ORGANIC CHEMISTRY I
LABORATORY MANUAL
By
STEPHEN ANDERSON
And
ROBERT SHINE
RAMAPO COLLEGE OF NEW JERSEY
MAHWAH, NEW JERSEY
SEPTEMBER 2009
© 2009 Stephen Anderson and Robert Shine
ORGANIC CHEMISTRY I
LABORATORY MANUAL
DR. STEPHEN ANDERSON
DR. ROBERT SHINE
RAMAPO COLLEGE OF NEW JERSEY
CONTENTS AND SCHEDULE:
LAB 1
CHECK IN, SAFETY AND PROCEDURES
LAB 2
MELTING POINTS
LAB 3
CRYSTALLIZATION
LAB 4
EXTRACTION
LAB 5
DISTILLATION - SIMPLE AND FRACTIONAL
LAB 6
THIN LAYER AND GAS CHROMATOGRAPHY
LAB 7
COLUMN CHROMATOGRAPHY
LAB 8
OPTICAL RESOLUTION
LAB 9
NUCLEOPHILIC SUBSTITUTION REACTIONS
LAB 10
INFRARED AND ULTRAVIOLET SPECTROSCOPY
LAB 11
NUCLEAR MAGNETIC RESONANCE SPECTROSCOPY
LAB 12
MICROWAVE SYNTHESIS OF ASPIRIN
LAB 13
CHECK OUT AND FINAL EXAMINATION
2
ORGANIC CHEMISTRY I
LABORATORY MANUAL
APPENDICES
APPENDIX LIST:
APPENDIX A
SAFETY ISSUES
APPENDIX B
COMMENTS ABOUT WRITING STYLE
APPENDIX C
MEASUREMENTS AND SIGNIFICANT FIGURES
APPENDIX D
PERCENT YIELD CALCULATION METHOD
APPENDIX E
CHEMICAL DATA
ACKNOWLEDGEMENTS: The Authors wish to thank the
following people for their assistance with this
manual: Carol Ichinco (Laboratory Coordinator),
Thomas Drwiega (Laboratory Technician), Gurpreet Kaur
(Honors Research Student), Benjamin Barrios (Research
Student), and all the other Instructors who have
taught Organic Chemistry Laboratory sections for the
past few years.
© 2009 Stephen Anderson and Robert Shine
3
LAB 1:
LABORATORY PROCEDURES
SCOPE OF THE COURSE:
This course offers a comprehensive introduction to
laboratory techniques in Organic Chemistry for science majors.
Chemistry is a mature science that continues to expand and
evolve in step with recent developments in science and
technology. Students will perform experiments that put into
practice the ideas discussed in lecture.
E-MAIL AND INTERNET USE:
Use will be made of e mail (Luminis) and the Internet to
provide important instructions to students concerning
laboratory material. Students are encouraged to contact the
Instructor in person or by e mail for additional help, if
needed. Information about the course may be transmitted
periodically by the Instructor via a web site for the course.
Students should check this web site at least once a week.
Students may not submit lab reports via e mail or e mail
attachments or by fax. Lab reports must be submitted as paper
copy.
ATTENDANCE:
Each lab session will begin at the scheduled time.
Attendance at all laboratory sessions is mandatory and will be
recorded. If you are late, it will be noted and recorded. A
missed laboratory experiment cannot be made up at another time
and any missed laboratory session will result in a grade of
zero (0) for that experiment unless an alternate library
report is completed. During the course of the semester you
will be allowed to substitute one alternate report for any one
lab period you cannot attend. You must discuss the content
and deadline of the alternate report with the Instructor as
soon as possible after your absence.
SPECIAL NEEDS:
Students with special needs who are registered with the
Office of Specialized Services and require special
accommodations should notify the Instructor as soon as
possible.
GRADING:
Each laboratory session will be graded by a procedure
explained to you by your laboratory Instructor. Your in-lab
work and your report will be used to determine your grade.
The final laboratory exam will account for about 25 % of your
final course grade. In grading laboratory sessions, the
following will be considered: meeting deadlines, submitting
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the necessary forms and reports on time and in good order,
being on time for lab, and working in a safe and efficient
manner. Your practical bench skills and laboratory results
will also be analyzed, along with your ability to effectively
explain and discuss them in your lab report.
SAFETY:
As with many activities in daily life, working in the
Organic Chemistry laboratory poses some hazards that you must
recognize. Many of the chemicals used can be toxic if not
used correctly. Also, many substances you will use are
flammable and must be kept away from sources of heat that
could cause them to ignite and start a fire. Glassware can be
broken and cause cutting hazards. You must always come to the
laboratory prepared so that you will be aware of the unique
hazards you will confront. You must consult the MSDS forms
for the chemicals you will use so that you will know their
toxic and flammability properties. You can find MSDS data on
the internet by typing the chemical name followed by msds in a
Google search box. If you have any questions about the
experiment you are performing or the chemicals you are using,
please ask the Instructor before you begin work. When you are
done with the experiment, the chemicals you have should be
discarded as hazardous waste or reclaimed material and placed
in the proper containers. If you are pregnant or expect that
you may become pregnant, you should consult with a medical
doctor about the potential hazards involved with exposure to
organic chemicals. See Appendix A for more information about
safety.
EYE PROTECTION:
Students are required to have department approved lab
goggles that must be worn at all times in the laboratory when
experimental work is being done. Students who fail to wear
eye protection could be asked to leave the lab by any employee
of the College. It is not recommended that you wear contact
lens in the laboratory.
PRE-LABORATORY PREPARATION:
It is imperative that students be prepared to perform the
scheduled experiments. An unprepared student is a hazard in a
chemistry laboratory. In order to prepare for the experiment,
students are expected to read the appropriate laboratory
experiment prior coming to the lab. Students may be tested at
the beginning of the lab period to ensure they are prepared to
begin work.
LABORATORY NOTEBOOK:
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Students should have a notebook that will be used in the
lab for recording all raw data. A title should appear at the
top of the page for each new experiment. Data obtained during
an experiment should always be neatly and clearly recorded in
this notebook. You should use the department approved
notebook for this course. You will be required to turn in
your experimental data sheet before you leave the lab at the
end of the period. You must include your name, the name of
the experiment, and the date of the lab period on the data
page you submit which will be graded as part of your report
grade. Any measurement data must have the numerical value
recorded to the correct number of significant figures and the
units of that measurement, such as 1.234 g for a weight
measurement.
LAB REPORTS:
Before leaving the laboratory, each student must submit a
laboratory data sheet as described above. Lab reports can
only be submitted for labs that have been attended and must be
submitted at the beginning of the lab session following the
lab that is being reported. Points will be deducted for
lateness or for an incomplete lab report. No lab report for a
given experiment will be accepted after graded reports for
that lab have been returned to the class. Generally, graded
lab reports will be returned in the next class meeting after
the report is submitted. A score of zero is assigned for a
report if the lab report is not submitted.
FORMAT OF LAB REPORT:
Laboratory reports must be typed on 8.5 x 11 inch white
paper. The report should have a professional appearance and
it must demonstrate that much thought and care went into its
preparation. Some comments about report writing style appear
in Appendix B. Spelling and grammar count. All laboratory
reports must be written in the following format that conforms
to the guidelines set forth by the American Chemical Society:
COVER PAGE
Place the following information in the upper right
hand corner of the title page or first page of the report:
title of the experiment, your name, date of submission
ABSTRACT
The abstract consists of two to five sentences that
concisely inform the reader of the nature of the experiment
that was performed and a brief summary of your final results.
Note that the abstract is read immediately after the title and
hence need not repeat any information that already appears in
the title. Though the abstract appears at the beginning of
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the lab report, it should be the last section of the report
that is written. The abstract should be about 40 to 80 words
and should give a concise description of the experiment. It
should not be too general or too specific.
INTRODUCTION
The introduction consists of three to six good
paragraphs that relate only the most essential elements of
theory to the reader. Consult other sources such as a
chemistry textbook for this information. Include only enough
theory so that the reader is made to understand the basic
physical and chemical principles involved in the experiment.
If applicable, any pertinent mathematical and chemical
equations must be briefly given in this section. The
introduction should conclude with a one- or two-sentence
paragraph that explains the objective or goal of the
experiment.
SAFETY
Students shall describe all relevant safety
precautions that were observed during the course of the
experiment. Information about the hazards that may be
encountered is particularly important. Material Safety Data
Sheet information should be included. You should give a
summary of the pertinent data (such as: toxicology,
flammability, and physical properties) and not just a copy of
the MSDS.
EXPERIMENTAL
For our purposes, the experimental section will
consist of a short paragraph that includes a sentence that
refers the reader to some source for the procedure. For
example, the student may write: "The procedure for this
experiment appears in the web page for the course (1)." The
number in brackets refers to the citation number. This number
is used to refer the reader to the citation in the References
section where the full reference (including information such
as the name of the author, title of book or web site, date of
publication, and page number or exact web address) will appear
next to number 1. In addition to the reference citation, any
deviations from the published procedure and any experimental
hints or tips that may aid the reader in understanding and
repeating the experiment should be included. For example, ‘In
this experiment, the instructor requested that 0.05M HCl be
used instead of 0.1M as specified in the lab module.’
RESULTS
This section consists of a written paragraph that
refers the reader to tables, graphs, data sheets, and figures
that contain your data. It is especially important to inform
your reader how your raw data were used to calculate your
final results. You must explain how you determined your final
results and include sample calculations with an accompanying
7
explanation is necessary. The use of a spreadsheet program
(such as Microsoft Excel) for calculations, tabulation of
results, and graphing is encouraged.
DISCUSSION
This section is used to indicate to the reader how
the results relate to the theory and whether or not the
objective was met. In addition, the final results must be
compared to literature values, if available. Reasonable
sources of error should be listed and discussed with respect
to their contribution to the final results. The discussion
provides a good indication of the student's comprehension of
the material. A good discussion should show that the student
was able to correctly interpret the data and to relate the
results to the scientific principles being tested by the
experiment. If the experiment was not successful, then the
discussion is equally important in stating the reasons for the
outcome. A good discussion can be written regardless of the
success of the experiment.
CONCLUSION
This section should include a summation of the
Results and Discussion. It should only be about 1 paragraph
long and is intended to draw together all the pertinent
information that has been determined from the experiment.
REFERENCES
This section consists of a numbered listing of
literature or internet references that were used to perform
the experiment and that were used to write the lab report.
This includes a full reference to the lab module and any other
publications or correct web page URLs that you may have used
to obtain literature values and supplemental theory. Note
that the reference numbers must correspond to the reference
citations used in the text of the report.
PROPER USE OF THE AUTOPIPETTE:
When an experiment calls for an accurate amount of a
liquid reagent, it is generally more convenient to measure its
volume than weight. Multiplying the measured volume by the
density of the substance gives the weight. The weight can
then be converted to moles by dividing the weight by the
molecular weight of the substance. The auto pipette is an
accurate, fast, easy and convenient method to measure the
exact volume of liquid reagents. In this experiment we will
learn to use the auto pipette.
Generally, the correct volume setting for a given
experiment will already be set by the Chemistry Technician.
You should check to be sure the setting is correct and, if it
isn't, please tell the Instructor. The disposable tip, which
the Instructor will discard when everyone is done, will be on
the pipette. To properly use the pipette, hold it by wrapping
8
your fingers around the barrel with your thumb on the knob at
the top of the pipette. Press the plunger knob to the first
"stop", that is, where you feel it resist. While holding the
knob down, dip the tip into the liquid to be measured.
Release the plunger slowly and smoothly so liquid is drawn
gently into the disposable tip. If the liquid "jumps" into
the tip, the volume measurement will be inaccurate. Do not
lift the tip out of the liquid until you have fully released
the plunger knob. When removing the tip, slide the tip on the
side of the vessel to avoid the extra drop on the outside of
the tip. Transfer the liquid to your reaction vessel by
pressing the plunger knob all the way to the "stop" and giving
an extra gentle push to dispense the final drop in the tip.
Do not dip the tip into any liquid or solid already in the
reaction vessel as it will contaminate the tip and, therefore,
the starting material.
PRACTICE USING THE AUTOPIPETTE:
Several auto pipettes will be supplied with tips and set
for 0.500 ml. Use the small Erlenmeyer flasks in your drawer.
Determine and record on the data sheet provided the exact mass
of the empty flask to three decimal places. Using the
autopipette and following the directions above, now dispense
0.500 ml of de-ionized water into the pre-weighed flask.
Determine the mass of the flask and water. Subtract the
weight of the empty flask from the weight of the flask and
water to find the exact mass of the water. Since the density
of water is 1.00 g/ml, the 0.500 ml has a mass of 0.500 g.
Repeat the process two more times, making sure that each
partner has a turn. Then complete the data sheet provided and
return it to the Instructor.
© 2009 Stephen Anderson and Robert Shine
9
LAB 2: MELTING POINTS
PURPOSE:
You will learn why the melting point of a solid is
important in organic chemistry and you will practice taking
melting points.
BACKGROUND INFORMATION:
The melting point of a solid substance is the temperature
range where the substance goes from the solid to the liquid
state. It is an important physical property of a substance
because it is accurate, reproducible, and easily done. The
melting point range for a pure sample is usually small if the
measurement is made carefully. Usually the range is 1 to 2
degrees centigrade (or Celsius). The lower temperature of the
melting point range is the temperature at which the solid
first begins to liquefy. The upper temperature is the
temperature at which the entire sample is in the liquid state.
For the range to be small the sample size should be small and
the heating should be done slowly. Usually, the temperature
should increase about one degree per minute near the melting
point. Unlike the boiling point of a substance (see the
distillation experiment), the melting point does not vary much
with changes in the atmospheric pressure.
The melting point can be used as a criterion for the
identity of a compound and a qualitative indication of its
purity. The observed melting point of a substance can be
compared to the melting point of a substance given in the
literature. If the two temperatures are different, the
compounds are not the same. If the two temperatures are the
same, the compounds may be the same but not necessarily. Many
different compounds have the same melting point. One could do
a mixed melting point (described below) to further help
identify an unknown. Pure samples generally have a very sharp
melting point range while impure samples have a broad (roughly
3 to 15 degree) melting point range. Further, impure samples
melt lower than expected.
A mixture melting point is done when one wishes to fully
show whether two compounds with the same melting point are
indeed the same. This can help identify an unknown. A good
example would be the two compounds urea and cinnamic acid
(shown on the next page) both of which have a melting point
close to 132-133 degrees C. Let’s say you have an unknown
that melts at 132-133 degrees C. If the unknown is urea and
you mix it with some pure urea from another source, the
melting point of the mixture would be 132-133 as the entire
sample is urea. However, if the unknown were cinammic acid
and you mixed it with urea the melting point of the mixed
10
sample would now be lower and broader than 132-133 as the
mixture was impure. So, if you suspect the identity of an
unknown, you can test this by doing a mixed melting point by
mixing the unknown with a sample of the pure compound and
measuring the melting point of the mixture.
Melting point composition diagram for a mixture of the solids X and Y
Liquid X + Y
mpt. Y
mpt. X
Liquid +
solid Y
T
em
p
Co
Liquid +
solid X
ET
(Eutectic point)
Solid X + Y
Mole % X
100
75
60
0
Mole % Y
0
25
40
100
H
COOH
C
C
O
H
C
H2N
NH2
Urea
mp 132.5-133 oC
trans-cinnamic acid
mp 132.5-133 oC
Since the melting point merely records a change in state
(solid to liquid), it should be possible to solidify the
liquid and retake the melting point again as a second check of
the measurement. This is never done. If the measurement is
11
to be repeated, a fresh sample of the solid is taken and a new
melting point capillary tube is used. Melting point capillary
tubes are only used once and then discarded in the broken
glass container. The reason a second melting point is not
taken on a given sample is that the solidification of the
liquid rarely gives pure crystals (as some decomposition may
have occurred) or may give a different crystal structure which
often has a different melting point.
Sometimes, a solid decomposes as it melts. This is true
of certain classes of compounds or compounds with high melting
points. In such cases, the melting point will have the letter
d after it to indicate decomposition. Melting points with
decomposition are not as accurate as a true melting point.
Often you will observe the decomposition by a strong gas
evolution or the charring of the sample when melting.
An error that can cause your melting points to be wrong
is that the thermometer or temperature measuring device may be
badly calibrated. For very careful work, the temperature
measuring device is calibrated before measurements are made.
This is done by taking the melting points of very pure
standard substances which melt at different temperatures so
the accuracy of the device can be noted at various
temperatures. A correction value is then added or subtracted
from your later measurements as needed. In this course,
temperature corrections will not be done. Please keep in mind
that our actual observations may vary from the literature
values by a degree or two for this reason.
EXPERIMENTAL PROCEDURE:
Melting point measurements are easy to do. Good eyesight
and patience are needed. Measurements can be made in any one
of a number of commercial melting point apparatuses. Some use
oil baths (silicon oil is the best to use as it is thermally
stable and can be heated to 350 degrees safely). Thermometers
are often used to measure the temperature but the mercury in
the thermometer poses a safety hazard as mercury is very
toxic. We use Mel-Temp devices in our laboratory. This
device uses a heated metal block in place of the oil bath. An
advantage of the heated block is that it cools quickly and
hence speeds up the process when many melting points need to
be done. Temperature measurement in our equipment uses a
thermocouple with a digital readout device which works very
nicely.
To make a melting point determination you need to:
1. Put the sample in a melting point capillary tube. Do
this by taking a melting point capillary tube (be sure to
use the capillaries that are sealed at one end - do not
use the capillaries that are open at both ends) and
placing a small amount of solid in the tube. Tap the
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2.
3.
4.
5.
solid down so it is tightly packed in the tube. The
solid should be about 1 millimeter or 1/8 inch in the
bottom of the tube.
Place the tube in the melting point apparatus and turn
the device on.
Observe the sample as the temperature rises and record
the melting point range of the sample noting both the
temperature at which the first drop of liquid appears (be
careful to differentiate sample softening from melting)
and the temperature at which the entire sample is liquid.
This means that your melting point should always be
reported as a temperature range.
Record in your notebook the sample identification and its
melting point range.
Discard the capillary tube in the broken glass container.
You should determine the melting points of:
Salicyclic acid
Caffeine
Urea
Cinnamic acid
Urea/Cinnamic acid mixtures: 1:4, 2:3, 3:2, 4:1, 1:1
Unknowns assigned by the Instructor
IMPORTANT INFORMATION ABOUT THE REPORT:
Record the melting point ranges you observed and compare
your results with the literature values. Explain any
differences you noted. The purpose of recording melting point
of the salicylic acid and caffeine in this experiment is to
test the calibration of the apparatus.
Diagram the urea - cinnamic acid data on a piece of graph
paper or via an Excel spreadsheet. Your graph should have 7
data points (which will not be enough to give careful
results). A sample diagram appears earlier in this document.
END OF EXPERIMENT.
© 2009 Stephen Anderson and Robert Shine
13
LAB 3: CRYSTALLIZATION
PURPOSE:
You will learn why crystallization is done in organic
chemistry and how to perform the necessary tasks to
recrystallize an impure sample.
BACKGROUND INFORMATION:
Organic solid substances often have impurities in them
which make them less desirable for their intended purposes.
Crystallization is an important technique that can be used to
increase the purity of a solid sample if certain conditions
are met. First, the impurities cannot be a large percentage
of the sample and, secondly, the properties of the impurities
should be such that the technique would be useful. An
important use of crystallization in industry is the refining
of sugar. Raw cane sugar is put through a series of
crystallization steps which results in commercial products of
varying degrees of purity. The table sugar you use is highly
refined sucrose that is obtained from cane sugar.
You will recrystallize naphthalene and benzoic acid in
this experiment. Their structures are as follows:
O
C
OH
Naphthalene
Benzoic Acid
The full procedure to recrystallize a sample involves the
following steps:
Step 1: Find a suitable solvent for the
recrystallization process.
Step 2: Dissolve the sample in the minimum amount of
solvent at its boiling point.
Step 3: Decolorize the solution (if needed) by using
decolorizing carbon.
Step 4: If needed, quickly and carefully filter the
hot solution to remove any insoluble impurities.
Step 5: Slowly cool the hot solution to room
temperature and then to 0 degrees C in an ice bath.
Step 6: Filter the crystallized solid using a
suction filtration apparatus.
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Step 7: Dry the collected solid.
Step 8: Weigh and determine the melting point of the
solid that was obtained and compare it to the reported melting
point for that compound. If the melting point is low or if
the melting range is broad, dry the sample further or
recrystallize the sample again.
Depending on the nature and purity of the starting
sample, one or more of the above steps may be omitted. In
order to guide you in this, the following considerations are
made:
a. The ideal solvent is one in which the solid is very
soluble at the boiling point of the solvent and very
insoluble at 0 degrees. Generally, the rule "like
dissolves like" is a good starting point to determine
suitable solvents. Solubility data found in reference
works may be useful.
b. Sometimes when a suitable solvent cannot be found one
can use a mixture of two miscible solvents to achieve
the desired solubility properties. The mixture can be
prepared before the solid is dissolved. Alternatively,
the solid can be dissolved in the solvent in which it
is more readily soluble and the other solvent is added
drop wise at an elevated temperature until
precipitation just begins.
c. The solvent should be volatile so that it can be
removed from the solid easily.
d. When dissolving the solid in the chosen solvent, heat
the sample in a small amount of the solvent and add
more solvent as needed. Do not rush this process as
the dissolution of some solids may take some time.
Generally this procedure is done in an Erlenmeyer
flask or beaker and the heating is done on a hot
plate.
e. A very small amount of decolorizing carbon or Norit
can be used to remove colored impurities from the
solution. The carbon adsorbent must be carefully
removed from the hot solution using gravity
filtration. If this is not done carefully, often the
final product will be gray due to traces of carbon in
the final crystals.
f. If carbon was used and/or if there are insoluble
impurities, the hot solution must be filtered by
gravity to remove these insoluble substances. This
procedure must be done as quickly as possible as some
of the desired solid may crystallize at this time
which will adversely affect the process. A wide bore
glass funnel (powder funnel) and fluted filter paper
are used in this step.
15
g. The hot solution is set aside and not disturbed as the
crystallization occurs. If the container is touched,
the supersaturated solution may crystallize quickly
giving small crystals (which generally contain
impurities). A slow undisturbed crystallization
generally results in larger more pure crystals.
h. If crystallization does not occur in a reasonable time
(as determined by your Instructor), you can sometimes
hasten the process by introducing a seed crystal or
scratching the side of the glass container with a
stirring rod.
i. Final cooling to 0 degrees C is done in an ice water
bath to maximize the amount of solid that will
precipitate.
j. The solid is filtered using suction to remove the
solvent and the crystals are moved around the funnel
using a spatula while suction is continued to dry the
crystals as much as possible.
k. The crystals which are left in the glass container can
be removed by using a spatula or a very small amount
of cold solvent. Students generally lose material at
this stage because they use too much solvent which
will then dissolve some of the desired product.
l. The choice of which size funnel to use is made by
determining the amount of solid obtained and the size
of the funnels which are available. Do not use a
funnel that is too small or the solid may spill out of
the funnel. Do not use a funnel that is too large as
there will be mechanical loses that will lower your
percent recovery.
m. The melting point of the final solid will determine
the success of your work. The melting point range
(see prior experiment) should be small and the melting
point should be very close to the reported value.
Students often note that there is solid material in the
filter flask below the funnel in which the solid has been
collected. They often think the funnel leaked some of the
solid. Recall that the filtrate is a saturated solution of
the substance you are trying to obtain. Also recall the
suction generally causes some solvent to evaporate. Hence,
the filtrate becomes supersaturated and solid will
precipitate. In the whole crystallization process, a 100 %
recovery is generally not possible due to this phenomenon.
Sometimes in important work a second crop of (less pure)
crystals is obtained by further working on the filtrate.
EXPERIMENTAL PROCEDURE:
16
PART 1: Crystallization of naphthalene. Crystallize
150 mg of naphthalene from an 80% aqueous methanol mixture (80
parts methanol and 20 parts water) in a small beaker or
Erlenmeyer flask. Weigh the naphthalene accurately both at
the start and end of the process. Completely dissolve the
solid in a minimum amount of hot methanol/water solution. Set
the hot solution aside to cool to room temperature and then
cool it in an ice bath. Quickly filter your solid using
suction and a Hirsch funnel. Dry and weigh your product and
determine its melting point. Be sure to accurately record the
weights of materials you used and the observations you made in
your notebook. Calculate and report your percent recovery.
PART 2: Crystallization of benzoic acid in water.
Using the procedure noted above for naphthalene, crystallize
150 mg of benzoic acid using water as the solvent. Determine
the melting point of the crystallized material. As above, be
sure to accurately record the quantities of materials you used
and the observations you made in your notebook.
PART 3: Crystallization of benzoic acid in
methanol/water. Dissolve 150 mg of benzoic acid in a
minimum amount of hot methanol. Then add water drop-wise
until precipitation just begins. Then repeat the
crystallization process as described above.
IMPORTANT INFORMATION ABOUT THE REPORT:
In your report, be sure to show your percent recovery of
each sample used. The percent recovery is the (amount of
final product divided by the starting amount) times 100%.
Show the correct number of significant figures which are
allowed by the measuring equipment you used and show your
calculations. Discuss the reasons why your percent recovery
is the value you obtained. Discuss which measurements that
you made had to be done carefully and which did not contribute
to the precision of your percent recovery calculations.
END OF EXPERIMENT.
© 2009 Stephen Anderson and Robert Shine
17
LAB 4: EXTRACTION
PURPOSE:
You will learn how to separate complex mixtures using
differential distribution of substances in two immiscible
liquids and then you will perform a separation of a strong
acid, a weak acid and a neutral substance via acid-base
liquid-liquid extraction.
BACKGROUND INFORMATION:
Some solid mixtures are hard to purify by crystallization
techniques especially when the mixture is complex and no one
substance is present in large excess. Liquid mixtures can
sometimes be separated by extraction techniques when
distillation methods (see a later experiment) would be more
difficult. Extractions are very often done as the first steps
in the purification of a product in a synthetic reaction.
This is generally called a "workup procedure".
Extractions can separate components of a mixture because
the substances distribute themselves differently in the two
immiscible liquids depending on their distribution
coefficients. The distribution coefficient is the ratio of
the concentration of the substance in solvent A divided by the
concentration of the substance in solvent B. It can be shown
mathematically that multiple smaller extractions are better in
separating components of a mixture than one larger extraction.
Hence, we will do two "washes" (another term often used for
extraction) instead of one. The term wash is often used in
older literature to mean extraction. Another consideration is
mechanical loss. This is the loss of material by not
completely placing it in the desired container. Any solution
that is left in the separatory funnel will have some solute
dissolved in it which will be lost in the process. To
minimize mechanical loss, one often adds a small amount of
pure solvent to rinse the container and then adds this rise to
the solution that contains the bulk of that substance.
Since there will be two layers, you must determine which
layer is which. The less dense layer is the upper layer and
the denser layer is the lower layer. So a mixture of water
and diethyl ether will have the ether as the upper layer and
the water as the lower layer. However, when substances are
added to a solvent, the density is changed (for example, salt
water is denser than distilled water). Smelling a layer is
not an effective way to determine which layer is which as each
solvent will be partially miscible in the other. In the above
case, both layers will smell like ether. A good method to
determine which layer is the aqueous layer is to add a few
drops of one layer to a few milliliters of water. If the
18
drops dissolve, that layer is the aqueous layer.
don’t dissolve, that layer is the organic layer.
If they
O
OH
C
OCH3
OH
H3C
H3C
C
H3CO
CH3
4-tbutylphenol
mp 101 oC
Benzoic Acid
mp 120 oC
1, 4-dimethoxybenzene
mp 57 oC
First Neutralization Reaction:
O
C
OH
O
C
+ NaHCO3
O - Na+
+ H2CO3
(aqueous layer)
(organic ether layer)
Second Neutralization Reaction:
O - Na+
OH
+ NaOH
H3C
H3C
C
CH3
(organic ether layer)
H3C
H 3C
+ H2O
C
CH3
(aqueous layer)
19
Very often changing the pH of the aqueous solvent can
have dramatic effects on the distribution coefficients of
substances as will be noted in the experiment below. For
example, benzoic acid is reasonably soluble in ether and not
very soluble in water. However, benzoic acid is converted to
sodium benzoate in a 5% sodium bicarbonate solution (see the
first neutralization reaction on the previous page). Sodium
benzoate is very soluble in this aqueous solution and
insoluble in ether. Hence, we can effectively move benzoic
acid between ether and water by pH adjustment as will be done
in today’s experiment.
An extraction is usually done using a separatory funnel
which is a pear shaped glass container with a stopper at the
top and a stopcock at the bottom. The mixture to be separated
is placed in the funnel along with the two extracting
solvents. The closed funnel is then shaken (with periodic
venting) to thoroughly mix all the contents and allowed to
stand so that the layers can separate. If the density of the
two solvents is close or if there are any emulsifying
materials present, the layers may not separate which poses an
experimental difficulty. Once the layers separate, the
stopper is removed and the lower layer is drained through the
stopcock. If the upper layer is to be removed it is poured
from the top of the funnel. Students often forget to remove
the stopper before opening the stopcock. The stopper keeps
air from entering the top and hence very little or nothing
will drain from the bottom. Students then think the stopcock
is clogged but this in not the case. You must always remember
to remove the stopper before the stopcock is opened.
EXPERIMENTAL PROCEDURE:
An extraction will be done by dissolving 300 mg of a
1:1:1 mixture of benzoic acid, 4-t-butylphenol and 1,4dimethoxybenzene in 25 ml of a suitable solvent such as
diethyl ether. The solution is then placed in a separatory
funnel. Be sure the bottom stopcock is closed before pouring
the solution in the funnel. About 10 ml of an aqueous 5%
sodium bicarbonate solution (immiscible in ether) is added
through the top of the separatory funnel. The glass stopper
is put in place to seal the container. Holding the stopper
tightly, invert the funnel and open the stopcock to vent the
carbon dioxide and volatile ether that have contributed to a
pressure buildup. Close the stopcock and shake to mix the
contents. Open the stopcock to vent the pressure buildup
again. Repeat this process of shaking and venting a few times
and then place the separatory funnel in a ring stand in the
hood to allow the layers to separate. Once an interface is
seen between the two layers (generally 10 to 50 seconds)
remove the stopper and drain off the lower aqueous layer into
20
a marked container. Leave the ether (upper) layer in the
separatory funnel and repeat this process with a fresh 10 ml
of 5% sodium bicarbonate solution. When this solution is
removed from the separatory funnel combine the two 10 ml
portions to become a 20 ml portion of a sodium bicarbonate
solution of sodium benzoate. This solution will be used later
to recover the benzoic acid that was present in the original
mixture which has been now separated from the other two
components which remain dissolved in the ether. Again, be
sure to leave the ether solution in the separatory funnel.
To separate the 4-t-butylphenol which is still in the
ether layer with the 1,4-dimethoxybenzene, add 10 ml of
aqueous 5% sodium hydroxide solution (see the second
neutralization reaction shown earlier) to the ether layer and
extract the two layers as done above. Place the 10 ml of the
sodium hydroxide solution in a suitable marked container.
Repeat this process with another 10 ml of fresh sodium
hydroxide and then combine the two 10 ml portions as was done
above. The 20 ml solution will contain sodium 4-tbutylphenoxide which will be used later to get the pure
sample. The ether solution now has only the 1,4dimethoxybenzene in it (along with some dissolved water which
will have to be removed).
At this point, you have separated the three substances
from each other but you must do some work to purify them. To
obtain the benzoic acid, very carefully and slowly add
concentrated hydrochloric acid to the 5% sodium
bicarbonate/sodium benzoate solution from above. Much carbon
dioxide gas will be evolved as the white benzoic acid
precipitates. Test the pH with pH paper to be sure the
solution is strongly acidic and then set the flask in an ice
bath. After a while suction filter the benzoic acid, dry it,
weigh it, and determine its melting point. The reported
melting point for benzoic acid is 122-123 degrees C.
To obtain the 4-t-butylphenol, carefully add concentrated
hydrochloric acid to the 5% sodium hydroxide/sodium 4-t-butyl
phenoxide solution from the extraction. The product may ‘oil
out’ (which means it is separated from the aqueous solution as
a liquid instead of as a solid) and crystallize slowly over
time. After it solidifies and has been cooled in an ice bath,
filter by suction, dry it, weigh it, and determine its melting
point which should be about 100-101 degrees C.
The final compound, 1,4-dimethoxybenzene, is obtained by
adding solid anhydrous sodium sulfate (a drying agent) to the
ether solution to remove dissolved water. Add the drying
agent until it not longer clumps together but do not add too
much. Your Instructor will assist you in determining the
appropriate amount. After letting this stand for 5 to 10
minutes, gravity filter the solid sodium sulfate through a
21
small piece of cotton or glass wool into a pre weighed 50 ml
beaker. The clear ether solution is then subjected to a slow
stream of air to evaporate the ether. The pure 1,4dimethoxybenzene remains in the glass container. Weigh the
container and solid to obtain the weight of the solid. Take
the melting point of the solid which should be about 56-57
degrees C.
IMPORTANT INFORMATION ABOUT THE REPORT:
From the weight of starting material used and the weights
of each of the pure compounds obtained, calculate and report
the percent recovery of each component (assume there are 100
mg of each component). The percent recovery of component A
would be the weight of component A divided by the weight of
the starting component in the mixture times 100%. Be sure to
report the correct measurements and the correct number of
significant figures. In an ideal world, the sum of your three
percent recoveries should add up to 300%. Add your three
percents and discuss how your results compare to 300%. Also,
do a total percent recovery for all three components. Explain
any reasons why your results may be what they are. Also,
report the melting points of each recovered substance and
compare to the literature values. What do your melting points
suggest about the purity of the individual compounds you
obtained?
END OF EXPERIMENT.
© 2009 Stephen Anderson and Robert Shine
22
LAB 5: DISTILLATION
PURPOSE:
You will learn how to separate a complex mixture of
liquids based on their boiling point differences and then do a
simple distillation and a fractional distillation to determine
the advantages and disadvantages of each method.
BACKGROUND INFORMATION:
A distillation is an experimental procedure in which a
liquid is heated to its boiling point and the vapors are
condensed and collected in a suitable container. This
technique can be useful in separating the components of a
mixture if those components have different boiling points.
Boiling point composition curve for a mixture of the liquids X and Y
Vapor
B
T
em
p
.
Co
A
D
C
F
Liquid
E
Mole % X
100
Mole % Y
0
80
60
40
20
0
40
60
80
100
The vapor above such a mixture would be richer in the
lower boiling component as it is more volatile. If this
process involves only one vaporization and condensation cycle,
the process is called a simple distillation. If equipment is
23
used in which many vaporization and condensation cycles are
achieved before the final liquid is collected, the process is
called a fractional distillation. In today’s laboratory
experiment you will do both types of distillation on a liquid
mixture of ethanol and water.
When a simple relationship exists between the liquid
components of a mixture and the difference in boiling points
is large, a simple distillation (see a diagram of this setup
on the next page) often works well and will do a good job at
separating the components. However, when the relationship is
more complex or the difference in boiling points is not large,
as in the case of ethanol - water, the separation is not as
good. Fractional distillation can be used to better separate
the components. Ethanol - water is more complex in that it
forms an azeotropic mixture, that is, a mixture that forms a
constant boiling mixture. The temperature of this azeotrope
is lower than the boiling point of either component so it is
impossible to obtain a 100 % pure sample in such cases. A
fractional distillation of any mixture of ethanol - water will
give at best a 95 % ethanol, 5 % water mixture. One would
need to use other techniques to obtain 100 % ethanol (which is
also called 200 proof ethanol). We will disregard this
complication in today's experiment.
There are other types of distillations that are useful in
research environments and are beyond the scope of this course.
A spinning band distillation is used if one needs to separate
two liquids with similar boiling points. A vacuum
distillation (one that is done under reduced pressure) is
useful to distill high boiling components as many organic
substances decompose near their atmospheric pressure boiling
point. A steam distillation is useful when one wants to
purify a high boiling material which is not soluble in water.
In this case, since the vapor pressures of both liquids are
separate, the material distills when the total of all vapor
pressures equals the atmospheric pressure (which will always
be below 100 degrees Celsius). Solid substances that have an
appreciable vapor pressure can also undergo a similar process
to distillation which is called sublimation.
Since many organic liquids have a tendency to superheat
(that is, achieve a temperature higher than the boiling point)
before boiling action starts, it is very important to add a
few boil easers (an inert solid material made from clay plates
or anthracite coal) to the distillation flask before heat is
applied. The boil easers give off a continual stream of
bubbles that break the surface of the liquid and guarantee a
smooth boiling process. Boil easers can only be used for one
heating cycle and then lose their effectiveness. Fresh boil
easers must be added if the liquid is reheated later.
24
Simple Distillation Apparatus:
Water Out
Part B
Water In
Insert
fractionating
column here
HEAT
The thermometer placement to measure the temperature of
the vapor is very important and often leads to confusion among
students. The bulb of the thermometer must be placed just
below the sidearm of the distilling head. This way, the
entire mass of mercury in the bulb will be in the hot vapor
and accurately measure the vapor temperature. Some confusion
results when a student sees the liquid boiling and yet the
thermometer reads room temperature. This is due to the fact
that hot vapor has not reached the thermometer yet and the
25
thermometer is not reading the temperature of the boiling
liquid. An interesting question that you should answer and
discuss in your report is: ‘Is the temperature of the boiling
liquid the same as the temperature of the hot vapor?’
EXPERIMENTAL PROCEDURE:
Using a 100 ml round bottom flask and other equipment
from the semi-micro glassware kit, do a simple distillation of
75 ml of a 50 % ethanol - 50 % water mixture. Be sure to use
about 5 boiling chips in the flask and carefully place the
thermometer so that accurate readings will be obtained.
Collect about 25 ml of distillate into a graduate cylinder
recording the temperature reading at 5 ml intervals.
Using the 25 ml of distillate you obtained from the
simple distillation, do a fractional distillation. The
fractionating column is prepared by filling a condenser with
glass beads, and inserting it at position B as shown in the
distillation diagram above. Water is not passed through the
jacket of the fractionating column. Be sure to use fresh
boiling chips. Measure the temperature of the vapor at 1 ml
intervals and collect the distillate into a pre-weighed 10 ml
graduate cylinder. Stop collecting distillate when you have 8
to 10 ml. Do not collect more than 10 ml and be sure to
accurately measure the exact volume you collected. Weigh the
graduated cylinder with the distillate in it and calculate the
weight and density of distillate. You will be able to
determine the percent ethanol in the distillate from these
data as described below.
IMPORTANT INFORMATION ABOUT THE REPORT:
From the data you collected from each distillation
discuss the effectiveness of each method. Which method gave
the better separation. Tabulate the data from each
distillation. Graph the data in the table below and, using
that graph and the density you obtained, show the percent of
ethanol in your fractional distillation condensate.
WEIGHT % EtOH
0.0
10.0
20.0
30.0
40.0
50.0
60.0
DENSITY (g/ml @ 20 degrees C)
1.000
0.982
0.969
0.954
0.935
0.914
0.891
26
70.0
80.0
90.0
100.0
0.860
0.838
0.818
0.780
END OF EXPERIMENT.
© 2009 Stephen Anderson and Robert Shine
27
LAB 6: THIN LAYER AND GAS CHROMATOGRAPHY
PURPOSE:
You will learn how to separate complex mixtures of small
samples of solids using thin layer chromatography and to
separate complex mixtures of solids, liquids, and/or gases
using gas chromatography. Thin layer chromatography
experiments of an analgesic drug mixture will be performed and
gas chromatography will be demonstrated.
BACKGROUND INFORMATION:
Historically, the term chromatography referred to the
separation of colored substances, particularly plant pigments.
Over time, many variations of chromatography were developed
and the technique extended to colored and colorless substances
in any physical state (solid, liquid or gas). In general, the
technique is based on the different absorption characteristics
of substances on some non reactive material. A series of
absorptions and de-absorptions occur as some eluting material
is passed through the chromatography container.
The many types of chromatography include:
1. Thin Layer (TLC) - best for very small samples of solids or
liquids.
2. Gas (GC) - best for small samples of volatile liquids or
gases.
3. Column - best for larger samples of solids or liquids (see
next experiment).
4. High Pressure Liquid (HPLC) - best for larger samples of
liquids.
5. Paper - best for biological or polar samples.
6. Ion Exchange - best for ionic substances.
7. Electrophoresis - best for large molecule biological
samples.
THIN LAYER CHROMATOGRAPHY - TLC started to be used in
many laboratories in the late 1950's or early 1960's. This
technique (which is generally qualitative and not preparative)
uses a stationary phase of silica gel or alumina which is
adhered to a backing material of plastic, glass or (less
frequently) metal. Often an ultraviolet absorbing substance
is placed in the binder material on the plate so that spots
can be more easily visualized when the chromatogram is
finished. The TLC plate is usually 2 by 4 inches to 8 by 8
inches. A concentrated solution (using methanol as the
solvent) of the mixture to be separated is prepared and
spotted on the TLC plate about 3/4 of an inch from the bottom
using a spotting capillary pipette made from a melting point
tube. The spot is made as small as possible. After the
methanol solvent has evaporated, the TLC plate in placed in a
28
developing chamber which has an eluting solvent that is
appropriate to separate the components of the mixture.
Capillary action causes the eluent to rise up the TLC plate.
The plate is removed from the developing chamber when the
eluent is about 1/2 inch from the top of the plate. The
b
a
Rf =
distance spot moved
distance solvent moved
=
a
b
eluting solvent is allowed to evaporate in a fume hood and the
chromatogram is studied in an ultraviolet light chamber to see
the separation that has been obtained. The location of a
given spot is recorded as an Rf value (retardation factor)
which is given by the formula shown above. An Rf value has no
units and always has a value between 0 and 1.0. In today's
experiment, you will try to separate a mixture of analgesics
whose structures appear below using TLC.
GAS CHROMATOGRAPHY - GC is a technique that can be used
to separate the components of a volatile (liquid or gas)
mixture. Usually the method is qualitative in that components
can be identified but not collected for further use. To
perform a GC one uses an instrument made for the purpose.
Such instruments can cost from $1,000 to $50,000 dollars
depending on the abilities and complexity of the instrument.
In order to use this instrument it must be set up correctly.
The temperature of three areas needs to be set and come to
equilibrium. These areas are the injection port, the column
oven and the detector area. The injection port is usually set
at a temperature above the highest boiling component. The
oven temperature (which can be programmed in more expensive
equipment) is set to give the best separation in the quickest
time.
29
Known Analgesics:
CH3
CH
CH3
CH2
O
C
OH
H
O
O
CH3
H3C
C
COOH
Ibuprofen
2-(4-Isobutylphenyl)propionic acid
Aspirin
Acetylsalicylic acid
H
O
N
C
HO
O
CH3
H3C
O
CH3
N
N
N
N
CH3
Acetaminophen
4-acetamidophenol
Caffeine
Usually this takes some experimentation to determine the best
values. Generally, the higher the oven temperature, the
faster the substance will be removed from the column and the
worse the separation of components. The detector temperature
is set at a value to reduce condensation in that chamber.
The separation of the components of the mixture depends
on many factors that can be controlled by the operator of the
instrument. The stationary phase in the column must be chosen
so that a good separation is possible. Generally, columns are
chosen and purchased for a particular type of separation. The
oven temperature is very important as noted above. The
pressure of the carrier gas (usually helium or nitrogen) is
also important. Higher gas pressures move the substances
along more quickly. One must achieve a balance between speed
of the chromatogram and the separation that is desired.
The detector identifies when a substance has been eluted
from the column. In some cases, the detector can give further
information about the nature of the eluted material. The
simplest (and least expensive) detector is a thermal
30
conductivity detector and can only identify that a substance
has passed the detector. More sensitive detectors include
flame ionization and electron capture detectors. The best
detectors are the mass spectrometry detectors which can give
detailed information about the eluted substance.
Usually in GC one obtains a strip chart recording of the
chromatography results on a piece of paper. A series of peaks
of varying heights is obtained. Peaks appear at certain time
intervals called the retention time. The retention time is
determined by many experimental factors and is not recorded in
the literature as a physical property of a substance. Rather,
they are used to infer some information about the nature of
the mixture. If you suspect the identity of a component of
the mixture, you should run a chromatogram of a known sample
and compare the results. The retention times should be the
same if the substances are the same. Also, keep in mind that
two different compounds may have the same retention time and
will appear as one peak on the chromatogram. This can give
misleading results.
In today's laboratory period the Instructor may
demonstrate the use of a gas chromatograph.
EXPERIMENTAL PROCEDURE:
Obtain 2 TLC strips (keep the strips away from moisture),
5 spotting capillaries and 1 developing chamber. Place 2 to 6
ml of the developing solution (47.5% ethyl acetate, 2.5%
acetic acid and 50% hexane) in the developing chamber and
place a cover on the chamber. Be sure there is enough
developing solution in the bottom of the chamber to wet the
entire bottom of TLC strip and for capillary action to wet the
entire strip during development. Do not allow the chamber to
be open to the atmosphere any longer than absolutely necessary
as volatility will alter the composition of the solvent
mixture. You may wish to use a fresh sample of the developing
solution for your second run. Be careful not to mix spotting
point capillaries and place them in the discarded glass
container when you are finished.
Since the silica gel strips are reasonably expensive,
each team will receive a practice strip and 2 TLC strips for
the experiment. Practice your spotting technique on the
practice strip. Then, carefully spot pure samples and the
unknown on the silica gel TLC strips, one sample at a time.
Be sure to place each spot high enough so that it will be
above the developing solvent in the bottom of the chamber when
the strip is placed in the chamber. Be sure each spot is as
small as possible (about 1 mm in diameter) and accurately
record what you are placing on each spot. On each strip,
place three known samples and your unknown sample. Use
acetaminophen, aspirin, and caffeine as your choice for known
31
samples. Develop the chromatogram and remove the strip when
the solvent front stops moving up the strip or when the
solvent front is about ½ inch from the top of the strip.
Evaporate the eluting solvent and visualize the spots in the
UV chamber. Carefully draw a light pencil circle around each
spot so they can be seen in normal room light. Compare the Rf
values of your unknown sample to those of your known standards
in order to identify the unknown, if possible.
IMPORTANT INFORMATION ABOUT THE REPORT:
Write your report about the thin layer chromatography
experiment only. Do not include anything in your report about
gas chromatography. In your report, give the Rf values of
each substance and discuss the conclusions you are able to
make about the unknown. Were you able to identify the unknown
you used?
END OF EXPERIMENT.
© 2009 Stephen Anderson and Robert Shine
32
LAB 7: COLUMN CHROMATOGRAPHY
PURPOSE:
You will learn how to separate complex mixtures of solids
or liquids and collect the various components using column
chromatography. In this experiment, the separation of
ferrocene and acetylferrocene will be done.
BACKGROUND INFORMATION:
Column chromatography is a technique in which an
absorbing material (usually alumina or silica gel) is placed
in a glass column and a mixture is separated by passing it
down the column under the influence of an eluting solvent.
The solution is collected as it drains from the column and the
pure component is collected by evaporating the eluting
solvent. It is important to never let the column of absorbent
become dry during the process. Alumina comes in four grades
of activity depending on its moisture content.
To prepare the column, one generally fills the column
(which looks like a burette without graduation markings) to
about 2/3 full with a suitable solvent. The alumina is then
added slowly with the stopcock partially opened so that the
adsorbent settles evenly in the chromatography column. It is
important that the alumina be nicely distributed in the column
for best separations.
Next, the mixture to be separated must be carefully added
to the top of the column. This is usually done by making a
very concentrated solution of the mixture and adding it drop
wise from a dropper to the top of the column. It is important
to have the minimum amount of solvent above the column at this
point so that a concentrated band of mixture can be applied to
the top of the column. Again, be careful that no part of the
alumina column becomes dry.
Once the sample is placed on the column, an eluting
solvent is added and the solvent dripping from the bottom of
the column is collected. If the components are colored (as in
our experiment) it is easy to see how they move down the
column. Fractions are collected in a pre-weighed flask or
beaker, the solvent is evaporated and the resulting material
is weighed and its melting point is determined.
Components of a mixture are separated based on
differences in polarity and/or molecular size. Ferrocene is
less polar than acetylferrocene due to the added polarity of
the acetyl group. The acetylferrocene ‘sticks’ to the alumina
more than the ferrocene when one elutes with a non polar
solvent such as hexane and this is what causes the separation.
33
O
C
Fe
Fe
Ferrocene
M.Mass 186, mp 172-174oC
Acetylferrocene
M.Mass 228, mp 85-86oC
CH3
EXPERIMENTAL PROCEDURE:
A chromatography column is prepared in the hood by
placing enough hexane in the column to fill it about 2/3 full.
Then, with the stopcock partially opened, add enough powdered
alumina to this column to fill the column 2/3 full. Never let
any of the alumina become dry. At this point, drain the
hexane from the column until it is about 1/2 inch above the
top of the alumina. Now take 100 mg of the ferrocene acetylferrocene mixture and mix it with about 100mg of
alumina. Carefully open the stopcock to a very slow drip
rate. We may wish to reuse some of the solvents used in this
experiment so carefully follow your Instructor's directions if
this is to be done. As this is happening, add the
alumina/sample mixture to the top of the column. Add one ml
of hexane at a time to keep the column from running dry and to
concentrate the sample on the top of the column. Once the
sample is fully adsorbed on the column add about 5 - 10 ml of
hexane and continue the elution process at a reasonable rate.
Add more hexane to the top of the column as needed. Watch the
colored material as it moves down the column and when it
reaches near the bottom of the column place an empty preweighed glass container below the stopcock to collect the
eluting solution. Continue to collect the solution until all
the colored material has been collected. Set this container
(called container A) aside for now. Place any container under
the stopcock to collect eluting solvent and now start to add a
50/50 mixture of hexane - diethyl ether as the eluting
solvent. At this point, the other colored substance will
begin the move down the column. Collect it as you collected
the other sample in a separate pre-weighed container (called
container B).
Using a slow stream of air in the hood, evaporate the
hexane from container A and the hexane - ether from container
B. Weigh each of these containers, determine the mass of each
product and calculate the percent recovery for each. Take the
34
melting point of each recovered sample. Be sure to save your
samples for the upcoming infrared spectroscopy lab experiment.
IMPORTANT INFORMATION ABOUT THE REPORT:
Be sure to calculate and report the percent recovery of
each substance. Show your calculations and use correct
significant figures. Give your observed melting points and
compare them to the literature values. Discuss your results.
END OF EXPERIMENT.
© 2009 Stephen Anderson and Robert Shine
35
LAB 8: OPTICAL RESOLUTION
PURPOSE:
You will learn how to resolve an enantiomeric mixture and
to use a polarimeter to measure optical rotation. A
resolution of two optical isomers will be performed using a
recrystallization technique.
BACKGROUND INFORMATION
A racemic mixture of two enantiomers can be difficult to
separate since the two isomers differ only in their sign of
rotation of plane polarized light (a physical property) and
their reactivity with other chiral compounds (a chemical
property). One can separate a racemic mixture by using a
chiral stationary phase in a column chromatogram. Another
method (which will be used in this laboratory experiment and
shown in the reaction below) is to convert the racemic mixture
to a diastereomeric mixture by reacting the racemic mixture
with an optically active reagent.
OH
NH2
NH2
+
H2N
H2N
H
H
HO
L-(+)-tartaric acid
(M.Mass 150, mp 170-172 oC)
racemic mixture of
trans-1,2-diaminocyclohexene
(M.Mass 114, mp 14-15 oC)
NH3
NH3+
+
_
O
CO2H
CO2H
_
O
O
O
OH
OH
H
(S,S) -1,2-Diammoniumcyclohexane
mono-(+)-tartrate
soluble in water
OH
H
H
O
H
OH
O
_
O
+
_ H3 N+
O
H3 N
(R,R) -1,2-Diammoniumcyclohexane
mono-(+)-tartrate
Insoluble in water
36
The diastereomers can be separated from one another
because their physical and chemical properties differ more
than the original enantiomers. After the separation has been
achieved, the original chiral compound must be regenerated by
some chemical means. You should review the chapter on
stereochemistry in your organic chemistry lecture course
textbook before you begin this experiment.
Often the differences in properties of the diastereomers
are not great so good technique must be used to achieve
successful results. Further, the conversion to the
diastereomers and the regeneration of the chiral starting
material must be easily accomplished or mechanical loses will
render the technique useless.
In this experiment a racemic mixture of a diamine called
trans-1,2-diaminocyclohexane (basic substances) will be
converted to a diastereomeric salt mixture using optically
pure L-(+)-tartaric acid. This salt mixture will be separated
using crystallization techniques learned in an earlier
experiment. The salt from the (R,R) diamine is less soluble
than the salt from the (S,S) diamine and can be separated by a
vacuum filtration. The degree of success of your separation
will be shown by the angle of rotation (alpha) you will
measure for this salt in a polarimeter and comparing it to the
specific rotation ([alpha]) by using the formula
[alpha] = alpha / c x l
where c is the concentration in g/100ml and l is the length of
the polarimeter tube in decimeters.
EXPERIMENTAL PROCEDURE:
Working in a fume hood, dissolve 0.750 g (5.0 mmol) of L(+)-tartaric acid in 2.5 ml of distilled water. Stir the
solution and add 1.140 g (10.0 mmol) of trans-1,2diaminocyclohexane. Warm the solution for a brief period.
Allow the solution to cool. Add 0.5 ml of glacial acetic acid
and cool the reaction mixture in an ice bath for 20 to 30
minutes. Collect the precipitate that forms by vacuum
filtration and rinse it with 0.5 ml of ice cold water followed
by four 0.5 ml portions of room temperature methanol. Dry your
product, record its mass and calculate the percent yield. If
needed, recrystallize the crude product using boiling water.
You will not regenerate the liquid diamine in this experiment.
You will use the optically pure salt for the polarimeter
reading described in the next paragraph.
To check the success of your separation, measure the
optical resolution of the solid salt you obtained. To do
this, dissolve 0.200 g of the salt in 10.00 ml of distilled
37
water and place this solution in a 1 decimeter polarimeter
tube. The salt is not very soluble in water, so do not use
more than 0.200 g of the salt. Use a 10.00 ml volumetric
flask for an accurate volume measurement. Record the angle of
rotation that is observed. Be sure all weight and volume
readings are accurately measured and recorded. From your
data, compute the specific rotation and compare your value to
the literature value which is 12.4 degrees.
IMPORTANT INFORMATION ABOUT THE REPORT:
Be sure you accurately record all important measurements
that you make in this experiment. You will need to compute
your percent yield (see Appendix D) of the optically pure salt
you obtained and its specific rotation which you will compare
to the literature value.
Also, in your report discuss in detail how you would
convert your salt back to the free diamine. Describe the
experimental details that would be needed to regenerate your
optically pure diamine. Using actual physical properties,
tell what difficulties this procedure would cause.
END OF EXPERIMENT
© 2009 Stephen Anderson and Robert Shine
38
LAB 9: NUCLEOPHILIC SUSTITUTION REACTIONS
PURPOSE:
You will learn and use the techniques of classical
qualitative organic analysis to study the chemical differences
between Sn1, Sn2, E1, and E2 reactions of organic halide
compounds.
BACKGROUND INFORMATION:
Substituted alkanes can undergo displacement or
elimination reactions to give new compounds. These reactions
can be useful in helping to identify the nature of the
compound or to synthesize new classes of compounds. Next
semester, you will synthesize an alkene by an E1 elimination
reaction of an alcohol. In this experiment you will study the
differences among the displacement reactions of alkyl halides.
Nucleophilic Substitution Reaction:
R
X
+
Nu-
R
Nu
+
X-
As a background for this experiment, you should read the
relevant chapter in the lecture course textbook. There are
many facets of these reactions that you should be aware of.
Hopefully, this experiment will help you better understand
these classes of reactions.
The four types of reactions we will consider include two
substitutions (Sn1 and Sn2) and two eliminations (E1 and E2).
There are four different mechanisms (reaction pathways) and
the mechanism that will predominately occur is dependent on
many factors such as: (1) the nature of the alkyl group, (2)
the nature of the leaving group, (3) the nature of the
solvent, and (4) the temperature of the reaction. It is
possible to have mixtures of products depending on how the
reaction is carried out. Recall before we start that
substituted alkanes can undergo these reactions readily but
substituted alkenes or alkynes generally do not undergo these
reactions. Further, substituted aromatic compounds are not
reactive.
In this experiment we will use reactions of alkyl and
aryl halides. We will vary the halide between chloride,
bromide and iodide. We will use various primary, secondary,
and tertiary alkyl compounds. We will use a reaction (silver
nitrate in aqueous alcohol) that favors the Sn1 mechanism and
a reaction (sodium iodide in acetone) that favors the Sn2
mechanism. A positive test for the silver nitrate reaction
39
will be the precipitation of silver chloride, silver bromide
or silver iodide depending on the halide being tested. Silver
halides darken on exposure to light. Also, silver nitrate
will cause a black stain if it touches your skin. This stain
will go away in a few days. A positive test for the sodium
iodide reaction will be the precipitation of sodium chloride
or sodium bromide depending on the alkyl halide being tested.
Note that sodium iodide has reasonable solubility in acetone
while sodium chloride and sodium bromide are insoluble in
acetone. We will not directly study elimination reactions in
the experiment but elimination may occur as an unwanted side
reaction you should be aware of. You should address and
discuss this complication in your report.
SN1 Mechanism: Occurs in two steps, rate dependent on [R-X] only
R
R
R
C
R
X
R
R
R
C+
R
+ X-
C+
R
+ Nu-
R
R
C
Nu
R
Relative stability of Carbocation generated in step one is important
SN2 Mechanism: Occurs in one concerted step, rate dependent on both [R-X] and Nu-
R
R
Nu- + R
C
R
X
Nu-
C
R
X
R R
Nu
C
R
+ X-
R
Alkyl halide must be relatively free of steric hindrance
Bromide is a better leaving group than chloride so the
reaction of alkyl bromides should occur more rapidly than
alkyl chlorides. Primary alkyl groups favor the Sn2 mechanism
and tertiary alkyl groups favor the Sn1 mechanism. Both
40
mechanisms may apply in the case of a secondary alkyl group.
Alkyl groups that can give a stable carbocation (carbenium
ion) such as the allyl or benzyl group tend to show an Sn1
mechanism even though they are strictly primary groups.
In your tests, you should note whether or not you observe
formation of a precipitate and, if you do, how many minutes
were needed at room temperature and whether heat needed to be
applied. No test should take more than 20 minutes and many
tests can be done concurrently if you carefully label your
test tubes.
EXPERIMENTAL PROCEDURE:
Halogenated organic waste is usually kept separate from
non-halogenated waste. The disposal of halogenated organic
wastes is more costly. Be sure to use the proper waste
containers in this experiment.
Carefully label a number of test tubes and begin heating
a beaker of water on a hotplate in the hood for use as a hot
water bath. You should test the following substances with
silver nitrate in aqueous alcohol in one test tube and sodium
iodide in acetone in a second test tube: 1-chlorobutane, 1bromobutane, 2-chlorobutane, 2-bromobutane, 2-methyl-2chloropropane (t-butyl chloride), benzyl chloride,
bromobenzene, iodoethane. The test is done by placing 1.0 ml
of the test solution (either silver nitrate solution or sodium
iodide solution) in a labeled test tube and adding 10 drops of
the halide to be tested. Be sure to shake the tube to mix the
reagents. Let the tube sit at room temperature and observe
the time needed for precipitation to occur and the approximate
amount obtained. If no precipitation is noted in 5 minutes
place the tube in the hot water bath and observe the time
needed for precipitation to occur. If no precipitation is
noted in 10 more minutes, record the alkyl halide as
unreactive. Be careful to note any evaporation of solvents
which may give you erroneous results. Record your results
noting the formation of any precipitate and whether time or
heat was needed for any precipitation to occur.
You must dispose of halogenated waste in a different
container than other organic waste.
IMPORTANT INFORMATION ABOUT THE REPORT:
Carefully tabulate the results of all your tests in a
nice table format. Discuss your findings and note how your
results support the facts you have learned about substitution
reactions. Note any discrepancies you have found and give
41
reasons for these anomalies. Comment on the use of the sodium
iodide reaction in acetone for use with alkyl iodides.
END OF EXPERIMENT.
© 2009 Stephen Anderson and Robert Shine
42
LAB 10: INFRARED AND ULTRAVIOLET SPECTROSCOPY
PURPOSE:
You will learn how to take and interpret infrared (ir)
and ultraviolet (uv) spectra.
BACKGROUND INFORMATION:
Spectroscopy is used in the organic chemistry laboratory
to aid in the identification of compounds. The major types of
spectroscopy used in organic chemistry are: infrared (IR),
ultraviolet (UV), nuclear magnetic resonance (NMR or PMR for
proton magnetic resonance) and mass spectroscopy (MS). All of
these types of spectroscopy are described in the textbook for
the lecture part of this course and that book should be
consulted in your study of these topics. Please be sure to
bring your Solomons text book to the two lab sessions on
spectroscopy. Infrared and ultraviolet spectroscopy will be
covered in this experiment and nuclear magnetic resonance will
be covered in the next experiment. Mass spectroscopy will be
only briefly mentioned. The types of spectroscopy most useful
to the organic chemist are IR and NMR.
Infrared spectroscopy uses electromagnetic radiation in
the infrared region. This IR energy is higher than microwave
and lower than visible energy. If you were hit with IR
energy, you would feel warm. IR lamps are often used in fast
food establishments to keep food warm. In the IR instrument,
a glowbar or Nernst glower was often used as the source of IR
energy. Since glass absorbs IR energy, glass cannot be used
in the instrument or in sample holders. Salt (NaCl or other
salts) is most often used as a sample holder (so water must
not be in the sample) and plastic diffraction gratings are
often used in the optics of the instrument to disperse the IR
energy into a spectrum.
Infrared energy is absorbed by organic molecules which
causes vibrational and rotational changes in the various
chemical bonds. This absorption is quantized which means that
only certain wavelengths are absorbed depending on the nature
of the bonds in the given molecule. Hence, different
molecules give different spectra. As such, IR spectra can be
interpreted in two useful ways: (1) a certain functional group
will often give a characteristic absorption in the spectrum
and (2) the sum total of all the absorptions in the spectrum
serves as a ‘fingerprint’ for a given molecule. No two
different molecules will have the same IR spectrum. Computer
matching of an unknown sample with a database of spectra can
be used to tell what the unknown is likely to be. This is
easier said than done as IR spectra depend on sample size and
impurities will affect the spectrum.
43
Most often, IR is used to help identify the presence or
absence of functional groups in a molecule. You will make use
of this fact in today’s experiment. The two most easily
identified functional groups are -OH (found in alcohols and
carboxylic acids) and C=O (carbonyl found in aldehydes,
ketones, carboxylic acids, esters, and amides). The -NH group
(found in primary and secondary amines and amides) is also
easily seen.
In ultraviolet spectroscopy electromagnetic radiation in
the ultraviolet region is used. This energy is more energetic
than visible light and causes the electronic excitation in
molecules. Often a UV instrument can also produce the visible
spectrum of a sample. Visible spectra are useful if the
sample appears colored to the human eye. Good quality lamps
are used in these spectrometers. Ultraviolet energy is
absorbed by glass so more expensive quartz sample holders are
used. Usually the sample cell has a light path of 1
centimeter.
Since UV energy is more energetic than IR, the spectra
often show very broad bands as there are many energy overtone
absorptions in the band. As such, UV does not give much
specific information. Ultraviolet spectroscopy is useful in
identifying broad classes of compounds by identifying so
called chromophore groups. Usually compounds rich in pi
electrons are best identified by this technique. These
include aromatic compounds and polyunsaturated compounds.
EXPERIMENTAL PROCEDURE:
In the summer of 2008 a new Thermo Scientific Nicholet
6700 FT-IR spectrometer was installed at Ramapo. Your
instructor will show you how to take an IR spectrum using this
instrument. You will then take spectra of known and unknown
samples. You will then try to determine the nature of the
unknowns. Further, if time is available, you should also
record the infrared spectra of your separated ferrocene and
acetylferrocene samples from a prior lab experiment to test
the effectiveness of your separation.
The instructor or the Laboratory Coordinator may
demonstrate how to take a UV spectrum. You will not do any
work with the UV spectrometer.
IMPORTANT INFORMATION ABOUT THE REPORT:
No formal laboratory report will be due for this lab.
However, the use and analysis of IR and NMR spectra will be a
significant part of the laboratory final examination.
44
END OF EXPERIMENT.
© 2009 Stephen Anderson and Robert Shine
45
LAB 11: NUCLEAR MAGNETIC RESONANCE SPECTROSCOPY
PURPOSE:
You will learn how to take and interpret hydrogen nuclear
magnetic resonance (NMR or PMR) spectra.
BACKGROUND INFORMATION:
Nuclear magnetic resonance spectroscopy uses certain
quantized properties of nuclei. The irradiation energy used
in this technique is in the radiofrequency (FM) region which
is generally recognized as safe. But, in order to obtain a
meaningful spectrum, the sample must be placed in a very
strong magnetic field. If you had an MRI scan you had an NMR
taken of a part of your body. The MRI scan shows contrast
between aqueous and fatlike (lipid) tissue in a body.
Only certain isotopes are NMR active. There has to be a
certain ratio of protons and neutrons in a nucleus for a
nucleus to be NMR active. Fortunately for organic chemists,
the most abundant isotope of hydrogen (H-1) and C-13 (present
in about 1.1% of all carbon atoms) are NMR active. To
distinguish the isotopes being used, hydrogen NMR is often
called proton magnetic resonance (PMR). Recall that the
hydrogen nucleus is a proton. We will use NMR and PMR
interchangeably in our discussion.
A reference material is used to calibrate the recorder
paper to the spectrum. This substance in called
tetramethylsilane and is abbreviated as TMS. It appears at 0
delta on the chart and is most often the rightmost peak on the
spectrum. TMS has a boiling point around 25 degrees C and is
often sold in deuteriochloroform which is used as the solvent.
Since hydrogen absorbs in the NMR region, the solvent should
not have any hydrogen atoms in it. The most common solvent
used is deuteriochloroform. Special glass tubes are used as
the sample holder. For us, the sample should be evenly
distributed in the magnetic field so a spinning liquid sample
works best. Since the technique is not very sensitive, a
concentrated solution should be used. If a pure liquid is
used as the sample, it is often referred to as ‘neat’.
You will spend much time in this lab trying to interpret
NMR spectra. This technique is very powerful in identifying
small organic molecules. In your interpretation you try to
build up a hydrogen structure for the molecule and, from those
pieces, you can often deduce the overall structure of the
compound.
There are three different aspects of the NMR spectrum you
will need to consider: (1) the chemical shift, (2) the
splitting pattern, and (3) the integral trace. The chemical
shift is measured in parts per million (ppm or delta) from the
46
TMS peak. The chemical shift is affected by the size of the
electron cloud around the nucleus. Hence, highly shielded
(much electronic cloud) nuclei appear ‘upfield’ near TMS.
Deshielded nuclei (low electron cloud) appear ‘downfield’ from
TMS.
The splitting pattern is often difficult for beginners to
understand but gives the most information about a given
hydrogen atom’s nearby structure. A H absorption is split by
the number of hydrogen atoms on the nearby carbon atoms.
A
hydrogen that has no hydrogen atoms on the neighboring carbons
is not split. If there is 1 hydrogen on the adjoining
carbons, the single peak is split into a doublet, and so on.
Hence, the splitting pattern is always one more than the
number of hydrogen atoms on the adjoining carbon atoms.
The integral trace gives the ratio of the various
different types of hydrogen atoms. Hence, an ethyl group
gives a 2:3 integral trace.
The above discussion only briefly touches on the NMR
technique. You should carefully study the course textbook for
more information and many problems. You should work on as
many NMR and IR problems in those chapters as you are able.
EXPERIMENTAL PROCEDURE:
The Instructor will show you how to take an NMR spectrum.
You will then take the spectrum of a known sample and an
unknown sample. You will practice identifying unknowns from
handout sheets and then you will try to determine the nature
of the unknowns.
IMPORTANT INFORMATION ABOUT THE REPORT:
No formal laboratory report will be due for this lab.
However, the use and analysis of IR and NMR spectra will be a
significant part of the laboratory final examination.
END OF EXPERIMENT.
© 2009 Stephen Anderson and Robert Shine
47
LAB 12:
SYNTHESIS OF ASPIRIN
PURPOSE:
In this experiment, the student will use the technique of
microwave heating to synthesize acetylsalicylic acid
(aspirin).
IMPORTANT REACTION:
O
O
C
OH
OH
Salicylic acid
M.Mass 138.12
mp 159oC
O
H+
O
C
+
OH
O
H3C
O
CH3
Acetic anhydride
M.Mass 102.09
bp 140oC
O
CH3
Acetylsalicylic acid
M.Mass 180.15
mp 128-137oC
BACKGROUND INFORMATION:
In this experiment the use of microwave heating will be
investigated. The synthesis of acetylsalicylic acid from
salicylic acid and acetic anhydride was chosen for its
simplicity and reliability when other methods of heating are
used.
The use of microwave energy for heating dates back to the
1960’s and became widespread in the 1980’s. Since microwave
energy can penetrate deep within a molecule, the entire
molecule can be heated at the same time. Conventional heating
is generally done by conduction and heats from the outside to
the inside of the sample. Therefore, microwave heating can be
faster than conventional heating. However, for heating to
occur, the microwave energy needs to be absorbed by the bonds
in the molecule. Absorption of energy is better in polar
bonds and poor in nonpolar bonds. Hence, water heats rapidly
in a microwave oven. It is said that alkanes do not heat well
with microwave energy.
The design of a microwave oven presents some challenges
to the designer. It is hard to have a uniform distribution of
microwave energy in the cabinet. Further, if no absorption of
microwave energy occurs a reflection of energy in the unit can
destroy the apparatus. Some manufacturers warn that operating
a microwave oven when empty can ruin the oven.
Microwave ovens can also cause fires. Many people who
have made popcorn in a microwave oven have seen burnt popcorn
48
kernels at one time or another. It is important to remember
that three things are needed for a fire to start: a source of
ignition such as microwave energy, a combustible substance
such as many organic compounds, and a source of oxygen such as
air. It is important that one follows the experimental
details carefully to work in a safe manner.
Microwave energy can penetrate the human body and can
interfere with certain electronic devices that may be
implanted for health reasons (such as cardiac pacemakers). If
you have implanted electronic devices, let the Instructor know
and/or avoid the lab while the microwave unit is turned on.
Please make careful observations during this experiment
and comment on any success or failure in your report.
EXPERIMENTAL PROCEDURE:
You will use the laboratory microwave in this experiment.
Be especially careful when handling the reaction tubes as they
are delicate and expensive to replace.
Place salicylic acid (1 g, 7.24 mmol) in a microwave
reaction tube and add acetic anhydride (2 ml) and concentrated
sulfuric acid (2 drops). Add a magnetic stirrer bar, then
screw on the cap and place your tube in the microwave
turntable. Record the position number of your reaction tube
on the turntable. When all of the tubes have been added, your
Instructor will place the turntable into the microwave unit,
close the door and select the correct file-setting for your
reaction (Aspirin.rot). The program will take 12 minutes to
run, followed by a 5 minutes vent and cool down program. On
completion, your instructor will open the microwave door.
Carefully remove your reaction tube, allow the reaction
mixture to completely cool to room temperature then empty the
contents into a beaker containing approximately 5 ml of water.
If the reaction mixture is not properly cooled before adding
the water, an oily product may be produced. Cool the contents
in an ice bath, and scratch with a glass rod if necessary.
Filter your product using a Buchner funnel, washing the
reaction flask with water. Dry your product, record the mass
and melting point of your crude product, and calculate the
percent yield. If required, recrystallize your product from
hot 2-propanol (isopropyl alcohol), re-recording the mass,
melting point and percent yield of your purified product.
IMPORTANT INFORMATION ABOUT THE REPORT:
The report for this experiment will follow the format for
synthesis reactions. Be sure the percent yield calculation
(see Appendix D) is carefully done. Also, record the melting
point range of the final product and compare that melting
49
point to the reported melting point of acetylsalicylic acid.
Using these data, discuss the relative success on the
experiment. Also, be sure to critique the effectiveness of
using microwave heating to cause the reaction to occur.
END OF EXPERIMENT.
© 2009 Stephen Anderson and Robert Shine
50
APPENDIX A:
SAFETY ISSUES
INTRODUCTION:
A very important aspect of laboratory work is working
safely. It is your responsibility to yourself and the other
people in the laboratory to work in a serious and careful
manner. If you are not sure about the dangers involved in
what you are assigned to do, be sure to ask the Instructor
before you begin. You should consult the material safety data
sheets (msds) for those chemicals you will be working with
before coming to the lab. You can find msds data by typing
the chemical name followed by msds in a Google search box.
For example msds data for acetone could be found by typing
acetone msds in the search box.
If you have allergies, you should consult your Allergist
for advice in working with organic chemicals. If you are
pregnant or become pregnant, you should consult your doctor
for advice in working with organic chemicals.
There are risks involved with all activities you do. By
knowing what those risks are and taking prudent actions, you
can lessen the dangers you face and thereby lead a relatively
safe life. The dangers you are likely to face in the organic
chemistry laboratory are due to:
1. Equipment
2. Toxic chemicals
3. Flammable chemicals
EQUIPMENT:
You will use glassware, plastic ware, and electrical
heating equipment in the lab. As you know, glass items can
break and cause cuts that can be minor or severe. You should
exercise care when handling glassware to minimize breakage.
If a glass item breaks, notify the Instructor who will then
use a broom and dustpan to clean up the area. Broken or
discarded glassware must be placed in the separate container
that is marked for broken glass. No other items other than
glassware should be placed in that container.
Plastic ware generally has no safety concerns and is safe
to use.
Heating in the lab will generally be done using
electrical heating equipment such as hot plates and heating
mantles. We will not use Bunsen burners in the laboratory.
We may use steam baths on occasion. Steam can cause a bad
scalding burn. Hotplates and heating mantles do not change
appearance when they are hot so you should always assume they
are hot until proven otherwise. There can also be an
electrical shock hazard with any electrical equipment. This
laboratory is equipped with ground fault circuit breakers to
51
minimize electric shock hazard. Sometimes these breakers trip
so if you notice that an outlet is not working, notify the
Instructor so it can be reset.
Also be aware that heating equipment can be hot enough to
exceed the flash point of a chemical and can start a fire.
Chemicals that have a low flash point can be ignited by a warm
hot plate. Evaporating diethyl ether from a beaker on a
hotplate caused a fire in our organic chemistry laboratory
many years ago.
TOXIC CHEMICALS:
Eye protection must be worn in the laboratory whenever
anyone in the lab is working with chemicals. Only department
approved eye goggles can be used. If, at any time, you
suspect that a chemical got onto your skin or eyes,
immediately wash the affected area with plenty of cold water
for a significant period of time. Be sure to notify the
Instructor after you began the washing process. Chemicals
that are immediately washed from the skin will cause less
damage than those that are allowed to remain on the skin until
you feel pain. So, even if you feel no pain, wash the
affected area immediately.
Many chemicals have toxic properties which vary from
compound to compound. The best source of information about
the toxic properties of chemicals you will use is the material
safety data sheet for that chemical. Be aware that chemicals
can enter the body by inhalation, ingestion, absorption or
injection through the skin. You should use gloves to protect
the skin and should rinse your hands after handling chemicals.
When you leave the lab you should thoroughly wash your hands.
Ingestion of chemicals can be avoided by not putting
anything in your mouth during the laboratory period. Never
pipette by mouth; use approved suctioning devices. Never eat
or drink any food in the lab, including chewing gum, as it
could have been contaminated.
Inhalation of volatile chemicals can be avoided by
working in the fume hood. Try to minimize the time a volatile
chemical will be exposed to the air outside the hood. Weigh
such substances quickly and move them immediately to the hood.
Some chemicals used in the lab are corrosive. Generally
these include acids and bases. The more concentrated the acid
or base, the more corrosive it will be. So use extreme care
with concentrated sulfuric acid, hydrochloric acid, nitric
acid, phosphoric acid, and sodium hydroxide solutions. Nitric
acid will cause orange patches to appear on the affected skin
in one to three days. This orange patch will slowly wear off.
Silver nitrate will cause black stains to appear in the
affected area within a few hours after exposure and will wear
off in a few days. Volatile organic solvents that enter the
52
body can cause systemic poisoning. Often, the liver or
kidneys are affected. Be sure to read the material safety
data sheet for specific information about the chemicals with
which you are working.
FLAMMABLE CHEMICALS
Many organic chemicals can burn. Halogenated organic
compounds are usually less flammable than others but they tend
to be more toxic, especially to the liver. Compounds with a
high percentage of carbon (such as aromatic hydrocarbons) burn
with a very black sooty flame. Other hydrocarbons burn with a
yellow flame and as the oxygen content of the compound
increases, the flame becomes more bluish.
In order to have a fire start, three things are needed.
You must have a compound that can burn, an oxidizing material
such as air, and an ignition source that will raise the
temperature of the flammable substance above its flash point.
Of real concern is whether or not a compound can cause a
fire. There are terms whose meaning you need to know. These
terms are: nonflammable, inflammable, flammable, and
combustible. Nonflammable means the compound will not burn
under reasonable conditions. They pose no fire hazard. Water
is such a substance. Inflammable is a very bad choice for a
word. The dictionary defines inflammable as "capable of being
set on fire; combustible." Inflammable does not mean "not
flammable." The word inflammable should not be used. A
flammable substance is "one that is easily set on fire" and a
combustible substance is "capable of being set on fire."
Flammable substances, with their lower flash points, pose a
more serious fire hazard than combustible substances.
Gasoline would be considered flammable and wood would be
considered combustible. So, use care when working with
flammable and combustible substances, especially with
flammable substances.
© 2009 Stephen Anderson and Robert Shine
53
APPENDIX B:
COMMENTS ABOUT WRITING STYLE
Before you begin writing lab reports for this course, you
should carefully study the following statements which are
presented to help you achieve good grades. Failure to use
this advice could result in a lower lab grade.
1. On all written work you submit, place your name, the full
descriptive name of the experiment, and the day of the week
(that is, Monday, Tuesday, Wednesday, Thursday or Friday) that
you have lab. Place this information on the first page.
2. Reports are due at the beginning of the next class meeting.
If you computer dies, submit the report on time, neatly
handwritten. Reports cannot be submitted electronically
unless allowed by your Instructor.
3. Use correct grammar and spelling. Use adjectives, nouns and
verbs correctly. Use scientific terms correctly. Proofread
your report and check all spelling as spell checkers do not
recognize the misuse of words (such as to, too and two).
4. Use terms (especially scientific terms) correctly. Do not
misspell the name of a chemical. Failure to use technical
terms correctly will have a definite negative impact on your
report grade. For example, do not report the boiling point of
a solid; report its melting point.
5. Do not be vague or overly explicit. Give only the
important data, observations and conclusions.
6. Your report should have a correct, logical flow. For
example, give the data before the conclusions that are made.
7. When giving a measurement, be sure to include both the
number (with the correct number of significant figures) and
correct units.
8. Pay attention to the correct use of significant figures.
See Appendix C for more information about the correct use of
significant figures.
9. Do not start a sentence with a digit. Spell out any
numbers that start a sentence.
10. Use the third person, passive voice in science report
writing.
11. The data page that you submit when you leave the lab is
part of your lab grade. The data page must be clearly
presented, as neat as possible, and must give all the critical
measurements (numbers and units) you made in the lab.
Computations of results may be left until the final report is
written if they are difficult to do. Simple computations may
be included with the data sheet.
12. Be sure that each section in the report has an
introductory sentence.
13. Cite all references used and be exact in giving each
reference.
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14. The introduction section must have structures of important
compounds or a chemical equation drawn early in the section.
15. The introduction section should deal with the theory of
the experiment and not give experimental details.
16. The experimental section can be brief with a reference to
the lab manual. You need add only the changes you made that
differed from the lab manual.
17. As much as possible, tabulate data and place them near the
text that describe it. Be sure the table has a meaningful
title and that the data are presented in a clear manner.
18. Do not describe a math formula in text. Draw the formula
and separate it from the text by at least one line.
19. For results that are computed from lab data, be sure to
show how the computation was done. That is, show the math
used. When doing computations in the lab report, show the
general formula and then show the formula with your actual
data to show how you obtained your result. Set the
mathematics away from text so it stands out.
20. Fully explain the data. Do not leave any critical data
out of the report. If you calculated the amount of a
substance by measuring the weight of a container empty and
then again with the sample in it, both weights must be in the
report along with the final weight.
21. Be sure data, calculations and conclusions are tied
together. If these items are in different parts of the
report, be sure to tell the reader where they can be found.
22. All data in the lab report should be able to be traced
back to the experimental data you recorded on the data sheet
you submitted. Do not take any shortcuts in data handling.
You must show all relationships in a logical, orderly manner,
and you must show how results were calculated.
23. When giving melting point or boiling point data, be sure
to report these data as a range, and be sure to give the
accepted literature values in parentheses immediately after
your experimental data.
24. The conclusion section may be brief but should tell what
the objectives of the experiment were and if they were
achieved.
25. Do not report a percent error unless told to do so. We
measure percent recovery or percent yield. Be sure you know
the difference between these terms.
26. Repeating errors, such as the misspelling of a chemical,
will be noted only once when reports are graded.
27. Do not capitalize a chemical name unless it is the first
word in the sentence. However, often Microsoft Word will do
this automatically.
28. Mass is not used as a verb in this course. A chemical can
be weighed but not massed. Also, precipitates are filtered,
not funneled.
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29. Two words are often misused in this course. The first is
'separate' which is often misspelled with an 'e' after the
'p'. Fortunately, spell checkers fix this error. The other
misused word is 'vial' which many students write as 'vile'. A
spell checker does not repair this mistake. The dictionary
defines 'vile' as: "mean, worthless, unclean, repulsive, bad."
The same dictionary defines the homonym 'vial' as: "a small
vessel." Be sure to use these two terms correctly.
30. Do not use the first or second person in scientific
writing. Do not use I, we, or you.
31. A rewrite of a previously graded lab report will not be
accepted unless it was authorized and agreed upon by the
Instructor before it was done.
32. More information about good report writing style for
chemists can be found on the Web by doing a Google search on
‘writing style for chemists’.
© 2009 Stephen Anderson and Robert Shine
56
APPENDIX C:
FIGURES
MEASUREMENTS AND SIGNIFICANT
INTRODUCTION:
A very important aspect of laboratory work is making and
reporting measurements carefully. A significant part of your
grade will be based on how you gather and report data. You
will need to know which measurements will need to be made
carefully and which are less important and can be made quickly
with less accuracy so as not to waste time.
There are three terms that you will need to know and
understand: precision, accuracy, and significant figures.
These terms have similar meaning and are often confused.
precision--The ability of repeated measurements to
have the same value. If you weigh an item three different
times, the number of grams should be exact or very close for
all the measurements.
accuracy--The agreement of a measurement with the
accepted value reported for that value. If you measure the
density of water at 4 degrees Celsius, it should be 1.00 g/ml.
significant figures--The number of digits that can
be reported for a measurement. If you weigh some item on the
top loading balance in our laboratory, it can be recorded to
the third decimal place if using units of grams, such as,
1.234 g.
In this laboratory course, the correct use of significant
figures will be the most important of the three. Accuracy
will be the next most important and precision will generally
be disregarded as you will usually not make repeated
measurements of the same data item.
SIGNIFICANT FIGURES:
The number of significant figures you give in your lab
report is governed by the limits of the measuring devices you
use and the care you exercise in making a measurement.
Mathematical manipulation of data can never increase the
number of significant figures in your result.
You will need to know what measurements are important and
which are not. Do not waste time in making a very careful
measurement on some less significant piece of data. For
example, the exact amount of a reagent is needed but the
precise volume of solvent in which it is dissolved is often
not needed.
Be sure to record important data to the level of
precision the measuring instrument allows. If the directions
tell you to measure out exactly 0.500 g of a compound, then
you will have to spend some time to carefully get 0.500 g.
However, often the directions will say to measure
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approximately 0.500 g of a compound. Here, you can quickly
weigh out an amount that is approximately 0.500 g but you must
record your exact weight. For example, you may have weighed
0.476 g. You would then use that figure (0.476 g) in your
report and further calculations.
It is important that you know the difference between 0.05
and 0.050. The first of these (0.05) has one significant
figure while the second (0.050) has two significant figures.
A zero used to place the decimal point is not a significant
figure. So, if you carefully weighed out 0.500 g and wrote it
as 0.5 g or 0.50 g, you would be wrong. You must show all the
significant figures allowed by the measurement you made along
with the correct units such as grams (g), milligrams (mg),
milliliters (ml), etc.
In any measurement, the last significant figure has an
uncertainty associated with it. Often you will see the
balance going up and down over time in the last decimal place.
So a measurement such as 0.476 g really is 0.476 +/- 0.001 g.
In using conversion factors such as 1 pound/453.6 g, the
number 1 is exact and can be assumed to have the same number
of significant figures as the other value. So here, the
conversion factor is 1.000 pound/453.6 g. Also note that
every conversion factor equals one so the reciprocal also is
true. That is, 453.6 g/1 pound is also true. In all your
calculations, be sure to show the numerical data and the
correct units. Both are required.
It is important to consider how to use significant
figures when doing mathematical operations. Using a
calculator which can report 8 digits does not make your result
more accurate. You must report your calculated result to the
correct number of significant figures allowed by the
laboratory measurements and not the calculator.
If the mathematical operation is addition or subtraction,
your result can only be reported to the decimal place of the
least accurate measurement. The following examples may
clarify this. If you add 123.4 g to .123 g, the result is
123.5 g. If you add 1.2 g to .123 g, the result is 1.3 g. If
you subtract 1.234 g from 1.236 g, the result is .002 g. Note
that in addition and especially subtraction, you can appear to
lose significant figures. So, in such cases, if one
measurement can only be made to the first decimal place, it is
a waste of time to make the other measurement carefully to the
third decimal place.
If the mathematical operation involves multiplication and
division, the result will have the number of significant
figures equal to that of the least significant piece of data.
If you multiply 123 by 12, the result is 1500 which has 2
significant figures and 2 zeros to place the decimal point.
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Such results are better written in scientific notation (1.5 x
103) to avoid confusion.
Be sure you fully understand the meaning and use of
significant figures as a major part of your laboratory grade
will be based on your lab measurements and data manipulation
in your lab reports.
© 2009 Stephen Anderson and Robert Shine
59
APPENDIX D:
PERCENT YIELD CALCULATION METHOD
The calculation of a per cent yield is a very important
part of those labs where a chemical synthesis has been done.
In your report be sure to correctly do the percent yield
calculation (showing the correct significant figures and
units) and present it in a very clear manner. Do not forget
the show the correct units with each measurement as you
perform the calculation. A percent yield calculation is an
application of a weight-weight problem which you learned in
Fundamentals of Chemistry.
You must show how the calculations are done in an area
that is separated from text. Do not put data details in
paragraph text. Present data details in tabular form and be
sure this data is near the text that makes reference to it for
clarity of presentation or, if the calculations have been
attached to the end of your report as an appendix, give a
clear sentence that refers the reader to this page.
Calculations should be close to where the final data are
presented in your report. Briefly, a percent yield
calculation is an application of a weight-weight problem in
General Chemistry. You calculate the number of moles (or
millimoles) of each reagent. Determine which reagent is the
limiting reagent from the balanced equation. Catalysts and
solvents are not usually limiting reagents. Then, determine
from the balanced equation the maximum number of moles of the
desired product that could be obtained and this is your
theoretical yield in moles. Convert this number of moles of
product to a theoretical weight of product. Then, divide the
weight of product that you obtained by the theoretical weight
of product you calculated and multiply by 100 % to get the
percent yield. Alternatively, you could determine the actual
moles of product obtained by taking the mass of the product
isolated and dividing by its molecular weight. The percent
yield in that case would be the actual moles of product
divided by the theoretical moles of product times 100 %. Be
sure to identify which method you use by showing numerical and
unit values for all data in your calculations.
Percent yield
calculations will be a major part of your report grade in
those labs where percent yield data need to be reported.
SAMPLE CALCULATION:
Below is just a
are done presented in
process. This is not
for your lab report.
reaction of a diamine
appeared in Lab 8.
guide to show you how the calculations
a manner to help you understand the
an acceptable percent yield calculation
The method below was adapted from a
with an acid to give a salt which
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The balanced equation showed that 1 mole of diamine reacts
with 1 mole of acid to give 1 mole of salt.
You started with 1.20 ml of diamine.
gives 1.14 g of diamine
1.20 ml * .951 g/ml
1.14 g of diamine / 114 g/mole of diamine gives .0100 mole of
diamine used.
You also weighed some tartaric acid that varied from group to
group. Say your group weighed .732 g.
The number of moles of acid in the .732 g is .732 g / 150
g/mole of tartaric acid or 0.00488 mole of tartaric acid.
From the balanced equation, the acid is the limiting reagent
and the theoretical yield of the product would be .00488 moles
of salt.
From the molecular weight of the product, the theoretical
weight of the product is 0.00488 mole * 264 g/mole of salt or
1.29 g.
If you obtained .847 g of product, your percent yield is
(.847 / 1.29 g) * 100% or 65.7 %.
© 2009 Stephen Anderson and Robert Shine
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APPENDIX E:
CHEMICAL DATA
INTRODUCTION:
This appendix will present physical data and safety data
about chemical substances that are mentioned in the
experiments described in this manual. The data are presented
in an Excel spreadsheet format. Table 1 gives all the data in
one location but the printing is small and may be difficult to
read. Table 2 gives safety information while Table 3 gives
the physical properties of the chemicals. Data were obtained
from chemical handbooks, supplier catalogs, and material
safety data sheets. In many cases, the value for a given
piece of data such as a melting point varied from source to
source within a narrow range. Therefore, you may find a data
value from some source you check that does not agree with what
is reported in this appendix. However, if you find any
serious error in the data tables, please send an e mail to
Robert Shine (bshine@ramapo.edu) informing him of the nature of
the error. A blank cell in the data tables indicate the value
for that data point was not found or does not apply. For
example, the boiling point of an inorganic solid is generally
not useful in this course.
Also, the density for a solid
substance could be meaningless for us.
Remember that you can obtain more complete information
from the Material Safety Data Sheet (msds) for a given
substance. You can obtain a material safety data sheet for a
compound by doing a Google search. For example, an msds for
acetone can be obtained by typing "msds acetone" or "acetone
msds" (don't include the quote marks) in the search box.
Usually this search will result in many hits so you can just
link to one of the sources. If you do not get any results
from a Google search, try to search using an alternate name
for the chemical. With practice, you will find your favorite
source for an msds.
EXPLANATION OF COLUMN HEADINGS:
The ChemicalName column gives the usual name for the
substance listed. All tables are arranged alphabetically
according to the ChemicalName. The AlternateName column gives
another name for a given substance. The Formula column gives
the general chemical formula for that substance. If you
compute the molecular weight for that chemical formula, you
would obtain the data value in the Mol.Wt. column. The units
for these data are grams/mole. The next column, Melt.Pt.,
gives the melting point in degrees Centigrade (Celsius).
After this, comes the column Boil.Pt. which gives the boiling
point for the substance in degrees Centigrade (Celsius). The
Density column gives the density for the substance in
grams/ml. The density for a solid is not useful in this
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course but the density for a liquid could give a useful method
when measuring liquid reagents. One can easily and accurately
measure a liquid by volume. The weight of that sample would
then be equal to the volume times the density. If you know
the weight you wish to measure, the volume would be equal to
the desired weight divided by the density. Expressed in
mathematical terms, the formulae are:
Volume = Weight / Density
Weight = Volume x Density
The remaining three columns give some safety data. The
CAS # is the Chemical Abstract Service number for the
substance. The CAS# is a unique identifier for a chemical
assigned by the American Chemical Society. The FlashPt gives
the temperature at which a substance can be readily ignited.
Substances with a flash point below 100 degrees Centigrade are
classified as flammable while those above 100 degrees are
called combustible. Be especially cautious with chemicals
with a flash point below 0 degrees. The final column gives
the NFPA Rating of the substance. NFPA stands for the
National Fire Protection Association and their web site can be
found at www.nfpa.org (not .com) which is a for profit
organization. You can find more information about this rating
system by doing a Google search on “nfpa diamond”. The rating
assigns values from 0 to 4 for each of three categories
(Health, Fire, and instability which is sometimes called
reactive). A value of 0 indicates little or no hazard whereas
a value of 4 indicates an extreme hazard.
© 2009 Stephen Anderson and Robert Shine
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