Experiment: salt analysis, cations

How are cations identified in salt analysis?

Where this experiment comes from. The JEE Main 2027 and NEET (UG) 2026 syllabuses list the chemical principles of qualitative salt analysis in Unit 20, for the cations Pb²⁺, Cu²⁺, Al³⁺, Fe³⁺, Zn²⁺, Ni²⁺, Ca²⁺, Ba²⁺, Mg²⁺ and NH₄⁺, with insoluble salts excluded. The NCERT textbooks give some of the theory: the colour reactions of metal ions with ligands such as dimethylglyoxime (NCERT Class 12 Section 5.7, page 136), and solubility equilibria and the common ion effect (NCERT Class 11 Section 6.13, pages 204–206). They do not describe the method, so this page follows the NCERT Laboratory Manual (Chemistry, Class 12), Unit 7. That unit also covers As³⁺, Mn²⁺, Co²⁺ and Sr²⁺, which these syllabuses leave out.

In plain words: each group reagent precipitates only the cations whose compound is least soluble under those conditions. Sulphide is the clearest case: in acid there is so little sulphide that only the least soluble sulphides come down (group II), while in ammonia there is more, and further sulphides precipitate (group IV).

• Principle: a salt precipitates when its ionic product exceeds its solubility product, and the ionic product is controlled through the common ion effect (Lab Manual, page 50; see the common ion effect).
• Original solution: dissolve the salt in water; if it does not dissolve, try dilute HCl, then concentrated HCl, then dilute nitric acid, then aqua regia (concentrated HCl and concentrated HNO₃, 3 : 1), heating where needed (Lab Manual, page 69).
• Preliminary tests give clues but are not conclusive (pages 63–68). For example:
  • a blue salt suggests Cu²⁺;
  • in the flame test (a platinum wire, concentrated HCl, a non-luminous flame), copper gives a green flame with a blue centre, calcium brick red and barium apple green.

The groups and reagents are the NCERT Lab Manual’s Table 7.11 (page 71), limited to the syllabus cations. Group I comes down as the chloride, groups II and IV as sulphides, group III as hydroxides and group V as carbonates (Lab Manual, page 70).

Group zero: NH_₄^⁺

• Heat 0.1 g of the salt with 1–2 mL of NaOH solution. A smell of ammonia shows ammonium ions; a glass rod dipped in HCl gives white fumes of NH₄Cl (Lab Manual, page 69).
• The gas gives a brown precipitate with Nessler’s reagent (Lab Manual, page 69).

• Note (Lab Manual, page 70): the brown product of the Nessler test is named “basic mercury(II) amido-iodine”. It contains iodide, as its formula HgO·Hg(NH₂)I shows, so “amido-iodide” is the better name (reasoned).

Group I: Pb^²⁺

• Dilute HCl gives a white precipitate of PbCl₂. It dissolves in hot water (Lab Manual, pages 71–72).
• Divide the hot solution into three parts (Lab Manual, pages 71–72):
  1. potassium iodide gives yellow PbI₂, which dissolves in boiling water and comes back on cooling as shining crystals;
  2. potassium chromate gives yellow PbCrO₄, which dissolves in NaOH;
  3. alcohol and dilute H₂SO₄ give white PbSO₄, which dissolves in ammonium acetate.
• Lead appears in both group I and group II. Lead chloride is fairly soluble, especially in hot water, so the lead that stays in solution is caught as PbS in group II (reasoned from Lab Manual, pages 70–72).

Group II: Pb^²⁺ and Cu^²⁺

• If group I is absent, add water, warm, and pass H₂S for 1–2 minutes. A black precipitate shows Cu²⁺ or Pb²⁺ (Lab Manual, page 72).
• Lead: boil the precipitate with dilute nitric acid, then add a few drops of alcohol and dilute H₂SO₄: white PbSO₄. In ammonium acetate solution acidified with acetic acid, this gives yellow PbCrO₄ with potassium chromate and yellow PbI₂ with potassium iodide (Lab Manual, page 73).
• Copper: if no precipitate forms with H₂SO₄, add excess ammonium hydroxide. A deep blue solution of [Cu(NH₃)₄]²⁺ forms; see naming such complexes. Acidify with acetic acid and add potassium ferrocyanide: a chocolate-brown precipitate of Cu₂[Fe(CN)₆] (Lab Manual, pages 73–74).
• NCERT slip (Laboratory Manual, page 74): the equation for dissolving basic copper sulphate in ammonia is printed as Cu(OH)₂·CuSO₄ + 8NH₃ → 2[Cu(NH₃)₄]SO₄ + 2OH⁻ + SO₄²⁻. That has one sulphate on the left and three on the right, and the charges do not balance either: 0 on the left, −4 on the right. Balanced: Cu(OH)₂·CuSO₄ + 8NH₃ → 2[Cu(NH₃)₄]²⁺ + SO₄²⁻ + 2OH⁻ (reasoned: count the atoms and charges). In an exam, write the balanced form.
• Note (NCERT Laboratory Manual, page 72): the H₂S step refers to “Fig. 7.6”, but Kipp’s apparatus for H₂S is Fig. 7.8; Fig. 7.6 is the charcoal cavity test.

Group III: Al^³⁺ and Fe^³⁺

• Take the original solution and add 2–3 drops of concentrated HNO₃, which oxidises any Fe²⁺ to Fe³⁺, and heat. Cool, add a little solid NH₄Cl and excess NH₄OH until the solution smells of ammonia. A gelatinous white precipitate shows Al³⁺, and a reddish-brown one shows Fe³⁺ (Lab Manual, page 75).
• The NH₄Cl, through its common ion NH₄⁺, holds back the ionisation of NH₄OH, so the hydroxide ion concentration stays just high enough to precipitate the group III hydroxides but not those of later groups (reasoned from Lab Manual, page 50).
• Aluminium (dissolve the precipitate in dilute HCl) (Lab Manual, pages 75–76):
  • NaOH gives white gelatinous Al(OH)₃, soluble in excess as sodium meta-aluminate;
  • in the lake test, add blue litmus, then NH₄OH drop by drop: Al(OH)₃ takes up the blue dye, and a blue mass floats in a colourless solution.
• Iron(III) (dissolve in dilute HCl) (Lab Manual, pages 75–76):
  • potassium ferrocyanide gives Prussian blue, Fe₄[Fe(CN)₆]₃;
  • potassium thiocyanate gives a blood-red colour, [Fe(SCN)]²⁺.

Group IV: Zn^²⁺ and Ni^²⁺

• If group III is absent, pass H₂S into the ammoniacal solution. A white precipitate is ZnS, and a black one is NiS (Lab Manual, pages 76–77).
• Zinc: dissolve ZnS in dilute HCl by boiling (Lab Manual, pages 77–78):
  • NaOH gives white Zn(OH)₂, soluble in excess, as sodium zincate;
  • after neutralising with NH₄OH, potassium ferrocyanide gives a bluish-white precipitate of zinc ferrocyanide.
• Nickel: dissolve NiS in aqua regia, evaporate to dryness and dissolve the residue in water. Make it alkaline with NH₄OH and add dimethylglyoxime: a bright red precipitate (Lab Manual, pages 77–78).

Group V: Ca^²⁺ and Ba^²⁺

• If group IV is absent, add solid NH₄Cl, excess NH₄OH and solid ammonium carbonate to the original solution: a white precipitate of carbonate. Dissolve it by boiling with dilute acetic acid, and keep a little precipitate for the flame test (Lab Manual, page 79).
• Barium: potassium chromate gives yellow BaCrO₄; the flame is grassy green (Lab Manual, pages 79–80).
• Calcium (if barium is absent): ammonium oxalate gives white calcium oxalate; the flame is brick red (Lab Manual, pages 79–80).
• Note on flame colours (NCERT Laboratory Manual, pages 65 and 79–80): Table 7.8 calls the barium flame apple green, while the group V tests call it grassy green. Through blue glass, Table 7.8 gives green for calcium and the group V test gives greenish-yellow. These are different words for the same flames, and both are NCERT’s own (reasoned). A flame colour is only a clue (Lab Manual, page 49), so give the confirmatory test with it.

Group VI: Mg^²⁺

• If group V is absent, take the original solution, add ammonium hydroxide and then disodium hydrogenphosphate, and scratch the inside of the tube with a glass rod: a white crystalline precipitate of magnesium ammonium phosphate (Lab Manual, pages 70, 81 and 85). It may appear only after a while, so warm and scratch (Lab Manual, page 81).
• Magnesium carbonate stays dissolved in group V because the ammonium salts keep the carbonate ion concentration low: NH₄⁺ + CO₃²⁻ ⇌ NH₃ + HCO₃⁻ (Lab Manual, page 81).
• NCERT slip (Laboratory Manual, page 81): the equation is printed as Mg²⁺ + Na₂HPO₄ → Mg(NH₄)PO₄ + NH₄OH + 2Na⁺ + H₂O, with no nitrogen on the left, so it cannot balance. The ammonia comes from the ammoniacal solution and belongs on the left: Mg²⁺ + Na₂HPO₄ + NH₄OH → Mg(NH₄)PO₄ + 2Na⁺ + H₂O (reasoned: count the atoms). In an exam, write the balanced form.

Safety

• The Lab Manual’s general precautions (pages 81–82):
  • wear an apron, eye protection and gloves;
  • read every label, and never use an unlabelled reagent;
  • never taste chemicals, and waft vapours towards the nose rather than sniffing;
  • always add acid to water;
  • never point a test tube at anyone while heating it.
• It marks mercury salts (Nessler’s reagent) as toxic (page 70), hydrogen sulphide as flammable and toxic (page 72), the alcohol of the lead tests as flammable and toxic (page 73), and nitric acid, used before group III, as oxidising and corrosive (page 61). Pass H₂S only in a fume cupboard (reasoned).

Solved examples

Common mistakes to avoid

Quick revision summary

• Order of the groups:
  • zero: NH₄⁺, with NaOH and heat;
  • I: PbCl₂, with dilute HCl;
  • II: PbS and CuS, with H₂S in acid;
  • III: Al(OH)₃ and Fe(OH)₃, with NH₄Cl and NH₄OH;
  • IV: ZnS (white) and NiS (black), with H₂S in NH₄OH;
  • V: CaCO₃ and BaCO₃, with (NH₄)₂CO₃;
  • VI: Mg²⁺, with NH₄OH and Na₂HPO₄.
• Confirm: Nessler brown (NH₄⁺); KI and K₂CrO₄ yellow (Pb²⁺); deep blue with NH₃, chocolate-brown with ferrocyanide (Cu²⁺); lake test (Al³⁺); Prussian blue and blood red (Fe³⁺); bluish-white ferrocyanide (Zn²⁺); red with dimethylglyoxime (Ni²⁺); yellow BaCrO₄, green flame (Ba²⁺); white oxalate, brick red flame (Ca²⁺); white MgNH₄PO₄ (Mg²⁺).

Practice questions

Try each question first, then open the answer to check your working.

1. Q1. Why is concentrated nitric acid added to the original solution before the group III reagent?

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   To oxidise any Fe²⁺ to Fe³⁺. Iron is then precipitated completely as Fe(OH)₃ in group III (Lab Manual, page 75; the reason for the completeness is reasoned).

2. Q2. How do you tell a white group IV precipitate from a black one, and how is the black one confirmed?

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   White is ZnS (zinc); black is NiS (nickel, among the syllabus ions). Dissolve the black precipitate in aqua regia, evaporate, take up in water, make alkaline with NH₄OH and add dimethylglyoxime: a bright red precipitate confirms nickel (Lab Manual, pages 77–78).