The Colour of a Precipitate Can Name the Metal Hidden in a Solution
Learn the colours of iron, copper, zinc, calcium, magnesium and lead salts, follow what sodium hydroxide and ammonium hydroxide do when added drop by drop and then in excess, and use the colour and behaviour of each hydroxide precipitate to identify the metal ion in an unknown solution.
How can a few drops of alkali reveal which metal is in a solution?
Line up seven test tubes of clear or lightly coloured solutions, each containing a different metal salt. Most look almost the same — several are completely colourless. Now add a few drops of sodium hydroxide solution to each.
One turns a dirty green, one a rusty reddish-brown, one a pale blue, and the rest turn cloudy white. Keep adding the alkali, and two of the white ones clear up again while the others stay cloudy. Switch to ammonium hydroxide and the pattern shifts once more.
Every one of those changes is a hydroxide precipitating out of solution. When an alkali meets a salt of a metal whose hydroxide does not dissolve, the metal ions and hydroxyl ions combine into a solid:
The colour of that solid, and whether it dissolves in more alkali, is a fingerprint of the metal. This is the core of qualitative analysis — finding out what is present, rather than how much.
Analysis of this kind is everyday work.
- Water-testing laboratories check drinking water for dissolved metals such as iron and lead
- The brown stains in a bucket left under an iron-rich tap are iron hydroxide formed the same way
- Farmers' soil-testing centres identify the metal ions present in a soil sample before recommending fertilisers
So this part builds the fingerprint table.
- The colours of the salts and their solutions, which give the first clue
- The action of sodium hydroxide, drop by drop and then in excess
- The action of ammonium hydroxide, drop by drop and then in excess
- Using the observations together to identify the metal in an unknown solution
The two reagents are chosen because they behave differently. Sodium hydroxide is a strong alkali and ammonium hydroxide a weak one, and ammonia can also bond to some metal ions. Where one reagent leaves two metals looking identical, the other usually separates them — which is why both are always used.
This page covers the first part of the ICSE Class 10 Chemistry chapter on analytical chemistry: colours of salts, the action of sodium hydroxide and ammonium hydroxide on salt solutions, and identifying cations from hydroxide precipitates.
One turns a dirty green, one a rusty reddish-brown, one a pale blue, and the rest turn cloudy white. Keep adding the alkali, and two of the white ones clear up again while the others stay cloudy. Switch to ammonium hydroxide and the pattern shifts once more.
Every one of those changes is a hydroxide precipitating out of solution. When an alkali meets a salt of a metal whose hydroxide does not dissolve, the metal ions and hydroxyl ions combine into a solid:
The colour of that solid, and whether it dissolves in more alkali, is a fingerprint of the metal. This is the core of qualitative analysis — finding out what is present, rather than how much.
Analysis of this kind is everyday work.
- Water-testing laboratories check drinking water for dissolved metals such as iron and lead
- The brown stains in a bucket left under an iron-rich tap are iron hydroxide formed the same way
- Farmers' soil-testing centres identify the metal ions present in a soil sample before recommending fertilisers
So this part builds the fingerprint table.
- The colours of the salts and their solutions, which give the first clue
- The action of sodium hydroxide, drop by drop and then in excess
- The action of ammonium hydroxide, drop by drop and then in excess
- Using the observations together to identify the metal in an unknown solution
The two reagents are chosen because they behave differently. Sodium hydroxide is a strong alkali and ammonium hydroxide a weak one, and ammonia can also bond to some metal ions. Where one reagent leaves two metals looking identical, the other usually separates them — which is why both are always used.
This page covers the first part of the ICSE Class 10 Chemistry chapter on analytical chemistry: colours of salts, the action of sodium hydroxide and ammonium hydroxide on salt solutions, and identifying cations from hydroxide precipitates.
What colours are the salts of iron, copper, zinc, calcium, magnesium and lead?
Salts of zinc, calcium, magnesium and lead are white and give colourless solutions, while iron(II) salts are pale green, iron(III) salts yellowish-brown and copper(II) salts blue.
1. Iron(II) or ferrous salts — such as iron(II) sulphate
- Salt: light green crystals,
- Solution: pale green
2. Iron(III) or ferric salts — such as iron(III) chloride
- Salt: yellowish-brown
- Solution: yellowish-brown
3. Copper(II) salts — such as copper(II) sulphate and nitrate
- Salt: blue crystals,
- Solution: blue; copper(II) chloride solution is bluish-green
- Related compounds: anhydrous copper sulphate white, copper(II) oxide black, copper carbonate green
4. Zinc salts — such as zinc sulphate
- Salt: white
- Solution: colourless
- Zinc oxide is white when cold and yellow when hot, turning white again on cooling
5. Calcium salts — such as calcium nitrate and calcium chloride
- Salt: white
- Solution: colourless
6. Magnesium salts — such as magnesium sulphate
- Salt: white
- Solution: colourless
7. Lead salts — such as lead nitrate
- Salt: white
- Solution: colourless
- Lead(II) oxide is yellow
Worked example 1. A solution is pale green. Which metal ion is most likely present? Iron(II), since it is the only ion in this list with a pale green solution.
Worked example 2. A solution is colourless. Which metal ions from this list could it contain? Zinc, calcium, magnesium or lead — four possibilities, and colour alone cannot choose between them.
That second result is the whole reason the reagent tests exist. Colour narrows the field; it cannot finish the job for the colourless ions.
Everyday examples from Indian shops. Blue copper sulphate crystals, neela thotha, are sold as a fungicide for crops, and green iron(II) sulphate crystals, hara kasis, as a source of iron for plants. Their colours are exactly those in the list above.
The link to what comes next. Iron and copper give coloured salts and, as the next sections show, coloured hydroxides as well. The four colourless ions all give white hydroxides, so they must be told apart by how those white solids behave with more reagent.
One boundary case about iron. A pale green iron(II) sulphate solution left standing in air slowly turns yellowish-brown at the surface, as iron(II) ions are converted to iron(III). An old iron(II) solution can therefore give misleading colours, and freshly prepared solutions are used in analysis for that reason.
1. Iron(II) or ferrous salts — such as iron(II) sulphate
- Salt: light green crystals,
- Solution: pale green
2. Iron(III) or ferric salts — such as iron(III) chloride
- Salt: yellowish-brown
- Solution: yellowish-brown
3. Copper(II) salts — such as copper(II) sulphate and nitrate
- Salt: blue crystals,
- Solution: blue; copper(II) chloride solution is bluish-green
- Related compounds: anhydrous copper sulphate white, copper(II) oxide black, copper carbonate green
4. Zinc salts — such as zinc sulphate
- Salt: white
- Solution: colourless
- Zinc oxide is white when cold and yellow when hot, turning white again on cooling
5. Calcium salts — such as calcium nitrate and calcium chloride
- Salt: white
- Solution: colourless
6. Magnesium salts — such as magnesium sulphate
- Salt: white
- Solution: colourless
7. Lead salts — such as lead nitrate
- Salt: white
- Solution: colourless
- Lead(II) oxide is yellow
Worked example 1. A solution is pale green. Which metal ion is most likely present? Iron(II), since it is the only ion in this list with a pale green solution.
Worked example 2. A solution is colourless. Which metal ions from this list could it contain? Zinc, calcium, magnesium or lead — four possibilities, and colour alone cannot choose between them.
That second result is the whole reason the reagent tests exist. Colour narrows the field; it cannot finish the job for the colourless ions.
Everyday examples from Indian shops. Blue copper sulphate crystals, neela thotha, are sold as a fungicide for crops, and green iron(II) sulphate crystals, hara kasis, as a source of iron for plants. Their colours are exactly those in the list above.
The link to what comes next. Iron and copper give coloured salts and, as the next sections show, coloured hydroxides as well. The four colourless ions all give white hydroxides, so they must be told apart by how those white solids behave with more reagent.
One boundary case about iron. A pale green iron(II) sulphate solution left standing in air slowly turns yellowish-brown at the surface, as iron(II) ions are converted to iron(III). An old iron(II) solution can therefore give misleading colours, and freshly prepared solutions are used in analysis for that reason.
What happens when sodium hydroxide is added drop by drop and then in excess?
Iron(II), iron(III), copper, calcium and magnesium give precipitates that do not dissolve in excess sodium hydroxide, while zinc and lead give white precipitates that do dissolve in excess.
1. Iron(II) salt.
- Drop by drop: dirty green precipitate of iron(II) hydroxide
- In excess: insoluble
2. Iron(III) salt.
- Drop by drop: reddish-brown precipitate of iron(III) hydroxide
- In excess: insoluble
3. Copper(II) salt.
- Drop by drop: pale blue precipitate of copper(II) hydroxide
- In excess: insoluble
4. Zinc salt.
- Drop by drop: white gelatinous precipitate of zinc hydroxide
- In excess: dissolves, giving a colourless solution of sodium zincate
5. Calcium salt.
- Drop by drop: white precipitate of calcium hydroxide, seen clearly only in a fairly concentrated solution because calcium hydroxide is slightly soluble
- In excess: insoluble
6. Magnesium salt.
- Drop by drop: white precipitate of magnesium hydroxide
- In excess: insoluble
7. Lead salt.
- Drop by drop: chalky white precipitate of lead hydroxide
- In excess: dissolves, giving a colourless solution of sodium plumbite
Why zinc and lead hydroxides dissolve. Their hydroxides are amphoteric — they react with alkalis as well as with acids. Excess sodium hydroxide reacts with the precipitate itself, forming a soluble salt. This is developed fully in Part 2.
The ionic view of every precipitation. Only the metal ion and the hydroxyl ions take part; the other ions stay dissolved:
Charge check: , matching the neutral solid.
Worked check — balancing the sodium zincate equation. Zinc ; sodium ; oxygen on the left, on the right; hydrogen on the left, on the right. Balanced.
Why drop by drop first. Adding the reagent a little at a time lets you see the precipitate form before judging whether it dissolves. Pouring in a large excess at once can dissolve a zinc or lead precipitate the instant it forms, so you would see a clear solution and wrongly conclude that no precipitate was produced.
An everyday parallel. Adding sugar to tea a spoonful at a time shows the moment it stops dissolving; tipping in the whole jar hides that point entirely. The drop-by-drop rule in analysis exists for the same reason — the observation happens at the boundary, and you have to approach it slowly to see it.
1. Iron(II) salt.
- Drop by drop: dirty green precipitate of iron(II) hydroxide
- In excess: insoluble
2. Iron(III) salt.
- Drop by drop: reddish-brown precipitate of iron(III) hydroxide
- In excess: insoluble
3. Copper(II) salt.
- Drop by drop: pale blue precipitate of copper(II) hydroxide
- In excess: insoluble
4. Zinc salt.
- Drop by drop: white gelatinous precipitate of zinc hydroxide
- In excess: dissolves, giving a colourless solution of sodium zincate
5. Calcium salt.
- Drop by drop: white precipitate of calcium hydroxide, seen clearly only in a fairly concentrated solution because calcium hydroxide is slightly soluble
- In excess: insoluble
6. Magnesium salt.
- Drop by drop: white precipitate of magnesium hydroxide
- In excess: insoluble
7. Lead salt.
- Drop by drop: chalky white precipitate of lead hydroxide
- In excess: dissolves, giving a colourless solution of sodium plumbite
Why zinc and lead hydroxides dissolve. Their hydroxides are amphoteric — they react with alkalis as well as with acids. Excess sodium hydroxide reacts with the precipitate itself, forming a soluble salt. This is developed fully in Part 2.
The ionic view of every precipitation. Only the metal ion and the hydroxyl ions take part; the other ions stay dissolved:
Charge check: , matching the neutral solid.
Worked check — balancing the sodium zincate equation. Zinc ; sodium ; oxygen on the left, on the right; hydrogen on the left, on the right. Balanced.
Why drop by drop first. Adding the reagent a little at a time lets you see the precipitate form before judging whether it dissolves. Pouring in a large excess at once can dissolve a zinc or lead precipitate the instant it forms, so you would see a clear solution and wrongly conclude that no precipitate was produced.
An everyday parallel. Adding sugar to tea a spoonful at a time shows the moment it stops dissolving; tipping in the whole jar hides that point entirely. The drop-by-drop rule in analysis exists for the same reason — the observation happens at the boundary, and you have to approach it slowly to see it.
What happens when ammonium hydroxide is added drop by drop and then in excess?
Ammonium hydroxide precipitates the same hydroxides as sodium hydroxide except calcium, dissolves the zinc and copper precipitates in excess, and leaves the lead precipitate undissolved.
1. Iron(II) salt. Dirty green precipitate, insoluble in excess.
2. Iron(III) salt. Reddish-brown precipitate, insoluble in excess.
3. Copper(II) salt. Pale blue precipitate, which dissolves in excess to give an inky, deep blue solution of tetraamminecopper(II) sulphate.
4. Zinc salt. White gelatinous precipitate, which dissolves in excess to give a colourless solution of tetraamminezinc(II) sulphate.
5. Calcium salt. No precipitate, even with excess.
6. Magnesium salt. White precipitate, insoluble in excess.
7. Lead salt. Chalky white precipitate, insoluble in excess.
Why calcium gives no precipitate here. Ammonium hydroxide is a weak alkali, so it supplies far fewer hydroxyl ions than sodium hydroxide. Calcium hydroxide is slightly soluble, and the small concentration of hydroxyl ions is not enough to push it out of solution. Magnesium hydroxide is much less soluble, so even ammonium hydroxide precipitates it.
Why zinc and copper precipitates dissolve. This is not the amphoteric behaviour that dissolves them in sodium hydroxide. Ammonia molecules donate their lone pairs to the metal ions, forming complex ions that stay in solution — the coordinate bond from the chemical bonding chapter, four times over. Lead does not form such a complex, so its precipitate remains.
Worked check — balancing the copper complex equation.
- Copper: ; sulphur:
- Nitrogen: on the left; on the right
- Hydrogen: on the left; on the right
- Oxygen: on the left; on the right
Balanced.
Comparing the two reagents side by side — the differences are what make them useful together:
- Calcium: white with sodium hydroxide; none with ammonium hydroxide
- Zinc: dissolves in excess of both
- Lead: dissolves in excess sodium hydroxide; stays in excess ammonium hydroxide
- Copper: stays in excess sodium hydroxide; dissolves deep blue in excess ammonium hydroxide
An everyday note. The ammonium hydroxide used in the laboratory is a solution of ammonia gas in water — the same sharp-smelling liquid as household ammonia for cleaning glass, only more carefully measured. The smell is a reminder of the free ammonia molecules that do the complex-forming.
1. Iron(II) salt. Dirty green precipitate, insoluble in excess.
2. Iron(III) salt. Reddish-brown precipitate, insoluble in excess.
3. Copper(II) salt. Pale blue precipitate, which dissolves in excess to give an inky, deep blue solution of tetraamminecopper(II) sulphate.
4. Zinc salt. White gelatinous precipitate, which dissolves in excess to give a colourless solution of tetraamminezinc(II) sulphate.
5. Calcium salt. No precipitate, even with excess.
6. Magnesium salt. White precipitate, insoluble in excess.
7. Lead salt. Chalky white precipitate, insoluble in excess.
Why calcium gives no precipitate here. Ammonium hydroxide is a weak alkali, so it supplies far fewer hydroxyl ions than sodium hydroxide. Calcium hydroxide is slightly soluble, and the small concentration of hydroxyl ions is not enough to push it out of solution. Magnesium hydroxide is much less soluble, so even ammonium hydroxide precipitates it.
Why zinc and copper precipitates dissolve. This is not the amphoteric behaviour that dissolves them in sodium hydroxide. Ammonia molecules donate their lone pairs to the metal ions, forming complex ions that stay in solution — the coordinate bond from the chemical bonding chapter, four times over. Lead does not form such a complex, so its precipitate remains.
Worked check — balancing the copper complex equation.
- Copper: ; sulphur:
- Nitrogen: on the left; on the right
- Hydrogen: on the left; on the right
- Oxygen: on the left; on the right
Balanced.
Comparing the two reagents side by side — the differences are what make them useful together:
- Calcium: white with sodium hydroxide; none with ammonium hydroxide
- Zinc: dissolves in excess of both
- Lead: dissolves in excess sodium hydroxide; stays in excess ammonium hydroxide
- Copper: stays in excess sodium hydroxide; dissolves deep blue in excess ammonium hydroxide
An everyday note. The ammonium hydroxide used in the laboratory is a solution of ammonia gas in water — the same sharp-smelling liquid as household ammonia for cleaning glass, only more carefully measured. The smell is a reminder of the free ammonia molecules that do the complex-forming.
How do the hydroxide colours identify the cation in an unknown solution?
A coloured precipitate names iron(II), iron(III) or copper at once; a white precipitate is sorted by whether it dissolves in excess sodium hydroxide, and then by what ammonium hydroxide does.
The colour table of hydroxide precipitates:
- Iron(II) hydroxide, — dirty green
- Iron(III) hydroxide, — reddish-brown
- Copper(II) hydroxide, — pale blue
- Zinc hydroxide, — white, gelatinous
- Lead hydroxide, — chalky white
- Magnesium hydroxide, — white
- Calcium hydroxide, — white, with sodium hydroxide only
The reasoning, step by step.
Step 1 — add sodium hydroxide drop by drop and note the colour.
- Dirty green **
- Reddish-brown** **
- Pale blue** **
- White** go to step 2
Step 2 — add sodium hydroxide in excess to the white precipitate.
- Dissolves ** or ** go to step 3a
- Does not dissolve ** or ** go to step 3b
Step 3a — to a fresh sample, add ammonium hydroxide drop by drop, then in excess.
- White precipitate that dissolves in excess **
- White precipitate that stays** **
Step 3b — to a fresh sample, add ammonium hydroxide.
- No precipitate** **
- White precipitate** **
Worked example 1. Solution P gives a white precipitate with sodium hydroxide that dissolves in excess. With ammonium hydroxide it gives a white precipitate that does not dissolve in excess.
White and soluble in excess sodium hydroxide means zinc or lead; insoluble in excess ammonium hydroxide rules out zinc. P contains .
Worked example 2. Solution Q gives a white precipitate with sodium hydroxide, insoluble in excess, and no precipitate with ammonium hydroxide.
Calcium or magnesium; no precipitate with ammonium hydroxide means calcium. Q contains .
Worked example 3. Solution R is blue. With ammonium hydroxide it gives a pale blue precipitate that dissolves in excess to an inky blue solution.
The blue solution, the pale blue precipitate and the deep blue complex all agree. R contains .
Worked example 4. Solution S gives a white precipitate with both reagents, and the precipitate dissolves in excess of both.
Only zinc dissolves in excess of both reagents. S contains .
Worked example 5 — a reasoning trap. Solution T gives a green precipitate with sodium hydroxide that, after standing for a while, turns brown at the top. Is it iron(II) or iron(III)?
Iron(II). The precipitate was green when it formed; the browning is iron(II) hydroxide being slowly converted to iron(III) hydroxide by oxygen from the air. Always record the colour as the precipitate first appears.
An everyday parallel to the scheme. A doctor narrowing down an illness does not run every test at once; each result decides which test comes next. The cation scheme works the same way** — the answer to one question chooses the next reagent, so that three steps at most identify any of seven metal ions.
The colour table of hydroxide precipitates:
- Iron(II) hydroxide, — dirty green
- Iron(III) hydroxide, — reddish-brown
- Copper(II) hydroxide, — pale blue
- Zinc hydroxide, — white, gelatinous
- Lead hydroxide, — chalky white
- Magnesium hydroxide, — white
- Calcium hydroxide, — white, with sodium hydroxide only
The reasoning, step by step.
Step 1 — add sodium hydroxide drop by drop and note the colour.
- Dirty green **
- Reddish-brown** **
- Pale blue** **
- White** go to step 2
Step 2 — add sodium hydroxide in excess to the white precipitate.
- Dissolves ** or ** go to step 3a
- Does not dissolve ** or ** go to step 3b
Step 3a — to a fresh sample, add ammonium hydroxide drop by drop, then in excess.
- White precipitate that dissolves in excess **
- White precipitate that stays** **
Step 3b — to a fresh sample, add ammonium hydroxide.
- No precipitate** **
- White precipitate** **
Worked example 1. Solution P gives a white precipitate with sodium hydroxide that dissolves in excess. With ammonium hydroxide it gives a white precipitate that does not dissolve in excess.
White and soluble in excess sodium hydroxide means zinc or lead; insoluble in excess ammonium hydroxide rules out zinc. P contains .
Worked example 2. Solution Q gives a white precipitate with sodium hydroxide, insoluble in excess, and no precipitate with ammonium hydroxide.
Calcium or magnesium; no precipitate with ammonium hydroxide means calcium. Q contains .
Worked example 3. Solution R is blue. With ammonium hydroxide it gives a pale blue precipitate that dissolves in excess to an inky blue solution.
The blue solution, the pale blue precipitate and the deep blue complex all agree. R contains .
Worked example 4. Solution S gives a white precipitate with both reagents, and the precipitate dissolves in excess of both.
Only zinc dissolves in excess of both reagents. S contains .
Worked example 5 — a reasoning trap. Solution T gives a green precipitate with sodium hydroxide that, after standing for a while, turns brown at the top. Is it iron(II) or iron(III)?
Iron(II). The precipitate was green when it formed; the browning is iron(II) hydroxide being slowly converted to iron(III) hydroxide by oxygen from the air. Always record the colour as the precipitate first appears.
An everyday parallel to the scheme. A doctor narrowing down an illness does not run every test at once; each result decides which test comes next. The cation scheme works the same way** — the answer to one question chooses the next reagent, so that three steps at most identify any of seven metal ions.
Exam tip
What makes an analytical chemistry answer complete?
Give the reagent, how it was added, the colour and nature of the precipitate, its behaviour in excess, the balanced equation, and the ion identified.
- **Say drop by drop and in excess separately, with an observation for each stage
- Use the exact colour words: dirty green, reddish-brown, pale blue, white gelatinous, chalky white
- Describe the excess result fully**: dissolves to a colourless solution or dissolves to an inky blue solution, not just dissolves
- Write both equations where a precipitate dissolves — its formation and its dissolving
- Name the soluble products: sodium zincate, sodium plumbite, tetraamminecopper(II) sulphate
- State that calcium gives no precipitate with ammonium hydroxide, and give the reason if asked
- Give the salt and solution colours of iron(II), iron(III) and copper salts
- Record the first colour of an iron(II) precipitate, not the brown it becomes later
- Use a fresh sample for each reagent in an identification
- End with the ion identified, written with its charge
The misconception to name. The zinc precipitate dissolves in excess of both reagents, but for two different reasons. In sodium hydroxide it dissolves because zinc hydroxide is amphoteric; in ammonium hydroxide it dissolves because ammonia forms a complex ion. Giving the amphoteric reason for both loses the mark.
A second trap. Confusing the pale blue copper hydroxide precipitate with the inky blue complex solution. One is a solid that settles; the other is a clear, intensely coloured liquid — the wording must make clear which you mean.
- **Say drop by drop and in excess separately, with an observation for each stage
- Use the exact colour words: dirty green, reddish-brown, pale blue, white gelatinous, chalky white
- Describe the excess result fully**: dissolves to a colourless solution or dissolves to an inky blue solution, not just dissolves
- Write both equations where a precipitate dissolves — its formation and its dissolving
- Name the soluble products: sodium zincate, sodium plumbite, tetraamminecopper(II) sulphate
- State that calcium gives no precipitate with ammonium hydroxide, and give the reason if asked
- Give the salt and solution colours of iron(II), iron(III) and copper salts
- Record the first colour of an iron(II) precipitate, not the brown it becomes later
- Use a fresh sample for each reagent in an identification
- End with the ion identified, written with its charge
The misconception to name. The zinc precipitate dissolves in excess of both reagents, but for two different reasons. In sodium hydroxide it dissolves because zinc hydroxide is amphoteric; in ammonium hydroxide it dissolves because ammonia forms a complex ion. Giving the amphoteric reason for both loses the mark.
A second trap. Confusing the pale blue copper hydroxide precipitate with the inky blue complex solution. One is a solid that settles; the other is a clear, intensely coloured liquid — the wording must make clear which you mean.
Did you know
Why does a green iron(II) hydroxide precipitate turn brown if you leave it?
Make a dirty green precipitate of iron(II) hydroxide in a test tube and leave it on the rack. Within a short time the top layer, where it touches the air, turns brown, while the precipitate lower down stays green. Shake the tube and more of it browns.
The precipitate is reacting with oxygen dissolved from the air. Iron(II) is easily converted to iron(III), and the green hydroxide becomes the reddish-brown one:
Checking the balance: iron ; oxygen on the left and on the right; hydrogen on the left and on the right. Balanced.
The brown creeping down from the top is the oxygen working its way in. Shaking brings more of the precipitate into contact with air, which is why it speeds up the colour change.
The same chemistry is visible all around you.
- Green iron(II) sulphate crystals, left in an open jar, develop yellowish-brown patches as the surface changes to iron(III)
- Water from an iron-rich borewell often comes out clear, then turns cloudy and orange-brown in a bucket as dissolved iron(II) meets air and precipitates as iron(III) compounds
- The brown stains on a washbasin under such a tap are the same iron(III) hydroxide, left behind as the water dries
Rust itself is closely related. When iron objects corrode in moist air, the iron is converted through iron(II) to hydrated iron(III) oxide — the familiar reddish-brown flaky coating. The colour of rust and the colour of the iron(III) hydroxide precipitate in your test tube are no coincidence.
This is exactly why the analytical scheme insists on noting the first colour. A student who glances at the tube a few minutes later might see brown and wrongly identify iron(III). The green that appears the instant the alkali is added is the genuine fingerprint; the brown is a second reaction happening afterwards, driven by the air in the laboratory.
The precipitate is reacting with oxygen dissolved from the air. Iron(II) is easily converted to iron(III), and the green hydroxide becomes the reddish-brown one:
Checking the balance: iron ; oxygen on the left and on the right; hydrogen on the left and on the right. Balanced.
The brown creeping down from the top is the oxygen working its way in. Shaking brings more of the precipitate into contact with air, which is why it speeds up the colour change.
The same chemistry is visible all around you.
- Green iron(II) sulphate crystals, left in an open jar, develop yellowish-brown patches as the surface changes to iron(III)
- Water from an iron-rich borewell often comes out clear, then turns cloudy and orange-brown in a bucket as dissolved iron(II) meets air and precipitates as iron(III) compounds
- The brown stains on a washbasin under such a tap are the same iron(III) hydroxide, left behind as the water dries
Rust itself is closely related. When iron objects corrode in moist air, the iron is converted through iron(II) to hydrated iron(III) oxide — the familiar reddish-brown flaky coating. The colour of rust and the colour of the iron(III) hydroxide precipitate in your test tube are no coincidence.
This is exactly why the analytical scheme insists on noting the first colour. A student who glances at the tube a few minutes later might see brown and wrongly identify iron(III). The green that appears the instant the alkali is added is the genuine fingerprint; the brown is a second reaction happening afterwards, driven by the air in the laboratory.
Exam relevance
How does identifying cations prepare you for JEE and NEET?
This is foundation work for the qualitative analysis section of JEE Main practical chemistry and for Class 12 The d- and f-Block Elements and Coordination Compounds, which are examined in both JEE Main and NEET Chemistry.
Where the scheme leads. JEE Main practical chemistry includes the chemical principles of qualitative salt analysis, with cations such as lead, copper, iron(III), zinc, calcium and magnesium in the list. At that level the cations are sorted into groups by group reagents — dilute hydrochloric acid, hydrogen sulphide, ammonium hydroxide with ammonium chloride, and others — but the logic is identical to the scheme here: one reagent, one observation, and the result chooses the next step.
Where the colours lead. Class 12 The d- and f-Block Elements explains why iron and copper compounds are coloured while zinc, calcium, magnesium and lead compounds in the same tests are white. The colour comes from partly filled d orbitals, which iron(II), iron(III) and copper(II) have and zinc, with a completely filled set, does not. Predicting whether an ion is coloured from its electronic configuration is a standard question in both exams.
Where the deep blue solution leads. Class 12 Coordination Compounds studies the tetraammine complexes of copper and zinc directly — their names, the coordination number of four, and the bonding of ammonia through its lone pair. The inky blue test result is used as a named example in questions on complex formation.
Where the calcium result leads. Class 11 Equilibrium explains precisely why weak ammonium hydroxide cannot precipitate calcium hydroxide, through the solubility product and the common ion effect. The same principle explains why ammonium chloride is added with ammonium hydroxide in group analysis — to keep the hydroxyl ion concentration low enough that only the least soluble hydroxides precipitate.
Question types to expect. At this level: colours, observations with equations, and identifying an ion from a set of results. In competitive papers: group reagents, colour from d-electron configuration, complex ion names and coordination numbers, and solubility-product reasoning, often as match-the-column questions pairing ions with precipitate colours.
The single trap that costs marks. Treating every dissolution in excess reagent as amphoterism. Copper and zinc hydroxides dissolve in ammonia by complex formation, and a question that asks why copper hydroxide dissolves in ammonia but not in sodium hydroxide can be answered only by keeping the two mechanisms separate.
A second trap. Assuming all transition metal ions are coloured. Zinc sits in the d-block but its compounds are white, because its d orbitals are full — a detail both exams test directly.
Board versus competitive emphasis. The ICSE paper marks the exact colour words, the observations at both stages and balanced equations; a competitive paper marks a group assignment, an electronic explanation of colour or a complex's name. The transferable habit is linking every observation to a cause — insoluble hydroxide, amphoterism or complex formation — because that is what turns memorised colours into chemistry.
Where the scheme leads. JEE Main practical chemistry includes the chemical principles of qualitative salt analysis, with cations such as lead, copper, iron(III), zinc, calcium and magnesium in the list. At that level the cations are sorted into groups by group reagents — dilute hydrochloric acid, hydrogen sulphide, ammonium hydroxide with ammonium chloride, and others — but the logic is identical to the scheme here: one reagent, one observation, and the result chooses the next step.
Where the colours lead. Class 12 The d- and f-Block Elements explains why iron and copper compounds are coloured while zinc, calcium, magnesium and lead compounds in the same tests are white. The colour comes from partly filled d orbitals, which iron(II), iron(III) and copper(II) have and zinc, with a completely filled set, does not. Predicting whether an ion is coloured from its electronic configuration is a standard question in both exams.
Where the deep blue solution leads. Class 12 Coordination Compounds studies the tetraammine complexes of copper and zinc directly — their names, the coordination number of four, and the bonding of ammonia through its lone pair. The inky blue test result is used as a named example in questions on complex formation.
Where the calcium result leads. Class 11 Equilibrium explains precisely why weak ammonium hydroxide cannot precipitate calcium hydroxide, through the solubility product and the common ion effect. The same principle explains why ammonium chloride is added with ammonium hydroxide in group analysis — to keep the hydroxyl ion concentration low enough that only the least soluble hydroxides precipitate.
Question types to expect. At this level: colours, observations with equations, and identifying an ion from a set of results. In competitive papers: group reagents, colour from d-electron configuration, complex ion names and coordination numbers, and solubility-product reasoning, often as match-the-column questions pairing ions with precipitate colours.
The single trap that costs marks. Treating every dissolution in excess reagent as amphoterism. Copper and zinc hydroxides dissolve in ammonia by complex formation, and a question that asks why copper hydroxide dissolves in ammonia but not in sodium hydroxide can be answered only by keeping the two mechanisms separate.
A second trap. Assuming all transition metal ions are coloured. Zinc sits in the d-block but its compounds are white, because its d orbitals are full — a detail both exams test directly.
Board versus competitive emphasis. The ICSE paper marks the exact colour words, the observations at both stages and balanced equations; a competitive paper marks a group assignment, an electronic explanation of colour or a complex's name. The transferable habit is linking every observation to a cause — insoluble hydroxide, amphoterism or complex formation — because that is what turns memorised colours into chemistry.
Key takeaways
What must you be able to do from this part?
One table of colours, two reagents at two stages each, and a four-step identification.
- Iron(II) salts: pale green solution. Iron(III) salts: yellowish-brown. Copper(II) salts: blue, copper chloride bluish-green
- Zinc, calcium, magnesium and lead salts: white solids, colourless solutions
- Zinc oxide is white cold and yellow hot; lead(II) oxide yellow; copper(II) oxide black
- With sodium hydroxide: dirty green, reddish-brown, pale blue, and white — all insoluble in excess
- Zinc and lead with sodium hydroxide: white precipitates that dissolve in excess as sodium zincate and sodium plumbite
- With ammonium hydroxide: dirty green and reddish-brown, insoluble; and white, insoluble
- Copper with excess ammonium hydroxide: pale blue precipitate dissolves to inky blue
- Zinc with excess ammonium hydroxide: white precipitate dissolves to colourless
- Calcium gives no precipitate with ammonium hydroxide, because the weak alkali supplies too few hydroxyl ions
- Zinc dissolves in excess of both reagents — by amphoterism in sodium hydroxide, by complex formation in ammonium hydroxide
- Identification: coloured precipitate names Fe(II), Fe(III) or Cu; white and soluble in excess NaOH means Zn or Pb, then NH4OH separates them; white and insoluble means Ca or Mg, then NH4OH separates them
- Iron(II) hydroxide turns brown in air: — record the first colour
- Add reagents drop by drop first, or a soluble precipitate can be missed
The sharpest self-test is seven imaginary test tubes. Label them with the seven metal ions, predict every observation with sodium hydroxide and ammonium hydroxide at both stages, and then cover the labels and try to identify each tube from your own predicted results.
- Iron(II) salts: pale green solution. Iron(III) salts: yellowish-brown. Copper(II) salts: blue, copper chloride bluish-green
- Zinc, calcium, magnesium and lead salts: white solids, colourless solutions
- Zinc oxide is white cold and yellow hot; lead(II) oxide yellow; copper(II) oxide black
- With sodium hydroxide: dirty green, reddish-brown, pale blue, and white — all insoluble in excess
- Zinc and lead with sodium hydroxide: white precipitates that dissolve in excess as sodium zincate and sodium plumbite
- With ammonium hydroxide: dirty green and reddish-brown, insoluble; and white, insoluble
- Copper with excess ammonium hydroxide: pale blue precipitate dissolves to inky blue
- Zinc with excess ammonium hydroxide: white precipitate dissolves to colourless
- Calcium gives no precipitate with ammonium hydroxide, because the weak alkali supplies too few hydroxyl ions
- Zinc dissolves in excess of both reagents — by amphoterism in sodium hydroxide, by complex formation in ammonium hydroxide
- Identification: coloured precipitate names Fe(II), Fe(III) or Cu; white and soluble in excess NaOH means Zn or Pb, then NH4OH separates them; white and insoluble means Ca or Mg, then NH4OH separates them
- Iron(II) hydroxide turns brown in air: — record the first colour
- Add reagents drop by drop first, or a soluble precipitate can be missed
The sharpest self-test is seven imaginary test tubes. Label them with the seven metal ions, predict every observation with sodium hydroxide and ammonium hydroxide at both stages, and then cover the labels and try to identify each tube from your own predicted results.