Zinc Hydroxide Dissolves in Both Acid and Alkali, and That Gives It Away
See why ammonia turns copper solutions inky blue and why warm alkali drives ammonia out of ammonium salts, understand amphoteric oxides and hydroxides of aluminium, zinc and lead, write their reactions with caustic alkalis, and run a complete scheme that identifies eight common cations.
How can one hydroxide dissolve in an acid and in an alkali as well?
A base dissolves in an acid; that is what a base does. An acid reacts with an alkali; that is what an acid does. So a substance that dissolves in hydrochloric acid and in sodium hydroxide seems to be breaking the rules.
Zinc hydroxide does exactly that. Put the white precipitate in dilute hydrochloric acid and it dissolves as zinc chloride. Put a fresh portion in excess sodium hydroxide and it dissolves again, this time as sodium zincate. Aluminium and lead hydroxides do the same.
Such compounds are called amphoteric — they behave as a base towards acids and as an acid towards alkalis. The elements that form them sit near the dividing line between metals and non-metals in the periodic table, so it is fitting that their oxides and hydroxides show a little of both characters.
This one property is what separated zinc and lead from calcium and magnesium in Part 1. All four give white precipitates with sodium hydroxide, but only the amphoteric two dissolve in excess.
Two special reactions complete the analytical toolkit.
- Ammonium hydroxide with copper salts gives an intense inky blue solution that no other common ion produces
- Sodium hydroxide with ammonium salts gives no precipitate at all, but on warming releases ammonia gas with an unmistakable smell
Where these reactions appear in daily life.
- Aluminium cooking vessels are damaged both by strongly alkaline cleaning soda and by sour foods left standing in them
- Farmers are advised not to mix ammonium fertilisers with lime, because the alkali drives off the nitrogen as ammonia gas
- The purification of bauxite ore before aluminium is extracted relies on aluminium oxide dissolving in sodium hydroxide — a use you will meet in the metallurgy chapter
The final section puts everything together into a single scheme using sodium hydroxide and ammonium hydroxide that identifies eight cations: iron(II), iron(III), copper, zinc, lead, calcium, magnesium and ammonium.
This page covers the second part of the ICSE Class 10 Chemistry chapter on analytical chemistry: the special reactions of copper and ammonium salts, amphoteric oxides and hydroxides, their reactions with caustic alkalis, and a cation identification scheme.
Zinc hydroxide does exactly that. Put the white precipitate in dilute hydrochloric acid and it dissolves as zinc chloride. Put a fresh portion in excess sodium hydroxide and it dissolves again, this time as sodium zincate. Aluminium and lead hydroxides do the same.
Such compounds are called amphoteric — they behave as a base towards acids and as an acid towards alkalis. The elements that form them sit near the dividing line between metals and non-metals in the periodic table, so it is fitting that their oxides and hydroxides show a little of both characters.
This one property is what separated zinc and lead from calcium and magnesium in Part 1. All four give white precipitates with sodium hydroxide, but only the amphoteric two dissolve in excess.
Two special reactions complete the analytical toolkit.
- Ammonium hydroxide with copper salts gives an intense inky blue solution that no other common ion produces
- Sodium hydroxide with ammonium salts gives no precipitate at all, but on warming releases ammonia gas with an unmistakable smell
Where these reactions appear in daily life.
- Aluminium cooking vessels are damaged both by strongly alkaline cleaning soda and by sour foods left standing in them
- Farmers are advised not to mix ammonium fertilisers with lime, because the alkali drives off the nitrogen as ammonia gas
- The purification of bauxite ore before aluminium is extracted relies on aluminium oxide dissolving in sodium hydroxide — a use you will meet in the metallurgy chapter
The final section puts everything together into a single scheme using sodium hydroxide and ammonium hydroxide that identifies eight cations: iron(II), iron(III), copper, zinc, lead, calcium, magnesium and ammonium.
This page covers the second part of the ICSE Class 10 Chemistry chapter on analytical chemistry: the special reactions of copper and ammonium salts, amphoteric oxides and hydroxides, their reactions with caustic alkalis, and a cation identification scheme.
What is special about ammonium hydroxide with copper salts and sodium hydroxide with ammonium salts?
Ammonium hydroxide first precipitates pale blue copper hydroxide and then dissolves it to an inky blue complex, while sodium hydroxide gives no precipitate with ammonium salts but releases ammonia gas on warming.
1. Ammonium hydroxide and copper salts.
Drop by drop — a pale blue precipitate of copper(II) hydroxide:
In excess — the precipitate dissolves to an inky, deep blue solution of tetraamminecopper(II) sulphate:
What actually dissolves the precipitate is ammonia bonding to the copper ion:
Each ammonia molecule donates the lone pair on its nitrogen to the copper ion — four coordinate bonds forming one complex ion.
Why this is special.
- Sodium hydroxide cannot dissolve copper hydroxide, however much is added, because copper hydroxide is not amphoteric
- Only ammonia can, because only ammonia forms the complex
- The inky blue colour is so intense that even a small amount of copper is easy to see, which makes it the confirmatory test for
Worked check — balancing the ionic equation. Copper ; nitrogen ; hydrogen on the left and on the right; oxygen ; charge on the left and on the right. Balanced.
2. Sodium hydroxide and ammonium salts.
No precipitate forms, because ammonium hydroxide is soluble. On warming, ammonia gas is given off:
The ionic essence — the hydroxyl ion takes a proton from the ammonium ion:
Tests that confirm the gas is ammonia:
- A sharp, pungent smell
- Moist red litmus paper turns blue — ammonia is the only common alkaline gas
- Dense white fumes form when a glass rod dipped in concentrated hydrochloric acid is held near it:
Why this is special. It is the only test in the scheme that needs heat and the only one identified by a gas rather than a precipitate. The ammonium ion has no insoluble hydroxide to show you, so it has to be driven out as ammonia.
The everyday consequence. Ammonium sulphate spread on a field that has just been limed loses nitrogen to the air as ammonia — the fertiliser and the lime react exactly as in the second equation above, with calcium hydroxide in place of sodium hydroxide.
One boundary case. A test for ammonium ions done in the cold can give nothing. Always warm the mixture before concluding that no ammonium ion is present.
1. Ammonium hydroxide and copper salts.
Drop by drop — a pale blue precipitate of copper(II) hydroxide:
In excess — the precipitate dissolves to an inky, deep blue solution of tetraamminecopper(II) sulphate:
What actually dissolves the precipitate is ammonia bonding to the copper ion:
Each ammonia molecule donates the lone pair on its nitrogen to the copper ion — four coordinate bonds forming one complex ion.
Why this is special.
- Sodium hydroxide cannot dissolve copper hydroxide, however much is added, because copper hydroxide is not amphoteric
- Only ammonia can, because only ammonia forms the complex
- The inky blue colour is so intense that even a small amount of copper is easy to see, which makes it the confirmatory test for
Worked check — balancing the ionic equation. Copper ; nitrogen ; hydrogen on the left and on the right; oxygen ; charge on the left and on the right. Balanced.
2. Sodium hydroxide and ammonium salts.
No precipitate forms, because ammonium hydroxide is soluble. On warming, ammonia gas is given off:
The ionic essence — the hydroxyl ion takes a proton from the ammonium ion:
Tests that confirm the gas is ammonia:
- A sharp, pungent smell
- Moist red litmus paper turns blue — ammonia is the only common alkaline gas
- Dense white fumes form when a glass rod dipped in concentrated hydrochloric acid is held near it:
Why this is special. It is the only test in the scheme that needs heat and the only one identified by a gas rather than a precipitate. The ammonium ion has no insoluble hydroxide to show you, so it has to be driven out as ammonia.
The everyday consequence. Ammonium sulphate spread on a field that has just been limed loses nitrogen to the air as ammonia — the fertiliser and the lime react exactly as in the second equation above, with calcium hydroxide in place of sodium hydroxide.
One boundary case. A test for ammonium ions done in the cold can give nothing. Always warm the mixture before concluding that no ammonium ion is present.
What are amphoteric oxides and hydroxides, and why are aluminium, zinc and lead amphoteric?
Amphoteric oxides and hydroxides react with both acids and alkalis to form a salt and water, and aluminium, zinc and lead form them because these metals are not strongly metallic enough to give purely basic oxides.
The definition. An amphoteric oxide or hydroxide is one that reacts with both acids and bases to form a salt and water.
The metals themselves — aluminium, zinc and lead — also react with both acids and strong alkalis, releasing hydrogen in each case.
Classifying oxides by how they behave:
- Basic oxides react only with acids: sodium oxide, calcium oxide, magnesium oxide, copper(II) oxide
- Acidic oxides react only with bases: carbon dioxide, sulphur dioxide
- Amphoteric oxides react with both: aluminium oxide, zinc oxide, lead(II) oxide
Their behaviour as bases — reactions with acids.
Their behaviour as acids — reactions with alkalis — is the subject of the next section.
Why these three elements. In the periodic properties chapter, metallic character decreased across a period. Aluminium sits in group 13, next to the metalloid silicon, and zinc and lead are metals whose oxides are much less basic than those of sodium or calcium. Their oxides are basic enough to neutralise acids but acidic enough to neutralise strong alkalis — the in-between character of elements near the metal–non-metal boundary.
Worked example — classify each oxide.
- ** — reacts with acids only: basic
- — reacts with alkalis only: acidic
- — reacts with both: amphoteric
- — reacts with acids only: basic
- — reacts with both: amphoteric
Worked check — balancing the aluminium oxide equation.** Aluminium ; oxygen ; hydrogen ; chlorine . Balanced.
An everyday consequence. Aluminium vessels in an Indian kitchen should not be scrubbed with washing soda, which is alkaline, and should not be used to store tamarind or sour pickles for long, which are acidic. The protective layer of aluminium oxide on the surface is attacked from both sides, precisely because it is amphoteric.
The boundary case in the lead equations. Lead oxide is shown reacting with nitric acid rather than hydrochloric acid. With hydrochloric or sulphuric acid, the lead salt formed is insoluble and coats the solid, slowing the reaction. Nitric acid gives soluble lead nitrate, so the basic behaviour is seen clearly.
The definition. An amphoteric oxide or hydroxide is one that reacts with both acids and bases to form a salt and water.
The metals themselves — aluminium, zinc and lead — also react with both acids and strong alkalis, releasing hydrogen in each case.
Classifying oxides by how they behave:
- Basic oxides react only with acids: sodium oxide, calcium oxide, magnesium oxide, copper(II) oxide
- Acidic oxides react only with bases: carbon dioxide, sulphur dioxide
- Amphoteric oxides react with both: aluminium oxide, zinc oxide, lead(II) oxide
Their behaviour as bases — reactions with acids.
Their behaviour as acids — reactions with alkalis — is the subject of the next section.
Why these three elements. In the periodic properties chapter, metallic character decreased across a period. Aluminium sits in group 13, next to the metalloid silicon, and zinc and lead are metals whose oxides are much less basic than those of sodium or calcium. Their oxides are basic enough to neutralise acids but acidic enough to neutralise strong alkalis — the in-between character of elements near the metal–non-metal boundary.
Worked example — classify each oxide.
- ** — reacts with acids only: basic
- — reacts with alkalis only: acidic
- — reacts with both: amphoteric
- — reacts with acids only: basic
- — reacts with both: amphoteric
Worked check — balancing the aluminium oxide equation.** Aluminium ; oxygen ; hydrogen ; chlorine . Balanced.
An everyday consequence. Aluminium vessels in an Indian kitchen should not be scrubbed with washing soda, which is alkaline, and should not be used to store tamarind or sour pickles for long, which are acidic. The protective layer of aluminium oxide on the surface is attacked from both sides, precisely because it is amphoteric.
The boundary case in the lead equations. Lead oxide is shown reacting with nitric acid rather than hydrochloric acid. With hydrochloric or sulphuric acid, the lead salt formed is insoluble and coats the solid, slowing the reaction. Nitric acid gives soluble lead nitrate, so the basic behaviour is seen clearly.
How do aluminium, zinc and lead and their oxides and hydroxides react with NaOH and KOH?
With caustic alkalis, the metals give a soluble salt and hydrogen, while the oxides and hydroxides give a soluble salt and water — sodium zincate, sodium aluminate and sodium plumbite with sodium hydroxide.
1. The metals — with hot, concentrated alkali, releasing hydrogen.
2. The oxides — giving salt and water.
3. The hydroxides — giving salt and water.
The salts formed:
- ** — sodium zincate
- — sodium aluminate, also called sodium meta-aluminate
- — sodium plumbite
With potassium hydroxide**, the equations are identical with potassium in place of sodium:
giving potassium zincate, potassium aluminate and potassium plumbite.
A pattern that makes the equations easy to remember. Zinc and lead form salts with two alkali-metal atoms and two oxygens, . Aluminium forms a salt with one of each, . Every equation follows from those two formulae plus balancing.
Worked check — balancing the aluminium metal equation.
- Aluminium: ; sodium:
- Oxygen: on the left; on the right
- Hydrogen: on the left; on the right
Balanced — and notice that the water on the left is essential. Without it, the oxygen and hydrogen cannot balance, which is why this equation is the one most often written wrongly.
Worked example — comparing the hydrogen produced. How many molecules of hydrogen come from each atom of metal?
- Zinc: atom gives molecule
- Lead: atom gives molecule
- Aluminium: atoms give molecules, so per atom
Aluminium gives the most hydrogen per atom, matching its valency of three against zinc's and lead's two.
The boundary case — why ammonium hydroxide cannot do this. Ammonium hydroxide is a weak alkali. It does not dissolve aluminium hydroxide or lead hydroxide at all. It does dissolve zinc hydroxide — but, as Part 1 showed, by forming a complex ion with ammonia, not by amphoteric reaction. Only the caustic alkalis, sodium and potassium hydroxide, show amphoterism.
An industrial use of the same reaction. Bauxite, the ore of aluminium, contains iron oxide as an impurity. Heating the powdered ore with sodium hydroxide dissolves the aluminium oxide as sodium aluminate while the iron oxide, which is basic, stays behind and is filtered off — amphoterism used to purify an ore.
1. The metals — with hot, concentrated alkali, releasing hydrogen.
2. The oxides — giving salt and water.
3. The hydroxides — giving salt and water.
The salts formed:
- ** — sodium zincate
- — sodium aluminate, also called sodium meta-aluminate
- — sodium plumbite
With potassium hydroxide**, the equations are identical with potassium in place of sodium:
giving potassium zincate, potassium aluminate and potassium plumbite.
A pattern that makes the equations easy to remember. Zinc and lead form salts with two alkali-metal atoms and two oxygens, . Aluminium forms a salt with one of each, . Every equation follows from those two formulae plus balancing.
Worked check — balancing the aluminium metal equation.
- Aluminium: ; sodium:
- Oxygen: on the left; on the right
- Hydrogen: on the left; on the right
Balanced — and notice that the water on the left is essential. Without it, the oxygen and hydrogen cannot balance, which is why this equation is the one most often written wrongly.
Worked example — comparing the hydrogen produced. How many molecules of hydrogen come from each atom of metal?
- Zinc: atom gives molecule
- Lead: atom gives molecule
- Aluminium: atoms give molecules, so per atom
Aluminium gives the most hydrogen per atom, matching its valency of three against zinc's and lead's two.
The boundary case — why ammonium hydroxide cannot do this. Ammonium hydroxide is a weak alkali. It does not dissolve aluminium hydroxide or lead hydroxide at all. It does dissolve zinc hydroxide — but, as Part 1 showed, by forming a complex ion with ammonia, not by amphoteric reaction. Only the caustic alkalis, sodium and potassium hydroxide, show amphoterism.
An industrial use of the same reaction. Bauxite, the ore of aluminium, contains iron oxide as an impurity. Heating the powdered ore with sodium hydroxide dissolves the aluminium oxide as sodium aluminate while the iron oxide, which is basic, stays behind and is filtered off — amphoterism used to purify an ore.
How does a NaOH and NH4OH scheme identify Fe2+, Fe3+, Cu2+, Zn2+, Pb2+, Ca2+, Mg2+ and NH4+?
Add sodium hydroxide first and read the colour and the effect of excess and of warming; then use ammonium hydroxide on a fresh sample to split the pairs that sodium hydroxide leaves together.
Step 1 — to the solution, add sodium hydroxide drop by drop, then in excess. If there is no precipitate, warm.
- No precipitate; on warming a pungent gas turns moist red litmus blue **
- Dirty green precipitate, insoluble in excess** **
- Reddish-brown precipitate, insoluble in excess** **
- Pale blue precipitate, insoluble in excess** **, confirmed in step 2
- White precipitate, soluble in excess** ** or , go to step 2
- White precipitate, insoluble in excess** ** or , go to step 2
Step 2 — to a fresh sample, add ammonium hydroxide drop by drop, then in excess.
- Pale blue precipitate dissolving to inky blue** confirms **
- White precipitate dissolving in excess** **
- White precipitate insoluble in excess** ** (if the pair was zinc or lead) or (if the pair was calcium or magnesium)
- No precipitate** **
Worked example 1. Solution A gives no precipitate with sodium hydroxide in the cold. On warming, a gas is given off that forms dense white fumes with a rod dipped in concentrated hydrochloric acid.
The gas is ammonia, so A contains .
Worked example 2. Solution B gives a white precipitate with sodium hydroxide that dissolves in excess. With ammonium hydroxide the white precipitate does not dissolve in excess.
Zinc or lead from step 1; not dissolving in ammonium hydroxide rules out zinc. B contains .
Worked example 3. Solution C gives a white precipitate with sodium hydroxide, insoluble in excess, and a white precipitate with ammonium hydroxide.
Calcium or magnesium from step 1; calcium gives no precipitate with ammonium hydroxide. C contains .
Distinguishing pairs — the form most questions take.
- Iron(II) and iron(III) salts: sodium hydroxide gives dirty green with iron(II) and reddish-brown with iron(III)
- Zinc sulphate and lead nitrate: excess ammonium hydroxide dissolves the zinc precipitate but not the lead one
- Calcium nitrate and magnesium nitrate: ammonium hydroxide gives a white precipitate with magnesium and none with calcium
- Zinc nitrate and calcium nitrate: excess sodium hydroxide dissolves the zinc precipitate but not the calcium one
- Copper sulphate and zinc sulphate: either reagent — pale blue precipitate with copper, white with zinc
- Ammonium chloride and sodium chloride: warm with sodium hydroxide — only ammonium chloride gives ammonia
The boundary case worth knowing.** Aluminium ions, , are not in this list, but they behave exactly like lead ions with both reagents: a white precipitate soluble in excess sodium hydroxide and insoluble in excess ammonium hydroxide. If aluminium might be present, add dilute hydrochloric acid to a fresh sample — lead gives a white precipitate of lead chloride and aluminium gives none.
An everyday parallel. A ticket checker on a train first sorts passengers by coach and only then checks seat numbers within a coach. The scheme sorts ions by colour and solubility first, then settles each remaining pair with a second reagent — the most efficient order for the checks.
Step 1 — to the solution, add sodium hydroxide drop by drop, then in excess. If there is no precipitate, warm.
- No precipitate; on warming a pungent gas turns moist red litmus blue **
- Dirty green precipitate, insoluble in excess** **
- Reddish-brown precipitate, insoluble in excess** **
- Pale blue precipitate, insoluble in excess** **, confirmed in step 2
- White precipitate, soluble in excess** ** or , go to step 2
- White precipitate, insoluble in excess** ** or , go to step 2
Step 2 — to a fresh sample, add ammonium hydroxide drop by drop, then in excess.
- Pale blue precipitate dissolving to inky blue** confirms **
- White precipitate dissolving in excess** **
- White precipitate insoluble in excess** ** (if the pair was zinc or lead) or (if the pair was calcium or magnesium)
- No precipitate** **
Worked example 1. Solution A gives no precipitate with sodium hydroxide in the cold. On warming, a gas is given off that forms dense white fumes with a rod dipped in concentrated hydrochloric acid.
The gas is ammonia, so A contains .
Worked example 2. Solution B gives a white precipitate with sodium hydroxide that dissolves in excess. With ammonium hydroxide the white precipitate does not dissolve in excess.
Zinc or lead from step 1; not dissolving in ammonium hydroxide rules out zinc. B contains .
Worked example 3. Solution C gives a white precipitate with sodium hydroxide, insoluble in excess, and a white precipitate with ammonium hydroxide.
Calcium or magnesium from step 1; calcium gives no precipitate with ammonium hydroxide. C contains .
Distinguishing pairs — the form most questions take.
- Iron(II) and iron(III) salts: sodium hydroxide gives dirty green with iron(II) and reddish-brown with iron(III)
- Zinc sulphate and lead nitrate: excess ammonium hydroxide dissolves the zinc precipitate but not the lead one
- Calcium nitrate and magnesium nitrate: ammonium hydroxide gives a white precipitate with magnesium and none with calcium
- Zinc nitrate and calcium nitrate: excess sodium hydroxide dissolves the zinc precipitate but not the calcium one
- Copper sulphate and zinc sulphate: either reagent — pale blue precipitate with copper, white with zinc
- Ammonium chloride and sodium chloride: warm with sodium hydroxide — only ammonium chloride gives ammonia
The boundary case worth knowing.** Aluminium ions, , are not in this list, but they behave exactly like lead ions with both reagents: a white precipitate soluble in excess sodium hydroxide and insoluble in excess ammonium hydroxide. If aluminium might be present, add dilute hydrochloric acid to a fresh sample — lead gives a white precipitate of lead chloride and aluminium gives none.
An everyday parallel. A ticket checker on a train first sorts passengers by coach and only then checks seat numbers within a coach. The scheme sorts ions by colour and solubility first, then settles each remaining pair with a second reagent — the most efficient order for the checks.
Exam tip
What do examiners look for in amphoteric and cation-scheme answers?
Define amphoteric with both halves, write reactions with acid and with alkali, include water where it belongs, and in identifications state each observation before its conclusion.
- Define amphoteric as reacting with both acids and bases to form salt and water
- Give one equation with an acid and one with an alkali whenever you call a substance amphoteric
- Name the products: sodium zincate, sodium aluminate, sodium plumbite
- Include water on the left in the aluminium metal equation:
- Say hot, concentrated alkali for the metals
- Use nitric acid for lead oxide and hydroxide with acids
- Mention warming in the ammonium ion test, and give a confirmatory test for ammonia
- Explain the copper result by complex formation, not amphoterism
- Use a fresh sample for the second reagent in any scheme
- For distinguishing pairs, name the reagent and the different result for each substance
The misconception to name. Copper hydroxide dissolving in excess ammonium hydroxide does not make copper amphoteric. Copper hydroxide does not dissolve in sodium hydroxide at all; it dissolves in ammonia only because ammonia forms a complex ion. Calling it amphoteric confuses two unrelated reasons for a precipitate dissolving.
A second trap. Stating that ammonium hydroxide dissolves aluminium and lead hydroxides. It is too weak an alkali to show amphoterism, and neither metal forms an ammonia complex, so both precipitates remain.
- Define amphoteric as reacting with both acids and bases to form salt and water
- Give one equation with an acid and one with an alkali whenever you call a substance amphoteric
- Name the products: sodium zincate, sodium aluminate, sodium plumbite
- Include water on the left in the aluminium metal equation:
- Say hot, concentrated alkali for the metals
- Use nitric acid for lead oxide and hydroxide with acids
- Mention warming in the ammonium ion test, and give a confirmatory test for ammonia
- Explain the copper result by complex formation, not amphoterism
- Use a fresh sample for the second reagent in any scheme
- For distinguishing pairs, name the reagent and the different result for each substance
The misconception to name. Copper hydroxide dissolving in excess ammonium hydroxide does not make copper amphoteric. Copper hydroxide does not dissolve in sodium hydroxide at all; it dissolves in ammonia only because ammonia forms a complex ion. Calling it amphoteric confuses two unrelated reasons for a precipitate dissolving.
A second trap. Stating that ammonium hydroxide dissolves aluminium and lead hydroxides. It is too weak an alkali to show amphoterism, and neither metal forms an ammonia complex, so both precipitates remain.
Did you know
Why does a drain cleaner with aluminium bits fizz and get hot?
Some drain-cleaning powders contain two ingredients that look an unlikely pair: granules of sodium hydroxide and small pieces of aluminium. Pour the powder into a blocked drain with a little water and the mixture bubbles furiously and becomes very hot. Both effects help clear the blockage, and both come from this lesson.
The bubbles are hydrogen, from aluminium reacting with the alkali:
Aluminium is amphoteric, so it dissolves in a strong alkali just as zinc does. The rising bubbles churn the mixture and help break up the plug of grease and hair.
The heat comes from two sources.
- Dissolving solid sodium hydroxide in water releases a great deal of heat on its own
- The reaction of aluminium with the alkali releases more
The hot, concentrated alkali then attacks the grease itself, turning fats into soap-like substances that wash away — the same kind of reaction used to make soap from oil and caustic soda.
The chemistry also explains the warnings on the packet.
- Hydrogen is flammable, so no flame should be near the drain while the mixture is working
- Hot, concentrated sodium hydroxide burns skin and eyes
- The products must not be used on aluminium pipes or fittings, which would dissolve by exactly the same reaction
And the equation gives a satisfying check on why aluminium is used rather than zinc. Two atoms of aluminium release three molecules of hydrogen, while one atom of zinc releases one. Aluminium is also a much lighter metal, so a small mass of it produces a large volume of bubbling gas — a practical advantage worked out from nothing more than an amphoteric reaction and its balanced equation.
The bubbles are hydrogen, from aluminium reacting with the alkali:
Aluminium is amphoteric, so it dissolves in a strong alkali just as zinc does. The rising bubbles churn the mixture and help break up the plug of grease and hair.
The heat comes from two sources.
- Dissolving solid sodium hydroxide in water releases a great deal of heat on its own
- The reaction of aluminium with the alkali releases more
The hot, concentrated alkali then attacks the grease itself, turning fats into soap-like substances that wash away — the same kind of reaction used to make soap from oil and caustic soda.
The chemistry also explains the warnings on the packet.
- Hydrogen is flammable, so no flame should be near the drain while the mixture is working
- Hot, concentrated sodium hydroxide burns skin and eyes
- The products must not be used on aluminium pipes or fittings, which would dissolve by exactly the same reaction
And the equation gives a satisfying check on why aluminium is used rather than zinc. Two atoms of aluminium release three molecules of hydrogen, while one atom of zinc releases one. Aluminium is also a much lighter metal, so a small mass of it produces a large volume of bubbling gas — a practical advantage worked out from nothing more than an amphoteric reaction and its balanced equation.
Exam relevance
How are amphoteric compounds and cation tests examined in JEE and NEET?
This is foundation work for Class 11 Classification of Elements and Periodicity in Properties, Class 12 Coordination Compounds and the qualitative analysis section of JEE Main practical chemistry.
Where amphoteric oxides lead. The periodicity chapter examined in both JEE Main and NEET describes how oxides change from basic to amphoteric to acidic across a period. Identifying which oxide in a list is amphoteric — aluminium oxide, zinc oxide, beryllium oxide or tin oxide — is a recurring objective question, and it is answered by exactly the classification on this page.
Where the aluminium reactions lead. The reactions of aluminium and its oxide with both acids and alkalis are asked in both exams as examples of amphoteric behaviour, often with the formula of the aluminate product. The water on the left of the aluminium equation is a detail that separates correct and incorrect options.
Where the copper complex leads. Class 12 Coordination Compounds treats tetraamminecopper(II) as a model complex: ammonia as a ligand donating a lone pair, a coordination number of four, and the naming of the complex. Questions asking why copper hydroxide dissolves in ammonia but not in sodium hydroxide are answered by complex formation — the distinction drawn in this lesson.
Where the scheme leads. JEE Main practical chemistry includes the chemical principles of qualitative salt analysis, with cations including lead, copper, aluminium, iron(III), zinc, calcium, magnesium and ammonium. The ammonium test with warm alkali is used unchanged, and the idea of choosing a second reagent to split a pair is the logic of group analysis.
Where the ammonium test leads. In Class 11 Equilibrium, the reaction of an ammonium ion with a hydroxyl ion is a proton transfer between a conjugate acid and a base. **Recognising as an acid and as its conjugate base is a standard question in both exams.
Question types to expect. At this level: amphoteric definitions with equations, reactions with caustic alkalis, and cation identification or distinguishing tests. In competitive papers: nature of oxides across a period, aluminate formation, complex formation and coordination number, salt-analysis reasoning, and conjugate acid-base pairs, often as match-the-column or assertion-reason items.
The single trap that costs marks. Treating every precipitate that dissolves in excess reagent as amphoteric. Zinc hydroxide dissolves in sodium hydroxide by amphoterism but in ammonia by complex formation; copper hydroxide dissolves only in ammonia. Assertion-reason questions are built precisely on this.
A second trap. Missing that aluminium and lead ions give identical results with both reagents. A salt-analysis question including aluminium needs a further test, and a candidate who stops at the two-reagent scheme gives an ambiguous answer.
Board versus competitive emphasis. The ICSE paper marks balanced equations with named products and a clear observation-conclusion sequence; a competitive paper marks a classification, a complex's coordination number or a group assignment. The transferable habit is naming the mechanism behind each observation** — amphoterism, complex formation or proton transfer — because the later chapters are organised by mechanism, not by colour.
Where amphoteric oxides lead. The periodicity chapter examined in both JEE Main and NEET describes how oxides change from basic to amphoteric to acidic across a period. Identifying which oxide in a list is amphoteric — aluminium oxide, zinc oxide, beryllium oxide or tin oxide — is a recurring objective question, and it is answered by exactly the classification on this page.
Where the aluminium reactions lead. The reactions of aluminium and its oxide with both acids and alkalis are asked in both exams as examples of amphoteric behaviour, often with the formula of the aluminate product. The water on the left of the aluminium equation is a detail that separates correct and incorrect options.
Where the copper complex leads. Class 12 Coordination Compounds treats tetraamminecopper(II) as a model complex: ammonia as a ligand donating a lone pair, a coordination number of four, and the naming of the complex. Questions asking why copper hydroxide dissolves in ammonia but not in sodium hydroxide are answered by complex formation — the distinction drawn in this lesson.
Where the scheme leads. JEE Main practical chemistry includes the chemical principles of qualitative salt analysis, with cations including lead, copper, aluminium, iron(III), zinc, calcium, magnesium and ammonium. The ammonium test with warm alkali is used unchanged, and the idea of choosing a second reagent to split a pair is the logic of group analysis.
Where the ammonium test leads. In Class 11 Equilibrium, the reaction of an ammonium ion with a hydroxyl ion is a proton transfer between a conjugate acid and a base. **Recognising as an acid and as its conjugate base is a standard question in both exams.
Question types to expect. At this level: amphoteric definitions with equations, reactions with caustic alkalis, and cation identification or distinguishing tests. In competitive papers: nature of oxides across a period, aluminate formation, complex formation and coordination number, salt-analysis reasoning, and conjugate acid-base pairs, often as match-the-column or assertion-reason items.
The single trap that costs marks. Treating every precipitate that dissolves in excess reagent as amphoteric. Zinc hydroxide dissolves in sodium hydroxide by amphoterism but in ammonia by complex formation; copper hydroxide dissolves only in ammonia. Assertion-reason questions are built precisely on this.
A second trap. Missing that aluminium and lead ions give identical results with both reagents. A salt-analysis question including aluminium needs a further test, and a candidate who stops at the two-reagent scheme gives an ambiguous answer.
Board versus competitive emphasis. The ICSE paper marks balanced equations with named products and a clear observation-conclusion sequence; a competitive paper marks a classification, a complex's coordination number or a group assignment. The transferable habit is naming the mechanism behind each observation** — amphoterism, complex formation or proton transfer — because the later chapters are organised by mechanism, not by colour.
Key takeaways
What must you be able to do from this part?
Two special reactions, one property with three metals, nine alkali equations and an eight-ion scheme.
- Copper salts with ammonium hydroxide: pale blue that dissolves in excess to inky blue , by complex formation
- Copper hydroxide does not dissolve in sodium hydroxide — copper is not amphoteric
- Ammonium salts with sodium hydroxide: no precipitate; on warming, ammonia gas —
- Ammonia tests: pungent smell, moist red litmus blue, dense white fumes with concentrated HCl
- Amphoteric oxides and hydroxides react with both acids and bases to give salt and water
- Basic oxides: , , . Acidic oxides: , . Amphoteric: , ,
- With acids: ; use nitric acid for lead compounds
- Metals with hot concentrated alkali: ; ;
- Oxides with alkali: , and give the same salts plus water
- Hydroxides with alkali: ; ;
- Products: sodium zincate, sodium aluminate, sodium plumbite; potassium versions with KOH
- Ammonium hydroxide is too weak to show amphoterism
- Scheme step 1, sodium hydroxide: no precipitate but ammonia on warming ; dirty green ; reddish-brown ; pale blue ; white soluble in excess or ; white insoluble or
- Scheme step 2, ammonium hydroxide: dissolves in excess ; stays ; white precipitate ; none
- ** mimics ** with both reagents; dilute HCl separates them
The sharpest self-test is six pairs of bottles. Take the distinguishing pairs from this lesson, cover the answers, and for each pair name one reagent and the two different results — then write the balanced equation for whichever result involves a precipitate dissolving.
- Copper salts with ammonium hydroxide: pale blue that dissolves in excess to inky blue , by complex formation
- Copper hydroxide does not dissolve in sodium hydroxide — copper is not amphoteric
- Ammonium salts with sodium hydroxide: no precipitate; on warming, ammonia gas —
- Ammonia tests: pungent smell, moist red litmus blue, dense white fumes with concentrated HCl
- Amphoteric oxides and hydroxides react with both acids and bases to give salt and water
- Basic oxides: , , . Acidic oxides: , . Amphoteric: , ,
- With acids: ; use nitric acid for lead compounds
- Metals with hot concentrated alkali: ; ;
- Oxides with alkali: , and give the same salts plus water
- Hydroxides with alkali: ; ;
- Products: sodium zincate, sodium aluminate, sodium plumbite; potassium versions with KOH
- Ammonium hydroxide is too weak to show amphoterism
- Scheme step 1, sodium hydroxide: no precipitate but ammonia on warming ; dirty green ; reddish-brown ; pale blue ; white soluble in excess or ; white insoluble or
- Scheme step 2, ammonium hydroxide: dissolves in excess ; stays ; white precipitate ; none
- ** mimics ** with both reagents; dilute HCl separates them
The sharpest self-test is six pairs of bottles. Take the distinguishing pairs from this lesson, cover the answers, and for each pair name one reagent and the two different results — then write the balanced equation for whichever result involves a precipitate dissolving.