Hello, and welcome to today's chemistry lesson. In this session, we will explore one of the most important nitrogen compounds in industry and daily life: ammonia. We will examine its molecular structure, how it occurs in nature, various methods of preparation, its physical and chemical properties, and its wide-ranging applications. Let us begin our journey into the fascinating world of this pungent, alkaline gas.
Ammonia has the molecular formula NH₃, with a relative molecular mass of 17. In this molecule, one nitrogen atom forms three covalent bonds with three hydrogen atoms. The nitrogen atom carries a lone pair of electrons, which gives ammonia its characteristic basic properties.
Let us consider where ammonia occurs. In the free state, you will find small amounts of ammonia in air and traces in natural water. In the combined state, ammonia occurs in many compounds such as ammonium chloride, ammonium sulphate, and numerous other ammonium salts. Because ammonia and its compounds are highly soluble in water, they do not occur as minerals in the earth's crust.
You may have noticed a sharp, pungent smell near decaying organic matter or toilets; this is ammonia produced by bacterial decomposition of urea present in urine.
Ammonia exists in several forms that you should know. Gaseous ammonia refers to dry ammonia gas. Liquid ammonia is formed when dry ammonia is liquefied under high pressure. Liquor ammonia fortis is a saturated solution of ammonia in water, also called eight-eighty ammonia because of its relative density of zero point eight eight zero. This must be stored in tightly stoppered bottles in a cool place. In the laboratory, you will typically use a dilute solution of liquor ammonia as your reagent.
Now let us examine how ammonia is prepared. The general principle involves warming an ammonium salt with a caustic alkali such as slaked lime, caustic soda, or caustic potash. The reaction follows this pattern: an ammonium salt plus an alkali, when heated, produces a salt, water, and ammonia gas.
For laboratory preparation from ammonium chloride, you mix ammonium chloride with excess calcium hydroxide. The balanced equation is: 2NH₄Cl plus Ca(OH)₂ yields CaCl₂ plus 2H₂O plus 2NH₃. The apparatus consists of a round-bottom flask fitted in a slanting position, mouth downwards. This tilting is crucial: it prevents water formed during the reaction from trickling back into the heated flask and cracking the glass.
To obtain dry ammonia, you must pass the gas through quicklime, CaO. Other drying agents are unsuitable because ammonia, being basic, reacts with them. Concentrated sulphuric acid forms ammonium sulphate. Phosphorus pentoxide forms ammonium phosphate. Anhydrous calcium chloride forms an addition compound, CaCl₂·4NH₃.
Ammonia is collected by downward displacement of air because it is lighter than air, with a vapour density of eight point five compared to air's fourteen point four, and because it is highly soluble in water, making collection over water impossible.
Ammonia can also be prepared from metal nitrides, though this method is costly. Magnesium nitride, Mg₃N₂, reacts with warm water to form magnesium hydroxide and ammonia. Similarly, aluminium nitride, AlN, reacts with water to form aluminium hydroxide and ammonia.
To prepare aqueous ammonia, you dissolve ammonia gas in water. Because ammonia dissolves extremely rapidly, there is risk of back suction. To prevent this, use a funnel attached to the delivery tube with rubber tubing. Only a small portion of the funnel mouth dips into water. When pressure drops momentarily, water rushes into the funnel, breaking contact with the water surface. The incoming gas then pushes the water down, re-establishing contact. This ingenious arrangement allows continuous dissolution without dangerous back suction.
Now we turn to the industrial manufacture of ammonia through the Haber process. The reactants are nitrogen and hydrogen in a 1:3 ratio by volume. Nitrogen is obtained by fractional distillation of liquid air. Hydrogen comes from water gas through the Bosch process or from natural gas.
The reaction is: N₂ plus 3H₂ in equilibrium with 2NH₃ releasing heat. This reaction is reversible, exothermic, and proceeds with decrease in volume.
The favourable conditions are: temperature of 450 to 500 degrees Celsius, pressure above 200 atmospheres, a catalyst of finely divided iron, and a promoter of molybdenum or aluminium oxide. Only about fifteen percent of the reacting gases convert to ammonia in one pass.
Ammonia is separated from unreacted gases either by liquefaction, since ammonia liquefies more easily than nitrogen or hydrogen, or by absorption in water. The unreacted nitrogen and hydrogen are recirculated, eventually achieving 98% yield. The heat evolved maintains the temperature, so external heating is needed only initially. Purification of reactants is essential because impurities like carbon dioxide, carbon monoxide, and hydrogen sulphide poison the catalyst.
Let us examine the physical properties of ammonia. It is a colourless gas with a strong, pungent, choking odour and a slightly bitter, alkaline taste. It is non-poisonous in small amounts, but inhalation affects the respiratory system and causes tears. Large quantities can stimulate the heart dangerously and may cause death. Its vapour density is eight point five, making it lighter than air. It liquefies easily at 10 degrees Celsius under 6 atmospheres pressure. Liquid ammonia boils at minus thirty-three point five degrees Celsius and freezes at minus seventy-seven point seven degrees Celsius. Most remarkably, one volume of water dissolves about 702 volumes of ammonia at 20 degrees Celsius and 1 atmosphere pressure.
The fountain experiment beautifully demonstrates this extraordinary solubility. A round-bottomed flask filled with ammonia is inverted over a trough of red litmus solution. When a few drops of water enter through a dropper, the ammonia dissolves instantly, creating partial vacuum. External pressure forces the red litmus solution up through a jet tube, producing a spectacular blue fountain as the basic ammonia changes the indicator colour. Always cool liquor ammonia bottles in ice before opening, as the high internal pressure drops with cooling, preventing sudden flushing of gas.
Now we explore the chemical properties of ammonia. At high temperatures or with electric sparks, ammonia dissociates reversibly into nitrogen and hydrogen. Dry ammonia is neutral, but its aqueous solution is a weak base due to the lone pair on nitrogen accepting protons. Ammonium hydroxide dissociates partially to produce hydroxyl ions, OH⁻, making the solution basic. It turns red litmus blue, methyl orange yellow, and phenolphthalein pink.
Ammonia does not support combustion and extinguishes a burning splint. However, in pure oxygen, it burns with a greenish-yellow flame, producing nitrogen and water vapour, proving ammonia contains nitrogen and hydrogen. The reaction is: 4NH₃ plus 3O₂ yields 2N₂ plus 6H₂O. This reaction is irreversible and strongly exothermic.
Catalytic oxidation occurs when ammonia reacts with oxygen in the presence of platinum at 800 degrees Celsius. The reaction produces nitric oxide and water: 4NH₃ plus 5O₂ yields 6H₂O plus 4NO plus heat. The nitric oxide immediately oxidizes to brown nitrogen dioxide. The platinum catalyst continues glowing even after heating stops because this exothermic reaction sustains the temperature.
Ammonia acts as a reducing agent. When passed over heated copper II oxide, it reduces the black solid to reddish-brown copper metal, forming water vapour and nitrogen gas. With chlorine, ammonia acts as a reducing agent, reducing chlorine to hydrogen chloride. With excess ammonia, the hydrogen chloride produced immediately reacts with more ammonia to form dense white fumes of ammonium chloride. With excess chlorine, yellow nitrogen trichloride forms, a highly explosive liquid.
Ammonia reacts with acids to form ammonium salts. With hydrogen chloride gas, it forms ammonium chloride directly from two colourless gases. With nitric acid, it forms ammonium nitrate. With sulphuric acid, it forms ammonium sulphate.
Aqueous ammonia reacts with soluble metal salts to precipitate insoluble metal hydroxides, except for sodium and potassium. Iron II sulphate gives a dirty green precipitate of iron II hydroxide. Iron III chloride gives a reddish-brown precipitate of iron III hydroxide. Lead nitrate gives a white precipitate insoluble in excess. Zinc nitrate gives a white gelatinous precipitate soluble in excess, forming a complex ion. Copper sulphate gives a pale blue precipitate that dissolves in excess ammonia to form a deep azure blue solution of tetraamminecopper II sulphate. These characteristic colours and solubilities make ammonium hydroxide invaluable for identifying cations in qualitative analysis.
You can identify ammonia gas by its sharp characteristic odour, its ability to turn moist red litmus blue, moist turmeric paper brown, and phenolphthalein solution pink. It produces dense white fumes with concentrated hydrochloric acid. With copper sulphate solution, it first forms a blue precipitate that becomes deep blue with excess ammonia. All ammonium salts release ammonia when warmed with alkali. With Nessler's solution, ammonium salts produce brown colour or precipitate.
Finally, let us consider the many uses of ammonia. Liquid ammonia serves as an excellent refrigerant in ice plants because it requires 5700 calories per mole to vaporize at minus 33 degrees Celsius, taking this heat from surrounding bodies and cooling them. Unlike chlorofluorocarbons that deplete the ozone layer through free radical chain reactions, ammonia is environmentally compatible, does not deplete ozone, and does not contribute to global warming. It has superior thermodynamic properties, making ammonia refrigeration systems more energy efficient. Its recognizable odour ensures leaks are detected quickly, and being lighter than air, it rises and disperses rather than accumulating at ground level. The main disadvantages are incompatibility with copper piping and toxicity at high concentrations.
Aqueous ammonia emulsifies fats and grease, making it useful for removing perspiration stains from woollen clothes and for cleaning tiles and windows.
Industrially, ammonia is essential for manufacturing nitrogenous fertilizers including ammonium sulphate, diammonium hydrogen phosphate, ammonium nitrate, and urea. Urea forms when ammonia reacts with carbon dioxide at 150 degrees Celsius and 150 atmospheres pressure, a valuable nitrogenous fertiliser. Ammonia is also used to produce explosives like ammonium nitrate, which decomposes explosively leaving no residue. Other ammonium salts include ammonium carbonate used as smelling salts and ammonium chloride used in dry cells. Ammonia further serves in producing nylon, rayon, sodium cyanamide, plastics, dyes, organic chemicals, and wood pulp. It is crucial in the Solvay process for sodium carbonate and the Ostwald process for nitric acid manufacture.
Let us recap the key points from today's lesson. First, ammonia has the formula NH₃ with a trigonal pyramidal structure and a lone pair of electrons on nitrogen. Second, it is prepared in the laboratory by heating ammonium salts with alkalis, collected by downward displacement of air, and dried using quicklime only. Third, the Haber process manufactures ammonia industrially from nitrogen and hydrogen in 1 to 3 ratio by volume at 450 to 500 degrees Celsius and 200 to 900 atmospheres pressure, practically about 250 atmospheres, using finely divided iron catalyst with molybdenum or Al₂O₃ promoter. Fourth, ammonia is extremely soluble in water, forming a weakly basic solution that turns red litmus blue and produces characteristic coloured precipitates with metal salt solutions. Fifth, it acts as a reducing agent, burning in oxygen with a greenish-yellow flame and reducing heated metal oxides. Sixth, ammonia finds extensive use as a refrigerant, in fertilizer manufacture, in producing explosives and numerous chemicals, and as a laboratory reagent for identifying cations.
I hope this comprehensive exploration of ammonia has given you a solid understanding of this remarkable compound. Remember to review the preparation methods, the conditions for the Haber process, and the characteristic reactions that make ammonia so useful in both laboratory and industry. Keep practising the equations and observations, and you will master this topic with confidence. Until next time, stay curious and keep exploring the fascinating world of chemistry.