cp(unknown) (J/gC)
Calc’d molar mass
Proposed identity of unknown: __________________
Literature values for density = __________________, specific heat = __________________, and molar
mass = __________________.
%error in density = __________________, specific heat = __________________, and molar mass =
__________________.
75
HEAT OF FORMATION FOR MAGNESIUM OXIDE USING HESS’ LAW
Objective: To give the student additional experience in the application of the heat additivity law and to
illustrate how the heat term can be obtained for a reaction that is inconvenient to perform. Students will
experimentally determine the heat of formation of magnesium oxide by doing two temperature change
determinations, calculating the amount of heat involved in these changes and applying Hess' Law to find
the ∆H of a reaction which is the sum of the others.
Introduction: When a reaction can be expressed as the algebraic sum of a sequence of two or more
other reactions, the heat of reaction is the algebraic sum of the heats of these other reactions. This
generalization has been found to be true for every reaction that has been tested and is known as Hess’
Law.
In this experiment you will use Hess’ Law to determine the heat for a reaction that is difficult to measure
directly. Magnesium metal burns rapidly, releasing light and heat, as you have observed in photo
flashbulbs or in burning magnesium ribbon. The reaction is represented by the equation:
Mg + ½O2  MgO
(1)
This equation can be obtained by combining equations 2, 3, and 4.
MgO + 2 HCl  MgCl2 + H2O
(2)
Mg + 2 HCl  MgCl2 + H2
(3)
H + ½ O2  H2O
(4)
Heats of reaction for equations 2 and 3 will be experimentally determined using a coffee cup
calorimeter. The heat for reaction 4 can be obtained from a table of values for previously measured
reactions in your text or other reference source such as the CRC.
Materials: Safety glasses, Styrofoam calorimeter, Triple-beam balance, 100-mL graduated cylinder,
Magnesium oxide (powdered), 0.5 g Magnesium ribbon, 200 mL of 1 M Hydrochloric acid solution
Safety: Hydrochloric acid
1. Reactivity: Contact with metals produces hydrogen gas which may form explosive mixtures with
air. Keep away from strong alkaline solutions. Maintain adequate ventilation. Neutralize with
chemically basic substances such as soda or slaked lime. Use rubber gloves and aprons. Use
protective eye-wear.
2. Health: Causes severe burns! May be FATAL if swallowed! Do not get in eyes, on skin, or clothing.
Immediately flush skin or eyes with plenty of water for at least 15 minutes.
3. When diluting, add acid to water. Keep out of the reach of children.
Procedure--Reaction 2:
1. Weigh out about 1.00 gram MgO to the nearest 0.01 gram and record its mass = __________
2. Measure out 100 mL 1.0 M HCl solution into your styrofoam calorimeter.
3. Use thermometer to measure initial temperature of solution. Be sure it has stabilized before you
start. Ti = __________
4. Add the MgO to the acid solution quickly while stirring.
5. Measure and record the final (maximum) temperature reached as the MgO reacts with the acid
solution. Tf = __________
6. Dispose of the solution as directed by your instructor.
76
Procedure--Reaction 3:
7. Weigh out the strip of magnesium ribbon provided by your instructor to the nearest 0.01 gram and
record its mass = __________
8. Measure out a fresh 100 mL 1.0 M HCl solution into your styrofoam calorimeter.
9. Measure and record the temperature of the acid solution. Ti = __________
10. Add the magnesium ribbon to the acid solution while stirring.
11. Measure and record the final (maximum) temperature reached as the Mg reacts with the acid
solution. Tf = __________
12. Dispose of the solution as directed by your instructor.
Procedure--Reaction 4:
13. Obtain the heat of formation for liquid water from the appropriate table in your text or reference
source. Hf(H2O(l)) = __________
Data Collection: In your lab notebook, create a data table to contain the data from this experiment. Be
sure to include space on the table for your tabulated values of ∆H.
Data Analysis: Assume the density of HCl(aq) is 1.0 g/mL and specific heat is 4.18 J/g K
1. Calculate the ∆H for reaction 2.
2. Calculate the ∆H for reaction 3. Be sure to take the calorimeter constant into account.
3. Use those answers, together with the literature value you researched to obtain the heat of formation
for MgO(s). Show your work in its entirety.
4. Find the value for the heat of formation of MgO(s) in your textbook or CRC.
5. Calculate the percent error of your value compared to the literature value.
77
Heat of Reaction Lab
Please read all instructions BEFORE conducting these experiments. It is very important to get all of the
steps completed in the right order AND in a timely fashion.
This lab consists of TWO separate experiments: (1) Determining the heat of neutralization of HCl and
NaOH, and (2) Determining the Hrxn between HCl and NaHCO3.
Materials
Styrofoam coffee cup
400 mL beaker
100 mL graduated cylinder
Thermometer
Stirring rod
NaHCO3
1.0 M HCl
1.0 M NaOH
Part A: Heat of Neutralization Procedure
1. Make sure your coffee cup is dry and then obtain its weight. (Q: To how many decimal places
should you be weighing???)
2. Place the weighed coffee cup into a 400 mL beaker for stability.
3. Measure out 50 mL of 1.0 M HCl in the 100 mL graduated cylinder and add the acid to the
weighed coffee cup.
4. Rinse the acid out of the graduated cylinder and then measure out 50 mL of 1.0 M NaOH.
5. Record the initial temperature of the HCl in the coffee cup. (Q: To how many decimal places
should you be measuring the temperature???)
6. Quickly, but without splashing, add the 50 mL of NaOH in to the acid in the coffee cup.
7. Mix gently with the stirring rod and monitor the temperature with the thermometer. DO NOT
STIR WITH THE THERMOMETER!!
8. Record the MAXIMUM temperature of the solution in the coffee cup.
9. Weigh the coffee cup and its contents again. If you do not get this final weight, you will have to
start over again.
10. Rinse out the coffee cup and redry it before proceeding to step 11.
Part B: Heat of Reaction Procedure
11. Place the weighed coffee cup from step 1 back into the 400 mL beaker for stability.
12. Weigh out ~4 g of NaHCO3, record the actual mass of the NaHCO3, then transfer the solid
NaHCO3 to the coffee cup.
13. Measure out 50 mL of 1.0 M HCl in the 100 mL graduated cylinder and record its initial
temperature.
14. Carefully add the acid to the solid NaHCO3 in the weighed coffee cup. (Don’t add it so fast that
it bubbles over).
15. Mix gently with the stirring rod and monitor the temperature with the thermometer.
16. Record the LOWEST temperature of the solution in the coffee cup.
17. Weigh the coffee cup and its contents. If you do not get this final weight, you will have to start
over again.
78
Questions (Answer on a separate piece of paper please)
Part A: Heat of Neutralization
1. Write the chemical reaction equation for the neutralization of HCl with NaOH.
2. How many moles of HCl and NaOH did you use in the reaction? (remember that Molarity = #
moles solute/ L of solution)
3. What equation do you need to use to calculate the total heat added to the water by the
neutralization reaction?
4.
5.
6.
7.
8.
What was your T? (T = Tf – Ti, Remember: signs are very important when dealing with heat)
What was the total mass of the solution in the coffee cup?
How much heat was absorbed by the water?
Use the appropriate sign here
How much heat was given off by the reaction? and include units
What is the molar heat of neutralization for the reaction between HCl and NaOH? (Hint: the
units are kJ/mol)
Part B: Heat of Reaction
9. Write the chemical reaction equation for the reaction of HCl with NaHCO3. (Hint: CO2 (g) is
evolved)
10. How many moles of HCl and NaHCO3 did you use in the reaction? Which one is the limiting
reactant?
11. What was your T?
12. What was the total mass of the water in the coffee cup? (subtract out the weight of the NaHCO3)
13. How much heat was absorbed by the water?
Use the appropriate sign here
and include units
14. How much heat was given off by the reaction?
15. What is the molar heat of reaction for the reaction between HCl and NaHCO3? Be sure to use the
molar amount of the limiting reactant when doing this calculation.
79
IODINE CLOCK LAB:
Data Analysis
In the Iodine Clock reaction, the kinetics of two competing reactions determine the overall rate of the
reaction.
2 I-1 (aq) + S2O8-2 (aq) → I2 (aq) + 2 SO4-2 (aq)
slow step
I2 (aq) + 2 S2O3-2 (aq) → 2 I-1 (aq) + S4O6-2 (aq)
fast step
As soon as there is excess I2 present in the solution, it reacts with a starch indicator and turns the
solution a blue-black color. Because the general rate law for this reaction is:
rate = k[S2O8-2]m [I-1]n
it is possible to determine the reaction order of the two reactants (m and n) and the specific rate constant
(k) by varying the concentrations of the iodide (I-1) and the persulfate (S2O8-2) ions. Once the specific
rate constant at several temperatures is known, the energy of activation (Ea) for the reaction can be
determined by plotting:
ln(k) vs 1/Temperature
where k is the specific rate constant and the temperature is in Kelvin. The slope of this line will be equal
to –Ea/R. (Note: ln(k) is the y-variable and 1/T is the x-variable)
Objective:
 Use the data provided to determine the reaction order and specific rate constant at five different
temperatures.
 Generate an ln(k) vs. 1/T plot to determine Ea for this reaction.
Data: Fill in the experimental data that was given to you, and then calculate the remaining values. NO
CREDIT will be given for this lab if the DATA SET# is not included on the line below.
Initial [S2O8-2] = ________ M
Data set #: ________
Temp
(K)
VS2O8
(mL)
VI(mL)
VH2O
(mL)
Time
(s)
[S2O8-2]
(M)
Initial [I-1] = ________ M
[I-1]
(M)
Rate
(M/s)
k
298
298
298
318
318
318
80
Temp
(K)
VS2O8
(mL)
VI(mL)
VH2O
(mL)
Time
(s)
[S2O8-2]
(M)
[I-1]
(M)
Rate
(M/s)
k
338
338
338
358
358
358
378
378
378
Directions:
-2
-1
-2
1. Calculate [S2O8 ] and [I ] for all reactions. Example: [S2O8 ] =
Initial [S O ]V
2
2
8
S2O8

total volume
[S O 2 ]
2. Calculate the rate of each reaction: rate = 2 8
time
3. Determine the values of m and n (will be the same for all reactions). m = _____
n = _____
4. Use the rate law equation and your values of m and n to solve for the specific rate constant, k, for all
reactions (will be different for each temperature).
5. Using Excel, plot ln(k) on the y-axis and 1/Temperature on the x-axis. Be sure to give your graph a
title and to label the axes.
6. Use the Trendline option to get a linear, best-fit line of your data. Be sure to select “Display equation
and R-squared value on graph” option.
7. Print out your graph and attach it to your lab sheet.
8. The slope of your best-fit line is
 Ea
J
where R = 8.314
. Use this information to calculate the
R
mol  K
activation energy for the Iodine Clock reaction. Note that the units of R are in Joules while Ea is in
kJ.
Ea = __________________ kJ/mol
9. Given that Hrxn = -370 kJ/mol for the Iodine Clock reaction, sketch an appropriately scaled reaction
energy diagram at a temperature =298 K. Include arrows to show Hrxn and Ea, and be sure to write
in the formulas for the reactants and products on the appropriate line.
81
[S2O8-2] = 0.20 M
[I -1] = 0.20 M
Temp (K)
298
298
298
VS2O8
(mL)
4
4
8
VI(mL)
8
4
4
VH2O
(mL)
8
12
8
1
time
(s)
39
78
78
318
318
318
4
4
8
8
4
4
8
12
8
338
338
338
4
4
8
8
4
4
358
358
358
4
4
8
378
378
378
4
4
8
[S2O8-2] = 0.20 M
[I -1] = 0.20 M
Temp (K)
298
298
298
VS2O8
(mL)
2
2
5
VI(mL)
5
2
2
VH2O
(mL)
8
11
8
2
time
(s)
52
104
104
36
73
73
318
318
318
2
2
5
5
2
2
8
11
8
48
97
97
8
12
8
34
68
68
338
338
338
2
2
5
5
2
2
8
11
8
45
91
91
8
4
4
8
12
8
32
64
64
358
358
358
2
2
5
5
2
2
8
11
8
43
86
86
8
4
4
8
12
8
31
61
61
378
378
378
2
2
5
5
2
2
8
11
8
41
82
82
Temp (K)
298
298
298
VS2O8
(mL)
3
3
6
VI(mL)
3
2
2
VH2O
(mL)
14
16
13
4
time
(s)
104
208
208
[S2O8-2]] = 0.10 M
[I -1] = 0.20 M
[S2O8-2] = 0.10 M
[I -1] = 0.20 M
Temp (K)
298
298
298
VS2O8
(mL)
3
3
6
VI(mL)
3
2
2
VH2O
(mL)
9
11
8
3
time
(s)
78
156
156
318
318
318
3
3
6
3
2
2
9
11
8
73
145
145
318
318
318
3
3
6
3
2
2
14
16
13
97
194
194
338
338
3
3
3
2
9
11
68
136
338
338
3
3
3
2
14
16
91
182
338
6
2
8
136
338
6
2
13
182
358
358
358
3
3
6
3
2
2
9
11
8
64
129
129
358
358
358
3
3
6
3
2
2
14
16
13
86
172
172
378
3
3
9
61
378
3
3
14
82
378
378
3
6
2
2
11
8
123
123
378
378
3
6
2
2
16
13
163
163
82
[S2O8-2] = 0.20 M
[I -1] = 0.10 M
Temp (K)
298
298
298
VS2O8
(mL)
5
5
10
VI(mL)
20
10
10
VH2O
(mL)
0
10
5
5
time
(s)
39
78
78
318
318
5
5
20
10
0
10
318
10
10
338
338
338
5
5
10
358
358
[S2O8-2] = 0.20 M
[I -1] = 0.20 M
Temp (K)
298
298
298
VS2O8
(mL)
6
6
13
VI(mL)
13
6
6
VH2O
(mL)
6
13
6
6
time
(s)
31
62
62
36
73
318
318
6
6
13
6
6
13
29
58
5
73
318
13
6
6
58
20
10
10
0
10
5
34
68
68
338
338
338
6
6
13
13
6
6
6
13
6
27
55
55
5
5
20
10
0
10
32
64
358
358
6
6
13
6
6
13
26
52
358
10
10
5
64
358
13
6
6
52
378
378
378
5
5
10
20
10
10
0
10
5
31
61
61
378
378
378
6
6
13
13
6
6
6
13
6
25
49
49
Temp (K)
298
298
298
VS2O8
(mL)
7
7
14
VI(mL)
5
2
2
VH2O
(mL)
23
25
18
8
time
(s)
78
156
156
[S2O8-2] = 0.25 M
[I -1] = 0.20 M
[S2O8-2] = 0.10 M
[I -1] = 0.30 M
Temp (K)
298
298
298
VS2O8
(mL)
2
2
4
VI(mL)
5
2
2
VH2O
(mL)
9
11
9
7
time
(s)
52
104
104
318
2
5
9
48
318
7
5
23
73
318
318
2
4
2
2
11
9
97
97
318
318
7
14
2
2
25
18
145
145
338
338
338
2
2
4
5
2
2
9
11
9
45
91
91
338
338
338
7
7
14
5
2
2
23
25
18
68
136
136
358
2
5
9
43
358
7
5
23
64
358
358
2
4
2
2
11
9
86
86
358
358
7
14
2
2
25
18
129
129
378
378
378
2
2
4
5
2
2
9
11
9
41
82
82
378
378
378
7
7
14
5
2
2
23
25
18
61
123
123
83