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In the practical exam you get a small heap of powder and limited time. Treat salt analysis as a decision tree, where each test rules something in or out, not as a list of colours. This post walks through the scheme, explains why the groups separate as they do, and works two salts from start to finish. Your school's lab manual and your teacher's safety instructions always come first.
Which ions can you be given?
The CBSE Chemistry curriculum for 2026-27 asks Class 12 students to determine one anion and one cation in a given salt, with insoluble salts excluded. The cations listed are Pb2+, Cu2+, Al3+, Fe3+, Mn2+, Ni2+, Zn2+, Co2+, Ca2+, Sr2+, Ba2+, Mg2+ and NH4+. The anions are CO3 2-, S2-, SO3 2-, NO3-, NO2-, Cl-, Br-, I-, SO4 2-, PO4 3-, CH3COO- and C2O4 2-.
The JEE (Main) 2026 syllabus lists "chemical principles involved in the qualitative salt analysis" for a shorter list: Pb2+, Cu2+, Al3+, Fe3+, Zn2+, Ni2+, Ca2+, Ba2+, Mg2+ and NH4+, with CO3 2-, S2-, SO4 2-, NO3-, NO2-, Cl-, Br- and I-. JEE asks why a test works; the board practical asks you to carry it out. Check both documents for your exam year.
Step 1: preliminary observations
Note colour and smell before adding anything. Blue suggests copper, green suggests nickel, and pale pink suggests manganese or cobalt. These are hints only; let the tests decide.
Step 2: how do you find the anion?
Work in a fixed order so that one test does not hide another.
Dilute H2SO4 with the solid. Brisk effervescence of a gas that turns lime water milky points to carbonate. A rotten-egg smell and a gas that blackens lead acetate paper points to sulphide. A gas that turns acidified potassium dichromate green points to sulphite. Light brown fumes point to nitrite. A smell of vinegar on rubbing the salt with the acid points to acetate.
If nothing happens, warm with concentrated H2SO4. Colourless fumes that give dense white fumes with a glass rod dipped in ammonia point to chloride. Reddish-brown vapour points to bromide, violet vapour to iodide. Brown fumes that get stronger with copper turnings point to nitrate. A colourless gas mixture in which carbon monoxide burns with a blue flame points to oxalate.
If neither acid reacts, use independent tests. With the salt solution, barium chloride gives a white precipitate insoluble in concentrated HCl for sulphate. Ammonium molybdate with nitric acid gives a yellow precipitate for phosphate.
Confirm each anion with its confirmatory test, such as silver nitrate for halides or the brown ring test for nitrate.
Step 3: how do you find the cation?
Test the original salt for ammonium first: heat with NaOH. Ammonia that turns moist red litmus blue and gives a brown precipitate with Nessler's reagent confirms NH4+. Then add group reagents to the salt solution in this order, stopping at the first precipitate:
Group I, dilute HCl: Pb2+ (white PbCl2).
Group II, H2S in dilute HCl: Cu2+ (black CuS), and any Pb2+ left over.
Group III, NH4OH in the presence of NH4Cl: Fe3+ (reddish-brown) and Al3+ (white gelatinous) hydroxides.
Group IV, H2S in ammoniacal solution: Zn2+ (white ZnS), Mn2+ (buff), Ni2+ and Co2+ (black).
Group V, (NH4)2CO3 with NH4OH and NH4Cl: Ba2+, Sr2+ and Ca2+ carbonates.
Group VI: Mg2+, tested with disodium hydrogen phosphate.
PbCl2 is slightly soluble in cold water, so some Pb2+ can reach group II as black PbS. Check before you write "copper".
Why does the group order work?
The scheme applies the common ion effect from the NCERT Class 11 equilibrium chapter: adding an ion already present in an equilibrium shifts it.
Group II: HCl adds H+, suppressing the ionisation of H2S and keeping sulphide very low. Only sulphides with tiny solubility products, such as CuS, precipitate; ZnS and NiS stay dissolved.
Group IV: in ammoniacal solution H+ is removed, sulphide rises, and ZnS, MnS, NiS and CoS precipitate.
Group III: NH4Cl suppresses the ionisation of NH4OH, so there is enough hydroxide for Fe(OH)3 and Al(OH)3 but not for Zn(OH)2 or Mg(OH)2.
JEE asks exactly this: why NH4Cl comes before NH4OH, or why group II is done in acid.
Worked salt 1: a blue crystalline salt
Blue, no smell. No gas with dilute or concentrated H2SO4.
Barium chloride: white precipitate, insoluble in concentrated HCl. Anion: SO4 2-.
No ammonia with NaOH. No precipitate with dilute HCl.
H2S in dilute HCl: black precipitate. Group II.
Dissolve the precipitate in hot dilute nitric acid and add excess NH4OH: deep blue solution. With potassium ferrocyanide in acetic acid: chocolate-brown precipitate. Cation: Cu2+.
Result: copper sulphate.
Worked salt 2: a white salt
White, no smell. Dilute H2SO4: no gas.
Concentrated H2SO4: colourless fumes, dense white fumes with an ammonia rod. With dilute nitric acid and silver nitrate: curdy white precipitate, soluble in NH4OH. Anion: Cl-.
Heating with NaOH: gas turning red litmus blue, brown precipitate with Nessler's reagent. Cation: NH4+.
Result: ammonium chloride.
Common mistakes in the practical
Adding the next reagent without filtering off the last precipitate.
Forgetting NH4Cl before NH4OH in group III.
Calling the cation from colour alone, or writing a result without its confirmatory test.
Working with H2S or concentrated acid outside the fume cupboard. H2S is toxic; follow your lab's rules exactly.
A practice routine before the exam
Draw the anion and cation flowcharts from memory, with reagent and observation on each branch.
For each group, write one sentence on why its reagent works.
Do three full mock analyses on paper with salts your teacher chooses.
Rehearse the viva: why dilute HCl, why NH4Cl, why test NH4+ on the original salt.
Our guides to the board practical exam and viva and to ionic equilibrium, pH and buffers go further on presentation and on the equilibria behind the groups. If the theory behind a test is the gap, the Fermi Concept Library covers every concept in chemistry, each with its own lesson.
Tonight, draw the cation group flowchart from memory and write, next to each reagent, the one-line reason it works.
