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Date:
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Spring Grove Area High School
Chemistry I
Chemistry I Honors
Laboratory Manual
Miss Frey
Mr. Henning
Mrs. Kimber
Student Name:___________________________________________
Homeroom:______________Chemistry Period:________________
Chemistry Teacher:________________________________________
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Date:
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SPRING GROVE AREA SCHOOL DISTRICT
SCIENCE DEPARTMENT SAFETY
From: The High School Science Department
To: All Science Students
The laboratory is a safe place to experiment if you are careful and safety conscious. You must
assume responsibility for the safety of yourself and your classmates. The following safety rules
will guide you in protecting yourself and others from injury in the lab. In addition your instructor
may have a list of additional rules which must be followed in the laboratory while under his/her
supervision. When necessary, your instructor will also give you specific instructions for a
particular laboratory exercise.
1. The laboratory is to be used for serious work. No inappropriate behavior of any type is permitted
at any time.
2. Do not handle the equipment in the laboratory unless you have been authorized to, or you are
performing an experiment, which makes use of that equipment.
3. Only experiments, which your instructor authorizes, are permitted. Always obtain your instructor’s
permission if you would like to modify an experiment or perform an unassigned experiment.
4. Study your laboratory assignment prior to coming to class. If you are in doubt about any
procedure, ask your lab instructor.
5. Use safety equipment provided for you. Know the location of available safety equipment, such as
fire extinguishers, showers, etc.
6. Handle and use all instruments and equipment properly and with care so as not to injure yourself,
others, or damage the equipment.
7. Report any accident, injury, or incorrect procedure to you instructor at once, when/if it occurs.
8. Handle toxic, combustible, or radioactive materials only under the direction of your instructor.
9. Never take any chemical substance or draw poisonous materials into a glass tube with your
mouth. At no time should you have any foreign material in your mouth during the lab.
10. If you spill toxic, combustible, or radioactive materials notify your instructor at once. He/she will
give you directions for cleaning up the spill.
11. Use gas, electrical, and water outlets properly.
12. Use electrical equipment only under the supervision of your instructor.
13. Students should be aware of loose clothing, long hair or other possible fire catches and
appropriate corrections should be made.
14. Keep your work area orderly and clean at all times. When your investigation is complete, return
all materials and apparatuses to their proper locations.
15. When your experiment is complete, clean up and dispose of materials properly.
16. Be sure to wash your hands after handling any toxic materials in the lab.
WE HAVE READ, DISCUSSED, AND UNDERSTAND THE SAFETY REGULATIONS
FOR THE SCIENCE CLASSROOM.
Student’s Name (Printed):
_________________________________________
Date:___________________
Parent’s/Guardian’s Signature:
_________________________________________
Date:__________________
Student’s Signature:
_________________________________________
Date:___________________
2
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3
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4
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5
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6
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EVIDENCE OF CHEMICAL CHANGE
Objectives:
1. Observe the types of evidence that indicate that a chemical change has taken place.
2. Infer from observation of a chemical change that a new substance has been formed.
3. Identify and record data that shows how heat is involved in a chemical change.
4. Explain why a substance can be either a reactant or a product in a chemical reaction
and how this relates to recycling.
Introduction:
Substances may be created by a chemical change and may also undergo chemical
change. If a substance is produced as a result of a chemical change or a reaction, it is called
the product. If the substance is subject to a chemical change, it is the reactant.
One way to indicate if a chemical change has occurred is to observe the properties of
both the reactants and the products, since they will be different. In this experiment you will
observe a series of reactions beginning with copper (II) nitrate. All of the reactions will take
place in the same test tube, with the reactants becoming products and those products
becoming the reactants for the next reaction.
This chemical conversion is useful when we recycle materials. For instance, the
copper (II) nitrate we are using in this experiment was created from a reaction between
elemental copper and nitric acid. Through a series of reactions, we can recover the copper.
Another example is with metals that are changed into alloys and shaped into soft-drink cans,
and by recycling the cans the aluminum can be recovered.
YOU MUST WEAR GOGGLES AND APRONS FOR THIS LABORATORY EXERCISE!
Equipment:
Materials:
1 large test tube
Test tube rack
1 – 100mL beaker
10mL Graduated cylinder
Beaker tongs
Bunsen burner
Stirring rod
Ring stand
Iron ring
Wire gauze
Forceps
1.0 M copper (II) nitrate, Cu(NO3)2
3.0 M sodium hydroxide, NaOH
3.0 M hydrochloric acid, HCl
8 cm piece of aluminum (Al) foil, rolled
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Procedure:
1. Place 100 mL of water into the 150 mL beaker and place on the wire gauze.
2. Heat the water to a boil to use in step 5.
3. Measure out 2mL of copper (II) nitrate and pour into the large test tube.
4. Measure out 2mL of sodium hydroxide and pour into the large test tube containing the
copper (II) nitrate.
5. Mix the two solutions slowly with the stirring rod. You have now created two new
products called copper (II) hydroxide and sodium nitrate.
a. Record any observations.
b. Touch the outside of the test tube. Is it warm or cold?
6. Remove the stirring rod and place the test tube in the water bath from step 1 & 2. Make
sure the test tube is facing the wall and not out towards you or any partners.
7. Heat until no more changes occur. The new products are copper (II) oxide and water.
a. Record any observations.
8. Remove the test tube from the water bath. Turn off the Bunsen burner and gas. Cool the
test tube for two minutes in the test tube rack.
9. Place the test tube in a second beaker that contains cool or room temp water. Allow to
cool an additional two minutes.
10. Remove the boiling water from the wire gauze and discard the water.
11. Once cool, remove the test tube from the water bath and place back in the test tube rack.
12. Measure out 4 mL of hydrochloric acid and pour into the large test tube containing the
copper (II) oxide and water.
13. Mix the solutions slowly with the stirring rod. If the solution does not turn clear, add 2
more mL of hydrochloric acid and continue stirring until it clears. You have now created
two new products called copper (II) chloride and water.
a. Record any observations.
14. Place an 8 cm piece of aluminum foil into the test tube and leave it until no more reaction
takes place. This may take a few minutes to occur. Two reactions are taking place: one
will make copper and aluminum chloride, and the other will make hydrogen and
aluminum chloride.
a. Record any observations.
b. Touch the outside of the test tube. Is it warm or cold?
15. Use the forceps to remove the aluminum foil/copper from the test tube. Compare it to a
sample of copper you receive from your teacher.
a. Record any observations.
16. Dispose of the aluminum foil/copper in the trash can and place the liquid contents that
remain in the labeled beaker in the fume hood.
17. Clean up your lab station, putting away all materials where they belong. Clean the test
tube with the soap solution before putting it away.
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Observations & Data:
5a. ________________________________________________________________________
5b. ________________________________________________________________________
7a. ________________________________________________________________________
13a. ________________________________________________________________________
14a. ________________________________________________________________________
14b. ________________________________________________________________________
15a. ________________________________________________________________________
Reactants
Copper (II) nitrate
+
Data Table
Observations of
Precipitate Formed
Products
Observations of
Supernate Formed
+
Sodium hydroxide
Copper (II)
hydroxide
+
+
Heat
Copper (II) oxide
+
+
Hydrochloric acid
Aluminum
+
+
Copper (II) chloride
Aluminum
+
+
Hydrochloric acid
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Post Lab Questions:
1. What are some causes of chemical changes?
2. In what two ways is heat involved in a chemical change? Cite two specific instances from
this experiment.
3. What is the color of the copper solutions in this experiment?
4. Which substances used or formed in the experiment are elements?
5. Which substances used or formed in the experiment are compounds?
6. List four types of observations that indicate that a chemical change has occurred.
7. Describe the advantages and disadvantages of recycling metals as was done in this
experiment.
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Quantitative Analysis of an Ocean Water Sample
Objective:
To design and carry out the separation of a mixture.
Background: After all of that oil spilled in the Gulf of Mexico, it’s finally hit the fan. Aquatic Life
is altered: fish and shellfish are dying. The shrimp population in particular is vastly diminished,
and BP is denying any responsibility for the population decrease. NOAA, the National Oceanic
and Atmospheric Administration has contracted your company to conduct a quantitative analysis
of the ocean water in the affected area. NOAA requires that you report the mass % of sodium
chloride, sand, iron, and oil (in the form of tar balls) in your water sample. NOAA officials will
then use your findings to decide if they have enough evidence to prosecute BP officials.
Since you must conduct a quantitative analysis, you must separate the ocean water into its
components, which include water, sodium chloride, iron, sand, and tar balls. You must then
obtain the mass of each individual substance, and calculate the mass % of each in the mixture.
Procedure:
1. Observe your ocean water sample. Using your observations, and your knowledge of each
of these substances, make a list of the physical properties you could use to help you separate
the components of this mixture. Then describe how you could separate these substances from
the mixture.
Substance
Physical Properties
Possible Separation Method
Water
Sodium
Chloride
Sand
Iron
Tar Balls
2. Next, make a flow chart of your separation process, and have your instructor initial it before
you begin the physical separation.
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3. Make a list of materials and lab equipment you will need to carry out the separation. If there
is anything on your list that is not in your lab station, see your instructor.
4. Write the procedure you will carry out to separate the components. Use the correct names
for materials and equipment. Use literate sentences and correct spelling and grammar.
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5. Carry out your procedure and complete the data table below.
substance
Mass (g)
sample
% of sample
------
Oil (tar balls)
salt
sand
Iron
Water
Results:
Ocean Water Sample Number: _________________
% NaCl
__________________________________
% Sand __________________________________
% Iron ____________________________________
% Tar Balls ________________________________
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Scientific Method
Purpose: To use the observations of chemical reactions to understand the stops of the
scientific method. 1 – State the problem, 2 – State a hypothesis, 3 – Experiment, 4 – Collect
data, 5 – Organize data, 6 – State a conclusion.
Procedure:
1) Using pipets, that are labeled to indicate the solution, place 4 drops of the solution
indicated into one of the wells on the microtiter plate.
a. USE ONE DROP OF SOLTUION # 5.
2) To the same well you just used, add 4 drops of the second solution indicated. Carefully
stir with a stirring rod to complete the reaction.
3) Record all observations!
4) If a third solution is listed, add 4 - 8 drops of that solution to the same well, to ensure a
permanent change is evident after stirring.
5) Record all observations!
6) Use a new well, for each new mixing.
7) Clean microtiter plate thoroughly when finished.
A) SAFETY GOGGLES AND APRONS MUST BE WORN!!
B) ALL DATA IS PERTINENT TO THE SUCCESS OF THE LAB ~ BE SPECIFIC WITH ALL
RECORDED OBSERVATIONS!
Solutions:
Names of compounds and chemical formulas are given
Barium chloride
Sulfuric acid
Sodium carbonate
Sodium hydroxide
Phenolphthalein
Sodium sulfate
Hydrochloric acid
Potassium carbonate
Potassium hydroxide
Ferric chloride
Ammonium hydroxide
Ammonium sulfate
Copper (II) sulfate
Ammonium carbonate
BaCl2
H2SO4
Na2CO3
NaOH
Na2SO4
HCl
K2CO3
KOH
FeCl3
NH4OH
(NH4)2SO4
CuSO4
(NH4)2CO3
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Solution Mixing Chart:
MIX SOLUTIONS
OBSERVATIONS
MIX SOLUTIONS
1+2
10 + 4
2+3
10 + 11
4+5
1 + 12
4+5+2
1+3+7
1+6
5 + 11
7+8
10 + 9
5+9
2+ 8
5+9+7
5 + 11 + 2
13 + 11
13 + 4
1 + 14 + 7
8+1+7
OBSERVATIONS
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Organize Data:


Study the observations made until you see combinations of solutions that have the
same observations. Groups of three solutions should not be compared with groups
of two solution mixes.
Record the observations that are the same and the solutions (by their number) that
caused the same observations. Six are repeated, and two are not. (Hint- the 2 that
are not are blue…)
Repeated Observations
Solution Combinations
1
2
3
4
5
6
Not Repeated Observations
1
2
Making Hypotheses:


Similar observations indicate the same or a similar reaction. In this experiment you
may assume that three similar observations represent a generalization. For
example: (4 + 5), (5 + 9), and (5 + 11) all turned pink. The generalization would be
that phenolphthalein turns pink in the presence of hydroxides. (The -OH group or
the hydroxide, are what solutions 4, 9 and 11 all have in common).
Using the chemical names and your organized data discover 5 OTHER
generalizations.
1
2
3
4
5
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Predict & Test:



From the five generalizations that you have made or discovered, write 2 procedures
to show examples of two different generalizations. At least one of the solutions in
your procedure can not be used once you have used it before. (There are different
solutions on the side shelf). Use the solutions to test your procedures recording
your observations and results.
If your results agree with your rules they are considered to be positive; if the
disagree they are negative.
Indicate positive or negative for your conclusion.
Procedure 1
Observations
1
Conclusion:
Procedure 2
Observations
2
Conclusion:
Questions:
1) To identify a sulfate solution the first step is to add barium ions from a barium
solution. What will happen once the barium ions are added?
2) A carbonate solution will react the same way when barium ions are added to it.
(Check you data to confirm this). What kind of solution could be added after the
barium ions to distinguish between a sulfate and a carbonate solution?
3) What would you observe if you added this solution (answer to Q2) to the carbonate
solution?
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Introduction to Graphing in Chemistry
There are various types of graphs that can be made. However, it depends on the
information you are working with, what type of graph best suits the data. If you are
comparing two variables to each other a line graph works best. If you are comparing
unrelated things, a bar graph would be well suited. A pie graph works best when you are
comparing parts of a whole.
When making a graph it is critical that the graph accurately depicts the data. This means
you have to place the correct variable on the correct axis. The x-axis is for the independent
variable or the variable that you can manipulate. The y-axis is for the dependent variable,
or the variable whose outcome is dependent upon the other variable.
The axes should always be labeled with the variables and include units when appropriate.
Additionally, it is important that the increments on the axes are even. For example, you
would want to use numbers such as 2,4,6,8,10 rather than 3,5,9 because the latter are not
in even increments. The graph should fill the page provided. If it doesn’t, rescale the axes
so that the graph is larger and will be a better predictive tool for analyzing the data.
Graphs should always be shown as smooth lines or curves – NEVER play connect the
dots. We want to show a smooth progression not a sawtooth pattern. For this reason, we
tend to use lines of best fit, or curved lines which show a general trend in the data and can
allow for us to analyze and interpret our data.
Titles should be written in statement form demonstrating what the graph is about. Avoid
the word ‘versus’ in your statement, and write it in a sentence where the independent
variable is stated before the dependent variable. For example, ‘The effect of [independent
variable] on [dependent variable]’ would work where the variables you are working with are
inserted into the brackets.
Lastly, USE A RULER! Sloppiness is acceptable.
Equipment:
1) Broken Pieces of Spaghetti at random lengths
2) Electronic or Triple Beam Balance
3) Graph Paper
Procedure:
1) Make a data table to record the lengths and masses of small pieces of spaghetti.
2) Measure the lengths of the pieces of spaghetti at your lab table.
3) Record your data that you obtain in your data table.
4) Measure the masses of the pieces of spaghetti at your lab table.
5) Record your data that you obtain in your data table.
6) Use the data you collected and construct a graph that best suits the data.
Remember to include all the pertinent information you would find on a graph.
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Density Is a Periodic Property
Introduction:
Dmitri Mendeleev proposed the periodic law for the classification of elements in 1869-1871.
After observing trends in the properties of elements when they were arranged in order of
increasing atomic mass, Mendeleev made a startling prediction. He predicted the existence
and properties of at least three undiscovered elements. Mendeleev saw what other
scientists before him had missed-he saw what wasn’t there!
Chemical Concepts:
•Periodic law
•Density
• Group 14 elements
• Period number
Background:
At the time Mendeleev proposed the periodic law, the foundation of the modern periodic
table for the classification of elements, 63 elements were known. Their physical and
chemical properties had been studied and their atomic masses measured. Mendeleev
arranged the known elements in a calendar-like table of rows and columns in order of
increasing atomic mass and repeating chemical properties. It is at this point; however, that
Mendeleev made a giant leap of discovery-he suggested that there were some gaps or
missing elements in the list of known elements.
Among the Group 14 elements in Mendeleev’s classification scheme, carbon appeared in
the second row, followed by silicon in the third row. Both tin and lead shared similar
chemical properties with carbon and silicon and were also known at this time. Because of
their high atomic masses, however, these metals were placed in later rows of Mendeleev’s
Group 14 column of elements. In 1871, Mendeleev proposed that there existed an as-yetunknown element beneath silicon in the Group 14 elements. He named the missing
element ekasilicon and predicted its physical properties (atomic mass, melting point,
density, and specific heat). In 1886 the element germanium was discovered by the German
chemist Clemens Winkler. In his report of the discovery, Winkler stated: “. . . There can be
no longer any doubt that the new element is no other than the ekasilicon prognosticated
fifteen years ago by Mendeleev.”
Within 15 years of Mendeleev’s prediction of the existence of missing elements, three of
the elements had been discovered, their properties in excellent agreement with those
predicted by Mendeleev. Is it possible to recreate some of the excitement that followed the
prediction and discovery of Mendeleev’s missing elements?
Experiment Overview:
The purpose of this experiment is to measure mass and volume data for silicon, tin, and
lead, calculate their densities, and use these results to predict the density of germanium,
Mendeleev’s “undiscovered” element in the Group 4 family of elements. The volume of the
elements will be measured by water displacement (see Figure 1).
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Name:
Pre-Lab Questions:
Date:
Period:
1. One of the elements Mendeleev predicted was ekaaluminum, corresponding to a gap in
the fourth row or period of the Group 13 elements, between aluminum and indium. The
density of aluminum (period 3) is 2.70 g/cm3, that of indium (period 5) 7.31 g/cm3, and that
of thallium (period 6) 11.85 g/cm3 Make a graph of period number on the x-axis versus
density on the y axis for each of these elements.
2. Use your graph to predict the density of ekaaluminum. What known element in the
modern Periodic Table corresponds to ekaaluminum? Look up the density of the modern
element in a reference source and record its actual and predicted density values.
3. How do the actual and predicted density values compare? Calculate the percent error
between the predicted and actual values for the density of ekaaluminum.
Materials:
Lead shot, Pb, 35 g
Paper towels
Silicon lumps, Si, 8 g
Water
Tin shot, Sn, 25 g
Balance, centigram (0.01 g precision)
Beakers, 50-mL, 3
Graduated cylinder, 25 mL
Forceps or tongs
Marking pencil or pen
Tape
Safety Precautions:
Lead powder is extremely toxic by inhalation and ingestion, lead fumes and dust are
possible carcinogens. Using lead shot does not present a powder or dust hazard. Do not
work with lead powder. Silicon is flammable in powder form and is slightly toxic. Do not
breathe or handle any fine silicon powder remaining on the bottom of the reagent bottle.
Wear chemical splash goggles and chemical-resistant gloves and apron. Wash your hands
with soap and water before leaving the laboratory.
Procedure
1. Label three 50 mL beakers or small containers Si (silicon), Sn (tin) and Pb (lead).
2. Obtain approximately 8 g of silicon chunks in the appropriately labeled beaker. Measure
the combined mass of the beaker plus solid to the nearest 0.01 g and record the value in
the Data Table. (Note: This value is the initial mass for sample 1.)
3. Fill a 25 mL graduated cylinder approximately half-full with water. Measure the initial
volume of water and record the value to the nearest 0.1 mL in the Data Table.
4. Using forceps or tongs, carefully add about one-third of the silicon chunks to the
graduated cylinder. Add the solid slowly, so as to avoid splashing or breaking the glass
cylinder.
5. Measure and record the new (final) volume of water plus solid in the graduated cylinder.
6. Measure and record the combined mass of the labeled beaker and remaining solid in the
Data Table. (Note: This value is the final mass for sample 1.)
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Period:
7. Repeat steps 4—6 twice with some of the remaining amount of solid in the beaker. Do
NOT empty the graduated cylinder between samples. The final volume of the previous
sample becomes the initial volume for the next sample.
8. Record all initial and final mass and volume data in the Data Table. There should be a
total of three sets of mass and volume data (samples 1-3).
9. After all three trials have been completed, empty the water from the graduated cylinder.
Carefully pour all the silicon chunks onto a paper towel and allow them to dry. Do not allow
any of the solid to go down the drain.
10. Rinse the graduated cylinder with water.
11. Obtain approximately 25 g of tin shot in the appropriately labeled beaker. Measure the
initial mass of the beaker plus solid to the nearest 0.01 g and record the value in the Data
Table.
12. Repeat steps 3-10 using tin. Record all initial mass, final mass and volume data in the
Data Table.
13. Obtain approximately 35 g of lead shot in the appropriately labeled beaker. Measure
the initial mass of the beaker plus solid to the nearest 0.01 g and record the value in the
Data Table.
14. Repeat steps 3-10 using lead. Record all initial mass, final mass and volume data in the
Data Table.
15. Return the correctly labeled solids to your instructor for reuse.
Data Table:
Element
Sample
Initial
Mass
(g)
Final
Mass
(g)
Mass of
Solid
(g)
Initial
Volume
(mL)
Final
Volume
(mL)
Volume
of
Solid
(mL)
1
Silicon
2
3
1
Tin
2
3
1
Lead
2
3
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Post-Lab Calculations:
Date:
Period:
1. Complete the Data Table: Calculate both the mass (initial mass -final mass) and volume
(final volume -initial volume) for each trial. Record these results in the Data Table.
2. It would be advisable to add to the table above or construct a new table for the
information from your calculations. The table should include density for each trial, an
average density for each element and a plus or minus value for each average density.
Note: The density of a solid is usually reported in units of g/cm 3.
3. Using the mass and volume data, calculate the density of each tria.
4. Calculate the average value (mean) of the densities for each element, silicon, tin, and
lead. Record all results in the Results Table. Use the range of density values for each
element to estimate “plus-or-minus” (±) error for each average (e.g., if your densities are
6.8 g/cm3, 7.0 g/cm3 and 7.2 g/cm3 your average would be 7.0 g/cm3. Your plus or minus
value would be ±0.2 g/cm3).
5. On a graph, plot the period number of Si, Sn, and Pb on the x-axis versus the average
density of each element on the y axis. Using a ruler or straightedge, draw a “best-fit”
straight line through the data points. Use this “best-fit” straight line to predict the density of
germanium.
6. Look up the actual density of germanium in a reference source and calculate the percent
error between the predicted and actual values (see Pre-Lab Question #3).
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Period:
Law of Conservation of Mass Experiment
The experiment will explore whether matter is created or destroyed during a chemical
reaction.
Materials
solutions of 0.1M - NaOH, CuSO4, NH4OH, and Na2CO3
1 graduated cylinder
4 plastic cups
balance
Procedure
Label the four cups to contain the solutions (one each for NaOH, CuSO 4, NH3 (aq), and
Na2CO3). You can do this by placing a paper towel under each cup and labeling it with the
chemical formula.
Trial 1
1. Use a graduated cylinder to measure about 60 mL of the NaOH solution, and then
pour the solution into its designated cup.
2. Clean out the graduated cylinder to measure about 60 mL of the CuSO4 solution,
and then pour the solution into its designated cup.
3. Carefully place the two solutions (in their cups) on the balance. Mass the solutions
and their containers together and record their combined mass in the data table.
4. Pour the NaOH solution into the container with the CuSO4 solution. Allow the
solutions to mix. Describe what happens in the data table.
5. Mass both containers and the mixture again. Record the new mass.
a. Did the mass change?
6. Clean out the cups & dispose of solutions by following the directions indicated by
your teacher.
Trial 2
7. Use a graduated cylinder to measure about 60 mL of the NH4OH solution, and then
pour the solution into its designated cup.
8. Clean out the graduated cylinder to measure about 60 mL of the CuSO4 solution,
and then pour the solution into its designated cup.
9. Carefully place the two solutions (in their cups) on the balance. Mass the solutions
and their containers together and record their combined mass in the data table.
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Period:
10. Pour the NH4OH solution into the container with the CuSO4 solution. Allow the
solutions to mix. Describe what happens in the data table.
11. Mass both containers and the mixture again. Record the new mass in the data table.
a. Did the mass change?
12. Clean out the cups & dispose of solutions by following the directions indicated by
your teacher.
Trial 3
13. Use a graduated cylinder to measure about 60 mL of the Na2CO3 solution, and then
pour the solution into its designated cup.
14. Clean out the graduated cylinder to measure about 60 mL of the CuSO4 solution,
and then pour the solution into its designated cup.
15. Carefully place the two solutions (in their cups) on the balance. Mass the solutions
and their containers together and record their combined mass in the data table.
16. Pour the Na2CO3 solution into the container with the CuSO4 solution. Allow the
solutions to mix. Describe what happens in the data table.
17. Mass both containers and the mixture again. Record the new mass in the data table.
a. Did the mass change?
18. Clean out the cups & dispose of solutions by following the directions indicated by
your teacher.
Data and Observations
Solutions Mixed
Mass prior to mixing
solutions (g)
Mass after mixing solutions
(g)
NaOH and CuSO4
Description of Reaction:
Solutions Mixed
Did the mass change?
Mass prior to mixing
solutions (g)
Mass after mixing solutions
(g)
NH4OH and CuSO4
Description of Reaction:
Did the mass change?
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Name:
Solutions Mixed
Date:
Mass prior to mixing
solutions (g)
Period:
Mass after mixing solutions
(g)
Na2CO3 and CuSO4
Description of Reaction:
Did the mass change?
Analysis Questions:
1) What evidence did you see in each reaction to know for certain that a chemical
reaction occurred? (Recall the evidence of a chemical change).
2) How did the final mass of each system compare to the initial mass of each system?
Be specific for each reaction.
3) Based on how you answered question 2, for each reaction – do any of the reactions
(appear to) violate the Law of Mass Conservation? If yes, state which mixtures (be
specific).
4) What could be a possible reason why, one or more of the reactions appeared to
violate the Law of Mass Conservation? (Come up with a reason even if all of your
reactions did not appear to violate the law).
Extension Question:
5) When a log burns, the resulting ash material has less mass than the unburned log
did. Explain whether this loss of mass violates the Law of Mass Conservation.
25
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Period:
Artifact Ages
26
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Date:
Period:
27
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Date:
Period:
28
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Date:
Period:
Mole Lab
Directions:
1. Carefully examine each vial – write a good descriptive physical description of the
element in the space provided in the data table.
2. To calculate the mass of the sample in each vial, determine the mass of the sealed
vial and its contents then subtract the mass of the empty vial and its cap (written at
the balance) from the total mass. Write this final value in the data table.
a. Since each vial contains 1 mole of an element, the difference is the mass of
one mole of the element (or you could say it’s the molar mass of the element).
3. Take the determined mass of the element and look up the mass on the periodic
table. Record the element name and symbol for the element that has the mass that
is closest to the mass you received in the lab.
4. Take the “molar mass” of each element that you determined by obtaining the mass
and determine the amount of atoms in the sample. Show your work stations 1 -6.
Place your answers in the data table. For the remaining stations, you do not have to
show your work, but you must have the answer in the data table.
Data Table:
Station
#
Physical Description
Final
Mass of
Element
Element Name
Element
Symbol
Number of
Atoms
1
2
3
4
5
6
7
29
Name:
Station
#
Date:
Physical Description
Final
Mass of
Element
Element Name
Period:
Element
Symbol
Number of
Atoms
8
9
10
11
12
Answer the following questions:
a. Which of the elements are metals?
b. Which of the elements are metalloids?
c. Which of the elements are non-metals?
d. How do the physical properties of metals differ from non-metals?
e. Does each sample have the same mass? Why or why not?
30
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Period:
Atomic Mass of Vegium
Purpose:
This lab is designed for students to determine the relative
isotope abundance and the average atomic mass of the
fictitious element, vegium.
Background information:
Isotopes are atoms of the same element that have different masses due to different
numbers of neutrons. The average atomic mass reported on the periodic table is the
weighted average of all isotopic masses of that element. The weighted average takes into
account both the mass and the relative abundance of each isotope as it occurs.
In the scientific community, scientists use a mass spectrometer to determine the average
atomic mass of an element. The mass spectrometer works by separating particles by mass
and then measuring the relative abundance of each. (The
relative abundance allows them to compare the amount of
each isotope against the other isotopes of an element). From
this data, the weighted average is calculated and reported as
the average atomic mass.
The weighted average allows for the fact that the amount of
each isotope of an element is not equal and there is a
difference in the mass of each atom due to a change in the
number of neutrons present in the nucleus.
Lab:
In this lab you will examine the element known as “Vegium”. Vegium consists of three
isotopes: cornium, beanium, and peaium. You will simulate the work of the mass
spectrometer by separating and counting counting the amount of each isotope you have in
a sample. Then you will determine the mass of each individual isotope, and eventually use
the data and calculations to determine the average atomic mass of Vegium.
Procedure:
Step 1:
Separate the beans, peas, and corn kernels. Mass all the
beans, then all the peas, and finally all the corn and record
your data for each on the data table.
Step 2:
Count all the beans, peas, and corn kernels. Record your data for each on the data table.
31
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Period:
Step 3:
Divide the total mass of each isotope by the number of each isotope to get the average
mass of each isotope. For the total column for step 3, calculate it just like you calculated
the average mass of each isotope.
Step 4:
Divide the number of each isotope by the total number of “atoms” (total in step 2), to get the
relative abundance.
Step 5:
Multiply the relative abundance of each isotope by 100 to get the percent abundance.
Step 6:
Multiply the relative abundance from step 4 by the average mass of each isotope (obtained
in step 3) to get the relative weight of each isotope.
Step 7:
Add the relative weights to get the average mass of all particles in vegium, the atomic
mass.
Step 8:
Return your sample of vegium in its weigh-boat.
Vegium:
1) Cornium
2) Beanium
3) Peaium
Interesting isotopes – if I do
say so myself.
32
Name:
Date:
Period:
Data Table:
BEANS
PEAS
CORN
TOTAL (add the
previous 3 columns
together to get this
value; except in step 3)
Mass of each
isotope
Number of
each isotope
Average
mass of each
“atom”
Relative
abundance
Percent
abundance
Relative
weight
Analysis Questions:
1) Which of your data (in the table) can be obtained through measuring? Which had to
be calculated? Be specific.
2) What is the difference between the percent abundance and the relative abundance?
33
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3) What is the difference between the average mass you calculate in step 3 and the
relative weight you calculated in step 6?
4) Compare the average atomic mass (relative weight) you received for your vegium
sample with two other lab groups. Record their values below:
a. Group 1: _____________
b. Group 2: _____________
c. Your value: ____________
Are they similar?
How would the difference be affected if larger samples were used?
5) If the accepted average atomic mass of vegium is 0.158g. Determine your percent
error:
% error = (|Experimental Value - Accepted Value| / Accepted Value) x 100
6) Describe sources of error in the procedure/sample.
34
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Date:
Period:
Flame Tests
Fingerprints are unique to each person, so is the color of light emitted by an element
heated in a flame. You will explore the characteristic color that is emitted by calcium,
copper, potassium, sodium, strontium, lithium, and barium.
When substances are exposed to heat, the atoms absorb the heat energy. When this
occurs the absorbed energy excites the electrons and allows them to “jump” to excited
energy levels. The natural tendency when electrons jump to an excited energy level is to
return to their ground state. When this transition occurs a particle of light, a photon, is
emitted. Electrons may return to their ground state in a single step or they may take
multiple steps, in which they emit a photon each time they take a step towards their original
ground state. The energy of each emitted photon is equal to the difference in energy
between the excited state and the state to which the electron returns. The energy emitted
determines the color of light observed in the flame.
We can use the following equation to determine the energy of a photon:
∆E = hc
λ
Where ∆E is the difference in energy between the two energy levels (unit – joules) and h is
Planck’s constant at 6.626x10-34 J/s, and c is the speed of light at 2.998x108 m/s. Lambda
(λ) is the wavelength of light in meters.
The visible light spectrum falls between 700nm and 400nm, where 700nm is red and
400nm is violet. There are regions in which a certain color can be obtained and we can
estimate a representative wavelength for the color. See table.
Wavelength Region (nm)
400-425
425-480
480-500
500-560
560-580
580-585
585-650
650-700
Representative Wavelength
(nm)
410
470
490
520
565
580
600
650
Color
Violet
Blue
Blue-Green
Green
Yellow-Green
Yellow
Orange
Red
Materials: A wooden splint, soaked, in each of the following chemical salts: calcium,
copper (II), lithium, potassium, sodium, strontium, barium, unknown 1. Bunsen burner,
match book, beaker with water
35
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Period:
Pre-lab:
Fill in the blanks: When an atom absorbs energy, the electrons move from their
___________________ state to an ___________________ state. When an atom emits
energy, the electrons move from a(n) ___________________ state to their
___________________ state and give off ___________________.
Is a flame test a qualitative and quantitative test? ___________________
Procedure:
1) Obtain a tray and a paper towel. On the paper towel, list each salt across the
bottom.
2) Obtain one wooden splint, soaked in its salt, and place it on the paper towel above
the label.
3) Set up your burner, and light it.
4) Wait until all students are ready to begin. The instructor will turn off the lights so that
the results are more visible.
5) Carefully place a wooden splint into the flame. Carefully move the wooden splint
back and forth, and watch for a unique color flame to be produced.
6) Record the color of the splint.
7) Place the splint in the beaker containing water.
8) Continue steps 5 – 7 until all splints have been burnt.
9) Discard wooden splints and matches in the trash can, and empty water into the sink.
10) Please put all materials used away properly.
TABLE of RESULTS
Metal Ion
Color of
Flame
λ (nm)
λ (m)
∆E (J)
Calcium
Barium
Strontium
Copper (II)
Sodium
Lithium
Potassium
Unknown 1
In the above table, when you have completed the flame tests complete the following:
Obtain the wavelength in nanometers from table 1 on the first page of the lab. Fill in the
wavelength based on the color you observed in the lab. Convert the wavelength from
36
Name:
Date:
Period:
nanometers into meters and place that value in the next column. Finally, calculate the
change in energy by using the equation on the first page of the lab and your wavelength
(m). Place that answer in the chart in the column labeled ∆E.
Post Lab Questions
1) In the space below, show the calculation (your work with units) for ∆E of calcium.
2) What element(s) makes up the unknown substance(s) you tested?
3) What evidence can you cite that indicates that the color produced was due to the
metal ion in the compound?
4) The alkali metals, cesium and rubidium, were discovered based on their
characteristic flame color. Cesium is named after the sky and rubidium after a
beautiful gemstone. What colors of light do you think these colors give off when
heated in a flame?
5) How does the information obtained from flame tests relate to real world applications
of these elements? (Think about entertainment).
37
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Period:
REACTIVITY OF ALKALINE EARTH METALS
Introduction:
The elements in group 2A of the periodic table are called the alkaline earth metals.
The name was given to these elements because they were first isolated from compounds in
which they were combined with oxygen. They were called “earths” by early chemists, and
later the word “alkaline” was added because they formed basic or alkaline solutions in
water. This can be tested in the lab by adding a solution of phenolphthalein which turns
pink when in the presence of a base.
The alkaline earth metals have two valence electrons, which they tend to give up
easily; therefore they are quite reactive. They are never found free in nature and they must
be protected from both air and water, in order to remain in their unreacted state.
Magnesium and calcium are obtained in their elemental state through a process called
electrolysis, and then are stored in air tight containers.
The purpose of the lab is to test the reactivity of magnesium and calcium, along with
two other alkaline earth metals. You will then compare the reactivity of each with aluminum,
which is part of group 3A on the periodic table.
Pre – Lab:
1) What are some common properties (characteristics) of the alkaline earth metals?
2) In a solution how would know if the solution is alkaline (or basic)?
3) Why are alkaline earth metals very reactive in nature?
Materials:
Goggles
Aprons
3 test tubes
Forceps
Distilled water
Micropipets
Bunsen burner
Wire gauze square
Scoopula
10 ml Graduated
Cylinder
Tongs
Matches
Steel wool
Phenolphthalein
Magnesium ribbon
Calcium metal
Aluminum foil
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SAFETY:




DO NOT TOUCH THE CALCIUM WITH YOUR HANDS. IF YOU ARE
EXPOSED IMMEDIATELY FLUSH WITH WATER.
DO NOT HAVE AN OPEN FLAME WHEN EXPOSING THE CALCIUM TO
WATER, THIS COULD CREATE AN EXPLOSIVE GAS.
DO NOT LOOK DIRECTLY AT THE MAGNESIUM WHEN BURNING – USE
PERRIFERAL VISION.
KEEP LONG HAIR TIED BACK!
Procedure:
1) Put on your safety gear: goggles and apron.
2) Retrieve one piece of calcium, magnesium, and aluminum from the front lab
table. Place each in a test tube.
3) Place the test tubes in the test tube rack when you return to your lab table.
Calcium
4)
5)
6)
7)
Measure out 3 mL of distilled water in a graduated cylinder.
Slowly pour the 3 mL of distilled water into the test tube with the calcium.
Observe all reactions, and record.
When the reaction stops, check to see if any metal remains. If it does, add 1 mL
of water to the test tube again.
8) When all the metal is reacted, place 1 drop of phenolphthalein into the test tube.
9) Record all observations.
Magnesium
10) Retrieve the magnesium strip from the test tube, and clean it thoroughly with the
steel wool. (It should be really shiny)
11) Place the magnesium strip back into the test tube.
12) Measure out 3 mL of distilled water in a graduated cylinder.
13) Slowly pour the 3 mL of distilled water into the test tube with the magnesium.
14) Observe all reactions, and record.
15) Place 1 drop of phenolphthalein into the test tube.
16) Record all observations.
17) Remove the magnesium strip from the test tube, and dry it off thoroughly.
18) Set aside for later.
Aluminum
19) Measure out 3 mL of distilled water in a graduated cylinder.
20) Slowly pour the 3 mL of distilled water into the test tube with the aluminum.
21) Observe all reactions, and record.
22) Place 1 drop of phenolphthalein into the test tube.
23) Record all observations.
24) Remove the aluminum strip from the test tube, and dry it off thoroughly.
39
Name:
25) Set aside for later.
Date:
Period:
Clean Up
26) Empty the contents of your calcium test tube into the beaker marked waste by
your instructor. Rinse all test tubes in the sink. Flush with lots of water. Clean
test tubes using detergent solution, and a brush.
PART B
27) In two separate test tubes, measure 3 mL of distilled water in a graduated
cylinder and place in each.
28) Light the Bunsen burner.
Magnesium
29) Using the forceps, pick up the magnesium metal strip.
30) Place the strip into the flame, and look away.
31) As soon as it lights, pull it out of the flame and let it continue to burn. When the
flame goes out, drop it into one of the test tubes with distilled water.
32) Place 1 drop of phenolphthalein into the test tube.
33) Record all observations.
Aluminum
34) Using the forceps, pick up the magnesium metal strip.
35) Place the strip into the flame. Attempt to light. It may not catch on fire, but it may
start to charcoal after a while. This is okay.
36) After a periods of time (when it starts to turn black) drop it into one of the test
tubes with distilled water.
37) Place 1 drop of phenolphthalein into the test tube.
38) Record all observations.
40
Name:
Date:
Period:
Data Table:
Step
Calcium + Water
Observations
Calcium + Water +
Phenolphthalein
Magnesium + Water
Magnesium + Water +
Phenolphthalein
Aluminum + Water
Aluminum + Water +
Phenolphthalein
Magnesium burning
Burned Magnesium +
Water +
Phenolphthalein
Aluminum burning
Burned Aluminum +
Water +
Phenolphthalein
Post Lab Questions:
1) Was the reaction of calcium in water an exothermic reaction, meaning energy
was released? If yes, how do you know?
2) Did any of your solutions turn pink after you added a drop of phenolphthalein to
the solution? If yes, which ones?
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Name:
Date:
Period:
3) Looking at the metals you used. Are the alkaline earth metals more or less
reactive than aluminum? Support your answer with details.
4) Look at your observations between magnesium and calcium, in relation to their
position on the periodic table. As you move from top to bottom on the periodic
table, does the reactivity of the metals increase or decrease? Support your
answer.
5) Based on your answer to question 4, would strontium (Sr) be more or less
reactive than calcium?
6) Which metal, Calcium or Aluminum would be more likely to be used in structural
materials? Explain why.
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Name:
Date:
Period:
Exploring the Halides
The elements in Group 7A of the periodic table are nonmetals called the halogens. The
word Halogen, comes from two Greek words that mean “salt” and “former.” There are a
total of five members of this group; fluorine, chlorine, bromine, iodine, and astatine.
Each halogen can react with a number of different metals to form compounds, called
halides. For example sodium reacts with chlorine to form the compound called sodium
chloride. The first four halogens from classes of halides with easily recognizable
names:
Fluorides – metal + fluorine
Chlorides – metal + chlorine
Bromides – metal + bromine
Iodides – metal + iodine
In this investigation you will observe chemical reactions of small quantities of four
halides with specific solutions. You will use your observations of these reactions to
determine which halide is present in a solution containing one or more unknowns.
Pre-lab Questions:
1) What are the five halogens and what is the chemical symbol of each?
2) What are fluorides, chlorides, bromides, and iodides?
3) What method will you use to study halide reactions?
Materials:
goggles
apron
microtitration plate (24 well)
micropipets with the following solutions:
sodium chloride (NaCl)
sodium bromide (NaBr)
sodium iodide (NaI)
silver nitrate (AgNO3)
sodium thiosulfate (Na2S2O3)
ammonium hydroxide (NH4OH)
starch sol’n
bleach (NaClO)
toothpicks (about 20 qty)
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Name:
Date:
Period:
Caution:
** Wear goggles and aprons at all times!
** Handle all chemicals carefully!
** Do not inhale fumes from bleach or ammonium hydroxide!
** Never mix bleach and ammonium hydroxide!
**Silver nitrate causes brown stains on hands and clothing so be careful not to
touch, spill, or splash it.
Procedure:
Step 1:
In the microtitration plate drawn below, place five drops of each solution in each row, as
indicated by the drawing.
A
B
C
D
1
2
3
4
5
6
44
Name:
Date:
Period:
In column B, see arrow, add 2 drops of silver nitrate to each well (1 – 5). Mix
thoroughly with a toothpick.
Record the observations on the data table provided on page 6. Detail color and if
a precipitate forms. If no reaction occurs, write NR in the data table.
In the same column, add 2 drops of ammonium hydroxide solution, and mix thoroughly
with a toothpick.
Record the observations on the data table provided on page 6. Detail color and if
a precipitate forms. If no reaction occurs, write NR in the data table.
In column C, see arrow, add 2 drops of silver nitrate to each well (1 – 5). Mix
thoroughly with a toothpick.
Record the observations on the data table provided on page 6. Detail color and if
a precipitate forms. If no reaction occurs, write NR in the data table.
In the same column, add 5 drops of sodium thiosulfate solution, and mix thoroughly with
a toothpick.
Record the observations on the data table provided on page 6. Detail color and if
a precipitate forms. If no reaction occurs, write NR in the data table.
In column D, see arrow, add 2 drops of starch solution to each well (1 – 5), and then
add 2 drops of bleach to each well (A – E). Mix thoroughly with a toothpick.
Record the observations on the data table provided on this page. Detail color
and if a precipitate forms. If no reaction occurs, write NR in the data table.
45
Name:
Date:
Data Table
Well Solution
1
2
3
4
5
NaF
NaCl
NaBr
NaI
Unknown
Period:
Column 1
Column 2
Column 3
Column 4
Ca(NO3)2
AgNO3 +
NH4OH
AgNO3 +
Na2S2O3
Starch +
Bleach
White
precipitate
N.R.
N.R.
N.R.
N.R.
N.R.
N.R.
NO
REACTION
NO
REACTION
NO
REACTION
NO
REACTION
Post Lab Questions
Answer each of the following by using the lab data, pre-lab information and text book.
1) Look at the reaction results for sodium fluoride. Do they follow the same pattern
as the other halides? Explain any differences.
2) Based on your observations, describe the precipitates formed by the reactions of
the halides with silver nitrate. Be specific.
46
Name:
Date:
Period:
3) Which of the precipitates formed with silver nitrate dissolved when ammonium
hydroxide was added?
4) Which of the precipitates formed with silver nitrate dissolved when sodium
thiosulfate was added?
5) How did the halides react with the starch solution? Be specific.
6) How did the halides react with the bleach solution? Be specific.
7) What is the unknown halide?
8) How did you determine this?
47
Name:
Date:
Period:
CHEMICAL BONDS
Chemical compounds are combinations of atoms held together by chemical bonds.
These chemical bonds are of two basic types—ionic and covalent. Ionic bonds result
when one or more electrons from one atom or group of atoms is transferred to another
atom. Positive and negative ions are created through the transfer of these electrons. In
covalent compounds no electrons are transferred; instead electrons are shared by the
bonded atoms.
The physical properties of a substance, such as melting point, solubility, and
conductivity, can be used to predict the type of bond that binds the atoms of the
compound. In this experiment, you will test six compounds to determine these
properties. Your compiled data will enable you to classify the substances as either ionic
or covalent compounds.
OBJECTIVES





Compare the melting points of six solids.
Determine the solubility of the solids in water and in ethanol.
Determine the conductivity of water solutions of the soluble solids.
Classify the compounds into groups of ionic and covalent compounds based on
the properties observed.
Summarize the properties of each group (ionic vs. covalent).
SAFETY




Always wear safety goggles, gloves, and a lab apron to protect your eyes
and clothing. If you get a chemical in your eyes, immediately flush the chemical
out at the eyewash station while calling to your teacher. Know the location of the
emergency lab shower and eyewash station and the procedures for using them.
Do not touch any chemicals. If you get a chemical on your skin or clothing,
wash the chemical off at the sink while calling to your teacher. Do not taste any
chemicals or items used in the laboratory. Never return leftovers to their original
container; take only small amounts to avoid wasting supplies.
Do not heat glassware that is broken, chipped, or cracked. Use tongs or a
hot mitt to handle heated glassware and other equipment because hot glassware
does not always look hot.
When using a burner, confine long hair and loose clothing. If your clothing
catches on fire, WALK to the emergency lab shower and use it to put out the fire.
48
Name:
Date:
Period:
MATERIALS
• 24-well microplate
• Burner
• Calcium chloride
• Citric acid
• Conductivity tester (not needed)
• Ethanol
• Goggles
• Iron ring
• Lab apron
• Lauric acid
• Potassium iodide
• Ring stand
• Safety goggles
• Sodium chloride
• Sucrose
• Thin-stemmed pipets (2)
• Wire gauze
• Metal tray
49
Name:
Date:
Period:
PROCEDURE
1. Put on safety goggles and a lab apron.
2. Before you begin, write a brief description of each of the six substances in Table 1.
3. Place a metal tray on an iron ring attached to a ring stand. Position the ring so that it is just above
the tip of a burner flame. Light the burner for a moment to check that you have the correct height.
Then extinguish it.
4. Place a few crystals of sucrose, sodium chloride, lauric acid, calcium chloride, citric acid, and
potassium iodide in separate locations on the metal tray atop the wire gauze. Do not allow the
samples of crystals to touch.
5. For this experiment, it is not necessary to have exact values for the melting point. The lid will
continue to get hotter as it is heated, so the order of melting will give relative melting points. Light the
burner and observe. Note the substance that melts first by writing a 1 in Table 1. Record the order of
melting for the other substances.
6. After 3 min, record an n in Table 1 for each substance that did not melt. Extinguish the burner.
Allow the metal tray and wire gauze to cool while you complete the remainder of the experiment.
7. Put a few crystals of each of the white solids in the top row of your microtitration plate. Repeat with
the second row. Add 10 drops of water to each well in the top row. Do not stir. Record the solubility of
each substance in Table 1. Use ‘yes’ and ‘no’ as your observations.
8. Add 10 drops of ethanol to each well in the second row of the microtitration plate. Do not stir.
Record the solubility of each substance in Table 1. Use ‘yes’ and ‘no’ as your observations.
9. THIS WILL BE DONE BY THE TEACHER. Test the conductivity of each water solution in the top
row by dipping both electrodes into each well of the microtitration plate. Be sure to rinse the
electrodes and dry them with a paper towel after each test. If the bulb of the conductivity apparatus
lights up, the solution conducts electric current. Record your results in Table 1.
10. Clean the microtitration plate with soap and water with a test tube brush. Wash your hands
thoroughly before you leave the lab and after all work is finished.
50
Name:
Date:
Period:
TABLE 1: CHARACTERISTICS OF COMPOUNDS
Substance
Description
of
Compound
Melting
Point
Solubility in
H2O
Solubility in
Ethanol
Conductivity
Calcium
chloride
Citric acid
Lauric acid
Potassium
iodide
Sodium
chloride
Sucrose
ANALYSIS & CONCLUSION
1. Organizing Results: Group the white substances (write their names down) into two groups
based on the properties (Description of Compound, Melting Point, Solubility in H 2O, Solubility
in Ethanol, Conductivity) you observed.
2. Organizing Results: Keeping the groups in mind from number one, make a list the properties
of each group. (See page 193 of your text for assistance).
51
Name:
Date:
Period:
3. Inferring Conclusions: Use your textbook and your experimental data to determine which of
the groups consists of ionic compounds and which consists of covalent compounds. State the
substances which belong in each type of compound.
4. Relating Ideas: Write a statement to summarize the properties of ionic compounds and
another statement to summarize the properties of covalent compounds using the information
you received in your lab and from your textbook.
52
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Period:
53
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Date:
Period:
54
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Date:
Period:
55
Name:
Date:
Period:
56
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Date:
Period:
57
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Date:
Period:
58
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Date:
Period:
Weighing: A Means of Counting
Objectives:
1) Measure masses of common compounds, objects, and minerals.
2) Calculate the number of moles and atoms from experimental masses.
Introduction:
You can often measure how much of something you have by counting the individual objects
present. For instance, if you wanted to know how many pennies are in your pocket or how
many pencils are in your book bag, you would simply count them. In chapter 7 you are
learning that in chemistry there is a specific name for the number of atoms, ions or molecules
present. One mole of a substance is equal to 6.022x1023 atoms, ions or molecules of that
substance. You also know that you can “count” the number of moles by massing the
substance.
Purpose:
In this lab you will measure the masses of samples that are commonly found. You will use
your results as a means of counting the atoms, ions, and molecules in your samples. You will
also extend your base knowledge by considering common objects that can be considered
pure, like chalk and polystyrene peanuts. You will measure the masses of various substances
and find the number of atoms, ions or molecules in each.
Equipment:
Balance, plastic spoon, sodium chloride, water, sucrose, silicon dioxide, calcium carbonate,
polystyrene peanut, sulfur, plumbic sulfide, and ferric oxide.
Procedure:
1) Mass one level teaspoon of sodium chloride, water, and sucrose. Record the value in the data
table.
2) Mass the silicon dioxide, calcium carbonate, and polystyrene peanut. Record the value in the
data table.
3) Mass the sulfur, plumbic sulfide, and ferric oxide. Record the value in the data table.
4) Determine the molar mass for each substance and record the value in the data table.
5) Calculate the number of moles, the number of moles for each element in the sample, and the
number of atoms present in each element for all samples. SHOW YOUR WORK!! You may
use the last page and additional sheets of notepaper.
59
Name:
Date:
Data Table:
Name
Formula
Sodium
chloride
NaCl
Water
H2O
Sucrose
C12H22O11
Silicon
dioxide
SiO2
Calcium
carbonate
CaCO3
Mass
(g)
Molar Mass
(g/mol)
Moles
in
sample
Period:
Moles of each
element
Grams of each element
60
Name:
Name
Formula Mass
(g)
Polystyrene
peanut
C7H8
Sulfur
S8
Plumbic
sulfide
PbS2
Ferric
oxide
Fe2O3
Molar Mass
(g/mol)
Date:
Moles Moles of each
in
element
sample
Period:
Grams of each element
Data Table
Analysis Question:
1) Looking at sodium chloride, water, and sucrose, which had the greatest number of moles
present in one teaspoon?
61
Name:
Date:
Period:
Ions and Solubility
Introduction
What do stalagmites and stalactites found in caverns have in common with the deposits found on old
water faucets? How were many minerals, now mined as ores, originally formed? The answers to both
questions can be found in a study of precipitates. If a positive ion (cation) of a dissolved salt reacts
with the negative ion (anion) of a different compound to form a new salt with low solubility, chemists
say that a precipitate has formed.
Purpose
To determine which ions react to produce precipitates by analyzing data regarding mixtures of ionic
compounds.
Safety Considerations
Wear protective glasses and an apron at all times. Avoid skin contact the solutions. Dispose of all
solutions in the containers provided by your teacher. Wash your hands before leaving the laboratory.
Pre-Lab
Write the formulas for the following compounds:
1. Silver nitrate
______________________
2. Barium nitrate
______________________
3. Calcium nitrate
______________________
4. Copper (II) nitrate
______________________
5. Zinc nitrate
______________________
6. Lead (II) nitrate
______________________
7. Potassium nitrate
______________________
8. Iron (III) nitrate
______________________
9. Sodium carbonate
______________________
10. Sodium chloride
______________________
11. Sodium hydroxide
______________________
62
Name:
12. Sodium sulfate
Date:
______________________
13. Sodium phosphate
______________________
14. Sodium oxalate
______________________
15. Sodium iodide
______________________
16. Sodium thiosulfate
______________________
Period:
Procedure
1. Place a clean plastic sheet over the data grid similar to the one at the end of this laboratory
activity.
2. To each section of horizontal Row A (A1 through A8) add two drops of 0.1 M silver nitrate
solution.
3. As in Step 2, add FIVE drops of the solutions listed to each of the sections indicated:
To Row B, add 0.1 M barium nitrate.
To Row C, add 0.1 M calcium nitrate.
To Row D, add 0.1 M copper(II) nitrate.
To Row E, add 0.1 M zinc nitrate.
To Row F, add 0.1 M lead(II) nitrate.
To Row G, add 0.1 M potassium nitrate.
To Row H, add 0.1 M iron(III) nitrate.
4. To the vertical Column 1 (A1 through H1), add two drops of 0.1 M sodium carbonate.
5. As in Step 4, add FIVE drops of the solutions listed to each of the vertical columns indicated:
To Column 2, add 0.1 M sodium chloride.
To Column 3, add 0.1 M sodium hydroxide.
To Column 4, add 0.1 M sodium sulfate.
To Column 5, add 0.1 M sodium phosphate.
To Column 6, add 0.1 M sodium oxalate, Na2C204.
To Column 7, add 0.1 M sodium iodide.
To Column 8, add 0.1 M sodium thiosulfate, Na2S2O3.
63
Name:
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Period:
6. Observe carefully to determine which combinations of solutions produced precipitates. In your
data table, record each combination of ions that showed the formation of a precipitate. Record
the color, texture, and other observations for each.
7. Dispose of chemical wastes and clean your plastic sheet as instructed by your teacher.
8. Wash hands thoroughly before leaving the laboratory.
Data Grid
Ag+
A
Ba2+
B
Ca2+
C
Cu2+
D
Zn2+
E
Pb2+
F
K+
G
Fe3+
H
CO32-
Cl-
OH-
SO42-
PO43-
C2O42-
I-
S2O32-
1
2
3
4
5
6
7
8
Data Analysis, Concept
1. How would you tell another student how to decide whether a precipitate has formed when two
solutions are mixed? Use examples from the activity to help you in this explanation.
2. Which cation(s) formed precipitates with all its anions horizontally across in Columns 1 through
8?
3. What anion is found in all solution combinations in this activity? What cation is common to all
the solution mixtures?
64
Name:
Date:
Period:
4. List the names of the pairs of reagents that produced each precipitate in vertical Column 5.
5. Analyze your data regarding the solubility behavior of carbonate CO 32-; sulfate, SO42-; and
iodide, I- State your conclusion to match the following sample conclusions: All chlorides are
soluble except silver chloride.
65
Name:
Date:
Period:
Law of Definite Composition
Purpose: To determine the formula and the percent composition of a copper sulfide compound
formed through synthesis.
Procedure:
LAB TABLE
1) Clean and thoroughly dry a porcelain crucible and its cover over a hot flame (blue) by holding
them with the crucible tongs.
2) Cool for two minutes.
3) Mass the crucible and cover to the nearest .01 gram. Record this in the data table.
4) Mass a 10 cm piece of fine copper wire. Record the mass in the data table to the nearest .01
gram.
5) Place the copper in the crucible.
6) Re-mass the crucible, copper, and cover. Record the mass in the data table to the nearest .01
gram.
7) Add .50 gram of flowers of sulfur to the crucible and copper wire. This mass does not need to
be recorded because it is in excess and will burn off as sulfur dioxide.
FUME HOOD
8) Place the crucible with contents and cover on a pipestem triangle that is set up in the fume
hood.
9) Heat the covered crucible and contents for as long as you can see flames from the burning
sulfur coming out from under the lid of the crucible.
10) After the flames cease continue to heat for two to three minutes.
11) Using the crucible tongs, carefully remove the lid and check it for sulfur residue or burning. If
this is present continue to burn until all sulfur has burned off.
12) Using the crucible tongs, carefully remove the lid and then the crucible from the pipestem
triangle. Place the crucible and cover on a watch glass that has been covered with aluminum
foil.
13) Let stand until cool. (may have to sit overnight)
14) Mass the crucible, contents and lid to the nearest .01 gram and record on the data table.
Data:
Mass (g)
Crucible & Cover
Copper Wire
Crucible, Cover & Cu Wire
Crucible, Cover & Cu Cmpd.
66
Name:
Date:
Calculations: Show all work, with units and labels.
Period:
1) Calculate the mass of the copper in the compound.
2) Calculate the mass of the sulfur in the compound.
3) Calculate the moles of copper in the compound.
4) Calculate the moles of sulfur in the compound.
5) Calculate the percent by mass of the copper in the compound.
6) Calculate the percent by mass of the sulfur in the compound.
7) What is the ratio of moles of copper to moles of sulfur in the compound?
8) Use the mole ratio to determine the formula and name of the compound.
Questions: Write in complete sentences.
1) Which element (Cu or S) was the limiting reagent in this lab experiment? How do you know?
2) What happened to the excess element?
3) How did you increase the potential for acid rain by doing this experiment?
67
Name:
Date:
Period:
Percent Composition and Formula of a Hydrate
Wear safety glasses/goggles and aprons when performing this lab activity.
PROCEDURE:
1. Mass a clean, dry crucible to maximum accuracy of the balance. (Data # 1)
2. By placing it into the crucible, mass out 2.50 g - 3.50 g of cupric sulfate hydrate. It should be a
fine granular consistency, not in chunks. Record new mass. (Data # 2)
3. Place the crucible on a pipestem triangle which is supported by a small ring on a ring stand.
4. Heat gently until the cupric sulfate hydrate turns all one color. If the sample starts turning
brown or dark gray around the edges, reduce the flame height or move the wire ring two
inches higher. Record all observations. (Data # 3)
5. Once all of the cupric sulfate hydrate is all one color. Turn off the Bunsen burner and allow the
crucible to cool for 2 minutes.
6. Remove the crucible from the pipestem triangle, using the crucible tongs and place it in the
dessicator. Allow the crucible to cool to room temperature (about 7-10 minutes).
7. Remove crucible from the dessicator and mass. Record the mass. (Data # 4)
8. IF TIME PERMITS, repeat steps 3 – 6.
9. Remove the anhydrous cupric sulfate and place it into a small beaker. Hold the beaker in your
hand and add about five drops of water to the beaker. Record your observations. ( Data # 5)
DATA:
#1) Mass of crucible: __________________
#2) Mass of crucible & cupric sulfate hydrate: __________________
#3) Observations: ________________________________________________________________
_______________________________________________________________________________.
#4) Mass of anhydrous cupric sulfate: __________________
#5) Observations: ________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________.
68
Name:
Date:
CALCULATIONS. Show all work with units.
Period:
1. Calculate the percent of water in the hydrate.
2. From the mass of the water lost in heating and the mass of the anhydrous cupric sulfate
determine the mole ratio of the water to the anhydrous cupric sulfate.
3. Write the formula for the hydrate of the cupric sulfate, from your results.
________________________________________________________________
QUESTIONS:
1. What is a hydrate? ________________________________________________
2. Use a reliable sources to determine the most likely formula of the hydrate. ____________
3. Which is more stable? (anhydrous cupric sulfate or the hydrate)
___________________________________________________________
4. How did part 9 of the procedure & data # 5 support your answer for question 5?
___________________________________________________________________
___________________________________________________________________
5. What does anhydrous mean? ________________________________________
6. Blue cobalt chloride turns pink when the relative humidity of the air is high. Explain why.
69
Name:
Date:
Period:
Balancing Chemical Equations Lab Activity
Objectives:
1) Write chemical equations from the given chemical reaction.
2) Balance each chemical reaction.
3) Make observations based on each chemical reaction completed.
Directions:
1) Write and balance the chemical equations for each of the following chemical reactions.
2) Verify your reaction with your teacher.
3) Go to the indicated lab station to test your reaction.
a. Take the pipette containing each chemical and place five drops of the first chemical
listed in the reaction in the watch glass.
b. To the watch glass, add five drops of the second chemical in your reaction.
c. Record your observations. Be descriptive with each of your observations. If you do not
see any change indicate there were no visual indicators seen.
d. Rinse the watch glass off under the water faucet, rinsing the solution down the drain
thoroughly. Dry the watch glass and make sure everything is in order for the next lab
group.
Materials:
 Watch Glass
 Chemicals
o Aluminum sulfate
o Ammonium hydroxide
o Barium nitrate
o Calcium chloride
o Cobalt (II) nitrate
o Copper (II) nitrate
o Copper (II) sulfate
o Iron (III) chloride
o Lead (II) nitrate
o Phosphoric acid



o Potassium carbonate
o Silver nitrate
o Sodium bicarbonate
o Sodium chloride
o Sodium hydroxide
o Sodium iodide
o Sodium phosphate
o Sodium sulfate
o Sulfuric acid
o Zinc metal
Safety goggles
Aprons
70
Name:
Date:
Period:
Pre – Lab Question:
1) Given the following chemical reaction, which is partially balanced:
K4Fe(CN)6 + 6 H2SO4 + 6 H2O 
K2SO4 +
FeSO4 + 3 (NH4)2SO4 + 6 CO
How many of each of the following components do you have present on each side:
Reactants:
Products:
_____ K
_____ K
_____ Fe
_____ Fe
_____ C
_____ C
_____ N
_____ N
_____ H
_____ H
_____ O
_____ O
_____ SO4
_____ SO4
To balance the reaction, what coefficient should be used, and in front of which chemical
formula would you place it? Write both the coefficient and the chemical formula in the
space below.
Chemical Equations:
1) Sodium bicarbonate and phosphoric acid react to form sodium phosphate, carbon dioxide, & water.
a. Observations of chemical reaction:
2) Barium nitrate reacts with sulfuric acid to produce barium sulfate and nitric acid.
a. Observations of chemical reaction:
3) Sodium hydroxide & copper (II) sulfate react to produce copper (II) hydroxide & sodium sulfate.
a. Observations of chemical reaction:
4) Aluminum sulfate and sodium hydroxide react to form aluminum hydroxide and sodium sulfate.
71
Name:
Date:
Period:
a. Observations of chemical reaction:
5) Sodium phosphate and calcium chloride react to form calcium phosphate and sodium chloride.
a. Observations of chemical reaction:
6) Potassium carbonate and sulfuric acid react to form potassium sulfate, carbon dioxide, and water.
a. Observations of chemical reaction:
7) Cobalt (II) nitrate reacts with sodium hydroxide to form cobalt (II) hydroxide and sodium nitrate.
a. Observations of chemical reaction:
8) Iron (III) chloride reacts with sodium sulfate to form sodium chloride and iron (III) sulfate.
a. Observations of chemical reaction:
9) Ammonium hydroxide and copper (II) sulfate to form copper (II) hydroxide and ammonium sulfate.
a. Observations of chemical reaction:
10) Sodium phosphate and iron (III) chloride react to form sodium chloride and iron (III) phosphate.
a. Observations of chemical reaction:
11) Iron (III) chloride reacts with sodium hydroxide to form iron (III) hydroxide and sodium chloride.
a. Observations of chemical reaction:
12) Zinc metal reacts with lead (II) nitrate to form zinc nitrate and lead.
72
Name:
Date:
Period:
a. Observations of chemical reaction:
13) Silver nitrate reacts with sodium chloride to produce sodium nitrate and silver chloride.
a. Observations of chemical reaction:
14) Sodium iodide and copper (II) nitrate react to produce copper iodide and sodium nitrate.
a. Observations of chemical reaction:
73
Name:
Date:
Period:
Classifying Chemical Reactions
Objective: In this experiment, you will observe example reactions representing the four
main classifications of chemical reactions. You will record your observations and answer
questions regarding each of the reactions.
Part I:
Procedure:
A. Obtain a piece of magnesium ribbon, approximately 10cm long. Roll the Mg into a
loose ball and place it in a clean, dry crucible. Record the mass of both the Mg and
crucible. Record a description of the Mg ribbon.
B. Set up the ring stand, ring, and pipestem triangle. Place the crucible in the
pipestem triangle. Light the Bunsen burner and heat the crucible and Mg. When
the Mg starts to glow, turn out the burner and allow the Mg continue to burn until it
stops burning.
C. Record your observations of the product. Once the crucible cools, mass the
product and the crucible.
D. Add a few drops of water to the product in the crucible. Waft the smell from the
crucible to detect an odor.
E. Clean out the crucible and dry it thoroughly.
Observations/Data:
Mass of crucible/Mg ribbon: ______________
Description of Mg:
Mass of crucible/product: _____________
Description of product:
Did you observe a change in mass? _________ Why or why not?_________________
_______________________________________________________________________
Did you detect an odor? ____________
What substance is responsible for the odor? __________________________
What is the chemical formula for the substance? __________
There are two primary gases present in the air. What are they? _______________
If each gas reacts separately with the Mg, two different substances are produced.
74
Name:
Date:
Period:
What are the first two chemical reactions in this experiment?
Once the products were formed, these products react with water.
What are the final two balanced chemical reactions in this experiment?
Part 2:
Procedure:
A. Obtain a piece of 16cm glass tubing with a rubber
stopper attached and two test tubes. Set up the
ring stand, test tube clamp, and burner as shown
in the diagram.
B. Place a very small amount of cupric carbonate
into one of the test tubes. Insert the rubber
stopper from the glass tubing into the test tube.
Place the test tube in the test tube clamp. Note
the color of the cupric carbonate.
C. In the second test tube add limewater (Calcium
hydroxide) to the test tube until it is about 2/3 the
way full. Place the open end of the glass tubing
into the test tube with the limewater. Hold this in
place with your hand.
D. Light the burner under the test tube with the
cupric carbonate in it. Heat the bottom of the test
tube. Record all observations.
E. Once the cupric carbonate completely changes
color, turn off the burner. Remove the limewater test
tube from the glass tubing.
F. Empty the contents, and clean thoroughly. Set aside
for the experiments in part 3.
Observations/Data:
Color of cupric carbonate: ______________
Observations of the reactions:
75
Name:
Date:
Period:
Two reactions occurred with this experiment. The cupric carbonate broke down into its
smaller components, and the calcium hydroxide reacted with the carbon dioxide produced
by the decomposition of the cupric carbonate. Write the two balanced chemical reactions.
Part 3:
Procedure:
A. Using the same materials from part 2 of the
experiment, obtain 5mL of 3.0M hydrochloric acid
and pour into one of the test tubes.
B. Place the test tube with the hydrochloric acid into the
test tube holder as shown in the diagram.
C. Obtain a few small pieces of “mossy” zinc.
D. Prior to placing the zinc into the test tube containing
the hydrochloric acid, make sure you have the glass
tube and empty test tube ready.
E. Drop the pieces of the zinc into the test tube with the
hydrochloric acid.
F. Quickly attach the rubber stopper/glass tube into this
test tube with the open end of the glass tube going
upwards. Cap the open end of the glass tube by
sliding the empty test tube down over it.
G. Touch the bottom of the test tube with the
hydrochloric acid and zinc. You will also see a gas
being emitted through the glass tube and into the
empty test tube.
H. After two minutes, slowly lift the test tube (collecting
the gas) off the glass tube, and quickly place your
thumb over the open end. Light a match and slowly
move your hand and the test tube close to the match.
Remove your thumb from the test tube.
I. Clean up thoroughly.
Observations/Data:
Description of the reaction:
76
Name:
Date:
Period:
Was the reaction endothermic or exothermic? _______________________
Write the chemical reaction that occurred.
What gas formed? _____________
What happened with the lit match and gas filled test tube?
Write the chemical reaction that occurred with that test.
Part 4:
Procedure:
A. Using two clean test tubes, place 4mL of ferric chloride hexahydrate into one of the
test tubes.
B. In the second test tube, place 4mL of sodium hydroxide.
C. Slowly pour the ferric chloride hexahydrate into the test tube containing the sodium
hydroxide. Record your observations.
D. Place the products formed in the waste container in the fume hood.
E. Clean up thoroughly, and put all materials away properly.
Observations/Data:
Description of the reaction:
77
Name:
Date:
Period:
Write the chemical reaction that occurred in this experiment.
Conclusion:
What indicators that a chemical reaction occurred are present in these four experiments?
78
Name:
Date:
Period:
Double Displacement Reactions
This experiment will enable you to identify a few unknown solutions by noting their
characteristic reactions with other known solutions. The seven solutions are:
silver nitrate
barium nitrate
sodium phosphate
copper (II) nitrate
sodium sulfate
sodium carbonate
ferric chloride
In order to learn certain characteristics of reactions of the preceding solutions, mix them in
pairs using the microtitration plate. For each chemical use 10 drops; once you add the
second chemical, mix the solution and record your results.
One good way to do this is to fill in the observation on page 3. For each reaction, write a
balanced equation. (See page 4 for this). Indicate if the reactants and products are in
solution (aq) or precipitates (s) by placing these subscripts behind each formula in every
equation. All reactions are double displacement reactions.
Use the table above to identify the unknown solution or solutions your instructor will give
you. Record all observations you observe from mixing part of the unknown solution with
portions of the solutions above. Record which solutions you used. Someone else should
be able to read your observations and repeat your work. Record you data, unknown
number or letter, and your conclusion at the bottom of the 1st page.
Along with writing balanced equations for each reaction you will identify the precipitate
formed in the reaction. You will do this by using an arrow pointing downward or marking it
with the subscript (s) for solid. That means that the chemical marked is insoluble in water
therefore does not dissolve, it comes out of solution. All soluble chemicals will be labeled
with the subscript (aq) and the equation will be balanced, example below:
Ba(NO3)2(aq) + K2SO4(aq)  BaSO4(s) + 2KNO3(aq)
Ions are separated by water molecules. The
ions are so small that they are invisible to the
human eye but are capable of reflecting light
in such a way that some of the solutions
have color.
79
Name:
Date:
Period:
DATA TABLE
AgNO3
Ba(NO3)2
Cu(NO3)2
FeCl3
Na2SO4
Ni(NO3)2
AgNO3
X
Ba(NO3)2
X
X
Cu(NO3)2
X
X
X
FeCl3
X
X
X
X
Na2SO4
X
X
X
X
X
Ni(NO3)2
X
X
X
X
X
X
X
X
X
X
X
Na2CO3
X
Na2CO3
X
Unknown # ______ Observations: ________________________________________
Conclusion: What chemical substance is it? __________________
Unknown # ______ Observations: ________________________________________
Conclusion: What chemical substance is it? __________________
80
Name:
Date:
Period:
Reactions: Solve the double displacement reactions, use the solubility rules &
show phases, and balance each reaction.
1)
AgNO3(aq) + FeCl3(aq) 
2)
AgNO3(aq) + Na2SO4(aq) 
3)
AgNO3(aq) + Ni(NO3)2 (aq) 
4)
AgNO3(aq) + Na2CO3(aq) 
5)
Ba(NO3)2(aq) + FeCl3(aq) 
6)
Ba(NO3)2(aq) + Na2SO4(aq) 
7)
Ba(NO3)2(aq) +
8)
Ba(NO3)2(aq) + Na2CO3(aq) 
9)
Cu(NO3)2(aq) +
10)
Cu(NO3)2(aq) + Na2SO4(aq) 
11)
Cu(NO3)2(aq) + Ni(NO3)2(aq) 
12)
Cu(NO3)2(aq) + Na2CO3(aq) 
13)
FeCl3(aq) + Na2SO4(aq) 
14)
FeCl3(aq) + Ni(NO3)2 (aq) 
15)
FeCl3(aq) + Na2CO3(aq) 
Ni(NO3)2(aq) 
FeCl3(aq) 
81
Name:
Date:
Period:
Colored Precipitates – Net Ionic Equations
Binary ionic compounds exhibit a variety of properties. Some dissolve easily in water, while
others are insoluble. Some are brightly colored compounds; many others are white. You
can use characteristic properties of known substances to determine the identity of an
unknown by comparing the properties of a known substance with those of the unknown.
In this experiment, you will carry out some double-displacement reactions and observe the
characteristic colors of precipitates. You will use your experimental data to identify
unknown substances.
OBJECTIVES
1. Observe displacement reactions in which precipitates are formed.
2. Compare chemical and physical properties of substances.
3. Relate observations to the identification of unknown solutions.
4. Infer a conclusion from experimental data.
MATERIALS
• 0.1 M Co(NO3)2
• 0.5 M CuCl2
• 0.5 M FeCl3
• 0.2 M Ni(NO3)2
•
•
•
0.1 M NaOH
mystery solutions 1 and 2
microtitration plate
1. Always wear safety goggles and a lab apron to protect your eyes and
clothing. If you get a chemical in your eyes, immediately flush the chemical out at
the eyewash station while calling to your teacher. Know the locations of the
emergency lab shower and eyewash station and the procedures for using them.
2. Do not touch any chemicals. If you get a chemical on your skin or clothing, wash
the chemical off at the sink while calling to your teacher. Make sure you carefully
read the labels and follow the precautions on all containers of chemicals that you
use. If there are no precautions stated on the label, ask your teacher what
precautions you should follow. Do not taste any chemicals or items used in the
laboratory. Never return leftovers to their original containers; take only small
amounts to avoid wasting supplies.
3. Call your teacher in the event of a spill. Spills should be cleaned up promptly,
according to your teacher’s directions.
PROCEDURE
1. Using your microtitration plate, place 20 drops of sodium hydroxide, NaOH, in seven
wells in the plate. For best results, all the drops should be the same size.
2. To the first well, add 5 drops of cobalt (II) nitrate solution, Co(NO 3)2. Observe what
happens. Record your observations and the physical properties of the substances
formed in the Data Table.
3. To the second well, add 5 drops of copper (II) chloride solution, CuCl 2. Observe
what happens. Record your observations and the physical properties of the
substances formed in the Data Table.
82
Name:
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Period:
4. To the third, add 5 drops of iron (III) chloride, FeCl3. Observe what happens. Record
your observations and the physical properties of the substances formed in the Data
Table.
5. To the fourth, add 5 drops of nickel (II) nitrate, Ni(NO3)2. Observe what happens.
Record your observations and the physical properties of the substances formed in
the Data Table.
6. To the fifth well, add 5 drops of mystery solution 1. Observe what happens. Record
your observations and the physical properties of the substances formed in the Data
Table.
7. To the sixth well, add 5 drops of mystery solution 2. Observe what happens.
Record your observations and the physical properties of the substances formed in
the Data Table.
DISPOSAL
8. Clean all apparatus and your lab station. Return equipment to its proper place.
9. Dispose of the solutions in the microtitration plate by rinsing them thoroughly down
the drain with a large amount of water. You can let the water run for a few minutes
while you wash out the microtitration plate with the soap solution. Dry the plate to
the best of your ability.
10. Wash your hands thoroughly before you leave the lab and after all work is finished.
DATA TABLE:
Co(NO3)2
Ni(NO3)2
CuCl2
FeCl3
Mystery 1
Mystery 2
Observations:
ANALYSIS & CONCLUSIONS:
The above reactions are all double displacement reactions, in which they were reacted
with sodium hydroxide.
1. Write a balanced double displacement reaction for each. Use the solubility rules to
determine aqueous and solid (precipitate) for your reaction and indicate using the
phase symbols in your equation.
2. Then write the total ionic and the net ionic equation for each reaction.
Reaction 1:
Total Ionic Equation:
Net Ionic Equation:
83
Name:
Date:
Period:
Reaction 2:
Total Ionic Equation:
Net Ionic Equation:
Reaction 3:
Total Ionic Equation:
Net Ionic Equation:
Reaction 4:
Total Ionic Equation:
Net Ionic Equation:
1. Which of the net ionic equations that you wrote describe the reactions of the two
mystery solutions?
a. Mystery Solution 1:
b. Mystery Solution 2:
2. How did you make the determination for each of the mystery solutions?
3. Were the characteristic colors caused by the positive metal ions or the negative ions
present in the compounds?
a. How did you come to this conclusion?
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Factors Affecting Reaction Rate
OVERVIEW
Chemical reactions occur at different
rates. In this experiment you will
consider some of the key factors that
influence the rate of a reaction:




nature of reactants - particle
size
temperature
concentration
catalysts
SAFETY


Portions of this lab may only be
carried out under the supervision
of a teacher
Safety goggles must be worn when
working with acids.
According to the collision theory, the
rate of a reaction depends on the
frequency of collisions between
reacting particles. The more frequent
the collisions, the faster the rate of the
reaction. However, in order for the
collisions to be effective, the particles
must collide with sufficient energy
(activation energy). Furthermore, the
particles must collide with the proper
orientation.
The factors that will be examined in
this lab influence reaction rate by
either increasing how often collisions
occur or by making collisions more
effective.
PURPOSE

To examine factors that increase
reaction rate
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EQUIPMENT AND MATERIALS
PROCEDURES
Part 1. Effect of Particle Size
Part 1. Effect of Particle Size on
Reaction Rate





solid marble chips (calcium
carbonate)
calcium carbonate powder
balance
2 test tubes
1M HCl (approx 10 mL per group)
Part 2. Effect of Temperature



3 Alka Seltzer tablets
3 250-mL beakers
water at three temperatures –
with ice, room temperature,
warm (around 70C)
Part 3. Effect of Concentration





1M HCl, 5 mL per group
3M HCl, 5 mL per group
6M HCl, 5 mL per group
3 pieces of zinc metal, each
approx 1 cm  1 cm
3 test tubes
Part 4. Effect of a Catalyst










3% hydrogen peroxide, H2O2 –
30 mL per group
0.1 M iron(III) nitrate, Fe(NO3)3
0.1 M iron (III) Chloride, FeCl3
0.1 M sodium chloride, NaCl
0.1 M calcium chloride, CaCl2
0.1 M potassium nitrate, KNO3
Solid Manganese (IV) Oxide
100-mL graduated cylinder
10-mL graduated cylinder
6 test tubes per group
1. Obtain a piece of solid calcium
carbonate, approximately 0.5 cm  2
cm. Find the mass of this sample, and
place it in a test tube.
2. Using the balance obtain a sample of
powered calcium carbonate that is
close to the mass of your piece of
solid calcium carbonate. Place this
sample in the second test tube.
3. Place both test tubes in a test tube
rack. Add 5 mL of 1M HCl to both test
tubes. Be sure to wear your safety
goggles.
4. Observe both test tubes and record
your observations in the data table.
Part 2. Effect of Temperature
1. Half fill three 250-mL beakers with
water. In one beaker add several ice
cubes. A second beaker will contain
water at room temperature. In the third
beaker add water that has been
heated to about 70C.
2. Record the water temperature in the
three beakers, and then add an Alka
Seltzer tablet to each.
3. Record the time it takes for the Alka
Seltzer tablet to completely dissolve.
4. If after 10 minutes your alka seltzer
tablet is not dissolved, record “did not
dissolve” in the data table.
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Part 3. Effect of Concentration
1. Pour 5 mL of each of the three HCl
solutions into separate test tubes.
Place the test tubes in a test tube
rack.
4. Mix each tube by swirling the test tube
or gently stirring with a clean stirring
rod.
2. Add one piece of zinc to each test
tube.
5. Observe each solution, noting the
production of any gas bubbles that
form. Record each reaction rate as
fast, slow, very slow, or none in your
data table.
3. Record the time you added the zinc to
the tubes, and the time each reaction
stops. Also record your observations
for each tube.
Part 4. Effect of a Catalyst
In this part of the lab you will determine
which substance/substances act as a
catalyst for the decomposition of
hydrogen peroxide.
RESULTS
Record your results for each part of the
lab in the data tables provided on the
following page.
1. Place 5-mL of 3% H2O2 solution into
each of 6 test tubes.
2. Add 5 drops of each of the following
solutions to separate test tubes:
0.1 M FeCl3
0.1 M NaCl
0.1 M Fe(NO3)3
0.1 M CaCl2
0.1 M KNO3
3. Place several grains of solid
Manganese (IV) oxide into the last test
tube.
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Questions
1. a. What effect does particle size have on the rate of a chemical reaction?
b. Write the reaction you carried out in part 1 and site evidence from your
experimental observation (in your data table) to support your answer to 1a.
2. a. What effect does temperature have on the rate of a chemical reaction?
b. Site evidence from your experimental observation to support your answer to
2a.
3. a. What effect does concentration have on the rate of a chemical reaction?
b. Write the reaction you carried out in part 3 and site evidence from your
experimental observation to support your answer to 3a.
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4. a. Which substances in part 4 acted as catalysts in the reaction?
b. What affect does a catalyst have on the rate of a chemical reaction?
5. List 4 factors that affect the rate of a chemical reaction.
RESULTS
Table 1. Effect of Particle Size on Reaction Rate
Substance Tested
Observations
powdered calcium
carbonate
marble chips
Table 2. Effect of Temperature
Water Condition
Water Temperature (C)
Time to Completion
cold
room temperature
warm
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Table 3. Effect of Concentration
Acid
Concentration
Start Time
Time at
Completion
Observations
1 M HCl
3 M HCl
6 M HCl
Table 4. Effect of a Catalyst
Possible Catalysts
FeCl3
NaCl
Fe(NO3)3
CaCl2
KNO3
MnO2
Reaction
Rate
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MOLE RATIO IN A CHEMICAL REACTION
Procedure:
1. Mass a clean, dry 205mL beaker to the nearest .01 grams. Record the mass.
2. Place about 9 grams of the cupric sulfate hydrate into the beaker. Record mass.
(This is your excess amount).
3. Add approximately 200mL of hot/warm water to the beaker; stir to dissolve the
cupric sulfate. Record your observations.
4. Mass about 1.5 grams of steel wool (iron) to the nearest .01 grams. Record the
mass.
5. Carefully add the steel wool to the cupric sulfate solution. Stir the contents
intermittently until it disappears and another substance takes its place. Record
all changes observed.
6. Set up the filtering apparatus with the funnel, ring stand, and filter paper. Mass
the filter paper. Record mass.
7. Place a 250mL beaker under the funnel to catch the drainage.
8. Allow all of the precipitate to settle to the bottom of the beaker. Decant (drain)
the solution off the precipitate and into the funnel. The liquid that goes through
the filter paper is called the filtrate. The solid particles that remains in the filter
paper is called residue. When all the solution has been decanted off – rinse the
precipitate into the filter paper using the wash bottle and rubber policeman on the
stirring rod. Make sure to get as much of the precipitate into the filter paper
as possible.
9. Remove filter paper from funnel and place on a labeled watch glass. Place
watch glass and filter paper in the fume hood to dry overnight.
10. Day 2 – Mass the filter paper and residue to the nearest .01 grams. Record the
mass.
Data Table:
Mass of dry beaker:
Mass of cupric sulfate
Mass of steel wool:
Mass of filter paper:
added to beaker:
Observations (Pro 3):
Observations (Pro 5):
Mass of filter
paper & residue:
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Questions:
Show all work, equations, and units for full credit. Circle your answers in your
calculations.
1. Write the balance chemical reaction that occurred.
2. Determine the moles of iron used in the experiment (mass-mole).
3. Determine the grams of copper produced (subtraction of data).
4. Determine the moles of copper produced (mass-mole).
5. Based on your calculations, above, did you make ferric or ferrous ions? Give
evidence to support your answer.
6. What visible evidence in the experiment indicated that copper sulfate was in
excess and the iron what the limiting reagent?
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Synthesis of a Compound by Double Displacement
Purpose of Lab:
1) To demonstrate synthesis of a compound.
2) To use stoichiometric relationships to calculate percent yield.
Procedure:
1) Measure out 1 gram of sodium carbonate monohydrate to the nearest .01 grams.
Record the mass in the data table.
2) Dissolve the sodium carbonate monohydrate in 40 ml of water in a beaker.
3) Calculate the amount of calcium nitrate tetrahydrate needed to react with the
mass of sodium carbonate monohydrate you have measured out. Perform
calculation in the data table.
4) Add 10% more of the calcium nitrate tetrahydrate to the calculated amount and
mass it out. Record new mass in data table.
5) Dissolve the calcium nitrate tetrahydrate in 40 ml of water in a beaker.
6) Combine the two aqueous solutions of calcium nitrate and sodium carbonate by
pouring the calcium nitrate into the beaker with the sodium carbonate. Write the
balanced chemical reaction for the double displacement that occurred in the data
table.
7) Set up the filtering apparatus, while the precipitate settles to the bottom of the
beaker.
8) Mass the filter paper to the nearest .01 grams before you place it in the funnel.
Record mass in data table.
9) Decant the liquid off of the precipitate by gently pouring it into the funnel. Be
careful not to fill past the top level of filter paper.
10) Using the rubber “policeman” on the stirring rod decant the rest of the solution
including the precipitate into the funnel.
11) Rinse out the beaker using the wash bottle to transfer the residue into the filter
paper. Repeat as necessary.
12) Remove the filter paper from the funnel and place on a watch glass. Place the
watch glass and filter paper in the fume hood or on your lab top for drying.
13) Clean up all lab materials, and wipe up any water. 
14) When the filter paper is dry, re-mass the filter paper and the product to the
nearest .01 grams. Record the mass on the data table.
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Data Table:
Mass of Sodium Carbonate
monohydrate
Calculate the mass of the calcium nitrate tetrahydrate
needed:
Mass of filter Paper
Mass of filter paper and product
Add 10% to calculated amt of calcium nitrate tetrahydrate:
Double Displacement
Reaction: (2 pts)
Questions/Calculations:
1) Calculate the actual yield of the product. Show work. (2 pts)
2) Calculate the expected yield of the product (calcium carbonate) based on the
mass of the sodium carbonate monohydrate given. Show all work and units. (5
pts)
3) Calculate percent yield. Show all work. (3 pts)
4) Based on your percent yield, list three possible reasons for error. (3 pts)
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MASS-VOLUME & IDEAL GAS LAW CONSTANT LAB
Procedure:
1. Cut a 3-4cm piece of magnesium ribbon and clean it with steel wool. Coil the
magnesium ribbon. Mass the magnesium on the analytical balance to the tenth
of a milligram. (.0001 g) Record your mass in the data table.
2. In the fume hood, measure 5mL of concentrated (12M) hydrochloric acid into a
50mL eudiometer tube. Be careful not to inhale the vapors of the acid. Rinse any
excess acid that spills off the side of the eudiometer tube.
3. At your lab table, fill the tube to the top with room temperature water. Pour the
water from a beaker very slowly so as not to mix the acid and water to any great
extent.
4. Fill a tall (2L) cylinder with room temperature water and place the cylinder in the
sink. Record the temperature of the water in the data table.
5. Tie a piece of thread to the magnesium ribbon, and tie this to a 1-hole stopper,
leaving about 1 -2 inches of string between the stopper and the magnesium.
6. Place the stopper on the tube, so that the magnesium is immersed in the water.
There should be no air trapped in the tube. If there is, remove the stopper and
completely fill it with water so that there is no air pocket. Any air in the tube will
cause a large error in your data.
7. Hold one finger over the hole of the stopper, invert the tube and place it in the
cylinder filled with water. Do not remove your finger from the hole until the
stopper is immersed. Observe the acid as it falls down through the water and
comes in contact with the magnesium. Describe your observations in the data
table.
8. When the reaction is completed record the temperature of the water in the
cylinder. (This temp will be used at the temp of the gas in the eudiometer since it
bubbled through this water). Record the temperature in the data table.
9. Raise or lower the eudiometer tube in the cylinder until the level of the water in
the cylinder is equal to the level of the water in the eudiometer. (The water is
slightly acidic since there was excess acid and the magnesium was the limiting
reagent.) When the levels are equal the pressure of the atmosphere is equal to
the pressure of the gas in the eudiometer tube. Record the volume of the gas in
the eudiometer tube by reading the liquid level in the tube. Note which way the
scale runs on the tube and then record the volume of the gas in the data table.
10. Read the barometer on the wall for the atmospheric pressure and record it in the
data table.
11. Using a water vapor pressure table in your text book, record the water vapor
pressure at the temperature of the gas in the data table.
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Data Table:
Mass of Mg ribbon:
Temp of gas:
Temp of water:
Volume of gas:
Observations of rxn:
Atmospheric pressure:
Water vapor pressure:
Calculations:
1. Calculate the pressure of the “dry” hydrogen gas.
2. Calculate the volume that the “dry” hydrogen gas would occupy at STP.
a.
b.
c.
d.
e.
f.
P1=
V1=
T1=
P2=
V2=
T2=
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3. Using the volume of dry hydrogen calculated in #2, calculate the moles of
hydrogen produced.
4. From the moles of hydrogen produced, and the balanced chemical equation,
calculate the moles of magnesium that would be needed to produce this
hydrogen.
5. Using the moles of magnesium calculated in # 4, calculate this mass of
magnesium.
6. Compare the mass of magnesium you calculated above with the mass of
magnesium you actually used. To do this, calculate percent error using the mass
of magnesium you actually weighed out as the accepted value and the mass of
magnesium you calculated in #5 as the experimental value.
Accepted value - Experimental value x 100
Accepted value
7. List several possible sources of error in this experiment.
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Endothermic and Exothermic Reactions
Many chemical reactions give off energy. Chemical reactions that release energy are
called exothermic reactions. Some chemical reactions absorb energy and are called
endothermic reactions. You will study one exothermic and one endothermic reaction in
this experiment.
In Part I, you will study the reaction between citric acid solution and baking soda. An
equation for the reaction is:
 3 CO2(g) + 3 H2O(l) + Na3C6H5O7(aq)
H3C6H5O7(aq) + 3 NaHCO3(s) 
In Part II, you will study the reaction between water and calcium chloride. An equation
for this reaction is:
 Ca 2+ (aq) + 2Cl- (aq)
CaCl2(s) + H2O(aq) 
Another objective of this experiment is for you to become familiar with Logger Pro, a
program you will use with nearly every experiment in this manual. In this experiment,
you will use Logger Pro to collect and display data as a graph or table, analyze your
experimental data values, and print a graph or data table.
Figure 1
MATERIALS
IBM-compatible computer
Lab Pro Interface
Logger Pro
Stainless Steel Temperature Probe
50-mL graduated cylinder
balance
Styrofoam cup
250-mL beaker
Citric Acid, H3C6H5O7, solution
Baking Soda, NaHCO3
Water, H2O
Calcium Chloride, CaCl2
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PROCEDURE
1. Obtain and wear goggles.
Part I Citric Acid plus Baking Soda
2. Place a Styrofoam cup into a 250-mL beaker as shown in Figure 1. Measure out 30
mL of citric acid solution into the Styrofoam cup. Place a temperature probe into the
citric acid solution.
3. Prepare the computer for data collection by opening “Exp 01” from the Chemistry
with Computers experiment files of Logger Pro. The vertical axis has temperature
scaled from -10 to 40°C. The horizontal axis has time scaled from 0 to 300 seconds.
4. Weigh out 10.0 g of solid baking soda on a piece of weighing paper.
5. The temperature probe must be in the citric acid solution for at least 45 seconds
before this step. Begin data collection by clicking Collect . After about 20 seconds
have elapsed, add the baking soda to the citric acid solution. Gently stir the solution
with the temperature probe to ensure good mixing. Collect data until a minimum
temperature has been reached and temperature readings begin to increase. You
can click on Stop to end data collection or let the computer automatically end it
after 300 seconds.
6. Dispose of the reaction products as directed by your teacher.
7. To analyze and print your data:
• Click the Statistics button, . In the statistics box that appears on the graph,
several statistical values are displayed for Temp 1, including minimum and
maximum. In your data table, record the maximum as the initial temperature and
the minimum as the final temperature. Close the statistics box by clicking the
upper-right corner of the box.
• To confirm the minimum and maximum temperatures, use the scroll bars in the
Table window to scroll through the table to examine the data. Compare the
minimum and maximum data points to those you recorded in the previous step.
• Print a copy of the Table window. Enter your name(s) and the number of copies of
the table.
• You will often want to change the scale of either axis of the graph. There are
several ways to do this. To scale the temperature axis from 0 to 25°C instead of
the present scaling, click the mouse on the “40” tickmark at the top of the axis. In
place of the “40”, type in “25” and press the Enter key. Notice that the entire axis
readjusts to the change you made. Use the same method to change the “-10”
tickmark to “0”. Note: A second option is to click the Autoscale button, . The
computer will automatically rescale the axes for you.
• You can also expand any portion of the graph by zooming in on it. Select the area
you want to zoom in on. Do this by moving the mouse pointer to the beginning of
this section of data—press the mouse button and hold it down as you drag across
the curve, leaving a rectangle. Then click the Zoom In button, . The computer
will now create a new, full-size graph that includes just the region inside the
rectangle. You can reverse this action by clicking the Undo Zoom button, .
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• When you again collect data in Part II of this experiment, the data will be collected
as Latest run, the most recent set of data you have collected. The original Latest
run will be lost if it is not saved or stored. Choose Store Latest Run from the Data
menu to store Latest as Run 1, then save or print it later. Note that the line for Run
1 is thinner than it was for Latest. To hide the curve of your first data run, click the
Temperature vertical-axis label of the graph, and uncheck Run 1. Click OK .
Part II Water Plus Calcium Chloride
8. Manually rescale the vertical axis to the original temperature scale of -10 to 40°C. To
do so, click the mouse on the bottom tickmark and type in “-10”. Then click on the
top tickmark and type in “40”.
9. Measure out 30 mL of water into the Styrofoam cup. Place the temperature probe
into the water.
10. Obtain 10 grams of calcium chloride from the teacher.
11. Note: The temperature probe must be in the water for at least 45 seconds before this
step. Begin data collection by clicking Collect . After about 20 seconds have elapsed,
add the CaCl2 to the water. Gently stir the solution with the temperature probe to
ensure good mixing. Caution: Do not breathe the vapors. Collect data until a
maximum temperature has been reached and the temperature readings begin to
decrease.
12. Dispose of the reaction products as directed by your teacher.
13. To analyze your Part II data:
• Change the appearance of the graph by double-clicking anywhere on the graph
bring up the Graph Options dialog. Check the box in front of Point Protector Every
1 Point—a point protector will now outline each data point on the graph. Click
OK
.
• Instead of scrolling through the Table window in this trial, click the Examine button,
. The cursor will become a vertical line. As you move the mouse pointer across
the screen, the temperature and time values corresponding to its position will be
displayed in the box at the upper-left corner of the graph. Scroll across the initial 34 points to determine the initial temperature. Record the initial temperature in the
data table. Move the mouse pointer across the peak of the temperature curve to
determine the maximum temperature, and record it as the final temperature in your
data table. To remove the examine box, click the upper-right corner of the box.
• It is also possible to calculate statistics just for a portion of your collected data. To
do so, you must first select the data you are interested in. For example, you might
want to find the average (or mean) of the first few data points to use as an initial
temperature, instead of using the minimum value. Select the flat portion of the
curve—move the mouse pointer to time 0 and drag across the flat part of the
curve. Now click the Statistics button, , and note the mean temperature value in
the statistics box on the graph. This value is the mean of only the selected data
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points. When you are done, click on the upper-right corner of the statistics box to
remove it.
14. To print a graph of temperature vs. time showing both data runs:
• Click the Temperature vertical-axis label of the graph. To display both temperature
runs, check the Run 1 and Latest boxes. Click OK .
• Label both curves by choosing Make Annotation from the Analyze menu, and
typing “Endothermic” (or “Exothermic”) in the edit box. Then drag each box to a
position near its respective curve.
• Print a copy of the Graph window. Enter your name(s) and the number of copies of
the graph you want.
15. Save the temperature and time data from both data runs. Choose Save As from the
File menu and give the file a distinct name. As directed by your teacher, choose a
location for the file, and click OK .
DATA AND CALCULATIONS
Part I
Part II
Final temperature, t2
________°C
________°C
Initial temperature, t1
________°C
________°C
Temperature change,
________°C
________°C
OBSERVATIONS
Processing the Data
1. Enter the values for T in the data table.
2. Tell which reaction is exothermic. Explain.
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3. Which reaction had a negative Δt value? Is the reaction endothermic or exothermic?
Explain.
4. For each reaction, describe three ways you could tell a chemical reaction was taking
place.
1._____________________________________________________________
2._____________________________________________________________
3._____________________________________________________________
5. Which reaction took place at a greater rate? Explain your answer.
6. A. Use the equation Qrxn = -(mc  T) surr to calculate the heat of reaction of Reaction
1 (citric acid and sodium bicarbonate reaction.) Consider the solution in the cup to be
the surroundings. Assume the density and specific heat of the citric acid solution are
the same as that for water.
6B. Use Qrxn to calculate  Hrxn for reaction 1. (Use the mass of NaHCO3 you massed
out, and the Molar Mass of NaHCO3 To convert from Q rxn to  Hrxn.)
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6C. Write the thermochemical equation for reaction 1. Use correct formulas and
subscripts, and place H rxn in the correct place.
7A. Use the equation Qrxn = -(mc  T) surr to calculate the heat of reaction of Reaction 2
(dissolving Calcium Chloride):
7B. Use Qrxn to calculate  H for reaction 2. (Use the mass of CaCl2 you massed out,
and the Molar Mass of CaCl2 ).
7C. Write the thermochemical equation for reaction 2. Use correct formulas and
subscripts, and place H in the correct place.
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