Chapter 10: Study of Compounds: C. Nitric Acid
In April 1940, German troops entered Copenhagen. Inside the Niels Bohr Institute, a Hungarian chemist named George de Hevesy had a problem sitting on his bench. Two solid gold Nobel medals had been sent there for safe keeping — one belonging to Max von Laue, the other to James Franck. Taking gold out of Germany was forbidden, and if the medals were found, the men who had smuggled them out would be in serious danger.
Burying them was too risky; the ground could be dug up. So de Hevesy did something only a chemist would think of. He mixed concentrated nitric acid with concentrated hydrochloric acid and lowered the medals into it. Gold is famous for resisting acids. It resists nitric acid on its own. It resists hydrochloric acid on its own. Yet in this mixture the metal slowly disappeared, leaving a clear orange solution in an unlabelled jar, which de Hevesy left standing on a shelf.
The soldiers searched the laboratory and walked past it.
When the war ended, de Hevesy returned, recovered the gold from that same jar and sent it to the Nobel Foundation. The medals were struck again from their own metal and returned to their owners.
A shelf, a jar, and a piece of chemistry you will meet in this chapter: what nitric acid can pull apart, and why.
Diagram
Fig. 10.0-1 — Gold Dissolved To Keep It Safe Simple scene illustration of a 1940s laboratory shelf: a plain unlabelled glass jar holding a clear orange-yellow solution, a gold medal shown half-dissolved inside it with fine bubbles rising from its surface; on the bench beside the shelf, two stoppered reagent bottles labelled concentrated nitric acid and concentrated hydrochloric acid, with a measuring cylinder between them. Add a small caption arrow to the jar reading ‘gold now in solution’.
10.1 Introduction
Nitric acid is one of the three mineral acids a chemistry laboratory keeps on its shelf, and it is by far the most restless of them. It attacks metals that other acids leave untouched, it stains skin yellow, and it slowly destroys itself in sunlight. All of that behaviour comes out of one small molecule, so it is worth looking at that molecule before anything else.
10.1.1 Formula, Structure And The Valency Of Nitrogen
The molecular formula of nitric acid is \(\mathrm{HNO_3}\). Its relative molecular mass is \(63\), which you can check by adding the relative atomic masses of the atoms in one molecule: \(1 + 14 + (3 \times 16) = 63\).
Inside the molecule, the hydrogen atom is not joined to nitrogen at all. It is joined to an oxygen atom, and that oxygen is joined to the nitrogen. The two remaining oxygen atoms are attached directly to the nitrogen — one by a double bond, and one by a coordinate bond, in which nitrogen supplies both the shared electrons. The structure is shown in Fig. 10.1-1.
Diagram
Fig. 10.1-1 — Structure Of The Nitric Acid Molecule Structural formula drawn left to right: H – O – N, with the nitrogen carrying a double bond (=) upward to one O and a coordinate bond drawn as an arrow (→) downward to a second O. Label the three parts: ‘hydroxyl oxygen carries the acidic hydrogen’, ‘double bond’, ‘coordinate bond, both electrons from nitrogen’. Below the structure add the line ‘nitrogen shows a valency of 5’.
Counting the bonds nitrogen has formed — one to the hydroxyl oxygen, two in the double bond, and one more in the coordinate bond — gives four bonds, and because the coordinate bond uses two electrons from nitrogen alone, nitrogen is said to show a valency of five in nitric acid. This is the highest valency nitrogen shows in any of its compounds.
That single hydrogen atom, held on the far side of an oxygen, is the only hydrogen the molecule has. It is the reason nitric acid is a monobasic acid, a point taken up in Section 10.4.3.
10.1.2 Occurrence In The Free And Combined States
A substance is said to occur in the free state when it is found in nature as itself, uncombined with anything else. Nitric acid does occur this way, though only in traces: rain water carries small amounts of it, formed in the air after a flash of lightning.
A substance occurs in the combined state when it is found locked up in a compound. Nitric acid occurs in this way as its salts, the nitrates — the compounds formed when the hydrogen of the acid is replaced by a metal. Three of these are worth naming.
- Chile saltpetre, sodium nitrate, \(\mathrm{NaNO_3}\)
- Nitre, potassium nitrate, \(\mathrm{KNO_3}\)
- Calcium nitrate, \(\mathrm{Ca(NO_3)_2}\)
The first two names are old ones, and both matter later in this chapter: nitre and Chile saltpetre are the very salts from which nitric acid is prepared in the laboratory.
10.1.3 Formation Of Nitric Acid In The Atmosphere
The air is roughly four-fifths nitrogen and one-fifth oxygen, yet the two gases sit side by side without reacting. Nitrogen molecules are held together by a very strong triple bond, and nothing in ordinary weather supplies enough energy to break it. A lightning discharge does. The sequence that follows is set out in FC 10.1.
Flowchart
FC 10.1 — Nitric Acid Formed In The Atmosphere Nitrogen and oxygen in air → lightning discharge at about 3000°C → nitric oxide (NO) → further oxidation by oxygen of the air → nitrogen dioxide (NO₂) → dissolves in rain water with oxygen present → nitric acid (HNO₃) in the free state → washed down by rain → combines with salts on the surface of the earth → nitrates in the soil
Each step of that chain can be written as an equation. First, during a lightning discharge the nitrogen of the air combines with oxygen to form nitric oxide.
nitrogen + oxygen → nitric oxide
\[\mathrm{N_2}\text{(g)} + \mathrm{O_2}\text{(g)} \xrightarrow{\ 3000^\circ\text{C}\ } 2\,\mathrm{NO}\text{(g)}\]
Nitric oxide is a colourless gas, and it does not stay as it is. It is oxidised further by the oxygen already present in the air, giving nitrogen dioxide.
nitric oxide + oxygen → nitrogen dioxide
\[2\,\mathrm{NO}\text{(g)} + \mathrm{O_2}\text{(g)} \rightarrow 2\,\mathrm{NO_2}\text{(g)}\]
Nitrogen dioxide is the reddish-brown gas you will meet again and again in this chapter. It dissolves in atmospheric moisture, or in the rain itself, and in the presence of the oxygen of the air it forms nitric acid in the free state.
nitrogen dioxide + water + oxygen → nitric acid
\[4\,\mathrm{NO_2}\text{(g)} + 2\,\mathrm{H_2O}\text{(l)} + \mathrm{O_2}\text{(g)} \rightarrow 4\,\mathrm{HNO_3}\text{(aq)}\]
The acid is washed down by the rain and combines with the salts present on the surface of the earth, so that what began as a flash in the sky ends as nitrate in the soil.
10.1.4 Fixation Of Atmospheric Nitrogen
Plants cannot take nitrogen out of the air. They can only take it up from the soil, as nitrate. The chain you have just followed therefore does something useful: it turns nitrogen the plant cannot use into nitrogen the plant can. Chemists have a name for that conversion.
Lightning is nature’s method of fixing nitrogen. Industry has its own, and you will meet it in Section 10.3, where nitric acid is manufactured from ammonia and air on a scale of millions of tonnes a year.
10.1.5 Aqua Fortis — The Older Name
Long before the acid had a formula, it had a name: aqua fortis, meaning strong water. The name was earned. Nitric acid reacts with nearly all metals and converts them into soluble nitrates, and it will even dissolve silver, which the other common acids leave alone.
10.2 Laboratory Preparation Of Nitric Acid
Nitric acid is not dug out of the ground. Its salts are, and a laboratory preparation therefore starts from a nitrate and takes the acid back out of it.
10.2.1 Reactants And The Principle Behind The Choice
Nitric acid is obtained by distilling concentrated sulphuric acid with either potassium nitrate, \(\mathrm{KNO_3}\), called nitre, or sodium nitrate, \(\mathrm{NaNO_3}\), called Chile saltpetre.
The choice of sulphuric acid is not accidental, and the reason lies in one word. A liquid is said to be volatile if it turns to vapour easily, at a fairly low temperature. Nitric acid is volatile; it boils at \(86^\circ\text{C}\). Sulphuric acid is non-volatile, boiling far higher, and this difference is what makes the preparation work.
When the two are heated together, the sulphuric acid stays behind in the flask while the nitric acid formed escapes as vapour. Because the nitric acid leaves the mixture as fast as it forms, the reaction keeps going in the forward direction. This is the general rule for such preparations: a less volatile acid drives a more volatile acid out of its salt.
10.2.2 Procedure And Apparatus
A mixture of equal parts by weight of potassium nitrate or sodium nitrate and concentrated sulphuric acid is placed in a glass retort. The mixture is heated gently, to a temperature between \(180^\circ\text{C}\) and \(200^\circ\text{C}\). The apparatus is shown in Fig. 10.2-1.
The retort’s long neck slopes downward into a receiver, which is kept cold by running water over it. Heating is gentle and even; a sand tray between the flame and the glass spreads the heat so that no part of the retort is heated fiercely. Reddish-brown fumes appear in the neck of the retort during the distillation, and Section 10.2.5 explains where they come from.
Diagram
Fig. 10.2-1 — Laboratory Preparation Of Nitric Acid Glass retort clamped on a stand over a sand tray heated by a burner, its long neck sloping downward into a round-bottomed receiver flask standing in a sink. Label: mixture of concentrated sulphuric acid and sodium or potassium nitrate inside the retort, sand tray, heat, brown fumes in the neck of the retort, tap delivering cold water running over the receiver, and the condensed nitric acid collecting in the receiver. Mark clearly that every joint and vessel is glass, with no rubber or cork anywhere.
10.2.3 Reactions
In words, the nitrate and the acid change partners: the sulphuric acid gives up one of its two hydrogen atoms to the nitrate, and takes the metal in exchange.
potassium nitrate + concentrated sulphuric acid → potassium hydrogen sulphate + nitric acid
\[\mathrm{KNO_3} + \mathrm{H_2SO_4} \xrightarrow{\ <200^\circ\text{C}\ } \mathrm{KHSO_4} + \mathrm{HNO_3}\]
sodium nitrate + concentrated sulphuric acid → sodium hydrogen sulphate + nitric acid
\[\mathrm{NaNO_3} + \mathrm{H_2SO_4} \xrightarrow{\ <200^\circ\text{C}\ } \mathrm{NaHSO_4} + \mathrm{HNO_3}\]
Notice that only one of the two hydrogen atoms of the sulphuric acid is used. The salt left behind is therefore a hydrogen sulphate — potassium hydrogen sulphate or sodium hydrogen sulphate — and not the normal sulphate. Table 10.1 sets the two reactions side by side.
| Table 10.1 — Reactions of the two nitrates with concentrated sulphuric acid | ||
|---|---|---|
| Nitrate taken | Residue left in the retort | Product distilled over |
| Nitre, \(\mathrm{KNO_3}\) | Potassium hydrogen sulphate, \(\mathrm{KHSO_4}\) | Nitric acid, \(\mathrm{HNO_3}\) |
| Chile saltpetre, \(\mathrm{NaNO_3}\) | Sodium hydrogen sulphate, \(\mathrm{NaHSO_4}\) | Nitric acid, \(\mathrm{HNO_3}\) |
10.2.4 Collection Of The Acid
The vapours of nitric acid pass down the neck of the retort into the receiver. There they are condensed to a light yellow liquid by chilling the receiver with running cold water.
The acid is collected simply by condensation, and no gas jar or water trough is needed, because what leaves the retort is the vapour of a liquid and not a gas.
10.2.5 The Yellow Colour And How It Is Removed
Pure nitric acid is colourless, yet the acid collected in the laboratory is slightly yellow. The colour is not the acid’s own. It comes from reddish-brown nitrogen dioxide gas dissolved in it — the same gas that showed as brown fumes in the neck of the retort.
That gas is produced because a portion of the nitric acid decomposes in the heat as it is being formed.
nitric acid → water + nitrogen dioxide + oxygen
\[4\,\mathrm{HNO_3} \rightarrow 2\,\mathrm{H_2O} + 4\,\mathrm{NO_2} + \mathrm{O_2}\]
The colour can be removed in either of two ways.
Dry air or carbon dioxide is bubbled through the yellow acid. The stream of gas drives the dissolved nitrogen dioxide out of the warm acid, which then turns colourless, and the nitrogen dioxide carried away is oxidised further to nitric acid.
Excess water is added. Nitrogen dioxide dissolves in water, so the yellow colour is removed. This reaction is simply the reverse of the decomposition of nitric acid written above.
10.2.6 Precautions
Three precautions govern this preparation, and each has a reason behind it.
All Glass, No Cork
- All-glass apparatus is used, because nitric acid vapours attack rubber and cork. A cork joint would be eaten away, the vapour would leak, and the acid collected would be contaminated.
Sulphuric Acid, Not Hydrochloric
- Concentrated hydrochloric acid is not used in place of concentrated sulphuric acid, because hydrochloric acid is itself volatile. It would boil off along with the nitric acid, and the vapours collected would carry hydrogen chloride into the receiver, so the product would not be pure nitric acid.
Keep Below 200 °C
- The temperature of the reaction should not exceed \(200^\circ\text{C}\). Above that temperature, the hydrogen sulphate reacts with a second portion of the nitrate, and the normal sulphate is formed instead.
\[\mathrm{NaNO_3} + \mathrm{NaHSO_4} \xrightarrow{\ >200^\circ\text{C}\ } \mathrm{Na_2SO_4} + \mathrm{HNO_3}\]
\[2\,\mathrm{NaNO_3} + \mathrm{H_2SO_4} \xrightarrow{\ >200^\circ\text{C}\ } \mathrm{Na_2SO_4} + 2\,\mathrm{HNO_3}\]
The yield of nitric acid is higher this way, so it may seem the better route. It is not, because sodium sulphate forms a hard crust that sticks to the walls of the retort and is very difficult to remove. The higher temperature brings three further troubles with it:
- it may damage the glass apparatus;
- decomposition of the nitric acid can also occur, lowering the purity of what is collected;
- fuel is wasted.
10.3 Manufacture Of Nitric Acid By The Ostwald Process
A retort and a burner will make enough nitric acid for a laboratory shelf. Industry needs it by the tonne, and it cannot afford to start from a mined nitrate. The Ostwald process starts from something far cheaper: ammonia and ordinary air.
The process was developed by the chemist Wilhelm Ostwald, who patented it in 1902 and saw it working on an industrial scale a few years later. It is carried out in three steps, in three separate chambers, and the whole plant is shown in Fig. 10.3-1. The route from ammonia to acid is summarised in FC 10.2.
Diagram
Fig. 10.3-1 — Manufacture Of Nitric Acid By The Ostwald Process Industrial plant drawn left to right in four linked units. Unit 1, tall catalytic chamber: two inlet pipes at the top labelled ‘ammonia gas, 1 volume’ and ‘dry air, 10 volumes, dust free’, inside it a horizontal layer of platinum gauze labelled ‘platinum catalyst’, marked 800°C; a heat exchanger drawn at its base. Unit 2, oxidising chamber, fed from the heat exchanger and from a separate air inlet, containing a cooling coil, marked 50°C. Unit 3, tall absorption tower packed with layers of quartz stone, water trickling in from a shower head at the top, marked ‘warm water’. Outlet at the base of the tower labelled ‘nitric acid’. Show arrows for the gas flow through all three units.
Flowchart
FC 10.2 — The Ostwald Process Dry ammonia (1 vol) + dry dust-free air (10 vols) → compressed → catalytic chamber, platinum gauze, 800°C → nitric oxide + steam (reaction exothermic) → heat exchanger cools the gases → more air added → oxidising chamber at 50°C → nitrogen dioxide → absorption tower packed with quartz, warm water trickling down → nitric acid (above 50%) at the base → distilled → concentrated nitric acid (68%)
10.3.1 Step I — Catalytic Oxidation Of Ammonia
A mixture of dry air and dry ammonia, in the ratio of \(10:1\) by volume, is first compressed. The air used must be free from carbon dioxide and from dust particles, because both would settle on the catalyst and spoil it. The compressed mixture is then passed into the catalytic chamber, which contains platinum gauze as the catalyst — a substance that speeds up a reaction without being used up in it — at a temperature of about \(800^\circ\text{C}\).
Here the ammonia is oxidised to nitrogen monoxide, better known as nitric oxide.
ammonia + oxygen → nitric oxide + water
\[4\,\mathrm{NH_3} + 5\,\mathrm{O_2} \xrightarrow{\ \mathrm{Pt},\ 800^\circ\text{C}\ } 4\,\mathrm{NO} + 6\,\mathrm{H_2O}\]
The reaction is exothermic, which means it gives out heat. That heat is not wasted. Once the reaction has started, the heat it releases is enough to keep the catalytic chamber at its working temperature, so no further heating is needed from outside.
The hot gases leaving the chamber are far too hot for the next step, so they are passed through a heat exchanger, which is simply a cooling chamber. They are then mixed with more air and passed into a second chamber, the oxidising chamber.
10.3.2 Step II — Oxidation Of Nitric Oxide
In the oxidising chamber, the nitric oxide combines with the oxygen of the added air to form nitrogen dioxide. The chamber is kept at about \(50^\circ\text{C}\).
nitric oxide + oxygen → nitrogen dioxide
\[2\,\mathrm{NO}\text{(g)} + \mathrm{O_2}\text{(g)} \xrightarrow{\ 50^\circ\text{C}\ } 2\,\mathrm{NO_2}\text{(g)}\]
10.3.3 Step III — Absorption Of Nitrogen Dioxide In Water
The nitrogen dioxide, together with the remaining air — and sometimes more air added at this stage — is passed into the absorption tower. This is a steel tower filled with quartz stone packed in layers, and warm water trickles down through it from the top.
As the gas rises through the packing and the water falls through it, the two meet over a very large area. The nitrogen dioxide and the oxygen present in the air react with the water to form nitric acid — the same change that produces nitric acid in the atmosphere after lightning, carried out here on purpose.
nitrogen dioxide + water + oxygen → nitric acid
\[4\,\mathrm{NO_2}\text{(g)} + 2\,\mathrm{H_2O}\text{(l)} + \mathrm{O_2}\text{(g)} \rightarrow 4\,\mathrm{HNO_3}\text{(aq)}\]
The nitric acid collected at the bottom of the tower is concentrated above \(50%\). On further distillation it gives \(68%\) nitric acid. This is the concentrated nitric acid that is supplied to laboratories.
10.3.4 Concentration Of Nitric Acid
Distillation alone will not take the acid past \(68%\), so a drying agent is used instead. Nitric acid of up to \(98%\) strength, called fuming nitric acid, is obtained by distilling the \(68%\) acid over concentrated sulphuric acid, which holds back the water.
Pure \(100%\) nitric acid is obtained by strongly cooling the \(98%\) acid. Colourless crystals of pure nitric acid separate out at \(-42^\circ\text{C}\).
10.3.5 Two Points To Note About The Process
A high ratio of air to ammonia is used. Oxygen makes up only about one-fifth of air, and oxygen is needed in all three chambers — to oxidise the ammonia, to oxidise the nitric oxide, and again in the absorption tower. Ten volumes of air per volume of ammonia supplies enough for all three.
Quartz is used as the packing because it is acid resistant, and when packed in layers it breaks the falling water into films and droplets. This helps the nitrogen dioxide dissolve uniformly in the water instead of rushing past it.
10.4 Properties Of Nitric Acid
Nitric acid leads a double life. In one set of reactions it behaves exactly as an acid should, neutralising bases and driving carbon dioxide out of carbonates. In another set it behaves as a powerful oxidising agent, and there its behaviour is nothing like that of hydrochloric or sulphuric acid. Both faces are described below, but first the acid has to be looked at as a liquid.
10.4.1 Physical Properties
Colour. Pure nitric acid of \(98%\) concentration is a colourless liquid. The commercial acid of \(68%\) concentration is yellowish brown, for a reason given in Section 10.4.2.
Odour. It has a suffocating smell.
Taste. It is acidic, that is, sour. No laboratory chemical is ever tasted; the property is stated because it is part of the description of an acid.
Nature. Nitric acid is hygroscopic — it absorbs moisture from the air — and it fumes in air. The bottle containing nitric acid should therefore always be kept stoppered.
Solubility. It is soluble in water in all proportions.
Physiological action. Nitric acid is not classed as a poison, but it is strongly corrosive. It attacks the skin and raises painful blisters, and it stains the skin a deep yellow. The stain appears because the acid reacts with the protein of the skin to form xanthoproteic acid.
Its fixed physical constants are collected in Table 10.2.
| Table 10.2 — Physical constants of nitric acid | |
|---|---|
| Property | Value |
| Density of the \(98%\) acid | \(1.54\ \text{g cm}^{-3}\) |
| Density of the commercial \(68%\) acid | \(1.42\ \text{g cm}^{-3}\) |
| Boiling point | \(86^\circ\text{C}\) |
| Melting point | \(-42^\circ\text{C}\) |
An aqueous solution of nitric acid of \(68%\) concentration forms a constant boiling mixture at \(121^\circ\text{C}\). A constant boiling mixture is one whose composition does not change as it boils, so the vapour coming off has the same strength as the liquid left behind. This is exactly why distillation cannot push the acid past \(68%\), and why the drying method of Section 10.3.4 has to be used instead.
10.4.2 Stability — Action Of Heat And Light
Pure nitric acid is unstable to heat and to sunlight. It decomposes, giving off reddish-brown nitrogen dioxide, and the acid takes on a yellow colour as some of that gas dissolves in it. This is the same decomposition you met while collecting the acid in Section 10.2.5.
nitric acid → nitrogen dioxide + water + oxygen
\[4\,\mathrm{HNO_3} \xrightarrow{\ \Delta\ \text{or sunlight}\ } 4\,\mathrm{NO_2} + 2\,\mathrm{H_2O} + \mathrm{O_2}\]
In the presence of sunlight, nitric acid decomposes even at room temperature. An acid standing in a clear bottle on a sunny bench slowly turns yellow, because the nitrogen dioxide formed dissolves in the acid around it. To avoid the decomposition, nitric acid is normally stored in coloured bottles.
This instability is also why the \(100%\) acid is not generally used. It decomposes so readily that it does not keep.
10.4.3 Acidic Properties Of Nitric Acid Solution
Nitric acid is a very strong monobasic acid — an acid whose molecule can give up only one hydrogen ion. Look back at Fig. 10.1-1 and you will see why: the molecule has just one hydrogen atom in it.
In aqueous solution it ionises almost completely, producing hydrogen ions and nitrate ions.
\[\mathrm{HNO_3}\text{(aq)} \rightleftharpoons \mathrm{H^{+}}\text{(aq)} + \mathrm{NO_3^{-}}\text{(aq)}\]
Those hydrogen ions are the acid’s acidity. Almost every molecule in the solution has released one, and that is what the word ‘strong’ means here. It has nothing to do with how much water the acid has been mixed with.
10.4.3.1 Action On Indicators
The hydrogen ions in the solution show themselves on any acid–base indicator.
- Blue litmus turns red.
- Methyl orange turns pink or red.
- Phenolphthalein remains colourless.
10.4.3.2 Reaction With Metallic Oxides And Hydroxides
All metallic oxides and hydroxides react with dilute nitric acid to form their respective nitrates and water only. Nothing is given off, because there is no gas to give off — a base and an acid simply neutralise each other.
Every nitrate formed dissolves. All nitrates are soluble in water, so the product of such a reaction is always a clear solution and never a precipitate. The reactions are set out in Table 10.3.
| Table 10.3 — Dilute nitric acid with metallic oxides and hydroxides | |
|---|---|
| Base | Equation |
| Potassium oxide | \(\mathrm{K_2O} + 2\,\mathrm{HNO_3} \rightarrow 2\,\mathrm{KNO_3} + \mathrm{H_2O}\) |
| Copper(II) oxide | \(\mathrm{CuO} + 2\,\mathrm{HNO_3} \rightarrow \mathrm{Cu(NO_3)_2} + \mathrm{H_2O}\) |
| Zinc oxide | \(\mathrm{ZnO} + 2\,\mathrm{HNO_3} \rightarrow \mathrm{Zn(NO_3)_2} + \mathrm{H_2O}\) |
| Sodium hydroxide | \(\mathrm{NaOH} + \mathrm{HNO_3} \rightarrow \mathrm{NaNO_3} + \mathrm{H_2O}\) |
| Iron(III) hydroxide | \(\mathrm{Fe(OH)_3} + 3\,\mathrm{HNO_3} \rightarrow \mathrm{Fe(NO_3)_3} + 3\,\mathrm{H_2O}\) |
| Magnesium hydroxide | \(\mathrm{Mg(OH)_2} + 2\,\mathrm{HNO_3} \rightarrow \mathrm{Mg(NO_3)_2} + 2\,\mathrm{H_2O}\) |
Read down the second column and you will see the pattern that decides the coefficients: a base with one hydroxide group needs one molecule of the acid, one with two hydroxide groups needs two, one with three needs three. The nitrate ion carries a single negative charge, so it must be supplied once for every unit of positive charge on the metal ion.
10.4.3.3 Reaction With Carbonates And Hydrogen Carbonates
Nitric acid reacts with carbonates and hydrogen carbonates to give the salt, water and carbon dioxide. The gas comes off as brisk effervescence, and it turns lime water milky. The reactions are given in Table 10.4.
| Table 10.4 — Dilute nitric acid with carbonates and hydrogen carbonates | |
|---|---|
| Carbonate or hydrogen carbonate | Equation |
| Calcium carbonate | \(\mathrm{CaCO_3} + 2\,\mathrm{HNO_3} \rightarrow \mathrm{Ca(NO_3)_2} + \mathrm{H_2O} + \mathrm{CO_2}\) |
| Sodium carbonate | \(\mathrm{Na_2CO_3} + 2\,\mathrm{HNO_3} \rightarrow 2\,\mathrm{NaNO_3} + \mathrm{H_2O} + \mathrm{CO_2}\) |
| Copper(II) carbonate | \(\mathrm{CuCO_3} + 2\,\mathrm{HNO_3} \rightarrow \mathrm{Cu(NO_3)_2} + \mathrm{H_2O} + \mathrm{CO_2}\) |
| Sodium hydrogen carbonate | \(\mathrm{NaHCO_3} + \mathrm{HNO_3} \rightarrow \mathrm{NaNO_3} + \mathrm{H_2O} + \mathrm{CO_2}\) |
| Calcium hydrogen carbonate | \(\mathrm{Ca(HCO_3)_2} + 2\,\mathrm{HNO_3} \rightarrow \mathrm{Ca(NO_3)_2} + 2\,\mathrm{H_2O} + 2\,\mathrm{CO_2}\) |
10.4.3.4 Reaction With Sulphites And Hydrogen Sulphites
Metallic sulphites and hydrogen sulphites react with dilute nitric acid to form their soluble metallic nitrates, water and sulphur dioxide gas. Sulphur dioxide has a sharp, choking smell, quite unlike the odourless carbon dioxide of the previous reaction. Table 10.5 gives the equations.
| Table 10.5 — Dilute nitric acid with sulphites and hydrogen sulphites | |
|---|---|
| Sulphite or hydrogen sulphite | Equation |
| Potassium sulphite | \(\mathrm{K_2SO_3} + 2\,\mathrm{HNO_3} \rightarrow 2\,\mathrm{KNO_3} + \mathrm{H_2O} + \mathrm{SO_2}\) |
| Copper(II) sulphite | \(\mathrm{CuSO_3} + 2\,\mathrm{HNO_3} \rightarrow \mathrm{Cu(NO_3)_2} + \mathrm{H_2O} + \mathrm{SO_2}\) |
| Calcium hydrogen sulphite | \(\mathrm{Ca(HSO_3)_2} + 2\,\mathrm{HNO_3} \rightarrow \mathrm{Ca(NO_3)_2} + 2\,\mathrm{H_2O} + 2\,\mathrm{SO_2}\) |
10.4.4 Nitric Acid As An Oxidising Agent
Everything so far could have been written about any strong acid. What follows could not. Nitric acid vigorously oxidises non-metals, metals, inorganic compounds and organic substances, and this is the property that sets it apart.
While it oxidises something else, the acid is itself reduced. Its reduction products differ — nitric oxide \(\mathrm{NO}\), nitrogen dioxide \(\mathrm{NO_2}\), dinitrogen oxide \(\mathrm{N_2O}\) — and which one appears depends on how concentrated the acid is and how hot the reaction is.
10.4.4.1 Where The Oxidising Power Comes From
The oxidising power is traced to nascent oxygen — oxygen in the instant of being set free, before it has paired up into \(\mathrm{O_2}\) molecules, and far more reactive in that state. Nitric acid gives nascent oxygen on decomposition, and it is written in square brackets as \([\mathrm{O}]\).
\[2\,\mathrm{HNO_3}\text{(conc.)} \rightarrow 2\,\mathrm{NO_2} + \mathrm{H_2O} + [\mathrm{O}]\]
\[2\,\mathrm{HNO_3}\text{(dil.)} \rightarrow 2\,\mathrm{NO} + \mathrm{H_2O} + 3[\mathrm{O}]\]
Notice the difference between the two lines. The concentrated acid gives nitrogen dioxide; the dilute acid gives nitric oxide. That single fact governs almost everything in the rest of this section.
That last sentence explains a puzzle you are about to meet. When a metal displaces hydrogen from nitric acid, the hydrogen does not escape as a gas. The nascent oxygen present oxidises it to water on the spot.
10.4.4.2 Action On Non-Metals
Concentrated nitric acid oxidises non-metals to their corresponding acidic oxides or oxyacids, and is itself reduced to nitrogen dioxide.
Carbon burns to carbon dioxide.
carbon + concentrated nitric acid → carbon dioxide + water + nitrogen dioxide
\[\mathrm{C} + 4\,\mathrm{HNO_3} \rightarrow \mathrm{CO_2} + 2\,\mathrm{H_2O} + 4\,\mathrm{NO_2}\]
Sulphur is oxidised all the way to sulphuric acid.
sulphur + concentrated nitric acid → sulphuric acid + water + nitrogen dioxide
\[\mathrm{S} + 6\,\mathrm{HNO_3} \rightarrow \mathrm{H_2SO_4} + 2\,\mathrm{H_2O} + 6\,\mathrm{NO_2}\]
In both reactions the non-metal gains oxygen and the nitrogen of the acid loses it. Dense reddish-brown fumes filling the tube are the visible sign that the acid has been reduced to nitrogen dioxide.
10.4.4.3 Action On Metals
Nitric acid reacts with all metals except gold and platinum. What is formed depends on the temperature and the concentration of the acid, and the two cases are compared in Table 10.6.
Cold and dilute nitric acid oxidises metals to their nitrates and liberates nitric oxide.
Concentrated nitric acid — or hot dilute nitric acid — liberates nitrogen dioxide instead.
Two details in that table are worth pausing over. With the dilute acid, iron ends up as iron(II) nitrate; with the concentrated acid, the stronger oxidising conditions carry it further, to iron(III) nitrate. And in no row is hydrogen among the products.
Nitric oxide is itself colourless, but it does not stay colourless for long. As soon as it reaches the air at the mouth of the tube it meets oxygen and turns brown, because it becomes nitrogen dioxide. Fig. 10.4-1 shows what the two reactions look like.
Diagram
Fig. 10.4-1 — Copper With Dilute And With Concentrated Nitric Acid Two test tubes side by side in a rack, each holding copper turnings. Left tube labelled ‘cold dilute nitric acid’: the solution is blue, a colourless gas rises from the metal, and the gas is shown turning brown just at the mouth of the tube where it meets air — label ‘nitric oxide, colourless; turns brown in air’. Right tube labelled ‘concentrated nitric acid’: vigorous reaction, blue-green solution, tube filled with dense reddish-brown fumes — label ‘nitrogen dioxide’. Both tubes shown inside a fume cupboard outline.
10.4.4.4 Passivity
Metals like iron, aluminium, cobalt and nickel become passive, that is inert, when treated with pure concentrated nitric acid. They stop reacting altogether, and the acid can then be stored in containers made of these metals.
The cause is an extremely thin layer of insoluble metallic oxide that the acid itself builds on the metal surface. This layer seals the metal off from the acid below it, and the reaction stops. The effect is called passivity.
Passivity can be removed by rubbing the surface layer away with sand paper, or by treating the metal with a strong reducing agent, which strips the oxide off chemically.
10.4.4.5 Aqua Regia — Action On Gold And Platinum
Concentrated nitric acid taken as one part by volume and mixed with concentrated hydrochloric acid taken as three parts by volume gives a mixture called aqua regia, meaning royal water. The name records what it does: it attacks the noble metals, gold and platinum, that no single acid will touch.
Neither acid on its own can do this. In the mixture, the nitric acid oxidises the hydrochloric acid to chlorine.
\[\mathrm{HNO_3} + 3\,\mathrm{HCl} \rightarrow \mathrm{NOCl} + 2\,\mathrm{H_2O} + 2[\mathrm{Cl}]\]
The \(\mathrm{NOCl}\) formed is nitrosyl chloride. What matters more is the \([\mathrm{Cl}]\) — nascent chlorine, and it is this that attacks the noble metals, converting them into soluble chlorides.
\[\mathrm{Pt} + 4[\mathrm{Cl}] \rightarrow \mathrm{PtCl_4}\]
\[\mathrm{Au} + 3[\mathrm{Cl}] \rightarrow \mathrm{AuCl_3}\]
Platinum(IV) chloride and gold(III) chloride are both soluble, so the metal passes quietly into solution and the liquid is left clear.
10.4.4.6 Reaction With Magnesium And Manganese
There is one case in which nitric acid does give hydrogen. Very dilute acid, of about \(1%\) strength, reacts with magnesium and with manganese at room temperature to give their nitrates and hydrogen gas. At this dilution the oxidising action of the acid is much reduced, so the hydrogen survives instead of being oxidised to water.
magnesium + very dilute nitric acid → magnesium nitrate + hydrogen
\[\mathrm{Mg} + 2\,\mathrm{HNO_3}\text{(v. dil.)} \rightarrow \mathrm{Mg(NO_3)_2} + \mathrm{H_2}\]
manganese + very dilute nitric acid → manganese nitrate + hydrogen
\[\mathrm{Mn} + 2\,\mathrm{HNO_3}\text{(v. dil.)} \rightarrow \mathrm{Mn(NO_3)_2} + \mathrm{H_2}\]
These two reactions prove that nitric acid contains hydrogen — a fact its other reactions with metals hide completely.
The whole of Section 10.4 can now be put in one sentence. Dilute nitric acid is generally considered a typical acid, except for its reaction with metals, since it does not liberate hydrogen from them. It is a powerful oxidising agent, and the nascent oxygen it forms oxidises that hydrogen to water.
10.5 Uses Of Nitric Acid
Every use of nitric acid rests on one of the properties you have just studied. It is worth naming the property alongside the use, because the two always go together. Table 10.7 sets them side by side.
In The Workshop And The Laboratory
It is used to etch designs on copper and brassware. The reason is that nitric acid acts as a solvent for a large number of metals, the noble metals excepted. Where the acid touches the metal, the metal dissolves away and a line is left behind; where it does not touch, the surface stays as it was. Fig. 10.5-1 shows how the design is protected.
It is used to purify gold. Gold may contain copper, silver, zinc or lead as impurities, and every one of these dissolves in nitric acid while the gold itself does not.
It is used in preparing aqua regia, the mixture that will dissolve even the noble metals.
Diagram
Fig. 10.5-1 — Etching A Design On A Copper Plate Three panels left to right. Panel 1: a flat copper plate coated with an even layer of wax, labelled ‘acid-resistant wax coating’. Panel 2: a fine steel needle scratching a pattern through the wax, labelled ‘design cut through to bare metal’. Panel 3: the plate lying in a shallow tray of dilute nitric acid, with the acid shown eating grooves only along the scratched lines, labelled ‘acid dissolves the exposed copper; the waxed surface is untouched’. Add a small enlarged inset of one etched groove in cross-section.
In Farming
- Nitric acid is used in the preparation of fertilisers. Four are named here, all of them nitrates:
- calcium nitrate, \(\mathrm{Ca(NO_3)_2}\)
- ammonium nitrate, \(\mathrm{NH_4NO_3}\)
- nitro chalk, a mixture of \(\mathrm{NH_4NO_3}\) and \(\mathrm{CaCO_3}\)
- basic calcium nitrate, \(\mathrm{CaO\!\cdot\!Ca(NO_3)_2}\)
In Industry
- Nitric acid acts as a rocket fuel oxidant. A rocket carries its own supply of oxygen, since there is none to be had outside, and an oxidising agent as powerful as nitric acid can provide it.
Nitric acid is also used industrially in the manufacture of:
I. explosives such as T.N.T.; II. synthetic fibres and materials such as artificial silk, nylon, celluloid, plastics and photographic film; III. important compounds such as the nitrates of potassium, ammonium and silver; IV. dyes, drugs and perfumes.
| Table 10.7 — Uses of nitric acid and the property behind each | |
|---|---|
| Use | Property of nitric acid involved |
| Etching designs on copper and brassware | Acts as a solvent for most metals |
| Purifying gold | Dissolves the base metal impurities but not gold |
| Preparing aqua regia | Oxidises hydrochloric acid to nascent chlorine |
| Rocket fuel oxidant | Powerful oxidising agent |
| Making nitrate fertilisers | Acidic character — forms nitrates with bases and carbonates |
10.6 Tests For Nitric Acid And Nitrates
Suppose you are handed an unlabelled bottle of a colourless liquid, or a jar of white crystals, and asked whether either contains nitric acid or a nitrate. Four tests will settle it, and every one of them depends on the same telltale product — reddish-brown nitrogen dioxide.
10.6.1 Brown Fumes From Concentrated Nitric Acid On Heating
Concentrated nitric acid gives brown fumes on heating. You met this decomposition as a nuisance in Section 10.2, where it spoiled the colour of the acid; here it is put to work as a test.
\[4\,\mathrm{HNO_3} \xrightarrow{\ \Delta\ } 2\,\mathrm{H_2O} + 4\,\mathrm{NO_2} + \mathrm{O_2}\]
The observation is the dense reddish-brown fumes of nitrogen dioxide. The conclusion is that the liquid is nitric acid.
10.6.2 Brown Fumes From Nitrates On Heating
Nitrates other than those of potassium, sodium and ammonium produce reddish-brown fumes of nitrogen dioxide when heated. The solid left behind is the metallic oxide.
metallic nitrate → metallic oxide + nitrogen dioxide + oxygen
\[\text{Metallic nitrate} \xrightarrow{\ \Delta\ } \text{Metallic oxide} + \mathrm{NO_2} + \mathrm{O_2}\]
The three exceptions matter. Potassium and sodium nitrates give off oxygen only, with no brown fumes at all, and ammonium nitrate leaves nothing behind. Section 10.7 takes these differences further.
10.6.3 Copper Added To Nitric Acid Or To An Acidified Nitrate
On adding copper to nitric acid, or to an acidified nitrate, and warming, dense reddish-brown fumes of nitrogen dioxide are evolved. This test works on a solution, where simply heating would tell you nothing.
\[\mathrm{Cu} + 4\,\mathrm{HNO_3} \xrightarrow{\ \Delta\ } \mathrm{Cu(NO_3)_2} + 2\,\mathrm{H_2O} + 2\,\mathrm{NO_2}\]
If the substance under test is a nitrate rather than the acid, concentrated sulphuric acid is added first. It sets nitric acid free from the nitrate, exactly as it did in the laboratory preparation, and the nitric acid then attacks the copper.
\[\mathrm{Cu} + 4\,\mathrm{NaNO_3} + 4\,\mathrm{H_2SO_4} \xrightarrow{\ \Delta\ } 4\,\mathrm{NaHSO_4} + \mathrm{Cu(NO_3)_2} + 2\,\mathrm{H_2O} + 2\,\mathrm{NO_2}\]
10.6.4 The Brown Ring Test
The brown ring test is the standard test for the nitrate ion, and it is the one an examiner is most likely to ask about. It is carried out on an aqueous solution of a nitrate, or of nitric acid itself. The steps are given in FC 10.4 and the result is shown in Fig. 10.6-1.
Flowchart
FC 10.4 — The Brown Ring Test Aqueous solution of the nitrate or of nitric acid in a test tube → add freshly prepared saturated solution of iron(II) sulphate → tilt the tube and run concentrated sulphuric acid gently down the inside wall, so that it does not fall through in drops → the heavy acid settles below as a separate layer → cool the test tube in water → a brown ring appears at the junction of the two liquids → nitrate confirmed
Diagram
Fig. 10.6-1 — The Brown Ring Test Two test tubes drawn side by side. Left tube, tilted, with a dropper resting against its inner wall delivering concentrated sulphuric acid down the side; label ‘concentrated sulphuric acid, run down the side of the tube’. Right tube, upright, showing three zones from top to bottom: an upper layer labelled ‘nitrate and iron(II) sulphate solution’, a distinct dark brown band at the junction labelled ‘brown ring, FeSO4.NO’, and a lower colourless layer labelled ‘concentrated sulphuric acid (denser, settles below)’. Add a small label ‘freshly prepared FeSO4 solution’ with an arrow to the upper layer.
The chemistry has two stages. The concentrated sulphuric acid sets nitric acid free from the nitrate, and the nitric acid then oxidises iron(II) sulphate to iron(III) sulphate, being itself reduced to nitric oxide.
\[6\,\mathrm{FeSO_4} + 3\,\mathrm{H_2SO_4} + 2\,\mathrm{HNO_3} \rightarrow 3\,\mathrm{Fe_2(SO_4)_3} + 4\,\mathrm{H_2O} + 2\,\mathrm{NO}\]
The nitric oxide so formed does not escape. It is absorbed by the iron(II) sulphate still present in the upper layer, and the two combine to give a brown compound called nitroso ferrous sulphate.
\[\mathrm{FeSO_4} + \mathrm{NO} \rightarrow \mathrm{FeSO_4\!\cdot\!NO}\]
It is this compound, formed only where the two liquids meet, that shows as the brown ring.
Three Points To Note
A freshly prepared ferrous sulphate solution is used, because on exposure to the atmosphere it is oxidised to ferric sulphate, and ferric sulphate will not give the brown ring. An old, standing solution simply fails to work.
The brown ring forms at the junction of the two liquids because the concentrated sulphuric acid is the heavier liquid and settles to the bottom, leaving the ferrous sulphate layer above it. The brown compound is produced where the two are in contact, and nowhere else.
The brown ring decomposes if the test tube is disturbed. Mixing the layers spreads the heat given out when concentrated sulphuric acid meets water, and that heat destroys the unstable brown compound. The tube is cooled in water beforehand for the same reason, and then left standing still.
10.7 Effects Of Heat On Nitrates
Every nitrate decomposes when it is heated strongly enough, and no two families of nitrate decompose in quite the same way. What is left in the tube — and what comes out of it — depends on how reactive the metal is. Four outcomes are possible, and they are set out as a branching scheme in FC 10.5.
Flowchart
FC 10.5 — Action Of Heat On Nitrates Heat a nitrate → branch 1: sodium or potassium nitrate → nitrite + oxygen only, no brown fumes → branch 2: nitrates of other metals such as calcium, zinc, lead, copper → metallic oxide + nitrogen dioxide + oxygen → branch 3: silver or mercury nitrate → the metal itself + nitrogen dioxide + oxygen → branch 4: ammonium nitrate → nitrous oxide + steam, no residue at all
10.7.1 Sodium And Potassium Nitrates
Sodium and potassium are the most reactive metals of the four groups, and their nitrates are the hardest to break down. On heating, they first melt into colourless liquids, and then decompose to give off oxygen gas. The nitrate is not destroyed completely; it loses one oxygen atom and becomes a nitrite.
sodium nitrate → sodium nitrite + oxygen
\[2\,\mathrm{NaNO_3} \xrightarrow{\ \Delta\ } 2\,\mathrm{NaNO_2} + \mathrm{O_2}\]
potassium nitrate → potassium nitrite + oxygen
\[2\,\mathrm{KNO_3} \xrightarrow{\ \Delta\ } 2\,\mathrm{KNO_2} + \mathrm{O_2}\]
The evidence is easy to see. When a glowing splinter is brought near the molten liquid, it bursts into flame, because the gas coming off is oxygen. No brown fumes appear at any stage, and this alone separates these two nitrates from every other nitrate in the chapter.
Sodium nitrite is slightly deliquescent, meaning it absorbs so much moisture from the air that it eventually dissolves in the water it has taken up. Potassium nitrite is a colourless crystalline solid.
10.7.2 Other Metal Nitrates
All other nitrates, except those of silver and mercury, decompose into their oxides, nitrogen dioxide and oxygen. The equations all follow one pattern, and the way to tell them apart is by the colour of the residue.
calcium nitrate → calcium oxide + nitrogen dioxide + oxygen
\[2\,\mathrm{Ca(NO_3)_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{CaO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\]
Calcium nitrate is colourless and crystalline, and it leaves a white residue of calcium oxide, with reddish-brown gas coming off.
zinc nitrate → zinc oxide + nitrogen dioxide + oxygen
\[2\,\mathrm{Zn(NO_3)_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{ZnO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\]
Zinc nitrate is very deliquescent and decrepitates slightly on heating. Its residue is the most distinctive of all: zinc oxide is yellow when hot and white when cold, so the colour changes as the tube cools.
lead nitrate → lead oxide + nitrogen dioxide + oxygen
\[2\,\mathrm{Pb(NO_3)_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{PbO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\]
Lead nitrate is colourless and crystalline, and it decrepitates on heating. Decrepitation is the sharp crackling you hear as crystals break apart in the heat — water and gas trapped inside them expand and split the crystal open. The residue is lead oxide, a yellow solid that fuses with the glass of the tube.
copper nitrate → copper oxide + nitrogen dioxide + oxygen
\[2\,\mathrm{Cu(NO_3)_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{CuO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\]
Copper nitrate is blue and crystalline, and the copper oxide left behind is black.
10.7.3 Silver And Mercury Nitrates
Silver and mercury are unreactive metals, and their oxides are so unstable that they do not survive the heating. These two nitrates therefore go one step further than the rest: they decompose into the metals themselves, along with nitrogen dioxide and oxygen. Both salts are white and corrosive.
silver nitrate → silver + nitrogen dioxide + oxygen
\[2\,\mathrm{AgNO_3} \xrightarrow{\ \Delta\ } 2\,\mathrm{Ag} + 2\,\mathrm{NO_2} + \mathrm{O_2}\]
Silver nitrate is white and corrosive, and it leaves a silver-grey residue of the metal.
mercuric nitrate → mercury + nitrogen dioxide + oxygen
\[\mathrm{Hg(NO_3)_2} \xrightarrow{\ \Delta\ } \mathrm{Hg} + 2\,\mathrm{NO_2} + \mathrm{O_2}\]
Mercuric nitrate is white and crystalline, and the mercury it leaves condenses on the cool upper walls of the tube as a silver mirror.
10.7.4 Ammonium Nitrate
Ammonium nitrate is the odd one out, because its cation is not a metal at all. Both the ammonium ion and the nitrate ion contain nitrogen, and on heating the two react with each other. Ammonium nitrate decomposes explosively, leaving behind no residue whatsoever.
ammonium nitrate → nitrous oxide + steam
\[\mathrm{NH_4NO_3} \xrightarrow{\ \Delta\ } \mathrm{N_2O}\text{(g)} + 2\,\mathrm{H_2O}\text{(vap.)}\]
The gas produced, nitrous oxide, is known as laughing gas. Because both products are gases, the tube is left completely empty — a result no other nitrate gives.
The residues are worth looking at as well as reading about, and Fig. 10.7-1 shows three of them side by side.
Diagram
Fig. 10.7-1 — Residues Left When Three Nitrates Are Heated Three hard-glass test tubes held in tongs over separate burner flames, drawn side by side. Tube 1 labelled ‘potassium nitrate’: molten colourless liquid, no coloured gas above it, and a glowing splinter held at the mouth bursting into flame; label ‘oxygen only’. Tube 2 labelled ‘lead nitrate’: crystals cracking, reddish-brown fumes filling the tube, yellow solid residue; labels ‘decrepitation’, ‘nitrogen dioxide’, ‘lead oxide, yellow’. Tube 3 labelled ‘silver nitrate’: reddish-brown fumes and a grey metallic residue at the bottom; label ‘silver’.
10.8 Summary
Nitric acid has the molecular formula \(\mathrm{HNO_3}\) and a relative molecular mass of 63. Its single hydrogen atom is joined to an oxygen atom, not to nitrogen, and nitrogen shows a valency of 5 in the molecule.
In the free state nitric acid occurs in traces in rain water after lightning. In the combined state it occurs as nitrates — Chile saltpetre \(\mathrm{NaNO_3}\), nitre \(\mathrm{KNO_3}\) and calcium nitrate \(\mathrm{Ca(NO_3)_2}\).
In the atmosphere, a lightning discharge makes nitrogen and oxygen combine to nitric oxide; this is oxidised to nitrogen dioxide, which dissolves in rain water with oxygen to form nitric acid. \[\mathrm{N_2} + \mathrm{O_2} \xrightarrow{\ 3000^\circ\text{C}\ } 2\,\mathrm{NO}\] \[2\,\mathrm{NO} + \mathrm{O_2} \rightarrow 2\,\mathrm{NO_2}\] \[4\,\mathrm{NO_2} + 2\,\mathrm{H_2O} + \mathrm{O_2} \rightarrow 4\,\mathrm{HNO_3}\]
The conversion of free atmospheric nitrogen into useful nitrogenous compounds in the soil is called fixation of atmospheric nitrogen.
Nitric acid was formerly called aqua fortis, meaning strong water, because it reacts with nearly all metals and converts them into soluble nitrates, and dissolves even silver.
In the laboratory, nitric acid is prepared by distilling concentrated sulphuric acid with potassium nitrate or sodium nitrate at \(180^\circ\text{C}\)–\(200^\circ\text{C}\) in a glass retort. Sulphuric acid is chosen because it is non-volatile and drives out the volatile nitric acid. \[\mathrm{KNO_3} + \mathrm{H_2SO_4} \xrightarrow{\ <200^\circ\text{C}\ } \mathrm{KHSO_4} + \mathrm{HNO_3}\] \[\mathrm{NaNO_3} + \mathrm{H_2SO_4} \xrightarrow{\ <200^\circ\text{C}\ } \mathrm{NaHSO_4} + \mathrm{HNO_3}\]
The vapours are condensed to a light yellow liquid by chilling the receiver with running cold water.
The acid collected is yellow because some of it decomposes in the heat and the nitrogen dioxide formed dissolves in it. The colour is removed by bubbling dry air or carbon dioxide through the acid, or by adding excess water.
Three precautions govern the preparation: all-glass apparatus, because the vapours attack rubber and cork; concentrated sulphuric acid rather than concentrated hydrochloric acid, because hydrochloric acid is volatile; and a temperature not above \(200^\circ\text{C}\), because sodium sulphate then forms a hard crust on the retort, the apparatus may be damaged, the acid may decompose, and fuel is wasted.
Nitric acid is manufactured by the Ostwald process in three steps. Dry ammonia and dry dust-free air, in the ratio 1 : 10 by volume, are passed over platinum gauze at about \(800^\circ\text{C}\) in the catalytic chamber; the reaction is exothermic and its own heat maintains the temperature. \[4\,\mathrm{NH_3} + 5\,\mathrm{O_2} \xrightarrow{\ \mathrm{Pt},\ 800^\circ\text{C}\ } 4\,\mathrm{NO} + 6\,\mathrm{H_2O}\]
The gases are cooled in a heat exchanger, mixed with air, and passed into the oxidising chamber at about \(50^\circ\text{C}\). \[2\,\mathrm{NO}\text{(g)} + \mathrm{O_2}\text{(g)} \xrightarrow{\ 50^\circ\text{C}\ } 2\,\mathrm{NO_2}\text{(g)}\]
In the absorption tower, a steel tower packed with layers of quartz down which warm water trickles, the nitrogen dioxide and oxygen react with water to give nitric acid of above \(50%\) strength; distillation raises this to \(68%\), which is concentrated nitric acid. \[4\,\mathrm{NO_2}\text{(g)} + 2\,\mathrm{H_2O}\text{(l)} + \mathrm{O_2}\text{(g)} \rightarrow 4\,\mathrm{HNO_3}\text{(aq)}\]
Distilling the \(68%\) acid over concentrated sulphuric acid gives \(98%\) fuming nitric acid; strongly cooling that acid to \(-42^\circ\text{C}\) gives colourless crystals of the \(100%\) acid.
A high ratio of air is used because oxygen is only about one-fifth of air and is needed in all three chambers. Quartz is used as packing because it is acid resistant and helps the nitrogen dioxide dissolve uniformly.
Physically, pure nitric acid is a colourless liquid of density \(1.54\ \text{g cm}^{-3}\), boiling at \(86^\circ\text{C}\) and melting at \(-42^\circ\text{C}\). It has a suffocating smell and a sour taste, is hygroscopic, fumes in air, and is soluble in water in all proportions. The commercial \(68%\) acid is yellowish brown, of density \(1.42\ \text{g cm}^{-3}\), and forms a constant boiling mixture at \(121^\circ\text{C}\).
Nitric acid is corrosive to skin, raising painful blisters and staining it yellow by forming xanthoproteic acid with the protein of the skin.
Nitric acid is unstable to heat and to sunlight, and is therefore stored in coloured bottles. \[4\,\mathrm{HNO_3} \xrightarrow{\ \Delta\ \text{or sunlight}\ } 4\,\mathrm{NO_2} + 2\,\mathrm{H_2O} + \mathrm{O_2}\]
Nitric acid is a very strong monobasic acid, ionising almost completely in water. \[\mathrm{HNO_3}\text{(aq)} \rightleftharpoons \mathrm{H^{+}} + \mathrm{NO_3^{-}}\]
As an acid it turns blue litmus red and methyl orange pink or red, and leaves phenolphthalein colourless. With metallic oxides and hydroxides it gives a nitrate and water only; with carbonates and hydrogen carbonates it gives a nitrate, water and carbon dioxide; with sulphites and hydrogen sulphites it gives a nitrate, water and sulphur dioxide. All nitrates are soluble in water.
Nitric acid is a powerful oxidising agent because it gives nascent oxygen on decomposition. The concentrated acid is reduced to nitrogen dioxide, the dilute acid to nitric oxide. \[2\,\mathrm{HNO_3}\text{(conc.)} \rightarrow 2\,\mathrm{NO_2} + \mathrm{H_2O} + [\mathrm{O}]\] \[2\,\mathrm{HNO_3}\text{(dil.)} \rightarrow 2\,\mathrm{NO} + \mathrm{H_2O} + 3[\mathrm{O}]\]
Concentrated nitric acid oxidises non-metals, giving nitrogen dioxide itself. \[\mathrm{C} + 4\,\mathrm{HNO_3} \rightarrow \mathrm{CO_2} + 2\,\mathrm{H_2O} + 4\,\mathrm{NO_2}\] \[\mathrm{S} + 6\,\mathrm{HNO_3} \rightarrow \mathrm{H_2SO_4} + 2\,\mathrm{H_2O} + 6\,\mathrm{NO_2}\]
Nitric acid attacks all metals except gold and platinum. Cold dilute acid gives the nitrate and nitric oxide; concentrated acid, or hot dilute acid, gives the nitrate and nitrogen dioxide. With iron, the dilute acid gives iron(II) nitrate and the concentrated acid iron(III) nitrate. \[3\,\mathrm{Cu} + 8\,\mathrm{HNO_3}\text{(dil.)} \rightarrow 3\,\mathrm{Cu(NO_3)_2} + 4\,\mathrm{H_2O} + 2\,\mathrm{NO}\] \[\mathrm{Cu} + 4\,\mathrm{HNO_3}\text{(conc.)} \rightarrow \mathrm{Cu(NO_3)_2} + 2\,\mathrm{H_2O} + 2\,\mathrm{NO_2}\]
Iron, aluminium, cobalt and nickel are rendered passive by pure concentrated nitric acid, because a very thin layer of insoluble oxide forms on the surface and stops the reaction. Passivity is removed by rubbing with sand paper or by treating with a strong reducing agent.
Aqua regia is one part of concentrated nitric acid to three parts of concentrated hydrochloric acid by volume. The nitric acid oxidises the hydrochloric acid to nascent chlorine, which converts gold and platinum into their soluble chlorides. \[\mathrm{HNO_3} + 3\,\mathrm{HCl} \rightarrow \mathrm{NOCl} + 2\,\mathrm{H_2O} + 2[\mathrm{Cl}]\]
Very dilute nitric acid, of about \(1%\) strength, is the one case that gives hydrogen with a metal, reacting so with magnesium and manganese and proving that the acid contains hydrogen. In all other reactions with metals the nascent oxygen oxidises that hydrogen to water, which is why dilute nitric acid is a typical acid in every respect except its reaction with metals. \[\mathrm{Mg} + 2\,\mathrm{HNO_3}\text{(v. dil.)} \rightarrow \mathrm{Mg(NO_3)_2} + \mathrm{H_2}\]
Nitric acid is used to etch designs on copper and brassware, to purify gold, to prepare aqua regia, as a rocket fuel oxidant, and in making nitrate fertilisers such as calcium nitrate, ammonium nitrate, nitro chalk and basic calcium nitrate. Industrially it is used to manufacture explosives such as T.N.T., synthetic fibres and plastics, nitrates of potassium, ammonium and silver, and dyes, drugs and perfumes.
Nitric acid and nitrates are identified by four tests: the concentrated acid gives brown fumes on heating; most nitrates give reddish-brown fumes of nitrogen dioxide on heating; copper warmed with the acid, or with an acidified nitrate, gives dense reddish-brown fumes; and the brown ring test gives a brown ring of nitroso ferrous sulphate at the junction of the liquids. \[6\,\mathrm{FeSO_4} + 3\,\mathrm{H_2SO_4} + 2\,\mathrm{HNO_3} \rightarrow 3\,\mathrm{Fe_2(SO_4)_3} + 4\,\mathrm{H_2O} + 2\,\mathrm{NO}\] \[\mathrm{FeSO_4} + \mathrm{NO} \rightarrow \mathrm{FeSO_4\!\cdot\!NO}\]
In the brown ring test the ferrous sulphate solution must be freshly prepared, because air oxidises it to ferric sulphate, which gives no ring; the ring forms at the junction because the denser sulphuric acid settles below; and the ring is destroyed if the tube is disturbed.
On heating, sodium and potassium nitrates give the nitrite and oxygen only, with no brown fumes. \[2\,\mathrm{KNO_3} \xrightarrow{\ \Delta\ } 2\,\mathrm{KNO_2} + \mathrm{O_2}\]
Other metal nitrates, such as those of calcium, zinc, lead and copper, give the metallic oxide, nitrogen dioxide and oxygen. Zinc oxide is yellow when hot and white when cold; lead oxide is yellow and fuses with glass; copper oxide is black; calcium oxide is white. \[2\,\mathrm{Pb(NO_3)_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{PbO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\]
Silver and mercury nitrates give the free metal, nitrogen dioxide and oxygen. \[2\,\mathrm{AgNO_3} \xrightarrow{\ \Delta\ } 2\,\mathrm{Ag} + 2\,\mathrm{NO_2} + \mathrm{O_2}\]
Ammonium nitrate decomposes explosively into nitrous oxide and steam, leaving no residue at all. \[\mathrm{NH_4NO_3} \xrightarrow{\ \Delta\ } \mathrm{N_2O}\text{(g)} + 2\,\mathrm{H_2O}\text{(vap.)}\]
Flowchart
FC 10.6 — The Nitrogen Trail Followed In This Chapter Atmospheric nitrogen → fixed by lightning, or converted to ammonia and fed to an Ostwald plant → nitric oxide → nitrogen dioxide → nitric acid → reacts as an acid with bases, carbonates and sulphites, and as an oxidising agent with metals and non-metals → nitrates → heated, the nitrates give back nitrogen dioxide and oxygen; spread on fields, they return nitrogen to the soil
10.9 Key Terms
- Absorption tower — the steel tower packed with layers of quartz, down which warm water trickles, in which nitrogen dioxide and oxygen are absorbed to give nitric acid.
- Aqua fortis — the old name for nitric acid, meaning strong water, given because it dissolves nearly all metals, silver among them.
- Aqua regia — a mixture of one part concentrated nitric acid to three parts concentrated hydrochloric acid by volume, which dissolves gold and platinum.
- Brown ring test — the test for a nitrate, in which a brown ring forms at the junction of the liquids when concentrated sulphuric acid is run down the side of a tube containing the solution and freshly prepared iron(II) sulphate.
- Catalyst — a substance that speeds up a chemical reaction without being used up in it.
- Catalytic chamber — the first chamber of the Ostwald process, holding platinum gauze at about 800 °C, in which ammonia is oxidised to nitric oxide.
- Combined state — the condition in which a substance occurs in nature locked up in a compound rather than as itself.
- Concentrated nitric acid — the 68% acid obtained by distilling the acid from the absorption tower, and the acid supplied to laboratories.
- Constant boiling mixture — a mixture whose composition does not change as it boils, so that it cannot be separated any further by distillation.
- Decrepitation — the crackling and splitting of crystals when they are heated.
- Deliquescent — describing a solid that absorbs so much moisture from the air that it dissolves in the water it has taken up.
- Exothermic — describing a reaction that gives out heat.
- Fixation of atmospheric nitrogen — the conversion of free atmospheric nitrogen into useful nitrogenous compounds in the soil.
- Free state — the condition in which a substance occurs in nature as itself, uncombined with anything else.
- Fuming nitric acid — nitric acid of up to 98% strength, obtained by distilling the 68% acid over concentrated sulphuric acid.
- Hygroscopic — describing a substance that absorbs moisture from the air without dissolving in it.
- Monobasic acid — an acid whose one molecule gives one hydrogen ion in aqueous solution.
- Nascent oxygen — oxygen set free as single atoms at the moment of a reaction, written [O], and far more reactive than ordinary oxygen gas.
- Nitrates — the salts of nitric acid, formed when its hydrogen is replaced by a metal; all of them are soluble in water.
- Nitric acid — the strong monobasic acid of formula HNO₃, which acts both as an acid and as a powerful oxidising agent.
- Nitroso ferrous sulphate — the brown compound FeSO₄·NO formed at the junction of the liquids in the brown ring test.
- Nitrous oxide — the gas N₂O, known as laughing gas, formed when ammonium nitrate is heated.
- Ostwald process — the industrial manufacture of nitric acid from ammonia and air, carried out in a catalytic chamber, an oxidising chamber and an absorption tower.
- Oxidising chamber — the chamber of the Ostwald process, kept at about 50 °C, in which nitric oxide is oxidised to nitrogen dioxide.
- Passivity — the loss of reactivity of a metal such as iron or aluminium when pure concentrated nitric acid lays a thin insoluble oxide layer over its surface.
- Valency — the combining capacity of an atom, measured by the number of hydrogen atoms it can combine with or replace.
- Volatile — describing a liquid that turns to vapour easily, at a fairly low temperature.
- Xanthoproteic acid — the yellow compound formed when nitric acid reacts with the protein of the skin, which is why the acid stains skin yellow.
10.10 Points To Ponder
Nitric acid must be kept in a coloured bottle because ordinary daylight is enough to decompose it, while sulphuric acid on the same shelf can stand in clear glass for years. What does that difference suggest about how firmly each acid holds the oxygen in its molecule, and what would you predict about an acid whose decomposition products are all gases?
The Ostwald process runs its three stages at about 800 °C, then about 50 °C, then merely warm water. Every stage is an oxidation, yet the temperature falls at each step. Why might a sequence of oxidations need less and less heat as it goes on, rather than more and more?
Dilute nitric acid is described as a typical acid in every respect except that it gives no hydrogen with metals. If a chemist were handed an unfamiliar liquid that never gave hydrogen with any metal, what other evidence would you want before deciding whether to call it an acid at all?
A thin layer of oxide makes aluminium safe to hold concentrated nitric acid, but the same metal is attacked by the dilute acid. Under what conditions would you expect a self-formed oxide layer to protect a metal, and under what conditions would you expect it to fail?
Every nitrate without exception dissolves in water. Follow that single fact through three different situations: separating a nitrate from a mixture of salts, testing an unknown solution for the nitrate ion, and the fate of nitrate fertiliser spread on a field before heavy rain.
The brown ring test relies on a compound so unstable that shaking the tube destroys it. Is it a weakness in a test that its positive result cannot survive being handled? What qualities would you look for in designing a better test for the nitrate ion?
10.11 Objective Questions
Choose the correct option from those given below each question.
- The residue left when silver nitrate is heated strongly is: [1]
- silver oxide
- silver
- silver nitrite
- no residue at all
- Nitric acid is stored in coloured bottles because it is: [1]
- hygroscopic
- decomposed by light
- a monobasic acid
- soluble in water in all proportions
- The nitrate which, on heating, leaves a residue that is yellow when hot and white when cold is: [1]
- lead nitrate
- copper nitrate
- zinc nitrate
- calcium nitrate
- In the Ostwald process, the ratio of ammonia to air taken by volume is: [1]
- 1 : 2
- 1 : 5
- 1 : 10
- 10 : 1
- Iron is rendered passive by: [1]
- very dilute nitric acid
- cold dilute nitric acid
- pure concentrated nitric acid
- dilute sulphuric acid
- The metal which liberates hydrogen with very dilute nitric acid is: [1]
- copper
- magnesium
- silver
- lead
- The temperature maintained in the oxidising chamber of the Ostwald process is about: [1]
- 50 °C
- 121 °C
- 200 °C
- 800 °C
- The compound responsible for the brown ring in the test for nitrates is: [1]
- iron(III) sulphate
- nitrosyl chloride
- nitroso ferrous sulphate
- nitrogen dioxide
For the questions that follow, choose the answer from the codes given below. (1) Both A and R are true and R is the correct explanation of A. (2) Both A and R are true but R is not the correct explanation of A. (3) A is true but R is false. (4) A is false but R is true.
Assertion (A): Concentrated sulphuric acid, and not concentrated hydrochloric acid, is used in the laboratory preparation of nitric acid. Reason (R): Sulphuric acid is a non-volatile acid. [1]
Assertion (A): The ferrous sulphate solution used in the brown ring test must be freshly prepared. Reason (R): On standing in air, iron(II) sulphate is oxidised to iron(III) sulphate. [1]
Assertion (A): Ammonium nitrate leaves no residue when it is heated. Reason (R): Ammonium nitrate decomposes into a metallic oxide, nitrogen dioxide and oxygen. [1]
Assertion (A): Quartz is packed in layers in the absorption tower of the Ostwald process. Reason (R): Quartz is acid resistant and helps nitrogen dioxide dissolve uniformly in water. [1]
Match each substance in Column I with the correct observation in Column II. [4]
| Column I | Column II |
|---|---|
| a. Potassium nitrate, heated | i. A silver-grey residue is left |
| b. Copper nitrate, heated | ii. Only a colourless gas is given off, with no brown fumes |
| c. Silver nitrate, heated | iii. Nothing at all is left in the tube |
| d. Ammonium nitrate, heated | iv. Brown fumes are given off and a black residue is left |
- Fill in the blanks. [4]
- The old name of nitric acid, meaning strong water, is ______.
- The catalyst used in the catalytic chamber of the Ostwald process is ______.
- Concentrated nitric acid oxidises sulphur to ______.
- Nitric acid stains the skin yellow because it forms ______ with the protein of the skin.
- State whether each of the following is true or false. [4]
- Nitric acid is a dibasic acid.
- All nitrates are soluble in water.
- Nitric acid reacts with gold to give gold nitrate.
- Nitric oxide is colourless but turns brown in air.
- Name the following. [4]
- The gas evolved when cold dilute nitric acid acts on zinc.
- The tower in which nitrogen dioxide is absorbed in the manufacture of nitric acid.
- The nitrate which, on heating, leaves the free metal as a mirror on the walls of the tube.
- The acid formed when concentrated nitric acid oxidises sulphur.
- Give one word or phrase for each of the following. [3]
- Oxygen set free as single atoms at the moment of a reaction.
- The crackling of crystals when they are heated.
- The conversion of free atmospheric nitrogen into useful nitrogenous compounds in the soil.
- Pick the odd one out in each set and give a reason for your choice. [3]
- Copper, zinc, iron, gold
- Sodium nitrate, potassium nitrate, lead nitrate, ammonium nitrate
- Catalytic chamber, oxidising chamber, glass retort, absorption tower
10.12 Check Your Concepts
Define fixation of atmospheric nitrogen, and name the natural agency that brings it about during a thunderstorm. [2]
State what you would observe when nitric acid that has been standing in a clear glass bottle in sunlight is examined, and explain the observation. [2]
Why is all-glass apparatus used in the laboratory preparation of nitric acid? [2]
Give a reason for each of the following. [3]
- Concentrated nitric acid may be stored and carried in aluminium vessels.
- A high ratio of air to ammonia is used in the Ostwald process.
- Nitric acid cannot be concentrated beyond 68% by distillation alone.
Differentiate between cold dilute nitric acid and concentrated nitric acid in their action on copper, stating clearly the basis of the difference. [3]
Explain, with an example in each case, why dilute nitric acid is called a typical acid in its reaction with carbonates but not in its reaction with metals. [3]
Distinguish between deliquescence and decrepitation, giving one substance from this chapter that shows each. [3]
What would happen if the mixture of sodium nitrate and concentrated sulphuric acid in a retort were heated well above 200 °C? Give three consequences. [4]
Why is nitric acid described as a powerful oxidising agent? Support your answer with its action on one non-metal and on one metal. [4]
Describe the three stages of the manufacture of nitric acid by the Ostwald process, naming the chamber, the conditions and the product of each stage. [5]
Four white solids are labelled P, Q, R and S. Each is a nitrate. On strong heating, P gives off only a gas that relights a glowing splint; Q gives brown fumes and a yellow residue that sticks to the glass; R gives brown fumes and a grey metallic residue; S leaves nothing behind. Identify the metal or radical present in each, and give a reason for every identification. [5]
10.13 Equations, Conversions & Identification
Write the balanced chemical equation for the laboratory preparation of nitric acid from Chile saltpetre, stating the condition required. [2]
Complete and balance each of the following. [4]
- \(\mathrm{Cu} + \mathrm{HNO_3}\text{(conc.)} \rightarrow\)
- \(\mathrm{C} + \mathrm{HNO_3}\text{(conc.)} \rightarrow\)
- \(\mathrm{NaHCO_3} + \mathrm{HNO_3}\text{(dil.)} \rightarrow\)
- \(\mathrm{Ca(NO_3)_2} \xrightarrow{\ \Delta\ }\)
Write balanced equations for the action of dilute nitric acid on: (a) copper(II) oxide, (b) iron(III) hydroxide, (c) potassium sulphite. Name the gas evolved in (c) and state the test for it. [4]
Name the gas evolved in each case and give one test by which you would identify it. [3]
- Zinc is added to cold dilute nitric acid.
- Calcium carbonate is treated with dilute nitric acid.
- Ammonium nitrate is heated.
State one observation in each of the following. [4]
- Concentrated nitric acid is warmed with copper turnings.
- A crystal of lead nitrate is heated strongly in a hard-glass tube.
- Concentrated sulphuric acid is run down the side of a tube containing a nitrate solution and freshly prepared iron(II) sulphate solution.
- A glowing splinter is held near molten potassium nitrate.
How would you obtain the following? Give the equation in each case. [4]
- Copper(II) nitrate from copper(II) carbonate.
- Nitric acid from nitre.
- Copper(II) oxide from copper(II) nitrate.
- Magnesium nitrate from magnesium, in a reaction that also gives hydrogen.
A colourless liquid A gives off dense reddish-brown fumes when copper turnings are warmed with it, and turns blue litmus red. A leaves a yellow stain on skin. Identify A, name the brown gas, and write the equation for its reaction with copper. [3]
Fig. 10.2-1 shows the apparatus used for the laboratory preparation of nitric acid. Name the vessel in which the mixture is heated, state why a sand tray is placed under it, and state the purpose of the cold water running over the receiver. [3]
Fig. 10.6-1 shows the brown ring test. Name the two liquids that form the upper and the lower layer, state where the brown ring appears, and explain why the ring vanishes if the tube is shaken. [4]
Write balanced equations to show how nitric acid is formed in the atmosphere from nitrogen and oxygen, giving all three stages in order. [3]