Why Preliminary Tests Come First
Qualitative analysis of an inorganic salt means identifying the cation and anion present. In a salt such as $\ce{CuSO4}$ the cation is $\ce{Cu^2+}$ and the anion is $\ce{SO4^2-}$; the cation is the part contributed by the base, the anion the part contributed by the acid. NCERT lays the work out in three stages: a preliminary examination of the solid salt and its solution, then wet tests for anions, then wet tests for cations.
The preliminary stage rests on reactions perceptible directly to sight and smell — formation of a precipitate, change in colour, evolution of a gas. NCERT is explicit that these dry tests are not conclusive but "sometimes they give quite important clues for the presence of certain anions or cations," and that preliminary tests should always be performed before starting the confirmatory tests. A salt that is snow-white, odourless and freely soluble has already ruled out the strongly coloured transition-metal cations before a single reagent is added.
Characteristic salt colours and the cations NCERT associates with them (Table 7.6).
Colour of the Salt and the Ions It Suggests
The colour of a dry salt is the first datum recorded. Coloured salts almost always carry a transition-metal cation whose partially filled d-subshell absorbs part of the visible spectrum. NCERT Table 7.6 maps the common colours to their cations; a white salt indicates that none of these strongly coloured ions is present, and the cation is likely one of the colourless ions such as $\ce{Na+}$, $\ce{K+}$, $\ce{Ca^2+}$, $\ce{Ba^2+}$, $\ce{Mg^2+}$, $\ce{Al^3+}$, $\ce{Zn^2+}$ or $\ce{NH4+}$.
| Observed colour of salt | Cation(s) indicated |
|---|---|
| Blue | $\ce{Cu^2+}$ |
| Bright green | $\ce{Ni^2+}$ |
| Light green, yellow or brown | $\ce{Fe^2+}$, $\ce{Fe^3+}$ |
| Light pink | $\ce{Mn^2+}$ |
| Blue, red, violet or pink | $\ce{Co^2+}$ |
| White / colourless | None of the above coloured cations |
Within this scheme a pale or light green points to ferrous iron $\ce{Fe^2+}$, while a deeper brown points to ferric iron $\ce{Fe^3+}$ — the oxidation state shifts the colour. Copper salts are the most reliable single colour cue: the blue of hydrated $\ce{CuSO4.5H2O}$ is unmistakable. Cobalt is the trickiest, since $\ce{Co^2+}$ salts span pink (hydrated) through red, violet and blue depending on the anion and degree of hydration.
Colour proves nothing on its own
Examiners frame colour as an assertion–reason or single-statement item. The safe stance: salt colour is only a clue. A white salt does not guarantee a colourless cation (some hydrates lose colour on dehydration), and a coloured salt must still be confirmed by group analysis.
NCERT wording: preliminary tests are "not conclusive" — never tick an option that treats colour as proof of identity.
Colour Cold vs On Heating
The dry heating test (about 0.1 g of salt heated for a minute) is read together with the original colour. A change of colour between the cold solid and the hot residue, and whether the colour returns on cooling, is itself diagnostic. NCERT Table 7.7 records the reversible colour shifts most often used.
| Colour when cold | Colour when hot | Inference |
|---|---|---|
| Blue | White | $\ce{Cu^2+}$ |
| Green | Dirty white or yellow | $\ce{Fe^2+}$ |
| White | Yellow | $\ce{Zn^2+}$ |
| Pink | Blue | $\ce{Co^2+}$ |
The classic case is zinc oxide formed on heating a zinc salt: white when cold, yellow when hot, returning to white on cooling. Copper(II) sulphate loses its water of crystallisation on heating, fading from blue to white, then turns blue again when re-hydrated. These shifts belong to the dry heating test, but they are read in the same breath as the cold colour during the preliminary survey.
Smell — Ammoniacal, Vinegar, Rotten-Egg
Smell is the second sense the analyst uses. Some salts are odourless as solids but release a tell-tale gas on contact with acid; others are recognised directly. NCERT's gas tables for dilute and concentrated acid action supply the diagnostic odours, which point to specific anions or to the ammonium cation.
| Smell detected | Gas / vapour | Ion suggested |
|---|---|---|
| Ammoniacal (pungent, like a cleaning agent) | $\ce{NH3}$ | Ammonium, $\ce{NH4+}$ |
| Vinegar-like | $\ce{CH3COOH}$ vapour | Acetate, $\ce{CH3COO-}$ |
| Rotten eggs | $\ce{H2S}$ | Sulphide, $\ce{S^2-}$ |
| Burning sulphur (suffocating, pungent) | $\ce{SO2}$ | Sulphite, $\ce{SO3^2-}$ |
An ammonium salt warmed with an alkali liberates ammonia, recognised by its sharp smell and by white fumes when a rod moistened with HCl is brought near. Acetates treated with dilute sulphuric acid give vapours that smell of vinegar and turn blue litmus red:
$$\ce{2CH3COONa + H2SO4 -> Na2SO4 + 2CH3COOH}$$
A sulphide warmed with dilute sulphuric acid gives the rotten-egg smell of $\ce{H2S}$, which blackens lead acetate paper, while a sulphite gives the choking smell of burning sulphur as $\ce{SO2}$ escapes:
$$\ce{Na2S + H2SO4 -> Na2SO4 + H2S ^}$$
$$\ce{Na2SO3 + H2SO4 -> Na2SO4 + H2O + SO2 ^}$$
Both $\ce{CO2}$ and $\ce{SO2}$ turn lime water milky
A carbonate and a sulphite both give a gas that turns lime water milky. NCERT flags the discriminator: $\ce{CO2}$ is odourless, whereas $\ce{SO2}$ has the characteristic suffocating smell of burning sulphur. Use the smell, not the lime-water test, to tell them apart.
Odourless milky gas → carbonate; pungent milky gas → sulphite.
The gas-evolution clues above feed straight into the wet anion scheme — see Anion tests with dilute acid for the confirmatory steps.
The Flame Test — Procedure and Colours
The flame test exploits the fact that chlorides of several metals are volatile in a non-luminous flame and impart a characteristic colour. NCERT prescribes a clean platinum wire (nichrome serves as a cheaper substitute) with a tiny loop at one end. The loop is cleaned by dipping in concentrated hydrochloric acid and holding it in the flame, repeating until it imparts no colour of its own.
A small quantity of the salt is then made into a paste with 2–3 drops of concentrated HCl on a watch glass, the clean loop is dipped in the paste, and the loop is introduced into the non-luminous (oxidising) flame. The colour is observed first with the naked eye and then through blue (cobalt) glass before assigning the metal ion.
Clean Pt/nichrome loop, conc. HCl paste, and the non-luminous Bunsen flame.
The colours imparted to the flame, in NCERT's order and supplemented by the standard alkali-metal colours, are summarised below. The "through blue glass" column matters because cobalt glass absorbs sodium's yellow and reveals colours otherwise masked by it.
| Cation | Flame colour (naked eye) | Through blue glass |
|---|---|---|
| $\ce{Na+}$ | Golden yellow | Invisible (absorbed) |
| $\ce{K+}$ | Violet / lilac | Crimson-violet |
| $\ce{Ca^2+}$ | Brick red | Green |
| $\ce{Sr^2+}$ | Crimson red | Purple |
| $\ce{Ba^2+}$ | Apple green | Bluish green |
| $\ce{Cu^2+}$ | Green with blue centre | Same colour as naked eye |
The conversion of the salt to its chloride is the key chemical step: even a sulphate or carbonate, once pasted with concentrated HCl, presents the metal as a volatile chloride that vaporises in the flame and emits. This is why NCERT specifies concentrated HCl rather than water for making the paste.
Why a Flame Test Works
The colour is an emission phenomenon. Heat from the flame supplies energy to the valence electrons of the metal atom, promoting them from their ground state to higher energy levels — the atom is said to be excited. The excited state is unstable; the electron falls back toward a lower level, and the energy difference is released as a photon of light.
The wavelength of that photon is fixed by the size of the energy gap, $\Delta E = h\nu = \dfrac{hc}{\lambda}$, and because each element has its own unique spacing of energy levels, each emits light of definite, characteristic wavelength. Sodium's intense doublet near 589 nm reads as golden yellow; potassium's transition lies in the violet. This is the same physics that underpins atomic emission spectroscopy.
Sodium contamination drowns potassium
Sodium's emission is so intense that the faintest trace masks potassium's violet flame. NCERT directs that the flame be viewed through blue (cobalt) glass, which absorbs the yellow sodium light and lets the crimson-violet of $\ce{K+}$ through.
If the question asks why blue glass is used → to filter out interfering yellow sodium light and reveal potassium.
Solubility, pH and Deliquescence
Solubility of the salt in water, and the pH of the resulting solution, carry real information. NCERT notes that if the solution is acidic or basic, the salt is being hydrolysed: a basic solution suggests a carbonate or sulphide, while an acidic solution suggests an acid salt or the salt of a weak base and a strong acid. In the acidic case it is best to neutralise the solution with sodium carbonate before testing for anions.
| Observation | Inference (NCERT) |
|---|---|
| Aqueous solution basic | Salt may be a carbonate, sulphide, etc. |
| Aqueous solution acidic | Acid salt, or salt of weak base + strong acid (hydrolysis) |
| Salt absorbs moisture, turns moist/pasty in air | Deliquescent salt (e.g. several chlorides, nitrates) |
| Salt insoluble in water | Excluded from this scheme; use sodium carbonate extract for anion tests |
Deliquescence is a physical clue rather than a confirmatory test: a deliquescent salt draws water vapour from the atmosphere and becomes moist or even forms a solution. Encountering a sample that is damp despite dry storage warns that it is hygroscopic and must be handled and weighed quickly. For NEET-relevant practical reasoning, the connected idea is that some hydrated salts must be acidified to prevent hydrolysis — the principle behind preparing Mohr's salt with dilute sulphuric acid.
Texture and Density Notes
Beyond colour, smell and solubility, the analyst notes the physical form of the sample — whether it is crystalline or amorphous (powdery), and its apparent density when handled. Well-formed crystals point to a single pure salt; a fluffy, light powder of low density and a dense, heavy granular solid behave differently when scooped and dissolved, and the difference can hint at the class of compound. These observations are weak clues, never decisive, and always read alongside the colour, smell and solubility data before moving to the systematic wet tests.
Preliminary tests in one pass
- Colour: blue → $\ce{Cu^2+}$; bright green → $\ce{Ni^2+}$; light green/yellow/brown → $\ce{Fe^2+}$/$\ce{Fe^3+}$; light pink → $\ce{Mn^2+}$; blue/red/violet/pink → $\ce{Co^2+}$; white → none of these.
- Cold vs hot: blue→white = $\ce{Cu^2+}$; white→yellow = $\ce{Zn^2+}$; pink→blue = $\ce{Co^2+}$.
- Smell: ammoniacal → $\ce{NH4+}$; vinegar → acetate; rotten egg → sulphide; burning sulphur → sulphite.
- Flame: Na yellow, K violet (lilac, crimson-violet through blue glass), Ca brick-red, Sr crimson, Ba apple-green, Cu green-blue — colour arises from electronic excitation and emission.
- Solution/solubility: basic → carbonate/sulphide; acidic → hydrolysing salt; deliquescent → hygroscopic. All clues, never proof.