Volume 3

Natural Occurrence and Physical Properties • The Exceptional Stability and Inertness of N2. Ammonia. Synthesis of Nitric Acid by Oxidation of Ammonia • General Properties of Nitric Acid • The Variety of Oxidation States of Nitrogen Nitric Acid as an Oxidizer. Reactions of HNO3 with Metals and Nonmetals: The Myths and the Reality • Thermal Decomposition of Nitrates • Phosphorus. Phosphorus Pentoxide • Phosphoric Acid • Exercises
3.5.1. Natural Occurrence and Physical Properties. Nitrogen and phosphorus belong to Group 15 of the periodic table. Besides N and P, group 15 has three more, heavier elements: arsenic (As), antimony (Sb), and bismuth (Bi). The elements of group 15 are often referred to as the pnictogens.

While air is nearly 80% nitrogen (N2), in the Earth's crust there is only about 0.002% of nitrogen by mass. The natural resources of the only two significant nitrogen minerals, KNO3 (Indian saltpeter or just saltpeter) and NaNO3 (Chilean saltpeter) have long been exhausted. Phosphorus is among top 15 most abundant elements in the Earth's crust. Unlike nitrogen, however, phosphorus does not occur in nature as a simple substance.

Nitrogen is a gas at standard temperature and pressure. The boiling and melting points of N2 are -196 and -210 oC, respectively. The three most common allotropes of phosphorus are white, red, and black phosphorus. All three are solids. The most common allotrope of the three is white phosphorus, sometimes called yellow phosphorus. White phosphorus melts at 44 oC and boils at 280 oC.

3.5.2. The Exceptional Stability and Inertness of N2. Ammonia. Synthesis of Nitric Acid by Oxidation of Ammonia. The two nitrogen atoms in a molecule of N2 are bonded to each other by a triple covalent bond, N≡N. This triple bond is the second strongest bond known (next to the C≡O triple bond) and is extremely challenging to break. At room temperature, N2 does not react even with fluorine, the most reactive substance known.

Why do we need to make N2 react in the first place? Why do we need low-cost, efficient methods to break the exceptionally strong N≡N bond? The answer is simple: because we need nitrogen compounds to make fertilizers, pharmaceuticals, crop protection agents, dyes, explosives, propellants, and various materials and polymers such as nylon and Kevlar®. And, atmospheric nitrogen is free to everyone!

Is there not a single substance that would react with N2 quite willingly under mild conditions? There are a few. Some metals (Li, Mg, Ca) react with N2 at ambient temperature, the fastest reaction being that with lithium to give lithium nitride, Li3N (Figure 3-40). Unfortunately, this reaction is too far from being economical for any industrial application.
Figure 3-40. Lithium reacts with nitrogen at room temperature.

The dire need to develop the industrial chemistry of N2 was experienced by Germany during World War I. The blockade of Germany by British submarines cut off the maritime supplies of many materials that were vital for the German economy, including Chilean saltpeter, NaNO3. Without the NaNO3 from Chile, not only was Germany unable to make explosives needed for the war, but the entire country was facing the grave risk of starvation from the severe shortage of nitrogen fertilizers, which back then were made exclusively from the saltpeter. The problem was solved by the development of the synthesis of ammonia from N2 and H2 (Figure 3-41).
Figure 3-41. Synthesis of ammonia from nitrogen and hydrogen.

The reaction of N2 with H2 is reversible, occurring only in the presence of a catalyst at high temperatures and pressures. We have already discussed this reaction in considerable detail to learn that the ammonia synthesis, also known as the Haber-Bosch process, is credited with feeding up to one-half of the current world population (Volume 2).

Ammonia is a colorless gas that has a very strong suffocating odor and is easily soluble in water, up to 32% at room temperature and atmospheric pressure. Ammonia is a base. The basicity of NH3, as we already know (Volume 2; 2.4.9), is due to the lone electron pair on the nitrogen atom. This lone pair becomes a shared one upon addition of a proton when ammonia is treated with an acid. Figure 3-42 illustrates the formation of ammonium chloride from ammonia and HCl.
Figure 3-42. Reaction of NH3 with H+ from HCl.

Ammonia burns in air to give N2 and H2O. However, the controlled oxidation of NH3 with O2 (air) over a platinum-rhodium catalyst furnishes nitric oxide, NO (Figure 3-43).
Figure 3-43. Combustion (top) and catalytic oxidation (bottom) reactions of ammonia.

The catalytic oxidation of NH3 to NO is the key step of the modern industrial synthesis of nitric acid, HNO3. There are three key steps involved in the nitric acid process (Figure 3-44), all of them being redox reactions.
Figure 3-44. Industrial synthesis of nitric acid.

Step 1 is the air-oxidation of ammonia to nitric oxide (NO) over a catalyst made of Pt and Rh at 800-950 oC. In this step, the oxidation state of the nitrogen atom changes from -3 to +2.

Step 2 is the oxidation of NO produced in the first step to nitrogen dioxide, NO2, with extra air. The oxidation state of the nitrogen atom in this step changes from +2 to +4. This is a facile reaction that occurs spontaneously in the absence of any catalyst. As the reaction occurs, the gaseous reaction mixture turns orange brown. Unlike NO and air that are colorless, NO2 is an orange-brown gas.

Step 3 is the reaction of the NO-derived NO2 with water in the presence of additional O2 from the air to give nitric acid. The change in the oxidation state of the nitrogen in this step is from +4 (NO2) to +5 (HNO3).

Before the development of the synthesis from NH3, nitric acid was made from naturally occurring nitrates (KNO3 and NaNO3) as well as from air in the arc process. The arc process takes its name from the use of an electric arc to reach the very high temperatures (> 2,000 oC) required for the reaction, N2 + O2 = 2 NO. The arc process is no longer used due to its low efficiency and high cost.

3.5.3. General Properties of Nitric Acid. Pure nitric acid is a colorless liquid that boils at 83 oC at 1 atm. Known as fuming nitric acid, undiluted HNO3 slowly decomposes in time (Figure 3-45). As the orange-brown NO2 produced in the decomposition is soluble in HNO3, most samples of fuming nitric acid are yellow, orange or even brown in color. Since the decomposition is promoted by light, fuming HNO3 is conventionally stored in dark brown glass bottles.
Figure 3-45. Decomposition of nitric acid.

What is called concentrated nitric acid is a 60-70% solution of HNO3 in water. Such concentrated solutions of HNO3 are considerably more stable than fuming nitric acid.

Nitric acid is a strong acid that exists in its dissociated form in aqueous solutions. Like other acids, HNO3 readily reacts with bases and basic oxides to form water and a salt (Figure 3-46).
Figure 3-46. Like other acids, HNO3 reacts with bases and basic oxides.

A notable feature of nitric acid is that virtually all of its salts (nitrates) are water-soluble. The vast majority of inorganic acids give rise to at least some insoluble salts. For example, HCl forms insoluble AgCl and H2SO4 insoluble BaSO4. In contrast, all metal salts of nitric acid are easily soluble in water.

3.5.4. The Variety of Oxidation States of Nitrogen. Besides NO and NO2, nitrogen forms other oxides, including N2O, N2O3, N2O4, and N2O5. In addition to nitric acid, HNO3, there is another nitrogen oxoacid, HNO2, called nitrous acid. The existence of all of these compounds shows that the oxidation state of nitrogen can vary in a broad range (Table 6).

Table 6. Oxidation states of nitrogen with examples.
The variety of oxidations states of nitrogen suggests that the redox chemistry of nitrogen compounds should be very rich. This is indeed the case, as we will see in the next subsection. However, of the compounds listed in Table 6, HNO2 (nitrous acid), N2O3 (dinitrogen trioxide), N2O4 (dinitrogen tetroxide), and N2O5 (dinitrogen pentoxide) are unstable and, consequently, are rarely formed as final products in chemical reactions. The most common nitrogen-containing substances involved and/or produced in redox reactions of HNO3 are NO2, NO, N2O, N2, and NH3 (or its protonated form, the ammonium cation, NH4+).
Digression. One stable nitrogen oxide is quite well-known to many non-chemists. This compound is nitrous oxide, N2O, also known as laughing gas. The common name "laughing gas" was coined by the outstanding British scientist Sir Humphry Davy (1778-1829) who did much experimentation with N2O as a euphoric effect substance. The satirical print from 1830 in Figure 3-47 is a humorous depiction of one of those experiments. Another hilarious 19th century print showing an early laughing gas party is presented in Figure 3-48. As an anesthetic and painkiller, N2O finds a broad use in medicine and dentistry, and is on the World Health Organization's list of essential medicines. To some individuals, laughing gas is better known as a recreational drug.
Figure 3-47. Humphry Davy administering laughing gas to a woman (source).
Figure 3-48. A laughing gas party in the 19th century (source).

3.5.5. Nitric Acid as an Oxidizer. Reactions of HNO3 with Metals and Nonmetals: The Myths and the Reality. We know that strong acids such as HCl react with metals to give H2 and a salt. These are redox reactions, in which the metal reduces the H+ of an acid, as exemplified by the reaction of Zn with HCl (Figure 3-49).
Figure 3-49. In the reaction of a metal with an acid, the metal is oxidized and the H+ is reduced.

Stronger electron-donating metals reduce the H+ more readily. The order of reactivity of selected metals toward acids, as already discussed in Volume 1, is presented in Figure 3-50. The least reactive metals, those to the right from the H in the series (Cu, Hg, Ag, Au, and Pt) are not powerful enough electron donors to reduce a proton.
Figure 3-50. Metal reactivity series.

While the least reactive metals are inert toward most typical strong acids such as HCl, some of them readily react with HNO3. As shown in this video, copper reacts vigorously with concentrated HNO3 to give blue Cu(NO3)2 and orange-brown NO2 (Figure 3-51).
Figure 3-51. Reaction of concentrated HNO3 with Cu metal.

Nitric acid dissociates to H+ and NO3-. Of these two ions, the H+ is not a strong enough oxidizer to pull electrons off copper metal. (Otherwise, HCl, which also produces H+ on dissociation, would react with Cu, too, but it does not.) The species that oxidizes the copper is the nitrate ion, NO3- in which the N atom has the oxidation state of +5. In the product, NO2, the oxidation state of N is +4. Silver and mercury react with concentrated HNO3 similarly (Figure 3-52). For a demonstration of the reaction of concentrated HNO3 with Hg, watch this video. Note the formation of large quantities of orange-brown NO2 in this experiment.
Figure 3-52. Reactions of concentrated HNO3 with silver and mercury.

Now, pay special attention, please. We just watched two demonstrations of the reactions of concentrated HNO3, one with copper metal and the other with mercury metal. Both reactions produced NO2, a brown gas (Figures 3-51 and 3-52). Now watch videos of two more reactions of the same metals, Cu (click here) and Hg (click here), this time with dilute, not concentrated HNO3. In these demonstrations, we also see the evolution of a gas, but the gas is barely yellow or pale brown. The dissimilarity is due to the difference in concentrations of the nitric acid used in these two pairs of experiments. Copper and mercury react with dilute HNO3 to give NO rather than NO2, and so does Ag (Figure 3-53).
Figure 3-53. Reactions of dilute HNO3 with Hg, Cu, and Ag.

Unlike orange-brown NO2, NO is a colorless gas. In air, NO undergoes oxidation to NO2. (Remember that this transformation is involved in the industrial synthesis of nitric acid, see Figure 3-44.) The air-oxidation of NO to NO2 is quite facile but not instantaneous. Before the NO released in the reactions (Figure 3-53) is oxidized to NO2 to a considerable extent, it dissipates into the air. The NO2 eventually produced is quite diluted with the surrounding air, which makes the brown color less intense or even barely noticeable.

Gold and platinum, which are even poorer electron donors than Cu, Hg, and Ag (Figure 3-50), do not react with HNO3 of any concentration. As for the metals positioned to the left from the hydrogen in the reactivity series (Figure 3-50), they all react with nitric acid of various concentrations. And that is where all hell breaks loose. Being more powerful electron donors than Cu, Hg, and Ag, these metals are capable of reducing both the H+ and NO3- of HNO3 (Figure 3-54).
Figure 3-54. Possible products of reduction of HNO3 with metals.

The reduction of the H+ would give rise to H2 (Figure 3-54, left), exactly as in the reaction of an active enough metal with conventional acids, such as Zn with HCl (Figure 3-49). The formation of H2 in reactions of nitric acid with some metals is well-documented. These reactions are rare though. We will touch on them toward the end of this subsection. In most cases, it is the NO3- of nitric acid, not the H+, that is predominantly reduced by a metal.

The reduction of the NO3- of HNO3 (oxidation state of N = +5) by a metal can lead to a variety of products (Figure 3-54, right). All metals to the left from the H atom in the reactivity series (Figure 3-50) are capable of reducing HNO3 not only to NO2 (oxidation state +4) or NO (oxidation state +2) – the way Cu, Hg, and Ag do – but also to N2O (oxidation state +1), N2 (oxidation state 0), and NH4+ (oxidation state -3). In most instances, more than one reduction product are formed due to the simultaneously occurring multiple redox processes. What particular nitrogen compounds are produced, and in what ratio, depends on many factors. The two particularly important ones are the nature of the metal and the concentration of HNO3 used for the reaction. A variation in either or both of these parameters can strongly affect the outcome of the reaction.

Figure 3-55 summarizes the general trends for reactions of HNO3 of different concentrations toward different metals. As a general rule ⁠—

Lower concentrations of HNO3 favor deeper reduction of its N atom in reactions with metals


The higher the reactivity (electron-donating ability) of the metal, the deeper it can reduce the N atom of HNO3.
Figure 3-55. Deeper reduction of HNO3 is favored by higher dilution and stronger electron-donating metals.

As an example, zinc can reduce HNO3 to NO2, NO, N2O, N2, and NH4+. What products are predominantly formed strongly depends on how concentrated HNO3 is used for the reaction (Figure 3-56).
Figure 3-56. Reactions of Zn metal with HNO3 of different concentrations.

Watch a demonstration of the reaction of Zn with 70% HNO3 here. Note the formation of the brown NO2 gas and – this is important – that the intense color of the gas bubbling off at the beginning fades as the reaction occurs. At the onset of the reaction, the main reduction product is NO2 (Figure 3-56, top). As the reaction progresses, water is produced, which dilutes the nitric acid, thereby lowering its concentration. As a result of this dilution, a different reaction pathway is triggered, leading to NO (Figure 3-56, 2nd line from top). The reduction to NO gives rise to more water and, consequently, further dilution, which favors the formation of N2O. Even further dilution of the reaction solution with the H2O product entails the formation of N2 and eventually NH4+ (Figure 3-56, two bottom lines).

Like zinc, iron can reduce nitric acid to NO2, NO, N2O, N2, and NH4+. Again, the distribution of the products is strongly dependent on acid concentration (Figure 3-57).
Figure 3-57. Product distribution for the reaction of nitric acid with iron. Modified from source.
Digression. Do you have a problem reading the graph in Figure 3-57? If you do not, skip this paragraph. If you do, continue reading. The horizontal line, called the X coordinate, shows the span of HNO3 concentrations, from about 10% (left) to about 70% (right). Let us find the highest concentration on the X coordinate (70%) and move up vertically until we hit the first point on a curve. This point belongs to the NO curve. From this point continue moving up until you hit the next point on a curve. The next point is on the NO2 curve. By moving further up we will not encounter any more curves. The two points belonging to the NO curve and to the NO2 curve indicate that at that concentration the reaction of HNO3 with Fe gives rise to two reduction products, NO and NO2. From these two points move horizontally to the left until you hit the vertical line, the Y coordinate. The points of intersection with the Y coordinate are the values showing the percentages of NO (~10%) and NO2 (~90%) produced in the reaction. If we do the same thing for 50% HNO3, we will intersect three curves. The first intercept point is on the N2 curve, the second on the NO2 curve, and the third on the NO curve. Going strictly horizontally from these three points to the left gives us the percentages of N2O (~5%), NO2 (~30%) and NO (~65%). In this way, we can find out what nitrogen reduction products are formed and in what quantities for any HNO3 concentration within the range given on the X coordinate.
Given the general complexity of many reactions of nitric acid with metals, there has been too much controversy regarding what particular products are formed in one or another particular case. People even argue over what single product of NO3- reduction is yielded in the reaction of a certain metal with either "dilute" or "concentrated" HNO3. Such disputes do not make sense as long as the precise acid concentration remains unspecified and reliable experimental data for the reaction are unavailable, which is most often the case.

Now let us go back to Figure 3-54 and the paragraph following it, to recall that the proton of nitric acid is also a candidate for reduction by any metal located to the left from the hydrogen in the reactivity series (Figure 3-50). There has been a myth stating that hydrogen (H2) is never produced in any reaction of HNO3 of any concentration with any metal. While the reaction of Mg with concentrated nitric acid (70%) gives rise to orange-brown NO2 (watch this), highly diluted HNO3 (5%) reacts with magnesium to give hydrogen (Figure 3-58).
Figure 3-58. Reactions of Mg with highly diluted (5%) and concentrated (70%) HNO3.

Watch these Russian youngsters demonstrate (0:45 – 2:05) elegantly and convincingly the formation of colorless and flammable H2 in the reaction of magnesium with 5% HNO3. If you do not understand Russian, magnesium turnings were treated with 5% HNO3, the gas produced was collected over water in a test tube, and subjected to the classic hydrogen flame test. The small "pop" observed in the flame test indicated that the gas produced in the reaction of Mg with 5% HNO3 was H2.

The two equations in Figure 3-58 are yet another illustration of the dramatic effect of concentration on the outcome of the reaction between a metal and HNO3. Figure 3-58 displays just the two extreme cases where HNO3 is either very dilute (5%) or very concentrated (70%). At intermediate HNO3 concentrations, the entire spectrum of reduction products (H2, NO2, NO, N2O, N2, NH4+) can be expected.

Although the formation of H2 upon treatment of a metal with HNO3 is known, these cases are rare. Usually, it is the N atom of nitric acid that is predominantly reduced by a metal.

It is also often stated that some metals such as aluminum and iron do not react with HNO3 at all. That is only partly true. Both iron and aluminum are quite unreactive toward highly concentrated HNO3 but do react with more dilute nitric acid. This is persuasively demonstrated for iron in this video. Also, watch this demonstration to see for yourself that aluminum does not react with concentrated HNO3 and this video, showing that with more diluted nitric acid it does. Both aluminum and iron are passivated by undiluted nitric acid. The chemistry behind the passivation is quite complex; in simplified terms, the passivation deals with the formation on the metal surface of solid oxides and other substances that block access of the acid to the underlying metal. Aluminum containers and steel drums and railway tanks are used for storage and transportation of highly concentrated nitric acid.
Digression. I have met school teachers and college professors who would give a student an F if he or she said that H2 can be formed on treatment of a metal with HNO3, or that iron and aluminum react with nitric acid. Apparently, those teachers do not know what they are teaching. It has been known for a long time that some metals such as Mg and Mn react with highly diluted (5-10%) HNO3 to give H2 as the main reduction product. As is clear from this research paper published as early as 1897, aluminum of 99.6% purity does react with 25%, 55%, and 82% HNO3, the faster reaction being observed for less concentrated nitric acid. In another article, published in 1948, the resistance of aluminum to HNO3 of 93-99% concentration is shown as convincingly.
Nitric acid can oxidize some nonmetals, such as sulfur and phosphorus (Figure 3-59). These reactions occur only with concentrated HNO3 and produce NO2 because nonmetals are weaker reductants (electron donors) than metals.
Figure 3-59. Oxidation of sulfur and phosphorus with concentrated HNO3.

3.5.6. Thermal Decomposition of Nitrates. As mentioned above, one notable feature of nitrates, salts of nitric acid, is that they are all soluble in water. Another interesting characteristic of all nitrate salts is that they all pretty easily decompose on heating. There are four main reaction pathways for the thermal decomposition of nitrates.

- Alkali metal nitrates decompose to the corresponding nitrites (MNO2) and oxygen (Figure 3-60). Nitrites are salts of nitrous acid, HNO2. Watch a demonstration of the thermal decomposition of KNO3 in this video. Note how the potassium nitrate first melts and then decomposes on continuous heating to give O2.
Figure 3-60. Thermal decomposition of potassium nitrate.

- Most other metal nitrates decompose to metal oxides, NO2, and O2 (Figure 3-61). Watch the decomposition of calcium nitrate and copper nitrate and note the formation of brown NO2 and oxygen.
Figure 3-61. Thermal decomposition of calcium, iron, and copper nitrates.

- Mercury and silver nitrates decompose to the corresponding metal, NO2, and O2 (Figure 3-62).
Figure 3-62. Thermal decomposition of silver and mercury nitrates.

- Ammonium nitrate decomposes to N2O and H2O on heating (Figure 3-63).
Figure 3-63. Thermal decomposition of ammonium nitrate.

3.5.7. Phosphorus. Phosphorus Pentoxide. The story behind the discovery of phosphorus is fascinating. In the 17th century, there were many alchemists who tried to find the Philosopher's Stone, the magic substance that they believed was capable of turning conventional metals such as iron and lead into gold. The Philosopher's Stone had another name among alchemists, the Elixir of Life, as it was also believed to be capable of rejuvenating human beings and even making them immortal. In the 1660s, one Hennig Brandt, a German alchemist tried to make the Philosopher's Stone from human urine. Brandt was evaporating enormous volumes of urine that were easily accessible to him, thanks to the many beer drinkers and soldiers residing in the army barracks in Hamburg, where the alchemist lived. Brandt's experiments with the residues left after the evaporation of the urine eventually lead to the discovery of phosphorus.

"On that dark night our lone alchemist was having no luck with his latest experiments to find the philosopher's stone. Like many before him he had been investigating the golden stream, urine, and he was heating the residues from this which he had boiled down to a dry solid. He stoked his small furnace with more charcoal and pumped the bellows until his retort glowed red hot. Suddenly something strange began to happen. Glowing fumes filled the vessel and from the end of the retort dripped a shining liquid that burst into flames. Its pungent, garlic-like smell filled his chamber. When he caught the liquid in a glass vessel and stoppered it he saw that it solidified but continued to gleam with an eerie pale-green light and waves of flame seemed to lick its surface. Fascinated, he watched it more closely, expecting this curious cold fire to go out, but it continued to shine undiminished hour after hour. Here was magic indeed. Here was phosphorus." — J. Emsley "The 13th Element: The Sordid Tale of Murder, Fire, and Phosphorus", Wiley, 2000.

Figure 3-64 displays a fragment of the famous painting by Joseph Wright of Derby, portraying the moment of the discovery of phosphorus.
Figure 3-64. The Alchemist Discovering Phosphorus (1771; fragment) by Joseph Wright of Derby (source).
Digression. Having read about the discovery of phosphorus at the age of 12, I was fascinated by the story and shared it with a group of four neighbor boys who were a few years younger. A couple of days later, when one of those kids was home alone, he invited the other three boys to his apartment and suggested that they try to make phosphorus while his single mother was away at work. The idea was met with great enthusiasm and in no time each of them took a leak into a large cooking pot. The pot was then put on the kitchen stove to evaporate the contents. Shortly after the "raw material" began to boil, the neighbors were disturbed by the horrible smell spreading around and growing stronger every minute. The source of the stench was eventually identified and the young experimenters were profoundly encouraged to terminate their work. One neighbor in the apartment next door called the 9-year-old research leader's mother at work, begging her to come home as soon as possible to liquidate the consequences of her son and his friends' research activity. She did, after which the team of the four young phosphorus makers was quickly dismissed and order restored. The smell lasted for some time though. This is a true story. Sometimes I wonder if the kitchen pot employed in the miserably failed attempt to make phosphorus was ever used for cooking foods afterwards.
The glow of white phosphorus, the most common allotrope of phosphorus, has intrigued and fascinated people for centuries. This glowing, known as chemiluminescence, originates from the slow oxidation of white phosphorus in air. It was not until the mid-1970s, however, that intimate details of this light-emitting oxidation process were revealed and reported.

White phosphorus is made in industry by the high-temperature reduction of naturally occurring phosphate rock minerals such as Ca5(PO4)3F and Ca3(PO4)2 with coke (carbon) in the presence of sand, SiO2 (Figure 3-65).
Figure 3-65. Reactions used in industry to make phosphorus.

The molecule of white phosphorus is made up of four P atoms, displaying the perfect tetrahedral geometry (Figure 3-66). Although the molecular formula of white phosphorus is P4, the empirical formula P is conventionally used for writing chemical equations.
Figure 3-66. The structure of white phosphorus.

White phosphorus is highly flammable and, if exposed to air, may catch fire spontaneously. To avoid that, white phosphorus is conventionally stored under water. As can be seen from these two (1 and 2) demonstrations, the highly exothermic reaction of phosphorus with oxygen gives off a lot of bright light and white fumes. These fumes are phosphorus pentoxide, P2O5, a white solid produced in the reaction in the form of fine particles (Figure 3-67).
Figure 3-67. Reaction of phosphorus with oxygen and the structure of P4O10.

The formula P2O5 is an empirical formula. The molecular formula of phosphorus pentoxide is P4O10, in accord with the established structure presented in Figure 3-67. The most notable property of P2O5 is its enormous hygroscopicity, the exceptional capacity to avidly absorb moisture from the air. The reaction of P2O5 with water produces phosphoric acid, H3PO4 (Figure 3-68), and is highly exothermic. The heat released in this reaction is enough to make the water boil (watch the demonstration; consider turning the audio off).
Figure 3-68. Reaction of P2O5 with water.

The extreme affinity of P2O5 for water makes it a highly efficient desiccant. A desiccant is a drying agent, a hygroscopic material that is placed in a certain area to bring about and/or maintain dryness in its vicinity.

3.5.8. Phosphoric Acid. Being an acidic oxide, P2O5 reacts with water to give phosphoric acid (Figure 3-68). This reaction is used to make very high purity, food grade H3PO4 that is added to soft drinks to give them the characteristic "tingly" taste. First, white phosphorus is produced (Figure 3-65) and thoroughly purified by sublimation. The refined pure phosphorus is reacted with oxygen to make P2O5 (Figure 3-67), which is then treated with water (Figure 3-68) to make H3PO4 of high purity. This method, however, is too expensive to manufacture vastly larger quantities of H3PO4 needed to make phosphorus fertilizers.

The largest scale process for making H3PO4 is based on the reaction of the mineral apatite, Ca5(PO4)3F (phosphate rock), with sulfuric acid (Figure 3-69). To make large quantities of phosphoric acid, large quantities of sulfuric acid are needed. It is the production of phosphorus fertilizers from H3PO4 that dictates the humongous scale of sulfuric acid manufacturing.
Figure 3-69. Reaction used to make phosphoric acid on the largest scale.

Phosphoric acid is a medium strength electrolyte and is the only tribasic inorganic acid we encounter in this course. Depending on how much base is used for neutralization of H3PO4, three types of salts can be formed. As an example, reacting 1 mol of H3PO4 with 3, 2, and 1 mol of KOH gives rise to K3PO4 (potassium phosphate), K2HPO4 (potassium hydrogen phosphate), and KH2PO4 (potassium dihydrogen phosphate), respectively (Figure 3-70).
Figure 3-70. Neutralization of H3PO4 with various quantities of KOH.

The phosphate anion plays a critical role in the human body. Calcium salts of phosphoric acid are the main construction material for our bones and teeth. Up to 70% of the bone is hydroxyapatite, Ca5(PO4)3(OH), a compound similar to the mineral apatite, Ca5(PO4)3F, that is used to make phosphoric acid (Figure 3-69). Phosphate is also one of the three components of nucleotides, the building blocks of DNA and RNA. Derivatives of phosphoric acid are involved in many biochemical transformations in the body, including energy production, regulation of proteins, cell growth, right pH maintenance, heart contraction, nerve and muscle activity, and hormone signaling.

3.5.9. Exercises.

1. Nitrogen as a simple substance (N2) is abundant on earth, whereas phosphorus occurs naturally only in the form of compounds. Why? Answer

2. Provide an explanation for the extreme stability of N2. Answer

3. Write a balanced chemical equation for the synthesis of ammonia from N2 and H2 (the Haber-Bosch process). This reaction is (a) irreversible; (b) endothermic; (c) catalytic (conducted in the presence of a catalyst); (d) performed on a large industrial scale. Answer

4. The synthesis of ammonia from N2 and H2 in the presence of a catalyst (the Haber-Bosch process) is run at the highest possible temperatures and pressures. True or false? Answer

5. Pure ammonia is a (a) colorless odorless liquid; (b) water-soluble colorless gas with a strong pungent odor; (c) greenish-yellow gas that has a strong pungent odor and is poorly soluble in water; (d) colorless liquid with a strong unpleasant smell; (e) water-insoluble colorless gas with a strong pungent odor. Answer

6. A solution of ammonia in water is (a) neutral (pH = 7); (b) acidic (pH < 7); (c) basic (pH > 7). Answer

7. How is nitric acid made in industry? Describe the process and write balanced chemical equations for the reactions involved. [Answer: See section 3.5.2]

8. Nitric acid (HNO3) and phosphoric acid (H3PO4) are ... and ... acids, respectively. (a) strong, weak; (b) weak, strong; (c) medium strength, strong; (d) strong, medium strength; (e) strong, strong; (f) weak, weak. Answer

9. Write a balanced chemical equation for the decomposition of HNO3. [Answer: See Figure 3-45]

10. Unlike other acids such as HCl, nitric acid never gives rise to H2 when reacting with metals. True or false? Answer

11. What are the two main factors that control the outcome of the reaction of HNO3 with a metal? Answer

12. What are the nitrogen-containing products of reactions of HNO3 with Cu, Hg, or Ag? Write balanced chemical equations for these reactions. Answer

13. Predict the main nitrogen reduction product(s) of the reaction of zinc with (a) highly concentrated HNO3, (b) medium concentration HNO3, and (c) highly diluted HNO3. Answer

14. Aluminum and steel containers are used to store and transport nitric acid because Al and Fe do not react with HNO3. True or false? Answer

15. Laughing gas is (a) NO2; (b) N2O; (c) NO. Answer

16. Write balanced chemical equations for the thermal decomposition of (a) NaNO3; (b) Mg(NO3)2; (c) AgNO3; (d) NH4NO3. [Answer: See Figures 3-60 through 3-63]

17. Draw the structure of white phosphorus, P4. [Answer: See Figure 3-66]

18. Write balanced chemical equations for the combustion of phosphorus (P) in oxygen and for the reaction of phosphorus oxide produced with water. [Answer: See Figures 3-67 and 3-68]

19. How is phosphoric acid made in industry? [Answer: See subsection 3.5.9]

20. To a solution of 49 g of H3PO4 in water, was added an aqueous solution of 40 g of NaOH. What salt was formed in this reaction? Answer

21. Name at least three critical roles that phosphorus plays in the human body. [Answer: See section 3.5.8]