Element of the Month - Phosphorus

"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 eighth element highlighted in our Element of the Month program is phosphous. We have found that the students are generally aware of phosphorus as an element but have little familiarity with its relevance to everyday life. Most often they will associate it with matches, which is true to the extent that red phosphorus is used in the striking strip of safety matches. The chemistry of elemental phosphorus is exciting to present and to observe; however, the relevance of phosphorus to the students’ lives actually centers upon the chemistry of phosphorus in its +5 oxidation state as phosphate ion. The authors were fortunate enough to be able to present the chemistry of elemental phosphorus as both its red and yellow allotropes using existing stockroom inventories. These experiments are described in the text of the article but will not be practical for many readers since the availability of elemental phosphorus has become very restricted. Nevertheless, the Element of the Month discussion of phosphorus focuses primarily upon the reactions of phosphoric acid and its salts.

Occurrence in Nature 

Phosphorus is not found free in nature, but is found in phosphate containing minerals, principally apatite, which is a calcium phosphate mineral with the formula Ca₁₀(PO₄)₆(OH)₂. We pass a sample of apatite in a clear plastic jar around the class. Phosphorus is also found in living organisms in DNA, RNA, and ATP, the latter of which is the energy currency of aerobic plants and animals. Animals with internal skeletons use apatite to form bones and teeth, and the enamel of teeth consists of fluorapatite Ca₁₀(PO₄)₆(F)₂. We ask the class who discovered phosphorus and when. This usually produces puzzled looks and whimsical answers. We note that phosphorus was discovered by the alchemist Hennig Brand, 1669, in present day Hamburg, Germany. It is the first element isolated that was not known in ancient times. We ask the class how the name phosphorus came about and explain that it is derived from the Greek, meaning “light bearer”. This is in recognition of the chemiluminescence of the reaction between white phosphorus and oxygen.

phosphorus_at_home.png

Figure 1: A few products containing phosphorus

Uses

Phosphorus is essential for life, as noted by its role in key molecules such as DNA, RNA and ATP. Phosphates are also used as regulators to turn the catalytic activity of enzymes on and off. Like nitrogen, phosphorus is incorporated in fertilizers to bring this element into the food chain. We show a typical fertilizer package and point out the composition is noted as three numbers, for example “28 – 5 – 10”. We note that the second number indicates the phosphorus content.¹ Phosphorus compounds are used in supplements for animal feed, to soften water, in detergents and cleaners, and in toothpastes. Its use in cleaning products has been limited more recently due to concerns about algal blooms in waterways. Phosphorus compounds are used in insecticides, such as malathion, and as chemical warfare agents (the “nerve gases”). Phosphorus compounds have many uses in foods, including the use of phosphoric acid as an acidifying agent for soft drinks and acidic phosphate salts in leavening. Phosphorus compounds are used in flame retardants, in matches (the striking strip for safety matches contains red phosphorus) and in some pyrotechnic compositions. We show on the lecture desk a package of fertilizer, some “TSP” cleaner,² a tube of toothpaste, a can of cola, box of matches, and a can of baking powder (figure 1). We note that phosphorus chemistry has many contrasts: it is used in foods and poisons, it is both a water purification agent and a pollutant, it is used both for pyrotechnics and fireproofing.

Physical Properties

Phosphorus is an excellent example of allotropes, different physical forms of the same element that result from different bonding patterns in the element. Phosphorus has three allotropes: yellow phosphorus (also known as white phosphorus), red phosphorus, and violet phosphorus. Yellow phosphorus is a waxy, brittle, low - melting (44 °C, about 110 °F), substance that dissolves in some organic solvents but not in water. Red phosphorous is a high melting (>400 °C) powder that is insoluble in all solvents. Violet phosphorus is relatively rare, requiring a difficult high temperature crystallization procedure to prepare.

Chemical Properties

Phosphorus is a non-metal. It is never found uncombined in nature. Elemental phosphorus, yellow or red, is a remarkable reducing agent. The yellow allotrope is more reactive than the red but both are powerful reducing agents. They react vigorously with oxidizing agents, even relatively weak oxidizing agents. Yellow phosphorus is remarkable for being spontaneously flammable in air and very poisonous. It must always be stored under water. Red phosphorus is not spontaneously flammable in air and is relatively non-toxic. Red phosphorus is much safer to handle than the yellow allotrope and it is much more commonly used.

We burn a small sample, about 0.1 g, of red phosphorus in a deflagrating spoon and we draw the students attention to the peculiarly colored flame and the rather abundant white smoke. These are very characteristic of phosphorus and combustible phosphorus compounds when they burn. We use a large test tube to hold the deflagrating spoon before the experiment and return the spoon to a large test tube of water afterward to extinguish the phosphorus.³ Next, we demonstrate the spontaneous combustion of yellow phosphorus using a solution of yellow phosphorus in carbon disulfide.⁴ We perform the experiment once using a filter paper on a tripod, and once as the “barking dog” variant with a filter paper circle on a 1000 mL cylinder. Again, we note the white smoke of phosphorus pentoxide P₂O₅ that was produced.

We then explain that phosphorus pentoxide, produced by burning phosphorus in air, reacts with water to form phosphoric acid, H₃PO₄. We note that the reaction is very vigorous and ask the class to listen carefully while we add about five grams of phosphorus pentoxide to about 500 mL of water in a 1000 mL beaker. The P₂O₅ dissolves instantly with an audible hiss and a puff of steam. Next, we add universal indicator to show solution is acid. Phosphoric acid has many uses, but one that the students encounter every day is its use in cola drinks to make them tart. When phosphoric acid is neutralized the salts formed are called phosphates. Because phosphoric acid has three protons per molecule, you can neutralize them one at a time. We write out the reactions for the neutralization of phosphoric acid on the board, using different color dry erase markers (red, orange, blue, purple) for each step as shown below.⁵

H₃PO₄ → H₂PO₄^- → HPO₄^2- → PO₄^3-

These different phosphate ions have different pH values when dissolved in water. We add a spoonful of NaH₂PO₄, Na₂HPO₄, and Na₃PO₄ to each of three beakers of water, one compound to each beaker, stir briefly, and then we add universal indicator to each beaker. We contrast colors produced in that experiment by adding universal indicator solution to a beaker containing dilute acetic acid, dilute NaOH, and plain water. We point out that phosphoric acid and phosphates are non-toxic and can be combined in different proportions to give almost any pH desired, making them very useful in controlling pH in foods, beverages, and other products. This pH control is brought about by the formation of buffers. And we remind the class that buffers are critically important in biology.

We show the class a beaker of water, add universal indicator pH and then add a few drops of dilute hydrochloric acid and observe the immediate pH change. We repeat the process with another beaker of water and add dilute sodium hydroxide dropwise. In contrast, phosphate solutions resist pH changes when treated with hydrochloric acid or sodium hydroxide because a buffer is formed. We take the NaH₂PO₄ solution containing universal indicator that we prepared earlier and add dilute sodium hydroxide dropwise. Next, we take the Na₃PO₄ solution containing universal indicator that we prepared earlier and add dilute hydrochloric acid dropwise. In both cases the solutions resist pH change. We explain that when phosphoric acid is neutralized with sodium hydroxide, the first product is sodium dihydrogen phosphate, but that’s still acidic. When hydrochloric acid is neutralized, the product is sodium chloride, which is neutral and doesn’t resist further pH changes.

Next, we ask if the buffer can be overwhelmed? Yes, of course, if enough acid or base is added. Again, we take the NaH₂PO₄ solution containing universal indicator that we prepared earlier and treated with dilute NaOH. Now we add NaOH pellets to the beaker and note the change in color of the indicator. Likewise, we take the Na₃PO₄ solution containing universal indicator that we prepared earlier and add treated with dilute hydrochloric acid. Now we add concentrated hydrochloric acid and note the change in color of the indicator.

We ask "what would happen if we combined solutions of NaH₂PO₄ and Na₃PO₄?" By now the class is fully on task and will immediately announce that the reaction will produce Na₂HPO₄. We dissolve 5.5 g of NaH₂PO₄ (monohydrate, 0.04 mol) in 400 mL of water in one beaker, and 15.2 g of Na₃PO₄ (dodecahydrate, 0.04 mol) in 400 mL of water in another beaker. We add universal indicator to each, then pour the two solutions together simultaneously into a 1000 mL beaker. As anticipated, once the solutions are mixed the indicator shows the pH is consistent with the presence of Na₂HPO₄.

These pH changes can be put to some interesting uses, for example to leaven bread and baked goods like cakes. Leavening is traditionally produced by carbon dioxide bubbles in the batter produced by the yeast fermentation of sugars. But yeast isn’t always practical, in particular because it is slow and not all batters have the appropriate consistency. Baking soda can be used in combination with acidic ingredients like buttermilk, but sometimes the baking soda isn’t all reacted in the final product and the food may taste bitter or soapy. For that reason, baking powder was invented, a carefully prepared mixture of baking soda with an acid. The mixture gives off carbon dioxide when water is added. Early baking powders used tartaric acid from wine making, but it reacted too quickly, and the carbon dioxide escaped from the batter. Slower - acting baking powders were invented using acidic phosphates. In most common formulas sold today, calcium dihydrogen phosphate Ca(H₂PO₄)₂ is used.⁶ We dissolve some calcium dihydrogen phosphate in water, add universal indicator, and note that the solution is acidic. Next we make a baking powder mix by combining some calcium dihydrogen phosphate and sodium bicarbonate in a beaker and stirring the powders together. Then we add water, followed by some dishwashing liquid to produce a fine beaker full of bubbles. Lastly, we note that calcium phosphates are used in toothpastes – just the right hardness for polishing teeth, which, we remind the class, are made of the calcium phosphate apatite.

In contrast to the fiery nature of elemental phosphorus, phosphates can be used to make things fireproof. Some fireproofing compositions are phosphate-based compounds, including the bright red fire retardant that is dropped from aircraft ahead of forest fires. Ammonium dihydrogen phosphate (NH₄)H₂PO₄ is a typical inexpensive fireproofing agent (but one that washes away with water!). It is sometimes used to fireproof one - use items like Christmas trees and disposable paper items like party decorations and crepe paper. We ignite a circle of filter paper with a burner, and of course it burns as expected. Then we show the class a similar circle of treated filter paper.⁷ The effect is remarkable. The filter paper circle blackens and chars, but no flame is produced.⁸

Phosphoric Acid and Buffers

Phosphoric acid has three acidic hydrogens. Each of these acidic hydrogens can be neutralized by a base such as sodium hydroxide. The complete neutralization of phosphoric acid H₃PO₄ to phosphate ion PO₄^3- occurs in three distinct steps. The intermediate ions dihydrogen phosphate H₂PO₄^1- and hydrogen phosphate HPO₄^2- can also be completely neutralized by adding two moles and one mole of sodium hydroxide, respectively. Each of the four species in the sequence has a characteristic pH in water and gives a characteristic color with universal pH indicator solution (figure 2). Note that when the neutralization is complete the solution pH is far from neutral. Phosphate ion PO₄^3- is a strong base in water because of its 3- charge. Trisodium phosphate (TSP, Na₃PO₄) is used in many cleaning products for this reason.

p_equations.png

Figure 2: Changing color with universal indicator

Phosphates are useful for making buffers because they are non – toxic, cheap, stable, don’t impart a taste to food products, and because they can be combined to give almost any pH value from acidic to basic. While phosphoric acid is a liquid, which limits its use to liquid products (like colas), the phosphate salts NaH₂PO₄, Na₂HPO₄ and Na₃PO₄ are all stable, solid compounds that can easily be obtained in a high degree of purity.

REFERENCES AND NOTES

1. The three-number sequence indicates the content of nitrogen, phosphorus and potassium. The phosphorus content is expressed as the weight percent of phosphorus pentoxide (P₂O₅) in the fertilizer. The phosphorus is present as phosphate salts, typically calcium, ammonium and / or potassium phosphates, and not as P₂O₅. The arithmetic conversion to P₂O₅ content by weight permits a comparison of fertilizers to be made regardless of the phosphate compounds actually present.

2. “TSP” stands for Tri Sodium Phosphate, Na₃PO₄, which is a hydrate of varying proportions in the commercial TSP sold as a cleaner in hardware stores. The PO₄^3- ion is strongly alkaline in water and therefore makes it a useful cleaner.

3. With phosphorus being largely unavailable now, one may obtain a similar result if the striking strip of a pack of safety matches is cut away, held in forceps, and burned. The small amount of red phosphorus on the striking strip will still generate a remarkable amount of smoke for the small sample size.

4. A solution of 2 g of yellow phosphorus in 10 mL of carbon disulfide was used. The experiment on a tripod is described in Conant, James Bryant; Black, Newton Henry New Practical Chemistry; Macmillan: New York, 1940; pp. 39. The “barking dog” variation is described in Chen, Philip S. Entertaining and Educational Chemical Demonstrations; Chemical Elements Publishing Co.: Camarillo, CA, 1974; pp 65. This demonstration works especially well in cylinders about 12” tall and 2” in diameter. Often one needs to use a little more solution than seems necessary in order to get a loud “woof”. The solution of yellow phosphorus was stored in a metal cylinder as described in Wright, S. W. J. Chem. Ed., 1996, 73, 818.

5. The dry erase marker colors are chosen to match the color of universal indicator solution when added to a solution containing the appropriate phosphorus compound.

6. For an excellent discussion of phosphates in baking powders and other food uses, see Toy, Arthur D. F.; Walsh, Edward N. Phosphorus Chemistry in Everyday Living; American Chemical Society: Washington, DC, 1987; pp. 23 – 61.

7. The filter paper circles are soaked in a solution of ammonium phosphate in water containing 60 g of (NH₄)H₂PO₄ in a liter of water and then allowed to air dry. For an excellent discussion of the mechanism of action and uses of ammonium phosphate as a flame retardant, see Toy, Arthur D. F.; Walsh, Edward N. Phosphorus Chemistry in Everyday Living; American Chemical Society: Washington, DC, 1987; pp. 135 – 139.