Microscale Gas Chemistry

Experiments with Oxygen


Carl Scheele
Joseph Priestley
Antoine Lavoisier
Carl Scheele
1772

Joseph Priestley
1774

Antoine Lavoisier
1789

THESE THREE MEN, Carl Scheele (Sweden), Joseph Priestley (England), and Antoine Lavoisier (France) all claimed credit for the discovery of the element that we now call oxygen.  Carl Scheele discovered fire air [oxygen] sometime before 1773.  He produced the gas several ways.  In one method, he reacted (using modern names) nitric acid with potash (KOH and/or K2CO3) which formed KNO3.  Distilling the residue with sulfuric acid produced both NO2 and O2.  The former was absorbed by limewater (saturated Ca(OH)2), leaving fire air.  He also obtained fire air from strongly heating HgO and MnO2 and by heating silver carbonate or mercuric carbonate and then absorbing the CO2 by alkali (KOH): 

AgCO3(s)  arrow Ag(s) + CO2(g) + O2(g)

    On August 1, 1774 Joseph Priestley first prepared oxygen by directing the sun's light with a 12-inch diameter burning lens onto a sample of red mercurius calcinatus per se (now HgO).  Thus, Priestley independently had discovered oxygen which he called dephlogisticated air.  His explanation of the reaction using was:  

mercurius calcinatus per se + heat yields quicksilver + dephlogisticated air

Today, we would describe the same reaction as follows:

 HgO(s)  arrow Hg(l) + O2(g)


    Priestley sampled his dephlogisticated air and wrote:

"[My] breast felt particularly light and easy for some time afterwards...... Who can tell, but that, in time, this pure air may become a fashionable article of luxury.  Hitherto only two mice and myself have had the privilege of breathing it." (A Short History of Chemistry, J. R. Partington, 3rd edition, (1957))


    Priestley shared his discovery with Antoine Lavoisier later that same year and also described the discovery in published form shortly thereafter.  Carl Scheele’s account of his discovery did not appear in print until 1777, by which time the scientific community had given credit for the discovery to Priestley.

    Antoine Lavoisier’s claim to the discovery of oxygen is based on the fact that he established that oxygen was an element.  He proposed the name oxygen and it became the cornerstone of his oxygen theory which debunked the prevailing phlogiston theory that had governed the minds of scientists for nearly one century. To spread his ideas and the Oxygen Theory, Lavoisier published Traité élémentaire de chemie in 1789.  In his book, Lavoisier named a total of 33 elements, most of which are still in use today.  It has been said that the book would be recognizable to a student of chemistry as it reads "like a rather old edition of a modern textbook."
(A Short History of Chemistry, J. R. Partington, 3rd edition, (1957))

    Oxygen, O2, is a colorless, odorless, and tasteless gases that has a very low solubility  in water.  Oxygen can be condensed to a pale-blue liquid by cooling to -183 °C (90 K) at a pressure of one atmosphere.  Oxygen’s melting point is -218.4 ºC (55 K).

    Oxygen is the second largest component of the Earth's atmosphere (21%).  It occurs as O2 as well as the allotrope O3, called ozone, in the atmosphere.  Oxygen represents 89% be mass of water molecules so that the Earth's water supplies are largely oxygen.  Much of the Earth's lithosphere (rocks, solid parts of the crust) is composed of silicates and other oxides.  Taken together, over 46% of the mass of the lithosphere is oxygen.  Photosynthesis accounts for virtually all of the oxygen present in the Earth's atmosphere.  Water and carbon dioxide are converted by chloroplasts to oxygen and plant carbohydrates, "CH
2O" respectively:

H2O(l) + CO2(g)  arrow O2(g) + "CH2O"


The reaction is endothermic and occurs only because of the energy supplied by the sun. 

    The density of O2 is 1.308 g/L at 25 
°C and 1 atm — 10% greater than that of air. At 20 °C oxygen dissolves in water to the extent of 30.8 cm3 per liter.  It is even more soluble in some non-aqueous solvents.

    Oxygen has numerous important uses in society. Oxygen is obtained by the fractional distillation of air. The single largest use of oxygen is in the manufacture of steel.  Oxygen is blasted under high pressure onto molten, impure iron to burn out impurities.  Oxygen has a variety of other uses including respiratory oxygen in medical facilities, welding, sewage treatment, production of TiO
2 (used to make paints opaque and white), and ethylene oxide to name a few.

Suitability
    The following experiments are included in this chapter. 

Part 1. Experiments with Oxygen
Experiment 1. Traditional test for oxygen
Experiment 2. Oxygen supports combustion
Experiment 3. Dynamite soap
Experiment 4. Hydrogen-oxygen rockets


Part 2. Demonstrations and Advanced Experiments with Oxygen
Experiment 5. Steel wool burns in oxygen
Experiment 6. The Blue Bottle experiment
Experiment 7. Oxygen makes the flame hotter
Experiment 8. Mini-sponge shooter
Experiment 9. Chemiluminescence


    The first four experiments are suitable as laboratory experiments for a wide variety of grade levels from middle school up through university-level.  Due caution is required with hydrogen/oxygen mixtures and students must not be allowed to try different experiments without the express consent of the instructor.  

    In most high school settings, the first four experiments can be used very early in the year — at about the time that chemical formulas and reactions are being introduced.  As a laboratory activity, these experiments are appropriate when discussing chemical compounds, chemical formulas, and chemical reactions.  Experiments 1 and 2 demonstrate oxygen’s ability to support and enhance combustion.  The chemical reactions in Experiments 3 and 4 are very simple: hydrogen and oxygen produce water.  Mole calculations may be done, but are optional.  The law of combining volumes can be determined experimentally by playing with the H2/O
2 mixture ratio in Experiments 3 and 4.  Experiment 4 requires a piezoelectric sparking device that can be constructed from a piezoelectric lighter.

    Experiments 5 - 9 are suited for use as classroom demonstrations.  

    Experiment 5 involves glowing hot steel wool and is quite impressive as a demonstration, but steel wool is messy and hot metal is dangerous. This experiment can be used to demonstrate that metals can burn and undergo rapid oxidation.  A discussion of corrosion and rust formation can help students compare rapid oxidation with slow oxidation of metals.  The experiment can be used when combustion reactions are discussed or oxidation as a type of chemical reaction is introduced.  

    The Blue Bottle experiment (Experiment 6) features a clear solution that turns blue upon shaking.  This demonstration can be used repeatedly for over 24 hours!  It can be used when discussing oxygen’s solubility in water, including the depletion of dissolved oxygen.  It makes a good demonstration when discussing natural waters, biological oxygen demand, water stagnation, etc.  In a traditional chemistry course, we use it to as a demonstration of LeChâltelier’s principle.

    Experiment 7, Oxygen makes the flame hotter, is used as a demonstration because the ensuing discussion about the nature of a flame is best done as a teacher-led discussion.

    Experiment 8, Mini-sponge shooter, in which a sponge projectile is shot up to 10 m requires a special syringe in which a hole has been drilled through the barrel.  Use this demonstration when discussing explosions and explosive mixtures.  

    Experiment 9, Chemiluminescence uses a relatively toxic chemical, dimethylsulfoxide (dmso), that is inappropriate for handling by students.  Use this demonstration when discussing types of energy — heat, light and electrical.  

Background skills required
    Students should be able to:

generate a gas using the In-Syringe method
measure quantities of liquid reagents
use a balance
light matches and work with small flames in a responsible and appropriate manner for a science laboratory

Time required

    Students should be able to generate hydrogen and perform the four experiments in Part 1 in a single 45 minute laboratory period.

Preparation of oxygen in a “gas bag
    Large samples of O2(g) can be prepared conveniently in 1 L food storage bag.

Student Instructions
    For classroom use by teachers, one copy per student in the class may be made free of charge and without further permission.  Student instructions and questions only (without teaching tips, suitability information, etc.) can be downloaded free of charge as a Microsoft Word document from the website. Download now.

Answers to the questions, lists of materials and chemicals, and additional reference information.
    This page is fairly similar in content to Chapter 3 in our book Microscale Gas Chemistry.   Our 500-page book includes all of the information included at this website and much more!  Answers to all of the questions, chapter-by-chapter lists of the equipment and chemicals needed to conduct the experiments as classroom demonstrations or laboratory activities for the entire class, construction instructions for various pieces of equipment, information for the preparation of stock solutions, ordering information, and a complete index — are all available in the book, but not at the website.  The book,
Microscale Gas Chemistry, can be ordered from Educational Innovations (Part Number BK-590) from  their website.



 
PREPARATION OF OXYGEN


General Safety Precautions
    Always wear safety glasses.  Gases in syringes may be under pressure and could spray liquid chemicals.  Follow the instructions and only use the quantities suggested.  CAUTION! Hydrogen forms explosive mixtures with air.

Toxicity
    Oxygen is non-toxic in normal quantities.  Pure oxygen can be toxic if inhaled in large quantities as the pure gas, but this is not a concern with these experiments.  Do not intentionally inhale oxygen samples produced in these experiments.


Equipment (Vendors and Part Numbers)
    Microscale Gas Chemistry Kit:

two 60 mL plastic syringes with a LuerLOK fitting
two Latex LuerLOK syringe caps
two plastic vial caps
one 15 cm length of Latex tubing
one 3 cm length of Latex tubing
one small bottle of silicone oil
one plastic pipet
one clear plastic beverage cup (250 mL/9 oz)
one small plastic weighing dish
one small test tube (12 x 100 mm)
one medium test tube (18 x 150 mm)
one birthday candle



Chemicals (needed for each syringe full of hydrogen produced)

0.05 g solid KI powder
5 mL 6% H2O
2(aq)

The production of O2 is slow and it typically takes a minute or more to fill a syringe.  Assist the plunger in its outward movement.  To speed up the reaction, hold the plunger so that the contents inside the syringe are under reduced pressure, and while doing so, tap or shake the syringe.  This process drives oxygen bubbles out of the solution.  Potassium iodide is the catalyst in the reaction:

2 H2O2(aq) arrow 2 H2O(l) + O2(g)


The actual mechanism has two steps: 

Step 1.  H2O2(aq) + I-(aq) arrow H2O(l) + IO-(aq)

Step 2. IO-(aq) + H2O2(aq) arrow   I-(aq) + H2O(l) + O2(g)


Oftentimes the solution takes on a yellow color due to I3
-(aq) that results form a competing side reaction: 

2 H+(aq) + IO-(aq) + 2 I-(aq) arrow   I3-(aq) + H2O(l)


Generating oxygen gas samples

    Samples of oxygen are generated by the In-Syringe Method.  A summary of these steps is provided here:

1. Wear your safety glasses!

2. Lubricate the seal.  

    Lubricate the black rubber seal of the plunger with silicone oil. 
     
3. Measure out 0.05 g potassium iodide, KI. 
 
    Place the KI(s) directly into the vial cap.

measure out solid reagent      

4. Fill the syringe barrel with water.
    Fill the barrel completely with water.  Place your finger over the hole to form a seal.    

filling syringe with water  

5. Float the vial cap
    Float the vial cap containing the solid reagent on the water surface.   

float the vial cap    

6. Lower the cap by flotation
    Release the seal made by finger to lower the cap into the syringe barrel without spilling its contents.    

lower the vial cap by flotation  

7. Install the plunger
    Install the plunger while maintaining the syringe in a vertical position.  The plunger has a plastic “rib” near the rubber seal that snaps past the “catch” — a small ridge just inside the mouth of the syringe.  Usually it takes a firm push to move the rib past the catch.  After that, the plunger should move smoothly.

insert plunger

     
8. Draw 5 mL 6% H2O2(aq) into syringe
    Pour the 6% H2O2(aq) into a small weighing dish.  Draw 3 – 5 mL of the solution into the syringe. 

draw up agueous reagent      

9. Install syringe cap
    Push the syringe into the syringe cap. 

push syringe into cap      

10. Generate the gas
    Shake the device up and down in order to mix the reagents.  Gently help the plunger move up the barrel.   

shake reagents to generate gas     

11. Remove cap to stop gas collection
    Remove the syringe cap with the syringe held “cap-up” as shown.  Assume contents are under positive pressure. 

remove the syringe cap

12. Discharge reagents
    Discharge the liquid reagent into the plastic cup.  Immediately cap the syringe to prevent loss of gas.   

drain spent reagents     



Wash away contaminants
    Oxygen-filled syringes must be washed in order to remove traces of unwanted chemicals from the inside surfaces of the syringe before the gases can be used in experiments.  Follow the procedure summarized here.

1. Remove the syringe cap,
2. draw 5 mL water into the syringe,
3. cap the syringe,
4. shake syringe to wash inside surfaces,
5. remove cap,
6. discharge water only, and finally
7. recap the syringe.
8. Repeat?

Repeat these Washing Steps if necessary.
(All traces of the reactants should be washed away.)


Disposal of oxygen samples
    Unwanted oxygen samples can be safely discharged into the room.

Teaching tips

1. The generation of oxygen is a slow reaction and students should be warned to be patient.  Show them how to speed up the process by tapping the syringe while contents are under reduced pressure.

2. CAUTION!  The amount of H2O2(aq) used (5 mL) is capable of generating more than 60 mL O2(g) and syringes left unattended will eventually “pop” their plungers.  This must be avoided because it sprays reagents that easily stain clothing and skin.

3. Using 10 mL 3% H2O
2(aq) also works, although it is slower .

4. Aqueous sodium bisulfite (1 M) can be used to remove iodine stains from clothing.

Introductory Questions
1. What is the formula for oxygen gas?

2. What are the formulas for each of the following: (a) potassium iodide; (b) hydrogen peroxide; and (c) water?

3. Did the reaction speed up when you tapped the syringe while holding the plunger slightly outward?  Why does this work?  What happens when you tap a bottle of carbonated beverage?  Is this somehow similar?

Questions
4. What would happen if some of the potassium iodide were to spill out of the cap before you were able to draw up the hydrogen peroxide?  Specifically, what would happen when you tried to draw up the hydrogen peroxide?

5. Write the balanced chemical equation for the reaction occurring inside the syringe.

6. What is a catalyst? 

Advanced Questions
7. What mass of hydrogen peroxide is present in 5 mL of 6% H2O2(aq)?  Assume the density of the solution is 1.00 g/mL.  Convert this mass to moles. 

8. Using the ideal gas law and your answer to the previous question, what volume of gas is predicted? (Assume the temperature is 25 ºC and standard pressure)


PART 1. OXYGEN EXPERIMENTS FOR STUDENTS


EXPERIMENT 1. TRADITIONAL TEST FOR OXYGEN

transfer oxygen to test tube Equipment

Microscale Gas Chemistry Kit
match or lighter
wooden splint


Chemicals

O2(g), 30 mL
limewater, 2 mL     
Suitability
    middle school lab, high school lab, university lab, and classroom demonstration


Applications, Topics, Purpose
    combustion reactions, chemical properties of gases, oxygen supports combustion, characterization of gases.  To demonstrate the traditional test for oxygen and to repeat a classical experiment of Joseph Priestley. 

Instructions
    Transfer 30 mL O2(g) to the kit’s larger test tube (18 x 150 mm) using the long piece of tubing so that the gas can be discharged near the bottom in order to displace the air.  Ignite and blow out the wooden splint and immediately plunge the burning splint into an oxygen-filled test tube.  The splint will re-ignite.  Hold the splint inside the test tube until it goes out (just a few seconds).  Add about 2 - 3 mL of limewater to the test tube and shake the test tube to mix gas and limewater.  Note the results. 

Teaching tips

1. The larger test tube holds approximately 30 – 35 mL gas  A test tube larger than the two in the kit (e.g. 25 x 200 mm) works well, too.  Discharge all of the gas in the syringe into the test tube.

2. Joseph Priestley (1774) noted that a candle burns with greater brightness in oxygen.


Introductory Questions

1. Did the glowing splint re-ignite with oxygen?  What would happen if you were to discharge carbon dioxide onto the glowing splint instead?

2. Why does wood burn with a flame sometimes and just glow red other times?  Why does blowing gently on a glowing piece of wood often cause it to burst into flames?

3. Suppose you were given three test tubes and only one of them contained oxygen.  What experiment would you use to determine which test tube contained oxygen?


Questions

4. As a continuation to Question 3, suppose that the other two test tubes contained carbon dioxide and hydrogen.  What experiments would you use to determine the contents of the other two test tubes?

5. What type of chemical reaction is occurring in the experiment?

6. What gas is being produced when the wood splint burns?

7. Describe the results of the limewater test.

Advanced Questions
8. Write the chemical reaction that occurs during the limewater test.

9. When oxygen displaces air during the first part of the experiment, sketch how you perceive the air being displaced.  Is it like pouring water through oil?

EXPERIMENT 2. OXYGEN SUPPORTS COMBUSTION

expt 2 figure Equipment

Microscale Gas Chemistry Kit
test tube, large (25 x 200 mm)
glass rod
tape
matches


Chemicals

O2(g), 50 mL
Limewater, 2 mL

Suitability
    middle school lab, high school lab, university lab, and classroom demonstration

Applications, Topics, Purpose
    combustion reactions, chemical properties of gases, oxygen supports combustion


Instructions
          Tape the candle (in your kit) to the glass rod as shown.  Transfer the oxygen gas to the large test tube using the long piece of tubing so that the gas can be discharged near the bottom in order to displace the air.  Light the candle and lower it into the test tube.  It should burn brightly for a moment before the oxygen is depleted.  Test the gaseous contents of the test tube with limewater.  Repeat the experiment with a test tube full of air: Either use a fresh test tube or fill the previous test tube with water to displace the gaseous contents of the test tube; drain the test tube and it is now filled with air.

Teaching tips

1. Try the experiment in the dark.

2. Joseph Priestley in the 1770s noted that a candle burns with greater brightness in oxygen.


Introductory Questions

1. Did the candle burn more brightly in the oxygen?

2. Did the candle burn longer in air or oxygen?

3. Do you think that the candle burned hotter in oxygen than in air?

4. Why does the candle go out?



Questions

5. Did anything unusual happen to the candle wax?  What does this indicate?

6. What are the two products of combustion?

7. Discuss and interpret the results of the limewater test.

8. What are the three components needed to sustain a fire?


Advanced Question

9. Calculate the ratio of oxygen in the oxygen-filled test tube to the oxygen in air, given that air is 21% oxygen.


EXPERIMENT 3. DYNAMITE SOAP

Equipment

Microscale Gas Chemistry Kit
match or lighter   

Chemicals

O2(g), 30 mL
H
2(g), 30 mL
3% dish soap, 10 mL


Suitability
    middle school lab, high school lab, university lab, and classroom demonstration

Applications, Topics, Purpose
    combustion reactions, kinetics, stoichiometry of reactions, activation energy, explosive mixtures


Instructions
    1. Connect a short piece of tubing to the oxygen syringe.

connect tubing

    2. Connect the other end of the tubing to the hydrogen syringe:

connect second syringe


    3. Transfer O2 to the H2

transfer gas  


  Form a 1:2 mixture.  Re-label the syringe mixture. Use the short tube to form a mound of bubbles in a small weighing dish as shown at right.  Remove the syringe and ignite the bubbles with a match.  Caution! LOUD! (Unused gas mixture can be saved for the next experiment.)   
mound of bubbles  


Dynamite Soap A
just before the soap in ignited
Dynamite Soap B
photo of hydrogen-oxygen explosion taken in the dark


Teaching tips

1. Students need to tell everyone in the room before igniting their soap bubbles. 

2. The name 'dynamite soap' was coined by “Weird Science,” (Western Chicago Area Chemistry Teachers’ Alliance)

3. You may wish to generate a bag of hydrogen for distribution to the students in order to save time in the laboratory. 



Introductory Questions

1. Describe your observations in writing for someone who did not see this reaction.

2. What is the purpose of the soap solution?

3. Why should extreme care be exercised when working with hydrogen-oxygen mixtures?

4. This experiment is an example of a mini-explosion.  What are some dangers associated with scaling reactions such as this one up in size? 

5. Were you aware that it is illegal to scale up any reaction to a size that can hurt others? 



Questions

6. Would you expect a louder or softer BANG if you used the same volume of air instead of the oxygen you used in the experiment?

7. Why did you need a flame or spark to get the reaction to go?


Advanced Questions

8. What volume of oxygen gas is needed to react with 10 mL hydrogen?

9. Is energy absorbed or released by the reaction in the soap bubbles?

   


EXPERIMENT 4. HYDROGEN-OXYGEN ROCKETS

Equipment

Microscale Gas Chemistry Kit
piezoelectric sparker
wide-stem disposable pipet

   
Chemicals:

O2, 30 mL
H2, 30 mL

Suitability
    middle school lab, high school lab, university lab, and classroom demonstration

Applications, Topics, Purpose
    combustion reactions, kinetics, stoichiometry of reactions, activation energy, explosive mixtures, rocketry, types of chemical reactions


Instructions
     Prepare a mixture of O2 and H2 via the method of syringe-to-syringe transfer as described in the previous experiment.  The procedure for filling and launching rockets is provided in a photo sequence and summarized here:

1. Cut the end off of a pipet leaving at about 2 cm of stem attached to the bulb.
   

2. Completely fill the “rocket” with water.    

3. Slip the water-filled “pipet rocket” into the syringe fitting.  Slowly displace the water with the hydrogen-oxygen mixture.  The water will dribble out onto the bench top or floor.  Holding the syringe and bulb at a 45º angle works best.  

    
4. Slip the pipet rocket over the wire leads of the piezoelectric sparker.  Remember that hydrogen is lighter-than-air!  Never tip the gas-filled rocket open-end up — the gas will escape.    

5. Draw some water into the stem — without this, the rocket will not fly far when “launched”.  The ends of the wire leads of the piezoelectric lighter must be above the water in the gas-filled region of the rocket.    


6. With some water in the stem, launch the rocket by triggering a spark.  DO NOT aim the rocket at anyone!  If the water leaks out of the stem while positioning the rocket over the wires, immediately fill the stem again by holding the wires and rocket in a cup of water and drawing a very small amount of water into the stem.  The rocket should fly up to 10 m!     

rocket ready to launch
rocket ready to launch
at the instant of launching
time exposure photo taken in the dark captures the instant of launching


Teaching tips

1. Use large bulb plastic pipets for the rockets.

2. Construction of the piezoelectric launcher is given in Appendix C.  

3. Award prizes for the rockets traveling the greatest distances.

4. One objective of this experiment is to encourage students to try various mixtures of hydrogen and oxygen.  They will empirically discover that the best mixture is 2 hydrogen : 1 oxygen.

5. This reaction is used by NASA in their Saturn launch rockets.

6. The rocket idea comes from David Ehrenkrantz and John Mauch.  Design for piezoelectric igniter is modified from a model developed by Bob Becker.

7. You may wish to generate a bag of hydrogen for distribution to the students in order to save time in the laboratory. 



Introductory Questions

1. How far, in meters, did your rocket fly?

2. Why did you start by filling the rocket with water ?

3. Which rocket would fly further: (a) a rocket filled with pure hydrogen; (b) a rocket filled with a mixture of hydrogen and air; (c) a rocket filled with a mixture of hydrogen and oxygen?


Questions

4. Why must some water be left in the stem of the rocket in order for the launch to be successful?

5. What is the reaction occurring inside the rocket?

6. Which hydrogen-oxygen rocket is expected to fly farther, a rocket that is mostly filled with oxygen and some hydrogen or one mostly filled with hydrogen and some oxygen?  

7. Would rockets filled with hydrogen and air fly at all?


Advanced Questions

8. What ratio of hydrogen to oxygen is optimal?  Use a balanced chemical equation to answer this question.

9. Real rockets such as the NASA’s Saturn launch vehicles use liquefied gaseous hydrocarbons and liquid oxygen for the rocket’s first stage and liquid hydrogen and liquid oxygen for the second stage.  What sort of design feature would keep liquid hydrogen and liquid oxygen from reacting until they are supposed to?

Clean-up and storage
    At the end of the experiments, wipe excess lubricant off of rubber seal. Clean all syringe parts (including the seal), caps and tubing with soap and water.  Use plenty of soap to remove oil from the rubber seal.  This extends the life of the plunger.  It may be necessary to use a 3 cm diameter brush to clean the inside of the barrel.  Rinse all parts with distilled water.  Be careful with the small parts because they can easily be lost down the drain. Important: Store plunger out of barrel.






PART 2.  DEMONSTRATIONS AND ADVANCED EXPERIMENTS WITH OXYGEN


steel wool burns in oxygen EXPERIMENT 5. STEEL WOOL BURNS IN OXYGEN


  Equipment

Microscale Gas Chemistry Kit
large test tube, 25 x 200 mm
tweezers or metal hemostat or pinch-nosed pliers
lighter or match



Chemicals

O2(g), 60 mL
Steel wool, ‘000’ grade
    


Suitability

    high school lab, university lab, and classroom demonstration


Applications, Topics, Purpose
    corrosion, oxidation, rust, household chemicals, combustion

 

 Instructions
    Transfer oxygen to the large test tube using the long piece of tubing so that the gas can be discharged near the bottom in order to displace the air.  Grip a small ball of '000' grade steel wool with a tweezers/hemostat/pliers.  Light the steel wool with a match and immediately plunge the glowing steel wool into the test tube of oxygen.  The steel wool will burn with a bright light.

Long-term Variant
    Generate O2 as before.  Wash twice.  Place a small ball of '000'-grade steel wool (cleaned with alcohol to remove oil film) in a vial cap and float the cap in a plastic cup (250 mL) half-filled with water.  Remove the plunger from the O
2-filled syringe and place it over the floating cap.  Clamp the syringe in this position.  The reaction is extremely slow.  Within a week the water level (and the cap) will start to rise within the syringe as the iron reacts with the O2.  

Teaching tips

1. Steel wool is iron and it is a reactant with the oxygen.

2. This experiment demonstrates that metals can burn and undergo rapid oxidation.

3. The experiment can be used when discussing types of chemical reactions.  Combustion reactions are all oxidation reactions.


Questions

1 What would happen if you used coarse steel wool instead of the fine steel wool?

2. Assuming that the product is Fe2O3, write the balanced chemical equation for the reaction



the Blue Bottle EXPERIMENT 6. THE BLUE BOTTLE EXPERIMENT

Equipment

Microscale Gas Chemistry Kit

Chemicals

O2(g), 50 mL
KOH(s), one pellet
Dextrose, 0.9 g
methylene blue, 1 drop

 
Suitability
    classroom demonstration


Applications, Topics, Purpose
    LeChatelier's principle, oxidation/reduction, equilibrium, limiting reagents. This demonstration is often used to show how a reaction at equilibrium can be disturbed and will return to equilibrium.



  Instructions
    Prepare a solution of 0.01 g KOH (one pellet) and 0.9 g dextrose in 50 mL water.  Add 1 drop methylene blue.  Draw 25 mL of the solution into the syringe. Draw another 25 mL of the solution into another syringe — this one containing air — for use as a control.  Apply a single shake to the solution whenever the color fades.  Another shake will return the blue color.  This process can be repeated dozens of times and will last for at least one day.  

Teaching tips

1. This demonstration can be scaled up everything by 10 times and performed as a demonstration in a 500 mL flask enriched with one syringe of oxygen.

2. This experiment is based on one published by J. A. Campbell, Journal of Chemical Education, 40, 578 (1963)  A full explanation is given in this reference.

3. Review LeChatelier's principle.  Adding O2(g) to the solution by shaking the syringe results in the position of the equilibrium to shift to the right and the blue color of the solution is seen.  Shaking the syringe causes more O2(g) to dissolve in the water.

4. This makes a good demonstration when discussing natural waters, biological oxygen demand, water stagnation, etc.

5. This reaction destroys the syringe by discoloring it.  Use an old syringe.  


Questions

1. Does the concentration of dissolved oxygen, O2(g), increase or decrease when you shake the syringe?

2. When the solution in the syringe becomes colorless, what must have happened to the amount of O2(aq)?



EXPERIMENT 7. OXYGEN MAKES THE FLAME HOTTER

Equipment

Microscale Gas Chemistry Kit
Bunsen burner, small
glass Pasteur pipet   
Chemicals
O2(g), 60 mL (or a bag of oxygen)

Suitability
    high school lab, university lab, and classroom demonstration

Applications, Topics, Purpose

    Role of oxygen in combustion, torches, temperature of flames

Instructions
    Generate either (a) a syringe filled with O2 using the general method and wash the O2-filled syringe, or (b) a gas bag filled with O2 using the gas bag method

option A  Option B
Option A. Use of an O2–filled syringe    Option B. Use of an O2–filled gas bag

Option A. Connect the O2-filled syringe to a glass pipet with a 15 cm piece of tubing (the tubing forms a snug fit inside the pipet). Option B. Connect the tubing from the gas bag to the glass pipet.  Light a small Bunsen burner. Position the other end of the pipet near (but not in) the burner flame as shown in figures.  Slowly discharge the oxygen into the flame.  A very hot, intensely blue flame will be produced. 

Teaching tips
1. Flames have hot and cool regions.  The inner cone is a cooler endothermic region where gas molecules are broken into fragments.  There is relatively little air in the inner cone.  One can demonstrate the relative coolness of the inner cone of a flame by sticking a pin through a match near the head of a match and then balancing the pin over the opening of the Bunsen burner. Center the match in the barrel of the burner.  Light the burner and the match will remain unlit while the flame of the Bunsen burner burns around it.

2. When oxygen is introduced discharged into the flame, the hot gas immediately reacts in a very exothermic process giving a bright blue flame.

3. Try the reaction in a darkened room.
Questions
1. What are the colors of the various parts of a flame?  What color is associated with the hottest region?

2. What real-world applications take advantage of oxygen’s ability to increase the temperature of a flame?

3. Balance the chemical reaction that takes place when methane burns in oxygen.  How might you confirm that water is a product? How might you confirm that carbon dioxide is a product?

Advanced Question
4. In the inner part of the flame, fuel molecules are broken into fragments due to the abundant energy nearby.  If the fuel were methane, list the two most common fragments expected assuming that in most cases, only one bond breaks per methane. 



EXPERIMENT 8. MINI-SPONGE SHOOTER

Sponge shooter Equipment

Microscale Gas Chemistry Kit
sponge
scissors
electric drill
electricians tape or equivalent
piezoelectric sparker
ring stand and clamp

Chemicals
O2(g), 60 mL
methanol or ethanol  

Suitability
    classroom demonstration

Applications, Topics, Purpose

    Explosive mixtures, rocketry, activation energy, flash point

equipment Instructions
    Begin by cutting a "projectile" from a 1.5 cm thick sponge.  The projectile should be round and disk-shaped with a diameter about 20% larger than that of the syringe.  Moisten the sponge and see if it fits snugly into the syringe barrel as shown in the figure.  Remove the sponge and set it aside.   Use an electric drill to make a 5 mm diameter hole through the syringe barrel near the 10 mL mark (or melt a hole with a hot nail).  Remove any traces of the burr.  Tape over the hole by strapping electrician's tape all the way around the syringe. 

    Insert the plunger.  Transfer O2 and to the syringe.  Next, draw up 2 – 3 mL methanol or ethanol.  Shake the syringe to vaporize some of the alcohol.  Note: The syringe must be at room temperature or slightly above; the vapor pressure of alcohol is much lower at low temperatures. 

    Drain the excess alcohol, remove the plunger and replace it with the moistened sponge.  Clamp the syringe in position.  Remove the electrician's tape and hold a piezoelectric lighter into the hole in the syringe and pull the trigger.  The sponge will fly 5 m or more.

Teaching tips
1. Provide the balanced equation that occurs between methanol (or ethanol) and oxygen.

2. Explain the role of vapor pressure and temperature

3. Flammable liquids have flash points, the temperature at which they explode in air with a spark.  In pure oxygen, the flash point is much lower and explosive mixtures are a much bigger problem.
Questions
1. Why is it necessary for the liquid to evaporate before the mixture is “explosive”?

2. Do liquids or gases react more explosively with air?

3. Would the reaction work better at higher or lower temperatures?  Would it be possible that at some temperature at which the explosion does not take place? 




chemiluminescence EXPERIMENT 9. CHEMILUMINESCENCE

Equipment

Microscale Gas Chemistry Kit

Chemicals
O2(g), 60 mL
KOH pellets, 4 g
Luminol, 0.1 g
DMSO (dimethylsulfoxide), 10 mL
Suitability
    classroom demonstration


Applications, Topics, Purpose
    forms of energy, chemiluminescence

Precaution
    Read the safety information available for DMSO, dimethylsulfoxide! Wear gloves.  Avoid dermal contact. 

Instructions
    Measure out 4 g KOH pellets and 0.1 g luminol and set them aside for use later.  You will also need 10 mL DMSO. Generate oxygen and wash the gas twice.  Remove the plunger, dump out the vial cap, add the solid KOH and luminol through the syringe mouth and reinsert the plunger to its previous mark (about 60 mL.)  Because O2 has a similar density to that of air, O2 loss is not excessive if you work quickly and replace the plunger as soon as possible.  Remove the syringe cap and draw 10 mL of DMSO into the syringe and 2 - 5 mL H
2O.  Darken the room.  Shake the solution and it will emit a blue chemiluminescent glow which will last for a very long time, depending on the amount of shaking.  We have had the system last 48 hours. 

Clean-up:  Do not reuse this syringe.  DMSO must be disposed of properly. One effective method that we have used is to remove the plunger and dump in 30 mL of absorbent cat litter, reinsert the plunger, remove the excess air, and recap.  Place capped syringe containing dmso and cat litter in a sealed food storage bag and place in the trash.  Check state regulations.

Teaching tips
1. Relate the demonstration to light sticks.

Questions
1. What are the three ways in which a chemical reaction can give off energy?

2. What are other examples of chemiluminescence?

3. Why does shaking the syringe make the light brighter?

4. Why does the reaction eventually stop?



Clean-up and Storage.

    At the end of the experiments, wipe excess lubricant off of rubber diaphragm. Clean all syringe parts (including the diaphragm), caps and tubing with soap and water.  Use plenty of soap to remove oil from the rubber seal.  This extends the life of the plunger.  It may be necessary to use a 3-cm diameter brush to clean the inside of the barrel.  Rinse all parts with distilled water.  Be careful with the small parts because they can easily be lost down the drain. Important: Store plunger out of barrel.



 
 
This article first appeared in Chem13 News in January, 1997.  The authors of the original Chem13 article are: 

Bruce Mattson, faculty member, principal investigator,  Department of Chemistry, Creighton University, Omaha, Nebraska 68178 USA

Joseph Lannan, Blair High School, Blair, NE


(This page last updated on 7 June 2003)