Study of Compounds Class 10 ICSE — Chemistry Chapter 9

The study of compounds class 10 ICSE chapter on ammonia explains how this pungent gas occurs in nature, how it is prepared, and why industry values it. This Chemistry chapter covers the Haber process, the physical and chemical properties, the tests for the gas and its ion, and the uses. Work through the sections below in order, with the equations you will need to reproduce in exams.

Download Study of Compounds Ammonia Class 10 ICSE Chemistry Chapter 9 PDF

Download the Study of Compounds Ammonia Class 10 ICSE Chemistry chapter PDF to revise offline or print for practice. It is the full chapter as printed in the book — 14 pages, 11 sections, 30 questions.

What is inside Chapter 9

14 pages · 11 sections · 30 exercise questions

• Intext Questions: 10 questions
• EXERCISE: 20 questions

9.1 OCCURRENCE

Ammonia is found in two states in nature. In the free state, it is present in small amounts in air and in traces in natural water. In the combined state, it appears in compounds such as ammonium chloride, known in Hindi as Nausadar, and ammonium sulphate.

Ammonium compounds do not occur as minerals. The reason is their high solubility in water. Rain and groundwater dissolve them, so they cannot stay locked in rock.

The pungent smell near a decaying heap of organic matter comes from ammonia. Bacteria break down urea, \mathrm{NH_2CONH_2}, present in urine:

\mathrm{CO(NH_2)_2} + 2\mathrm{H_2O} \rightarrow 2\mathrm{NH_3} + \mathrm{H_2O} + \mathrm{CO_2}

Ammonia is used in four common forms:

• dry ammonia gas
• liquid ammonia, made by liquefying the dry gas under high pressure
• 880 ammonia, a concentrated solution with relative density 0.880, stored in tightly stoppered bottles in a cold place
• a dilute laboratory bench reagent

These forms appear throughout the Class 10 Chemistry hub, which links this chapter to the other compounds you study.

9.2 PREPARATION OF AMMONIA GAS

This section of the ICSE chapter sets up the central idea behind making ammonia. Every laboratory route rests on one pattern. An ammonium salt warmed with an alkali gives a salt, water, and ammonia, with heat over the arrow.

The chapter develops this idea in two places. Section 9.2.1 shows the general methods using different ammonium salts and alkalis. Section 9.2.2 gives the standard laboratory preparation you must be able to describe. Keep the pattern in mind as you read the equations that follow.

9.2.1 General methods of preparation

Any ammonium salt will give ammonia when warmed with a caustic alkali. The alkalis used are slaked lime, \mathrm{Ca(OH)_2}, caustic soda, \mathrm{NaOH}, and caustic potash, \mathrm{KOH}.

The general pattern is:

\text{Ammonium salt} + \text{alkali} \xrightarrow{\Delta} \text{salt} + \text{water} + \text{ammonia}

The book gives four balanced examples. Note the \Delta over each arrow, which means the mixture must be heated.

In each case the metal from the alkali swaps with the ammonium group. The product is a metal salt, water, and ammonia gas. The pattern holds for any ammonium salt, which is why the choice of salt matters little in principle.

9.2.2 Laboratory preparation

Reactants and ratio. The laboratory uses ammonium chloride and calcium hydroxide in a 2:3 ratio by weight. The extra alkali offsets the loss of ammonium chloride by sublimation. Calcium hydroxide is chosen because it is cheap and not deliquescent.

The reaction:

2\mathrm{NH_4Cl} + \mathrm{Ca(OH)_2} \xrightarrow{\Delta} \mathrm{CaCl_2} + 2\mathrm{H_2O} + 2\mathrm{NH_3}

The flask is slanted. The tilt stops condensed water from running back and cracking the hot glass.

Drying. The gas passes through a drying tower packed with lumps of quicklime, \mathrm{CaO}. Three drying agents are rejected because ammonia reacts with each:

2\mathrm{NH_3} + \mathrm{H_2SO_4} \rightarrow (\mathrm{NH_4})_2\mathrm{SO_4}

6\mathrm{NH_3} + \mathrm{P_2O_5} + 3\mathrm{H_2O} \rightarrow 2(\mathrm{NH_4})_3\mathrm{PO_4}

\mathrm{CaCl_2} + 4\mathrm{NH_3} \rightarrow \mathrm{CaCl_2} \cdot 4\mathrm{NH_3}

Collection. The gas is collected in inverted jars by downward displacement of air. Ammonia is lighter than air, with a vapour density of 8.5 against 14.4 for air. It is also too soluble in water to collect over it.

Salts to avoid. Ammonium nitrate and ammonium nitrite are not used here. Both decompose on heating instead of giving ammonia:

\mathrm{NH_4NO_3} \xrightarrow{\Delta} \mathrm{N_2O} + 2\mathrm{H_2O}

\mathrm{NH_4NO_2} \xrightarrow{\Delta} \mathrm{N_2} + 2\mathrm{H_2O}

From metal nitrides. Ammonia also forms when warm water acts on magnesium or aluminium nitride, but this route is costly:

\mathrm{Mg_3N_2} + 6\mathrm{H_2O} \rightarrow 3\mathrm{Mg(OH)_2} + 2\mathrm{NH_3} \uparrow

\mathrm{AlN} + 3\mathrm{H_2O} \rightarrow \mathrm{Al(OH)_3} + \mathrm{NH_3} \uparrow

In the nitride method the jar is held mouth downwards over the delivery tube. The gas is dried through quicklime packed in a drying tower.

9.3 PREPARATION OF AQUEOUS AMMONIA

Ammonia dissolves in water far faster than the flask can produce it. Because of this, water can be sucked back up the delivery tube into the apparatus. This is called back suction.

To stop it, a funnel is attached to the outer end of the delivery tube with rubber tubing. The same method is used to change HCl gas into hydrochloric acid.

Procedure. Water sits in a container, and only a small part of the funnel mouth dips below the surface. As ammonia dissolves, pressure above the water in the funnel drops for a moment. Water rushes up into the funnel, and the funnel rim loses contact with the water. Fresh ammonia then pushes the water back down, and the funnel re-seals.

The result is that ammonia dissolves steadily, with no water drawn back into the flask. The funnel breaks the water seal the instant suction starts, then restores it as gas pressure builds again.

9.4 MANUFACTURE OF AMMONIA (HABER’S PROCESS)

Ammonia is made industrially from nitrogen and hydrogen in a 1:3 ratio by volume. Nitrogen comes from the fractional distillation of liquid air. Hydrogen comes from the Bosch process or from natural gas.

The reaction is:

\mathrm{N_2} + 3\mathrm{H_2} \rightleftharpoons 2\mathrm{NH_3} + \text{heat}

The reaction is reversible, exothermic, and proceeds with a decrease in volume. Four volumes of reactants become two volumes of product.

Because the reaction is exothermic, low temperature favours ammonia. But low temperature slows the reaction, so the optimum is 450 to 500 °C. Because four volumes become two, high pressure favours the product. The heat evolved maintains the temperature, so external heating stops after the start.

Mnemonic for the four values: Fifty and Iron-Moly. That is temperature 450–500 °C, pressure above 200 atm, catalyst finely divided iron, and promoter molybdenum or alumina.

Recovery and recirculation. Ammonia is separated by liquefaction, since it condenses more easily than nitrogen or hydrogen. It is also absorbed in water, in which it dissolves readily. The unreacted gases are recirculated, lifting the eventual yield to about 98%.

Impurities must be removed first. Carbon dioxide, carbon monoxide, and hydrogen sulphide poison the catalyst.

9.5 PHYSICAL PROPERTIES OF AMMONIA

Ammonia has a set of physical properties you should recall quickly. The table lists the ten the book gives.

The fountain experiment demonstrates the high solubility. A flask is filled with ammonia gas. A dropper squeezes a few drops of water into it. Those drops dissolve enough ammonia to create a partial vacuum. Outside pressure then forces red litmus solution up the jet tube, where it emerges as a blue fountain.

The blue colour shows ammonia is basic. The rush of water shows its solubility.

Worked example. At 20 °C and 1 atm, one volume of water absorbs about 702 volumes of ammonia. So 250 mL of water absorbs about:

250 \times 702 = 175{,}500\ \text{mL}

That is roughly 175.5 L of ammonia gas dissolved into a quarter-litre of water.

Safety note. Open a bottle of liquor ammonia only after cooling it in ice or cold water. The pressure inside is high. Cooling lowers it and prevents sudden flushing of the gas.

9.6 CHEMICAL PROPERTIES OF AMMONIA

Thermal dissociation

At high temperature or under electric sparks, ammonia splits into nitrogen and hydrogen. The reaction is reversible:

2\mathrm{NH_3} \rightleftharpoons \mathrm{N_2} + 3\mathrm{H_2}

This is dissociation because the products recombine. Do not confuse it with thermal decomposition, where the change is not reversible.

Basic nature and action on indicators

Dry ammonia is a covalent molecule and is neutral even when liquefied. Its aqueous solution is a weak base. The basicity comes from a lone pair of electrons on the nitrogen atom.

\text{Lone pair} \rightarrow \begin{array}{c} \mathrm{H} \\ | \\ :\mathrm{N}-\mathrm{H} \\ | \\ \mathrm{H} \end{array}

The aqueous solution dissociates partially to give hydroxyl ions:

\mathrm{NH_3} + \mathrm{H_2O} \rightleftharpoons \mathrm{NH_4OH} \rightleftharpoons \mathrm{NH_4^+} + \mathrm{OH^-}

The alkaline behaviour is due to the hydroxyl ions. On indicators, red litmus turns blue, methyl orange turns yellow, and phenolphthalein turns pink.

Reaction with acids

Ammonia accepts protons to form ammonium salts. The book gives three examples:

Ammonium chloride is special. Two colourless gases react directly to form a white solid. Both ammonium chloride and ammonium nitrate show thermal dissociation:

\mathrm{NH_4Cl} \rightleftharpoons \mathrm{NH_3} + \mathrm{HCl}

\mathrm{NH_4NO_3} \rightleftharpoons \mathrm{N_2O} + 2\mathrm{H_2O}

Reaction with soluble metal salts

Ammonium hydroxide gives insoluble precipitates of metallic hydroxides with soluble metal salts, except sodium and potassium. The precipitates differ in colour and solubility. That is why ammonium hydroxide is used in qualitative analysis to identify cations.

Reaction with oxygen

Ammonia does not support combustion. It extinguishes a burning splint and does not burn in air. In pure oxygen it burns with a yellowish green flame:

4\mathrm{NH_3} + 3\mathrm{O_2} \rightarrow 2\mathrm{N_2} + 6\mathrm{H_2O}

The reaction is irreversible and strongly exothermic. A mixture of ammonia and oxygen is explosive and therefore dangerous.

Catalytic oxidation

Over platinum at 800 °C, ammonia reacts with oxygen to give nitric oxide and water vapour:

4\mathrm{NH_3} + 5\mathrm{O_2} \xrightarrow{\mathrm{Pt},\ 800^\circ\mathrm{C}} 4\mathrm{NO} + 6\mathrm{H_2O} + \text{heat}

The nitric oxide then oxidises to brown nitrogen dioxide:

2\mathrm{NO} + \mathrm{O_2} \rightarrow 2\mathrm{NO_2}

Reddish brown vapours appear in the flask. The platinum keeps glowing after heating stops because the oxidation is exothermic.

Reaction with carbon dioxide

At 150 °C and 150 atm, ammonia and carbon dioxide give urea, a valuable nitrogenous fertiliser:

2\mathrm{NH_3} + \mathrm{CO_2} \xrightarrow{150^\circ\mathrm{C},\ 150\ \text{atm}} \mathrm{NH_2CONH_2} + \mathrm{H_2O}

Ammonia as a reducing agent

Ammonia reduces chlorine to hydrogen chloride. The HCl then combines with excess ammonia to form ammonium chloride:

8\mathrm{NH_3} + 3\mathrm{Cl_2} \rightarrow \mathrm{N_2} + 6\mathrm{NH_4Cl}

The yellowish green colour of chlorine disappears, and white fumes of ammonium chloride form. With chlorine in excess, the products are hydrogen chloride and nitrogen trichloride:

\mathrm{NH_3} + 3\mathrm{Cl_2} \rightarrow 3\mathrm{HCl} + \mathrm{NCl_3}

Ammonia also reduces heated metallic oxides. Black copper oxide becomes reddish brown copper:

2\mathrm{NH_3} + 3\mathrm{CuO} \rightarrow 3\mathrm{Cu} + 3\mathrm{H_2O} + \mathrm{N_2} \uparrow

Yellow lead monoxide becomes greyish metallic lead in the same way. The table compares ammonia with oxygen in three situations.

9.7 TESTS FOR AMMONIA GAS AND AMMONIUM ION

Ammonia gas is recognised by six characteristics. Each test gives a clear observation.

1. It has a sharp characteristic odour.
2. It turns moist red litmus blue, moist turmeric paper brown, and phenolphthalein solution pink.
3. It gives dense white fumes with conc. hydrochloric acid:

\mathrm{NH_3} + \mathrm{HCl} \rightarrow \mathrm{NH_4Cl}

4. With copper sulphate solution it gives a blue precipitate, which changes to a deep blue solution in excess ammonia.
5. All ammonium salts give ammonia when warmed with an alkali such as sodium hydroxide.
6. An ammonium salt gives a brown colour or precipitate with Nessler’s solution, \mathrm{K_2HgI_4}.

These tests distinguish ammonium salts from other salts in the board questions.

9.8 USES OF AMMONIA

As a refrigerant. Liquid ammonia takes 5700 calories per mole to vaporise at −33 °C. That heat is drawn from the surroundings, which are consequently cooled. This makes ammonia a good refrigerant.

The main refrigerants used are chlorofluorocarbons, or CFCs. They deplete the ozone layer and contribute to global warming. Ultraviolet rays decompose them into reactive chlorine atoms:

\mathrm{CF_2Cl_2(g)} \xrightarrow{\text{Ultraviolet}} \mathrm{CF_2Cl(g)} + \mathrm{Cl(g)}

\mathrm{Cl(g)} + \mathrm{O_3(g)} \rightarrow \mathrm{ClO(g)} + \mathrm{O_2(g)}

\mathrm{ClO(g)} + \mathrm{O(g)} \rightarrow \mathrm{Cl(g)} + \mathrm{O_2(g)}

Ammonia avoids this. It does not deplete ozone, does not add to global warming, uses less electricity, and has a recognisable odour so leaks are not missed. Its disadvantages are that it is incompatible with copper and poisonous in high concentration.

Other uses. Ammonia serves as a laboratory reagent, giving coloured hydroxide precipitates. Its aqueous solution emulsifies fats and grease, so it removes stains and cleans tiles and windows.

Ammonia is a feedstock for:

• nitrogenous fertilisers such as ammonium sulphate, diammonium hydrogen phosphate, ammonium nitrate, and urea
• explosives, since ammonium nitrate decomposes leaving no residue: \mathrm{NH_4NO_3} \xrightarrow{\Delta} \mathrm{N_2O} + 2\mathrm{H_2O}
• ammonium carbonate, used as smelling salt, and ammonium chloride, used in dry cells and to prepare nitrogen gas
• nylon, rayon, plastics, dyes, and wood pulp
• sodium carbonate by the Solvay process and nitric acid by the Ostwald process

In the Solvay process, ammonia, carbon dioxide, and sodium chloride give sodium hydrogen carbonate:

\mathrm{NaCl} + \mathrm{NH_3} + \mathrm{CO_2} + \mathrm{H_2O} \rightarrow \mathrm{NaHCO_3} \downarrow + \mathrm{NH_4Cl}

2\mathrm{NaHCO_3} \xrightarrow{\Delta} \mathrm{Na_2CO_3} + \mathrm{H_2O} + \mathrm{CO_2} \uparrow

9.9 MANUFACTURE OF NITRIC ACID

The Ostwald process, developed in 1914, makes nitric acid from ammonia in three steps.

Step I — catalytic oxidation. Dry air and dry ammonia in a 10:1 volume ratio pass over platinum gauze at about 800 °C in the catalytic chamber:

4\mathrm{NH_3} + 5\mathrm{O_2} \xrightarrow[katex display="true"]800^\circ\mathrm{C}{\mathrm{Pt}} 4\mathrm{NO} + 6\mathrm{H_2O} + \text{heat}[/katex]

The reaction is exothermic, so the heat maintains the temperature. The hot gases pass through a heat exchanger to the oxidising chamber.

Step II — oxidation of nitric oxide. Nitric oxide combines with oxygen at about 50 °C:

2\mathrm{NO(g)} + \mathrm{O_2(g)} \xrightarrow{50^\circ\mathrm{C}} 2\mathrm{NO_2(g)}

Step III — absorption. Nitrogen dioxide, with excess air, meets warm water trickling down the absorption tower packed with quartz:

4\mathrm{NO_2(g)} + 2\mathrm{H_2O(l)} + \mathrm{O_2(g)} \rightarrow 4\mathrm{HNO_3(aq)}

Concentration. The tower delivers acid above 50%. Distillation gives 68% acid. Distilling over conc. sulphuric acid gives 98% fuming nitric acid. Cooling to −42 °C gives pure 100% acid as colourless crystals.

Air is used in excess because oxygen is only one fifth of air and is needed in all three chambers. Quartz is acid resistant and helps dissolve nitrogen dioxide uniformly. This route leads into the next chapter, the nitric acid chapter, which builds on the product you have just made.

What to practise from the Study of Compounds Class 10 ICSE chapter

The chapter closes with two blocks of questions. Work through both in order.

• Intext Questions: 10 questions covering sections 9.1 through 9.5, with the weight on 9.2.2 and 9.4.
• EXERCISE: 20 questions testing 9.2.2, 9.4, 9.5, 9.6, 9.7, 9.8, and 9.9.

Board question map. Past board questions from 2008 to 2014 are embedded in the exercise. They map to sections like this:

• The 2011 diagram question on the laboratory set-up tests 9.2.2.
• The 2013 process-table question tests 9.4.
• The reducing-agent questions test 9.6.
• The Ostwald questions test 9.9.

Common mistakes to avoid. The table lists the errors students repeat most often.

Verify the Class 10 Chemistry syllabus and exam pattern for this chapter on the official CISCE site.

Study of Compounds Class 10 ICSE frequently asked questions

What are the key reactions in the study of compounds class 10 ICSE chapter on ammonia?

The core reactions are the Haber synthesis, \mathrm{N_2} + 3\mathrm{H_2} \rightleftharpoons 2\mathrm{NH_3}, and the laboratory preparation from ammonium chloride and slaked lime. Also know the catalytic oxidation that gives nitric oxide for the Ostwald process, the reduction of copper oxide, and the formation of urea.

Why is ammonia collected by downward displacement of air and not over water?

Ammonia is lighter than air, with a vapour density of 8.5 against 14.4 for air, so the inverted jar collects it by downward displacement. It cannot be collected over water because it dissolves too readily, with one volume of water absorbing about 702 volumes of the gas.

Which drying agents cannot be used for ammonia and why?

Conc. sulphuric acid, phosphorus pentoxide, and anhydrous calcium chloride all react with ammonia. The acid forms ammonium sulphate, the pentoxide forms ammonium phosphate, and calcium chloride forms the addition compound \mathrm{CaCl_2} \cdot 4\mathrm{NH_3}. Quicklime, \mathrm{CaO}, is used instead.

What is the difference between thermal decomposition and thermal dissociation of ammonium salts?

In thermal dissociation, the change is reversible. Ammonium chloride splits into ammonia and hydrogen chloride, which recombine on cooling. Thermal decomposition is not reversible, as with ammonium nitrite breaking down into nitrogen and water.

How does the fountain experiment prove that ammonia is highly soluble in water?

A few drops of water dissolve enough ammonia to create a partial vacuum in the flask. Outside atmospheric pressure then drives red litmus solution up the jet tube as a blue fountain. The rush of water shows solubility, and the blue colour shows the gas is basic.

Reference: ICSE Class 10 Chemistry textbook material, chapter 9: Study of Compounds Ammonia — CISCE-aligned.

More from this book

• Chapter 2: Chemical Bonding
• Chapter 5: Mole Concept and Stoichimetry
• Chapter 8: Study of Compounds Hydrogen Chloride
• Chapter 10: Study of Compounds Nitric Acid
• Chapter 11: Sulphuric Acid