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NANYANG TECHNOLOGICAL UNIVERSITY

SCHOOL OF CHEMICAL AND BIOMEDICAL

ENGINEERING

DIVISION OF CHEMICAL AND BIOMOLECULAR

ENGINEERING

CH 2071

ORGANIC LABORATORY MANUAL

PREPARED BY: DR. ZAHER JUDEH

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Table of Contents

Page

Introduction 03

Safety in the Chemical Laboratory 05

Laboratory Notebooks 08

Notes on Laboratory Equipments and Techniques 09

Notes on Laboratory Quantitative Techniques 16

Laboratory Experiments:

Experiment 1 Synthesis of Aspirin from Salicylic Acid 20

Experiment 2 Characterization of Aspirin Synthesis Product 25

Using IR and NMR Spectroscopy

Experiment 3 Determination of Aspirin Content Using Visible 30

Spectroscopy

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INTRODUCTION

The CH 2022 practical course consists of 3 experiments which are designed to familiarize

you with organic synthesis, characterization and quantitative analysis. You will also learn

different techniques during the experimental course e.g. TLC, crystallization, melting

point determination...etc. You are required to purchase an exercise book before you

attend the first experimental session. This book will serve as a laboratory notebook where

you are required to record your experimental work.

Important points to note:

(A) Report writing

A formal report is required for each experiment. Reports are due at the beginning of the

next lab period after you complete an experiment. Your final report must consist of the

following:

1. Title of the experiment, your name (and your partner name, if any) and the date of

the experiment.

2. Introduction: This section should state the objectives of the experiment. Presents a

description of what to be done and the general approach to the problem. Present

structures and balanced equations.

3. Experimental Section: this section must include-

a) Procedure: a summary of how the experiment was conducted and the

equipment used. Describe in your own words how you performed the

experiment and the difficulties you met and how you solved them.

b) Raw Data: Record all experimental data and observations as it is obtained.

c) Results: Present your data in tables and graphs when possible. Results

(quantitative observations / properties of product, physical appearance,

m.p., b.p., % yield) must be reported. All graphs, tables and sections must

have a title/caption and should be referenced in your text.

d) Discussion and Conclusions: Discuss your findings; make comparisons

with known values if available. Elaborate on possible sources of errors,

selectivity and sensitivity of the technique, detection limits, interferences,

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accuracy, precision, applicability, etc. Suggest any possible improvements

in the experiment and present your summary conclusion.

e) Questions: Answer all questions at the end of each experiment clearly.

(B) Grading Guidelines

Your grade will depend on: a) your understanding of the experiment and the results

obtained; b) your experimental technique and attitude in the lab; c) presentation of the

results in the final formal laboratory report and

Your technique will be evaluated at the end of the laboratory course. Technique points

will be awarded at the discretion of the instructors. The maximum number of points will

be awarded to students who:

- Show up on time

- Are fully prepared for lab

- Follow all safety rules

- Properly dispose of waste

- Clean up their bench and common work areas

(C) Product Labels

Each product you synthesize should be submitted in a vial with the report. The vial

should be labeled with student’s name, product name, product amount (in grams), %

yield and measured physical data such as melting or boiling point range. The quality of

your products will be graded base on yield and good quality.

(D) Time Management

It is expected that students will come to the lab fully prepared and will work at a

reasonable pace. All of our experiments have been designed and tested so that they may

be completed in less than two and one-half hours if the student is adequately prepared. If

for some un-foreseen circumstance a fully prepared student who works efficiently is

unable to complete an experiment, the student may make arrangements to complete the

experiment on another day. This option is available at the discretion of the instructor.

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SAFETY IN THE CHEMICAL LABORATORY

Safety comes first: Safe practices in the chemical laboratory are of great importance and

they are everyone’s responsibility. Some chemicals are toxic, flammable, explosive

carcenogenous and thus one must acquire a basic knowledge and understanding of the

chemical properties and equipments he/she is dealing with in the lab. One should realize

the types of hazards that exist and the accidents and injuries that can result from

ignorance or irresponsibility on the part of the student or a neighbor from bad planning,

ignorance or simply by having indifferent attitude.

(A) Safety regulations

1. Wear safety glasses or goggles at all times while in the laboratory. Take your

contact lenses out prior to entering the laboratory.

2. Gloves must be worn when performing experiments.

3. Never eat or drink in the laboratory.

4. Caution must be taken to prevent chemical contact with the skin or clothing.

5. Never leave an experiment unattended. Inform the instructor or lab assistant if

you must leave the lab.

6. Clean up all chemical spills immediately.

7. To avoid contamination of laboratory supplies, do not use personal equipment

such as spatulas in shared chemicals and replace all lids after use.

8. Be fully aware of what you do in the lab. Ask you instructor if in doubt.

9. Use all electrical and heating equipment carefully to prevent shocks and burns.

10. Become familiar with the location of safety equipment, such as eye wash, fire

extinguisher, fire blanket, emergency shower.

11. Never handle broken glassware with your bare hands; use a broom and a dust pan.

12. After the experiment is completed, turn all equipment off, making sure it is

properly stored, and clean your area.

13. Wash your hands well at the end of the laboratory.

14. Report all accidents to the instructor or lab assistant immediately.

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(B) Safety Facilities

The laboratory is equipped with various safety facilities. Students working in the lab must

familiarize themselves with the location and proper use of these facilities.

1. Eye wash: In the event of a chemical splash in the eye, the eye, with the eyelids

widely open, should be washed immediately with copious amounts of water for at

least 15 minutes. The casualty must seek medical attention immediately.

2. Shower: If a chemical spill on the body, immediately flush the body with

generous quantity of water for at least 15 minutes. All contaminated clothes

should be removed.

3. Fire extinguisher: Use CO2 fire extinguisher to put out small fires resulting from

chemical incidents. Halohydrocarbons extinguishers should be used only when no

chemicals are involved in the fire. Never use water to extinguish a laboratory fire.

4. First aid boxes: Use when necessary.

5. Fire blanket: cover the affected area of the body to prevent the fire from spreading.

Never try to blow out a fire; severe facial burns may result.

(C) Personal Attire

Proper clothing must be worn in the laboratory and students must wear a lab coat in the

lab. Lab coats can be bought from the division office. A lab coat is to help keep clothes

protected and close to the body. Long, floppy and loose clothes can easily come into

contact with chemicals. If your hair is long enough to interfere with you practical

operation it should be tied back. Jewelry on your hand can be ruined if they come in

contact with chemicals and therefore must be removed.

Open toed shoes do not adequately protect you against chemical spills. No open footwear

or high heeled shoes are allowed in the lab.

(D) Equipment and Glassware

Equipment should be assembled in the most secure and convenient manner. Consider the

safe location of the equipments you use e.g. hot plate. Keep it away from the bench edge.

Handle all laboratory glassware with care since they are usually fragile. Serious injuries

may result if they are not properly handled. Apparatus that can roll should be placed

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between two immobile objects away from the edge of the bench. Chipped or broken

glassware should never be used. After the experiment is completed, all glassware should

be emptied, rinsed, cleaned and returned to your drawer.

(E) Personal belonging

Personal belonging must not be placed on working benches or floor of the laboratory.

They should be kept in their designated place on the shelves.

Chemistry Lab Sketch

I have read and understood the safety in the chemical laboratory guidelines.

Name: Date:

Signature:

Please sketch the laboratory below and indicate the location of all safety items in the lab.

Also indicate first aid kit and phone locations in relation to the lab. Then, have your

instructor check and initial before starting any experimental work.

Instructor initials:

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LABORATORY NOTEBOOK

You are required to use a notebook in the lab to record all primary experimental data and

observations during the lab period. Take note of any special instructions given by your

lab instructor at the start of each experiment. Your notebook should be a complete record

of the work you do in the lab. All data must be recorded immediately in ink; recording

data or observations on scrap paper is not acceptable. You or other chemists should be

able to understand the notes in the future, not just during the current experiment.

A copy of all spectra should be placed permanently (with glue) into the notebook. Be sure

to include a caption.

General guidelines to be followed:

1. Always bring your notebook with you to lab.

2. Record directly, in ink, on your notebook and not on loose scraps of paper.

3. Write down all observations such as color and phase changes - don't rely on your

memory.

Before coming to the lab, your notebook should have:

1. Title, date and objectives of the experiment

2. A short procedure for the experiment

3. Any reaction equations

4. Chemical formulas and calculation of molecular masses/weights

After the lab is finished, the notebook should contain:

1. Raw data from the experiment you performed

2. All observations made during the experiment, e.g.: color change, temperature

change, physical state (solid, liquid, gas) change….etc.

3. Calculations and final answers

4. Write up a short conclusion for the lab - your conclusion should address the

objectives of the lab - did you achieve the objective? Any difficulties you met and

how you solved them?

5. All information needed for your report sheets should be in your notebook first and

then transferred to the report sheet.

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NOTES ON LABORATORY EQUIPMENTS AND TECHNIQUES

I. The electronic balance

The balances to be used during this laboratory course are electronic and have a button to

"tare" them to “Zero”. To weigh an object, simply press the tare button to “zero” the

balance, then place the object you wish to weigh on the pan and read the weight directly

off the LCD screen. ALL Balances have limits. If you exceed the maximum weight that it

can measure it will post an error message on the readout. Keep the lids of the balance

closed when weighing to cut down on the air current that causes the readout value to

fluctuate. In case you spill a chemical inside or next to the balance make sure you clean it

immediately.

Never place chemicals directly on the weighing pan! Always use weighing paper, beaker,

or flask, to contain the chemical you wish to weigh. When using weighing paper to weigh

a powder or a crystalline chemical, fold the paper in half first before placing it on the

weighing pan and “tarring” the balance. The chemical to be weighted can be placed near

the fold on the paper. The fold will make pouring the powder off the paper much easier as

you transfer it to a container. Sometimes, it is easier to weigh directly into a beaker or a

flask to be used.

II. Special set-ups for filtration

Filtration is needed to isolate the product of interest whether it is liquid or solid. If you

wish to isolate a liquid from a solid, you would use a gravity filtration. It is less likely

that small pieces of solid will be pulled through the filter paper by gravity. Likewise, if

you are trying to isolate a pure clean solid from a reaction mixture you would be better

off using a vacuum filtration set-up which speeds up the passage of liquid through the

filter paper and makes it easier to wash contaminants out of your solid product.

Gravity filtration:

A filtration apparatus for obtaining pure liquids, gravity filtration can be set up using the

following components: stand, ring, glass funnel, beaker and filter paper

Vacuum filtration:

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For obtaining pure solids, vacuum filtration, can be set up using stand, clamp, Buchner

funnel, Buchner/Sidearm flask, vacuum tubing and filter paper (Make sure the filter

paper paper fits in the funnel, covering all the holes but not touching the sides.)

III. Crystallization:

Crystallization is a fundamental purification technique used to purify solid compounds

from impurities. It is based on the principles of solubility: compounds (solutes), tend to

be more soluble in hot solvents than they are in cold solvents. If a saturated hot solution

is allowed to cool, the solute is no longer soluble in the solvent and forms crystals of pure

compound. Impurities are excluded from the growing crystals and the pure solid crystals

can be separated from the dissolved impurities by filtration.

Crystallization Technique:

To crystallize an impure solid compound, dissolve it completely in hot solvent in a flask.

The flask then contains a hot solution, in which solute molecules (both the desired

compound and impurities) move freely among the hot solvent molecules. As the solution

cools, the solvent can no longer “hold” all of the solute molecules, and they begin to

leave the solution and form solid crystals. During this cooling, each solute molecule in

turn approaches a growing crystal and rests on the crystal surface. If the geometry of the

molecule fits that of the crystal, it will be more likely to remain on the crystal than it is to

go back into the solution. Therefore, each growing crystal consists of only one type of

molecule. After the solution has come to room temperature, it is carefully set in an ice

bath to complete the crystallization process. The chilled solution is then filtered to isolate

the pure crystals.

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Vacuum filtration Gravity filtration

Summary of the recrystallization steps:

1. Selection of solvent: The choice of solvent for crystallization is based on "Like

dissolves like" principle. Use the solvent specified in the experimental section to

crystallize your compound. Solid

2. Dissolution: Crush or ground the solute in a small Erlenmeyer flask to a fine

powder. Make sure you have 2-3 boiling chips in the Erlenmeyer flask to prevent

the solution from boiling over. Add in small portions of hot solvent then keep the

solution hot on a hot plate. Add enough solvent to obtain a clear solution a part

from some insoluble impurities. If crystallization begins fairly rapidly after

removing from the heat, add more hot solvent then heat till all of the crystals

dissolve. The solution is now ready for filtration. See step 3 below before you go

ahead and filter according to step 4.

3. Decolorizing the solvent with charcoal: Sometimes, the crude product may

contain colored impurities which must be removed before recrystallization. Cool

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the above solution somewhat and carefully add charcoal to it using a spatula. If

charcoal is added to a superheated solution, the solution may foam excessively

and boil over. Therefore, you must cool the solution somewhat before adding

charcoal to it. Add a small portion (about 10% of the amount of sample) of

charcoal and swirl the flask then heat gently for a few minutes. Proceed to step 4.

4. Filtering the hot solution: Use gravity filtration to filter the hot solution through

a fluted filter paper supported in a preheated short stemmed glass funnel. To

prevent the solute from crystallizing in the filter paper and funnel, add excess (5-

10%) hot solvent to the original solution. The solution is filtered and then

concentrated by evaporation. If any material crystallizes in the filtration assemble,

scrap it back into the conical flask, redissolve it and then filter as before.

5. Crystallizing the solute: Let the hot saturated solution cool slowly and remain

undisturbed until it reaches room temperature. If no crystals appear try to induce

crystallization by adding a seed crystal to the solution. Crystallization also may be

induced by scratching the insides of the flask with a glass rod. After the solution

has cooled to room temperature, immerse it in an ice bath for few minutes to

complete crystallization.

6. Collecting and washing the crystals: Before vacuum filtration, make sure that

the filter paper is moistened with the solvent used for crystallization and also that

the funnel’s perforated disc is sealed. Swirl the contents of the flask and pour on

to a moisten filter paper in a Buchner or Hirsch funnel. Make sure to transfer all

the crystals to the funnel. This can be done with the help of a spatula or by adding

a small amount of the cold recrystallization solvent, or some of the filtrate to the

flask, swirling it and then quickly pouring the mixture into the funnel. Turn off

the vacuum while you add a small amount of cold wash solvent (same solvent

used for crystallization) and then turn the vacuum again. Remove the solvent from

the filtration flask, reassemble the filtration apparatus, then turn on the vacuum

and continue to dry the crystals by passing air across the crystals.

7. Drying the crystals: Crystals should be dried from the recrystallization solvent. If

water was part of the recrystallization solvent, it may be removed by placing the

crystals with the filter paper on a watch glass to dry under reduced pressure in a

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vacuum desiccator over silica gel desiccant. If the melting point of the crystals is

above 120°C, you can dry it under a heat lamp or in an oven kept at 110°C. Don't

scrape the filter paper until it's dry. If organic solvents, other than ethanol and

methanol, have been used for recrystallization, it is sufficient to spread the

crystals on a piece of paper or watch glass and leave them for few minutes to dry.

8. Sample Labels: Store all samples in glass vials that have been correctly labeled.

The label should include at least the following information: identity of the

material, amount present (most conveniently given in grams), molar mass, the

experimentally determined melting or boiling point (include range), approximate

purity, the date and your name.

III. Melting point

The melting point of a solid is the temperature at which the first crystal just starts to melt

until the temperature at which the last crystal just disappears. Melting point of a

substance can provide useful information about its purity and identity. Most pure solids

typically melt at a sharp and defined, single temperature value and not over a range of

temperature values. A small amount of impurity can cause the melting point to spread out

over a range of several degrees. An impure substance will start melting at a temperature

that is lower than the melting point of the pure substance and may stop melting at a

temperature that is higher than for the pure substance. The greater the amount of

impurities in a substance, the wider the range over which it will melt. For this reason the

determination of the melting point is used to help in assessing purity of a given substance.

Pure substances have their own unique melting points. By determining the melting point

of a substance and then verifying it with the literature values, it is possible to identify an

unknown substance.

Melting Point Technique

To determine a melting point of a substance, a small amount of it is introduced into a

melting point capillary tube which is sealed at one end and open at the other. The solid is

placed in the tube by thrusting the open end down into the substance placed on a filter

paper. Some solid should be trapped in the tube. The solid is then tapped down to the

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sealed end by dropping the melting point capillary tube down a long narrow hollow tube

allowing the sealed end to bounce several times on a hard surface. In this way the

material will be vibrated down to the sealed end of the tube. Make sure that you have 2-3

mm of the material packed well. Once the capillary tube has the appropriate amount of

substance packed in it, the melting point can be determined using the melting point

apparatus. This is usually done in a two part operation. The melting point is initially

determined quickly to obtain an approximate melting point. A second melting point is

then determined with a second sample of the substance in a second capillary tube. During

the second melting point determination, heating is slowed considerably near the melting

point determined initially in order to give more accurate results.

IV. Thin layer chromatography (TLC)

TLC is a simple technique used to check the purity of non-volatile materials. A TLC plate

is made up of an adsorbent, which is a thin layer of silica, and a support which is

normally glass, plastic or aluminum support. Some TLC plates have a fluorescent

component added to them which will emit light when placed under a source of UV light.

This helps in viewing compounds which do not have color.

Chromatography works based on polarity where a solvent or solvent system is selected

for the particular compounds to be separated. In the ideal solvent system the compounds

of interest are soluble to different degrees. As the solvent moves up the TLC plate, by

capillary action, the compounds that are spotted on the plate are pulled up the plate with

it. The ability of compound to move up the plate depends on several factors:

1. Solubility: the more soluble a compound is in a solvent, the faster it will move up

the plate.

2. Attractions: the more the attraction between the compound and the silica, the

slower the movement of the compound is.

3. Size of the compound: the larger the compound is, the slower it will move up the

plate.

Once you have found a solvent or a solvent system that will separate the mixture of

compounds you are after, you can then identify them by running controls (known

compounds) along side the compound you are testing. You will be able to compare their

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individual movement up the plate. The most common way to compare them is by their

individual Rf values. The Rf value is calculated by dividing the distance the compound

moved up the plate by the distance the solvent moved up the plate. To accurately measure

the distance traveled by each compound and also by the solvent, you will need to mark a

line on your TLC plate, with pencil (do not use pen, Why?), and spot all of your controls

and samples on the line, in spots that are about the same size. After the TLC plate is

developed, quickly mark (with pencil) how far up the plate the solvent traveled. Measure

the distance the solvent traveled from your original line to the top of the wet area.

Similarly, for your controls and samples, circle the spots and measure from the line to the

center of the spot for the distance each compound traveled.

Rf = distance compound traveled

distance solvent traveled

If you use the same solvent, type of TLC plates, and method of spotting your compounds

on the plates, the Rf value for each known compound should remain the same from one

TLC run to the next. Ideally you should spot all of your controls and samples to be

analyzed on the same TLC plate to guarantee that they are run under the exact same

conditions.

Procedure:

1. Sample preparation: Place a small amount (1-2 mg) of your sample in a test tube

and then add a few drops of a solvent to dissolve it completely. Usually

dichloromethane is an effective solvent to dissolve a wide range of substrates.

2. TLC plate preparation: Using your ruler measure 0.5cm up the long side and

draw a line. On this line mark a spot (X).

3. Applying sample to plate: Put the thin end of the capillary spotter into the above

solution. The solution should rise into the capillary. Touch the capillary to the

spot (X) on the plate briefly. The compound will run out and form a small spot on

the plate. Repeat the process again. Ideally you want the spot (wet area) to be as

small as possible for the best results. Allow the spot to dry completely before

adding the “second coat”. You will need to spot each sample several times to get

enough compound on the plate. Use the UV lamp to make sure you have enough

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sample on the spot (X). CAUTION: do not look at the UV light to see if it is on!!

Hold it over your TLC plate!!

4. Running the TLC: To run your TLC set up a chromatography chamber in the

hood. Use a 100ml beaker, place 5ml of the eluting solvent in the beaker, then

carefully place your TLC plate in the beaker and cover the beaker with a filter

paper to keep the eluting solvent from evaporating too quickly. It will only take 1-

2 minutes for the solvent to rise up the plate. Mark the solvent front with pencil.

Then allow the plate to dry completely in the hood.

5. Analysis: Measure the distance the solvent traveled in mm, from the line to the

solvent front mark. Then using UV lamp carefully circle each spot on the plate,

record what color the spots are under the UV light and calculate the Rf values.

NOTES ON LABORATORY QUANTITATIVE TECHNIQUES

Cleaning volumetric glassware

There are three main types of volumetric glassware: pipettes, burettes, and volumetric

flasks. All glassware should be washed with soap solution and rinsed thrice with tap

water then thrice with de-ionized water. After using NaOH solution in a burette, pour 10

ml of 6M HCI into your burette, fill with tap water, run some of the solution through the

tip, then rinse thoroughly with tap, and then de-ionized water, 3 times each.

Reading the Meniscus

The meniscus is the curvature exhibited by a liquid confined

in a narrow tube. For accurate results one need to be

consistent in reading the meniscus. Common practice is to

read the bottom of the meniscus when using volumetric

glassware. It is often advantageous to highlight the meniscus

by placing a white paper behind the container (pipettes,

burettes, or volumetric flask). The eye of the observer should

be horizontal with the meniscus, noting the nearest calibration line that goes completely

around the container, the observer can ascertain if the eye is at the correct level.

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The pipette

Solutions are drawn into the pipette with rubber bulbs. Always keep the inside of

the bulb dry, it must not become contaminated with solutions. Standardized

technique to fill and drain a pipette is necessary to minimize error due to

irreproducible drainage. First, if not dry, the pipette is rinsed with two or three

small portions of the solution to be pipetted so the entire inner wall is wetted with

the solution before pipetting begins. The outside of the tip of the pipette should

be kept free of drops of solution, as there is the danger that drops might be

transferred to the receiver.

The bulb is then put on the pipette top as loosely as possible. The solution is

drawn to a point at least an inch above the calibration line, the rubber bulb is

withdrawn and the index finger (not the thumb) is placed over the opening. The

outside of the long delivery tube is wiped off with a tissue paper and the solution allowed

to drain down to the mark by varying the angle of the index finger. Both the finger and

the opening in the pipette should be dry for best results.

The solution is allowed to flow freely but without spattering into the receiver. Wait five

seconds after flow has ceased for the pipette to drain, and then touch the tip of the pipette

to the moist side of the receiver. A small residue will remain in the pipette. Do not blow

this out.

The burette

Burettes provide a means for accurately delivering a volume of liquid. They consist of a

calibrated tube and stopcock arrangement, which allows a flow from the tip to be

controlled. The burette is to be rinsed with the solution it is to contain. The stopcock plug

should be liquid-tight, so that it must not leak. Pay particular attention to air bubbles.

Teflon stopcock plugs are made of very chemically inert materials but a few rules for

care are in order:

1. Never use abrasive materials to clean either plug or barrel. This includes

brushes.

2. The Teflon washer (white) is placed adjacent to the end of the barrel so

minimal friction is created on turning the barrel.

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3. Teflon plugs can be easily scored around the bore if rotated when solid

particles are lodged between plug and barrel. Once scored, the plug may leak.

4. If Teflon plugs are used with liquids corrosive to glass, such as alkalis, rinse

stopcock thoroughly with water after use. Do not allow the liquid to

evaporate. The concentrated solution remaining will attack the glass surface;

and the eventual solids may also mar the Teflon surface if plug is then rotated.

The volumetric flask

The main use of the volumetric flask is to prepare solutions of a specified

concentration, either in preparation of standard solutions or in dilution of

samples of known volumes.

After transferring the solute, fill the flask about half-full and swirl to mix.

Add more solvent and mix again. Bring the liquid level almost to the mark,

allow time for drainage down the neck of the flask, and then using a dropper,

make the necessary additions of solvent. Firmly stopper the flask and invert

repeatedly to assure uniform mixing. Loosen the stopper to allow solution to

drain back into flask before removing stopper completely. The contents

must be homogeneous to achieve any valid results.

Preparing a standard solution often requires that a known weight of solid solute be

introduced into a volumetric flask. This can be accomplished in two manners. In the first,

insert a powder funnel into the neck of the flask and add the solid sample. After transfer,

the solid is washed off the funnel into the flask with de-ionized water, and the resulting

solution is diluted to the correct volume.

The second procedure involves dissolving the sample in a clean beaker or flask with

small portion of water and then transferring the solution into the volumetric flask. The

beaker is then rinsed three times with small amounts of water and each rinse is added to

the volumetric flask. The advantage of using this procedure is that solids that do not

easily dissolve in room-temperature water can be heated in the beaker.

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EXPERIMENT 1 SYNTHESIS OF ASPIRIN FROM SALICYLIC ACID

I. Objectives

The aims of this experiment are to:

1. Synthesize acetylsalicylic acid (Aspirin) from salicylic acid. 2. Ascertain its purity by melting point determination and thin layer chromatography

(TLC). 3. Recrystallize Aspirin from ethanol

II. Introduction Aspirin, or acetylsalicylic acid, was first synthesized in 1893 by Felix Hofmann. Aspirin is a derivative of salicylic acid that is used as analgesic (painkiller, reduces and prevents pain) and antipyretic (reduces or prevents fever) drug. It is probably the most commonly used over-the-counter drug. When ingested, acetylsalicylic acid remains intact in the acidic stomach, but in the basic medium of the upper intestinal tract, it hydrolyzes forming the salicylate and acetate ions according to scheme 1:

Scheme 1

The analgesic effect of aspirin is probably due to the salicylate ion. Salicylic acid has the same therapeutic effects as aspirin, but it causes more severe stomach upset since it is too acidic to be safely consumed by itself.

Aspirin can easily be synthesized from salicylic acid by reacting salicylic acid with acetic anhydride as shown in the following equation:

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Scheme 2 Though esters can be produced from the direct esterification of an alcohol and a carboxylic acid in the presence of an acid catalyst, typically sulfuric acid, the present method uses a derivative of acetic acid (acetic anhydride) to form more quickly an acetate ester with salicylic acid (the alcohol donor of the ester). Acetic anhydride is used because it is cheap and forms a by-product, acetic acid, which can easily be removed from the product. In this experiment you will synthesize aspirin using the above reaction (Scheme 2) and then determine the purity of the initial product by melting point and thin layer chromatography (TLC). You will then attempt to purify the crude product by recrystallization and again use the melting point and TLC to determine the purity of the final product. III. Materials and apparatus: Salicylic acid, acetic anhydride, concentrated sulfuric acid, ethanol, dropper, Erlenmeyer flask (125 mL), beakers (400 mL, 100 mL, 20 mL), graduated cylinders (10 mL, 25 mL), watch glass, stirring rod, vial to store Aspirin sample, ring, stand, clamp (to hold 125 mL Erlenmeyer flask), Buchner funnel, filter paper to fit Buchner funnel, vacuum filtration flask, rubber tubing for vacuum flask, ice, thermometer 150°C, melting point capillary tube, melting point apparatus, dropper, rubber gloves. SAFETY PRECAUTIONS Acetic anhydride: corrosive and lachrymator. It is destructive to tissue of the mucous membranes and upper respiratory tract, eyes and skin its vapor is irritating to the respiratory system. Avoid skin contact and inhalation of the vapors. In the event of skin contact, rinse well with cold water. If the vapors are inhaled, move to an area where fresh air is available. Use only in a ventilation hood. Sulfuric acid: corrosive and toxic. Avoid skin contact. In the event of skin contact, rinse well with cold water. Causes severe irritation or burns. Avoid contact with skin or breathing of vapors. Extreme care should be taken when handling this material. Use only in a ventilation hood. Salicylic acid: causes irritation to the skin, eye, and respiratory tract. Avoid contact with skin or inhalation of dust.

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IV. Procedure

1. Using an electronic pan balance, weigh out 3 g (to ± 0.01 g) of salicylic acid and transfer it into a clean and dry 125 mL Erlenmeyer flask.

2. Add 6.0 mL of acetic anhydride to the flask (Must be done in the fume hood, Caution: severe eye irritant, avoid skin and eye contact) and gently swirl the flask for a minute to wet the crystals and CAREFULLY add 6-7 drops of concentrated sulfuric acid. (Caution: very corrosive).

3. Place the Erlenmeyer flask in a hot water bath in the fume hood and let it heat for 10 minutes while swirling the flask occasionally. During this time period all of the salicylic acid should dissolve.

4. Remove the flask form the hot water bath and add 30 mL of deionized ice water to decompose any excess acetic anhydride. Place the flask in an ice bath and stirr occasionally using a glass rod to decompose residual acetic anhydride. Keep stirring until crystals of aspirin no longer form. If an oily liquid forms instead of a solid, use the end of a glass rod to scratch the bottom of the flask underneath the oil. This is usually sufficient to initiate crystallization. If no solid appears then reheat the flask in the hot water bath until the oil disappears and chill again.

5. Set up a vacuum filtration apparatus. Wet the filter paper in the Buchner funnel with 1-2 mL of distilled water. Decant the liquid onto the filter paper.

6. Add 20 mL of ice-cold water to the flask, swirl, and chill again. Repeat until the transfer of the crystals to the vacuum filter is complete; maintain the vacuum for 5 minutes to dry the crystals as much as possible. If aspirin forms in the filtrate, transfer the filtrate and aspirin to a beaker, chill in an ice bath, and vacuum filter as before, using a new piece of filter paper. Dump filtrate in the sink

7. Weigh the crude aspirin and calculate the percent yield of the reaction (see below). Save a small portion of the crude product (~0.1g) for a melting point determination and TLC then recrystallize the rest.

8. The major impurity in aspirin is salicylic acid. It can be removed by recrystallization. Transfer the crystals from the filter paper to a 100 mL beaker. Dissolve crystals in a minimum volume of ethanol (≤ 13 mL). Warm the mixture in a 60 °C water bath (use a hot plate). Pour 30 mL of hot water (about 60 °C) into the solution. If a solid forms, continue warming until the solid dissolves.

9. Cover the beaker with a watchglass, remove it from the heat, and set it aside to cool slowly to room temperature then set the beaker in an ice bath. Determine the mass (to ± 0.01 g) of a piece of filter paper. Collect the solid aspirin using vacuum filtration on this filter paper. Wash the crystals with two 15 mL portions of ice water. Place the filter paper and aspirin sample on a watchglass and allow them to air-dry. The time for air-drying the sample may require that it be left until the next laboratory period. Dump filtrate in the sink.

10. After the crystals have dried, transfer the dry aspirin crystals to a pre-weighed sample container or vial. The aspirin sample should be labeled with your name.

Steps 11-14 will be done in the next lab!

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11. Determine the mass of the aspirin crystals by weighing and correcting for the fact that you didn't use the entire sample, calculate a percent yield for the entire reaction.

12. Determine the melting point of your recrystallized and unrecrystallized aspirin samples. Fill a capillary melting point tube to a depth of 0.2 cm with the aspirin. Place the capillary tube in the melting point apparatus. Determine its melting point. (Your instructor will demonstrate the use of this apparatus).

13. Run a TLC on your recrystallized and unrecrystallized aspirin. Circle the spots and stick the TLC plate in your report.

14. The aspirin sample should be labeled with your name, the mass of the aspirin, the percent yield, and its melting points and submitted to the lab technician.

V. Calculations

Determine the % yield of your product.

Based on the balanced equation for the reaction, you should obtain 1 mole of aspirin for every mole of salicylic acid. The percent yield (%yield) is determined by

%Yield = (actual mass of product) X 100 (theoretical mass of product)

Sample Data and calculations

Mass of salicylic acid 12.05 g

Volume of acetic anhydride 13.2 mL

Mass of "aspirin" 18.33 g

1. Calculate the moles of salicylic acid.

Salicylic acid (SA) is C7H6O3 with a molar mass of 138.12 g/mol.

2. Calculate the moles of acetic anhydride given that its density is 1.08g/mL.

Acetic anhydride (AA) is C4H6O3 with a molar mass of 102.09 g/mol.

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3. Calculate the theoretical yield of aspirin

Acetyl salicylic acid (ASA) is C9H8O4 with a molar mass of 180.16 g/mol. Use the balanced chemical equation to decide the grams of aspirin that could be made from each starting material. The theoretical yield is the smallest of these amounts. Once the smallest amount is made, one of the chemicals is used up, and no more of the product can be produced.

C7H6O3 + C4H6O3→ C9H8O4 + C2H4O2

SA + AA → ASA + acetic acid

The theoretical yield of aspirin is 15.72 g.

4. Calculate the percent yield of aspirin.

The percent yield of any reaction is never greater than 100%. The percent yield may appear to be greater because of impurities in the final product.

VII. Report Sheet PREPARATION OF ASPIRIN (A) DATA Mass of crude aspirin: Mass of purified aspirin: Theoretical yield of aspirin (show calculation): Percent yield of aspirin (show calculation): Melting point of aspirin: Comment on the purity of your aspirin sample based on melting point and TLC results. Write a mechanism for the reaction of formation of Aspirin. (B) Discussion: In your discussion section address the following questions:

25

1. Which reactant, acetic anhydride or Salicylic acid is the limiting reagent? Why? 2. What is the most likely impurity in the final aspirin product? How is its presence

detected? 3. Was there a difference in the melting points of the crude & pure aspirin? What

can you conclude? 4. Was there a difference in the TLCs of the crude & pure aspirin? What can you

conclude? 5. What are the possible sources of error in this experiment? 6. Write, in your own words, a one-page summary of the experiment. State what you

think the purpose of the experiment was and what you think you accomplished including a brief summary of the final results.

7. Compare the form and texture of the crude and recrystallized materials. Give comments.

Reference: www.chymist.com/aspirin.pdf

EXPERIMENT 2 PART A: DETERMINATION OF ASPIRIN CONTENT USING VISIBLE

SPECTROSCOPY Materials: Acetylsalicylic Acid: Irritant 1 M NaOH: Corrosive 0.02 M FeCl3: Corrosive/Hygroscopic Commercial Aspirin 1. Purpose In this laboratory experiment, you will assess the purity of a sample of aspirin and become acquainted with the concept of visible spectroscopy. Introduction Spectroscopy is the study of the interaction of electromagnetic radiation (EMR), 'light', with matter. EMR is a set of oscillating electric and magnetic fields which is characterized by a frequency (ν) and a wavelength (λ). They are related through the speed of light, c, as shown in the following equation: λ*ν = c In the visible region of the spectrum the color of the sample indicates which regions of the electromagnetic spectrum are interacting with the sample, and which regions are passing through the sample with little or no interaction. The former are said to be 'absorbed' by the sample while the latter are said to be 'transmitted'. The color and amount of light absorbed can be detected and measured using a spectrophotometer. The major components of a spectrophotometer operating in the visible portion of the spectrum are a white-light source; a monochromator, a device that separates the white-light into its

26

component colors; a detector which measures the amount of light passing through the sample; and various associated mirrors, lenses, and filters to change the light path within the instrument. The instrument is constructed in such a way that the output signal, called the absorbance, that it provides is directly proportional to the concentration of the light absorbing species. This relationship is known as the Beer-Lambert Law and is expressed mathematically as (1)

The constant of proportionality, ε, is called the molar absorptivity, and must be experimentally determined. The aspirin which you previously synthesized is probably not pure, despite your best efforts. The most likely impurities are salicylic acid (unreacted starting material) or acetic acid. Even commercially prepared aspirin tablets are not 100 percent acetylsalicylic acid. Most aspirin tablets contain a small amount of binder which helps prevent the tablets from crumbling. Moreover, moisture can hydrolyze acetylsalicylic acid. Thus, aspirin which is not kept dry can decompose. Acetic acid is the hydrolysis product formed by the reaction of water with acetylsalicylic acid: You may have noticed the smell of vinegar (acetic acid) when opening an old bottle of aspirin, or a bottle which has not been properly sealed. Aspirin, acetylsalicylic acid, is a white crystalline powder that forms a colorless solution when dissolved in water, i.e. it does not absorb visible light. It is also a weak acid, which can react with a strong base to form the salicylate ion as shown in equation (2):

O

C OH

C CH3

O

O

+ 3 OH-(aq)

O-

C O-

O

+ H3C CO

O- + 2 H2O (l)

(aq) (aq)(aq)

(equation 2)

The salicylate ion will react with Fe (III) in the presence of acid forming a reddish-purple complex (Fe (III)-SA) which absorbs light at a wavelength of 530 nm (equation 3).

O-

C O-

O

+ [Fe(H2O)6]+3 (aq)

O

C OO(aq)

(aq)

Fe(H2O)4

+

+ 2 H2O (l)H+

(euation 3)

The concentration of the complex can easily be determined by measuring the amount of light absorbed by the complex and rearranging equation (1).

[analyte]*=Absorbance ε (1)

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The concentration of the salicylate ion and hence the amount of acetylsalicylic acid are determined through the use of the appropriate stoichiometric factors (equations (2) and (3)). EXPERIMENTAL In order to use equation (4) we must first determine ε. This is done by reacting a known amount of aspirin (as the salicylate ion) with an excess of Fe (III) in order to drive reaction (3) to completion. Since the salicylate ion is the limiting-reagent, the amount of the complex formed can be readily calculated and the value of ε determined by measuring the absorbance and applying equation (4). A more common practice is to prepare several standards, solutions of different but known concentrations, and to plot the measured absorbance of each of the standards against the concentration of the Fe (III)-SA complex in each. The slope of the resulting line is equal to ε.

1. Accurately weigh out approximately 0.400 g (400 mg) of reagent grade acetylsalicylic acid and quantitatively transfer it to a 125 mL erlenmeyer flask. Record the weight of the aspirin used in Table I of your lab book.

2. Go to the hood and using a graduated cylinder, very carefully add approximately 10 mL of 1 M NaOH.

3. Warm the solution to near boiling on a hot plate taking care not to spatter the corrosive liquid outside the flask.

4. Remove the flask from the hot plate and let it cool for several minutes then cool it to room temperature by very carefully swirling it under a stream of cold tap water.

5. Quantitatively transfer the cooled solution to a 250 mL volumetric flask. Using a distilled water bottle, rinse the erlenmeyer flask several times, transferring the rinsings to the 250 mL volumetric flask.

6. Fill the volumetric flask to the mark with distilled water, cap it tightly with parafilm and mix thoroughly. This will be your 'stock' solution from which all of your standards will be prepared.

7. Obtain five 100 mL volumetric flasks and label them STD 1 through STD 4. 8. Using a calibrated glass pipette, transfer 1.00 mL of the 'stock' solution to the

flask marked STD 1, 2.00 mL of the 'stock' solution to flask STD 2, 3.00 mL of 'stock to STD 3, 4.00 mL to flask STD 4.

9. Carefully fill each of the flasks to the mark with the provided 0.02 M Fe (III) solution. You should see the reddish-purple color of the Fe (III)-SA form immediately.

10. Cap each flask tightly then thoroughly mix each of them. Measure the absorbance of each of the solutions at 530 nm using water as a blank and record these values in Table I.

ε

Absorbance=SA]III[Fe −)( (4)

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ASPIRIN ANALYSIS

1. Accurately weigh out approximately 325 mg of your synthesized aspirin into a 125 mL Erlenmeyer flask and treat with 10 mL of 1 M NaOH as before.

2. Quantitatively transfer the solution to a 250 mL volumetric flask and dilute to the mark with distilled water. Cap and mix thoroughly.

3. Using a pipette, transfer 5.0 mL of this solution to a 100 mL volumetric flask and fill to the mark with the 0.02 M Fe (III) solution. Cap tightly and mix thoroughly.

Repeat the above procedure using a commercial aspirin product.

1. Accurately weigh one tablet and transfer it to a 125 mL Erlenmeyer flask. 2. Add the 1 M NaOH as before and heat and swirl until the table completely

disintegrates (it probably will not completely dissolve). 3. Quantitatively transfer the mixture to a 250 mL volumetric flask and fill to the

mark with water. Cap tightly then mix thoroughly. Let any solids settle to the bottom.

4. Using a pipette, transfer 5.0 mL of the solution to a 100 mL volumetric flask and fill to the mark with the 0.02 M Fe (III) solution.

5. Measure the absorbance of each of the two solutions at 530 nm and record these values in Table 2.

ANALYSIS The concentrations of the solutions in Table 1 can be calculated using dilution factors

Plot the measured absorbance (y-axis) versus the calculated concentration of Fe (III)-SA (X-axis). Draw the best straight line through these points and determine the slope, which is ε. The mass of aspirin in each of the two samples can be calculated using the equation below.

Compare your calculated amounts of aspirin with the expected or known amounts using the following equations: % difference (synthetic) = |exp value - mass sample| x 100 mass of sample % difference (commercial) = |exp value - stated value| x 100 stated value

100.0mL

stock(mL)ofvolume*250.0mL

)reagent(mgaspirinmass=)mLmgSA](-[Fe(III) ⋅⋅⋅⋅

(250.0mL)*)5.0mL

100.0mL(*)(530nm)

(530nm)Absorbance(=sample

aspirinmgε

⋅

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In the above equations, 'exp values' are the values that you obtained in this experiment and the 'stated value' is the value on the aspirin bottle. TABLE 1. DETERMINATION OF ε

weight of aspirin = _____________________ g

Flask

Volume of Salicylate Stock (mL)

Total Volume (mL)

Calculated Concentration Fe (III)-SA mg/mL

Measured Absorbance λ =

STD 1

STD 2

STD 3

STD 4

CALCULATED VALUE FOR ε TABLE 2. ASPIRIN SAMPLE

SAMPLE MASS ABSORBANCE CALCULATED MASS

PERCENT DIFFERENCE

Reference: Pinnell, Robert P.; Motz, Leonard P. “A Colorimetric Determination of Aspirin in Commercial Preparations”, Modular Laboratory Program in Chemistry Anal-126, Neidig, H. A., Ed. Willard Grant Press, Boston, MA, 1973, 1.

PART B. 1. Continue steps 11-14 from EXPERIMENT 1 2. Prepare your NMR sample according to page 33, NMR procedure steps 1-4

and label the NMR tube from the top. Hand it to the lab technician.

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EXPERIMENT 3 CHARACTERIZATION OF ASPIRIN SYNTHESIS PRODUCT USING IR AND

NMR SPECTROSCOPY Materials: Salicylic Acid: Irritant Acetylsalicylic Acid: Irritant CDCl3: Highly Toxic/Cancer Agent Ethyl Acetate: Flammable Liquid/Irritant I. Objectives

1. Develop a general understanding of using IR and NMR spectroscopy to characterize organic compounds

2. Determine the purity of aspirin synthesis product using IR and NMR spectroscopy 3. Estimate the % of aspirin in your product from its NMR spectra

II. Introduction Spectroscopy is the study of the interaction of electromagnetic radiation (EMR) or light, with matter. EMR is a set of oscillating electric and magnetic fields which are characterized by a frequency (ν) and a wavelength (λ) that are related to the speed of light (c) as follows: (c = λ*ν). The following experiment that you will perform is designed to introduce you to two different regions of the EMR spectrum which cannot be observed by the human eye. Since energy is related to the frequency of EM-radiation, each region of the electromagnetic spectrum interacts with molecules very differently based on the energy, E = hν. For example, microwaves cause polar molecules to rotate, whereas, infrared radiation (IR) causes molecules to vibrate. In this experiment, we will use IR and radio frequencies to determine the purity of the aspirin you made. IR Spectroscopy: A molecule can absorb radiation when some parts of the molecules vibrate at the same frequency of the IR light. After absorbing the IR light, the molecules vibration increases. It has been established that functional groups absorb about the same frequency of radiation even if these functional groups are found in different molecules. This allows us to detect the functional groups that are present in different molecules based on their absorbance. It has also been shown that each molecule has a unique IR pattern; therefore, molecules can be identified from their IR spectra. For example, look at the structures of salicylic acid & acetylsalicylic acid shown below.

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COH

CO

CCH3

O

Salicylic Acid Acetylsalicylic Acid(Aspirin)

OHO OHO

Functional Groups

CO

OH (Carboxyl)

OH (Hydroxyl)

CO

OH (Carboxyl)

O CO

CH3 (Ester) Both contain the same functional group, carboxyl, but both also contain a different functional group. From the vibrations of these groups, the compounds can be distinguished. A list of the vibrations is provided in the table below. Functional Group Vibration Frequency (cm-1)

CO

OH (Carboxyl)

CO

1700-1650

OH (Hydroxyl)

OH

3500-2700

O CO

CH3 (Ester)

CO

1750-1700

EXPERIMENTAL

(A) Infrared measurement:

1. Obtain from your Laboratory officer KBr, spatula and mortar and pestal. 2. The instrument is designed to look for differences between spectra; therefore, you

must take a blank spectrum of the pure KBr. This need to be done once at the beginning of the lab.

3. Sample Preparation: using a clean, dry mortar and pestal grind up your product (~ 100mg) and KBr (~ 20 mg) into a fine powder

4. Obtain IR Spectra: with the help of the lab assistant take a spectra of your sample 5. The displayed spectrum is the IR pattern for synthesized aspirin. Save the

32

spectrum. 6. Set the parameters for your spectral graph and print it out 7. Make a copy of your spectra so both you and your lab partner will have a copy for

your notebooks 8. Using the reference spectra Analyze your Aspirin Synthesis Product

IR discussion:

1. Can you identify peaks in your Aspirin Synthesis Product spectra which are found only in acetylsalicylic acid?

2. Can you identify peaks in your Aspirin Synthesis Product spectra which are found only in salicylic acid?

1. From this analysis using Infrared Spectroscopy, estimate the purity of Aspirin Synthesis Product?

2. Are there any major differences between the IR of salicylic acid and your aspirin? 3. Which peak(s) did you use to determine if you had converted the salicylic acid to

aspirin? 4. How does your aspirin spectrum compare to that of the pure aspirin? 5. Are there peaks present which are not found in either acetylsalicylic acid or

salicylic acid? What could these peaks represent?

(B) NMR spectroscopy

Nuclear Magnetic Resonance (NMR) is a way of characterizing an atom based on the interaction between its nucleus and the magnetic field. It is used in organic chemistry to identify known organic chemicals and to determine the structure of novel compounds. NMR spectroscopy is based on the measurement of radio frequency absorption. In contrast to other absorptions, the nuclei of atoms absorb the energy, hence the name nuclear. However in order for the nuclei to absorb the radio waves, the sample must be placed in a strong magnetic filed. The advantage of this technique is that structures of compounds can be determined. The most common use of this technique is proton (H1) NMR. Proton NMR allows one to determine to which atoms the H atom is bonded and therefore, the structure of the molecules can be determined. Look at salicylic acid and aspirin and notice the different number of hydrogens. There are many similarities in the

33

COH

CO

CCH3

O

Salicylic Acid Acetylsalicylic Acid(Aspirin)

OHO OHO

CO

OH 1-Hydrogen

OH 1-Hydrogen

CO

OH 1-Hydrogen

O CO

CH3 3-Hydrogens

4-Hydrogens 4-Hydrogens

compounds, but the differences can be seen in the NMR spectra. Each hydrogen will absorb a different radio frequency depending on what type of atom it is bonded to. As a standard, the hydrogens are compared to a reference (TMS) that is set at 0 ppm. The table below gives an approximate range in which the hydrogen will absorb. Hydrogens Reference to TMS

CO

OH 1-Hydrogen 9.5-10 ppm

OH 1-Hydrogen

11-12 ppm

4-Hydrogen

6.5-8 ppm

O CO

CH3 3-Hydrogen 2.5-2.8 ppm

NMR Procedure:

1. Obtain one NMR tube from your lab assistant. 2. Place a small amount (~ 20 mg) of Aspirin on the end of a spatula and transfer it

carefully to the NMR tube. 3. Add CDCl3 drop wise into the NMR tube until the height is approximately 1 cm. 4. Cap the NMR tube and invert a few times until all the Aspirin is dissolved.

34

5. Place the tube in the NMR and obtain a spectrum. The instructor or lab assistant will help you run your sample and set up the spectra for printing.

6. Compare this spectra with an NMR of pure aspirin. Using your spectra, compare it to the reference spectra provided to determine how pure it is and make a rough estimate of the % aspirin in your product. Keep in mind that the intensity of the pattern is tied to its concentration in your sample. To calculate the percent of aspirin in your sample use the following formula: [ peak height aspirin ÷ (peak height aspirin + peak height salicylic acid) ] x 100 = % aspirin

NMR discussion:

1. Can you identify any of the patterns of the reference compounds in your sample spectrum? If so, which of these compounds can you identify in your sample?

2. On your spectrum, identify all peaks associated with the correct product, aspirin (acetylsalicylic acid).

3. On your spectrum, identify all peaks associated with the salicylic acid or acetic anhydride starting materials, if any.

4. On your spectrum, identify all peaks associated with the acetic acid byproduct, if any.

5. Based on the intensities of these patterns, roughly estimate the % of each of the above components in your product.

6. From this analysis how pure is your Aspirin Synthesis Product? What is the percentage of aspirin in your product?

7. Are there any peaks in your spectrum that you cannot identify? If so, what do these peaks represent, and how could you get rid of them?

8. Which peak(s) did you use to determine if you had converted the SA to aspirin? 9. Why would you find acetic acid, ethanol or water in your product?

Reference: Houston Byrd and Stephen E. O’Donnell, Journal of Chemical Education, 2003, Vol 80 (2), 174-176.