The Organic Chemistry II Laboratory Course
Chem 302, Spring 2008
The experiments in this course have been chosen to illustrate important organic
reactions that we will discuss in lecture, and to utilize many of the skills that you
learned last semester.
Laboratory Safety
During this class we will work with a range of solvents, organic materials, acids,
and bases which could be harmful if you were to splash them into your eye or
onto your skin. For your own protection, we must require you to wear laboratory
goggles, which you can purchase at the campus store in Blanchard. Yes, they
are ugly and uncomfortable, but they are also very, very important. We also
strongly advise you to wear disposable gloves whenever appropriate. You may
not wear open sandals, but instead must wear closed–toe shoes. Furthermore,
we advise you wear clothing that fully covers your torso. (Shirts with spaghetti
straps and bare midriffs might be cute, but burns on your chest and stomach are
not).
On a related note, these lab handouts clearly indicate which chemicals should be
used in the hood because of possible inhalation hazards. Please read the labs
carefully and follow the directions. You will find that it is easiest, and safest, to do
everything you can in the hood.
Finally, it is extremely important to us that we be considerate of the environment.
Solvents may not under any circumstances be poured down the sink. Use the
appropriate designated waste bottles for halogenated and non–halogenated
waste. If you are unsure of which waste container to use, please ask your
instructor or TA.
Laboratory Reports
In place of lengthy laboratory reports, you will be asked to keep a careful record
of your work in a laboratory notebook containing carbon–copy duplicate pages
(just as you did last semester). The original copy will remain in your notebook for
reference later in the course, and the carbon copy will be turned in for grading.
Try to make your lab notebook as legible and organized as possible, but don’t
worry if you need to cross something out or change it—real science is
sometimes, in fact often, messy! Prepare your reports in ink.
Your laboratory notebook should include each of the following sections for each
and every lab, under clear sub–headings. The first three sections should be
completed before you come to lab and turned in to your instructor or teaching
assistant as a pre-lab before you begin the experiment. You should work on the
last three sections as you go along and finish as much of your write up as
possible before you leave.
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Organic Chemistry II Laboratory Experiments, Spring 2008(Nunez, Browne, Hamilton, edited 1/08 JW)
1
(a)
Title, date, and experiment number. Also, the name of your laboratory
partner.
(b)
Reaction: A brief summary of the reaction(s) involved, including the
structures of the reactants, intermediates, and products, and a reasonable
mechanism. (You will probably have to look at your textbook and your
lecture notes and perhaps do some more extensive research).
(c)
Methods: A brief summary of the process followed in the experiment. It
may be in the form of a flow–chart, or schematic outline, or some similarly
visually–accessible form. You need not repeat all of the details in the lab
handout, but you should make it clear that you have thought about the
experiment ahead of time. The intention here is that these additional
instructions will serve as an additional guide to your experimental work.
(d)
Observations and Data: A careful record of what happened as you did the
experiment. Did a precipitate form, liquids separate into layers, or the color
of a solution change? At what stage of the experiment did this happen?
Include measurements, descriptions of the solutions, sketches of
chromatographic separations, and all spectra. You should also indicate
any changes you made in the experimental procedure.
(e)
Discussion and Conclusions: Comment here on the overall level of
success and the ease (or otherwise) with which success was achieved.
Include any conclusions you have drawn about the purity or yield of your
material, and discuss what your results may have revealed about the
mechanism of the reaction.
(f)
Answers to Questions: The questions are usually related to practical
elements of the experiments. You should look at the questions during the
laboratory sessions and talk over your ideas with your instructor or TA.
Reports are due at the beginning of the following week’s lab. If you submit a lab
report more than a week after the completion of an experiment, the grade you
receive will be halved for each week of delay. And, you must complete and
submit all 12 lab reports in order to pass this course.
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Organic Chemistry II Laboratory Experiments, Spring 2008(Nunez, Browne, Hamilton, edited 1/08 JW)
2
Chemistry 302: Organic Chemistry II, Spring 2008
Week of
Jan 30th
Feb 4th
Feb 11th
Feb 18th
Feb 25th
Mar 3rd
Mar 10th
Chapter and Topic
Lab Experiment
Intro to class, review Diels-Alder
Ch 13 NMR
Ch 13 NMR
Ch 13 NMR
Ch 13 NMR
Ch 13, 14 NMR, Benzene
Ch 14 Reactions of Benzene
Ch 14, 15 Reactions of substituted benzenes
Ch 15 Reactions of substituted benzenes
Ch 15 Reactions of substituted benzenes
Ch 16 Reactions of carbonyls
Ch 16 Reactions of carbonyls
Ch 16 Reactions of carbonyls
review
Exam #1
Ch 17 Aldehydes and Ketones
Ch 17 Aldehydes and Ketones
Ch 17 Aldehydes and Ketones
Ch 17 Aldehydes and Ketones
Ch 18 Intro to reactions at the alpha carbon
None
#1 The Diels–Alder Reaction
(set up Spectroscopy)
#2 Spectroscopy
#3 Tropylium
#4 Nitration of Methyl Benzoate
#5 Banana Oil
#6 Preparation of a Naphthalene
Diimide
☺ Spring Break!!! Yeah!!! ☺
Mar 24th
Mar 31st
Apr 7th
Apr 14th
Apr 21st
Apr 28th
May 5th
Ch 18 Substitutions at the alpha carbon
Ch 18 Reactions at the alpha and beta carbons
Ch 18 Condensations
Review of carbonyls
reveiw
Exam #2
Ch 19 Reductions
Ch 19 Oxidations
Ch 19 Oxidations
Ch 20 Amines
Ch 20 Amines
Selected topics in biological chemistry
Selected topics in biological chemistry
Selected topics in biological chemistry
Selected topics in biological chemistry
Exam #3
Ch 29 Pericyclic Reactions
Ch 29 Pericyclic Reactions
Last day of classes
#7 The Wittig Reaction
#8 Chemiluminesence
#9 The Aldol Reaction
#10 The Grignard Reaction
(set up Ethanol lab)
#11 Biosynthesis and Distillation
of Ethanol
#12 Hydrolysis of Esters
Check-out
-----------------------------------------
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Organic Chemistry II Laboratory Experiments, Spring 2008(Nunez, Browne, Hamilton, edited 1/08 JW)
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#1: Diels–Alder Reaction of Anthracene with Maleic Anhydride
Purpose of the experiment
You will prepare the Diels–Alder adduct, sometimes called the cycloaddition
product, formed from the reaction of anthracene with maleic anhydride.
Introduction
As you will remember from last semester, the Diels-Alder reaction couples two
molecules, a conjugated diene and an alkene nicknamed a “dienophile”, via a
concerted [4+2] cycloaddtion reaction. In this reaction, three pi bonds (two on
the diene and one on the dienophile) are transformed into two new sigma bonds
and a new pi bond. The result is a new 6-membered ring containing a double
bond. In the molecular orbital description of this reaction, the HOMO of one
molecule reacts with the LUMO of the other in-phase to create the new bonds
through what is called a pericyclic reaction.
Today you will couple an electrophilic alkene, maleic anhydride, with the aromatic
compound anthracene (which will act as the diene) to create a new bridged
compound:
O
O
O
O
O
O
What you will do
This experiment provides a straightforward introduction to organic synthesis. You
will prepare and isolate a Diels–Alder addition product by simply heating a
mixture of the two components—other than a reaction solvent, no other materials
are required.
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Organic Chemistry II Laboratory Experiments, Spring 2008(Nunez, Browne, Hamilton, edited 1/08 JW)
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Physical constants for the materials you will use or prepare
Material
Mol. Wt. (g/mol)
Mp (°C)
Anthracene
178
216–218
Maleic anhydride
98
54-56
m–Xylene
106
–48
Product
276
261–262
The experiment, step–by–step
Care! None of the reagents you will use in this experiment require special
precautions but you should get in the habit of handling all chemicals with respect.
Wear your laboratory glasses AT ALL TIMES. You will only use small amounts of
liquids and solids and these can be weighed and dispensed with little fuss.
Disposable gloves are available in the laboratory to use if you wish. Perform this
entire experiment in the hood.
1. Weigh out 1.0g of anthracene and 0.5g of maleic anhydride. Transfer both to
a clean, dry, 25mL round bottomed flask and add 10mL of m–xylene and a
boiling chip. Attach a reflux condenser—it is not necessary to connect water
hoses to the condenser for this experiment.
2. Arrange the apparatus on a sand bath for heating. For even heating it is best
if the flask is bedded down into the sand to around the same level as the
liquid contents. Turn up the controller to around 50% power and watch
(patiently) for bubbles to begin to form around the boiling chip—shortly after
the first bubbles form gentle refluxing (boiling) should commence. If boiling
has not commenced after 10 minutes then turn up the power slowly, and
incrementally, until it does so. Once you have reached this point let the
reaction boil gently for 30 minutes.
3. Turn off the power and carefully raise the reaction from the sand. Let the
mixture cool for at least 10 minutes and watch carefully—the reaction product
crystallizes as the mixture cools (avoid the temptation to cool the flask too
quickly, the best crystals are always obtained by slow cooling). Place the
cooled flask in a beaker of ice–water to complete the cooling and
crystallization process. Place 5mL of m-xylene in a small Erlenmeyer flask
and cool this on ice also.
4. Arrange the small Hirsch funnel and flask for suction filtration; you will need
to clamp the neck of the flask to keep it stable. The vacuum supply (yellow
taps on the hood) is quite strong and you will not need to have the tap fully
open (your Instructor or TA will show you how to set the vacuum at an
appropriate level). Inspect your reaction—if the crystals appear clumped
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Organic Chemistry II Laboratory Experiments, Spring 2008(Nunez, Browne, Hamilton, edited 1/08 JW)
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together, or adhere to the sides of the flask, gently poke around in the flask
with a clean spatula to free them. Swirl the flask and pour the contents into
the Hirsch funnel—try to be bold as otherwise you will simply drain the
solvent away and leave the crystals in the flask! Use a little of the extra cold
m–xylene to wash any remaining crystals out of the reaction flask, and use
the remaining cold m–xylene to wash the collected crystals on the filter
funnel. Let the crystals dry with the passage of air for a few minutes.
Disconnect the vacuum hose carefully and then shut off the vacuum. Spread
the crystals on a filter paper to complete the drying process.
5. Weigh the dried crystals and determine their melting point.
Cleaning up
Once you have recorded the yield and melting point of the final product you can
discard the material in the solid organic waste container. The filtrate from
collection of the crystals should be discarded in the non–chlorinated waste
solvent container. Rinse all of the glassware you have used with water and then
acetone, discarding the acetone washings in the non–chlorinated waste solvent
container.
Questions
1. The m–xylene reaction solvent, though it is not part of the chemical reaction,
plays two crucial roles in this experiment. Identify and explain both.
2. What are the structures of o–xylene, m–xylene and p–xylene? Why would p–
xylene have been a poor choice as a solvent for the procedure you followed
in this experiment? [Hint: what is the melting point of p–xylene?].
3. Why, in this particular case, is it not necessary to have cold water flowing
through the condenser in order for the solvent to condense and return to the
flask?
4. What is the purpose of washing the final product with cold xylene? [Hint:
think about the differing solubilities of the reactants and product in both hot
and cold xylene].
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Organic Chemistry II Laboratory Experiments, Spring 2008(Nunez, Browne, Hamilton, edited 1/08 JW)
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#2: Identification of an Unknown using Spectroscopy and
Spectrometry
Purpose of the experiment
You will use what you have learned about infrared spectroscopy, mass
spectrometry, and nuclear magnetic resonance spectroscopy to determine the
identity of an unknown compound.
Introduction
As you have seen in lecture, infrared spectroscopy (IR), mass spectrometry
(MS), and nuclear magnetic resonance (NMR) spectroscopy each provide
valuble information about the structure of a molecule. When considered together
they can often be used to determine the structure of a molecule about which little
or nothing else is known. Infrared spectroscopy provides information about the
types of functional groups present (or absent) in a molecule, though it tells little
about the way in which these groups are connected. Mass spectrometry
provides information about the mass of a molecule as well as the fragments it
forms when ionized. These fragments provide vital clues about the shape of the
molecule. Finally, NMR tells you about the chemical environment of all of the
hydrogen (1H NMR) or carbon (13C NMR) atoms in a molecule, allowing you to
connect the functional groups and fragments postulated from MS and IR properly
into a whole molecule.
What you will do
Each person will select an unknown compound of her own and obtain IR, 13C
NMR, 1H NMR, and mass spectra of it, using these to determine its structure.
You may work in groups of any size you like, but you must write up all the
unknowns for your group.
Advice regarding experimental or instrumental aspects
NMR: place one to two drops of unknown into an NMR tube and then add
solvent (CDCl3) up to the mark that is on the poster in the lab.
IR: Place one or two drops onto the sodium chloride cell then place the other cell
on top of the first, then run spectrum. Clean cells off with acetone.
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Organic Chemistry II Laboratory Experiments, Spring 2008(Nunez, Browne, Hamilton, edited 1/08 JW)
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GC/MS: Place one to two drops into 10 mL of pentane. Mix well before running
the spectrum. Please do not forget to clean the syringe when you are done with
the spectra.
Cleaning up
All NMR solvents must go in halogenated waste because of the CDCl3. Acetone
rinsings from the IR should go in the acetone waste bottle. All GC/MS samples
will go in non halogenated waste, UNLESS your GC/MS shows that you have a
halogenated compound.
Questions
When proposing the structure of your unknown(s), include copies of all spectra
and discuss the following:
IR: What types of functional groups are present, and what band(s) at what
wavelength(s) leads you to that conclusion? What types of functional groups are
absent, and what band(s) at what wavelength(s) lead you to that conclusion?
MS: What is the mass of your whole molecule, and how does that match the
chemical formula of your molecule? How many degrees of unsaturation does
that correspond to? What fragments are formed, and how does that correpsond
to the molecule you have proposed?
1
H NMR: How many non-equivalent protons are present on your molecule, and
how many peaks are present on your spectrum? Identify each peak on the
spectrum and explain how its chemical shift, integration, and splitting are
consistent with your proposed structure.
13
C NMR: How many non-equivalent carbons are present in your molecule, and
how many peaks are present on your spectrum? Are any carbons missing a
peak, and if so, which and why? Identify each peak on the spectrum and explain
how its chemical shift is consistent with the 1H NMR spectrum and your proposed
structure.
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Organic Chemistry II Laboratory Experiments, Spring 2008(Nunez, Browne, Hamilton, edited 1/08 JW)
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#3:Preparation of Triphenylmethyl Fluoroborate and Tropylium
Iodide
Purpose of the experiment
To achieve a practical demonstration of the validity of Hückel’s 4n+2, π–electron
rule by preparing a non–benzenoid aromatic system.
Introduction
In this experiment you will prepare two relatively stable organic carbocations.
First you will prepare triphenylmethyl (a.k.a. “trityl”) carbocation from
triphenylmethanol. The acid protonates the –OH, making it a better leaving
group. When the water departs and is “soaked up” by acetic anhydride, what is
left behind is a stable tertiary carbon that is stabilized by resonance.
In the second step, the trityl carbocation will abstract a hydride from
cyclohepatriene, generating a tropylium carbocation (and triphenylmethane).
The tropylium carbocation is even more stable than the trityl carbocation and
thus can give up a hydride because it is aromatic.
The mechanisms for the reactions you will carry out are shown below:
What you will do
You will prepare sequentially two relatively stable organic carbocations. After
preparing the second carbocation you will demonstrate its ionic character by
confirming the presence of iodide with aqueous silver nitrate.
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Organic Chemistry II Laboratory Experiments, Spring 2008(Nunez, Browne, Hamilton, edited 1/08 JW)
9
Physical constants for the materials you will use or prepare
Material
Mol. Wt. (g/mol)
Mp (°C)
Triphenylmethanol
260
160–163
Tetrafluoroboric acid:dimethyl ether complex
88
Cycloheptatriene
92
Tropylium Iodide
218
The experiment, step–by–step
Care! Both acetic anhydride and tetrafluoroboric acid are corrosive and will case
painful burns if allowed to come into contact with your skin. Wear disposable
gloves when dispensing these materials.
Work in pairs for this experiment. Perform the entire experiment in the hood.
Part I. Preparation of the triphenylmethyl and tropylium carbocations.
1. Place a magnetic stirrer bar into a clean and dry 10ml round bottomed flask.
It is essential that the flask is absolutely clean and dry, the reaction will
otherwise likely fail. If you have any doubts scrub the flask in detergent and
water, rinse with water and then acetone and place in the drying oven for at
least 15 minutes.
2. Using the plastic graduated dropper provided transfer 3.5mL of acetic
anhydride to the flask and cool it on an ice bath for at least 10 minutes.
3. Using the syringe provided with the reagent transfer 0.25mL of
tetrafluoroboric acid:dimethyl ether complex to the flask and stir (gently) with
continued cooling.
4. Weigh out 0.39g of triphenylmethanol. Add the triphenylmethanol in one
batch to the stirred solution in the flask. Try to avoid having any of this solid
material stick to the neck of the flask and ensure that the solid and liquid mix
completely by adjusting the stirring rate and agitating the entire flask if
necessary. Remove the ice bath and stir for an additional 5 minutes.
5. Arrange a sand bath on top of the magnetic stirrer and bed your flask a little
way into the sand. Recommence stirring and heat the reaction very gently
(set the rotary temperature controller to around 30% for ten minutes and then
at 45% for five minutes) until you obtain a homogeneous red solution. It is
better to heat for a shorter time than a longer one so consult your Instructor
or TA for advice on when to stop. If you remove your solution from the heat,
it should stay the homogenous red color and not get cloudy. If it does, put it
back on the heat for a few more minutes.
6. Remove the heat source and allow the mixture to cool for 5 minutes. Using
the syringe provided add 0.2mL of cycloheptatriene to the stirred mixture. A
precipitate should soon begin to form. Continue to stir the suspension for a
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Organic Chemistry II Laboratory Experiments, Spring 2008(Nunez, Browne, Hamilton, edited 1/08 JW)
10
further 10 minutes. The solution has a pungent odor; it would be best to keep
your hood closed as much as possible.
7. Cool the flask in a beaker of ice and continue to stir for another 5 minutes.
8. Add 3.0mL of anhydrous t–butyl methyl ether to the flask in 1mL increments,
and after each addition, stopper, shake gently and then immediately release
any pressure build up by gently removing the stopper. Then let the mixture
stir with continued cooling for a further 5 minutes.
Part II. Isolation of tropylium tetrafluoroborate. Conversion to, and isolation
of, tropylium iodide–confirmation of the presence of iodide.
9. Filter the cold suspension at the pump using the small Hirsch funnel and
flask—ensure that both funnel and flask are clean and dry. Wash the
precipitate with two 1mL portions of t–butyl methyl ether (need not be
anhydrous), leave the material on the filter for around 2 minutes to allow it to
dry.
10. You will save this filtrate and transfer it to a 10mL round bottom flask. Use
1.0mL of t-butyl methyl ether to wash out the Hirsch flask. This solution
should be rotary evaporated and then TLC plated with commercially-made
triphenylmethane in hexane. Prepare the triphenylmethane in 2mL of
acetone in a sample vial.
11. Transfer the dry filtered material to a small, clean, test tube. Add 0.5mL of
water and heat the tube gently on the sand bath to dissolve the crystals. DO
NOT allow the water to boil—if the crystals do not all dissolve within a couple
of minutes then add additional tiny (0.1mL) volumes of water until they do.
Occasionally flick the tube to mix the contents.
12. Add 0.5mL of saturated sodium iodide solution to the tube. Flick the tube a
couple of times to ensure that the two solutions mix thoroughly. Cool the tube
in a beaker of ice, dark red crystals of tropylium iodide should form within
minutes. Allow the tube to sit on ice for 10 minutes, put 1.0mL of methanol in
a second tube and cool this on ice also. Clean up the Hirsch assembly you
used previously while cooling proceeds.
13. Swirl the contents of the tube and try to ensure that all the crystals are
moving relatively freely (a boiling stick is useful for dislodging affixed
crystals). Prepare the Hirsch assembly for filtration. With one swift motion try
to pour the whole suspension into the Hirsch funnel—the more bold you can
be the less material will get stuck on the sides of the tube! If necessary use
small batches of the ice cold methanol to transfer the final crystals to the
Hirsch funnel. Allow the crystals to dry with the passage of air for a few
minutes. Weigh the crystals. Discard the filtrate in the non–chlorinated waste
solvent container.
14. To confirm the presence of iodide ion in the final product transfer a few
crystals (around a spatula tip’s worth) to a small test tube, dissolve them in a
few drops of water (heating if necessary), and then add one drop of 2%
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Organic Chemistry II Laboratory Experiments, Spring 2008(Nunez, Browne, Hamilton, edited 1/08 JW)
aqueous silver nitrate solution. A cloudy precipitate of silver iodide should
form.
Clearing up
Once you have recorded a yield of the final product you can discard the material
in the solid organic waste container. Ensure that any remaining liquids from
earlier filtrations are transferred to the non–chlorinated waste solvent container.
Rinse all of the glassware you have used with water and then acetone,
discarding the acetone washings in the non–chlorinated waste solvent container.
Questions
1. What is the reasoning behind TLC plating the filtrate from the first filtration,
after the addition of anhydrous t-butyl methyl ether?
2. Draw resonance structures to show how the positive charge is delocalized
in the triphenylmethyl and tropylium carbocations (one set of structures for
one of the three phenyl rings will suffice for the triphenylmethyl cation).
13
3. C NMR Spectra of cycloheptatriene and tropylium iodide are available in
the laboratory. The relevant peaks are numbered on each spectrum and
you can use the table of peak positions to record accurate chemical shift
values. Explain the appearance of these spectra.
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Organic Chemistry II Laboratory Experiments, Spring 2008(Nunez, Browne, Hamilton, edited 1/08 JW)
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#4: Nitration of Methyl Benzoate
Purpose of the experiment
In this experiment you will perform a nitration of an aromatic ring, an important
and common transformation in organic chemistry. A substituent on the aromatic
ring will ensure that nitration is directed to only one position on the ring–ortho,
meta or para with regard to this group. The outcome will be proven with a melting
point determination.
What you will do
You will nitrate methyl benzoate using a mixture of nitric and sulfuric acids,
isolate and purify the product, and determine the orientation of substitution by
measuring the melting point of the purified material.
O
O
O
H2SO4
O2N
O
?
+HNO3
Physical constants for the materials you will use or prepare
Material
Mol. Wt. (g/mol)
Mp (°C)
Methyl benzoate
136 (d 1.094)
Concentrated H2SO4
98 (d 1.84)
Concentrated HNO3
63 (d 1.42)
Methyl o–nitrobenzoate
181
-13
Methyl m–nitrobenzoate
181
78–80
Methyl p–nitrobenzoate
181
94–96
The experiment, step–by–step
CAREFUL! Both sulfuric and nitric acids are extremely corrosive and will rapidly
burn holes in clothing and gloves if spilt or splashed. There is little danger of this
happening if they are measured using the pipettes provided with the reagent
bottles. Wait for your turn to use the communal reagents and then bring the
required reagent to your hood for use. DO NOT carry your reaction around the
laboratory.
Perform this entire experiment in the hood.
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1. Place a magnetic stirrer bar in a 5mL round bottomed flask. Measure 2.0mL
of concentrated sulfuric acid into the flask and stand it in a small beaker of
ice on top of the stirrer unit. Test the stirrer to check that the bar is able to
spin slowly—DO NOT stir the mixture at high speed, this is unnecessary and
there is a risk of splashing.
2. Measure 1.0mL of concentrated nitric acid and 1.0mL of concentrated
sulfuric acid into a small test tube and cool this mixture in a separate beaker
of ice for at least 5 minutes.
3. Add 1.0mL of methyl benzoate to the round bottomed flask while stirring
slowly and continuing to cool—add more ice if necessary. Using a glass
pipette add the cooled acid mixture in the test tube to the stirred solution in
the round bottomed flask dropwise over a period of a couple of minutes.
Once addition is complete let the reaction warm to room temperature by
removing the ice, but keep stirring for at least the next 20 minutes.
4. Prepare about 10mL of an ice–water slush in a small beaker. Carefully (use
gloves and beware of spitting as the acids hit the water) pour the reaction
mixture into the slush, this process will precipitate your product as a solid.
Carefully rinse the round bottomed flask with around 2mL of water and empty
into the same beaker.
5. While the ice in your slush is melting arrange the Hirsch funnel and flask for
suction filtration. Measure around 5mL of methanol into a test tube and place
to cool in a beaker of ice for at least 10 minutes.
6. Once the ice in your slush has FULLY melted filter the solution on the Hirsch
funnel to collect the crude product. Wash the solid with two 2mL portions of
ice–cold methanol (from the test tube you have been cooling—it is crucial
that this methanol is ICE COLD).
7. The product left on the filter will almost certainly be sticky and needs to be
recrystallized. You first need to transfer the sticky solid to a 10mL conical
flask—this is best achieved with a clean spatula and a lot of patience. Don’t
worry about using the neck of the conical flask to wipe off the solid from the
spatula, just get all the material you can into the flask. Your filtrate should be
disposed of in the aqueous acid waste container.
8. Add around 4mL of methanol to the conical flask and heat the suspension of
material gently on the hotplate until it begins to boil very gently (set the
control to no more than around 30% to begin, turn up a little more if
necessary). Your material may all dissolve at this point, though if your
reaction has gone really well and you have a lot of product you may need to
add extra methanol (in batches of 0.5mL) to get all the solid to dissolve. You
may also need to poke material on the sides of the flask down into the
solvent with a spatula—the aim, as with all crystallizations, is to end up with a
hot, concentrated, clear solution.
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9. Remove the flask from the heat, crystallization will soon commence.
Complete the crystallization process by cooling the flask in a beaker of ice for
10 minutes. Meanwhile, prepare the Hirsch apparatus for a second suction
filtration.
10. Ensure that the bulk of the crystals are loose in solution, dislodge and break
up with a spatula tip if necessary. Collect the crystals by suction filtration on
the Hirsch funnel, returning some of the filtrate to the flask if necessary to
wash out the final few crystals. Dry the product by drawing air through the
filter for 10 minutes.
11. Weigh the crystals and determine their melting point.
Cleaning up
The filtrate from the first filtration should be discarded in the aqueous acid waste
container. Once you have recorded the yield and melting point of your final
product you can discard the material in the solid organic waste container. The
filtrate from collection of the final product should be discarded in the non–
chlorinated waste solvent container. Rinse all of the glassware you have used
with water and then acetone, discarding the acetone washings in the non–
chlorinated waste solvent container.
Questions
1. One of the three possible products from this reaction could be ruled out
before sitting down at the melting point apparatus. Which one, and why?
2. Would you be confident distinguishing the three possible products by (i) IR
spectroscopy, (ii) mass spectrometry, (iii) 1H NMR spectroscopy and (iv) 13C
NMR spectroscopy? Fully justify your claim in each case.
3. Is the original ester substituent on methyl benzoate activating or deactivating?
What effect would you expect this to have upon the yield of your reaction?
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15
#5: Preparation
Esterification
of
Synthetic
Banana
Oil
by
a
Fisher
Purpose
To prepare a fragrant ester from an alcohol and a carboxylic acid and purify it by
distillation.
Introduction
The flavors and scents used in foods, beverages, and perfumes for millennia
were purified from fruits, flowers, and other natural products. Now it is just as
common for these pleasantly-scented compounds and analogs to be chemically
synthesized. Artificial flavors might resemble vanilla, strawberries, pineapples,
oranges, and other complex plant products. Interestingly, the characteristic
strong flavors or odors are commonly due to organic compounds called esters.
Esters can be prepared by Fisher esterification, a nucleophilic substitution
reaction in which a carboxylic acid is heated with an alcohol to make the ester:
The acid in this case functions as a catalyst.
In this experiment, you will prepare synthetic banana oil, an ester also known as
isopentyl acetate or isoamyl acetate, by heating isopentyl alcohol (3-methyl-1butanol) with acetic and sulfuric acid:
Because several of the components are volatile, you would lose reactants and
products if you were to just heat the reaction in an open flask. As a result, we will
reflux the mixture and condense the reactants and products as they try to
escape. After refluxing for an hour and allowing the mixture to cool, our desired
banana oil product must be separated from the leftover acids. This separation
will involve extraction of the alcohol with water and then with base using a
separatory funnel.
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Once the acids have been removed and the product has been dried, the product
must be separated from other side products and starting material using
distillation. Distillation is a method for separating two or more liquids based on
differences in their boiling points. All pure, stable liquids have characteristic
boiling points at atmospheric pressure, reflecting the intermolecular interactions
holding molecules together in the liquid. As a general rule, polar compounds
have higher boiling points than nonpolar compounds, and larger molecules have
higher boiling points than smaller ones with similar polarities. Since different
compounds in a mixture will boil at different temperatures, this behavior can be
used to separate them. When a mixture is heated to the point where the most
volatile (the lowest boiling) component starts to boil, the vapor can be separated
using an appropriate arrangement of glassware, and subsequently condensed
back to liquid form and collected in a separate flask.
It should be noted that this method of distillation is “simple”, meaning that at the
end of the distillation there will be a little of the banana oil left in the distillation
flask, and perhaps a little of the starting material in the receiving flask. However,
you will likely achieve both good separation and a good recovery of the banana
oil if you carefully monitor the temperature during your distillation. To really purify
a liquid by this means it is often necessary to distil several times in succession,
which we will do later this semester.
To complete the experiment you will assess the purity of your isolated oil by
recording a gas chromatogram.
What you will do
Important information about the substrates, catalyst, and product are included
below:
Acetic acid
Isopentyl alchohol
Isopentyl acetate
Sulfuric acid
Density
1.049
0.809
0.876
1.84
Boiling point, ˚C
118
130
142
290
Carefully measure 16.3 mL of of isopentyl alcohol with a graduated cylinder and
pour into a 100 mL round bottom flask. In the hood, slowly add 17 mL of glacial
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17
acetic acid and 1.0 mL of concentrated sulfuric acid, with gentle swirling. Be very
careful with the sulfuric acid and use the pipette provided to make the addition
(there are markings on the pipette to allow you to add 1.0mL without danger of
any spills). Then, add boiling chips.
Connect a condenser to the round-bottom flask, ensure that the cold water hoses
are securely attached and that the outflow hose is safely in the sink, then (gently)
turn on the cold water. Heat the mixture slowly using a sand bath until it begins
to boil and allow it to continue refluxing (i.e. boiling gently) for about an hour. At
the end of an hour turn off the sand bath, allow the reaction to cool completely to
room temperature, then turn off the condenser.
Transfer the reaction mixture to a 250mL separatory funnel, leaving the boiling
chips behind in the round bottom flask. Add 100 mL of distilled water, thoroughly
mix the layers (remembering to release the pressure frequently), and isolate the
organic layer. Now, extract twice with 50 mL portions of 5% aqueous sodium
carbonate. (Use the table of densities to ensure you know which layer is which
and check with your instructor if at all unsure). Set aside the aqueous layers from
each of your extractions for safe disposal at the end of the experiment.
Dry the isopentyl acetate with sodium sulfate and gravity filter using a narrowstem filter to remove the drying agent. Next, purify your product by distillation:
Don’t forget to clamp the round bottom flasks AND condenser onto ring
stands, and to use Keck clips on both ends of the condenser to hold the
glassware together! Also, make sure the cold water is running through the
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condenser. Position the thermometer as shown, turn on the sand bath, and
monitor the temperature of the vapor. Once again we are looking to bring the
liquid to a gentle boil so be sure to increase the temperature slowly and carefully
until you reach this point. Collect any liquid that distills below 136˚C and discard
it. Then using a fresh flask or round-bottom, collect the liquid that distills
between 136 and 143˚C: this is your product. Weigh it.
Identify the number of products and their distribution by gas chromatography and
mass spectrometry.
Cleaning Up
Acid fractions after extraction should be carefully neutralized with solid sodium
bicarbonate before washing down the sink with lots of water.
Your Report
The important issues in this lab are yield and purity.
Using the table above, calculate the moles of each starting material. Which was
the limiting reagent, and which was in excess? What reason might you have for
adding more of one reagent than the other? What is your maximum theoretical
yield, and therefore your percentage yield of purified banana oil?
Based on the gas chromatography, how many compounds were present at the
end, and how much of each (area under the peaks)? If you are told that the
stationary phase is moderately polar, which compounds would you guess are
which? What techniques might you use to further purify your product?
Question
1. Which would be separated best by a simple distillation, two liquids with
boiling points 10˚C apart, or two liquids with boiling points 50˚ apart?
Why?
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#6: Preparation of a Naphthalene Diimide
Purpose of the experiment
To prepare a naphthalene diimide derivative from the condensation of a
naphthalene dianhydride with two molar equivalents of a primary amine. The
resulting diimide is a strong electron acceptor, i.e. it is readily reduced, and
molecules of this type are currently being used in emerging areas of chemical
technology, including molecular scale electronic components and the selective
binding and modification of DNA.
What you will do
Heating of the naphthalene dianhydride with the primary amine in a high boiling
solvent (dimethylformamide, DMF) leads to elimination of water and imide
formation:
O
O
O
O + 2 NH2CH2(CH2)4CH3
O
CH3(CH2)4H2C
O
O
N
N CH2(CH2)4CH3
O
O
O
The product is only soluble in hot DMF so once the reaction is allowed to cool the
desired product crystallizes and may be collected by filtration.
Physical constants for the materials you will use or prepare
Material
Mol. Wt. (g/mol)
Mp (°C)
Naphthalenetetracarboxylic dianhydride
268
>300
n–Hexylamine
101 (d 0.766)
bp 131–132
DMF
–
bp 153
Naphthalene diimide
434
200–202
The experiment, step–by–step
Care! The materials used in this experiment should not be allowed to come into
contact with your skin or respiratory system. Weigh and measure all materials in
the hood.
Work in pairs for this experiment. Perform the entire experiment in the hood.
1. Place 0.25g of naphthalenetetracarboxylic dianhydride in a small round
bottomed flask. Add 4mL of DMF using the pipette provided, and a small
stirrer bar. Clamp the flask in a sand bath placed on top of the stirrer unit.
Start the stirrer unit and ensure that the reaction mixture is stirring gently.
2. Using the pipette provided with the reagent transfer 0.25mL of n–
hexylamine to the flask. At this point your reaction will become thick with
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20
precipitate. Attach an air condenser and turn on the power to the sand
bath (about 50% power to begin).
3. Watch for the first signs of condensation to form on the inside surface of
the condenser. This is the point at which you should try to maintain the
reaction, you do not need to have this reaction refluxing vigorously. You
will find that all the precipitate in the reaction flask will dissolve, giving a
clear red solution.
4. After 1 hour of heating and stirring turn off the power to both the heater
and stirrer units and allow the reaction to cool to room temperature—as
usual, allowing this process to happen slowly will lead to the highest
quality crystals. Complete precipitation of the product by cooling the
reaction flask in a beaker of ice–water. Cool 2mL of DMF in a test tube for
use later.
5. Collect the product on the Hirsch funnel, using extra small portions of
water to rinse remaining material from the flask. Wash the crystals with the
cold DMF from the test tube, and then with 2mL of cold water. Dry the
crystals by continuing to draw air through them for at least a further 10
minutes.
6. Measure your isolated yield and record a melting point.
Clearing up
Once you have recorded a yield of the final product you can discard the material
in the solid organic waste container. Ensure that any remaining liquids from
earlier filtrations are transferred to the non–chlorinated waste solvent container.
Rinse all of the glassware you have used with water and then acetone,
discarding the acetone washings in the non–chlorinated waste solvent container.
Questions
1. The π–system of the product you have made is often described as being
electron deficient. What is the origin of this electron deficiency? (Hint: think
about the electronic nature of those carbons directly attached to the π–
system).
2. Derivatives of the molecule you have prepared bind to DNA (in other words,
they form stable complexes with DNA). Apply what you know about the
structure of double stranded DNA to propose a model for the ability of these
systems to form stable complexes.
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#7: Wittig Reaction: Synthesis of trans 9–(2–Phenylethenyl)
anthracene
Purpose of the experiment
You will synthesize an attractive, yellow colored, crystalline material using a
Wittig reaction. This process involves the formation of a new carbon–carbon
double bond and exemplifies a common and powerful reaction in organic
chemistry. The product you prepare, in addition to absorbing visible light (hence
its color), is potentially fluorescent. This property is exploited in next week’s
experiment.
What you will do
You will prepare trans 9–(2–phenylethenyl)anthracene, purify your material by
crystallization, and confirm its identity with a melting point determination.
(C6H5)3PHCH2
50% NaOH
(C6H5)3P CH
(C6H5)3P CH
The Wittig reagent,
an ylide
Cl
O
Benzyltriphenylphosphonium
chloride
H
9-Anthraldehyde
(C6H5)3P
O
Triphenyphosphine
oxide (TPPO)
H C
C H
(C6H5)3P CH
O CH
trans-9-(2-Phenylethenyl)anthracene
Physical constants for the materials you will use or prepare
Material
Mol. Wt. (g/mol)
Mp (°C)
Benzyltriphenylphosphonium chloride
389
9–Anthraldehyde
206
104–105
trans 9-(2-Phenylethenyl)anthracene
280
131–132
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The experiment, step–by–step
CAREFUL! Aqueous sodium hydroxide is caustic and will cause painful burns if it
comes into contact with your skin.
Work in pairs for this experiment. Conduct the entire experiment in the hood.
1. Weigh 200mg of benzyltriphenylphosphonium chloride and 115mg of 9–
anthraldehyde into a 5mL round bottomed flask. Add a magnetic stirrer bar.
Arrange the flask above the magnetic stirrer unit using a small clamp.
2. Using the pipette provided add around 2mL of dichloromethane to the flask—
try to wash any solid material down into the bottom of the flask during this
addition of solvent. Turn on the power to the stirrer unit and stir gently for a
few minutes.
3. Turn up the speed until the reaction is stirring quite vigorously. Add,
dropwise, 0.25mL of 50% aqueous sodium hydroxide using the pipette
provided. Continue to stir at a fairly high speed for 30 minutes.
4. Turn the stirrer speed down to a gentle spin and add around 2mL of water to
the flask. Stir for a few minutes and then turn off the power. You should find
that you have a two–layer system (your yellow reaction product is currently
dissolved in the lower dichloromethane layer, the upper aqueous layer
contains inorganic, water soluble, material we no longer need).
5. Pour the contents of the flask into a small separatory funnel. Add around 5mL
of dichloromethane and 5mL of water. Swirl the two layers gently and allow
the them separate. Run the lower organic (yellow) layer into an Erlenmeyer
flask. Add another 5mL of dichloromethane to the separatory funnel, swirl,
and allow to separate once again. Run the lower organic (probably pale
yellow this time) layer into the Erlenmeyer flask containing the first organic
layer. Discard the aqueous layer in the aqueous waste container provided.
6. Add a few calcium chloride pellets to the yellow organic solution and swirl
gently. These anhydrous pellets will dry the organic solution by removing final
traces of water, the source of the cloudiness. Try not to swirl too vigorously
as the pellets will break up and make determination of the “dry point” tricky. If
the pellets seem clumpy and the solution remains cloudy then add a few
more pellets.
7. You will eventually generate a relatively clear solution. Decant the clear
solution carefully into a round bottomed flask. Add a little dichloromethane to
the pellets in the Erlenmeyer, swirl to dissolve any traces of yellow product,
and carefully decant these washings into the round bottomed flask. Use a
rotary evaporator to remove the solvent from the flask. Consult your TA or
instructor if necessary.
8. Add 3mL of 2–propanol to the yellow residue in the flask and heat gently on a
sand bath (set the control to around 50% to begin). Not much will happen
until the solution gets pretty hot, then the material will rapidly dissolve—swirl
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Organic Chemistry II Laboratory Experiments, Spring 2008(Nunez, Browne, Hamilton, edited 1/08 JW)
the solution a little to wash down any material on the sides of the flask. Once
dissolution is complete remove from the flask from the heat and let it cool.
After 10 minutes cooling in air cool the flask thoroughly in a beaker of ice.
Cool 2mL of 2–propanol on ice also.
9. Collect the product on the Hirsch funnel, dislodging the crystals from the
inside of the flask prior to filtering with a clean spatula. Wash out any
remaining crystals with a little ice cold 2–propanol. Dry with the passage of
air for 10 minutes.
10. Weigh the crystals and determine their melting point. Transfer them to glass
vial for use in next week’s experiment. Store the vial in your drawer.
Cleaning up
The aqueous solution from the extraction process should be discarded in the
aqueous waste container. The soggy calcium chloride pellets from the drying
process should be transferred to the solid organic waste container. The filtrate
from collection of the final product should be discarded in the non–chlorinated
waste solvent container. Rinse all of the glassware you have used with water and
then acetone, discarding the acetone washings in the non–chlorinated waste
solvent container.
Questions
1. Extraction of an aqueous solution with an immiscible organic solvent is a
crucial step in most organic preparations. With reference to the overall
reaction for this experiment, predict the location of ALL of the components of
the process (reactants, reagents, and products) after step 4.
2. Assuming the reaction were to proceed with 100% efficiency would an
extraction process still be required as part of the product purification
process? Justify your answer (Hint: you will need to look carefully at the
molar ratios of the reactants to work out what will be used, and what might be
left behind).
3. In very general terms, what molecular property of your product leads to its
ability to absorb visible light? (Hint: join up with a neighboring group and build
a molecular model using the kits provided in the laboratory).
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#8: Synthesis of a Cyalume and Chemiluminesence
Purpose of the experiment
To achieve a laboratory demonstration of chemiluminesence by preparation of a
cyalume, a molecule whose chemical decomposition releases energy that can be
transferred to a fluorescent molecule and emitted in the form of visible light. The
light sticks sold at fairgrounds use this exact chemical technology.
Introduction
The cyalume is essentially a bis–ester and is prepared from an alcohol and a
bis–acid chloride as shown below. The preparation is relatively straightforward.
When treated with peroxide, it decomposes to release energy.
O
O
O
H
Cl
Cl
Cl
Cl
Cl
Cl
Et3N
O
Cl
Cl
O
2
O
toluene
Cl
Cl
Cl
Cl
Cl
Cl
Cl
H2O2
C
C
O
O
O
C
C
O
Cl
Cl
This happens agin to the other
acid chloride
O
O
Cl
C
O
Cl
OOH
O
Cl
Cl
Cl
2HCl
+
Cl
OH
O
O
+
O
O
2
Cl
Cl
The chemiluminesence part of the experiment uses the cyalume product of this
experiment and the yellow product from the Wittig reaction. The latter acts as the
receiver of the chemical energy released on decomposition of the cyalume,
subsequently emitted as visible light. Different flourophores will give different
colors of light. If you prepare sufficient of the cyalume you may perform additional
fluorescence experiments using either 9,10–diphenylanthracene and/or rubrene
as the fluorophore.
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O
O
O
2CO2 + excited fluorophore
charge transfer complex
+
O
light
+
ground state fluorophore
fluorphore
Physical constants for the materials you will use or prepare
Material
Mol. Wt. (g/mol)
Mp (°C)
2,4,6–Trichlorophenol
197.5
64–66
Oxalyl chloride
127 (d 1.455)
Triethylamine
101 (d 0.726)
The Cyalume (product)
449
190–192
What you will do
Perform this entire experiment in the hood.
Part I. Preparation of a cyalume
Care! The entire preparative sequence should be performed in the hood. Oxalyl
chloride is corrosive and lachrymatory. 2,4,6–Trichlorophenol is a carcinogen and
this solid material should be weighed out in the dispensing hood at the side of the
laboratory. The other reagents are liquids and should be dispensed using their
accompanying syringes in your fume hood.
1. Weigh out 0.4g of 2,4,6–trichlorophenol and transfer it to a clean, dry 5mL
round bottomed flask (as with previous experiments, the reaction flask MUST
be both clean and dry). Add a stirrer bar and 3mL of toluene. Arrange the
flask above the magnetic stirrer unit and stir at a reasonable rate until the
solid has dissolved.
2. Using the pipette provided transfer 0.28mL of triethylamine to the stirred
mixture in the flask. Attach an air condenser (as before, use the water
condenser without the water hoses) and place a beaker of ice under the flask
and cool for 5 minutes while continuing to stir.
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3. Arrange your sand bath on top of the stirrer unit in order that the reaction
may be heated and stirred simultaneously, but don’t turn on the power to the
sand bath at this stage. Wipe any moisture from the bottom of the flask and
clamp the reaction assembly securely in position in the sand—check that the
stirrer still operates.
4. Withdraw 0.6mL of oxalyl chloride(2M in dichloromethane) in the syringe
provided with this reagent. Remove the air condenser and add the liquid
directly from the syringe into the stirred reaction mixture. DO NOT immerse
the tip in the mixture, or run the liquid down the side of the flask, rather add
dropwise with the tip of the syringe above the stirred solution. You should see
immediate formation of a fluffy precipitate that will rapidly thicken the solution.
5. Filter the solid material from the reaction mixture on the Hirsch funnel. If there
is a lot of material it will help to break some of this up with a spatula prior to
filtering. Use 1mL portions of hexane to wash remaining material out of the
reaction flask, and to wash the solid on the filter funnel. This process will
remove some of the brown coloration, if present, from your material. Draw air
through the solid for a few minutes and press it down on the filter with a
spatula to squeeze out the last traces of solvent.
6. Measure 5mL of water into a small conical flask and add your filtered material
(a useful method is to scoop a portion onto a spatula and then swirl this in the
water until the solid falls off). Once all the solid has been transferred add a
stirrer bar to the suspension and stir vigorously until all the solid is thoroughly
mixed with the water. While stirring clean up the Hirsch assembly, disposing
of the filtrate from the first filtration in the halogenated waste container.
7. Filter the aqueous suspension on the Hirsch funnel. Wash out the conical
flask with 2mL of water to collect any remaining solid and wash the collected
solid on the filter with a further 2mL of water. Leave the product on the funnel
to dry with the passage of air for a few minutes.
8. Transfer the solid to a small conical flask and add 2mL of toluene. Heat the
flask on a sand bath until the material dissolves, this often requires a gentle
boil. If there is a significant amount of a dark impurity visible on the inside
surface of the flask then you should pour the hot solution, carefully but
quickly, into a second conical flask. The dark oily impurities will remain stuck
to the sides of the original vessel.
9. Allow the solution to cool. Crystallization will soon begin. Complete the
crystallization process by cooling the flask in a beaker of ice. Clean up the
Hirsch apparatus while cooling, dispose of the filtrate from the last filtration in
the non–halogenated waste container.
10. Collect the crystals on the Hirsch funnel, using a little hexane to transfer the
last few crystals and to wash them on the filter. Dry with the passage of air
for a few minutes.
11. Weigh the crystals and determine their melting point.
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Part II. Chemiluminesence
1. Weigh around 50mg (0.05g) of the cyalume and 3mg of 9–(2–phenylethynyl)anthracene (the product from the Wittig reaction) into a glass vial.
2. In the hood, add 5mL of diethylphthalate and warm the entire mixture on a
sand bath (start at 50% power) until all the solids have dissolved. This
process can take a few minutes, be patient and resist the temptation to turn
the power up.
3. Take the vial (while still warm if possible) to the darkroom around the corner
from the laboratory. Your instructor or TA will be ready with a suspension of
hydrogen peroxide in diethylphthalate (0.2mL in 5mL) which they will show
you how to add to your solution to initiate the chemiluminesence reaction.
4. If you have extra of the cyalume you may return to point 1 of this section and
repeat the process with another 50mg of the cyalume and 3mg of an
alternate fluorophore, either 9,10–diphenylanthracene or rubrene. Give any
excess cyalume or 9–(2–phenylethynyl)-anthracene to your instructor.
Cleaning up
The solid materials from this week’s and last week’s reactions should be
discarded in the solid waste container after melting point determination–IF you
have spare cyalume than please check with your instructor before throwing it
away. The filtrate from the first collection of solid material should be transferred
to the halogenated waste container, that from the second to the non–
halogenated waste container. The filtrate from the final filtration should also be
discarded in the non–halogenated waste container. The contents of the vials
from the chemiluminesence experiment should be discarded in the separate
container marked for these samples. Rinse all of the glassware you have used
with water and then acetone, discarding the acetone washings in the non–
chlorinated waste solvent container.
Questions
1. It is stated at the beginning of this experiment that one of the materials you
will use is a carcinogen, another is lachrymatory. What do these terms
mean?
2. The chemiluminesence reaction you have performed will run for many hours
under normal conditions. If the reaction is run at an elevated temperature the
chemiluminesence is far brighter, but lasts for a much shorter time. Explain.
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#9: The Aldol Reaction: Synthesis of Dibenzalacetone
Purpose of the experiment
Design and carry out an aldol condensation reaction, analyzing the yield, purity,
and structure of the product.
Introduction
The Aldol reaction is among the most powerful and versatile methods of carbon–
carbon bond formation available to the synthetic chemist. This reaction, and its
variants, have made possible the syntheses of many familiar pharmaceuticals.
In this lab we will guide you to design and carry out an aldol condensation
reaction between acetone and benzaldehyde to form dibenzalacetone. As part of
your pre-laboratory preparation, work through the questions below. These will
help you to figure out ahead of time what you will be doing in the laboratory.
The general reaction for this aldol condensation can be written thus:
O
O
O
NaOH
2
+
CH 3
C
CH3
CH
H
C
H
C
CH
CH
Which carbonyl compound is the nucleophile, and which is the electrophile?
Draw out the mechanism for this reaction. Why is it called a mixed aldol
condensation?
What you will do
You will aim to prepare 2 mmoles of dibenzylacetone. You first must figure out
how much of acetone and benzaldehdye you will need. Do some research to fill
in the table of physical constants below, then use the molecular weights and
densities you found to calculate appropriate amounts of your starting materials.
Physical constants for the materials you will use or prepare
Material
Mol. Wt. (g/mol)
density
Mp (°C)
NaOH
-----------------------
------------------------
Dibenzalacetone
-----------------------
Acetone
Benzaldehyde
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Next, write out a flow chart of directions that you will follow when carrying out this
experiment. Several issues that you may wish to consider in your design are the
following:
1. How much acetone, benzyaldehyde, and NaOH will you need? How will
you measure it out? Think about state (liquid or solid?), volume/mass,
volatility, and causticity.
2. What kind and size of reaction vessel will you use?
3. What kind of solvent will you use? Do you even need a solvent?
4. Does this reaction need to be heated? If so, how will you heat it?
5. How will you know when the reaction is done?
6. How will you separate your product from the reactants and purify it?
7. How will you characterize your product?
Do not worry if you cannot answer all of these questions in advance, but do take
some time to think about them.
Cleaning up
Please dispose of your waste properly in the appropriate non-halogenated solid,
liquid, and aqueous waste containers.
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#10: The Grignard Reaction: Synthesis of Triphenylmethanol
Purpose of the experiment
To prepare a tertiary alcohol, triphenylmethanol, a material you will use again
later in the laboratory course. The Grignard reaction is one of the oldest methods
of carbon–carbon bond formation and remains a much used reaction in modern
synthetic chemistry.
Introduction
The reaction splits into three distinct steps, (i) preparation of the Grignard
reagent, (ii) action of this reagent on a ketone, (iii) isolation of the product. Parts
(ii) and (iii) are pretty straightforward. Part (i) can be tricky, simply because the
reaction involved in forming the Grignard reagent is difficult to initiate and
maintain in the presence of even TINY amounts of moisture. You will almost
certainly need assistance from your instructor or TA in this first stage.
Physical constants for the materials you will use or prepare
Material
Mol. Wt. (g/mol)
Mp (°C)
Bromobenzene
157 (d 1.491)
Magnesium
24
Benzophenone
182
48
Triphenylmethanol
260
160–163
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What you will do
Note! For the first two parts of the experiment the glassware you use MUST be
dry. This does not mean a quick rinse with acetone and a wipe with a paper
towel! YOU MUST USE ONLY the tubes, rods and vials that you placed in the
drying oven last week, and the rubber septa and syringes from the desiccators at
the back of the laboratory. Once underway keep tubes and vials capped
whenever possible, read ahead so that you don’t uncap before you are ready for
the next manipulation. There are THREE different bottles of ether, each used at a
different time in the experiment. DOUBLE CHECK that you are using the correct
reagent/grade. Some of these bottles have a rubber septum in place of a cap to
keep the contents dry. You must syringe from these bottles, do not remove the
septum. Ask if you are in any doubt whatsoever concerning any of the above.
Work in pairs for this experiment. Perform the entire reaction in the hood.
Part I. Synthesis of the Grignard Reagent:
1. Cap an oven dried reaction tube with a rubber septum (from the desiccator).
Cap two oven dried vials with caps (desiccator).
2. Weigh 50mg of magnesium powder and transfer to the reaction tube,
recapping with the septum as quickly as possible. Pierce the septum with a
syringe needle (from the desiccator) to allow for release of pressure.
3. With a dry syringe and needle (from the desiccator) transfer 0.5mL of
anhydrous diethyl ether to the reaction tube via the septum. Keep the syringe
ready.
4. Using the dropper provided with the reagent bottle transfer around 0.22mL of
bromobenzene to a dry vial (from the desiccator) and rapidly cap. With the
same syringe you used for point 3 add 0.7mL of anhydrous diethyl ether to
the bromobenzene in the vial. Immediately withdraw the solution (around
0.9mL total volume) into the syringe and then transfer the whole syringe to
the reaction tube, piercing the septum next to the pressure release needle.
5. Carefully add around 0.1mL of the bromobenzene solution to the reaction
tube and flick to mix the contents. Look for the small bubbles of gas that
indicate that the Grignard formation reaction is underway. It is very likely that
the reaction will not begin spontaneously, collect a glass rod from the oven
and call your TA or instructor. They will show you how to initiate the reaction
by grinding the surface of the magnesium powder—this process can be
repeated numerous times if the reaction, once begun, slows again within a
few minutes. Once underway the exothermic reaction may cause the small
volume of solvent in the tube to boil—a water moistened pipe cleaner
wrapped around the top of the reaction tube will act as a mini condenser (but
be careful to keep the damp pipe cleaner clear of the septum and syringe!).
6. Once all the bromobenzene solution has been added, and the reaction
appears to have calmed down, remove the syringe, add a stirrer bar
(recapping afterward with the septum), arrange the tube above the stirrer
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unit, and stir gently for a few minutes. A little more reaction may occur. The
Grignard reagent is now ready.
Part II. Action of the Grignard Reagent on Benzophenone:
1. Weigh 0.36g of benzophenone into a clean, dry vial.
2. Using a fresh syringe and needle (from the desiccators) add 0.5mL of
anhydrous t–butyl methyl ether to the vial, cap, and shake gently to dissolve
all the solid. Withdraw the solution into the same syringe and transfer it to the
reaction tube, puncturing the septum next to the pressure release needle.
3. With stirring, add the benzophenone solution dropwise to the Grignard
reagent. The mixture will become thick and you will probably see a deep
red/purple coloration that will fade as the reaction reaches completion.
4. Once all the benzophenone solution has been added, stir for a further 10
minutes.
Part III. Isolation of the Triphenylmethanol Product
1. The reaction is no longer air sensitive so you may remove the septum and
breathe a sigh of relief.
2. Transfer the sticky contents of the reaction tube, including the stirrer bar, to a
small Erlenmeyer flask. The bulk of the reaction mixture should come out
relatively easily with the use of a glass rod. Use a little more diethyl ether (not
anhydrous) and a little more water to transfer remaining residues from the
tube.
3. With gentle stirring, cautiously add 2mL of 3M aqueous HCl to the mixture.
There will likely be a few remaining bits of magnesium metal that will react
vigorously with the acid solution—let the reaction calm down a little before
you add more acid. Once all the acid has been added and the excess
magnesium destroyed turn off the stirrer—you should have a clear lower
aqueous layer and a yellow upper organic layer.
4. Transfer the two layer system to a small separatory funnel. Carefully run off
the lower aqueous layer, leaving the ether layer in the funnel. Add an equal
volume of saturated aqueous NaCl, swirl the two layers gently together, and
allow them separate. Run off the lower aqueous layer. Collect the organic
layer in a small Erlenmeyer flask and dry with a few anhydrous calcium
chloride pellets.
5. Decant the dried ether extracts into a small round bottomed flask. Wash the
drying agent with a few mLs of ether and decant these into the round
bottomed flask. Remove the solvent on the rotary evaporator. Consult your
instructor or TA if necessary.
6. Dissolve the crude solid in the minimum volume of hot diethyl ether by
warming gently on a sand bath—start with around 1mL and add tiny extra
volumes until dissolution is complete. Pour the solution into a small conical
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flask and add 1.5mL of ligroin. Heat gently on the sand bath to maintain
gentle boiling, watch very carefully for the point at which the solution
becomes ever so slightly cloudy. Remove from the heat and allow the flask to
cool. Complete the crystallization process by cooling the flask in a beaker of
ice for at least 10 minutes.
7. Collect the crystals on the Hirsch funnel, returning a little of the filtrate to the
flask to wash out any remaining crystals. Dry with the passage of air for a few
minutes.
8. Weigh the dried crystals and determine their melting point.
Cleaning up
The aqueous layers from the two extraction processes should be discarded in the
aqueous waste container. The filtrate from collection of the final recrystallized
material should be discarded in the non–chlorinated waste container. Rinse all of
the glassware you have used with water and then acetone, discarding the
acetone washings in the non–chlorinated waste solvent container.
Questions
1. You took elaborate precautions to exclude moisture from your reaction.
Why?
2. After you added the benzophenone to your Grignard reagent you will have
seen the reaction mixture become thick with precipitate. What is this
precipitate? After the addition of aqueous acid the whole reaction cleans up
to give you two clear liquid layers, one organic and one aqueous. What
chemistry has occurred to prompt this change? At this stage of the procedure
in which layer are the various materials you have used, and products you
have made, situated?
3. With reference to your answer to question 2, what do you think the purpose
of the final recrystallization step might be?
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#11 Biosynthesis and Distillation of Ethanol
Purpose of the experiment
To set up a simple fermenter and prepare ethanol; to isolate ethanol by
distillation; to analyze the purity of the final product.
Introduction
With the price of petroleum reaching record high levels due to increasing demand
in Asia and decreasing supplies worldwide, production of ethanol for fuel from
renewable biological sources is becoming increasingly common. Ethanol is
generally produced from the fermentation of a sugar source (corn, sugarcane,
sugarbeets, etc.) by microbes. After purification, the ethanol thus produced can
be blended with gasoline or can replace gasoline entirely as a fuel for cars.
Because the ethanol is produced from a renewable biological source, its impact
on atmospheric carbon dioxide levels is theoretically less than gasoline.
However, concerns about deforestation and the significant energy demands of
fertilizer production must be taken into account when calculating the net effect of
burning biofuels.
Part I: Biosynthesis of Ethanol
What you will do
We will use yeast to synthesize ethanol through the process of fermentation. In
lieu of corn we will use sucrose (table sugar).
Today we will set up the fermentation, which will be allowed to run until next
week when we will distill the reaction solution to isolate the ethanol.
Mix half an envelope of dry yeast in 50 mL of water in a beaker. Add 0.35 g of
disodium hydrogen phosphate to the beaker and transfer the contents to a 500
mL Florence flask. Dissolve 51.5 g of sugar in 150 mL of water and add to the
500 mL flask. Attach the flask to the apparatus shown below, adding a saturated
solution of calcium hydroxide to the test tube. This test tube allows CO2 to
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Organic Chemistry II Laboratory Experiments, Spring 2008(Nunez, Browne, Hamilton, edited 1/08 JW)
bubble out without allowing oxygen to enter the flask. In the presence of oxygen
acetobacter bacteria can convert the alcohol to vinegar (acetic acid).
Come back 3 times during the week to see when the bubbling ceases. Since
CO2 is a product of the reaction, when no more CO2 is being evolved the reaction
is over. The ideal temperature for fermentation is 35˚C.
Part II: Distillation of Ethanol
Introduction
As you will remember from the Banana Oil lab, distillation is a method for
separating two or more liquids based on differences in their boiling points. The
boiling point is the temperature at which the vapor pressure of the liquid equals
the vapor pressure of the atmosphere, i.e. ~760 torr or 1 atm. Every pure liquid
has a characteristic boiling point, which reflects its structure and intermolecular
interactions in the liquid phase. These intermolecular interactions must be
overcome for the molecules to escape the liquid phase and vaporize.
For a mixture of two miscible “ideal liquids” (similar to “ideal gases” in that there
are no complicating intermolecular interactions, and that they also don’t really
exist, but we talk about them for the sake of simplification), the observed boiling
point of the mixture will be somewhere between the boiling point of each of the
two component liquids because the vapor pressure above the liquid has
contributions from both liquids. That is,
Ptotal = Pliq1 + Pliq2
Here is a nice graph to demonstrate this principle for a mixture of pentane and
hexane:
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As you can see, the total pressure and boiling point depend on the relative
amounts of each liquid in the original mixture (or the “mole fraction” of each
liquid). The pressure of each individual liquid in the vapor is the product of the
vapor pressure of the pure liquid scaled down by the fraction of that liquid:
Pliq1=P°liq1 x mol fraction
This little rule, called Raoult’s law, makes sense if you think about it: there will
be less pentane in the gas if only half of the liquid mixture is made up of pentane
than if all of the liquid was pentane.
You will notice something important if you look carefully at the graph above.
Even when the liquid mixture is half pentane and half hexane, the vapor is mostly
pentane because pentane has a higher vapor pressure and lower boiling point.
Thus we can begin to separate the liquids by boiling them and collecting the
vapor, the process of distillation.
Fractional (or Complex) Distillation
One problem that Raoult’s law makes clear is that when the boiling points and
vapor pressures of the two liquids are similar (such as in the example above),
both of them are vaporized to some extent during a distillation and the distillate is
still a mixture— enriched in one liquid, but a mixture nonetheless.
To improve this, we can distill again, and again, and again, each time improving
the fraction of our desired liquid in the distillate and getting closer to its pure
boiling point. Instead of doing the entire distillation from scratch every time,
which would be horribly tedious, we can use a special distillation column with a
lot of surface area that allows for repeated vaporizations and condensations all at
once. There are different kinds of these columns, but mostly they look like your
regular condenser but with little prongs and tongues sticking out or a long helical
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interior or wire mesh shoved inside. You should have one or two of these in your
lab drawer.
The efficiency of the distillation process is expressed as its number of
theoretical plates. A standard, simple distillation might be described as having
one theoretical plate; by using a fractionating column, the number of theoretical
plates might be increased to two, four, eight, or more (i.e. the equivalent of two,
four, eight or more simple distillations).
Not only does the efficiency of separation dramatically improve with fractional
distillation, but the temperature at which the distillation occurs is sharper because
each liquid emerges into the receiving flask more or less independently.
Azeotropes
Real liquids don’t always behave like ideal liquids. The water-alcohol mixture
we’re working with is not ideal because of significant intermolecular interactions.
(What kind?) We can illustrate the problem graphically using what are called
temperature-composition diagrams. The temperature-composition diagram
for a benzene-toluene mixture (which is close to ideal) looks something like this:
This diagram is admittedly a little hard to read, but it is very useful. It has two
lines, one for the liquid composition (bottom) and one for the vapor (top). You
start on the x-axis at the starting composition of the liquid, let's say 30%
benzene. As you heat the liquid, you go upwards from 0.30 mole fraction of
benzene (that is, 30%) until you hit the liquid line at point a, around 371 K. This
means that your mixture will boil at 371 K. Then, you draw a horizontal line,
called a tie line, until you hit the vapor line at point b. Point b tells you the
composition of the vapor phase at the same temperature, 371 K, about 50%
benzene. As you would expect for a simple distillation, the vapor is enriched in
benzene relative to the liquid. This process corresponds to one theoretical plate.
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If you then cool the vapor to a liquid (point c) and distill again, the new boiling
point is about 365 K. You go across the new tie line and your new vapor (point
d), about 70% benzene. This is your second simple distillation and your second
theoretical plate. Note the "staircase effect" as you do successive distillations,
moving toward 100% benzene. A fractional distillation would be expected to
accomplish multiple steps all at once.
Now let's look at the temperature-composition diagram for the non-ideal mixture
of ethanol and water, below. It has a minimum where the two lines meet, giving
the graph a bow-tie like appearance, and making it an azeotrope.
What you will do
When you return for the second lab, carefully decant the liquid in your fermenter
(do not mix!) and filter through a ~1/2 inch layer of filter aid using a large buchner
funnel. This wet, mud-like substance will help the filtration. You will then distill
the filtrate using a simple or complex distillation apparatus of your choice. Put
the fractionating column above the round bottom flask but before the condenser,
like so:
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Another
fractionating
column
Water OUT
Water IN
Fractionating column
Receiving Flask
Warm Sand Bath
You will need to determine when to stop the distillation, i.e. when no more
alcohol remains. Before you begin the distillation, determine the temperature
and concentration that you expect for the azeotrope and for pure water.
To determine your final yield of ethanol, you will need to determine what
percentage of your product is ethanol. You may do this in one of two ways: by
measuring the density of your product (i.e. the weight per 1 mL of volume), or by
using GC to measure the fractions of ethanol and water in your product. In either
case, you will first want to construct a standard curve against which you can
calibrate your results, using a set of mixed ethanol:water standards.
Cleaning Up
Any of the alcohol product or standards should go into the non-halogenated
organic waste. The debris from the cells (which are now dead) along with the
filter-aid can go into the solid organic waste. The limewater (calcium hydroxide
solution) can go into the base waste.
Questions
1. What is the boiling point of pure water? Of pure ethanol? Of the mixture?
2. Above a pentane/hexane mixture is used as an example of an ideal mixture,
but water and ethanol is a non-ideal mixture because of additional intermolecular
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interactions that complicate the distillation. What intermolecular interactions are
those?
3. Use the temperature-composition diagram for the non-ideal mixture of ethanol
and water (given above) to explain your expected and observed results. Why are
azeotropes also called "constant boiling" mixtures? Include an explanation of
what an azeotrope is, what was your boiling point, and what was the percentage
of ethanol in your final mixture. What is the best result you might have expected
for your distillation, and how would you have predicted that from the graph? Why
might you not have observed even this best possible result?
For the next two questions you may have to do a bit of research:
4. At the end of even the most careful distillation, ethanol still contains some
water. However, water is very bad for internal combustion engines. How is the
remaining water removed industrially to prepare ethanol for fuel use?
5. What is “cellulosic ethanol” and what are the promises and challenges of this
technology?
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#12: Hydrolysis of Biologically-Derived Triacylglyceride Esters
Purpose of the experiment
To utilize the base-catalyzed hydrolysis of esters to make soap from biologicallyderived oils and fats.
Introduction
“Lipid” is a general term for biological molecules that are more soluble in organic
solvents than in water. This broad term includes fatty acids, cholesterol, fatsoluble vitamins, steroid hormones, and other dissimilar organic molecules.
Fatty acids are composed of long alkane or alkene chains (the “fatty” part)
terminated by a carboxylic acids (the “acid” part). These fatty acids are then
linked as esters to glycerol to form triacylglycerols, molecules used primarily for
energy storage. The formation of a generic triacylglycerol is shown below:
fatty acid
glycerol
triacylglycerol
O
O-
3
n
OH
+
OH
O
n
O
n
O
O
OH
O
O
n
esters
Alternatively, fatty acids can be linked to other small hydrophilic molecules to
make phospholipids, ceramides, and other biologically-important cellular
membrane components.
In water at neutral pH, the reverse reaction, i.e. the hydrolysis of esters to
carboxylic acids, is thermodynamically favorable but kinetically slow (why?). In
biological systems, the hydrolysis reaction to break down triacylglycerols to
glycerol and fatty acids is catalyzed by enzymes called lipases. Alternatively, we
catalyze this reaction with base. Please draw out for yourself in your
notebook the mechanism for base-catalyzed ester hydrolysis based on
what you have learned in class. Interestingly, this reaction has been used for
centuries to produce soap from fats using lye (otherwise known as KOH).
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What you will do
You will choose a source of fatty acids from those available in the laboratory,
which may include lard, olive oil, vegetable oil, corn oil, and/or cheese. Please
feel free to bring something from home if you would like.
You will then hydrolyze ~10 g of this fat source with ~ 5 g of sodium hydroxide in
40 mL of a 50-50 mixture of water and ethanol. Heat in a “double-boiler,” stirring
frequently; neutralize, and cool. Filter your solid product, which is your “soap.”
Please be sure to describe the progress of this reaction in your lab notebook.
Remember that sodium hydroxide is corrosive; you must wear splash goggles
and gloves when working with this material.
Cleanup
All acids and bases must be fully neutralized before disposal. Please put solid
waste into appropriate containers and not down the sink where they would
assuredly clog the drain.
Questions
1. What do you think of your product? Does it at all resemble soap? Inspect your
neighbors’ products, and within your class decide which fat sources made the
most appealing soaps.
2. Determine your yield of fatty acid (in grams per gram of starting material), and
hypothesize about the missing mass.
3. Where did the glycerol go?
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