recrystallization and melting points

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4. Recrystallization and Melting Points PreLab - All PreLabs must be completed before coming to lab. You cannot start any

experimental work until the PreLab white pages are turned in! There are no exceptions to

this rule!This first experiment is quite lengthy because it involves two techniques, recrystallization and

melting point. Therefore, be forewarned that the PreLab is quite extensive and because it’s your

first one, will probably take about two hours to do. First, read through the Introduction and all

the Experimental sections given in this chapter. Then, read through sections 1 and 2 of Chapter

3 describing the PreLab writing format with examples. Finally, in your lab notebook, using

carbonless copies, write an introductory PreLab section for the experiments to be run including:

1. A brief Summary of what will be done and why it will be done.

2. Four Learning Objectives or Goals for the experiment.

3. Diagrams of Apparatus. There are no chemical reactions or equations in this experiment.

4. Complete Chemical Data Tables including ALL possible solvents and unknowns used in

Procedures 1 to 5, but not Procedure 6. Fill in the data on the Common Shelf Chemical

Data Table for the ten solvents designated Chapter 4 in column 3 of that table (includingwater). On the blank Chemical Data Table, fill in data for the 6 possible organic solid

unknowns, phthalic acid, and methylene blue (0.01% solution in water).

5. Do the PreLab Exercise given here.

a) If cooling fails to produce crystals, give two methods for inducing crystallization and

describe how they work.

b) Define immiscible and give one common household example of this phenomenon.

c) When choosing a solvent for recrystallization, what factors should be considered? (List

at least four criteria for choosing a solvent and explain where appropriate.)

d) What is decantation and when should you employ this technique?

The PreLab and Final Report for this first experiment will be graded on a, + ,- basis.

This means that the report will not be graded in the strict sense, but errors will be indicated and

comments for improvement will be given by your TA. Almost everyone will receive a or 100

points each on the PreLab and Final Report if they do a reasonable job. A few may receive a-

or 90 points if blatantly sloppy or inferior work is turned in. Also, a few may receive a+ or 110

points if their work is truly exceptional. The intent here is to use these first lab reports as feedback

to help your TA communicate to you what he or she is looking for in a typical PreLab and Final

Report . It is obviously very important for you to look over your TA’s analysis and comments and

discuss any areas you don’t understand so that you are better prepared to do a good job in writing

subsequent PreLabs and Final Reports which will be graded more rigorously.

The white pages of this section including the Chemical Data Table will be collected by your

TA at the beginning of the lab period. Tear out the appropriate Grading Sheet from the back of

the Lab Guide in (Appendix A.10) and attach it to the front of the PreLab. The point assignments

for each portion of the PreLab are shown below.

Before coming to

Lab

PreLab Exercise

Grading

GRADING FOR RECRYSTALLIZATION/MELTING POINT EXPERIMENT:

Possible

PreLab: Points

Date, Name, Desk #,Experiment # & Title (abbreviatedafter 1st page),Section& TA Name 8

Summary 16

Goals 16

Reactions and/or Diagrams of SpecialApparatus 20

Common Shelf Chemical DataTableand experimentChemicalData Table 20

PreLabExercise 20

TotalforPreLab 100



Solubility

Introduction

Italian salad dressing is a classic example of an

everyday solubility problem. It is clear from

observing the interaction between vegetable oil

and water, they are not compatible and do not mix

well. This phenomenon can be understood by the

simple principle “like-dissolves-like”. Onlymolecules of similar structure will be soluble in

each other and mix easily. Water (a small polar

molecule) and vegetable oil (a long saturated alkane

chain) are very different compounds and thus do

not mix well. However, methanol (CH3OH) or

ethanol (CH3CH

2OH) which are small organic

molecules containing a functional group that can

H-bond, are readily soluble in water. We would

find that the vegetable oil, which contains molecules

with large hydrocarbon chains are readily soluble

in hydrocarbon solvents such as ligroine or paint

thinner. These are examples of like-dissolves-like.

Table 4.1: List of common solvents by

decreasing polarity.

Solubility depends on the polarities and interactions between the solute molecules (the solid

to be purified) and the solvent molecules. For recrystallization, a solvent must be picked with a

polarity similar to the polarity of the compound you are attempting to purify. In general, the more

heteroatoms, such as nitrogen or oxygen, a molecule contains, the more polar the molecule. Table

4.1 lists some common organic solvents in order of polarity. This table can be used to select

solvents with polarities closely matching the polarities of your solute. It is quite unlikey that a

solvent with polarity that is very different from your solute will work at all, since you need to have

a least partial solubility in a solvent for it to be a possible recrytallizing solvent.

Chemicals are

found everywhere!!

Like-dissolves-like!Most Polar

Water (H20)

Acetic Acid (CH3COOH)

Methanol (CH3OH)

Ethanol (CH3CH2OH)Acetone (CH3COCH3)

Dichloromethane (CH2Cl2)

Chloroform (CH3Cl)

Ether (CH3CH2OCH2CH3)

Benzene (C6H6)

Toluene (C6H5CH3)

Hexane (CH3(CH2)4CH3)

Ligroin (mix of hydrocarbons)

Least Polar

Most of chemistry deals with the production of PURE chemicals for use in products we

purchase and use everyday. As you look around your kitchen or bathroom you see many of these

products. Can you tell which are pure compounds and which are mixtures? You should conclude

that most of the foods, cosmetics, cleaning agents or drugs are mixtures. Very likely, the only pure

single-compound materials in the kitchen are water (H2O), salt (NaCl), baking soda (Na

2HCO

3),

and sugar (sucrose, C12

H22

O11

). In the bathroom you might have some crystals of Epsom salts

(MgSO4) and some liquid rubbing alcohol (2-propanol), but these are about the only common

examples. Pills and capsules are all mixtures of a pharmaceutical compound and an inert binder

such as starch or gelatin. Many materials are really solutions, for example vinegar (5% acetic acid

in water) or hydrogen peroxide (usually 3% in water). Even though most things are mixtures,

many are formulated from pure compounds. For example, Anacin is a mixture of pure crystalline

aspirin and pure crystalline caffeine mixed with starch and pressed into tablets. The Food and

Drug Administration (FDA) sets very rigorous standards for purity of all the chemical components

that go into these mixtures or “formulations” sold as consumer products.

How do drug firms or the FDA measure the purity of organic compounds used in drugs? More

importantly, how do they purify compounds to meet these purity standards? This experiment will

answer these questions by demonstrating two common techniques: recrystallization for purifying

organic solids and melting point determination for testing the purity of organic solids. Mastering

these techniques will be critical to the success of your synthetic experiments in the second half

of this course. As with all techniques, mastery means not just learning the motions, but also

understanding the underlying principles involved. For this experiment, the principles are:

solubility (like-dissolves-like) including its variation with temperature and the melting behavior

of mixtures versus pure compounds.



The second factor affecting solubility is temperature. The temperature of the solvent directly

relates to the amount of material that can be dissolved. One example is making rock candy from

sugar. In order to obtain the saturated sugar solution needed to make rock candy, the water is

brought to a vigorous boil. More sugar will dissolve in the water when the temperature is

increased. This is a very important factor in recrystallization because these differences in

solubility at different temperatures will be used to our advantage.

The best solvent for any particular recrystallization should only sparingly dissolve thecompound at room temperature and easily dissolve the complex at higher temperatures. Figure

4.1 shows the solubility of a typical organic compound in three different solvents as a function

of temperature.

Temperature (°C)

A

B

C

S o l u b i l i t y

( w e i g h t i n g r a m s o

f s o l i d d i s s o l v e d p e r v o l u m e o f s o l v e n t )

Figure 4.1: Solubility differences of a solid in three solvents, A, B & C verses temperature.

At any given temperature, the compound has a very low solubility in solvent B. Solvent C is

exactly the opposite, its solubility is too great at all the temperatures. However, solvent A seems

to be a good choice because it has a low solubility at moderate temperatures, and the solubility

of the compound changes significantly as the temperature increases. This difference in solubilitywith small changes in temperature is an ideal solvent for recrystallization.

By understanding the solubility principles and the temperature factors involved, recrystallization

can be a useful method for purifying solid organic compounds in the laboratory. This easy,

convenient, and inexpensive purification technique can be used to remove small amounts of

impurities from solid compounds. Essentially the process of recrystallization breaks down into

five simple steps:

• Choosing a suitable solvent

• Dissolving the compound in a minimum amount of solvent

• Removing insoluble and/or colored impurities

• Crystallization

• Collecting and washing the collected solids

• Drying the crystals.

1. Choosing a Suitable Recrystallization SolventThe success of a recrystallization is based on the amount of pure solid crystals that can be

obtained using this technique. Ideally one would like to obtain 100% of the original weight of

material. This is rarely possible, but recovery of 80 to 90 % of the original weight can be obtained

if: 1) the “correct” solvent is chosen and 2) the recrystallizing technique is carried out correctly.

Selecting an appropriate recrystallization solvent is the first and normally the most difficult

step in the purification process. A successful recrystallization depends on the solubility

differences between the substance and the impurity. Recrystallizing a known substance makes

Six easy steps to

recrystallization.

Temperature plays

a large role in

recrystallization.



choosing a solvent of similar polarity easier. Reference textbooks, such as the Handbook of

Chemistry and Physics, can be used to find information on the solubility of organic compounds.

If the material you will be recrystallizing is unknown, than the process of choosing a solvent is

not as straight-forward and is mainly trail and error. Dissolve the unknown compound in a variety

of solvents, heat and observe the results.

Based on the principle of “like-dissolves-like”, which solvent, benzene or water, would be a

better candidate solvent for recrystallizing naphthalene? Which solvent would be better forbenzoic acid? *

Solids:

Solvents:

Benzene

O

HH

Water

Napthalene

C

O

OH

Benzoic Acid

* Benzene would be a better choice for recrystallizing naphthalene. Both of these compounds

are nonpolar organic molecules and the naphthalene should dissolve readily in hot benzene.

Water would be a better solvent for organic molecules having polar substituents such as alcohol

groups (sugar) or carboxylic acid groups (benzoic acid).

A Problem!

There are also several practical concerns when choosing a recrystallization solvent. For

obvious reasons, the solvent should not react with the solute. The solvent must have a boiling point

lower than the melting point of the compound. If the boiling point of the solvent is higher than

the melting point of the solid it will tend to “melt” instead of dissolve. This will lead to a very poor

crystallization. Chemists use the term “oiling out” to describe when a solid turns to an oil instead

of well-defined crystals. Oils are hard to handle and often difficult to crystallize. By choosing

a different recrystallization solvent, oiling out can normally be avoided. In addition, a non-toxic,

inexpensive, volatile and nonflammable solvent is the best choice. However, since most solvents

do not fit all these requirements, as the chemist you will have to find the best compromise for your

specific recrystallization.

In an ideal system, the compound should be completely soluble in the hot solvent and quite

insoluble in the cold solvent and the impurities should be soluble at all temperatures. This idealsituation is never obtainable. A single solvent system is always the best for recrytallization, but

sometimes this cannot be found and therefore a two solvent system may be necessary. When

choosing a two solvent system, the two solvents must be miscible. Miscible means capable of

being mixed in all proportions. Table 4.2 summarizes the miscibility of some common solvents.

Notice that with only one exception (can you find it?) organic solvents are miscible with each

other, but many are not miscible with water.

Some typical two solvent systems include; ethanol-water, methanol-water, acetone-ligroin,

ethanol-toluene and acetic acid-water. (Notice that each of these combinations are miscible.)

When using two solvents, it is best to find one solvent that will completely dissolve the compound

to be recrystallized at high temperatures and the second that dissolves it poorly at high

temperatures. Dissolve the compound in a minimum amount of the good solvent (dissolves

compound well at the boiling point), then add the bad solvent (dissolves the compound poorly)

dropwise until the solution becomes slightly cloudy. Reheat the solution until clear and then cool

slowly. Upon cooling, the compound should crystallize. If you are having problems crystallizing

the compound, reheat and add more of the poor solvent. The complex should eventually

precipitate, although this can often lead to the production of an oily non-crystalline mass.

The Answer!



Miscibility

A c e t i c a c i d

A c e t o n e

B e n z e n e

C h l o r o f o r m

D i c h l o r o m e t h a n e

E t h a n o l

E t h e r

H e x a n e

L i g r o i n

M e t h a n o l

T o l u e n e

W a t e r

Acetic Acid - M M M M M M M M M M MAcetone M - M M M M M M M M M M

Benzene M M - M M M M M M M M I

Chloroform M M M - M M M M M M M I

Dichloromethane M M M M - M M M M M M I

Ethanol M M M M M - M M M M M M

Ether M M M M M M - M M M M I

Hexane M M M M M M M - M I M I

Ligroin M M M M M M M M - M M I

Methanol M M M M M M M I M - M M

Toluene M M M M M M M M M M - I

Water M M I I I M I I I M I -

Table 4.2: Miscibility of Common Solvents (M = miscible, I = immiscible)

2. Dissolving the Compound

Miscible means

capable of being

mixed.

Once a suitable single solvent or solvent pair has been found, the impure material is dissolved

in the solvent with heat. The most common mistake when dissolving the compound is adding too

much solvent. For an easy purification, adding the minimum amount of hot solvent is essential.

The minimum amount is completely dependent on the quantity of material and the solubility of

that material in the solvent used. If the compound’s identity is known, the solubility data can be

used to calculate the approximate amount of solvent needed. For example, when recrystallizing

phthalic acid, use the data in Table 4.3.

Water Ethanol Other

0.5414 11.715 0.6915, Ether

0.7025 - 20.421, MeOH

18.0100 27.478 insol, CHCl3

Table 4.3: Solubility data for phthalic acid (in grams of solute per 100 mL of solvent.)

The superscripts refer to temperature in °C.

This solubility data can be found using the Beilstein Crossfire searching program or the

Handbook of Chemistry and Physics (CRC). Later in this chapter, we will use this data to

understand the more quantitative aspects of recrystallization.This data is convenient for a known complex, but if your substance is unknown or has been

synthesized for the first time, this data would not be available. In this case, the amount of solvent

needed to dissolve the compound is determined experimentally. Begin with just enough solvent

to cover the material and heat. Once the solute-solvent mixture comes to a boil, continue to add

hot solvent. (Heating a small amount of the recrystallization solvent at the same time will avoid

cooling the solution when adding additional solvent.) If a good recrystallization solvent was

selected, the solid will completely dissolve at the solvent’s boiling point. The substance is

completely dissolved when the solution is translucent . If the complex is white, then it will be a

clear translucent solution. However, if the compound is yellow, then the solution will be a

Use Beilstein or the

CRC to find

solubility data.



translucent yellow solution. Recognizing when the solid just dissolves completely is very

important. DO NOT add too much solvent. This will cause problems when attempting to collect

the solid. The goal when heating this mixture is to create a saturated solution in hot solvent. (Just

like our rock candy example.) By creating a completely saturated solution, the less soluble

compound will crystallize upon cooling. If an appropriate solvent was selected, your compound

will crystallize with ease. Too much solvent will hold more solid on cooling and it may not

crystallize. However, if too much solvent is added, there is a simple solution. Heat the solutionto the boiling point and evaporate some of the solvent to reduce the volume of the solution. It is

very easy to add too much solvent while learning to recrystallize organic compounds. This is a

way to correct your mistakes without starting over or wasting material.

One practical aspect of dissolving the compound in heated solvent is how to avoid a

phenomenon called “superheating” or “bumping”. Superheating is when the temperature of the

solvent reaches a temperature greater than that of the boiling point of the solvent. When this

occurs, the solvent will very rapidly and sometimes violently bump or splash. Bumping is the

sudden explosive boiling of the liquid near the bottom of the flask. Avoid bumping by using a

Teflon boiling chip (small white chips found on the Common Shelf ) or a wooden boiling stick (also

found on the Common Shelf ). These two methods allow the solution to heat evenly and boil

smoothly.

3. Removing Insoluble and Colored ImpuritiesOnce a hot saturated solution is obtained, there may still be some remaining particles, including

dirt, sand, or insoluble impurities. If these particles are dense and settle to the bottom, the liquid

can often be poured off or decanted from the solid. If they don’t settle to the bottom, there are

several methods to remove these insoluble materials. When working on small scale, pipet

filtration is a good method. To make a filter pipet, carefully stuff a small portion of cotton

(Common Shelf) into a disposable glass pipet as shown in Figure 4.2. Some chemists also add

about 1/2” of silica or alumina.

Figure 4.2: Pipet filtration

for insoluble materials.

Transfer the hot solution into the filter pipet with another pipet. The insoluble materials collect

in the glasswool. Be careful though; since the filter pipet cools the solution, crystallization can

occur inside the pipet. One method to avoid precipitation inside the pipet is to heat some pure

recrystallization solvent to the boiling point and run this hot solution through the pipet first. This

will increase the temperature of the pipet and glasswool decreasing the chance of precipitation.

Another way to insure that the compound will not crystallize inside the pipet is to use extra solvent

and create a more dilute solution. This will decrease the chance of crystallization at cooler

temperatures. After collecting the dilute solution, you will have to evaporate the extra solvent to

initiate crystallization.

If you are working on a larger scale, use the hot filtration apparatus shown in Figure 4.3. As

in the microscale pipet filtration, it is essential to have the entire flask and funnel heated. This can

be accomplished by heating a small amount of recrystallization solvent in the bottom of the

Erlenmeyer flask. This allows the vapor to rise warming the flask, funnel, and filter paper. If the

Use boiling chips or

a boing stick to

avoid bumping!!

Pipet filtration of

insoluble particles.

Reaction tube to c olle ct mother liquor

Pour in dilute solu tio n

Cotton used to collect insoluble materials

Silica gel or alumina plug or layer on top of cotton



Figure 4.3: Hot filtration set-up

4. Crystallization

Once you have obtained a clear saturated solution in hot solvent free of insoluble or colored

impurities, it is time to crystallize your solid product. Crystallization can sometimes occur within

seconds and other times it may take days. A perfect crystal needed for X-ray crystallography

requires very slow crystallization in order to obtain a crystal large enough to analyze. It should

also be mentioned that crystallization is different from precipitation. Precipitation is defined as

the formation of a solid product from a reaction that separates from solution. Precipitation

normally occurs very rapidly and results in very fine crystals. Crystallization is similar to

precipitation in that you separate a solid from the remaining solution; however, no chemical

reaction occurs and crystallization normally refers to the slow growth of crystals. Slow crystal

growth yields purer crystalline formations. There are three ways to initiate crystallization;cooling, scratching and seeding.

Slow cooling is the easiest method and works in most cases. Simply allow the hot saturated

solution to cool slowly without agitating or disturbing the solution in any way. Allow the hot

solution to cool to room temperature by placing the flask on a book or paper towel, then move it

to the cooler bench top and finally put it in a beaker of ice. Setting a beaker on an insulating surface

such as a book or placing a test tube in a beaker of glasswool will slow the heat losses and the

cooling process. The slower the solution cools, the more likely it will form perfect crystals which

can exclude the impurities since these will not fit in the crystal structure. Alternatively, you may

cool the solution slowly to room temperature and then place the solution in the refrigerator until

entire apparatus is not warm, the product will crystallize along the inside of the filter paper.

These filtration methods will remove insoluble materials from the solution. However, if you

have soluble colored impurities, Norit® pellets can be used to remove them. Norit® is the

commercial name for activated carbon. When mixed with the solution, the colored materials

physically absorb onto the Norit® pellets. You may have to heat the mixture and repeat the

procedure more than once. (Note that even very small quantities of colored impurities can produce

highly colored solutions.) The benefit of using the Norit® pellet form of activated carbon insteadof finely powdered charcoal when removing colored impurities is that the pellets can then be

easily removed by pipet or gravity filtration as described above. In other words, the soluble

colored impurities stuck to the Norit® pellets are now an easily removable insoluble impurity.

Remember, however, that some of your product will be absorbed so don’t use too much Norit.

Obviously filtration and/or treatment with activated carbon are only necessary when there are

insoluble or colored impurities present. Most recrystallizations will not require either step.

Heating the entire

hot filtration

apparatus is

essential to the

process.

Slow colling leads

to the best crystals.



the next lab period. (Note: All solutions placed in the refrigerator must be in capped and labeled

containers!) This slow cooling process normally yields beautiful crystals.

Sometimes it is difficult to obtain crystals even after the solution is cooled in an ice bath. In

this situation crystallization can be induced by scratching the inside of the beaker or flask with

a glass stirring rod. This will produce microscopic fragments of glass that may act as surfaces on

which crystal growth can begin.

If slow cooling and scratching does not produce crystals, crystallization may be induced byseeding. Seeding is accomplished by taking a small crystal from the original solid and dropping

it into the solution. This will provide a nucleation site for similar crystals to grow. This method

works well if you are having problems with a solid that separates without crystallization or “oils”

out. However, by adding some of the original material, you are inevitably increasing the amount

of impurity in the resulting solid. Seeding is normally used only as a last resort.

One crystallization method that is not encouraged is evaporation. Naturally, if you are having

trouble crystallizing a solid, evaporating the solvent completely will recover the material.

However this method has defeated the purification process. Not only have you recovered the pure

material, but also the impurities. When this is done, the recrystallization procedure must be

carried out again.

5. Collecting and Washing the CrystalsThe filtration method used to separate the crystals from the remaining solution and “collect”

the solid depends largely on the amount of crystals and solution. For recrystallizations involving

1 g to 100 g or more, gravity filtration through filter paper can be used. For recrystallizations

involving at least 0.1 g up to 100 g or more, organic chemists use a fast and effective filtration

method called vacuum filtration. But for the small quantities of 10 - 100 mg in a few mL of

solution, pipet filtration is quick and effective.

The first method, gravity filtration is used when there is at least a half a gram of large, well-

defined crystals. A glass funnel fitted with fluted filter paper is used to collect the crystals. With

very large amounts of solid, this process can consume a lot of time.

A quicker and more efficient method for this same quantity of material is vacuum filtration.

Vacuum filtration is probably the most common filtration type used by chemists. The funnel used

for this method called a Hirsch funnel if small or a Buchner funnel if large. These funnels have

a large flat plate inside which can either be made from ceramic materials that contains lots of smallholes through it or from a porous polyethylene disk.

Figure 4.4: Vacuum filtration

apparatus using a Hirsch funnel.

To use these funnels, a round piece of filter paper of

a matching diameter is placed over the perforated or

porous plate. The Hirsch or Buchner funnel is put into

a vacuum filter flask with a rubber adapter in between to

assure a vacuum tight seal. The filter flask has a

connecting tube or “side arm” which is connected to the

house vacuum or water aspirator. Wet the filter paper

with a small amount of solvent or recrystallization

solution to prohibit loss of product under the filter paper.

Once the vacuum is turned on, the filter paper is sucked

down on the plate so that it seals uniformly. The crystalsand solution are swirled or agitated and quickly poured

onto the filter funnel. The vacuum quickly draws the

liquid through the filter paper disk, leaving the crystals

behind. Any crystals remaining behind in the beaker are

scraped out with a spatula and/or washed out with some

of the recrystallizing solution from the filter flask or with

small amount of cold recrystallizing solvent. This step

requires great care. Using too much solvent that is not

cold enough can dissolve a lot of solid product, thus

decreasing your recovery and yield significantly.

The filtration

method used

depends largely on

the amount of

crystals collected.

Wash your

crystals with care!



One variation on a traditional vacuum filtration apparatus is shown in Figure 4.5. The mother

liquor is the filtrate or liquid left behind after removing solids from a solution. In this set-up, the

mother liquor is collected in a clean beaker and can be used again. Often during a recrystallization,

the first collection process will not crystallize all the material from the solution. By collecting the

filtrate, the remaining crystals can be retrieved. After the first filtration, concentrate the mother

liquor by evaporating the solvent to approximately half the original volume. Upon cooling,

additional crystals, called a second crop, may appear.

Figure 4.5: Bell Jar with vacuum filtration used to collect the mother liquor.

To use this set-up, first apply a light coating of stopcock grease to the bottom edge of the bell

jar. Seat the Hirsch funnel in a rubber filter adapter and place in the top of the bell jar. Connect

the side arm of the bell jar to the house vacuum using heavy-walled vacuum tubing. Pour the

solution carefully onto the filter. Once you have collected the solid, turn off the vacuum and allow

air to leak slowly into the jar. Carefully disconnect the tubing being careful not spill the contents

of the filter apparatus or beaker. If there is not sufficient time to concentrate the mother liquor

by evaporation, than simply setting the beaker in the back of the hood for slow evaporation. This

Figure 4.6: Pipet Filtration of Solid Product.

will often yield crystals without effort. If crystals are obtained from this second

crop, the crystals may not be as pure as

the first crop because there were a higher

percentage of impurities present during

the second attempt. A melting-point

determination can be used to determine

the purity of both crops.

For small quantities of materials, pipet

filtration is the preferred method. Figure

4.6 shows the pipet filtration method for

collecting solid products.

When you have solid material left-

over in a reaction tube or small test tube,there will be less product loss if you

remove the solvent and leave the solid

behind. This can be easily achieved by

pipet filtration. Select a pipet with a

perfectly square tip (when looked at from

the side), and without chips on the tip.

Push out all the air by squeezing the pipet

bulb. Slowly lower the pipet into the

solution and ease your way to the very

Filtration method

used to obtain a

second crop of

crystals.

Square tip pipet

Pipet filtration of a

solid product.



bottom of the reaction tube. Once the pipet tip is squarely and firmly pressed against the bottom

of the tube, slowly release the pipet bulb while keeping the pipet tip pressed against the bottom

of the tube. The solvent will be slowly drawn up into the pipet and the product should remain

behind. Initially you may see a small amount of product move up the pipet. Don’t panic! This

is a normal loss for this technique. This technique is a little trickier than vacuum filtration and

normally requires some practice. Practice with some inert materials from theCommon Shelf , such

as sodium sulfate and hexane. The separation is easier if you have big, chunky crystals. Very finecrystals will be difficult to separate using this method.

After the crystals have been successfully separated from the solvent, washing the crystals with

cold recrystallization solvent will put the finishing touches on your purification. This will wash

away any last traces of soluble impurity that have been left on the surface of the crystals. It is

essential to use cold solvent; hot or room temperature solvent will dissolve too many of your

crystals. If you used gravity filtration or vacuum filtration to collect the crystals, drop the cold

liquid directly onto the crystals with a pipet. If you performed a pipet filtration, add the cold

solvent directly to the tube while immersing it in ice, stir the solvent and then remove the solvent

using pipet filtration as you did before. Careful pipet filtration will increase your percent recovery

of the product.

6. Drying the CrystalsThe easiest way to dry the collected crystals is to be patient and allow nature to take its course.

Leave the crystals in a container open to the air in your desk. Over a period of a few days,

depending on the boiling point of the solvent, they will dry. If there is not sufficient time to allow

this slow drying process or if you are using test tubes, reaction tubes or solvents with higher boiling

points, such as toluene or water, there are several other methods including vacuum drying or using

a stream of nitrogen. Pressing the solid between pieces of filter paper or paper towel is a simple

way to remove a major proportion of a solvent.

Vacuum drying can be easily accomplished with the desk equipment provided. If your product

is in a reaction tube from a pipet filtration, you can connect the yellow thermometer adapter to the

reaction tube and then on to a piece of thick-walled rubber tubing using the cut off barrel from a

1-mL plastic syringe. Connect the tubing to the house vacuum and allow it to dry. (Figure 4.7)

Alternatively, the reaction tube can be pushed directly into the opening of the thick-walled tubing.

Eveaporation is speeded up by clamping the reaction tube in a beaker of warm water. This methodshould only be used to dry residual solvent. Using a vacuum to remove large amounts of solvent

will be slow and could eventually destroy the vacuum pumps.

For volatile or low boiling solvents, placing your product under a stream of nitrogen may be

sufficient to dry your product. Each hood has its own nitrogen line. The outlet should be fit with

a plastic pipet attached to the nitrogen line by a piece of Tygon tubing. Simply place the end of

the plastic pipet over the solid to be dried or directly into the reaction tube. Be careful not to “blow”

all the product away with a very fast flow rate. Drying your product under a stream of nitrogen

works very well for very volatile solvents such as ether and dichloromethane.

Pipet filtration takes

practice and

patience.

Air Drying

Vacuum Drying

Nitrogen

Figure 4.7 Vacuum

drying solids in a

reaction tube.



Melting Behavior of Mixtures vs. Pure Compounds

Once you have purified the solid by recrystallization, its purity can be verified by determining

its melting temperature which is commonly called a melting point, but is really a melting range.

To better understand this behavior, the theory of melting point determination must be examined.

Butter is a mixture of many different types of fat or triglyceride molecules. When stored at 5°C

in a refrigerator, it is fairly solid. On warming to room temperature (~ 25°C) it becomes soft and

if allowed to stand in the warm sun (40°C), it will turn to a liquid. Thus, this mixture of organic

compounds “melts” over a very broad range of temperatures. Pure water containing only one type

of molecule, H2O, freezes (if we remove heat from the liquid) or melts (if we add heat to the solid)

at a single temperature, 0°C, which is defined as the melting point (mp) of pure water. Until all

the solid ice is melted to liquid water, the temperature will remain at 0°C no matter how hard we

heat the container. What would happen if we add an impurity, such as salt, to water? Everyday

experience with the effects of spreading salt on highways and sidewalks tells us that the melting

temperature of a water-salt mixture is lower than 0°C causing the ice to melt as it dissolves the

salt. This behavior is calledmelting point depression. Hopefully you are familiar with this from

general chemistry.

Of the 18 million organic compounds known today, only a few thousand are liquid or gaseous

organic compounds. The remainder are all solids, many of which will melt from just above room

temperature to about 350°C. Above this temperature, organic compounds just decompose and

char rather than melt. (Also, the glass in a thermometer starts to get red hot and soften.) Solids

with normal melting behavior are in the molecular weight range from 59 (e.g., acetamide

CH3CONH

2, mp 79-81°C) to roughly 1500 ( e.g. the antibiotic bacitracin, C

66H

103N

17O

18S, mp

221-225°C). Notice that while we talk about melting point and use the abbreviation mp for it, we

almost always report a melting temperature range of a few degrees as seen in the examples just

cited. This is an anachronism that chemists have perpetuated and we will continue this peculiar

habit.

As with water, we should expect that adding impurities to organic compounds would lower

their melting point also. In addition we observe that the melting range is broadened. Even small

amounts of impurities, ~1%, will lower and broaden the melting temperature range. This

behavior of solids is consistent and can therefore be used as a test of whether we have a pure

compound or one mixed with some impurities. Thus, a melting point (really melting range)determination is a quick and simple test of a compounds purity and drug companies will use this

and other analyses to assure the purity of any drug compounds they produce. For example, the

true melting point range for caffeine is 234 - 237 ° C. If the caffeine is obtained directly from tea

and is not purified, the observed melting point is much lower. In fact, crude caffeine extracted

from tea bags melts as low 180 - 220 ° C. This melting range is significantly different from the

true melting point and can be used to determine that it is impure.

Melting Point (Range) Determination

The melting point or freezing point of a compound is the physical state change from a solid to

a liquid or a liquid to a solid respectively. Each pure organic compound has a characteristic

melting point or range. This determination can be used to correctly identify and/or demonstratethe purity of the compound.

In addition to traditional melting points, compounds can sometimes decompose before

actually melting. Instead of a solid clearly turning into a liquid, the solid may turn gray, brown

or black at a certain temperature. This does not mean that you have obtained the wrong compound,

just that this particular compound is not stable at the observed temperatures. Saccharin for

example does not melt, but decomposes at 229°C. This is symbolized as 229 d . This may happen

to some of the compounds you purify throughout the semester.

There are over 18

million organic

compounds knowntoday.

Not all compounds

melt.



Purity can be determined by evaluating the melting point range. When an impurity is soluble

in the compound, the melting point will be depressed. Figure 4.8 is a Melting Point Composition

Diagram that shows how the melting temperatures are affected in a two component mixture.

Either component may be regarded as an impurity in the other component.

Mole percent of AMole percent of B

100 50 0 0 50 100

MP in °C of pure A

MP in °C of pure B

Ep

Et

Increasing MeltingTemperature

Figure 4.8: Melting Point Composition Diagram of a Two Component Mixture.

The left and right y-axis intercept the known melting temperatures for pure compound A and

pure compound B respectively. However, upon mixing A with B, the melting temperature of the

mixture is decreased. The amount the melting point is depressed is a function of the molar percent

of each component and can be read directly from the diagram. There are two special points on

this diagram. The first is the Eutectic temperature, designated by Et. This is the lowest

temperature at which the mixture will begin to melt. The Eutectic point, Ep is the mole percentage

where the mixture of two compounds are dissolved equally in each other. At this particular point,

even though the mixture is not pure, it will appear to have a sharp melting point. As discussedearlier, a sharp melting point indicates a pure material. A broad melting point is normally a

mixture of compounds. The Eutectic point is an exception to this rule and can lead to

misidentification of the sample. It should also be noted that insoluble impurities, although very

undesirable, will not affect the melting point. One example of an insoluble impurity would be a

piece of sand. The sand will have not affect on the melting point.

If you are determining the melting point of a known compound, the results can be compared

to literature values. Aldrich catalog has an enormous amount of physical data for organic

compounds including melting points. Use these and other sources to verify your melting points.

However, if the compound is unknown or has been synthesized for the first time, the melting range

may still help in determining the purity of the substance. The melting point range is the

temperature at which the solid begins to melt to the temperature when the solid has completely

been converted into a liquid. Melting points should always be recorded as a range. As previously

mentioned, the purer the compound is the narrower the melting point range. In contrast, the moreimpure the compound is the broader the melting point range. Melting point ranges can vary, but

a good range is typically 1-3 °C. Poor ranges can be as broad as 20°C.

One of the biggest and most common mistakes when obtaining a melting point, is not allowing

the sample to dry completely. Traces of solvent in the sample act just as any other impurity. It

will depress the melting point and broaden its range. This will lead to incorrect identification of

the unknowns you will be recrystallizing in this chapter. Another common mistake is to heat the

sample too fast. In general it is best to heat the sample slowly, no faster than 2°C per min in the

range of the melting point. Use this rate to expedite the use of the few melting point units we have.

Melting Point

Depression.

Solids must be free

of solvent in order

to obtain an

accurate melting

point.



In our laboratory, we use two different types of melting point apparatuses: Mel-Temps, located

at the end of every other laboratory bench, and Thomas Hoover’s located at the North end of Lab

205. Both are electrically heated melting point apparati. The directions for the Thomas-Hoover

and Mel-Temp Melting Point Apparati can be found in Step 3 of the experiment found in this

chapter. Since both the Mel-Temp and the Thomas Hoover machine use thermometers, it is very

important to calibrate the thermometers. Obviously, the identification of your unknown will be

incorrect if the thermometer reading is wrong. This is a very common and serious mistake duringthe determination of unknown samples. Please check to see that the thermometer is giving

accurate temperature readings by recording the temperatures of boiling water and ice. It is a

simple and straight-forward test you must complete before the melting point determination.

One advantage of a melting point analysis is the very small amount of sample needed to obtain

good results. Each machine uses glass capillary tubes to obtain the melting point. To load to

capillary, take a few milligrams of the material to be analyzed and push the solid together on a

watch glass or other clean surface. Use the open end of the capillary tube (found in your desk

equipment) and push the open end onto the pile of solid material. Once a small portion of the

compound is in the tube, turn the capillary over and rap the tube gently on the desk top to pack

the sample in the botton. Be careful not to break the tube accidentally. Even though the tubes are

small, the glass can be very dangerous and puncture the skin very easily. While rapping on the

desktop, the compound (if thoroughly dry) will move to the closed end of the tube. If this does

not occur, drop the mp capillary through a 3 foot long plastic tube held vertically on the benchtop.These plastic tubes can be found clamped to the side shelves. You only need a few millimeters

of sample in the end of the tube. Too much sample will lead to uneven melting and possibly

increase the range of your melting point. Insert the tube, closed-end down, into the Mel-Temp

or Thomas Hoover, heat it slowly, and observe the melting point range.

Although melting points can be used as an identification tool, there are many compounds that

have coincident melting points. There may be an occasion when you have two compounds that

have the same melting point. They could be the same compound or it may be a coincidence. A

mixed-melting point determination can verify whether they are the same compound or two

different compounds. Mix equal amounts of the two solids and determine the melting point of this

mixture. If the mixed-melting point is the same for both the mixed materials and the individual

compounds, they are the same compounds. However, since mixtures of different materials melt

at lower than either compound, if the combined melting point is depressed compared to the

individual samples, then the materials are not the same. This simple method can easily verify if these compounds with the same melting point are the same.

Mixed-melting

points.

Mel-temps &

Thomas Hoovers

only 1 to 2

mm of sample

in mp capillary



Procedures 1,2, and 3 provide challenging tests of your experimental and observational skills.

You will be given a small amount of an impure unknown organic compound whose identity can

be determined from its melting point. However, since its melting point is depressed by impurities,

you will need to purify the compound by recrystallization before the melting point determination.

Procedure 4 involves a more quantitative analysis of solubility in cold and hot solutions.Procedure 5 demonstrates how colored impurities might be removed from an organic solid.

Procedure 6 shows how the melting point can be used to help identify an unknown solid.

All students should complete Procedures 1, 2 & 4 in the first lab. Procedure 3,

determination of unknown melting points and Procedure 5, decolorization can be done in

the following lab period. Procedure 6 may be done in subsequent labs or in the instrument

room. In order to use your time efficiently, whenever possible try to do more than one

operation or experiment at a time.

CH3

O

O

OH

C

NH2

C

OH

CH2C

OH

CH3

O

O

Benzamidem.p. 127-130°C

6 5 4

3 21o-Toluic Acid

m.p. 104-107°Cm-Toluic Acid

m.p. 109-113°CFluorene

m.p. 114-116°C

E-Cinnamic Acidm.p. 133-134°C

E-Stilbenem.p. 122-124°C

Procedure 1 • Finding a Suitable Recrystallization Solvent for your Unknown

You will use a sand bath, not a Bunsen burner or steam bath, to heat solvents. As soon as you

arrive and while you’re waiting for your TA to collect your PRELAB pages, fill your heating

mantle half full of sand, mount it to a ring stand in your hood, and plug it into the heating controller

(often called a varistat or powerstat) which is mounted on the outside of the hood. Plug the

controller into the electrical socket above the bench and set the dial to 50.

To effect smooth boiling and avoid explosive boiling or bumping of the solvent when heating

it, be sure to use boiling sticks (wooden applicator sticks). Boiling sticks also can serve as stirrers.

Your instructor will provide you with approximately 100 mg of an impure unknown organic

solid (record the number of your unknown immediately into your lab notebook) which is one of

the six compounds shown below:Unknown

Compounds

NOTE: Set up and

heat your sand bath

as soon you arrive

in lab.

Recrystallization Experiments

Solubility Test

You will test the solubility of your unknown in different

solvents as follows:

Weigh out between 8 and 13 mg (0.008 to 0.013g) of your

unknown and transfer it to a tared reaction tube or micro test tube.

(This is about the amount that can be picked up on the last half inch

of the sharply pointed end of your stainless steel micro-spatula.)

Using a Pasteur pipet, add to this solid about 0.25 mL of distilled

water. You can use the graduations on the side of the reaction

tubes to estimate the 0.25 mL of solvent.

Graduated

Reaction Tube



Always make

careful

observations.

Cleaning Up

Alternatively, use your 1 mL x 1/100 mL pipette to measure out liquid volumes less than 1 mL.

To draw liquids into the pipet in a controllable manner, attached your 2.5-mL plastic syringe to

the pipet with a small rubber connector available on the Common Shelf . (Ask your TA for a

demonstration if you are unsure of how to use the rubber connector.) Volumes can be approximate

for the solubility tests; for 0.25 mL of water, for example, you can use a volume between 0.20

and 0.30 mL.

Agitate the tube vigorously for 10-30 sec by hitting or tapping the bottom with your fingers,or by crushing the large crystals with a stirring rod. If the solid dissolves, the unknown is too

soluble in the solvent and the solvent can be eliminated as suitable for recrystallization. If the solid

does not dissolve after mixing for a short period, place a boiling stick (wooden applicator) in the

tube and push the bottom of the tube part-way down into the hot sand of your sand bath. The

solvent should start to boil within 10-30 seconds. If the boiling becomes too vigorous, pull the

tube partially out of the sandbath. Continue gently boiling the solvent for about a minute, noting

how the vapors condense on the tube walls and run back down into the liquid. This is called

refluxing.

Carefully observe the mass of crystals while you’re refluxing the liquid. If the solid dissolves,

place the tube in a small beaker of ice water for one or two minutes to cool the solution. If you

observe solid crystals coming out of solution you have found a suitable solvent for recrystallization.

Go directly to Step 2 and recrystallize your unknown according to the instructions there.

Note: Be alert to the difference between melting and dissolving! Compounds that melt just

above the boiling point of water when pure may melt just below the boiling point of water when

small amounts of impurities are present (the maximum solvent temperature in these experiments

is approximately that of the boiling point of the solvent). Solids that melt but do not dissolve will

form two immiscible liquids in the tube; upon crystallization, little impurities will stay in the

solvent, so the degree of purification of the solid from this step will be rather small.

If refluxing in the solvent leads to an observable decrease in the mass of the solid, but not

complete dissolution, try adding more solvent, a few drops at a time, heating the solution to boiling

(refluxing) for 30 to 45 seconds after each addition to see whether all of the solid dissolves. If it

finally does, place the tube in an small ice-filled beaker for one or two minutes to cool the solution.

If you observe solid crystals coming out of solution you have found a suitable solvent forrecrystallization. Go directly to Step 2 and recrystallize your unknown according to the

instructions there.

If the mass of solid remains unchanged after one or two minutes in the refluxing solvent, then

the unknown is obviously insoluble in both cold and hot solvent, and the solvent is unsuitable for

recrystallization. Set the tube aside. With a clean tube, carry out the solubility test procedure as

described above with another 8 to 13 mg sample of your unknown and a different solvent. Try

the other extreme in solvent polarity, for the second test, ligroin (ligroin is the common name for

a mixture of hexane and cyclohexane isomers with boiling points between 60 and 90°C). If this

doesn’t work, try a solvent with polarities intermediate between water and ligroin such as

methanol. If you don’t find a suitable recrystallization solvent in 3 or 4 trials, talk with your

instructor.

Cleaning Up Place organic solvents and solutions of the compounds in the Nonhalogenated

Organics Disposal container. Dilute the aqueous solutions with water and flush down the drain.

Procedure 2 •• Microscale Recrystallization of your Unknown

Having found a suitable solvent for recrystallization in Step 1 above, recrystallize the

remainder of your unknown as follows:

Refluxing



Obtain approximately 80 mg of your unknown solid by weighing your unknown on the

electronic balance. Transfer the weighed unknown into a tared reaction tube. (See the Check-

in instructions in the back of the laboratory notebook for information on how to permanently mark

tare weights on your reaction tubes, if they haven’t been marked already.) Make sure you have

about 5 mg of the impure solid left to be used in a melting point determination in Step 3. Use the

plastic funnel from your red glassware kit can be used as a powder funnel to make the transfer of

the solid into the narrow neck of the reaction tube easier. If you do not have 80 mg, you’ll haveto obtain more from the stockroom. Please be sure to tell the stockroom personnel your unknown

number if you need more.

To the solid in the tube, add about 1.0 mL of the solvent you’ve found to be suitable for

recrystallization, place a boiling stick in the tube and heat to refluxing in the sand bath. If the solid

has not completely dissolved after a minute of boiling, add additional solvent, 5 or 6 drops at a

time, bringing the solution back to a gentle reflux for a brief period after each addition. Be sure

to allow time for as much undissolved solid to go into solution as possible at the boiling point of

the solvent. After all the solid has dissolved, set the tube into a test tube rack and allow the solution

to cool slowly. This will lead to slower crystal growth, yielding larger, more perfect, and purer

crystals. While this is cooling, put approx. 0.5 mL fresh recrystallization solvent in a small test

tube and cool in an ice bath. You will use this later to wash your purified crystals.

After about five to ten minutes of slow cooling, crystals should have started to form. If theyhaven’t scratch the inside wall of the tube with a glass stirring rod. When you are sure that the

solution is close to room temperature, cool it further by placing the tube in a small beaker of ice.

After a few minutes in ice, if necessary, break up any solid crystalline mass that may have formed

and rap the reaction tube bottom on a book or paper towel to force the liquid to the bottom. Remove

the solvent using the pipet filtration method described in Figure 4.6 and demonstrated to you by

your instructor. Add 6 to 10 drops of previously-cooled pure solvent to the crystals in the ice bath

and agitate to wash the crystals. Remove this wash solvent using the pipet filtration method.

If your recrystallization solvent was water, you should remove a small clump (about 5 mg, just

enough to fit on the tip of the spatula) of crystals and place them on an open watch glass in your

locker so that they can dry thoroughly by the next lab period. They will be used to obtain and

accurate melting point. If it was ligroin or other more volatile solvent, the solids can be left in the

open reaction tubes in your locker until the next lab period and weighed directly in the tube(subtracting the tare weight of the tube). This will reduce losses in of solid during transfers which

can be a problem at microscale. Alternatively, you can speed up the removal of the residual

solvent adhering to the solids by applying vacuum to the reaction tube while warming in a beaker

of warm water as shown in Figure 4.7.

When dry, simply weigh the tube containing the sample and subtract the reaction tube tare

weight as determined previously. Calculate and report the percent recovery in your final report.

A small portion of this solid will be used to determine the melting point of your purified unknown

in Step 3.

Cleaning Up Place organic solvents and solutions in the halogenated or nonhalogenated

organics waste container, whichever is appropriate. Dilute the aqueous solutions with water and

flush down the drain. Carry your sandbath to the side shelf and dump the HOT! sand back into

the sand supply bowl. Do NOT dump it in the waste bins.

Scratch or Seed

Your compound

must be dry to

determine the

correct melting

point.

Cleaning Up

Procedure 3 • Melting Points of Your Impure and Recrystallized Unknown

In this experiment, you will determine the melting temperature ranges of both your impure

unknown and your recrystallized, pure unknown. These will be determined using melting point

capillary tubes (found in your locker) and one of the “Mel-Temp” electrically heated melting

point apparati found on the shelf above your desk (1 per 4 students) or “Thomas-Hoover” type

found on the side shelves. (Please don’t move the ones on the side shelves.) Two students can

Obtain the weight of

your unknown.



use one Thomas-Hoover apparatus at the same time, since 4 melting point tubes can be placed in

the heating oil bath of this apparatus, and the melting points determined for all 4 samples by

heating from low temperature, say 70°C, to high temperature, say over 130°C, in one pass. It

should take 20-30 min for each temperature “scan”, including the time necessary for the apparatus

to cool down. All students should be able to determine their mp’s in one three-hour lab.

Turn the power on, but turn the heating rate control

dial all the way counter-clockwise to zero. If the

temperature of the aluminum block is more than 70°C

as read on the thermometer, wait for it to cool see the

temperature on the thermometer rise. At this point you

need to keep an eye on both the mp tubes (to watch for

any signs of melting) and the thermometer (to monitor

the rate at which the temperature is rising). [It helps to

have chameleon eyes!] You may need to adjust theheating rate control dial so that the temperature does not

rise too rapidly.

It is best to heat the samples slowly, no faster than

1°C per min. However, it is possible to get useable mp’s

heating at 2 to 3°C per minute, and to expedite the use

of the few mp units we have, try this faster rate. Watch

the samples, and when you start to see any one of them

start to melt, record the temperature. When the solid in

Mel-Temp

It is of utmost importance that your sample be thoroughly dry and free from all traces of solvent

before trying to determine its melting point. Otherwise, the solvent, like any other impurity, willcause the melting point to be depressed, and you won’t obtain an accurate melting point.

Normally, most solvents evaporate completely from crystals in an open reaction tube in the two

to five days between labs, but higher boiling solvents such as water (bp 100°C) and toluene (bp

110°C) may require drying on an open watch glass for a day. If your solid still seems wet and

sticky after this period of time, you may have to allow additional time for drying by leaving it open

in your locker for another few days, or you may have to speed up the evaporation process by

applying vacuum. Assuming your crystals are dry, proceed as follows:

Remove a few mg of the purified unknown and place it on a watch glass. If it is lumpy, grind

it into a fine powder using the flat end of a spatula. Push the crystals into a small pile and then

push the open end of a melting point capillary tube into the crystals so that some solid is forced

into the opening of the tube. The sample is then shaken down into the closed end of the capillary

either by: 1) wrapping the capillary tube with a paper towel and rapping sharply onto a hardsurface; or by: 2) dropping it down a 2 to 3-ft length of 5-10 mm I.D. tube onto a hard surface.

(there are a few of these tubes on the side shelves) If necessary, push more crystals into the open

end and pack it down into the closed end so that you have no more than 2-3 mm column of solid

at the bottom of the mp. tube.

In a second melting point capillary, load a small amount of the unrecrystallized and impure

unknown that you retained from Step 2 in the snap-top plastic vial. Mark the tubes so they can

be differentiated.

When a melting point apparatus is available, check that it has cooled to below 70°C. Then

you (and another student in the case of the Thomas Hoover units) can insert two samples and take

their melting points using one of the following procedures, depending on the melting point

apparatus. Be sure to bring your notebooks with you to the mp apparatus so that you can record

the mp data

Loading the melting

point capillary.

Your unknown must

be solvent free.

Figure 4.12: Mel-Temp Apparatus

Mel-Temp Melting Point Apparatus Procedure

only 1 to 2

mm of sample

in mp capillary



Thomas-Hoover Melting Point Apparatus Procedure

Turn the power on, if it isn’t already. Turn the heating rate control dial all the way counter-

clockwise to zero. Make sure the stirrer speed control dial is set at a little more than half of full

speed. If the temperature of the oil bath is more than 70°C as read on the thermometer, wait for

it to cool below this. Once below 70°C, you and another student can place your two mp tubes

gently into the holes above the oil bath. Now start heating the oil by turning the heating rate control

dial to about 3. After a moment, you should start to see the temperature on the thermometer rise.

At this point you need to keep an eye on both the mp tubes (to watch for any signs of melting) and

the thermometer (to monitor the rate at which the temperature of the oil bath is rising). [It helps

to have chameleon eyes!] You may need to adjust the heating rate control dial so that the

temperature does not rise too rapidly.

Thomas-Hoover

Figure 4.13: Thomas Hoover Apparatus

It is best to heat the oil bath slowly, no faster

than 1°C per min. However, it is possible toget useable mp’s heating at 2 to 3°C per

minute, and to expedite the use of the few mp

units we have, try this faster rate. Watch all the

samples, and when you start to see any one of

them start to melt, you or your co-worker

should record the temperature and the position

of the tube, for example #3 (from right) - 86.5°.

If that is one of your partner’s tubes, it would

be nice if you could let them watch the tubes

while you keep tabs on the temperature and

rate of heating and then change places again

when yours starts to melt. When the solid in a

particular tube has completely melted, record

this temperature also (87-88°C, for example).

Once all samples have melted, you should

have recorded melting ranges for each. Turn

the heating rate control dial to zero, but do not

turn off the power or stirrer, since stirring

helps to cool the oil bath down faster.

a particular tube has completely melted, record this temperature also (87-88°C, for example).

Once both samples have melted, you should have recorded melting ranges for each. Turn the

heating rate control dial to zero and turn off the power. Cooling of the aluminum block can be

expedited by holding a moist (not sopping wet!) paper towel on it, but be careful not to burn your

fingers.

Proof that an unknown compound is identical to known compound can be obtained by taking

a mixed melting point. If two organic solids are thoroughly mixed together, two possibilities are

observed:

1. If they are identical, the mp will be unaffected.

2. If they are different, each acts as an impurity in the other, and the mp will be considerably

depressed.

For example, you could verify that your unknown is fluorene by grinding it together with an

authentic sample of fluorene using a glass rod on a watch glass. If the melting point is still 114-

116°C, your unknown is indeed fluorene; if lower, it must be something else. Occasionally, the

melting point range of your purified unknown is still more than a few degrees and in between two

possibilities. If this is the case, ask you TA to provide you with small amounts of pure samples

of each of the possibilities, grind a small amount of your unknown together with each, and take

the mp’s. Your unknown is then identical to the authentic compound that doesn’t effect the mp

of your unknown when mixed with it.

Mixed Melting Points If your mp data

isn’t definitive, try

this!



Once cooled close to room temperature, cool the flask in an ice bath with occasional swirling.

If the crystals have formed a solid mass, break this up with a spatula or boiling stick. Swirl to make

sure everything is in suspension and then pour the solid-liquid slurry into the Hirsch funnel. Thewater solution should be rapidly sucked through the filter. Rinse out any crystals remaining in

the flask into the Hirsch funnel with about 1 mL of ice water. Press down on the crystals in the

Hirsch funnel with a spatula to squeeze out as much water as possible. Using a stainless steel

spatula, scrape the crystals and filter paper disc out of the funnel onto a tared beaker or onto a tared

watchglass. Remove the 1.3-cm filter paper disc and discard it. Allow the crystals to dry in your

locker until the next lab period and then weigh them. For your Final Report, compare the

calculated volume of water with the volume of water actually used to dissolve the acid. Also,

calculate the percent recovery of dry, recrystallized phthalic acid.

Collecting the

purified crystals.

The large difference in solubility in water as a function of temperature suggests this as the

solvent of choice. The solubility in alcohol is too high at room temperature. Ether is difficult to

use because it has such a low boiling point and boils away rapidly; the compound is insoluble inchloroform (insol. CHCl

3).

Crystallize 1.0 g of phthalic acid from the minimum volume of water, using the above data to

calculate the required volume. (This calculation should appear in your Observation and Data

Section.) Add the solid to a 25-mL Erlenmeyer flask and then, using a Pasteur pipet, add the

amount of water calculated above to the flask from a full 10-mL graduated cylinder. A boiling

stick (a wooden applicator stick) facilitates even boiling and will prevent bumping. After a portion

of the water has been added, gently heat the solution to boiling on a shallow sand bath or a hot plate.

(A hot plate can be shared between 3 or four people.) As soon as boiling begins, continue to add

water in small portions, heating to boiling each time, until all the solid just dissolves. Record the

total amount of water used to do this.

Place the flask on a book or other insulator and allow it to cool undisturbed to room

temperature, during which time the crystallization process can be observed. Slow cooling favors

larger and purer crystals. While this is taking place, set up for suction filtration by clamping your

25 mL filter flask to a ring stand, and connecting the vacuum hose from the house vacuum outlet

to the side arm of the filter flask. Place the white plastic Hirsch funnel into the filter flask and cover

the porous disc with a 1.3 cm filter paper disc from the Common Shelf (see Figure 4.4) The filter

paper keeps the porous plastic frit from getting clogged with small particles. Open the vacuum

valve and place the palm of your hand firmly over the top of the Hirsch funnel to see if you can

feel a vacuum. If you don’t, there may be an air leak or the vacuum system isn’t working properly.

Consult your TA if you can’t figure out the problem. An alternative filtration method is also

demonstrated above in Figure 4.5.

Don’t forget to

collect a second

crop of crystals.

Water Ethanol Other

0.5414 11.715 0.6915, Ether

0.7025 - 20.421, MeOH

18.0100 27.478 insol, CHCl3

Table 4.3: Solubility data for phthalic acid (in grams of solute per 100 mL of solvent).

The superscripts refer to temperature in °C.

Procedure 4 • Macroscale Recrystallization of Phthalic Acid

The more quantitative aspects of recrystallization will be studied using phthalic acid. Using

the Beilstein Crossfire as a reference the following solubility data (in grams of solute per 100 mL

of solvent) can be summarized. The superscripts refer to temperature in °C:

Quantative aspects

of recrystallization.



More phthalic acid crystals (a second crop) can be obtained from the filtrate (mother liquor)

by concentrating the solution. Pour the filtrate from the first crystallization above into a 100-mL

beaker and either let the solution evaporate in your desk until the next laboratory period or boil

the solution to at least half the volume. Collect the second crop of crystals by suction filtration,

wash the crystals with about 1 mL of ice water, and air dry the crystals on tared watch glass until

the next lab period. Weigh the watchglass and determine the yield and percent recovery based

on the original sample of phthalic acid.

The technique used for recovering the second crop of crystals is a general approach in synthetic

chemistry. Occasionally, you may find that you’ve formed an unsaturated solution that is so

dilute, that you have to boil away a lot of solvent before it becomes concentrated enough that

cooling gives a supersaturated solution. You may also find that upon crystallization, there is a

great loss of product; as long as you do not throw out your product, the solvent can always be

removed and repeat the crystallization procedure using less solvent.

Cleaning Up Phthalic acid crystals can be recycled in the specially marked recycle containers

on the hooded side shelves. Please don’t put filter paper in the recycling bottle. The water

solutions from the phthalic acid recrystallization can go down the drain. Phthalic acid is not

considered toxic to the environment. Carry your sandbath to the side shelf and dump the HOT

sand back into the sand supply bowl. Do NOT dump it in the waste bins. (Why is this a bad idea?)

Cleaning Up

Into a reaction tube place 1.0 mL of a 0.01% solution of methylene blue dye. Add to the tube

a few pieces (8 or 10) of pelletized Norit decolorizing charcoal. Shake, and observe the color over

a period of a minute or two. Add a boiling stick, heat the contents of the tube to boiling (reflux)

and observe the color by holding the tube in front of a piece of white paper from time to time. How

rapidly is the color removed at room temperature versus heating to higher temperature? If

the color is not removed after heating for a minute or so, add more pelletized Norit charcoal, a few

pieces at a time, until the color does disappear.

Cleaning Up Place the Norit in the waste bin (the nonhazardous solid waste container.) The

water solutions can go down the drain.

Procedure 5 • Removing Colored Impurities Using Activated Carbon

Cleaning Up

Decolorization

Obtaining a second

crop.

Procedure 6 • Melting Point Determination of your Spectral Unknown

Obtain the melting point of your spectral unknown. Assume that the sample is pure and does

not need to be recrystallized. You do not necessarily need to do this during the two lab periods

scheduled for this recrystallization/melting point technique experiment. If the melting point

apparati are tied up, you can take your melting point during the distillation technique lab periods

or in the instrument room any time it is open. If you plan to do it outside your regularly scheduled

lab time, put your unknown sample tube in a labeled ziplock bag (available on the Common Shelf)

and place it by the melting point apparatus in the instrument room or carry it with you.

Use only as much sample as necessary. Remember that you need most of your sample for NMR

and IR analysis.

Once you have determined the melting point, look at the Unknown Possibilities by Melting

Point list located on the Web at:

http://courses.chem.psu.edu/chem36/HTML/Experiments/Unknown_Lists/Unknown.html

In you Final Report, give the name and draw the structures for all compounds that have

melting points within plus or minus 5 °C of your spectral unknown’s melting point.

Spectral Unknown

melting point



Summary of In-Lab and Post-Lab Points for Final Report

OBSERVATION /DATA - Accuracy and completeness 8

Solubility test data; Weighing data; Melting point data, Calculated water volume for

Phthalic Acid Recrystallization, Decolorizing Ex data and observations

10

RESULTS/DISCUSSION - Overall organization, readability, completeness. 8

Identification of Unknown with explanation 12

Yield and % recovery for unknown 8

Yield and % recovery of phthalic acid (firstcrop) 8

Yield and % recovery of phthalic acid (second crop) 4

Calculate solubility of phthalic acid in hot waterfromyour data and Table 4.3 8

Discussion of the volumeof water used vs calculated 8Solution decolorization results and discussion 6

SpectralUnknownmeltingpoint and possiblecompounds' names andstructures 8

PostLab Questions 12

Grading

Follow the guidelines given below in the Summary for In-Lab and Post-Lab Points to write a

thorough discussion of recrystallization, solubility, and melting point determination. Include all

in-lab observations and data and calculations or answers to questions shown in bold to the above

experiments. Answer the PostLab Questions below.

PostLab Questions:

1) Under what circumstances might it be necessary to use a mixture of solvents to carry out

crystallization? What characteristics should these two solvents have?

2) What is the effect of aninsoluble impurity, such as sodium sulfate, on the observed melting

point of a compound?

3) Strictly speaking, why is it incorrect to speak of a melting point?

Final Report

A copy of the In-Lab and Post-Lab portions of the Recrystallization/Melting Point Grading

Sheet is givenbelow

PostLab Questions

recrystallization and melting points

Documents