Ammonia Will Not Burn in Air, yet Burns in Pure Oxygen and Glows Over Platinum
Write the equations for ammonia burning in oxygen and for its catalytic oxidation over platinum, see ammonia act as a reducing agent with hot copper oxide and lead oxide, compare its reactions with chlorine in excess and in short supply, and review aqueous ammonia with acids and metal salts, its uses and its tests.
What kinds of reactions does ammonia take part in?
Part 2 treated ammonia as a gas to be made and collected. This part treats it as a reactant, and it turns out to play three quite different roles.
1. Ammonia as a fuel. Ammonia does not burn in air, but in pure oxygen it burns with a yellowish flame. Over a hot platinum catalyst, it reacts with oxygen in a completely different way, giving nitric oxide — the first step in making nitric acid.
2. Ammonia as a reducing agent. The hydrogen in ammonia can take oxygen away from hot metal oxides such as copper oxide and lead oxide, leaving the metal and releasing nitrogen. With chlorine, ammonia gives up its hydrogen to form hydrogen chloride.
3. Ammonia as a base. Its lone pair captures protons from acids to form ammonium salts, and its solution precipitates metal hydroxides from salt solutions — sometimes dissolving them again as deep blue or colourless complexes.
Knowing which role applies explains every equation in this part.
- Burning and catalytic oxidation — ammonia is oxidised by oxygen
- Copper oxide, lead oxide and chlorine — ammonia reduces the other substance
- Acids and metal salts — ammonia acts as a base
The uses of ammonia follow the same roles. It is turned into nitric acid and fertilisers, it serves as a cleaning agent and a laboratory reagent, and it is the working fluid in large refrigeration plants.
An everyday reason to know one of these reactions well. Household cleaners based on ammonia and those based on chlorine bleach must never be mixed, because ammonia and chlorine compounds react to give toxic gases. The chlorine reactions in this part show the kind of chemistry responsible.
A thread from earlier chapters. The lone pair on nitrogen explains the base reactions, the complex ions from analytical chemistry reappear with copper and zinc, and Gay Lussac's law gives the volume relationships in the chlorine reactions. Very little here is new in principle; what is new is the detail of conditions and observations.
This page covers the ICSE Class 10 Chemistry study of ammonia, properties and uses: burning and catalytic oxidation, reactions with hydrogen chloride, copper oxide and lead oxide, reactions with chlorine, aqueous ammonia with acids and salts, uses and tests.
1. Ammonia as a fuel. Ammonia does not burn in air, but in pure oxygen it burns with a yellowish flame. Over a hot platinum catalyst, it reacts with oxygen in a completely different way, giving nitric oxide — the first step in making nitric acid.
2. Ammonia as a reducing agent. The hydrogen in ammonia can take oxygen away from hot metal oxides such as copper oxide and lead oxide, leaving the metal and releasing nitrogen. With chlorine, ammonia gives up its hydrogen to form hydrogen chloride.
3. Ammonia as a base. Its lone pair captures protons from acids to form ammonium salts, and its solution precipitates metal hydroxides from salt solutions — sometimes dissolving them again as deep blue or colourless complexes.
Knowing which role applies explains every equation in this part.
- Burning and catalytic oxidation — ammonia is oxidised by oxygen
- Copper oxide, lead oxide and chlorine — ammonia reduces the other substance
- Acids and metal salts — ammonia acts as a base
The uses of ammonia follow the same roles. It is turned into nitric acid and fertilisers, it serves as a cleaning agent and a laboratory reagent, and it is the working fluid in large refrigeration plants.
An everyday reason to know one of these reactions well. Household cleaners based on ammonia and those based on chlorine bleach must never be mixed, because ammonia and chlorine compounds react to give toxic gases. The chlorine reactions in this part show the kind of chemistry responsible.
A thread from earlier chapters. The lone pair on nitrogen explains the base reactions, the complex ions from analytical chemistry reappear with copper and zinc, and Gay Lussac's law gives the volume relationships in the chlorine reactions. Very little here is new in principle; what is new is the detail of conditions and observations.
This page covers the ICSE Class 10 Chemistry study of ammonia, properties and uses: burning and catalytic oxidation, reactions with hydrogen chloride, copper oxide and lead oxide, reactions with chlorine, aqueous ammonia with acids and salts, uses and tests.
What happens when ammonia burns in oxygen, and what is its catalytic oxidation?
Ammonia burns in pure oxygen with a yellowish-green flame to give nitrogen and water, but over a platinum catalyst at about 800 °C it is oxidised instead to nitric oxide and water, releasing so much heat that the platinum keeps glowing.
1. Ammonia and air. Ammonia does not burn in air, and it does not support combustion — a burning splint is extinguished in it.
2. Burning in oxygen. A jet of ammonia burns in pure oxygen with a yellowish-green flame:
3. Catalytic oxidation. A mixture of ammonia and air or oxygen is passed over platinum at about :
The nitric oxide immediately combines with more oxygen, forming reddish-brown nitrogen dioxide:
Observations in the laboratory demonstration:
- A red-hot platinum wire held just above concentrated ammonia solution, with oxygen bubbled through, continues to glow without further heating
- Reddish-brown fumes of nitrogen dioxide appear in the flask
The difference between the two reactions. Without a catalyst, the nitrogen of ammonia ends up as nitrogen gas. With platinum, it ends up as nitric oxide — a product that can be turned into nitric acid. The catalyst decides which reaction happens, not merely how fast.
Worked check — balancing the catalytic equation.
- Nitrogen:
- Hydrogen: on the left; on the right
- Oxygen: on the left; on the right
Balanced.
Worked example — oxygen needed in each case. What volume of oxygen, under the same conditions, reacts with of ammonia when it burns, and when it is catalytically oxidised?
The catalytic reaction uses more oxygen, because the nitrogen is also oxidised, to nitric oxide, instead of being released as nitrogen gas.
An everyday connection. Every nitric acid plant begins with the catalytic equation above, carried out on fine platinum gauze. The fertilisers made from that nitric acid — such as ammonium nitrate — trace back to this single step.
The boundary case — ammonia as a fuel. Because burning ammonia gives only nitrogen and water, with no carbon dioxide, it is being studied as a carbon-free fuel. Its reluctance to burn in air is the main practical difficulty, which is why it needs pure oxygen or special conditions.
1. Ammonia and air. Ammonia does not burn in air, and it does not support combustion — a burning splint is extinguished in it.
2. Burning in oxygen. A jet of ammonia burns in pure oxygen with a yellowish-green flame:
3. Catalytic oxidation. A mixture of ammonia and air or oxygen is passed over platinum at about :
The nitric oxide immediately combines with more oxygen, forming reddish-brown nitrogen dioxide:
Observations in the laboratory demonstration:
- A red-hot platinum wire held just above concentrated ammonia solution, with oxygen bubbled through, continues to glow without further heating
- Reddish-brown fumes of nitrogen dioxide appear in the flask
The difference between the two reactions. Without a catalyst, the nitrogen of ammonia ends up as nitrogen gas. With platinum, it ends up as nitric oxide — a product that can be turned into nitric acid. The catalyst decides which reaction happens, not merely how fast.
Worked check — balancing the catalytic equation.
- Nitrogen:
- Hydrogen: on the left; on the right
- Oxygen: on the left; on the right
Balanced.
Worked example — oxygen needed in each case. What volume of oxygen, under the same conditions, reacts with of ammonia when it burns, and when it is catalytically oxidised?
The catalytic reaction uses more oxygen, because the nitrogen is also oxidised, to nitric oxide, instead of being released as nitrogen gas.
An everyday connection. Every nitric acid plant begins with the catalytic equation above, carried out on fine platinum gauze. The fertilisers made from that nitric acid — such as ammonium nitrate — trace back to this single step.
The boundary case — ammonia as a fuel. Because burning ammonia gives only nitrogen and water, with no carbon dioxide, it is being studied as a carbon-free fuel. Its reluctance to burn in air is the main practical difficulty, which is why it needs pure oxygen or special conditions.
How does ammonia react with hydrogen chloride, hot copper oxide and lead oxide?
Ammonia combines with hydrogen chloride to form white fumes of ammonium chloride, and it acts as a reducing agent towards hot copper(II) oxide and lead oxides, reducing them to the metal while it is oxidised to nitrogen and water.
1. With hydrogen chloride — a base with an acid.
Observation: dense white fumes of solid ammonium chloride form where the two gases meet. This is the standard test for either gas.
2. With hot copper(II) oxide — ammonia as a reducing agent. Dry ammonia is passed over heated copper(II) oxide:
Observations:
- Black copper(II) oxide turns reddish-brown, as copper metal forms
- Droplets of a colourless liquid collect in the cooler part of the tube; they turn white anhydrous copper sulphate blue, showing they are water
- A colourless gas, nitrogen, is given off; it does not burn and does not support combustion
3. With lead oxides — the same reducing action. Heated lead(IV) oxide, a dark brown solid, is reduced to lead:
And yellow lead(II) oxide reacts similarly:
Observation: the coloured oxide turns to greyish, metallic lead, with water and nitrogen formed.
What is oxidised and what is reduced. In :
- Copper(II) oxide is reduced — it loses oxygen and becomes copper
- Ammonia is oxidised — its hydrogen gains oxygen to form water, and its nitrogen is released as nitrogen gas
- Ammonia is the reducing agent
Worked check — balancing the lead(IV) oxide equation.
- Lead:
- Oxygen: on the left; on the right
- Nitrogen: on the left; on the right
- Hydrogen: on the left; on the right
Balanced.
Worked example — products from copper oxide. What mass of copper and what volume of nitrogen at STP form when of ammonia reacts completely with hot copper(II) oxide? Cu .
- Copper: mol
- Nitrogen: mol
A laboratory link. This reaction is a convenient laboratory method for preparing small samples of nitrogen — a reminder that ammonia contains nitrogen that can be released simply by taking away its hydrogen.
The boundary case. Ammonia behaves as a base with hydrogen chloride but as a reducing agent with copper oxide. The same molecule plays different roles depending on its partner: an acid takes its lone pair, while a hot metal oxide takes its hydrogen.
1. With hydrogen chloride — a base with an acid.
Observation: dense white fumes of solid ammonium chloride form where the two gases meet. This is the standard test for either gas.
2. With hot copper(II) oxide — ammonia as a reducing agent. Dry ammonia is passed over heated copper(II) oxide:
Observations:
- Black copper(II) oxide turns reddish-brown, as copper metal forms
- Droplets of a colourless liquid collect in the cooler part of the tube; they turn white anhydrous copper sulphate blue, showing they are water
- A colourless gas, nitrogen, is given off; it does not burn and does not support combustion
3. With lead oxides — the same reducing action. Heated lead(IV) oxide, a dark brown solid, is reduced to lead:
And yellow lead(II) oxide reacts similarly:
Observation: the coloured oxide turns to greyish, metallic lead, with water and nitrogen formed.
What is oxidised and what is reduced. In :
- Copper(II) oxide is reduced — it loses oxygen and becomes copper
- Ammonia is oxidised — its hydrogen gains oxygen to form water, and its nitrogen is released as nitrogen gas
- Ammonia is the reducing agent
Worked check — balancing the lead(IV) oxide equation.
- Lead:
- Oxygen: on the left; on the right
- Nitrogen: on the left; on the right
- Hydrogen: on the left; on the right
Balanced.
Worked example — products from copper oxide. What mass of copper and what volume of nitrogen at STP form when of ammonia reacts completely with hot copper(II) oxide? Cu .
- Copper: mol
- Nitrogen: mol
A laboratory link. This reaction is a convenient laboratory method for preparing small samples of nitrogen — a reminder that ammonia contains nitrogen that can be released simply by taking away its hydrogen.
The boundary case. Ammonia behaves as a base with hydrogen chloride but as a reducing agent with copper oxide. The same molecule plays different roles depending on its partner: an acid takes its lone pair, while a hot metal oxide takes its hydrogen.
How does ammonia react with chlorine when chlorine is in excess and when ammonia is in excess?
With excess ammonia, chlorine's colour disappears and dense white fumes of ammonium chloride form along with nitrogen; with excess chlorine, a yellow, highly explosive oily liquid, nitrogen trichloride, forms together with hydrogen chloride.
1. Ammonia in excess.
Observations:
- The greenish-yellow colour of chlorine disappears
- Dense white fumes of ammonium chloride form
- Colourless nitrogen is left
How it happens, in two steps. Chlorine first takes hydrogen from ammonia; the hydrogen chloride formed then reacts with the extra ammonia:
Adding the two gives the overall equation — molecules of ammonia in all.
2. Chlorine in excess.
Observations:
- A yellow, oily liquid, nitrogen trichloride, forms
- Nitrogen trichloride is highly explosive
- Hydrogen chloride is also formed
Why the product depends on which gas is in excess.
- With plenty of ammonia, any hydrogen chloride formed is immediately captured as ammonium chloride, and chlorine removes the hydrogen from ammonia completely, releasing nitrogen
- With plenty of chlorine, chlorine atoms replace all three hydrogen atoms of each ammonia molecule, giving nitrogen trichloride
Worked check — balancing the excess-ammonia equation.
- Nitrogen: on the left; on the right
- Hydrogen: on the left; on the right
- Chlorine: on the left; on the right
Balanced.
Worked example — volumes of gas. In the excess-ammonia reaction, of chlorine reacts completely. What volume of ammonia is used and what volume of gas is produced, all at the same temperature and pressure?
Ammonium chloride is a solid, so it adds no gas volume — only of nitrogen remains as gas from the reaction.
The comparison in one list:
- Excess ammonia: products and ; white fumes; chlorine colour vanishes
- Excess chlorine: products and ; yellow explosive oily liquid
An everyday safety example. Bleaching powder and many toilet and floor cleaners release chlorine compounds, while others contain ammonia. Mixing the two kinds releases toxic gases formed by reactions of the kind shown here, which is why cleaning products carry warnings never to combine them.
The boundary case. The same two reactants give a harmless white solid in one proportion and a dangerous explosive in another. In chemistry, the ratio of reactants can change not just the amount of product but its identity — which is why the excess reactant must always be named.
1. Ammonia in excess.
Observations:
- The greenish-yellow colour of chlorine disappears
- Dense white fumes of ammonium chloride form
- Colourless nitrogen is left
How it happens, in two steps. Chlorine first takes hydrogen from ammonia; the hydrogen chloride formed then reacts with the extra ammonia:
Adding the two gives the overall equation — molecules of ammonia in all.
2. Chlorine in excess.
Observations:
- A yellow, oily liquid, nitrogen trichloride, forms
- Nitrogen trichloride is highly explosive
- Hydrogen chloride is also formed
Why the product depends on which gas is in excess.
- With plenty of ammonia, any hydrogen chloride formed is immediately captured as ammonium chloride, and chlorine removes the hydrogen from ammonia completely, releasing nitrogen
- With plenty of chlorine, chlorine atoms replace all three hydrogen atoms of each ammonia molecule, giving nitrogen trichloride
Worked check — balancing the excess-ammonia equation.
- Nitrogen: on the left; on the right
- Hydrogen: on the left; on the right
- Chlorine: on the left; on the right
Balanced.
Worked example — volumes of gas. In the excess-ammonia reaction, of chlorine reacts completely. What volume of ammonia is used and what volume of gas is produced, all at the same temperature and pressure?
Ammonium chloride is a solid, so it adds no gas volume — only of nitrogen remains as gas from the reaction.
The comparison in one list:
- Excess ammonia: products and ; white fumes; chlorine colour vanishes
- Excess chlorine: products and ; yellow explosive oily liquid
An everyday safety example. Bleaching powder and many toilet and floor cleaners release chlorine compounds, while others contain ammonia. Mixing the two kinds releases toxic gases formed by reactions of the kind shown here, which is why cleaning products carry warnings never to combine them.
The boundary case. The same two reactants give a harmless white solid in one proportion and a dangerous explosive in another. In chemistry, the ratio of reactants can change not just the amount of product but its identity — which is why the excess reactant must always be named.
How does aqueous ammonia react with acids and metal salts, what is ammonia used for, and how is it tested?
Aqueous ammonia neutralises acids to form ammonium salts and precipitates metal hydroxides from salt solutions, some of which dissolve in excess; ammonia is used for fertilisers, nitric acid, refrigeration and cleaning, and identified by its smell, red litmus and white fumes with hydrogen chloride.
1. With acids — ammonium salts.
2. With metal salt solutions — hydroxide precipitates.
Iron(III) chloride — reddish-brown precipitate, insoluble in excess:
Iron(II) sulphate — dirty green precipitate, insoluble in excess:
Lead nitrate — chalky white precipitate, insoluble in excess:
Zinc nitrate — white gelatinous precipitate, soluble in excess, giving a colourless solution:
Copper sulphate — pale blue precipitate, soluble in excess, giving an inky blue solution:
3. Uses of ammonia:
- Manufacture of fertilisers — urea, ammonium sulphate, ammonium nitrate and ammonium phosphate
- Manufacture of nitric acid by catalytic oxidation
- Refrigerant in ice plants and cold storages, because liquid ammonia absorbs much heat as it evaporates
- Cleansing agent — dilute ammonia solution removes grease and stains from glass and fabrics
- Laboratory reagent for precipitating metal hydroxides
- Manufacture of other chemicals, including some plastics and synthetic fibres
4. Tests for ammonia:
- Characteristic pungent smell
- Turns moist red litmus blue
- Dense white fumes with a rod dipped in concentrated hydrochloric acid
- With copper sulphate solution, a pale blue precipitate dissolving in excess to an inky blue solution
Worked check — balancing the zinc complex equation.
- Zinc:
- Nitrogen: on the left; on the right
- Hydrogen: on the left; on the right
- Oxygen: on the left; on the right
Balanced.
Worked example — fertiliser from ammonia. What mass of ammonium sulphate can be made from of ammonia?
An everyday example. Bags of ammonium sulphate and urea sold at agricultural supply shops across rural India are the end products of ammonia made in fertiliser plants — the most important single use of the gas.
The boundary case — two blue results. The pale blue precipitate and the inky blue solution with copper sulphate are different substances. **The first is copper hydroxide; the second is the soluble complex ion .** Describing only blue loses the distinction a test answer needs.
1. With acids — ammonium salts.
2. With metal salt solutions — hydroxide precipitates.
Iron(III) chloride — reddish-brown precipitate, insoluble in excess:
Iron(II) sulphate — dirty green precipitate, insoluble in excess:
Lead nitrate — chalky white precipitate, insoluble in excess:
Zinc nitrate — white gelatinous precipitate, soluble in excess, giving a colourless solution:
Copper sulphate — pale blue precipitate, soluble in excess, giving an inky blue solution:
3. Uses of ammonia:
- Manufacture of fertilisers — urea, ammonium sulphate, ammonium nitrate and ammonium phosphate
- Manufacture of nitric acid by catalytic oxidation
- Refrigerant in ice plants and cold storages, because liquid ammonia absorbs much heat as it evaporates
- Cleansing agent — dilute ammonia solution removes grease and stains from glass and fabrics
- Laboratory reagent for precipitating metal hydroxides
- Manufacture of other chemicals, including some plastics and synthetic fibres
4. Tests for ammonia:
- Characteristic pungent smell
- Turns moist red litmus blue
- Dense white fumes with a rod dipped in concentrated hydrochloric acid
- With copper sulphate solution, a pale blue precipitate dissolving in excess to an inky blue solution
Worked check — balancing the zinc complex equation.
- Zinc:
- Nitrogen: on the left; on the right
- Hydrogen: on the left; on the right
- Oxygen: on the left; on the right
Balanced.
Worked example — fertiliser from ammonia. What mass of ammonium sulphate can be made from of ammonia?
An everyday example. Bags of ammonium sulphate and urea sold at agricultural supply shops across rural India are the end products of ammonia made in fertiliser plants — the most important single use of the gas.
The boundary case — two blue results. The pale blue precipitate and the inky blue solution with copper sulphate are different substances. **The first is copper hydroxide; the second is the soluble complex ion .** Describing only blue loses the distinction a test answer needs.
Exam tip
What earns full marks on the properties of ammonia?
Give each reaction its condition, its balanced equation and its observation, and state the role ammonia plays — fuel, reducing agent or base.
- Say ammonia does not burn in air before describing burning in oxygen
- Give the flame colour for burning in oxygen — yellowish-green
- Name platinum and the temperature for catalytic oxidation, and note the glowing wire and brown fumes
- Write both oxidation products correctly: nitrogen when burning, nitric oxide over platinum
- With copper oxide, give the colour change black to reddish-brown, the water test and nitrogen
- Identify ammonia as the reducing agent in the metal oxide reactions
- For chlorine, name the excess reactant before writing the equation
- Excess ammonia: , white fumes
- Excess chlorine: , yellow explosive liquid
- For salt solutions, give the precipitate colour and its behaviour in excess ammonia
The misconception to name. Catalytic oxidation of ammonia does not give nitrogen dioxide directly. The catalyst gives nitric oxide, which then combines with oxygen to form nitrogen dioxide. Writing as the product of the platinum step loses the equation mark.
A second trap. Swapping the chlorine equations. White fumes of ammonium chloride can form only when ammonia is left over to capture the hydrogen chloride — which is the excess-ammonia case, not the excess-chlorine one.
- Say ammonia does not burn in air before describing burning in oxygen
- Give the flame colour for burning in oxygen — yellowish-green
- Name platinum and the temperature for catalytic oxidation, and note the glowing wire and brown fumes
- Write both oxidation products correctly: nitrogen when burning, nitric oxide over platinum
- With copper oxide, give the colour change black to reddish-brown, the water test and nitrogen
- Identify ammonia as the reducing agent in the metal oxide reactions
- For chlorine, name the excess reactant before writing the equation
- Excess ammonia: , white fumes
- Excess chlorine: , yellow explosive liquid
- For salt solutions, give the precipitate colour and its behaviour in excess ammonia
The misconception to name. Catalytic oxidation of ammonia does not give nitrogen dioxide directly. The catalyst gives nitric oxide, which then combines with oxygen to form nitrogen dioxide. Writing as the product of the platinum step loses the equation mark.
A second trap. Swapping the chlorine equations. White fumes of ammonium chloride can form only when ammonia is left over to capture the hydrogen chloride — which is the excess-ammonia case, not the excess-chlorine one.
Did you know
Why does a platinum wire keep glowing over ammonia solution without being heated?
Heat a coil of platinum wire until it glows red, then lower it into a flask containing a little concentrated ammonia solution through which oxygen is gently bubbled. Take the flame away. Instead of cooling and going dark, the wire goes on glowing, and reddish-brown fumes begin to fill the flask.
Nothing is heating the wire from outside. The heat comes from the reaction happening on the surface of the platinum itself.
- Ammonia and oxygen molecules settle on the hot platinum surface
- There they react quickly, forming nitric oxide and water
- The reaction releases a great deal of heat, which keeps the wire hot enough to continue catalysing
The platinum is not used up, which is what makes it a catalyst. It simply provides the surface on which the reaction runs, and the heat released pays for keeping the surface hot.
The reddish-brown fumes are a second reaction. The colourless nitric oxide formed at the wire meets more oxygen in the flask:
and nitrogen dioxide is brown, so the colour builds up as the experiment continues.
This small demonstration is a model of a large industrial process. In a nitric acid plant, ammonia and air pass through fine gauzes of platinum, which are heated only to start the process. After that, the heat of the reaction itself keeps the gauze at working temperature, just as it keeps the wire glowing in the flask.
And it shows the key idea of catalysis in a single glance. Without platinum, ammonia would not react with oxygen at that temperature in this way; with platinum, it reacts fast enough to sustain its own heat. The catalyst does not change what is possible, but it changes which reaction actually happens and how quickly — here, the difference between ammonia sitting inert and a wire glowing on its own.
Nothing is heating the wire from outside. The heat comes from the reaction happening on the surface of the platinum itself.
- Ammonia and oxygen molecules settle on the hot platinum surface
- There they react quickly, forming nitric oxide and water
- The reaction releases a great deal of heat, which keeps the wire hot enough to continue catalysing
The platinum is not used up, which is what makes it a catalyst. It simply provides the surface on which the reaction runs, and the heat released pays for keeping the surface hot.
The reddish-brown fumes are a second reaction. The colourless nitric oxide formed at the wire meets more oxygen in the flask:
and nitrogen dioxide is brown, so the colour builds up as the experiment continues.
This small demonstration is a model of a large industrial process. In a nitric acid plant, ammonia and air pass through fine gauzes of platinum, which are heated only to start the process. After that, the heat of the reaction itself keeps the gauze at working temperature, just as it keeps the wire glowing in the flask.
And it shows the key idea of catalysis in a single glance. Without platinum, ammonia would not react with oxygen at that temperature in this way; with platinum, it reacts fast enough to sustain its own heat. The catalyst does not change what is possible, but it changes which reaction actually happens and how quickly — here, the difference between ammonia sitting inert and a wire glowing on its own.
Exam relevance
How are ammonia's reactions examined in JEE and NEET Chemistry?
This is foundation work for Class 12 The p-Block Elements, Class 11 Redox Reactions, Class 12 Coordination Compounds and Class 11 Equilibrium, all examined in JEE Main and NEET Chemistry.
Where ammonia's reactions lead. Class 12 The p-Block Elements covers the properties of ammonia among the compounds of group 15, including its basic nature, its reactions with metal salt solutions and its complex formation with copper ions. The catalytic oxidation to nitric oxide is the first step of the manufacture of nitric acid treated in the same chapter.
Where the reducing reactions lead. Class 11 Redox Reactions assigns oxidation numbers: nitrogen in ammonia is , in nitrogen gas and in nitric oxide . So ammonia is oxidised both when it burns to nitrogen and when it is catalytically oxidised to nitric oxide, and the reduction of copper oxide can be analysed the same way. Balancing such equations by oxidation-number change is a standard skill in both exams.
Where the chlorine reactions lead. The dependence of the product on which reactant is in excess is a limiting-reagent idea from Class 11 Some Basic Concepts, and the volume relationship worked above uses Gay Lussac's law. Questions combining stoichiometry with a named reaction appear as numericals.
Where the complexes lead. Class 12 Coordination Compounds studies and directly — naming, coordination number and the role of ammonia as a ligand.
Where aqueous ammonia leads. Class 11 Equilibrium treats ammonium hydroxide as a weak base and the mixture of ammonia with ammonium chloride as a basic buffer. pH and buffer calculations on this system appear in both exams.
Question types to expect. At this level: equations with conditions and observations, reducing-agent identification, and uses and tests. In competitive papers: oxidation numbers of nitrogen, balancing redox equations, limiting-reagent products, complex naming, and buffer calculations.
The single trap that costs marks. Writing nitrogen dioxide as the direct product of catalytic oxidation. Platinum gives nitric oxide; nitrogen dioxide forms only when that nitric oxide meets more oxygen, and competitive questions on the steps of nitric acid manufacture hinge on it.
A second trap. Treating ammonia only as a base. Its reducing action, shown with copper oxide and chlorine, is tested separately, often through assertion-reason items.
Board versus competitive emphasis. The ICSE paper marks each reaction's condition, equation and observation; a competitive paper marks an oxidation state, a product under stated proportions or a complex's name. The transferable habit is naming ammonia's role in every reaction — fuel, reducing agent or base — because the role predicts the product.
Where ammonia's reactions lead. Class 12 The p-Block Elements covers the properties of ammonia among the compounds of group 15, including its basic nature, its reactions with metal salt solutions and its complex formation with copper ions. The catalytic oxidation to nitric oxide is the first step of the manufacture of nitric acid treated in the same chapter.
Where the reducing reactions lead. Class 11 Redox Reactions assigns oxidation numbers: nitrogen in ammonia is , in nitrogen gas and in nitric oxide . So ammonia is oxidised both when it burns to nitrogen and when it is catalytically oxidised to nitric oxide, and the reduction of copper oxide can be analysed the same way. Balancing such equations by oxidation-number change is a standard skill in both exams.
Where the chlorine reactions lead. The dependence of the product on which reactant is in excess is a limiting-reagent idea from Class 11 Some Basic Concepts, and the volume relationship worked above uses Gay Lussac's law. Questions combining stoichiometry with a named reaction appear as numericals.
Where the complexes lead. Class 12 Coordination Compounds studies and directly — naming, coordination number and the role of ammonia as a ligand.
Where aqueous ammonia leads. Class 11 Equilibrium treats ammonium hydroxide as a weak base and the mixture of ammonia with ammonium chloride as a basic buffer. pH and buffer calculations on this system appear in both exams.
Question types to expect. At this level: equations with conditions and observations, reducing-agent identification, and uses and tests. In competitive papers: oxidation numbers of nitrogen, balancing redox equations, limiting-reagent products, complex naming, and buffer calculations.
The single trap that costs marks. Writing nitrogen dioxide as the direct product of catalytic oxidation. Platinum gives nitric oxide; nitrogen dioxide forms only when that nitric oxide meets more oxygen, and competitive questions on the steps of nitric acid manufacture hinge on it.
A second trap. Treating ammonia only as a base. Its reducing action, shown with copper oxide and chlorine, is tested separately, often through assertion-reason items.
Board versus competitive emphasis. The ICSE paper marks each reaction's condition, equation and observation; a competitive paper marks an oxidation state, a product under stated proportions or a complex's name. The transferable habit is naming ammonia's role in every reaction — fuel, reducing agent or base — because the role predicts the product.
Key takeaways
What must you be able to do from this part?
Two oxidations, three reductions, two chlorine reactions, and a set of base reactions, uses and tests.
- Ammonia does not burn in air and does not support combustion
- Burning in oxygen: , yellowish-green flame
- Catalytic oxidation: , platinum, about ; then , brown fumes; wire keeps glowing
- ** of ammonia** needs of oxygen to burn but for catalytic oxidation
- With HCl: , dense white fumes
- With hot CuO: — black to reddish-brown, water, nitrogen
- With lead oxides: and
- Ammonia is the reducing agent in the metal oxide reactions
- ** of ammonia** with CuO gives of copper and of nitrogen
- Excess ammonia with chlorine: — chlorine colour vanishes, white fumes
- Excess chlorine: — yellow explosive oily liquid
- With acids: ammonium sulphate, ammonium nitrate, ammonium chloride
- With salts: reddish-brown, dirty green, chalky white — insoluble in excess; white and pale blue — soluble in excess
- Uses: fertilisers, nitric acid, refrigeration, cleaning, laboratory reagent
- Tests: pungent smell, red litmus blue, white fumes with HCl, inky blue with copper sulphate
The sharpest self-test is a sorting exercise. List every reaction on this page and sort it under three headings — ammonia oxidised, ammonia reducing something else, ammonia acting as a base — then check that each equation's products match the heading you chose.
- Ammonia does not burn in air and does not support combustion
- Burning in oxygen: , yellowish-green flame
- Catalytic oxidation: , platinum, about ; then , brown fumes; wire keeps glowing
- ** of ammonia** needs of oxygen to burn but for catalytic oxidation
- With HCl: , dense white fumes
- With hot CuO: — black to reddish-brown, water, nitrogen
- With lead oxides: and
- Ammonia is the reducing agent in the metal oxide reactions
- ** of ammonia** with CuO gives of copper and of nitrogen
- Excess ammonia with chlorine: — chlorine colour vanishes, white fumes
- Excess chlorine: — yellow explosive oily liquid
- With acids: ammonium sulphate, ammonium nitrate, ammonium chloride
- With salts: reddish-brown, dirty green, chalky white — insoluble in excess; white and pale blue — soluble in excess
- Uses: fertilisers, nitric acid, refrigeration, cleaning, laboratory reagent
- Tests: pungent smell, red litmus blue, white fumes with HCl, inky blue with copper sulphate
The sharpest self-test is a sorting exercise. List every reaction on this page and sort it under three headings — ammonia oxidised, ammonia reducing something else, ammonia acting as a base — then check that each equation's products match the heading you chose.