In the Classroom edited by
Tested Demonstrations
Ed Vitz Kutztown University Kutztown, PA 19530
Combustion of White Phosphorus† submitted by:
Richard L. Keiter* and Chaminda P. Gamage Department of Chemistry, Eastern Illinois University, Charleston, IL 61920-3099; *
[email protected] checked by:
Paul E. Smith Department of Chemistry, Purdue University, West Lafayette, IN 47907-1393
The story of Brandt’s 1669 production of white phosphorus from a urine paste heated over coke has been retold many times (1). This discovery was depicted in a rather famous painting by Joseph Wright in 1771 (The Alchymist in Search of the Philosophers’ Stone: http://www.levity.com/alchemy/ wright.html). The method of production was known by only a few until it was studied by the French government in 1737 (2). Those with knowledge of the secret traveled about giving demonstrations for profit. The brilliant light that results when phosphorus reacts with oxygen was so captivating that audiences were enthralled in both Europe and the United States. P4(s) + 5O2(g) → P4O10(s) + heat + light Over the years a variety of phosphorus demonstrations have made their way into classrooms. Among the more popular ones are those in which phosphorus, dispersed in oil, undergoes slow oxidation (chemiluminescence) in a dark room on the gloves of the demonstrator (3) and phosphorus ignitions resulting from the evaporation of carbon disulfide solutions on filter paper (barking dogs) (4 ) or paper towels (5). Chemistry instructors have become acutely aware of the risks associated with demonstrations involving toxic, flammable materials (6 ) and look for ways to make working with these materials safer. At the same time they try to retain some of the more dramatic results. A recent report describing a safer method for burning phosphorus under water is typical of such efforts (7). Taking our cue from Brandt’s alchemy, we have developed a demonstration that shows the spectacular reaction of phosphorus with oxygen without exposing students or demonstrator to volatilized, toxic materials. The use of a retort in this demonstration easily provides historical connection to Brandt’s original experiment. It also provides a good example of using a water barrier to create a semi-closed system. The P4O10 generated reacts with water to produce nontoxic H3PO4.1
Experimental Procedures
Materials 0.5 g white phosphorus oxygen tank or cylinder 200-mL retort with glass stopper (thin-walled or nonPyrex retorts should not be used for this demonstration; very old retorts should be avoided)2 Keck joint clip for glass stopper (demonstrations without Keck clips have not resulted in the stopper popping out; nevertheless their use is recommended) rubber stopper with glass tube hot-plate 250-mL crystallizing dish 200 mL deionized water ring stand forceps extension clamp methyl red indicator 1.0 M copper(II) sulfate solution 5.25% sodium hypochlorite solution 6 g calcium oxide transparent shield
Apparatus A ring stand and extension clamp are used to mount the retort on the hot-plate with the retort exit tube suspended below the surface of the water in the crystallizing dish (Fig. 1). A 200-mL flask with a distilling head and a condenser may be substituted for the retort. A transparent shield should be placed in front of the apparatus to protect the audience.
P4O10(s) + 6H2O(aq) → 4H3PO4(aq) Those familiar with phosphorus from reactions involving carbon disulfide solutions might imagine that placing it in pure oxygen would be hazardous, but in fact at room temperature it does not react readily (8). White phosphorus lumps, unlike finely divided phosphorus obtained from CS2 solutions, ignite at 10 to 15 degrees above room temperature. Therefore, transporting a small lump of phosphorus from a water reservoir to a retort filled with oxygen is not a highrisk procedure. † This demonstration was first presented on October 13, 1999, in a Presidential Inauguration Lecture honoring Eastern Illinois University’s new president, Carol D. Surles.
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Figure 1. Apparatus for combustion of white phosphorus.
Journal of Chemical Education • Vol. 78 No. 7 July 2001 • JChemEd.chem.wisc.edu
In the Classroom
Procedure The glass stopper of the retort (or distilling head) is replaced with a one-hole rubber stopper containing a glass tube. The tube is connected to an oxygen tank or cylinder by means of rubber tubing. The retort is flushed with oxygen for two minutes. During this time oxygen is seen bubbling into the water of the crystallizing dish. The oxygen gas is shut off and the rubber stopper and oxygen inlet tube are replaced with a glass stopper. Several drops of methyl red indicator are added to the water of the crystallizing dish. No more than 0.5 g of white phosphorus is cut under water.3 Forceps are used to transfer the pea-sized lump of white phosphorus to the retort. The room lights are turned off and the hot-plate is turned on (highest setting). At this point the demonstrator should join the audience, where the transparent shield affords protection.4 Gas expansion before ignition causes some bubbles to appear in the crystallizing dish. After one to two minutes the phosphorus ignites, producing a brilliant light (Fig. 2). At the same time gas expansion leads to rapid bubbling into the water-filled crystallizing dish. After about five seconds, water from the crystallizing dish is drawn back into the retort and the reaction is extinguished. The hot-plate is turned off immediately, and the characteristic acid color of methyl red is noted in the water in the retort and crystallizing dish. Disposal In addition to the phosphoric/phosphorous acid solution generated from the combustion reaction, small amounts of solid residue (primarily red phosphorus, with perhaps traces of unreacted white phosphorus) are present at the conclusion of the demonstration. For those who want to avoid disposing of the excess phosphorus by allowing it to dry and oxidize in a hood, the following method has been suggested (9). To the reaction mixture is added 32 mL of 1.0 M copper(II) sulfate.5 The white phosphorus is oxidized to phosphoric acid and the copper(II) is reduced to copper metal. In addition, some copper phosphides are precipitated. After two days a small amount of blue-black residue forms, which crumbles when pushed upon with a stirring rod. No evidence of phosphorus can be visually detected at this point. The mixture is treated with 20 mL of 5.25% sodium hypochlorite solution (freshly prepared) to oxidize phosphides and other phosphorus species to phosphate. A slurry prepared from 6 g of calcium oxide in 20 mL of water is added to this mixture and the mixture is stirred for two hours to yield a black precipitate of calcium and copper phosphates and hydroxides. The solid is collected by filtration and deposited in the solid waste container. The filtrate is allowed to evaporate in the hood and the resulting residue is also added to the solid waste. Hazards White phosphorus is extremely hazardous. Ingestion of even small amounts may produce severe gastrointestinal irritation, bloody diarrhea, liver damage, skin eruptions, circulatory collapse, coma, convulsions and death. The approximate fatal dose is 50 to 100 mg. External contact may cause severe burns. Chronic poisoning resulting from ingestion or inhalation causes deterioration of bones (especially the jaw bone), spontaneous fractures, anemia, and weight loss. Keep white phosphorus under water and handle only with forceps.
Figure 2. Photographs of phosphorus reacting with oxygen, (top) initially and (bottom) at maximum brilliance.
Results and Discussion In this demonstration the oxidation of white phosphorus may be carried out safely with essentially no escape of P4, P4O10, or H3PO4. The retort not only provides a useful reaction vessel but also allows an easy historical connection to be made to the original Brandt reaction.6 Gas expansion during the course of the reaction is clearly visible as rapid bubble formation in the crystallizing dish. When the reaction ceases, water from the crystallizing dish is drawn back into the retort. This occurs because a partial vacuum exists in the retort owing to consumption and expulsion of oxygen during the course of the reaction. The vacuum is enhanced because the temperature of the retort falls at the end of the reaction, and the initial water drawn into the retort lowers the temperature even more. The red color from methyl red that appears in the solution in the crystallizing dish demonstrates that an acid has been generated. Acknowledgments We thank David Ebdon, Ken Osborne, and Ellen Keiter for helpful discussions, Ellen Keiter for 31P NMR spectra, Bev Cruse for photographs, and Ed Vitz and the checker for suggestions for improvement of the demonstration. Notes 1. The limiting reagent in this reaction is oxygen. Post-reaction examination of the water solution with 31P NMR (D2O) showed the presence of both H3PO4 (major component; δ = 0.0 ppm) and HP(O)(OH)2 (δ = 4.95 ppm, 2JPH = 664 Hz), indicating that both
JChemEd.chem.wisc.edu • Vol. 78 No. 7 July 2001 • Journal of Chemical Education
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In the Classroom P4O10 and P4O6 (and perhaps P4O7, P4O8, and P4O9) were present prior to reaction with water. Two minor signals at ᎑10.8 ppm and ᎑23.8 (multiplet) ppm were not identified. 2. Pyrex retorts may be custom-ordered from any of the major laboratory glassware companies. Although this demonstration has been repeated many times in our labs without incident, a thin-walled retort cracked during one of the checker’s demonstrations. While the reaction does not generate much heat, the retort does become somewhat hot where it is in contact with the hot-plate. Thus the cold water drawn into the hot retort may have caused the crack. Chances of this occurring will be minimized if the hot-plate is turned off immediately upon ignition. An extension cord allows this to be accomplished from a distance. If there is any doubt about the quality of the retort, it is recommended that the demonstration be carried out with a round-bottom flask, distilling head, and condenser. 3. White phosphorus samples of 0.5 g are quite adequate for large lecture rooms. Larger amounts of white phosphorus do not enhance the value of this demonstration and should be avoided. 4. An advantage of using a hot-plate for ignition rather than a Bunsen burner is that the demonstrator can be far removed from the reaction site at the time of ignition. 5. The amount of copper(II) sulfate used is sufficient to react with all of the phosphorus present even if no combustion had occurred. 6. On one occasion the reaction temperature was monitored with
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a thermometer inserted into the top of the retort. The temperature did not exceed 50 °C, affirming Brandt’s description, “cold fire”.
Literature Cited 1. Kee, T. P. Chem. Br. 1996, 32 (1), 41–45. Emsley, J. The 13th Element; Wiley: New York, 2000; p 4. 2. Toy, A. D. F.; Walsh, E. N. Phosphorus Chemistry in Everyday Living, 2nd ed.; American Chemical Society: Washington, DC, 1987; pp 1–2. 3. Shakhashiri, B. Z. Chemical Demonstrations: A Handbook for Teachers of Chemistry, Vol. 1; University of Wisconsin Press: Madison, WI, 1983; pp 186–189. 4. Ibid.; pp 74–76. 5. Silverman, L. P.; Bunn, B. B. J. Chem. Educ. 1993, 70, 405. 6. Worley, J. D. J. Chem. Educ. 1992, 69, 241. 7. Taylor, L. C. J. Chem. Educ. 1997, 74, 1074. 8. Tested Demonstrations in Chemistry, 5th ed.; Alyea, H. N.; Dutton, F. B.; Eds.; Journal of Chemical Education: Easton, Pennsylvania, 1963; p. 39. 9. Lunn, G.; Sansone, E. B. Destruction of Hazardous Chemicals in the Laboratory; Wiley-Interscience: New York, 1990; pp 215-217. Roesky, H. W.; Möckel, K. Chemical Curiosities; VCH: New York, 1996; pp 41–44.
Journal of Chemical Education • Vol. 78 No. 7 July 2001 • JChemEd.chem.wisc.edu
