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o I have started the question that required calculations. USE THE DATA PROVIDED TO
o Feel free to ask questions
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o I have started question #1 (not sure if its right)
o Read the lab before answering all questions
1. What is the unknown painkiller in the pill? Explain your evidence for its identity, and
how you were able to rule out the other possible unknown. (5 points)
- DMS melting point report for phenacetin: 134˚
- DMS melting point for acetanilide: 114.3˚
In the lab we learned that both phenacetin and acetanilide are compounds reported with melting
points. The reported melting point for phenacetin is 134 degree Celsius, whereas the melting
point for acetanilide is 114.3 degree Celsius. Based on the reported melting point of these
compounds, it was able to compare it with the unknown’s melting point to determine which of
the two compounds had been substituted for acetaminophen. According to the melting points
obtained in the lab, the unknown painkiller in the pill is phenacetin because the melting point of
the recrystallized sample was 136˚ and for the crude, it was 135.8˚ which is close to the reported
melting point for phenacetin.
2a. Based on the sum of your weights of your separated compounds, calculate the total
percent recovery of the painkiller. Use the crude weight of the unknown for this calculation.
Discuss your results – how well were you able to collect each component during the separation
Initial mass of pain killer = 1.501g
Sucrose mass = 0.146 g
Aspirin mass = 0.3977 g
Unknown (crude) mass = 0.54 g
Total mass collected = 1.0837 g
(sucrose mass + aspirin mass + unknown mass)/initial mass of painkiller x 100%
(0.146 +0.3977 + 0.54)/ 1.501 x 100% = 72.19%
Total percent recovery = 72.19%
The separation of the unknown painkiller didn’t go as expected. The 72.19% recovery is a low
amount of solid recovered. It needs to be taken in consideration that after the filtration, and the
solution exposed to heat
2b. Calculate the percent composition of your painkiller. Again, use the crude weight of the
unknown for this calculation. Does your percent composition match the label? You may
assume a range of 8-12% sucrose, 35-45% aspirin, and 45-55% unknown is a match to the
label. Discuss your results – explain any experimental errors that could have led to deviations
of your values from the label values.
o Total mass collected = 1.0837 g
Each component mass/total mass collected x 100%
Sucrose 0.146/1.0837 x 100% = 13.5%
Aspirin 0.3977/1.0837 x 100% = 36.7%
Unknown (crude) 0.54/1.0837 x 100% = 49.8%
13.5 + 36.7 + 49.8 = 100%
3a. Calculate a percent recovery for the recrystallization of the unknown component of the
painkiller. Discuss your results – what experimental factors would cause a low percent
recovery of your sample? (5 points)
Recrystallized/crude x 100%
0.3855/0.54 x 100% = 71.38%
3b. Based on your crude and recrystallized melting points, was your recrystallization a
success? Did you see any improvement to your melting point after recrystallization?
Crude Melting point : 135.8˚
Recrystallized melting point : 136˚
4. While acetanilide and phenacetin are not very acidic, acetaminophen (like aspirin) has an
acidic hydroxyl group. Based on this information, what problem would you encounter during the
separation steps if the unknown component were acetaminophen rather than acetanilide or
5. Equation 1 shows the balanced reaction equation for the deprotonation of aspirin by sodium
hydroxide. What is the balanced reaction equation when sodium acetylsalicylate is treated with
concentrated hydrochloric acid?
6. Assuming the reaction from question #5 is spontaneous under standard conditions, label the
stronger acid, stronger base, weaker acid, and weaker base.
INFORMATION YOU WILL NEED TO ANSWER QUESTIONS CAN BE FOUND
HERE* READ CAREFULLY
Experiment - SEPARATION AND IDENTIFICATION OF THE COMPONENTS OF A
Analgesic drugs reduce pain and antipyretic drugs reduce fever; some drugs, such as aspirin, do
both. Most of the common over the counter analgesic/antipyretic drug preparations contain
aspirin, acetaminophen, or combinations of these substances with other ingredients. You can see
from the chemical structures of acetaminophen, phenacetin, and acetanilide (Figure 1) that they
are structurally related.
Figure 1. Chemical structures of analgesic drugs
As is often the case for scientific discoveries, the pain-killing and fever reducing effects of
acetanilide were discovered by chance. In 1886, Arnold Cahn and Paul Hepp were trying to treat
patients with intestinal worms. They tried everything in their stock room. One patient was treated
with a substance labeled as naphthalene. The chemical did not affect the worms, but it did
effectively treat the patient’s fever and pain. Since the substance did not have the traditional
moth ball smell, it was sent to Hepp’s cousin for analysis. Tests indicated that the new drug was
Acetanilide was soon marketed under the name antifebrin. It is a pain reliever and fever reducer,
but it has a nasty side effect. Some patients who took antifebrin developed methemoglobinemia,
a condition in which hemoglobin molecules are altered in a way that reduces their ability to
transport oxygen. Acetanilide is considered too toxic for medicinal use today, but its discovery
led to the development of safer and more effective analgesic and antipyretic drugs.
Just a few months after the discovery of antifebrin, Carl Duisberg had to get rid of 50 tons of paminophenol, a by-product of dye manufacturing. Duisberg knew that many compounds with a
hydroxyl (-OH) group attached directly to a benzene ring were toxic, so he decided to hide the
hydroxyl group by converting the alcohol to an ether. Converting the primary amine (-NH2) to
the amide generated phenacetin (Figure 2). Phenacetin was used for years as an effective
analgesic/antipyretic drug. Only recently was it banned in the United States since it can cause
kidney damage, hemolytic anemia, or even cancer in some patients.
If you look carefully at the synthesis of phenacetin, you might notice that if Duisberg had not
masked the hydroxyl group, he would have synthesized acetaminophen, which has proved to be
safer than either acetanilide or phenacetin as an analgesic.
Organic chemistry lab techniques can be used to solve a variety of everyday “real” problems.
You will be asked to solve such a problem in this experiment. We have discovered an over-thecounter painkiller that looks as if it might be a counterfeit. The label says the product is
manufactured in the United States, but a savvy consumer noticed minor discrepancies in the
labeling and noticed that the pills in the bottle seemed to crumble more easily than the previously
purchased bottle of the same painkiller. These discrepancies caused her to contact us. We have
been charged with determining if the pills in the bottle are what the bottle claims them to be.
The label indicates that each tablet contains 200 mg of aspirin, 250 mg of acetaminophen, and 50
mg of sucrose. The sucrose is simply added to make the pills taste better. The other components
are painkillers. Initial analysis does indicate that the tablets contain sucrose and aspirin, but the
third component is still a mystery. The third component of the mixture is probably very similar
in structure to acetaminophen, being either acetanilide or phenacetin. Both acetanilide and
phenacetin are effective painkillers, so a substitution would be unnoticeable to the consumer, but
they are both banned in the United States because of their toxicity. If either of the banned
substances is present in the over-the-counter product, then we will need to have the product
removed from the shelves and turn the case over to the FDA. It is your job to find out what
percentages of aspirin, sucrose, and the unknown painkiller component are present in the tablets,
and to determine the true identity of the third component of the painkiller.
You will be separating the three components of the painkiller by making use of their solubilities
and acid-base properties. The painkiller contains sucrose, aspirin, and an unknown third
component. You will find these solubility characteristics to be of utmost importance:
1. Sucrose is soluble in water but insoluble in the organic solvent dichloromethane (also
called methylene chloride, with chemical formula CH2Cl2).
2. Aspirin is soluble in dichloromethane but relatively insoluble in water. Sodium
hydroxide, a strong base, converts aspirin to a salt that is insoluble in dichloromethane but
soluble in water.
3. Acetanilide and phenacetin, like aspirin, are soluble in dichloromethane and insoluble in
water. They are not converted to salts by sodium hydroxide.
Dichloromethane will dissolve aspirin and the unknown component, but not sucrose. The sucrose
can then be removed from the solution of the other two components by gravity filtration
To separate the aspirin from the unknown component, you will take advantage of the acid/base
properties of aspirin and the fact that the solubility of aspirin changes when it is deprotonated
Equation 1. Formation of sodium acetylsalicylate from aspirin
A dichloromethane solution of aspirin and the unknown component can be placed in a separatory
funnel and extracted with an aqueous solution of sodium hydroxide. The salt of aspirin will move
to the water layer. The unknown component will stay in the dichloromethane. The layers can
then be separated from one another.
A separatory funnel is shown in Figure 3. The funnel is fitted with a stopcock and a glass
stopper. The separatory funnel is top heavy, especially when filled with liquid, and should be
supported in an iron ring of proper size attached to a ring stand. Before adding any liquid to the
separatory funnel, make sure that the stopcock is closed.
To perform a separation, add the two immiscible liquids for the extraction into the separatory
funnel, never filling the separatory funnel more than about 3⁄4 full to allow room for mixing.
Place the stopper in the separatory funnel and, holding the stopper firmly in place with the index
finger of one hand, remove the separatory funnel from the iron ring (Figure 4). Invert the funnel,
pointing the stem up away from you and away from everybody else, and carefully open the
stopcock to vent any gases. You may hear a hissing sound as the gases are released through the
stopcock. Close the stopcock and gently shake the two liquids together to mix them for several
minutes. Since the reaction between aspirin (in the dichloromethane layer) and the sodium
hydroxide (in the aqueous layer) only occurs where the two liquids meet, the mixing step is key
to the extraction process, and should not be rushed. It is common for pressure to build up in the
separatory funnel during mixing, especially when using a volatile organic solvent, so it is
necessary periodically to stop mixing and open the stopcock to vent any pressure that has built
After mixing is complete, close the stopcock and place the separatory funnel in the iron ring,
remove the stopper and allow the layers to settle and separate; a clean interface should form
between the two layers. Note: if bubbles have formed at the interface, these can be broken up by
using a glass rod or spatula to disturb the interface. Place an Erlenmeyer flask or a beaker under
the funnel, slowly open the stopcock and drain the lower layer into the beaker or flask. As the
interface between the two solvents approaches the stopcock, slow the rate of draining by
adjusting the stopcock. If you are going to perform another extraction on the top layer, leave it in
the separatory funnel and add more of the next extracting solvent to the separatory funnel. If,
however, you are going to perform another extraction on the lower layer, pour the remaining
upper layer out the top of the separatory funnel into a separate flask and return the lower layer to
the separatory funnel for the next extraction.
A common problem that you may face in doing extractions is trying to determine which one of
the two layers in the separatory funnel is the aqueous layer and which one is the organic layer.
The heavier layer (i.e., the more dense liquid) is the lower layer, of course, but some organic
liquids (e.g., benzene, diethyl ether, ethyl acetate) are lighter than water and some (e.g.,
chloroform, dichloromethane) are heavier than water, so, depending on the solvents being used,
the organic phase might be the upper or the lower layer in the separatory funnel. If you are
unsure about which layer is which, carry out a simple test: mix a small sample of each layer with
a few drops of water in a test tube and see if the two liquids are miscible; the aqueous solution
will dissolve the water drops but the organic solution will be immiscible with the added water.
The GOLDEN RULE of extraction is to never throw away a layer until the end of the experiment
or until you are absolutely certain that you no longer need it. The most common mistake made
during extraction is throwing the wrong layer down the drain or in the waste bottle, which means
loss of material and starting experiment over from the beginning.
Once the aspirin and the unknown component have been separated from each other using the
separatory funnel, they must be recovered. One of the simplest ways to recover the aspirin is to
add acid, thereby converting the sodium acetylsalicylate back to aspirin and changing the
solubility so that the aspirin precipitates from the aqueous solution. Cooling the aqueous layer in
ice will help with precipitation. The aspirin can be isolated by vacuum filtration. The unknown
component can be isolated by drying the organic solution, removing the solution from the drying
agent, and evaporating the solvent. What remains in the flask is the unknown component of the
To determine the identity of the unknown painkiller, you will find its melting point. The melting
point of a pure solid organic compound is one of its characteristic physical properties along with
molecular weight, boiling point, refractive index, and density. A pure solid will melt
reproducibly over a narrow range of temperatures, typically 1-2 °. Note that a melting “point” is
actually a misnomer, and melting points of compounds are usually recorded as a range.
Melting points are determined for three reasons. If the compound is known one, the melting
point will help to characterize the sample in hand. If the compound is new, then the melting point
is recorded in order to allow future characterization by others. Finally, the range of melting point
is indicative of the purity of the compound; an impure compound will melt over a wide range of
temperatures with lower than recorded melting range.
The process of determining this melting point range is done on a truly micro scale using less than
1 mg of the material. The apparatus is simple, consisting of a thermometer, a capillary tube to
hold the sample, and a heating bath. The DigiMelt and Mel-Temp apparatuses (Figure 5) consist
of an electrically heated aluminum blocks that accommodates three capillaries. The sample is
illuminated through a viewing lens.
As the sample is heated to its melting point, two temperatures are recorded. The first temperature
in the range is recorded when the sample just begins to barely melt. The second temperature in
the range is recorded when the entire sample has become liquid.
The accuracy of the melting point depends on a on the accuracy of the thermometer and on a rate
of heating. At the melting point the temperature rise should not be greater than 1 °C per minute.
This may seem extraordinary slow, but it is necessary in order for heat from the bath to be
transferred equally to the sample and to the thermometer, and to obtain an accurate reading.
Both acetanilide and phenacetin are known compounds with reported melting points. By
analyzing your unknown’s melting point, you can determine which of the two compounds has
been substituted for acetaminophen in the counterfeit tablets.
This lab takes place over two weeks. In the first week, you will separate the three components of
the painkiller. In the second week of the experiment, you will recrystallize your unknown
compound and determine the identity of your unknown compound by melting point comparisons.
Separation of Sucrose
Accurately weigh approximately 1.5 g of the painkiller and transfer it to a clean, dry 100 mL
beaker, then add 35 mL of dichloromethane. Stir the mixture thoroughly to dissolve as much of
the solid as possible, using a stirring rod to break up any lumps or granules. Flute and preweigh a
piece of filter paper, then filter the mixture into a 125 mL Erlenmeyer flask by gravity to collect
the sucrose. When not actively pouring the mixture, cover your funnel with a small watch glass
to deter solvent evaporation. Set the filter paper aside in a small beaker and allow it to dry. At the
end of the lab period, determine the weight of the sucrose. If some sucrose remains in the
original flask, allow it to dry, then transfer it into the filter paper before obtaining the weight at
the end of the lab period.
Note: If a significant amount of crystals collects at the top of the filter paper as you are
performing the gravity filtration, you will need to wash the crystals with small portions of
dichloromethane to get them to dissolve again. Some dichloromethane evaporates from the top of
the filter paper as the gravity filtration is taking place. As the solvent evaporates, it leaves the
aspirin and unknown component at the top of the paper. Placing the watch glass over the filter
paper immediately after pouring your solution can inhibit evaporation and crystal formation. If
you leave significant crystals behind on the filter paper, your sucrose weight will be too high, the
other two component weights will be too low, and you will not effectively separate the
components of your mixture.
Separation of Aspirin
Transfer the filtrate to a separatory funnel and extract the aspirin with a 15 mL portion of
aqueous 1 M NaOH by mixing the layers thoroughly. The density of CH2Cl2 is 1.34 g/mL. Stop
and think, which layer will be on top? Collect each layer separately, then return the
dichloromethane layer to the separatory funnel and extract it again with a fresh 15 mL portion of
aqueous NaOH. Collect the aqueous layer with the first one. Save the dichloromethane solution
for the next part of the experiment.
To the combined aqueous extracts, add up to 5 mL of 6 M HCl slowly, with stirring. Once you
see significant aspirin precipitate that does not redissolve upon stirring, test the pH of the
solution with pH paper to confirm a pH of 2 or lower. Do not dip the pH paper directly into your
solution; use a glass rod to transfer a drop of the solution to the pH paper. If your pH is above 2,
add more acid dropwise, and retest the pH. Cool the mixture in an ice bath, collect the aspirin by
vacuum filtration and use a minimal amount of ice-cold water to aid in the transfer. Allow the
aspirin to dry while pulling vacuum. Once it is completely dry, determine the weight of the
Isolation of the Unknown Component
Dry the dichloromethane solution with anhydrous sodium sulfate (Na2SO4), then decant the
solution into a dry, preweighed 125 mL Erlenmeyer flask.
Evaporate the solvent from the dichloromethane solution with minimal heat using a hot plate in
your hood (remember to add a boiling stick!). If the unknown is quite impure, it may remain a
liquid after all the solvent is removed, but it should solidify upon cooling. This is your crude
painkiller (acetanilide or phenacetin), which you will save for the following week.
Purification of the unknown component
Weigh your flas ...
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