Chemistry · Chemistry of the elements · Grade 8

Nitrogen, ammonia and nitric acid

How does nitrogen from the air become ammonia, saltpetre and fertiliser for cotton?

The situation

Three quarters of the air that enters our lungs with every breath is nitrogen. Everybody needs nitrogen for proteins, but neither cotton nor wheat can take it directly from the air: the N₂ molecule is extremely strong. The Navoiyazot, Ferganaazot and Maxam-Chirchiq plants make ammonia from the nitrogen in the air, and from it nitric acid and ammonium nitrate. How does inert nitrogen become reactive compounds, and how do they reach the field?

Topic

Nitrogen, ammonia and nitric acid. Navoiyazot and Maxam-Chirchiq: ammonia from atmospheric nitrogen, nitric acid from ammonia, and ammonium nitrate

Key ideas

The N≡N triple bond
Ammonia fountain
Ostwald process

Formulas

2 NH_{4} Cl + Ca (OH)_{2} \overset{\overset{}{t}}{\to} CaCl_{2} + 2 NH_{3} ↑ + 2 H_{2} O

Preparing ammonia in the laboratory (lab work 9).

NH_{3} + H_{2} O \overset{}{⇌} NH_{4}^{+} + OH^{-}

Aqueous ammonia is a weak base: phenolphthalein turns crimson.

4 NH_{3} + 5 O_{2} \overset{\overset{}{Pt - Rh}}{\to} 4 NO + 6 H_{2} O

Ostwald process, step 1: catalytic oxidation of ammonia.

3 Cu + 8 HNO_{3} \overset{}{\to} 3 Cu (NO_{3})_{2} + 2 NO ↑ + 4 H_{2} O

Dilute nitric acid dissolves copper: NO is released, not hydrogen.

NH_{3} + HNO_{3} \overset{}{\to} NH_{4} NO_{3} ω (N) = 35 %

Ammonium nitrate is a nitrogen fertiliser.

Theory

Nitrogen and the N≡N bond

Nitrogen (Z = 7, $2 s^{2} 2 p^{3}$ ) is a Group V A element; in its compounds its oxidation state ranges from −3 to +5. Nitrogen makes up 78.08 % of air by volume. In the N₂ molecule the atoms are joined by a triple bond (one σ and two π): 945 kJ/mol is needed to break it — almost twice as much as for O=O (498). So under ordinary conditions nitrogen is inert: crisp packets and food stores are filled with nitrogen. It boils at −195.8 °C (cooled to that temperature, it liquefies); industrially it is obtained by distilling liquid air, and in the laboratory by heating ammonium nitrite: $NH_{4} NO_{2} \overset{\overset{}{t}}{\to} N_{2} ↑ + 2 H_{2} O$ . At high temperature nitrogen becomes active: $N_{2} + O_{2} \overset{}{⇌} 2 NO$ (in lightning), $3 Mg + N_{2} \overset{\overset{}{t}}{\to} Mg_{3} N_{2}$ , $N_{2} + 3 H_{2} \overset{}{⇌} 2 NH_{3}$ (catalyst, pressure).

The nitrogen cycle in nature

Plants cannot take up N₂ directly — nitrogen must first be fixed: 1) lightning turns N₂ and O₂ into NO, and rain turns it into nitrate; 2) the nodule bacteria on the roots of legumes (lucerne, chickpeas, mung beans) turn N₂ into ammonia; 3) factories synthesise ammonia by the Haber–Bosch process. A plant absorbs NO₃⁻ and NH₄⁺ and builds protein; when remains decay, the nitrogen turns into NH₄⁺, and with the help of bacteria into NO₃⁻; denitrifying bacteria reduce nitrate back to N₂. Choose a process in front of the diorama.

Ammonia

NH₃ is a trigonal pyramid (H–N–H 106.7°); nitrogen has a lone pair of electrons. It is a colourless gas with a sharp smell, lighter than air (M = 17 g/mol < 29), boils at −33.3 °C and is easily liquefied under pressure (used in refrigerators). It is very soluble in water: at 20 °C 1 volume of water dissolves ≈ 700 volumes of NH₃ — the “fountain” experiment is based on this. Its solution is aqueous ammonia (the 10 % solution is “ammonia solution”), a weak base: $NH_{3} + H_{2} O \overset{}{⇌} NH_{4}^{+} + OH^{-}$ . With acids it forms ammonium salts: $NH_{3} + HCl \overset{}{\to} NH_{4} Cl$ — white “smoke” of tiny NH₄Cl crystals appears in the air. It burns in oxygen: $4 NH_{3} + 3 O_{2} \overset{}{\to} 2 N_{2} + 6 H_{2} O$ , and over platinum it is oxidised to NO. Lab work 9: a mixture of NH₄Cl and slaked lime is heated in a test tube (mouth slightly downward), and the ammonia is collected in a dry flask held mouth downward; damp red litmus turns blue, and a rod dipped in HCl gives white smoke. Detecting ammonium salts: when heated with alkali, NH₃ is released.

Nitrogen oxides and the Ostwald process

Nitrogen forms five oxides: N₂O (+1) and NO (+2) are neutral (they form no salts), and N₂O₃ (+3), NO₂ (+4) and N₂O₅ (+5) are acidic oxides. NO is colourless and in air at once turns into brown NO₂: $2 NO + O_{2} \overset{}{\to} 2 NO_{2}$ . NO₂ is poisonous and a cause of acid rain and smog; it forms in engines as it does in lightning. Nitric acid is made by the Ostwald process: 1) $4 NH_{3} + 5 O_{2} \overset{}{\to} 4 NO + 6 H_{2} O$ — on a Pt–Rh gauze at ≈ 850–900 °C; 2) $2 NO + O_{2} \overset{}{\to} 2 NO_{2}$ ; 3) $4 NO_{2} + O_{2} + 2 H_{2} O \overset{}{\to} 4 HNO_{3}$ . Navoiyazot and Maxam-Chirchiq make nitric acid and, from it, ammonium nitrate. In the laboratory a heated platinum spiral over ammonia vapour keeps itself glowing red — the reaction releases heat.

Nitric acid and nitrates

HNO₃ is a strong acid and a strong oxidising agent: with metals hydrogen is not released, and nitrogen is reduced. Concentrated: $Cu + 4 HNO_{3} \overset{}{\to} Cu (NO_{3})_{2} + 2 NO_{2} ↑ + 2 H_{2} O$ — a brown gas; dilute: $3 Cu + 8 HNO_{3} \overset{}{\to} 3 Cu (NO_{3})_{2} + 2 NO ↑ + 4 H_{2} O$ — colourless NO, which turns brown at the mouth of the test tube. Cold concentrated HNO₃ passivates iron and aluminium — it is carried in aluminium tankers. It decomposes in light ( $4 HNO_{3} \overset{}{\to} 4 NO_{2} + O_{2} + 2 H_{2} O$ ), so it is stored in dark glass. Ammonium nitrate: $NH_{3} + HNO_{3} \overset{}{\to} NH_{4} NO_{3}$ , ω(N) = 28 : 80 = 35 % (≈ 34 % in commercial fertiliser). Practical 4: the NH₄⁺ ion — when heated with alkali it gives NH₃ (damp litmus turns blue); the NO₃⁻ ion — when heated with copper and concentrated H₂SO₄ it gives brown NO₂.

Safety and the model

Safety: ammonia and nitrogen oxides are poisonous — work in a fume cupboard; when smelling a gas do not hold the test tube to your face, but waft the vapour towards you with your hand; concentrated HNO₃ and H₂SO₄ burn the skin — wear goggles and gloves; do not point the mouth of a heated test tube at yourself or others. Model: the rates of gas evolution are schematic (time is shortened); volumes are reduced to STP ( $V_{m}$ = 22.4 L/mol), and the gas in the flask is at 20 °C; the flask fills by displacement, the air mixing with the ammonia; in the fountain the ammonia dissolves completely, and the volume of water entering the flask equals the volume occupied by NH₃; the pH of aqueous ammonia is calculated with $K_{b}$ = 1.8·10⁻⁵; the NO → NO₂ process and the platinum temperature are schematic.

Did you know?

Nitrogen was discovered in 1772 by the Scottish chemist Daniel Rutherford. Lavoisier called it “azote” — Greek for “lifeless”: in this gas a candle goes out and animals cannot live.
The method of making nitric acid by oxidising ammonia over platinum was devised in the early 20th century by the German chemist Wilhelm Ostwald, who received the Nobel Prize in 1909 for his work on catalysis.

Inside the experiment

The experiment starts from a real-life situation: you predict the result, change the controls, run the experiment and draw a conclusion. Inside are missions, theory, a quiz and a guided tour in which the camera moves and explains the key points.

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