Download Chemistry I - State College Area School District

Laboratory Manual
Introductory Experiments and Procedures
Chemistry I
Name ________________________________
Period
________
Lab Station _____
Mr. Gallagher
Room 132
0
Contents:
Introduction
Lab #1 Observations of a Burning Candle
Lab #2 Percent Oxygen in the Atmosphere
Lab #3 Using the Bunsen Burner
Lab #4 Separation of a Simple Mixture
Lab #5 Classification of Matter
Lab #6 Density Lab
Lab #7 Percent Sugar in Soda
Lab #8 Testing the Mettle of Metal
Lab #9 Striking It Rich
Lab #10 Emission Spectroscopy
Lab #11 Types of Chemical Bonds
Lab #12 Polar and Nonpolar Molecules
Lab #13 Microscale Crystallization
Lab #14 Pyrolysis of Wood
Lab #15 Formula of a Hydrate
Lab #16 Types of Chemical Reactions
Lab #17 The Iron Chemist
Lab #18 Copper Cycle Lab
Lab #19 Serial Dilutions
Lab #20 Acid-Base Properties of Com. Sub.
Lab #21 Acid-Base Titration
Lab #22 Heat of Combustion Lab
Lab #23 Radioactive Decay and Half-Life
Common Laboratory Equipment List
Lab Drawer Equipment List
Interpreting Chemical Hazard Labels
Unit Conversions and Formulas
Rules for Proper Graphing Technique
Chemistry I Course Guidelines
Student Safety Contract Rules
Student Safety Contract Sign-off Sheet
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Lab Partners:
Names
Phone #
email
________________________
_________________
__________________
________________________
_________________
__________________
1
Introduction:
The study of chemistry is vital to understanding the natural world around you.
The principles of chemistry are at the heart of nearly every aspect of life that you perhaps
take for granted, but nevertheless, depend upon every day. The agricultural, automotive,
cosmetics, energy, food service, pharmaceutical, and plastics industries, just to name a
few, are directly dependent upon an in-depth understanding of chemistry. New advances
in chemistry will eventually have a positive impact on the quality of your life. Chemists
will play a major role in finding ways to clean up the environment, develop alternative
sources of fuel, and increase crop yields around the world. In addition to all of these
practical benefits, chemistry is a fascinating subject, worthy of study for curiosity’s sake.
Chemistry is also an experimental science. Experimentation is the key to
discovery and valid experiments depend on accurate measurements and detailed
observations. Therefore, we will be focusing on measurement and observation
techniques in our first unit. One of the main objectives of this course is for you to
develop the thinking skills and habits of a scientist. Here is a partial list of some of these
characteristics.
Scientists:
 Pay attention to detail.
 Are careful to record their observations and results
 Look for connections.
 Never stop asking questions.
 Are not afraid to try new things.
 Talk to other scientists about their results and ideas.
The laboratory activities in this unit are designed to challenge you to utilize these skills.
As the year goes on, you will need to employ these skills in order to succeed in this class,
and hopefully you will utilize them for the rest of your life.
“The most important of my discoveries have been suggested
to me by my failures.”
~ Sir Humphrey Davy
2
Lab #1
Observations of a Burning Candle
“The Greatest enemy of knowledge is not ignorance, it is the illusion of
knowledge.”
~Stephen Hawking
Objectives:
1. To improve your observation skills.
2. To learn to ask questions that lead to discovery.
3. To develop a hypothesis of the process of the burning of a candle based on evidence
derived from your observations.
Note: Include labeled drawings with all of your observations.
Part A – Initial Observations
Procedure:
Light a candle and record your observations with as much detail as you can. Here are
some suggestions to help you observe more thoroughly: What are the colors of the
flame? What is the shape of the flame? Is the candle’s flame hollow? What is the color
and shape of the wick? Does the wick give visible evidence of being hot? Does the wick
get longer as the candle wax is consumed? Is the wick combustible? Does the wick burn
right at the surface of the candle wax?
Observations:
Unanswered questions:
Part B - What’s in a Flame?
Blow gently at the candle flame through a horizontal straw that is held about one inch
from the flame. Start at the top and work down. What do you see? Record your
observations in the form of several labeled diagrams. You may want to include this
method of recording observations throughout this lab.
Observations:
3
Part C – Where There’s Smoke!
Blow out the flame. What do you see? Describe the smoke. From what part of the
candle does the smoke originate? What is the last site of smoke generation? Hypothesize
as to the composition of the smoke. What other event(s) occurs with the cessation of
smoke generation?
Observations:
Relight the candle and allow it to burn for about a minute. With a lighted match in one
hand, blow out the candle and hold the burning match in the rising column of smoke
about an inch above the candle’s wick. What happens? Does this support your
hypothesis about the composition of the smoke?
Observations:
Part D – Scorch Marks
Move a horizontally held notecard (that has been moistened with water on the underside)
quickly down over the flame to a height where the card almost touches the wick. Hold
the notecard in this position just long enough for the flame to scorch through to the upper
side of the card, but not long enough to cause it to burst into flames. (If the card begins
to burn, toss it into the sink and run water to extinguish the flame.) Draw a pattern of the
scorched area on the top of the card. Note the shading.
Observations:
4
Part E – The Chimney Effect
Wind a length of copper wire (16 gauge or heavier) around a pencil about ten times,
leaving enough uncoiled to serve as a handle. Carefully slip the coiled wire off the
pencil. Lower the coiled portion over the flame so that the flame is passing through the
center of the coils and the copper coils wrap around the flame. You may need to vary the
height of the coiled wire to achieve the desired effect. Record your observations.
Observations:
Vary the length of the coiled portion and the spacing between the individual coils. Note
how each affects the candle flame.
Observations:
Crush the head of a burnt match into a fine powder and place some of this material into
the liquid in the bowl of the burning candle. Observe the motion of these particles in the
liquid.
Observations:
Part F – Wikipee?
Lengthen the wick to about 1 cm by removing some of the wax. Light the candle and
allow it to burn for a minute or so. Place the edge of a clean scoopula against the wick
about halfway up the wick. Do not allow the scoopula to touch the wax. Observe for at
least one minute and record your observations.
Observations:
5
Analysis and Conclusions:
Based on all of your observations propose a hypothesis for how a candle burns. Describe
the entire process. Be specific. Include all of the components of the candle; the solid
wax, the melted wax, the wick, the flame, and any other relevant facts. Be sure to refer to
your observations in parts A-F to support your hypothesis.
Grading:
Observations (Completion) Part A
Part B
Part C
Part D
Part E
Part F
Observations (Quality)
Excellent - Thoughtful & Detailed
Good - requires more detail/thought
Average – be more observant
Poor – superficial
Conclusion
Total
(2)
(2)
(3)
(2)
(3)
(2)
___
___
___
___
___
___
(5)
6
4
3
2
___
(25)
___
6
Lab #2 Percent Oxygen in the Atmosphere
Objective: In this lab you will be determining the percentage of oxygen in the air you
breathe.
Procedure:
1. Clean the outside of a large ignition tube with alcohol and let dry.
2. Obtain a metric ruler and glass marking pencil. Starting from the open-end of the
tube mark every centimeter on the outside of the tube with the marking pencil.
3. Fill a 400 mL beaker about ¾ full with tap water.
4. Obtain a small piece of steel wool. Prepare a 50:50 mixture of vinegar and water
by mixing 10 mL of vinegar with 10 mL of water in a beaker. Soak the steel
wool in a 50:50 mixture of vinegar/water for 1 minute. Remove the steel wool
and carefully squeeze out any excess solution.
5. Pull the steel wool apart to increase the surface area. Push the steel wool into the
marked ignition tube with a stirring rod.
6. Cover the end of the tube with your finger and invert tube into beaker of water.
7. Every 5 minutes, move the test tube so that the water level inside the tube is equal
to the water level in the beaker. Record the height of the water.
8. Continue to measure and record the height every five until the water level stops
changing. Allow the tube with steel wool to sit over night.
9. Carefully measure the height of the water and record.
10. Record the color of the steel wool.
Analysis: (Complete on separate paper)
1. Complete the following calculation;
distance water traveled up in the tube
----------------------------------------------- x 100% =
total length of tube
2. Where did the oxygen go?
3. Were the changes that you observed physical or chemical? Explain.
4. Air is about 20.8% oxygen. How does your value compare to this value?
Calculate the percent error for your value and propose some reasons as to why
your value is different.
5. Write a procedure to determine the mass of oxygen that was removed from the air
in the tube.
Grading Rubric:
Lab procedure and technique ___ (5)
Observations and results
___ (3)
Analysis Questions (2,3)
___ (3)
Analysis Questions (4,5)
___ (4)
Total
___ (15)
7
“All truths are easy to understand once they are discovered; the point is to
discover them.”
~ Galileo Galilei
8
9
10
Lab #4 Separation of a Simple Mixture
Objectives:
1. To separate a mixture of iron, sulfur and salt.
2. To recover as much of the original mass as possible.
3. To learn important laboratory techniques.
Pre-Lab:
Define the following terms:
1. Pure substance –
2. Element –
3. Compound –
4. Homogeneous mixture –
5. Heterogeneous mixture –
Procedure:
1. Mass a clean, dry watch-glass.
2. Place a scoop of the iron-sulfur-salt mixture on the watch-glass, and determine the
mass of the mixture.
3. Use a magnet to remove the iron from the mixture. Transfer the iron to a premassed piece of weighing paper. You may need to go over the mixture more than
once. Determine the mass of the iron.
4. Transfer the sulfur-salt mixture to a 50 mL beaker. Add 25 mL of water and stir
with a glass stirring rod to dissolve the salt.
5. Place a piece of pre-massed filter paper in a funnel and place the funnel in a 125
mL Erlenmeyer flask.
6. Filter the mixture and collect the filtrate – the liquid that passes through the filter.
7. Wash the residue in the filter with 15 mL of water and collect the rinse water with
the filtrate. Set the filter paper and residue aside to dry.
8. Transfer the filtrate to a pre-massed evaporating dish. Set up a ring stand to hold
the evaporating dish and evaporate the water from the filtrate using a #2 flame
with the Bunsen burner. Try to avoid spattering during evaporation.
9. Allow the evaporating dish to cool and then determine the mass of the recovered
salt.
10. Determine the mass of sulfur once it has completely dried.
Science is built up of facts, as a house is built of stones; but an
11
accumulation of facts is no more a science than a heap of stones is a house.
Henri Poincaré
Data:
Table 1 – Mass of original mixture
Mass of original mixture + watch-glass
Mass of watch-glass
Mass of original mixture
______________ g
______________ g
______________ g
Table 2 – Mass of iron
Mass of iron + weighing paper
Mass of weighing paper
Mass of iron
______________ g
______________ g
______________ g
Table 3 – Mass of sulfur
Mass of filter paper + sulfur
Mass of filter paper
Mass of sulfur
______________ g
______________ g
______________ g
Table 4 – Mass of salt
Mass of evaporating dish + salt
Mass of evaporating dish
Mass of salt
______________ g
______________ g
______________ g
Table 5 – Summary of results
Mass of original mixture
Mass of recovered iron + sulfur + salt
_______________ g
_______________ g
12
Mass difference
_______________ g
Analysis:
1. What properties did you observe in each of the components of the mixture?
2. How did these properties help you to separate the components of the mixture?
3. What types of changes (chemical / physical) were involved in the separation of
the components of the mixture?
4. List at least 5 specific factors that may have contributed to the mass difference
that you observed.
5. What changes could be made to this procedure to increase the accuracy of the
results?
Grading Rubric:
Pre-Lab
Laboratory technique
Data
Analysis
Safety and Clean-up
____
____
____
____
____
Total ____
(2)
(5)
(5)
(6)
(2)
(20)
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Lab #5 Classification of Matter
“Science is organized knowledge. Wisdom is organized life.”
~Immanuel Kant:
Introduction: Matter can be classified according to its composition and its properties.
Upon careful inspection, matter can first be divided into categories based on whether the
composition is uniform throughout or non-uniform and therefore, can be separated into
distinct parts. Matter that is uniform in composition and properties can be further
categorized as either pure substances or mixtures. Finally, those that are pure substances
may be either elements or compounds. Elements are pure substances that cannot be
broken down into simpler substances by physical or chemical means. Compounds, on the
other hand, are pure substances made of two or more elements that can be broken down
into simpler substances by chemical means.
Objective: In this lab you will distinguish between elements and compounds based on
mass changes that occur during a chemical reaction.
Materials:
crucibles (2)
crucible lid (1)
ring stand
clay triangle
electronic balance
sand paper
crucible tongs
iron ring
Bunsen burner
PreLab:
Fill in the each box with one of the terms listed below to create a
classification scheme that describes all matter.
( Compounds, Elements, Heterogeneous matter, Homogeneous Matter,
Homogeneous mixtures / Solutions, Matter, Pure Substances)
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Procedure:
Part 1
1.
Set up the ring stand, iron ring, and clay triangle as shown by your teacher.
2.
Obtain a clean, dry crucible and record its mass. If you need to use water to clean
it, then you will need to heat the crucible for about 3 minutes to dry it completely.
Allow it to cool before proceeding.
3.
Place about 0.5 g of the white powder into the crucible and record the mass.
4.
Place the crucible on the clay triangle and heat it with a #3 flame for 6 minutes.
Allow the crucible to cool for 5 minutes, or until it is cool enough to touch.
5.
Re-mass the crucible. Calculate the change in mass of the white powder.
6.
Dispose of the white powder in the trash.
Part 2
1.
While you are waiting for the crucible in part 1 to cool, obtain a second clean, dry
crucible.
2.
Record the mass of the crucible and a lid.
3.
Obtain a metal ribbon and lightly sand its entire surface. Wipe the surface with a
Kim-wipe to remove any dust.
4.
Loosely coil the metal ribbon and place it in the bottom of the crucible.
5.
Mass the crucible and lid with the metal ribbon.
6.
Place the crucible on the clay triangle, and place the lid on the crucible so that
only a small gap is visible.
7.
Heat the crucible for 12 minutes with the hottest flame. Allow it to cool for at
least 5 minutes.
8.
Re-mass the crucible and lid with the contents. Calculate the change in mass of
the metal ribbon.
Data:
Data Table 1: Mass of White Powder Before and After Heating
Mass of crucible and powder before heating
Mass of crucible
Calculated mass of white powder before heating
Mass of crucible and powder after heating
Calculated mass of white powder after heating
Change in mass of white powder
________ g
________ g
________ g
________ g
________ g
________ g
Data Table 2: Mass of Metal Ribbon Before and After Heating
Mass of crucible, lid and metal before heating
Mass of crucible and lid
Calculated mass of metal ribbon
________ g
________ g
________ g
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Mass of crucible, lid and metal after heating
________ g
Calculated mass of metal (remaining)
________ g
Change in mass of metal
________ g
Analysis: [Refer to the stated objective for the lab to answer these questions.]
1.
Based on your calculation for the change in mass that you observed in part 1,
classify the white powder that you placed in the crucible according to the
classification scheme that you created in the pre-lab.
2.
Classify the product that was obtained after heating the white powder.
3.
Based on your calculation for the change in mass that you observed in part 2,
classify the metal ribbon that you placed in the crucible.
4.
Classify the product that was obtained after heating the metal ribbon.
5.
Hypothesize as to what may have happened in each of the crucibles to cause the
changes in mass. Be as specific as possible.
Grading Scale:
PreLab
Data Table 1
Data Table 2
Analysis
#1
2
3
(3)
(3)
(3)
(1)
(1)
(1)
____
____
____
____
____
____
16
4
5
Total
(1)
(4)
(17)
____
____
____
Lab # 6 Density Lab
An experiment is a question, which science poses to Nature, and a measurement
is the recording of Nature's answer. ~Max Planck
Objectives:
1. To learn the proper use of significant digits in calculations involving measurements.
2. To identify seven unknown metals from their measured densities.
3. To learn proper graphing techniques and analysis.
Part A - Identifying Unknown Metals
Procedure:
1. Mass each metal object on the balance that is provided at that station.
2. Use the measuring instrument available at each station to determine the volume of the
metal object.
3. Calculate the density of the metal and report it to the correct number of significant
digits.
4. Look up the densities of the metals listed in reference table A.
5. Identify each of the metals A – G.
Reference Table A1
Accepted Densities of Metals
Density (g/cm3)
Aluminum
Chromium
Copper
Gallium
Iron
Lead
Nickel
Tin
Titanium
Zinc
Use your periodic table to find the densities of the metals in Table A.
17
Table A2 – Mass and Volume Measurements of Various Metals
Metal
Mass (g)
Volume Measurements
(cm3)
Calculated Density
(g/cm3)
A
B
C
D
E
F
G
Volume formulas:
Rectangular solids:
V = length x width x height
Cylinders:
V= r2 x height
18
Table A3 – Summary of Results
Metal
Balance
Precision
(0.1-0.001)
Volume
Method
Volume
precision
(0.1 – 0.001)
Density
(g/cm3)
Sig figs
Identity of
Metal
A
B
C
D
E
F
G
Calculations:
Part B - Densities of Plastic Cylinders
Procedure:
1. Mass each cylinder on the electronic balance.
2. Measure the length and diameter of each cylinder using a ruler.
3. Calculate the volume of each cylinder.
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4. Create a graph of mass vs. volume and plot the data points from the 12 cylinders
that you measured. (Follow the guidelines listed for proper graphing technique in
Appendix B on p. 79.)
Table B1 – Mass and Volume Measurements of Plastic Cylinders
Cylinder #
1
Mass (g)
Length (cm)
Diameter (cm)
Volume (cm3)
2
3
4
5
6
7
8
9
10
11
12
Analysis:
1. Determine the proper type of line or curve that should be drawn to fit your data.
2. What does this line or curve represent?
3. What are the densities of each of the plastic cylinders?
20
4. What conclusions can be drawn from your graph pertaining to the composition of
the plastic cylinders?
Grading Rubric:
Accepted densities of metals
Data Table A2
Summary Table A3
Correct Sig Figs
Correct Identities
Data Table B1
Graph
(2)
(5)
(5)
(5)
(7)
(6)
(6)
___
___
___
___
___
___
___
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Analysis questions
Total
Lab #7
(4)
___
(40)
___
Percent Sugar in Soft Drinks
Introduction:
In this lab exercise you will gain skill in producing proper graphs from laboratory
data. Graphs are important tools used to express a set of data visually. They enable us to
quickly identify trends and relationships among experimental variables. Graphs can also
be used to determine the value of an unknown variable such as concentration. Therefore,
it is important that the graphs that you create are both neat and accurate.
Refer to the guidelines for proper graphing technique found in Appendix B as you
construct your graphs from the data given and obtained. You will be expected to follow
these guidelines for all of the graphs you create in this course.
Pre-Lab:
Define the following terms and explain how these techniques are used in
graphing.
1. extrapolate 2. interpolate 3. Describe a real-world situation where extrapolation or interpolation is used.
Objective:
To determine the percent sugar found in 2 different soft drinks by
graphical analysis.
Materials:
Graduated Cylinder (10mL) Sugar solutions: (1%, 5%, 10%, and 20%)
Graph paper
Soft Drinks
Balance
Procedure:
Part 1
1. Tare a dry 10 mL graduated cylinder.
2. Pour out 8-10 mL of 1% sugar solution. Do not pour out more than 10 mL!
3. Record the volume of the sugar solution.
4. Mass the graduated cylinder and sugar solution. Record the mass.
5. Calculate the density of the sugar solution and record on the table.
6. Use a separate graduated cylinder for each solution or rinse the graduated cylinder
with distilled water and dry before going on.
7. Repeat steps 1-5 for the 5%, 10%, and 20% sugar solutions.
22
Research is to see what everybody else has seen, and to think what nobody else
has thought.
Albert Szent-Györgi
Part 2
8. Obtain a piece of graph paper.
9. Graph the density and the % solution. Refer to proper graphing technique to
determine which variable should go on the x-axis, and which should go on the y-axis.
10. Draw the appropriate type of line for these data points. Attach your graph to the lab.
Part 3
11. Repeat steps 1-5 using three of the soft drinks. Record the data on the table.
12. Calculate the density of each soft drink.
13. Using the density, determine the percent sugar in each soft drink by interpolating on
the graph from Part 1. Draw the interpolation lines on the graph.
Data Table A: Sugar Solution Densities
Sugar Solution
mass (g)
volume (mL)
1%
5%
10%
20%
Soda:_____________
Soda:_____________
Soda:_____________
density (g/mL)
Density Calculations- (show your work!)
Analysis:
1. What is the percent sugar in each soft drink?
% Sugar
__________________________
__________
__________________________
__________
__________________________
__________
2. Which measurement, mass or volume, limits the precision of the value of the %
sugar? Explain your choice.
3. List at least 4 very specific sources of error that could affect your calculated % sugar
in your soda samples. Simply saying incorrect mass or volume measurements does
23
not count! Specify what could have contributed to incorrect mass or volume
measurements.
Lab #8 Testing the Mettle of Metal
A Controlled Experiment
Introduction:
How closely matched are a series of measurements or observations is called
reproducibility. Scientific measurements are usually repeated to show reproducibility. If
measurements or observations are not reproducible, there must be a reason. Anything
that affects reproducibility is called a variable. Experiments are usually designed to
study only one variable at a time. All other variables are held constant so that they do not
affect the experimental results with repeated measurements. Such experiments are called
“controlled experiments”. Conclusions that are drawn from experiments that are not
controlled, and/or that have not demonstrated reproducible results are invalid.
In this experiment you will study the effects that various methods of heating and
cooling have on the physical properties of a certain metal. In order to establish the effect
that is caused by each method of heating and cooling you will first perform the bending
test on metal that has not been heated. This will serve as the experimental control. You
will also perform multiple trials for each experiment so that you can demonstrate
reproducibility and therefore draw valid conclusions from your results.
Metals are used for a wide variety of applications. In certain uses, the metal must
be able to bend easily without breaking, and in other applications we may need the metal
to resist bending. We may be able to find the desired property by choosing different
metals or by mixing metals to create an alloy. However, another option available to
metallurgists is to heat treat a metal to create the desired properties required for a specific
application. In this experiment you will investigate the effects of the heat treatment
techniques known as annealing, quenching, and tempering.
Objective:
To determine the effects of annealing, quenching, and tempering on metals.
Pre-Lab:
A.
Define the following terms:
1. Annealing –
2. Quenching –
3. Tempering –
B.
Create a hypothesis as to which process will enable the metal to be bent the most
without breaking and which process will cause the metal to be the most brittle.
24
C.
Read the procedure and prepare the appropriate data tables.
Procedure:
Control Sample
1. Straighten a bobby pin and determine the number of times it must be bent in
order to break it in two. Record this on the data table. Repeat this procedure
two more times.
Annealing
1. Heat a bobby pin to red-hot by holding it over a #3 flame with crucible tongs.
It must remain red hot for thirty seconds. Then gradually lift it straight up
until it is about 12 inches out of the flame. Let the sample cool gradually for
about three minutes. This process of strong heating and slow cooling is
called annealing.
2. After it has cooled, bend it back and forth until it breaks and record the
number of bends that it takes to break the metal.
3. Repeat this procedure two more times.
Quenching
1. Heat a bobby pin to red-hot by holding it over a #3 flame with crucible tongs.
When it is red hot, immediately place it in a beaker of water. This process of
strong heating and quick cooling is called quenching.
2. Bend the pin back and forth until it breaks and record the number of bends
needed to break the pin.
3. Repeat this procedure two more times.
Tempering
1. Heat a bobby pin to red-hot by holding it over a #3 flame with crucible tongs
and keep it red-hot for thirty seconds. Then place it in a beaker of water.
Reheat the pin until it glows with a dull redness and remove it gradually from
the flame as you did in the annealing process. This process of strong heating,
quick cooling, strong heating and then slow cooling is called tempering.
2. Bend the pin as before and determine the number of bends needed to break
the pin.
3. Repeat this procedure two more times.
Repeat these four procedures (control, annealing, quenching, and tempering) for any
other metals that are available for you to test.
Those who have an excessive faith in their theories or in their ideas are not only
poorly disposed to make discoveries, but they also make very poor observations.
25
Claude Bernard
Analysis:
1. Describe the effects of annealing on a metal.
2. What applications of metals would benefit from the annealing process?
3. Describe the effects of quenching on a metal.
4. What applications of metals would benefit from the quenching process?
5. Describe the effects of tempering on a metal.
6. What applications of metals would benefit from the tempering process?
7. Based on the requirements for controlled experiments, evaluate the validity of
the conclusions that you drew from results of your experiments.
26
Lab # 9
Striking It Rich!
Safety Precautions:
Sodium hydroxide is very caustic! It will burn your eyes and your skin. Keep your
safety goggles on at all times and wear a lab apron. Be sure to clean up any spills and
wash your hands when you are finished.
Refer to pp. 89-91 in the ChemCom textbook for the procedure for this experiment. Note
that there are a few additional steps that you will be required to perform.
Additions to lab procedure:
1. Mass each of the pennies before beginning the experiment.
2. Mass each of the pennies after you have completed the experiment. Be sure to rinse
and dry the pennies that you altered.
Data / Observations:
Penny
Year
Initial
Mass
(g)
Final
mass
(g)
Initial
appearance
Appearance
after NaOH
treatment
Appearance
after heating
1
2
3
Analysis:
1. Describe the changes in mass for each penny.
2. What tests could you perform to determine if penny #2 was really silver and penny
#3 was really gold?
3. Were the changes that you observed physical changes or chemical changes? Explain
your answer.
27
4. Do you think that it would be possible to convert the pennies back to their initial
condition? Explain your answer.
5. What is the Law of Conservation of Matter?
6. Explain how your lab results relate to this law.
7. Think about what happened to the copper atoms as the pennies went through their
various changes. Also consider any other atoms that may have been involved in these
changes. Propose a hypothesis to explain what happened to all of the atoms involved
to cause the visible changes that you saw. Pay particular attention to the sequence of
events involving the atoms and their changes. Draw a labeled diagram to help with
your written explanation.
Grading:
Lab technique
Observations
Analysis #1
#2
#3
#4
#5
#6
#7
Total
(5)
(3)
(1)
(1)
(2)
(2)
(1)
(2)
(5)
(22)
___
___
___
___
___
___
___
___
___
___
Science progresses best when observations force us to alter our
preconceptions.
28
Vera Rubin
Lab #10 EMISSION SPECTROSCOPY
Introduction:
When you look at a rainbow you are looking at an example of a continuous
spectrum, which is produced when white light is refracted through a prism or water
droplets. In a continuous spectrum all the wavelengths of light are present and so the
colors of light appear to gradually blend from one to the next. However, a bright-line
spectrum seems to have many of the colors missing and only a few or in some cases
dozens of bright colored lines are visible. What causes these bright-line spectra and what
do they have to do with atomic theory or the structure of the atom? When atoms of
various gases are excited with heat or electrical energy, they emit light of characteristic
color. The color that is visible is actually a composite of light of specific wavelengths
that can be seen as bright lines when viewed through a spectroscope. The color and
therefore the wavelength of each bright line are determined by the energy of the light that
is emitted from the atom. Each bright line corresponds to a specific amount of energy
that is emitted (hence the term “emission spectroscopy”) when an electron loses the
energy that it absorbed and returns to a lower energy level. This series of bright lines or
bright-line spectrum is unique for each gas, and serves as an atomic fingerprint that can
be used to identify the gas. This method actually led to the discovery of Helium on the
sun before it was discovered on earth.
But why are there only a few distinct bright lines generated instead of a
continuous spectrum? The reason is that the electrons moving around in an atom may
only possess certain specific amounts of energy. Think of electrons as occupying
positions on a staircase. You may find an electron on the first step or the second step or
the fifth step, but you’ll never find an electron hanging out in between two steps.
Therefore, the difference in energy between two given steps will be a set value for a
given staircase. Now think about how this analogy relates to the atom. The steps on the
staircase represent the allowable energy levels or quantum shells. When an electron
moves from a higher energy level to a lower energy level the excess energy is given off
in the form of light. The energy of the light and therefore its color will depend on the
difference in energy between the two levels.
Before you perform this experiment, you are going to calculate the wavelengths
of light that you should observe when you look at the bright-line spectrum of Hydrogen.
You may want to refer to pp. 92-94 in your textbook for more background on this topic.
Pre-Lab Assignment:
Show all of your work, including the units for your calculations!
1. What is the difference between a continuous spectrum and a bright-line spectrum?
29
2. The energies of the four visible bright lines in the Hydrogen spectrum are as follows:
(3.027 x 10-19 J), (4.086 x 10-19 J), (4.576 x 10-19 J), and (4.843 x 10-19J). Remember
as you work on this lab that these energies represent two related things; the energy of
the light given off and the difference in energy (E) between two different energy
levels. Calculate the frequency () of each of the bright lines of Hydrogen. Use
Planck’s equation, which states that the energy of light is equal to Planck’s constant
times the frequency of the light. E = h .  where Planck’s constant,
h = 6.626 x 10-34 J. s The units for frequency are 1/s or a Hertz (Hz). One Hertz is a
cycle per second.
3. Once you have calculated the frequency of each bright line, you will now be able to
calculate the wavelength of each bright line. The frequency of a wave and the
wavelength are related by the equation c =  .  , where c is the speed of light and 
(lambda) is the wavelength. The speed of light is a constant and is equal to
2.9979 x 108 m/s. Convert the wavelengths from meters to Angstroms.
1 m = 1010 A
5. Now that you have the wavelengths of each line, use the chart below to predict the
color of each of the bright lines.
Color
wavelength range
Violet
Indigo
Blue
Green
Yellow
Orange
Red
3800 – 4500 A
4200 – 4500 A
4500 – 4950 A
4950 – 5700 A
5700 – 5900 A
5900 – 6200 A
6200 – 7500 A
30
Summary of Predicted Results:
Line
Energy (J)
Frequency (1/s)
Wavelength (A)
Predicted Color
1.
2.
3.
4.
Objectives:
You will observe the bright-line or emission spectra for several gases. You will use the
emission spectra to identify 3 unknown gases. You will also calculate the energy
involved in producing the bright lines of certain colors for hydrogen.
Procedure:
1.
Create four charts on graph paper like the one below. You will use colored pencils
to record the emission spectra for each of the gases observed. Be as precise as you
can when recording the location of each line.
Identity of Gas _________________
wavelength (Angstroms)
4000
5000
6000
7000
|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_|_
|
|
|
|
___________________________________________
2. Use the reference chart provided to identify the unknown gases.
Analysis: (Show your work for all calculations, including the units!)
1. Compare the wavelengths of the spectral lines that you observed for hydrogen and
compare them with the wavelengths that you calculated in the pre-lab assignment.
Calculate the percent error for each of the lines that you observed.
2. Which of the values were more precise; your observed wavelengths or the calculated
wavelengths? How do you account for this difference in precision?
3. Now you will be performing the same calculations that you did in the Pre-lab, but in
reverse order. You will start with your observed wavelengths and calculate the
energies of each observed bright line. First calculate the frequency of each line and
31
then calculate the energy of each line. Remember to report your answer to the correct
number of significant digits!
4. The visible lines in the Hydrogen spectrum all result when electrons return to the n=2
energy level. From which of the energy levels (n=3, n=4, etc.) did the electrons that
created each of the four bright lines originate? Explain on what you base your
answers.
Summary of Calculated Results:
Line
Color
Wavelength (A)
Frequency (1/s)
Energy (J)
1.
2.
3.
4.
5.
Read section 3.3 pp. 92-94 in your textbook (Holt). Write a significantly detailed
summary that explains your observations from the lab. This will serve as your
conclusion for the lab.
Grading Rubric:
Pre-Lab Assignment
Spectra of Elements (4)
Identities of unknowns
Percent Errors
Analysis
Summary
Conclusion
____
____
____
____
____
____
____
(12)
(12)
(3)
(3)
(8)
(2)
(4)
32
Total
____
(44)
Lab #11 Types of Chemical Bonds and Physical Properties
Introduction:
The type of chemical bonds present in a compound will determine the general
physical properties of that compound. In this experiment the properties of solubility,
melting point and conductivity will be compared for typical ionic and covalent
compounds. Each of these properties is primarily determined by how easily the particles
that make up these compounds can be separated under various circumstances. However,
it is important to note that whenever an ionic substance undergoes one of these physical
changes, the individual ions that make up the crystal must be separated. Therefore, with
ionic compounds the ionic bonds must, in essence, be broken. On the other hand, when a
covalent compound undergoes one of these physical changes, the chemical bonds
between the atoms are not broken, rather the forces of attraction that exist between
molecules must be broken.
Objectives: To compare the properties of solubility, conductivity, and melting point for
six compounds and to classify these compounds as ionic or covalent on the basis of these
properties.
Materials:
24-well plate
capillary tubes
melt-temp apparatus
stirring rod
conductivity apparatus
Prelab:
Use the CRC Handbook, the Merck Index, your textbook, or the internet to
find the chemical formulas for each of the six compounds to be tested.
Name
oxalic acid
iron (II) sulfate
potassium iodide
sodium chloride
p-dichlorobenzene
sucrose
cyclohexane in dropper
distilled water bottles
Formula
Oxalic acid
Iron (II) sulfate
1,4-dichlorobenzene
potassium iodide
Sodium chloride
Sucrose
33
Safety Precaution: While some of these chemicals are common and harmless, some of
them are toxic and/or quite reactive. Therefore, do not touch any of them, and wash your
hands afterwards as a precaution.
Procedure:
Part 1 – Solubility Tests
1. Place the plastic well-plate on the 24 well-plate guide.
2. Half fill the first row of wells with distilled water.
3. Use a wood-splint to place a small amount of each chemical into its
designated well, and stir with a plastic stirrer.
4. Record your observations, and note if the chemical dissolves completely,
mostly, slightly, or not at all. Do not discard yet, set the well-plate aside for
part 2.
5. Repeat this procedure using the ceramic well-plate and the cyclohexane as a
solvent, and record your observations. [Note: The cyclohexane will ruin the
plastic well-plates, so be sure to use the ceramic plates!]
6. Use a plastic Beral pipet to remove the cyclohexane mixtures from the wells
and place them in the organic waste container.
Part 2 – Conductivity Tests
1. Test the conductivity of each compound in the wells containing the distilled
water only. Rinse the prongs with distilled water and dry them with a paper
towel after each test.
2. Record your observations, and note if the chemical is a strong conductor
(bright light), a weak conductor (dim light), or a nonconductor.
3. Discard the distilled water mixtures in the aqueous waste container.
Part 3 – Melting Point Tests
1. Six mortars and pestles are set out for each of the six compounds.
2. Grind up a small amount of each compound using the designated mortar and
pestle.
3. Place a small amount of the compound into a capillary tube.
4. Four Mel-Temp’s are stationed around the lab. Each one is set at a specific
setting. Starting with the one set at the lowest setting, place the capillary tube
into the apparatus and watch the substance for about two minutes. If the
compound melts, then record the melting point for that substance as less than
or equal to the set temperature.
5. If the compound does not melt, then remove the capillary tube and place it in
the apparatus at the next highest setting, and repeat the procedure.
6. If the compound does not melt after a couple of minutes at the highest
temperature, then record its melting point as above that temperature.
7. Repeat this procedure for the remaining compounds
34
Data:
Substance
Solubility in
water
Solubility in
cyclohexane
Electrical
conductivity in
H2 O
Melting point
(oC)
Oxalic acid
Iron (II)sulfate
p-dichlorobenzene
Potassium
iodide
Sodium
chloride
sucrose
Results :
Summarize the results of the tests for each chemical:
Oxalic
Acid
iron(II)
sulfate
p-dichlorobenzene
potassium sodium
iodide chloride
sucrose
Solubility*
Conductivity#
Melting point+
*For solubility results indicate whether the compound was soluble in only water, only cyclohexane, both
solvents or neither solvent.
#For conductivity results indicate whether the compound was very conductive, slightly conductive, or nonconductive.
+For melting point results indicate whether the compound melted at a high temperature
(T >350oC), a moderate temperature (350oC >T > 200oC), or a low temperature (T < 200oC).
35
Analysis:
1.
Based on the results of your tests, classify each of the six chemicals as either
possessing ionic or covalent bonds.
2.
Summarize how the properties of ionic compounds differ from the properties of
covalent compounds.
3.
Which of the properties that you tested does not seem to give a clear-cut
difference between ionic and covalent compounds? Suggest a possible reason for
why this ambiguity exists.
4.
Write a paragraph describing the difference between how ionic and covalent
bonds are formed, and explain how these differences may account for the physical
properties of ionic and covalent compounds.
Ref: Chemistry Text, Myers
Chapter 5 pp.170-173
Chapter 6 pp.197-198
Grading Scale:
Prelab
Data / Results
Lab Technique
Analysis:
1. correct classification
2. summary of properties
3. not clear-cut/reason
4. conclusion
Total
(3)
(6)
(5)
____
____
____
(3)
(2)
(2)
(4)
(25)
____
____
____
____
____
Science, at bottom, is really anti-intellectual. It always
distrusts pure reason, and demands the production of
objective fact. H.L. Mencken
36
Lab #12 Polar and Nonpolar Molecules
Model Kit Lab
Introduction:
At this point in your study of chemical bonding, it is important that you learn
how to make the distinction between polar / nonpolar bonds vs polar / nonpolar
molecules. A polar bond is formed between two atoms if they do not share electrons
equally. However, in order to determine if a molecule is polar or nonpolar, you must
consider the shape of the molecule in addition to the polarity of the bonds that make up
the molecule. Some molecules that contain polar bonds are also polar molecules, but it is
possible for a molecule to be nonpolar even though it contains polar bonds. In general,
polar molecules will have an asymmetrical shape and nonpolar molecules tend to be
symmetrical. (See your note packet).
Pre-Lab Study:
Refer to: Chapter 6, Sec. 1 pp. 194-198 in your textbook (Myers)
Chapter 4, Sec. 4.4 pp. 106-112 in Miller/Lygre
1-4 Define the following terms:
1.
electronegativity –
2.
polar covalent bond –
3.
nonpolar covalent bond –
4.
dipole –
5.
If a molecule has covalent bonds between different atoms, the shared electrons
are pulled toward the (more / less) electronegative atom.
6.
Use either your periodic table or figure 4.15 on p. 108 in Miller/Lygre to calculate
the electronegativity difference between the following pairs of atoms. Use the
chart on the back of this sheet to indicate the type of bond that would be formed
between these pairs.
Electronegativity
type of bond
Difference
a.
H, Cl
_____________
________________
b.
Na, Cl
_____________
________________
c.
C, O
_____________
________________
d.
H, O
_____________
________________
e.
C, H
_____________
________________
f.
N, O
_____________
________________
g.
N, H
_____________
________________
h.
Mg, O
_____________
________________
37
Electronegativity Difference and Bond Type
Electronegativity difference
type of bond
0 – 0.49
nonpolar covalent
(equal sharing of electrons)
0.50 – 2.1
polar covalent
(unequal sharing of electrons)
2.1 – 4.0
ionic
(transfer of electrons)
Procedure:
Use the model kit to build the following molecules. You must fill every hole in
each “atom” with a peg or spring when you build your molecules. Nitrogen is an
exception; you only need to use 3 of the holes. Use the following terms to describe the
shape of each molecule: linear, bent, trigonal planar, trigonal pyramid, tetrahedral,
ring, then determine whether each molecule is polar or nonpolar.
Data:
Shape
Polarity
1.
CH4
_________________
__________________
2.
CH3Cl
_________________
__________________
3.
O2
_________________
__________________
4.
CH3OH
_________________
__________________
5.
H2O
_________________
__________________
6.
HBr
_________________
__________________
7.
CO2
_________________
__________________
8.
N2
_________________
__________________
9.
C2H6
_________________
__________________
10.
CCl4
_________________
__________________
11.
C2H4O
_________________
__________________
12.
C6H6
_________________
__________________
38
Analysis Questions:
1.
How does the shape of a molecule affect its polarity?
2.
Why do water and carbon dioxide have different polarities? Draw a diagram to
help explain your answer.
3.
If four pairs of negatively charged electrons are located on the surface a sphere
they will repel each other to reach maximum separation. What is the angle
between each pair of electrons and the center of the sphere?
4.
Explain why CH3Cl is polar but CCl4 is nonpolar.
5.
What practical difference does it make whether a molecule is polar or nonpolar?
Give an example. See p. 111 in Miller/Lygre.
6.
Draw the Lewis-dot structures for molecule #’s 1, 3, 5, 7, 8 and 11.
Grading:
PreLab
Data
Analysis:
1.
2.
3.
4.
5.
shape
water/CO2
angle
CH3Cl v CCl4
practical diff.
(5)
(10)
____
____
(2)
(2)
(1)
(1)
(2)
____
____
____
____
____
39
6. dot structures
Total
(3)
(26)
____
____
Lab #13 Microscale Crystallization
Lab
Introduction:
Crystallization is usually the best technique for purifying a solid. Solubility
differences allow the separation of one type of molecule from another, or from various
contaminants. Usually the sample is dissolved in an appropriate hot solvent. As the
solvent cools, the solution becomes saturated with respect to the substance; further
cooling results in crystallization of the solute. In the crystallization process, molecules
gradually deposit from solution and attach to each other in an orderly array, or lattice. As
the deposit grows, it takes on a characteristic appearance, unique to that substance. The
crystal lattice has a high degree of symmetry which excludes molecules with different
geometries, size, or intermolecular force characteristics.
Impurities with different
shapes or sizes are excluded.
Molecules packing in orderly array.
The purity of the substance may be checked by its melting point. The key to
obtaining an accurate melting point is to raise the temperature slowly. Use the CRC or
Merck Index to find the melting point of the pure substance. Impurities usually lower the
melting point. Impurities may be solvent, water, by-products of reaction, or unreacted
starting material. Recrystallization helps eliminate these impurities.
Selecting the proper solvent is critical. The best solvent is one in which the
material is insoluble at room temperature but completely soluble when heated.
Remember the rule, “like dissolves like.” In the event that you cannot find a single
solvent to use the use of solvent pairs may work.
40
Note: Be sure to include a detailed drawing or photograph of all the crystals that you
obtain in this lab.
Theory guides. Experiment decides.
An old saying in science, seen attributed to many different persons
Pre-Lab Assignment:
Use the CRC, Merck Index or other references to prepare a table of solubilities,
melting points, molecular and structural formulas, and crystalline forms for acetanilide,
benzoic acid, caffeine, citric acid, glycine, malonic acid, salicylic acid, and urea.
Read over the information sheets carefully. You will need to these throughout the
experiment.
Procedure:
Note: Keep a detailed record of your procedure and data and observations on
separate paper.
A.
Crystallization of Benzoic Acid
Prepare a hot sand bath by half-filling a 250 mL beaker with sand. Place a wire
gauze on the hot plate and place the beaker on top. Turn the hot plate to setting 5 and
allow the sand to heat up while you are Place 0.050 g of benzoic acid in a medium test
tube. Add approximately 10 drops of water to the test tube. Gently heat the solution to
boiling on a hot sand bath. The deeper the tube is placed in the sand, the hotter it will be.
As soon as boiling begins, continue to add water drop-wise until all the solid just
dissolves. Then remove the tube from the sand bath and cork the tube and clamp it; as it
cools, observe the phenomenon of crystallization.
After the tube reaches room temperature, cool it in ice, stir the crystals with a
Pasteur pipette. Expel the air from the pipette as the tip is pushed to the bottom of the
tube. When the tip is firmly and squarely seated in the bottom of the tube, release the
bulb and withdraw the water. (See figure 1) Rap the tube sharply on a wood surface to
compress the crystals and remove as much of the water as possible with the pipette.
Using a stainless steel spatula, scrape the crystals onto a piece of filter paper, fold the
paper over the crystals, and squeeze out excess water before allowing the crystals to dry
to constant mass. Mass the dry crystals and calculate the percent recovery of product.
Determine the melting point with the Mel-Temp apparatus.
B.
Crystallization of An Unknown
Obtain a sample of an unknown crystal from your teacher. Follow the procedure
for selecting a proper solvent to determine the best solvent to use for recrystallizing your
unknown. You do not need to mass out precise amounts of your unknown when you are
simply trying to determine the best solvent to use. Once you have determined which
solvent to use, then dissolve another 0.050 g sample of your assigned unknown in the
minimum volume of the appropriate solvent. Heat the solvent in the reaction tube until
the unknown crystals completely dissolves. Once the crystals have dissolved, then
41
remove the tube from the sand bath, cork it and place it in the ice bath. Observe the
crystallization process. Dry the crystals as before and determine the melting point with
the Mel-Temp apparatus. You must submit your purified crystals in a clean, labeled glass
vial to Mr. Gallagher for grading.
Analysis:
Part A
1. Determine the melting point of your benzoic acid crystals.
2. Calculate the percent error for the melting point of your benzoic acid crystals.
3. Calculate the percent recovery of your benzoic acid crystals.
% recovery = mass of recovered crystals/ mass of original x 100
4. Describe the structure of your benzoic acid crystals.
Part B
4. Determine the melting point of your unknown crystals.
5. Find the compound from the list in the Pre-lab that has the closest melting point to
your unknown crystals and calculate the percent error for your crystals.
6. Determine the identity of your unknown crystals.
Grading Rubric
Pre-Lab Assignment
Laboratory technique
Data and observations
Percent error mp BA
Percent recovery
Crystal description
Percent error mp unknown
Identity of unknown
Submission of crystals
Total
____
____
____
____
____
____
____
____
____
(10)
(10)
(5)
(2)
(2)
(2)
(2)
(5)
(4)
____
(42)
42
Lab #14
Pyrolysis of Wood
Introduction:
Everyone knows that wood will burn. Burning is a process that we call
combustion. Combustion is a process in which a substance is combined with oxygen to
produce new compounds and heat is typically released as a by-product. In some cases,
such as those involving metals, the combustion process will produce more complex
substances. For example, the elemental metal magnesium will be oxidized to produce the
compound magnesium oxide. However, many organic compounds, such as those found
in wood, also burn but the products of these combustion reactions result in the formation
of smaller and simpler substances. This occurs because the combustion process also
involves some sort of thermal decomposition in which heat is used to break apart large
complex molecules into smaller simpler ones that will readily combine with oxygen.
This process of thermal decomposition is known as pyrolysis. Is it possible to perform
pyrolysis without combustion? What if you were to remove or limit the amount of
oxygen that is available while heating a substance? What happens if you heat wood
without actually burning it? What do you think will happen?
Objectives:
1. To investigate the differences between direct combustion and pyrolysis.
2. To discover what substances are produced from the pyrolysis of wood.
Hypothesis:
1. Predict how you think the results of burning wood directly will differ from
simply heating the wood without catching the wood on fire.
2. Predict which process you think will produce the most residue and state your
reasons for your answer.
Pyrolysis vs. Direct Combustion
Procedure:
1.
2.
3.
4.
5.
Obtain a large test tube and attach it to the ring stand with a utility clamp.
Adjust the clamp so that the test tube is set at a 45o angle.
Obtain 3 or 4 wooden splints and break them into thirds.
Record the mass of the wood and place the pieces in the bottom of the large test tube.
Adjust the height of the test tube so that it is about 5 or 6 cm above the top of the
Bunsen burner.
43
6. Light the Bunsen burner and heat the bottom of the test tube with a #3 flame. Move
the burner periodically to ensure that the heat is distributed evenly to all of the wood
pieces.
7. Record your observations.
8. After you have been heating for a minute or so, bring a lit match near the mouth of
the test tube. If the vapors do not ignite, then try again.
9. Continue heating the wood until the vapors no longer burn.
10. Allow the set-up to cool and then transfer the charcoal into a pre-massed crucible.
11. Determine the mass of the remaining charcoal.
12. Place the crucible containing the charcoal on a clay triangle that is resting on a small
iron ring attached to the ring stand.
13. Heat the crucible strongly with a #3 flame until all of the charcoal is consumed.
Record your observations. Allow the crucible to cool and then mass the crucible and
its contents to determine the mass of the remaining ash.
14. Obtain a single wood splint and record its mass.
15. Obtain a watch glass and record its mass.
16. Hold the wood splint over the watch glass using crucible tongs and light the wood
splint. Try to burn the wood splint as completely as possible. Allow the remaining
ash to fall onto the watch glass. Relight any unburned portions of the wood splint if
necessary and collect all of the ash. Record your observations.
17. Record the mass of the ash.
Teacher Demonstration:
1. Your instructor will conduct an experiment similar to your first procedure. However,
instead of burning the vapors, we will attempt to collect them by condensation.
2. Draw a diagram of your instructor’s experimental set-up.
3. Record your observations as your instructor performs the experiment.
Analysis:
1. Calculate the percent of charcoal that remains after the pyrolysis process.
2. Calculate the percent of ash that remains after the pyrolysis process.
3. Charcoal is primarily composed of carbon. Approximately what percent of the wood
do you think is made of carbon? Is it the same as the percent of charcoal, or is it
more or less. Explain your answer.
4. How does the ash differ from the charcoal? What do you think the ash is made up of?
5. Calculate the percent of ash that remains after the direct combustion process.
6. In what ways does pyrolysis differ from direct combustion? Base your answer at
least partly on your results and observations.
7. What did you learn about some of the products of the pyrolysis of wood from your
teacher’s experiment?
Conclusion:
1. Evaluate the hypotheses that you made at the beginning of this experiment.
44
2. Summarize what you learned from this experiment.
Scientific principles and laws do not lie on the surface of nature. They are hidden, and
must be wrested from nature by an active and elaborate technique of inquiry.
John
Dewey
Chem I
Lab # 15 Formula of a Hydrate
Introduction:
Some compounds are hygroscopic, which means that they have the ability to
absorb water. When a hygroscopic compound is dissolved in water and then the water
evaporates, the remaining crystal contains some of the water and is known as a hydrate.
For example, copper (II) sulfate readily dissolves in water to form a blue solution. When
the water is evaporated, some of the water molecules remain “trapped” as part of the
crystalline structure, and they do so in a specific ratio. The formula of this hydrate would
be written as CuSO4 . X H2O where X represents the ratio of water molecules to one
CuSO4 formula unit. The objective of this lab is to determine the specific formula of the
hydrate of copper (II) sulfate. You will do this by massing the hydrate and then heating it
vigorously to drive off the water molecules that are incorporated within the crystalline
structure. This process will leave behind the “dry” copper (II) sulfate, known as
anhydrous copper (II) sulfate.
Procedure:
Your teacher will demonstrate the proper handling of crucibles.
1. Obtain a crucible and matching lid. Clean if necessary.
2. Place the crucible on the clay triangle with the lid slightly ajar. Heat with a #3 flame
for 3 to 4 minutes, gently at first, then strongly, to drive off any water that is adsorbed
to the crucible.
3. Allow the crucible and lid to cool on the triangle. When cool to the touch, mass the
crucible and lid to the nearest milligram. Record the mass of the empty crucible and
lid in your data table.
4. Transfer approximately 3 grams of copper (II) sulfate hydrate to the crucible. Mass
to the nearest milligram and record the mass of crucible + lid + hydrate. Describe
the original appearance of the compound.
5. Calculate and record the mass of copper (II) sulfate hydrate before heating.
6. Replace the crucible on the clay triangle. With the lid slightly ajar as before, heat the
contents gently at first, then strongly for 8 to 10 minutes.
7. Slide the lid to completely close the crucible and allow it to cool. When cool to the
touch, mass the crucible + lid + anhydrous copper (II) sulfate to the nearest
milligram. Calculate and record the mass of the anhydrous copper (II) sulfate.
45
8. Calculate and record the mass of water lost. Describe the appearance of the
compound after heating.
9. After massing the anhydrous compound, carefully transfer it to a clean test tube. Add
about 20 drops of distilled water. Observe (visually and by touch) the behavior of the
mixture. You may dispose of this solution down the drain with lots of water.
Data:
Mass Data for Copper (II) Sulfate
Mass of crucible + lid + hydrate (before)
Mass of empty crucible + lid
Mass of hydrate
Mass of crucible + lid + anhydrous CuSO4
Mass of anhydrous copper (II) sulfate
Mass of water lost
(g)
(g)
(g)
(g)
(g)
(g)
Observations:
Initial appearance before heating:
Final appearance after heating:
Results of water addition to anhydrous crystals:
Analysis: (Show all of your work including units for all calculations.)
1. Calculate the number of moles of anhydrous copper (II) sulfate.
2. Calculate the number of moles of water lost.
3. Calculate the ratio of moles of water to moles of anhydrous copper (II) sulfate.
4. Write the formula you determined for this hydrate.
5. Ask the teacher about the correct formula and then calculate your percent error.
% error = calculated value – actual value x 100
actual value
6. Discuss at least three specific sources of experimental error that could have affected
your results.
7. Propose a hypothesis to explain your observations in step 9 of the procedure. Think
in terms of what was required to create the anhydrous compound and what you did to
re-create the hydrate.
46
Grading Rubric:
Data and lab technique
Observations (quality)
Calcualtions (1-3)
Accuracy of formula
(8)
(4)
(3)
(4)
___
___
___
___
% error calculation
Sources of error
Hypothesis step 9
Total
(1)
(2)
(2)
(24)
___
___
___
___
Chemical Reactions Unit
Lab # 16 Types of Chemical Reactions Lab
Verifying Products
Introduction:
In class you are learning about the different types of chemical
reactions, and now is your chance to see them in action. The objective of this lab is for
you to conduct two successful reactions for four of the five types. You will conduct the
experiments while making careful observations. If a reaction is successful, then you will
have to verify that your predicted reaction did indeed take place by matching your
observations with the known properties of the chemicals that you claim have been
produced. This will require that you look up the properties of the suspected products in
the CRC Handbook or the Merck Index, or you can use the reference charts provided at
the end of this lab handout.
Pre-Lab Assignment: (See page 7)
1. Give the general patterns, i.e., (A + B  C + D) for the following reaction types.
Composition
________________________________________
Complete Combustion
Decomposition
__________________________________
________________________________________
Double Replacement
__________________________________
Single Replacement
__________________________________
2. What determines if a single replacement reaction will occur spontaneously or not?
3. What determines if a double replacement reaction will occur spontaneously or not?
4. What are the seven diatomic elemental molecules?
47
5. How do you test for the presence of hydrogen gas?
6. How do you test for the presence of oxygen gas?
7. Which reagent provided can act as a catalyst for certain reactions?
Reagent List:
Metals: Cu, Mg, Zn
Aqueous solutions:
Solids: MnO2
Liquids: H2O2, H2O
AgNO3, BaCl2, CuSO4, NaCl, Na2CO3, HCl, H2SO4, Na2SO4
Safety Precautions: You must wear your safety goggles at all times!!!
Procedure:
Composition 1
1. Obtain a strip of copper metal and hold it in the hottest part of a #3 Bunsen burner
flame with crucible tongs for about 1 minute.
2. Remove the copper from the flame and allow it to cool in the air.
3. Record your observations.
4. Reheat the copper strip in the flame and then cool it by placing it in a beaker of tap
water. Record your observations.
5. Write the balanced equation for the reaction that you think has occurred.
6. Provide physical evidence to verify that your predicted reaction did in fact take place.
Refer to pages 7 and 8 of this lab for all reactions.
Composition 2
1. Obtain a 5 cm strip of magnesium metal.
2. Using crucible tongs, hold the magnesium in a #3 Bunsen burner flame until it
ignites. (Caution: Do not stare directly at the flame!)
3. Record your observations and write the balanced equation for this reaction.
4. Provide evidence to verify this reaction. See pages 7 and 8 of this lab.
Observations for Composition Reactions:
1.
48
2.
Single Replacement 1
1. Place about 5 mL of the copper (II) sulfate solution into a medium test tube.
2. Place a piece of zinc metal into the solution and observe for approximately 5 minutes.
3. Write the balanced equation for the reaction that occurred and provide evidence to
verify your prediction.
Single Replacement 2
1. Obtain a small test tube to serve as a reaction vessel and a Beral pipet bulb to serve as
a gas collection chamber. See figure 1 below.
2. Completely fill the gas collection chamber with water.
3. Fill the test tube (reaction vessel) 1/3 full with 1 M hydrochloric acid (HCl).
4. Fold a 3-cm strip of magnesium metal, and place it into the test tube.
5. Immediately attach the cork with the plastic tube on the reaction vessel and place the
collection chamber over the tube to begin collecting gas.
6. Stop collecting gas when the collection chamber is 7/8 full of gas.
7. Remove the collection chamber but keep the opening of the chamber pointed
downwards to prevent the gas from escaping.
8. Squeeze the collected gas into the flame of a Bunsen burner and record your
observations. If nothing happens, then repeat the procedure.
Observations for Single Replacement Reactions:
1.
2.
Figure 1
Gas Collection Reaction Vessel
49
Double Replacement 1
1.
2.
3.
4.
5.
Place about 10 drops of the silver nitrate solution into a well plate.
Add 10 drops of the sodium chloride solution to the same well.
Record your observations and write the balanced equation for the predicted reaction.
Supply evidence from your observations that your predicted reaction occurred.
Dispose of the reaction mixture in the aqueous waste container.
Double Replacement 2
1.
2.
3.
4.
5.
Place about 10 drops of the barium chloride solution into a well plate.
Add 10 drops of the sodium sulfate solution to the same well.
Record your observations and write the balanced equation for the predicted reaction.
Supply evidence from your observations that your predicted reaction occurred.
Dispose of the reaction mixture in the aqueous waste container.
Observations for Double Replacement Reactions:
1.
2.
Decomposition 1
1. Obtain a small test tube to serve as a reaction vessel and a Beral pipet bulb to serve as
a gas collection chamber. See figure 1 on p. 3.
2. Completely fill the gas collection chamber with water.
3. Half fill the test tube (reaction vessel) with hydrogen peroxide.
4. Place a small scoop of MnO2 into the test tube.
50
5. Immediately attach the cork with the plastic tube on the reaction vessel and place the
collection chamber over the tube to begin collecting gas.
6. Stop collecting gas when the collection chamber is 7/8 full of gas.
7. Remove the collection chamber but keep the opening of the chamber pointed
downwards to prevent the gas from escaping.
8. Test the gas by lighting a wood splint, allow it to burn for several seconds and then
blow it out.
9. Immediately squeeze a puff of the collected gas toward the glowing splint. Record
your observations. If nothing happens, then repeat the procedure.
Decomposition 2
1. Obtain an electrolysis apparatus as described by your teacher. See Figure 2 below.
2. Fill the reservoir ¾ full with tap water and then add about 10 mL of 3M sulfuric acid
solution (H2SO4) to the reservoir and stir to mix.
3. Fill two medium test tubes with water from the reservoir.
4. Invert the test tubes into the reservoir without allowing any air to fill the test tubes.
5. Attach the test tubes to the electrolysis apparatus and connect the wires to the power
source provided.
6. Turn on the power source and record your observations. You may need to look
closely to see the tiny gas bubbles that are forming. Allow the apparatus to run long
enough to collect at least 5 cm of gas in one of the test tubes.
7. Record your observations. Pay particular attention to the relative amounts of gas that
is collected in each test tube.
Observations for Decomposition Reactions:
1.
2.
Figure 2
Electrolysis Apparatus
51
Composition Reactions:
Predicted Balanced Equation
Spontaneous
(Y / N)
Evidence
Spontaneous
(Y / N)
Evidence/Source
Spontaneous
(Y / N)
Evidence/Source
1.
2.
Single Replacement Reactions:
Predicted Balanced Equation
1.
2.
Double Replacement
Predicted Balanced Equation
1.
2.
Decomposition Reactions:
52
Predicted Balanced Equation
Spontaneous
(Y / N)
Evidence/Source
1.
2.
Additional Notes: You must read this before beginning the lab!
1. Record detailed observations for each reaction. Remember that you must verify that
the substances that you predicted to form were actually produced.
2. Small amounts of MnO2 can be used as a catalyst for a decomposition reaction
involving hydrogen peroxide. A catalyst is not used up during a chemical reaction
and therefore, it is not listed as a reactant but it is written above the arrow in a
chemical equation.
3. To test for the presence of oxygen gas, light a wooden split and blow it out. Quickly
place the splint near the source of the gas. If oxygen is present in sufficient quantity,
the splint will re-ignite or at least glow more brightly.
4. To test for the presence of hydrogen gas, bring a burning wooden splint or Bunsen
burner flame near the source of the gas. If hydrogen is present in sufficient quantity,
you will hear a sudden pop.
5. Remember that seven elements form diatomic molecules when they occur by
themselves. (H2, N2, O2, F2, Cl2, Br2, I2)
6. Single replacement reactions will occur spontaneously if the lone metal reactant is
more active than the metal ion present in the compound.
7. Double replacement reactions will be spontaneous if either of the following occurs:
- the formation of an insoluble precipitate.
- the formation of a stable covalent compound such as H2O or CO2
8. The reaction of an element with oxygen gas can often occur if the element is heated
strongly in a #3 flame and then removed from the flame and exposed to the air.
9. A chemical reaction is considered to be spontaneous in this lab if it occurs to a
significant extent at room temperature without the input of additional energy.
Physical Properties of Selected Substances:
Name
Copper (II) hydroxide
Copper (I) oxide
Copper (II) oxide
Copper (I) sulfide
color
solubility in water
blue gel, or light blue crystals
yellow, red, or brown powder
black, to brownish black powder
blue, to grayish black powder
insoluble
insoluble
insoluble
insoluble
53
Copper (II) sulfide
Magnesium hydroxide
Magnesium oxide
Zinc sulfide
black powder
colorless crystals
very fine white powder
grayish-white to yellowish powder
insoluble
insoluble
insoluble
insoluble
Types of Chemical Reactions Lab
Grading Rubric
Pre-Lab Assignment
Observations (quality)
Predicted Reactions
Evidence (verification)
Lab Technique
Clean-up
Total
(5)
(12)
(20)
(8)
(10)
(5)
____
____
____
____
____
____
(60)
____
54
Lab # 17 The Iron Chemist
Objective - To produce 0.750 grams of iron (III) oxide from iron (III) chloride
hexahydrate.
Reactions:
FeCl3
Fe(OH)3
NH4OH  Fe(OH)3
+

Fe2O3
+
+
NH4Cl
H2O
Materials: Iron (III) chloride hexahydrate, 250 mL beaker, 3.0 M HCl, 3.0 M
NH4OH, distilled water, filter paper, funnel, clay triangle, crucible, crucible tongs, and
Bunsen burner.
Procedure:
1. Based on the reactions, determine the mass of iron chloride hexahydrate needed to
produce 0.750 grams of iron (III) oxide.
2. Dissolve the iron compound in a minimal amount of 3.0 M HCl.
3. This is a very important step! Add 3.0 M NH4OH to the solution until it turns
basic. Test with red litmus paper. Stir thoroughly and let the solution sit for a
minute and test with litmus again. If it is not blue add several drops of 7 M
NH4OH.
4. Gently boil the solution for approximately five minutes.
5. Fold a piece of filter paper and place in a funnel. Transfer all of the precipitate to
the filter paper by washing with distilled water and filter. Make sure to transfer
all of the precipitate to the filter paper. The effluent should be colorless and clear.
If is not then go back to step 3.
6. Mass a clean, dry crucible. Record the mass.
7. Remove the filter paper from the funnel and gently squeeze out the water.
8. Place the wet filter into the crucible and gently heat to char the paper. Once the
paper has begun to burn, place the lid on the crucible and increase the heat.
Oxygen can be added by gently and carefully blowing into the crucible.
9. Once the paper has completely burned, heat strongly for 10 more minutes.
10. Place the crucible on a wire gauze and allow to cool.
11. Once cool, mass the crucible and its contents. Record the mass.
12. Determine the mass of the iron (III) oxide by subtracting the mass of the crucible.
55
Observations:
Analysis
1. Calculate the percent yield.
Experimental mass of Fe2O3
---------------------------------- x 100%
Theoretical mass of Fe2O3
= % yield
2. What are some possible sources of error in this lab? Be specific and describe how
these errors would affect you results.
3. What chemical reaction is happening in procedure step 10 as the contents of the
crucible are cooling? Refer to your observations to help answer this question.
Conclusion
Write a short paragraph that addresses the following;
Did you meet your objective for the lab? How was this accomplished?
What types of reactions did you observe?
What mass of iron (III) oxide did you produce?
What was your percent yield? How can you improve your yield?
56
Lab # 18 Copper Cycle Lab
Introduction:
In this experiment you will take a sample of pure copper, follow it through a
series of reactions, and recover the copper in the end. You will need to record your
observations in detail as you attempt to identify the products as well as the type of
reaction that occurred at each step. You will then determine the percent of the original
copper that you recovered.
Safety:
You must wear safety goggles at all times while in the laboratory. Several
corrosive and dangerous chemicals will be used. Handle concentrated nitric acid
(HNO3), 3M sodium hydroxide (NaOH), 3M sulfuric acid (H2SO4), and 3M hydrochloric
acid (HCl) with care. If you spill any chemicals on your hands, wash them thoroughly.
The gas produced when copper reacts with nitric acid is toxic – perform this reaction
under the fume hood. Read the procedure for each section through completely before
beginning that section!
Procedure:
Part A
Disappearing Copper
1. Obtain a sheet of copper and mass it to the nearest milligram.
2. Place the copper in a clean 250 mL beaker. Label the beaker with your initials.
3. UNDER DIRECT SUPERVISION OF THE TEACHER, place the beaker in
the fume hood and add 10 mL of concentrated nitric acid. DO NOT BREATHE
THE FUMES! Leave the beaker in the hood until the reaction is complete.
4. Record all your observations.
Part B
Basic Blue Goo
1. Be aware that the contents of the beaker are still very acidic!
2. Set the beaker in an ice bath.
3. Very slowly, add about 20 mL of 3M NaOH to the beaker, stirring constantly.
4. Use a glass stirring rod to place a small drop of the liquid from the beaker
onto a piece of red litmus paper. (Do not place the litmus paper directly on
the lab bench.) If the litmus paper does not turn dark blue, then continue to
add NaOH several drops at a time until the litmus paper turns dark blue.
Note: Some of the light blue precipitate may stick to the litmus paper. This
does not mean that the litmus paper has turned blue. A positive result can be
confirmed when the paper that absorbs the liquid has turned dark blue.
Part C
Muddying the Water
(Safety Note: Potential Explosion Hazard!)
57
1. Set up a ring stand to hold the beaker over the Bunsen burner. Use the small
iron ring to hold the wire gauze on which you will place the beaker. Use the
large iron ring to place around the beaker to keep it from tipping over.
2. Use a small, light blue flame to gently heat the contents of the beaker.
3. Continue gentle heating until the reaction is complete. Aggressive heating
will cause the contents to spatter out of the beaker in a violent manner. You
should periodically remove the Bunsen burner from beneath the beaker to
slow the heating process.
Part D
Clearing Things Up
1. Allow the product from part C to settle.
2. Carefully decant and discard the clear portion down the drain with lots of
water. Be careful not to lose any of the solid.
3. Add 25 mL of 3M H2SO4.
4. Stir until the reaction is complete. The change is very obvious.
Part E
Completing the Cycle
1. Add a few pieces of mossy zinc to the product of part D. Do not breathe the
fumes given off during this reaction. The gas is a product of the reaction
between the zinc and the excess sulfuric acid from part D.
2. This reaction is complete when all of the blue color has been removed from
the solution. If the reaction appears to stop before it is complete, your teacher
will add some powdered zinc to the beaker to speed up the reaction.
3. This reaction will produce a lot of heat which may convert the free copper
back to an oxide, an undesirable result. Therefore, as the reaction nears
completion, place the beaker in a cool water bath.
4. After the reaction is complete, remove any large chunks of un-reacted zinc.
Before you remove the zinc from the beaker, use a distilled water bottle to
squirt off any copper that is sticking to the zinc back into the beaker.
Part F
Mopping Up
1. Add 20 mL of 3M HCl to the beaker. The purpose of the HCl is to use up the
excess zinc that could not be removed by hand. The HCl reacts readily with
zinc, but not with copper.
2. When the reaction is complete (no more bubbles), decant the liquid carefully,
so as not to lose any of the copper.
3. Rinse the product with distilled water at least 3 times, decanting each time.
4. Mass a piece of filter paper, fold it, and place it in a plastic funnel. Wet the
filter paper to keep it in place. Place the funnel in a 250 mL Erlenmeyer flask.
5. Add about 50 mL of distilled water to the copper in the beaker. Swirl the
beaker to stir up the copper, and carefully pour it into the funnel.
6. Add more water to the copper as needed to transfer it to the funnel.
7. Use the rubber policeman to transfer any remaining copper to the filter paper.
8. After the water has dripped through, carefully remove the filter paper, unfold
it, place it on a paper towel, and allow it to dry overnight.
9. Mass the dry copper and filter paper.
Data and Observations:
58
Table 1 - Mass Data:
Mass of copper metal (part A)
Mass of filter paper + recovered copper
Mass of filter paper
Mass of recovered copper (part F)
Percent recovery of impure copper
Part A
Balanced
Equation
Chemical
name
+

________
________
________
________
________
+
+
State of
matter
(s, l, g, aq)
Grams
Molar mass
Moles
Detailed
Observations:
Type of reaction: _________________________ + _________________________
(combination of 2 types)
Calculations: Show a sample calculation for each type of calculation. Show all of your
work, including units and significant digits.
59
Part B
Balanced
Equation
Chemical name

+
+
State of matter
(s, l, g, aq)
Grams
Molar mass
Moles
Detailed
Observations
Type of reaction: _______________________
Part C
Balanced Equation

+
Chemical name
State of matter
(s, l, g, aq)
Grams
Molar mass
Moles
60
Detailed
Observations:
Type of reaction: ________________________
Part D
Balanced
Equation
+

+

+
Chemical name
State of matter
Grams
Molar mass
Moles
Detailed
Observations:
Type of reaction: ___________________________
Part E
Balanced
Equation
+
Chemical name
State of matter
Grams
Molar mass
Moles
61
Detailed
Observations:
Type of reaction: ___________________________
Part F
Balanced
Equation
+

+
Chemical name
State of matter
Observations:
Type of reaction: ____________________________
Analysis:
1. Show a sample calculation for each of the different calculations that you had to
perform. Be sure to include all of the proper units and report your answers to the
proper number of significant digits.
2.
Calculate the percent yield for the recovered copper.
% yield = (mass of recovered copper/mass of original copper) x 100
3. Evaluate the relative purity of your recovered copper and discuss at least three
specific substances that are likely candidates that could be contaminating your
recovered copper, and describe at which part of the lab these contaminants originated.
4. Create a detailed diagram showing each step of this experiment. Be sure to illustrate
the cyclic nature of the entire process. This diagram should be drawn neatly on a
separate sheet of paper / poster, and it should be done in color.
Grading Rubric:
Mass Data
(2)
____
62
Balanced equations
Chemical names
Calculations
Observations
Percent yield
Purity evaluation
Diagram*
Total
(6)
(3)
(15)
(9)
(2)
(3)
(10)
(50)
____
____
____
____
____
____
____
____
Lab # 19 Serial Dilutions
Introduction:
Making accurate dilutions of an original stock solution is a very
important laboratory skill. Many systematic errors that occur during the course of an
experiment occur while making dilutions. Since the original solution is usually much
more concentrated than the diluted solution you are making, any excess from the original
solution will have a profound impact on the concentration of the new solution.
Therefore, it is imperative that you work carefully and accurately. This is especially true
when you are making “serial dilutions”. It is a good practice to use a new pipette each
time you are making a new dilution. If this is not possible, then you should thoroughly
rinse the pipette several times with distilled water. Also, note that if you dilute a solution
in a volumetric flask above the graduation line, it is impossible to reverse that mistake by
removing the excess solution, since some mixing has inevitably occurred.
Objectives: You will make a stock solution of known concentration*. You will then
perform a series of dilutions of this solution as accurately as possible. You will be
graded on technique and accuracy.
Pre-Lab:
Read step 1 of the procedure, and calculate the mass of KMnO4 that you will need
to weigh in order to create a 0.0150 M solution using a 50.00 mL volumetric flask. You
must show the instructor your calculation before you may begin the lab.
Procedure:
1. Make a 0.01500 M solution of KMnO4. (Toxic! Do not touch!)
a. You will first need to calculate the mass of KMnO4 that you will
need to place in a total volume of 50.00 mL to create a 0.01500 M
solution. Use a plastic weigh boat to mass this amount of solute.
63
2.
3.
4.
5.
b. Refer to the figure at the bottom of p. 463 in your textbook, Holt,
for the proper technique for making solutions.
c. This first solution is referred to a the Stock solution.
Make a ten-fold (1:10) dilution of stock solution.
a. Transfer the stock solution to a clean and dry beaker.
b. Use a 10.00 mL pipet to transfer 10.00 mL of the stock solution
from the beaker into a 100.00 mL volumetric flask.
c. Be certain that all of the solute is completely dissolved before
proceeding to the next step.
Make a 12.5-fold (1:12.5) dilution of this new solution.
a. Transfer the solution from the 100.00 mL volumetric flask to a
clean and dry 250 mL beaker.
b. Use a 5.00 mL pipet to transfer 2.00 mL of the solution from the
beaker into a 25.00 mL volumetric flask.
c. Be sure to shake the solution with the stopper on to ensure proper
mixing.
Place approximately 5 mL of this final dilution into a clean cuvette and
bring this solution to your instructor.
Record the absorbance from the Spec 20 of your final solution.
Calculations:
1. Create a standard concentration curve of Absorbance vs. [KMnO4]
from the standard data provided by your teacher. From your standard
curve, determine the molarity of your final solution.
2. Compare your actual molarity with the calculated molarity, and
calculate the percent error of your dilution.
3. Calculate the mass of KMnO4 that can be obtained upon the
evaporation of the water from your final solution based on the
concentration determined from the standard curve.
4. Discuss the specific sources of possible errors that may have
contributed to the inaccuracy of your results. Indicate whether these
sources would have caused the measured concentration to be too high
or too low.
[The accuracy of your solution will be determined spectrophotometrically.]
Grading:
Pre-Lab calculation
Lab technique and clean-up
Standard Concentration curve
Calculation of final molarity
Calculation of mass of KMnO4
Accuracy of final solution
Sources of error
Total
(2)
(6)
(3)
(2)
(2)
(10)
(5)
(30)
____
____
____
____
____
____
____
____
+/- 2% = 10/10
+/- 5% = 8/10
+/- 8% = 7/10
+/- 12%= 6/10
+/- 15%= 5/10
+/- 20%= 4/10
+/- 25%= 2/10
>25%=1/10
64
* When massing a solid solute to make a solution, you must determine if the solid is
anhydrous, a hydrate or if it is hygroscopic. If it is a hydrate, you must take into account
the additional mass from the water. If the solid is hygroscopic, such as sodium
hydroxide, you can only make a solution of approximate concentration. To determine the
concentration more precisely, you would need to standardize the solution by titrating it
with a solution of known concentration. KMnO4 is anhydrous.
Lab #20 Acid – Base Properties of Common Substances
Sample
Predict
A/B/N
Predict
Rank (1-16)
Red
Blue
Litmus Litmus
pH
paper
pH
meter
conducts
? Y/N/S
Rank
(1-16)
Ammonia
Baking soda
Bleach
Cocoa
Coffee
Distilled water
Milk
Orange juice
Rain water
Soda
Shampoo
Tap water
Tea
Vinegar
0.10M HCl
0.10M NaOH
Analysis Questions:
1. Which substance is the most acidic?
________________
2. Which substance is the most basic?
________________
3. Rank all 16 substances from most acidic (1) to most basic (16) to the left of the table.
4. What is the purpose of red litmus paper?
5. What is the purpose of blue litmus paper?
6. List at least 4 properties of acids.
7. List at least 4 properties of bases.
65
8. When you take the pH of something, what do you think it is actually measuring?
9. How is it possible for us to consume substances with such a wide range of pH’s?
10. What may account for the differences in properties between tap water, rain water,
and distilled water?
Lab #21 Acid – Base Titration
Objective: To determine the concentration of an acid.
Pre-Lab: Look up definitions for the following terms: (See pp. 550-555, Holt)
Titration –
End point –
Equivalence point –
Titrant –
Standard solution –
Safety Precautions: Wear your safety goggles at all times while in the lab!!
Procedure: [See pp. 552-553 in your textbook (Holt)]
1. Fill each buret with the appropriate solution. Open the valve to allow a small amount
of solution to flow out into a waste beaker. This will allow each solution to fill the tip
of the buret. Be sure to accurately record the starting volume of both the acid and the
base. Remember that a buret is calibrated to measure the volume of liquid that has
been dispensed.
2. Place approximately 20 mL of HCl into a clean 125 mL Erlenmeyer flask. Be sure to
record the precise starting and final volumes. Also record the precise concentration
of the NaOH.
3. Place several drops of phenolphthalein into the flask with the HCl and swirl the flask.
4. Begin adding NaOH to the flask approximately 1 mL at a time. Note the rate of the
color change. When the change starts to take longer to return to its original color,
begin adding base at a slower rate. Your goal is to add one drop of base that just
causes the solution to form a faint pink color. The color should persist for at least one
minute.
66
5. If you add too much base and the solution turns a dark pink color, you over shot the
equivalence point. Therefore, you will have to add more acid to return to the
equivalence point.
6. Record the final volumes of both the acid and the base added.
7. Repeat this titration two more times if time permits.
8. You may dispose of these solutions down the drain with lots of water.
Data:
Reported concentration of NaOH
__________ (M)
Part A Titration of HCl Solution
HCl
Trail 1
Trial 2
Trial 3
Final buret reading (mL)
Initial buret reading (mL)
Volume of HCl added (mL)
______
______
______
______
______
______
______
______
______
Final buret reading (mL)
Initial buret reading (mL)
Volume of NaOH added (mL)
______
______
______
______
______
______
______
______
______
Trail 1
Trial 2
Trial 3
Final buret reading (mL)
Initial buret reading (mL)
Volume of H2SO4 added (mL)
______
______
______
______
______
______
______
______
______
Final buret reading (mL)
Initial buret reading (mL)
Volume of NaOH added (mL)
______
______
______
______
______
______
______
______
______
NaOH
Part B Titration of H2SO4 –
Solution
H2SO4
NaOH
Analysis:
1. Calculate the average molarity of the HCl solution. Show all of you work, including
units. Your grade will be determined on how close you are to the actual
concentration of the HCl. Report your answer to the correct number of significant
digits.
67
2.
Calculate the average molarity of the H2SO4 solution. Show all of you work,
including units. Your grade will be determined on how close you are to the actual
concentration of the H2SO4. Report your answer to the correct number of significant
digits.
3.
If while performing the titration there was an unseen air bubble in the tip of the
buret, describe how your calculated molarity of the acid would differ from the actual
molarity. Explain.
Lab # 22 Heat of Combustion Lab
Objective:
You will investigate the relationship between the thermal energy released
when a fuel burns and its molecular structure.
Hypothesis:
Prior to beginning the experiment, predict which of the fuels that you
think will produce the highest heat of combustion and which will produce
the lowest heat of combustion,
Procedure:
Follow the same procedure for the following fuels:
Methanol, Ethanol, Kerosene, Biodiesel, and Paraffin Wax
1. Determine the mass of the fuel before burning. (For the candle, mass the
candle and the glass plate together.)
2. Hang a clean soda can on the iron ring attached to the ring stand as shown by
your instructor.
3. Adjust the height of the soda can so that the can is about 1 inch above the wick
of the fuel source.
4. Place 100.0 mL of cold tap water into the soda can.
5. Record the temperature of the water in the can to the nearest 0.2oC.
6. Light the fuel source and heat the water for 2 minutes. Gently stir the water as
it heats up and monitor the temperature. Do not allow the water to boil.
7. Extinguish the flame and record the highest temperature the water attains.
8. Re-mass the fuel and determine the mass of fuel that was burned.
9. Pour the water out of the can and clean off any soot from the bottom of the
can.
Data:
Fuel
Final H2O
Initial
o
Temp ( C) H2O Temp
(oC)
Temp
Change
(oC)
Final mass
of fuel (g)
Initial
mass of
fuel (g)
Mass of
fuel
burned (g)
Biodiesel
Ethanol
68
Kerosene
Methanol
Paraffin
Calculations:
Calorimetry a method in which you calculate the thermal energy released from a
combustion reaction by measuring the amount of heat absorbed by a given mass of water.
Water was used to capture the heat form the burning fuel because water has such a high
specific heat. The specific heat of the substance is defined as the amount of energy
required to raise one gram of that substance one degree Celsius. The specific heat of
water is 1.00 calorie/g . oC or 4.184 J/g . oC.
To calculate the amount of heat that the water absorbed you need to know three values:
1. The mass of the water that was heated. (Remember that the density of water is
1.00 g/mL, so 100.0 mL of water is the same as 100.0 g of water.)
2. The specific heat of water; you will be using the 4.184 j/g.oC value.
3. The change in the temperature of the water as it was heated.
The formula for calculating the amount of heat absorbed by the water is:
H = mw x Cp x T
Where, mw = mass of water, Cp = specific heat of water, T = change in temp of water
The units cancel to give the heat absorbed in units of Joules. It is typical to divide
the heat by 1000 to give your answer in kilojoules (kJ)
At this point you have simply calculated the amount of heat absorbed by the water. In
order to properly compare the different fuels that you tested, you need to calculate the
heat of combustion of each fuel in units of kilojoules per gram (kJ/g). To do this, simply
divide the heat absorbed by the water in each trial by the mass of the fuel burned. This
will allow you to determine which fuels released the most heat for every gram of fuel that
was burned.
Calculate the heat of combustion in kJ/g for each of the fuels that you tested. Be sure to
show all of your work, including the units and report your answers to the proper number
of significant digits.
Results for Heat of Combustion Calculations (kJ/g)
Biodeisel
Ethanol
Kerosene
Methanol
Paraffin
69
Show work here:
Questions:
1. Why was water used in this lab?
2. Why is it important that you do not allow the water to boil during this lab?
3. Rank the fuels from the lowest heat of combustion to the highest.
4. What other factors besides heat of combustion would you consider when choosing the
best overall fuel? (List at least 3 additional factors and explain your reasons.)
5. Discuss the assumptions that you made in calculating the heats of combustion and
explain how these assumptions may have affected your final results. Be specific!
Grading Rubric:
70
Lab Technique and Clean-up
Data
Calculations
Results
Questions
Total
Lab # 23
_____
_____
_____
_____
_____
(10)
(10)
(7)
(2)
(8)
_____ (37)
Radioactive Decay and Half-Lives Lab
Procedure
1. Spread out a few paper towels on your table and count the total number of M&M’s
provided. (This represents the original number of radioactive atoms.)
2. Place the M&M’s in the cup and then dump them onto the paper towels on the table.
3. Separate the candies with the M&M side up from the ones that are face down.
Record the number that are face up (decayed) and face down (undecayed). Place the
face down candies back into the cup.
4. Continue dumping the M&M’s and removing the “decayed atoms” until they are all
decayed. Each toss represents one half-life.
5. Convert the numbers of undecayed atoms to percentages in your data table.
6. Collect class data for the percentages of undecayed atoms vs. half-lives.
7. Create a graph of % undecayed atoms vs. half-lives. Use different colors to draw two
lines; one for your data and one for the class data.
Data:
HalfLives
(tosses)
Number
decayed
Number
undecayed
Percent
undecayed
Average
Percent
undecayed
class data
0
1
2
3
4
5
6
7
8
9
10
11
71
12
Percent Undecayed Atoms vs. Number of Half-Lives
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Analysis Questions:
1. Which line on your graph do you think more accurately represents the decay of a
radioactive isotope? Why?
2. If the half-life of M&Msium is 1268 years, how long would it take for the sample to
decay to 8.5% of its original amount? Use your graph to answer this question.
3. If you started with a 1000 gram sample of M&Msium, what mass would remain after
6974 years? Show your work.
4. In this simulation is there a way to predict when a particular “atom” will decay?
Explain.
5. In what ways do you think that this lab simulation is a good representation of
radioactive decay? In what ways do you think that this simulation differs from the
process of radioactive decay?
6. What other ways (not including coins or other candies) could you use to model the
concept of radioactive decay and half-lives? Explain how this would work.
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Grading Rubric:
Data
Graph
Questions
(5)
(5)
(10)
____
____
____
Total
(20)
____
74
75
Interpreting NFPA Hazard Identification System Labels
(Blue)
Health Hazard
Type of Possible Injury
0
Material that on exposure under fire conditions would
offer no hazard beyond that of ordinary combustible
material.
Example: peanut oil
1
Material that on exposure would cause irritation but
only minor residual injury.
Example: turpentine
2
Material that on intense or continued but not chronic
exposure could cause temporary incapacitation or
possible residual injury.
Example: ammonia gas
3
Material that on short exposure could cause serious
temporary or residual injury.
Example: chlorine gas
4
Material that on very short exposure could cause
death or major residual injury.
Example: hydrogen cyanide
(Red)
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Flammability
Burning
Susceptibility of Material to
0
Material will not burn.
1
Material must be pre-heated before ignition can occur.
Example:
corn oil
2
Material must be moderately heated or exposed to relatively
high ambient temperature before ignition can occur.
Example: diesel fuel
3
Liquids and solids that can be ignited under almost all
ambient temperature conditions.
Example:
gasoline
4
Materials that will rapidly or completely vaporize at
atmospheric pressure and normal ambient temperature, or
that are readily dispersed in air and that will burn
readily.
Example: propane gas
(Yellow)
Reactivity
Burning
Example: water
Susceptibility of Material to
0
Material that in itself is normally stable, even under fire
exposure conditions, and is not reactive with water.
Example: liquid nitrogen
1
Material that in itself is normally stable, but which can
become unstable at elevated temperatures and pressures.
Example: phosphorus (red or white)
2
Material that readily undergoes violent chemical change at
elevated temperatures and pressures or which reacts
violently with water or which may form explosive mixtures
with water.
Example: calcium metal
3
Material that in itself is capable of detonation or
explosive decomposition or reaction but requires a strong
initiating source or which must be heated under confinement
before initiation or which reacts explosively with water.
Example: fluorine gas
4
Material that in itself is readily capable of detonation or
of explosive decomposition or reaction at normal
temperatures and pressures.
Example: trinitrotoluene (TNT)
(White)
Special Precautions
77
W
Material shows unusual reactivity with water (i.e. don't
put water on it).
Example: magnesium metal
OX
nitrate
Material possesses oxidizing properties.
Example: ammonium
Other symbols commonly used:
ACID
Material is an acid.
ALK
Material is a base (alkaline).
COR
Material is corrosive.
Material is radioactive.
Unit Conversions and Formulas
Unit Conversions
1 inch = 2.54 cm
1 mile = 5280 ft
1 mile = 1.609 km
1 lb = 453.6 g
1 kg = 2.205 lb
1 mL = 1 cm3
1 L = 1000 mL
1 L = 1 dm3
1 kg = 1000g
1 g = 1000 mg
1 gallon = 3.785 L
1 fluid ounce = 29.575 mL
speed of light = 2.9979 x 108 m/s
Standard SI Units
Length
Mass
Time
Temperature
Electric current
Amount of substance
Luminosity
meter (m)
kilogram (kg)
second (s)
Kelvin (K)
ampere (A)
mole (mol)
candela (cd)
Standard SI Prefixes
78
Prefix
Symbol
Decimal notation
tera
giga
mega
kilo
hecto
deca
deci
centi
milli
micro
nano
pico
T
G
M
k
h
da
d
c
m

n
p
1,000,000,000,000
1,000,000,000
1,000,000
1,000
100
10
0.1
0.01
0.001
0.000001
0.000000001
0.000000000001
Exponential notation
1012
109
106
103
102
101
10-1
10-2
10-3
10-6
10-9
10-12
Formulas
Volume of a sphere = 4/3 r3
Volume of a cylinder = r2 . h
Percent error = [(experimental value – accepted value) / accepted value] x 100
Proper Graphing Technique
1. Use as much of the paper as possible. Do not scrunch all of the data points into a
small corner of the paper. A larger graph is more precise than a smaller graph.
2. Use a ruler to create your lines. A neat graph is more precise than a messy one.
3. Create a scale on each axis that is consistent. Every block must represent the
same interval on a given axis, and you cannot have gaps in your scale. There is
no rule that says the graph must start at 0,0.
4. Be sure to label the axis with the appropriate variables including the units. The
independent variable should be plotted along the x-axis and the dependent
variable should be plotted along the y-axis.
5. Do not plot your data points as simple dots. It is easy to confuse stray marks with
intended data points. Rather, circle or box your dots or use an “X” to represent
your data points.
79
6. Consider the best method for interpreting your data. Options include; best-fit
lines, exponential curves, quadratic curves, sigmoidal curves, or scatter plots.
7. Always include a descriptive title for your graph. A common title would read the
dependent variable on the y-axis vs. the independent variable on the x-axis.
8. Create a key if your graph involves more than one line on the same plot.
Chemistry I Course Guidelines
Course Description - Chem I is an introductory general chemistry course designed for
students who plan to attend college, but may not plan to pursue a career in the sciences.
Basic chemical principles will be explored from an experimental and theoretical
perspective. An emphasis will be placed on thought processes and the practical
application and relevance of the principles covered.
Expectations - Chemistry is a challenging subject for most students. Most of the
concepts will be new to you, and you will be required to do much more than memorize
facts. In addition to learning basic concepts, you will need to organize your thoughts and
observations, think critically, and solve complex problems using a systematic approach.
Furthermore, you will have the opportunity to participate in a variety of labs and
activities in which you will be expected to strive for accuracy and precision. While these
may seem like very high expectations, be assured that if you are willing to put forth the
effort, I will do my best to help you succeed.
Student Responsibilities 1.
Arrive to class on time and bring all required materials.
2.
After any excused absences, you should obtain any notes that
you may have missed, and make up any missed assignments.
3.
Do your homework! Homework will not be accepted late.
4.
Obey all safety rules when in the lab. I take this very seriously!
80
5.
Read the procedure for an experiment, and know the objective
before you begin working in the lab.
6.
Make every effort not to be absent on lab days. If you should
miss a lab, you must make it up within two weeks from the day
you return. It is your responsibility to make the arrangements.
7.
If you have relevant questions during class, do not be afraid to ask them!
Thoughtful questions will only serve to make the class more interesting.
8.
However, do not expect the teacher to simply give you the answers. You
are here to learn chemistry, but more importantly, you are here to learn to
how to think! Be prepared to be challenged.
9.
Cell phone use is not permitted during class! Failure to comply will result
in the confiscation of your cell phone. The use of a cell phone, camera or
similar electronic device during a test will result in a grade of zero!
10.
You may not use the hall pass during a test or quiz. If you must, use the
bathroom prior to taking the test or quiz.
Grading - Your grade will be based on the following formula:
Tests and Quizzes
45%
Labs and Projects
35% [-10% for each day late]
Homework and Activities
20% [Not accepted late]
Final Exam - There will be a comprehensive final exam given at the end of the year.
The final exam will be worth 10% of your final grade.
Student Safety Contract
Laboratory Safety Rules
General Conduct
1. Conduct yourself in a responsible manner at all times while in the lab. Never fool
around in the lab. Inappropriate behavior will result in the loss of lab privileges
for a specified length of time. Intentional dangerous behavior will result in a
grade of zero for the lab.
2. Follow all written and verbal instructions carefully.
3. Read all procedures thoroughly before entering the lab. If you do not understand
any part of the procedure, ask the instructor before proceeding.
4. Never work alone in the lab. No student may work in the lab or chemical storage
area without an instructor present.
5. Do not touch any equipment, chemicals, or other materials until you are instructed
to do so.
6. Never eat, drink, or chew gum during lab. Some materials you will be using are
toxic!
7. Perform only those experiments authorized by the instructor.
8. Notify the instructor immediately of any unsafe conditions that you observe.
9. Experiments must be monitored at all times. You will be assigned a lab station at
which to work. Do not wander around the room, distract other students, or
interfere with the experiments of others.
81
10. Work areas should be kept clean and tidy at all times. You could lose points for
failing to clean up your lab area.
11. Never taste any chemical in the lab!!
12. Know the locations and operating procedures of all safety equipment including
the first-aid kit, eye-wash station, safety shower, fire extinguisher, fire blanket,
and emergency gas-shut-off button. Know where the fire alarm and exits are
located.
13. During fire drills close all containers and turn off all gas valves, fume hoods, and
electrical equipment.
14. Read all labels and equipment instructions carefully before use.
15.You must wear safety goggles any time you are told to and any
time glassware, flame, heat, or chemicals are used.
16. Contact lenses present special dangers and should not be worn in the lab without
the instructor’s permission.
17. Long hair must be tied back, and loose or baggy clothing must be secured. It is
best to wear durable, close-fitting clothing. Lab aprons are provided for your use.
18. You must wear closed-toed shoes in the lab. Sandals and flip-flops are not
permitted.
19. Always wash your hands after working in the lab.
20. Report any accident (spill, breakage, etc.) or injury (cut, burn, etc.) to the
instructor immediately, no matter how trivial it may seem.
Handling Glassware and Lab Equipment
21. Examine glassware before each use. Never use chipped, cracked, or dirty
glassware – report it to the instructor.
22. Do not immerse hot glassware in cold water; it may shatter.
23. Do not use your hands to handle hot glassware. Remember, hot glassware looks
the same as cold glassware.
24. Never handle broken glass with your bare hands. Use a brush and dustpan to
clean up broken glass. Place broken glass in the designated glass disposal
container.
25. Fill wash bottles only with distilled water and use only as intended. Squirting
other students is considered dangerous and disruptive behavior and will result in
disciplinary action.
26. If you spill any chemicals on a balance, it is your responsibility to clean up your
mess. Do not leave it for another group or for the instructor to clean up.
Heating Substances
27. Keep flammable and volatile substances away from heat sources and open flames!
28. Never leave a lit burner unattended.
29. Never leave unattended anything that is being heated or is visibly reacting.
Always turn the burner or hot plate off when not in use.
30. Never look directly into a container that is being heated.
31. Never reach over an exposed flame, and be certain that long hair and loose
clothing is secured.
82
32. Do not point the open end of a test tube being heated at yourself or anyone else.
33. Do not place hot apparatus directly on the lab bench. Always use an insulating
wire gauze pad.
34. Rinse lit matches with water before placing them in the trash. Do not leave them
in the sink.
Handling Chemicals
35. All chemicals in the lab are to be considered dangerous.
36. Double check labels to make sure that you are using the proper chemical.
37. Take only as much chemical as you need, and never return a chemical to a reagent
bottle once it has been placed in another container.
38. Take great care when transferring acids and other chemicals from one part of the
lab to another. Hold them securely and walk carefully.
39. Keep backpacks out of the traffic isles near the lab.
40. Keep your hands away from your face, eyes, and mouth while using chemicals.
41. Never use mouth suction to fill a pipet.
42. When diluting an acid, always add acid to water, never the other way around.
43. Dispose of all chemical waste in their designated waste containers. Do not put
solids or unauthorized chemicals in the sink. If you are unsure, ask the instructor.
44. Never remove chemicals or other materials from the lab.
Name _________________________
Period ________
Student Laboratory Safety Contract
Chemistry is a hands-on laboratory class. You will be performing many
laboratory activities which require the use of hazardous chemicals. Safety in the
laboratory is the #1 priority for students, teachers, and parents. To ensure a safe
chemistry laboratory experience for you and your classmates, the list of safety rules given
to you must be followed at all times. It is your responsibility to read all of these rules and
you will be held accountable for any violations.
In order to be given permission to participate in the laboratory activities you must
read and sign this contract and have a parent or guardian read and sign it as well. You
will also be tested on this information and you will be required to pass this test with a
score of 80% or better before you may enter the lab.
Personal Information



Do you wear contact lenses?
Yes ____
No ____
Do you have allergies?
Yes ____
No ____
If so, list specific allergies _______________________________________
Do you have any medical conditions that I should be aware of? Yes ___ No ___
If so, list specific conditions _____________________________________
Agreement
83
I have read and agree to follow all the safety rules set forth in this contract. I realize that I must
obey these rules to ensure my own safety and that of my fellow students and instructors. I will cooperate to
the fullest extent with my instructor and fellow students to maintain a safe lab environment. I will also
closely follow the oral and written instructions provided by the instructor. I am aware that any violation of
this safety contract that results in unsafe conduct in the laboratory or misbehavior on my part, may result in
my being removed from the laboratory, receiving a failing grade, detention, or other disciplinary action in
accordance with the school code as outlined in the student handbook.
______________________________________________
Student signature
____________
Date
Parent or Guardian
We feel that you should be informed regarding the school’s effort to create and maintain a safe
chemistry laboratory environment. Our gola is to eliminate, prevent and correct possible hazards. Please
read the list of safety rules outlined in this contract. No student will be permitted to perform laboratory
experiments unless the contract is signed by both the student and parent/guardian and is on file with the
teacher. Your signature on this contract indicates that you have read the laboratory safety rules, are aware
of the measures taken to insure the safety of your son/daughter in the chemistry lab, and will instruct your
son/daughter to uphold his/her agreement to follow these rules and procedures in the laboratory.
______________________________________________
Parent/Guardian signature
_______________
Date
84