CHEM 102 lab packet_W13

Chemistry 102 – Laboratory Experiments

Table of Contents

Chemistry 102: Introduction to Organic Chemistry

Experiment 1:

Natural Indicators…………………………………………………….…..……2

Experiment 2:

Equilibrium……..………………………………………………………….…..7

Experiment 3: Nuclear Chemistry…………………………………...……………………….15

Experiment 4: Hot Stuff……………………………………………………………………...23

Experiment 5: Like Dissolves Like………………………………………………...………...24

Experiment 6: Identification of an Unknown Liquid………………………………………...27

Experiment 7: Identification of Plastics…………………………………………………...…31

Experiment 8: Biodiesel……………………………………………………………………...38

Experiment 9:

Analysis of Vinegar…………………………………………………………..46

Experiment 10: Thin-Layer Chromatography of Analgesics…………………………...……..51

1


Experiment #1 – Natural Indicators

Learning Objectives

The objectives of this lab are:

-To learn the relationship between pH and the acidity/alkalinity of a solution

-To prepare a natural indicator (Red Cabbage) and a color chart showing the color of the indicator for various pH values

-To determine the pH of unknown solutions using the prepared indicator and color chart

Introduction

Review sections 8.1 – 8.2 in your text prior to coming to lab. Acids and bases are types of chemicals that we encounter on a daily basis. Cleaners, cosmetic products, food products, etc. are some examples of substances that are often acidic or basic. In this lab we will create an indicator solution from a natural product (Red Cabbage) and add it to standards of known pH. We will then create a chart that shows the color of the indicator along with the pH of the solution. Finally, we will use the indicator and color chart to determine the pH of unknowns.

Materials

Red cabbage leaves

Watch glass

250 mL beaker

Wellplate

Toothpicks

Drano

Sugar water

Household ammonia

Chlorox

Medicine dropper

Colored pencils

2


Buffered solutions of pH 2, 4, 6, 7, 8, 10, and 12

Litmus paper

Dilute HCl, 0.1 M and 1.0M

Dilute NaOH, 0.1 M and 1M

Vinegar

Baking soda

Washing soda

Liquid soap

Spic and span

Lemon juice table salt

Experimental Procedure a) Part 1: Extraction of indicator and generation of color chart

Break a leaf of red cabbage into small pieces and press them against the bottom of a 250 mL beaker. Just enough to cover the bottom of the beaker should be enough.

Add enough DI water to just cover the leaves. Do not add a large quantity of water as this would require a longer period of boiling to concentrate the solution. Cover the beaker with a watch glass so that the center of the watch glass will cause condensing water vapor to drip back into the center of the beaker.

Gently boil the cabbage until the solution is a deep purple color.

Using the plastic wellplate, place 1 mL of buffer solutions of pH 2, 4, 6, 7, 8, 10 and 12 into separate wells (be sure to keep track of which solution is in each well). Use the 1 M NaOH solutions for pH 14 and the 1M HCl solution for pH 0. You should have 9 wells filled in all.

Add (dropwise) enough of the cabbage juice indicator to each of the wells to produce a pronounced color. Use a toothpick to swirl the contents of each well and obtain a uniform color throughout the solution. Make sure to rinse the toothpick with DI water between wells to prevent contamination.

With the colored pencils match colors and construct a color chart in the report sheet. It may be helpful to put a white sheet of paper under the wellplate to better visualize the color changes.

Save the colored solutions for reference in the second part of the experiment.

3

Chemistry 102 – Laboratory Experiments b) Part 2: Application of indicator and color chart to determine pH

Place about 1 mL of the solution labeled 0.1 M HCl into a clean well and add enough cabbage juice indicator to give a pronounced color. Compare its color with the color of the buffer solutions you prepared as a reference. Judging by the color, what is the pH? Record your observations.

Determine the pH of each of the materials and record in the table provided. If the material is in solid form, you will need to make your own saturated solution. To do so, transfer a small amount of solid into one of the wells and dilute with DI water. Mix thoroughly. There should be a little undissolved material left in the bottom of the well, if not add more solid. Next, add a few drops of the cabbage juice indicator and observe the color of the solution. Record the pH.

For each of your standard solutions and unknowns, you will also determine the concentration of hydronium ions (H

3

O + , or simply H + ). The equation below can be used to calculate [H + ]: pH = -log(H

+

)

This equation can be rearranged to solve for H

+

:

10 -pH = [H + ]

4


Name ______________________________ Section ______________ Date _______________

Partner_____________________________ TA initials_____________________ pH of buffers pH of buffer

0

2

4

6

7

8

10

12

14

Color Color Chart [H + ] Condition of solution

(acidic/basic/neutral) pH of common household products

Solution pH [H + ] Condition of solution

Observations

0.1 M HCl

0.1 M vinegar

Baking soda

Washing soda

Drano

Soap

Salt water

Sugar water

Spic and span

Household ammonia

Chlorox

Lemon juice

0.1 M NaOH

1. If the hydrogen ion concentration, [H + ], in an aqueous solution is 10 -8 , what is the pH of the solution?

5


2. In question #1, would the solution be acidic or basic?

3. What is the pH for a solution with a [H

+

]=5.0 x 10

-4

?

4. If the pH is 7.34, what is [H

+

]?

5. If the purple cabbage juice is not dark enough, what should you do to deepen its intensity?

6. What could happen to the colors of the pH solution for the first part of the experiment if they were not buffered?

6


Experiment #2 – Equilibrium



-To investigate how reactant/product concentrations change with time as equilibrium is reached

-To explore Le Chatelier’s Principle

-To predict shifts in equilibrium according to Le Chatelier’s Principle

Introducion

Review section 7.4 in your text prior to coming to class. Chemical equilibrium and Le

Chatelier’s principle are typically difficult concepts for students to grasp, but extremely important in understanding chemistry. In this experiment, you will perform two different exercises to help reinforce the concept of chemical equilibrium. The first is a modeling exercise in which you will allow a “chemical” system to reach equilibrium. In the second exercise, you will explore Le Chatelier’s Principle using an actual chemical system.

Materials

100 macaroni pieces

Paper (“reaction flask”)

Spreadsheet software

Well plate

2 test tubes

0.01 M FeCl

3

0.1 M FeCl

3

0.01 M KSCN

0.1 M KSCN

1M NaOH

0.1 M Na

2

HPO

4

Hot plate

7


Ice

Droppers

Experimental Procedure

Part 1: Equilibrium simulation

When a system is at equilibrium, the reaction will have equal rates of reaction in both the forward and reverse directions (i.e. reactants will be converting to products at the same rate as products convert to reactants). At equilibrium, the concentrations of reactants and products doesn’t change, however, the actual molecules that are reactants and products does change. A simple analogy is people on a dance floor. The number of couples dancing may not change, but the actual dancers changes with each new song.

Count out 100 pieces of macaroni. Obtain a piece of paper towel; draw a line dividing the paper in half length-wise. Label the left side of the paper A (reactants) and the right side of the paper B

(products).

The reaction you will model is A



B. When a piece of macaroni is on the A side of the paper, it is a reactant molecule. When it is on the B side of the paper, it is a product molecule.

Place all of the macaroni pieces on side A. At time 0 seconds, there are 100 A molecules and 0 B molecules. At each step (one second) 10% of A converts to B and then 10% of B converts back to A.

For example: at one second, 10% of A (10 pieces of macaroni) converts to B, giving 90 pieces of macaroni on the A side and 10 on the B side. Remember, however, that B converts back to A at a rate of 10%, so 1 macaroni piece of B goes back to A. The final count after 1 second is 91 pieces of macaroni on the A side and 9 pieces of macaroni on the B side (remember that macaroni represents molecules).

Complete the table in the data sheet for the remaining trials. Do not record fractional numbers of molecules (round to the nearest whole number), just make sure that your total (A + B) is always

100 molecules (pieces of macaroni).

Once you have completed the table, open up a spreadsheet and enter your data. Plot A and B concentrations on the Y-axis and time on the X-axis. Print your curve and attach to your data sheet.

Part 2: Le Chatelier’s Principle

According to Le Chatelier’s Principle, when a stress is applied to a system at equilibrium, the system will adjust itself to reduce this stress and re-establish equilibrium. Using the example of dancing couples, let’s consider the equilibrium reaction:

8


Men + Women = Dancing couples

We can write the equilibrium constant expression from this reaction:

Keq = [Dancing couples]/[Men][Women]

If we have 20 men and 20 women at the dance and at equilibrium half of the men and women go on the dance floor, we can then calculate the equilibrium constant (Keq).

Keq = [10 dancing couples]/[10 non dancing men][10 non dancing women] = 0.1

If 20 more men arrive, a stress has been applied to the equilibrium of the dance. The poor ladies who were without a dancing partner now have a greater selection of dancing partners. What will happen to the number of dancing couples? Intuitively, we would expect there to be more dancing couples, but how many more?

Let’s do the numbers….At the start of the stress, we have 10 couples, 30 non-dancing men and

10 non-dancing women.

Keq = 10/(30)(10) = 0.033, this value doesn’t match our above value of 0.1, so something must happen to get the dance floor back in equilibrium.

To get back to equilibrium we need more dancing couples, which will reduce the number of nondancing men and women. If we had 14 dancing couples, then we will have 26 non-dancing men and 6 non-dancing women.

Keq = 14/(26)(6) = 0.09, this value is much closer to our original Keq value from above. So, we have a system in equilibrium, we add more men (which is a reactant in our system) and the result is an increase in dancing couples that is dictated by the equilibrium constant.

Things to remember about LeChatelier’s Principle, when a stress is introduced to a system at equilibrium, the system will shift to relieve this stress. For example:

-if more reactants are added, more product is formed

-if more product is added, reactants are formed

-if we remove reactant, more reactants will form

-if we remove products, more products will form

-heat can be treated as a reactant in an endothermic reaction and a product in an exothermic reaction

For this exercise, you will be investigating the effects of stresses applied to the following chemical system at equilibrium:

FeCl

3

(aq) + 3KSCN(aq)



3KCl(aq) + Fe(SCN)

3

This reaction can be condensed (by removing spectator ions) to:

Fe

3+

(aq) + 3SCN

-

(aq)



Fe(SCN)

3

(aq)

9


It is well suited for this exercise because the reactants are a very distinct yellow color, while the products are red, allowing you to visibly see how the equilibrium shifts in response to different applied stresses.

Obtain a clean, dry well plate and place it on a white piece of paper. Obtain two test tubes. Label the paper corresponding to wells 1 – 7 as follows:

#1 – Blank (control)

#2 – FeCl

3

#3 – KSCN

#4 – NaOH

#5 – Na

2

HPO

4

#6 – Heated

#7 – Cooled

Add 10 drops of 0.01 M FeCl

3

, 10 drops of 0.01 M KSCN, and 10 drops of DI water to each well except #6 and #7. Label your test tubes #6 and #7 and add the same amount of reagents as above.

Mix all of the wells using your stirring rod (cleaning it after every mix).

Add an additional 5 drops of DI water to well #1 and test tubes #6 and #7. Well #1 represents the blank, or control and will be used to compare to the rest of the wells.

Set up a hot water bath (almost boiling). Place test tube #6 in the bath for 5 minutes.

Set up an ice bath. Place test tube #7 in the ice bath for 5 minutes.

Add 5 drops of 0.1 M FeCl3 to well #2, mix and compare to color in well #1.

Add 5 drops of 0.1 M KSCN to well #3, mix and compare to color in well #1.

Add 5 drops of 1 M NaOH to well #4, mix and compare to color in well #1. Note any cloudiness that may have developed. Let sit for at least 10 minutes before disposing.

Add 5 drops of 0.1 M Na

2

HPO

4

to well #5, mix and compare color to well #1. Again, look for cloudiness after 10 minutes.

Remove test tube #6 and pour into well #6. Compare to well #1.

Remove test tube #7 and pour into well #7. Compare to well #1.

Discard of your solutions in the appropriate waste container and clean and dry your well plate.

10



Partner______________________________ Instructor Initials_________

Part 1. Equilibrium Simulation

Time

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

16

17

# of A molecules

100

91

# of B molecules

0

9

Print and attach a copy of the curve you generated.

1. When does the net reaction (i.e. rate at which reactants become products) proceed at the fastest rate? When is the rate the slowest?

2. How does the reversibility of the reaction affect the concentrations of reactants and products?

How would your graph look if the reaction was not reversible?

3. Did your reaction reach equilibrium? How do you know when equilibrium is finally reached?

11


4. What does the term “dynamic equilibrium” mean?

5. What would your graph look like if the rate of conversion from A to B was 25% per second instead of 10%?

Part 2. LeChatelier’s Principle

Record observations in table below:

Well Number

5

6

7

1

2

3

4

Stress added Observations compared to well #1

Explanation according to Le

Chatelier’s Principle

6. Write Keq for the reaction observed in part 2 of the experiment. a. If the color turned red, which side of the reaction is favored? b. If the color turned yellow, which side of the reaction is favored?

7. What is the purpose of having a control or blank? Why were 5 drops of water added to well #1 and test tubes #6 and #7?

12


8. Based on your observations, is this reaction endothermic or exothermic? Explain.

9. In the body, an equilibrium exists between calcium, phosphate and calcium phosphate as shown below. Kidney stones are composed mainly of calcium phosphate. Explain why a large increase in calcium intake may pose a health risk.

3Ca

2+

(aq) + 2PO

4

3-

(aq)



Ca

3

(PO

4

)

2

(s)

10. Hydrogen gas production can be accomplished by reacting carbon monoxide with water vapor. An equilibrium is established:

CO(g) + H

2

O(g)



H

2

(g) + CO

2

(g) + 9.9 kcal

Predict how the production of H

2

(g) would be affected by the following conditions: a. Increase in CO(g)concentration. b. Decrease in temperature. c. Removal of water vapor. d. Removal of CO

2

(g).

13


14


Experiment #3 – Nuclear Chemistry



-To use a Geiger-Muller tube and associated equipment

-To graph a half-life using a coin toss model

-To investigate the relationship between radiation exposure and distance

-To learn about the different types of radiation

-To be able to characterize the type of radiation from a common kitchen plate

Introduction

Review sections 9.1 – 9.3, and 9.7 in your text prior to coming to lab. Radiation is something we are exposed to every day. There are cosmic as well as rock sources that are constantly undergoing nuclear transitions and thus releasing radioactivity. Along with these natural sources, there are also anthropogenic sources due to human demand for energy. The Hanford Nuclear

Plant, located just a couple of hours away from LaGrande, is an example. Nuclear isotopes are also finding broader use in the medical field. Nuclear chemistry is therefore a ubiquitous type of chemistry that is important for everyone to understand. In this lab we will model radioactive decay to determine the half-life, we will use a Geiger-Muller tube to measure levels of radiation, and we will investigate the shielding properties of different materials and different distances from a source.

Materials

Geiger-Muller tube, scalar and timer

Meter stick

Fiesta ware plate

Ring stand and clamp

Various filtering materials

Clay and beads

Pennies

Box with lid

15



Part 1: Modeling half-life

This simple exercise will help you in understanding half- lives of radioactive isotopes.

Obtain 100 pennies; lay them all head side down in a box. Put the lid on the box and give it one good shake. Remove the lid and remove the pennies which show head side up and count them.

Record the number in your data sheet.

Replace the lid and repeat until you have fewer than 5 pennies in the box. Record all your numbers in your data sheet.

Open a spreadsheet program and enter your numbers, plot the toss number on the x-axis and the number of tails remaining on the y-axis.

On the y-axis, locate the point at which 50 tails remain, from your graph determine the toss number that corresponds to this point. Record in data sheet.

On the y-axis, locate the point at which 25 tails remain, from your graph determine the toss number that corresponds to this point. Subtract the first interval half-life from your second toss number, this is your second interval half life. Record in data sheet.

On the y-axis, locate the point at which ~12 tails remain, from your graph determine the toss number that corresponds to this point. Subtract the second interval half-life from your third toss number, this is your third interval half-life. Record in data sheet.

Part 2: Modeling radioactive decay

For this part of the experiment you will build clay models of atomic nuclei to discover two types of radioactive decay. You will make a hypothetical nucleus called “mysterium” (My).

From the clay, roll 5 pea-sized balls. These represent the protons in the nucleus of My.

Now roll an additional 10 pea-sized balls and insert one colored bead into each. These represent neutrons in the nucleus of My.

Loosely pack the protons and neutrons together to form your final My nucleus which is represented by

15

5

𝑀𝑦 .

You will now perform a radioactive decay by removing 2 protons and 2 neutrons from the nucleus and placing them aside. This represents alpha decay, or the loss of a He nucleus.

Determine what is left over in the original My nucleus and denote it as “X”. Complete the balancing of the radioactive disintegration shown below by filling in the atomic mass and number of “X”.

15

5

𝑀𝑦  _____?

_____?

𝑋 +

4

2

𝐻𝑒

16


X is the disintegration product of My (also called the daughter nuclide of My), it will also undergo radioactive decay by losing a beta particle (electron). This is done by removing one of the beads from one of the neutrons in the nucleus of X.

Determine what is left over in the original X nucleus and denote it as “Y”. Complete the balancing of radioactive disintegration shown below by filling in the atomic mass and number of

“Y”.

?

?

𝑋  ______?

______?

𝑌 +

0

−1 𝑒

Part 3: Distance and intensity of nuclear radiation

In this portion of the lab, you will use the Geiger-Muller tube, scalar and timer to determine the relationship between exposure and distance from a radioactive source.

Turn on the power to the Geiger-Muller tube, scalar and timer. Follow the directions posted on the top of the scalar. Practice using the COUNT toggle and reading the timer.

Measure the “background” radiation with the Fiesta Ware plate well away from the counter.

Read the background for 30 seconds. It is unnecessary to re-zero the timer after each use. It will automatically stop after the COUNT toggle is turned off. Record the background in your data sheet. Measure the background twice and average your results.

Place a Fiesta Ware plate in the clamp on the ring stand. Center it in front of the Geiger-Muller tube. Place the plate so that it is 5 cm from the counter. Count for 30 seconds, record your reading.

Move the plate to 10, 20, and 40 cm from the Geiger counter. Measure the radiation counts at each interval and record your reading.

Remove the plate from the clamp and hold it directly in front of the Geiger counter, this will be your 0 cm reading, record in the data sheet.

Correct all of your readings by subtracting the average background reading.

Open a spreadsheet program and plot counts/minute (y) vs. distance in cm (x). Be sure to label axis.

Part 4: Absorption of nuclear radiation

In this part of the lab, you will investigate how different types of shielding materials protect the

Geiger counter from exposure. From this information, you should also be able to determine the types of radiation emitted by the Fiesta Ware.

Return the plate to the 10 cm mark and repeat the counts for 30 seconds. Record in the data sheet.

17


Place a piece of paper in front of the Geiger counter and measure for 30 seconds. Record in the data sheet.

Repeat the previous step using the glass, wood and lead disc. Record results of all in the data sheet.

Calculate the percent of shielded radiation/total radiation for each shielding material. Use the results with no shield as the total radiation.

18



Partner_____________________________ TA initials_____________________

Part 1. Modeling half life:

Toss number Heads removed

Print and attach the graph generated from this data.

Tails remaing

1 st

t1/2 = ____________________

2 nd t1/2 = ____________________

3 rd t1/2 = ____________________

Part 2. Modeling radioactive decay

Give the identities of the nuclei you created in part 2 based on the number of protons:

My:______________

X:_______________

Y:_______________

19


Part 3. Distance and intensity of nuclear radiation:

Background reading: __________ Average: ___________

Distance (cm)

0

10

20

40

60

Counts Counts - Background

Print the graph you obtained from plotting Counts/minute vs. Distance in cm.

1. Describe how the behavior in the count rate changes as a function of distance from a radioactive source.

2. Using the plot of counts vs. distance, estimate the number of counts at 15 cm from the source.

Part 4. Absorption of nuclear radiation:

Shielding material

No shield

Paper

Glass

Wood

Lead

Counts/0.5 minutes Counts - background

3. Which type of radiation is most easily stopped?

% total counts

100

4. Which of the shielding materials provided the greatest protection from radioactivity?

20


5. What type of radiation does this material stop?

6. Which types of radiation are stopped by the following: a. no shield b. paper c. wood d. lead

10. According to your data, which types of radiation are being emitted by the Fiesta Ware plate?

Which appears to be prevalent? Explain.

11. During nuclear reactions, the sums of atomic masses and atomic charges must be conserved.

Following these rules, balance the following nuclear reactions that occur in the synthesis of americium-241 by filling in the boxes. Am-241 is the isotope used in smoke detectors and further decays by alpha emission. Plutonium is the starting material in this series (particles are all shown over the reaction arrow, those in parentheses are reactants those that aren’t are products).

21


12. Fiesta Ware, brightly colored glazed pottery, is now considered “collectible.” The pre-60’s orange colored dishes contained uranium oxide which is radioactive. Fill in the blanks of the following nuclear decay reaction for uranium.

22


Experiment #4 – Hot Stuff


The objectives of this experiment are:

-To design and execute a novel experiment

-To determine the temperature of hot water without the use of a thermometer that goes high enough.

Intoduction

This is a very different kind of experiment from most of the other experiments in this lab packet.

It poses a scientific problem but gives no suggestions of how to solve it. In groups of 3-4 students you will plan your own approach, collect data and report your answers to the problem.

Assorted glassware and other equipment will be available to carry out your experiment. The problem in this case is to determine the temperature of hot water when there is no thermometer that goes that high.

Prior to conducting your experiment you must get approval for your procedure from your lab instructor.

Following the completion of this experiment you will turn in a brief description of your experiment, the data you collected and your conclusions.

23


Experiment #5 – Like Dissolves Like



-To investigate the solubilities of different substances in polar and nonpolar solvents.

-To correlate structures of different substances to their solubilities.

-To use the “like dissolves like” rule to make solubility predictions

Introduction

The idea behind the “like dissolves like” rule is that molecules with similar chemical properties

(polarity in this case) will interact to a greater degree. To put it another way, polar substances will dissolve in polar solvents, while nonpolar substances will dissolve in nonpolar solvents. This relationship provides chemists another tool at their disposal for understanding different materials.

Polarity is determined largely by the intermolecular forces present in substances. Intermolecular forces are those forces that lead to attraction or repulsion between particles or molecules when they are close together. Compared to intramolecular forces (the forces holding atoms in a molecule, e.g. covalent and ionic forces), intermolecular forces are quite weak. These forces include: 1) London dispersion forces – weak attractive forces due to temporary dipole formation, these occur for all molecules, 2) Dipole-dipole forces – attractive forces between molecules with permanent dipoles (look for Cl, N, O, or F – highly electronegative atoms create polar bonds and hence permanent dipoles), and 3) Hydrogen bonding – attractive force between –OH or –NH groups. You will want to refer back to this description when answering the questions in the data sheet.

In this lab you will experiment with solvents and solutes with varying polarities to determine if the “like dissolves like” rule holds true. You will also identify whether or not a few unknown compounds are polar or nonpolar. Finally, you will use the structures of known compounds to determine which intermolecular forces are active.

Materials

-10 test tubes and rack

-Dropper

-Stirring Rod

-Spatula

-NaCl

24


-Sucrose

-Paraffin wax

-Crisco

-Unknown solutes (A, B, and C)

-water


Label 5 test tubes “water” and 5 “hexane.” Add 20 drops of water and 20 drops of hexane to the appropriately labeled test tubes.

To one “water” test tube and one “hexane” test tube, add a small amount (0.25 gram or so…) of

NaCl. Stir each tube using your stirring rod and record your observation. Not all of the solute has to dissolve for the substance to be considered soluble.

Repeat the process with the remaining substances, record your observations in the table provided in your data sheet.

25




Substance

NaCl

Sucrose

Paraffin wax

Crisco

Unknown A

Unknown B

Unknown C

1. Define soluble and insoluble.

Water Hexane Polarity of substance

2. Draw the structure of water and hexane. What intermolecular forces are present in each?

Which is polar and which is nonpolar?

3. Draw the structure of the 4 known substances and list the intermolecular forces (you may have to do a bit of research to find their structures…). Predict their polarity based on these forces. Are your predictions consistent with your observations?

4. The unknown B substance is known to be an electrolyte. Coupling this information with the solubility data, what can you conclude about the unknown?

26


Experiment #6 – Identification of an Unknown Liquid



-To be able to measure different physical properties of liquids (density, solubility in polar/nonpolar solvents, boiling point).

-To be able to use these properties to identify an unknown liquid.

Introduction

Review density, polarity and boiling point prior to coming to lab. Often times it is the chemists’ job to identify an unknown substance. This can be done using a number of different analytical techniques.

In this lab, you will be the chemist, your job is to identify an unknown liquid. You will use three different physical properties of liquids to determine the identity of your unknown. You will measure the solubility in water vs. hexane (polar vs. nonpolar), density, and boiling point of your unknown. By comparing the results from your unknown with established values for known compounds, you should be able to determine the identity of your unknown.


Part 1: Solubility of a Liquid

Place about 1 mL of water in a small test tube, add 1 mL of your unknown to the test tube.

Stopper the test tube and shake for about 10-15 seconds. Set the tube aside for a few minutes and reexamine, record your observations. If there is a visible boundary between the solvent and the unknown then your unknown is insoluble. Dispose of liquid mixture in appropriate waste container.

Part 2: Density of a Liquid

The following equation can be used to calculate density:

Density (g/mL) = mass (g)/volume (ml)

Pre-weigh a small beaker that is covered with a square of aluminum foil. Attach a pipette pump to a clean, dry 2.00 mL volumetric pipette. Hold the pipette vertically in the unknown liquid and aspirate the liquid by using the thumb wheel. Bring the meniscus exactly even with the etched mark on the pipette.

27


Position your pipette over your pre-weighed beaker and remove the pipette pump (the fluid will drain into the beaker). Cover the beaker with the aluminum foil. *A small amount of liquid will remain in the pipette, do not try to remove it.

Re-weigh the covered beaker. Perform the calculation using the equation shown above. Save the liquid for the next determination.

Part 3: Boiling Point of a Liquid

The boiling point of many liquids are well established. The boiling point of your unknown will serve as the final piece of information that you will use to determine its identity. This experiment will be performed in the hood.

Obtain your largest test tube and pour the liquid from Part 2 into it. Place 1-2 boiling stones in the test tube.

Set a thermometer into the tube using a slotted rubber stopper. Slide the thermometer into the tube so it is about 1 cm above the liquid level.

Using a ring stand, clamp and hot plate, attach your test tube with the thermometer onto the ring stand. Place a 200 – 400 mL beaker that is about ½ full of tap water under the test tube. Adjust the height of the test tube so that it is about 3 cm under the surface of the water in the beaker.

Heat the water gradually and look for changes. Record the temperature at which the liquid in the test tube is boiling vigorously.

Discard the unknown liquid into the waste container provided.

Part 4: Identification of the Liquid

Use the table below to help you identify your liquid unknown.

Compound

Acetone

Cyclohexane methanol

Ethyl Acetate

Dichloromethane

Hexane

Toluene

Diethyl Ether

Cyclopentane

Density (g/mL)

0.79

0.78

0.79

0.89

1.34

0.66

0.86

0.71

0.74

Boiling Point (C)

56

81

65

77

40

69

111

35

49

Solubility in water

Soluble

Insoluble

Soluble

Insoluble

Insoluble

Insoluble

Insoluble

Insoluble

Insoluble

28




1 st

Unknown:_______________

Part 1: Solubility of a Liquid

1 st Unknown:

Observation:

Conclusion:

2 nd Unknown:

Observation:

Conclusion:

Part 2: Density of a Liquid

Mass of beaker + foil + 2.00 mL liquid

Mass of beaker + foil

Mass of Liquid

Density of Liquid

Show calculations:

2 nd

Unknown:_______________

1 st Unknown 2 nd Unknown

29


Part 3: Boiling point of a Liquid: ________________

Part 4: Identification of a Liquid

Solubility

Property

Density

Boiling Point

1 st

Your Values

Unknown 2 nd

Unknown

Final Identification of 1 st

unknown liquid: _________________

Chemical Structure of the Liquid:

Final Identification of 2 nd unknown liquid: _________________

Chemical Structure of the Liquid:

Table Values

1. The mass of 5.00 mL of carbon tetrachloride is 7.97 grams. What is the density of carbon tetrachloride?

2. An unknown liquid is found oto be insoluble in water, has a density of 0.73 g/mL and a boiling point of 35 C. The most likely identity of the liquid is….

3. Using the like dissolves like rule and chemical structure of your unknown liquid, explain its solubility/insolubility in water. Describe the type of bonding that your unknown has and locate any dipoles.

30


4. The atmospheric pressure in La Grande is significantly lower than that of Portland. How would the boiling points of liquid compare in these two locations?

5. Explain and illustrate why ethanol (molar mass = 46 g) has a higher boiling point than diethylether (molar mass = 74 g)

31


Experiment 7 – Identification of Plastics



-To learn the differences between the 6 main classifications of plastics

-To investigate the different properties of the different classes of plastics

-To use the different properties of the classes of plastics to identify an unknown plastic

Introduction

Of the more than 60,000 different plastics manufactured, six of them make up more than 70% of all plastics used in this country. These six are low-density polyethylene (LDPE), high-density polyethylene (HDPE), polyethylene terephthalate (PETE), polyvinyl chloride (PVC), polystyrene

(PS) and polypropylene (PP). They are encountered by most people every day in familiar products, yet most people are unaware of their specific uses and properties. In this experiment you will investigate some of these properties and develop a classification and identification scheme.

The plastic industry has adopted a code for packaging materials that can be used to quickly identify the type of plastic involved. This code is a symbol on the bottom of many containers of a triangle of arrows with a number in the middle of the triangle. The number corresponds to a certain type of plastic, allowing easy recognition for recycling purposes. See table below.

No.

1

2

3

4

5

6

7

Name

Polyethylene terephthalate

High-density polyethylene

Polyvinyl chloride

Low-density polyethylene

Polypropylene

Polystyrene

Other (mixtures)

Abbreviation

(PET)

(HDPE)

(PVC)

(LDPE)

(PP)

(PS)

Compliance in labeling is voluntary, and not all plastics have a symbol. Without these symbols, plastics become difficult to identify and recycle.

In this experiment, you will use four simple tests to aid in classification and identification of the six types of plastics. The first test will be to determine the relative densities of the plastics by checking to see whether the samples float or sink in three liquids of differing densities.

32


The second test will be to melt the plastic. All six of these common plastics melt reversibly, which means that when they are cooled, they harden and may regain their original properties. If a plastic sample does not melt, it is a thermosetting plastic. Thermosetting plastics, such as melamine (used in dinner ware) or the plastic used in electrical components, do not melt cleanly and reversibly, but tend to char instead.

The third test , burning the plastic, must be performed in the fume hood. All common plastics burn, but with slightly different characteristics and fumes. The vapors given off from the burning plastic may have different properties depending on the plastic.

The last test, the copper-wire test, will be used to determine if a halogen, such as chlorine, is part of the polymer.

Materials

Known samples of all six types of plastic

Plastic samples of unknown composition

1:1 mixture of 95% ethanol and water

10% NaCl solution

3 test tubes with test tube rack

Bunsen burner

Stricker

6 inch piece of copper wire and a cork

Metal spatula

Few strips of blue litmus paper

Pair of forceps

Scissors

Glass stirring rod


Part 1: Density test

Three liquids of differing densities will be used, as shown in the following chart.

33


Liquid

1:1 95% ethanol/water

Water

10% NaCl

Density (g/mL)

0.94

1.0

1.08

Pour about 5 mL of each of the above liquids into three separate, labeled test tubes. Obtain two narrow strips of each of the six different types of plastic. Cut one strip of each plastic into 3 small pieces.

For one of the plastics, place a piece into each of the three test tubes with the density test liquids.

Push each piece under the liquid surface with a glass stirring rod or the end of a pencil. If the sample floats, it has a lower density than that of the liquid. If it sinks, it has a density greater than that of the liquid. Repeat with all of the other types of plastics. Record your observations.

Finally, use your observations to sort the six plastics by their densities, filling in the table and the ranking list on the data sheet.

Part 2: Melt test

Place a small sample of each plastic, one at a time, on the end of a metal spatula and hold the end of the spatula over a light blue microburner flame. Heat slowly and observe the plastic as it warms and finally melts. Do not heat strongly enough for the plastic to catch fire. Enter your observations in the data sheet. Cool the sample and examine it for appearance and flexibility by bending it. Enter your observations in the data sheet. The melted plastic can be used for the tests that follow.

Part 3: Ignition test – To be done in the fume hood only!!

Place a Bunsen burner in the fume hood. Light it and adjust to a small flame. Place a large beaker of water in the hood. Hold one end of a small strip of plastic in a pair of forceps and place it directly in the flame. Observe the color of the flame and its characteristics (e.g. is there a lot of smoke or visible vapor? Does the plastic continue to burn when removed from the flame, etc.).

Record your observations.

Test the vapors given off for acid/base character by holding a piece of wet litmus paper in the vapors from the burning plastic. If the paper turns red, acidic fumes are being released. Record your observations.

Extinguish the burning plastic by dropping it into the beaker of water. Repeat the procedure for the remaining plastics.

34


Part 4: Copper-wire test – To be done in the fume hood only!!

Push the end of a 6-inch length of copper wire into a small cork (this apparatus may already be prepared for you). Use the cork as a handle and heat the free end of the wire in a burner flame until the flame has no green color.

Touch the hot copper wire to the plastic you are testing and then return the wire end to the flame.

A tiny bit of plastic should be picked up by the hot wire. Return the wire end to the flame. When the tip of the wire is put in the flame, watch for a slight flash of luminous flame. This indicates that you have correctly picked up a little bit of plastic on the wire.

Watch for the appearance of a green flame or green color in the flame when the plastic is heated.

The green color indicates the presence of chlorine in the plastic. Repeat for your remaining plastics.

Part 5: Identify an unknown plastic

For this part of the experiment, you will develop a scheme to identify two unknown plastic samples using the properties investigated above. Ideally, you will find the minimum number of tests to do to make the identification. After you have devised your scheme, use it to identify your unknown plastic samples.

35



TA initials_____________________ Partner_____________________________

Unknown 1 number:__________

Unknown 2 number:__________

Density test (sinks/floats/can’t tell):

Type of plastic 1:1 95% ethanol/water (0.94 g/mL)

PET

HDPE

PVC

LDPE

PP

PS

Unknown 1

Unknown 2

Place each plastic into one of these categories:

Water (1.0 g/mL) 10% NaCl solution

(1.08 g/mL)

Less than 0.94 g/mL Less than 1.0 g/mL Less than 1.08 g/mL Greater than 1.08 g/mL

Rank the densities of the six types of common plastics:

Lowest _________ __________ __________ __________ __________ __________ highest

36


Melting/ignition test:

Plastic Melting test Ignition test (acidic smoke?)

PET

HDPE

PVC

LDPE

PP

PS

Unknown 1

Unknown 2

Identification of unknown samples:

Unknown 1 number __________ Identity ___________

Unknown 2 number __________ Identity ___________

Copper wire test

37


1. Outline your scheme for identifying your unknown plastic sample.

2. Suppose you had to add two other plastics to your scheme, polymethymethacrylate (density between 1.18 and 1.20 g/mL) and poly-4-methyl-1-pentene (density 0.83 g/mL). Where would they fit in your scheme?

3. Polyethylene terephthalate (PET) is the most valuable waste plastic at the present time.

Suggest a way to separate it commercially from other waste plastics (based on your observations).

4. Why are plastic recyclers very concerned about identifying the different polymers and not mixing them together?

5. Since waste plastic consists mostly of hydrocarbon compounds, it has been suggested that waste plastic could be used as fuel. Based on your observations, do you think this is a reasonable suggestion? Would some be more dangerous to burn than others? Justify your answer.

38




Experiment #8 – Biodiesel

-To learn how to convert vegetable oil into a form of fuel (biodiesel)

-To prepare biodiesel from common vegetable oil

-To investigate some physical and chemical properties (viscosity, effects of temperature, and heat content) of biodiesel

Introduction

Due to concern over finite fossil fuel resources (i.e. a limited amount of gas), many alternative fuels have been investigated. One of these fuels, biodiesel, is a clean burning alternative fuel produced from domestic, renewable resources. It contains no petroleum but can be blended with petroleum diesel to create a biodiesel blend. It can be used in standard diesel engines with little to no modifications. It is easy to use, biodegradable, non-toxic, and free from sulfur and other polluting aromatic compounds.

Biodiesel is produced from any fat or oil such as soybean oil through a process called transesterification. This process reacts the oil with an alcohol to remove glycerin as shown in the reaction below:

In this lab you will use standard vegetable oil to produce your own biodiesel. You will then characterize the biodiesel using a few standard chemical techniques.

Materials

50 mL of vegetable oil

10 mL of methanol

0.5 mL 9.0M NaOH

Salt

39


Ice

Stirring hotplate with stir bar

250 mL beaker

100 mL graduated cylinder

10 mL graduated cyinder

Centrifuge

4 centrifuge tubes

2 disposable pipets, labeled

20 mL beaker

Candle wick on paper clip

1 empty soda can without top

1 thermometer

1 ring stand with iron ring

1 glass rod

Tongs

Matches

Balance


Part 1: Preparation of Biodiesel

Choose and record one of the oils provided. Measure out 25 mL of the oil you chose. Pour oil into 250 mL beaker.

Use a 10 mL graduated cylinder to measure out 5 mL of methanol into a small beaker. Add 0.25 mL (5 drops) of 9M NaOH to the methanol ( *NaOH at this concentration is extremely caustic, be very careful ). Set this mixture aside.

Add a magnetic stir bar to your oil and place the beaker on a stirring hotplate. Adjust the stirring rate so that the solution is being well mixed without splashing. Turn the heat on low.

Heat your oil to 50°C. Hold the thermometer in the oil, but away from the stir bar (the stir bar could break the thermometer, costing you $$).

40


Once your oil reaches 50°C, turn off the heat but continue stirring. Slowly pour your methanol/NaOH mixture into the oil. Be sure the mixture continues to stir so that layers don’t form. Record your observations.

After 20 minutes, remove your beaker from the hotplate. Record your observations. *While waiting for your mixture to cook, you should prepare glass pipets for part 2 below.

Obtain 2 centrifuge tubes. Fill each ¾ full with your biodiesel mixture. Place the tubes opposite of each other in the centrifuge and spin for several minutes.

Turn off centrifuge and allow it to stop before trying to remove your tubes. You should see two distinct layers, if not, put your tubes back in the centrifuge and spin for several more minutes.

The top layer is your biodiesel and the bottom layer is glycerol.

Use a disposable transfer pipet to transfer your biodiesel into a small beaker (you will need about

20 mL of biodiesel in total, if you don’t have this much add more reaction mixture to centrifuge tubes and centrifuge for a few minutes).

Keep the small beaker with your biodiesel. Discard the glycerol in the appropriate waste container.

Part 2: Viscosity of oil and biodiesel

Obtain two 6-inch disposable glass pipets that have been marked with lines on the barrels 4 cm apart. You’ll also need a matching dropper bulb and a timer that can measure in seconds.

We will estimate viscosity by measuring the time required for the liquid level in the pipet to drop from the upper line to the lower line. Viscosity is proportional to this time.

Use the bulb to fill one of your pipets with biodiesel, such that the liquid level is above the upper line. Hold the pipet vertically over the beaker of biodiesel. Quickly remove the bulb and place your index finger firmly over the top of the pipet. This will stop the biodiesel from dripping.

Have your partner ready the timer. Remove your finger from the top of the pipet. Allow the liquid to drain freely back into the beaker. Measure the time required for the liqiuid level to drop from the upper line to the lower line. Record your time.

Repeat the measurement. Record your data and calculate the average. Discard the pipet.

Use the test tube to get about 5 mL of the oil you started with to make your biodiesel. Use your second marked pipet to measure the viscosity of the oil. Repeat the measurement twice, record your data and calculate an average. Save this oil for part 3.

41


Part 3: Temperature effect on oil and biodiesel

Fill a 250 mL beaker most of the way with ice. Add two large scoops of salt to the ice and then add enough water to make the mixture slushy. Stir briefly to mix the salt (the salt lowers the temperature of the ice bath).

Place a thermometer inside your test tube of oil. Note how cloudy or clear the oil appears. Note how viscous the liquid appears as it drips from the thermometer or flows down the inside wall of the test tube. Now place the test tube of oil in your ice bath.

Observe the oil as it cools. Periodically remove it from the ice long enough to see if the oil is becoming cloudy or more viscous. If so, record the oils temperature. Keep the test tube in the ice until its temperature is no longer dropping. Again note the oil’s appearance and apparent viscosity. Record your observations.

Wipe the oil from your thermometer with a paper towel or tissue.

Now put some of your biodiesel in a test tube, add the thermometer, repeat the cooling process and record your observations.

Part 4: Heat content of biodiesel

Prepare a burner for your biodiesel. Obtain a 20 mL beaker, a candle wick, and a paper clip. The end of the candle wick should be wrapped around one wire of the paper clip. Spread the paper clip out slightly so that it makes a stable bottom support for the wick. The wire inside the wick will allow it to stand vertically. Put the paper clip and wick in your beaker. The wick should extend to the top of the beaker. Now fill the beaker most of the way with your biodiesel.

The wick should stand freely in the middle of the biodiesel. Use matches to light the wick and ensure that your burner is working well. Once you know that your burner is working, extinguish its flame.

Obtain a dry soda can with the top removed and two holes punched on opposite sides near the top. Slide a glass rod through the holes in the can so that it can be suspended from the ring attached to a ring stand.

Take the empty soda can and your data sheet to a balance. Tare the balance, then weigh the can to the nearest 0.1 grams.

Using a graduated cylinder, add approximately 100 mL of water to the can. Re-weigh the can plus water to the nearest 0.1 gram. Calculate the mass of the water by subtraction.

Put a thermometer in the can, stir the water for a few moments, and then measure the temperature of the water. Estimate to the nearest 0.1°C if possible.

Take the biodiesel burner to a balance. Weigh the burner to the nearest 0.01 grams, and record its mass in your data sheet.

42


Place the biodiesel burner under the soda can and light the burner. Observe the flame. If necessary, cautiously adjust the height of the can so that the top of the flame is just below the bottom of the can.

Stir the water occasionally and continue heating the water until the temperature has increased about 20 C, then extinguish the flame.

Continue stirring the water gently until the temperature stops rising; then record the highest temperature, again estimating it to the nearest 0.1 degree C. Calculate the temperature change of the water in the can by subtracting the initial temperature from the final temperature.

Your burner is probably too hot to touch near its top, so handle the beaker at its base. Take the burner and your data sheet to a balance. Re-weigh the burner (again to the nearest 0.01 grams) and calculate the mass of the fuel consumed by subtracting the initial mass from the final mass.

Discard the water from your can and repeat the procedure (you will only need to measure the mass of the can with water in it, no need to re-weigh the dry can).

Calculations

To calculate the amount of heat liberated by 1 gram of your biodiesel, you will need the following: the mass of water heated, the change in temperature of your water, and the mass of biodiesel burned.

It takes 1.00 calories of energy to raise the temperature of 1 gram of water by 1 degree C, thus the total heat absorbed by the water is equal to the temperature change times the grams of water times 1.00 cal/gram degree.

Heat absorbed = temp change x grams of water x 1.00 cal/gram degree C

For each trial, calculate the total calories of heat absorbed by the water. This will be assumed to be the same as the amount of heat liberated by burning your fuel.

Then calculate the calories of heat per 1 gram of fuel burned.

Average your results.

43




Part 1. Preparation of biodiesel

Type of oil used: __________

Observations:

Part 2. Viscosity

Drip time (in seconds):

Liquid

Trial 1

Trial 2

Average

Part 3. Temperature effects

Observations for oil:

Observations for biodiesel:

Oil Biodiesel

44


Part 4. Heat content of biodiesel

Trial number

Mass of can + water (grams)

Mass of empty can (grams)

Mass of water (grams)

Final temp of water (C)

Initial temp of water (C)

Temperature change (C)

Initial mass of burner

(grams)

Final mass of burner

(grams)

Mass of fuel burned (grams)

Total heat absorbed by water (calories)

Heat per gram of fuel

(calories/gram)

Average heat per gram of fuel (calories/gram)

1 2

Questions

1. Compare the viscosities of your oil and biodiesel. Which would flow more easily in an engine?

2. When a liquid flows, its molecules get tangled up. More tangling makes the liquid flow more slowly. Look at the structures of triglyceride oil and biodiesel. Which structure looks like it should have a higher viscosity? Explain. Does this match your observations?

45


3. What did you observe when you cooled the oil and biodiesel? How would this effect the use of your biodiesel during the winter?

4. Do you expect your calculated heat content to be a little low or a little high? Explain.

46


Experiment 9 – Analysis of Vinegar



-To introduce students to the titration method of chemical analysis.

-To use this method to determine the acid content of household vinegar.

Introduction

Review sections 8.1 – 8.3, and 14.1 in your text prior to coming to lab. Acids and bases are among the most common types of chemical substances. They are literally everywhere. Solutions of acids contain excess H

+

ions and have a pH below 7, while bases have excess OH

-

and pH values above 7. Acids and bases readily react with each other in a neutralization reaction. This type of reaction is the basis for the titration method of analysis. In a titration, a known amount of acid solution is measured out; then a solution of base (with a known concentration) is slowly added to the acid solution until the reaction is complete. Based on the volume and the concentration of the base, one can calculate the concentration of the acid in solution.

In a titration, an indicator is required. This indicator is a chemical species that changes color in response to changes in pH. The color change indicates the end of the titration, i.e. the acid has completely reacted with the base. In this particular experiment the indicator that is used is phenolphthalein which turns from colorless to pink. The pink color indicates the end of the titration. Once the endpoint is reached, one can use the volume and concentration of base to calculate the concentration of acid (for accurate results the volume of base and acid must be known).

Materials

-24-well plate

-vinegar samples (vinegar contains acetic acid – H

2

C

2

H

3

O

2

)

-NaOH solution of known concentration

-indicator solution (phenolphthalein)

-4 plastic transfer pipettes

-small stirring rod or tooth pick

-piece of white paper

47


-tape or labels for the pipets


NaOH is corrosive, goggles are required for all of this experiment.

Obtain a plastic well plate and a small stirring rod or tooth pick. Place the well plate on a white piece of paper so that the color change can be easily observed.

You should have 3 labeled pipets: vinegar, NaOH, and indicator.

Fill the vinegar pipet from a bottle of vinegar and record the brand of vinegar.

Fill the NaOH pipet from the supply bottle. Record the concentration of the NaOH solution.

Partially fill the indicator pipet with phenolphthalein solution.

Fill the unlabeled plastic pipet with water and try dispensing drops into a well of the well plate.

Your goal is to be able to confidently add a known number of drops to the bottom of the well, one drop at a time. Squeeze the bulb gently, while the pipet is held vertically and directly over the center of the well. You may find it helpful to use two hands to steady the pipet. Plastic transfer pipets easily acquire air bubbles in their stems that lead to frustrating half-drops and quarter drops that introduce errors. The best solution is to reserve one well for solution waste.

During titration, half drops are not added to the titration well or counted. They are simply added to the waste well. When you can confidently add a known number of drops to a well, proceed to the practice titration.

Practice titration: add 10 drops of vinegar (any type of vinegar will do) to well A1 in the well plate. Then add 1 drop of indicator and mix (the mixture should be colorless). Start adding

NaOH, one drop at a time while counting the drops, with gentle stirring. Carefully observe what happens. As each drop is added, a slight pink color should appear which should then disappear upon stirring. After enough drops have been added the pink color will persist. Note how many drops were used. Then add a few more drops with stirring. The color should become darker with more NaOH.

Real titration – Part 1: white vinegar: your goal is to catch the point where one drop of NaOH produces the first persistent color change. This is the end point of the titration, or the point where all of the acid has been neutralized by the base.

Carefully add 20 drops of white vinegar to each of 6 wells. Make sure that the drops fall directly to the bottom of the well and are not trapped along the side. If you make a mistake, simply start again in another well. Record the position numbers for the wells you are using on the data sheet.

Add 1 drop of indicator to each well.

Start adding NaOH one drop at a time with stirring. Stop adding NaOH when one drop produces a permanent pink color. Record the number of drops you used. If you make a mistake, skip this trial and go on to the next.

48


Repeat this titration for each of the other 5 wells. Be sure to record the number of drops each time.

You should now have 6 results. They should agree with each other to within a few drops. If they do not, consult your lab instructor.

Dispose of your solutions in the waste container provided (do not pour anything down the sink!).

Use the following equation to calculate the molarity of acid in your vinegar sample:

Molarity of HC

2

H

3

O

2

= molarity of NaOH X (# of drops of NaOH/# of drops of vinegar)

Real titration – Part 2: any vinegar: Choose another vinegar sample (white wine vinegar, apple cider vinegar, etc.) and repeat the above procedure.

49




Molarity of the NaOH solution: __________ mole/liter

Part 1:

Type of vinegar sample:_________________________________

Trial #

Well #

Drops of vinegar

Drops of

NaOH

Notes

1 2 3 4 5 6

Average number of drops of NaOH:________________(per 20 drops of vinegar)

Calculated molarity of acetic acid:________________mole/liter

Percent acidity in the white vinegar:________________%

7

Part 2:

Type of vinegar sample:_________________________________

Trial #

Well #

Drops of vinegar

Drops of

NaOH

Notes

1 2 3 4 5 6

Average number of drops of NaOH:________________(per 20 drops of vinegar)

7

8

8

50


Calculated molarity of acetic acid:________________mole/liter

Percent acidity in the white vinegar:________________%

1. During the experiment, you were cautioned about making sure that all drops of solutions were at the bottom of the well, not along the sides. Why?

2. Suppose that during the titration, you added sufficient NaOH to change the color of the indicator to a darker pink. How would this change the results? Explain.

3. What is the advantage of doing several trials, especially since you are doing exactly the same thing each time?

4. Could you have done the titration “backwards”, i.e., start with 20 drops of NaOH, add phenolphthalein and then add the vinegar one drop at a time? Explain. What would you observe?

5. Can you think of ways to improve this method of analysis? What changes might be possible to give you more confidence in the accuracy of the results?

6. Why do you suppose that the acetic acid content in vinegar is reported as percent acidity rather than as molarity?

51


7. Suppose you were asked to determine the exact acetic acid concentration in an acetic acid solution that was about 50% acid. How would you change the procedure you used to determine the acetic acid content of vinegar so it could be used to solve this problem?

52


Experiment 10 – Thin-Layer Chromatography of Analgesics



-To become familiar with Thin Layer Chromatography.

-To learn about the active molecules in different analgesics (pain killers).

-To understand how TLC can be used to separate compounds in a mixture.

Introduction

Review the meaning of analgesic and some of the compounds that are considered analgesics, as well as chromatography and the different types and applications prior to coming to lab.

Chromatography is the branch of chemistry that is focused on separating complex mixtures.

There are many forms of chromatography: gas chromatography, thin layer chromatography, liquid chromatography, size exclusion chromatography, etc. They all operate by the same basic principle, an analyte (molecule of interest) is unequally distributed between a mobile phase and a stationary phase. An analyte is initially in the mobile phase, as the mobile phase moves past the stationary phase, the analyte spends some time interacting with the stationary phase. The amount of time the analyte spends in the stationary phase is dictated by its chemistry and every analyte will spend a specific amount of time in the stationary phase. As the mobile phase continues to move it separates compounds from a mixture due to the time spent by any particular analyte interacting with the stationary phase.

In this lab you will be performing Thin-Layer Chromatography (TLC) on different analgesic tablets. This means you will essentially be separating the different molecules in these tablets to identify what is in them. TLC is one of the easiest chromatographic techniques. A thin layer of solid substance (stationary phase) is coated on glass or plastic. By immersing one end of the sheet in an appropriate liquid solvent, the solvent is drawn up the sheet by capillary action, and the analytes are carried along at different rates. This is known as developing the plate.

Many of these compounds are colorless, therefore we need a way to make them visible to us. In this lab, we will use UV light to identify the location of different compounds on the TLC plate.

Most analgesics today contain aspirin (acetylsalicylic acid), ibuprofen, or acetaminophen. Some analgesics also contain caffeine to fight drowsiness and starch and sugars as binders. Buffered aspirin contains a base such as magnesium hydroxide or calcium hydroxide to neutralize the acid. You will use TLC to see what compounds are in what analgesic tablets.

53


Materials

-methanol

-analgesic tablets (aspirin, acetaminophen, ibuprofen, caffeine)

-solvent mixture

-large beaker, 400 – 600 mL

-wrap to cover beaker while plate is developing

-TLC plates (do not touch with your hands, these are very fragile, a chip will skew results)

-glass capillary tubes

-small ruler and pencil

-mortar and pestle

-UV lamp


Obtain large beaker and a piece of plastic wrap or foil to cover it. Beaker must be large enough for TLC plate.

Add enough of the solvent mixture to give a thin layer of solvent in the bottom of the container.

To provide an atmosphere saturated with solvent inside the container, place a piece of filter paper around the inside surface of the container, extending into the solvent. Then cover the container with the wrap or foil and set it aside while preparing the chromatographic sheet.

Obtain an unknown analgesic tablet. Using a mortar and pestle, crush a portion of the tablet into a fine powder. Add a few mLs of methanol to the powder and stir, allow mixture to settle for a few minutes.

Obtain a chromatography plate (handle carefully!). Using a pencil, draw a very light line across the sheet (the short side) about 1 cm from one end. Then make 5 small light marks at 1 cm intervals along the line. These are the points where the samples will be spotted. Label each point with a penciled number or abbreviation and record in your data sheet.

Samples will be spotted on the sheet using glass capillary tubes. Before spotting you should practice a few times, the idea is to get the smallest possible spot, consistently.

Carefully place small spots of the four known solutions at the appropriate marks on the chromatography sheet. Finally, spot a sample of the clear solution from your unknown tablet onto the plate, using a separate capillary tube. Allow the solvent to evaporate. A hair dryer may be used to expedite drying.

54


When the spots are dry, place the sheet in the developing container. Make sure the bottom edge

(near the spots) is in the solvent but the spots are above the solvent. Also, be sure that the filter paper does not touch the chromatographic sheet. Cover with wrap or foil.

When the front edge of the liquid has moved about 80-90% of the way to the top of the sheet, remove the sheet. Immediately, while the sheet is still wet, draw a pencil line on the sheet to show the top edge of the liquid. Then lay the sheet on a clean surface in a fume hood and allow the solvent to evaporate.

The spots are unlikely to be visible to the naked eye, but they should be visible under a UV lamp

(UV light is bad for the eyes, don’t look directly into the light). While observing under UV light, use your pencil to circle each spot.

It is likely that you can already see if each was a pure compound or a mixture, as well as how many compounds are in each mixture.

A more scientific means of interpreting results is to calculate a retention factor (R f

) value for each spot. These values are then directly comparable to each other.

R f

= distance traveled by substance/distance traveled by the solvent = distance to center of spot/distance to solvent front

Measure the center of each spot in mm from the starting line, also measure the starting line to the solvent front (i.e. the distance the solvent traveled). Calculate a R f

value for each spot, record each value in the table provided (these values should all be less than 1 since nothing traveled farther than the solvent).

55




Sample Label on TLC sheet

Distance to center of spot

Distance to solvent front

Calculated R f

Aspirin

Acetaminophen

Ibuprofen

Caffeine

Mixture

Unknown

What compounds are present in the unknown sample?

Tape the TLC sheet here.

56


1. Why do you think it is necessary to maintain a saturated vapor atmosphere in the container used for development? What do you predict would happen if this were not done?

2. Suggest possible advantages and disadvantages of using longer TLC sheets.

3. Why do you think it was important to use a very small amount of sample when spotting the plate?

4. The movement of compounds in this case is dictated by polarity. The TLC sheet is very polar while the solvent has less polarity. From your results, list the compounds in analgesics from lowest to highest polarity.

5. Do you expect that changing the solvent will change the R f

value for a given compound?

Explain your reasoning.

6. If two components have identical R f

values, does this mean they necessarily have the same structure? Why or why not?

57

CHEM 102 lab packet_W13

Table of Contents

Chemistry 102: Introduction to Organic Chemistry

Experiment #1 – Natural Indicators

Experiment #2 – Equilibrium

Experiment #3 – Nuclear Chemistry

Experiment #4 – Hot Stuff

Experiment #5 – Like Dissolves Like

Experiment #6 – Identification of an Unknown Liquid

Experiment 7 – Identification of Plastics

Experiment #8 – Biodiesel

Experiment 9 – Analysis of Vinegar

Experiment 10 – Thin-Layer Chromatography of Analgesics

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CHEM 102 lab packet_W13

Table of Contents

Chemistry 102: Introduction to Organic Chemistry

Experiment #1 – Natural Indicators

Experiment #2 – Equilibrium

Experiment #3 – Nuclear Chemistry

Experiment #4 – Hot Stuff

Experiment #5 – Like Dissolves Like

Experiment #6 – Identification of an Unknown Liquid

Experiment 7 – Identification of Plastics

Experiment #8 – Biodiesel

Experiment 9 – Analysis of Vinegar

Experiment 10 – Thin-Layer Chromatography of Analgesics

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