Download HONORS LAB MANUAL - Tenafly High School

CHEMISTRY HONORS
LABORATORY MANUAL
2009/2010
1
2
TABLE OF CONTENTS
NUMBER TITLE
Laboratory Format and Policy
1
Safety Lab
2
Percent Composition of a Mechanical
Mixture
3
Comparing Physical and Chemical Changes
4
Alchemy
5
Determining an Empirical
6
Formula Composition of Hydrates
7
Determining the Thickness of Aluminum Foil
8
Types of Chemical Reactions
9
Relating Moles to Coefficients in an Equation
10
Particle Size from Collision Probabilities
11
Flame Tests
12
Investigation on the Hydrogen Spectrum
13
Determination of the Half-life of Ba-137
14
Periodic Table
15
Preparation of Oxygen
16
Model Construction
17
Paper Chromatography
18
Determination of Absolute Zero
19
Molar Volume of a Gas
20
Construction of an Air bag
21
A study of Phase Change Using Lauric Acid
22
Application of Solubility Rules
23
The 6-Solution Puzzle
24
Rate of a Chemical Reaction
25
A Study of Equilibrium
26
Change in Enthalpy of a Reaction
27
Determination of Hydrogen Ion
Concentration
28
Acid-Base Titration
29
Decomposition of Baking Soda
30
The Ice Cream Lab
3
PAGE
5
7
9
11
13
17
19
21
23
27
29
31
33
35
37
39
43
45
47
49
51
53
55
57
59
61
63
65
67
69
71
DATA PAGES (To be printed and brought to class)
1
3
5
8
9
10
11
12
13
14
15
16
17
19
20
21
22
23
24
25
26
27
Safety Lab
Comparing Physical and Chemical Changes
Determining an Empirical
Types of Chemical Reactions
Relating Moles to Coefficients in an Equation
Particle Size from Collision Probabilities
Flame Tests
Investigation on the Hydrogen Spectrum
Determination of the Half-life of Ba-137
Periodic Table
Preparation of Oxygen
Model Construction
Paper Chromatography
Molar Volume of a Gas
Construction of an Air Bag
Application of Solubility Rules
The 6-Solution Puzzle
Rate of a Chemical Reaction
A Study of Equilibrium
Change in Enthalpy of a Reaction
Determination of Hydrogen Ion Concentration
Acid-Base Titration
4
71
75
77
79
81
83
85
87
89
91
93
95
97
99
101
103
105
107
109
111
113
115
LABORATORY POLICY AND REPORT FORMAT
INTRODUCTION:
Since each laboratory activity is different the requirements for the report will vary.
Follow the directions as given for each lab in the manual. Lab activities are often done in
pairs or groups. Depending on the lab activity, reports will be written individually or by
pairs or group.
PRELAB
The lab prep will be given the day before the lab, whenever possible. The first part of the
lab report will then be written before you start the lab and will be checked while you are
doing the activity. This should be typed. Your pre-lab should include the following:
NAME
DATE
GRADING RUBRIC
TITLE: Use the title from the manual.
PURPOSE: This can be obtained from the prelab discussion and from the
introduction.
ANSWERS TO PRELAB QUESTIONS (When required)
PROCEDURE: For some labs you will be required to write a description for the
procedure that you will use. These should be in the form of a list.
DATA TABLE: For some labs you will be required to generate you own table; for
many others the table is contained in the Data Section of the manual and must be
printed and brought to class.
CALCULATIONS: Some activities require some pre-calculations. These may be
hand written in ink.
If you are absent the day before a lab, it is your responsibility to find out what lab
will be done and write your pre-lab, since it is included as part of the grade. If you
do not have the prelab done at the beginning of the lab, the grade on that part is
ZERO.
THE LABORATORY REPORT
The report must be typed and will be due the following day at the beginning of class.
When labs were done as partners or groups, everyone must submit a completed report,
unless indicated otherwise. The reports should be attached and handed in as a packet.
One will be graded using the Grading Points. Grades for members of the team may vary
because of the prelab, lab behavior, organization, etc.
DATA/OBSERVATION:
All data taken during the experiment must be written in ink. Nothing is to be written on
paper and transferred to the tables. If you make an error or repeat the procedure, draw a
single line through the data and write the corrected date next to it. There should be no
erasing or white-out. I will initial the data at the end of the lab.
CALCULATIONS: Show all set-ups using proper units and significant digits. Each
calculation should appear on a separate line. Percent error calculations should be shown
here.
5
GRAPHS: Graphs must have a title and have both axes labeled. It may be computer
generated or done on graph paper and stapled or taped to the report.
CONCLUSION: Some labs will require a conclusion, while in others the conclusion will
be covered by the questions. It should be in paragraph style and in formal English (as if
you were writing for your English class). Keep in mind that the conclusion section
communicates whether you have accomplished the PURPOSE. The paragraph should
be brief. The following information should be included:
1. YOUR FINDINGS; REFER TO YOUR DATA FOR SUPPORT. THE
CONCLUSION ALWAYS REFERS BACK TO THE PURPOSE. If there is an
unknown, the number of the unknown must be included in the identification.
2. A discussion of errors; how they could have been avoided; what could be
improved in the lab procedure to reduce the errors.
3. For some labs, the scientific theory on which the lab is based; how the
theory was applied in this lab.
QUESTIONS: Number the questions, answer in complete sentences, and skip a line
between questions. Reports may require both conclusions and questions or only one or
the other.
GRADING
Laboratory reports will be graded on a point system. The grade will be based on the
following:
1.Pre-lab.
2. During the laboratory activity:
Experimental technique
Safety
Clean-up
3. Organization – neatness, format, completeness.
4. Data and results.
5. Calculations.
6. Conclusion and answers to questions.
Graded lab reports must be saved in your binders. They will be collected as a packet
at the end of each quarter and graded for completeness. Be prepared for lab quizzes
and/or questions pertaining to labs on tests.
Points will be deducted for reports handed in after the due date.
6
1. SAFETY LAB
INTRODUCTION
One of the most important considerations during any laboratory activity is safety. Labs
provide a valuable learning experience (and can be fun as well) as long as common sense
and safe behavior are practiced.
This activity will help to remind you of safe lab practices that you will use all year.
PROCEDURE
1. Work in groups. Write the names of the group members. Each member will complete
the lab Write-up and one will be selected for grading.
2. At each station, write the question as it appears. Proceed to the next available station.
Be sure that the station number corresponds with the number on the report page.
3. List the important materials at the station, make observations, answer the question,
and/or carry out the activity where an activity is required.
4. Clean up materials as directed.
CONCLUSION
Write AT LEAST one safety rule associated with each station at the end of the report
page.
7
2. PERCENT COMPOSITION OF A MECHANICAL
MIXTURE
INTRODUCTION
In a mixture whether homogeneous or heterogeneous, the components may be present in
any proportions. In order to determine the percent composition, the components must be
separated by some means, usually physical. In a mixture of components A and B, the
percent by mass can be determined by the relationship:
% A in mixture AB = mass of component A
Mass of mixture AB
PURPOSE
To determine the % composition of a mixture of sand (silicon dioxide) and table salt
(sodium chloride).
MATERIALS
Sample of mixture, any equipment needed to carry out procedure
PROCEDURE
1. Working in teams of two, design an experiment you will use. Include equipment you
will need, a data table and safety precautions.
2. Submit it for approval.
3. Carry out the experiment, recording data on the prepared data table.
4. Perform the necessary calculations
CONCLUSION
Write a brief conclusion which should include the results of your experiment and an error
discussion..
GRADING
Procedure
Safety/clean-up
Calculations
Organization
8
4
4
4
20
9
3. COMPARING PHYSICAL AND CHEMICAL CHANGES
Changes in which the physical or chemical properties of a substance are altered can be
considered physical or chemical changes, respectively. Some chemical and physical
changes are so closely related that it is difficult to tell the difference between the two.
The test of a physical or chemical change is this: Does the change alter the nature of the
substance: is something new formed? If the shape, size or physical state is altered but
the chemical composition remains the same, the change is a physical change. In a
chemical change or a chemical reaction, the atoms of a substance are rearranged. A
chemical change requires that the new substance have a chemical composition that is
different from the composition of the original substance. Usually physical changes also
occur.
In this laboratory activity you will carry out seven procedures record observations and
determine in each case whether a chemical or physical change is brought about.
GOGGLES MUST BE WORN. EXERCISE CAUTION TO PREVENT BURNS.
PRELAB
1. Copy the Grading.
2 State the purpose of the lab in your own words.
3 List five observable indications of a physical change
4 List two observable indications that a chemical change has occurred
5. Complete the first column on the data page.
PROCDURES
1. Demonstration. Break several toothpicks into small pieces and put the pieces in a
large test tube. Heat the tube strongly over a laboratory burner for several minutes.
Dispose in the solid waste container.
2. Mass a piece of copper wire. Record the mass. Place it in a small test tube and add
silver nitrate until it is just covered. Place the test tube in a rack and let it stand until
fifteen minutes before the end of the lab since this experiment will take time to obtain any
measurable change. Meanwhile, go on to the experiment 2 and return to the next step at
the END OF THE LAB.*
* To be done at the END OF THE LAB. Remove the wire with tweezers and scrape and
wash off any material sticking to the wire. Dry it with a paper towel and mass it.
3. Place A few grams of sodium chloride (tip of the spatula) in a small test tube. Add
enough water to dissolve the salt. Stir.
4. Place an evaporating dish on a wire mesh. Pour the salt solution into the
evaporating dish. Heat carefully until all the water evaporates.
11
5. Remove the evaporating dish using tongs to a ceramic tile. Let it cool for a few
minutes then dissolve the crystals in a little water. Add silver nitrate several drops at
a time until a change is observed. Pour the material in the waste vessel in the hood.
Wash the dish for a later experiment.
6. a) Place about 1 g (tip of the spatula) of cobalt(II) chloride hydrate, CoCl2 . 6H2O, in
the evaporating dish. Heat the solid slowly at first, then more strongly.
b) Allow the dish to cool then add a few drops of water.
7. a) Place about 25 mL of water in a 100- mL beaker. Add a level teaspoon (use plastic
spoon) of copper (II) chloride dihydrate, CuCl2 . 2H2O to the beaker. Stir until the
crystals disappear. Measure and record the temperature of the solution. Use the
thermometer clamp and leave the thermometer in the solution.
b) After CuCl 2 has dissolved, add a piece of Aluminum to the solution. Make sure
the solution covers the aluminum. Continue to observe the mixture and measure
change in temperature after five minutes. *While you are waiting, go back to
procedure 2 to complete that experiment.
QUESTIONS
1. List three physical properties you observed in the lab. Be specific.
2. Explain the difference between qualitative and quantitative data.
3. Answer the following question in chart form.
 Which experiments were qualitative and which were quantitative?
 Which changes are physical and which are chemical?
Experiment # qualitative/quantitative
physical/chemical
CONCLUSION
Explain the difference between a physical and chemical change. Use two examples of
physical changes and two of chemical changes from this lab to illustrate the difference.
GRADING
Prelab
Safety/cleanup
Observation table
Organization
Questions
Conclusion
3
2
4
2
12
4
27
12
4. ALCHEMY
Copper to Silver to Gold
INTRODUCTION: Many of the processes and substances that we know today were
discovered by alchemists. They discovered alcohol, hydrogen, phosphorus, and gun
powder as well as the processes of distillation, evaporation, and filtration.
Alchemy was a mixture of science, medicine, magic, and religion. One of the main goals
of alchemy was to change a lesser metal into gold. Producing gold was thought by the
alchemists to be a major step toward everlasting life.
Now let’s take an imaginary trip back to the Middle Ages when knights were in fashion,
fair ladies were to be saved, dragons were to be slain (well maybe there weren’t any
dragons), and alchemists were at work. The King has just called you for advice. The
local alchemist has devised a way to make copper into silver and then into gold. Your
instructions are to perform the experiment, test the gold, and give the King your advice.
Does he reward the alchemist or hang him as a cheat?
The alchemists’ formulas are:
13
Now it’s up to you. Remember the King is very gracious in his rewards for good work
and harsh in his punishments for wrong answers.
THESE ARE THE INSTRUCTIONS FROM THE ALCHEMISTS NOTE BOOK.
You will follow them to transform the copper. Once it is transformed you will
devise a procedure to determine the validity of the alchemist’s claim.
GOGGLES AND APRONS MUST BE WORN. Hair tied back.
NaOH will cause burns on the skin. Do not inhale fumes.
MATERIALS: Tong, evaporating dish, beaker, ring stand ,burner, wire gauze, beaker
with water
6M NaOH, Zinc dust, copper pennies
ALCHEMIST’S PROCEDURE
1. Clean a penny with an eraser.
2. Place a small spoonful of zinc in an evaporating dish. Cover the zinc with the
solution of 6M NaOH and heat just until some steam is visible. Move the burner
away. DO NOT ALLOW IT TO BOIL. (FUMES WILL IRRITATE YOUR NOSE
AND THROAT.)
3. Immerse the penny in the solution. When bubbles of gas can be seen to escape, the
penny should have changed appearance. Use tongs to remove the penny and place it
in a beaker of water.
4. Remove it when cool and blot it carefully with a paper towel.
5. .Hold the penny on the edge with the tongs and heat it in the outer cone of a burner
flame. The second change will be quite sudden. Remove it as soon as a change occurs
and again, place it in a beaker of water. Overheating will destroy the penny.
6. After it cools, show it to the instructor. Carry out your procedure to determine if
the penny is really gold.
7. WITH YOUR GOGGLES ON, dispose of all materials in the waste container in
the hood.
PRELAB SHOULD INCLUDE:
PROCEDURE*
Describe the procedure you will use to test whether the penny was “transmuted”
into gold. The alchemist’s procedure should NOT be included.
OBSERVATION AND DATA
What data will you collect to verify that the penny was changed to gold? Make data table
based on your procedure.
14
LABORTORY REPORT SHOULD INCLUDE:
PROCEDURE (modified if necessary)
DATA TABLE
CALCULATIONS
CONCLUSION
Did the alchemist cheat the king? The king’s decision will be based on the outcome of
your experiment. Use your data and calculations to support your conclusion.
QUESTION
What is a possible explanation for the changes to the penny?
* It is possible that you may have to modify or change the procedure you use. If you do,
then you must write the new procedure as part of your final report.
GRADING
Prelab
Lab safety/clean-up
Procedure
Organization
Calculations
Conclusion
Question
4
4
4
4
2
5
2
25
15
5. DETERMINING AN EMPIRICAL FORMULA
In a sample of a compound, regardless of the size of the sample, the number of moles of
one element in the sample divided by the number of moles of another element in the
sample will form a small whole-number ratio. These ratios can be used to determine the
subscript in the empirical formula of the compound.
In this experiment a pure metal, magnesium, will react with oxygen in the air to form a
binary compound. From the masses before and after the reaction, you will calculate the
moles of each element and find the empirical formula
GOGGLES MUST BE WORN. USE TONGS TO HANDLE CRUCIBLE AND
COVER. DO NOT PLACE METAL DIRECTLY IN THE FLAME.
PRELAB QUESTIONS
1. What are the reactants?
2. What product is formed?
3. How will the mass of the contents in the crucible change after heating? Explain.
4. Prepare a data table. Use the sample at the end as an example.
PROCEDURE
1. Clean a crucible and cover. Dry them by heating in the hottest part of a burner flame
for a few minutes. Allow them to cool. Mass and record.
2. Clean a 35-cm length of magnesium ribbon and cut it into small pieces. Place the
pieces in the crucible. Mass the crucible, cover and contents.
3. Cover the crucible and place it in a clay triangle. Heat for a few minutes. Using the
tongs to carefully tilt the cover and check that the magnesium is glowing. Leave the
cover in this position to provide an opening for air to enter the crucible. Heat for about
10 minutes or until all the magnesium has reacted (turned to powder). Make sure that no
magnesium remains.
4. Use the tongs to remove the crucible from the stand to a tile. Add a few drops of water
slowly, then continue to add water to just cover the contents of in the crucible. Observe
any odor. Record.
5. Holding the burner in your hand, gently heat the contents of the crucible by moving
the burner slowly back and forth. Avoid splattering by partially covering the crucible.
Observe an odor.
6. When the liquid has boiled off, repeat steps 4 and 5.
7. When all the liquid has boiled off a second time, strongly heat the uncovered crucible
for about 5 minutes to make sure the compound is completely dry.
17
9. Turn off the burner and allow the crucible and contents to cool. Measure the mass of
the crucible, cover and contents.
CALCULATIONS
Show all work.
Include a % error in your calculations.
QUESTIONS
1. In addition to the magnesium oxide, a small amount of another compound of
magnesium forms during the heating in air. In step 5, water was added to the contents
in the crucible to convert this compound into magnesium oxide. The odor indicated
that ammonia was also formed in the reaction with water. What was the formula of
the other magnesium compound? Answer by writing and balancing the equation.
___?______ + water 
ammonia
+ magnesium oxide
2. There are three main sources of error in this lab. How would the ratio of Mg:O be
different than 1:1 if:
a) some of the product formed was lost in heating at the end?
b) all the Mg had not completely reacted?
c) all the water had not been completely evaporated?
CONCLUSION
What was your ratio? Discuss the error in YOUR experiment.
GRADING
Prelab
Safety/cleanup
Organization
Calculations
Conclusion
Questions
Results
3
2
2
4
3
7
1
22
18
6. COMPOSITION OF HYDRATES
Hydrates are ionic compounds (salts) that have a definite amount of water (water of
hydration) as part of their structure. The water is chemically combined with the salt in a
definite ratio. Ratios vary in different hydrates, but are specific for any given hydrate. The
formula of a hydrate is represented in a special manner. The hydrate of Cobalt (II) chloride
(which you worked with in a previous lab) has the formula CoCl2 . 6H2O. The unit formula
for the salt appears first, and the water formula is last. The dot means that water is loosely
bonded to the salt. The coefficient 6 stands for the number of molecules of water bonded to
one unit of salt. This special formula illustrates the law of definite composition.
When hydrates are heated, the chemically bonded "water of hydration" is released as vapor.
The remaining solid is known as the anhydrous salt. The general reaction is
heat
hydrate

anhydrous salt + water
The percent of water in a hydrate can be found experimentally by accurately determining the
mass of the hydrate and the mass of the anhydrous salt. The difference in mass is due to the
water lost by the hydrate. The percentage of water in the original hydrate can easily be
calculated.
In this investigation you will determine the % of water in an unknown hydrate and identify
the hydrate from a given list of hydrates.
PRELAB ASSIGNMENT
1. STATE A PURPOSE
2. ANSWER THE QUESTION
How will the mass of the material in the evaporating dish compare before and after
heating?
3. WRITE A PROCEDURE
You will be given a sample of unknown hydrate. Describe in detail the procedure
you will use to determine the % water in the hydrate to identify the unknown from
the list below.
4. LIST MATERIALS NEEDED
5. PREPARE A DATA TABLE
Make a NEAT data table based on your procedure
6. PERFORM THE CALCULATIONS TO DETERMINE THE THEORETICAL
% WATER IN EACH OF THE POSSIBLE UNKNOWN HYDRATES:
Cobalt (II) chloride hexahydrate
Copper (II) sulfate pentahydrate
Strontium chromate monohydrate
Calcium chloride dihydrate
Barium chloride dihydrate
Magnesium sulfate heptahydrate
Copper(II) chloride dihydrate
19
GOGGLES AND APRONS MUST BE WORN. BE CAUTIOUS IN HANDLING
HOT GLASSWARE
POST LAB
CALCULATIONS
Show calculations for your unknown including a % error.
CONCLUSION
Identify your hydrate
Support your identification using data
Briefly discuss the principle used to accomplish the lab
CONCLUSION QUESTIONS (error analysis)
1. If all the water of hydration was not released, how would your % of water
compare to the accepted value? Explain.
2. What could cause the opposite results (from the above question) to occur?
3. When water is added back to the anhydrous salt is the reaction endothermic or
exothermic? Explain.
GRADING
Prelab
Procedure
Safety/cleanup
Organization
Calculations
Identification
Conclusion
Questions
5
3
2
3
3
2
4
6
28
20
7. DETERMINING THE THICKNESS OF ALUMINUM
FOIL
This activity will be done cooperatively. All members must be involved and one grade
will be given to the group. (If however, I observe that a member of the group is not
participating, that person will receive a lower grade.)
OBJECTIVE: To determine the thickness of a piece of aluminum foil to three
significant digits.
To determine the thickness in atoms.
To determine the number of atoms of aluminum in the piece of foil.
SOME DIRECTIONS:
1. Write the names of the group members on the top of the card.
2. Decide on a procedure.
3. You may use any book or piece of equipment that is presently in the room.
SOME THINGS THAT YOU ARE NOT PERMITTED TO DO:
1. You may not destroy, crumble, tear or fold your piece of foil. It must be
returned in exactly the same condition that you received it.
2. You can only confer with members of your group. You are not permitted
to spy or eavesdrop on other groups.
3. You cannot remove any papers from this lab. All materials must be turned
in by the end of the period.
WHAT IS TO BE SUBMITTED BY THE END OF THE PERIOD:
1. The aluminum foil
2. The Report with
a) Procedure
b) Data table
c) Calculations with answers boxed.
d) Grading
Grading
Procedure
Calculations
Organization
2
8
4
21
8. TYPES OF CHEMICAL REACTIONS
Have you ever thought about the fact that aluminum, unlike iron doesn't seem to rust? Actually,
aluminum does react with the oxygen in the air to form, the white, powdery compound aluminum
oxide, Al2O3. The aluminum oxide remains on the aluminum as a protective coating, and no
further reaction occurs.
Have you ever wondered why hydrogen peroxide (H2O2) is sold in brown bottles? Hydrogen
peroxide decomposes gradually into water and oxygen, but light hastens the reaction.
In 1776 Henry Cavendish reacted hydrochloric acid with zinc to produce hydrogen gas. Zinc
chloride remained behind in the vessel he used.
Hydrochloric acid reacts with sodium hydroxide to produce table salt and water.
Once you have learned to write a balanced equation with its correct reactants, products, formulas,
and coefficients, you have crossed one of the major hurdles or chemistry.
In this laboratory you will carry out six different reactions and then identify each type of reaction
and write a balanced equation for it.
GOGGLES AND APRONS MUST BE WORN.
PRELAB
1. Read the introduction. Write balanced equations for each of the reactions
described in the introduction above and identify the type of reaction
2. Complete the first column on the Table by listing the reactants for each
experiment.
MATERIALS: right-angle glass bend with stopper and rubber tubing, burner, large
beaker, watch glass, large test tubes small test tubes, test tube holder, ring stand, solid
stoppers, splints forceps, tongs, litmus paper, test tube rack, evaporating dish.
PROCEDURE
1. Record all observations.
2. Carry out the following reactions.
REACTION A
1. Place the copper(II) carbonate in a small test tube to a height of about 1cm.
2. Note the appearance of the sample.
3. Using the test tube holder heat the tube strongly until a change occurs. Immediately
hold a burning splint in the mouth of the test tube. Observe the effect on the splint
and changes in the appearance of the residue in the test tube.
4. Discard contents of tube in a paper towel and dispose into the waste basket.
23
REACTION B
1. Add ammonium carbonate to a small test tube . Heat as in REACTION A.
2. Test with a burning splint again. Then remove the test tube from flame and hold a
piece of moistened litmus paper in the mouth of the test tube. Identify any odor
that is readily apparent by wafting the fumes. DO NOT SNIFF THE TEST
TUBE DIRCTELY.
3. Discard any remaining material down the sink, flushing with water.
REACTION C
1. Place a small amount of calcium carbonate in a dry test tube.
2. Add about 20 drops of hydrochloric acid to the test tube.
3. As soon as the reaction is observed, hold a burning splint in the mouth of the test
tube.
4. Discard the contents down the sink with running water.
REACTION D
1. Obtain a piece of magnesium ribbon, tongs and a watch glass.
2. Using the tongs, place the ribbon in the flame. Remove it as soon as begins to burn
and hold it over the watch glass. DO NOT LOOK DIRECTLY AT THE
MAGNESIUM.
3. Examine the residue.
REACTION E
1. Place about a mL of sodium chloride solution into the very small test tube
provided.
2. Add silver nitrate solution to the sodium chloride a few drops at a time.
3. Pour contents in the beaker in the hood labeled "waste".
REACTION F
1. Add about 20 drops of sodium hydroxide solution into a small test tube.
2. Add one drop of phenolphthalein and mix by gently swirling the tube.
3. Add hydrochloric acidone drop at a time to the test tube. Count the number of
drops of acid required fro a permanent color change.
REACTION G (See fig)
1. Place 10 ml of hydrochloric acid in a large test tube. Insert the one-hole
stopper with the glass tube into the test tube. CAUTION: HCl CAN CAUSE
SEVERE BURNS.
2. Clamp on a ring stand.
3. Fill a small test tube with tap water, and invert the tube in trough filled with
enough tap water, so your hand can fit under the test tube. Stand the inverted test
tube in the trough.
4. Obtain some pieces of mossy zinc. Open the test tube containing the HCl and
carefully drop in the zinc. Record your observation.
5. Allow the collection tube to fill with gas. Collect another if gas is still being
produced.
6. Stopper the tubes of gas with solid stoppers and remove from the water. Keep
them inverted.
24
REACTION H
1. Dry the tubes on the outside. Remove the stopper from one of the test tubes and
carefully bring a burning splint to the mouth of the gas-filled tube. Notice any
change in the tube.
2. Record observation.
3. Carefully dispose of the acid in the beaker labeled "waste" in the hood. BE SURE
TO KEEP GOGGLES ON DURING CLEANUP.
DEMONSTRATION
REACTION I
1. Place about I mL ethanol in a clean evaporating dish. Place the dish on a tile.
2. Fill a test tube about one third full with cold tap water. Set aside.
3. Ignite the alcohol.
4. Hold the test tube with water (in a test tube clamp) above the burning alcohol.
Observe the outside of the test tube for evidence of product formation.
5. Allow the alcohol to burn until it is completely consumed.
ANALYSIS
1. Study the observations you recorded and then identify the type of reaction that took
place in each case.
2. Write balanced equations for each reaction.
QUESTIONS
1. What gases were generated in this lab?
2. What tests were used to help identify the gases in each experiment? (4 pts)
2. Name a characteristic they have in common. What characteristic distinguishes them
from each other? (4 pt)
3 Identify the reactions that were redox (oxidation/reduction). (3 pts)
4. What was The Hindenburg and what is significant about it? Relate the event
concerning the Hindenburg to a reaction in this lab? Write the balanced chemical
equation. (4 pt)
GRADING
Prelab
Equations
Table
Safety/cleanup
Equations
Questions
Organization
2
8
2
8
15
2
27
25
26
9. RELATING MOLES TO COEFFICIENTS OF A
CHEMICAL EQUATION
Coefficients in a chemical equation indicate the number of moles of each substance;
therefore, the ratio of moles of a substance to moles of any other substance in the reaction
can be determined at a glance
In this experiment, iron filings will be added to an aqueous solution of copper (II) sulfate
and a reaction will take place. The copper produced will be separated from the solution,
dried and massed. From the known masses of reactants and products, the moles of each
can be calculated and the equation written.
You will relate the ratio of the moles to the coefficients of the balanced equation. Using
the balanced equation you will calculate the theoretical yield and from the measured data,
the % yield of copper.
GOGGLES AND APRONS MUST BE WORN.
PRELAB CALCULATION
1. Calculate how many grams of the hydrate of copper sulfate (pentahydrate) must be
massed to give 8 grams of anhydrous copper sulfate. This reactant must be in excess to
ensure that all the iron reacts.
PROCEDURE
1. Obtain a 100-mL beaker. Clean it to remove any material clinging to the walls. (The
beakers are used only for this experiment and there may be stain that is not
removable). Dry the beaker, label it and mass it.
2. Measure the proper amount of the copper sulfate pentahydrate (amount you
calculated) to give about 8 grams of copper sulfate and add to the beaker. Record
the exact amount.
3. Mass about 2.24 grams of iron filings on a piece of paper. Record the exact mass.
4. Measure about 50 mL distilled water and add to the beaker.
5. Heat the beaker using a burner to just below boiling, stirring the mixture until the
crystals completely dissolve. DO NOT ALLOW THE LIQUID TO BOIL.
6. Using beaker tongs remove the hot beaker and place on a tile. Add the iron filings,
SMALL AMOUNTS AT A TIME TO AVOID A SUDDEN BURST OF
VAPOR, to the hot solution. Stir continuously. After all the iron has been added
and mixture stirred, allow the beaker to sit for at least 10 minutes or until the
reaction is complete.
7. Decant the liquid into a waste beaker as demonstrated. Try not to disturb the solid at
the bottom of the beaker.
27
8. Add about 10 mL of water to the solid remaining in the breaker. Stir vigorously to
wash the solid. Let it settle and decant the liquid into the waste beaker. Repeat
the washing at least 6 times to ensure that all that only copper and some water
remain in the beaker.
9. Spread the solid out on the bottom of the beaker and place it on the tray provided to
dry over night.
10. When it is dry, find the mass of the beaker and copper metal. Scrape out contents in
waste basket. Return the beaker to container as directed.
CALCULATIONS
1. Write the two possible balanced equations.
2. Calculate moles of reactants and copper.
3. Calculate the theoretical yield of copper and the % yield.
QUESTIONS
1. How do calculated moles of reactants and products compare with the
coefficients of the balanced equation? Refer to your calculations.
2. Would this reaction occur with copper metal and iron sulfate as the reactants?
3. How do you know whether iron (II) or iron (III) sulfate forms? Use your data
and calculations to explain.
4. Discuss the factors that caused YOUR yield to be either too high or too low?
GRADING
Prelab
Safety/clean-up
Organization
Calculations
Questions
2
2
3
6
8
22
28
10. PARTICLE SIZE FROM COLLISION PROBABILITIES
There are situations where it is necessary to use indirect means to gather information or
make measurements. For example, scientists use indirect means to study galaxies, the
size of the universe, and the atom.
In this lab activity, you will determine the size of a marble by counting how often you
succeed in hitting one when another marble is shot at a group of them. From the data
collected you will apply a formula to calculate the size. You will then actually measure
the marble and compare the experimental with the actual size.
PRELAB
Write the purpose.
PROCEDURE
1. Obtain ten target marbles and one marble with which to bombard them.
Arrange the target marbles inside a three sided enclosure (walls) as
demonstrated. The target marbles may be placed in any random position BUT
a) there must always be room for the bombarding marble to get through
between any two target marbles and
b) no marble should be shielded from being hit; that is each target marble
must be a potential target.
2. Roll the bombarding marble toward the target marbles. Attempt to roll the
bombarding marble along a path perpendicular to the target marbles, parallel to
the walls of the enclosure. For good results, the bombarding marble must be
rolled randomly. DO NOT AIM. It may help to close your eyes.
3. Continue rolling the bombarding marble toward the target marble, counting
both the total number of rolls and the number of hits. A “hit” occurs when the
bombarding marble hits another marble before hitting the wall. If more than
one marble is hit during a roll, it is counted as one hit. Results improve with
the number of rolls: no fewer than 300.
4. Line up your ten target marbles so that they touch each other in a straight line
along a meter stick and measure the length.
CALCULATIONS
D
=
Hxd
2N x Tr
1. Calculate the diameter of the marbles from the experimental data.
2. Find the actual diameter of one marble from the measurement of ten.
3. Calculate the percent error. Assume the measured diameter to be the actual.
29
QUESTIONS
1. Describe Rutherford’s experiment. What did he conclude about atomic
structure?
2. How is this experiment a model of Rutherford’s experiment?
3. Besides the obvious, in what ways does this model of Rutherford’s experiment
differ from the actual experiment?
GRADING
Prelab
Safety/cleanup
Organization
Calculations
Questions
Results
2
2
2
2
7
3
30
11. FLAME TESTS AND EMISSION SPECTROSCOPY
All atoms give off electromagnetic radiation if their gases or ions are energized by heating or by
high voltage electric discharge. If the light emitted by a gas is passed through a spectroscope, a
pattern of narrow lines of light is produced. Each element produces its own distinct pattern that
differs from the pattern of every other element. The particular pattern of frequencies of light
emitted by an atom is referred to as its emission spectrum or bright-line spectra. The emission
spectrum of an element can be used as a means of identification , just as fingerprints (or DNA)
can be used to identify a human being.
The unique pattern of frequencies of light emitted by an atom corresponds to the set energy
given off as electrons drop from higher allowed energy states or energy levels to lower energy
levels. The energy of each transition is given by Max Planck’s equation, E = hv,
where E = energy in Joules, v = frequency, and h = a constant (Planck’s constant). An electron
can be raised from its lowest allowed energy level (ground state) to other higher allowed energy
levels by absorbing certain set amounts of energy. This “excited” electron cannot remain at any
of these higher allowed energy levels if there are unoccupied lower energy levels closer to the
nucleus. The electron is attracted back to one of the lower allowed energy levels. As the
electron drops back it emits a photon, a set amount of energy, in the form of electromagnetic
radiation. The amount of energy that the electron emits is equal to the energy difference between
the higher energy level of the excited state and the lower energy level to which the electron
dropped.
In Part I flame tests will be used to demonstrate the emission spectrum of a metal. Using this
method, a small amount of a metallic salt is heated. Each element gives off a unique set of
wavelengths, which our eyes recognize as a particular color. You will perform a flame test on
several known metallic salts, and record the flame color for each metal ion. You will then
perform a flame test on an unknown metallic salt. By comparing the color of the unknown to the
known samples, you will be able to identify the metal ion in the unknown. You will use cobalt
glass as a tool for reducing the “masking” effect of one element upon another element in a
mixture.
In Part II, you will use a spectroscope to examine continuous spectra from light sources and the
bright-line spectra of some gaseous elements and attempt to identify one of the sources by
comparing the observed spectrum to the spectrum chart.
PRELAB
1. State the purpose of this lab activity.
2. What is meant by ground state?
3. How can electrons become “excited”?
4. When is the energy absorbed by electrons released?
5. What is the form of this energy?
6. State the equation that is used to determine the energy content of a packet of light of
specific frequency.
7. How should the burner flame be adjusted for best results.
31
GOGGLES MUST BE WORN.
Part I
PROCEDURE
1. Obtain a set of wooden splints, each saturated with a solution of one metallic salt.
Take two splints with sodium, potassium, and the mixture of sodium and potassium.
2. Adjust the Bunsen burner to a nonluminous flame. Hold a splint in the flame. Record
the color. Repeat with all the splints. Reserve the duplicate splints for the test with
cobalt glass.
3. Obtain a sample of unknown and repeat the flame test.
4. Observe the flame produced with sodium, potassium, and the mixture through the
cobalt glass.
Part II
PROCEDURE
Using the diffraction slides, view the discharge tubes located in the room. Record the
bands of light observed. Compare the observed bands to the chart showing the known
spectra of several elements and identify the tube that contains neon.
WARNING: DO NOT TOUCH POWER SUPPLIES OR TUBES.
QUESTIONS
1. Give two reasons why a flame test might not always be valid for identifying an
element.
2. Explain the reason potassium was visible in the mixture only when using the cobalt
glass.
3. Based on the observations made in this lab, is it possible to identify EACH element in
a mixture. Explain.
4. What is the difference between an absorption and an emission spectrum? Which is
observed from the discharge tubes?
5. What would be observed if a spectroscope were used during a flame test?
6. Which tube contained the element to be identified?
CONCLUSION:
Identify your unknown and support your conclusion with your data.
GRADING
Prelab
Safety/cleanup
Questions
Conclusion
Identification
Organization
7
2
12
3
1
2
32
12. INVESTIGATION ON THE HYDROGEN SPECTRUM
Bohr’s model of the atom assumed that the single electron in the hydrogen atom could be
located in certain location (an orbit) and would remain there unless excited by energy
from outside the atom. Since each orbit has a particular energy associated with it, Bohr’s
restriction meant that energies associated with electron motion in the permitted orbits are
fixed in value; that is they are quantized. Bohr knew that the correct model had to
account for the experimental spectrum of hydrogen. The emission of radiation by an
excited hydrogen atom could then be explained in terms of the electron dropping from a
higher energy orbit to a lower one and giving up a quantum of energy (a photon) in the
form of light. These changes in energy levels are called transitions.
The energies that the electron in the hydrogen atom can possess can be calculated by the
equation:
En = -RH(1/n2)
where n = principal quantum number (an integer from 1….7) and RH is a constant,
called the Rydberg constant equal to 2.18 x 10-18. The negative value is a convention; it
says that the energy of the electron in the atom is lower than the energy of a free electron,
which is assigned a value of 0. So as the electron gets closer to the nucleus (as n
decreases), En becomes more negative. When n = 1, this is the ground state for the
hydrogen atom.
The difference in the energy levels then is
ΔE = Ei - Ef
where Ei and Ef represent the energies of the initial energy level and the final energy
level respectively during a transition.
In this activity, you will
1. Calculate the energies of the possible energy levels (orbits) of the hydrogen
atom shown on the diagram.
2. Measure the wavelengths of the lines on the hydrogen spectrum.
3. Calculate the energy of the photon for each line.
4. Determine which energy level transition was responsible for each of the three
lines.
33
13. DETERMINING THE HALF-LIFE OF Ba-137m
(A class experiment)
INTRODUCTION
The half-life of a radioactive isotope (radioisotope) is the time required for the activity
(or mass) of a sample to be reduced to one half of its original activity (or mass). Today,
the class will determine the half-life of the radioisotope Ba-137m. The m stands for
“metastable”, a condition of temporary stability. A Ba-137 generator will be used as the
source of the Ba-137m. The generator contains Cs-137, which decays by beta emission
to Ba-137m with a half-life of over 30 years. The Ba-137m that is produced possesses
more energy than is normally possessed by a stable nucleus and “cools” (decays) to stable
Ba-137 by gamma emission with a half-life that you will determine.
PRELAB QUESTIONS
1. What is the purpose of this activity?
2. If you are taking a 5 minute background reading, how do you calculate the counts per
minute (cpm) ?
3. If your readings are taken for 30 seconds, how do you calculate the cpm (gross)?
4. How do you calculate the cpm (net) from the cpm (gross)?
MATERIALS
A Cs-137/Ba-137m “minigenerator, Geiger tube and sample holder, and Scalar or Geiger
counter
PROCEDURE
1. The Geiger counter will be turned on by your teacher and allowed to run without a
sample for 5 minutes so that a background radiation count can be taken. Record the
counts.
2. Your teacher will “milk” the minigenerator to wash out a 7 or 8 drop sample of Ba137m into the small metal dish. A Geiger tube will be placed over the sample in a
sample holder.
3. The scalar will be turned on for 30 seconds and then stopped. You will quickly
record the counts. The Counter will be set to zero and restarted 30 seconds after it was
turned off. In this manner you will collect several 30 second readings with a 30 second
“rest” between each reading. This will continue for about 10 minutes.
4. Fill out your data sheet as the experiment proceeds.
5. Construct a graph of the data. Your cpm (net) should be plotted on the y-axis and the
time (either in seconds or minutes) on the x-axis. Use the beginning of each interval as
your time. (For example: for the interval 0-30 sec, plot your activity versus 0 seconds.)
Remember to draw the best fit smooth curve.
35
6. Use your graph to calculate the half-life of Ba-137m. Choose any convenient cpm
(net) for your first activity and find the time that corresponds to that activity. Then take
half of that first activity (this is your second activity) and find the time that corresponds
to that. Repeat using half of the second activity. Be sure to do at least two trials and
average the results. Show all your work on the graph paper.
7. Calculate a % error on the graph.
CONCLUSION QUESTIONS
1. Based on your graph, what is the half-life of Ba-137m? How does it compare to the
actual value?
2. Some Cs-137 is always mixed into the Ba-137m sample. Does this affect the
measurement of the half-life of the Ba-137m sample? Explain why or why not.
3. Why is it safe to throw the remaining Ba-137 down the drain after 10 or 15 minutes?
GRADING:
Prelab
Graph
Calculations
Questions
Organization
4
5
4
6
2
36
14. PERIODIC TABLE
A STUDY OF REACTIVITY OF METALS
The relative reactivity of an element can be predicted by its position on the periodic table.
In this activity, you will observe reactions of several metals in Group I, II and III with
water and determine the trends in reactivity within a Group and across a period.
PRELAB: Write the purpose of the experiment.
GOGGLES AND APRONS MUST BE WORN FOR THESE EXPERIMENTS.
PART I
MATERIALS
test tubes, test tube rack, calcium piece, magnesium ribbon, aluminum pellet,
phenolphthalein solution, splint, match, forceps, Bunsen burner, test tube holder
PROCEDURE
1. The reactivity of Alkali metals will be observed by demonstration.
2. Pour about 5 mL of distilled water into a clean, dry test tube and place it in the
test tube rack. Using the forceps, add a piece of calcium to the water. DO NOT
TOUCH THE CALCIUM WITH YOUR HANDS. IT IS CORROSIVE TO
SKIN. Observe the reaction. Collect the gas being released by holding an
inverted dry test tube over the reactant tube.
3. Test for hydrogen gas by inserting a burning splint into the mouth of the upper
test tube.
4. Add a drop of phenolphthalein solution to the reactant tube. Record your
observations. (Phenolphthalein indicates the presence of a base.)
5. Break the magnesium ribbon into small pieces and place the pieces in a test
tube. Repeat the above procedure.
6. If there is no visible reaction, gently heat the water in the test tube just to
boiling. BE SURE TO POINT THE TUBE AWAY FROM YOURSELF AND
OTHERS WHILE HEATING.
7. Once it is boiling, turn off the burner, return the test tube to the rack and
carefully observe if a gas is being formed. (Look on the surface of the metal
pieces for the presence of small bubbles.) Collect the gas and test for hydrogen.
Add phenolphthalein. Note any color change in the test tube.
8. Repeat the procedure with one aluminum pellet.
37
PART II
The more active a metal, the more reactions will be observed. Formation of a precipitate
will indicate that a reaction has occurred.
SALTS OF BARIUM AND STRONTIUM ARE EXTREMELY TOXIC. AVOID
CONTACT WITH THESE CHEMICALS AND WASH YOUR HANDS
THOROUGHLY AFTER USE,
MATERIALS
spot plate, dropper pipets containing solutions of Mg(NO3)2, Ca(NO3)2, Sr(NO3)2,
Ba(NO3)2, H2SO4, Na2CO3, Na2CrO4.
PROCEDURE
1. To each well of a spot plate place about 10 drops of one of the following
solutions containing alkaline earth metals: Mg(NO3)2, Ca(NO3)2, Sr(NO3)2,
Ba(NO3)2. To each solution add a few drops of H2SO4.
Record whether a precipitate appears, the relative amount and color.
If no precipitate is formed leave the box blank.
2. Repeat the procedure but add a few drops Na2CO3(aq) to each solution.
3. Repeat the procedure again but add Na2CrO4(aq) to each solution.
EQUATIONS
1. Write balanced equations for the reactions of the Alkaline earth metals and water
(Part I).
2. Write balanced equations for the precipitation reactions that occurred in Part II
CONCLUSION QUESTIONS
1. What trend in reactivity was observed within metals in Group I? Group II?
2. Predict what you would have observed if the following elements had been tested:
Beryllium, Strontium, Francium
3. What trend was observed among the three metals tested in Period 3?
4. If zinc and iron had been tested, how would their reactivity compare to the metals in the same
period that you observed in this lab?
5. Describe any relationship that you can determine between the number of precipitates formed by
each compound and the location of the alkaline earth metal on the periodic table.
GRADING:
Prelab
1
Safety/cleanup 2
Tables
4
Equations
3
Questions
10
38
15. PREPARATION AND PROPERTIES OF OXYGEN
Oxygen is the most abundant element in nature. While oxygen is 21% of the atmosphere, it is
not readily isolated from the atmosphere under laboratory conditions. It can, however be
readily generated in the laboratory from oxygen containing compounds. There are several
methods which are generally used to prepare oxygen: the pyrolysis of a chlorate, the
electrolysis of water, the pyrolysis of a heavy metal binary oxide such as mercuric oxide, and
the decomposition of hydrogen peroxide using enzymes. Pyrolysis is the destructive
decomposition of a compound by heat. All methods are useful in that they are quantitative.
The method to be used in this experiment is the pyrolysis of potassium chlorate. A catalyst,
MnO2 will be used.
This experiment will involve both quantitative and qualitative data. You will prepare oxygen,
collect it by water displacement and determine the % yield and the % oxygen in potassium
chlorate. You will then study some physical and chemical properties of the oxygen you have
collected. You will write equations for all the reactions.
PRELAB
1. Write the purpose.
2. What is the purpose of a catalyst? What will be used to catalyze this reaction?
3. Prepare Data tables for Part I and Part II.
CAUTION: POTASSIUM CHLORATE IS EXPLOSIVE. ALL DIRECTIONS MUST
BE CAREFULLY FOLLOWED. GOGGLES AND APRONS MUST BE WORN
THROUGHOUT THIS EXPERIMENT.
NOTE; You will work in pairs to prepare oxygen. To complete the reactions in Part II, you
will need 4 bottles of oxygen. Some of you will not collect this much, therefore it may be
necessary to combine groups for Part II.
PROCEDURE
PART I: THE PREPARATION AND COLLECTION OF OXYGEN GAS
1. Mass a clean dry large test tube.
2. Add approximately three grams of potassium chlorate. TRY NOT TO GET
ANY ON THE SIDES OF THE TEST TUBE. WIPE ANY FROM
AROUND THE TOP. Mass the test tube and solid and record.
3. Add approximately 1.5 g of MnO2. Mass the test tube and the contents again
and record.
4. Break up any lumps and mix the contents THOROUGHLY.
39
5. Connect the test tube to an apparatus for the collection of a gas by water
displacement. HAVE YOUR SET-UP APPROVED BEFORE YOU
BEGIN THE NEXT STEP.
6. Heat gently holding the burner in your hand. Control the heating. Do not
allow the chemicals to rise in the test tube. If the reaction appears to be going
too fast, slow the heating.
7. Do not collect the first bubbles that appear, since this will contain mostly air.
Collect as many bottles as you can. Cover the bottles with glass plates.
8. When the bubbling stops, remove heat and IMMEDIATELY DISCONNECT
THE TEST TUBE to prevent water from backing up into the apparatus.
9. Place the test tube in a safe place out of the way to be massed when cool.
Meanwhile proceed with Part II.
PART II: PROPERTIES OF OXYGEN
Record observations about the physical properties. Record all
observations of the following reactions.
REACTION I
1. Place a small amount of charcoal into a deflagrating spoon. Heat until red
hot and quickly lower into a bottle of oxygen. DO NOT ALLOW THE
SPOON TO TOUCH THE SIDES OR BOTTON OT THE GLASS. Keep
the bottle covered as much as possible.
2. When the reaction stops, remove the spoon and discard the ash in the waste
pail. Keep the jar covered, add about 10 mL of water and shake. Test the
water with litmus paper.
REACTION 2
3. Heat a small piece of magnesium in a deflagrating spoon until it begins to
glow. Lower it into another bottle following the directions above. Do not
look directly at the jar.
4. When the reaction stops remove the spoon. Allow it to cool then add a
small amount of water to the SPOON to dissolve the material. Test the
solution in the spoon with litmus.
40
REACTION 3
5. Take a small amount of steel wool in the spoon. Heat until glowing and
again lower into the bottle as in step 2. (No other tests need to be done
with the steel wool).
REACTION 4
6. Place a glowing splint into the last bottle and see how many re-lights you
can get. Record.
CALCULATIONS
1. Calculate the theoretical yield of oxygen from the amount of potassium
chlorate used.
2. Calculate the actual amount of oxygen produced using your data.
3. Calculate the % yield.
CONCLUSION
The conclusion should relate back to the purpose. There should be two brief
paragraphs, one for Part I which involved quantitative data, and one for Part II
which was qualitative. In Part II use the observations to describe physical and
chemical properties of oxygen.
QUESTIONS
1. Write the 6 chemical equations for reactions that occurred in this
laboratory.
2.
What does litmus paper indicate? Specifically in this experiment what
information did you obtain with litmus paper?
3.
Identify the following reactions as exothermic or endothermic:
a) pyrolysis of KClO3
b) reaction of carbon and oxygen
c) reaction of magnesium and oxygen
GRADING
Prelab
Safety/cleanup
Organization
Calculations
Conclusion
Questions
5
2
2
3
4
8
24
41
16. MODEL CONSTRUCTION: POLARITY AND SHAPE
The most common type of bond between two atoms is a covalent bond. A covalent bond
is formed when two atoms share a pair of electrons. If both atoms have the same
electronegativity the BOND is nonpolar covalent. When atoms have different
electronegativities, the electrons are attracted to the atom with the higher
electronegativity. The BOND that forms is polar covalent.
MOLECULES made up of covalently bonded atoms may themselves be polar or
nonpolar. If the polar bonds are symmetrical around the central atom, the bonds offsest
each other and the molecule is nonpolar. If the polar bonds are not symmetrical, the
electrons will be pulled to one end of the molecule and the molecule will be polar.
Many physical properties of matter are the result of the shape and polarity of molecule.
Water, for example, has unusual properties that can only be explained by the shape of its
molecule and the distribution of charge on the molecule.
In this lab activity, you will build models of molecules, name their shape, predict their
polarity and the type of hybridization.
PRELAB
Prepare a data table for the 20 models using the headings listed below.
.
formula electron dot bond type molecule type sketch shape hybridization
structure
(p/np)
(p/np)
PROCEDURE
1. Complete the table for the following compounds
1. H2
6.CO2
11. C2H6
2. HCl
7. NH3
12. C2H4
3. O2
8. CH4
13. C2H2
4. N2
9. CH3Cl
14. HClO
5. H2O
10. CCl4
15. HCN
16. H2CO
17. HCOOH
18. H2O2
19. PO4320. NO3-
2.a) Construct a model of pentane and draw the structural formula.
b) Construct isomers of pentane and draw structural formulas for each.
3. Construct 2 models of pentene and draw structural formulas.
4. Construct 2 models of dichloropentane and draw structural formulas.
5. Construct a model of bromochlorofluoromethane. Draw the structural formula. Construct an
isomer of this compound. (Hint: Is your left hand identical to your right?)
43
QUESTIONS
1. Classify each of the following using one of the following: ionic bonding, polar covalent in
polar molecules, polar covalent in nonpolar molecules, nonpolar covalent in nonpolar
molecules: (4 pts)
a) Cl2
b) CH2Cl2
c) CCl4
d) KF
2.
a) What are isomers?
b) How many isomers were possible for pentane?
c) What type of isomerism is shown with the various models of
pentane
pentene
dichloropentane
bromochlorofluoromethane ? (6 pts)
3. Both water and carbon dioxide are triatomic molecules with 2 bonding pairs. What factor
explains the difference in their shapes? (2pts)
4. Account for CHCl3 being polar while CCl4 is not. (1 pt)
5. Paradichlorobenzene and sodium chloride are both white crystals at room temperature.
Paradichlorobenzene melts at 53C while sodium chloride melts at 1253C. Explain this
difference in melting point. (2pts)
6. Although methane, water and ammonia have approximately the same molecular mass, their
melting points and boiling points differ greatly. How you think molecular shape and polarity can
affect boiling and melting points. (2pts)
7. The polarity of a substance can have a great effect on its solubility. A rough rule of thumb is
“like dissolves like”. Using this general rule, what can you predict about the polarity of ethyl
alcohol if you know that alcohol dissolves in water? (1 pt)
GRADING
Prelab
Chart
Questions
2
5
18
44
17. PAPER CHROMATOGRAPHY
INTRODUCTION:
Chromatography is a technique for separating and identifying components in a mixture. The name
of the method derives from the fact that it was first applied to the separation of colored substances
(Greek: chromos, color). Although the effectiveness of the separation is most easily detected if the
components of the mixture are colored, the method is applied today to all kinds of mixtures. All
types of chromatography employ two different immiscible phases in contact with each other. One
of the phases is moving, the mobile phase, and the other non-moving phase is the stationary phase.
In Paper Chromatography, a small amount of mixture to be separated is placed near the bottom
edge of the paper. This end of the paper strip is then placed into the solvent, which then travels
up the paper by capillary action. Separation occurs because different chemicals in the mixture
travel different distances up the paper due to their attraction to the paper and solubility in the
solvent. Those substances that are not tightly adsorbed on the paper will move at the same rate as
the solvent. Substances that are bound tightly to the paper will not move at all. Those
substances that are weakly attracted to the paper will move, but more slowly than the solvent.
When the solvent has moved near the end of the paper, the paper is removed from the solvent
and dried. Once developed, the paper called a chromatogram, will contain different chemicals
located at different positions on the paper.
This ratio of the distance each component of the mixture moves with respect to the distance that
the solvent moves can be calculated as follows:
Rf = Ds
Df
where Ds = distance traveled by a spot, and Df = distance traveled by the solvent (the solvent
front).
The Rf value is characteristic of a substance under a specified set of conditions, and can sometimes
be used in the identification of substances in an unknown mixture when compared to known pure
substances or used to detect the same components in different mixtures.
MATERIALS:
Chromatography paper, Marker pens, Pencil and ruler, Cup
PROCEDURE:
1. Plac2 a small amount of water (to a depth of about 1 cm) in a plastic cup.
2. Draw a line with pencil about 2 cm from the edge of the paper. This marks the origin.
Spot markers on the line as directed.
3. Place the paper (spot down) in the cup containing the solvent. Make sure the origin is
above the level of water.
4. Observe, and when the solution has nearly reached the top, remove the paper and draw
a line across the paper to mark the distance the solvent has traveled.
45
5. Measure the distance traveled by solvent for each chromatogram (Df). Measure the
distances traveled by each spot (Ds) on each chromatogram (measure from the origin to
the center of the spot) and calculate the Rf for each spot.
6. Calculate Rf values for each spot.
7. Share the data with other others at your table in order to answer question 5.
CONCLUSION QUESTIONS
1. What were the mobile phase and the stationary phase in this lab ?
2. What causes the separation to occur?
3. Which ink(s) appear(s) to be composed of only one compound? Explain.
4. Which ink(s) appear(s)to be made up of more than one compound? Explain.
5. Do any of the markers appear to contain the same inks? Refer to Rf values.
GRADING
Chromatograms
Calculations
Format
Questions
Clean-up
1
2
1
14
1
46
18. DETERMINATION OF ABSOLUTE ZERO
The relationship between temperature and volume is a direct proportion and can be
shown by the equation of a straight line in which the independent variable is the
temperature and the dependent variable is the volume. When the volume of a fixed
amount of gas at a given temperature is cooled, the volume is decreased. If the volumes
and temperatures at both conditions are accurately measured, these can be plotted. The
straight line that results can be extrapolated to find the point where the line would cross
the x-axis. This point should correspond to theoretical absolute zero. Since the value of
y at any point on the x axis is zero, the volume at absolute zero will be zero if the gas has
the properties of an ideal gas.
PRELAB QUESTIONS
1. State the purpose.
2. What variables are being measured?
3. What gas is used in this lab to find absolute zero?
4. What is an ideal gas?
5. Prepare a data table.
1. Set up a boiling water bath using the 1-L beaker. While waiting for the water to boil prepare
a cold water bath in the Styrofoam container with ice and cold water. The temperature should
be less than 10C or as close to 0 as possible.
2. Tightly stopper a DRY 250-mL flask with a one- hole rubber stopper and submerge it in the
hot water bath so the air in the flask is completely submerged in the water. The water should
not go above the bottom of the rubber stopper to ensure that no water gets into the flask.
3. As soon as the water appears to boil, measure the temperature of the water (T1). Do not
allow the thermometer to touch the bottom of the beaker.
4. Leave the flask in the boiling water for at least one minute, then remove it and quickly invert
it in the ice bath. Push the flask down so it is completely submerged under the ice water.
Allow it to remain there until the flask has reached the temperature of the water. Measure
the temperature of the ice bath when the flask is cold (T2)
5. Remove the flask by placing your finger over the hole so no more water can enter or leave
the flask. Measure the volume of water drawn into the flask (Vd).
6. To obtain the volume of the gas when it completely occupied the flask at the higher
temperature, completely fill the flask with water then insert the rubber stopper.
Measure this volume, (V1).
7. Return the large beaker, stoppers and Styrofoam container to the side of the room.
47
CALCULATIONS
1. Find V2 by calculating the difference betweenV1 and Vd.
2. Calculate the theoretical V2 using your data and the Charles’ Law equation.
3. Calculate % error using the theoretical V2 as the accepted value.
4. Prepare a graph. The temperature values should range from -300C to 100C.
5. Find absolute zero (x-intercept) using the equation derived from your graph.
6.
Calculate the % error for the experimental value of absolute zero.
QUESTIONS
1. Explain how this experiment illustrates Charles’ Law?
2. Explain why water was drawn into the flask at the lower temperature?
3. The experiment assumes that at absolute zero the volume is zero. Is that correct? Explain.
4. Would you expect different results if another gas were used in the flask?
CONCLUSION
Write a brief conclusion. Remember to refer to the purpose of this lab.
48
19. MOLAR VOLUME OF A GAS
The molar volume of any gas is defined as that volume occupied by one mole of gas at standard
temperature and pressure (STP). In this laboratory activity, you will be experimentally finding
the molar volume of a gas. Hydrogen gas will be used, since it is easily generated and collected.
Hydrogen gas can be generated when an active metal reacts with acid solution. The moles of
hydrogen can be obtained from the mass of magnesium reacting with an excess of hydrochloric
acid. The hydrogen gas will be collected by water-displacement in a eudiometer, a gas collection
tube which is calibrated in mL.
To determine the molar volume, the data must be corrected for laboratory conditions.
1. The collected hydrogen gas will be saturated with water vapor; this pressure of water must be
subtracted from the total pressure to obtain the pressure of hydrogen alone. The pressure of the
water vapor at the temperature of the reaction can be obtained from a chart. 2. The volume of
Hydrogen collected at lab conditions must be converted to STP. 3. The molar volume then is
simply the ratio of the volume at STP to the number of moles of hydrogen produced.
PRELAB QUESTIONS
1. Write the equation for the reaction used to generate the gas in the eudiometer.
2. What two gases will be in the eudiometer?
3. Define STP.
4. What information is obtained from a barometer? Why is that needed?
5. What information must be obtained from the chart? How will that information be used?
HCl IS CORROSIVE. GOGGLES AND APRONS MUST BE WORN.
PROCEDURE
1. Carefully mass the strip of magnesium ribbon. Roll the magnesium into a coil and secure
it to the copper wire on the stopper.
2. Set up collecting apparatus as demonstrated, using de-ionized water
3. Insert the rubber stopper with the magnesium into the tube again making sure that there
are no air bubbles. Place your finger over the hole at the bottom of the stopper and invert
the eudiometer into the beaker. Once the stopper is under the water, remove your finger.
There is no danger from the HCl, since it is at the other end of the tube. You will be able to
observe the HCl migrating down the tube and mixing with the water.
4. After the metal has reacted completely, allow the system five minutes to reach
equilibrium.
5. While waiting, record the barometric pressure and the vapor pressure of water at room
temperature using the chart.
49
6. Gently tap the bottom of the tube on the bottom of the beaker to dislodge any bubbles
remaining near the wire. As directed, transfer the eudiometer to one of the long containers
to read the volume at atmospheric pressure. Measure and record the temperature of the
water in this container.. Assume that the collected gas is at this temperature.
7. Repeat the procedure for an additional trial.
CALCULATIONS
Calculations 1-5 will be done twice, using data from each trial.
Number the calculations.
1. From your data calculate the moles of magnesium used.
2. Using the equation you wrote for the reaction, calculate the moles of hydrogen
produced.
3. Use the vapor pressure of water to calculate the pressure of dry hydrogen gas.
4. Correct the volume* to STP conditions.
5. Calculate molar volume of hydrogen using the ratio of corrected volume to.
moles of hydrogen calculated above.
6. Average the molar volumes obtained from both trials.
7. Calculate the % error using the average.
*You can convert to Liters at any point.
CONCLUSION QUESTIONS
1. Consider how the following conditions would affect the results in terms of the volume
of hydrogen collected and the calculated molar volume.
a) the magnesium was not pure.(assume the other materials are inert)
b) some hydrogen dissolved in the collecting water.
c) some air bubbles remained in the eudiometer after water was added
2. Which reactant is limiting?
CONCLUSION
Your conclusion should state what you found and what principles (gas laws) were applied.
Discuss errors and how they affected your results. (Be specific depending on whether your
results were higher or lower than accepted value.)
GRADING
Prelab
5
Safety/cleanup 2
Calculations 7
Questions
4
Conclusion 5
Organization
50
LAB 20.
51
52
21. A STUDY OF A PHASE CHANGE USING LAURIC ACID
INTRODUCTION
Matter can exist in three different physical states- gas, liquid, or solid. In a pure substance,
changes of physical state take place at discrete temperatures, which are constant and
characteristic for each substance. The temperature at which a substance changes state from solid
to liquid at standard pressure is called the normal melting point of that substance. The change
back to the solid state occurs at the same temperature and is called the normal freezing point.
In Part I you will examine what happens when a pure substance undergoes a change in physical
state, specifically the freezing behavior of an organic compound called lauric acid (C12H24O2 )
also known as dodecanoic acid and determine the freezing point.
In Part II you will explore the energy change when lauric acid freezes, by determining the heat
released to water in a calorimeter. You will then calculate the heat of fusion (Hf ) of the
substance and compare it to the literature value.
PRELAB
1. Write a purpose for Part I and Part II.
2. Prepare a Data Table for Part II
3. Define calorimetry.
4. Define heat of fusion .
5. What is the literature value for the Hf of lauric acid.
PROCEDURE
GOGGLES AND APRONS MUST BE WORN DURING THE EXPERIMENT.
EXERCISE CAUTION IN HANDLING THE THERMOMETER ANDTHE HOT
PLATE.
PART I DETERMINATION OF FREEZING/MELTING POINT (Do with a partner)
A FREEZING POINT
1. Make a data table of time and temperature.
2. Obtain a test tube containing Lauric acid. Remove the stopper.
3. Melt the lauric acid in a beaker of water that has been heated on a hot plate. There
must be enough water so that the part of the test tube containing lauric acid is under
the water. KEEP THE TEMPERATURE BELOW 60C. DO NOT ALLOW
WATER TO CONTAMINATE THE LAURIC ACID. When the lauric acid has
completely melted, place a dry thermometer in the test tube and remove the test tube
from the hot water. Make sure the thermometer is in the center of the lauric acid.
5. One partner should stir and read the temperature while the other partner
times and records. Take temperature readings every 30 seconds until the
temperature of the lauric acid falls below 40C.
53
PART II DETERMINATION OF HEAT OF FUSION USING CALORIMETRY (Do
individually)
1. Place enough water in the calorimeter to cover the lauric acid when the test tube is
placed in the calorimeter (at least 100 mL of water). Measure the volume of water
carefully. Record the temperature of the water (Ti).
2. Melt the lauric acid and remove it from the water bath when it just melted. DO
NOT ALLOW IT TO GO ABOVE 50C. Remove the thermometer and clean it with
a paper towel. Let the lauric acid cool in the air while carefully observing for the
appearance of crystals. As soon as the crystals appear immediately place the test tube
in the calorimeter. Gently swirl. Record the highest temperature the water reaches. (
Tf ).
3. Dry the test tube and mass it. The number written on it is the mass of the empty test
tube. Record this.
4. Replace the stopper and return the test tube with lauric acid to the side of the room.
CALCULATIONS
1.Use data from Part I to Plot a graph of the cooling curve.
2. Calculate the heat of fusion in J/g of lauric acid assuming
Heat lostlauric acid = heat gained water
3. Calculate % error for Hf.
CONCLUSION
The conclusion should contain 2 paragraphs, one for Part I and one for Part II. Each
should include a statement of your results, a BRIEF discussion of the theory and how it
was used in this lab. Discuss sources of error in your experiment Part II only.
GRADING
Prelab
5
Clean-up/safety 2
Graph
4
Calculations 4
Conclusion
8
Organization 2
54
22. APPLICATION OF SOLUBILITY RULES
Solubility can be thought of as the tendency of a substance (the solute) to dissolve in another
substance (the solvent). For qualitative purposes, such terms as "soluble", "insoluble", and
"slightly soluble" can be used to describe these tendencies.
Ionic compounds (salts and bases) dissolve in water by a process known as dissociation. In this
process, the crystal lattice of the solid breaks down, and free ions move throughout the solution.
The total number of positive charges is equal to the total number of negative charges in the
solution.
If aqueous solutions of two different ionic compounds are mixed, one of two things will occur. If
all of the ions remain free no precipitate will form and the solution will remain clear. However, if
two oppositely charged ions are attracted to each other strongly enough, they will bond together
to form an ionic compound that is insoluble in water. In such a case, a precipitate forms.
Using the solubility rules, you will predict whether aqueous solutions will form precipitates and
then carry out the experiment to confirm your predictions. Equations will be written for those
reactions which form precipitates.
PRELAB
On the data table in ink, using the solubility rules predict whether or not the pairs of
solutions will produce a precipitate when mixed. Assume all are double replacement
reactions. Enter yes if a precipitate should form or no if both products are soluble.
PROCEDURE
WEAR GOGGLES DURING THE EXPERIMENT. WASH YOUR HANDS WHEN
FINISHED.
1. Obtain a set of solutions.
2. Add a few drops of the first reagent down the column. Repeat down the other columns.
3. Add a few drops of the first reagent across the row. DO NOT LET THE TIP OF THE
PIPET TOUCH THE DROP OR FALSE DATA WILL RESULT
4. Use a different color pen to record yes if there is a reaction and no if there is no visible
precipitate. (If it is weak, indicate this).
5. a) Write double replacement equations for assigned reaction that occurred. Label the
precipitate.
b) Write net ionic equations for the reactions above.
55
CONCLUSION QUESTIONS
1. Use your data to answer the following: (4 pts)
a) Which anions were always spectator ions.
b) List the anions with the number of precipitates each formed.
c) Which cations were always spectator ions?
d) List the cations with the number of precipitates each formed.
3. Compare your results to the predictions based on the solubility rules. Were
there cases where a precipitate formed when it should not have, or a
precipitate did not form where predicted?
In each case, what might
have caused the discrepancy? (2 pts)
4. For each of the following pairs of mixtures of salts, determine whether the
cations can be separated easily from each other. Explain why or why not. For
those pairs that can be separated, describe what reagent you could use to separate
them. (4 pts)
a) Pb(NO3)2 and AgNO3
b) MgCl2 and BaCl2
c) Na2CO3 and K2CO3
d) Li2SO4 and CuSO4
GRADING
Prelab
Clean-up
Equations
Questions
4
2
6
10
56
23. THE 6-SOLUTION PUZZLE
INTRODUCTION
The art of logical thinking is essential to mastering chemistry. Although it is important to be familiar with
many chemical facts so that you can utilize the language of chemistry effectively, it is even more important
to be able to assimilate these facts and organize them in logical ways so they can be used to solve problems.
The art of logical thinking is not difficult to acquire but it does require practice. Try out your own ability on
the following logic problem.
PROBLEM
During a chemistry class, five pairs of students are working on five different experiments at a row of lab
stations numbered 1-5 consecutively. From the information given below, can you tell what each
student’s experiment is, at which lab station he or she is working, and who his or her lab partner is?
Clues:
1. Halogens experiment and all the other experiments are performed by lab partners of the opposite sex.
For example Ann Ion’s lab partner is a boy.
2. Milli Liter and Ben Zene work together.
3. Charles Law does not work at lab station #2.
4. Molly Cool and her partner work between station # 3 and the station occupied by Barry Um.
5. Phyllis Pipet, building a spectroscope, is no at #4.
6. Milli Liter does not work next to Molly Cool.
7. Earl N. Meyer works at #4.
8. Ionic reactions are being carried out at #5.
9. Hal Ogen is building a small-scale balance.
10. Only one of the Um twins, Francie or Barry, works at station #1 doing an acid-base titration.
57
In this activity will attempt to identify unknown solutions by noting their characteristic
reactions. You will be given 6 solutions to react with each other using micro scale
chemistry. The reactions, if they occur will be precipitation reactions. One combination
should produce bubbles. Some solutions have a characteristic color. The activity will be
carried out in groups of two.
MATERIALS
Droppers containing 0.1 M solutions of each of the following:
AgNO3
BaCl2
CuSO4
Na2CO3
HCl
Cu(NO3)2
PROCEDURE
1. Carry out micro-scale reactions with the known compounds. Record
observations.
2. Repeat experiments with unknown compounds.
3. Identify the unknown compounds.
4. Write equations for the reactions.
58
24. RATE OF A CHEMICAL REACTION
The rate of a chemical reaction is the rate at which reactants are converted to products. Some
reactions are fast and some reactions are slow. The rate of a specific reaction can be found only
by experiment. Reaction rates vary with the temperature and the concentrations of the reactants.
In this laboratory activity, you will study the effects of concentration of reactants and of
temperature on the rate of reaction between the iodate ion (IO3- )and the sulfite ion (SO32-). The
reaction is known as the Iodine Clock reaction, because a measurable amount of time elapses
before the reaction reaches completion. The reactions are as follows:
1) IO3- (aq) + SO32-(aq)
I- (aq) + SO42-(aq)
2) I- (aq) + IO3-
(aq) I2 (aq)
+ H2 O (l)
3) I2 + starch dark blue color
. The end of the reaction is signaled by the formation of the of the colored starch-iodine
complex.
You will run several experiments in which the concentration of the iodate ion is varied and
accurately time the reactions. You will then measure the time of reaction at two temperatures.
PRELAB
Have a watch that can time seconds.
1. Write a purpose.
2. Answer the questions:
1.What are the variables and the controls in
Part I
Part II
3. What is the total volume when both reactants are mixed in each trial?
4. What is the purpose of the starch in the reaction?
GOGGLES AND APRONS ARE REQUIRED DURING THE COURSE OF THE
EXPERIMENT. THE SOLUTIONS CONTAIN SULFURIC ACID WHICH IS
CORROSIVE.
PROCEDURE
You will be working in pairs.
PRACTICE
1. Obtain solutions A (iodate ion) and B (sulfite- starch). Clean two 10-mL graduate
cylinders and label them “A” and “B” respectively.
2. Measure 3 mL of solution A in one graduate and 3 mL of solution B into the other
graduate. Pour into individual test tubes and mix as demonstrated . Begin timing just as the
two liquids come in contact. As soon as the color appears, note the time.
3. Repeat this procedure until you get two close readings (within 2 seconds). Rinse test
tubes. It is not necessary to rinse graduate cylinders.
59
PART I Effect of concentration on time
1. Prepare different concentrations of solution A as shown in the Data table at the end of the
procedure. Add the water to solution A in the graduate so the volume with each dilution is 7
mL. Proceed as you practiced in PART I for each dilution.
PART III Effect of temperature on time (Demonstration)
The concentrations used in trial 3 will be run at 3 different temperatures and the times compared.
PART IV Effect of a catalyst (Time permitting)
CALCULATIONS
1. Calculate the concentration of the iodate solution for each dilution. The stock solution
contains 4.3 g KIO3 in one Liter of solution. Copy calculations onto Results table
2. Prepare the following graphs.
a) time vs [IO3-].
b) rate vs [IO3-]. Rate = 1/s.
3. Calculate the slope from graph 2.
QUESTIONS
1. What is the order of the reaction with respect to iodate ion? Explain referring to the
graph and data.
2. Write the rate law expression as determined by data in this lab.
3. Can the overall order of the reaction be determined from this experiment? Explain why
or why not?
4. How should the temperature affect the rate? Was this effect observed in this lab?
5. What does the slope represent?
GRADING
Prelab
Clean-up/safety
Graphs
Calculations
Questions
Organization
5
2
6
4
10
1
60
25. A STUDY OF CHEMICAL EQUILIBRIUM
Le Chatelier's principle says that if a system at equilibrium is stressed, the equilibrium balance
will shift in a direction that will relieve the stress. For example, if we add a reactant, the
equilibrium will shift toward products so that there is a different balance of reactants and
products. Similarly, if we add products, it will shift toward reactants. Keep in mind that ALL
SPECIES OF REACTANTS AND PRODUCTS ARE PRESENT IN THE REACTION
TEST TUBES.
Color is an easily observed macroscopic property that can be used to indicate shifts in
equilibrium concentrations. In this lab you will investigate chemical systems at equilibrium.
You will disturb them by adding or subtracting reactants or products and observe color changes
that indicate how the equilibrium systems respond. You will explain those changes in terms of
Le Chatelier's principle.
PRELAB:
1. Write a purpose.
2. Paste the Data Table in your lab book.
3. Write the equilibrium expressions (mass action expressions) for each reaction.
GOGGLES AND APRONS MUST BE WORN. CAUTION: HCl and NaOH cause burns;
avoid skin contact.
PROCEDURES
PART I Acid-Base Equilibrium
1. Add 2 drops of bromthymol blue (HBB )to about 5 cm3 of distilled water
2. Add 0.1 M HCl a drop at a time until a change occurs.
3. Add 0.1 M NaOH, a drop at a time, until a color change occurs.
4. Now add 0.1 M HCl a drop at a time until a color change again occurs.
PART II Iron (III)- thiocyanate Equilibrium
CAUTION: KSCN (potassium thiocyanate) is poisonous. Use with care. .
1. Place about four 4 mL portions of the iron (III) thiocyanate solution in 4 separate test
tubes. Note the color.
2. Leave one as a control to compare color changes.
1) To one test tube add 10 to 15 drops of Fe(NO3)3 (Fe+3 ) solution.
2) To the second test tube add 10 to 15 drops of the KSCN (SCN-) solution.
3) To the third tube add a few drops of 6 M NaOH. What is the solubility of
Fe(OH)3
3. Compare the color in each test tube to the original solution.
61
PART III
Solubility equilibrium for NaCl
1. In a clean test tube obtain 5 cm3 of saturated NaCl solution.
2. Add a few drops of concentrated HCl and record the results.
PART IV Complex ions of Cobalt II
1. Place about 0.5 cm of CoCl2 in a small test tube. Add about 1 mL distilled water.
2. Place the test tube in a hot water bath.
3. Add acetone until the color changes. (Acetone dissolves the water.)
4. Add water until the color changes.
5. Add solid CaCl2 until color changes, shaking after each addition
6. Add AgNO3.
CONCLUSION
1. Explain LeChatelier’s Principle. Use two systems from the lab as examples to illustrate it.
QUESTIONS
1. What is the most common chemical method for removing H3O+ ions in aqueous
solution? Write a net ionic equation which describes this method. (2 pts)
2. PART IV, on which side of the chemical reaction should the “heat” term be placed? (2
pts)
3. Predict the shift if the following were added to the systems at equilibrium. Answer on
the table provided:
PART I: H2SO4, KOH(aq)
PART II: Ca(SCN)2 (aq), K2CO3(aq), NaBr(aq), Ca(OH)2(aq), NaNO3 (aq)
PART III: NaNO3(aq) ,
solid NaCl
PART IV: Pb(NO3)2, NaBr(aq), NaCl(aq)
GRADING
Prelab
Questions
Conclusion
Organization
4
16
4
2
62
26. CHANGE IN ENTHALPY OF A REACTION
INTRODUCTION: When a chemical reaction takes place, chemical bonds in the
reactants are broken and new chemical bonds in the products are formed. Energy is
always absorbed in the breaking of bonds and always released when bonds are formed. If
the energy required to break old bonds is less than the energy released in forming new
bonds, the difference in energy is given off and the reaction is said to be exothermic. The
enthalpy change for an exothermic reaction is given a negative sign to indicate that
energy flows from the system.
In this experiment you will measure the amount of energy released by the decomposition
of hydrogen peroxide, H2O2, into water and oxygen and compare the experimentally
determined result with a calculated value.
The common disinfectant, 3% hydrogen peroxide (H2O2) solution, will be used with
manganese dioxide as a catalyst for the decomposition. The amount of energy released in
the reaction can be calculated from the mass of the water in the solution, the change in
temperature, and the specific heat of water.
PRELAB QUESTIONS:
1. State the purpose of this experiment.
2. Write the equation for the decomposition of hydrogen peroxide into water and oxygen
gas.
3. Use Hess’ Law to calculate the net energy released. Then calculate the energy in
kJ/mol. This will be the theoretical value.
H2 + ½ O2  H2O ΔH = -286 kJ/mol
H2 + O2 
H2O2 ΔH = -191 kJ/mol
PROCEDURE:
GOGGLES MUST BE WORN AT ALL TIMES
Each person will do the experiment independently and the results will be shared
with the class. Questions should be answered jointly with a partner.
1. Place two small scoops of manganese dioxide into the calorimeter.
2. Measure about 45 mL of 3% hydrogen peroxide and record the exact volume on the
data table.
3. Measure the temperature of the hydrogen peroxide in the graduated cylinder and record
it.
63
4. Pour the hydrogen peroxide into the calorimeter, put the lid on and push the
thermometer through the lid. Swirl gently and note the temperature change. Record the
highest temperature.
5. Open the calorimeter and look inside. Record your observation.
6. Empty the calorimeter and rinse it. Return it.
CALCULATIONS: Number the calculations and show set-up for each in your
report.
1. Calculate the change in enthalpy for the amount of H2O2 in your data using MCΔT.
Assume the density of the H2O2 is the same as water.
2. Convert the Joules to kilojoules.
3. Calculate the mass of H2O2 in the solution of 3% hydrogen peroxide.
4. Calculate the moles of H2O2.
5. Calculate the ΔH in kJ/mol.
6. Calculate what % of the theoretical ΔH you were able to measure in this experiment.
(Refer to question 4 in the Prelab for theoretical value.)
7. The values for the class will be collected and analyzed. Calculate a class average and
use this average to find your % error.
CONCLUSION QUESTIONS:
1. How will the mass of the system change throughout the experiment? Explain.
2. Compare your results to the theoretical value. How do they compare? Refer to your
data in your answer.
3. Suggest two reasons for the difference between your results and the theoretical value.
4. Although the reaction is exothermic, explain why it doesn’t feel warm when hydrogen
peroxide is put on a cut.
5. The decomposition of hydrogen peroxide is slow. If poured on the table there is no
evidence of decomposition. On the other hand when 3% hydrogen peroxide is applied to
a cut, the decomposition begins immediately (bubbles appear). Use your knowledge of
biology to suggest a reason why this occurs.
GRADING
Prelab
4
Safety clean-up
Calculations
Questions
Organization
2
7
10
2
64
27. DETERMINATION OF HYDROGEN ION
CONCENTRATION (pH) USING INDICATORS
The pH scale represents the hydronium ion concentration of a solution and is used to indicate how acidic
or basic a solution is. The pH can be measured with instruments or with indicators.
Acid-base indicators are organic compounds (usually weak acids) with characteristic vivid colors. Upon
addition of base, the acid is converted into its conjugate base, which is a different color. When the weak
acid and its conjugate base exist in solution in approximately equal concentrations, an intermediary color
is evident. For example, you are already familiar with the common acid-base indicator bromthymol blue,
or BTB. This is a weak acid that is yellow; its conjugate base is blue. It changes colors around pH 7. At
pH 7, approximately equal concentrations of BTB's conjugate acid and conjugate base exist in solution,
resulting in a green mixture.
HBTB + H2O  BTB- + H3O+
yellow
blue
Some indicators are mixtures of compounds, each of which changes color at a different pH.
Indicators are chosen for specific applications according to the pH range in which they show a color
change.
In this experiment you will investigate the color changes of various indicators as a function of pH.
You will also determine the identity of an unknown weak acid by observing its pH, calculating its Ka,
and comparing the Ka to that of known acids.
PROCEDURE
WEAR GOGGLES AND APRONS FOR THIS EXPERIMENT. SOLUTIONS
USED MAY CAUSE BURNS AND STAIN CLOTHING.
PART A Standard solutions
Each person at a lab table will do one or two indicators. You will then exchange data
with the other people at your table.
1. Place the template in the plastic envelope that has been cleaned. Any contamination
will interfere with your data.
2. Standard solutions range in pH from 1 to 12. Place about three drops of each solution
on the correct box. Take care not to mix the solutions.
3. Add two drops of your indicator to each drop of solution. DO NOT ALL0W THE
TIPS OF THE PIPETS TO BECOME CONTAMINATED. Record the color. Be sure to
indicate where there is a color change.
4. KEEP THE SHEETS FOR COMPARISON WITH UNKNOWN SOLUTIONS.
PART B Unknown solutions
EACH PERSON WILL DO A DIFFERENT UNKNOWN. Obtain an unknown
sample and record the number on your data table.
65
1. Repeat above procedure, testing the unknown with each of the indicators. USE THE
SAME NUMBER OF DROPS AS THE KNOWNS, SO THAT THE COLORS CAN BE
MATCHED.
2. Record the colors and pH values.
CALCULATIONS
Using the data and the concentration calculate the Ka and identify the acid from its Ka.
The possible unknowns and Ka values will be listed on the board. All unknown solutions
are 0.1 M.
QUESTIONS
1. Answer in TABLE FORM. Make a list of the indicators used:
a) State the lower color and the higher color.
b) State the range of pH over which the indictor changes.
2. Of the indicators used, which one(s) were useful in determining your unknown.
Which one(s) were not useful. Explain.
3. How could the pH of a solution be more accurately determined?
4. How do indicators work?
5. Which indicator would be best to use, if you ONLY needed to know whether a
solution was acidic or basic?
CONCLUSION
The conclusion should include:
1) An identification (and justification for the selection) of your unknown. Be
specific.
2) Sources of error.
GRADING
Clean-up/safety
Calculations
Questions
Conclusion
Identification
2
3
12
4
1
66
28. ACID-BASE TITRATION
In a neutralization reaction one mole of an acid (H3O+ ) reacts with one mole of a base (OH- ) to form a
salt and water. Thus the following relationship exists.
xVaMa = yVbMb
If three of the above quantities are known, the fourth may be calculated. In order to determine the
volumes a titration will be performed. An indicator, phenolphthalein, will be used to show the endpoint.
In this experiment you will titrate a standard solution of NaOH against a determined volume of acid of
unknown concentration. In the first part acid concentration will actually be known in order to
practice the technique. In the second part you will determine the concentration and the percent of
acetic acid in a sample of vinegar of unknown concentration.
MATERIALS:
2 Burets, 250-mL Erlenmeyer flask, beaker, wash bottle, small funnels
Solutions: phenolphthalein indicator, 0.10 M NaOH & 0.10 M HCl;
0.50 M NaOH & unknown solution of vinegar
PROCEDURE
GOGGLES AND APRONS MUST BE WORN DURING THIS LABORATORY
EACH STUDENT WILL PERFORM THIS ACTIVITY INDEPENDENTLY.
PART I: PRACTICE TECHNIQUE. (You will use 0.10M NaOH as the standard
base to find the molarity of HCl. (The molarity of HCl is 0.10M. Compare your value to
this to practice the technique.)
1. The proper use of burets and will be demonstrated.
2. Label burets “A” and “B”. Rinse the burets with water. Then rinse “A” with a
little acid and “B” with a little base.
3. Fill each buret with the proper solution. Let a little run out into the beaker so
the tip of the burets are filed. Record the initial volume to 0.1 mL. It is not
necessary to start at exactly 0.
4. Let about 8.0 to 10.0 mL of acid flow into an Erlenmeyer flask. Record the
final volume exactly.
5. Add about two drops of indicator to the flask.
6. Begin adding base. Swirl the flask after each addition. At the beginning you
can let several mL run into the flask at a time. As the color takes longer to fade
you will have to add less each time. The endpoint is reached when the color
changes to a LIGHT PINK.
67
7. Calculate molarity of acid. Repeat for another trial. If your results are
consistent and accurate proceed with Part II. If not, do another trial with the
known acid to perfect technique.
PART II FIND CONCENTRATION OF ACETIC ACID IN VINEGAR
1. Switch NaOH solutions to 0.50 M NaOH as the standard base for the vinegar
titration.
2. Repeat the above procedure with unknown vinegar solution and the 0.50 M
NaOH. Do three trials.
CALCULATIONS
1. Calculate molarity for each practice trial (on separate paper).
2. Calculate molarity for each trial of unknown.
3. Calculate average.
4. For each trial calculate % using the formula below. Use the molarity you
calculated (mole/L), to find the mass of acetic acid in the vinegar solution.
Assume the density of vinegar to be the same as that of pure water to get the mass
of the vinegar solution used in that trial.
% =
Mass (g) acetic acid
Mss (g) of vinegar solution
x 100
QUESTIONS
1. Write the equations and net ionic equations for the neutralization of each acid.
2. Suppose the unknown contained a solution of a polyprotic acid like sulfuric
acid. How would the results differ from this experiment?
3. Why would the indicator bromthymol blue (BTB) not be appropriate for
determining the endpoint of the vinegar/sodium hydroxide titration?
68
29. DECOMPOSITION OF BAKING SODA
PROBLEM:
Baking soda is a substance commonly used in baking to cause cakes to rise. They rise
because a gas is released either when the baking soda reacts with an acid or when it
decomposes. (What gas is released?) You will investigate the decomposition reaction of
baking soda to determine the solid product of the reaction.
Working in teams, devise a laboratory procedure that will experimentally determine
whether NaOH, Na2CO3, or Na2O is the solid product formed in the correct
decomposition reaction.
Write balanced equations for each of the three possible reactions.
2. Write a purpose.
3. Describe the procedure you will use to carry out the experiment.
4. Include safety directions, equipment and data tables.
5. Explain what calculations will be needed.
(Suggestion: You may wish to refer to other decomposition reactions you have carried
out this year.) Submit your prelab today. It will be checked for safety and feasibility;
therefore, it may be necessary to make some revisions.
In a later lab period, you will actually use your procedure (or revised procedure) to solve
the problem. Each person on the team will carry out the experiment. A final report will
be submitted by each team containing data and calculations from each member and a
conclusion based on the results of the team’s experiments.
GRADING
PRELAB
Procedure
Equations
Organization
5
3
4
POSTLAB
Laboratory technique/safety 3
Cooperation among
team members
2
Calculations
Results/conclusion
Organization
2
8
2
69
70
30. THE ICE CREAM LAB:
A PRACTICLE APPLICATION OF COLLIGATIVE PROPERTIES
Properties of solutions determined only by the number of dissolved particles and the nature
of the solvent are called Colligative properties. You will change milk into ice cream (ice
milk really) by dissolving sodium chloride in ice and placing the ice cream mixture to be
frozen (in a separate container) into this ice solution.
MATERIALS
1- gallon plastic bag
Rock salt`
Ice
1- quart plastic bag
Recipe:
2 tablespoons sugar*
½ cup of milk
¼ teaspoon vanilla
Any other flavors (chocolate, strawberry)
*If you add flavors other than vanilla, cut back on the sugar.
PROCEDURE
Place ingredients in the 1-quart bag. Expel the air and zip the bag. Place ice in the 1gallon bag and check the temperature. Add about 9 tablespoons of rock salt to the ice.
Place the small bag into the larger bag, expel the air and seal. Wrap a towel around the
bags and vigorously message for about five minutes or until the ingredients are frozen to
custard ice cream consistency.
Open the bag and check the temperature of the ice. Remove the small bag, wipe it off, eat
the ice cream.
QUESTIONS
1. What were the temperatures of the pure ice and the ice with salt?
2. Why was salt added to the ice?
3. Identify the changes as endothermic or exothermic:
a) in the inner bag containing milk.
b) in the outer bag containing ice and salt
4. Which colligative property was applied in this activity? Explain.
5. Would using calcium chloride instead of sodium chloride change the reaction? Explain.
6. If you did not add sugar would the ice cream have frozen faster? Why?
7. Why is salt spread on the roads during a winter storm
71
DATA TABLES
72
1. SAFETY LAB
GRADING
Safety/cleanup
Organization
Observations
Conclusion
4
5
3
6
NAMES_______________________
________________________
________________________
________________________
SAFETY LAB REPORT PAGE
1.
QUESTION:
MATERIALS:
OBSERVATIONS:
2.
QUESTION:
MATERIALS:
OBSERVATIONS:
73
3.
QUESTION:
DESCRIPTION/ SKETCH:
4.
QUESTION:
MATERIALS:
OBSERVATIONS/COMMENTS:
5.
QUESTION:
MATERIALS:
OBSERVATIONS/COMMENTS:
74
6.
QUESTION:
MATERIALS:
OBSERVATIONS/COMMENT:
SAFETY RULES:
75
3. COMPARING PHYSICAL AND CHEMICAL CHANGES
TABLE
EXPERIMENT
1.
OBSERVATION
2.
3.
4.
5.
6. a
b
7 a
b
77
78
5. DETERMINING AN EMPIRICAL FORMULA
DATA TABLE
Mass of crucible and cover __________
Mass of magnesium + crucible and cover__________
Mass after reaction _________
79
80
8. TYPES OF CHEMICAL REACTIONS
OBSERVATIONS
REACTANTS
OBSERVATIONS
TYPE OF
REACTION
A
B
C
D
E
F
81
CHEMICAL EQUATION
G
H
I
82
9. Relating Moles to Coefficients in an Equation
CALCULATION OF COPPER SULFATE HYDRATE REQUIRED:
DATA TABLE
Mass of beaker
Mass of copper sulfate
Mass of iron filings
Mass of dried copper
83
84
10. Particle Size from Collision Probabilities
DATA TABLE
Total number of marbles ROLLED, Tr
Total number of hits, H
Distance between walls, d
Number of target marbles used, N
Total length of 10 marbles lined up
85
86
11. FLAME TESTS AND EMISSION SPECTROSCOPY
Part I
Metal ion
Barium, Ba2+
Flame color
Calcium Ca2+
Copper, Cu2+
Lithium, Li+
Potassium, K+
Sodium, Na+
Strontium, Sr2+
Sodium &
potassium, Na+ & K+
Unknown #_____
+
Flame with cobalt glass
Potassium, K
Sodium, Na+
Sodium &
potassium, Na+ & K+
Part II
Fluorescent spectrum
Gas tube
1
2
3
4
5
red
orange
yellow
87
green
blue
violet
88
12. INVESTIGATION ON THE HYDROGEN SPECTRUM
NAMES_____________________________________________________
1. Using the equation given calculate the energies for each energy level and write it on the
appropriate place on the diagram below.
2. Calculate the energy of each line observed on the hydrogen spectrum.
a) RED ______________________ nm = _________________ m
b) BLUE ________________ nm = _____________ m
c) VIOLET ______________ nm = _____________ m
d) VIOLET* _____________ nm = _____________ m
89
3. Using the electron energies you calculated on the diagram and the energies
of the three (four) spectrum lines, determine which energy level transition
was responsible for each of the lines.
RED from n=______ to n=____
BLUE from n=______ to n=_____
VIOLET from n=______ to n=____
*VIOLET from n= _____ to n = ____
There are many more possible transitions for the electron than the ones you observed.
They are grouped into series and fall into different regions of the electromagnetic
spectrum, depending on the amount of energy released. The diagram below shows all the
possible transitions and the names of the series.
Which series did you observe?____________________________
90
13. DETERMINING THE HALF-LIFE OF Ba-137m
Background radiation ___________counts/______ = _______________cpm
Time Interval
Counts/30sec
Cpm (gross)
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
91
Cpm (net)
92
14. PERIODIC TABLE
A STUDY OF REACTIVITY OF METALS
PART I: Reactions of metals with water
METAL
Reaction with water
Room temperature
Li
Na
K
Rb
Cs
test Tes Hot ho
Ca
Hot water
Mg
Al
PART II: Precipitation reactions
Mg(NO3)2
Mg2+
Ca(NO3)2
Ca2+
Sr(NO3)2
Sr2+
H2SO4
SO42Na2CO3
CO32K2CrO4
CrO42-
93
Ba(NO3)2
Ba2+
Unknown
#______
94
15. PREPARATION AND PROPERTIES OF OXYGEN
PART I
Mass of test tube
Mass of test tube and potassium chlorate
Mass of test tube and potassium chlorate and
catalyst
Mass after reaction
PART II: OBSERVATIONS
1.
2.
3.
# re-lights
95
96
16. MODEL CONSTRUCTION: POLARITY AND SHAPE
97
98
17. PAPER CHROMATOGRAPHY
Marker
COLOR/NAME
Df (cm)
Ds(cm)
COLOR/NAME
COLOR/NAME
COLOR/NAME
COLOR/NAME
99
Rf
100
19. MOLAR VOLUME OF A GAS
Trial 1
Trial 2
Mass of Mg ribbon
Barometric
pressure(mmHg)
Temperature
Volume of H2 in
tube (adjusted for
atmospheric
pressure)
Vapor pressure of
H2O(mmHg)
101
102
20. AIR BAG LAB PROJECT
Scoring page
Every member of the group will receive the same grade.
Attach your report sheet to the back of this page.
Names_____________________
Table #___________
___________________________
___________________________
Success
score
Bag failed Bag
to inflate inflated
very little
or inflated
too much
and opened
0
7
Bag
inflated
with a lot
of wrinkles,
but did not
open
8
Cleanup
Cooperation
Part I
Part II
Part III
Organization/neatness
4
4
5
3
5
4
Possible score
35
Team’s score
_________
103
Bag inflated
without
opening, but
had some
wrinkles
Bag
inflated
without
wrinkling
or opening
9
10
I. PROCEDURE: List the steps of your plan.
II. MATERIALS: List all equipment and chemicals.
III. CALCULATIONS: Show your work for any calculations you need to make. Don’t
forget to include units.
104
21. APPLICATION OF SOLUBILITY RULES
K2CO3
1
AgNO3
Na3PO4
2
KI
3
Ba(NO3)2
4
NaCl
9
10
11
12
FeCl3
16
17
18
Na2SO4
22
23
Pb(C2H3O2)2
27
Ba(NO3)2
Na2SO4
6
7
13
14
15
19
20
21
24
25
26
28
29
30
31
32
33
KI
34
35
Na3PO4
36
105
Pb(C2H3O2)2
5
FeCl3
NaCl
8
106
22. THE 6-SOLUTION PUZZLE
NAMES_______________
GRADING:
Clean-up/safety
Cooperation
Results
Organization
________________
AgNO3
Na2CO3
HCl
CuSO4
Cu(NO3)2
BaCl2
PREDICTED REACTIONS
BaCl2
Cu(NO3)2
CuSO4
x
x
x
UNKNOWN SOLUTIONS
107
x
x
x
2
2
12
2
HCl
x
x
x
x
RESULTS
UNKNOWN #
COMPOUND
____________
_____________
____________
_____________
____________
_____________
____________
_____________
____________
_____________
____________
_____________
EQUATIONS:
108
23. RATE OF A CHEMICAL REACTION
Trial
Vol A (mL)
Vol H2O
(mL)
Vol B (mL)
Total vol.
(mL)
1
6
1
3
10
2
5
2
3
10
3
4
3
3
10
4
3
4
3
10
5
2
5
3
10
6
1
6
3
10
DATA TABLE PART II
Temperature
Time
sec
Time sec Average
time
*Room temp ____C
Low temp ____C
High temp ____C
* You may use data from Trial 3 in Part I.
Trial
RESULTS TABLE
Average time Rate 1/s
(s)
[IO3-] M
1
2
3
4
5
6
7
109
Time (s)
Time (s)
110
24. A STUDY OF CHEMICAL EQUILIBRIUM
Stress
Species added
or removed
observations
Change in concentration (or amount): indicate a
increase , decrease , no change n.c.
PART I
Step1
yellow
HBB
Step 2_______
Step 3_______
Step 4_______
_____
_____
_____
PART II
step 1______
step 2______
step 3______
pale brown
Fe3+
+ SCN- 
_____
_____
_____
_____
_____
_____
PART III
_________
NaCl(cr)  Na+
___
+

H
+
_____
_____
_____
+
Co(H2O)6
_____
_____
_____
_____
_____
step 3______
step 4______
step 5______
step 6______
111
2+
(aq)
+
_____
_____
_____
dark red
FeSCN2+
_____
_____
_____
Cl___
Pink
PART IV
step 1
step 2______
Blue
BB-
4Cl-(aq) 
_____
_____
_____
_____
_____
blue
CoCl42-(aq) + 6H2O(l
_____
_____
_____
_____
_____
Table for question 3.
Substance added
Visible change observed
Shift to reactants or
products
Part I H2SO4(aq)
KOH(aq)
Part II Ca(SCN)2 caq)
K2CO3(aq
NaBr(aq)
Ca(OH)2(aq)
NaNO3(aq)
Part III NaNO3(aq)
Solid NaCl
Part IV Pb(NO3)2(aq)
NaBr(aq)
NaCl(aq)
112
25. CHANGE IN ENTHALPY OF A REACTION
DATA TABLE
Volume of 3% hydrogen peroxide
Ti
Tf
Observations after highest temperature is reached
CALCULATIONS:
113
114
26. DETERMINATION OF HYDROGEN ION CONCENTRATION
pH
1
2
3
4
5
6
7
8
9
10
11
12
Unknown
#
115
116
27. ACID-BASE TITRATION
NAME________________________
Technique
Clean-up
Safety
Calculations
Questions
Identification
Organization
Total
DATE_____________
4
2
2
8
6
4
2
28
DATA TABLES
PART I: Practice trials with 0.1M NaOH
Buret
Trial 1 volume
Trial 2 volume
reading
HCl
NaOH
HCl
NaOH
NaOH
Final
Trial 3 volume (if
needed)
HCl
NaOH
Initial
Volume
used
PART II: Determine concentration of acetic acid in a vinegar solution with 0.50 M
NaOH.
UNKNOWN #_________
Buret
reading
Final
Trial 1 volume
Vinegar
NaOH
Trial 2
volume
Vinegar
NaOH
Initial
Volume
used
117
Trial 3 volume
Vinegar
NaOH
CALCULATIONS:
Molarity of vinegar solution.
Trial 1
% acetic acid in vinegar sample
Trial 1
Trial 2
Trial 2
Trial 3
Trial 3
Average molarity :
Average %:
QUESTIONS
118
119