The Organic Chemistry II Laboratory Course Chem 302, Spring 2008

________________________________________________________________ Organic Chemistry II Laboratory Experiments, Spring 2008(Nunez, Browne, Hamilton, edited 1/08 JW)

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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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(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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Chemistry 302: Organic Chemistry II, Spring 2008 Week of Chapter and Topic Lab Experiment

Intro to class, review Diels-Alder Jan 30th

Ch 13 NMR None

Ch 13 NMR Ch 13 NMR Feb 4th

Ch 13 NMR

#1 The Diels–Alder Reaction (set up Spectroscopy)

Ch 13, 14 NMR, Benzene Ch 14 Reactions of Benzene Feb 11th

Ch 14, 15 Reactions of substituted benzenes

#2 Spectroscopy

Ch 15 Reactions of substituted benzenes Ch 15 Reactions of substituted benzenes Feb 18th

Ch 16 Reactions of carbonyls

#3 Tropylium

Ch 16 Reactions of carbonyls Ch 16 Reactions of carbonyls Feb 25th

review

#4 Nitration of Methyl Benzoate

Exam #1 Ch 17 Aldehydes and Ketones

Mar 3rd Ch 17 Aldehydes and Ketones

#5 Banana Oil

Ch 17 Aldehydes and Ketones Ch 17 Aldehydes and Ketones Mar 10th

Ch 18 Intro to reactions at the alpha carbon

#6 Preparation of a Naphthalene Diimide

☺ Spring Break!!! Yeah!!! ☺

Ch 18 Substitutions at the alpha carbon Ch 18 Reactions at the alpha and beta carbons Mar 24th

Ch 18 Condensations

#7 The Wittig Reaction

Review of carbonyls reveiw Mar 31st

Exam #2

#8 Chemiluminesence

Ch 19 Reductions Ch 19 Oxidations Apr 7th

Ch 19 Oxidations

#9 The Aldol Reaction

Ch 20 Amines Ch 20 Amines Apr 14th

Selected topics in biological chemistry

#10 The Grignard Reaction (set up Ethanol lab)

Selected topics in biological chemistry Selected topics in biological chemistry Apr 21st

Selected topics in biological chemistry

#11 Biosynthesis and Distillation of Ethanol

Exam #3 Ch 29 Pericyclic Reactions Apr 28th

Ch 29 Pericyclic Reactions

#12 Hydrolysis of Esters Check-out

May 5th Last day of classes -----------------------------------------


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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:

OO 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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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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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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#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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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? 1H 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. 13C 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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#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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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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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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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).

3. 13C 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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#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

+HNO3

H2SO4

O

OO2N

?



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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#5: Preparation of Synthetic Banana Oil by a Fisher Esterification 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-1-butanol) 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:

Density Boiling point, ˚C Acetic acid 1.049 118 Isopentyl alchohol 0.809 130 Isopentyl acetate 0.876 142 Sulfuric acid 1.84 290

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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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 narrow-stem 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:

OO

O

O

O

+ 2 NH2CH2(CH2)4CH3 NN

O

O

O

CH2(CH2)4CH3CH3(CH2)4H2C

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.



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


21

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.


22

#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.



Benzyltriphenylphosphonium chloride 389

9–Anthraldehyde 206 104–105

trans 9-(2-Phenylethenyl)anthracene 280 131–132

(C6 H5 ) 3 PHCH2

Cl

50% NaOH(C6H5)3P CH

The Wittig reagent,an ylide

( C6 H5) 3 P CH

O H

(C6H5)3 P CH

O CH

Benzy ltripheny lphosphoniumc hloride

9-Ant hraldehyde

CC H

H

(C6 H5)3P

O

trans-9-(2-Phenylethenyl)anthracene

Tripheny phosphineox i de (TPPO)


23

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


24

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.

Cl Cl

Cl

O

Cl

O

Cl

OO

Cl

Cl Cl

C

O

O

Cl Cl

Cl

O

2Et3N

toluene

HCl

Cl

This happens agin to the other acid chloride

O

Cl

Cl Cl

C

O

C

O

O

Cl

Cl

Cl

2HCl

H2O2

O O

O O

+

+

O

Cl

Cl Cl

C

O

C

O

OOH

2

OH

Cl

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.


26

2CO2 + excited fluorophore

O O

O O

+charge transfer complex

lightfluorphore

ground state fluorophore+



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.


27

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.


28

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:

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)

Acetone

Benzaldehyde

NaOH ----------------------- ------------------------

Dibenzalacetone -----------------------

CHC H 3 C CH3

O

+ 2 O

H

NaOHCH

C C H

O

C H


30

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.


31

#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.



Bromobenzene 157 (d 1.491)

Magnesium 24

Benzophenone 182 48

Triphenylmethanol 260 160–163


32

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


33

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


34

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?


35

#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


36

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:


37

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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Water IN

Water OUT

Warm Sand Bath

Receiving Flask

Fractionating column

Another fractionating column

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 biologically- derived 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, fat-soluble 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:

O

O-

n3

+

OH

OH

OH

glycerolfatty acid

O

O

O

O

nO

nO

n

esters

triacylglycerol

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?

Documents

The Organic Chemistry II Laboratory Course Chem 302, Spring 2008