Practical chemistry questions in NEET are "free marks" if you know the reagent, the observation and the reason behind it — they test recognition of standard reactions and colour changes, not memorization of long procedural steps. A handful of named reagents, colour changes, and detection schemes cover almost every question this unit produces.
Volumetric (Titrimetric) Analysis
Most titration questions reduce to unit conversions plus one balanced equation. Concentration can be expressed as molarity ($M = \text{mol solute}/\text{L solution}$), normality ($N = M \times n\text{-factor}$, where the $n$-factor is the number of replaceable H⁺/OH⁻ or electrons transferred), or molality (mol solute/kg solvent — temperature-independent, unlike molarity). At the equivalence point of any titration, milliequivalents of titrant equal milliequivalents of analyte: $N_1V_1 = N_2V_2$.
- Acid–base titrations use indicators whose own colour-change range must straddle the equivalence-point pH: phenolphthalein (colourless→pink, range 8.3–10) for a weak-acid/strong-base titration where the equivalence point is alkaline; methyl orange (red→yellow, range 3.1–4.4) for a strong-acid/weak-base titration where it is acidic. Using the wrong indicator gives a visibly wrong end point even with perfect technique.
- Redox titrations often self-indicate: KMnO₄ is its own indicator (colourless→persistent pale pink at the end point, since MnO₄⁻ is intensely purple and Mn²⁺ is almost colourless in acidic medium). Iodometric titrations use starch, added only near the end point (a deep blue-black complex with I₂) — adding it too early can trap iodine in the starch and give a fading, indistinct end point.
- Always rinse a burette with the solution it will hold (not just water) and a pipette similarly, to avoid dilution errors — a classic "why did my titre come out high" trap.
Qualitative Salt Analysis
Cation analysis proceeds group-wise by selectively precipitating cations with a sequence of reagents, exploiting differences in solubility product:
- Group I (Pb²⁺): precipitated by dilute HCl as PbCl₂ (soluble in hot water, distinguishing it from other group precipitates).
- Group II (Cu²⁺): precipitated as sulfides by H₂S in dilute HCl (acidic medium keeps [S²⁻] low, precipitating only the least-soluble sulfides).
- Group III (Al³⁺, Fe³⁺): precipitated as hydroxides by NH₄OH in the presence of NH₄Cl (which suppresses NH₄OH's own ionisation via the common-ion effect, keeping [OH⁻] just high enough for these hydroxides but not others).
- Group IV (Zn²⁺, Ni²⁺, Co²⁺): precipitated as sulfides by H₂S in ammoniacal (basic) medium — the same H₂S reagent as Group II, but the higher pH now pushes [S²⁻] high enough to precipitate more-soluble sulfides too.
- Group V (Ca²⁺, Ba²⁺): precipitated as carbonates by (NH₄)₂CO₃.
- Group VI (Mg²⁺, NH₄⁺): no group precipitate; confirmed by specific spot tests (e.g., Mg²⁺ with magneson reagent giving a blue lake).
Each cation then has a specific confirmatory test — e.g., Cu²⁺ gives a deep blue colour with excess NH₄OH (tetraamminecopper(II)), Fe³⁺ gives blood-red with KSCN (ferric thiocyanate).
Anion tests are typically run directly on the solid salt or its solution, grouped by the reagent that detects them: dilute H₂SO₄ liberates a characteristic gas from carbonates (CO₂, turns lime water milky), sulfides (H₂S, rotten-egg smell), and nitrites (brown fumes); AgNO₃ gives characteristic-coloured precipitates with halides (AgCl white, AgBr pale yellow, AgI yellow) that differ in their solubility in ammonia; BaCl₂ precipitates sulfate as white BaSO₄, insoluble in dilute acids (this insolubility-in-acid detail is what distinguishes it from a sulfite or carbonate precipitate).
Organic Qualitative Analysis
Detection of extra elements (N, S, halogens) uses Lassaigne's test: fusing the organic compound with sodium metal converts these elements to soluble ionic sodium salts (NaCN, Na₂S, NaX), which are then tested in the aqueous extract — Prussian blue formation (with FeSO₄ then FeCl₃, under acidic conditions) confirms nitrogen via cyanide; a violet colour with sodium nitroprusside confirms sulfur; a curdy precipitate with AgNO₃ (white/pale-yellow/yellow, as above) confirms halogens. When both N and S are present together, they can combine as NaSCN in the fusion, giving a blood-red FeSCN³⁺ colour instead of Prussian blue — a well-known false-negative trap for nitrogen unless this interference is anticipated (fusing with excess sodium, or adding excess FeSO4, is used to avoid it).
Functional groups have their own named spot tests: Tollens' reagent (ammoniacal AgNO₃) gives a silver mirror with aldehydes, not ketones; Fehling's solution gives a brick-red Cu₂O precipitate with aliphatic aldehydes, not aromatic ones; the iodoform test (I₂/NaOH) gives a yellow precipitate for compounds with a CH₃–CO– group or ethanol; neutral FeCl₃ gives a violet/purple colour with phenols.
Common Mistakes / Exam Traps
- Confusing the indicator range with the equivalence point itself — the indicator must change colour at the equivalence pH, not at pH 7 by default.
- Assuming any white precipitate with AgNO₃ is AgCl — colour and behaviour in ammonia (soluble/partially soluble/insoluble) are what actually distinguish Cl⁻, Br⁻, and I⁻.
- Forgetting that Lassaigne's test can silently fail to detect nitrogen when sulfur is also present (NaSCN forms instead of NaCN) unless excess sodium/FeSO₄ is used.
- Mixing up Tollens' and Fehling's selectivity: Fehling's does not oxidise aromatic aldehydes (e.g. benzaldehyde) even though Tollens' does — a frequently tested distinction.