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⁺).

Figure 3-47. Humphry Davy administering laughing gas to a woman (source).

Figure 3-48. A laughing gas party in the 19^th 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 Video 3-20, 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 Video 3-21. Note the formation of large quantities of orange-brown NO₂ in this experiment.

Figure 3-52. Reactions of concentrated HNO₃ with silver and mercury.

Video 3-21. Reaction of mercury with concentrated nitric acid (source).


Now, pay special attention, please. We just watched two demonstrations of the reactions of concentrated HNO₃, one with copper metal (Video 3-20) and the other with mercury metal (Video 3-21). 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 (Video 3-22) and Hg (Video 3-23), 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

and

— 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 in Video 3-24. 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, 2^nd 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).

Video 3-24. Reaction of zinc with concentrated nitric acid (source).


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.

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 above 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 misconception, a myth 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 (Video 3-25), highly diluted HNO₃ (5%) reacts with magnesium to give hydrogen (Video 3-26). The two Russian youngsters in Video 3-26 (0:45 – 2:05) elegantly and convincingly demonstrate the formation of colorless and flammable H₂ in the reaction of magnesium with 5% HNO₃. In their experiment, magnesium turnings were treated with 5% HNO₃ and 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.

Video 3-26. Reaction of magnesium with highly diluted (5%) nitric acid gives rise to colorless and flammable H₂ (source).


Balanced chemical equations for both reactions are presented in Figure 3-58.

Figure 3-58. Reactions of Mg with highly diluted (5%) and concentrated (70%) HNO₃.


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 Video 3-27. Also, watch the demonstration in Video 3-28 to see for yourself that aluminum does not react with concentrated HNO₃ and Video 3-29 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.

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 Video 3-30. Note how the potassium nitrate first melts and then decomposes on continuous heating to give O₂.

Video 3-30. Decomposition of potassium nitrate on heating (source).


- Most other metal nitrates decompose to metal oxides, NO₂, and O₂ (Figure 3-61). When watching the decomposition of calcium nitrate (Video 3-31) and copper nitrate (Video 3-32) note the formation of brown NO₂ and oxygen.

Figure 3-61. Thermal decomposition of calcium, iron, and copper nitrates.

Video 3-32. Thermal decomposition of copper nitrate (source).


- 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 17^th 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).

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 the two demonstrations (Videos 3-33 and 3-34), 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 (Video 3-35).

Video 3-35. Highly exothermic reaction of P₂O₅ with water (source).


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]