Organic Lab Manual

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LABORATORY MANUAL

FOR

CHEMICAL EXPERIMENTATION I

(COURSE NO. CHEM F242)

DEPARTMENT OF CHEMISTRY

BIRLA INSTITUTE OF TECHNOLOGY AND SCIENCE, PILANIK. K. BIRLA

GOA CAMPUS, GOA


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Team of Chemical Experimentation I, 2014-2015

Department of chemistry,

BITS, Pilani- K K Birla Goa Campus

All rights reserved.


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Table of Contents

INDEX Page no.

PART A Introduction 5

General Laboratory Safety 7

Preparation of notebook 9

Common laboratory techniques 12

PART B

Experiment 1: Elemental analysis of organic compounds. 38

Experiment 2: Functional group analysis of organic compounds 49

Experiment 3: Qualitative analysis of unknown compounds including derivatization. 60

Experiment 4: Chemical separation and identification of mixture of two compounds 80

Experiment 5: Single step synthesis:

a) reduction of benzophenone by NaBH4

83

b) reduction of camphor by NaBH4 85

c) paracetamol to another analgesic drug phenacitin 88

d) synthesis of 3-Methylpyrazol-5-one 90

e) synthesis of Methyl Orange 91

f) synthesis of benzylideneacetophenone 93

g) Preparation of Dinitrobenzene 94

h) Diels-Alder reaction between furan and maleic acid 98

i) Synthesis of dibenzalpropanone (Dibenzylideneacetone) 100

j)

Preparation of 1, 1-bis-2-naphthol 102

k)

Synthesis of dihydropyrimidinone 104

l)

Synthesis of biodiesel: Glycerol 106
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Experiment 6: Multistep syntheses:

a)

IBX from anthranilic acid

108

b)

Benzaldehyde to Dilantin (i.e. Phenytoin) 110

c)

Benzaldehyde to benzilic acid 115

Experiment 7: Extraction of natural products: extraction of caffeine from tea leave 117

Experiment 8: Purification techniques: Separation of known mixture by

chromatography, purification of one of compound from single

step/multistep synthesis, resolution of enantiomers

120

Experiment 9: Quantitative analysis: Estimation of Glucose 130

Experiment 10: Operation of software related to organic chemistry: chemdraw, chem3D

etc.

135

Appendix 136


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PART A

INTRODUCTION

Welcome to the Organic Chemistry lab. You are about to embark on a journey that will be both

challenging and interesting. The principle objective of this course is to give students hands-on

experience on various basic and advance laboratory methods of organic chemistry. This is an

attempt to integrate students theoretical knowledge and concept on organic chemistry to

practical experience. Here, you will be introduced to techniques chemists use to carry out in

day-to-day operations all over the world from the academy to industry. In this course you will

be introduced to chemical methods as well as instrumentation used by the organic chemist. The

sequence of experiments in this Laboratory Manual is designed to follow the lecture

curriculum. Prior to each lab period, you will need to spend some time reading the Laboratory

Manual. This reading will provide background information and an outline of the procedures to

be performed. Precision is key as many of the outcomes of the reaction will depend on how

well you measure, stir, heat, observe, etc. As such, in this course a large part of your grade will

depend on your technique (how well you carry out the reaction). If you are careful and come to

lab prepared, you will do just fine.

We hope you find this laboratory manual helpful in your study of chemistry.


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ACKNOWLEDGMENTS

This manual is the culmination of the efforts of many individuals. Many faculty members of the

Department of Chemistry have provided inputs in preparation and successful implementation

of this laboratory course. Particularly, the efforts of the core team members of Chemical

Experimentation I (Drs. Mainak Banerjee (IC), Amrita Chatterjee, Tincy Lis Thomas and

Subhadeep Banerjee) are sincerely acknowledged.

The experiments described in this laboratory manual are mainly variations of similar

experiments that may be found described in the laboratory manuals of other universities or in

commercially produced lab texts. Most of the experiments have been modified and rewritten,

keeping the particular needs of BITS, PilaniK. K. Birla Goa Campus students in mind.

Ph.D. students of the departments and laboratory staffs have had a role in helping to develop

these experiments and, in particular, helping to ensure that the experiments are tailored to our

laboratories here at BITS, PilaniK. K. Birla Goa Campus. The team of CHEM F242 thanks all of

them.

This is the 2nd

edition of the earlier version of 2012-2013. In this version, many new Green

experiments have been added to enrich the content and to give a hands-on exposure of one of

the prime needs of our society, what will call the Sustainable Chemistry.

The following online resources are gracefully acknowledged.

http://www.umsl.edu/~orglab/

http://www.bluffton.edu/~bergerd/classes/CEM221/Handouts/LabManual.pdf

http://ocw.mit.edu/courses/chemistry/5-301-chemistry-laboratory-techniques-january-iap-

2012/labs/MIT5_301IAP12_comp_manual.pdf

http://www.aircleansystems.org/ductlessfumehoods.htm

http://www.dst.gov.in/green-chem.pdf

Mainak Banerjee

Instructor-In-Charge

Semester II 2014-2015

Department of Chemistry
http://www.umsl.edu/~orglab/http://www.bluffton.edu/~bergerd/classes/CEM221/Handouts/LabManual.pdfhttp://ocw.mit.edu/courses/chemistry/5-301-chemistry-laboratory-techniques-january-iap-2012/labs/MIT5_301IAP12_comp_manual.pdfhttp://ocw.mit.edu/courses/chemistry/5-301-chemistry-laboratory-techniques-january-iap-2012/labs/MIT5_301IAP12_comp_manual.pdfhttp://www.aircleansystems.org/ductlessfumehoods.htmhttp://www.aircleansystems.org/ductlessfumehoods.htmhttp://www.aircleansystems.org/ductlessfumehoods.htmhttp://ocw.mit.edu/courses/chemistry/5-301-chemistry-laboratory-techniques-january-iap-2012/labs/MIT5_301IAP12_comp_manual.pdfhttp://ocw.mit.edu/courses/chemistry/5-301-chemistry-laboratory-techniques-january-iap-2012/labs/MIT5_301IAP12_comp_manual.pdfhttp://www.bluffton.edu/~bergerd/classes/CEM221/Handouts/LabManual.pdfhttp://www.umsl.edu/~orglab/


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LABORATORY SAFETY

Safety First, Last, and Always

The organic chemistry laboratory is potentially one of the most dangerous of undergraduate

laboratories. That is why you must have a set of safety guidelines. Do consult/inform your

instructor when you have any doubts regarding safety.

Disobeying safety rules will be penalized by deduction of performance marks.

Attire:

1. Eye injuries are extremely serious and can be mitigated or eliminated if you keep your

safety goggles on at all times. The wearing of contact lenses in the laboratory is strongly

discouraged.

2. Lab coats and shoes must be worn in the Lab. Avoid very loose fitting clothes. Long hair

must be tied back.

Safety Equipment

A set of safety rules is written below. Careful observation of these rules will help to prevent

accidents in the laboratory. However, from time to time accidents can occur. Therefore, safety

equipment is installed for his eventuality in the laboratory. Safety equipment should include:

An eye wash

A safety shower

Fire extinguishers

First-aid kit.

a) Eye Wash

The eyewash is designed to flush irritating chemicals from the eyes. It should be capable of

providing a stream of water for at least 15 minutes. In the event of an eye accident, you


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should proceed to the eyewash at once and wash the eye for at least 15 minutes. During

this process, the eye should be kept open. The eyes are the most vulnerable part of the

body. In the event of any eye injury notify the instructor at once. All eye injuries should be

immediately examined by a health professional.

Never use the eyewash for anything other than its intended purpose.

b)

Safety Shower

The safety shower is designed for two purposes, namely, to extinguish clothing fires and to

provide a whole body wash if a large chemical spill occurs.

i. Clothing Fires: If your clothing catches fire, perhaps the best rule is to fall and roll. Never

run to a shower with your clothes on fire, it will only fan the flames. Use the shower

afterwards to squelch any residual embers.

ii. Large Chemical Spills: Large chemical spills on clothing or exposed parts of the body

should be removed at once using the deluge shower. Contaminate clothing should be

removed, and the affected body areas should be thoroughly washed to remove any

chemical traces. Do not reuse contaminated clothing until it has been completely washed!

Serious and avoidable injuries have resulted from wearing contaminated clothing.

c)

Fire Extinguishers

In the laboratory, you will sometimes work with flammable materials. Several fire

extinguishers should be placed in the laboratory. Learn their location. In case of any

accidental fire you should use it or immediately inform your instructor.

d) First Aid Kits

First aid kits are used for minor injuries. Report all cuts and bums to the instructor, and athis/her discretion, visit the school physician for further treatment.


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Handling of Chemicals and Equipment:

1. Consider all chemicals to be hazardous. Know what chemicals you are using and their

properties.

2.

Avoid contact of chemicals with your skin or eyes. In case such contact does occur, flush

immediately with profuse amounts of water, and inform the instructor immediately.

3. Do not bring flammable regents near open flames.

4. Never point a test tube are heating towards yourself or a neighbor, or vertically

upwards.

5. Always pour acids into water, and not the other way around.

6. Never taste or directly smell chemicals. To detect odor, by means of your cupped hand,

waft a small sample of vapour towards your nose.

7. Dispose off chemical waste properly as directed by the instructor.

8. Follow directions carefully while using instruments.

9. Do not leave burners unattended. Turn them off when you leave. If the burner goes off,

turn off the supply valve immediately. Open again only while relighting the burner.

Never use paper torches for lighting burner.

10.Beware, hot glass looks just the same as cold glass.

Conduct:

1. No food or beverages are allowed in the laboratory.

2. No unauthorized experiments are to be performed.

3. Keep your work area clean. Put paper trash and broken glass, if any, in the dustbins.

Clean your work area before you leave.

4.

Avoid spills. If you do spill something, clean up the area immediately taking adequate

precautions. Inform the instructor.


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5. All instruments are very sensitive. Handle them with care. Keep the area around

instruments clean and free of trash paper.

6. If, during the course of a laboratory, you are injured, no matter how minor, you should

bring this to the attention of the instructor or teaching assistant.

7. Always wash your hands thoroughly before you leave the laboratory.

ACCIDENTS WILL NOT HAPPEN

That's an attitude you should hold while working in the laboratory.

ADVANCED PREPARATION AND THE NOTEBOOK

Before coming to the laboratory you will find that some advanced preparation can save

significant time and effort in the laboratory. The advanced preparation will vary depending on

the nature of the experiment, but some general rules will be applicable to all experiments. Read

the experiment first and identify all of the reagents and substances that will be used and

prepared during the course of the experiment. Next, collect the physical properties of these

materials before coming to the laboratory. Physical properties that will be useful include

melting point and boiling points (if available) of all the reagents and expected products, their

densities (if liquid) and some general solubility properties. Molecular weights should be

calculated for all the reagents. This information should be recorded in your notebook which will

be described below.

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"Aldrich Catalog Handbook of Fine Chemicals", Aldrich Chemical Co.

"Merck Index", Windholz, M. editor, Rahway N.J.

The laboratory notebook is the permanent record documenting what you did in the

laboratory during a specific period of time. It should contain sufficient details and

documentation that an individual with similar training could repeat your work months or years

later. The specific requirements for entries will vary from experiment to experiment and from

instructor to instructor. The following is intended to serve as a guideline of what to include in

the notebook. A substantial part of the grade earned in this course will be derived from the

contents of your notebook.

i)

Use a bound notebook. Notebooks with a spiral binding are not acceptable nor

are any that allow for insertion or removal of pages.

ii) Write in ink. It is expected that you will make mistakes. Just draw a line through the

material you wish to delete. Neatness is always appreciated.

iii) Leave the first few pages of the notebook blank for a table of contents so that each

experiment can be readily located. Pages should be numbered consecutively.

iv) Some notebook preparation should be completed before you come to the

laboratory.

What is required will vary from instructor to instructor. However, the following will apply to all.

1. Title: Each reaction must contain an appropriate title. The title for each experiment can be

found in this manual.

2. Equation: This section only applies to those experiments where you are carrying out a

reaction. Not every experiment will have an associated reaction.


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3. Table of Reagent: This table will contain all of the amounts of reagents used in the reaction

as well as their physical properties (molecular weight, melting point (solids), boiling point

(liquids), density (liquids), etc.).

4. Experimental Procedure: This is where you will write out the procedure so you can follow it.

That means you need to include appropriate details without including everything contained in

the manual. Do not simply copy the procedure from the manual. Again, you should follow

closely the style of the example.

5. Data/Observations: You should record all pertinent observations that you note for each step

of a reaction. If you add a reagent and the color changes, you should note the change of colors.

In addition to any observable changes during the reaction, you will need to record any

calculated values (mass, percent yield/recovery, retention factor (Rf), melting point, boiling

point, etc.) Any calculations that you carry out need to be written in the notebook.

6. Signature/Date: Each experiment must have your signature and the date it was completed.

This is standard procedure and authenticates your work.

Notebook Calculations

The number of calculations you will need to perform in the organic laboratory is quite small but

those that you will be required to perform are very important. The calculations are similar to

the calculations you performed in introductory chemistry and be sure to review them if you

have forgotten the details and theory behind them. What follows is a brief summary of some

important definitions.


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The calculations associated with conversion of the starting materials to product are

based on the assumption that the reaction will follow simple ideal stoichiometry. For example,

in the preparation of aspirin from salicylic acid and acetic anhydride, in calculating the

theoretical and actual yields, it is assumed that all of the starting material is converted to

product, even though some of the starting material actually forms a polymer as a consequence

of the reaction conditions and catalyst that is used.

The first step in calculating yields is to determine the limiting reagent. The limiting

reagent in a reaction that involves two or more reactants is simply the reagent which is present

in lowest molar amount based on the stoichiometry of the reaction. This reagent will be

consumed first and will limit any additional conversion to product. In the reaction of salicylic

acid with acetic anhydride, the latter reagent is used in excess. There are several reasons for

doing so. An important consideration whenever any reagent is used in excess is to determine

how this reagent will be separated from the product at the end of the reaction. In this case, the

acetic anhydride which is not water soluble can be allowed to stand in water a while during

which time it will slowly hydrolyze to acetic acid. The acetic acid which is water soluble can be

separated from the aspirin by filtration.

The salicylic acid will be the limiting reagent. As far as the calculation of the theoretical

yield is concerned, we assume that every mole of salicylic acid will be converted to aspirin or

acetylsalicylic acid. According to the stoichiometry of the reaction, one mole of salicylic acid will

be converted to one mole of aspirin. Therefore, from the number of moles of salicylic acid we

used, we evaluate the maximum amount of aspirin that can be obtained. Multiplying the

number of moles of aspirin by its molecular weight results in the theoretical yield which is

usually reported in grams. In all cases, the calculation is performed similarly, regardless of how

many products are formed. However, we can now evaluate the actual yield by determining

how much aspirin we have actually isolated, experimentally. The % yield is simply the ratio of

the actual yield divided by the theoretical yield times 100.


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COMMON LABORATORY TECHNIQUES

Crystallizationis used to purify a solid. The process requires a suitable solvent. A suitable solvent is one

which readily dissolves the solid (solute) when the solvent is hot but not when it is cold. The best

solvents exhibit a large difference in solubility over a reasonable range of temperatures. (eg, Water can

be a crystallization solvent between 0-100oC).

Characteristics of a solvent:

a. chosen for solubilizing power-- solubility usually increases with increasing temperature.

b. polarity is important--like dissolves like; polar compounds are more soluble in polar solvents;

nonpolar compounds in nonpolar solvents

c. almost all solvents are COMBUSTABLE--avoid flames

d. mixed solvents (eg; 1:1 water/methanol) provide a huge range of possible solvents but they must be

soluble in one another

Use solvent to get solids into solution but to get them out of solution:

a. lower the temperature--solute will be less soluble

b. concentrate the solution by removing solvent with a hot plate, heating mantle (flasks), steam bath

(use in hood) or with the Roto-Evaporator.

To remove solvent:

1. You must have ebullation to concentrate at atmospheric pressure--use a boiling stone.
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2. If you used reduced pressure to concentrate solution, use the water aspirator with a TRAP in the line.

DO NOT turn off the water until the pressure is released. In general, CLAMP any flask that could

conceivably trip over.

Crystallization (review)

Used to obtain pure crystalline solid

Use the proper solvent or solvents--test if necessary; a proper solvent will exhibit a big solubility

difference over a small temperature range.

Re-crystallization or crystallization

a. use an Erlenmeyer flask, it is specifically designed for this purpose

b. dissolve solid in minimum amount of boiling solvent - add solvent in small amounts. For example, if

you add 5 mL and approx. half of the solid dissolves, it should take only another 5 mL to dissolve the

remaining half. If some of the solid does not dissolve then....

c. are remaining particles your compound or insoluble material (eg, sand, old boiling stones)?

d. to determine this, add ca. 10% more hot solvent. If insoluble material, you can decant (carefully

transfer solution into another flask leaving the insoluble material behind) or filter.

e. if filtering is necessary, do so to remove suspended solids, the faster the better, keep solution warm

so crystallization does not occur (this may require filtering on a hot plate or other heating device).

f. to decolorize, use a small amount of charcoal and filter with filter aid (see below). For both (e) and (f),

rinse filter paper with a small amount of hot solvent.

g. let the filtered liquid (filtrate) cool to room temperature slowly in the Erlenmeyer flask


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h. cool the filtrate in an ice-water bath

i. if crystals have not formed

1. seed with a small crystal of product, or

2. scratch the flask with a glass rod which has not been fire polished at the end (ask for a

demonstration), or

3. add a second solvent dropwise until the cloud pointis reached; the cloudiness suggests that the

solute has reached a saturation point in this new mixed solvent and will start to come out of solution.

j. if material oils out, you must redissolve by heating the solution and then proceed again from part (g)

k. if material precipitates out, it is time to filter.

Decolorization:

Most organic compounds are colorless. Highly conjugated compounds (e.g., polar polymers) will absorb

light in the visible region of the spectrum and thus be colored. If these highly polar, large molecules are

impurities, they can be removed by use of finely granulated activated charcoal (Norit). Polar compounds

(eg, polar impurities) adsorb to the charcoal which is insoluble in the solvent and can be filtered away

from solution. Unfortunately, some of your compound will also adsorb if there is enough charcoal so

the trick is to use just the right amount. Usually, a very small amount of charcoal will suffice (there is a

lot of surface to these particles). The Norit is added in small amounts to the hot (but not boiling)

solution until sufficient decolorization has occurred. CAUTION--trapped air in the Norit can cause rapid

frothing when it hits the hot solution. The Norit can be filtered from the hot solution using a fluted filter

paper.

Filtering

Used to remove insoluble solids suspended in solution.


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Use a GRAVITY FILTER FUNNEL when you DONOT want the solid

Use a BUCHNER funnel with vacuum when you do want the solid but....

Never use a Buchner funnel with a hot solution unless suggested by your instructor.

For the Buchner funnels, use a piece of round filter paper which fits the funnel. A proper fit is a piece of

paper that just covers the holes but does not touch the sides of the funnel (ask for a demo). You will

need to use reduced pressure, usually via a water aspirator (be certain that the TRAP is clean). For

aqueous solutions, you can use the house vacuum line (be certain that the TRAP is clean).

Just prior to pouring the precipitate onto the filter paper, wet the filter paper with a few mL of solvent

and apply the vacuum. This will cause the filter paper to stick to the filter.

For the gravity funnel, you can prepare a suitable filter from round filter paper by folding into quarters

or by folding it into more than quarters (fluted like a fan). Ask for a demo. Another trick is to use a

piece of cotton loosely wedged into the cone of the funnel. It must be tight enough to not move and to

trap the solid particles but loose enough to allow the solution to flow through. This procedure is often

used to remove drying agents from the solution.

Receiver Flasks: For gravity filtration, you can collect the filtrate in any receiver, including the

concentration flask you plan to use immediately thereafter. Be certain to secure the receiver since it

will contain your important compound.

For a Buchner funnel, you will need a filter flask with a sidearm for a rubber tubing connection to the

TRAP and the reduced pressure source. Use an appropriately sized filter flask, one that will not fill more

than halfway and secure this flask so that it does not tip over with that expensive funnel containing your

valuable crystals.

TRAP: For suction filtration, you want a clean glass trap in between your filter flask and the suction

source. One reason is obvious--in the event that your filtrate is sucked out of the filter flask, it can be

trapped and recovered before it goes down the drain or into the house vacuum line. Another reason is

that a changing flow of water affects the pressure in the water aspirator so that water can back up and

flow towards your filter flask. This way, the trap will fill first and prevent dilution of your filtrate.


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Rinsing: You should rinse the gravity funnel with a small amount of solvent to wash down the remaining

solution that adheres to the filter paper and funnel. Remember, you want the solution; you do not want

the solid. Use a small amount of solvent to prevent excess dilution of the filtrate.

You should rinse the crystals in the Buchner and Hirsch funnel with a minimum amount of cold solvent.Remember, you want the crystals and do not want to dissolve the crystals with excess solvent which will

wash them into the filtrate.

The proper way to wash the crystals is to SHUT OFF the vacuum, add a minimum amount of cold solvent

so that the crystals are barely sitting in solvent for about 5 seconds (the solvent will not drip through

quickly) and then apply the vacuum. The solvent will be sucked into the filtrate. Do these one or two

times for each solid. You can air dry the solid by sucking air through the solids. To more rapidly dry a large

amount of solid, press another piece of filter paper on top of the solid.

Filtrate from Buchner filtration: The filtrate will probably still contain some of your desired compound

and is called the mother liquor. Often, by concentrating this solution further and cooling the solution, one

obtains more crystals. This may happen while you are drying the original crystals under reduced pressure.

To collect this second batch of crystals, filter as before but do not combine with the first batch of crystals.

The purity of this second crop may be different from the first batch. Use melting point or other methods

(eg, thin layer chromatography) to determine whether the purity of the second crop is equal to that of the

first. If so, they can be combined. If not, keep them separate.


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Solvent Polarity in decreasing order:

water

amides (N,N-dimethylformamide)

alcohols (methanol, ethanol)

ketones (acetone, methyl ethyl ketone)

esters (ethyl acetate)

chlorocarbons (methylene chloride, chloroform)

ethers (diethyl ether)

aromatics (toluene)

alkanes (hexanes, petroleum ether)

Heating:

There are different methods used for heating material in the laboratory. Flames are used only when it is

instructed. Electric hot plates and heating mantles are most commonly used. Be careful not to turn this

equipment to its highest setting which can burn it out. It does take several minutes for these instruments

to reach the desired temperature. The heating mantles are plugged into a variable rheostat which

provides a temperature control. Heating mantles are used for round-bottom flasks (rbf); choose an

appropriate size to fit the flask you plan to use.

Steam is often used for heating volatile, non-aqueous, flammable solvents in the fume hoods. This is

preferable to using hot plates which can lead to flash fires. The 100oC maximum temperature from steam

is sufficient for most commonly used organic solvents including many mixed aqueous solvents.

Ask your instructor which technique is to be used for a particular point of time


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Melting Point Determination:

The standard physical property of a solid is its melting point. The melting point is actually a melting point

range. It is used to help determine the purity of a solid and to help verify the identity of the compound. A

pure compound should melt over a narrow temperature range. Impurities usually cause the melting pointrange to widen and lower in value. To obtain the melting point range, you record the temperature at

which the first crystals begin to melt (solid to liquid phase) and the temperature at which the last crystal

melts, eg; m.p. 124-126oC.

Preparing the sample in a capillary tube: The sample must be dry or else you cannot get it into the

capillary tube. The capillary tube should be open at one end only. Tamp the open end onto your crystals

until a few collect inside the mouth of the capillary. Now tap the crystals to the closed bottom end by

bouncing the capillary (bottom end down) through the open space of a large piece of glass tubing. Ask fora demo. Follow the instructions given above for the rate of heating.

Extraction

Extraction is a method for moving a compound from one medium to another. For example, if you make

coffee from coffee beans, you are extracting some flavorful components of the bean and some caffeine

into the water. The remainder of the beans (grounds) are left behind and discarded. This is called a solid-

liquid extraction. If you are trying to move a compound from one liquid phase (solvent 1) into another

liquid phase (solvent 2), this is liquid-liquid extraction but the two solvents must be immiscible or

insoluble to the extent that they form two distinct layers. The compound is now distributed mostly into

one layers which can be separated.

Liquid-liquid extractions are common in organic chemistry. Usually, one of the solvents is water and the

objective is to remove a component from an aqueous solution into a solvent such as ether, methylene

chloride, or ethyl acetate (all of which have low water solubility). Often, water-insoluble organic solvents,

such as ether, methylene chloride and hexane, may contain some undesirable water soluble components

(like HCl). In that case, we would extract those components out of the organic solvent by using water as

the second solvent. That is often called a water wash. You will have an opportunity to do several

extractions and water washes. For these purposes, you will use a separatory funnel.


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The rule of thumb for liquid-liquid extractions is that several small extractions are more efficient than one

big extraction.

Separatory Funnel: This glass equipment is very cleverly designed to carry out the task of separating two

immiscible liquids (which form two distinct layers). Work with this equipment in a proper fashion and itwill perform remarkably well. However, read the instructions first and follow the steps carefully.

Use of the Separatory funnel: First, check that the stopper fits and that the stopcock works properly.

Glass stopcocks must be greased; Teflon stopcocks do not need grease.

2. Close the stopcock and support the funnel using a ring clamp.

3. Place a beaker underneath the funnel to catch any spills or leaks.

4. Fill the funnel with the two solvents. Do not fill the funnel more than 75% of capacity (so plan ahead in

choosing the proper size funnel)

5. Stopper. Remove the funnel from the stand, hold properly and invert (IMPORTANT, ask for a demo),

release any pressure buildup by opening the stopcock repeatedly BEFORE shaking (and then after

shaking).

6. After returning the funnel to the stand where it is secured, loosen the stopper immediately--shaking

builds up pressure

7. When the two phases separate, draw off the lower layer.

8. Pour out the upper layer if necessary. NOTE: The funnel is designed to retain the last few drops.

9. Save both upper and lower layers until you are certainthat you have the compound you want. If you do

not throw it away, it is not lost. I repeat, neverthrow away a layer unless you are certain that you will

never need it.

Percent Yield: Although you may have obtained the product you desired, the amount of material you

obtained (in grams), compared to the amount you could have obtained (in grams), is a valuable piece of

information. It can suggest that you did or did not run the reaction efficiently or that other products may


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have been formed via other reactions. To calculate % Yield you must first calculate the theoretical yield.

You then take actual yield/theoretical yieldx 100 = % yield. The % yield can never exceed 100%; if it does,

it may mean that your sample is wet or contains extraneous material, such as a boiling stone.

Labels

Properly label your sample vials when you submit your products for grading.

e.g.

BENZOIC ACID (appropriate, unambiguous name of the compound)

mp 120-121oC (melting point observed for this sample)

3.2 g (amount in vial)

Students name and ID

Date of submission

Distillation

Distillation is an important commercial process that is used in the purification of a large variety of

materials. All substances regardless of whether they are liquids or solids are characterized by a

vapor pressure. The vapor pressureof a pure substance is the pressure exerted by the substance

against the external pressure which is usually atmospheric pressure. Vapor pressure is a measure

of the tendency of a condensed substance to escape the condensed phase. The larger the vapor

pressure, the greater the tendency to escape. When the vapor pressure of a liquid substance

reaches the external pressure, the substance is observed to boil. If the external pressure is

atmospheric pressure, the temperature at which a pure substance boils is called the normal

boiling point.


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Simple Distillation

Although all of us have brought water to a boil many times, some of us may have not

realized that the temperature of pure boiling water does not change as it distills. This is why

vigorous boiling does not cook food any faster than a slow gentle boil. The observation that

the boiling point of a pure material does not change during the course of distillation is an

important property of a pure material. The boiling point and boiling point range have been

used as criteria in confirming both the identity and purity of a substance. For example, if we

synthesized a known liquid that boiled at 120-122 C, this value could be used to confirm

that we prepared what we were interested in and that our substance was reasonably pure.

Of course, additional criteria must also be satisfied before the identity and purity of the

liquid are known with certainty. In general, a boiling point range of 1-2 C is usually taken as

an indication of a pure material. You will use both of these properties later in the semester

to identity an unknown liquid.

Occasionally, mixtures of liquids called azeotropes can be encountered that mimic

the boiling behavior of pure liquids. These mixtures when present at specific concentrations

usually distill at a constant boiling temperature and can not be separated by distillation.

Examples of such mixtures are 95% ethanol-5% water (bp 78.1 C), 20% acetone-80%

chloroform (bp 64.7 C), 74.1% benzene, 7.4% water, 18.5 % ethanol (bp 64.9). The

azeotropic composition sometimes boils lower the than boiling point of its components and

sometimes higher. Mixtures of these substances at compositions other than those given

above behave as mixtures.

Returning to our discussion of boiling water, if we were making a syrup by the

addition of sugar to boiling water, we would find that the boiling point of the syrup would

increase as the syrup begins to thicken and the sugar concentration becomes significant.

Unlike pure materials, the boiling point of an impure liquid will change and this change is a

reflection of the change in the composition of the liquid. In fact it is this dependence of

boiling point on composition that forms the basis of using distillation for purifying liquids.


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We will begin our discussion of distillation by introducing Raoult's Law, which treats liquids in

a simple and ideal, but extremely useful manner.

Figure 1. The apparatus used in a simple distillation. Note the position of the

thermometer bulb in the distillation head and the arrangement of the flow of the cooling

water.

PAob s = PA

o , where PAob s

is the observed vapor pressure of A 1

= nA/(nA + nB + nC +...)

PAo= vapor pressure of pure A

nA, nB,nC ... : number of moles of A, B, C, ...


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This relationship as defined is capable of describing the boiling point behavior of

compound A in a mixture of compounds under a variety of different circumstances. Although

this equation treats mixtures of compounds in an oversimplified fashion and is not applicable

to azeotropic compositions, it does give a good representation of the behavior of many

mixtures.

Let's first consider a binary system (2 components) in which only one of the two

components is appreciably volatile. Our previous discussion of sugar + water is a good

example. Raoult's law states that the observed vapor pressure of water is simply equal to the

product of the mole fraction of the water present and the vapor pressure of pure water at

the temperature of the mixture. Once the sugar-water mixture starts to boil, and continues

to boil, we know that the observed vapor pressure of the water must equal one atmosphere.

Water is the only component that is distilling. Since the mole fraction of water in a mixture

of sugar-water must be less than 1, in order for the observed vapor pressure of water ( PAob s

)

to equal one atmosphere, PAo must be greater than one atmosphere. As the distillation of

water continues, the mole fraction of the water continues to decrease thereby causing the

temperature of the mixture to increase. Remember, heat is constantly being added. If at

some point the composition of the solution becomes saturated with regards to sugar and the

sugar begins to crystallize out of solution, the composition of the solution will become

constant; removal of any additional water will simply result in the deposit of more sugar

crystals. Under these set of circumstances, the composition of the solution will remain

constant and so will the temperature of the solution although it will exceed 100 C.

During the course of the distillation, the water vapor which distilled was initially at

the temperature of the solution. Suspending a thermometer above this solution will record

the temperature of the escaping vapor. As it departs from the solution, the temperature ofthe vapor will cool by collisions with the surface of vessel until it reaches 100 C. Cooling

below this temperature will cause most of the vapor to condense to a liquid. If cooling to 20

C occurs in the condenser of a distillation apparatus, then by using the appropriate

geometry as shown in Figure 1, it would be possible to collect nearly all of the liquid. The


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only vapor that would be lost to the environment would be that small amount associated

with the vapor pressure of water at 20 C. Since the vapor pressure of water at 20 C is

roughly 2.3 kPa, then 2.3/101.325 or 0.023 would be the fraction of water that would not

condense and would pass out of the condenser. This is why the distillate is frequently chilled

in an ice bath during the distillation.

The distillation of a volatile material from non-volatile is one practical use of

distillation which works very well. However, often there may be other components present

that although they may differ in relative volatility, are nevertheless volatile themselves. Let's

now consider the two component system you will be using in the distillations you will

perform in the laboratory, cyclohexane and methylcyclohexane. The vapor pressures of

these two materials in pure form are given in Table 1. As you can see from this table,

although cyclohexane is more volatile than methylcyclohexane, the difference in volatility

between the two at a given temperature is not very great. This means that both materials

will contribute substantially to the total vapor pressure exhibited by the solution if the

distillation is carried out at 1 atmosphere. The total pressure, PT, exerted by the solution

against the atmosphere according to Dalton's Law of partial pressures, equation 2, is simply

the sum of the observed vapor pressures of cyclohexane, Pcob s, and methylcyclohexane, Pm

ob s:

PT = Pcob s

+ Pmob s

. 2

As before, boiling will occur when the total pressure, PT , equals an atmosphere. However

since we have two components contributing to the total pressure, we need to determine the

relative contributions of each. Again we can use Raoult's Law but we need more information

about the system before we can do so. In particular we need to know the composition of

the solution of cyclohexane and methylcyclohexane. For ease of calculation, let's assume

that our original solution has equal molar amounts of the two components. What we would

like to determine is whether it would be possible to separate cyclohexane from

methylcyclohexane by distillation. By separation, we would like to determine if it would be

possible to end up with two receiver flasks at the end of the experiment that would contain


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mainly cyclohexane in one and mainly methylcyclohexane in the other. It is clear that at

some point we will need to intervene in this

Table 1. Vapor pressures of cyclohexane and methyl cyclohexane as a function of

temperature.

cyclohexane methylcyclohexane

T/K P/kPa T/K P/kPa

300 14.1 300 6.7

305 17.6 305 8.5

310 21.7 310 10.6

315 26.5 315 13.2

320 32.2 320 16.2

325 38.8 325 19.8

330 46.5 330 24.0

335 55.3 335 28.9

340 65.4 340 34.6

345 77.0 345 41.2

350 90.0 350 48.7

354 101.3 354 55.4

360 121.3 360 66.9

362 128.5 362 71.1

365 139.9 365 77.9

370 160.5 370 90.2


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373 174.0 373 101.3

380 208.8 380 119.3

385 236.7 385 136.4

390 267.3 390 155.3

395 300.9 395 176.2

400 337.5 400 199.1

experiment. Otherwise, if we were to collect the entire contents of the original distilling

flask, called the pot, into one receiver flask, we would end up with the same composition as

we started. Initially the mole fractions of both cyclohexane and methylcyclohexane are 0.5.

From Raoult's Law (equation 1), Dalton's Law (equation 2) and the information in Table 1, we

can estimate that boiling will occur at approximately 362 K when the total pressure of the

two components equals one atmosphere or 101.3 kPa.:

PT = Pcob s+ Pm

ob s= 0.5(128.5 kPa) + 0.5(71.1 kPa) 101.3 kPa

The first thing that we should note is that the initial boiling point is higher than the lowest

boiling component and lower than the highest boiling component. Next, we should inquire

about the composition of the vapor. Is the composition of the vapor the same as the initial

composition of the pot or is it enriched in the more volatile component? If the composition

of the vapor is the same as that of the original mixture, then distillation will not be successful

in separating the two components. However, we should ask, "What is the composition of the

vapor?" Before the vapor is cooled and condenses on the condenser, we can treat the vapor

as an ideal gas. Recalling that: PV = nRT, where P is the pressure of the gas or vapor, V is the

volume it occupies, n is the number of moles of gas, R is the Gas Constant (0.0821 L .atm.K-

1.mol

-1) and T is the temperature, we can determine the composition of the vapor by taking

advantage of the following factors. First we note that:


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PTV = nTRT so that Pcob sV = ncRT and Pm

ob sV = nmRT where nT refers to the total

number of moles, since (Pcob s+ Pm

ob s) = (nc + nm)RT.

If the total vapor can be treated as an ideal gas, then according to Dalton's Law, so can each

of the components. Since the two components are in thermal contact and are distilling

together, we can expect them to be at the same temperature. We don't necessarily know

the volume of the container, but since it is assumed that the volumes of the molecules are

very small in comparison to the total volume the gas occupies, whatever the value of V, it is

the same for both components. This means we can establish the following ratio:

Pcob s/Pm

ob s= nc/nm

which in turn allows us to determine the composition of the vapor from the observed partial

pressures of the two components. If we use the experimental values found in Table 1, we

conclude that the composition of the vapor is 1.8/1, and is indeed enriched in the more

volatile component.

This simple treatment allows us to understand the principles behind distillation.

However it is important to point out that distillation is far more complex than our simple

calculation indicates. For example, we just calculated the composition of the vapor as soon

as the solution begins to boil and we have correctly determined that the vapor will be

enriched in the more volatile component. This means that as the distillation proceeds, the

pot will be enriched in the less volatile component. Since the composition of the pot will

change from the initial 1:1 mole ratio and become enriched in the less volatile component;

the new composition in the pot will introduce changes in the composition of the vapor. The

composition of the vapor will also change from the initial ratio we just calculated to a new

ratio to reflect the new composition of the pot. The consequences of these changes are that

the temperature of both the pot and the distillate will slowly increase from the initial value

to a value approaching the boiling point and composition of the less volatile component. If

we are interested in separating our mixture into components, we are left with the task of


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deciding how much material to collect in each receiver and how many receivers to use.

Obviously this will depend on the quality of separation we are interested in achieving.

Generally, the more receivers we use the less material we will have in each. It is possible to

combine fractions that differ very little in composition but this requires us to analyze each

mixture. While it is possible to do this, in general, we really want to end with three receivers,

one each enriched in the two components of our mixture and a third that contains a

composition close to the initial composition.

It is difficult to describe how much material to collect in each receiver since the

volume collected will depend on the differences in the boiling points of the components. As

a general rule, the receiver should be changed for every 10 C change in boiling point. Each

fraction collected can be analyzed and those with compositions similar to the initial

composition can be combined. The main fractions collected can then be fractionated a

second time if necessary.

The experiment we have just discussed is called a simple distillation. It is an

experiment that involves a single equilibration between the liquid and vapor. This distillation

is referred to as involving one theoretical plate. As you will see, it is possible to design more

efficient distillation columns that provide separations on the basis of many theoretical

plates. Before discussing these columns and the advantages offered by such fractionating

columns, it is important to understand the basis of the advantages offered by columns with

many theoretical plates. The following is a simplified discussion of the process just described

involving a column with more than one theoretical plate.

Vacuum Distillation

Elevation of the boiling point with an increase in external pressure, while important

in cooking and sterilizing food or utensils, is less important in distillation. However, it

illustrates an important principle that is used in the distillation of many materials. If the

boiling point of water is increased when the external pressure is increased, then decreasing

the external pressure should decrease the boiling point. While this is not particularly


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important for the purification of water, this principle is used in the process of freeze drying,

an important commercial process. In addition, many compounds cannot be distilled at

atmospheric pressure because their boiling points are so high. At their normal boiling points,

the compounds decompose. Some of these materials can be distilled under reduced

pressure however, because the required temperature to boil the substance can be lowered

significantly. Rewording the "rule of thumb" described above so that it is applicable here

suggests that the boiling point will be lowered by 10 C each time the external pressure is

halved. For example, if the external pressure above a substance is reduced to 1/16 of an

atmosphere by mean of a mechanical pump, the boiling point will have been reduced four

times by 10 C for a total reduction of 40 C (1 atm x (1/2)(1/2)(1/2)(1/2) =1/16 atm).

A nomograph is a useful device that can be used to estimate the boiling point of a

liquid under reduced pressure under any conditions provide either the normal boiling point

or the boiling point at a some given pressure is available. To use the nomograph given the

normal boiling point, simply place a straight edge at on the temperature in the central

column of the nomograph (b). Rotating the straight edge about this temperature will afford

the expected boiling point for any number of external pressures. Simply read the

temperature and the corresponding pressure from where the straight edge intersects the

first and third columns. As an example lets choose a normal boiling point of 400 C. Using

the nomograph in Figure 2 and this temperature for reference, rotating the straight edge

about this temperature will afford a continuous range of expected boiling points and the

required external pressures necessary to achieve the desired boiling point. At a pressure of 6

mm, the expected boiling point would be 200 C. Likewise, our compound boiling at 400 C

at 1 atm would be expected to boil at 145 C at 0.1 mm external pressure.


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(a) (b) (c)

Figure 2. A nomograph used to estimate boiling points at reduced pressures. To use, place a

straight edge on two of the three known properties and read out the third. Column c is in

mm of mercury. An atmosphere is also equivalent to 101.3 kPa and will support a column of

mercury, 76 cm (760 mm).

Fractional Distillation

We have just seen that starting with a composition of 1:1, cyclohexane:

methylcyclohexane, the composition of the vapor was enriched in the more volatile

Pressure-Temperature Nomograph for Vacuum Distillations

Boiling Point

at pressure PBoiling Point

at 1 atm

Pressure (mm)

1 atm=760

mm;

F C


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component. Suppose we were to collect and condense the vapor and then allow the

resulting liquid to reach equilibrium with its vapor. Lets call this liquid, liquid 2. The

properties of liquid 2 will differ from the original composition in two ways. First, since the

composition of liquid 2 is higher in cyclohexane than the initial one; the temperature at

which liquid 2 will boil will be lower than before (what is the approximate boiling point of a

1.8/1 mixture of cyclohexane/methylcyclohexane? see Table 1). In addition, the composition

of the vapor, vapor 2, in equilibrium with liquid 2 will again be enriched in the more volatile

component. This is exactly what happened in the first equilibration (first theoretical plate)

and this process will be repeated with each new equilibration. If this process is repeated

many times, the vapor will approach the composition of the most volatile component, in this

case pure cyclohexane, and the liquid in the pot will begin to approach the composition of

the less volatile component, methylcyclohexane. In order for this distillation to be successful,

it is important to allow the condensed liquid which is enriched in the less volatile component

relative to its vapor, to return to the pot. In a fractional distillation, the best separation is

achieved when the system is kept as close to equilibrium as possible. This means that the

cyclohexane should be removed from the distillation apparatus very slowly. Most fractional

distillation apparati are designed in such a way as to permit control of the amount of

distillate that is removed from the system. Initially the apparatus is set up for total reflux,(i.e. all the distillate is returned back to the system). Once the distillation system reaches

equilibrium, a reflux to takeoff ratio of about 100:1 is often used (about 1 out of every 100

drops reaching the condenser is collected in the receiver).

A column which allows for multiple equilibrations is called a fractionating column and

the process is called fractional distillation. An example of a fractionating column is shown in

Figure 4. Each theoretical plate is easy to visualize in this column. The column contains a

total of 4 theoretical plates and including the first equilibration between the pot and

chamber 1 accounts for a total of 5 from pot to receiver. As you might expect, a problem

with this column is the amount of liquid that is retained by the column. Many other column

designs have been developed that offer the advantages of multiple theoretical plates with


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low solvent retention. Typical spinning band columns often used in research laboratories

offer fractionating capabilities in the thousand of theoretical plates with solute retention of

less than one mL. Commercial distillation columns have

Figure 4. A fractionating column which contains four chambers, each with a center opening

into the chamber directly above. The vapor entering the first chamber cools slightly and

condenses, filling the lower chamber with liquid. At equilibrium, all chambers are filled with


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distillate. A portion of the liquid condensing in the first chamber is allowed to return to the

pot. The remaining liquid will volatilize and travel up the column condensing in the second

chamber and so on. As discussed in the text, the composition of the vapor at each

equilibration is enriched in the more volatile component. The heat necessary to volatilize the

liquid in each chamber is obtained from the heat released from the condensing vapors

replacing the liquid that has been removed. The vacuum jacket that surrounds the column

ensures a minimum of heat loss.

Columns have been designed for gasoline refineries that are multiple stories high and are

capable of separating compounds with boiling points that differ by only a few degrees.

In addition to performing a fractional distillation at one atmosphere pressure, it is

also possible to conduct fractional distillations at other pressures. This is often avoided

when possible because of the increased difficulty and expense in maintaining the vacuum

system leak free.

Steam Distillation

The concentration and isolation of an essential oil from a natural product has had a

dramatic impact on the development of medicine and food chemistry. The ability to

characterize the structure of the active ingredient from a natural product has permitted

synthesis of this material from other chemicals, resulting in reliable and often cheaper

sources of the essential oil. The process often used in this isolation is called steam

distillation. Steam distillation is an important technique that has significant commercial

applications. Many compounds, both solids and liquids, are separated from otherwise

complex mixtures by taking advantage of their volatility in steam. A compound must satisfy

three conditions to be successfully separated by steam distillation. It must be stable and

relatively insoluble in boiling water, and it must have a vapor pressure in boiling water that is

of the order of 1 kPa (0.01) atmosphere. If two or more compounds satisfy these three


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conditions, they will generally not separate from each other but will be separated from

everything else. The following example, expressed as a problem, illustrates the application of

steam distillation:

Suppose we have 1 g of an organic compound present in 100 g of plant material composed

mainly of macromolecular material such cellulose and related substances. Let's assume that

the volatile organic material has a molecular weight of 150 Daltons, a vapor pressure of 1

kPa and is not soluble in water to an appreciable extent. Examples of such materials

characterized by these properties include eugenol from cloves, cinnamaldehyde from

cinnamon bark or cuminaldehyde from cumin seeds. How much water must we collect to be

assured we have isolated all of the natural oil from the bulk of the remaining material?

We can simplify this problem by pointing out that the organic material is not

appreciably soluble in water. We know from previous discussions that boiling will occur

when the total pressure of our system equals atmospheric pressure. We can also simplify the

problem by assuming that the essential oil in not appreciably soluble in the macromolecular

material. While in reality this does not have to be correct, this assumption simplifies our

calculation. Boiling of our

PT = Pwaterob s + Porgobs; PT = waterPwatero + orgPorgo

mixture will occur close to 100C. Remember that very little oil is soluble in water which

makes the mole fraction of water near unity. Similarly for the volatile oil, its mole fraction is

also close to one according to our assumption. The total pressure, PT, is the sum of the vapor

pressure of water, 100 kPa, and the essential oil, Porgo

,1 kPa. Boiling will occur very close to

the boiling point of pure water. Treating the water vapor and the organic vapor which are

miscible as ideal, the PV ratio for both vapors is given by the following:

PwaterV/ PorgV = nwaterRT/norgRT;

Pwater/ Porg = nwater/norg and


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nwater= wtwater/18; norg = 1/150; rearranging:

wtwater = (100/1)(18/150) = 120 g water or 120 mL

Our calculation suggests that we can be assured that most of the 1 g of the organic mater

has been transferred by the steam if we condense and collect 120 mL of water. The basis of

the separation by steam distillation is that while the water and organic condensed phases

are immiscible, the vapors of both are miscible. Once condensed, the two separate again

allowing for an easy separation. As noted above, both liquids and solids can be distilled by

steam.


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PART B

Here, you will be made familiar with three major parts of practical organic chemistry namely

a) qualitative analysis of known and unknown organic samples, b) quantitative estimation of

a few organic substances and c) single and multi step syntheses of several organic

compounds. Each of these parts consists of number of experiments and will take several lab

hours. The following section will describe the experimental procedure.


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EXPERIMENT - 1

ELEMENTAL ANALYSIS OF ORGANIC COMPOUNDS

Objective:

To carry out the detection of inorganic elements in an organic compound i.e. nitrogen, sulfur

and halogens (chlorine, bromine and iodine).

Principles:

In an organic compound carbon, hydrogen and oxygen are assumed to be present. Elements

other than these elements i.e. nitrogen, sulfur and halogens (chlorine, bromine and iodine)

may also be present in an organic compound. As you have known earlier that chemical

analysis is of two types: Qualitative analysis and Quantitative analysis. The detection of extra

elements in a given organic compound is a type of qualitative analysis since here we are

dealing with the composition of the compound. This experiment is to be done very carefully

as further analysis of the organic compound is based on the extra element present in it.

LASSAIGNES TEST

These extra elements are usually detected by Lassaignes test. In this test, the organic

compound is fused with metallic sodium to convert these elements into water soluble

sodium salts. And then this extract is used to perform the tests.

Experimental Procedure

Preparation of Lassaignes extract

1. Take a small piece of sodium metal with a metallic spatula and dry it by pressing between

the folds of filter paper and put this piece of sodium metal into an ignition tube.

2. The ignition tube is heated slowly till the sodium metal turns into shining globule.


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3. Remove it from flame and add the compound to be tested (a pinch of solid compound or

2-3 drops of liquid compound). First heat it gently after compound addition and then

strongly until it becomes red hot.

4. Plunge it into 10 ml of distilled water contained in a china dish and cover the china dish

immediately with wire gauze to stop the flames. The same procedure is repeated by 2

more ignition tubes. The ignition tubes if not broken completely, are crushed with the

help of a glass rod.

5.

Boil the contents of the china dish for about 5 minutes and filter. The filtrate is called

Lassaignes extract or sodium extract. The filtrate should be colorless. Filtrate will be

colored when fusion is incomplete. So, if it is colored, whole procedure should be

repeated again.

The reactions involved in the preparation of Lassaignes extract are:

Nitrogen, Na + C + N NaCN (sodium cyanide). So, nitrogen will be present in the form of

sodium cyanide in the Lassaignes extract.

Sulfur, 2Na + S Na2S (sodium sulfide). Thus, sulfur will be present in the form of sodiumsulfide in the Lassaignes extract.

Nitrogen and sulfur both, Na + C + N + S NaSCN(sodium thiocyanate). So, when nitrogen

and sulfur are present together in the organic compound, they will be present in the form of

sodium thiocyanate in the Lassaignes extract.

Halogens, Na + X NaX (X=Cl,Br and I) (sodium halide). So, halogens will be present in theform of sodium halide (if X=Cl, then sodium chloride, if X= Br, then sodium bromide and if X=

I, then sodium iodide) in the Lassaignes extract.


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Reactions of Nitrogen test:

1. Ferrous sulfate reacts with the NaOH present in the extract to give a dirty green color

precipitate. FeSO4+ 2NaOH Fe(OH)2 + Na2SO4

(Ferrous hydroxide, dirty green ppt)

NaOH is formed in the Lassaignes extract by the reaction between unreacted sodium

metal and water. Thus, if no green precipitate is obtained that means NaOH is not

present in the extract so few drops of it is added in the test tube to get the precipitate.

2.

Ferrous ions react with sodium cyanide (ionic form of nitrogen in the extract) to give

sodium ferricyanide as per the equation Fe(OH)2+ 6NaCN Na4Fe(CN)6+ 2NaOH

(sodium ferrocyanide)

On boiling the alkaline solution of ferrous salt, some ferric ions are inevitably produced by

the aerial oxidation. And ferrous and ferric hydroxides are dissolved on the addition of dilute

sulfuric acid. Sodium ferrocyanide reacts with ferric salt to give ferriferrocyanide complex

which is Prussian blue color complex.

3Na4Fe(CN)6+ 2Fe2(SO4)3 Fe4[Fe(CN)6]3+ 6Na2SO4

(Ferri-ferrocyanide, Prussian blue)

Rection of Sulfur test:Sodium sulfide (ionic form of sulfur in the extract) reacts with sodium nitroprusside to give

sodium sulphonitroprusside which is violet in color.

Na2S + Na2[Fe(CN)5NO] Na4[Fe(CN)5NOS]

(sodium nitroprusside) (sodium sulphonitroprusside) (violet color)

Reaction of Nitrogen and Sulfur test together:

Sodium thiocyanate (ionic form of nitrogen and sulfur together in the extract) reactswith ferric chloride to give ferric thiocyanate, a blood red color complex.

3NaSCN + FeCl3 Fe(SCN)3+ 3NaCl

(ferric thiocyanate) (blood red color)


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During the preparation of Lassaignes extract, if the organic compound containing both

nitrogen and sulfur is fused with an excess of sodium metal, instead of sodium

thiocyanate, sodium cyanide and sodium sulfide are formed.

NaSCN + 2Na NaCN + Na2S

In such cases, the tests for NaCN and Na2S can be performed. Thus in such cases

individual test for nitrogen and sulfur will be positive.

Reaction of Halogens:

Sodium halide (ionic form of halogen in extract) reacts with silver nitrate to give the

precipitate of silver halide. NaX + AgNO3 AgX + NaNO3

If X=Cl, then AgCl (silver chloride) will be formed which is white color precipitate soluble in

Ammonium hydroxide.

If X=Br, then AgBr (silver bromide) will be formed which is pale yellow color precipitate

soluble in excess of Ammonium hydroxide.

If X=I, then AgI (silver iodide) will be formed which is yellow color precipitate insoluble in

Ammonium hydroxide.

Reaction of Layer test:

Chlorine water is used as an oxidizing agent and Br-(ionic form of bromine in extract)

gets converted into Br2as per the following equation

2NaBr + Cl2 2NaCl + Br2and Br2gives orange color to organic layer.

Chlorine water oxidises I- (ionic form of iodine in extract) into I2

2NaI + Cl2 2NaCl + I2and I2gives violet color to organic layer.

SUMMARY

We have learnt about the detection of following extra elements in the given organic

compound:


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nitrogen, sulfur, chlorine, bromine and Iodine.

HOW TO RECORD

AIM:To detect the presence of extra element(s) in the given organic compound.

APPARATUS REQUIRED:China dish, funnel, test tubes, ignition tubes, capillary tubes, glass

rod, wire gauze, pair of tongs, test tube holder, dropper, and spatula.

CHEMICALS REQUIRED: Organic compound whose extra element is to be determined,

sodium metal, ferrous sulfate, sodium nitroprusside, ferric chloride solution, silver nitrate,

aq. NaOH, dil. Sulfuric acid, dil. Nitric acid, dil. Hydrochloric acid, conc. Nitric acid

OBSERVATION TABLE:

Sl.No. Experiment Observation Inference

1 Test for Nitrogen

To the Lassaignes extract (1-2 ml),add ferrous sulfate (50-75mg).

[If no precipitate is formed, add a

few drops of dilute sodium

hydroxide solution to get the

precipitate.]

Now heat this mixture gently with

shaking for 1 minute and add

dilute sulfuric acid into it.

Dirty green

precipitate

Prussian blue color

Nitrogen is

present

2 Test for Sulfur Violet color Sulfur is


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To the sodium extract (1-2 ml) add

1ml of freshly prepared sodium

nitroprusside solution.

formation present

3 Test for Nitrogen and Sulfur

present together

Take 1-2 ml of sodium extract and

acidify it with dilute hydrochloric

acid. Add 2-3 drops of ferric

chloride solution.

Blood red color Nitrogen and

Sulfur are

present

together

4 Test for Halogens:Silver Nitrate Test When nitrogen

and / or sulfur are absent

Acidify 1-2ml of sodium extract

with dilute nitric acid. Add silver

nitrate solution (0.5ml)

Silver Nitrate Test When nitrogen

and / or sulfur are present

White precipitatesoluble in

Ammonium

hydroxide solution

or

Pale yellow

precipitate soluble

in excess ofAmmonium

hydroxide solution

or

Yellow precipitate

insoluble in

Ammonium

hydroxide

White precipitate

soluble in

Chlorine ispresent

Bromine is

present

Iodine is

present

Chlorine is

present


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PRECAUTIONS:

1) Do not touch the sodium metal with your fingers; always use forceps to handle it.

2) Never throw the sodium metal into sink.

3) Unreacted sodium metal should be carefully decomposed by adding small amount of

ethyl alcohol into it.

4) Always use freshly prepared ferrous sulfate and sodium nitroprusside solution.

5)

Fuse 3-4 ignition tubes so that the extract should be concentrated.

6) Remove sodium cyanide and sodium sulfide before performing the test for halogens by

using nitric acid before performing the silver nitrate test.

7) In layer test, shake the test tube by putting thumb on the mouth of the test tube.

8) When the liquid compound gets evaporated very rapidly before its fusion with sodium

metal, a paste of the compound with sodium carbonate may be used or after the

addition of liquid compound in the ignition tube, solid sodium carbonate may be added.

Alternative Green Procedures of LASSAIGNES TEST

In the present practice, the use of metallic sodium for fusion with organic compound is

terribly hazardous and is a cause of great worry and concern in a student laboratory. The

idea of fusion with sodium metal is to convert the water insoluble organically bound extra

elements to water-soluble sodium salts which can be easily detected by various tests.


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A non hazardous and safe procedure:

Use of zinc and sodium carbonate instead of metallic sodium

Organic sample (about 10 mg) is thoroughly mixed with an intimate mixture of Zn dust (40

mg) and Na2CO3 (60 mg) powder in a fusion tube [no need to weigh any of the chemicals,

take very small amount of sample and approximately 5 times of Zn and 5 times of Na2CO3],

heated first gently and then strongly in the flame till it becomes red hot and kept at red-hot

condition for two minutes. The bottom part of the fusion tube is plunged into 5 ml of

distilled water taken in a mortar, ground well with the pestle and filtered. With the filtrate

tests for S, N and Cl / Br / I are carried out as usual as in the case of Lassaignes Test.

Notes: 1. The fusion tube must be heated VERY STRONGLY, KEEPING AT RED HOT CONDITION

THROUGHOUT FOR AT LEAST TWO MINUTES. If not properly heated, fusion is not properly

done (as in case of sodium also), and thus expected observation (colour change) may not bemade. In that case, it is advised to repeat the fusion.

2. The amount of water taken in the mortar must be within 5 ml.; otherwise, the

solution will be too dilute to respond to tests, described below.

3. While carrying out the test for nitrogen, ferrous sulphate crystals are to be added; not the

solution. This is to avoid excessive dilution.

4. Acidification must be carried out with dilute sulfuric acid, not with HCl.

5. No ferric chloride should be added.

6. This method works well with covalently bound nitrogen of aromatic compounds (except p-

nitrophenol and p-aminophenol).

S. No. Experiment Observation Inference

1. 0.5 ml filtrate + FeSO4

crystal heat + dil. H2SO

4

Prussian blue color N present

2. a) 0.5 ml filtrate + Sodium

nitroprusside

b) filtrate + dil acetic acid +

Lead acetate

Violet color

Black ppt

S present

S present

3. filtrate + FeCl3 Blood red N+S both present


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4. filtrate + 2 drops

Conc HNO3

boil

cool + AgNO3

Curdy white ppt

(Soluble in

NH4OH)

Pale yellow ppt

(Partly soluble in

NH4OH)

Yellow ppt

(Insoluble in

NH4OH)

Chlorine present

Bromine present

Iodine present

Green context:

This experiment totally eliminates the risk of explosion and fire hazard which are

often met while carrying out the same experiments using metallic sodium.

The aforesaid zinc-alkali mixture (prepared by intimately mixing 2 parts by weight ofzinc dust and 3 parts by weight of sodium carbonate can be stored in a stoppered

bottle for more than a month.


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EXPERIMENT - 2

DETECTION OF FUNCTIONAL GROUPS IN KNOWN ORGANIC COMPOUNDS

I Objective:

To carry out the detection of reactive functional groups in a known organic compound

containing functional groups such as alcohol, aldehyde, ketone, nitro, amine, etc.

II Principles:

In an organic compound one or more reactive functional groups could be present giving the

compound its characteristic properties. Generally these are present due to the heteroatoms

(nitrogen and oxygen) found in the molecules giving rise to functional groups like alcohols,

aldehydes, ketones, amines, etc. These functional groups impart a particular reactivity to the

molecule towards oxidizing/reducing agents or can for addition products using which, we will

detect them.

III Procedures

The following procedures describe the methods of detection of various functional groups

in a molecule based on their reactivities.

Detection of unsaturation in a molecule using Bromine test: Alkenes and alkynes will add

bromine across the multiple bonds unless there are electron withdrawing groups on the

multiple bonds. One observes the rapid disappearance of the red-brown bromine color.

Aromatic compounds can react with bromine more slowly to give bromine substitution andthe formation of HBr, which can sometimes be observed by placing a piece of wet litmus

paper over the mouth of the test tube. Warning: the reagent deteriorates with the

formation of HBr so compare the results with a blank.


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Experiment Observation Inference

Into a dry, clean test tube, dissolve 0.1 mL of a

liquid (or 50 mg of a solid) in 1 mL of

methylene chloride. Add a 2% solution of

bromine in a dichloromethane dropwise with

agitation.

A positive test requires five or more drops of

bromine solution to reach a persistent red-

brown color.

The bromine color

should disappear. To

continue addition

until the bromine

color disappears to

ensure all

unsaturation has

reacted.

Detection of unsaturation using Permanganate test: Aqueous permanganate rapidlyoxidizes double and triple bonds while being reduced to MnO2, a brown precipitate.

Therefore, disappearance of the purple color and formation of a brown precipitate in

minutes is a positive test. However, other compounds react slowly with the reagent

including alcohols, aldehydes, phenols, and aromatic amines so interpret your results

carefully and look for corroboration from the other tests.


Into a clean test tube, dissolve 0.1 mL of a liquid (or

50 mg of a solid) in 1 mL of 95% ethanol or 1,2-

dimethyoxyethane. Add a 1% solution of aqueous

potassium permanganate dropwise with agitation.

The purple

color of

permanganate

disappears

Detection of Aldehydes and Ketones using 2,4-DNP: Hydrazines such as 2,4-

dinitrophenylhydrazine react with the carbonyl group of aldehydes and ketones to give

colored precipitates. Normally the reaction is fast but heating may be necessary. The test

solution is prepared using sulfuric acid and 95% ethanol. Later, if you wish to make a

derivative of your compound, you can use a different 2,4-DNP solution prepared with HCl


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and methanol. This usually gives a slower forming precipitate which often provides a

derivative of higher purity (and higher mp). However, the slow formation of the precipitate

is not desirable when looking for a qualitative test signal


Dissolve only 1 drop of your liquid compound (or

10 mg of your solid) in a minimum number of

drops of 95% ethanol in a test tube. Add 1 mL of

the 2,4-DNP test solution and agitate. If a

precipitate does not form in 10 minutes, heat on a

water or steam bath for a few minutes.

Yellow to

orange or red

precipitate

formation

Tollens Test for Aldehydes and other easily oxidized functional groups. In this test, a

stabilized silver ion is reduced to elemental silver by an easily oxidized compound, such as an

aldehyde. The aldehyde is oxidized to a carboxylic acid. Prepare the Tollens reagent

immediately before you plan to use it


Preparation of Tollens reagent:

Mix 3 mL of Tollens solution A (aqueous silver

nitrate) with 3 mL of Tollens solution B (10%

aq. NaOH) resulting in the formation of solid

silver oxide. Add 10% ammonium hydroxide

solution dropwise, with agitation, until the

silver oxide just dissolves. This produces a

silver ammonium complex and is the Tollens

solution you will use for the test.

Into each of 3 clean, dry test tubes, add 2 mL


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of the Tollens reagent which is freshly

prepared as above. Dissolve 10 mg of a solid

(or 1 drop of a liquid) unknown in the

minimum amount of bis(2-ethoxyethyl)ether

required to give a clear solution (less than 1

mL). Add the test compound solution

dropwise, with agitation, to the first test tube.

Mix vigorously and allow the solution to

stand. For the blank, simply add 0.5 mL of

bis(2-ethoxyethyl)ether.

A positive test is the

formation of a silver

mirror as the elemental

silver adheres to the

wall of the glass tube.

This test is to be

performed only after

getting a positive result

for the 2,4-DNP test.

Warning: Wash any minor amounts of residual Tollens reagent into a sink and flush with

water. The reagent forms silver fulminate which is very explosive. The test solutions can be

disposed of in a jar labeled for that purpose. The silver mirror can usually be washed clean

with soapy water and a scrub brush. If not, see your instructor.

Detection of alcohols using Chromic Acid Test: In this test, a chromium oxidant is used to

oxidize alcohols and aldehydes and is visually detectable by a color change of the chromium

oxidant. Aliphatic aldehydes are oxidized in less than a minute, aromatic aldehydes take a bit

longer. Since the condition of the acetone is critical, it is wise to carefully run the blank with

just acetone to be certain that the acetone itself is not giving a false positive. Warning:

Cr(VI) compounds are considered suspect carcinogens and should be handled carefully.


Dissolve 10 mg of a solid (or 1 drop of a liquid)

compound in reagent grade acetone in a clean,

dry test tube. Add a few drops of chromic acid

The orange color of the

Cr(VI) ion is replace by

the green color of Cr(III)


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solution one drop at a time with shaking.

Aldehydes and primary and secondary alcohols

are oxidized very quickly. Tertiary alcohols are

not oxidized.

The chromic acid solution is prepared by

dissolving 1.0 g of CrO3 in 1.0 mL of

concentrated sulfuric acid and then carefully

diluting with 3 mL of water.

as the chromium is

reduced indicating that

oxidation of the

functional group has

occurred.

Detection of esters using Ferric Hydroxamate Test: If you have a carbonyl compound whichis not an aldehyde or ketone or carboxylic acid, it could be an ester. One test for esters is the

ferric hydroxamate test whereby the ester is converted to a hydroxamic acid (HOHN-C=O)

which will give a positive ferric chloride test. Since enols can give a positive ferric chloride

test, first test your compound with ferric chloride solution.


Dissolve 10 mg of solid (or 1 drop of liquid)

compound in 1 mL of 95% ethanol, add 1 mL of 1

N HCl, and then a 1-2 drops of 5% ferric chloride

solution. If you obtain a color other than yellow,

the test cannot be used. Otherwise, the test is

conducted as follows: dissolve 50 mg of solid or 2

drops of liquid in 1 mL of 0.5N hydroxylamine

hydrochloride in 95% ethanol and 0.2 mL 6N

NaOH. Heat to boiling for 2-3 minutes, then cool

and add 2 mL 1N HCl. If the solution becomes

cloudy, add 1-2 mL of 95% ethanol to clarify. Add

1 drop of 5% ferric chloride solution. If a color

A deep burgundy color

is positive.


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forms and then fades, add additional drops of 5%

ferric chloride until the color persists. The color is

due to a complex between the hydroxamic acid

and the ferric ion.

Detection of Phenols using Ferric Chloride test: Just as enols can form colored complexes

with ferric ion, phenolate ions can as well. Therefore, this test is designed to convert the

weakly acidic phenols to their conjugate base which can then complex with ferric ion. Not all

phenols will give a positive test.


If the phenol is water soluble, add a few drops of2.5% aqueous ferric chloride solution to a 3%

aqueous solution of the phenol.

If the phenol is not water soluble, dissolve 20 mg of

the solid (or 1 drop of the liquid) in 1 mL of

methylene chloride and add 1 drop of pyridine. Add

3 drops of 1% ferric chloride in methylene chloride.

A deep red, green,or blue color is

positive.

An intense color is

a positive test.

Use phenol as a

known.

Detection of methyl ketones using Iodoform test : In this test you will convert the methyl

ketone to a triidomethyl ketone which is then cleaved to form iodoform, HCI3, a yellow solid.

Acetone gives a nice positive test so be certain that no traces of acetone are in your

glassware. The color of the iodine will disappear more slowly in the later additions.


In a large clean test tube or a vial, place 100 mg

of a solid or 5 drops of a liquid compound. Add 2

mL bis(2-ethoxyethyl)ether, 5 mL of water, and 1

After standing for 15

minutes, a pale yellow

precipitate of


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mL of 10% NaOH, and mix well. Add a total of 3

mL of iodine-potassium iodide solution in six

equal portions, stopper and shake well after each

addition. Caution: seal the tube carefully and

avoid skin contact with the iodine solution.

The solution should be slightly yellow. Heat if

necessary and shake again to force the iodine to

react. When the color is slightly yellow, add

water to nearly fill the test tube or container,

stopper, and shake vigorously.

The iodine-potassium iodide solution is preparedfrom 10 g of iodine and 20 g of potassium iodide

in 100 mL of water.

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