Hello, and welcome to today's chemistry lesson. We are going to explore one of the most important industrial acids: nitric acid. By the end of this session, you will understand how nitric acid is prepared in the laboratory and on an industrial scale, its remarkable chemical properties as a powerful oxidising agent, and how we can identify it through simple tests.
Let us begin with the basics. Nitric acid has the molecular formula HNO₃ and a relative molecular mass of 63. In this compound, nitrogen exhibits its maximum valency of five.
Interestingly, nitric acid occurs naturally in trace amounts in rainwater after lightning strikes. Here is how this happens. During a lightning discharge, the intense heat—around 3000 degrees Celsius—causes nitrogen and oxygen from the atmosphere to combine, forming nitric oxide. This nitric oxide further oxidises to nitrogen dioxide, which then dissolves in atmospheric moisture in the presence of oxygen to form dilute nitric acid. The rain washes this acid down, where it combines with salts on the earth's surface to form nitrates. This natural conversion of atmospheric nitrogen into usable nitrogen compounds is called fixation of atmospheric nitrogen.
Nitric acid was historically known as aqua fortis, meaning strong water. This name reflects its remarkable ability to react with nearly all metals, converting them into soluble nitrates. Significantly, it can even dissolve silver—a metal that resists other common acids.
Now, let us examine how we prepare nitric acid in the laboratory.
The laboratory preparation uses a metal nitrate—either potassium nitrate, KNO₃, also called nitre, or sodium nitrate, NaNO₃, known as Chile saltpetre. These are heated with concentrated sulphuric acid in an all-glass apparatus.
The procedure involves mixing equal parts by weight of the metal nitrate and concentrated sulphuric acid in a glass retort. The mixture is gently heated to between 180 and 200 degrees Celsius. Sulphuric acid is non-volatile, so it displaces the volatile nitric acid, which distills over.
The chemical reactions are as follows. Potassium nitrate reacts with concentrated sulphuric acid below 200 degrees Celsius to produce potassium bisulphate and nitric acid. Similarly, sodium nitrate yields sodium bisulphate and nitric acid.
The nitric acid vapours are condensed by chilling the receiver with cold running water, yielding a light yellow liquid. Here is an important point: pure nitric acid is actually colourless. The yellow tint in laboratory preparations comes from dissolved nitrogen dioxide gas, which forms due to slight thermal decomposition of the acid.
This yellow colour can be removed in two ways. First, by bubbling dry air or carbon dioxide through the warm acid, which drives out the nitrogen dioxide. Second, by adding excess water, which dissolves the nitrogen dioxide.
Several precautions are essential in this preparation. All apparatus must be glass because nitric acid vapours attack rubber and cork. Concentrated hydrochloric acid cannot substitute for sulphuric acid because it is volatile and would contaminate the product. Most critically, the temperature must not exceed 200 degrees Celsius. At higher temperatures, sodium sulphate forms a hard crust that sticks to the retort walls, and the acid itself decomposes.
For large-scale production, we turn to Ostwald's process, developed by the German chemist Wilhelm Ostwald in 1914. This industrial method proceeds in three distinct steps.
Step one is the catalytic oxidation of ammonia. A mixture of dry ammonia and air in a ten-to-one ratio by volume is compressed and passed over platinum gauze catalyst at approximately 800 degrees Celsius. The ammonia oxidises to nitric oxide, releasing considerable heat. This exothermic reaction maintains the catalytic chamber temperature without external heating.
Step two occurs in the oxidation chamber. The hot gases are cooled, mixed with additional air, and passed through at about 50 degrees Celsius, where nitric oxide combines with oxygen to form nitrogen dioxide.
Step three is absorption in water. The nitrogen dioxide, along with excess air, enters a steel absorption tower packed with quartz stones. Warm water trickles from the top while the gases rise from the bottom. The nitrogen dioxide and oxygen react with water to form nitric acid. The acid collected at the bottom is concentrated above 50 percent.
Further distillation yields 68 percent nitric acid, known as concentrated nitric acid. By distilling this over concentrated sulphuric acid, we obtain 98 percent fuming nitric acid. Pure 100 percent nitric acid can be obtained as colourless crystals by cooling 98 percent acid to minus 42 degrees Celsius.
Let us now examine the properties of nitric acid, beginning with its physical characteristics.
Pure nitric acid is a colourless liquid with a suffocating odour and acidic taste. It is hygroscopic and fumes in moist air, so bottles must always be kept tightly stoppered. The density of 98 percent acid is 1.54 grams per cubic centimetre, while 68 percent commercial acid has a lower density of 1.42 grams per cubic centimetre. It boils at 86 degrees Celsius and melts at minus 42 degrees Celsius. The 68 percent aqueous solution forms a constant boiling mixture at 121 degrees Celsius. Nitric acid is miscible with water in all proportions.
Regarding physiological effects, nitric acid is non-poisonous but highly corrosive. It causes painful blisters on skin and produces a characteristic yellow stain by reacting with skin proteins to form xanthoproteic acid.
The chemical properties of nitric acid are particularly fascinating.
First, consider its stability. Pure nitric acid is unstable to heat and sunlight, decomposing to produce nitrogen dioxide, water, and oxygen. This decomposition explains why 100 percent acid is rarely used, and why nitric acid is stored in amber or brown bottles to exclude light.
Second, nitric acid is a strong monobasic acid. In aqueous solution, it ionises almost completely to produce hydrogen ions and nitrate ions. It turns blue litmus red, methyl orange pink, and leaves phenolphthalein colourless. It neutralises alkalis to form salts and water, reacts with metallic oxides and hydroxides to produce soluble nitrates, and liberates carbon dioxide from carbonates and bicarbonates.
Third, and most importantly, nitric acid is a powerful oxidising agent. This property arises from its decomposition to release nascent oxygen. Concentrated acid yields nitrogen dioxide and one atom of nascent oxygen, while dilute acid produces nitric oxide and three atoms of nascent oxygen.
With non-metals, concentrated nitric acid oxidises carbon to carbon dioxide and sulphur to sulphuric acid, itself being reduced to nitrogen dioxide in both cases.
The action on metals depends critically on concentration and temperature. Cold, dilute nitric acid oxidises metals like copper, zinc, and iron to their nitrates, liberating nitric oxide. Concentrated or hot dilute acid produces nitrogen dioxide instead.
Iron, aluminium, cobalt, and nickel exhibit an interesting phenomenon called passivity. When treated with concentrated nitric acid, these metals become inert due to formation of an extremely thin, insoluble oxide layer that prevents further reaction. This passivity can be removed by sanding the surface or using strong reducing agents.
Very dilute nitric acid—about one percent—reacts with magnesium and manganese to liberate hydrogen gas, because the oxidising action is sufficiently reduced. This reaction proves that nitric acid contains hydrogen.
Fourth, concentrated nitric acid mixed with concentrated hydrochloric acid in a one-to-three ratio by volume forms aqua regia, or royal water. Nitric acid oxidises hydrochloric acid to liberate nascent chlorine. This mixture can dissolve even noble metals like gold and platinum, forming soluble chlorides.
Nitric acid finds extensive practical applications. It etches designs on copper and brassware due to its solvent action on most metals. It purifies gold by dissolving metallic impurities. It serves as an oxidant in rocket fuels. It is essential in manufacturing nitrogen fertilisers including calcium nitrate, ammonium nitrate, and nitrochalk. Industrial applications include explosives like TNT, synthetic fibres, plastics, photographic film, dyes, drugs, and perfumes.
Finally, let us discuss how we identify nitric acid and nitrate salts.
Concentrated nitric acid produces brown fumes of nitrogen dioxide when heated. Most metal nitrates, when heated, decompose to give reddish-brown nitrogen dioxide. Adding copper to nitric acid or acidified nitrate solutions evolves dense reddish-brown nitrogen dioxide fumes.
The most distinctive test is the brown ring test. To perform this, add freshly prepared saturated iron two sulphate solution to an aqueous nitrate solution. Carefully pour concentrated sulphuric acid down the side of the test tube so it forms a separate layer. Cool the tube in water. A brown ring appears at the junction of the two liquids.
This brown ring is nitroso ferrous sulphate, formed when nitric oxide reacts with iron two sulphate. Fresh iron two sulphate is essential because atmospheric oxidation converts it to iron three sulphate, which will not give the test. The ring decomposes if the tube is disturbed due to heat evolution.
Let us briefly examine what happens when nitrates are heated. Sodium and potassium nitrates melt and decompose to nitrites with oxygen evolution. Most other nitrates decompose to metal oxides, nitrogen dioxide, and oxygen. Silver and mercury nitrates are exceptional, decomposing to the free metals instead of oxides. Ammonium nitrate decomposes explosively to nitrous oxide and water vapour, leaving no residue.
To summarise the key points from today's lesson.
First, nitric acid is prepared in the laboratory by heating potassium or sodium nitrate with concentrated sulphuric acid below 200 degrees Celsius, using all-glass apparatus.
Second, Ostwald's process manufactures nitric acid industrially through catalytic oxidation of ammonia over platinum at 800 degrees Celsius, followed by oxidation of nitric oxide and absorption in water.
Third, nitric acid is a powerful oxidising agent due to nascent oxygen release; its reaction with metals produces nitrogen oxides rather than hydrogen, except with very dilute acid and reactive metals.
Fourth, the brown ring test using iron two sulphate and concentrated sulphuric acid specifically identifies nitrate ions.
Fifth, heating behaviour of nitrates varies: alkali metal nitrates give nitrites and oxygen; most others give metal oxides, nitrogen dioxide, and oxygen; silver and mercury nitrates yield the free metals.
Sixth, aqua regia—a mixture of concentrated nitric and hydrochloric acids—can dissolve noble metals through nascent chlorine liberation.
That concludes our exploration of nitric acid. Remember to review the preparation methods, the oxidation reactions with different metals, and the distinctive tests for identification. Understanding the conditions that determine whether nitrogen monoxide or nitrogen dioxide is produced will help you predict reaction outcomes confidently. Keep practising the equations, and you will master this fascinating compound. Until next time, stay curious and keep experimenting safely.