Inorganic Salt Analysis

Before any reagent is added, record the physical appearance of the salt: colour, crystalline form, and any odour. Colour gives an immediate clue — blue or green solids often contain Cu²⁺ or Fe²⁺, yellow-brown solids may contain Fe³⁺, and white or colourless solids narrow the field to many common cations such as Na⁺, K⁺, Ca²⁺, Mg²⁺, Zn²⁺, Al³⁺, NH₄⁺, and Pb²⁺.

A clean platinum or nichrome wire is dipped into the powdered salt and held in the hottest part of a Bunsen flame. Each cation imparts a characteristic emission colour: Li⁺ → crimson, Na⁺ → persistent yellow, K⁺ → lilac (viewed through blue cobalt glass to remove sodium masking), Ca²⁺ → brick-red, Sr²⁺ → scarlet, Ba²⁺ → pale green, Cu²⁺ → blue-green. The colours arise because electrons in metal ions are excited by the flame's thermal energy and emit photons of specific wavelengths as they return to lower energy levels.

Heating the solid in a hard-glass test tube provides further clues. Carbonates (CO₃²⁻) of metals other than Group 1 decompose to give the metal oxide and CO₂, which turns limewater milky: MCO₃ → MO + CO₂. Nitrates of Group 1 metals (except Li) decompose to nitrites and O₂; nitrates of most other metals decompose further to the oxide, NO₂ (brown gas), and O₂. Ammonium salts leave no residue. Hydrated salts lose water and may show colour changes (e.g., blue CuSO₄·5H₂O becomes white anhydrous CuSO₄).

Adding dilute HCl to the solid or its solution tests for several anions. Effervescence with CO₂ (turns limewater milky) confirms CO₃²⁻ or HCO₃⁻. A pungent smell of SO₂ may indicate SO₃²⁻ or S²O₃²⁻. A white precipitate of PbCl₂ or AgCl suggests Pb²⁺ or Ag⁺ respectively. A white precipitate of BaSO₄ (with dilute HCl and then BaCl₂) confirms SO₄²⁻ even after the carbonate has been acidified away.

Dissolve the salt in deionised water or dilute acid. Add dilute NaOH dropwise and then in excess to test for metal hydroxide precipitates. Al³⁺ gives a white precipitate soluble in excess NaOH (amphoteric). Cu²⁺ gives a pale-blue precipitate insoluble in excess NaOH but soluble in excess NH₃ to give the deep-blue tetraamminecopper(II) complex [Cu(NH₃)₄]²⁺. Fe²⁺ gives a dirty-green precipitate; Fe³⁺ gives a red-brown precipitate; Zn²⁺ gives a white precipitate soluble in excess NaOH; Mg²⁺ gives a white precipitate insoluble in excess. NH₄⁺ is confirmed by warming with NaOH — the released NH₃ gas turns damp red litmus blue.

Halides (Cl⁻, Br⁻, I⁻) are confirmed by adding dilute HNO₃ followed by AgNO₃ solution: Cl⁻ → white AgCl (soluble in dilute NH₃); Br⁻ → pale-yellow AgBr (sparingly soluble in concentrated NH₃); I⁻ → yellow AgI (insoluble in NH₃). Sulfate (SO₄²⁻) is confirmed by a white precipitate of BaSO₄ with acidified BaCl₂. Phosphate (PO₄³⁻) gives a yellow precipitate of ammonium phosphomolybdate with ammonium molybdate reagent in HNO₃. Nitrate (NO₃⁻) is confirmed by the brown-ring test: add FeSO₄ solution and carefully layer concentrated H₂SO₄; a brown ring of [Fe(NO)(H₂O)₅]²⁺ forms at the interface.

• \ce{Cu^{2+}} — Copper(II) ion (mol dm⁻³): The central metal ion acting as a Lewis acid, accepting lone pairs from the four ammonia ligands
• \ce{NH_3} — Ammonia ligand (mol dm⁻³): A Lewis base donating a lone pair of electrons to the copper centre; four molecules coordinate in a square-planar arrangement
• \ce{[Cu(NH_3)_4]^{2+}} — Tetraamminecopper(II) ion (mol dm⁻³): The deep-blue complex ion formed; its intense colour arises from d–d electronic transitions in the visible region

Problem. An unknown white solid is dissolved in deionised water to give a colourless solution. A flame test gives a persistent bright yellow colour. When dilute HNO₃ and then AgNO₃ solution are added to the solution, a white precipitate forms that dissolves readily in dilute aqueous ammonia. Adding BaCl₂ solution acidified with dilute HCl produces no precipitate. Identify the cation and anion present, and write the net ionic equation for the silver salt precipitation.

Answer. The salt is sodium chloride (NaCl). The cation is Na⁺ (yellow flame) and the anion is Cl⁻ (white AgCl precipitate soluble in dilute NH₃).

• Solubility Rules and Precipitation Reactions — Understanding which ionic compounds are soluble or insoluble in water is essential for predicting the precipitates formed during anion and cation tests in salt analysis.
• Acid–Base Reactions and Salt Formation — Many preliminary tests in salt analysis (e.g., reaction with HCl or NaOH) rely on acid–base chemistry to release anions as gases or to precipitate metal hydroxides.
• Transition Metal Complex Ions — Confirmatory tests for Cu²⁺, Fe²⁺, and Fe³⁺ involve the formation of coloured coordination complexes, making knowledge of ligand substitution and d–d transitions directly relevant.
• Flame Emission and Atomic Spectroscopy — Flame tests are interpreted using the theory of electronic transitions and quantised energy levels; understanding why each metal emits a distinct colour deepens appreciation of the test's mechanism.
• Identification of Gases Produced in Chemical Reactions — Several anion tests produce characteristic gases (CO₂, SO₂, NH₃, NO₂, HCl) that must be correctly identified using standard test procedures such as limewater, litmus paper, and dichromate paper.