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Chemistry: Phosphorus


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  • Type: Video Tutorial
  • Length: 12:48
  • Media: Video/mp4
  • Use: Watch Online & Download
  • Access Period: Unrestricted
  • Download: MP4 (iPod compatible)
  • Size: 137 MB
  • Posted: 07/14/2009

This lesson is part of the following series:

Chemistry: Full Course (303 lessons, $198.00)
Chemistry: Nonmetals (12 lessons, $19.80)
Chemistry: Group 15: Nitrogen and Phosphorus (2 lessons, $3.96)

This lesson was selected from a broader, comprehensive course, Chemistry, taught by Professor Harman, Professor Yee, and Professor Sammakia. This course and others are available from Thinkwell, Inc. The full course can be found at http://www.thinkwell.com/student/product/chemistry. The full course covers atoms, molecules and ions, stoichiometry, reactions in aqueous solutions, gases, thermochemistry, Modern Atomic Theory, electron configurations, periodicity, chemical bonding, molecular geometry, bonding theory, oxidation-reduction reactions, condensed phases, solution properties, kinetics, acids and bases, organic reactions, thermodynamics, nuclear chemistry, metals, nonmetals, biochemistry, organic chemistry, and more.

Dean Harman is a professor of chemistry at the University of Virginia, where he has been honored with several teaching awards. He heads Harman Research Group, which specializes in the novel organic transformations made possible by electron-rich metal centers such as Os(II), RE(I), AND W(0). He holds a Ph.D. from Stanford University.

Gordon Yee is an associate professor of chemistry at Virginia Tech in Blacksburg, VA. He received his Ph.D. from Stanford University and completed postdoctoral work at DuPont. A widely published author, Professor Yee studies molecule-based magnetism.

Tarek Sammakia is a Professor of Chemistry at the University of Colorado at Boulder where he teaches organic chemistry to undergraduate and graduate students. He received his Ph.D. from Yale University and carried out postdoctoral research at Harvard University. He has received several national awards for his work in synthetic and mechanistic organic chemistry.

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Unlike nitrogen, dinitrogen, which is the principle component of air, phosphorous is most definitely not a diatomic molecule. In fact in nature it appears mostly in the form of phosphate rock, of which calcium phosphate is an example. And to get elemental phosphorous from phosphate rock, we have to reduce the phosphate rock using carbon. So this is the principle industrial process for making phosphorous in its elemental state. It involves using silicon dioxide or sand and carbon to form P[4], which is one of the allotropes of phosphorus along with calcium silicate and carbon monoxide.
Now the reason why this reaction works so well is because everything else is a solid or a liquid, except for the P[4] and the carbon monoxide and so it is possible to isolate this by distillation essentially directly out of the reaction mixture.
Now I will remind you that phosphorous is one of the principle components in fertilizer. Remember it is NPK for nitrogen, phosphorous and potassium. And so that is going to be a recurring theme in our discussion right now.
Now the bad thing about white phosphorous, which is the P[4] allotrope, is that it consists of phosphorous tetrahedral with a phosphorous atom in each one of the corners. And that implies that you have 60 degree PPP bond angle and there is a great deal of strain there. Remember phosphorous doesn't want to have such small bond angles, because it has p orbitals to make the bonds. And so white phosphorous is air sensitive and the more stable allotropes that you can see in the graphic are red phosphorous and black phosphorous. So white phosphorous is converted to red phosphorous and black phosphorous by heating in the absence of air. Remember allotropes are different physical forms of the crystalline forms of the element.
Now, reactions of phosphorous include reactions with the halogens. For instance phosphorous reacts with chlorine to form PCL[3] and then in the excess of chlorine it will go all the way to PCL[5]. What we have here is phosphorous in the three plus oxidation state and then PCL[5], phosphorous in the five plus oxidation state. That is one of the differences between phosphorous and nitrogen of course as you go down the column in any of the columns, you can have higher oxidation states. So phosphorous[5] whereas nitrogen[5] is a non-existent oxidation state. These sorts of reactions also work with bromine and iodine, but fluorine is much more oxidizing, and so when you take phosphorous and you react it with fluorine you go all the way to phosphorous pentafluoride, so in order to get phosphorous trifluoride you have to sneak up on it a little bit.
Now the reason why we would talk about the phosphorous halides is that the hydroiodic acid and hydrobomic acid, remember you cannot get by reacting sodium bromide or sodium iodide with sulfuric acid. So you get to hydroiodic and hydrobomic acid commercially by the hydrolysis of PI[3] or PPR[3] or PI[5] or PPR[5]. You can see that if you hydrolyze phosphorous trioxide you make three moles of hydroiodic acid and phosphorus acid as the byproduct H[3]PO[4]. If you start with phosphorous pen iodide you get five equivalents of hydroiodic acid and phosphoric acid. Where the difference here is in the oxidation state of the phosphorous.
Now, phosphorous reacts in oxygen to form two oxides P[4]O[10] and P[4]O[6], tetra phosphorous deca oxide and tetra phosphorous hex oxide and they are obviously somewhat cumbersome names, so I will just call them either P[4]O[10] or P[4]O[6]. In the laboratory you will also sometimes hear P[4]O[10] referred to as phosphorous-pent oxide, but we understand that the molecular structure is actually P[4]O[10]. Where phosphorous is five plus oxidation state here and the three plus oxidation state here.
Now both of these compounds are hydroscopic. They absorb water and they obviously the reason why they absorb water and react with water is that P[4]O[10] reacts with water to form phosphoric acid. Another way to say this is that P[4]O[10] is the anhydride of phosphoric acid. And then P[4]O[6] reacts with water to form phosphorous acid. So P[4]O[6] is the anhydride of phosphorous acid.
Now both phosphoric acid and phosphorous acid have the tendency to concatenate, that is if you take two phosphoric acids, so here is an example, molecules, you can condense them to form a POP bond and what gets spit out is water. So this is an example of a condensation reaction and because this happens with phosphoric acid, you'll see phosphoric acid in several different forms, commercially available. What this mono phosphorous compound is called is ortho phosphoric acid. And then the diamer is known as pyrophosphoric acid, but you can see that if you take either of these two things and put them into water, you are going to get a solution of phosphoric acid. The only thing to consider is how many phosphorous equivalents you're going to get, depending on what you start with. Orthophosphoric acid is typically purchased as a solution.
Now, if you extend this idea to infinity then you get what is known as metaphosphoric acid where you've condensed a bunch of phosphoric acid molecules and made these infinite chains of OPOPOP going down. And this has the empirical formula of HPO[3] and we can illustrate the fact that its empirical formula is HPO[3], then it's molecular formula is N, with N being the number for repeat units. And this is known as metaphosphoric acid. But again, just another form of phosphoric acid. And what I forgot to mention is that nature actually uses this reaction. The idea that phosphates can condense and kick out water. The reaction going in this direction is a way that nature stores energy. You may have heard of the phrase ATP or the compound ATP, which stands for adenosine triphosphate. So nature makes ATP and that is how it stores up energy and then when it wants to get some energy back, it runs this reaction in reverse. Reacting ATP with water to form ADP, which is adenosine diphosphate, spitting out inorganic phosphate. So again, nature uses the formation or breaking of OP bonds in phosphates to store or utilize energy.
Now the congregate base of phosphoric acid or the ultimate congregate base of phosphoric acid after it has been deprotonated three times is the phosphate and ion. When you go to the supermarket you can buy sodium phosphate, it goes by the name trisodium phosphate in the store. It has the abbreviation TSP, and it turns out that TSP is pretty good at cleaning things. It is very alkaline, very basic. So be sure to use gloves when you use it. It is really good for cleaning up greasy things actually. In the detergent industry, like laundry detergents and automatic dishwashing detergents, these trisoduim phosphates salts and the various other biphosphates and things like that are used as builders. And what a builder is, it's something that makes the detergent work better. So for instance, one of the properties of phosphates, sodium phosphate is that it is going to make the solution alkaline. The pH is going to be above seven. And it turns out that soaps work better in high pH. It is not really high, but above seven. And the reason is that a soap, you'll recall, is a fatty acid carboxylate salt, like stearic acid carboxylate salt. And if the sodium stearate gets protonated in low pH, it become insoluble in water and it precipitates out and so it can't be in the solution to do the cleaning that it is going to do. And so that is why we basify the water. But the other thing is that hard water contains calcium and magnesium ions and these calcium and magnesium ions will react with the stearate and precipitate out. Bathtub ring consists of calcium and magnesium stearate. And what happens is that if these precipitate out the soap, then again the soap isn't there to do what it is supposed to do, to do it's cleaning. And sort of head them off at the pass. The phosphates will react with the calcium or magnesium first, form insoluble calcium phosphates or magnesium phosphates, take those ions out of solution and that leaves the soap to do what it wants.
Now the problem is, that these phosphates again are fertilizers and so when we run them down the drain, they eventually make it into our water supplies, our surface water supplies and if those things are used to irrigate fields or something like that. That is okay, but so long as they are in the water supply, what they do is fertilize the production of algae. And the problem with algae is that not only does algae gum up the water supply, because it just floats on a mat and makes things really ugly, but because they are biological the consume oxygen. And so they take oxygen out of the water and that means that there is less water to allow the fish to grow. So the fish will die very often times when you have algae blooms. And again algae blooms can be the cause of dumping a bunch of phosphate into the water. So there is a trend in the United States, actually a pretty significant trend to get rid of the phosphates in detergents. They still appear in automatic dishwashing detergents, however laundry detergents are typically phosphate free. And you will see that written on the box. What people have done instead is to use other builders like sodium carbonate. Sodium carbonate still is a base so it does the pH control and calcium carbonate and magnesium carbonate are insoluble. And so what they do is they precipitate out these salts and they do basically what the phosphate does.
Now another place where phosphorous appears in life is, strike anywhere matches, you will know a strike anywhere match because the tip is two different colors. It's typically white and blue or white and red or blue and red. The other component is potassium chlorate and these two react in the presence of oxygen to oxidize the phosphorus and the sulfur. This is an exothermic reaction and so that is what gets the match going. This is an unbalanced reaction, because obviously what we have are two separate supplies of oxygen, so it wouldn't make any sense to balance this, unless we knew exactly what the relative proportions of these things were. But that is an example of where phosphorous is in your everyday life.
Finally, let me just say a few words about the other elements in the group 15, in the nitrogen family. Arsenic and antimony, bismuth is not a nonmetal, bismuth is a metal, but antimony and arsenic, you'll recall, are what we call metalloids. They are right on the borderline between metal and nonmetals. Arsenic appears in you life in the form of gallium arsenide is a semi conductor. If you think of where gallium and arsenic are in the periodic table, they are on either side of silicon. And silicon is the semi-conductor that gives rise to the semi-conductor industry. Well it turns out that gallium arsenide does something that silicon can't do, it is a semi-conductor, but gallium arsenide can give off light and silicon can't. It just can't. So all of the lights emitting diodes, at least the early light emitting diodes are gallium arsenide. The way you get different colors is you mix solid solutions, you put aluminum on the gallium site, and it changes the band gap. When you change the band gap, you get different colors. So gallium arsenide is the light emitting diode material. And it has the zinc blend structure which is face centered cubic latticed of gallium with arsenic in half of the tetrahedral holes, so it is a lot like the silicon or diamond structure, except instead of having the same element in both sights, it has gallium in one side and arsenic in the other side.
Finally, the alloy pewter, which is mostly, tin, but with amounts of copper, bismuth and antimony is an alloy that was very popular in the 1800's, 1700's. If you ever go and visit Professor Harmon in Virginia, be sure to stop by Colonial Williamsburg, which is a recreated 18^th century town, and the dinner plates that you will eat off of are typically pewter. You will recognize it as sort of a dull silvery colored material and it is an alloy, again, of tin with small amounts of bismuth and antimony. So the members the group 15 family play a role in your life that, probably most important is that in the phosphate fertilizer, so if you are a gardener, now you know where phosphorous comes from and how we process it to make it useful to making your plants grow.
The Nonmetals
Group 15: Nitrogen and Phosphorus
Phosphorus Page [1 of 3]

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