Phosphorus Flamethrower: A Demonstration Using Red and White

In the Classroom edited by

Todd P. Silverstein Willamette University Salem, OR 97301-3922

Phosphorus Flamethrower: A Demonstration Using Red and White Allotropes of Phosphorus Melissa L. Golden,* Eric C. Person, Miriam Bejar, Donnie R. Golden, and Jonathan M. Powell Department of Chemistry, California State University Fresno, Fresno, California 93740-8034 *[email protected]

The phosphorus flamethrower demonstration was developed to spotlight the chemistry of elemental phosphorus. It is an adaptation of other white phosphorus synthesis demonstrations (1, 2) and of a simple test used by many forensic chemistry laboratories as part of the identification of red phosphorus in methamphetamine laboratory casework. The demonstration highlights the properties and reactivity of the white and red allotropes of elemental phosphorus. Solid elemental phosphorus may be obtained as one of four known categories of allotropes: white, red, black, and violet. The two most common types of allotropes are red and white phosphorus (3) and are the focus of this article. A discussion and comparison of black and violet phosphorus is included in the supporting information. Allotropes are different forms of a single element with distinct molecular units (4). Each allotrope has unique physical properties and chemical reactivities due to different bonding arrangements even though they are made entirely of the same element. White phosphorus, also known as yellow phosphorus (P4), exists as a volatile solid with a yellowish, waxy appearance and is the least thermodynamically stable solid form of phosphorus at room temperature (3, 5). When the other phosphorus allotropes are heated, they convert to the white form (6), which is the most stable at higher temperatures. The white allotrope is composed of discrete P4 tetrahedrons (Figure 1). Each P atom contains a lone pair of electrons and is singly bound to 3 adjacent phosphorus atoms. The four equilateral triangle faces of the tetrahedron result in phosphorus bond angles of 60°, which are significantly more constrained than the typical 100° (:PR3) (7). This high level of bond strain causes white phosphorus to be extremely reactive. White phosphorus must be stored in an anaerobic atmosphere to prevent its violent oxidation (6, 7). Red phosphorus, Pred, is considerably more stable at room temperature and is an amorphous, polymeric form of phosphorus. It was formerly thought to consist of a chain of P4 units (Figure 2A) (8). Currently, it is believed to consist of repeating P2(P10) units (Figure 2B) (9, 10). The polymeric structure makes this allotrope insoluble in most solvents. To synthesize Pred from P4, white phosphorus must be slowly heated because of the exothermic nature of the reaction (-17.7 kJ/mol) (7, 11). Heating white phosphorus between 200-400 °C for 2 days results in the formation of Pred where a P-P bond in a white phosphorus P4 unit is broken to significantly relieve the bond strain and P2 linkers between P10 units are formed. Red phosphorus is essentially air stable, although slow surface oxidation may be observed (12). Subliming red phosphorus results in conversion to the white allotrope (1, 2, 6), which will subsequently ignite in the presence of oxygen. 1154

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Figure 1. Molecular structure of white phosphorus, P4.

Figure 2. Repeating polymeric unit of amorphous red phosphorus. (A) Initially proposed repeating unit, (P4)n and (B) currently accepted repeating unit, (P2(P10))n. The P2 linkers are highlighted as well as the 5-membered ring similar to violet phosphorus.

Phosphorus chemistry is a part of everyday life and plays a role in a variety of applications: biological (teeth, bones, DNA), chemical (starting material, reducing agent, drying agent), industrial (detergents, fertilizer, carbonated beverages), military (chemical weapons, bombs, smoke screens), and criminal (illicit drug synthesis) (13-15). Many of the applications of elemental phosphorus center on its role as a reducing agent in chemical reactions. One such application of Pred is its use in the chemical reduction of pseudoephedrine to manufacture the illegal drug methamphetamine (13, 14). A second application is its use in matchbooks to provide a safe, single use source of fire (16). The same chemistry that is utilized in the combustion reaction of safety matches and for the identification of phosphorus samples in methamphetamine laboratory casework provides an interesting demonstration of the chemical differences of phosphorus allotropes. The red striking strip from a matchbook contains a mixture of ground glass and red phosphorus held together with adhesives (7, 16). Friction in the striking process supplies a source of heat to generate white phosphorus, which subsequently reacts with oxygen to release even more heat. This highly exothermic oxidation forms a white cloud of smoke composed of tetraphosphorus decaoxide and initiates reactions of fuels with potassium chlorate in the match head (7). This highly exothermic oxidation reaction of the white phosphorus

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In the Classroom

allotrope has also led to its use in mortars and shells as an incendiary and smoke producing agent since World War I despite its extreme toxicity (12, 17). In the phosphorus flamethrower demonstration and the Pred test used by forensic scientists, heat converts the relatively nontoxic red phosphorus allotrope to the more reactive white allotrope. The white phosphorus sample reacts dramatically in air to form a variety of phosphorus oxides. These phosphorus oxides then react with water to form phosphorus oxoacids. This low-cost, small-scale demonstration has the benefit of presenting the dramatic reaction of white phosphorus in air without many of the hazards associated with the storage and handling of white phosphorus (18). An additional advantage of this demonstration is that small samples of red phosphorus can be easily obtained by extraction using a matchbook if red phosphorus is not readily available. Experimental Details Materials • • • • • • • • • • • • •

Glass wool Disposable pipet Matchbook Acetone Scissors 100 mL beaker Metal microspatula Tweezers or tongs Large pipet bulb Bunsen burner Litmus paper Water Watch glass

Figure 3. Phosphorus flamethrower demonstration: (A) prepared pipet with red phosphorus, (B) synthesis apparatus, (C) sample after heating, containing white phosphorus, (D) rapid oxidation of P4 to produce bright flame, and (E) litmus paper turning red as evidence of phosphoric acid being produced.

Procedure The safety hazards associated with these reactions are controlled in part by conducting this demonstration on a small scale. Careful consideration to safety should be taken before increasing the quantity of red phosphorus used in this demonstration or deviating from the tested procedure. If chemical-grade red phosphorus is not readily available, a red phosphorus sample of sufficient quality can be easily obtained from a matchbook striking strip. It should be noted that a matchbox is not suitable for this application. The dark redbrown striking strip was cut away from the matchbook with scissors. The strip was then placed in a 100 mL beaker, and a sufficient volume of acetone to cover and soak the strip was added. The acetone dissolved enough of the adhesives after 510 min, and the red phosphorus was acquired by scraping the strip with a metal microspatula. Roughly 15 mg of a red-brown colored paste was obtained that contained red phosphorus, ground glass, and some adhesive. The sample was of sufficient purity for the purposes of this demonstration. The paste was completely dried by either letting the acetone evaporate or passing a stream of air across the surface. The tapered end of a 5.25 in. disposable pipet was lightly packed with glass wool while the paste dried. Approximately 5 mg of the red-brown solid containing red phosphorus was transferred to the center of the disposable pipet. Then the other end of the pipet was loosely packed with glass wool leaving roughly 1 in. of space between

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each glass wool plug and the sample. This gap aided in seeing the transformation from red to white phosphorus (Figure 3A). The window between the two glass wool plugs in the prepared pipet was then heated evenly over the flame of a Bunsen burner (Figure 3B). If the demonstrator does not feel comfortable handling the pipet while being heated in the flame, securing the pipet with a three-finger clamp is an option; however, this makes rotating the pipet to achieve efficient conversion difficult. Shortly after the flame showed the orange coloring of sodium from the glass, the melting of red phosphorus and transformation to white phosphorus vapor was observed (Figure 3C).(6) The white phosphorus that was made immediately consumed the oxygen in the pipet to form phosphorus oxides as evidenced by the characteristic white smoke inside the tube (12). The glass wool served as a barrier to minimize the quantity of oxygen that could enter the system while the Pred was converted to P4. This helped to prevent a premature flame. It was common to see black soot due to the decomposition of the glues that may remain in the sample. Once the white smoke was seen, the pipet was removed from the flame. A large pipet bulb was used to force air through the system. The influx of oxygen created a quick, bright flame that shot out of the pipet (Figure 3D). Care was taken so that the end of the pipet was not pointed toward any flammable items. A large pipet bulb maximized the quantity of air that passed through the pipet. After the flame self-extinguished, the bulb was used once again to bubble air from the pipet into a shallow pool of water (∼1 mL) on a watch glass that contained a piece of litmus paper (Figure 3E). The red color change substantiated the presence of acid resulting from the hydrolysis of tetraphosphorus decaoxide to form phosphoric acid. Alternatively, the air from the extinguished pipet can be bubbled through an aqueous solution of universal indicator in a 100 mL beaker. The increased visibility compared to the litmus paper makes this portion of the demonstration more suitable for larger audiences.

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Common first time difficulties performing this demonstration include properly packing the pipet with glass wool and knowing when to stop heating. Similar to packing a column with glass wool, packing the pipet with glass wool requires some practice. If it is packed too tightly, the air will not be efficiently passed through the pipet to create a spectacular flame. However, if it is packed too loosely, oxygen will be more likely to diffuse into the pipet during heating and consume the white phosphorus as it is being created. This can sometimes be seen as white smoke leaking from the pipet during the heating process. In addition, the force of the air from squeezing the large pipet bulb may cause the solids to be projected from the pipet, which is a danger that should be avoided. Thus, scaling up this reaction is not recommended. The optimal time for heating is between when smoke appears in the pipet and before the pipet begins to melt. This coincides with the pipet glowing bright orange in the flame. The pipet should be immediately removed from the flame and air forced through it to create the best flame. Because this is a very quick demonstration, dimming the lights and having multiple prepared pipets (as shown in Figure 3A) can be useful. Practicing this demonstration will help to determine the optimal time to push air through the system as well as how tightly to pack the glass wool in the pipet. To aid in this process, a video of this demonstration can be found in the supporting information. Disposal

Results and Discussion

The waste hazard associated with this demonstration is minimal. The pipet may contain traces of both white and red phosphorus after this demonstration. The first step in disposal of unreacted starting material is completely converting the red and white phosphorus to its oxides. Because of limited oxygen within the pipet, there may be unreacted white phosphorus after the end of the demonstration. Excess air is blown through the pipet to oxidize any remaining white phosphorus, and the pipet is allowed to cool before disposing as solid chemical waste. The phosphorus oxides produced can be easily added to water to form the corresponding oxoacid acids and neutralized for disposal (19, 20). As always, because of variances in local policies, precautions should be taken to comply with institutional, local, and state waste disposal rules and regulations. The demonstration produces only a small quantity of white phosphorus that should be completely used. Thus, making and storing white phosphorus for this demonstration is not recommended. It should be prepared and used immediately. Hazards The phosphorus allotropes differ considerably in their reactivity and their chemical hazards. Though the red allotrope of phosphorus is considered to be nontoxic, common trace impurities of white phosphorus result in the classification of red phosphorus as highly toxic (21). Heating red phosphorus has the potential to generate both the white allotrope and the extremely poisonous phosphine gas (22). Red phosphorus is a reducing agent and should be stored away from oxidizing agents and ignition sources. It can also react with alkaline solutions (19). The white allotrope of phosphorus is both highly toxic and reactive. The OSHA permissible exposure limit and the ACGIH threshold limit value are 0.1 mg/m3 (22). It spontaneously combusts upon exposure to oxygen and forms tetraphosphorus 1156

decaoxide, which is extremely hygroscopic, corrosive, and toxic. Target organs for white phosphorus include skin, respiratory system, digestive system, liver, muscular system, blood, kidneys, and bone (22). Long-term exposure results in necrosis of the jawbone (also known as “phossy jaw”). Should an uncontrolled fire be started, water can be used to quench the flames. To extinguish burning phosphorus on the skin, a 1% (w/w) copper sulfate solution may be used (22). The phosphorus is consumed by reduction of the Cu(II) to metallic copper. If contact of white phosphorus is made with skin, the affected area should be washed immediately. Combination of white phosphorus with moisture in the air results in trace quantities of phosphine and corrosive phosphoric and phosphorus acids (22). Standard precautions, such as having a fire extinguisher nearby, should be employed while performing this demonstration. The significant hazards of white phosphorus and phosphine gas are controlled in this demonstration by the small scale of this reaction and conducting the reaction in a properly ventilated area. When this demonstration is conducted as described, the white phosphorus is generated only transiently before it reacts in air. The small scale limits the quantity of phosphine produced to safe levels in a well-ventilated area. For more detailed information, the reader should consult the materials and safety data sheets or the chemical laboratory information profiles (21, 22).


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The phosphorus flamethrower demonstration can be presented at a wide variety of levels depending on the information included and the level of detail presented. The authors have used this demonstration for multiple purposes such as to break down the way matches ignite for grade school students, to illustrate fundamental principles of chemistry in introductory courses, and to facilitate discussions of chemical safety with first responders at crime scenes. When presented with a more detailed discussion of the chemical reactions and reactivity, it is a valuable tool to engage students in descriptive inorganic chemistry. Descriptive Inorganic Chemistry: Reactions and Reactivity The interconversion between the various elemental forms of phosphorus is complex and varies with temperature and pressure. This is further complicated by the lack of a definitive structure for red phosphorus. Red phosphorus is described as amorphous and has multiple values reported for density (2.0-2.4 g/mL) and melting point (585-610 °C) (12). Early hypothesized structures are based on a polymeric chain of P4 units (Figure 2A) (8). This structure is still reported in some introductory chemistry (23, 24) and descriptive inorganic textbooks (3). However, current textbooks used in upper-level inorganic courses tend not to describe the molecular structure of Pred other than it having an amorphous polymeric configuration (11, 25-27). The currently accepted structure is based on theoretical calculations as well as X-ray and neutron diffraction experiments and is consistent with reactivity studies (9, 10, 28, 29). The molecular arrangement of Pred is thought to range between a disordered Hittorf's phosphorus and an unsystematic arrangement of small clusters. It is hypothesized that a predominate feature of the amorphous red phosphorus is the P2(P10) polymeric chain (Figure 2B) (9). At room temperature, the amorphous red allotrope is more thermodynamically stable than the white allotrope, allowing


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In the Classroom

Scheme 1. Synthesis of Phosphorus Oxoacids from White Phosphorus

and identifying types of reactions. The first reaction in the phosphorus flamethrower demonstration is an endothermic reaction in which heat is supplied so that one elemental form of phosphorus is converted to another (11, 30): Pred f 1=4 P4ðwhiteÞ

ΔH ¼ 17:6 kJ=mol

ð1Þ

Because both the red and white phosphorus allotropes are elemental forms, the oxidation number of phosphorus in both forms is zero. In the second reaction, reactive white phosphorus oxidizes to form P4O10: P4 þ 5O2 f P4 O10

long-term storage under standard conditions. However, at higher temperatures, the white allotrope is most stable. In our demonstration, the red phosphorus is sublimed when the prepared pipet is placed in the flame. The temperature at which Pred sublimes at ambient pressure is around 450 °C (12, 30). It might be noteworthy to the reader that above 900 °C an equilibrium begins to be established between P4 and P2. The higher the temperature, the more favorable P2 becomes (12). As a reducing agent, phosphorus has the potential to react with a variety of oxidizing agents. Damp red phosphorus is known to undergo slow surface oxidation in air (12). The white phosphorus allotrope is considerably more reactive due to the strained bond angles. It oxidizes dramatically upon contact with oxygen gas to primarily form two distinct phosphorus oxides. Which oxide is formed is dependent on the relative availability of oxygen during the burning process (Scheme 1). In a shortage of oxygen, tetraphosphorus hexaoxide (P4O6) is formed on burning (5, 7). In excess oxygen, white phosphorus will burn to form tetraphosphorus decaoxide, also known as phosphorus pentoxide (P4O10) (5, 7). Both of these oxidation-reduction reactions are highly exothermic and proceed spontaneously on contact of white phosphorus with oxygen in air. Evidence of these reactions are observed as the flame that is projected when the white vapor is forced out of the pipet into the air and the white smoke that contains the phosphorus oxides (Figure 3D). As there is a significant excess of oxygen once air is blown through the tube, P4O10 can be expected to be the major product in this demonstration. In addition, any P4O6 that is formed can be further oxidized to P4O10 (7). Both of these phosphorus oxides, similar to other nonmetal oxides (e.g., CO2), will react with water to form oxoacids. Tetraphosphorus hexaoxide forms phosphorous acid, and tetraphosphorus decaoxide forms phosphoric acid (3). Evidence of these reactions is observed when the remaining smoke in the pipet, after it self-extinguishes, is bubbled through a shallow pool of water containing a strip of pH indicator paper (Figure 3E). In this step, the phosphorus oxides remaining in the pipet come in contact with the water, allowing the formation of the oxoacids, which exhibit a strongly acidic pH. General Chemistry: Redox, Heats of Reactions, and Oxoacids The three main reactions taking place during this demonstration can be used as examples for a variety of concepts and skills including balancing chemical equations, determining oxidation numbers, distinguishing oxidizing and reducing agents,


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ΔH ¼ - 2986 kJ=mol

ð2Þ

The oxidation numbers for phosphorus and oxygen as reactants are both zero. In the P4O10 product, phosphorus has an oxidation number of þ5 and oxygen has the typical -2. This reaction is a tremendously exothermic oxidation-reduction reaction with the elemental phosphorus oxidized and the oxygen gas reduced. Correspondingly, the elemental phosphorus is acting as the reducing agent whereas oxygen gas is acting as an oxidizing agent. In the final reaction, the P4O10 remaining in the pipet is bubbled through water to form phosphoric acid: P4 O10 þ 6H2 O f 4H3 PO4

ð3Þ

This is a hydration reaction typical of nonmetal oxides. The subsequent dissociation of phosphoric acid increases the hydronium ion concentration resulting in the observed color change of the pH paper. It can also be interesting to explore the hydration reactions of the phosphorus oxides in more detail. First the students are to assign the oxidation number to phosphorus in each of the two potential phosphorus oxides (P4O6 and P4O10) and then in the two oxoacids that are produced (H3PO3 and H3PO4). Students are then asked to predict which phosphorus oxide forms which oxoacid. The students should recognize that both P4O6 and H3PO3 acid have þ3 oxidation states of phosphorus as well as that both P4O10 and H3PO4 have þ5 oxidation states of phosphorus. The discussion of phosphorous acid also affords the chance to stress the importance of proper uses of suffixes in chemistry and provides an opportunity to help students avoid common spelling errors related to phosphorus suffixes. For instance, phosphorus is a noun and refers to the element, an example being the white phosphorus allotrope. Phosphorous is the adjective form of phosphorus in the þ3 oxidation state, such as phosphorous acid. Acknowledgment This work is supported in part by the CSU-LSAMP program funded by NSF under grant #HRD-0802628-515291, the Chancellor's Office of the California State University, and the Camille and Henry Dreyfus Foundation. The authors thank Angela Person Photography (www.angelaperson.com) for photography and video work. In addition, we are grateful to Tim McKibben at the Colorado Bureau of Investigation for helpful discussions pertinent to the forensic test for red phosphorus, which he developed with J. S. Chappell, H. Evans, and N. Mausolf. We thank Abdul K. Mohammed (Winston Salem State University) and Paul Krause (University of Central Arkansas) for testing this demonstration.


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Literature Cited 1. Brodkin, J. J. Chem. Educ. 1960, 37, A93. 2. Jackman, K. V.; Slabaugh, W. H. J. Chem. Educ. 1968, 45, A673. 3. Rayner-Canham, G.; Overton, T. Descriptive Inorganic Chemistry; 4th ed.; W. H. Freeman and Company: New York, 2006. 4. Sharma, B. D. J. Chem. Educ. 1987, 64, 404–407. 5. Greenwood, N. N.; Earnshaw, A. Chemistry of the Elements, 2nd ed.; Elsevier: Amsterdam, 2005. 6. Cotton, F. A.; Wilkinson, G. Advanced Inorganic Chemistry, 5th ed.; John Wiley & Sons: New York, 1988. 7. Holleman, A. F.; Wiberg, E. Inorganic Chemistry; Academic Press: San Diego, CA, 2001. 8. Pauling, L.; Simonetta, M. J. Chem. Phys. 1952, 20, 29–34. 9. Bocker, S.; Haser, M. Z. Anorg. Allg. Chem. 1995, 621, 258– 286. 10. Hartl, H. Angew. Chem., Int. Ed. Engl. 1995, 34, 2637–2638. 11. Housecroft, C. E.; Sharpe, A. G. Inorganic Chemistry, 2nd ed.; Pearson Education Limited: Edinburgh Gate, 2005. 12. Corbridge, D. E. C. Phosphorus An Outline of Its Chemistry, Biochemistry and Technology, 3rd ed.; Elsevier: Amsterdam, 1985. 13. Lee, J. S.; Han, E. Y.; Lee, S. Y.; Kim, E. M.; Park, Y. H.; Lim, M. A.; Chung, H. S.; Park, J. H. Forensic Sci. Int. 2006, 161, 209–215. 14. Skinner, H. F. Forensic Sci. Int. 1990, 48, 123–34. 15. Toy, A. D. F.; Walsh, E. N. Phosphorus Chemistry in Everyday Living, 2nd ed.; American Chemical Society: Washington DC, 1987. 16. Crass, M. F. J. Chem. Educ. 1941, 18, 280–282. 17. Emsley, J. The 13th Element The Sordid Tale of Murder, Fire, and Phosphorus; John Wiley & Sons: New York, 2000. 18. Keiter, R. L.; Gamage, C. P. J. Chem. Educ. 2001, 78, 908–910.

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19. Armour, M.-A. Hazardous Laboratory Chemicals Disposal Guide, 3rd ed.; Lewis Publishers: Boca Raton, FL, 2003. 20. Lunn, G.; Sansone, E. B. Destruction of Hazardous Chemicals in the Laboratory, 2nd ed.; John Wiley & Sons, Inc.: New York, 1994. 21. Young, J. A. J. Chem. Educ. 2004, 81, 945. 22. Young, J. A. J. Chem. Educ. 2004, 81, 946. 23. Tro, N. J. Chemistry A Molecular Approach; Pearson Prentice Hall: Upper Saddle River, NJ, 2008. 24. Silberberg, M. S. Chemistry The Molecular Nature of Matter and Change, 5th ed.; McGraw-Hill Higher Education: Boston, MA, 2009. 25. Atkins, P.; Overton, T.; Rourke, J.; Weller, M.; Armstrong, F.; Salvador, P.; Hagerman, M.; Spiro, T.; Stiefel, E. Shriver & Atkins Inorganic Chemistry, 4th ed.; W. H. Freeman and Company: New York, 2006. 26. Huheey, J. E.; Keiter, E. A.; Keiter, R. L. Inorganic Chemistry Principles of Structure and Reactivity, 4th ed.; HarperCollins College Publishers: New York, 1993. 27. Miessler, G. L.; Tarr, D. A. Inorganic Chemistry, 3rd ed.; Pearson Prentice Hall: Upper Saddle River, NJ, 2004. 28. Pfitzner, A.; Freudenthaler, E. Angew. Chem., Int. Ed. Engl. 1995, 34, 1647–1649. 29. Scherer, O. J. Acc. Chem. Res. 1999, 32, 751–762. 30. Stephenson, C. C.; Potter, R. L.; Maple, T. G.; Morrow, J. C. J. Chem. Thermodyn. 1969, 1, 59–76.

Supporting Information Available A discussion and comparison of black and violet phosphorus; a video of this demonstration. This material is available via the Internet at http://pubs.acs.org.


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Phosphorus Flamethrower: A Demonstration Using Red and White

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