Element of the Month - Oxygen

"In honor of the International Year of the Periodic Table this series of articles details the Element of the Month project developed by Stephen W. Wright (SWW), Associate Research Fellow at Pfizer Inc., and Marsha R. Folger (MRF), chemistry teacher (now retired) at Lyme – Old Lyme High School in Connecticut. Read The Element of the Month - An Introduction for an overview of the project and links to the other articles in the series." - Editor

The second element highlighted by our Element of the Month program is OXYGEN. By this point the students understand the Element of the Month program and have prepared a poster. A full class period is devoted to the discussion of oxygen and the accompanying demonstrations.

Occurrence in Nature 

Students will usually immediately respond correctly that oxygen is found in the atmosphere. When prodded “Where else?” they will answer that it occurs in water. Further “What else?” prodding usually leads to looks of puzzlement. They are likely to be unaware that most of the planet’s oxygen is to be found chemically combined in the rocks in the earth’s crust.

Uses

Clearly oxygen is essential for life! Further, it should be noted that combustion in its many forms powers human society, and that oxygen can be considered a global energy currency. One can derive energy from combustion reactions anywhere on the planet. However, oxygen is also one of the highest production volume commodity chemicals. It is used in the manufacture of steel (to remove sulfur, carbon, and phosphorus impurities that make iron brittle), paper, chemicals, rocket propellants, and in hospitals.

Figure 1: Container of oxygen gas

Physical Properties 

Students will know that oxygen is a colorless, odorless gas (figure 1). We ask the class if these properties can be used to identify oxygen. Often there will be agreement until we ask what other colorless, odorless, tasteless gases there might be. Very quickly the students will realize that carbon monoxide is colorless, odorless and tasteless.

Chemical Properties

We ask the class how oxygen is made in order to be used for all the purposes we described. Usually the immediate answer that is offered is electrolysis. We decline that answer and note that the energy costs of such a process are prohibitive. What is the cheapest way to obtain elemental oxygen? We lead the students to conclude that photosynthesis provides large amounts of oxygen and that air is a free feedstock. All that must be done is to separate the oxygen from the nitrogen. While this could be done by distillation of liquid air, membrane separation technologies are more modern and energy efficient.¹

Figure 2: Reactions yielding oxygen gas

Since we don’t have a membrane separation plant at hand, how might we prepare oxygen in the high school laboratory? Writing the reactions on the board, we note that there are several possible ways, and each method involves decomposing certain oxygen - containing compounds (see figure 2). For example, there is the thermal decomposition of red mercuric oxide that led Joseph Priestley to discover oxygen in 1774, the thermal decomposition of potassium chlorate (which is generally regarded as too dangerous to practice these days), the electrolysis of water (too slow), and the decomposition of hydrogen peroxide. We note, however, that the decomposition of sodium chlorate is used in the oxygen generators in commercial aircraft.

Video 1: Electrolysis of a (neutral) sodium sulfate solution (accessed March 2019, subscription required). Derived from Jerrold J Jacobsen and John W. Moore. Chemistry Comes Alive! Vol. 1: Abstract of Special Issue 18, a CD ROM. Journal of Chemical Education 1997 74 (5), p 607-608. DOI: 10.1021/ed074p607.

At this point, if an electrolysis apparatus is available, the electrolysis of water may be started and permitted to continue through the remainder of the period (video 1 and 2). 

Video 2: Electrolysis of a neutral solution animation (accessed March 2019, subscription required). Derived from Jerrold J Jacobsen and John W. Moore. Chemistry Comes Alive! Vol. 1: Abstract of Special Issue 18, a CD ROM. Journal of Chemical Education 1997 74 (5), p 607-608. DOI: 10.1021/ed074p607.

We show a bottle of drugstore hydrogen peroxide and note that hydrogen peroxide has an expiration date. Why is that? We ask the class why the dark bottle is used and the students will correctly guess that it can be decomposed by light. How can we decompose it to oxygen gas? We solicit heat or the use of a catalyst as answers. We demonstrate the preparation of oxygen from hydrogen peroxide using some aqueous potassium permanganate solution as a catalyst, and remark that we will encounter catalysts later in the course.²

Figure 3: Test for oxygen gas

We note that the classic test for oxygen is the so-called glowing splint test, in which a wooden splint such as a coffee stirrer is ignited, blown out, and the glowing ember is placed into the oxygen gas that we just prepared (see figure 3).³ The ember bursts into flame. We ask the class what would happen with nitrogen or helium? Next we ask the class why this may not be a great test for a sample of an unknown gas. What would happen if we tried this test with propane or hydrogen? Next we generate additional portions of oxygen and demonstrate that oxygen violently accelerates the combustion of a piece of charcoal⁴ and of iron in the form of steel wool.⁵ A video showing the combustion of steel wool may be found in Tom Kuntzleman's ChemEd X blog post, Put a Spark in Your Stoichiometry Lesson. These two demonstrations are naturally invariably well received by the class and are repeated with the room darkened. With the room still darkened, the combustion of a match box striker strip (which contains red phosphorus) is then shown, which burns with a brilliant light.⁴ We put the lights on and show the class a "No Smoking" sign and ask why one might see signs like these in areas where oxygen is in use. Using a rolled up piece of tissue paper as a surrogate for a cigarette, we repeat the glowing splint test.

Figure 4: The ozone generator apparatus⁶ 

If time permits, we ask the class if they can name another form of oxygen, and solicit the allotrope ozone as the answer. We explain to the class what allotropes are and note that other elements such as phosphorus and carbon also form allotropes. We ask how ozone is formed, and students will usually correctly answer that it is formed from diatomic oxygen in the presence of high energy UV light or electrical energy, such as lightning. We prepare ozone using a high voltage discharge and show that it is more reactive than diatomic oxygen by passing the ozonized oxygen through a flask of potassium iodide solution (see figure 4).⁷ No formation of iodine is observed until the high voltage discharge is turned on.

References and Notes 

1. A brief discussion of the properties of liquid air and the concentration of oxygen from liquid air may be found in Tom Kuntzleman's ChemEd X blog post, Collection of and Experiments with Liquid Air. See also Abstract 2-9 in Alyea, Hubert N. Tested Demonstrations in Chemistry, 6th ed.; Journal of Chemical Education: Easton, PA: 1965; pp 9.
2. Oxygen may also be generated from hydrogen peroxide and laundry bleach according to the following equation: NaOCl (aq) + H2O2 (aq)  O2 (g) + H2O (l) + NaCl (aq) An excess of hydrogen peroxide is used to ensure complete reaction of the sodium hypochlorite. A ratio of 2 teaspoons (10 mL) of laundry bleach and 3 teaspoons (15 mL) of 3% hydrogen peroxide for each 4 ounces (about 120 mL) of volume in the test bottle works well, generating enough oxygen to displace the air from the bottle without adding too much liquid to the bottle. The bleach must be added in small portions (1 to 2 teaspoons or 5 to 10 mL) at a time, to control the bubbling that occurs. See Wright, S. W. J. Chem. Educ. 2003, 80 (10), 1160A-B.
3. See Conant, James Bryant; Black, Newton Henry New Practical Chemistry; Macmillan: New York, 1940; pp. 23.
4. See Conant, James Bryant; Black, Newton Henry New Practical Chemistry; Macmillan: New York, 1940; pp. 34.
5. Summerlin, Lee R.; Borgford, Christie L.; Ealy, Julie B. Chemical Demonstrations: A Sourcebook for Teachers Volume 2, 2nd ed.; American Chemical Society: Washington, DC, 1988; pp 43. 
6. The ozone generator is set up with oxygen gas entering the glass three neck flask and passing through steel wool and out the exit tubing. The steel wool is merely an electrode surface. The center neck of the flask is closed with a stopper through which a wire is passed. The wire is in electrical contact with the steel wool inside the flask, while the other end of the wire is bonded to the ring stand. This results in the steel wool, and the inside of the flask, being electrically grounded. To generate ozone, the Tesla coil is turned on and the probe is touched to the under side of the glass flask. This creates a high voltage static field on the outside of the glass which is negated by the electrically grounded steel wool inside the flask. It is this invisible static field that converts the oxygen to zone. The probe is not brought into contact with the wire, the steel wool, or the ring stand at any time. It is only brought into contact with the glass surface. The percent conversion of oxygen to ozone is very low, less than one percent, but it is plenty to be able to smell the ozone and to produce a positive reaction with KI solution. Note that the exit tubing must be made of plastic tubing since rubber tubing reacts with ozone. The Tesla Coil (sold by Flinn Scientific) is used to apply the static field to the glass surface. The high voltage probe is also sold as a vacuum leak detector made by Electro Technic Products and sold by Fisher Scientific. (Note that this is a high voltage discharge apparatus and not a combustion apparatus.)  
7. Ransford, J. E., J. Chem. Educ., 1951, 28 (9), 477. See also Abstract 2-33s in Alyea, Hubert N. Tested Demonstrations in Chemistry, 6th ed.; Journal of Chemical Education: Easton, PA: 1965; pp 58.