Ammonia Is Collected in an Upside-Down Jar and Dried Over Quicklime
Prepare ammonia from ammonium chloride and slaked lime with the right apparatus, drying agent and collection method, make it from nitrides and other ammonium salts, understand the conditions chosen for Haber's process, and explain the density, solubility and fountain experiment of ammonia.
Why does ammonia need such different handling from hydrogen chloride?
Hydrogen chloride is collected in an upright jar and dried by bubbling it through concentrated sulphuric acid. Try either of those with ammonia and the experiment fails completely. The ammonia would escape out of the top of an upright jar, and the sulphuric acid would absorb every bit of it.
Both failures come from two properties of ammonia:
- It is lighter than air, so it must be collected in an inverted jar, where it rises and pushes the air out downwards
- It is a base, so it reacts with acidic drying agents, and only a basic drier such as quicklime can be used
A third property it shares with hydrogen chloride — it is extremely soluble in water — rules out collection over water and gives a fountain experiment of its own, this time turning red litmus blue.
Ammonia matters far beyond the laboratory. Most of it is made industrially from nitrogen and hydrogen and turned into fertilisers such as urea and ammonium sulphate, on which a great deal of the world's food production depends. India operates many large fertiliser plants built around this single reaction.
This part covers:
- Laboratory preparation from ammonium chloride and slaked lime, with the apparatus, drying agent, collection and identification
- Other routes — from nitrides and water, and from ammonium salts and alkalis
- Haber's process — raw materials, conditions, catalyst and the reasoning behind each choice
- Density, solubility and the fountain experiment
Part 3 then takes up ammonia's chemical properties and uses.
The idea to carry through this part. Every practical choice — the slope of the flask, the drying agent, the direction of the jar — follows from one property of ammonia or of the reaction that makes it. Learning the reason makes the procedure impossible to forget.
This page covers the ICSE Class 10 Chemistry study of ammonia, preparation: the laboratory method, preparation from nitrides and ammonium salts, Haber's process, and the density, solubility and fountain experiment of ammonia.
Both failures come from two properties of ammonia:
- It is lighter than air, so it must be collected in an inverted jar, where it rises and pushes the air out downwards
- It is a base, so it reacts with acidic drying agents, and only a basic drier such as quicklime can be used
A third property it shares with hydrogen chloride — it is extremely soluble in water — rules out collection over water and gives a fountain experiment of its own, this time turning red litmus blue.
Ammonia matters far beyond the laboratory. Most of it is made industrially from nitrogen and hydrogen and turned into fertilisers such as urea and ammonium sulphate, on which a great deal of the world's food production depends. India operates many large fertiliser plants built around this single reaction.
This part covers:
- Laboratory preparation from ammonium chloride and slaked lime, with the apparatus, drying agent, collection and identification
- Other routes — from nitrides and water, and from ammonium salts and alkalis
- Haber's process — raw materials, conditions, catalyst and the reasoning behind each choice
- Density, solubility and the fountain experiment
Part 3 then takes up ammonia's chemical properties and uses.
The idea to carry through this part. Every practical choice — the slope of the flask, the drying agent, the direction of the jar — follows from one property of ammonia or of the reaction that makes it. Learning the reason makes the procedure impossible to forget.
This page covers the ICSE Class 10 Chemistry study of ammonia, preparation: the laboratory method, preparation from nitrides and ammonium salts, Haber's process, and the density, solubility and fountain experiment of ammonia.
How is ammonia prepared in the laboratory, dried, collected and identified?
A mixture of ammonium chloride and excess slaked lime is heated in a sloping flask; the ammonia is dried over quicklime and collected by downward displacement of air in inverted jars, and identified by its smell, red litmus and white fumes with hydrogen chloride.
Reactants: ammonium chloride and slaked lime, calcium hydroxide, mixed intimately, with the slaked lime in excess.
Why these reactants:
- Slaked lime is cheap and not deliquescent, unlike sodium hydroxide, which absorbs moisture from air
- The excess of lime ensures that all the ammonium chloride reacts
- Ammonium nitrate is never used, because it can decompose explosively on heating
Apparatus and setting:
- A hard-glass round-bottom flask, or a boiling tube, containing the mixture
- The flask is clamped with its mouth tilted slightly downward — so that water formed in the reaction does not run back onto the hot glass and crack it
- A delivery tube leads the gas to a drying tower and then to the collecting jar
Drying agent: quicklime, calcium oxide, in a drying tower. The common alternatives cannot be used:
- Concentrated sulphuric acid reacts with ammonia:
- Phosphorus pentoxide, an acidic oxide, also reacts with it
- Anhydrous calcium chloride absorbs ammonia, forming a compound with it
Collection: by downward displacement of air — the delivery tube goes up into an inverted gas jar, because ammonia is lighter than air. It is not collected over water, because it is extremely soluble.
Observation: a colourless gas with a sharp, pungent smell.
Identification:
- Pungent smell — detected cautiously, by wafting
- Turns moist red litmus paper blue — ammonia is the only common alkaline gas
- Dense white fumes with a glass rod dipped in concentrated hydrochloric acid:
- With copper sulphate solution, a pale blue precipitate that dissolves in excess ammonia to an inky blue solution
Worked example — quantities. What mass of ammonia is obtained from of ammonium chloride, what volume does it occupy at STP, and what mass of slaked lime reacts? N , H , Cl , Ca , O .
- Ammonia: mol
- Slaked lime reacting: mol — and more than this is used, so lime is in excess
How to tell a jar is full. Hold moist red litmus paper at the mouth of the inverted jar; when it turns blue, ammonia has filled the jar and is escaping downward.
An everyday example. The sharp smell in a poorly cleaned public toilet is largely ammonia, released as urea in urine breaks down. It is the same gas, recognised by the same smell used in the laboratory test.
The boundary case — the sloping flask. In the preparation of hydrogen chloride, no water forms in the flask, so the flask stands upright. Here water is a product, and the tilt is the precaution that protects the hot glass from it.
Reactants: ammonium chloride and slaked lime, calcium hydroxide, mixed intimately, with the slaked lime in excess.
Why these reactants:
- Slaked lime is cheap and not deliquescent, unlike sodium hydroxide, which absorbs moisture from air
- The excess of lime ensures that all the ammonium chloride reacts
- Ammonium nitrate is never used, because it can decompose explosively on heating
Apparatus and setting:
- A hard-glass round-bottom flask, or a boiling tube, containing the mixture
- The flask is clamped with its mouth tilted slightly downward — so that water formed in the reaction does not run back onto the hot glass and crack it
- A delivery tube leads the gas to a drying tower and then to the collecting jar
Drying agent: quicklime, calcium oxide, in a drying tower. The common alternatives cannot be used:
- Concentrated sulphuric acid reacts with ammonia:
- Phosphorus pentoxide, an acidic oxide, also reacts with it
- Anhydrous calcium chloride absorbs ammonia, forming a compound with it
Collection: by downward displacement of air — the delivery tube goes up into an inverted gas jar, because ammonia is lighter than air. It is not collected over water, because it is extremely soluble.
Observation: a colourless gas with a sharp, pungent smell.
Identification:
- Pungent smell — detected cautiously, by wafting
- Turns moist red litmus paper blue — ammonia is the only common alkaline gas
- Dense white fumes with a glass rod dipped in concentrated hydrochloric acid:
- With copper sulphate solution, a pale blue precipitate that dissolves in excess ammonia to an inky blue solution
Worked example — quantities. What mass of ammonia is obtained from of ammonium chloride, what volume does it occupy at STP, and what mass of slaked lime reacts? N , H , Cl , Ca , O .
- Ammonia: mol
- Slaked lime reacting: mol — and more than this is used, so lime is in excess
How to tell a jar is full. Hold moist red litmus paper at the mouth of the inverted jar; when it turns blue, ammonia has filled the jar and is escaping downward.
An everyday example. The sharp smell in a poorly cleaned public toilet is largely ammonia, released as urea in urine breaks down. It is the same gas, recognised by the same smell used in the laboratory test.
The boundary case — the sloping flask. In the preparation of hydrogen chloride, no water forms in the flask, so the flask stands upright. Here water is a product, and the tilt is the precaution that protects the hot glass from it.
How is ammonia made from nitrides and from ammonium salts?
Metal nitrides such as magnesium nitride and aluminium nitride react with warm water to release ammonia and form the metal hydroxide, and every ammonium salt releases ammonia when warmed with an alkali.
1. From nitrides, with warm water.
Where the nitride comes from. When magnesium ribbon burns in air, most of it forms magnesium oxide, but some combines with the nitrogen of the air:
Adding a little water to the ash therefore gives a faint smell of ammonia — evidence that nitrogen took part.
2. From ammonium salts, with alkalis, on warming.
The common ionic change in all of them:
The hydroxyl ion removes a proton from the ammonium ion, and the nitrogen gets its lone pair back — the coordinate bond from the bonding chapter, reversed.
Worked check — balancing the magnesium nitride equation.
- Magnesium:
- Nitrogen:
- Hydrogen: on the left; on the right
- Oxygen: on the left; on the right
Balanced.
Worked example — ammonia from magnesium nitride. What volume of ammonia at STP is released when of magnesium nitride reacts completely with water? Mg , N .
Worked example — ammonia from ammonium sulphate. How many moles of ammonia come from of ammonium sulphate warmed with excess sodium hydroxide? S .
An everyday example. Ammonium fertilisers should not be mixed with lime or spread on freshly limed soil. The lime is an alkali, and exactly the reactions above release the fertiliser's nitrogen into the air as ammonia instead of leaving it in the soil for crops.
The boundary case. The reaction with alkalis is the test for ammonium salts as well as a method of preparation. Any solid that gives ammonia when warmed with sodium hydroxide must contain the ammonium ion — no other common ion does that.
1. From nitrides, with warm water.
Where the nitride comes from. When magnesium ribbon burns in air, most of it forms magnesium oxide, but some combines with the nitrogen of the air:
Adding a little water to the ash therefore gives a faint smell of ammonia — evidence that nitrogen took part.
2. From ammonium salts, with alkalis, on warming.
The common ionic change in all of them:
The hydroxyl ion removes a proton from the ammonium ion, and the nitrogen gets its lone pair back — the coordinate bond from the bonding chapter, reversed.
Worked check — balancing the magnesium nitride equation.
- Magnesium:
- Nitrogen:
- Hydrogen: on the left; on the right
- Oxygen: on the left; on the right
Balanced.
Worked example — ammonia from magnesium nitride. What volume of ammonia at STP is released when of magnesium nitride reacts completely with water? Mg , N .
Worked example — ammonia from ammonium sulphate. How many moles of ammonia come from of ammonium sulphate warmed with excess sodium hydroxide? S .
An everyday example. Ammonium fertilisers should not be mixed with lime or spread on freshly limed soil. The lime is an alkali, and exactly the reactions above release the fertiliser's nitrogen into the air as ammonia instead of leaving it in the soil for crops.
The boundary case. The reaction with alkalis is the test for ammonium salts as well as a method of preparation. Any solid that gives ammonia when warmed with sodium hydroxide must contain the ammonium ion — no other common ion does that.
How is ammonia manufactured by Haber's process, and why are those conditions chosen?
Nitrogen and hydrogen in the ratio 1 : 3 by volume are passed at high pressure and about 450–500 °C over finely divided iron with a promoter; ammonia is liquefied out and the unreacted gases are recycled.
Raw materials:
- Nitrogen — from air, by fractional distillation of liquid air
- Hydrogen — from natural gas or from water gas
- Mixed in the ratio by volume, matching the equation
The reaction:
It is reversible and exothermic, and those two facts explain every condition.
Optimum conditions:
- Pressure: high — of the order of a few hundred atmospheres
- Temperature: about to
- Catalyst: finely divided iron
- Promoter: molybdenum, which makes the catalyst more effective
- Purified gases: impurities such as carbon monoxide, carbon dioxide and moisture are removed, because they poison the catalyst
Why high pressure. Four volumes of gas, , become two volumes of ammonia. Increasing the pressure favours the side with fewer gas molecules, so more ammonia forms.
Why a moderate temperature. The forward reaction releases heat, so a lower temperature gives a higher proportion of ammonia — but at low temperatures the reaction is far too slow. About 450–500 °C is a compromise between a good yield and a useful rate.
Why a catalyst. The iron catalyst speeds up the reaction so that equilibrium is reached quickly at the compromise temperature. It does not change the proportion of ammonia at equilibrium — it only gets there faster.
The plant, as a flow of stages:
- Compressor — brings the purified nitrogen-hydrogen mixture to high pressure
- Catalyst chamber — the heated gases pass over finely divided iron and partly combine
- Cooler or condenser — ammonia liquefies far more easily than nitrogen or hydrogen, so it condenses out and is drawn off
- Recycling — the unreacted nitrogen and hydrogen are returned to the compressor and passed over the catalyst again
Worked example — volumes. What volumes of nitrogen and hydrogen, measured under the same conditions, are needed to produce of ammonia if conversion is complete?
Worked example — masses. What mass of ammonia can form from of nitrogen with enough hydrogen?
Check: .
An everyday connection. Most ammonia from such plants becomes urea and other fertilisers, spread on fields across India to supply the nitrogen crops need for growth.
The boundary case. Because the reaction is reversible, only part of the gas mixture turns into ammonia in each pass over the catalyst. Recycling the unreacted gases is what makes the process efficient — nothing is wasted, even though no single pass converts it all.
Raw materials:
- Nitrogen — from air, by fractional distillation of liquid air
- Hydrogen — from natural gas or from water gas
- Mixed in the ratio by volume, matching the equation
The reaction:
It is reversible and exothermic, and those two facts explain every condition.
Optimum conditions:
- Pressure: high — of the order of a few hundred atmospheres
- Temperature: about to
- Catalyst: finely divided iron
- Promoter: molybdenum, which makes the catalyst more effective
- Purified gases: impurities such as carbon monoxide, carbon dioxide and moisture are removed, because they poison the catalyst
Why high pressure. Four volumes of gas, , become two volumes of ammonia. Increasing the pressure favours the side with fewer gas molecules, so more ammonia forms.
Why a moderate temperature. The forward reaction releases heat, so a lower temperature gives a higher proportion of ammonia — but at low temperatures the reaction is far too slow. About 450–500 °C is a compromise between a good yield and a useful rate.
Why a catalyst. The iron catalyst speeds up the reaction so that equilibrium is reached quickly at the compromise temperature. It does not change the proportion of ammonia at equilibrium — it only gets there faster.
The plant, as a flow of stages:
- Compressor — brings the purified nitrogen-hydrogen mixture to high pressure
- Catalyst chamber — the heated gases pass over finely divided iron and partly combine
- Cooler or condenser — ammonia liquefies far more easily than nitrogen or hydrogen, so it condenses out and is drawn off
- Recycling — the unreacted nitrogen and hydrogen are returned to the compressor and passed over the catalyst again
Worked example — volumes. What volumes of nitrogen and hydrogen, measured under the same conditions, are needed to produce of ammonia if conversion is complete?
Worked example — masses. What mass of ammonia can form from of nitrogen with enough hydrogen?
Check: .
An everyday connection. Most ammonia from such plants becomes urea and other fertilisers, spread on fields across India to supply the nitrogen crops need for growth.
The boundary case. Because the reaction is reversible, only part of the gas mixture turns into ammonia in each pass over the catalyst. Recycling the unreacted gases is what makes the process efficient — nothing is wasted, even though no single pass converts it all.
Why is ammonia lighter than air, how soluble is it, and what does its fountain experiment show?
Ammonia's molecular mass of 17 makes it lighter than air, it is extremely soluble in water, and its fountain experiment turns red litmus solution blue, showing both the solubility and the alkaline nature of the solution.
1. Density.
**Air has a vapour density of about , so ammonia is lighter than air — which is why it is collected by downward displacement of air in inverted jars.
2. Solubility. Ammonia is extremely soluble in water: a single volume of water can dissolve several hundred volumes of the gas at room temperature. Its solution is alkaline:**
A concentrated solution is called liquor ammonia. Bottles of it are cooled before opening, because dissolved gas escaping at room temperature builds up pressure inside.
3. The fountain experiment.
Apparatus: a dry round-bottom flask filled with dry ammonia, closed by a cork with a long jet tube, clamped upside down so that the tube dips into a trough of water containing red litmus solution (or phenolphthalein).
Procedure: a few drops of water are introduced into the flask to start the process.
Observation: water rushes up the jet and sprays into the flask as a fountain, and the red litmus turns blue (or phenolphthalein turns pink).
Explanation:
- The first drops dissolve a large volume of ammonia, so the pressure in the flask falls sharply
- Atmospheric pressure pushes water up the tube into the partly empty flask
- Each new spray dissolves more gas, keeping the fountain going
Inference: ammonia is extremely soluble in water, and its solution is alkaline.
Worked comparison — the two fountains in this chapter.
- Hydrogen chloride: blue litmus solution turns red — the solution is acidic
- Ammonia: red litmus solution turns blue — the solution is alkaline
- Both: water rises as a fountain, because both gases are extremely soluble
Worked example — comparing densities at STP.
A litre of ammonia weighs a little more than half as much as a litre of air, confirming why it rises.
An everyday example. Some glass-cleaning solutions contain dilute ammonia, and the sharp smell when a bottle is opened is gas escaping from the solution — the same tendency that makes liquor ammonia build up pressure.
The boundary case — why the flask must be dry and filled with dry gas. A damp flask would let the ammonia dissolve before the experiment starts, lowering the pressure too early. The drying tower of quicklime in the preparation is what makes this experiment possible, linking the two halves of this lesson.
1. Density.
**Air has a vapour density of about , so ammonia is lighter than air — which is why it is collected by downward displacement of air in inverted jars.
2. Solubility. Ammonia is extremely soluble in water: a single volume of water can dissolve several hundred volumes of the gas at room temperature. Its solution is alkaline:**
A concentrated solution is called liquor ammonia. Bottles of it are cooled before opening, because dissolved gas escaping at room temperature builds up pressure inside.
3. The fountain experiment.
Apparatus: a dry round-bottom flask filled with dry ammonia, closed by a cork with a long jet tube, clamped upside down so that the tube dips into a trough of water containing red litmus solution (or phenolphthalein).
Procedure: a few drops of water are introduced into the flask to start the process.
Observation: water rushes up the jet and sprays into the flask as a fountain, and the red litmus turns blue (or phenolphthalein turns pink).
Explanation:
- The first drops dissolve a large volume of ammonia, so the pressure in the flask falls sharply
- Atmospheric pressure pushes water up the tube into the partly empty flask
- Each new spray dissolves more gas, keeping the fountain going
Inference: ammonia is extremely soluble in water, and its solution is alkaline.
Worked comparison — the two fountains in this chapter.
- Hydrogen chloride: blue litmus solution turns red — the solution is acidic
- Ammonia: red litmus solution turns blue — the solution is alkaline
- Both: water rises as a fountain, because both gases are extremely soluble
Worked example — comparing densities at STP.
A litre of ammonia weighs a little more than half as much as a litre of air, confirming why it rises.
An everyday example. Some glass-cleaning solutions contain dilute ammonia, and the sharp smell when a bottle is opened is gas escaping from the solution — the same tendency that makes liquor ammonia build up pressure.
The boundary case — why the flask must be dry and filled with dry gas. A damp flask would let the ammonia dissolve before the experiment starts, lowering the pressure too early. The drying tower of quicklime in the preparation is what makes this experiment possible, linking the two halves of this lesson.
Exam tip
What earns full marks when describing the preparation of ammonia?
State reactants, equation, setting of the flask, drying agent with reasons for rejecting the others, collection with its reason, and a test — then give Haber's conditions with the reason for each.
- Write the balanced equation with ammonium chloride and slaked lime, and mention the lime in excess
- Explain the tilted flask — water formed must not run back onto hot glass
- Name quicklime as the drier, and give why concentrated sulphuric acid, phosphorus pentoxide and anhydrous calcium chloride are unsuitable
- Collect by downward displacement of air in inverted jars, because ammonia is lighter than air and very soluble
- Give two tests: moist red litmus turning blue, and white fumes with concentrated hydrochloric acid
- Write the nitride equations with warm water, balanced
- State Haber's raw materials, ratio, pressure, temperature, catalyst and promoter
- Explain each condition using fewer gas molecules, the exothermic reaction and reaction rate
- Say that ammonia is liquefied and removed and that unreacted gases are recycled
- For the fountain experiment, give apparatus, procedure, observation and inference
The misconception to name. A catalyst does not increase the yield of ammonia at equilibrium. It only makes equilibrium arrive faster; the proportion of ammonia is set by the pressure and temperature. Writing that iron increases the amount of ammonia formed loses the mark.
A second trap. Naming concentrated sulphuric acid as the drying agent for ammonia because it dried hydrogen chloride. Ammonia is a base and reacts with it, so the drier must be basic.
- Write the balanced equation with ammonium chloride and slaked lime, and mention the lime in excess
- Explain the tilted flask — water formed must not run back onto hot glass
- Name quicklime as the drier, and give why concentrated sulphuric acid, phosphorus pentoxide and anhydrous calcium chloride are unsuitable
- Collect by downward displacement of air in inverted jars, because ammonia is lighter than air and very soluble
- Give two tests: moist red litmus turning blue, and white fumes with concentrated hydrochloric acid
- Write the nitride equations with warm water, balanced
- State Haber's raw materials, ratio, pressure, temperature, catalyst and promoter
- Explain each condition using fewer gas molecules, the exothermic reaction and reaction rate
- Say that ammonia is liquefied and removed and that unreacted gases are recycled
- For the fountain experiment, give apparatus, procedure, observation and inference
The misconception to name. A catalyst does not increase the yield of ammonia at equilibrium. It only makes equilibrium arrive faster; the proportion of ammonia is set by the pressure and temperature. Writing that iron increases the amount of ammonia formed loses the mark.
A second trap. Naming concentrated sulphuric acid as the drying agent for ammonia because it dried hydrogen chloride. Ammonia is a base and reacts with it, so the drier must be basic.
Did you know
Why do cold storages and ice factories use ammonia to keep things cold?
Behind many large cold storages that keep potatoes and fruit fresh for months, and in the ice factories that supply fish markets, the working fluid in the cooling system is often ammonia. It is the same pungent gas prepared in this lesson, doing a completely different job.
The cooling depends on a change of state. A liquid absorbs heat from its surroundings as it evaporates, which is why water drying on your skin feels cool. Ammonia is especially good at this.
- It liquefies easily under moderate pressure, because its molecules attract one another fairly strongly
- When the pressure is released, the liquid boils at a very low temperature
- It absorbs a large amount of heat as it evaporates, cooling the pipes around it and everything nearby
The cycle then repeats. A compressor squeezes the ammonia vapour back into a liquid, releasing the absorbed heat outside the cold room, and the liquid is sent round again.
That same easy liquefaction appears in Haber's process. Of the three gases leaving the catalyst chamber — nitrogen, hydrogen and ammonia — only ammonia condenses to a liquid on cooling under pressure, which is exactly how it is separated from the unreacted gases before they are recycled.
Why ammonia is used despite its smell. Its pungent odour is actually a safety feature: even a tiny leak is noticed immediately, long before the concentration becomes dangerous. A gas that warns you of itself is easier to manage in a large plant than one that leaks silently.
And its high solubility provides a second safety measure. Because ammonia dissolves so readily in water, water sprays can be used to absorb escaping gas — the fountain experiment of this lesson, put to work on an industrial scale.
The cooling depends on a change of state. A liquid absorbs heat from its surroundings as it evaporates, which is why water drying on your skin feels cool. Ammonia is especially good at this.
- It liquefies easily under moderate pressure, because its molecules attract one another fairly strongly
- When the pressure is released, the liquid boils at a very low temperature
- It absorbs a large amount of heat as it evaporates, cooling the pipes around it and everything nearby
The cycle then repeats. A compressor squeezes the ammonia vapour back into a liquid, releasing the absorbed heat outside the cold room, and the liquid is sent round again.
That same easy liquefaction appears in Haber's process. Of the three gases leaving the catalyst chamber — nitrogen, hydrogen and ammonia — only ammonia condenses to a liquid on cooling under pressure, which is exactly how it is separated from the unreacted gases before they are recycled.
Why ammonia is used despite its smell. Its pungent odour is actually a safety feature: even a tiny leak is noticed immediately, long before the concentration becomes dangerous. A gas that warns you of itself is easier to manage in a large plant than one that leaks silently.
And its high solubility provides a second safety measure. Because ammonia dissolves so readily in water, water sprays can be used to absorb escaping gas — the fountain experiment of this lesson, put to work on an industrial scale.
Exam relevance
How does ammonia manufacture prepare you for JEE and NEET Chemistry?
This is foundation work for Class 11 Equilibrium, Class 12 Chemical Kinetics and Class 12 The p-Block Elements, all examined in JEE Main and NEET Chemistry.
Where Haber's conditions lead. Class 11 Equilibrium introduces Le Chatelier's principle, and the synthesis of ammonia is its most widely used illustration. Predicting the effect of pressure, temperature, adding reactants or removing product on the position of equilibrium is a recurring question type in both exams, and the reasoning for high pressure and a compromise temperature on this page is exactly that principle applied.
Where the equilibrium constant leads. The same chapter writes and for and relates them through the change in the number of gas moles, **which is here. Numericals built on this equation appear in both exams.
Where the catalyst leads. Class 12 Chemical Kinetics explains that a catalyst provides a path with lower activation energy, speeding up forward and reverse reactions equally. That is why the iron catalyst changes the rate but not the equilibrium yield — a statement tested repeatedly as an assertion-reason item.
Where the preparation and properties lead. Class 12 The p-Block Elements covers the preparation of ammonia, its manufacture by Haber's process, its basic nature and its reactions with metal ions. The laboratory equations and the fountain-experiment properties here are used there directly.
Where the stoichiometry leads.** Gas volumes in the ratio and masses from nitrides and ammonium salts are Class 11 Some Basic Concepts numericals.
Question types to expect. At this level: preparation details, Haber's conditions with reasons, and experiments. In competitive papers: Le Chatelier predictions, and relations, the role of catalysts, and ammonia's reactions, often as assertion-reason or statement-based questions.
The single trap that costs marks. Saying a higher temperature increases the yield of ammonia because it speeds up the reaction. Rate and yield are different: higher temperature increases the rate but, for this exothermic reaction, decreases the equilibrium yield.
A second trap. Claiming a catalyst shifts equilibrium towards ammonia. It does not; it only helps equilibrium be reached sooner, and both exams set options that confuse the two.
Board versus competitive emphasis. The ICSE paper marks the listed conditions, the catalyst and promoter, and a clear reason for each; a competitive paper marks a prediction about equilibrium position or a calculation with . The transferable habit is separating what changes the rate from what changes the yield — the distinction behind every question on this process.
Where Haber's conditions lead. Class 11 Equilibrium introduces Le Chatelier's principle, and the synthesis of ammonia is its most widely used illustration. Predicting the effect of pressure, temperature, adding reactants or removing product on the position of equilibrium is a recurring question type in both exams, and the reasoning for high pressure and a compromise temperature on this page is exactly that principle applied.
Where the equilibrium constant leads. The same chapter writes and for and relates them through the change in the number of gas moles, **which is here. Numericals built on this equation appear in both exams.
Where the catalyst leads. Class 12 Chemical Kinetics explains that a catalyst provides a path with lower activation energy, speeding up forward and reverse reactions equally. That is why the iron catalyst changes the rate but not the equilibrium yield — a statement tested repeatedly as an assertion-reason item.
Where the preparation and properties lead. Class 12 The p-Block Elements covers the preparation of ammonia, its manufacture by Haber's process, its basic nature and its reactions with metal ions. The laboratory equations and the fountain-experiment properties here are used there directly.
Where the stoichiometry leads.** Gas volumes in the ratio and masses from nitrides and ammonium salts are Class 11 Some Basic Concepts numericals.
Question types to expect. At this level: preparation details, Haber's conditions with reasons, and experiments. In competitive papers: Le Chatelier predictions, and relations, the role of catalysts, and ammonia's reactions, often as assertion-reason or statement-based questions.
The single trap that costs marks. Saying a higher temperature increases the yield of ammonia because it speeds up the reaction. Rate and yield are different: higher temperature increases the rate but, for this exothermic reaction, decreases the equilibrium yield.
A second trap. Claiming a catalyst shifts equilibrium towards ammonia. It does not; it only helps equilibrium be reached sooner, and both exams set options that confuse the two.
Board versus competitive emphasis. The ICSE paper marks the listed conditions, the catalyst and promoter, and a clear reason for each; a competitive paper marks a prediction about equilibrium position or a calculation with . The transferable habit is separating what changes the rate from what changes the yield — the distinction behind every question on this process.
Key takeaways
What must you be able to do from this part?
One laboratory method, two other routes, one industrial process and a fountain experiment.
- Laboratory preparation: , slaked lime in excess
- Flask tilted mouth-down so water formed does not crack the hot glass
- Dried over quicklime — not concentrated sulphuric acid, phosphorus pentoxide or anhydrous calcium chloride, which all react with or absorb ammonia
- Collected by downward displacement of air in inverted jars — lighter than air, very soluble
- Tests: pungent smell, moist red litmus turns blue, white fumes with concentrated HCl, inky blue with copper sulphate
- ** of ** gives , or , of ammonia
- Nitrides with warm water: ;
- Ammonium salts with alkalis: — also the test for ammonium ions
- Haber's process: , nitrogen from air, hydrogen from natural gas or water gas, ratio
- Conditions: high pressure, about –, finely divided iron catalyst, molybdenum promoter, purified gases
- High pressure favours fewer gas molecules; moderate temperature balances yield and rate; catalyst speeds equilibrium but does not change the yield
- Ammonia is liquefied out, and unreacted gases are recycled
- **Vapour density — lighter than air
- Extremely soluble**; solution alkaline:
- Fountain experiment: dry flask of dry ammonia, jet into red litmus water — fountain, litmus turns blue
The sharpest self-test is a side-by-side table in your head. Compare the preparation of hydrogen chloride and ammonia on reactants, flask position, drying agent, collection direction and fountain colour — and give the property that explains every single difference.
- Laboratory preparation: , slaked lime in excess
- Flask tilted mouth-down so water formed does not crack the hot glass
- Dried over quicklime — not concentrated sulphuric acid, phosphorus pentoxide or anhydrous calcium chloride, which all react with or absorb ammonia
- Collected by downward displacement of air in inverted jars — lighter than air, very soluble
- Tests: pungent smell, moist red litmus turns blue, white fumes with concentrated HCl, inky blue with copper sulphate
- ** of ** gives , or , of ammonia
- Nitrides with warm water: ;
- Ammonium salts with alkalis: — also the test for ammonium ions
- Haber's process: , nitrogen from air, hydrogen from natural gas or water gas, ratio
- Conditions: high pressure, about –, finely divided iron catalyst, molybdenum promoter, purified gases
- High pressure favours fewer gas molecules; moderate temperature balances yield and rate; catalyst speeds equilibrium but does not change the yield
- Ammonia is liquefied out, and unreacted gases are recycled
- **Vapour density — lighter than air
- Extremely soluble**; solution alkaline:
- Fountain experiment: dry flask of dry ammonia, jet into red litmus water — fountain, litmus turns blue
The sharpest self-test is a side-by-side table in your head. Compare the preparation of hydrogen chloride and ammonia on reactants, flask position, drying agent, collection direction and fountain colour — and give the property that explains every single difference.