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Transcript of Lab Manual
The University of Western Ontario Faculty of Engineering
DEPARTMENT OF CHEMICAL AND BIOCHEMICAL ENGINEERING
INDUSTRIAL ORGANIC CHEMISTRY I
CBE 2206A LABORATORY MANUAL
Instructor: E. R. Gillies
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TABLE OF CONTENTS
1 INTRODUCTION .................................................................................... 3 1.1 CONVERSION FACTORS .................................................................... 3 1.2 COMMON ITEMS OF GLASSWARE AND APPARATUS ..................... 5 1.3 A NOTE TO THE STUDENT .................................................................. 9 1.4 SAFETY GUIDELINES ........................................................................... 9 1.5 GENERAL HOUSEKEEPING .............................................................. 14 1.6 GUIDELINES FOR PREPARATION OF LABORATORY REPORTS ... 16 LABORATORY 1 - THE FRACTIONATION OF HYDROCARBONS AND GAS CHROMATOGRAPHY (GC)……………………………………….……..21 LABORATORY 2- ESSENTIAL OILS AND STEAM DISTILLATION...…….……..29 LABORATORY 3- PURIFICATION OF ORGANIC COMPOUNDS BY EXTRACTION AND RECRYSTALLIZATION…….……………………….….……...35 LABORATORY 4- SN1 AND SN2 REACTIONS……………..………………………46 LABORATORY 5- HYDROGENATION OF A FATTY ACID..………………………52 APPENDIX ….……………………………………………………………………..……58
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INTRODUCTION
1.1 CONVERSION FACTORS FOR SOME COMMONLY USED UNITS OF MEASUREMENT
Acceleration m/s2 ft/s2 3.281X100 3.048X10-1
g (accel. of gravity) 1.020X10-1 9.807X100 Area m2 ft2 1.076X101 9.290X10-2 inch2 1.550X103 6.452X10-4 yard2 1.197X100 8.361X10-1 Density kg/m2 g/cm 1.000X10-3 1.000X103 lb/gallon 8.345X10-3 1.198X102 lb/ft3 6.242X10-2 1.602X101 Energy J btu 9.484X10-1 1.054X103 (includes work) calorie 2.387X10-1 4.184X100
(thermochemical) erg 1.000X107 1.000X10-7 ft-lb 7.375X10-1 1.356X100 kW-hr 2.778X107 3.600X10-6
Force N dyne 1.000X105 1.000X10-5 pound force 2.248X10-1 4.448X100 Heat capacity J/kg*K btu/lb*oF 2.390X10-4 4.184X103 (includes entropy) cal/g*oC 2.390X10-4 4.184X103 Length m Å 1.000X1010 1.000X10-10 in 3.937X101 2.540X10-2
ft 3.281X100 3.048X10-1 micron 1.000X106 1.000X10-6 mile 6.213X10-4 1.609X103 yard 1.094X100 9.144X10-1 Mass kg ounce 3.527X102 2.835X10-1 lb 2.205X100 4.536X10-1 Power W btu/hr 3.414X100 2.929X10-1 btu/sec 9.484X10-4 1.054X101 cal/sec 2.390X10-1 4.184X100 ft-lb/sec 7.376X10-1 1.356X100 horsepower 1.341X10-1 7.457X102 (550 ft*lb/sec) Pressure Pa atm 9.869X10-6 1.013X105
(76cm Hg) bar 1.000X10-5 1.000X105 cm of Hg 7.506X10-3 1.333X103 dyne/cm2 1.000X101 1.000X10-1 in of Hg 2.961X10-4 3.337X103 kg force/cm2 1.020X10-5 9.807X104
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CONVERSION FACTORS FOR SOME COMMONLY USED UNITS OF MEASUREMENT CONT’D
lb/in2 (psi) 1.450X10-4 6.895X103 torr 7.501X10-3 1.332X102 (mm of Hg)
Volume m3 ft3 3.531X101 2.832X102 (includes capacity) in3 6.102X10-1 1.639X10-5 gallon (U.S.) 2.642X102 3.785X10-3 L 1.000X103 1.000X10-3 ounce (U.S.) 3.381X104 2.957X10-5
psig ≡ pounds per square inch gauge psia ≡ ponds per square inch absolute
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1.2 COMMON ITEMS OF GLASSWARE AND EQUIPMENT
Most of the apparatus used will be familiar to you, but the following notes may help
you in identifying and using specific pieces.
The ERLENMEYER or CONICAL FLASK is used for handling
solutions, and for titrations. It is designed with a narrow neck to
minimize loss of solution through splashing or evaporation.
The FILTER or BUCHNER FLASK is used in conjunction with the
Buchner funnel for vacuum-assisted filtration. It is heavy-walled to give pressure
resistance; for this reason, a Buchner funnel should never be
used to heat a solution. Attach it to the vacuum line or water
aspirator with heavy-wall rubber tubing. For any operation involving vacuum, always use heavy walled rubber tubing - never use soft tubes like Tygon. If a water aspirator is used,
an empty flask should come between the filter and the pump to
avoid ‘suck-back’ if the water pressure falls. The Buchner
assembly is top-heavy and should be supported when in
operation.
The BUCHNER FUNNEL is used for filtration. It fits through a rubber
bung or cone into the Buchner flask. To use, assemble the flask and
funnel, place a filter-paper flat across the perforated porcelain plate,
and wet the paper with the solvent being used (usually distilled
water). Turn on the vacuum and make sure the paper is correctly seated in the
funnel; the filter paper should be cut slightly smaller than the funnel, but make sure it
covers all the holes. Stir the suspension to be filtered and quickly pour it onto the
center of the paper, using a glass rod to guide it. Filtration will go more quickly if you
keep liquid in the funnel. If all of the liquid is filtered off, the residual solid will pack
down into a solid cake, slowing filtration. To empty the funnel after the solid cake has
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been washed and sucked as dry as possible, loosen the cake around the edge with
a spatula, carefully invert the funnel onto a watch-glass and tap it gently.
The EVAPORATING DISH is used to provide a large surface area to
speed up evaporation. It can be heated on the steam-bath but should
never be heated with a direct flame.
The BURETTE accurately measures volumes to 0.1ml accuracy. When titrating,
always fill the burette to the zero milliliter marking. Your eye should be level with the
bottom of the meniscus in order to take a proper reading of the liquid level.
The TRANSFER PIPETTE accurately delivers one volume (e.g. 5 or10 or25ml).
The PASTEUR or DROPPING PIPETTE is for transferring a few drops.
The SEPARATORY FUNNEL is used for the separation of
liquids with differing densities and for washing. The funnel
should have a properly working stopcock and a stopper of the
correct size. The solution to be separated is poured into the
funnel with the stopcock closed and the funnel stoppered. It is
then shaken vigorously with two hands; one holding the bottom
of the flask between first and second fingers and the other on
the stopper so it does not fall out. This maneuver is performed
with the flask upside down and the stem directed away from
people standing by in case of any splashing. The pressure
which may build up inside the flask is released by holding the funnel upside down
and opening the stopcock. Next, the funnel is placed in a proper size support ring.
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Enough time is given for the solutions to separate into two distinct layers after which
the bottom layer can be removed and the procedure repeated until necessary. (It is
assumed that one knows which layer is to be kept!)
GAS CYLINDERS with the safety cap off need to be securely strapped to the wall or
a desk in order to prevent them from falling. There is considerable pressure in these
cylinders and care must be taken to control the flow of gases from them. The main
cylinder or tank valve should be closed when not in operation. This valve measures
the pressure present in the tank (i.e. how much gas is left in the tank). The control
valve, usually the second one, gives the pressure reading present in the
line connected to the tank. This is usually a backwards valve, meaning
that to reduce pressure it needs to be turned in the counterclockwise
direction. A BUBBLER is usually inserted in the gas line between the
cylinder and the connection to the apparatus to be filled with the gas.
This is advantageous for two reasons: 1) the flow of gas is actually seen
as it bubbles through the oil in the bubbler and 2) this is an outlet for the gas if the
pressure becomes too high so that it exits via the bubbler rather than blowing the
glassware or connecting tubes. When working with gas cylinders it is very important
that you know and understand how everything is connected and what function each
piece of equipment has.
There are many different shapes of CONDENSERS available for
use. All of them serve the same purpose in conjunction with
distillation apparatus. Their purpose is to cool the vapors inside the
condenser usually with water as coolant. The condenser is placed
before and is tilted toward the receiving flasks. The glass is blown
so that the cooling liquid is separated from the vapors which are to
be condensed. To have good cooling cold water should flow
through the condenser at all times. This is achieved by connecting
the water inlet to the bottom end of the condenser and the outlet to
the top so the water flows from the bottom up the condenser and out the top.
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ERROR DATA Volumetric flasks Volumetric pipettes
Volume (ml) Error Volume (ml) Error
10 ±0.04 1 ±0.006
25 ±0.06 2 ±0.006
50 ±0.10 5 ±0.01
100 ±0.16 10 ±0.01
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1.3 A NOTE TO THE STUDENT
The objective of this laboratory is to introduce the student to basic organic reactions
and analytical instrumentation, as used in industrial operations and processes. By
performing the prescribed experiments the student will become familiar with a typical
organic chemical laboratory and the operation of typical analytical instruments.
She/he will also obtain a feeling and routine for generation of analytical results and
the technical capabilities of various instruments. The generated knowledge will
enable the student to better understand the basic chemical principles and control of
industrial processes, which is essential for proper operation of individual units in the
plant and the management of processes for optimum performance and product
quality, and environmental effects. Whether in management, processing, design or
laboratory, an engineer should have good knowledge and understanding of the
chemistry, measurements and instrumentation being used in the plant. The
important decisions and modifications that an engineer must make in industry will be
based on the results obtained from the laboratory. A good understanding of possible
errors in procedures and instruments is also required and particularly a good
understanding of variables that could affect a result. A lack of this understanding
very often results in erroneous judgments that can affect considerably both
production and quality of the final product.
1.4 SAFETY GUIDELINES
Although this laboratory does not involve extensive manipulation of hazardous
chemicals, some of the materials that will be used are often flammable and volatile
as well as toxic. Each student must therefore follow strictly all safety procedures and
not perform any unauthorized manipulations with these chemicals prior to
consultation with the instructor or the demonstrator. Although most of laboratory
safety is common sense, this is a general guideline, and therefore may be
incomplete. If you are ever unsure about safety, please ask.
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Remember: ACCIDENTS ARE CAUSED AND CAN BE PREVENTED!
VIOLATION OF ANY OF THE REGULATIONS DESCRIBED BELOW WILL MEAN THAT YOU WILL NOT BE PERMITTED TO WORK IN THE LABORATORY AND THEREFORE RECEIVE A MARK OF ZERO FOR THE LABORATORY REPORT.
1.4.1 Laboratory Apparel Rules
-Safety goggles are required in the laboratory AT ALL TIMES! Eyes are extremely
sensitive and delicate to minimum amount of most chemicals. You are responsible to
provide your own goggles.
-Laboratory coats must be worn at all times inside the laboratory. If you need to step
outside the laboratory for a while, your coat must be removed and left behind in the
laboratory.
-Sandals, open-toed shoes and high heels are not permitted in the lab. Shorts or
skirts cut above the knee are not permitted either. If a spill occurs, your clothing will
protect you from direct exposure. Open toe and shorts or skirts do not protect your
feet from splashes and spills. The restriction on high heels is for balance. If you must
wear some of this gear for a later appointment or situation you should consider
carrying with you a pair of sneakers and sweat pants to wear during the lab.
- Careful consideration should be given before wearing any jewelry into the lab.
Some chemicals can get beneath a ring, watch or some other form of jewelry. This
prevents them from evaporating and holds them against the skin increasing the risk
of injury. If you decide to wear jewelry to the laboratory be particularly mindful of
itching, burning or any other irritation under or around your jewelry. Some gems and
precious metals might be easily damaged by the laboratory environment (for
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instance silver and opal jewelry). If you decide to ear jewelry you do it at your own
risk.
- Never wear clothes that hang, such as loose sleeves. Ties and scarves must be
tucked inside your laboratory coat
- A good suggestion is to wear only very old clothes to the laboratory. Some
students might consider bringing old clothes with them in a gym bag and change
right before and after lab. If you have a very tight schedule (must be documented)
and decide to change into old clothes before the lab we can arrange for 10 minutes
for you to change clothes.
- Long hair is to be constrained at all times.
1.4.2 Safety rules
- Eating and drinking in the laboratory is strictly forbidden.
- No radios, tape players, CD players, iPods or any other devices of this type will be
permitted in the laboratory at any time.
-Use of cell phones is not permitted in the laboratory.
-Identify all of the laboratory safety equipment, and keep their location in
your mind at all times. We might ask you to close your eyes any time during a lab
and point to such safety equipment as the fire extinguisher, the emergency eyewash
stations, the safety shower, the nearest exit etc. This exercise might save you from
a big injury, for instance if you were to splash a chemical in your eyes, you'd better
be able to find that eyewash station without your eyes well before permanent
damage can occur (which can be seconds depending on the nature of the chemical).
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- Develop good working habits. Keep your area clean and tidy. Your working area
reflects your working habits and also the quality of your work and results. A clean
and tidy environment decreases the probability of a laboratory accident.
- ALL FLASKS, BEAKERS AND CONTAINERS WITH ANY CHEMICALS THAT
YOU LEAVE BEIND AFTER A LABORATORY SESSION MUST BE CLEARLY
LABELED WITH THE NAME OF THE CHEMICAL, THE OWNER AND DATE!
UNLABELED VIALS CONTAINING CHEMICALS WILLL BE DISCARDED.
1.4.3 Handling Chemicals and Equipment Students will work in groups of two or three and enough time will be provided to
finish all the prescribed experiments. If any piece of equipment fails or does not
function properly, students are required to report the problem immediately to the
demonstrator or the instructor and are not allowed to attempt to fix the instrument on
their own.
1. Do not taste chemicals.
2. Do not pipette any chemicals by mouth. Use the rubber bulb.
3. Do not pour liquids that are flammable or that do not mix with water into sinks.
Pour them into the provided and labeled containers.
4. Do not mix incompatible chemicals. If you do not know their compatibility ask.
5. Do not heat flammable liquids with an open flame or put them in the even.
6. Operations involving volatile or toxic materials are to be conducted in the fume
hood.
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7. Dispose of solid waste in the appropriate solid waste containers. Do not use the
sinks.
8. Clean up spills (solid or liquid) at once.
9. Return chipped or broken glassware to the laboratory TAs or technician.
10. Be sure the apparatus is placed properly. Do not move instruments without
proper consultation with the laboratory TAs or technician.
11. When heating a test tube make sure that it is not pointing towards yourself or
other people in the vicinity, so no damage will result if the contents suddenly
‘dump’ out.
12. Never apply force to any glass apparatus. Many serious cuts are caused by the
sudden fracture of glass under strain or from misuse. In particular, never use
force in an attempt to push a thermometer or glass tube through a hole in a cork
or rubber.
13. Never heat a tightly sealed flask even if it is empty-----it will explode.
14. Do not attempt to buttress a laboratory assembly with makeshift supports such
as books, pencils and the like. Use several ring stands if necessary. Round
bottom flasks cannot stand freely on the bench----use a special cork support or
place the flask into a beaker.
15. Do not attempt to break up a solid in the bottom of a flask by punching the solid
with a glass stirring rod. The rod may either fracture in your hand or puncture the
bottom of the flask.
16. Avoid shortcuts! If you have an idea for an improvement talk it over with your
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demonstrator; if no objections, try it; if it is successful, tell us about it, you will
get extra marks.
17. There are certain necessary precautions associated with particular chemicals or
experiments. Your demonstrator will point these out when required.
18. If you are not familiar with a piece of apparatus or an experimental procedure
ask for help. Don’t just try to muddle through without knowing what you are
doing.
19. Chemical waste should be disposed of in the labeled waste containers provided. Halogenated chemicals must be disposed of separately from non-
halogenated chemicals. Nothing should go down the drain! Dispose of all
chemicals in the bottles marked for the specific lab.
1.5 GENERAL HOUSEKEEPING AND LABORATORY WORK
ALL STUDENTS MUST HAVE A LABORATORY BOOK IN WHICH TO RECORD
ALL LAB DATA DURING THE LABORATORY PERIOD. The lab book used in
industry is a legal document. Loose papers for recording of data or comments are
not allowed and will be removed from the lab.
Each student must read the instructions for the particular experiment PRIOR to
coming to the laboratory. He/she should understand the whole procedure and what
must be done in the experiment. This knowledge will be checked periodically by the
instructor or the demonstrator and it will be evaluated.
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Equations for all reactions should be written out in your lab notebook before coming
to the laboratory. Also, any calculations required (e.g. theoretical yield, preparation
of solutions) should be written in full in your laboratory notebook before coming to
the laboratory. Record all observations in the lab book including any color changes,
unexpected events, smells, etc. The notebook will be marked from time to time
during the term. Plan your working time in the laboratory! By doing this your
laboratory will be a useful and pleasant experience rather than a frustrating one.
ALL LABORATORY EXPERIMENTS MUST BE DONE DURING THE ALLOCATED
TIME PERIOD. THERE WILL BE NO EXTENSION OF THE LAB AND NO
ADDITIONAL TIME PERIOD AVAILABLE FOR THE EXPERIMENTS, except under
special circumstances such as illness verified by a doctor’s note.
1. Keep benches clean and orderly and sinks clean. You must leave your portion of
the bench and all glassware and equipment clean at the end of the lab period.
2. Aisles and floors are to be kept free of obstructions. Keep cupboard doors and
drawers closed when not in use.
3. Hang coats on the rack. No coats are permitted on tables or benches.
4. Laboratory doors MUST be unlocked during lab period.
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1.6 GUIDELINES FOR PREPARATION OF LABORATORY REPORTS Laboratory reports should be written as though they were short technical reports.
Thus Tables and Figures should always be referred to in the prose text of the report,
i.e. they should not appear on their own.
The report should be written in the past tense, since it is a description and a
correlation of past observations. The present tense may be used in referring to laws
of nature, properties of materials etc. which are independent of time. Thus, for
instance, in a particular experiment “The ambient temperature was 22 °C”; on the
other hand, “The ambient temperature is 20 °C ”.
TITLE PAGE
Title of experiment
Name of person writing the report
Name of experimenters
Date when the experiment was performed
All Formal reports must be TYPED
ABSTRACT
The Abstract should summarize the entire report. It should state clearly and briefly
the objectives, methods, results and conclusions of the lab.
Objective: State the objective or purpose of the lab
Method: In one or two sentences, summarize the methods, including scientific
and common names of organisms and techniques used.
Results: Summarize what was found in the study
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Conclusions: State the significance of the results in relation to the objective.
INTRODUCTION
The Introduction should describe the scope and the purpose of the lab and include
any background information necessary to understand the experiment.
State the general problem. Give a brief statement of why the general topic is
relevant and important. Define any specialized terms or concepts (e.g. the concept
of distillation) likely to be encountered later in the lab report. Supply sufficient
background (historical and theoretical) information to allow the reader to evaluate
and understand the results of the study without needing to refer to other
publications.
State the specific objective or purpose of the lab and the approach to be used. The
purpose states what you are investigating and why; how you perform the
investigation should be described later in the Methods and Materials.
MATERIALS AND METHODS
The Methods section should describe what was done and how it was done. It should
be written in the past tense with passive voice, and in paragraph form.
The Materials and Methods section should provide only enough detail to permit a
competent worker to evaluate the validity of the experiment and to repeat it, if
necessary. It should not be simply a recipe of all the steps involved. State the names
(IUPAC if possible) of the chemicals used, the instruments, equipment and pattern of
replication. Describe any unusual numerical calculations and state the statistical
technique used to analyze the data.
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RESULTS
The Results section should present the data collected in a summarized form and
describe only the key features of these data, emphasizing trends or patterns that are
relevant to the hypotheses being tested. Interpretation of the data is reserved for the
discussion section.
Do not present the same data in both a table and figure i.e. place table of raw data in
an appendix and place figure in the results section. Titles of tables and figures
should contain enough information to understand the contents without reference to
the text. The number and title are placed at the top of a table, and at the bottom of
the figure.
Guide the reader through your figure (s) and table (s) in a logical and systematic
manner, pointing out trends and differences that pertain to the objective (s) of the
report. Simply state what you found in your study, without inference or reference to
"expected" results.
DISCUSSION
The Discussion section should provide an explanation and interpretation of your
results and indicate the significance of the results to the hypothesis being tested.
Results of previous studies on the same topic should be compared with yours, with
an explanation of why your results are different from previous studies, if necessary.
State how and why your results either support or do not support the objectives and
hypotheses. REMEMBER: results are results, they are never wrong simply by being
different from either your expectations or from other investigations.
CONCLUSIONS
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Draw conclusions about the hypotheses (objectives) of your study, based on all
data available in the current studies.
APPENDIX
Includes all raw experimental data, e.g. time vs. temperature data points, sample
calculations, and any other information or data used for the experiment and
calculations.
REFERENCES
The Reference section should be a list of all books, journals, and other materials
cited in the body of the paper.
The surname of the author(s) and the year of publication should be inserted in the
text at an appropriate place:
"Smith (1991) compared..." or "... have been recently compared (Smith, 1991)."
If the reference has more than 2 authors, include only the surname of the first
author, followed by "et al."
"Smith et al. (1991) compared..." or "... have been recently compared (Smith et al.,
1991)."
When listing more than one citation at a given point in the text, list them in
chronologically by first author, but for 2 (or more) papers, published in the same
year, list these alphabetically:
"(Jones, 1978; Black et al., 1989; Smith, 1989; Jones and Smith, 1991)"
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If an author or group of authors has published more than one article in a given year,
you can distinguish between these articles by placing a letter postscript after the
publication year:
"(Black and Smith, 1990a; Black and Smith, 1990b)"
List all references in alphabetical order, sorted by the author(s)' last name(s). In
cases where the same author or group of authors has/have published multiple
papers that you have cited, then arrange these references in chronological order.. All
authors must be given in the reference list - the abbreviation "et al." Is used only in
the text. The following are examples of the punctuation, style and abbreviations that
may be used for references (note: the headings given here are not to be included in
your reference list).
Journal article:
Jones, R.S., E.J. Gutherz, W.R. Nelson and G.C. Matlock. 1989. Burrow utilization
by yellowedge grouper, Epinephelusflavolimbatus, in the northwestern Gulf of
Mexico. Env. Biol. Fish. 26: 277-284.
Chapter in a Book:
Gross, M.T. 1984. Sunfish, Salmon and evolution of alternative reproductive
strategies and tactics in fishes. Pp. 55-57. In: G.W. Potts and R.J. Wooten (eds.)
Fish reproduction: strategies and tactics. Academic Press, London.
Book:
Siegel, S. 1956. Nonparametric statistics for the behavioral sciences. McGraw-Hill,
New York. pp. 312.
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LABORATORY 1
FRACTIONAL DISTILLATION OF HYDROCARBON SOLVENTS (2 lab periods)
INTRODUCTION Distillation is the process of heating a liquid until its more volatile components
pass into the vapour phase, and then cooling the vapour to recover such
constituents in liquid form by condensation
The main uses of distillation are to facilitate the separation of volatile materials
from non-volatile materials or to separate a mixture of miscible liquids by taking
advantage of their different volatilities. If the difference in volatility (and hence in
boiling point ) between two constituents is great, complete separation may be easily
accomplished by a single distillation. Seawater, for example, which contains about
4% dissolved solids can be readily purified by vapourizing the water, condensing the
steam thus formed, and collecting the product, distilled water. An industrial example
of this type of simple distillation is the removal of water from dilute food products
such as apple juice, orange juice or maple syrup to produce concentrates which are
more biologically stable and easier to ship. In the case of separating miscible liquids
the most common applications are the separation of ethanol from fermentation
mashes and the fractionation of crude oil to produce light naptha, gasoline,
kerosenes , and heavy oils.
SOME THEORY
The average energy of the particles in a liquid is governed by the temperature;
the higher the temperature the higher the average energy. This relationship is
described by the Clausius-Clapeyron Equation. Some of the more energetic
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particles on the surface of the liquid can be moving fast enough to escape from the
attractive forces holding the liquid together. This is known as evaporation. If this
takes place in a closed system some of the gaseous particles will hit the surface of
the liquid and rejoin the liquid phase. Rapidly an equilibrium will occur in which the
number of particles leaving the surface is exactly balanced by the number rejoining
it. In this equilibrium there will be a fixed number of gaseous particles in the space
above the liquid. When these particles hit the walls of the container they exert a
pressure. This is known as the saturated vapour pressure of the liquid. A liquid boils
when its saturated vapour pressure becomes equal to the external pressure on the
liquid.
In a mixture of two or more volatile liquids each liquid makes a partial contribution
to the overall vapour pressure;
Pmixture = PA + PB + … ( Dalton’s Law ) When the sum of these partial pressures equals atmospheric pressure, the mixture
boils. This law implies that if this mixture is heated to boiling and the vapours
collected and condensed they will be enriched in the component that is more volatile
(lower boiling point).
This concept was then extended to illustrate that the contribution of each
components vapour pressure is related to its mole fraction in the mixture;
Pmixture = XAPA + XBPB … ( Raoult’s Law )
The enrichment of a particular component in the condensed vapours of a boiling
mixture is related to both their volatility ( P ) and their concentration ( X ) in the
original mixture.
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SIMPLE OR BATCH DISTILLATION Separation of two different miscible liquids by simple distillation is only effective
when the liquids differ in boiling points by approximately 50 degrees or more. If the
liquids comprising the mixture have boiling points that are closer than 50 degrees to
one another the distillate collected will be richer in the more volatile compound but
not to the degree necessary for complete separation of the individual compounds.
Combining the three laws mentioned and using a series of toluene (B.P. 111°C)
and benzene (B.P. 80° C) mixtures we can determine their boiling points, calculate
the relative concentration of each component in the condensates and construct the
graphs shown in Figure 1.1.
Figure 1.1. A liquid-vapour phase diagram for a mixture of benzene and toluene
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Figure 1.2. An 80/20 mole % mixture of toluene and benzene will boil at 100° C
and the condensate will consist of 56 % toluene and 44 % benzene
Figure 1.3. Boil, condense, repeat, boil, condense, repeat, boil, condense …. After
12 cycles ( plates ) the condensate will be 100 % benzene.
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FRACTIONAL DISTILLATION
A fractionation column consists essentially of a long vertical tube through which
the vapour passes upward and is partially condensed; the condensate flows down
the column and is eventually returned to the boiling flask. Inside the column the
returning liquid is brought into intimate contact with the ascending vapour and a heat
exchange occurs whereby the vapour is enriched with the more volatile component
at the expense of the liquid, in an attempt to reach an equilibrium within the liquid-
vapour system. In order to work efficiently the following three conditions are
necessary ;
(a) comparatively large amounts of liquid returning through the column ( reflux )
(b) thorough mixing of the reflux liquid and the rising vapour
(c) a large active surface area of contact between liquid and vapour
Although there are many different styles of fractionation column we are using a
device called a Vigreux column. It consists of a glass tube with a series of
indentations such that alternate sets of indentations point downward at a 45° angle
in order to promote the redistribution of liquid from the walls to the centre of the
column. These indentations also produce a large increase the surface area of
liquid/vapour contact. In this lab, using the Vigreux distillation column you will distill,
fractionate and quantify a three component mixture (that has been prepared by the
lab demonstrator in advance).
GAS CHROMATOGRAPHY
To quantify and identify the components of your mixture and the resulting
fractions you will use an analytical technique called gas chromatography (GC). GC
involves a sample being vapourised and injected into the head of a chromatographic
column. The mobile phase is a carrier gas, usually an unreactive gas such as helium
or nitrogen, and the stationary phase of the column is a microscopic layer of liquid or
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polymer on an inert solid support. The rate of progression of molecules along the
column depends on their degree of adsorption to the stationary phase. For a given
stationary phase, the degree of adsorption depends on the type of molecule,
meaning that different molecules reach the end of the column at different times.
Molecules exiting the column are monitored by a detector. The time at which an
analyte exits the column (called the retention time) is characteristic of the analyte
and can therefore aid in its indentification, while the quantitative response of the
detector can be used to determine its concentration.
MATERIALS
• 1000 ml two-necked boiling flask
• Heating mantle with Variostat power controller
• Digital thermometer
• Vigreux fractionation column
• Claison adapter with glass thermometer
• West condenser with water cooling
• Vacuum take-off adapter
• 2 100 ml graduated cylinders
• 4 sample vials
• Green plastic glassware clamps
• Ceramic boiling chips
• 300ml of a solvent mixture
• Varian Gas Chromatograph with an FID detector, a 20 m non-polar
capillary column, and He as carrier gas
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METHODS Assemble your distillation apparatus. Labware with ground glass joints is both
fragile and expensive so work carefully and ask your TA for assistance if necessary.
A small amount of silicone grease applied to the top one third of the male joint will
assure a leak-free connection but secure each joint with a green plastic clip.
1. Place 300 ml of the solvent mixture in your boiling flask along with 3 or 4 ceramic
boiling chips and check to ensure that all the joints in your apparatus are tight. The
Variac voltage controller plugs into the mains and the heating mantle plugs into the
Variac. Do not switch anything on until your TA has had a chance to check your set-
up.
2. Set the dial on the Variac to 70 to begin heating your solvent mixture. When the
mixture begins boiling start recording the temperatures from both the boiling flask
and the vapour as it enters the condenser. Take a reading every three minutes.
3. Collect the condensate in a 100ml graduated cylinder. When you have collected
the first 75ml change to a clean cylinder. Transfer about 10ml of this first fraction into
a labeled sample vial. Discard the rest of the fraction in the waste solvent bottle in
the fume hood. At this point you will probably have to increase the temperature of
the heating mantle (Variac to 80) to maintain your distillation rate. Collect the second
75ml fraction and save an aliquot. When approximately 50ml of the third fraction has
been collected shut off the heating mantle and let the system cool down. Combine
the 50ml collected as fraction three with the solvent remaining in the boiling flask.
Give it a mix and take an aliquot as fraction three.
4. Gas chromatographic analysis of your three fractions will take about 45 minutes.
Book a time with the lab instructor during the next week for your group to come in
and run your samples.
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PRE-LAB QUESTIONS
1. Draw the structures of toluene, pentane, and hexane and list their physical
properties (molecular weights, density, boiling point).
2. Explain Lab safety (goggles, lab coats and environment) and its importance.
IN-LAB QUESTIONS (Hand in to Lab Supervisor at end of Lab 2)
1. What are the names and email addresses of your TAs and Lab Supervisors?
2. What is a Mobile phase and a Stationary phase in gas chromatography? Explain
their differences.
3. What kinds of detectors are used in the GC in the lab? How do they work?
4. Draw a picture of a GC and label its parts.
SUMMARY OF REQUIREMENTS FOR YOUR LAB REPORT Within the context of the laboratory report format outlined in the beginning of this
manual, for this experiment be sure to include the following:
1. Draw a picture of the distillation set-up you used and label all the individual parts.
2. Based on the retention times of the peaks in your GC traces, identify the
components of your mixture. In table form show the name of the individual
components of the mixture, the retention time of the peak, the area under the peak
as found on the GC report. Calculate the percentage of each compound in the
mixture.
3. Repeat this table for each of your fractions.
4. Plot your temperature/ time data.
5. Using your GC results, discuss the effectiveness of your distillation for the
separation of the components of the mixture.
6. Discuss ways in which the separation efficiency of the distillation might be
improved.
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LABORATORY 2 ESSENTIAL OILS AND STEAM DISTILLATION
(1 lab period)
INTRODUCTION
An essential oil is a concentrated, hydrophobic liquid containing a mixture of
volatile aromatic compounds derived from plants. These products are used in
perfumery, aromatherapy, cosmetics, incense, flavouring food and drink, and
household cleaning products. The term essential is intended to indicate that the oil is
the fragrant essence of the plant from which it is extracted and not in the more
common grammatical sense of being indispensable. Essential oils can be extracted
from various types of plant material such as berries (allspice, juniper), seeds
(almond, nutmeg), bark (cinnamon), wood (camphor, sandalwood), leaves (bay,
eucalyptus, lemongrass), resins (frankincense, myrrh), flowers (jasmine, lavender,
rose), and fruit peel (orange, lemon, lime). Extraction methods include distillation,
solvent extraction, pressing, and supercritical fluid extraction with liquid CO2.
Although most citrus oils are produced by cold-pressing of the peel ( a by-product of
juice production ) the most common method of extracting essential oils is by steam
stripping or as it is more commonly referred to Steam Distillation.
The fundamental nature of steam distillation is that it enables a compound or
mixture of compounds to be distilled (and subsequently recovered) at a temperature
substantially below that of the boiling point(s) of the individual constituent(s).
Essential oils contain substances with boiling points up to and exceeding 200° C. In
the presence of steam or boiling water, however, these substances are volatilized at
a temperature close to 100° C at atmospheric pressure. The mixture of hot vapours
will, if allowed to pass through a cooling system, condense to form a liquid in which
the oil and water comprise two distinct layers. Most, but not all essential oils are less
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dense than water and form the top layer. This layer can then be removed by
decantation or further extraction with a hydrophobic solvent such as hexane or
diethyl ether.
SOME THEORY
Recall from laboratory one, the average energy of the particles in a liquid is
governed by the temperature - the higher the temperature the higher the average
energy. Some of the more energetic particles on the surface of the liquid can be
moving fast enough to escape from the attractive forces holding the liquid together.
This is known as evaporation. If this takes place in a closed system some of the
gaseous particles will hit the surface of the liquid again and be trapped again. There
will rapidly be an equilibrium set up in which the number of particles leaving the
surface is exactly balanced by the number rejoining it. In this equilibrium there will be
a fixed number of gaseous particles in the space above the liquid. When these
particles hit the walls of the container they exert a pressure. This is known as the
saturated vapour pressure of the liquid. A liquid boils when its saturated vapour
pressure becomes equal to the external pressure on the liquid.
Let’s assume that two immiscible liquids A and B are mixed together and the
system is well stirred. Since A and B are immiscible the presence of one does not
affect the properties of the other. Their partial pressures in the mixture of vapors are
the vapor pressures of the pure liquids. The total vapor pressure of the system is the
sum of the vapor pressures of the pure components A and B at that temperature:
PT = P°A + P°B
The vapor pressures of pure A and pure B and the total vapor pressure, PT, as a
function of the temperature are shown in Figure 2.1.
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Figure 2.1
In this particular example, pure A has a lower boiling point than pure B. It can also
be observed from the figure that when A and B are mixed, the boiling point of the
mixture (the temperature at which the total vapor pressure equals 760 Torr) is lower
than the boiling points of both A and B. In steam distillation one of the components is
always water, and thus the boiling point is always less than 100ºC. Steam distillation
is particularly useful to distill compounds with very high boiling points, which may
decompose during normal distillation. In this experiment you will isolate the
essential oil of your choice from a selection of natural materials, and analyze its
components by gas chromatography.
MATERIALS Each group of students will be provided with a sample of crude essential oil from
different and diverse sources. You will have the opportunity to choose the one you
think smells the best to work with. Your choices are as follows:
Nutmeg crushed dried
seed pods
Myristica fragrans Indonesia
Bay Leaf dried leaves Laurus nobilis L Morocco
Lemongrass dried leaves Cymbopogon
flexuosus
India
Key Lime whole ripe fruit Citrus aurantifolia Mexico
Juniper Berry crushed dried Juniperus India
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berries communis
Eucalyptus wood and leaves Eucalyptus globus China
Blood Orange cold pressed peel Citrus sinensus Italy
Citronella dried leaves Cymbopogon
narous
Sri Lanka
• 1 L , round bottom flask
• Distillation glass ware
• Heating mantle
• Graduated cylinder
• Varian 3380 Gas Chromatograph with a 20m Biosil-5 column, an FID detector
and using He as the carrier gas. The column oven is programmed to rise from
75° C to 200° C at 5° C /min. Total run time is 28 min.
METHODS
1. Transfer 300 ml of water to a 1L round-bottom flask. Connect the flask to a
distillation apparatus as shown in Figure 2.2.
2. Heat the water to boiling and when approx. 25 ml of condensate has collected in
the graduated cylinder transfer the contents of your sample vial into the boiling flask
with a Pasteur pipette.
3. Collect approx. 25 ml of distillate (you will see the oil floating on the top). Rinse
the vial your sample came in with acetone, dry it with a stream of air and transfer the
oil layer from the top of the receiver with a Pasteur pipette. Add a small amount of
sodium sulfate to dry the oil sample.
4. Arrange with your TA to have your sample analyzed by gas chromatography.
When you receive your GC report try to identify some of the compounds in your
sample using the posted list of known retention times.
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Figure 2.2
PRE-LAB QUESTIONS
1. Pick two of the oils listed in the materials section and look up the two or three
major components
2. Draw the structures and find the formulas and boiling points of the following
compounds:
d-limonene
1,8 cineole
citral
lynalyl acetate
3. Identify one industry and give an example of a company where the steam
distillation of essential oils might be an important process.
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IN-LAB QUESTIONS
1. Why do you add sodium sulfate to your oil phase following the distillation?
SUMMARY OF REQUIREMENTS FOR REPORT
1. Draw the structure of the major components of the essential oil that your group
selected to distill. Identify the functional groups in each molecule.
2. Include your GC trace for your product in your report. Discuss the results. What
does it tell you about your product and about the success of your experiment?
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LABORATORY 3 PURIFICATION OF ORGANIC COMPOUNDS BY EXTRACTION AND RECRYSTALLIZATION (2 lab periods)
INTRODUCTION.
Compounds found in nature are rarely found in their pure form. In addition,
the chemical reactions that are used to prepare molecules, such as in the
pharmaceutical industry, often lead to mixtures of products. If we wish to obtain the
molecule of interest, without possible undesirable properties of the impurities we
must isolate the molecule of interest in its pure form. Therefore, methods must be
available for the isolation and purification of the molecule of interest. In general, the
most important methods of isolation and purification are based on three particular
physical properties of the molecules:
Volatility: the tendency of molecules from the liquid or solid state to escape to the
vapour state, “vapour pressure” is the quantitative measure
Solubility: the extent to which one compound (solid, liquid, or gas) will dissolve in a
second compound to form a single homogeneous phase
Adsorption: the tendency of foreign molecules to be attracted and held in a
molecular monolayer on the surface of a solid, involves weak van der Waals forces
These properties are all determined mainly by intermolecular forces, which
we have discussed in class and which are described in detail in your text (Ch. 2,
Wade). A summary of the main methods of purification using these properties is as
follows:
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Method Physical property Type of material
Distillation Volatility Liquids
Recrystallization Solubility Solids
Extraction Solubility Solids or liquids
Chromatography Adsorption/solubility Solids, liquids or gases
In previous experiments, you have used the property volatility to separate
molecules by distillation. On an analytical scale you have also used the property of
adsorption to separate molecules by chromatography. In this experiment you will
use the property solubility to separate molecules using two techniques – extraction
and recrystallization.
Basic principles of extraction
The technique of extraction is probably the most widely used method for
either the initial isolation of natural products from their source materials or the
preliminary separation of products from reaction mixtures. In general, it is based on
the principle of phase distribution and involves the selective transfer of one or more
components of a mixture to a second, immiscible phase in contact with it. In practice,
the mixture is either a solid or liquid and the second phase is always a liquid solvent.
Liquid-solid extraction
This simple form of extraction is well-known and is practiced by most of us
each day when we brew a cup of coffee or tea. A finely divided solid mixture is
stirred, usually with heating, with a suitable solvent to effect selective dissolution of
one or more components of the solid. After mechanical separation of the liquid from
undissolved solids by filtration, the resultant solution may be used for some purpose
(such as a drink). Alternatively the extracted components may be isolated by
evaporation of the solvent (eg. instant coffee powder). The extraction of alkaloids
from leaves and barks, flavouring extracts from seeds, perfume essences from
flowers and sugar from sugar cane are all examples of separations of this type.
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Liquid-liquid extraction
In a liquid-liquid extraction, the desired compound is distributed or partitioned
between two immiscible liquid phases. This partitioning is a consequence of the
differing relative solubility of the component in the two phases. In this lab, you will
attempt to separate an organic acid, organic base, and a neutral compound by
liquid-liquid extraction. This separation is based on the observation that acids and
bases can be interconverted from their uncharged (neutral) form to a charged form
upon treatment with strong acids or strong bases. In general, charged species (eg.
ions) are soluble in water and are insoluble in organic solvents (eg.
dichloromethane, diethyl ether). On the other hand, neutral uncharged species tend
to be water insoluble and soluble in organic solvents. Therefore, if a charged species
has a choice between water and an organic solvent, it will go into the water, whereas
the neutral species will choose the organic solvent over water.
Like all generalizations, there are certainly exceptions to this rule. For
example, while sodium chloride dissolves in water, silver chloride does not. There
are quite a few ionic compounds that are insoluble in water. However, there are very
few ionic compounds that dissolve in organic solvents. Similarly, there are several
neutral compounds that dissolve readily in water. Most of these are small polar
molecules that are capable of hydrogen bonding to water (eg. sugars, methanol,
ethanol, acetic acid, acetone).
There are very few types of organic compounds that are acids and the most
common contain either the carboxylic acid group (-CO2H) or the sulfonic acid group
(-SO3H). The one you will encounter in this experiment contains a carboxylic acid
group. Most carboxylic acids are weak acids with a pKa ≅ 5. The consequence of
this is that an excess of hydroxide will deprotonate the uncharged acid and convert it
to its charged conjugate base form. Treatment of the charged conjugate base with
an excess of hydrochloric acid will protonate it and convert it back to the uncharged
carboxylic acid form.
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R OH
O
R O-
OHO-
H+
uncharged(organic soluble)
charged(water soluble)
If you now recall the solubility behaviour discussed above, this means that an
organic acid can be moved back and forth from water to an organic solvent by
treatment with excess hydrochloric acid or excess sodium hydroxide. When excess
hydrochloric acid is used, the organic acid exists in the uncharged form and will
select an organic solvent over water. With excess hydroxide, the organic acid exists
in the charged form and will select water over an organic solvent.
Analogously, an organic base, typically an amine (RNH2, R2NH or R3N) can
exist as either a charged or uncharged species depending on the pH of the solution.
However, the behaviour is opposite that of an organic acid. In acidic solution, the
base is protonated and exists as a charged ammonium salt. If the solution is made
basic, the ammonium salt is deprotonated and the base becomes uncharged. Thus
under acidic conditions, the charged base will prefer to dissolve in water and under
basic conditions the base will choose an organic solvent.
R +NR
RH
H+
HO-
R NR
R
charged(water soluble)
uncharged(organic soluble)
The majority of organic compounds are neither acids or bases and these
“neutral” compounds will not be protonated or deprotonated. Consequently, these
neutral compounds remain uncharged species that greatly prefer to dissolve in
organic solvents in preference to water.
It is this differing behaviour between acids, bases and neutral organic
compounds that is the basis for their separation by extraction. You will be given a
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mixture of the following compounds and will be asked to separate them by solvent
extraction.
OH
O
H2N NO2
NN
Benzoic acid (mp 122.4 oC) 3-nitroaniline (mp 114 oC) Azobenzene (mp 69 oC)
More details regarding extraction can be found in the Appendix 1C of this manual.
You should review these details before coming to lab. Recrystallization
Following your extraction, your products will not be totally pure. Therefore,
you will further purify them by recrystallization. This is a technique that is commonly
used on both the laboratory and industrial scale for the purification of chemical
products, particularly pharmaceuticals. This process is based on the fact that most
compounds are more soluble in hot solvents than in cold ones and that the
impurities present have solubilities different from those of the desired compound. It
involves dissolving your impure compound in a minimum volume of boiling solvent,
then allowing this solution to cool, resulting in the formation of pure crystals of
compound which can be isolated by filtration. More details regarding recrystallization
can be found in the Appendix 1A of this manual. You should review these details before coming to lab.
Melting point determination
After you have separated and purified your mixture of the acid, base, and
neutral compound, you will need some means to identify the compounds and assess
their purities. There are many techniques available for accomplishing this, including
chromatography which you used in the previous experiments, as well as infrared (IR)
spectroscopy, nuclear magnetic resonance spectroscopy (NMR) and ultraviolet
spectroscopy (UV) which you will use in future experiments. In this experiment you
will use a technique called melting point determination. The temperature and
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temperature range over which a solid product melts can assist in determining its
identity and purity. More information on melting points can be found in Appendix 1B.
You should also review this material prior to coming to lab. MATERIALS
• dichloromethane
• 2 - 100 mL Erlenmyer flasks
• 125 mL separatory funnel with stopper
• 500 mL beaker
• 3 M hydrochloric acid
• 3 - 250 mL Erlenmyer flasks
• Sodium sulfate
• Glass funnel
• Filter paper to fit above glass funnel
• Hot plate in fume hood (to be shared by several groups)
• 6 M NaOH
• Litmus paper
• Ice bath
• Buchner funnel with filter paper to fit
• 6M HCl
• Distilled water
• Melting point tubes
• Melting point apparatus
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METHODS Lab one CAUTION: Gloves and safety glasses should be worn at all times. Perform all work
under the elephant trunks on in the fume hood. Exercise care in the use of
hydrochloric acid (corrosive, will cause burns), potassium hydroxide (corrosive, will
cause burns), and dichloromethane (toxic). You will be using a lot of different
solutions and flasks. Label them appropriately and be organized. Dissolving the mixture 1. You will be given approximately 2.0 g of a mixture that contains equal portions of
benzoic acid, azobenzene, and 3-nitroaniline. Weigh the mixture to obtain an
accurate mass and then dissolve it in 50 mL of dichloromethane in a 100 mL
Erlenmyer flask.
Extraction of the organic base
1. Place the dissolved mixture into a separatory funnel. Make sure the stopcock is
closed. Be sure a large beaker is placed below the separatory funnel in case of any
spillage or leaking from the stop-cock.
2. Add 30 mL of 3 M hydrochloric acid to the dichloromethane solution in the
separatory funnel. Invert the funnel with the stopcock closed and top corked. Point
the flask away from you and shake gently while regularly venting.
3. Remove the two layers into separate labeled, clean, 250 mL Erlenmeyer flasks.
4. Place the lower organic layer back into the funnel and extract it two more times
with additional 3 M HCl (30 mL portions). All aqueous layers should be combined in
a 250 mL Erlenmeyer flask and set this aside until “Isolation of the organic base”.
Extraction of the organic acid
1. To the dichloromethane solution in the separatory funnel (you may want to add a
bit more dichloromethane), add 30 mL of 3 M sodium hydroxide and extract the two
layers as described above.
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2. Wash the lower organic layer two more times with additional 3 M NaOH (30 mL
portions).
3. Remove the two layers into separate labeled, clean, 250 mL Erlenmeyer flasks.
4. Place the lower organic layer back into the funnel and extract it two more times
with additional 3 M NaOH (30 mL portions). All aqueous layers should be combined
in a 250 mL Erlenmeyer flask and set this aside until “Isolation of the organic acid”.
Extraction of the neutral organic 1. Extract the dichloromethane solution in the separatory funnel 3 times with 20 mL
of distilled water. The combined upper aqueous layers may be discarded in the
aqueous waste. Note that these aqueous layers are contaminated with toxic dichloromethane and should not be disposed of in the sink. 2. Place the dichloromethane solution in an Erlenmeyer flask (you may need to add
more dichloromethane). Add the appropriate amount of anhydrous sodium sulfate to
this solution (until you see the snow globe effect) and allow it to dry for about 5 min.
Gravity filter the solution through a fluted filter paper into another pre-weighed
Erlenmeyer flask. 3. Evaporated the dichloromethane solution on a hot plate (on low setting) in the fume hood. 4. Set the flask with the solid aside to dry until the following week. Note its colour.
Isolation of the organic base
1. To be sure you have the right flask, check the pH using litmus paper. It should be
acidic to start. You will neutralize it by adding base.
2. Neutralize the combined acidic aqueous extracts (containing the organic base)
by adding 6 M NaOH dropwise (with swirling) until the solution is basic. This may be
monitored using litmus paper. The litmus paper will change from red to blue. Do not add too much base! Notice the colour change! (What is going on?)
3. Cool the flask in an ice bath for 10 minutes and collect the solid precipitate by
vacuum filtration using a Buchner funnel. Wash with 2 mL of cold distilled water.
4. Set the solid aside to dry until the following week. Note the colour of the product.
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Isolation of the organic acid
1. To be sure you have the right flask, check the pH with litmus paper. It should be
basic to start. You will neutralize it by adding acid.
2. While on ice, neutralize the combined basic extracts (containing the organic acid)
by adding 6 M HCl dropwise (with swirling) until the solution is acidic. This may be
monitored using litmus paper. The litmus paper will change from blue to red.
3. Continue to cool the flask in an ice bath for 10 min and collect the solid
precipitate by vacuum filtration using a clean Buchner funnel. Wash with 2 mL of
cold distilled water
4. Set the solid aside to dry until the following week. Note the colour of the product.
Lab two 1. Record the mass of each solid in order to calculate its yield.
2. Set aside a small amount of each solid to perform a melting point determination.
3. Recrystallize the benzoic acid from water
4. Recrystallize the 3-nitroaniline from water
5. Recrystallize the azobenzene from ethanol/water (5/1)
6. While you are waiting for the molecules to recrystallize, perform melting point
determinations on each crude solid that you set aside prior to its recrystallization.
7. Isolate your recrystallized solids and allow them to dry.
8. Weigh the dried solids and perform melting point determinations.
PRELAB QUESTIONS
1. Of the three compounds you are trying to isolate, identify the acid, base and
neutral. Draw the relevant conjugate acid and conjugate base of these molecules.
2. When you extract an aqueous solution with dichloromethane which solvent forms
the layer on top?
3. Outline the properties of a good recrystallization solvent.
4. What information can a melting point provide?
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5. Draw the appropriate chemical structure (acid, base, neutral, conjugate base,
conjugate acid) in the boxes in the following flow chart:
Azobenzene (mp 69 oC) 3-nitroaniline (mp 114 oC)Benzoic acid (mp 122.4 oC)
NN
H2N NO2
OH
O
6 M HCl
upper aqueous layer lower organic layer
6 M NaOH 3 M NaOH
upper aqueous layer lower organic layer
3 M HCl
dichloromethane solution
IN LAB QUESTIONS 1. If you were extracting using diethyl ether-water instead of dichloromethane-water,
how would this change your procedure?
3. If your solid is not totally dry before you weigh it for yield determination, what
effect would this have on your perceived yield?
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SUMMARY OF REQUIREMENTS FOR LABORATORY REPORT
1. Based on the number of grams of the starting solid mixture, and the assumption
that equal masses of each of the three compound were in the starting mixture,
calculate your isolated yield of each molecule both before and after your
recrystallization. Where might you have gained/lost material?
2. What do your melting points tells you about the identity and purity of each solid?
3. Was the extraction effective in separating the molecules?
4. Was the recrystallization step effective in purifying the molecules?
5. Why do you perform multiple extractions each time to isolate your molecule? To
illustrate your answer consider the Appendix 1C on partition coefficients. Using the
partition coefficient KC/W of 10 for caffeine in a chloroform water extraction, calculate
the % recovery if 100 mL of water containing 1.0 g of caffeine is extracted with either
150 mL of chloroform in a single extraction compared with 3 consecutive extractions
each of 50 mL.
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LABORATORY 4 NUCLEOPHILIC SUBSTITUTION REACTIONS (1 lab period)
INTRODUCTION.
Some of the most well established mechanisms of organic chemistry are
those of nucleophilic substitution reactions. In these reactions a nucleophile is used
to displace a leaving group from a carbon atom of an organic substrate:
R LG + Nu- R Nu + LG-
Substrate Nucleophile Product Leaving group
The nucleophile provides both electrons of the new bond between the
nucleophile and the carbon, while the leaving group takes with it both electrons of its
bond to the substrate. It is not necessary that the nucleophile has a negative charge,
but it must possess at least one nonbonded electron pair at its nucelphilic atom. A
typical example of a nucleophilic substitution reaction is as follows:
HO- CH3 I HO CH3 + I-+
There are two limiting mechanisms for nuclephilic substitutions. The
nucleophile may directly displace the leaving group in an SN2 process, or a
carbocation intermediate may be involved in an SN1 process. SN2 stands for
substitution (S), induced by a nucleophile (N) in a bimolecular rate-limiting step (2).
SN1 similarly denotes a nucleophilic substitution reaction with only one molecule
involved in the rate limiting step. This is typically the step in which the carbocation
forms.
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Mechanism of SN2:
Nu-+R3C LG C
RR
R
LGNuδ− δ−
LG-+R3C Nu
Mechanism of SN1:
R3C LG R3C+ + LG-
R3C++ Nu- R3C Nu
The SN2 process involves a backside attack by the nucleophile, to displace
the leaving group in a single, concerted step. Therefore, good nucleophiles are very
important in SN2 reactions. Effective nucleophiles include strong bases such as
alkoxide ions, as well as polarizable ions such as bromide and iodide. Good leaving
groups include ions that have readily polarizable bonds to carbon, as well as other
weak bases. The ease with which the leaving group is displaced influences the rate
of an SN2 reaction. Any increase in steric hindrance about the reactive carbon center
makes the transition state more difficult to form and slows down the reaction.
Therefore in general, the following order of reactivity is generally observed for SN2
reactions:
Methyl > primary > secondary > tertiary
Some molecules which are very poor substrates for SN2 reactions are able to
undergo substitution by the SN1 mechanism, involving a carbocation intermediate
which is generated by the departure of leaving group. For example, tertiary alkyl
halides undergo substitution by the SN1 mechanism because tertiary carbocations
are relatively stable. As in the SN2 mechanism, the ability of the leaving group to
stabilize the electron pair that it carries away strongly influences the rate of an SN1
reaction. It is also necessary that the carbocation is relatively stable or other
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competing pathways leading to other products will win out. In general, the following
order of reactivity is observed for SN1 reactions:
Tertiary > secondary > primary > methyl
In this experiment you will use qualitative chemical tests to study structure-
reactivity relationships in SN1 and SN2 reactions. To investigate the SN2 reaction, a
solution of the alkyl halide and sodium iodide in acetone will be used. These
conditions favour the SN2 mechanism because the iodide anion is a strong
nucleophile and the solvent acetone has very limited ability to stabilize a carbocation
intermediate. Sodium iodide is soluble in acetone, but when a bromide or chloride
anion is the leaving group in an SN2 reaction, sodium bromide or sodium chloride
precipitates as an insoluble salt, thus driving the reaction to completion.
To study the SN1 reaction, a solution of the alkyl halide with silver nitrate in
ethanol will be used. These conditions favour the SN1 mechanism because ethanol
is a poor nucleophile (suppressing SN2 reactions) but it is a polar solvent which
effectively stabilizes charged species such as carbocations. The silver ion
coordinates with the halide ion in the substrate, enhancing the ability of the carbon-
halogen bond to break, producing a carbocation and a silver halide salt which
precipitates. Compounds that undergo rapid reactions with ethanolic silver nitrate
solutions are those from which the carbocations are stable enough to form.
In both the SN1 and SN2 cases, precipitation of the product salt provides an
indicator that the reaction has occurred, and the rate at which this solid appears
gives insight into the rate of the reaction. This allows you to compare the reactivity of
several substrates in SN1 and SN2 reactions. The following alkyl halides will be
tested under each conditions:
1-Chlorobutane
1-Bromobutane
Crotyl chloride
Bromobenzene
2-Chlorobutane
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2-Bromobutane
2-Chloro-2-methylpropane
2-Bromo-2-methylpropane
MATERIALS
• All alkyl halides listed above
• 16 clean dry 10 x 75 mm test tubes with corks
• 8 mL of 15 % (wt) NaI in acetone
• Water or sand bath at 50°C
• 8 mL of 1 % silver nitrate in ethanol
METHODS CAUTION: Safety glasses should be worn at all times. Perform all work under the
elephant trunks or in the fume hood. All alkyl halides are harmful if inhaled, ingested,
or absorbed through the skin. Dispense the compounds only in a hood and always
transport test tubes stoppered. Silver nitrate solutions will discolour the skin.
Acetone and ethanol are flammable and irritants. Do not use near open flames or
hot electronic devices. Avoid contact with skin, eyes, and clothing.
SN2 reaction conditions
5. Preheat a water bath to 50 °C.
6. Label 8 dry 10 x 75 mm test tubes before obtaining any reagents and fit each
with a cork.
7. Place 3 drops, measured with a Pasteur pipet, of each alkyl halide into its own
test tube. Immediately close each tube with a cork and keep the tubes stoppered
except while adding the test reagent.
8. Add 1 mL of the 15 % NaI in acetone solution to the first test tube. Recork the
tube, record the time, and shake the tube gently to ensure complete mixing. Monitor
the reaction and record the time required for the precipitate to form.
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9. In the meantime, continue adding 1 mL of the test reagent to the other tubes,
following the same procedure as above.
10. If no precipitate forms, arrange the cork loosely in the tube and heat it in the
warm water bath (50 °C). Check the reaction periodically and record the time that
any cloudiness or precipitate appears. Consider the substrate unreactive if no
change occurs after 15 minutes of heating.
11. Pour the contents of all test tubes into the halogenated organic waste. Discard
the test tubes in the glass waste.
SN1 reaction conditions
5. Label 8 clean test tubes and fit each with a cork.
6. Place 3 drops, measured with a Pasteur pipet, of each alkyl halide into its own
test tube. Immediately close each tube with a cork and keep the tubes stoppered
except while adding the test reagent.
7. Follow the same procedure as above for the SN2 tests, but this time adding 1 mL
of 1 % silver nitrate in ethanol solution to each tube.
4. When you are finished, pour the contents of each tube into the halogenated
organic waste. Excess silver nitrate solution should be disposed of in the inorganic
waste.
PRELAB QUESTIONS
1. Draw the structures of each of the test alkyl halides.
2. Predict which compound will be most reactive under SN2 conditions. Which will be
most reactive under SN1 conditions?
3. Is the concentration of the nucleophile important in SN2 conditions? What about in
SN1 conditions?
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SUMMARY OF REQUIREMENTS FOR LABORATORY REPORT
1. Write out a balanced equation for the reaction of one of the alkyl halides with
silver nitrate in ethanol and for another alkyl halide with sodium iodide in acetone.
2. Draw complete arrow pushing mechanisms for each of the reactions from
question 1.
3. Which two compounds had the highest reactivity with silver nitrate in ethanol?
Why?
4. Which two compounds had the lowest reactivity with silver nitrate in ethanol?
5. Which two compounds had the highest reactivity with sodium iodide in acetone?
Why?
6. Which two compounds had the lowest reactivity with sodium iodide in acetone?
7. In general, what was the effect of having a chloride versus bromide leaving
group? Why?
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LABORATORY 5
HYDROGENATION OF A FATTY ACID METHYL ESTER (1 lab period)
INTRODUCTION
Fatty acids are molecules consisting of a carboxylic acid attached to the
terminus of a long hydrocarbon chain. In nature, they serve many roles such as
biological fuels, in phospholipids which constitute cell membranes, and to modify
proteins, providing them with a lipid to anchor them to the cell membrane. In nature,
fatty acids are usually found to contain an even number of carbons between 14 and
24, with 16 and 18 being the most common. They can be either fully saturated,
containing only C-C single bonds or they can contain one or more C=C double
bonds. These double bonds are almost always in the Z (cis) configuration. From
Table 5.1, some trends concerning the melting points of fatty acids can be identified.
The melting points increase with increasing numbers of carbons in the chain, and
double bonds reduce the melting points relative to their unsaturated counterparts.
In most fats and oils, fatty acids occur as triglycerides (Figure 5.1), and fats
from different sources contain different proportions of the different fatty acids. For
example, olive oil contains approximately 85% oleic acid, while palm oil contains
about 40% palmitic acid and 43% oleic acid. Coconut oil contains 45% lauric acid.
The physical properties of the fatty acids, such as their melting points give an
indication as to how they behave in the human body. Fatty acids lacking double
bonds tend to build up in the body, while fatty acids with double bonds tend to be
more fluid and less likely to build up causing health problems. This is why olive oil is
considered to be more healthy than palm or coconut oils.
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Number
of
Carbons
Double
Bonds
Structure Name Melting
Point
(°C)
12 0 CH3(CH2)10COOH Lauric acid 44
14 0 CH3(CH2)12COOH Myristic acid 58
16 0 CH3(CH2)14COOH Palmitic acid 63
18 0 CH3(CH2)16COOH Stearic acid 70
20 0 CH3(CH2)18COOH Arachidic acid 77
22 0 CH3(CH2)20COOH Behenic acid 80
24 0 CH3(CH2)22COOH Lignoceric acid 122
16 1 CH3(CH2)5CH=CH(CH2)7COOH Palmitoleic acid -1
18 1 CH3(CH2)7CH=CH(CH2)7COOH Oleic acid 16
18 2 CH3(CH2)4(CH=CHCH2)2(CH2)7COOH Linoleic acid -5
18 3 CH3CH2(CH=CHCH2)3(CH2)7COOH Linolenic acid -11
20 4 CH3CH2(CH=CHCH2)4(CH2)7COOH Arachidonic acid -49
Table 5.1 Properties of Fatty Acids
O
O
O
O
O
O
O
O
O
O
O
O
Triglyceride of stearic acid
Triglyceride of oleic acid
Figure 5.1 Structures of example triglycerides (fats)
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The process of hydrogenation involves the addition of hydrogen across a
double bond to convert the alkene into an alkane. This will be the focus of this
experiment.
The production of margarine by the hydrogenation or partial hydrogenation of
vegetable oils is a major industrial process. This process accomplishes two things.
First, by reducing the unsaturated / saturated ratio of the product the melting point of
the mixture increases to the point that the margarine becomes a spreadable solid at
room temperature. The second improvement in the product is that unsaturated fats
exposed to air oxidize to create compounds that have rancid, stale or unpleasant
odours and flavours and so partial hydrogenation reduces the amount of unsaturated
material and thus promotes an increase in shelf life of the product.
Hydrogenation is a very slow reaction with a high energy barrier unless it is
catalyzed. Metal catalysts such as Ni, Pt, and Pd are often used. To facilitate the
handling of these metals, they are often supported on an inert material such as
charcoal (carbon). While supported catalysts such as Pd/C (Palladium on Carbon)
do not dissolve in the reaction mixture, they are still highly active and the reactions
are referred to as heterogeneous rather than homogeneous in which all reagents
are dissolved in one solvent.
To monitor your hydrogenation, you will use another common and more rapid
reaction of alkenes: the addition of bromine (Br2) to the double bond. A solution of
bromine appears brown in colour, but if bromine adds to the double bond the
solution becomes colourless. Therefore, you can identify the completion of your
reaction by the point at which the solution remains brown in the “bromine
unsaturation” test.
The material that we will use for our hydrogenation experiment is a sample
of biodiesel made from canola oil. The base-catalyzed reaction of canola oil with
methanol (trans-esterification) converts all the triglycerides of the starting material to
methyl esters of the individual fatty acids. Since the triglycerides of canola oil are
composed almost entirely (96%) of C18 fatty acids ( stearic C18:0, oleic C18:1,
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linoleic C18:2, and linolenic C18:3 ) the end product of complete hydrogenation will
be methyl stearate . This product can be crystallized from MeOH.
MATERIALS
• 50 mL round bottom flask
• Magnetic stirrer
• Magnetic stir bar
• 0.5 g of biodiesel (methyl canolate)
• 3 mL of methanol
• 15 mg of 10% Pd/C
• Balloon
• Rubber septum
• Syringe and needle
• 0.5 mL of dichloromethane
• Celite (filter aid)
• Ice bath
• Bromine solution (0.1 M in Methanol)
• Melting point tubes
• Melting point apparatus
METHODS Reaction procedure 2. Set up your hydrogenation apparatus consisting of a 50 mL round bottom flask
with a magnetic stirrer.
3. Add 0.5 g of methyl canolate to the round bottom flask, followed by about 3 mL of
methanol.
4. Transfer 1 drop of this reaction mixture to a 10 x 75 test tube and add 1mL of
bromine reagent. Observe colour change.
5. Ask your TA to introduce about 15 mg of 10% Pd/C catalyst to your reaction
mixture.
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6. Place a septum on your flask, followed by a balloon filled with hydrogen which
you will introduce via a needle and part syringe. Watch your TA demonstrate how to
set up the reaction vessel. Start the magnetic stirring. The reaction will take approx
45 min to reach completion.
7. The black catalyst colour will eventually change to grey and the volume of the
balloon will probably decrease. Remove the balloon and septum from the flask and
test 1 drop of the reaction mixture with bromine solution. If the reaction is complete
what kind of colour change will you observe?
8. Add 0.5 mL of dichloromethane to the mixture and heat to dissolve all of the
product.
9. Add a full spatula of Celite. Filter the hot methyl stearate/Celite slurry as fast as
possible using a Hirsch funnel, and wash with a minimal amount of methanol.
10. Cool the filtered solution, using an ice bath, then collect the resulting crystals by
filtration.
11. Dry the crystals, weigh them, and determine the melting point.
Note: the catalyst is highly pyrophoric and should not be exposed to air. The catalyst should remain covered with methanol at all times when exposed to air. The TA should wash down any catalyst clinging to the side of the flask with a little squirt of methanol. Do not start the magnetic stirrer until after the system is closed under hydrogen atmosphere.
PRELAB QUESTIONS
1. Why do unsaturated fatty acids have lower melting points than their saturated
analogs?
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2. Draw the mechanism and product for the reaction of bromine with methyl oleate.
3. What does pyrophoric mean?
4. Draw a reaction energy diagram for the hydrogenation, labeling all axes and
relevant energies. Use the diagram to compare the catalyzed reaction with the
uncatalyzed reaction.
IN LAB QUESTIONS 1. What would be the consequence if you add too much dichloromethane to your
reaction mixture at the end of your reaction?
3. How could you improve the procedure of this experiment?
SUMMARY OF REQUIREMENTS FOR LABORATORY REPORT
1. Calculate the yield for your product. What would be any sources of product loss in
this experiment?
2. Discuss your melting point result. What does this result tell you about the identity
and purity of your product?
3. Discuss the results of your bromine tests. What do they tell you about the result of
your reaction and the identity and purity of your product?
4. What was the purpose of adding Celite before your filtration?
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APPENDIX: RECRYSTALLIZATION, MELTING POINT DETERMINATION, AND EXTRACTION
A. RECRYSTALLIZATION
1. Theory
(a) General Methods of Recrystallization
Chemical transformations lead invariably to mixtures of products, and therefore
various techniques of separation or purification must be employed to isolate
individual components in pure form, from the crude reaction mixtures. In the case of
solid substances the most commonly employed technique, at least until the advent
of chromatographic methods, was that of recrystallization (or more simply
"crystallization").
As commonly practiced, purification by recrystallization depends upon the fact that
most solids are more soluble in hot than in cold solvents. The solid to be purified is
dissolved in the solvent at its boiling point, the hot mixture is filtered to remove all
insoluble impurities, and then crystallization is allowed to proceed as the solution
cools. In the ideal case, all of the desired substance separates nicely in crystalline
form and all the soluble impurities remain dissolved in the mother liquor. Finally, the
crystals are collected on a filter, washed and dried. If a single recrystallization
operation does not yield a pure substance, the process may be repeated with the
same or another solvent.
(b) Nature of Suitable Solvents
The single most important factor contributing to a successful recrystallization is the
proper choice of solvent. In general, the most "suitable" solvent for recrystallization
purposes is one in which the compound to be purified is only very slightly soluble at
low temperatures but very soluble at higher temperatures (e.g. at the boiling point of
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the solvent). Most compounds exhibit positive temperature coefficients of solubility.
Ideally of course, the impurities to be removed should be readily soluble in the cold
solvent in which they will remain after the crystallization process or almost
completely insoluble even at elevated temperatures.
It should be realized that the solvent selected must be inert and not enter into
chemical reaction with the sample. In addition, it is desirable that the solvent be
reasonably volatile (low boiling point) so that it can be fairly easily removed from the
crystals by evaporation. Where two or more solvents are comparable with respect to
the properties already cited, factors such as inflammability, toxicity and cost are to
be considered.
The lower the solubility of the compound to be purified in the cold solvent, the
greater will be the recovery of purified material from the crude mixture. The fact that
the solubility of the impurities may be comparable to that of the desired compound
does not preclude the use of a particular solvent, since most impurities are present
in relatively small amounts. As an example, consider the recrystallization of a
mixture of solids consisting of 10 g of A and 1 g of B from a solvent in which the
solubility of each is 1.5 g per 100 ml at room temperature and 10 g per 100 ml at the
boiling point. One hundred milliliters of hot solvent would be required to dissolve the
mixture, and upon cooling the solution would precipitate 8.5 g of A (i.e. 85%
recovery) and no B because the solubility of B had not been exceeded. Only if there
were more than 1.5 g of B and 10 g of A would any B crystallize, and even then a
second recrystallization would complete the separation of up to 2.5 g of B.
(c) Choosing a Suitable Solvent
If no information concerning the solubility characteristics of the substance to be
recrystallized is available, the choice of solvent becomes an experimental problem. It
is necessary to test various solvents for their suitability according to the criteria
outlined above. As a rule, solvents of decreasing polarity are tried in succession and
the solubility behavior in each case observed. To do this, small-scale trial
recrystallizations are carried out rapidly in micro (10 X 75 mm) test tubes. A few
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drops of various common solvents (see table below) are added to small portions of
the crude, finely divided solid mixture and the crystals are stirred and crushed under
the cold solvent with a stirring rod. If solution occurs at room temperature the solvent
is obviously unsuitable. If solution does not occur, the test tube is heated gently on a
steam bath or over a small flame with stirring or shaking. A few more drops of
solvent are added if only partial solution has occurred. (Transfer of solvent is most
conveniently done with small clean dropping tubes drawn out at the end like
pipettes). If a homogenous solution is obtained it is cooled, and the inside walls of
the test tubes are scratched if crystallization does not occur readily. If no crystals
can be obtained or if solution does not occur on warming, the solvent is unsuitable
and another should be tried. To avoid misleading observations, some care and
judgment must be exercised in choosing the relative amounts of solid and solvent to
be used in these solubility tests.
Choosing a Suitable Solvent
The ultimate proof of the suitability of a particular solvent is in achieving a separation
of the desired component from the unwanted impurities. This can be established by
collecting the crystals which precipitate from the solvent being tested and
determining their melting point.
In many cases it is difficult to predict a suitable solvent. In general, it is said that "like
dissolves like" -that is, a substance will dissolve in a solvent containing similar
groups -or better, that polar solvents will dissolve polar molecules and nonpolar
solvents will dissolve nonpolar molecules; but a good recrystallization solvent cannot
be too like the compound being purified. The accompanying table lists, in order of
decreasing polarity, some of the common solvents used for recrystallization.
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2. DETAILED DESCRIPTION OF EXPERIMENTAL STEPS AND APPARATUS
(a) Preparation of Hot Solution
At this stage the key words are saturated and minimum. Since the compound to be
purified will invariably be soluble in cold solvent to some extent, however small, the
recovery of pure material will be a maximum only by employing no more solvent than
is absolutely essential to obtain a complete solution at the elevated temperature. By
working at or near the boiling point of the solvent, full advantage is taken of the
temperature coefficient of solubility for that particular solute/solvent combination.
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Quantitative solubility data are not essential, and the following general approach
may be applied to any solute/solvent combination. The solid is placed in a flask of
suitable size and just covered with a small quantity of solvent (use a volume
comparable to that of the solid phase, but certainly less than will be required
ultimately). The flask and contents are heated gently on a steam bath, shaking or
swirling, to a temperature just below the solvent's boiling point. Heating may then be
interrupted, an additional small quantity of solvent added, and heating resumed. This
procedure is repeated until the last bit of solid just dissolves or until no further
decrease in the amount of undissolved material is apparent. In many instances it will
not be possible to obtain complete solution because of the presence of insoluable
impurities in the mixture.
For this and all subsequent operations in the recrystallization sequence, it is
convenient to use the conically shaped Erlenmeyer flask, but never beakers. This
particular design offers many practical advantages. It minimizes both solvent loss
(the upper walls acting as a condenser) and the distribution of crystals on the vessel
walls out of reach of the solvent phase. It is also particularly convenient for handling
in the transfer operations or for corking or fitting with a condenser. In this way, hot,
ascending solvent vapor does not escape but is condensed and continuously
returned to the solution flask. This is particularly important with solvents such as
ethyl ether, benzene, and petroleum ether, but in practice it is advantageous with
any solvent because loss of solvent over the period of time taken by the subsequent
filtration step will cause the solution to become supersaturated prematurely. In fact, it
is often desirable to have the solution slightly below saturation at this point to
minimize difficulties in the hot filtration (see below). This is especially true for highly
volatile solvents (e.g., b.p.<65°C).
(b) The Use of Decolorizing Charcoal
Frequently, the crude product of a reaction mixture is colored by the presence of
certain impurities which may have arisen through some oxidation, charring, or
polymerization process accompanying the main transformation. Such impurities are
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usually highly polar and, even though very soluble in the recrystallization solvent,
tend to be adsorbed or occluded by the growing crystals of the solute. In these
instances, a series of wasteful and repetitive recrystallizations can be avoided by the
addition of a small quantity of animal charcoal ("Norite", Darco", "Nuchar", etc.) to
the hot solution. This addition permits selective adsorption of the colored impurities
by the active carbon prior to the crystallization process.
It is important that the solution not be super-heated when the active carbon is to be
added, or excessive frothing and "boiling over" of the flask contents may occur.
Usually the flask is removed from the heat, and after a moment the carbon is added.
The contents of the flask are kept hot and shaken briefly to ensure wetting of the
carbon surface. Adsorption occurs very rapidly, and no advantage is gained by
boiling the suspensions for several minutes. Charcoal is actually less effective at
elevated temperatures, and the only reason for operating at the boiling point is to
keep the substance to be crystallized in solution. The effectiveness of the charcoal in
adsorbing the colored impurities is directly proportional to the solvent polarity and is
best in an aqueous solution. The smallest amount of decolorizing agent that will do
the job should be used because the desired solute may also be adsorbed (thus
reducing the yield of product recovered) if the surface of the adsorbent is not
saturated by the coloring matter. For this reason (and since the color may also be
due to the desired compound) the decolorization step is seldom repeated whatever
the results of the single trial. In practice, one seldom uses more than 20 mg of
charcoal per gram of dry compound.
(c) Hot Filtration
The hot solution must be filtered to achieve separation from any insoluble impurities
or other undissolved materials. If charcoal is used for decolorization the necessity for
filtration is obvious. If no undissolved material is evident (which is very seldom) this
step may be omitted. The chief difficulty encountered in this operation is that of
keeping the solution hot enough to avoid premature crystallization in the filter. This
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means that the filtration must be done as rapidly as possible with minimal cooling of
the solution.
Rapid filtration of small quantities of solution is best done by gravity through a fluted
filter paper supported in a stemless glass funnel. Fluting of the filter paper (the
technique of folding will be demonstrated in the laboratory) increases the rate of
filtration by presenting a much larger surface area to the solution. If a regular funnel
with a stem is used, there is a good possibility of filtrate cooling in the stem, with
crystallization resulting. The relatively narrow stem thus becomes clogged and
filtration is impeded. It is often advantageous to preheat the glass funnel simply by
briefly heating it in a flame or by pouring a quantity of hot solvent through it
immediately prior to filtration. If water is used as the solvent, the filter funnel may be
warmed conveniently on a steam bath.
When working with particularly volatile solvents or with solids having very large
temperature coefficients of solubility, it is particularly difficult to avoid premature
crystallization. In these cases it is usually better to prepare the hot solution with
excess solvent (i.e. the solution is not saturated at the boiling point). After the hot
solution has been filtered the excess solvent must, of course, be removed by
evaporation before inducing crystallization.
It should be emphasized that the hot filtration is done by gravity (at least in an
elementary laboratory) and not by suction filtration as described below for the
collection of the crystallized product. The use of suction for filtering a hot, nearly
saturated solution is nearly always highly unsatisfactory, because the reduced
pressure in the filter flask causes rapid evaporation of the hot solvent; consequently,
the solution is not only more concentrated but it is cooled by the heat of vaporization
and becomes supersaturated.
Crystallization in the funnel is then almost inevitable and the funnel may become
completely plugged by the deposited crystals.
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In carrying out the actual filtration, the fluted paper is inserted into the stemless
funnel so that the lower tip of the paper projects into the opening at the bottom of the
funnel. The hot solution is decanted quickly but carefully into the paper, keeping the
level of solvent well below the top of the paper. When all of the solution cannot be
put into the funnel at once, the remainder is kept warm on the steam bath or hotplate
until it can be transferred to the funnel.
After the solution has run through the paper, a crust of crystals often remains around
the tip of the funnel and ill-formed crystals often form in the body of the cooling
filtrate. It is common practice to rinse the original flask with a little hot solvent and to
filter this through the filter paper to redissolve the crystals adhering thereto. The
filtrate which has been collected in an Erlenmeyer flask should be reheated to
redissolve any material that has crystallized and, if significantly diluted below the
saturation point, should be concentrated to its original optimal volume prior to
cooling and crystallization.
If the various precautions outlined above fail to prevent excessive crystallization of
the solute in the filter paper, the simplest expedient is to return the complete filter
paper and its contents to the original flask, add additional solvent, boil briefly to
ensure complete solution, and begin a new filtration.
(d) Cooling/Crystallization
Crystallization is accomplished by allowing the hot filtrate to cool slowly, undisturbed,
to room temperature (or at least until crystallization has begun) and then chilling the
mixture in ice or cold water to complete the precipitation. The objective is, of course,
that the desired substance be deposited as pure crystals while any "insoluble"
impurities remain dissolved in the "mother liquor". The lower the temperature to
which the solution is cooled, the more the desired substance will crystallize;
however, at some point the impurities may also begin to separate from solution. The
size of the crystals which separate will vary with the rate of cooling and the degree of
agitation of the solution. Rapid cooling with stirring tends to produce small crystals,
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while slow cooling of an undisturbed solution tends to give larger crystals. In general,
either very large or very small crystals are undesirable. There are problems
associated with the collection of very fine crystals because of clogging of the pores
in the filter paper and of adhesion of the small particles to the walls of the
crystallization flask. Moreover, if the solubility of the impurities is comparable to that
of the desired compound, sudden chilling may result in the co-deposition of the
impurities; whereas, with slow, undisturbed cooling these tend to remain in
supersaturated solution and more complete separation is effected. On the other
hand, with very large crystals there is a tendency for the mother liquors to be
occluded within the crystals. In the subsequent drying operations, evaporation of the
solvent will leave a deposit of impurities on the crystals.
Yet another problem associated with too rapid chilling of the solution, especially in
the case of low melting solids, is the tendency for the solute to separate first from the
solution as an "oil" which subsequently solidifies to a crystalline cake. If this
happens, it is possible for impurities to be distributed between the solvent layer and
the "oily" layer (see discussion of principles of extraction). The impurities will then be
entrapped when the oil solidifies. For this reason it is often desirable to choose a
solvent for recrystallization whose boiling point is lower than the melting point of the
solid being purified.
(e) Collection of Crystals -Cold Filtration
The important objective in the collection of a purified product is complete separation
of the crystals from the "mother liquor" containing the dissolved impurities. This is
achieved most effectively by employing suction filtration. The necessary apparatus
consists of a Buchner funnel attached to a heavy-walled filter flask which is
connected through a "trap" bottle to the source of suction (water aspirator or vacuum
pump) -see accompanying illustration.
The Buchner funnel is prepared for filtration by attaching it to the filter-flask by
means of a cork or rubber adapter, inserting a piece of filter paper whose diameter is
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just sufficient to cover the holes in the filter plate (the paper must not fold up against
the sides of the funnel), wetting the paper with a-small quantity of the solvent being
used, then smoothing the paper snugly against the filter plate by the application of
gentle suction.
The cold contents of the crystallization flask are stirred to break up any lumps and
swirled to obtain suspension of crystals. The suspension is decanted quickly into the
funnel in such a way that a layer of uniform thickness is obtained across the whole
surface of the filter bed. This is essential for obtaining complete separation of the
mother liquor. It is important, particularly in the early stages, to use only sufficient
suction to obtain a steady flow of filtrate. Very strong suction at this stage will draw
the finer particles into the pores of the paper, clogging them, and slowing the rate of
filtration unnecessarily. The bulk of the crystals remaining in the flask may be
transferred to the funnel with the aid of a metal spatula. Any crystals still remaining
are most efficiently transferred to the funnel by rinsing the flask with a portion of the
filtrate (which is already saturated with solute) rather than using fresh solvent.
When the bulk of the mother liquor has drained through, the cake of crystals is
pressed down quickly with a spatula or glass stopper and the suction is interrupted
by removing the rubber tubing from the filter flask. It is particularly important at this
stage not to draw air through the crystal cake, because this will cause evaporation of
the mother liquor and the impurities that were dissolved therein will be deposited on
the surface of the crystals. If it is intended to use the mother liquor to obtain a
second "crop" of crystals after concentration to a suitable volume, it should be
transferred at this point to a separate vessel (or the Buchner funnel attached to a
clean filter flask).
(f) Washing the Crystals
To complete the separation of the mother liquor, the crystalline cake must be
washed with small quantities of fresh, clean solvent. This is done conveniently by
covering the filter cake completely with a thin layer of solvent, with the suction
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disconnected. If possible the crystalline cake should be loosened with a spatula to
ensure complete wetting by the wash solvent, but care must be exercised not to
disturb the filter paper. The suction is reapplied and the wash solvent drawn down
through the crystals, which are then pressed down firmly as before to remove the
wash liquid as completely as possible. For complete removal of the mother liquor
from the crystals two or three such washes are recommended; however, to minimize
loss of product, the wash portions should be small in volume and the solvent should
be cold. After the last wash, full suction is applied to draw air through the filter cake
to suck it as dry as possible.
(g) Drying the Crystals
The final operation in the recrystallization of a product, is the drying of the solid. If
the solvent employed in the recrystallization is rather volatile, it is possible that it will
have evaporated completely during the last stages of the suction filtration.
Otherwise, further drying is necessary. As used here the term "drying" refers to the
complete removal of solvent, be it organic or aqueous.
The cake of crystals is transferred, with the aid of a spatula, to a sheet of glazed
paper, a watch glass, or any suitable container having a relatively large surface
area. The solid sample should be spread out and permitted to stand in air with
periodic stirring with a spatula. In many instances, however, simple "air-drying" as
just described will be inadequate or much too slow. The last traces of solvent (and/or
atmospheric moisture) are removed most conveniently by using a drying oven, a
dessicator, or by evaporation under vacuum. A dessicator is simply a closed vessel
with a lower compartment containing an anhydrous salt such as phosphorus
pentoxide which can remove water vapor by forming a hydrated salt. When using
either an oven or a dessicator the rate of drying can be enhanced even further by
reducing the pressure in the system. The temperature at which the crystalline solid is
dried in an oven should, of course, be significantly lower than the melting point.
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B. MELTING POINT DETERMINATION
1. Use of Melting Points for Analysis
Most crystalline organic compounds have characteristic melting points that are
sufficiently low (50 – 300 °C) to be conveniently determined with simple equipment.
Organic chemists routinely use melting points to a) get an indication of the purity of
crystalline compounds and b) help identify such compounds.
Pure crystalline compounds usually have a sharp melting point. That is, the melting-
point range or the difference between the temperature at which the sample begins to
melt and the temperature at which the sample is completely melted, is relatively
small (narrow). Impurities, even when present in small amounts, usually depress the
melting point and broaden the melting point range. A wide melting-point range (more
than 5 °C) usually indicates that the substance is impure, while a narrow melting
point range of about 0.5 – 2 °C usually indicates that the substance is fairly pure.
However, there are some exceptions to both of these generalizations. Small
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differences in melting point (on the order of 2 – 3 °C) may also result from variations
in technique, thermometer accuracy, and in the amount of experience possessed by
the person doing the melting-point determination.
Melting points can be used in the following way to help identify a compound.
Suppose a sharp-melting, unknown substance X is suspected of being identical to
some unknown substance A. If the two are identical, they should have the same
melting point. Thus, if A is reported in the chemical literature to have a melting point
significantly different from that observed for X, we can be quite certain that X does
not have the same structure as A. On the other hand, if A is reported to have a
melting point within a few degrees of that observed for X, the two substances may
be identical (the small difference being due to variations in technique or purity).
To be certain that X and A are identical, a mixture melting point can be determined,
i.e. the melting point of a mixture of X and A. If X and A are identical, the mixture
should have the same melting point as X or A has alone. On the other hand, if X and
A are not the same substance, even though they separately have the same melting
point, then a mixture of the two usually have a lower melting point and a broader
range than either substance alone. This is because each substance acts as an
impurity in the other.
To summarize, if a crystalline substance is pure, its melting point is likely to be
narrow. If two samples have identical structures, their mixture melting point is not
depressed and the melting point range is not broadened.
2. General Technique for Melting Point Determination
To determine the melting point, introduce a small amount of the finely powdered
material into a thin-walled capillary tube that is sealed at one end. The capillary tube
is inserted into a melting point apparatus and heated. Your lab demonstrator will
instruct you on its use. Two temperatures are recorded: the temperature at which the
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substance begins to liquefy and the temperature at which it becomes completely
liquefied. The observed melting-point range is the interval between these two
temperatures.
The observed melting point range can be influenced not only by the purity of the
material but also be the size of the crystals, the amount of material, how densely it is
packed in the tube, and the rate of heating. A finite time is required to transfer heat
from the metal block through the walls of the capillary tube and throughout the mass
of the sample. When a block is heated too quickly, its temperature rises several
degrees during the time required for melting to occur. This can result in an observed
range that is higher than the true one.
When the temperature of the block approaches the melting point of the sample, it is
essential for good results to raise the temperature slowly and at a uniform rate,
usually about 2 °C per minute. The sample should be small, finely powdered, and
packed tightly with a consistent density in a thin-walled capillary tube of small
diameter. The column of solid in the capillary tube should be just high enough to be
seen clearly during the melting (about 1-2 mm). It is a good idea to carry out the
packing process by dropping each sample melting point tube the same number of
times, e.g. three times.
The behaviour of a material upon melting should be observed and recorded
carefully. Record the range of the melting, such as 89.0 – 89.5 °C, as well as any
observations (decomposition, colour change, etc.).
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C. LIQUID-LIQUID EXTRACTIONS
1. Solvent partitioning in liquid-liquid extractions
The physical process that governs liquid-liquid extractions is solvent-solvent
partitioning. This is the distribution of solutes between a pair of solvents. In organic
chemistry, most commonly one of the solvents is water and the other is an organic
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solvent that is immiscible with water, such that two layers are formed with the upper
layer containing the less dense solvent. There are many organic solvents such as
diethyl ether, dichloromethane, toluene, and ethyl acetate, which have very limited
solubility in water and thus form a two-layer system. Each layer is saturated with the
other solvent. Inorganic compounds can usually be easily separated from organic
compounds by extraction. The former dissolve in the aqueous phase and the latter in
the organic solvent. In such cases, a single extraction may be sufficient to effect a
satisfactory extraction.
Many organic compounds, such as aldehydes, alcohols, esters, and amines,
which can form hydrogen bonds, are partially soluble in water. They distribute
themselves between the aqueous phase and the organic phase. For example, if we
add solute A to a mixture of water and chloroform, shake the system vigorously to
reach equilibrium, then allow the system to settle, solute A will be present in both
layers, but with a higher concentration in solvent for which it has higher affinity. The
distribution of A between the two solvents is dictated by the partition coefficient of
A between the solvents. The partition coefficient is the equilibrium constant for the distribution of a solute between two immiscible layers. For example, for
solute A distributed between chloroform (C) and water (W), the partition coefficient K
is defined as follows:
KC/W = [A] in chloroform/[A] in water
To a first approximation, the partition coefficient can be estimated as the ratio of the
solubility of the compound in each solvent. For example, the solubility of caffeine in
chloroform is approximately 18 g/100 mL, while its solubility in water is
approximately 1.8 g/100 mL. Therefore, we expect a partition coefficient KC/W ≈ 10
and most of the caffeine will be in the chloroform layer. To illustrate how the partition
coefficient can be used, consider performing an extraction of 1.0 g of caffeine using
60 mL of chloroform and 100 mL of water. If we let x be the number of grams of
caffeine in the chloroform layer, then (1.0 – x) will be the number of grams in the
water. The equation for K will therefore be:
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KC/W = [caffeine]chloroform/[caffeine]water ≈ 10 = (x/60)/((1.0-x)/100)
Solving for x we find that 0.86 g of caffeine will be in the chloroform layer and 0.14 g
will remain in the aqueous layer. We could then recover the caffeine by evaporating
the chloroform and we should get 0.86 g, which would be an 86 % recovery. In
general, the efficiency of a liquid-liquid extraction will depend on the partition
coefficient, the volume of solvent used, and the number of extractions. In general, it
is desirable to have the highest possible Korg/water. Increasing the volume of the
organic phase will also increase the % recovery, as will performing several
consecutive extractions.
2. Choice of solvent
The selection of an appropriate extraction solvent is critical to the success of this
technique. There are several important criteria to consider:
• The chosen solvent must not react chemically in an irreversible manner with
any component of the mixture
• The chosen solvent must be immiscible or nearly immiscible with the original
solution
• The chosen solvent must be favoured by the distribution coefficient for the
component being extracted
• The chosen solvent must be easily separated from the desired component
following the extraction. This generally means it should be low boiling and
thus easily removed by distillation or evaporation
3. General Technique for Extraction
In this discussion, assume that one of the solvents is water. Place the solution to be
extracted in a separatory funnel supported on an iron ring with the stopcock closed
(see figure below). The volume of the funnel should be at least twice that of the
volume of the solution. Add the extraction solvent and then tightly stopper the flask.
As a guideline, the volume of the extracting solvent should be between 25 – 50 % of
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the volume of the solution being extracted. Next, grasp the funnel with the palm of
the right hand around the body and the thumb and first two index fingers around the
stopper. Vigorously shake the funnel to mix the phases well. Never point the stem towards your face or anyone nearby. With the funnel inverted, carefully open the
stopcock to release the built-up pressure. Close the funnel, shake it again, and
carefully vent it again. Repeat this process until the pressure subsides.
Approximately 1-2 minutes of shaking should be adequate to reach equilibrium.
After the shaking, return the funnel to the ring and remove the stopper. Allow the layers to separate and then collect the lower layer in an Erlenmeyer flask
by opening the stopcock. Swirling the funnel may help if the interface between the
two layers is unclear. Always try to predict which phase is your desired phase based
on the known densities of the solvent, but be aware that the densities of liquids can
change depending on what is dissolved in them. To be sure that you do not discard
the wrong phase, always save all phases until the end of the experiment. If you are
in doubt as to which phase is the aqueous, you can add a small amount of water to
the top of the funnel and observe which layer it is miscible with.
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APPENDIX 2: PERIODIC TABLE OF THE ELEMENTS
1 H
2 He
3 Li
4 Be
5 B
6 C
7 N
8 O
9 F
10 Ne
11 Na
12 Mg
13 Al
14 Si
15 P
16 S
17 Cl
18 Ar
19 K
20 Ca
21 Sc
22 Ti
23 V
24 Cr
25 Mn
26 Fe
27 Co
28 Ni
29 Cu
30 Zn
31 Ga
32 Ge
33 As
34 Se
35 Br
36 Kr
37 Rb
38 Sr
39 Y
40 Zr
41 Nb
42 Mo
43 Tc
44 Ru
45 Rh
46 Pd
47 Ag
48 Cd
49 In
50 Sn
51 Sb
52 Te
53 I
54 Xe
55 Cs
56 Ba
57 La
72 Hf
73 Ta
74 W
75 Re
76 Os
77 Ir
78 Pt
79 Au
80 Hg
81 Tl
82 Pb
83 Bi
84 Po
85 At
86 Rn
87 Fr
88 Ra
89 Ac
104 Rf
105 Db
106 Sg
107 Bh
108 Hs
109 Mt
110 110
111 111
112 112