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Full text of "General Chemistry Lab Spring"

General Chemistry Lab Spring



By:

Mary McHale



General Chemistry Lab Spring



By:

Mary McHale



Online:

< http://cnx.Org/content/coll0506/l.56/ >



CONNEXIONS

Rice University, Houston, Texas



This selection and arrangement of content as a collection is copyrighted by Mary McHale. It is licensed under the

Creative Commons Attribution 2.0 license (http://creativecommons.Org/licenses/by/2.0/).

Collection structure revised: April 3, 2009

PDF generated: February 4, 2011

For copyright and attribution information for the modules contained in this collection, see p. 86.



Table of Contents



1 Practical Examples of the Gas Laws 1

2 Colligative Properties and Ice Cream 5

3 Pervasive Polymers 11

4 Determine the Value of an Equilibrium Constant by Complex Ion Forma-

tion 15

5 indigestion? Which is the Best Commercial Antacid? 19

6 Acid and Bases to Buffers 25

7 Forensics 29

8 The Curious Case of Catalase 37

9 Organic Reactions 43

10 Kitchen Synthesis of Nanorust 53

11 Electrochemistry and Alchemy 59

12 From Cells and Electrodes to Golden Pennies 69

13 Amphoteric Aluminum 77

14 Crystal Violet Kinetics 81

Index 85

Attributions 86



IV



Chapter 1

Practical Examples of the Gas Laws 1



Practical Examples of the Gas Laws
Objectives

• Learn and understand physical properties of gases and explain observations in terms of the kinetic
molecular theory of gases.

• Plot and calculate the root mean square speed of the Carvone molecules. (Comparison with speed in
vacuum) .

• Estimate volume and volume change of a balloon when it goes from room temperature (RT) to liquid
nitrogen temperature.

• Observe and explain behavior of gas in: a soda can, a balloon in a flask, Cartesian diver, etc., when a
change in pressure or temperature is applied.

1.1 Grading

You grade will be determined according to the following:



• Pre-lab (10%)

• Lab Report Form (80%)



• TA points (10%)

Introduction

Expanding and contracting balloons, imploding soda cans, exploding marshmallows are just some of the
demonstrations that are often used to illustrate the empirical gas laws and the kinetic molecular theory
of gases. In this experiment, you will be performing these and other 'demonstrations' and using your
understanding of the physical properties of gases to explain your observations.

There will be two demonstrations laid out at each of the seven different stations around the room and you
will go as a group, half the group working at each station (you don't need to start with #2). If your group
starts with, for example station 5, you should then follow the following order: 5, 6, 7, 8, 2, etc. Your group
should spend no more than 15 minutes at each station, in some cases 5 minutes is sufficient. Perform the
experiment by following the instructions placed at each station. Then discuss your observations with your
group. For each of the activities, it is important to ask yourself what is going on; "how can our observations be
explained using the kinetic molecular theory of gases?" Remember that for some demonstrations calculations
may also be required. Be thorough and precise in your explanations.
Important Safety Notes:

Remember to use tongs, hot grips as appropriate when dealing with hot liquids, vapors and containers.



1 This content is available online at <http://cnx.Org/content/ml9475/l.4/>.



2 CHAPTER 1. PRACTICAL EXAMPLES OF THE GAS LAWS

Liquid nitrogen is extremely cold, with a boiling point of — 196 �C and if it comes into contact with skin
can result in severe frostbite.

The vacuum dessicator should be regarded as a potential implosion hazard when evacuated. Handle it
carefully.

When doing the egg experiment do not put the hot flask immediately in the water bath (let it for at least

3 minutes sitting on the bench) as it will crack.

Observe and record what happens in your laboratory report form and explain your observations in terms
of the Kinetic Molecular Theory of Gases.

You are encouraged to discuss among yourselves possible explanations to your observations.

Experimental Procedure

Diffusion:

• The goal of this experiment is to measure the rate of diffusion of carvone, a major component of
spearmint oil. We will do these trials altogether, with volunteers, at the end of the pre-lab lecture.
You will all stand in a line, with the first person in the group holding the bottle of carvone and several
paper towels. All four people should be 1 meter apart. You will need to know the distance each person
is from the bottle of carvone. The fourth person should act as the timekeeper.

• When the timekeeper gives the signal, the first person should place a few drops of carvone on the paper
towels. Record the time that it takes for each person to smell the carvone. Seal the paper towel in a
plastic bag when you are finished.

• After the odor has dissipated, we will repeat the experiment twice with more volunteers.

• Using Excel plot the data in distance traveled versus time. Obtain a least squares fit ( R 2 , R squared
value) for this data and determine from it the rate of diffusion of carvone in meters per second. Create
a graph for each trial. Calculate the average of the rates for the three trials. Calculate the root mean
square speed of carvone molecules at 25 �C. Your TA will help you with this equation. Compare the
result with the diffusion rate you measured.

• If they are significantly different, offer an explanation.

• Would the diffusion take place faster in a vacuum?

Note: The formula of carvone is C w H 14 0, MW=150.22 g/mol.

Since, PV = nRT

and nRT = (l/3)Nm fi 2 .

Solving for /i gives:

H = ^/3RT/M.

where M = mN/n or the molar mass in kg/mol, T is in k, and R is 8.3145 kg * m 2 / s 2 * mol * K. The
/i, we are using is the root mean square speed, as it is the root of the sum of the squares of the individual
velocities.

The plots can be prepared when you have finished the lab.

1. Gas Laws in a Soda Can:

1. Pour 15 mL of water into an aluminum soda can. Set the can on a hot plate and turn on to a high
temperature setting. While the can water heats, fill a 1000-mL beaker with cold water (You may have
a metal tin set out for this purpose). Continue heating the can until the water inside boils vigorously
and until steam escapes from the mouth of the can for about 20 seconds.

2. Using the hot grips to grip the can near the bottom, quickly lift the can from the burner and invert it
in the beaker of cold water so water covers the mouth of the can.

3. Describe what happens.

4. Explain why it happens. You may repeat this experiment using a second soda can if you wish.

5. Why is it necessary to invert the can in the water? What would happen if a rigid container were used?

6. Balloon in liquid nitrogen:



Review the safety notes above regarding the handling of liquid nitrogen.

1. Inflate a balloon and tie the end (several balloons may have already been inflated and tied). Using
tongs, place the balloon in a Dewar flask containing liquid nitrogen. After the balloon stops changing size,
remove it from the Dewar and allow it to warm to room temperature.

2. Observe and record the changes (you should be able to measure the radius and estimate volume).
Estimate the size of the balloon in liters.

3. What is the pressure inside the balloon before it is placed in the liquid nitrogen?

4. What is the pressure inside the balloon after it is placed in the liquid nitrogen?

5. Use the ideal gas law to calculate the percent change in volume expected on going from room temper-
ature to liquid nitrogen temperature.

6. Is the volume of the cold balloon consistent with what you calculated, or is it larger or smaller?

7. Suggest an explanation for your observation. Explain all of your observations in detail using the
kinetic molecular theory of gases.

8. How does the liquid nitrogen cool the gas in the balloon?

1. Kissell's tygon tube in liquid nitrogen:
Review the safety notes above regarding the handling of liquid nitrogen.

1. Place a 2 foot long tygon clear tube in a Dewar with liquid nitrogen.

2. Observe what happens and explain.

3. Balloon in a flask:

1. Place about 5 mL of water in a 125-mL Erlenmeyer flask. Heat the flask on a hot plate until the water
boils down to a volume of about 1 mL.

2. Meanwhile, inflate a balloon and then let the air out (this may not be necessary if balloons on table
have been previously used).

3. Remove the flask from the heat, hold it with a towel, and immediately place the open end of the balloon
over the mouth of the flask.

4. Observe the effect as the flask cools.

5. Can you get the balloon back out again?

6. If you can, How?

7. Cartesian diver:

The Cartesian diver is named for Rene Descartes (1596-1650), noted French scientist and philosopher. At
this station, you will find a plastic soda bottle containing a medicine dropper, water, and air. Squeeze the
bottle.

What happens? Why?

1. The Egg:

1. Lightly grease the inside of the neck of a 1 L Erlenmeyer flask with stopcock grease. Clamp the flask
onto the stand. Place about 5 mL H^O in the flask and gently warm it with a Bunsen burner until
the water vaporizes. Do not boil the water to dryness.

2. Meanwhile, prepare an ice water bath in an evaporating dish. While the flask is warm, seat the egg,
narrow end down, in the mouth of the flask. Unclamp the flask, allow to cool slightly sitting on the
bench and then immerse it in the ice water. (Read the safety notes above to avoid breaking the flask)

3. Can you get the egg back out again?

4. Assuming that the flask reaches the maximum vacuum (minimum pressure) possible before the egg is
drawn into the flask, calculate the minimum pressure reached in the flask.

5. Expanding balloon:



4 CHAPTER 1. PRACTICAL EXAMPLES OF THE GAS LAWS

1. Partially inflate a balloon. Place the balloon inside the vacuum chamber and close the chamber with
the black rubber circle and the top of the chamber carefully centered on the base (A partially inflated
balloon may already be in the dessicator).

2. Close the needle valve (at the bottom of the black rubber tubing) by turning it clockwise. Turn the
stopcock to the up position to connect the chamber to the vacuum pump.

3. What happens? Explain? To open the chamber, turn the stopcock to the left position and open the
needle valve.

Moore's bonus 2 points:

lpt to name a real life example of the physical properties of gases at work

lpt for a good explanation of how and why it works according to what you have learned in the lab.



Chapter 2

Colligative Properties and Ice Cream



Colligative Properties and Ice Cream
Objectives

• To record facile and fast data collection from the computer interface, ubiquitous in industry and, in
this case, to calculate the molecular weight of the unknown solute using freezing point depression

• To learn the definition of molality and the importance of molality in colligative property calculations

• To learn to calculate the molality of a solution

• To measure the freezing point depression caused when adding antifreeze to tert-butanol

• To calculate the molecular weight of the unknown solute using freezing point depression

2.1 Grading

You will be determined according to the following:

• Pre-lab (10%)

• Lab Report Form (80%) - including temperature plots







TA evaluation of lab procedure (10%)



Introduction

Although colligative properties involve solutions, they do not depend on the interactions between the
solvent and the solute molecules but rather on the number of solute particles dissolved in solution. Colligative
properties include vapor pressure lowering, osmotic pressure, boiling point elevation, and freezing point
depression. In this experiment you will explore freezing point depression using a solution of ethylene glycol
in tert-butanol. You will then use freezing point depression to calculate the molar mass of an unknown solute
that is dissolved in tert-butanol.

Ethylene glycol, (CH20H)2 the major component of antifreeze, is a large organic molecule that dissolves
easily in water. The structure of ethylene glycol is shown in Figure 1.



lr This content is available online at <http://cnx.Org/content/ml9545/l.4/>.



CHAPTER 2. COLLIGATIVE PROPERTIES AND ICE CREAM




Antifreeze keeps the water in a car's radiator from freezing because the ethylene glycol molecules get in
the way when water tries to crystallize into ice. It is more difficult for the ice crystals to form, due to the fact
that the water must be at a lower kinetic energy. Therefore, the water freezes at a lower temperature than
if the glycol molecules were not present. The effect of the ethylene glycol molecules present in a solutioncan
be quantified by the following equation:

AT = iKfm Equation 1

where A T = Tpure - Tsolution, the difference between the freezing temperature of the pure solute and
the freezing temperature of the solution. Kf is the freezing point depression constant of the solvent, having
units of �C/m, and m is concentration of the solution using units of molality. This equation reflects the fact
that a more concentrated solution results in a greater change in freezing temperature.

Most of the previous work that we have done with solutions probably has involved units of molarity, or
moles per liter of solution. Freezing point depression calculations (as well as those for boiling point elevation)
use molality, or moles of solute per kilogram of solvent. By definition, a freezing point depression or boiling
point elevation involves a change in temperature. When the temperature of a solution changes, its volume
also changes. Since molarity depends on the volume of the solution, a change in temperature will change the
solution's molarity. Molality depends on the mass of the solvent, and this does not change with temperature.

The solvent we will use in this experiment is tert-butanol (IUPAC name: 2-Methyl-2-propanol) also called
tert-butyl alcohol. It has a characteristic camphor type smell and is used in paint removers, to boost octane
in gasoline and in perfumes. Its structure is given in Figure 2.




Figure 2.2



In this experiment we will measure the freezing temperature of pure tert-butanol, then measure the
freezing point of a solution containing 3-5 grams of ethylene glycol added to tert-butanol. The difference in
freezing temperatures for the two solutions gives the A T in Equation 1. Since the purpose of this experiment



is to find the molecular weight of the solute, Equation 1 can be rewritten to include molecular weight of the
solute:



AT = iK f m = K,



' moles solute
kg of solvent ,



\



grams solute



v



molar mass of solute



kg of solvent



Equation 2



Figure 2.3



For this experiment, use a Kf for tert-butanol of 8.37 �C/m.

The only unknown in equation 2 is the molar mass of the solute. If you algebraically rearrange Equation
2, you can then solve for molar mass. This algebraic manipulation is left as an exercise for you to complete.

SAFETY PRECAUTIONS: Ethylene glycol and tert-butanol are safe if handled properly, but are mildly
poisonous if swallowed. These chemicals can also cause allergic reactions with skin contact. Wear plastic
gloves when pouring and measuring these chemicals. If you spill any on your hands, wash immediately with
soap and water. Be sure to wear safety glasses at all times during this experiment.

Experimental Procedure

Part 1: Freezing point of tert-butanol

1. Open the MicroLab Program by clicking on the shortcut to MicroLab.exe tab on the desktop.

2. On the "Choose an Experiment Type" tab, enter a name for your experiment, and then double click
on the MicroLab Experiment icon.

3. Click "Add Sensor", choose sensor = Temperature (thermistor).

4. To choose an input, click on the red box that corresponds to the port which the thermistor is connected.

5. Choose label = Thermistor, sensor units = � C, click next.

6. Click "Perform New Calibration".

7. Click "Add Calibration Point" and place the thermistor and a thermometer in an ice water bath. Wait
until the temperature is constant, then read the temperature on the thermometer and enter that value
into the "Actual Value" box in MicroLab and hit "ok".

8. Again, click "Add Calibration Point" and place the thermistor and a thermometer in a warm water
bath. Wait until the temperature is constant, then read the temperature on the thermometer and enter
that value into the "Actual Value" box in MicroLab and hit "ok".

9. Under Curve Fit Choices, click on "First order (linear)" and then "Accept and Save this Calibration",
when prompted to enter units, enter as "deg C". Save as your name-experiment data.

10. Click "Add Sensor", choose sensor = Time

11. Choose an input, click on the red box that corresponds any of the timers.

12. Label = Time 1, click next, click finish.

13. Left click on thermistor and drag to: the Y-axis over "data source two", column B on the spreadsheet,
and the digital display window.

14. Left click on time and drag to: the X-axis over "data source one", column A on the spreadsheet, and
the digital display window.

15. When ready to obtain data, click start.

16. Take a clean, dry 10cm test tube, and fill it half-way with tert-butanol, dispensed by your TA.



8 CHAPTER 2. COLLIGATIVE PROPERTIES AND ICE CREAM

NOTE: The tert-butanol must be distributed by your TA to avoid impurities that will cause tremendous
errors in the experiment. You will need a very clean and very dry test tube for each of your experiment
runs. If any impurities (especially water) mix with the tert-butanol, your data will be severely affected.

ANOTHER NOTE: Make sure that your solution in your test tube is below the level of the water in the
water bath

1. Warm the test tube to 30-35 �C by placing it in a warm water bath. Start your data acquisition
program, and place the test tube in an ice/water bath. You must constantly stir the tert-butanol to
prevent supercooling.

2. The temperature of the tert-butanol should steadily drop, then level off, as the liquid freezes. When
the tert-butanol is completely solid and the temperature starts decreasing again, you may stop your
experiment. This will be the freezing temperature you use as Tpure when you calculate A T. Record
this value in your data sheet. If your cooling curve does not flatten out very well and it is difficult
to determine the freezing temperature, warm your sample with the hot water bath, and measure the
freezing point again.

Note: If you do not get an acceptable curve (your TA can verify if it's acceptable) on your second try,
then you should ask for a new sample of tert-butanol. The rest of your results for this lab depend on this
measurement being accurate.

Part 2: Freezing point depression

1. After finding the freezing temperature of pure tert-butanol you will make a solution of tert-butanol
and antifreeze. Set your test tube inside a 50 or 100 mL beaker and then place the beaker on the
balance to weigh your sample.

2. First, weigh only the empty test tube and record this value in your data sheet.

3. Fill the test tube half full of tert-butanol (again from your TA), weigh again, and record this value.

4. Finally, add a few drops of antifreeze (3 or 4 drops is sufficient), weigh the beaker/test tube combination
a third time, and record this value. Use the same balance for all three weighings. Use subtraction to
find the masses of the tert-butanol and the added antifreeze. It is not critical how many drops you
add, but the mass that you measure is the important value.

5. Find the freezing temperature for your solution in the same way you found the freezing temperature
for pure tert-butanol. The flat portion of your temperature curve will be smaller and more difficult
to see for your mixed solution than for the pure tert-butanol. If you are running a live graph in your
program, you should be able to tell where the freezing point of your solution occurs. As before, you
may re- warm your solution and run it again if your cooling curve does not show a clear freezing point.
If time permits, you may want to perform more than one run on each solution to confirm your freezing
temperatures.

6. Subtract the freezing temperature your solution from the freezing temperature of pure tert-butanol.
Record this A T value in your data sheet.

7. Use the information recorded in your data sheet to calculate the molar mass of ethylene glycol. Calcu-
late the percent error for your experimentally determined molar mass. See your TA if you are unsure
how to make this calculation.

8. Repeat Steps 1 through 5 for the unknown solutes. After calculating the molar mass of the unknowns,
identify them using the following information:



Compound


Molar Mass (g/mol)


Acetone


58.08


Ethyl Acetate


88.10


Water


18.02



Table 2.1



Part 3: Chemistry of Life: Ice Cream

As we found above, adding a solute to a solvent lowers the freezing point of that solvent. This occurs
because as a substance freezes, a crystal is formed, but if a solute is added to the solvent more kinetic energy
must be removed from the solvent in order to freeze, since it's harder for the solvent molecules to form the
regular pattern of the crystal. Therefore, the more solute molecules you add, the lower the freezing point
becomes. We can use this to our advantage to lower the freezing point of water by enough to freeze ice
cream, since ice cream is mostly water.

This is a recipe that you could use at home:

Put 59.15 ml (J cup) of sugar, 118.29 ml (^ cup) of milk, 118.29 ml (^ cup) of whipping cream, and 1.23ml
(J teaspoon) vanilla (4-hydroxy-3-methoxybenzaldehyde) into a one-quart Ziploc ™ bag. Seal the bag and
mix well by carefully shaking.

Put this one-quart Ziploc ™ bag into a one-gallon Ziploc ™ bag that has 2 cups of ice..

1. However, we are going to cheat by using 6 packets of Junket ice cream mix + 7 ^ cups of whole milk
+ 4 2 cups of whipping cream in a one galloon jug.

2. Measure and record the temperature of the ice with your thermometer in the one-gallon Ziploc ™ bag .

3. Weigh and pour 177.44 ml (f cup) of sodium chloride into the gallon bag.

4. Place the smaller bag inside the larger bag and seal the large bag securely.

5. Holding the large bag by the zipper seal, carefully shake the bag back and forth.

NOTE: Do not touch the part of the bag with the ice as it could cause tissue damage.

1. Continue until your ice cream is solid, approximately 10 - 15 min.

2. Measure and record the temperature of the salt/ice mixture.

3. Remove the frozen ice cream and place into a Styrofoam cup and enjoy!



10 CHAPTER 2. COLLIGATIVE PROPERTIES AND ICE CREAM



Chapter 3

Pervasive Polymers 1



Pervasive Polymers

3.1 Objectives

In this laboratory you will become familiar with the classifications of polymers by synthesizing and examining
several of the following:

• a polyamide (nylon)

• a cross-linked condensation copolymer (Glyptal ™ resin)

• a cross-linked polyvinyl alcohol

• a loosely cross-linked silicon-based condensation polymer (a polymethylsiloxane)

Additional information about polymers can be found in Chapter 12 of your textbook.

3.2 Grading

Your grade will consist of the following:

• Pre-lab (10%)

• Correctness and thoroughness of your observations and the answers to the questions on the report form
(80%)

• TA evaluation of lab procedure (10%)

Before Coming to Lab . . .

• Complete the pre-lab exercise

• Read the introduction and any related materials provided to you

NOTE: If you wear contact lenses, for this week's lab, you may prefer to wear your prescription glasses.

Introduction

Approximately 50% of the industrial chemists in the United States work in some area of polymer chem-
istry, a fact that illustrates just how important polymers are to our economy and standard of living. These
polymers are essential to the production of goods ranging from toys to roofing materials. So just what exactly
are polymers? Polymers are substances composed of extremely large molecules, termed macromolecules, with
molecular masses ranging from 10 4 to 10 8 amu. Macromolecules consist of many smaller molecular units,



1 This content is available online at <http://cnx.Org/content/ml9573/l.4/>.

11



12 CHAPTER 3. PERVASIVE POLYMERS

monomers, joined together through covalent bonds. The molar mass of the polymer is quoted as an average
molar mass.

Both natural and synthetic polymers are ubiquitous in our lives: elastomers (polymers with elastic,
rubber-like properties), plastics (the first plastic was used in 1843 to make buttons), textile fibers, resins,
and adhesives. The more common polymers include acrylics, alkyds, cellulosics, epoxy resins, phenolics,
polycarbonates, polyamides, polyesters, polyfluorocarbons, polyolefins, polystyrenes, silicones, and vinyl
plastics, to name but a few.

Naturally occurring macromolecules are derived from living things: wood, wool, paper, cotton, starch,
silk, rubber and have provided us for centuries with materials for clothing, food, and housing. Starch,
glycogen, and cellulose are all polymeric versions of the monomer glucose. Again, we see that minor struc-
tural variations create chemicals with very different properties. Proteins are macromolecules composed of
monomeric units of alpha amino acids; nucleic acids are composed of subunits (nucleotides) containing a
nitrogeneous base, sugar and phosphate groups. Natural rubber is a latex exudate of certain trees and com-
posed of monomers called isoprene units. The usefulness of latex was first discovered by Lord Mackintosh
in Malaysia in the 19th century and provided the foundation of his waterproof rainwear empire.

When scientists linked the special properties of these substances (physical properties such as tensile
strength and flexibility) to the sizes of their molecules, the next logical step involved chemical modifications
of naturally occurring polymers.

Synthetic celluloid derives from natural cellulose and stems from an accident that Christian Schoenbein,
a chemistry professor, had in 1846. The age of plastic had begun, although the interest in cellulose nitrate
was initially more for its explosive properties. When cellulose (from wood chips or fiber) is treated with
a mixture of nitric acid, camphor, and alcohol, the resultant product is called Celluloid ™ and bears
very little resemblance to the starting material. Celluloid ™ possesses the ability to be molded into hard,
smooth billiard balls (replacing the original, very expensive ivory balls) and into thin sheets for making
movie pictures. Celluloid ™ is highly flammable and today has been replaced by greatly improved synthetic
polymers such as Bakelite discovered in 1907 by the Belgian- American Chemist, Leo H. Baekeland.

When cellulose is treated with sodium hydroxide and carbon disulfide (CS2), cellulose xanthate is formed.
A viscous (thick) solution of cellulose xanthate, forced through fine holes into dilute sulfuric acid, regenerates
the cellulose as fine, continuous, cylindrical threads called rayon. If the solution is forced instead through a
narrow slit, a thin transparent film or sheet is obtained called cellophane.

Experimental Procedure

A. Synthesis of Nylon (to be performed in the hood by Dr McHale): Two liquids are mixed in a small
beaker and nylon is formed at their interface. The nylon is pulled from the beaker; a continuous thread
10-15 ft long can be formed. CAUTION: Wear gloves while performing this experiment; do not touch the
nylon with your bare hands until it has been rinsed thoroughly with water.

1. Solutions A and B have been prepared for you:

• Solution A: A 0.5 M basic solution of hexamethylenediamine (or 2,6-diaminohexane, H —
NH(CH2) 6 NH — H) was prepared as follows: Weigh 5.81 g in a large beaker and dilute to 100 mL
with 0.5 M NaOH solution (20 g of NaOH per liter). Warm the solid until it melts. Wear gloves;
hexamethylenediamine is absorbed through the skin.

• Solution B: Adipoyl chloride (CI - (C = O) - (CH 2 )7 (C = O) - Cl) , 0.25 M: Weigh 4.58 g and dilute
to 100 mL with cyclohexane.

• Place 5 mL of solution A in a small beaker. Place 5 mL of solution B in a second beaker

• Slowly add solution A to solution B by gently pouring it down the side of the beaker. Do not stir or
mix. Solution A should form a separate layer on top of solution B.

• A film will form at the interface of the two solutions.

• Carefully hook the film with a bent paper clip and pull the film from the beaker.

• Continue pulling until the solutions are exhausted.

• If you want to keep the nylon, rinse it several times with water until it is free of all traces of amine.



13

B. Glyptal ™ Resin (linear condensation copolymer): A small aluminum dish is used to mold the polymer.
A coin, favorite small stone, or small flower may be placed on the bottom of the mold if you wish to make a
souvenir of this experiment. This will harden several hours after your lab has finished.

CAUTION: Use care when heating any of the solutions - if the solutions come in contact with your skin,
severe burns could result.

1. In a large test tube, mix 4.0 mL glycerol, 0.5 g sodium acetate, and 10.0 g phthalic anhydride.

2. Carefully heat the mixture with a low flame, starting at the top of the contents and moving down
toward the bottom as the mixture melts. Remember to point the test tube towards the back of the
fume hood for safety reasons. CAUTION Excessive heating during step 2 can cause the hot contents
to spurt out. If the hot liquid contacts the skin, severe burns result.

3. Continue heating until the melt appears to boil, and then continue heating for 5 minutes. Sufficient
heating is required to produce a nonsticky product but excessive heating turns the product into a
brittle amber material.

4. If desired, place a clean and dry coin into the aluminum dish used for a mold.

5. Carefully pour the hot liquid into the mold.

6. Allow the material to cool until the end of the laboratory period.

7. After the material has cooled, peel the mold from the cooled polymer and describe its appearance.

C. Cross-linked polyvinyl alcohol aka Slime!

CAUTION: The poly(vinyl alcohol) is a fine dust which you should avoid inhaling. For this reason, you
will only use it in solution.

1. Measure 50 mL of poly (vinyl alcohol) solution into a paper cup or small beaker and observe its proper-
ties. Vinyl alcohol does not exist. Poly (vinyl alcohol) is prepared by first forming poly (vinyl acetate)
from vinyl acetate following by hydrolysis to the alcohol.

2. Measure 7-8 mL of sodium tetraborate solution into another cup or beaker and observe its properties.
Add a few drops of food coloring at this step, if you wish.

3. Pour sodium tetraborate into the poly(vinyl alcohol) solution while stirring vigorously with a wooden
stick. The borate forms a complex structure called tetraborate, B4O5 (OH) ~~ , that links the poly(vinyl
alcohol) polymer strands together by hydrogen bonds.

4. Wearing safety gloves, examine the properties of the cross-linked polymer. See how far the polymer
will flow from your hand. Is the flowing endothermic or exothermic?

D. Silicone Plastic aka Silly Putty

Silicone plastic, commonly called Silly Putty, can be successfully approximated with Elmer's Glue instead
of silicone oil. You may each make your own silly putty if you wish and if you want to keep it, bring a Ziploc
bag - great finger exerciser, stress reliever, bouncy ball etc.

1. Mix the food coloring with 30 mL of 50% Elmer's Glue solution.

2. Add the 5 mL of 4% sodium tetraborate (Borax) solution and stir for 2 minutes in the cups provided.

3. Wearing your safety gloves, roll around the lump in your hands for two minutes, after which time it
will cease to be sticky.

4. Examine the properties of this inorganic polymer.

5. If you wish to keep your silly putty, bring a Ziploc bag to lab.



14 CHAPTER 3. PERVASIVE POLYMERS



Chapter 4

Determine the Value of an Equilibrium
Constant by Complex Ion Formation 1

Determine the Value of an Equilibrium Constant by Complex Ion Formation

4.1 Objectives

In this laboratory you will:

• Use MicroLab to take colorimetric measurements

• Use Beer's Law to measure the equilibrium concentration of a complex ion

• Review Le Chatelier's Principle

• Calculate the equilibrium constant for the formation of a complex ion

4.2 Grading

Your grade will consist of the following:

• Pre-lab (10%)

• Correctness and thoroughness of your observations and the answers to the questions on the report form
(80%)

• TA evaluation of lab procedure (10%)

Before Coming to Lab . . .

• Complete the pre-lab exercise

• Read the introduction and any related materials provided to you

Introduction

When two reactants are mixed, the reaction typically does not go to completion. Rather, they will
react to form products until a state is reached whereby the concentrations of the reactants and products
remain constant at which point the rate of formation of the products is equal to the rate of formation of the
reactants. The reactants and products are then in chemical equilibrium and will remain so until affected by
some external force. The equilibrium constant K c for the reaction relates the concentration of the reactants
and products.



lr This content is available online at <http://cnx.Org/content/ml9605/l.3/>.



15



16 CHAPTER 4. DETERMINE THE VALUE OF AN EQUILIBRIUM

CONSTANT BY COMPLEX ION FORMATION

In our experiment we will study the equilibrium properties of the reaction between iron (III) ion and
thiocyanate ion:

Fe 3+ (aq) + SCN" (aq) -> FeSCN 2+ (aq) Equation 1

When solutions containing Fe 3+ ion and thiocyanate ion are mixed, the deep red thiocyanatoiron (III)
ion ( FeSCN 2+ ) is formed. As a result of the reaction, the starting concentrations of Fe 3+ and SCN~ will
decrease: so for every mole of FeSCN 2+ that is formed, one mole of Fe 3+ and one mole of SCN~ will react.
The equilibrium constant expression K c , according to the Law of Chemical Equilibrium, for this reaction is
formulated as follows:

[FeSCN 2+ ] / [Fe 3+ ] [SCN - ] = K c Equation 2

Remember, square brackets ([]) are used to indicate concentration in mol/liter, i.e., molarity (M).

The value of K c is constant at a given temperature. This means that mixtures containing Fe 3+ and
SCN~ will react until the above equation is satisfied, so that the same value of the K c will be obtained no
matter what initial amounts of Fe + and SCN~ were used. Our purpose in this experiment will be to find
K c for this reaction for several mixtures made up in different ways, and to show that K c indeed has the
same value in each of the mixtures.

The reaction is a particularly good one to study because K c is of a convenient magnitude and the red
color of the FeSCN 2+ ion makes for an easy analysis of the equilibrium mixture using a spectrophotometer.
The amount of light absorbed by the red complex is measured at 447 nm, the wavelength at which the
complex most strongly absorbs. The absorbance, A, of the complex is proportional to its concentration, M,
and can be measured directly on the spectrophotometer:

A = kM Equation 3

We know it as the Beer-Lambert law which relates the amount of light being absorbed to the concentration
of the substance absorbing the light and the pathlength through which the light passes:

A = ebc Equation 4

In this equation, the measured absorbance (A) is related to the molar absorptivity constant ( e), the
path length (b), and the molar concentration (c) of the absorbing species. The concentration is directly
proportional to absorbance.

Thiocyanate ( SCN~) is an interesting ion and is widely used in a variety of industrial processes such as
the manufacturing of thiourea, photofinishing, metal separation, and electroplating. It is also found in gold
mining wastewater as a result of treating the cyanide rich ore with sulfur dioxide in order to produce the
less toxic thiocyanate( SCN~) ion. Iron, as you will see later on in the semester, has the unique ability to
inexpensively clean up and produce drinking quality water in Third World countries.

4.3 Procedure

4.4 Equipment

• 5 test tubes

• Burette (0.1 ml graduations) filled with 0.0200M Fe(N0 3 ) 3 in 0.5 M HN0 3

• Neighboring partners' burette filled with 6.00 x 10~ 4 M KSCN

• stirring rod

• small labels or china clay pencil

• 5 cuvettes for the spectrophotometers

• 250 ml bottle acetone for rinsing

4.5 Hazard: As always wear Safety glasses while performing this
experiment

Contamination Notes: If your flask is wet before you prepare your standard/sample solutions
ensure that the flask is wet with dilutant (in this case it is 0.5 M HNO3).



17

Calibration Of MicroLab/Spectrophotometer

N.B. Do not use test tubes from your drawer!
Find and open the MicroLab program.

1. Find and open the MicroLab program. This brings up a box that will enable you to select from a list
of experiments. Select colorimeter. Check that the accompanying box has power and is turned on, and
that it is connected to the laptop via the USB plug.

2. In the tab labeled "New" you will find the icon for the "Spectrophotometer", please double click this.
Make sure that you click on the absorbance tab.

3. This brings up the program, at which point you should take a reading of a blank, this is done by filling
a vial with 15 mL of 0.5 M HN03(not deionized water in this case) and placing in the appropriate
slot. Covering with the film case to exclude ambient light from entering the system. When the blank
sample is in place, click the button "Read Blank". This will generate a series of data points.

4. A solution of 0.0200M Fe (N0 3 ) 3 in 0.5 M HN0 3 has been prepared for you.

5. Dilute 1.5, 3.0, 4.5 and 6.0 mL portions of 6.00 x 10" 4 M KSCN to 20 mL with the 0.0200M Fe (N0 3 ) 3
in 0.5 M HNO3. For this you can use some of the smaller beakers in your lab drawer. These need
to be rinsed and dried thoroughly before each use. Use two clean, dry burettes to dispense the two
solutions.

6. This will give you 4 solutions which can be assumed to be 6.0 x 10~ 5 M, 1.2 x 10~ 4 M,1.8 x 10" 4 M
and 2.4 x 10" 4 M in FeSCN 2+

7. Take the most concentrated solution and collect a visible spectrum from 400 to 800 nm to determine

8. Measure the absorbance of all the solutions at A max , using 0.5 M HNO3 as the blank/reference solution.

9. Measure the absorbance of these solutions again at 430nm, under the same conditions.

10. Plot absorbance vs. [ FeSCN 2+ ] for both wavelengths using Excel and add a linear trendline passing
through the origin (under Options, set intercept equal to 0). Using Beer's law and the equation of the
trendline, find the molar absortivity.

4.6 Experimental Determination of Kc

The mixtures will be prepared by mixing solutions containing known concentrations of iron (III) nitrate,
Fe(N0 3 ) 3 , and potassium thiocyanate, KSCN. The color of the FeSCN + ion formed will allow us to de-
termine its equilibrium concentration. Knowing the initial composition of a mixture and the equilibrium
concentration of FeSCN 2+ , we can calculate the equlibrium concentrations of the rest of the pertinent species
and then determine K c .

1. Label five regular test tubes 1 to 5, with labels or by noting their positions in the test tube rack.

2. Dilute 0.02 M Fe (N0 3 ) 3 in 0.5 M HN0 3 by a factor of 10 using the volumetric flask that you have in
your drawer. Remember to do the dilution with 0.5 M HNO3 not water.

3. Pour about 25 mL 0.002 M Fe (N0 3 ) 3 in 0.5 M HN0 3 into a clean, dry burette.

4. Dispense 5.00 mL of that solution into each test tube.

5. Then pour about 20 mL 6.00 x 10 -4 M KSCN into another clean, dry burette.

6. Dispense 1,2,3,4, and 5 mL from the KSCN burette into each of the corresponding test tubes labeled
1 to 5.

7. Using a small graduated cylinder, dispense the proper number of milliliters of 0.5 M HN0 3 into each
test tube to bring the total volume in each tube to 10.00 mL.

8. The volumes of reagents to be added to each tube are summarized in the table.



18



CHAPTER 4. DETERMINE THE VALUE OF AN EQUILIBRIUM
CONSTANT BY COMPLEX ION FORMATION



TestTube #


1


2


3


4


5


Reagents (mL)


1


2


3


4


5


Fe(N0 3 ) 3


5.00


5.00


5.00


5.00


5.00


KSCN


1.00


2.00


3.00


4.00


5.00


0.5 M HN0 3


4.00


3.00


2.00


1.00


0.00



Table 4.1



1. Mix each solution thoroughly with a glass stirring rod. Be sure to dry the stirring rod after mixing
each solution to prevent cross-contamination.

2. Place the mixture in tube 1 in a spectrophotometer cell and measure the absorbance of the solution at

3. Determine the concentration of FeSCN 2+ from your calibration curve. Record the value on your Report
form. Repeat the measurement using the mixtures in each of the other test tubes.



Chapter 5

indigestion? Which is the Best
Commercial Antacid? 1

5.1 Indigestion? Which is the best Commercial Antacid?

5.1.1 Objectives

• Measure the quantity of commercial antacid required to neutralize a simulated stomach acid (0.15
M hydrochloric acid) and then compare the effectiveness of several brands of antacids in neutralizing
acids.

• Learn and practice the back-titration method.

5.1.2 Grading

You grade will be determined according to the following:



Pre-lab (10%)

Lab Report Form (80%)

TA points (10%)



5.1.3 Introduction

The parietal cells in the stomach secrete hydrochloric acid at a quite high concentration of about 0.155 M. The
flow of HC1 increases when food enters the stomach. When you eat or drink too much, your digestive system
may generate too much acid. You may develop a condition called "heartburn" or indigestion. Antacids are
swallowed to neutralize this excess acid and "relieve" but not eliminate the condition. The reaction that
takes place is an acid/base reaction. A small amount of NaOH might be as effective, although rough on the
rest of the digestive system, so antacids have been formulated to reduce acidity while avoiding physiological
side-effects. Many antacids use CaCC>3 for this purpose.

In addition to the active ingredient (base), tablets may also contain flavors, sweeteners, binders, fillers,
antifoam agents, pain relievers (aspirin), etc. In this experiment, the tablets will be analyzed only for their
ability to neutralize acids. The base in antacids varies with the brand. Below is an example of active agents
in several brands.



1 This content is available online at <http://cnx.Org/content/ml9622/l.5/>.



19



20



CHAPTER 5. INDIGESTION? WHICH IS THE BEST COMMERCIAL

ANTACID?



Brand Active Agent


Base


Pepto-Bismol


BiO (HOC 6j ff 4 COO)


Milk of Magnesia


Mg(OH) 2


Rolaids


NaAl (OH) 2 C0 3 (newer tablets: CaC0 3 )


Turns


CaC0 3


Alka-Seltzcr II


NaHC0 3 and KHCO3


Maalox


Mg(OH) 2 and Al(OH) 3


Gaviscon


Al(OH) 3



Table 5.1

Acids are neutralized by these bases as illustrated below.

BiO (HOC 6 ff 4 COO) + 3H + (aq) -� Bi 3+ (aq) + H 2 (I) + HOC 6 # 4 COOH (5)

Mg (OH) 2 (*) + 2H+ (aq) -> Mg 2+ (aq) + 2H 2 (I)

Al (OH) 3 (*) + H+ (aq) -� Al (OH) + (aq) + H 2 (I)



Al (OH)+ (aq) + H + (aq) -> Al (OH) + (aq) + H 2 (I)



CaC0 3 (s) + H+ (aq) -� Ca 2+ (aq) + HCO3 (aq)



HCO^ (aq) + H+ (aq) -♦ C0 2 (g) + H 2 (I)



(5.1)
(5.2)
(5.3)
(5.4)
(5.5)
(5.6)



In this simple experiment you will find the neutralizing capacity of various commercial antacid tablets.
To test their capacity to neutralize acid, we will first dissolve an appropriate and measured amount of the
sample in a simulated stomach environment. This is a solution containing a known quantity of HC1 that
will react with all of the antacid and still leave some extra HC1. Then we will determine how much of the
original HC1 remains by titrating it to neutrality with a standardized solution of NaOH. Simple subtraction
will reveal how much of the acid was neutralized by the antacid tablet. This general method of analysis is
called back-titration.

Note: The standardized solutions of NaOH and HC1 have been prepared for you. However
you need to understand how and why it is done this way. (See Supporting Information on
this) Also see notes on Titration Tips.

If Na is the total number of moles of HC1, and Nb is the number of moles of NaOH needed to back-titrate
the excess HC1, then, N samp i e is the number of moles neutralized by the sample:



N� :



N A -N E



(5.7)



v sample

Remember that the number of moles in a volume of solution equals the concentration multiplied by the
volume, (n = M x V)

Dissolved carbon dioxide is converted into the weak acid, carbonic acid (.H2CO3) which reacts with
NaOH. Although the deionized water is free of most impurities, it does contain dissolved carbon dioxide that
must be removed by boiling to give accurate titration results. All of the antacids that you titrate contain
carbonates. When you acidify the antacid sample with standard HC1, the carbonates are converted into
carbonic acid that must be boiled off in the form of CO2.



21

In this experiment you will titrate one commercial antacid twice. Each lab section will compile their
results and decide which antacid is the "best buy" in terms of neutralizing ability per dollar. The general
approach to this quantitative determination is volumetric.

Experimental Procedure

5.1.3.1 Titration of Antacids

Your TA will randomly assign you one of the commercial brands of antacid tablets which you will analyze
twice.

1. Weigh one whole tablet using the analytical balance. Record the mass of the intact tablet.

2. Break or cut the tablet to obtain pieces roughly the size indicated in the table below (e.g.. ^, 1/3,
etc.) and weigh two of these pieces separately on sheets of weighing paper. (Don't forget to subtract or tare
the weight of the paper.) After weighing, fold over the paper and gently crush the sample fragments using
the round portion of your padlock, so it dissolves faster.

3. Use the burettes around the periphery of the lab to dispense 25.00 mL of standardized HC1 into a clean
125 mL Erlenmeyer flask and add about 20 mL of deionized water. This solution will be used to dissolve
the fraction you weighed of the antacid tablet.

4. Transfer one of the weighted crushed samples (without any spilling) into the flask containing the HC1
solution. Repeat for the other weighed fraction in another flask and label each flask (1 and 2) to keep track
of which sample and which portion is being titrated.

5. Warm gently to dissolve the sample and then boil solution for about a minute. Some components of
the samples may remain undissolved, but these will not cause problems.

6. Let the flask cool down sitting on the bench for couple minutes, and then cool the outside of the flask
with tap water.

7. When room temperature is reached add a few drops of methyl purple indicator solution. The flask
should now be purple in color. (If it is green instead of purple, you have used too large a fragment and
have neutralized all of the HC1. (In this case add 5mL of HC1 and observe the color change. If needed add
measured amounts of HC1 little by little until solution is purple. Remember to record the amount of HC1
added and add to the first 25mL of HC1.)

8. It is important that you read Titration Tips in Supporting Information below before you titrate, or
you may have to repeat titration several times.)

9. Titrate with your standardized NaOH solution until you reach the endpoint, a change in color from
purple to green. Some samples may not give color changes as sharp as for the HC1 standardizations; for these
use your best judgment to estimate the endpoint. Endpoints will generally be sharper for quick titrations
than for slow ones.

10. Repeat the above procedure with a new sample of the same antacid. Enter your data in a wiki in
the designated laptop, check the web page later on the week and make a final comment on the overall cost
to neutralize one mole of HC1 for various brands. Use the data from your lab section (i.e. all of Tuesday
afternoon, Monday night, etc.). THIS IS PART OF YOUR REPORT AND IF DATA IS NOT FOUND ON
THE WEB YOU SHOULD CONTACT YOUR TA. (No excuses)



22



CHAPTER 5. INDIGESTION? WHICH IS THE BEST COMMERCIAL

ANTACID?



5.1.3.2 Retail Cost Maximum



Label


Antacid Brand


Price (cents per tablet)


Fraction


B


Walgreens Antacid Tablets


2.66


1/6


C


Titralac Plus


5.99


1/4


E


Rolaids


2.69


1/4


J


Turns Regular


2.86


1/6


K


Turns Plus


5.32


1/6


M


Eckerd Antacid


2.19


1/6



Table 5.2
*Active ingredient in all of these antacids is CaCC>3



5.1.3.3 Supporting Information on Standardization:

Molarity is the most commonly used concentration term when one is interested in the amount of materials
involved in a chemical reaction in solution. Molarity (M) is defined as the number of moles of solute per
liter of solution.



M



moles(solute)



(1)



Liters (solution)

The number of moles is calculated by dividing the mass of the sample in grams by the gram formula
weight (GFW or molar mass). One GFW is the same as one mole.

Number moles = Fap gffi) ute) (2)

For example in a 0.150 M HNO3 solution, there are 0.150 moles of HNO3 in one liter of this solution.
The following factors may then be used in chemical calculations:

0.150molHNO 3 lL(solution)

lL(solution) OI 0.150molHNO 3

In a chemical reaction that takes place in solution, the volume and the molarity of one reactant and the
molarity of the second reactant can be used, together with the stoichiometry of the equation, to find the
volume of the second reactant needed to react completely with the first reactant.

Titration is a process in which a solution of one reagent, usually the base, is added to an accurately
measured volume of another solution, usually the acid, until the reaction is complete. The concentration of
one of the reagents is known. From the known concentration and the measured volumes, the concentration
of the second solution can be calculated.

In acid-base reactions the end of the reaction or equivalence point is detected by adding a compound
that undergoes a color change as it changes from its acid form to its basic form. This compound is called
an indicator. An indicator is an organic dye that changes color at a characteristic H + ion concentration. A
dye can be an indicator if it has an intense color that changes when it gains or loses H + ions.

HIn + OH" -� In" + H 2 (3)

The change of an indicator dye from Color A to Color B depends upon the concentration of H + (or
OH - ) ions. Care must be taken in selecting an indicator to be sure that the color change (endpoint) occurs
at the H + ion concentration that corresponds to the equivalence point. Phenolphthalein, methyl red, methyl
orange, and litmus are examples of indicators.

A primary standard is usually a solid reactant that:

(1) is available in a high purity form

(2) does not change chemically when stored or exposed to air

(3) has a high formula weight to minimize errors in weighing

(4) is soluble in the solvent being used.



23

Sodium carbonate, Na 2 C03, is commonly used as a primary standard base for standardizing acids, while
potassium acid phthalate (KHP), KHCSH4O4, and oxalic acid dihydrate, H2C2O4 ■ 2H2O, are primary
standard acids used for standardizing bases.

Na 2 C0 3 (s) + 2HC1 -� H 2 + C0 2 + 2NaCl(4)

KRCsHiOi (s) + NaOH -> KNaC 8j ff 4 4 + H 2 0{5)

(COOH) 2 • 2H 2 (s) + 2NaOH -� Na 2 (COO) 2 + 4i7 2 (6)

The dry solid is carefully weighed on an analytical balance and then diluted in a volumetric flask to give
a known molarity. The molarity of an acid or base is determined by titrating a measured volume of the acid
or base with the primary standard. Then these acid or base solutions can in turn be used to determine the
molarity of another acid or base (this is the case with the standardized solutions of NaOH and HC1 used in
this experiment).

In the example below a known amount of KHP will be titrated with a solution of NaOH to determine
the NaOH solution concentration.

EXAMPLE:If 0.8168 g of KHP requires 39.35 mL of NaOH to reach the endpoint in a titration, what is
the molarity of the NaOH? (1 mol KHP = 204.2 g)

/ ImolKHP \ /lmolNaOH\ / 1 \ ZlOOOmlA

0.8168gKHP x = 0.1016mol/L = 0.1016M

B \204.2gKHPy V ImolKHP J V39.35mLNaOH7 \ 1L J '

(5.8)



5.1.3.4 Notes on titration method:

1. Clean the burette until it drains smoothly.

2. Make sure the stopcock doesn't leak. The level should hold for ~3 minutes.

3. Remove air bubbles from burette tip before beginning. Rapid spurts usually work.

4. Rinse the burette two or three times with 5 mL portions of titrant (the solution you will use to titrate).
Hold it horizontal and rotate to rinse.

5. Don't forget to record the initial reading. It does not need to be 0.00 mL.

6. Use a reading card to find the bottom of the meniscus.

7. Always read and record burette volumes to 0.01 mL.

8. Remember that the burette scale reads down, not up.

9. Put white paper under the flask to see color change easily.

10. Swirl the flask with one hand while turning the stopcock with the other or use stir bar and stir plate.

11. Add titrant slowly near the endpoint. Color that dissipates can be seen when getting close to
endpoint.

12. Drops can be "split" by quickly turning the stopcock through the open position.

13. Near the endpoint use deionized water from wash bottle to rinse the flask walls from any splashed
drops (adding water does not affect number of moles you have) .

14. Don't drain burette below 25 mL (You won't be able to determine final volume accurately). If
necessary, read the burette level, refill, read the level, and continue.

5.1.3.5 Calculation Clarification

Only part of the HC1 is neutralized by the fragment of the antacid tablet. Since the solution is still acidic,
NaOH is used to finally cause the change in color. The number of moles of NaOH is equal to the excess of
HC1 that was not neutralized by the antacid. We can use the formula in the procedure to fill in the part of
the report form where it asks for the number of moles of HC1 neutralized by the antacid:

"HClneutralizedbyAntacid = ^HCladded ~ "HClexcess (5-9)

So then, this would be the 2.5 x 10 _3 mol HC1 - the excess which is calculated by converting the volume of
NaOH used into moles, this number of moles then equals the number of moles of HC1.



24 CHAPTER 5. INDIGESTION? WHICH IS THE BEST COMMERCIAL

ANTACID?

25ml of HCLVolume titrated by NaOHNaOHVolume titrated byantacid



Figure 5.1



Chapter 6

Acid and Bases to Buffers 1

6.1 Acids and Bases to Buffers
6.1.1 Objective



To reinforce the importance of titration as an analytical tool.

To graphically verify the number of donated protons per molecule of phosphoric acid.

To prepare a phosphate buffer and realize the importance of buffers in our everyday life.



6.1.2 Grading

• Pre-Lab (10%)

• Lab Report Form (80%)

• TA Points (10%)



6.1.3 Background Information

Phosphoric acid (H3PO4) is a chemical that is commonly found in everyday products such as soft drinks
and cleaning agents. It is called a polyprotic acid because it can donate more than one proton (H + ion) per
phosphoric acid molecule. The released protons combine with water to form hydronium ions (H30 + ).

Phosphoric acid releases its protons in a step-wise manner:

H3PO4 + H 2 <-> H3O+ + H 2 PO- K al = 7.5 '10 3 (1)

H 2 P0 4 - + H 2 <-> H 3 0+ + HP0 2 4 - K a2 = 6.2 '10 8 (2)

HPO2-+ H 2 <-> H 3 0+ + POa- K a3 = 4.2 '10 13 (3)

4 4

For example, reaction (2) will not occur until reaction (1) is complete.

The K a values listed after each reaction are called acid ionization constants. They indicate the relative
ease with which each reaction occurs. A small K a value shows that a reaction does not occur easily. The K a
value for phosphoric acid's second donated proton is much smaller than for the first donated proton, while
the third K a is five orders of magnitude smaller than the second.

To determine the amount of acid in an unknown sample, you will need to add a known amount of base
until the acid and base are neutralized. This technique is known as titration, and it is widely used in
chemistry and other natural sciences.

During a titration, the pH of the solution is constantly monitored while the known acid or base (called the
titrant) is slowly added to the unknown solution. The pH of the unknown solution will stay fairly constant



1 This content is available online at <http://cnx.Org/content/ml5809/l.13/>.



25



26



CHAPTER 6. ACID AND BASES TO BUFFERS



until the moles of titrant added equals the moles of unknown acid or base. When the moles of acid and
base are the same, further additions of titrant will cause a dramatic change in pH until the pH eventually
stabilizes. A graph of pH versus added titrant is called a titration curve, and the point at which the pH
changes drastically is called the equivalence point.

The titration curve for a polyprotic acid will have more than one equivalence point. As the added base
completely removes each proton from the acid, the pH will jump significantly. Figure 1 shows the titration
curve for ascorbic acid, a polyprotic acid also known as Vitamin C:

Figure 1. Titration curve for ascorbic acid.




Figure 6.1



2 nd equiv. point

1 st equiv. point

By graphing the pH versus volume of base added during an acid-base titration, you can easily see the
successive ionization steps taking place. To find the concentration of a polyprotic acid, the volume of base
required to reach the first equivalence point is needed. The half-equivalence points on this graph can also
be used to obtain the K a value of each successive ionization.

In the third part of the lab, you will be making a buffer solution. Buffers are important in everyday life
because they regulate the pH in our blood, keeping the pH between 7.35 and 7.45; if pH values for our blood
go outside this range, death can result. A buffer is composed of a weak acid and its conjugate base (or a
weak base and its conjugate acid). When a strong acid or base is added to a buffer, one of the species will
react to maintain the pH within a small range.

To determine the amount of conjugate acid and base needed to make a buffer of a certain pH, the
Henderson-Hasselbach must be employed.

pH = pK + Iog(Q)(4)

With a given pH and known pK a , the solution of the Henderson-Hasselbach equation gives the logarithm
of a ratio which can be solved by performing the antilogarithm of pH/pK a .

10 pH-pK o = g5j(5)



27



6.1.4 Experimental Procedure

6.1.4.1 Materials Required

• pH electrode and pH 7 buffer for calibration

• burette

• 250 mL beaker

• magnetic stirrer

• 0.4 M and 0.1 M NaOH

• 0.2 M phosphoric acid

• buffer solutions (pH 4 and pH 7)



6.1.5 Part I. Demo During PreLab Lecture: Drink Anyone?

1. Six wine glasses are filled with the same "mystery" liquid.

2. Each glass takes on a different color of the rainbow, despite the fact that the same liquid was added
to each.



6.1.6 Part II. Titration of Phosphoric Acid

1. Obtain a pH probe and connect it to pH/mV 1 on the Microlab interface. Open the MicroLab program
and select "Microlab Experiment." Choose "Add Sensor." This will bring up a window where you need
to select "pH /D.O.", click on the appropriate port of the interface, and choose "pH" out of the two
options below. To calibrate the pH probe, click "next."

2. Take a sample of two buffer pH standards at pH 4.00 and 7.00. Calibrate the pH probe with these two
solutions. This is done by selecting "Add Calibration Point" and entering the correct pH value noted
on the bottle. Note that the pH probe should always be rinsed with deionized water and
carefully patted dry before being inserted into a solution so as to avoid cross contamina-
tion. A large waste beaker is useful to have for rinsing. After the two points are entered, select linear
calibration, and save the calibration data.

3. In a dry beaker, obtain 30 mL of a 0.2 M phosphoric acid solution. Use a graduated cylinder to add 50
mL of deionized water to a 250 mL beaker. Rinse your 10 mL volumetric pipette with the phosphoric
acid solution and pipette 10 mL of the acid into the water. Rinse and fill your 25 mL burette with 0.4
M NaOH. The initial burette reading should be mL. Remember to clear the air out of the tip of the
burette.

4. Place the beaker on the magnetic stirrer and add a stir bar. Position the burette ready for titration.
Insert the pH probe. Turn on the magnetic stirrer and adjust the stirring rate to moderate speed
(without allowing the stir bar to splash or hit the probe).

5. On the Microlab main screen choose "Add Sensor" and select "Keyboard" under the sensor drop box.
Click "Next." This will bring up a prompt in which you should enter "KBD" in the top box and "mL"
in the bottom "units" box, and then hit "Finish."

6. Drag the keyboard sensor from the top left of the screen to "Data Source 1" on the x-axis of the graph.
Drag the pH sensor to "Data Source 2" on the y-axis of the graph. Drag the pH sensor to the box in
the bottom right corner.

7. Click "Start." Enter your starting volume, mL, in the window that appears, and hit enter. The
window will not disappear. Slowly add a small volume of NaOH to the beaker, approximately 0.5 mL,
enter the reading on your burette into the box, and hit enter. Repeat this process until both peaks
have been observed and the pH has stabilized.

8. To save your data, choose "export data" under File, and select "comma separated value." For help with
plotting the data and derivative of the data see the "Data Analysis" section below.



28 CHAPTER 6. ACID AND BASES TO BUFFERS

6.1.7 Part III. Buffers (Use the same MicroLab program)

1. Using equations (4 and 5), calculate the ratio of concentrations of Na 2 HP0 4 and NaH 2 P0 4 to pro-
duce 100 mL of buffer solution with pH = 6.91. Show your calculations to your TA before proceeding.

2. Prepare your buffer solution from 0.1 M Na2HP04 and 0.1 M NaH2P04 solutions.

3. Insert the pH probe in your buffer solution and wait until the reading becomes stable and write down
the value in your report form. Don't worry if the pH reading isn't exactly 6.91. The important thing
is that there isn't a drastic change in pH upon addition of acid or base.

4. Pour 50 mL of the buffer solution into another beaker so that you have two beakers each with 50 mL
of your buffer solution.

5. Add 1 mL of 0.1 M NaOH to the first beaker and mix the solution with a glass rod. Wait until the pH
reading becomes stable and write down the value in your report form. If the pH of your buffer solution
changes by more than 0.3 pH units, you will need to redo the calculations and re-prepare the buffer
solution in order to get an acceptable result.

6. Add 1 mL of 0.1 M HC1 to the second beaker and mix the solution with the glass rod. Insert the pH
probe into the second beaker. Wait until the pH reading becomes stable and write down the value in
your lab report form. If the pH of your buffer solution changes by more than 0.3 pH units, you will
need to redo the calculations and re-prepare the buffer solution in order to get an acceptable result.

Data Analysis

Two plots need to be made from the data taken in Part II. One plot, pH vs. volume of NaOH, can be
made directly from the data that is initially present. The data needs to be further analyzed to make the
plot of the first derivative. This plot should be ( A pH / Avol NaOH) vs. volume of NaOH. Remember to
include a title and axis labels.



Chapter 7

Forensics

7.1 Forensic Lab

Objectives

To appreciate the variety of tests available to Forensic Scientists

• To observe latent fingerprinting development

• To study ink identification and observe invisible ink

• To do a breathalyzer analysis

• To use luminal to detect chemicals present in blood

• To perform Chemical Spot Tests and analyze your dollar bill for possible drug residues

The Crime Scene

You are part of the CSI (Chemistry Scientists Investigators) team and are called to a crime scene where
a kidnapped man has been murdered during a drug raid. Your task is to analyze the fingerprint evidence to
determine who the culprit is. In addition, you need to identify the pen that the kidnapper used to write the
ransom note. Determine whether alcohol and drugs were present at the scene. Finally, similar to a scene
out of 'The Da Vinci Code', the murdered man wrote the initials of the murderer before he expired.

Introduction

Part 1: Latent Fingerprint Development

The earliest recognition of the uniqueness of fingerprints and their suitability for personal identification
came from the ancient Chinese, who employed a thumbprint in lieu of a signature on legal conveyances and
even criminal confessions. Since literacy was uncommon, this proved a practical measure. The first scientific
recognition of fingerprints in the West came in the 17th century, when the first studies on fingerprints were
published in England and Italy. Two hundred years later Sir Francis Galton published a book, Finger Prints,
were he proposed that no two fingers have identical ridge characteristics and fingerprints remain unchanged
during the individual's lifetime. Today the practice of utilizing fingerprints as means of identification is an
indispensable aid to modern law enforcement. Fingerprint identifications have solved a vast amount of cases.
Crime scene fingerprints fall into three types:

1. Patent or visible impressions occur as the result of transferring a foreign material (paint, grease, blood
or ink) coating the skin of the fingers to the object. 2. Plastic or molded impressions are deposited when the
hands, fingers or feet are pressed into a soft rubbery type material (wax, putty, clay or tar) that will retain
the impression of the ridge pattern in this material. 3. Latent or hidden impressions are left on polished
surfaces such as wood, metal or glass by the sweat-moist ridges of the fingertips. Since latent fingerprints
are not visible to the naked eye, they need to be developed using one of the following techniques:



1 This content is available online at <http://cnx.Org/content/ml9954/l.7/>.



29



30 CHAPTER 7. FORENSICS

• 'Powder and brush' technique: The surface is dusted with a very fine powder that sticks to the oils and
perspiration that are left behind from the friction ridges of the skin. Some surfaces, however, absorb
this powder and the fingerprints are not identifiable.

• Laser luminescence: Involves illumination of fingerprints which fluoresce due to particles picked up
during everyday life such as paints, inks and oil. It can be used on metals, plastic, cloth and wood.

• Ninhydrin test: Indantrione hydrate is sprayed onto the fingerprint where it reacts with the amino
acids, giving a dark purple deposit. It can be used to develop very old prints (made over 30 years ago).

• Iodine vapor: Can be used to develop fingerprints on fabrics and rough surfaces. Iodine vapor alone
is useful only for prints up to 24 hours old; however a mixture of the vapor with steam allows this
method to be effective for up to two months. Prints developed by this method disappear rapidly, so it
works well in situations where you want to conceal your work.

• Silver nitrate: Silver nitrate reacts with chlorides in the fingerprints, to give the insoluble salt, silver
chloride, which rapidly turns black on exposure to light. This method is not suitable for fabrics or
rough surfaces.

After developing the latent impression it is photographed and lifted with a clear tape to be placed on a
backing card with a contrasting background. It can then be entered into a computer, which allows it to be
quickly and easily recalled and compared to the fingerprint of a suspect. Identification depends on showing
a minimum of twelve matching characteristics in the ridge pattern. When these points of comparison are
shown, it is considered that the proof of identity has been established. In this lab you will be developing
your fingerprints using the iodine vapor and silver nitrate methods.

Part 2a: Identification of Inks

The identification of inks often plays an important part in document examination. As a rule, the ex-
amination centers on the question as to whether the ink of certain passages or of alternations in the text is
identical with the ink found in the possession of the suspect. For this reason the examination of questioned
documents seldom consists of a complete determination of the inks in question but is usually restricted to a
comparative analysis of certain properties of these inks.

Many different nondestructive techniques of the examination of inks are available: reflected infrared
radiation, reflectance microspectrophotometry, lasers and scanning electron microscopy. Unfortunately, the
reflectance methods are often subject to interference effects from "bronzing" or "sheering" of the ink.

Semi-destructive methods involve high-performance liquid chromatography (HPLC) and thin layer chro-
matography (TLC). Most chromatographic techniques are based on the minute sampling of a single written
character representative of the questioned text. Small samples of ink bearing paper are removed from the
document, they are then extracted with a suitable solvent, and the components of the solution are separated
using HPLC and TLC. If the inks being compared show different composition, they did not come from the
same pen.

In this lab you will be separating dyestuffs of several ballpoint pens using thin layer chromatography.
Comparison of the dye composition will allow you to find out which pen was used by your TA to spot the
TLC plate.

Part 2b: Invisible Ink

Invisible ink has been used to conceal secret messages for a long time. Many different liquids can be used
as invisible inks such as lemon juice, milk, vinegar or a solution of phenolphthalein.

Part 3: Breathalyzer

To determine whether a driver is driving under the influence of alcohol, law enforcement officers perform
a Breathalyzer test to measure the blood alcohol content of the bloodstream. In the breath analyzer test, a
breath sample is passed through a solution containing acidified potassium dichromate (K 2 Cr 2 O7), which is
bright yellow. Potassium dichromate, a strong oxidizing agent, oxidizes ethyl alcohol to acetic acid (vinegar).
The chromium is consequently reduced from the VI to the III oxidation state, which is green. The unbalanced
equation for this reaction is

Cr 2 7 2 " + H++ C 2 H 5 OH — > Cr 3 + + CH 3 C0 2 H + H 2

The amount of alcohol in a breath analyzer sample is therefore proportional to the amount of potassium
dichromate that is used up and also therefore to the loss of yellow color.



31

The Blood Alcohol Concentration (BAC) may then be calculated from the equation

BAC = 0.8 A/WR

Where W is a body weight of the individual being tested, A is the amount of alcohol in the body (in
mL) and R is a "Widmark R Factor", approximately 0.68 for men and 0.55 for women. In most states, a
BAC of 0.1 percent is sufficient to be convicted for driving under the influence of alcohol; in some states the
threshold BAC is even lower.

Part 4: Blood Stain Analysis Using Chemiluminescence

Investigators often find bloodstains during their examination of a crime scene. They also find stains that
could be similar substance, something other then blood, such as red paint. How would you test a stain to see
if it is blood? Human blood contains a pigment called hemoglobin, which is used to transport oxygen through
our body. This pigment is used by number of tests to identify the presence of blood. One most common test
used by investigators that reveals the presence of blood is the Luminol Test. In this test the bloodstain can
be made to glow with a blue light due to chemoluminescent reaction of the luminol reagent with the iron
(Fe) in the hemoglobin. Chemoluminescence is the reversed case of photoreaction: by a chemical reaction,
an excited particle is formed, which looses its energy by producing a light quantum of light. The most
important characteristic is that the light is emitted in cold. In other words, chemoluminescence happens
when a molecule capable of fluorescing is raised to an excited level during a chemical reaction. Upon its
return to the ground state, energy in the form of light is emitted. Luminol is one of the most outstanding
molecules that emit appreciable amounts of light.

Part 5: Spot Tests and Analysis of a Dollar Bill

The urban legend that 3 out of 4 dollar bills, which usually have a life span of 9 years, are contaminated
with trace amounts of cocaine is not only true but they could also be silent carriers of diseases such as
Hepatitis C too (Note: SARS and AIDS cannot be spread by contaminated bank notes). The contamination
arises from a dollar bill being used to snort cocaine or handled by a cocaine user, since cocaine is excreted in
skin oils and the contamination is then spread in bank sorting machines. Interestingly, Australia introduced
plastic currency in 1988 to prevent the ability of the crystalline structure of illicit drugs to gain a 'foothold'
on paper money, to make it harder for counterfeiters to reproduce and to last four times longer. They now
print polymeric notes for many countries including: Israel, Malaysia, Mexico and Romania.

Investigators often use Chemical Spot Tests (CST) as simple chemical reagents to test unknown suspect
powders by observing a quick colour change that will confirm or deny the presence of a drug or a class of drugs.
A caveat needs to be sounded here, since multiple compounds of the same class of drugs may give the same
colour change, e.g benadryl with active ingredient, diphenhydramine, a first generation anti-histamine drug
and since it has sedative properties is also used in sleep aids such as Tylenol PM where it is combined with
acetaminophen (paracetamol). Consequently, this is by no means conclusive evidence for that, you would
need to use more expensive, sophisticated instrumentation: such as gas chromatograph - mass spectrometer
(GC-MS), Fourier transform - infrared (FT-IR) and UV- visible techniques but is obviously easier to do
these simple, screening tests in the field.

Interesting Facts

The highest amount of cocaine detected on a single one-dollar bill was 1327 micrograms.

One gram of cocaine = head of a thumbtack

Heroin and ecstasy are less often found on banknotes as they degrade more rapidly than cocaine.

Germany had a problem with their Euros cracking and disintegrating due to the presence of sulfates from
the manufacture of methamphetamines mixing with human sweat to form sulfuric acid.

Experimental Procedure

Part la: Latent Fingerprint Development using Iodine vapors

Caution! Iodine vapors are poisonous and should not be inhaled. Keep the jar with iodine in the fume
hood at all times.

1. Press you finger onto a piece of filter paper.

2. Using tweezers place the filter paper into a jar with iodine and recap it.

3. When you can see the fingerprint clearly, remove the filter paper using tweezers and keep it until the
end of the lab.



32 CHAPTER 7. FORENSICS

Part lb: Latent Fingerprint Development using fuming cyanoacrylate.




Figure 7.1



Materials

Large hotplate

Large glass dish for fuming chamber

Bottom of soda can cut off to hold superglue

Superglue

Flat surface with fingerprints (covered in colored tape works well (black electrical tape))

Your TA will fingerprint each person in their group twice on 2 glass microscope slides, one labelled with
their name, one unnamed for your TA to decide which one of you is the culprit.

Each student should also use the ink pad to fingerprint themselves on a piece of paper to include in their
lab report.

After your TA develops your slides and along with your ink prints, identify who stole the priceless
manuscript amongst your lab group.

Experimental Procedures

***This experiment should be conducted in the hood***

1. Place finger prints on large tape covered watch glass. Alternately, small objects that contain fingerprints
can be placed in the chamber.

2. Make sure vial has water in it.

3. Place a drop of superglue the size of a nickel onto the soda can top.

4. Place watch glass on top of the fuming chamber.

5. Turn the hotplate on low and allow the fingerprints to develop for approximately ten minutes.

6. Turn the hotplate off and remove the fingerprint specimen. Allow it to cure for approximately 5
minutes.



33

Caution! Overheating cyanoacrylate can produce highly toxic hydrogen cyanide gas. The reaction should
take place in a well ventilated area with the hot plate on low.

Caution! Superglue can bond instantly to fingers and other body parts. Take caution in using it. In the
event of an inadvertent bonding, use acetone to soften the glue.

Part lc: Ink Pad

1. Press your finger on the ink pad and place on your report form.

Part 2a: Ink Identification

Your TA has written a ransom note and given it to you. Your task: determine which pen was used.

1. Obtain a precut TLC plate. Do not touch the white surface and handle carefully only by the edges.

2. Using a pencil, draw a light line across the shorter dimension 1 cm from the bottom. Using a ruler as
a guide on the line, mark off five equally spaced intervals on the line as shown in Figure 1.





C







Compound 2 (R f =c/i?)
Compound 3 (R f =d/a)



Original Plate



Developed Plate




V



Distance traveled by substance
Distance traveled by solvent front



Figure 7.2



Figure 1. TLC plate

3. Cut a few letters out of the ransom note written in ink. Dissolve the pieces in about 1 mL of acetone.
Use this to spot your TLC plate as the unknown using a capillary tube.

4. In a similar manner, spot using capillary tubes, your TLC plate with three 'standards' diluted in test
tubes:

Pen #1 : Foray super comfort grip retractable black gel ink pen fine point (0.7 mm)
Pen #2 : Paper Mate gel bold black ink (1.0mm)
Pen #3 : Pilot G2 gel black ink rolling ball fine point
The applied spots should be no bigger than 1-2 mm.



34 CHAPTER 7. FORENSICS

5. Use a 400 mL beaker for the development chamber. Add the solvent mixture (ethyl ac-
etate/ethanol/water in a ratio of 50:30:20) to the beaker to a depth of about 0.5 cm. Remember: the
level of the solvent must be below the spots on the plate. Using tweezers or forceps, place the spotted plate
in the development chamber so that it rests in the solvent and against the beaker wall. Cover the beaker
(with a paper towel) and allow the plate to develop. It might take more than 10 min, so you can proceed
with Part 3 and Part 4.

6. When the solvent has risen on the plate to within 1 cm from the top, remove the plate from the beaker
with tweezers or forceps. Using a pencil, mark the position of the solvent front.

7. Allow the plate to air dry and observe the colored separation. Note how many dye components make
up the individual pens, along with their color.

8. Using a millimeter ruler, measure the distance that each spot (use the center of each spot for consis-
tency) has traveled relative to the solvent front. Calculate the Rf values for each spot.

Part 2b: Invisible Ink
Method 1

Caution! Iodine vapors are poisonous and should not be inhaled. Keep the jar with iodine vapor in the
fume hood at all times.

1. Mix 1 g of cornstarch with 10 mL of water in a beaker and stir until smooth.

2. Heat the mixture for several minutes. Soften the point of a toothpick in the mixture and write a letter
or a message with it. Let the paper dry.

3. To observe the message, place it in a jar with iodine vapor in the hood. Recap the jar and let it stand
for more than 1 minute.

Method 2

1. Pour about 5 mL of phenolphthalein solution in a beaker. Soften the point of a toothpick in the
solution and write a letter or a message with it. Allow the paper to dry.

2. To read the invisible writing, dip a small piece of cotton wool into the 1 M NaOH solution and carefully
wipe the paper (do not rub!).

Method 3

1. Pour about 5 mL of the milk solution in a beaker. Soften the point of a toothpick in the solution and
write a letter or a message with it. Allow the paper to dry.

2. To read the invisible writing, heat the paper using a heat gun or hair dryer
Part 3: Breathalyzer

Caution! Sodium dichromate is a strong oxidizing agent. Avoid contact with the skin.

You are given four unknown clear, colorless liquids. Use the sodium dichromate to determine which one
is ethyl alcohol.

1. Pour 2 mL of each unknown into the test tube with the same number label and add 2 mL of acidic
sodium dichromate solution to each test tube. Mix gently and record your observations.

Part 4: Luminol Analysis

The Luminol solution and fake blood are prepared as follows, but it will be made for you ahead of time
for this lab.

Solution 1: Prepare by dissolving O.lg of Luminol in 20 ml of 10% NaOH in a 50 ml beaker and diluting
this solution to 200ml in your spray bottle and set aside.

Solution 2 ("blood"): Prepare by dissolving 0.5g of K3[Fe(CN)g] in 20 ml of 5% hydrogen peroxide solution.
This solution is only stable for a short period of time so look ahead and know what you need to do next
before you make the solution.

1. Go to a hood which is blacked out by construction paper and observe the initial of the killer written
by his victim just before he expired. Your TA will have smeared some "blood" (solution 2) onto
the bench inside the hood with a paper towel for you.

2. Quickly spray the wet "blood" stain with the luminol solution (solution 1) and close the hood to the
point that you can barely see at the bottom and record your observations.

Part 5: Spot Tests and Analysis of a Dollar Bill

2 bottles Cobalt thiocyanate reagent A (Each bottle: 0.5 g cobalt thiocyanate in 25 mL deionized water)



35

2 bottles Cobalt thiocyanate reagent B (Each bottle: 1 g stannous chloride in 20 mL deionized water)

Ethanol

Spot wells

1. Your TA will give you your spot wells to analyze.

2. Put 1 or 2 drops of the cobalt thiocyanate solution in deionized water into each well, if cocaine is
present a blue precipitate will form, if not, a colourless solution will be observed.

3. Be aware that lidocaine hydrochloride will also test positive.

4. Then soak your dollar bill on a watch glass in 10 mL ethanol to which you will add 5 or 6 drops
of stannous chloride solution in deionized water. Then add up to 5 drops of cobalt thiocyanate in
deionized water and observe any precipitate.



36 CHAPTER 7. FORENSICS



Chapter 8

The Curious Case of Catalase 1

8.1 The Curious Case of Catalase

8.1.1 Objective

• To prepare the enzyme o-diphenoloxidase from a vegetable or fruit.

• To study the effects of temperature.

• To study the specificity of enzyme activity.

• To observe the change in enzyme activity due to changes in pH.

• To study the effect of an inhibitor on your prepared enzyme.

• To put you off potatoes for life!

8.1.2 Grading

• Pre-Lab (10%)

• Lab Report Form (80%)

• TA Points (10%)



8.1.3 Background Information

Many details of how catalysis occurs have been obtained from the study of enzymatic reactions in biological
systems, where specific protein molecules called enzymes function as homogeneous catalysts. They produce
an increase in the rate of reaction by providing an alternate lower-energy pathway for the formation of
products. This phenomenon of enzymatic catalysis makes biological reactions necessary for the maintenance
of life possible. As biological catalysts, enzymes retain the characteristics of chemical catalysts: they increase
the reaction rate, remain unchanged after the reaction, have no effect on the equilibrium constant (K eq ) or
on the ultimate equilibrium conditions for a reaction, and are highly efficient. Enzymes help orient the
reaction participants to be more likely to react, to discriminate between one possible reactant molecule and
another with uncanny specificity, and sometimes to provide a coupling mechanism that ensures one reaction
always is accompanied by another reaction in a specific sequence.

A molecule acted upon by an enzyme is referred to as the substrate of that enzyme. The presence or
absence of a single atom, or a single charge, may decide whether a molecule is the optimum substrate or is
rejected by the enzyme. The ability of the enzyme to select from among many possible molecules with which
it could react is called enzyme specificity.



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37



38 CHAPTER 8. THE CURIOUS CASE OF CATALASE

Although some molecules sufficiently resemble the optimum substrate of an enzyme to bind to the active
site, they cannot undergo chemical reaction: they simply sit there, blocking the site rather like a bump on
a log, preventing the enzyme from functioning with the true substrate. Such molecular impersonators are
termed competitive inhibitors. This competitive inhibition can be reversible, since the impersonators can
be flushed off the enzyme with a sufficient excess of true substrate. DFP (diisopropyl fluorophosphate, an
organophosphate) is a potent and lethal nerve gas, i.e. an irreversible inhibitor as it irreversibly inhibits the
enzyme acetylcholinesterase, which is essential for the conduction of nerve impulses.



— o



Structure of DFP

Many organophosphorus compounds used as insecticides are deadly nerve toxins for exactly the same
reason.

The ability of an enzyme to catalyze a specific reaction is termed its activity - a measure of the rate
at which the reaction proceeds. Enzyme activity depends on several variables such as pH, temperature,
concentration, and specificity of substrate, cofactors, and inhibitors. Vitamins and minerals, two important
factors of human nutrition, play an essential role in the proper function of certain enzymes. Approximately
one-third of known enzymes require a metallic ion for their activity. The term cofactor is used to group
coenzymes and minerals within a general category.

An oxidase is an enzyme that catalyzes the transfer of hydrogen from some compound to molecular
oxygen in order to form water. Polyphenoloxidase is a copper-containing enzyme that catalyzes the removal
of hydrogen, as in the oxidation of dihydroxyphenols to the corresponding quinones. This type of oxidation
is accompanied by a color change. Polyphenoloxidases, also known as o-diphenoloxidases, are distributed
widely in the plant kingdom (e.g. champignon mushrooms, potatoes, bananas). They are responsible for
the darkening of freshly cut surfaces of plants or fruits. For insects, o-diphenoloxidase is important both for
melanin formation and for browning and hardening (sclerotization) of the cuticle. In this experiment, the
observed color change is used as a measure of extent of the reaction.



8.1.3.1 Experimental Procedure

8.1.3.2 Part I. Preparation of the Enzyme

1. Peel and grind two to three potatoes in a blender.

2. Add 25-30 mL of water to the mixture, swirl the mixture occasionally, and let it stand for about 10
minutes.

3. Filter the solution in cheesecloth to obtain a clear solution.

4. While you are filtering the solution, prepare the water bath that you will use in the following sections.

5. Use this solution as a source of catalase and o-polyphenoloxidase.



8.1.3.3 Part II. The Effects of Temperature

1. The gas-collecting apparatus has been prepared for you, and it consists of a large test tube fitted with
a rubber tube leading from a one-hole rubber stopper. The other end of the tube should be directed
beneath an inverted, water-filled, graduated cylinder in a large beaker filled with water.



39

2. Add about 5 mL of hydrogen peroxide, at room temperature, to the tube.

3. Quickly add 3 mL of catalase extract.

4. Stopper the test tube, swirl it once to mix, and collect the gas in the graduated cylinder.

5. Note the time required to collect 5 mL of the gas.

6. Repeat the procedure using both extract and hydrogen peroxide cooled in an ice bath. Record the
temperature and the time required to produce 5 mL of oxygen.

7. Repeat using extract and hydrogen peroxide warmed in a water bath to 37 �C. Again, record the
temperature and the time required to collect 5 mL of oxygen.

8. From your results, answer the question: What effect does temperature have on enzyme activity?

8.1.3.4 Part III. The Specificity of Enzyme Activity

In this part of the experiment, you will evaluate the specificity of the substrate structure required for enzyme
activity. The substrate structure varies to demonstrate the need for:

1. an aromatic or aliphatic ring

2. one or two hydroxy groups

3. proximity of hydroxy groups on the ring

The substrates include the following:

1. deionized water

2. cyclohexanol



.^;




Figure 8.1



1. 1,4 cyclohexanediol



Figure 8.2



40



CHAPTER 8. THE CURIOUS CASE OF CATALASE



1. phenol




OH



Figure 8.3



1. catechol (o-hydroxyphenol)



Ci-




Figure 8.4



1. resorcinol (m-hydroxyphenol)



OH



^Sa�



Figure 8.5



1. hydroquinone (p-hydroxyphenol)



41




Figure 8.6



Caution: Some of the solutions used in the next step are toxic and can be absorbed through the skin. Avoid
skin contact.

1. Prepare seven labeled 6-inch (15-cm) test tubes with labels and contents as follows:

Tube 1: 1.0 mL distilled waterTube 2: 1.0 mL 0.01 M cyclohexanolTube 3: 1.0 mL 0.01 M 1,4-
cyclohexanediolTube 4: 1.0 mL 0.01 M phenolTube 5: 1.0 mL 0.01 M catecholTube 6: 1.0 mL 0.01 M
resorcinolTube 7: 1.0 mL 0.01 M hydroquinone

To each tube add an additional 3.0 mL of distilled water. (These are the substrate tubes.)

1. Place all the labeled substrate test tubes into a 37 �C water bath.

2. Transfer 3.0 mL of vegetable extract into each of another set of seven test tubes, and then place them
into the constant temperature water bath. (These are the enzyme tubes.)

3. Allow the two sets of test tubes to remain in the constant temperature bath for five minutes.

4. Quickly add the contents of the enzyme tubes to the substrate tubes.

5. Remove all the test tubes from the water bath and compare the intensity of any developed color to
Tube 1, which contains distilled water.

6. Arrange the test tubes in progressive intensity of color and record the sequence. Tube 1 (blank) is
assigned a value of 0; the other tubes are then ranked 1, 2, 3, and so on.

7. What can you conclude about the development of a color?

8. On your report form, list the substrates that do not react with the enzyme.



8.1.3.5 Part IV. The Change in Enzyme Activity Due to Change in pH

The final part of the experiment will involve an investigation into the effect of pH on enzyme activity.

1. Prepare five small test tubes with labels and contents as follows:

Tube 1: 4 mL of pH 4.0 buffer solutionTube 2: 4 mL of pH 6.0 buffer solutionTube 3: 4 mL of pH 7.0 buffer
solutionTube 4: 4 mL of pH 8.0 buffer solutionTube 5: 4 mL of pH 10.0 buffer solution

1. Add to each of these test tubes 15 drops of 0.01 M solution of the most active substrate and 15 drops
of extract solution (enzyme).

2. Mix gently, and then incubate the test tubes in the constant-temperature water bath set at 37 �C.
While you are waiting 15 minutes for the color to develop, do not start the final section involving
copper sulfate.

3. After 15 minutes, examine each test tube for color changes and record your results.

4. From your results answer the question: What is the optimum pH range for maximum activity of the
o-diphenoloxidase enzyme?



42 CHAPTER 8. THE CURIOUS CASE OF CATALASE

8.1.3.6 Part V. The Effect of Adding Copper Sulfate

1. Place 5 mL of hydrogen peroxide in a test tube.

2. Place 5 mL of the catalase solution in a second test tube.

3. Add 5-10 drops of copper sulfate solution to the second test tube.

4. Place both test tubes in a water bath at 20 �C for 5 min.

5. Mix the test tubes and measure the amount of oxygen produced in 5 min.

6. Record your observations on your report form.



Chapter 9

Organic Reactions 1

9.1 Organic Reactions
9.1.1 Objectives



Synthesis of some important esters.

Oxidation of a primary alcohol first to an aldehyde and then a carboxylic acid.

To saponify a typical vegetable oil.



9.1.2 Grading

You will be assessed on



detailed answers required in the lab report.

the correctness and thoroughness of your observations.



9.1.3 Introduction

Esters are an important class of organic compounds commonly prepared from the esterification reaction of
an organic acid with an alcohol in the presence of a strong mineral acid (usually H2SO4). They are chiefly
responsible for the pleasant aromas associated with various fruits, and as such are used in perfumes and
flavorings. Some esters also have useful physiological effects. The best known example is the analgesic
("pain killing") and anti-pyretic ("fever reducing") drug acetylsalicylic acid, otherwise known by its trade
name aspirin.

Liniments used for topical relief of sore muscles contain the ester methyl salicylate ("oil of wintergreen"),
which is prepared from the reaction of methyl alcohol with the acid group of salicylic acid. Methyl salicylate
acts as an analgesic and is absorbed through the skin; however, methyl salicylate is also a skin irritant
(like many organic substances), which in this instance provides the beneficial side effect of the sensation of
warming in the area of the skin where the liniment is applied.

Oxidation of a primary alcohol may yield either an aldehyde or a carboxylic acid, depending on the
reaction conditions. For example, mild oxidation of ethanol produces acetaldehyde, which under more
vigorous conditions may be further oxidised to acetic acid. The oxidation of ethanol to acetic acid is
responsible for causing wine to turn sour, producing vinegar.

A number of oxidising agents may be used. Acidified sodium dichromate (VI) solution at room temper-
ature will oxidise primary alcohols to aldehydes and secondary alcohols to ketones. At higher temperatures
primary alcohols are oxides further to acids.



1 This content is available online at <http://cnx.Org/content/ml5483/l.3/>.

43



44



CHAPTER 9. ORGANIC REACTIONS



RCH 2 OH



H



OH



K 2 Cr a 7 , H + (aq) / K a Ci 2 Q 7 , H + (aq) /
*�■ RC. v „_. *■ RC V \



room temperature







50�C



R 2 CHOH



K 2 Cr 5 Q 7 , H + (aq) RC Q



Figure 9.1



The dichromate solution turns from the orange color of the Cr20j~ (aq) to the blue color of the Cr +
(aq). This color change is the basis for the "breathalyser test". The police can ask a motorist to exhale
through a tube containing some orange crystals. If the crystals turn blue, it shows that the breath contains
a considerable amount of ethanol vapor.

Soaps are produced by the reaction of metallic hydroxides with animal fats and vegetable oils. The major
components of these fats and oils are triglycerides. Triglycerides are esters of the trihydroxy alcohol called
glycerol and various long-chain fatty acids. Tristearin is a typical triglyceride. Upon reaction with sodium
hydroxide, the ester bonds of tristearin are broken. The products of the reaction are the soap, sodium
stearate, and glycerol. This type of reaction is called saponification (Greek: sapon, soap) and it is depicted
below.



CH 2 O-

3NaOH+ CH Q-

im, — o-





//

o

//
-C(CH 2 ) 16 CH 5 -

o

H

- C (CH 2 ) 16 CH 3



CH 2 OH o

CHOH + 3CH 3 (CH 2 ) 16 C



CH'jOH



•0"Na +



Figure 9.2



Soap is made commercially by heating beef tallow in large kettles with an excess of sodium hydroxide.
When sodium chloride is added to this mixture (called the "saponified" mixture), the sodium salts of the
fatty acids separate as a thick curd of crude soap. Glycerol is an important by-product of the reaction. It is



45

recovered by evaporating the water layer. The crude soap is purified, and coloring agents and perfumes are
added to meet market demands.

9.1.3.1 EXPERIMENTAL PROCEDURE

CAUTION WEAR EYE PROTECTION!

CAUTION - Concentrated sulfuric acid will burn and stain the skin as well as damage clothing. In case
of skin or clothing contact, wash the area immediately with large amounts of water.

9.1.3.2 Synthesis of esters

1. Place approximately 2 g (or 2 mL if the substance is a liquid) of the organic acid and 2 mL of the
alcohol in a large test tube.

2. Add 5-7 drops of concentrated (18 M) sulfuric acid, mix the solution well with a glass stirring rod
and then place the test tube in a hot water bath (largest beaker in your drawer) (~ 80 �C) for 5-10
minutes.

3. Remove the test tube from the hot water bath and cautiously pour the mixture into about 15 mL of
saturated sodium bicarbonate contained in a small beaker. The sodium bicarbonate will destroy any
unreacted acid.

4. Observe the aroma produced from each of the following esterification reactions. Write the structure of
the esters produced, and the balanced equations for the esterification and the acid/sodium bicarbonate
reactions:



Complete the following reactions using the procedure above and record your observations.

(1) C 7 H 6 3 + CH 3 OH ->
salicylic acid + methyl alcohol

(2) CH 3 CH 2 CH 2 CH 2 CH 2 CH 2 CH2CH 2 OH + CH 3 COOH ->
1 - octanol + glacial acetic acid

(3) CH 3 CH 2 CH 2 CH 2 CH 2 OH + CH3COOH -�
amyl alcohol + glacial acetic acid

(4) C 2 H 5 0~K + CH 3 COOH -�
ethanol + acetic acid



9.1.3.3 Oxidation of an alcohol with acidified potassium dichromate(VI) solution

• Add 10 drops of dilute sulfuric acid (6M) and 5 drops of potassium dichromate(VI) solution (0.01M)
to 5 drops of ethanol. The oxidising agent is added slowly to the alcohol so that the temperature is
kept below that of the alcohol and above that of the carbonyl compound. (Carbonyl compounds are
more volatile than the corresponding alcohols). Usually the alcohol is in excess of the oxidant and the
aldehyde is distilled off to avoid further oxidation.

• Note the color and smell cautiously (Royal Wave).

• Warm the mixture and smell cautiously (Royal Wave).

• Repeat the experiment using first methanol and then propan-2-ol in place of ethanol.

Describe what happens and explain the color changes.

What conditions and techniques would favour the oxidation of ethanol to

a. ethanal rather than ethanoic acid.

b. ethanoic acid rather than ethanal?



46 CHAPTER 9. ORGANIC REACTIONS

9.1.3.4 Oxidation of an alcohol with acidified potassium permanganate (VII) solution

• Add 10 drops of dilute sulfuric acid and 5 drops of potassium permanganate (VII) solution (0.01M) to
5 drops of ethanol. Note the color and smell cautiously.

• Warm the mixture and smell cautiously (Royal Wave).

• Repeat the experiment using first methanol and then propan-2-ol in place of ethanol.

• Take the pH of your final mixture using Universal indicator paper

Describe what happens and explain the color changes.
What is your final product?

9.1.3.5 Saponification of a vegetable oil

CAUTION - Sodium hydroxide is a very caustic material that can cause severe skin burns. Eye burns caused
by sodium hydroxide are progressive: what at first appears to be a minor irritation can develop into a severe
injury unless the chemical is completely flushed from the eye. If sodium hydroxide comes in contact with
the eye, flush the eye with running water continuously for at least 20 minutes. Notify your TA immediately.
If sodium hydroxide is spilled on some other parts of the body, flush the affected area with running water
continuously for at least 2-3 minutes. Notify your TA immediately.

Never handle sodium hydroxide pellets with your fingers. Use weighing paper and a scoopula. Solid
sodium hydroxide will absorb water from the atmosphere. It is hygroscopic. Do not leave the container of
sodium hydroxide open.

Keep ethanol and ethanol-water mixtures away from open flames.

Aqueous iron chloride will stain clothes permanently and is irritating to the skin. Avoid contact with this
material.

In this experiment, you will saponify a vegetable oil

1. Pour 5 mL (5.0 g) of vegetable oil into a 250-mL beaker.

2. Slowly dissolve 2.5 g of NaOH pellets in 15 mL of the 50% ethanol/ water mixture in a 50- mL beaker.

3. Add 2-3 mL of the NaOH solution to the beaker containing the oil. Heat the mixture over a hot plate
with stirring. CAUTION: Keep your face away from the beaker and work at arm's length. Stirring
is required to prevent spattering. Every few minutes, for the next 20 minutes, add portions of the
ethanol/water mixture while continuing to stir to prevent spattering. After about 10 more minutes of
heating and stirring, the oil should be dissolved and a homogenous solution should be obtained.

4. Add 25 mL of water to the hot solution. Using the hot grips, pour this solution into a 250 mL beaker
containing 150 mL of saturated NaCl solution. Stir this mixture gently and permit it to cool for a few
minutes.

5. Skim the soap layer off the top of the solution and place it in a 50-mL beaker.

6. Into a test tube, place a pea-sized lump of your soap. Place a similar amount of laundry detergent in a
second tube and a similar amount of laundry detergent in a second tube and a similar amount of hand
soap in a third tube. Add 10 mL of water to each tube. Stopper each tube and shake thoroughly.

7. Estimate the pH of the solution using wide-range indicator solution or wide-range test paper. Record
the results. Pour the contents of the test tubes into the sink and rinse the tubes with water.



2



Figure 9.3



47

9.2 Pre-Lab: Introductory Organic Reactions

9.3 (Total 25 Points)

Hopefully here 2 for the Pre-Lab

Name(Print then sign):

Lab Day: Section: TA

This assignment must be completed individually and turned in to your TA at the beginning of lab. You

will not be allowed to begin the lab until you have completed this assignment.
For questions 1-4, draw the structural formulae of:

1) 2,2 - dimethylbutane

2) 3-ethyl-2,4-dimethylpentane

3) 2,3,4-trimethylhexane

4) 3-ethyl-2-methylheptane



For questions 5-8, give the names of
5) CH 3 CH 2 CH 2 CH 2 CH = CH 2



6) CH 3 CH = C = CH 2

7) CH 3 CH = CHCH 3

8) (CH 3 ) 2 C = CHCH 3



For questions 9-11, give the structural formulae for:
9) hex-3-ene



10) 3-methylhex-l-ene

11) 2,5- dimethylhex-2-ene



For questions 12-14, give the names of:

12)

2 http://cnx.org/content/ml5483/latest/PreLabOR07.doc



48



CHAPTER 9. ORGANIC REACTIONS



Figure 9.4



13)



Br



Br



Figure 9.5



14)



Figure 9.6



For questions 15-19, give the stuctural formulae of:
15) trans- 1,2-dibromoethene



16) trans- 1-chloroprop-l-ene

17) cis- hex-2-ene

18) pent-1-yne

19) 3-methylbut-l-yne



For questions 20-25, name the following compounds:
20)



Figure 9.7



21)



22)



23)



49



Figure 9.8



Figure 9.9



\ /



24)



Figure 9.10



Br-



25)



Figure 9.11




50



CHAPTER 9. ORGANIC REACTIONS



9.4 Report: Organic Reactions

Hopefully here 3 for the Report Form

Note: In preparing this report you are free to use references and consult with others. However, you may

not copy from other students' work (including your laboratory partner) or misrepresent your own data (see

honor code).

Name(Print then sign):

Lab Day: Section: TA



9.4.1 Observations:

9.4.1.1 Synthesis of esters



Reagents


Product


Observations


C-jH^Oy, (salicylic acid ) +
CH 3 OH(methyl alcohol)






(1 - octanol ) +CH3COOH
(glacial acetic acid)


2 OH




CH 3 CH2CH 2 CH 2 CH 2 OH(amyl
alcohol) +CH3COOH (glacial
acetic acid)






C 2j ff 5 OH (ethanol) +CH3COOH
(acetic acid)







Table 9.1



9.4.2 Oxidation of an alcohol with acidified potassium dichromate(VI) solution.

Remember to describe what happens and explain the color changes.

What conditions and techniques would favour the oxidation of ethanol to

a. ethanal rather than ethanoic acid.

b. ethanoic acid rather than ethanal?



Add 10 drops of dil. # 2 S0 4 to 5
drops of K 2 CT 2 0r to the follow-
ing alcohols


Observations: color/smellSmell
cautiously!


Observations: color/smell on
warmingSmell cautiously!


Ethanol






Methanol






continued on next page



5 http://cnx.org/content/ml5483/latest/ReportOR07.doc



51



Propan-2-ol







Table 9.2



9.4,3 Oxidation of an alcohol with acidified potassium permanganate (VII) solu-
tion

Remember to describe what happens and explain the color changes.
What is your final product?



Add 10 drops of dil. H 2 S0 4 to 5
drops of KMnC>4 to the following
alcohols


Observations: color/smellSmell
cautiously!


Observations: color/smell on
warmingSmell cautiously!


Ethanol






Methanol






Propan-2-ol







Table 9.3



9.4.4

9.4.5 Saponification of a vegetable oil



Reagent


pH from indicator paper


Your soap




Laundry Detergent




Hand Soap





Table 9.4

w from the File menu, and then double-click your template.



52 CHAPTER 9. ORGANIC REACTIONS



Chapter 10

Kitchen Synthesis of Nanorust 1



Kitchen Synthesis of Nanorust

J.T. Mayo, Courtney Payne, Lauren Harrison, Cafer Yavuz, Dr. Mary McHale, Professor
Vicki Colvin

Objectives

• To perform a kitchen synthesis

• To obtain functional iron oxide nanocrystals that can be used for water purification of arsenic in Third
World countries by using everyday items found in any kitchen.

• To appreciate the many forms of iron and the many uses of iron oxide nanocrystals.

• To appreciate the advanced analytical instruments that are used in the continuing research of nan-
otechnology.

• Remember: Thinking simple saves money and lives.

Grading

Your grade will be determined according to the following:

• Pre-lab (10%)

• Lab Report Form (80%)

• TA points (10%)

Background



1 This content is available online at <http://cnx.Org/content/m20813/l.3/>.



53



54



CHAPTER 10. KITCHEN SYNTHESIS OF NANORUST




SUSGS



EXPLANATION
Arsenic, in u<^L
| | >50
| | 10-50
I 15-10
I I <5



Figure 10.1



Having clean drinking water is one of the most fundamental necessities of life. There are many different
forms of possible contamination, but among the most well-studied and problematic inorganic contaminants
is arsenic. Arsenic is one of the oldest known carcinogens. In 1999, the US National Academy of Sciences
reported that arsenic can cause bladder, lung and skin cancer, and possibly cause liver and kidney cancer.
The physical symptoms of arsenic poisoning include: extreme fatigue, nausea, vomiting, partial paralysis,
and reproductive damage. 2 Arsenic is naturally occurring in water due to its abundance in certain types of
rocks, but it can have anthropogenic origins as well.

Arsenic can be found all over the world, but is currently a particular problem in Third World countries
due to the costly nature of water purification. It is especially abundant in Bangladesh, but arsenic has also
been found in the ground water of Argentina, Chile, India, Mexico, Taiwan and Thailand. Additionally,
closer to home, most states in the western US have levels of arsenic concentrations of greater than 10 parts
per billion (10 ppb). This was not a cause for concern until the Environmental Protection Agency (EPA)
in 2006 lowered the maximum allowable level of arsenic from 50 ppb to 10 ppb. In 2001, approximately 13
million Americans were drinking water that had elevated levels of arsenic in the water. 3

Previous methods for arsenic removal have included: manganese greensand columns that have been
pretreated with dilute acid, coagulation/microfiltration, iron oxide based filtration, and activated alumina.
The "Arsenic Removal Using Bottom Ash" or "ARUBA" method, invented by Ashok Gadgil of the Lawrence
Berkeley National Laboratory, involves coating the surface of the contaminants with bottom ash and ferric

2 National Research Council.Arsenic in drinking water. Washington, DC, National Academy Press, 1999.
3 http:// www.epa.gov/safewater/arsenic/index.html



55

hydroxide. Bottom ash is sterile waste material from coal-fired power plants which would make the cost of
remediation about 0.5 cents per kg ARUBA of which generally 4-5 grams of ARUBA is needed for 1 liter of
water, initially containing 400 ppb arsenic.

Nanomagnetite synthesis for arsenic removal has been hailed as Forbes: 'Top 5 Nanotech Breakthroughs of
2006' and Esquire listed it as 'Six Ideas That Will Change the World' in 2007. Basically, the technique entails
forming iron oxide nanocrystals that possess very unique and size-dependent characteristics for environmental
remediation of arsenic contaminated water.

Introduction




Figure 10.2



Both iron oxide nanocrystals: Fe3C>4 (magnetite) and Fe2C>3 (maghemite, as it is a cross between MAG-
netite and HEMatite) are ferrimagnetic materials which means they can behave as permanent magnets.
Additionally, those oxides below 10 nanometers in diameter, exhibit superparamagnetic properties and are
used as MRI contrast agents.

Remember that last semester, you prepared a solution of magnetite ferrofluid by mixing iron(II) chloride
and iron(III) chloride in the presence of tetramethylammonium hydroxide.

Magnetite is the most magnetic of all the naturally occurring minerals on Earth and has shown a lot of
promise in environmental remediation as it efficiently removes As (III) and As(V) from water, this efficiency
of the removal increases ~200 times when the magnetite particle size decreases from 300 to 12 nm. Since
arsenic contaminated drinking water is a major problem around the world, using magnetite as a sorbent
shows a great deal of promise.

Additionally, Fe(II) compounds have been used to oxidize organic contaminants such as trichloroethylene
(TCE), while inorganic contaminants such as arsenic, lead and uranium are separated out of solution.
Between 10 and 20 nanometers, the contaminants can be removed from water via handheld magnets, which is
an important consideration in purifying water in the Third World, where power is not a standard commodity.

We will produce nanocrystalline and functional iron oxides following a green approach by using everyday



56



CHAPTER 10. KITCHEN SYNTHESIS OF NANORUST



items and equipment found in kitchens worldwide. The nanocrystalline and functional oxides are produced
by thermal decomposition of the iron-precursors in order to form highly uniform, isolatable nanocrystals
of tunable size. The iron precursors will decompose into iron oxides in organic solvents (thermally stable
non-polar solvents, aka fatty acids such as oleic acid) at temperatures in excess of 200 Q C; the presence of
amphiphilic stabilizers, in this case fatty acids derived from soap, limits the growth of crystalline products
which are either magnetite, maghemite, or mixtures of both phases.

The beauty of this method lies in the ability to use inexpensive iron sources, such as rust, to form
iron carboxylate intermediates, that when scaled to the gram level effectively produces a relatively low cost
method for removal of arsenic from contaminated water (see Table 1).

Rust is a mixture of iron hydroxides, oxides, and in some cases even zero-valent iron, but is as effective
as FeOOH used in any laboratory method. The fatty acid used in conventional methods is oleic acid, an
unsaturated 18 carbon fatty acid. It can be replaced by many cooking oils that can be processed to create a
homemade soap through saponification, by the addition of a base such as lye. The soap is allowed to cure for
a few days and then dissolved in a weak acid, such as vinegar. The organic layer of the liquid can be collected
and used without further processing. The "fatty acid mixture" or FAM is an impure fatty acid whose exact
composition depends on the starting edible oil. Olive oil contains the most oleic acid; coconut oils contain
more lineolic acid. For this lab, the FAM is derived from vegetable oil, a standard starting reactant.



Pure lab chemicals




Everyday chemicals




Chemical


Price per kg


Chemical


Price per kg


FeOOH


$ 778.00


Rust


$ 0.20*


Oleic acid


$ 20.60


Edible oil (coconut oil)


$ 0.25


1-octadecene


$ 24.75


Crystal drain opener (NaOH)


$ 1.24






Vinegar


$ 0.65


Magnetite Nanocrystals


$ 2,624.00


Magnetite Nanocrystals


$ 21.7



Table 10.1

Table 1. Cost comparison of the materials needed for a FAM/rust synthesis of magnetite nanocrystals
with a conventional laboratory synthesis. Most of the savings results from the reduction in cost of the iron
source. *Cost of the rust is an estimate.

In the kitchen synthesis, the black product that forms can be separated from the solution by simply using
a handheld magnet rather than the expensive and large centrifuges used in a conventional laboratory setting.

Experimental Procedure

Caution!! While all of the following chemicals and utensils can be found in a kitchen, this
procedure is potentially dangerous (even the soap is caustic). Gloves and goggles must be
worn at all times!!

Materials



vegetable oil

lye or 100% NaOH drain opener

wooden spoon

glass bowl

5% vinegar

cooking pot

hot plate

turkey baster or plastic pipette

rust



57

Part 1: Soap making process

This step requires a week of advanced preparation and has been done for you

1. In a crystallization dish or a similar container, weigh 100 g. of the liquid oil (if not liquid gently melt
it and keep as melted).

2. In a 50 mL vial (or a cup) weigh 15 g. of crystal drain opener (or caustic soda, or sodium hydroxide,
or potash).

3. Add 30 mL of tap water and shake (or stir) until all solid is dissolved (CAUTION: solution gets hot!).
While still warm pour it into the liquid oil.

4. Stir with a spoon (or a magnetic stir bar) for about 15 minutes (or until tracing occurs - tracing is
the visible tracks of stirring).

5. Let it sit open to air in a hood (or ventilated area) to dry and cure for a few weeks. If a shorter time
span is allotted, the soap can be dried in an oven. To help the drying process, the excess oil can be
decanted after 48 hours. Use caution when decanting the soap as there may be excess unreacted NaOH
present.

Part 2: Oleic acid from soap with commercial vinegar

1. Grate the soap given to you and weigh it. You should have approximately 30 grams. If you have extra
soap, do not discard it; give it back to your TA.

2. Check the vinegar's acidity (i.e. 5%).

3. Use 1 mL of acid for every gram of soap (i.e. 30 mL of acid, 600 mL of commercial vinegar with 5%
acidity) .

4. Combine the vinegar and soap in a cooking pot.

5. Heat on med-high and stir with a wooden spoon until all of the chunks are dissolved (light boiling is
preferred). - This takes 15 to 30 minutes.

Caution!! This must be done a hood or other well- ventilated area!!

1. Turn off the heating and cool the solution down.

2. Pour your solution into the 1L beaker provided. You should see two layers separating from each other.

3. Separate the top yellowish layer using a turkey baster or plastic pipette into a 50 mL beaker. Make
sure you get as much of the organic layer that as you can out of the vinegar/acid mixture.

4. Separate the organic layer off again into another 50 mL beaker so that you are left with only the organic
layer. This is your fatty acid mixture (FAM)

5. Thoroughly clean your cooking pot in the sink with soap and water.

Part 3: Magnetite nanocrystals from rust and fatty acids

1. Carefully measure 0.5 grams of rust. Do not discard excess rust; put it back in the stock container.

2. Mix the rust with the fatty acid mixture in the cooking pot.

3. Cover the top of the container with a loose cap for proper ventilation. The reaction smokes and steams.
This method produces 50-90 nm nanocrystals.

4. Start heating and timing. The rust should be heated for roughly 1 hour, until the solution is dark
black with little or no smoking. Remember to continue stirring at regular intervals. Do not heat the oil
to the point of popping and spattering. Adjust the heat as necessary so that the solution only steams
and smokes.

Caution!! This must be done a hood or other well- ventilated area!!

1. If your rust solution looks like it might dry out, notify your TA immediately and they will provide
you with extra oleic acid to complete your reaction. DO NOT LET YOUR RUST DRY ONTO THE
FRYING PAN!!!!!!!



58 CHAPTER 10. KITCHEN SYNTHESIS OF NANORUST

2. Once the rust solution appears dark black and little to no smoke is being produced, pour the nanorust
solution into a 50 mL beaker.

3. Hold a magnet to the side of the beaker and observe what happens. Hold the magnet to the beaker
for several minutes.

4. Clean the cooking pot thoroughly again with soap and water. If there is any black or burnt crusting,
this needs to be scrubbed away.



Chapter 11

Electrochemistry and Alchemy

11.1 Electrochemistry and Alchemy:

11.2 Molar Mass of Cu and Turning Cu into Au
11.2.1 Objectives

• To learn Faraday's two laws of electrolysis

• To relate an electric current to the passage of an amount of electric charge

• To discuss electrolysis in molten salts and in aqueous solutions

• To determine the molar mass of copper by electrodeposition from an aqueous solution

• To turn copper into "gold" (please bring a penny that is pre-1982 to lab with you)



11.2.2 Grading

You grade will be determined according to the following:

• Pre-lab (10%)

• Report Form (80%)

• TA evaluation of lab procedure (10%) which includes bringing a pre-1982 penny to lab



11.2.3 Before Coming to Lab . . .

• Read and complete the pre-lab







Read the background information



• Read and be familiar with the Lab Instructions







Find a pre-1982 penny and bring it with you to lab



11.2.4 Introduction

Electrochemistry describes the interaction between electrical energy and chemical processes. Electricity
continues to intrigue us, as it has since people first observed the sky shattered by bolts of lightning. Elec-
trochemistry is of great practical value to contemporary living. Consider the number of batteries used for
powering the many portable items of pleasure and need - everything from cassette recorders to hearing



1 This content is available online at <http://cnx.Org/content/m21301/l.4/>.

59



60 CHAPTER 11. ELECTROCHEMISTRY AND ALCHEMY

aids, from calculators to pacemakers. Pure metals are produced from natural ores, inorganic and organic
compounds are synthesized, metal surfaces are plated with other metals or coated with paint to enhance
their value and utility - all through electrochemistry.

Electricity is a moving stream of electrical charges. This flow, or electric current, can occur as electrons
moving through a wire or as ions flowing through an aqueous solution. If the electrons lost and gained in a
spontaneous reaction can flow through a wire on their pathway from the substance oxidized to the substance
reduced, the energy of the reaction is released as electrical energy. Conversely, a non-spontaneous redox
reaction can be driven forward by the introduction into the system of electrical energy from another source.
Any device in which either process can occur is called an electrochemical cell.

There are two types of electrochemical cells. The first type generates electrical energy from a spontaneous
redox reaction. These are called voltaic or galvanic cells, common household batteries are classic examples.
An Italian physicist, Allesandro Volta in 1800 explained that electricity is generated by the connection of
two dissimilar metals separated by any moist body (not necessarily organic). A simple voltaic cell, similar
to that made by Volta, can be assembled using twelve pennies and twelve nickels (construct a column of
alternating pennies and nickels with each coin separated by disk-size pieces of wet filter paper soaked in salt
water) .

In the second type of electrochemical cell, called an electrolytic cell, a non-spontaneous redox reaction
is caused by the addition of electrical energy from a direct current source such as a generator or a storage
battery. The process of generating a non-spontaneous redox reaction by means of electrical energy is called
electrolysis.

Electrolysis can be used for purifying a metal through the electrolytic dissolution of an impure anode
and the subsequent re-crystallization of the pure metal on the cathode. The impurities are left behind in
solution. Copper is refined commercially by this electrolytic technique.

Electrolysis is often used for electroplating a metal to another material acting as the cathode. The other
material must also be electrically conducting. Non-conducting materials, such as leaves, can also be plated
by first being painted with a metallic conductive paint. Silver plating can be done with a silver anode and
the object to be plated as the cathode.

Electrolytic reduction (cathodic reduction) has developed into a useful technique for the restoration of
artifacts such as corroded nails and encrusted silver. In the case of silver, the degradation is usually due to
the surface formation of insoluble (black) silver sulfide ( Ag 2 S). The artifact (a silver coin, for example) is
attached to the negative electrode of the electrolysis cell. The Ag + ions of the silver sulfide pick up electrons
and are converted back to metallic silver:

Ag 2 S* (s) + 2e~ -� 2Ag (s) + S 2 ~ (aq) (11.1)

The sulfide ions are swept away by the water and the surface of the object is restored.
In this experiment, you will electroplate copper quantitatively to a copper cathode (the anode is also
composed of copper). The current is measured over an interval of approximately one hour so that the amount
of charge passing through the cell is known. The molar mass of copper is calculated from its equivalent mass
using Faraday's second law. In the second part of the experiment, you will use turn copper into "gold"!

11.2.5 Background Information

In the 1830s, Michael Faraday published his experiments using the recently discovered voltaic column to
decompose substances through the use of electric current. Electrolysis is an oxidation-reduction process
involving a conversion of electrical energy to chemical energy. The electrolytic cell is a galvanic cell operating
in reverse. The automobile battery is acts as a collection of galvanic cells when delivering electric current,
but acts as a collection of electrolytic cells when being recharged.

Faraday first described the quantitative relationships between the amount of electric charge (number of
electrons) that has passed through an electrolytic cell and the amount of materials that have formed at the
electrodes. These are summarized as Faraday's Laws of Electrolysis:



61

1. The mass of substance reacting at an electrode is directly proportional to the total amount of electric
charge that has passed through the cell.

2. The masses of the substances reacting at the electrodes are in direct ratio to their equivalent
masses. The equivalent mass of a reacting substance is defined as its mass that reacts with one mole
of electrons in the oxidation or reduction process. In the case of sodium and chlorine, the equivalent
masses of the sodium and chlorine are equal to their molar masses; the equivalent mass of copper is
equal to its molar mass divided by two. The second law is a consequence of the stoichiometry of the
balanced half- reactions.

Through exhaustive experimentation, the charge of a single electron has been determined to be 1.602 x
10~ 19 coulombs (C). The coulomb charge unit - defined as useful for much larger charged objects - is
inconvenient for expressing such a small charge, so other electrical charge units are commonly used. One
mole of electrons has a total charge calculated to be 96,485 C; this quality is defined as faraday (F):
1 F = 96,485 C/mol e"
Electric currents (I) are measured in amperes (A), amps for short, and defined in terms

I = Q/t
1 A = 1 C/s
For example, a constant current of .600 A (milliamperes) over a period of 2.00 x 10 2 seconds represents

Q = I x 5 = 0.600 A x 200s = 0.600 C/s x 200s = 120 C (11.2)

a movement of 120 coulombs. The number of moles of electrons (n) transported during the time interval is

120 C
n =



96,485 G/mole 1



Figure 11.1



= 1.24 x 10" 2 mol e" 1 (11.3)

Time intervals measured in minutes and hours must be converted to seconds in such calculations.



62



CHAPTER 11. ELECTROCHEMISTRY AND ALCHEMY



11.2.5.1 Experimental Procedure



Voltmeter



Using voltmeter

- — V\A —



resistor



Stir bar
Magnetic stirrer




Figure 1



Figure 11.2



CAUTION WEAR EYE PROTECTION!



63

CAUTION: The 6 M nitric acid used in the next step will burn and stain the skin as well
as damage clothing. In case of skin or clothing contact, wash the area immediately with large
amounts of water.

1. Obtain a piece of copper mesh (about 5 cm x 8 cm) and remove any loose pieces of copper. Clean and
rinse. Place the copper mesh on a watch glass in the drying oven. Be careful not to touch the cleaned
surfaces. This is the cathode.

2. Obtain 1 piece of copper foil (about 2 cm x 8 cm). Holding the foil with tweezers or tongs, dip it into
6 M nitric acid several times until its surface is bright and shiny. Do not allow tweezers or tongs to
touch the acid solution. Rinse the foil in de-ionized water and set it aside. This is the anode. Set the
nitric acid aside to use in the electroplating exercise.

3. Add 350 mL 1.0 M KN0 3 solution to a 400 mL beaker.

CAUTION: The copper sulfate used in the next step is toxic. Avoid skin contact!

1. To this solution, add about 5 mL of 1 M 7J2SO4 and 10 g of CUSO4 • 5H 2 <3. Stir until the copper
sulfate pentahydrate is fully dissolved.

2. Assemble the apparatus shown in Figure 1, but leave the copper mesh electrode in the oven. Add a
magnetic stirring bar to the beaker. If necessary, add additional 1.0 M KNO3 to bring solution level
in the beaker within 2 cm of the rim. You will either measure the electricity directly with an ammeter
in series with the electrolytic cell or you will measure the current indirectly by measuring the voltage
across a resistor of known value (about 10 ohms).

3. Remove the copper mesh electrode from the oven, let it cool, and determine its mass to the nearest
milligram.

4. Attach the copper mesh electrode to the negative terminal of your power supply using an alligator clip.
Turn on the magnetic stirrer.

5. Turn on the low voltage power supply and adjust the current until about 140 mA are flowing through
the cell. Record the time and current.

6. Record the time and current every five minutes for an hour.

7. After the last reading, gently remove the cathode from the solution while leaving it attached to the
power supply. After the copper mesh has cleared the solution, remove the wire and turn off the power
supply.

8. Gently dip-rinse the copper mesh electrode several times in a beaker of deionized water, and place it
on a watch glass in the drying oven. Be careful not to rinse the mesh too harshly because you don't
want any copper that has deposited on it to come off.

9. When dry, remove the electrode from the oven and let it cool. Reweigh the mesh electrode.
10. Remove the magnetic stirring bar from your beaker and dispose of the solution in the sink.

11.3 Alchemy: Copper into Gold

Place your pre-1982 copper penny in an watch glass and heat with a mixture that first turns it silver, then
suddenly turns it "gold" when the penny is then heated on a hot plate.

Caution: Wear safety goggles and gloves and do the reaction in the fume hood with the
sash down. Note step 10: special disposal.

1. Place approximately 2 g of zinc in an watch glass.

2. Add enough NaOH solution to cover the zinc and fill the dish about one-third.

3. Place the dish on a hot plate and heat until the solution is near boiling.

4. Prepare a copper penny (pre-1982) by cleaning it thoroughly with a light abrasive (steel wool pads
work well).

5. Using crucible tongs or tweezers, place the cleaned penny in the mixture in the dish.

6. Leave the penny in the dish for 3-4 min. You will be able to tell when the silver coating is complete.



64 CHAPTER 11. ELECTROCHEMISTRY AND ALCHEMY

7. Remove the penny, rinse it, and blot dry with paper towels. (Do not rub!) Remove particles of zinc.

8. Using crucible tongs or tweezers, place the coated penny on the hot plate. The gold color appears
immediately.

9. When the gold color forms, remove the coin, rinse it, and dry it with paper towels.

10. Special disposal procedures: Do not discard the waste zinc in the trash container. When zinc dries,
it forms a powder that may spontaneously ignite. Rinse the NaOH-zinc mixture several times with
water. Then add the solid to a beaker that contains 200 mL of 1 M H2SO4. When all of the solid
dissolves, flush the zinc sulfate solution down the drain.



11.3.1 THE GOLDEN PENNY EXPERIMENT.

1. Your TA will set this up for you by putting 8 g of 30 mesh zinc in the bottom of a 400-mL beaker. It
is best to weigh-out the zinc in a watch glass and pour the zinc into a tilted 400-mL beaker so as to
keep the zinc on one side of the beaker. Use a spatula and tilt and tap the beaker on the bench top in
order to get all the granular zinc to cover about half the bottom of the beaker (one-half not covered:
see Figure 5). Carefully pour 200 mL of 1 NaOH down the side of the beaker, being careful not to
disturb the distribution of zinc. Use a stirring rod or spatula to clear any remaining granules so that
half of the beaker bottom is completely free of zinc granules. Place the beaker on a hot plate in the
fume hood and turn the hot plate to medium heat. The solution should be heated to about 80-90 �C;
if it is heated to boiling the distribution of zinc granules will be disturbed. Continually monitor and
check the temperature to keep it in this range. 20-25min.

2. While waiting for the solution to heat, buff six copper pennies with steel wool until they are shiny.
Wash them with deionized water and dry. Solder 10-cm lengths of 20-gauge copper wire to two of the
pennies, overlapping the wire and penny about 2 to 3 mm from the edge. Solder the free end of one
of the copper wires to a 5 x 100 mm strip of zinc metal, as shown in Figure 5. Clean any rosin off the
soldered joints with steel wool, and rinse with water.



65



B




1 M NaOH
solution



30 mesh
granular
fine



Figure 11.3



Figure 2

1. When soldering, place the penny on a ceramic fiber square, with the end of a copper wire overlapping
the penny 2-3 mm. Ask a partner to apply pressure to the wire to hold it in place while you are
soldering the joint -use a rubber stopper as before to do this.

2. Four arrangements of copper pennies in the golden penny experiment.

The bottom of the beaker is half-covered with 30-mesh zinc metal. (A) Two pennies are lying on top of the
30-mesh zinc.

(B) Two pennies are lying on the bottom of the beaker but not in contact with the 30-mesh zinc. (C)
A penny soldered to copper wire is immersed in solution, the other end of the copper wire being soldered



66 CHAPTER 11. ELECTROCHEMISTRY AND ALCHEMY

to a strip of zinc metal in contact with 30-mesh zinc on the bottom of the beaker. (D) A penny soldered to
copper wire is immersed in solution. The solution in the beaker is 1 M NaOH.

Review the General Soldering Instructions in Part I. Place the freshly tinned tip of the penny
next to the wire angling the iron to get good thermal contact. Don't dab at the joint with the tip of the
iron while soldering.

1. When the solution has warmed, use forceps to place two copper pennies on top of the granular zinc
metal and two pennies in the area that is free of granular zinc (make sure that the pennies on the
uncovered side do not contact even one grain of zinc). Bend a small "foot" on the penny in the beaker
as shown in Figure 5. The "foot" of the zinc strip should rest on the granular size so that both are
in direct contact. The penny should be completely immersed in solution but should not contact any
granular zinc metal on the bottom of the beaker. Finally, hang the last penny (the one with only copper
wire soldered to it) over the edge of the beaker so that the penny is completely immersed in solution.
See Figure 5. HINT- use an empty 400-mL beaker to bend and shape your soldered metal
pieces to match what is pictured in Figure 5. Do this before attempting to put them in the
400-mL beaker containing your warm NaOH solution. 30min.

2. Leave the pennies in the beaker until some of them turn a silvery color. This may take anywhere from
5 to 30 min., depending on the temperature of the solution. (Some of the pennies will never turn silver
even after waiting an hour or more.)

3. Which pennies turn a silvery color? Is it the three pennies that are in contact with the
solution and with zinc, either directly or through the copper wire? Or is it the three
pennies in contact with the solution but not indirect or indirect contact with zinc metal?
5-10min.

4. Using a pair of forceps, remove the pennies that have turned a uniform silvery color, rinse them with
water, and put the pennies on a hot plate for a few seconds. Watch what happens to the silver-colored
pennies as they heat on the hot plate. Keep the solution warm in the beaker in case you need to repeat
some part of the experiment or try some new experiment, as described below. 5-10min.

5. Further Experiments. Solder another length of copper wire to a shiny clean penny. Actually, you can
use the lone penny soldered to the copper wire from the first part of the experiment, just clean it with
steel wool and deionized water. Connect one lead of a voltmeter to a zinc strip and the other lead
to the copper wire soldered to the penny. Using the same solution you prepared earlier, immerse the
zinc strip and penny in the solution. Is there a voltage difference between the zinc strip and
the copper wire soldered to the copper penny? Which metal is the electron source (the
negative terminal of this electrochemical cell)?5-10min.

6. Do you think a current flows in the copper wire connecting the zinc strip and copper
penny when both are immersed in the solution? If so, which direction will electrons
flow, and what are the anode and cathode reactions?Put the digital voltmeter into its current
measuring mode on its most sensitive (microampere) scale. Then see if any current is flowing when you
connect the meter in series between the copper wire soldered to the penny and to the zinc strip -both
the penny and the zinc strip should be immersed in the hot 1 M NaOH solution. The series connections
should look like this: Penny/copper wire/(+) ammeter(-)/zinc strip. How large a current flows? .
Is the current (charge flow) from penny to zinc strip or vice versa?

7. In wires the charge carriers are electrons. Current (defined as a flow of positive charge) is opposite to
the flow of electrons. In which direction are electrons flowing: from penny to zinc strip or
vice versa?



11.3.1.1 Example of the calculation of Molar Mass:

1. Graph the electric current (in amps) on the y-axis against time (in seconds) on the x-axis. The total
charge that passed through the electrolysis cell is given by the area beneath this curve. If the current
is constant, this area is:



67



Q = area = I x t (11-4)

Calculate this charge in coulombs.

1. Convert the coulombs of charge to mol electrons:

N =



Q(C)



96,4S5C/mol

Figure 11.4



The equation for the reduction half-reaction responsible for the plating at the cathode is

Cu 2+ (aq) + 2e" -> Cu (s) (11.5)

Use the mol ratios of the preceding balanced equation to calculate the number of moles of Cu plated out:

n(Cu) 1
n 2

Figure 11.5



n(Cu) = n/2

1. Use the initial and final masses of the copper mesh electrode to calculate the mass of copper plated
out:

m(Cu) = m(final) - m(initial)

1. Calculate the molar mass (M) of copper:

M =

m(Cu)
n

Figure 11.6



68 CHAPTER 11. ELECTROCHEMISTRY AND ALCHEMY



Chapter 12

From Cells and Electrodes to Golden
Pennies 1

Protocol adapted from 'Chemistry in the Laboratory'

12.1 ElectroGels - Metal Corrosion, Anodic Protection and the
Golden Penny

12.1.1 Objective

The goal of this experiment is:

• Observe the deposition of zinc metal on a penny in contact with metallic zinc and to develop a hy-
pothesis to explain how this happens.

• Observe how iron may be corroded or protected from corrosion depending on whether it is connected
to copper or to zinc.

• Observe the plating of a penny with zinc and subsequent color changes.

12.1.2 Grading

You will be assessed on:



• completion of the report form.

• TA evaluation of lab procedure.



12.1.3 Introduction

12.1.3.1 Electrochemistry is Everywhere



In general, whenever two condensed phases (solid or liquid) are brought into contact, a potential (or voltage)
difference develops across the interface. Because the interface region is very thin, even transfer of a small
amount of charge across the interface can create a very large electric field. For example, transferring about
one picomole ( 10 -12 mole) of electron charge per square centimeter of area will typically create a potential
difference of approximately 1 volt across an interface layer about one nanometer thick. The electric field in



lr This content is available online at <http://cnx.Org/content/ml5879/l.7/>.



69



70 CHAPTER 12. FROM CELLS AND ELECTRODES TO GOLDEN PENNIES

this interface region would be about 109 volts/meter. Electric fields this large can cause the transfer of elec-
trons across an interface layer or the transfer of ions between the inside and outside of ells in living organisms.
Because contacts between condensed phases are very common in nature, electrochemical phenomena are very
common, even though we are often unaware of them. At the cellular level, electrochemical phenomena are
crucial to the propagation of nerve impulses, the timing of muscular contractions of the heart, and activity
in your brain cells.

Most of the electrical technology created by humans involves the simplest kind of chemical change;
electron transfer across an interface. Often, the interface is between a good electron conductor, called an
electrode, and a solution containing molecules or ions. The electrode might be a solid (like platinum or
copper metal or graphite), or it could be liquid (like mercury metal). When electrons are transferred from
the electrode to a molecule, we say the molecule has been reduced. Electron transfer in the opposite sense
(from molecule to electrode) is called oxidation.

There are two parts to this lab the first, ANODIC PROTECTION, you will perform in pairs. By coating
steel (which is mostly iron metal) with a more active metal like zinc, a process called galvanizing, the steel's
corrosion can be retarded or entirely prevented. Often, simply making a good electrical connection between
a piece of iron and a piece of zinc is sufficient to keep the iron from corroding. We will study the acceleration
and the prevention of iron corrosion by connecting it to various metals.

In the second part of this lab, we will be revisiting THE GOLDEN PENNY EXPERIMENT, which will
be set up for you by your TA, so that you can make observations. As you remember, the golden penny
experiment involves the plating of a penny with zinc metal. First, the penny is immersed in a solution
containing 1 M NaOH and granular zinc. Subsequent heating of the penny for a few seconds on a hot plate
causes the silver color of the penny to turn a bright golden yellow. Explaining the details of the process
presents a challenge.

12.1.4 EXPERIMENTAL PROCEDURE

12.1.4.1 Special Supplies:

Part 1: Nominal 100 x 15 mm disposable polystyrene. Petri dishes (three per group); fine steel wool;
approximately one soldering kit for every six students consisting of 140-watt soldering iron, rosin-core solder,
and one 6x6 inch ceramic fiber square (available from Flinn Scientific Inc.); digital voltmeters with alligator
clip leads and 2 short lengths (3 cm) of Pt wire to use as voltage probes.

Part 2: Digital voltmeters with alligator clip leads, 6 pennies per group (preferably clean and bright),
hot plates, stainless steel forceps; approximately one soldering kit (see the description in Part 1) for every
six students.

12.1.4.2 Chemicals:

Part 1: Agar (powder), 1% phenolphthalein indicator, 0.1 M potassium ferricyanide [hexacyanoferrate)III],
K 3 Fe (CN) 6 ; two zinc metal strips, 6 x 40 mm, cut from 0.01-inch thick zinc foil; two copper metal strips 6
x 40 mm, cut from 0.01- (or 0.005)-inch thick copper foil; 2 ungalvanized finishing nails per group (before
use, clean by soaking briefly in 3 M H2SO4 acid, rinsing with deionized water, and drying in an oven).

Part 2: 30-mesh zinc metal; zinc metal powder, 6 x 100 mm strips of zinc metal (one per group) cut
from 0.01-inch thick zinc foil; 20 gauge copper wire; 1 M NaOH, 1 M HCL in dropper bottles, and a 1 M
NaOH/Zn (N0 3 ) 2 50:50 mix solution.

! SAFETY PRECAUTIONS WEAR EYE PROTECTION AT ALL TIMES. Sodium hydroxide is cor-
rosive. You may want to provide latex rubber gloves for handling pennies that have been in contact with 1
M NaOH.

WASTE COLLECTION: Your instructor may direct you to waste containers for NaOH solutions used in
this experiment. These substances can be disposed of down the drain only if they are neutralized by sodium
bicarbonate.
5-10 min.



71

METAL CORROSION AND ANODIC PROTECTION.

1. Obtain two 6 x 40 mm strips of zinc foil, two 6 x 40 mm strips of copper foil, and two 4-penny (40 mm
long) ungalvanized iron finishing nails (which should have been previously cleaned by immersion in 3
M 7J2SO4, then rinsed with deionized water, and dried in an oven).

2. Clean the zinc and copper strips with steel wool to produce a clean, shiny surface.

3. Use detergent (1% alconox located by each sink)to remove the film of oil on the strips, rinse them with
deionized water, and dry them with a tissue. 15 min. experienced, 30 min. novice

4. General Soldering Instructions. Go to a soldering station where you will find a ceramic fiber square,
soldering iron, and rosin core solder. When soldering, place the zinc strip on a ceramic fiber square
and position the copper strip so the ends of the two strips overlap by about 4-5 mm. Ask a partner to
apply pressure to the copper strip, holding it in place while you are soldering the joint. (Your partner
may use almost any tool for this except a bare finger, because the strip will get very hot- a rubber
stopper is recommended.) In the combination using a iron nail, have the nail ontop of the other metal,
heat the nail in order to make a good seal.

5. Plug the soldering iron in and let it heat up for a couple minutes. Then tin the tip of the iron with
solder, wiping off all excess with a damp sponge or damp paper towels. Place the freshly tinned tip on
the copper strip next to the zinc strip, angling the iron to get good thermal contact. Let the copper
strip heat up for a good 30 seconds.

6. Feed the solder into the area between the zinc strip and the tip of the soldering iron. When the copper
strip is hot enough, the solder will flow into the joint. Don't dab at the joint with the tip of iron while
soldering. The tip must be kept in continuous good thermal contact with the joint so that the copper
strip heats up. Using steel wool, remove any rosin remaining from the soldering operation; then rinse
the soldered joints, and dry the metal strips with a tissue.

7. If the metal strips no longer have a clean, shiny surface due to excessive touching and handling,
use detergent to remove the oily film as before and rinse them thoroughly with deionized water.
WARNING!- be very gentle cleaning your soldered metal strips to avoid breaking the joint you just
made!

8. If you are a novice at soldering or have never soldered before, check out the helpful hints at: How to
Solder. 2

9. Now you should have a Zn/Cu piece, a Cu/Fe piece, and a Zn/Fe piece. (Once these soldered bimetal
pieces have been made, they can be cleaned and reused several times, so don't throw them away at the
end of the experiment unless your instructor direct you to.) 10-15 min.

10. Add 3 g of agar to 225 mL of boiling deionized water in a 600- mL beaker. At this point it's best to
turn down the heat and stir with a large magnetic stir bar until the agar is dissolved. Be patient, the
agar tends to clump and complete dissolution can take 10 min. Be careful not to heat the agar so
much that it scorches or boils over. While the agar suspension is hot, add 2 mL of 1% phenolphthalein
indicator with a calibrated transfer pipette. Continue stirring for a couple minutes. 5-10min.

11. Get three nominal 100 x 15 mm polystyrene Petri dishes. Put each of the three soldered bimetal pieces
in the bottom (taller) half or a Petri dish. Protecting your hands from the beaker containing the hot
agar solution with "hot hands" or paper towels, pour the agar over the metal pieces in the Petri dishes.
Cover the metal pieces, but do not fill the dish beyond half its depth. There should be no voids or
bubbles underneath the metal strips.

12. After the agar cools and gels (approximately 20 min.), use a fine-tipped transfer pipette to place one
small (0.02 mL) drop of 0.1 M potassium ferricyanide K^Fe [III] [CN] 6 along side each iron nail, about
1 cm away from the middle. The ferricyanide salt will diffuse radially outward. If the ferricyanide
ions encounter any Fe 2+ formed by oxidation of Fe, they will react with the Fe(II) ions to form a dark
blue compound formulated as KFe (III) Fe (II) (CN) 6 . (It's reported that you get the same produce by
mixing Fe + with K^Fe [II] [CN] 6 ). Evidently in the final product the iron atoms have exchanged an
electron so that the Fe 2+ ion is oxidized to Fe(III), and the Fe(III) originally in the ferricyanide ion is



2 http://aaroncake.net/electronics/solder.htm



72 CHAPTER 12. FROM CELLS AND ELECTRODES TO GOLDEN PENNIES

reduced to Fe(II). WARNING!- the drop of potassium ferricyanide will not immediately "soak" into
the agar gel. It will remain on top like a bead of water. Be careful not to jar or move the Petri dish
too much after placing the drop.

13. During the course of the afternoon, make periodic observations of the Petri dishes. Look for evidence
of formation of hydroxide ion, which will turn the phenolphthalein pink; the formation of insoluble
metal ion-hydroxide salts, which will appear as a cloudy band; or the formation of a blue compound
in those dishes containing iron nails (with an added drop of 0.1 M potassium ferricyanide, i^Fe [CN] 6 ,
indicating oxidation of Fe to form Fe + .20-25min.

14. Obtain two short lengths (about 3 cm long) of platinum wire and a digital voltmeter with leads
connected to alligator clips to hold short lengths of wire that will be used as voltage probes. Adjust
the voltmeter to its most sensitive voltage range (200 millivolts). First, carefully clamp the alligator
clips to the platinum voltage probes and immerse the probes in the agar, one probe midway alongside
one metal and the other probe midway along the other jointed metal. (Support the probes at all
times with your hands and keep the probes upright, perpendicular to the Petri dish. Make sure the
probes do not touch the metal strips.) Note whether there is any voltage difference. Note the polarity.
Which metal is nearest the positive (+) end? Any voltage difference indicates an electric field between
the two points in the agar created by the formation of positive and negative ions in the two regions.
Considering the polarity of the measured field, what ions do you think might be responsible fore the
presence of the electric field? Write plausible reactions for the formation of positive ions (metal atom
oxidation) and negative ions (reduction of water or oxygen).

15. Next, touch the probes directly to the two metals at their midpoints and note any voltage reading.
(The voltage reading is expected to be zero volts because metals are such good electronic conductors
that only a tiny electric field can exist in the two metals together.) The metals soldered together are
said to form an equipotential surface (a surface where the potential is constant, so that the voltage
difference between any two points on the metal is zero.)

16. Put the top cover on your Petri dish, and tape the top cover in place with two or three short strips of
tape. Write your initials or other identifying marks on the tape.

17. Continue visual observations in your next lab period, looking for evidence of formation of any pink color
or any visible precipitates. Make sketches, and write verbal descriptions of the changes you observe.



12.1.5 REACTIONS THAT MIGHT FORM OH- ION.

When a pink color develops around a metal in a gel containing phenolphthalein indicator, it means that the
solution next to the metal is basic. In an aqueous gel, the pink color means that some hydroxide ions have
been formed.

Although any electrons given up when a reactive metal is oxidized might react at the spot where the
oxidation occurs, they can also readily travel to any other spot on the surface of the two joined pieces of
metal. That means, it is possible that the point where metals atoms are oxidized could be some distance
from the point where hydroxide ions are produced.

Now let's think about what might be most likely to accept these available electrons. Metal atoms
typically don't accept electrons to form negatively charged metal ions. Rather, metal ions tend to give up
electrons to form positive ions. Things that are easy to reduce have the most positive standard reduction
potentials, like halogens, but we don't have any halogens in our system. The gel surrounding the metal
consists mainly of water with about one percent of agar. Although water is not easy to reduce, because
water has a negative standard reduction potential in basic solution, this substance can be reduced when
the reaction is coupled to the oxidation of Zn metal in basic solution, as shown by the following standard
reduction potentials:

Zn (OH)^ + 2e" -> Zn (s) + 40H"



(��inlMOH" = -1.28volts) (12.1)



73

2H 2 + 2e~ -� H 2 (g) + 20H~

(�; o mlM0H" = -0.80volts) (12.2)

Agar is a polysaccharide (like starch), and polysaccharides are not easy to reduce. Finally we must not
forget that the Petri dishes are open to the air, so the agar gel also contains dissolved oxygen, a good acceptor
of electrons. At least two reactions involving oxygen deserve serious consideration:
2 (g) + H 2 + 2e~ -> H 2 + OH~



(��inlMOH~ = -0.065volts) (12.3)



2 (g) + 2H 2 + 4e" -> 40H"



(��inlMOH- = -0.40volts) (12.4)

The E � for reduction of oxygen in basic solution is considerably more positive than for the reduction of
water. So we definitely must consider the possibility that oxygen might be the species that could most easily
be reduced, with OH - (and possibly hydrogen peroxide) being the reduction product.

The reduction of either water or oxygen produces hydroxide ions, but the formation of a pink color with
phenolphthalein does not tell us which reaction might be responsible. Thermodynamics (as measured by the
standard reduction potentials) favors reduction of oxygen over reduction of water. However, the reduction
of oxygen on many metals is known to have a large activation energy, which usually causes the reaction to
be slow. Thus, kinetics may favor the reduction of water, particularly because the concentration of water is
much greater than the concentration of oxygen in the agar gel. Can you think of an experiment that might
allow you to distinguish if water or oxygen is the major species being reduced? 20-25min.

12.1.6 THE GOLDEN PENNY EXPERIMENT.

1. Your TA will set this up for you by putting 8 g of 30 mesh zinc in the bottom of a 400-mL beaker. It
is best to weigh-out the zinc in a watch glass and pour the zinc into a tilted 400-mL beaker so as to
keep the zinc on one side of the beaker. Use a spatula and tilt and tap the beaker on the bench top in
order to get all the granular zinc to cover about half the bottom of the beaker (one-half not covered:
see Figure 5). Carefully pour 200 mL of 1 NaOH down the side of the beaker, being careful not to
disturb the distribution of zinc. Use a stirring rod or spatula to clear any remaining granules so that
half of the beaker bottom is completely free of zinc granules. Place the beaker on a hot plate in the
fume hood and turn the hot plate to medium heat. The solution should be heated to about 80-90 �C;
if it is heated to boiling the distribution of zinc granules will be disturbed. Continually monitor and
check the temperature to keep it in this range. 20-25min.

2. While waiting for the solution to heat, buff six copper pennies with steel wool until they are shiny.
Wash them with deionized water and dry. Solder 10-cm lengths of 20-gauge copper wire to two of the
pennies, overlapping the wire and penny about 2 to 3 mm from the edge. Solder the free end of one
of the copper wires to a 5 x 100 mm strip of zinc metal, as shown in Figure 5. Clean any rosin off the
soldered joints with steel wool, and rinse with water.



74



CHAPTER 12. FROM CELLS AND ELECTRODES TO GOLDEN PENNIES



B




1 M NaOH
solution



30 mesh
granular
zinc



Figure 12.1



Figure 5

1. When soldering, place the penny on a ceramic fiber square, with the end of a copper wire overlapping
the penny 2-3 mm. Ask a partner to apply pressure to the wire to hold it in place while you are
soldering the joint -use a rubber stopper as before to do this.

2. Four arrangements of copper pennies in the golden penny experiment.

The bottom of the beaker is half-covered with 30-mesh zinc metal. (A) Two pennies are lying on top of the
30-mesh zinc.

(B) Two pennies are lying on the bottom of the beaker but not in contact with the 30-mesh zinc. (C)
A penny soldered to copper wire is immersed in solution, the other end of the copper wire being soldered
to a strip of zinc metal in contact with 30-mesh zinc on the bottom of the beaker. (D) A penny soldered to
copper wire is immersed in solution. The solution in the beaker is 1 M NaOH.



75

Review the General Soldering Instructions in Part I. Place the freshly tinned tip of the penny next to
the wire angling the iron to get good thermal contact. Don't dab at the joint with the tip of the iron while
soldering.

1. When the solution has warmed, use forceps to place two copper pennies on top of the granular zinc
metal and two pennies in the area that is free of granular zinc (make sure that the pennies on the
uncovered side do not contact even one grain of zinc). Bend a small "foot" on the penny in the beaker
as shown in Figure 5. The "foot" of the zinc strip should rest on the granular size so that both are
in direct contact. The penny should be completely immersed in solution but should not contact any
granular zinc metal on the bottom of the beaker. Finally, hang the last penny (the one with only
copper wire soldered to it) over the edge of the beaker so that the penny is completely immersed in
solution. See Figure 5. HINT- use an empty 400-mL beaker to bend and shape your soldered metal
pieces to match what is pictured in Figure 5. Do this before attempting to put them in the 400-mL
beaker containing your warm NaOH solution. 30min.

2. Leave the pennies in the beaker until some of them turn a silvery color. This may take anywhere from
5 to 30 min., depending on the temperature of the solution. (Some of the pennies will never turn silver
even after waiting an hour or more.)

3. Which pennies turn a silvery color? Is it the three pennies that are in contact with the solution and
with zinc, either directly or through the copper wire? Or is it the three pennies in contact with the
solution but not indirect or indirect contact with zinc metal? 5-10min.

4. Using a pair of forceps, remove the pennies that have turned a uniform silvery color, rinse them with
water, and put the pennies on a hot plate for a few seconds. Watch what happens to the silver-colored
pennies as they heat on the hot plate. Keep the solution warm in the beaker in case you need to repeat
some part of the experiment or try some new experiment, as described below. 5-10min.

5. Further Experiments. Solder another length of copper wire to a shiny clean penny. Actually, you can
use the lone penny soldered to the copper wire from the first part of the experiment, just clean it with
steel wool and deionized water. Connect one lead of a voltmeter to a zinc strip and the other lead to
the copper wire soldered to the penny. Using the same solution you prepared earlier, immerse the zinc
strip and penny in the solution. Is there a voltage difference between the zinc strip and the copper
wire soldered to the copper penny? Which metal is the electron source (the negative terminal of this
electrochemical cell)? 5-10min.

6. Do you think a current flows in the copper wire connecting the zinc strip and copper penny when both
are immersed in the solution? If so, which direction will electrons flow, and what are the anode and
cathode reactions? Put the digital voltmeter into its current measuring mode on its most sensitive
(microampere) scale. Then see if any current is flowing when you connect the meter in series between
the copper wire soldered to the penny and to the zinc strip -both the penny and the zinc strip should be
immersed in the hot 1 M NaOH solution. The series connections should look like this: Penny /copper
wire/(+) ammeter (-) /zinc strip. How large a current flows? . Is the current (charge flow) from penny
to zinc strip or vice versa?

7. In wires the charge carriers are electrons. Current (defined as a flow of positive charge) is opposite to
the flow of electrons. In which direction are electrons flowing: from penny to zinc strip or vice versa?



76 CHAPTER 12. FROM CELLS AND ELECTRODES TO GOLDEN PENNIES



Chapter 13

Amphoteric Aluminum 1



Properties of Aluminum and Its Compounds

13.1 Objectives

• To prepare common alum, KA1 (S04) 2 .I2H2O, from a discarded aluminum beverage can.

• To perform qualitative analysis on alum.

• To investigate the acid-base behavior of aluminum compounds.

Grading

You grade will be determined based on the following:
[U+F0B7] Pre-lab (10%)

[U+F0B7] Lab Report Form (80%)

[U+F0B7] TA evaluation of lab procedure (10%)

Before coming to lab

[U+F0B7] Read the lab instructions

[U+F0B7] Print out the lab instructions and report form







Complete the pre-lab, due at the beginning of the lab. See page 616 in Brown, LeMay and Bursten
for the definition of amphoteric behavior.



13.2 Introduction

Aluminum is the most abundant metal in the earth's surface (7.5% by mass). The abundance of aluminum,
coupled with its attractive combination of physical and chemical properties, accounts for the fact that it is
one of the principal industrial raw materials used by industrialized societies. Production of aluminum from
raw materials is an energy intensive process.

Since the metal is not consumed rapidly by corrosion, the amount of scrap aluminum grows rapidly while
the available supply of raw materials for the manufacture of aluminum decreases. The average predicted
longevity of an aluminum can along the roadside is 100 years.

Environmental problems thus created are typical of those of several different metals. One obvious solution
to the problem is to recycle the used fabricated aluminum into other useful metallic objects or into aluminum
compounds. Aluminum metal can be recovered from scrap by melting the metal and separating it from solids
and volatile impurities. This process uses a large amount of energy. The energy requirement to prepare an
aluminum can from recycling is only 5% of the energy required to produce the can from bauxite ore.



lr This content is available online at <http://cnx.Org/content/ml5790/l.10/>.



77



78 CHAPTER 13. AMPHOTERIC ALUMINUM

This experiment illustrates a chemical recovery process in which waste aluminum is converted chemically
into an aluminum compound, hydrated potassium aluminum sulfate, KA1 (SC>4) 2 .V2H2O, or common alum.
Although alum is an important industrial compound, the method of preparation in this experiment is not the
way alum is obtained for use in industry. Nevertheless, this experiment will illustrate an interesting example
of the reduction of environmental waste. "Alum" is a generic term that describes hydrated double salts of
certain metals having the generalized formula, [MM' (S04) 2 .I2H2O}, in which M (univalent) is commonly
Na+, K+, NH 4+ , or Rb + and M' (trivalent) is commonly Al 3+ , Ga 3+ , V 3+ , Cr 3+ , Mn 3+ , Fe 3+ , or Co 3+ .
True alums crystallize in well-defined octahedral shapes and many are beautifully colored, particularly those
containing d-block transition metals. The ancient Egyptians, Greeks and Romans used alum as a mordant
in dyeing cloth. A mordant contains metal ions that bind dyes to the fabric. Presently alum is used to
harden photographic film, to prepare pickles, as a mordant, and for other purposes.

13.3 Experimental Procedure

Wear safety goggles. Wear gloves when using concentrated acid and base.

13.4 Part 1: Preparation of Alum/Quantative Analysis (Calculation

of % Yield)

1. Cut a 5cm x 5cm square piece of aluminum from a scrap aluminum can and sand the paint off using

steel wool.

2. Cut this piece into smaller pieces (about 0.5 cm long) and weigh out ~lg. Record the weight to
3 decimal places.

3. Place the aluminum in a 400-mL beaker and add 50 mL of 4 M KOH.

4. Place beaker on the heating mantle and heat gently.

CAUTION - H 2 GAS (VERY EXPLOSIVE) IS PRODUCED. ENSURE THAT YOU ONLY HEAT THE
BOTTOM OF THE BEAKER AND DO NOT LET THE FLAMES GET NEAR THE TOP.

5. When the bubbles have stopped, remove from the heat, i.e., hydrogen is no longer evolved.

6. Vacuum filter the solution and save the filtrate, you may toss out the solid. (Ask your TA if you
need instruction in performing a vacuum filtration.)

7. Carefully rinse your graduated cylinder, pour out 25-30 mL of 9 M H2SO4 and then add slowly to
the filtrate.

CAUTION - # 2 S0 4 IS A STRONG ACID AND DEHYDRATOR. SEE TA IMMEDIATELY IF YOU
SPILL ANY!

8. Heat gently while stirring until the solution becomes clear. Boil the solution down to a volume of
about 45 mL.

7. While the solution is boiling, prepare an ice bath by filling a bowl half-way with ice and then
adding water until the bowl is three-quarters full.

8. After the "boiling off", let the beaker cool to room temperature and then place it in the ice bath.
Crystals of alum should form. Allow to cool for 15 minutes.

9. After 15 minutes, vacuum filter the product and wash with 10 mL of ethanol.

10. Allow to dry for a few minutes. Dry completely in the microwave at less than 50% power in 10 sec
intervals until the weight no longer changes.

11. Calculate the percent yield and describe the appearance of the crystals.

13.5 Part 2: Qualitative Analysis of Alum

1. Use a straw spatula to transfer a few of the alum crystals (about 5 mg) to a watch glass. Add 3 drops
of water to the crystals. Stir gently until the crystals dissolve.



79

2. Use a small piece of indicator paper to see whether the solution is acidic, basic, or neutral.

3. Now add 1 drop of 0.5 M BaCl2 (barium chloride) to the solution. Record your observations.

4. A really good test for the presence of potassium is a flame test. Using the hot grips, hold a stainless
spatula in the flame of a Bunsen burner to volatilize impurities from the spatula.

5. When one end of the spatula is red hot, remove it, and quickly touch it to a small cluster of
crystals. Several should stick.

6. Slowly bring the spatula (plus crystals) toward the flame and watch carefully. Hold the crystals
in the flame for at least 5 seconds (until the solid glows). Record your observations.

7. Remove the spatula and place on a non-asbestos mat.

13.6 Part 3: Acid-Base Properties of Aluminum Compounds

Clean 3 test-tubes and label them.

Tube 1:

Place 10 drops of 1M aluminum nitrate solution and 2 drops of 6.0 M sodium hydroxide solution, mix
well and record your observations on your lab report Then add more sodium hydroxide solution to tube 1
until a reaction is observed (around 7 drops). Mix well. Record your observations.

Tube 2

Place 10 drops of 1M aluminum nitrate solution and 2 drops of 6.0 M sodium hydroxide solution, mix well
and record your observations on your lab report Then add 4 mL of 6M HC1 and record your observations.

Tube 3

Add 3 mL of 6M ammonium hydroxide, i.e aqueous ammonia to your alum and record your observations.



80 CHAPTER 13. AMPHOTERIC ALUMINUM



Chapter 14

Crystal Violet Kinetics 1

14.1 Crystal Violet Kinetics
14.1.1 Objective



• To study the reaction rate of crystal violet with NaOH using the Lab Works Interface colorimeter.

• To determine the reaction order with respect to each of the reactants.



To calculate the room temperature rate constant for the reaction.



14.1.2 Grading

• Pre-Lab 10%

• Lab Report Form 80%

• TA Points 10%



14.1.3 Background Information

14.1.3.1 Reaction Chemistry

Chemical kinetics is the study of reaction rates. In this experiment, the kinetics of the reaction between
crystal violet and NaOH will be studied. The MicroLab Interface colorimeter will be used to monitor the
change in concentration of crystal violet as a function of time. The reactant and product structures and
the reaction stoichiometry are shown in Figure 1.

All of the reactants and products shown in Figure 1 are colorless except for crystal violet which has an
intense violet color. Thus, during the course of the reaction, the reaction mixture color becomes less and
less intense, ultimately becoming colorless when all of the crystal violet has reacted with the OH~ .



1 This content is available online at <http://cnx.Org/content/ml5811/l.3/>.



81



82



CHAPTER 14. CRYSTAL VIOLET KINETICS



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— S\FH,I ;



14.1.3.2 Figure 1. Net ionic reaction between Crystal Violet and NaOH

The color that crystal violet exhibits is due to the extensive system of alternating single and double bonds,
which extends over all three benzene rings and the central carbon atom. This alternation of double and
single bonding is termed conjugation, and molecules which have extensive conjugation are usually highly
colored. The color is due to continuous movement of electrons between single and double bonds. When
crystal violet reacts with a base (—OH), the conjugation is disrupted and the color is lost. Note that in the
reaction product, the three rings are no longer in conjugation with one another, and hence the material is
colorless. This reaction has kinetics slow enough that the change in color can be observed over time as the
molecules are being changed. In today's experiment, you will trace the loss of conjugation in the crystal
violet structure by using colorimetry.



14.1.3.3 Kinetic rate laws

The rate of the reaction of crystal violet with NaOH is given by the generalized rate expression

Rate = fc[OH - f[CV] y (l)

In Equation (1), k is the rate constant for the reaction, CV is an abbreviation for crystal violet,
C2$HzoN + , x is the order of reaction with respect to OH - , and y is the order of reaction with respect
to CV. The values of x and y will be determined experimentally. Possible values are 0, 1, or 2 (zeroeth order,
first order or second order).

In the experiment you will perform, the [ OH - ] will always be much greater than [CV]. Thus the change
in [ OH - ] has a negligible effect on the initial [ OH - ]. For this reason, [OH - ] can be treated as a constant
and Equation (1) can be rewritten

Rate = k' [CVf (2)

where k' = k [OH]*, k' is termed a pseudo rate constant.

The integrated form of the rate law depends on the order of reaction with respect to the concentration
of CV. The integrated rate laws for y = 0, 1, and 2 are given in Equations 3 through 5.

[C V] t = -k't+ [C V] (zero order) (3)

In [CV] t = -jfc't + In [CV] (first order) (4)

1 = k't+ 1 (second order) (5)

[CV] t [CV]

In Equations 3-5, [CV] is the concentration of crystal violet in the reaction mixture at time zero before
any reaction occurs; [CV] t is the concentration at any time during the course of the reaction. Equations 3,
4b, and 5 are each an equation of a straight line. If a plot of [CV] t versus time is linear, y = and the



83

reaction is zero order in CV. Similarly, a linear plot of In [CV] t versus time indicates a first order reaction
in CV, and a linear plot of 1 / [CV] t versus time indicates second order behavior. In every case, the slope
of the resulting straight line would be the pseudo rate constant, k'. All three of these plots will be made to
determine the actual value of y and the value of k'.

In order to do the graphing just described, we will need to have data showing how the concentration of
CV changes with time. This data will be obtained using the MicroLAB colorimeter at the 590 nm wavelength
and the Kinetics program. The light from the LED will pass through the solution containing CV and NaOH
and then fall on the system photocell. The photocell circuit will then produce a current in microamps (I)
which is proportional to the light intensity striking the photocell surface. This current is divided by the
current obtained with the blank, and the result is termed Transmittance.

Solutions of crystal violet obey Beer's Law. Thus, the relationship between the observed current and the
concentration of CV is given by

A t = -log (I t /I ) = ebc(6)

In Equation (6), A t is the reaction solution absorbance at any time t; I is the photocell current observed
for pure water (the blank value); It is the current observed for the CV reaction mixture at time t, e is the
molar absorptivity of crystal violet; b is the cell path length (2.54 cm for the MicroLAB colorimeter) vial;
and c is the molar concentration of CV at time t, [CV] t . Since e and b are constants at a given wavelength,
it should be clear that the absorbance, A t , is directly proportional to the concentration of CV at any time
during the reaction and can be used in place of [CV] t in preparing the graphs described above.

14.1.3.4 Experimental Procedure

14.1.3.5 Part I.

Measurements

Open the MicroLAB program by selecting it, then click on the Kinetics Experiment Icon. Enter the
experiment filename (CV then your name) when requested, then click OK. Data will come from the interface.
Be sure the MicroLAB interface is connected to the computer and turned on. From that point follow the
procedures listed below.

1. Before beginning kinetic measurements, the current reading for pure H2O, I , will be obtained. Fill
a clean, rinsed colorimeter vial about � full with distilled H2O, dry the outside of the cell thoroughly
with a KimWipe being careful to remove any finger prints, insert the cell into the colorimeter and place
the black cap over the cell. Note carefully and mark the positioning of the cell for future reference.

2. Click on the Blank button. The program will now measure I for each of the 10 wavelengths, divide
each by itself and multiply by 100 to get the 100 % transmittance value.

3. Empty the vial and dry it thoroughly inside and out.

4. Set the Time Interval to 5 seconds, and the Number of Points to 300.

5. Using the buret provided, dispense exactly 9.00 mL of 1.50 x 10~ 5 M crystal violet solution into a
clean, dry colorimeter vial

6. Using the calibrated plastic dropper provided, add 1.0 mL of 0.050 M NaOH to the CV solution as
rapidly as possible without splashing. Cap the vial, rotate it twice to mix the CV/NaOH, place the cell
in the colorimeter in exactly the same manner as was used for the blank and cap the colorimeter. All
of the operations in this step should be completed as quickly as possible so that the first measurement
will be made as close to the beginning of the reaction as possible.

7. As soon as the vial is in place and capped, press the Start button, the program will take readings at 5
second intervals for a period of 60 minutes and then automatically stop. If there is a need to stop data
collection prior to the end of 60 minutes, click on the Stop button and the program will terminate. It
300 points is insufficient, increase the number of points.

8. When the reaction is completed, save the file with the name CV.kin.XM.DH, where CV.kin defines
the type of data, XM indicates the NaOH concentration (0.10 or 0.05, etc), and DH is the student's
initials.



84 CHAPTER 14. CRYSTAL VIOLET KINETICS

14.1.3.6 Data Analysis

1. Retrieve your MicroLAB data file under the name you saved it.

Click on the Linear - Zero Order tab at the bottom of the graph. If the reaction you just did was zero order
on the concentration of crystal violet, this will show a horizontal straight line. Print this screen as follows:

• Press Ctrl-Print Screen to capture the screen image.

• Open Wordpad by clicking Start > Programs > Accessories > Wordpad.

• Press Ctrl-V to paste the screen image into Wordpad.

• Press Ctrl-P to print the item.

• Repeat step (2), clicking on the Logarithmic - First Order tab. A linear graph in this instance would
indicate first order dependence on the concentration of crystal violet. Print this screen also.

• Finally, click on the Inverse - Second Order tab to determine if the reaction is second order in crystal
violet. Also print this screen

• With the linear plot you have identified the value of y, i.e., the order of the reaction with respect to
CV. Record the value of y in your lab book. The slope of the straight line at the top of the linear
regression plot is the best value of k'. Record this value with proper units and to the correct number
of significant figures in your lab book. Attach your graphs to your lab book.



14.1.3.7 Part II.

14.1.3.8 Measurements

Recall that the k' just obtained is a pseudo rate constant, whose value depends upon the concentration of
OH - , i.e. k' = k [OH - ] . In this part of the experiment, the value of x will be determined as well as the
value of the true rate constant, k.

In Part I of the experiment, 9.00 mL of 1.50 x 10 -5 M crystal violet and 1.0 mL of 0.050 M NaOH were
combined to form the reaction mixture. A second kinetic run will now be made in exactly the same way
except that the concentration of NaOH will be doubled to 0.10 M.

1. Repeat each of the six experimental steps 4 through 8 using 1.0 mL of 0.10 M NaOH in place of 0.050
M NaOH.



14.1.3.9 Data Analysis

1. Repeat data treatment steps 1 through 6 above, and again record the value of k' on your data sheet.

2. From the ratio of the two k' values to one another, determine the order of reaction with respect to
OH - (the value of x). Clearly indicate your reasoning in evaluating x.

Note: The value of x should be an integer. If your value is not an integer, it is probably due to experimental
error (probably in measuring and adding the NaOH solutions). If necessary, round your value to the nearest
integer.

1. Calculate the value of the true rate constant k using each of the k' values. In the calculations, the
concentrations of OH - will have to be adjusted to account for the dilutions which occurred when the
NaOH and crystal violet solutions were mixed. Finally, average the two k values obtained. Again, be
sure to watch significant figures and use proper units.

2. Using the linear plot from your first kinetics experiment, calculate the value of the molar absorptivity,
e, for crystal violet under these experimental conditions. Include units in your answer. The colorimeter
vial is 2.54 cm thick. (Hint: The intercept of your linear plot is important.).



INDEX



85



Index of Keywords and Terms

Keywords are listed by the section with that keyword (page numbers are in parentheses). Keywords
do not necessarily appear in the text of the page. They are merely associated with that section. Ex.
apples, � 1.1 (1) Terms are referenced by the page they appear on. Ex. apples, 1



A antacids, � 5(19)

B Back-titration, � 5(19)

blood alcohol analysis, � 7(29)

C Complex, � 4(15)
Constant, � 4(15)
Copper, � 11(59)

D Depression, � 2(5)

E Equilibrium, � 4(15)

F fingerprints, � 7(29)
Formation, � 4(15)
Freezing, � 2(5)

G Gases, � 1(1)
Golden, � 11(59)



Ion, � 4(15)

K Kinetic, � 1(1)
Kitchen, � 10(53)

M Mass, � 11(59)
Molar, � 11(59)
Molecular, � 1(1)

N Nanorust, � 10(53)

O of, � 1(1), � 10(53)

P Pennies, � 11(59)

Playing with different polymers, � 3(11)
Point, � 2(5)

S spot tests, � 7(29)
synthesis, � 10(53)



I inks, � 7(29)



T Theory, � 1(1)



86 ATTRIBUTIONS

Attributions

Collection: General Chemistry Lab Spring

Edited by: Mary McHale

URL: http://cnx.org/content/coll0506/l-56/

License: http://creativecommons.Org/licenses/by/2.0/

Module: "Practical Examples of the Gas Laws"

By: Mary McHale

URL: http://cnx.org/content/ml9475/l-4/

Pages: 1-4

Copyright: Mary McHale

License: http://creativecommons.Org/licenses/by/2.0/

Module: "Colligative Properties and Ice Cream"

By: Mary McHale

URL: http://cnx.org/content/ml9545/l-4/

Pages: 5-9

Copyright: Mary McHale

License: http://creativecommons.Org/licenses/by/2.0/

Module: "Pervasive Polymers"

By: Mary McHale

URL: http://cnx.org/content/ml9573/l-4/

Pages: 11-13

Copyright: Mary McHale

License: http://creativecommons.org/licenses/by/2-0/

Module: "Determine the Value of an Equilibrium Constant by Complex Ion Formation"

By: Mary McHale

URL: http://cnx.org/content/ml9605/l-3/

Pages: 15-18

Copyright: Mary McHale

License: http://creativecommons.org/licenses/by/3-0/

Module: "indigestion? Which is the Best Commercial Antacid?"

By: Mary McHale

URL: http://cnx.org/content/ml9622/l-5/

Pages: 19-24

Copyright: Mary McHale

License: http://creativecommons.org/licenses/by/3-0/

Module: "Acid and Bases to Buffers"

By: Mary McHale

URL: http://cnx.org/content/ml5809/l-13/

Pages: 25-28

Copyright: Mary McHale

License: http://creativecommons.org/licenses/by/2-0/



ATTRIBUTIONS 87

Module: "Forensics"

By: Mary McHale

URL: http://cnx.org/content/ml9954/l-7/

Pages: 29-35

Copyright: Mary McHale

License: http://creativecommons.Org/licenses/by/3.0/

Module: "Enzyme Kinetics/Mr Potato"

Used here as: "The Curious Case of Catalase"

By: Mary McHale

URL: http://cnx.org/content/ml5868/l-3/

Pages: 37-42

Copyright: Mary McHale

License: http://creativecommons.Org/licenses/by/2.0/

Module: "Organic Reactions"

By: Mary McHale

URL: http://cnx.org/content/ml5483/l-3/

Pages: 43-51

Copyright: Mary McHale

License: http://creativecommons.Org/licenses/by/2.0/

Module: "Kitchen Synthesis of Nanorust"

By: Mary McHale

URL: http://cnx.org/content/m20813/l-3/

Pages: 53-58

Copyright: Mary McHale

License: http://creativecommons.org/licenses/by/3-0/

Module: "Electrochemistry and Alchemy"

By: Mary McHale

URL: http://cnx.org/content/m21301/l-4/

Pages: 59-67

Copyright: Mary McHale

License: http://creativecommons.org/licenses/by/3-0/

Module: "From Cells and Electrodes to Golden Pennies"

By: Mary McHale

URL: http://cnx.org/content/ml5879/l-7/

Pages: 69-75

Copyright: Mary McHale

License: http://creativecommons.org/licenses/by/2-0/

Module: "Amphoteric Aluminum"

By: Mary McHale

URL: http://cnx.org/content/ml5790/l-10/

Pages: 77-79

Copyright: Mary McHale

License: http://creativecommons.org/licenses/by/2-0/



88 ATTRIBUTIONS

Module: "Crystal Violet Kinetics"

By: Mary McHale

URL: http://cnx.Org/content/ml5811/l.3/

Pages: 81-84

Copyright: Mary McHale

License: http://creativecommons.Org/licenses/by/2.0/



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