Chapter 9: Study of Compounds: B. Ammonia
In 1774, in a room crowded with glass tubes and troughs, Joseph Priestley heated a mixture of sal ammoniac and lime in a small glass vessel and watched a gas come off. He had spent years collecting gases over water, letting the gas bubble up into an inverted jar and push the water down. This gas refused to be collected that way. It vanished into the water almost as fast as it was made.
So Priestley did something unusual. He set his jar over mercury instead, a liquid the gas would not dissolve in, and at last he had a jarful of it to study. It was colourless. It stung the nose. It turned a wet strip of red litmus paper blue, and no gas known at the time did that. He called it alkaline air.
What alkaline air was made of took another ten years or so to settle. In the 1780s the French chemist Claude Louis Berthollet passed electric sparks through it and broke it apart, and found nothing coming out but nitrogen and hydrogen — the two gases it had been built from all along.
That gas is ammonia. Today it is made by the million tonnes, and most of the nitrogen in the food you ate today passed through it first.
Diagram
Fig. 9.0-1 — Priestley Collecting Alkaline Air Over Mercury Eighteenth-century laboratory bench, drawn simply: a small glass vessel of sal ammoniac and lime heated over a flame on the left, a delivery tube leading right and bending upward under an inverted glass jar that stands in a shallow trough of mercury; the jar is partly filled with gas and the mercury inside it has been pushed down. Label the heated mixture, the flame, the delivery tube, the mercury trough, the inverted jar, and the collected alkaline air. Beside the jar, show a strip of moist red litmus paper turning blue.
9.1 The Ammonia Molecule
Before you follow ammonia through a flask, a factory and a dozen reactions, it is worth knowing what a single molecule of it looks like. Almost everything ammonia does — the sting in its smell, its enormous appetite for water, its behaviour as a base — comes from the way one nitrogen atom is joined to three hydrogen atoms.
9.1.1 Molecular Formula And Relative Molecular Mass
Ammonia is a compound of nitrogen and hydrogen only. One molecule of it contains one atom of nitrogen and three atoms of hydrogen, so its molecular formula is \(\mathrm{NH_3}\).
From the formula you can work out how heavy the molecule is compared with other molecules. The relative atomic mass of nitrogen is 14 and that of hydrogen is 1, so a molecule of ammonia weighs the same as 17 hydrogen atoms.
This number is small, and that turns out to matter a great deal. Ammonia is one of the lightest gases you will meet in this book, and later in the chapter you will see that its lightness decides how it is collected in the laboratory.
9.1.2 Covalent Bonding And The Lone Pair
A nitrogen atom has seven electrons, arranged 2, 5. The five electrons in its outermost shell are three short of the stable eight it needs. A hydrogen atom has a single electron and needs one more to fill its only shell. Neither atom can afford to give electrons away, so instead they share.
Nitrogen shares one electron with each of three hydrogen atoms. Each shared pair holds two atoms together, and a bond made by sharing a pair of electrons in this way is a covalent bond. Ammonia therefore has three covalent bonds, and the arrangement is drawn in Fig. 9.1-1.
Diagram
Fig. 9.1-1 — Covalent Bonding In Ammonia Two drawings side by side. Left: the electron-dot (Lewis) diagram — a central N with three H atoms around it, each N–H bond shown as a shared pair with the nitrogen’s electrons as dots and the hydrogen’s as crosses, and the fourth outer pair of nitrogen drawn on top of the N as an unshared pair; label this pair ‘lone pair’ with an arrow. Right: the overlapping-circle diagram — a large circle for the nitrogen shell overlapping three smaller hydrogen circles, with the shared pairs drawn in the overlaps and the unshared pair drawn at the top; label ‘lone pair’. Caption both drawings clearly as ‘electron-dot diagram’ and ‘overlapping-circle diagram’.
Now count the outer electrons of nitrogen again. Three of its five are used in bonding, which leaves two that are not shared with anybody. A pair of outer-shell electrons that is not used in bonding is called a lone pair.
That leftover pair is not a loose end to be ignored. It sits on the nitrogen atom, available to be offered to anything short of electrons, and it is the single feature of the ammonia molecule that explains why ammonia behaves as a base and why it can form the ammonium ion in water. You will meet both of these later in the chapter.
9.2 Occurrence Of Ammonia
9.2.1 Free State And Combined State
A substance is said to occur in the free state when it is found as itself, uncombined with anything else. Ammonia does occur this way, but only in very small amounts — a little in the air, and traces dissolved in natural water.
Far more of it is present in the combined state, that is, locked inside compounds. Nitrogen and hydrogen atoms that once belonged to ammonia are found in ammonium salts such as ammonium chloride and ammonium sulphate, which occur in soils and in the residues left where organic matter has decayed.
One thing you will not find is a mineral of ammonia. A mineral has to survive rain and groundwater, and both ammonia and its ammonium compounds are extremely soluble in water. Any deposit that formed would simply be washed away and carried off in solution.
9.2.2 Ammonia From Decaying Organic Matter
You have almost certainly smelt ammonia without being told its name. The sharp, choking smell near a heap of rotting organic matter, or in a neglected toilet, is ammonia gas escaping into the air.
Urine contains urea, \(\mathrm{CO(NH_2)_2}\), a nitrogen compound the body makes to get rid of surplus nitrogen. Bacteria break urea down in the presence of water, and the nitrogen leaves as ammonia while the carbon leaves as carbon dioxide.
urea + water → ammonia + carbon dioxide
\[\mathrm{CO(NH_2)_2} + \mathrm{H_2O} \rightarrow 2\,\mathrm{NH_3} + \mathrm{CO_2}\]
This is happening quietly and continuously wherever plant and animal matter is rotting, and it is one of the ways nitrogen is returned to the soil and the air.
9.2.3 Forms In Which Ammonia Is Used
Ammonia is not always handled as a gas. Depending on what it is wanted for, it is supplied in four different forms, which are set out in Table 9.1.
Dry ammonia gas is simply the gas itself, with all moisture removed. Squeeze that gas hard enough and it turns to a liquid, and this liquid ammonia is ammonia alone, with no water in it at all.
Dissolve the gas in water instead and you get a solution. A saturated solution of ammonia in water is called liquor ammonia fortis. It has a relative density of 0.880 and is often called 880 ammonia for that reason. Because it steadily gives off ammonia gas, it is stored in tightly stoppered bottles and kept in a cold place. Diluting it gives the ordinary bottle of ammonia solution that stands on a laboratory bench as a reagent.
Notice that the word ‘liquid’ and the word ‘liquor’ point to two completely different things here. Liquid ammonia contains no water; liquor ammonia is mostly water. Confusing the two is one of the commonest slips in this chapter.
9.3 Preparation Of Ammonia Gas
Ammonia is a gas at room temperature, so preparing it means making it inside a closed vessel and leading it away through a tube before it escapes into the room. There are two routes you need to know: one that starts from an ammonium salt, which is the usual laboratory route, and one that starts from a metal nitride.
9.3.1 The General Method: An Ammonium Salt With An Alkali
Every ammonium salt holds the ammonium ion, \(\mathrm{NH_4^{+}}\). That ion is an ammonia molecule that has picked up an extra hydrogen ion, so ammonia is already there inside the salt, waiting to be released.
A strong alkali does the releasing. The hydroxide ion of the alkali pulls the extra hydrogen ion away from the ammonium ion, leaving ammonia behind and forming water. Warming the mixture drives the ammonia out as a gas.
ammonium salt + alkali → salt + water + ammonia
The alkali may be slaked lime, \(\mathrm{Ca(OH)_2}\), caustic soda, \(\mathrm{NaOH}\), or caustic potash, \(\mathrm{KOH}\). Table 9.2 sets out the four combinations, each with its balanced equation.
| Ammonium salt | Alkali | Salt formed | Water | Ammonia |
|---|---|---|---|---|
| \(2\,\mathrm{NH_4Cl}\text{(s)}\) | \(\mathrm{Ca(OH)_2}\text{(s)}\), slaked lime | \(\mathrm{CaCl_2}\text{(s)}\) | \(2\,\mathrm{H_2O}\) | \(2\,\mathrm{NH_3}\uparrow\) |
| \((\mathrm{NH_4})_2\mathrm{SO_4}\) | \(\mathrm{Ca(OH)_2}\), slaked lime | \(\mathrm{CaSO_4}\) | \(2\,\mathrm{H_2O}\) | \(2\,\mathrm{NH_3}\uparrow\) |
| \((\mathrm{NH_4})_2\mathrm{SO_4}\) | \(2\,\mathrm{NaOH}\), caustic soda | \(\mathrm{Na_2SO_4}\) | \(2\,\mathrm{H_2O}\) | \(2\,\mathrm{NH_3}\uparrow\) |
| \((\mathrm{NH_4})_2\mathrm{SO_4}\) | \(2\,\mathrm{KOH}\), caustic potash | \(\mathrm{K_2SO_4}\) | \(2\,\mathrm{H_2O}\) | \(2\,\mathrm{NH_3}\uparrow\) |
Table 9.2 — Ammonium salts warmed with alkalis, all reactions carried out on heating.
Read down the last column and the pattern is unmistakable. Whichever ammonium salt and whichever alkali you choose, ammonia comes off. This is why the reaction is used as a test for the ammonium ion later in the chapter.
9.3.2 Laboratory Preparation From Ammonium Chloride
Of all the combinations in Table 9.2, the first is the one used in the laboratory. Ammonium chloride and slaked lime are both cheap solids, and neither needs special handling.
9.3.2.1 Reactants, Reaction And Apparatus
The reactants are ammonium chloride, \(\mathrm{NH_4Cl}\), and calcium hydroxide, \(\mathrm{Ca(OH)_2}\). The calcium hydroxide is taken in excess and mixed well with the ammonium chloride.
ammonium chloride + calcium hydroxide → calcium chloride + water + ammonia
\[2\,\mathrm{NH_4Cl}\text{(s)} + \mathrm{Ca(OH)_2}\text{(s)} \xrightarrow{\ \Delta\ } \mathrm{CaCl_2}\text{(s)} + 2\,\mathrm{H_2O}\text{(l)} + 2\,\mathrm{NH_3}\text{(g)}\uparrow\]
There are two reasons for the excess of calcium hydroxide. Ammonium chloride sublimes on heating — it turns straight from solid to vapour without melting — and some is always lost that way, so extra alkali is provided to react with what remains. Calcium hydroxide is also cheap, and unlike the other alkalis it is not deliquescent, meaning it does not absorb moisture from the air and turn damp on standing.
The apparatus is shown in Fig. 9.3-1. Follow it from left to right: the mixture is heated in a flask, the gas passes through a drying tower, and it is collected in an inverted jar.
Diagram
Fig. 9.3-1 — Laboratory Preparation Of Dry Ammonia A round-bottomed flask clamped on a stand in a slanting position with its mouth pointing downwards, heated from below by a burner; inside it the ground mixture of calcium hydroxide and ammonium chloride. A delivery tube leaves the mouth of the flask and enters the base of an upright drying tower packed with lumps of quicklime; a tube from the top of the tower leads into a gas jar held mouth downwards. Label: mixture of calcium hydroxide and ammonium chloride, burner, delivery tube, quicklime to dry ammonia, gas jar, and the collected ammonia. Show the slant of the flask clearly, with the mouth lower than the base.
9.3.2.2 Procedure And Precautions
The two solids are finely ground together, so that their particles are in close contact and the reaction is quick. The mixture is placed in a round-bottomed flask, which is clamped in a slanting position with its mouth pointing downwards, and heated. Ammonia is given off at once.
The slant is not for convenience — it is a precaution. Water is one of the products, and it condenses on the cooler upper walls of the flask. If the mouth pointed upwards, that water would trickle back onto the hot glass at the bottom and crack the flask. Tilting the flask mouth-downwards lets the water run away from the heated part.
The steps of the whole preparation are collected in FC 9.1.
Flowchart
FC 9.1 — Laboratory Preparation Of Dry Ammonia Grind ammonium chloride with excess calcium hydroxide → place in a round-bottomed flask clamped mouth-downwards → heat gently → ammonia gas passes into the delivery tube → gas passed up a drying tower packed with quicklime → dry ammonia collected in an inverted gas jar by downward displacement of air
9.3.2.3 Drying Of Ammonia Gas
The gas leaving the flask carries water vapour with it. To obtain dry ammonia it is passed through a drying tower packed with lumps of quicklime, \(\mathrm{CaO}\). A drying agent is a substance that removes water from a gas without reacting with the gas itself, and that second condition is the whole difficulty here.
The three drying agents used for most other gases all fail with ammonia, because ammonia is basic and each of them is either acidic or able to combine with it. Concentrated sulphuric acid neutralises it:
\[2\,\mathrm{NH_3} + \mathrm{H_2SO_4} \rightarrow (\mathrm{NH_4})_2\mathrm{SO_4}\]
Phosphorus pentoxide, an acidic oxide, does the same by way of phosphoric acid, giving ammonium phosphate:
\[6\,\mathrm{NH_3} + \mathrm{P_2O_5} + 3\,\mathrm{H_2O} \rightarrow 2\,(\mathrm{NH_4})_3\mathrm{PO_4}\]
Anhydrous calcium chloride does not neutralise ammonia, but it takes the gas up into an addition compound, which removes it just as effectively:
\[\mathrm{CaCl_2} + 4\,\mathrm{NH_3} \rightarrow \mathrm{CaCl_2\!\cdot\!4NH_3}\]
Quicklime is the exception. It is a basic oxide, so it has no reaction with a basic gas, and it takes up water readily. Quicklime is caustic and gives out heat as it absorbs moisture, so the lumps are handled with tongs and never with bare fingers.
9.3.2.4 Collection Of Ammonia
Dry ammonia is collected in gas jars held mouth downwards, a method called downward displacement of air. The gas rises into the jar and pushes the air in it downwards and out.
Two properties settle this choice. First, ammonia is lighter than air: its vapour density is 8.5, while that of air is 14.4. A gas lighter than air will not stay in an upright jar, but it collects perfectly well in an inverted one. Second, ammonia is extremely soluble in water, so it cannot be collected over water in the ordinary way — it would dissolve as fast as it was made.
9.3.2.5 Ammonium Salts That Cannot Be Used
Table 9.2 said that ammonium salts warmed with alkalis give ammonia, and that is true as a general rule. Two salts are exceptions, and the reason is that they never get as far as reacting with the alkali.
Ammonium nitrate and ammonium nitrite both decompose the moment they are warmed. Ammonium nitrate gives dinitrogen oxide and water:
\[\mathrm{NH_4NO_3} \xrightarrow{\ \Delta\ } \mathrm{N_2O} + 2\,\mathrm{H_2O}\]
Ammonium nitrite gives nitrogen and water:
\[\mathrm{NH_4NO_2} \xrightarrow{\ \Delta\ } \mathrm{N_2} + 2\,\mathrm{H_2O}\]
No ammonia appears in either equation. Ammonium nitrate is explosive in nature besides, which is a second and sufficient reason for keeping it out of a preparation that involves heating.
9.3.3 Preparation From Metal Nitrides
There is a second route to ammonia that uses no ammonium salt at all. A nitride is a compound of a metal with nitrogen alone, and certain nitrides react with warm water to give ammonia directly. The method is a costly one, so it is not the usual laboratory choice.
Magnesium nitride is the common example. Warm water breaks it up into magnesium hydroxide and ammonia.
magnesium nitride + water → magnesium hydroxide + ammonia
\[\mathrm{Mg_3N_2}\text{(s)} + 6\,\mathrm{H_2O}\text{(l)} \rightarrow 3\,\mathrm{Mg(OH)_2}\text{(s)} + 2\,\mathrm{NH_3}\text{(g)}\uparrow\]
Aluminium nitride behaves in the same way, giving aluminium hydroxide:
\[\mathrm{AlN}\text{(s)} + 3\,\mathrm{H_2O}\text{(l)} \rightarrow \mathrm{Al(OH)_3}\text{(s)} + \mathrm{NH_3}\text{(g)}\uparrow\]
The arrangement is shown in Fig. 9.3-2. The nitride is kept in a round-bottomed flask fitted with a thistle funnel and a delivery tube, and warm water is added very slowly through the funnel — slowly, because the reaction is brisk and a sudden rush of water would send gas and spray out through the funnel.
Diagram
Fig. 9.3-2 — Preparation Of Ammonia From Magnesium Nitride Round-bottomed flask on a stand holding magnesium nitride, closed with a two-holed stopper; a thistle funnel passes through one hole with its stem dipping below the level of the solid, and a delivery tube leaves the other hole. Water is being added slowly down the thistle funnel from a small flask above. The delivery tube leads to an upright drying tower packed with quicklime, and from its top a tube leads to a gas jar held mouth downwards. Label: magnesium nitride, thistle funnel, warm water, delivery tube, drying agent CaO, ammonia gas.
The metal hydroxide formed is insoluble and stays behind in the flask, so only ammonia escapes along the delivery tube. As before, the gas is dried by passing it through quicklime packed in a drying tower, and collected in a jar held mouth downwards.
9.4 Preparation Of Aqueous Ammonia
Aqueous ammonia is nothing more than ammonia gas dissolved in water, and preparing it sounds as though it should be the easiest job in the chapter — lead the gas from the flask into a beaker of water and wait. In practice, that simple arrangement wrecks the apparatus.
Why The Gas Cannot Simply Be Bubbled In
Ammonia dissolves in water astonishingly fast — faster than the flask can make it. The moment the gas reaches the water it disappears into it, and the pressure inside the delivery tube drops below the pressure of the atmosphere outside.
Air pressure then pushes water up the tube and back towards the flask. This drawing of water backwards along the delivery tube is called back suction, and it is dangerous. The flask is hot, the water is cold, and cold water meeting hot glass cracks it at once.
The Funnel Arrangement
The remedy is to fit a funnel to the outer end of the delivery tube with a piece of rubber tubing, and to dip only a small part of the rim of the funnel in the water. The arrangement is shown in Fig. 9.4-1. It is the same device used when hydrogen chloride gas is dissolved in water to make hydrochloric acid.
Diagram
Fig. 9.4-1 — Preparation Of Aqueous Solution Of Ammonia Round-bottomed flask on a stand, clamped mouth-downwards and heated by a burner, containing the mixture of ammonium chloride and calcium hydroxide; a delivery tube leaves the flask, runs across to the right and turns down, ending in a wide inverted funnel joined to it by rubber tubing. Below the funnel is a trough of water, with only the rim of the funnel’s wide mouth just touching the water surface. Label: mixture of ammonium chloride and calcium hydroxide, burner, delivery tube, rubber tubing, funnel, water. Show the water level meeting the funnel rim, not covering it.
The wide mouth of the funnel gives the gas a large surface of water to dissolve into, so the solution forms quickly.
More importantly, the funnel acts as a safety valve. When the gas dissolves and the pressure inside falls, water rushes up into the funnel — but as it does so, the rim rises clear of the water surface in the trough and the connection is broken. Ammonia pouring out of the flask now pushes that water back down, the rim touches the water again, and the cycle repeats.
Water therefore moves in and out of the funnel a few centimetres at a time and never travels up the narrow delivery tube to the flask. The gas dissolves steadily, and the hot flask stays safe.
9.5 Manufacture Of Ammonia By Haber’s Process
The laboratory methods of Section 9.3 will fill a gas jar. The world needs ammonia by the tonne, and for that it goes to the cheapest raw material there is — the air overhead, which is about four-fifths nitrogen. The industrial method that combines that nitrogen directly with hydrogen is Haber’s process.
9.5.1 Reactants And Their Sources
The two reactants are nitrogen and hydrogen, taken in the ratio of 1 : 3 by volume.
Nitrogen gas is obtained from the air, by the fractional distillation of liquid air. Air is first cooled until it liquefies, and the liquid is then allowed to warm slowly; its components boil off one after another at their own boiling points, and nitrogen is collected as one of them.
Hydrogen gas is obtained either from water gas, by the Bosch process, or from natural gas.
9.5.2 The Reaction And The Favourable Conditions
Nitrogen and hydrogen combine directly to give ammonia.
nitrogen + hydrogen ⇌ ammonia + heat
\[\mathrm{N_2}\text{(g)} + 3\,\mathrm{H_2}\text{(g)} \rightleftharpoons 2\,\mathrm{NH_3}\text{(g)} + \text{heat}\]
Three features of this equation decide everything that follows. The double arrow says the reaction is reversible — ammonia breaks back down into nitrogen and hydrogen as fast as conditions allow. The word ‘heat’ on the right says it is exothermic, giving out heat as it goes. And counting the volumes shows that four volumes of reactants make only two volumes of product, so the reaction proceeds with a decrease in volume.
The conditions used in the plant are these:
- Temperature: optimum temperature is 450–500 °C.
- Pressure: above 200 atm.
- Catalyst: finely divided iron.
- Promoter: traces of molybdenum or \(\mathrm{Al_2O_3}\).
Even under these conditions, only about 15% of the reacting gases are converted to ammonia in one pass through the plant.
9.5.3 The Process In The Plant
Nitrogen and hydrogen are first dried and purified, then mixed in the ratio of 1 : 3 and compressed to a high pressure. Purification is not a formality: carbon dioxide, carbon monoxide and traces of sulphur compounds such as hydrogen sulphide act as catalytic poisons, substances that destroy the activity of the catalyst, and removing them is essential.
The compressed mixture is passed into an electrically heated catalytic chamber packed with finely divided iron together with aluminium oxide, or a little molybdenum, and held at about 500 °C. Ammonia forms here.
The hot gases leaving the chamber are a mixture of ammonia with the nitrogen and hydrogen that did not react — about 15% ammonia in all. Before they leave, they give up their heat to the cold incoming gases, which warms those gases on their way in and saves fuel.
The plant is shown in Fig. 9.5-1, and the whole cycle is set out in FC 9.2.
Diagram
Fig. 9.5-1 — Manufacture Of Ammonia By Haber’s Process Vertical plant diagram. At the top, two inlet arrows: nitrogen from liquid air, 1 volume, and hydrogen by the Bosch process, 3 volumes, meeting and entering a compressor drawn as a box. From the compressor a pipe runs down into a tall catalytic chamber shown as a cylinder with a hatched bed of catalyst; label the bed ‘finely divided iron, 500 °C’ and the vessel ‘250 atmospheres pressure’. Below the chamber, a coiled pipe inside a cooling chamber, with arrows marked cooling liquid in and cooling liquid out. At the bottom, an outlet arrow labelled liquid ammonia. From above the cooling chamber, a recirculation pipe carrying uncombined gases runs up the left side, through a second compressor, and rejoins the inlet stream. Label every part named here.
Flowchart
FC 9.2 — The Cycle Of Haber’s Process Dry nitrogen 1 volume and dry hydrogen 3 volumes, purified → compression pump, above 200 atm → electrically heated catalytic chamber, finely divided iron with molybdenum, 450–500 °C → outgoing mixture, about 15% ammonia → condenser with cooling pipes → liquid ammonia drawn off and stored under pressure; uncombined nitrogen and hydrogen returned along the recirculation pipe to the compression pump and round again
9.5.4 Recovery And Recirculation Of Ammonia
Only one part in seven of the outgoing stream is ammonia, so it has to be separated from the unchanged nitrogen and hydrogen. Two properties allow this.
The first is liquefaction. Ammonia liquefies far more easily than nitrogen and hydrogen do, so when the mixture is passed through cooling pipes the ammonia condenses to a liquid and runs off, while the other two gases stay gaseous. This is the method used in the plant.
The second is solubility in water. Ammonia is highly soluble, while nitrogen and hydrogen are only very slightly soluble, so washing the mixture with water also strips the ammonia out.
The nitrogen and hydrogen left over are not wasted. They are recirculated through the plant and given another chance to react, and by sending them round again and again an eventual yield of 98% can be achieved. The ammonia collected is stored as a liquid under pressure.
9.5.5 Why Each Condition Is Chosen
Each figure in the list of conditions is a considered compromise, and the reasoning is worth following carefully.
Temperature
The reaction is exothermic, so a low temperature favours the formation of ammonia. But at a low temperature the reaction proceeds too slowly to be of any use — a good yield that takes weeks is no yield at all. In practice the optimum temperature has been found to lie between 450 °C and 500 °C, high enough for a workable speed and not so high as to destroy the yield.
There is a bonus. Since the reaction gives out heat, the heat evolved keeps the chamber hot once it is running, and external heating is not needed after the reactants have been heated at the start.
Pressure
Four volumes of reactants produce two volumes of product, so squeezing the gases pushes the reaction towards the side that occupies less space — the ammonia side. High pressure therefore favours the formation of ammonia. The optimum pressure lies between 200 and 900 atm, and in practice a pressure of about 250 atm is used, because vessels strong enough for still higher pressures become impossibly expensive.
Catalyst And Promoter
A catalyst does not change how much ammonia can be formed; it changes how quickly the mixture gets there. The catalyst used is finely divided iron, obtained by the reduction of iron oxide. It is used finely divided because a powder presents far more surface area than a lump, and the reaction happens on the surface.
Along with it goes a promoter — a substance which is not itself the catalyst but which increases the catalyst’s efficiency. Molybdenum or aluminium oxide is used in traces for this purpose.
9.6 Physical Properties Of Ammonia
You have now made ammonia three different ways. This section is about what the gas is actually like — how it looks, how it smells, how heavy it is, and how much of it a beaker of water can swallow.
- Colour: colourless.
- Odour: strong, pungent, choking smell.
- Taste: slightly bitter, alkaline taste.
- Physiological action: non-poisonous. If inhaled it affects the respiratory system and brings tears to the eyes. It is a powerful heart stimulant and may cause death when inhaled in large quantities.
- Density: vapour density = 8.5, so ammonia is lighter than air.
- Nature: alkaline.
- Liquefaction: easily liquefied at 10 °C by compressing it at 6 atm pressure.
- Boiling point: liquid ammonia boils at −33.5 °C.
- Freezing point: solid ammonia melts at −77.7 °C.
- Solubility: highly soluble in water. One volume of water dissolves about 702 volumes of ammonia at 20 °C and 1 atm pressure.
Two entries in that list deserve a second look. The fourth says ammonia is non-poisonous and then says it can kill — and both statements are true. Ammonia is not a poison in the sense that a small trace of it does lasting damage; but in quantity it attacks the lining of the airways, and enough of it will stop a person breathing. In a laboratory a gas is never inhaled directly, and nothing whatever is tasted.
The tenth is the extraordinary one. Seven hundred volumes of gas dissolving in one volume of water is a solubility unlike almost anything else you will meet. It explains why ammonia cannot be collected over water, why back suction is a danger, and why the experiment in the next subsection works at all.
9.6.1 The Fountain Experiment
The fountain experiment is set up for one purpose: to demonstrate the extremely high solubility of ammonia gas in water, and to show at the same time that the solution formed is alkaline.
Apparatus
- A round-bottomed flask filled with ammonia gas.
- A rubber stopper with two holes closing the mouth of the flask — one hole for a jet tube, the other for a dropper containing water.
- A trough containing red litmus solution.
- A retort stand, used to fix the flask in an inverted position as shown in Fig. 9.6-1.
Diagram
Fig. 9.6-1 — The Fountain Experiment A round-bottomed flask clamped to a retort stand in an inverted position, mouth downwards, filled with ammonia gas. A two-holed rubber stopper closes the mouth; through one hole a narrow jet tube runs from inside the flask down into a trough of red litmus solution standing below, and through the other hole a rubber dropper containing water projects. Inside the flask, draw a spray of liquid rising from the tip of the jet tube and spreading out as a fountain, coloured blue. Label: ammonia gas, jet tube, dropper, red litmus solution, blue fountain.
Procedure And Explanation
The dropper is squeezed, sending a few drops of water into the flask. Those few drops dissolve an enormous quantity of the ammonia around them, because the gas is so soluble.
Removing that gas leaves far fewer molecules pressing on the walls, so a partial vacuum forms inside the flask. The pressure outside is now much higher than the pressure inside.
The atmosphere therefore pushes down on the surface of the litmus solution in the trough and drives it up the jet tube, where it bursts into the flask as a fountain. As it enters, the ammonia turns the solution alkaline, and the red litmus changes to blue — so what you see is a blue fountain filling the flask.
Every stage of that chain has a cause: high solubility → gas removed from the flask → pressure inside falls → outside pressure pushes the liquid up → basic gas turns red litmus blue.
9.7 Chemical Properties Of Ammonia
Ammonia has a small number of reactions, and almost all of them come from two features of the molecule: the lone pair on the nitrogen atom, which makes it a base, and the three hydrogen atoms it is willing to give up, which make it a reducing agent. Keep those two ideas in view as you work through this section.
9.7.1 Thermal Dissociation
Ammonia is not a rugged molecule. Heat it strongly, or pass electric sparks through it, and it comes apart into the two gases it was built from.
ammonia ⇌ nitrogen + hydrogen
\[2\,\mathrm{NH_3}\text{(g)} \rightleftharpoons \mathrm{N_2}\text{(g)} + 3\,\mathrm{H_2}\text{(g)}\]
Notice the double arrow. Breaking up on heating and re-forming on cooling, both in the same vessel, is called thermal dissociation, and this is the reverse of the very reaction that Haber’s process is built on.
9.7.2 Basic Nature And Action On Indicators
Dry ammonia gas is a covalent molecule with no ions in it, so it is neutral — even when it has been liquefied under pressure. A dry gas jar of ammonia will not affect a dry piece of litmus paper.
Everything changes when water is present. The lone pair on the nitrogen atom, which you met in Section 9.1.2, is a pair of electrons with nothing to do, and it reaches out and takes a hydrogen ion from a water molecule. This produces ammonium hydroxide.
\[\mathrm{NH_3}\text{(g)} + \mathrm{H_2O}\text{(l)} \rightleftharpoons \mathrm{NH_4OH}\text{(aq)}\]
Ammonium hydroxide is a weak base — a base that dissociates only partly in water, so that at any moment most of it is still whole. The small fraction that does dissociate releases hydroxide ions.
\[\mathrm{NH_4OH}\text{(aq)} \rightleftharpoons \mathrm{NH_4^{+}}\text{(aq)} + \mathrm{OH^{-}}\text{(aq)}\]
Those hydroxide ions are the whole of the alkaline behaviour. Every indicator change in Table 9.3 is the response of the indicator to \(\mathrm{OH^{-}}\) in the solution, and none of them happens without water present.
| Indicator | Original colour | Colour with ammonia |
|---|---|---|
| Moist litmus | Red | Blue |
| Methyl orange | Orange | Yellow |
| Phenolphthalein | Colourless | Pink |
Table 9.3 — Action of ammonia on indicators.
9.7.3 Burning Of Ammonia In Oxygen
Ammonia is not a supporter of combustion — a burning splint pushed into a jar of it goes out. Nor does it burn in air. But in pure oxygen it does burn, with a greenish yellow flame.
The Experiment
A wide glass tube is fitted with a cork carrying two tubes, one of which reaches almost to the top of the wide tube while the other stops lower down. Ammonia is passed in through the longer tube and oxygen through the shorter one, and a lighted match is applied. The apparatus is shown in Fig. 9.7-1.
Diagram
Fig. 9.7-1 — Burning Of Ammonia In Oxygen A wide vertical glass tube, open at the top, closed at the bottom by a cork carrying two narrow tubes. The left-hand tube passes up through the cork and reaches almost to the top of the wide tube; the right-hand tube stops well below it. Show a flame at the open top of the wide tube. Label: dry ammonia gas entering the longer tube, oxygen entering the shorter tube, greenish yellow flame.
ammonia + oxygen → nitrogen + water
\[4\,\mathrm{NH_3}\text{(g)} + 3\,\mathrm{O_2}\text{(g)} \rightarrow 2\,\mathrm{N_2}\text{(g)} + 6\,\mathrm{H_2O}\text{(l)}\]
The observation is a greenish yellow flame, with water vapour and nitrogen produced. Since nothing was present but ammonia and oxygen, and the products contain nitrogen and hydrogen, the experiment shows once again that ammonia is made of nitrogen and hydrogen. The reaction is irreversible and strongly exothermic.
9.7.4 Catalytic Oxidation Of Ammonia
Change the conditions and the same two gases give an entirely different product. In the presence of platinum at 800 °C, ammonia reacts with oxygen to give nitric oxide and water vapour instead of nitrogen.
Dry ammonia and oxygen are passed through inlets over heated platinum placed inside a combustion tube, as in Fig. 9.7-2. The platinum glows red when hot.
Diagram
Fig. 9.7-2 — Catalytic Oxidation Of Ammonia A horizontal combustion tube on a stand, with a gauze or coil of platinum supported in the middle of it and glowing; a burner beneath the platinum. Two inlets enter at the left: one carrying dry ammonia and one carrying dry oxygen. At the right the tube leads into a round-bottomed flask in which reddish brown fumes are shown, with water droplets condensing on the walls of the tube before it. Label: dry oxygen, dry ammonia, platinum catalyst, heat, water droplets, nitrogen dioxide fumes.
\[4\,\mathrm{NH_3}\text{(g)} + 5\,\mathrm{O_2}\text{(g)} \xrightarrow{\ \mathrm{Pt},\ 800^\circ\text{C}\ } 4\,\mathrm{NO}\text{(g)} + 6\,\mathrm{H_2O}\text{(g)} + \text{heat}\]
Nitric oxide is colourless, but it does not stay that way. As soon as it meets more oxygen it is oxidised further to nitrogen dioxide, a brown gas.
\[2\,\mathrm{NO}\text{(g)} + \mathrm{O_2}\text{(g)} \rightarrow 2\,\mathrm{NO_2}\text{(g)}\]
Two observations follow. Reddish brown vapours of nitrogen dioxide are seen in the flask, and they are your evidence that nitric oxide was formed. And the platinum goes on glowing even after the burner has been taken away, because the catalytic oxidation is exothermic and its own heat keeps the metal hot.
This reaction is the first step in the industrial manufacture of nitric acid, which you will meet again among the uses of ammonia.
9.7.5 Ammonia As A Reducing Agent
A reducing agent is a substance that removes oxygen from another substance, or supplies hydrogen to it. Ammonia is a good one, because it carries three hydrogen atoms that it readily gives up, leaving harmless nitrogen behind.
9.7.5.1 Reduction Of Heated Metallic Oxides
Ammonia reduces heated metallic oxides to give the metal, water vapour and nitrogen.
Dry ammonia gas is passed over heated copper oxide in a hard glass tube. Copper(II) oxide is a black solid; the copper produced is reddish brown, so the change is easy to see along the tube.
copper(II) oxide + ammonia → copper + water + nitrogen
\[3\,\mathrm{CuO}\text{(s)} + 2\,\mathrm{NH_3}\text{(g)} \rightarrow 3\,\mathrm{Cu}\text{(s)} + 3\,\mathrm{H_2O}\text{(l)} + \mathrm{N_2}\text{(g)}\uparrow\]
The arrangement in Fig. 9.7-3 does more than carry out the reaction; it catches both products so that each can be identified.
Diagram
Fig. 9.7-3 — Reducing Action Of Ammonia On Copper Oxide Left to right on a bench: an inlet delivering dry ammonia into a hard glass tube containing copper oxide, heated from below by a burner; the tube leads into a U-tube standing in a trough of cold water, in which water droplets collect; from the U-tube a delivery tube runs to a gas jar inverted over water in a trough, filling with nitrogen. Label: dry ammonia, copper oxide, heat, U-tube, cold water, water collected, nitrogen.
The nitrogen is collected over water in a gas jar, which is possible because nitrogen is only very slightly soluble in water. The water formed condenses in the U-tube, and it is tested with anhydrous copper sulphate, which turns from white to blue.
Taken together, these two products prove the composition of ammonia once more: the nitrogen came from it, and the hydrogen that made the water came from it too. Ammonia is therefore a hydride of nitrogen.
Other metallic oxides behave in the same way. Lead monoxide, which is yellow, and lead(IV) oxide, which is dark brown, are both reduced to greyish metallic lead.
\[3\,\mathrm{PbO}\text{(s)} + 2\,\mathrm{NH_3}\text{(g)} \rightarrow 3\,\mathrm{Pb}\text{(s)} + 3\,\mathrm{H_2O}\text{(l)} + \mathrm{N_2}\text{(g)}\uparrow\]
\[3\,\mathrm{PbO_2}\text{(s)} + 4\,\mathrm{NH_3}\text{(g)} \rightarrow 3\,\mathrm{Pb}\text{(s)} + 6\,\mathrm{H_2O}\text{(l)} + 2\,\mathrm{N_2}\text{(g)}\uparrow\]
9.7.5.2 Reduction Of Chlorine To Hydrogen Chloride
When chlorine reacts with ammonia, the chlorine is reduced to hydrogen chloride — it takes hydrogen from the ammonia. What happens to that hydrogen chloride afterwards depends entirely on which of the two gases is present in excess.
When Ammonia Is In Excess
Chlorine first strips hydrogen from some of the ammonia:
\[2\,\mathrm{NH_3} + 3\,\mathrm{Cl_2} \rightarrow \mathrm{N_2} + 6\,\mathrm{HCl}\]
The hydrogen chloride so formed is at once neutralised by the ammonia still left over:
\[6\,\mathrm{NH_3} + 6\,\mathrm{HCl} \rightarrow 6\,\mathrm{NH_4Cl}\]
Adding the two gives the overall change:
\[8\,\mathrm{NH_3}\text{(excess)} + 3\,\mathrm{Cl_2} \rightarrow \mathrm{N_2} + 6\,\mathrm{NH_4Cl}\]
The observation is in two parts. The greenish-yellow colour of the chlorine disappears, and dense white fumes of ammonium chloride are formed.
When Chlorine Is In Excess
Now there is no spare ammonia to mop up the hydrogen chloride, and the excess chlorine attacks the nitrogen as well. The products are hydrogen chloride and nitrogen trichloride, a yellow, highly explosive liquid.
\[\mathrm{NH_3} + 3\,\mathrm{Cl_2}\text{(excess)} \rightarrow 3\,\mathrm{HCl} + \mathrm{NCl_3}\]
Because nitrogen trichloride is dangerously explosive, this reaction is never carried out as a demonstration.
Both reactions make the same two points: chlorine has a great affinity for hydrogen, and ammonia is a reducing agent.
9.7.6 Reaction With Acids
Being a base, ammonia accepts protons — that is, hydrogen ions — from an acid, and in doing so becomes the ammonium ion. The salts formed are the ammonium salts, and you have already used two of them as starting materials in Section 9.3.
The reaction happens between the dry gases as well as in solution. Table 9.4 gives the three you need.
| Ammonia | Acid | Ammonium salt formed |
|---|---|---|
| \(\mathrm{NH_3}\text{(g)}\) | \(\mathrm{HCl}\text{(g)}\) | \(\mathrm{NH_4Cl}\text{(s)}\), ammonium chloride |
| \(\mathrm{NH_3}\) | \(\mathrm{HNO_3}\) | \(\mathrm{NH_4NO_3}\), ammonium nitrate |
| \(2\,\mathrm{NH_3}\) | \(\mathrm{H_2SO_4}\) | \((\mathrm{NH_4})_2\mathrm{SO_4}\), ammonium sulphate |
Table 9.4 — Ammonia gas with acids.
The first line is worth dwelling on. Two colourless gases meet and a white solid appears directly out of the air between them — no liquid, no heating, nothing else. This is the dense white cloud you get when a rod moistened with hydrochloric acid is held near ammonia, and it is used later as a test.
Heating The Salts Formed
Ammonium chloride and ammonium nitrate both break up on heating, but not in the same way, and the difference has a name.
Ammonium chloride shows thermal dissociation: it splits into ammonia and hydrogen chloride on heating, and the two recombine on cooling, so the change is reversible.
\[\mathrm{NH_4Cl}\text{(s)} \rightleftharpoons \mathrm{NH_3}\text{(g)} + \mathrm{HCl}\text{(g)}\]
Ammonium nitrate shows thermal decomposition: it breaks down on heating into dinitrogen oxide and water, as you saw in Section 9.3.2.5, and the products do not recombine. The change is irreversible.
Aqueous Ammonia With Acids
In solution, ammonia is present as ammonium hydroxide, and its reaction with an acid is an ordinary neutralisation: salt plus water, as set out in Table 9.5.
| Aqueous ammonia | Acid | Ammonium salt | Water |
|---|---|---|---|
| \(\mathrm{NH_4OH}\) | \(\mathrm{HCl}\) | \(\mathrm{NH_4Cl}\) | \(\mathrm{H_2O}\) |
| \(\mathrm{NH_4OH}\) | \(\mathrm{HNO_3}\) | \(\mathrm{NH_4NO_3}\) | \(\mathrm{H_2O}\) |
| \(2\,\mathrm{NH_4OH}\) | \(\mathrm{H_2SO_4}\) | \((\mathrm{NH_4})_2\mathrm{SO_4}\) | \(2\,\mathrm{H_2O}\) |
Table 9.5 — Aqueous ammonia with acids.
9.7.7 Aqueous Ammonia With Soluble Metal Salts
Ammonium hydroxide supplies hydroxide ions, and hydroxide ions meeting a metal ion in solution give the metal hydroxide. Most metal hydroxides are insoluble, so they come out of the solution as a precipitate — sodium and potassium are the exceptions, their hydroxides being soluble. In this behaviour, aqueous ammonia acts much like sodium hydroxide or potassium hydroxide solution. Table 9.6 gives the five salts you need.
| Metallic salt solution | Ammonium hydroxide | Products | Hydroxide precipitated | Colour of precipitate | Solubility in excess \(\mathrm{NH_4OH}\) |
|---|---|---|---|---|---|
| \(\mathrm{FeSO_4}\) | \(2\,\mathrm{NH_4OH}\) | \((\mathrm{NH_4})_2\mathrm{SO_4}\) | \(\mathrm{Fe(OH)_2}\downarrow\) | Dirty green | Insoluble |
| \(\mathrm{FeCl_3}\) | \(3\,\mathrm{NH_4OH}\) | \(3\,\mathrm{NH_4Cl}\) | \(\mathrm{Fe(OH)_3}\downarrow\) | Reddish brown | Insoluble |
| \(\mathrm{Pb(NO_3)_2}\) | \(2\,\mathrm{NH_4OH}\) | \(2\,\mathrm{NH_4NO_3}\) | \(\mathrm{Pb(OH)_2}\downarrow\) | White | Insoluble |
| \(\mathrm{Zn(NO_3)_2}\) | \(2\,\mathrm{NH_4OH}\) | \(2\,\mathrm{NH_4NO_3}\) | \(\mathrm{Zn(OH)_2}\downarrow\) | White gelatinous | Soluble |
| \(\mathrm{CuSO_4}\) | \(2\,\mathrm{NH_4OH}\) | \((\mathrm{NH_4})_2\mathrm{SO_4}\) | \(\mathrm{Cu(OH)_2}\downarrow\) | Pale blue | Soluble |
Table 9.6 — Aqueous ammonia added to solutions of metal salts.
Read the table across and you will see that two things vary: the colour of the precipitate, and whether it dissolves when more ammonium hydroxide is added. Between them, those two observations tell one metal ion from another, which is why ammonium hydroxide is a standard reagent in qualitative analysis — the branch of analysis concerned with finding out which substances are present, rather than how much.
The copper case is the most striking. The pale blue precipitate of copper hydroxide dissolves in excess ammonium hydroxide to give a deep azure blue solution, because a soluble complex salt, tetraamminecopper(II) sulphate, is formed.
\[\mathrm{Cu(OH)_2} + (\mathrm{NH_4})_2\mathrm{SO_4} + 2\,\mathrm{NH_4OH} \rightarrow [\mathrm{Cu(NH_3)_4}]\mathrm{SO_4} + 4\,\mathrm{H_2O}\]
So the full observation with copper sulphate solution comes in two stages: a pale blue precipitate first, then an inky blue solution as more ammonia is added. Reporting only one of the two stages is an incomplete answer.
9.8 Uses Of Ammonia
Almost everything ammonia is used for follows from something you have already learnt about it — it evaporates with a strong cooling effect, it precipitates metal hydroxides, it dissolves grease, and it is the cheapest source of combined nitrogen there is.
9.8.1 Ammonia As A Refrigerant
Liquid ammonia is used as a refrigerant in ice plants. Anhydrous ammonia is a clear, colourless liquid under pressure, and when the pressure is released it evaporates rapidly and produces a cooling effect.
The reason is the energy the change of state demands. Vaporising liquid ammonia takes a great deal of energy — 5700 calories per mole at −33 °C — and that energy has to come from somewhere. It is drawn from the bodies surrounding the evaporating liquid, which are consequently cooled down.
The Refrigerants That Replaced It
For much of the twentieth century the main refrigerants used were the freons, or chlorofluorocarbons, usually written CFC. They are stable, non-flammable liquids, and for a long time that stability looked like an unmixed advantage.
It is exactly the stability that causes the trouble. CFC molecules survive long enough to drift up to the ozone layer, where ultraviolet rays decompose them and release chlorine in the atomic form.
\[\mathrm{CF_2Cl_2}\text{(g)} \xrightarrow{\ \text{ultraviolet}\ } \mathrm{CF_2Cl}\text{(g)} + \mathrm{Cl}\text{(g)}\]
The chlorine atom released is a free radical — a highly reactive particle carrying an unpaired electron. It attacks an ozone molecule and forms chlorine monoxide:
\[\mathrm{Cl}\text{(g)} + \mathrm{O_3}\text{(g)} \rightarrow \mathrm{ClO}\text{(g)} + \mathrm{O_2}\text{(g)}\]
The chlorine monoxide then reacts with atomic oxygen, and in doing so releases the chlorine atom again:
\[\mathrm{ClO}\text{(g)} + \mathrm{O}\text{(g)} \rightarrow \mathrm{Cl}\text{(g)} + \mathrm{O_2}\text{(g)}\]
That last equation is the sting. The chlorine atom comes back out unchanged, ready to destroy another ozone molecule, and then another — so a single chlorine atom goes on causing ozone depletion long after the CFC molecule that carried it has gone. The cycle is drawn in FC 9.3.
Flowchart
FC 9.3 — How One Chlorine Atom Destroys Ozone Repeatedly CFC molecule rises to the ozone layer → ultraviolet radiation splits off a chlorine free radical → chlorine radical attacks an ozone molecule → chlorine monoxide and oxygen formed → chlorine monoxide meets an atomic oxygen → oxygen released and the chlorine radical is regenerated → the regenerated radical returns to the start of the cycle and attacks a further ozone molecule
Why Ammonia Is Used Again
Set against the CFCs, ammonia has three advantages as a refrigerant, and they are compared in Table 9.7.
- It is environmentally compatible. It does not deplete the ozone layer and does not contribute to global warming.
- It has superior thermodynamic qualities, so ammonia refrigeration systems use less electricity.
- It has a recognisable odour, so leaks are not likely to escape notice. Being lighter than air, escaped ammonia rises in the atmosphere rather than collecting at ground level.
Ammonia also has disadvantages, and an honest comparison has to name them. It is not compatible with copper, so it cannot be used in any system with copper pipes, and it is poisonous in high concentration.
9.8.2 Ammonia As A Laboratory Reagent And A Cleansing Agent
Ammonia solution is an important laboratory reagent in qualitative analysis, because, as Table 9.6 showed, it gives characteristic coloured precipitates of metallic hydroxides.
Aqueous ammonia also emulsifies or dissolves fats and grease, breaking an oily film up into fine droplets that water can carry away. This makes it useful for removing grease and perspiration stains from clothes, and for cleaning tiles and windows.
9.8.3 Ammonia In Manufacture
Nitrogenous Fertilisers
The largest single use of ammonia is in making nitrogenous fertilisers — ammonium sulphate, diammonium hydrogen phosphate, ammonium nitrate and urea. A plant cannot take up nitrogen gas, but it can take up these soluble compounds through its roots.
Urea is made by reacting ammonia with carbon dioxide at 150 °C and 150 atm pressure.
ammonia + carbon dioxide → urea + water
\[2\,\mathrm{NH_3} + \mathrm{CO_2} \xrightarrow{\ 150^\circ\text{C},\ 150\ \text{atm}\ } \mathrm{NH_2CONH_2} + \mathrm{H_2O}\]
This is the reverse of what the bacteria in Section 9.2.2 were doing to urea, and it closes a neat circle: nitrogen taken from the air becomes ammonia, ammonia becomes urea, urea feeds a crop, and the nitrogen eventually returns to the soil.
Explosives And Other Ammonium Salts
Ammonium nitrate is used as an explosive. It decomposes explosively on heating, leaving no residue behind — the decomposition you met in Section 9.3.2.5, which produces nothing but gas and vapour, so a very large volume appears in an instant from a small mass of solid.
Other ammonium salts have quieter uses. Ammonium carbonate is the smelling salt used to revive a person who has fainted. Ammonium chloride is used in dry cells, and in the laboratory preparation of nitrogen gas, where it is first warmed with sodium nitrite:
\[\mathrm{NaNO_2}\text{(aq)} + \mathrm{NH_4Cl}\text{(aq)} \rightarrow \mathrm{NaCl}\text{(aq)} + \mathrm{NH_4NO_2}\text{(aq)}\]
\[\mathrm{NH_4NO_2}\text{(aq)} \xrightarrow{\ \Delta\ } \mathrm{N_2}\text{(g)} + 2\,\mathrm{H_2O}\text{(g)}\]
Industrial Products
Ammonia is a starting material for nylon and rayon, sodium cyanamide, plastics, dyes, organic chemicals, explosives and wood pulp.
It is also central to the Solvay process for making sodium carbonate. Ammonia, carbon dioxide and sodium chloride react in water solution, and sodium hydrogen carbonate is precipitated:
\[\mathrm{NaCl} + \mathrm{NH_3} + \mathrm{CO_2} + \mathrm{H_2O} \rightarrow \mathrm{NaHCO_3}\downarrow + \mathrm{NH_4Cl}\]
The sodium hydrogen carbonate is separated and ignited to obtain sodium carbonate:
\[2\,\mathrm{NaHCO_3} \xrightarrow{\ \Delta\ } \mathrm{Na_2CO_3} + \mathrm{H_2O} + \mathrm{CO_2}\uparrow\]
Finally, ammonia is the starting material for nitric acid, made by the Ostwald process. Its first step is the catalytic oxidation you studied in Section 9.7.4, and the process is treated in full in the chapter on nitric acid.
9.9 Tests For Ammonia Gas And The Ammonium Ion
Identifying a gas means collecting observations until only one substance fits them all. Ammonia is one of the easier gases to pin down, because several of its reactions give changes you can see at a glance. Each test below is an observation first; the conclusion follows from it.
Tests For Ammonia Gas
- It has a sharp, characteristic odour. A gas is never sniffed directly from the jar — the mouth of the jar is held away from the face and a little of the gas is wafted towards the nose with the hand.
- It turns moist red litmus paper blue, moist turmeric paper brown, and phenolphthalein solution pink. All three changes are responses to the hydroxide ions formed when the gas dissolves in the moisture present, which is why the paper must be moist.
- It gives dense white fumes with a glass rod moistened with concentrated hydrochloric acid. The fumes are solid ammonium chloride, formed directly from the two gases: \[\mathrm{NH_3}\text{(g)} + \mathrm{HCl}\text{(g)} \rightarrow \mathrm{NH_4Cl}\text{(s)}\]
- When the gas is passed through copper sulphate solution, a blue precipitate is formed, which changes to a deep blue solution as more gas is passed. This two-stage change, described in Section 9.7.7, is very nearly a fingerprint for ammonia.
Tests For The Ammonium Ion In A Salt
- All ammonium salts generally produce ammonia gas when warmed with an alkali such as sodium hydroxide solution. The gas coming off is then identified by the tests above — most conveniently by holding moist red litmus paper at the mouth of the tube. This is the standard test for the ammonium ion, and it is simply the general reaction of Section 9.3.1 used the other way round, as a means of identification rather than of preparation.
- An ammonium salt gives a brown colour, or a brown precipitate, with Nessler’s solution, \(\mathrm{K_2HgI_4}\). This is a slightly alkaline solution of potassium mercuric iodide, made alkaline by adding potassium hydroxide.
9.10 Summary
- Ammonia has the molecular formula \(\mathrm{NH_3}\) and a relative molecular mass of 17.
- The nitrogen atom is joined to three hydrogen atoms by three covalent bonds, and carries one lone pair of electrons. The lone pair is the source of ammonia’s basic character.
- Ammonia occurs in the free state in small amounts in air and in traces in natural water, and in the combined state in ammonium salts. Neither ammonia nor its compounds occur as minerals, because they are far too soluble in water.
- Bacteria decompose the urea present in urine, releasing ammonia: \(\mathrm{CO(NH_2)_2} + \mathrm{H_2O} \rightarrow 2\,\mathrm{NH_3} + \mathrm{CO_2}\)
- Ammonia is supplied in four forms: gaseous ammonia, liquid ammonia, liquor ammonia fortis (a saturated solution of relative density 0.880), and the dilute laboratory bench reagent.
- Any ammonium salt warmed with a caustic alkali — slaked lime, caustic soda or caustic potash — gives ammonia.
- Laboratory preparation uses ammonium chloride with excess calcium hydroxide: \(2\,\mathrm{NH_4Cl} + \mathrm{Ca(OH)_2} \xrightarrow{\ \Delta\ } \mathrm{CaCl_2} + 2\,\mathrm{H_2O} + 2\,\mathrm{NH_3}\)
- Excess calcium hydroxide is used because ammonium chloride is lost by sublimation; calcium hydroxide is chosen because it is cheap and not deliquescent.
- The flask is clamped in a slanting position, mouth downwards, so that the water formed does not run back onto the hot glass and crack it.
- Ammonia is dried over quicklime, \(\mathrm{CaO}\). Concentrated sulphuric acid, phosphorus pentoxide and anhydrous calcium chloride cannot be used, because all three react with the basic gas.
- Ammonia is collected by downward displacement of air, in an inverted jar, because it is lighter than air (vapour density 8.5 against 14.4) and far too soluble to be collected over water.
- Ammonium nitrate and ammonium nitrite cannot be used to prepare ammonia, because both decompose on heating instead of reacting with the alkali.
- Metal nitrides give ammonia with warm water: \(\mathrm{Mg_3N_2} + 6\,\mathrm{H_2O} \rightarrow 3\,\mathrm{Mg(OH)_2} + 2\,\mathrm{NH_3}\) and \(\mathrm{AlN} + 3\,\mathrm{H_2O} \rightarrow \mathrm{Al(OH)_3} + \mathrm{NH_3}\)
- Aqueous ammonia is prepared using an inverted funnel dipped just at the water surface, which gives a large surface for dissolution and prevents back suction of water into the hot flask.
- Haber’s process combines nitrogen and hydrogen in the ratio 1 : 3 by volume: \(\mathrm{N_2} + 3\,\mathrm{H_2} \rightleftharpoons 2\,\mathrm{NH_3} + \text{heat}\)
- Nitrogen comes from the fractional distillation of liquid air; hydrogen from water gas by the Bosch process, or from natural gas.
- The conditions are 450–500 °C, above 200 atm (about 250 atm in practice), finely divided iron as catalyst, and traces of molybdenum or \(\mathrm{Al_2O_3}\) as promoter.
- A low temperature favours the exothermic reaction but is too slow; high pressure favours it because four volumes of reactants give two volumes of product.
- About 15% of the gases are converted in one pass. Ammonia is recovered by liquefaction, or by absorption in water, and the unchanged gases are recirculated to give an eventual yield of 98%.
- Carbon dioxide, carbon monoxide and sulphur compounds must be removed from the reactants, as they act as catalytic poisons.
- Ammonia is a colourless gas with a pungent choking smell, lighter than air, alkaline in nature, boiling at −33.5 °C and freezing at −77.7 °C, and extremely soluble — one volume of water dissolves about 702 volumes of the gas at 20 °C and 1 atm.
- The fountain experiment demonstrates this high solubility: water entering the flask dissolves the gas, the pressure inside falls, and red litmus solution is pushed up the jet tube and enters as a blue fountain.
- Ammonia dissociates thermally at high temperature or on sparking: \(2\,\mathrm{NH_3} \rightleftharpoons \mathrm{N_2} + 3\,\mathrm{H_2}\)
- Dry ammonia is neutral. In water it gives ammonium hydroxide, a weak base, which dissociates partly to give hydroxide ions: \(\mathrm{NH_3} + \mathrm{H_2O} \rightleftharpoons \mathrm{NH_4OH} \rightleftharpoons \mathrm{NH_4^{+}} + \mathrm{OH^{-}}\)
- Aqueous ammonia turns red litmus blue, methyl orange yellow and phenolphthalein pink.
- Ammonia does not burn in air, but burns in oxygen with a greenish yellow flame: \(4\,\mathrm{NH_3} + 3\,\mathrm{O_2} \rightarrow 2\,\mathrm{N_2} + 6\,\mathrm{H_2O}\). The mixture of the two gases is explosive.
- With platinum at 800 °C, ammonia is catalytically oxidised to nitric oxide: \(4\,\mathrm{NH_3} + 5\,\mathrm{O_2} \rightarrow 4\,\mathrm{NO} + 6\,\mathrm{H_2O}\), and the nitric oxide is further oxidised in air to brown nitrogen dioxide.
- Ammonia is a reducing agent. It reduces heated metallic oxides to the metal: \(3\,\mathrm{CuO} + 2\,\mathrm{NH_3} \rightarrow 3\,\mathrm{Cu} + 3\,\mathrm{H_2O} + \mathrm{N_2}\), black copper oxide giving reddish brown copper.
- Lead monoxide and lead(IV) oxide are similarly reduced to greyish lead.
- With excess ammonia, chlorine gives nitrogen and white fumes of ammonium chloride: \(8\,\mathrm{NH_3} + 3\,\mathrm{Cl_2} \rightarrow \mathrm{N_2} + 6\,\mathrm{NH_4Cl}\)
- With excess chlorine, the products are hydrogen chloride and explosive nitrogen trichloride: \(\mathrm{NH_3} + 3\,\mathrm{Cl_2} \rightarrow 3\,\mathrm{HCl} + \mathrm{NCl_3}\)
- Ammonia reacts with acids to form ammonium salts, giving dense white fumes with hydrogen chloride gas: \(\mathrm{NH_3} + \mathrm{HCl} \rightarrow \mathrm{NH_4Cl}\)
- Ammonium chloride shows thermal dissociation, which is reversible; ammonium nitrate shows thermal decomposition, which is not.
- Aqueous ammonia precipitates insoluble metallic hydroxides from solutions of their salts: dirty green \(\mathrm{Fe(OH)_2}\), reddish brown \(\mathrm{Fe(OH)_3}\), white \(\mathrm{Pb(OH)_2}\), white gelatinous \(\mathrm{Zn(OH)_2}\) and pale blue \(\mathrm{Cu(OH)_2}\).
- The zinc and copper precipitates dissolve in excess ammonium hydroxide; the copper one gives an inky blue solution of tetraamminecopper(II) sulphate. These differences make ammonium hydroxide a standard reagent in qualitative analysis.
- Liquid ammonia is used as a refrigerant in ice plants, because vaporising it draws a large quantity of heat from the surroundings.
- Chlorofluorocarbon refrigerants release chlorine free radicals under ultraviolet light; each radical destroys ozone repeatedly and is regenerated, causing ozone depletion. Ammonia neither depletes ozone nor contributes to global warming, but it is poisonous in high concentration and cannot be used with copper pipes.
- Aqueous ammonia dissolves fats and grease, and is used as a cleansing agent.
- Ammonia is used to manufacture nitrogenous fertilisers, including urea: \(2\,\mathrm{NH_3} + \mathrm{CO_2} \xrightarrow{\ 150^\circ\text{C},\ 150\ \text{atm}\ } \mathrm{NH_2CONH_2} + \mathrm{H_2O}\)
- It is also used for explosives, ammonium salts, nylon, rayon, dyes and plastics, for sodium carbonate by the Solvay process, and for nitric acid by the Ostwald process.
- Ammonia gas is identified by its pungent smell, by turning moist red litmus blue and moist turmeric paper brown, and by the dense white fumes it gives with concentrated hydrochloric acid.
- The ammonium ion is identified by warming the salt with an alkali to release ammonia, and confirmed by the brown colour or precipitate given with Nessler’s solution.
9.11 Key Terms
- Back suction — the drawing of water backwards along a delivery tube towards the flask, caused when a very soluble gas dissolves faster than it is produced and lowers the pressure inside the tube.
- Catalytic poison — a substance which destroys the activity of a catalyst, such as carbon monoxide or a sulphur compound in Haber’s process.
- Deliquescent — describes a substance that absorbs moisture from the air and becomes damp on standing.
- Drying agent — a substance that removes water vapour from a gas without reacting with the gas itself; quicklime is the drying agent used for ammonia.
- Free radical — a highly reactive particle carrying an unpaired electron, such as the chlorine atom released when ultraviolet light splits a chlorofluorocarbon molecule.
- Haber’s process — the industrial manufacture of ammonia by combining nitrogen and hydrogen in the ratio 1 : 3 by volume over a heated iron catalyst at high pressure.
- Liquid ammonia — dry ammonia gas that has been liquefied by applying high pressure; it contains no water.
- Liquor ammonia fortis — a saturated solution of ammonia in water, of relative density 0.880, also called 880 ammonia.
- Lone pair — a pair of electrons in the outermost shell of an atom that is not used in bonding; ammonia has one, on its nitrogen atom.
- Nessler’s solution — an alkaline solution of potassium mercuric iodide, \(\mathrm{K_2HgI_4}\), which gives a brown colour or precipitate with the ammonium ion.
- Promoter — a substance added in traces to increase the efficiency of a catalyst, such as molybdenum in Haber’s process.
- Qualitative analysis — the branch of analysis concerned with finding out which substances are present in a sample, rather than how much of each.
- Reducing agent — a substance which removes oxygen from, or adds hydrogen to, another substance, and is itself oxidised in the process.
- Thermal decomposition — the irreversible breaking up of a compound on heating, the products of which do not recombine on cooling.
- Thermal dissociation — the reversible breaking up of a compound on heating, the products of which recombine on cooling.
- Vapour density — the density of a gas compared with that of hydrogen; a ratio with no unit, equal to half the relative molecular mass.
- Weak base — a base that dissociates only partly in water, so that only a small fraction of it releases hydroxide ions at any moment.
9.12 Points To Ponder
- Haber’s process is run at 450–500 °C even though a lower temperature would give a greater proportion of ammonia. What does this compromise tell you about the difference between how much of a product can form and how quickly it forms? Where else in chemistry would such a compromise be needed?
- Ammonia is listed as non-poisonous, and in the same list as a gas that may cause death when inhaled in quantity. What does this suggest about how carefully the word ‘poisonous’ has to be used in chemistry, and about the part played by the amount of a substance?
- Ammonia and oxygen give nitrogen in one experiment and nitric oxide in another. If naming the reactants is not enough to fix the products, what else must be stated before a chemical equation means anything?
- A chlorine free radical destroys an ozone molecule and is then released again, unchanged, to destroy another. Compare this with the way finely divided iron behaves in a catalytic chamber. Is it reasonable to call the chlorine atom a catalyst, and what would be gained or lost by doing so?
- A single lone pair on one nitrogen atom accounts for ammonia’s alkalinity, for the formation of the ammonium ion, and for the deep blue solution it gives with copper. How far should one structural feature be expected to explain the whole chemistry of a substance, and what does the chapter contain that the lone pair does not explain?
- Chlorofluorocarbons were adopted as refrigerants precisely because they were stable and unreactive — the same property that later made them a problem in the upper atmosphere. What does this suggest about the tests a new material should be put through before it comes into wide use?
9.13 Objective Questions
Choose the correct answer from the options given below.
- The gas leaving the flask in the laboratory preparation of ammonia is dried by passing it through a tower packed with: [1]
- concentrated sulphuric acid
- quicklime
- anhydrous calcium chloride
- phosphorus pentoxide
- The promoter used in Haber’s process is: [1]
- finely divided iron
- platinum
- molybdenum
- vanadium pentoxide
- A white gelatinous precipitate which dissolves in excess ammonium hydroxide is obtained from a solution of: [1]
- iron(II) sulphate
- lead nitrate
- zinc nitrate
- iron(III) chloride
- When ammonia is passed in excess over chlorine, the white solid formed is: [1]
- nitrogen trichloride
- ammonium chloride
- hydrogen chloride
- ammonium nitrite
- The vapour density of ammonia is: [1]
- 8.5
- 14.4
- 17
- 34
- Ammonia is catalytically oxidised to nitric oxide in the presence of: [1]
- iron at 500 °C
- platinum at 800 °C
- copper at 300 °C
- nickel at 200 °C
The following questions are assertion–reason based. Choose the answer using the codes given below.
- Both A and R are true, and R is the correct explanation of A.
- Both A and R are true, but R is not the correct explanation of A.
- A is true but R is false.
- A is false but R is true.
- Assertion (A): Ammonia cannot be collected over water. Reason (R): Ammonia is lighter than air. [1]
- Assertion (A): In the laboratory preparation of ammonia, the flask is clamped with its mouth pointing downwards. Reason (R): Water formed during the reaction would otherwise run back onto the heated glass and crack it. [1]
- Assertion (A): Dry ammonia gas turns dry red litmus paper blue. Reason (R): Ammonia is basic in nature. [1]
- Assertion (A): Ammonium nitrate is not used to prepare ammonia in the laboratory. Reason (R): Ammonium nitrate decomposes on heating instead of reacting with the alkali. [1]
- Match the substance in Column I with its role in Column II. [4]
| Column I | Column II |
|---|---|
| a. Quicklime | i. Promoter in Haber’s process |
| b. Molybdenum | ii. Confirmatory test for the ammonium ion |
| c. Nessler’s solution | iii. Nitrogenous fertiliser |
| d. Urea | iv. Drying agent for ammonia |
- Fill in the blanks. [4]
- Ammonia is collected by the …………… displacement of air.
- The catalyst used in Haber’s process is …………… .
- Ammonia turns moist turmeric paper …………… .
- A saturated solution of ammonia in water is called …………… .
- State whether each of the following is true or false. [4]
- Liquid ammonia is a solution of ammonia in water.
- Ammonia supports combustion.
- Ammonia reduces heated copper(II) oxide to copper.
- Ammonium hydroxide is a strong base.
- Name the following. [5]
- The drying agent used for ammonia.
- The nitride of a divalent metal which gives ammonia with warm water.
- The gas produced when excess chlorine reacts with ammonia, other than hydrogen chloride.
- The ammonium salt used in dry cells.
- The coloured precipitate obtained when ammonium hydroxide is added to iron(III) chloride solution.
- Give one word or phrase for each of the following. [4]
- The breaking up of a compound on heating, the products of which recombine on cooling.
- A substance added in traces to increase the efficiency of a catalyst.
- A pair of outer-shell electrons not used in bonding.
- The drawing back of water along a delivery tube towards a hot flask.
- Pick the odd one out in each set and give a reason for your choice. [4]
- Ammonium chloride, ammonium sulphate, ammonium nitrate, sodium chloride.
- Copper(II) oxide, lead(II) oxide, lead(IV) oxide, calcium oxide.
9.14 Check Your Concepts
Define a reducing agent, and state why ammonia can act as one. [2]
State the two properties of ammonia which decide how it is collected in the laboratory. [2]
Give a reason why excess calcium hydroxide is taken in the laboratory preparation of ammonia. [2]
Why is concentrated sulphuric acid not used to dry ammonia, although it is a good drying agent for many other gases? [2]
Differentiate between liquid ammonia and liquor ammonia fortis, stating clearly the basis of difference. [2]
Explain why a bottle of liquor ammonia is cooled in ice before it is opened. [2]
State three conditions used in Haber’s process, and for each one give the reason it is chosen. [3]
With reference to the fountain experiment, explain how the fountain is produced and why it is blue. [3]
Differentiate between thermal dissociation and thermal decomposition, stating the basis of difference, and give one example of each from this chapter. [3]
What would happen if the funnel used in preparing aqueous ammonia were replaced by a plain delivery tube dipping into the water? Explain your answer. [3]
Ammonia is described in this chapter as both a base and a reducing agent. Explain, with one example of each behaviour, which feature of the molecule is responsible in each case. [4]
Only about 15% of the gases are converted in one pass through a Haber plant, yet the eventual yield is 98%. Explain how this is achieved, and state the two ways in which the ammonia is separated from the unchanged gases. [4]
Explain why a chlorofluorocarbon molecule that has long since broken up can go on destroying ozone for a long time, and state two advantages and one disadvantage of using ammonia in its place. [5]
Describe the laboratory preparation of dry ammonia from ammonium chloride, under the headings: reactants, reaction, precaution taken while heating, drying, and collection. [5]
9.15 Equations, Conversions & Identification
Write balanced chemical equations for the following. [5]
- Ammonium sulphate is warmed with caustic soda.
- Aluminium nitride is treated with warm water.
- Ammonia burns in an atmosphere of oxygen.
- Ammonia is passed over heated lead(II) oxide.
- Ammonium hydroxide is added to zinc nitrate solution.
Complete and balance the following equations. [4]
- \(\mathrm{NH_3} + \mathrm{H_2SO_4} \rightarrow\)
- \(\mathrm{CuO} + \mathrm{NH_3} \rightarrow\)
- \(\mathrm{NH_3} + \mathrm{Cl_2}\ \text{(chlorine in excess)} \rightarrow\)
- \(\mathrm{NH_4Cl} \xrightarrow{\ \Delta\ }\)
State what you would observe in each of the following. [4]
- Dry ammonia is passed over heated copper(II) oxide in a hard glass tube.
- A glass rod moistened with concentrated hydrochloric acid is held at the mouth of a jar of ammonia.
- Ammonium hydroxide solution is added, first a little and then in excess, to copper sulphate solution.
- Ammonia and oxygen are passed over platinum heated to 800 °C, and the products are allowed to meet air.
Name the gas evolved and give one test for it in each case. [4]
- Ammonium sulphate is warmed with slaked lime.
- Magnesium nitride is treated with warm water.
- Ammonia is passed over heated lead(IV) oxide.
- Ammonium nitrite is heated.
How would you obtain the following? Give balanced equations and state the conditions where required. [5]
- Nitrogen from ammonia, using a heated metallic oxide.
- Ammonium sulphate from ammonia.
- Urea from ammonia.
- Nitric oxide from ammonia.
- Ammonia from magnesium nitride.
Identify the substance in each case, giving a reason for your answer. [4]
- A white crystalline solid which, on warming with sodium hydroxide solution, gives a gas that turns moist red litmus blue.
- A solution which gives a dirty green precipitate, insoluble in excess ammonium hydroxide.
- A yellow, highly explosive liquid formed when ammonia meets excess chlorine.
- A colourless gas which, on meeting air, at once turns reddish brown.
With reference to Fig. 9.3-1, name the part of the apparatus that holds the quicklime and state its purpose. State also why the flask in the same figure is not clamped upright. [3]
With reference to Fig. 9.6-1, name the liquid taken in the trough, state the purpose of the dropper, and explain why the flask is clamped in an inverted position. [3]