Analytical Chemistry

Analytical chemistry is the detective work of the subject: qualitative tests ask what is present, separations ask how can it be isolated, and titrations ask how much. It carries about 3 questions, usually as a colour, precipitate, gas, separation choice or a short titration calculation. The winning habit is to treat every observation as evidence with a chemical reason — a white precipitate is only useful when you also know where it dissolves, what gas came before it and which rival ion could imitate it.

22.1 Tests for acid and basic radicals; unsaturation and functional groups; distinction tests; Lassaigne's test; fats, proteins and carbohydrates

Acid radicals (anions): let the gas speak first

Qualitative salt analysis begins with the sample's appearance and solubility, but the systematic anion hunt starts with dilute H₂SO₄, warm if needed:

Observation with dil. H₂SO₄	Ion indicated	Confirmation
Brisk effervescence; gas turns lime water milky, excess gas clears it (soluble bicarbonate)	CO₃²⁻ / HCO₃⁻	CO₂ has no smell and no effect on dichromate paper
Choking gas of burning sulfur; turns acidified K₂Cr₂O₇ paper green	SO₃²⁻	SO₂ also decolorizes dilute KMnO₄
Rotten-egg gas; lead-ethanoate paper turns black (PbS)	S²⁻	Sodium nitroprusside on alkaline extract → violet
Colourless gas browning at the mouth of the tube (NO → NO₂ in air); brown fumes	NO₂⁻	FeSO₄ gives a dark brown/black ring even with dilute acid
Vinegar smell, stronger on warming	CH₃CO₂⁻	Neutral FeCl₃ gives a deep-red colour; esterification gives fruity ethyl ethanoate
No gas	Try conc. H₂SO₄ tests and specific tests below	—

Concentrated H₂SO₄ (handle as the dehydrating oxidant of Ch. 13) interrogates a second tier: chlorides fume HCl (white smoke with an NH₃ rod; with MnO₂, greenish Cl₂); bromides give HBr partly oxidized to orange Br₂ vapour; iodides give violet I₂ vapour; nitrates with Cu turnings give brown NO₂; oxalates evolve CO + CO₂ without charring; fluorides generate HF, which etches glass. These are old but logically ordered tests — the strong acid first expels a weaker/more volatile acid, then may oxidize it.

Independent confirmations settle the common ions:

• Sulfate: BaCl₂ to an acidified solution gives white BaSO₄ that stays insoluble in dilute HCl — acid solubility separates it from BaCO₃/BaSO₃ impostors.
• Nitrate (brown ring): fresh FeSO₄ solution, then conc. H₂SO₄ poured down the wall; a brown [Fe(H₂O)₅(NO)]²⁺ ring forms at the junction. Do not shake the layers away.
• Halides: AgNO₃ after acidifying with dilute HNO₃ — AgCl white and soluble in dilute NH₄OH, AgBr pale yellow and only sparingly soluble in concentrated NH₄OH, AgI yellow and insoluble. The organic-layer test agrees: chlorine water liberates Br₂ (orange) or I₂ (violet) into CHCl₃/CCl₄.
• Phosphate: ammonium molybdate in conc. HNO₃, warmed, gives a canary-yellow precipitate of ammonium phosphomolybdate.
• Chromate: yellow solution; BaCl₂ gives yellow BaCrO₄ soluble in strong acid; acidifying shifts CrO₄²⁻ (yellow) toward Cr₂O₇²⁻ (orange), Ch. 7 in colour.

Basic radicals (cations): separation by controlled precipitation

The classical scheme precipitates groups in order, each reagent exploiting a solubility cliff. Ammonium is hunted first (Group 0) because later reagents add it: warm the original solid with NaOH — NH₃ turns moist red litmus blue and gives a brown precipitate with Nessler's reagent.

Group reagent	Ions precipitated	Discriminating observations
I: dilute HCl	Ag⁺, Hg₂²⁺, Pb²⁺ as white chlorides	PbCl₂ dissolves in hot water; AgCl dissolves in NH₄OH; Hg₂Cl₂ turns black with NH₄OH
II: H₂S in dilute HCl	Cu²⁺/Pb²⁺/Hg²⁺ (black sulfides), Cd²⁺ (yellow CdS), As³⁺ (yellow), Sb³⁺ (orange), Sn²⁺ (brown)	Acid suppresses S²⁻ (common-ion logic), so only the least soluble sulfides fall
III: NH₄OH + NH₄Cl	Fe³⁺, Al³⁺, Cr³⁺ as hydroxides	Fe(OH)₃ reddish-brown, Al(OH)₃ white gelatinous (amphoteric), Cr(OH)₃ green
IV: H₂S in NH₄OH	Zn²⁺, Mn²⁺, Ni²⁺, Co²⁺ as sulfides	ZnS white, MnS flesh/pink, NiS and CoS black
V: (NH₄)₂CO₃ + NH₄Cl/NH₄OH	Ba²⁺, Sr²⁺, Ca²⁺ as white carbonates	Flame: Ba apple-green, Sr crimson, Ca brick-red
VI: filtrate	Mg²⁺, Na⁺, K⁺	Mg²⁺ white crystalline MgNH₄PO₄ with Na₂HPO₄/NH₄OH; flames settle Na (golden) and K (lilac through cobalt glass)

Spot tests finish what groups begin: Fe³⁺ with KSCN gives an intense blood-red complex; Fe³⁺ with K₄[Fe(CN)₆] gives Prussian blue; Cu²⁺ turns deep blue in excess NH₄OH after a pale-blue Cu(OH)₂ stage (Ch. 14); Ni²⁺ with dimethylglyoxime in faintly ammoniacal solution gives a rosy-red chelate precipitate; Ca²⁺ gives white calcium oxalate with ammonium oxalate. Modern laboratories generate H₂S in situ from thioacetamide or avoid the classical scheme altogether, but its solubility logic remains superb equilibrium teaching.

Organic qualitative analysis: unsaturation and functional groups

Unsaturation. Shaking with bromine in CCl₄ discharges the orange colour when a C=C/C≡C adds Br₂; cold dilute alkaline KMnO₄ (Baeyer's reagent) loses its purple colour to brown MnO₂. Two cautions earn marks: benzene, with DU = 4, passes neither test (Ch. 17); and phenol/aniline decolorize bromine water by electrophilic substitution, throwing a white tribromo precipitate — evidence about the ring, not about an alkene.

Suspected group	First-line test(s)	Positive sign
Carboxylic acid	NaHCO₃ solution	Brisk CO₂ effervescence (phenol fails)
Phenol	Neutral FeCl₃; bromine water	Violet colour; white 2,4,6-tribromophenol
Alcohol	Na metal (dry sample); Lucas reagent; oxidation behaviour	H₂ evolved; 3° > 2° > 1° cloudiness in Lucas; see Ch. 18
Aldehyde	Tollens; Fehling; Schiff	Silver mirror; red Cu₂O (aliphatic); magenta restored
Ketone	2,4-DNP positive but Tollens negative; NaHSO₃ for methyl ketones	Orange hydrazone without a mirror
CH₃CO– or CH₃CH(OH)–	I₂/NaOH, warm	Yellow CHI₃, antiseptic smell
Terminal alkyne	ammoniacal AgNO₃; ammoniacal Cu₂Cl₂	White Ag acetylide; red Cu(I) acetylide
Primary amine	CHCl₃/alcoholic KOH; HNO₂	Foul isocyanide; aliphatic N₂↑, aromatic diazonium coupling
Amide	Boil with NaOH	NH₃ evolved (damp red litmus → blue)
Ester	Hydroxamic-acid test (NH₂OH, then FeCl₃)	Wine-red/violet ferric hydroxamate; fruity smell supports
Alkyl halide	Warm with alcoholic AgNO₃; sodium fusion then AgNO₃	AgX precipitate (reactive halides faster)

High-yield distinctions are one-reagent decisions: ethanol vs methanol by iodoform (ethanol +); propanal vs propanone by Tollens; pentan-2-one vs pentan-3-one by iodoform (only the methyl ketone); formaldehyde vs acetaldehyde by iodoform (acetaldehyde +); benzaldehyde vs acetophenone by Tollens; formic vs acetic acid by Tollens/KMnO₄ (formic reduces); phenol vs benzoic acid by NaHCO₃ (acid fizzes); ethylamine vs aniline by cold HNO₂ then β-naphthol (aniline couples to a dye; ethylamine just releases N₂); glucose vs sucrose by Benedict before and after acid hydrolysis.

Lassaigne's test: make covalent elements ionic

Nitrogen, sulfur and halogens locked in covalent molecules do not answer ionic reagents. Sodium fusion solves the problem violently and elegantly: a small piece of sodium is melted in an ignition tube, the organic compound added and heated to redness, and the hot tube plunged into distilled water. Boiling and filtering gives the alkaline sodium fusion extract, in which the elements now sit as ions:

Na + C + N → NaCN; 2Na + S → Na₂S; Na + X → NaX; and if N and S coexist with insufficient carbon/sodium balance, NaSCN may form instead.

• Nitrogen: add fresh FeSO₄, boil, cool, then FeCl₃ and acidify. A Prussian-blue colour/precipitate, Fe₄[Fe(CN)₆]₃, confirms N. (Fe²⁺ first traps CN⁻ as ferrocyanide; Fe³⁺ then precipitates it.)
• Sulfur: sodium nitroprusside gives an intense violet thionitroprusside colour, or lead ethanoate in acetic acid gives black PbS.
• Nitrogen + sulfur together: FeCl₃ alone gives a blood-red thiocyanate colour — and the Prussian-blue test then fails unless excess sodium was used in fusion.
• Halogens: acidify a separate portion with dilute HNO₃ and boil to expel HCN/H₂S if N or S is present (otherwise AgCN/Ag₂S contaminate the result), cool and add AgNO₃. The AgX colours/solubilities match the inorganic table above. Fluorine escapes this test because AgF is soluble.

The test fails when the compound is too volatile, when nitrogen is present without carbon (no CN⁻ can form, as in hydrazine), or when N₂ is lost from diazo-type compounds. Carbon and hydrogen are detected separately by combustion with CuO: CO₂ turns lime water milky and water turns white anhydrous CuSO₄ blue.

Tests for carbohydrates, proteins and fats

Test	Detects	Positive result / chemistry
Molisch (α-naphthol, then conc. H₂SO₄ down the wall)	Essentially all carbohydrates	Violet ring at the junction (furfural derivatives condense with α-naphthol)
Fehling/Benedict, heat	Reducing sugars (free anomeric carbon: glucose, fructose, maltose, lactose)	Blue Cu²⁺ → brick-red Cu₂O; sucrose negative until acid-hydrolysed
Barfoed (weakly acidic Cu²⁺)	Monosaccharides faster than disaccharides	Red Cu₂O within a couple of minutes
Seliwanoff (resorcinol/HCl)	Ketoses such as fructose	Rapid cherry-red colour
Iodine solution	Starch (amylose helix)	Blue-black complex; colour fades on heating, returns on cooling
Biuret (NaOH + very dilute CuSO₄)	Two or more peptide bonds (proteins/peptides)	Violet Cu²⁺–peptide complex; free amino acids essentially negative
Ninhydrin, heat	α-Amino acids (and protein N-termini)	Purple (Ruhemann) colour; proline gives yellow
Xanthoproteic (conc. HNO₃, then alkali)	Aromatic amino-acid residues	Yellow nitro products turning orange in alkali — the yellow stain when acid touches skin
Millon	Tyrosine phenolic group	Brick-red colour on warming
Grease spot; Sudan III/IV	Fats/oils	Translucent spot on paper; red-stained oil droplets
Acrolein (heat fat/glycerol with KHSO₄)	Glycerol backbone of fats	Pungent acrolein vapour (do this only as a demonstration)

22.2 Separation techniques: filtration, sublimation, evaporation, precipitation, crystallization, distillation, paper chromatography and atmolysis

Every separation buys purity by spending a difference in a physical property:

Technique	Property exploited	Choose it when…	School example
Filtration (gravity/vacuum)	Particle size vs pore size	An insoluble solid must leave a liquid	Sand from salt solution; the filtrate still needs evaporation for the salt
Sublimation	A solid with high vapour pressure below its melting point	One solid sublimes, the rest does not	NH₄Cl, iodine, camphor or naphthalene away from sand/salt; collect on a cold surface
Evaporation	Solvent volatility	The dissolved solid is heat-stable and the solvent is not wanted	NaCl from seawater (identity yes, high purity no)
Crystallization	Solubility varies with temperature/solvent	A purer solid is required; impurities stay in mother liquor	Benzoic acid from hot water; impure CuSO₄ recrystallized. Fractional crystallization separates salts of differing solubility curves
Precipitation	Selective formation of a sparingly soluble compound	One ion/component can be thrown down quantitatively	BaSO₄ gravimetry; softening water by precipitating CaCO₃
Simple distillation	Volatility/boiling point	Volatile liquid + non-volatile impurity, or boiling points far apart (roughly more than 25–30 K)	Pure water from salt water; ethanol from wash (partly)
Fractional distillation	Small boiling-point difference; repeated condense–revaporize in a column	Close-boiling miscible liquids	Benzene/toluene; petroleum fractions; ethanol–water stops at the 95.6% azeotrope
Steam distillation	Immiscible liquids boil when their vapour pressures sum to atmospheric	A heat-sensitive, water-immiscible organic must distil below 100 °C	Aniline, essential oils (eucalyptus, clove)
Vacuum distillation	Lowering external pressure lowers boiling point	High-boiling or heat-sensitive liquids	Glycerol purification; vacuum evaporation of milk/sugar liquors
Paper chromatography	Partition between water held on cellulose (stationary) and a moving solvent	Tiny, coloured or colourless mixtures must be fingerprinted	Amino acids (ninhydrin spray), plant pigments, ink dyes
Atmolysis	Graham's law: diffusion/effusion rate ∝ 1/√M through a porous barrier	Gases of different molar mass, classically isotopes	H₂ escapes 4× faster than O₂; historical UF₆ isotope enrichment

In paper chromatography the sample is spotted on a pencil baseline above the solvent level (ink would run; submerging the spot would wash it away). As solvent climbs by capillarity, components partition repeatedly and travel different distances. The retardation factor,

is characteristic only for the stated paper, solvent, temperature and saturation — which is why standards run on the same sheet convict or acquit an unknown. Colourless spots are located with UV light, iodine vapour or a spray reagent.

Atmolysis deserves its one calculation: at the same temperature and pressure, Graham's law gives rate(H₂)/rate(O₂) = √(32/2) = 4. A porous pot therefore enriches the faster, lighter gas on the far side; repeating the stage multiplies a tiny enrichment — historically the hard road to ²³⁵UF₆/²³⁸UF₆ separation, now largely replaced by gas centrifuges.

22.3 Acid–base, redox and complexometric titrations; choice of indicator

A titration delivers a standard solution (the titrant, in the burette) to a measured volume of analyte until the equivalence point — exact stoichiometric consumption. The end point is what the indicator shows; good method design makes the two coincide within a drop. A primary standard can be weighed and trusted (pure, stable, non-hygroscopic, ideally high molar mass: anhydrous Na₂CO₃, oxalic acid, potassium hydrogen phthalate, K₂Cr₂O₇). NaOH, KMnO₄ and Na₂S₂O₃ solutions are secondary standards and must be standardized. At equivalence, equivalents of titrant = equivalents of analyte (N₁V₁ = N₂V₂), or work in moles with the equation ratio — the safer modern habit.

Acid–base indicators follow the equivalence pH

Titration	Salt at equivalence	Equivalence pH	Indicator (range, colour change)
Strong acid–strong base	Neutral	7, with a huge vertical jump	Phenolphthalein (8.2–10, colourless→pink) or methyl orange (3.1–4.4, red→yellow/orange end point)
Strong acid–weak base (for example HCl vs NH₄OH)	Acidic (NH₄⁺ hydrolyses)	About 5	Methyl orange or methyl red (4.2–6.3)
Weak acid–strong base (CH₃CO₂H vs NaOH)	Basic (CH₃CO₂⁻ hydrolyses)	About 8.7	Phenolphthalein
Weak acid–weak base	Depends on Kₐ/K_b	No steep jump	No satisfactory single visual indicator

Sodium carbonate titrated with HCl shows two end points: phenolphthalein fades when CO₃²⁻ has become HCO₃⁻ (half-neutralization), and methyl orange then marks HCO₃⁻ → CO₂ + H₂O. Comparing the two volumes analyses NaOH/Na₂CO₃ mixtures — an old examination favourite built from Ch. 7's hydrolysis.

Permanganometric titrations

In acid, permanganate is a five-electron oxidant:

MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O.

Its own purple colour is the indicator: the end point is the first permanent pale pink that survives about 30 seconds of swirling. The medium must be dilute H₂SO₄ — HCl would itself be oxidized to Cl₂, and HNO₃ is an oxidant in its own right. Oxalate titrations are run at 60–70 °C and start slowly until Mn²⁺ autocatalyses them:

2KMnO₄ + 5H₂C₂O₄ + 3H₂SO₄ → K₂SO₄ + 2MnSO₄ + 10CO₂ + 8H₂O.

Fe²⁺, oxalate and H₂O₂ are standard analytes; KMnO₄ is standardized against sodium oxalate or oxalic acid. Read the upper meniscus of the dark solution, and remember that in neutral or alkaline solution permanganate stops at MnO₂ — a different electron count and a different equivalent weight.

Iodimetry versus iodometry — similar names, opposite directions

Iodimetry (direct) titrates a reducing analyte with a standard iodine solution; starch turns blue at the first excess of I₂. Vitamin C and sulfite are classic analytes. Iodometry (indirect) is far commoner: an oxidizing analyte liberates I₂ from excess KI, and that iodine is titrated with standard sodium thiosulfate,

I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻.

Copper estimation is the model: 2Cu²⁺ + 4I⁻ → 2CuI↓ + I₂; the I₂ found equals the Cu²⁺ taken, one-to-one in moles of electrons. Dichromate, chlorine/bleaching powder and dissolved oxygen (Winkler method) are handled the same way. Add starch only when the iodine colour has faded to pale straw: added early, the starch–iodine complex forms too strongly/adsorbs I₂ and the end point drags and overshoots. The end point is blue → colourless; in copper work a little thiocyanate near the end frees iodine adsorbed on CuI.

Complexometric (EDTA) titrations

EDTA (used as its disodium salt) is a hexadentate chelator that binds metal ions in a strict 1:1 ratio regardless of the metal's charge: M²⁺ + H₂Y²⁻ → MY²⁻ + 2H⁺. Because protons are released, the solution is buffered. At pH 10 (NH₃/NH₄Cl buffer) with Eriochrome Black T, Ca²⁺/Mg²⁺ first hold the indicator in a wine-red complex; EDTA strips the metals away and the free indicator turns blue — the standard total-hardness titration. At pH about 12, Mg(OH)₂ precipitates and murexide lets Ca²⁺ alone be found; Mg²⁺ follows by difference. Interfering Fe³⁺/Cu²⁺ can be masked (triethanolamine, cyanide where permitted) — masking is selectivity by competitive complexation. Results are conventionally reported as mg L⁻¹ of equivalent CaCO₃. The same chelation principle, scaled up, is water softening, detergent building and the lead-poisoning therapy of Chapters 14–15.

Formula and fact sheet

Topic	Must-know
Anion gases	CO₂ lime-water milky→clear; SO₂ dichromate green; H₂S PbS black; NO₂⁻ brown fumes
Confirmations	SO₄²⁻ BaSO₄/acid-proof; NO₃⁻ brown ring; PO₄³⁻ yellow molybdate; AgCl soluble NH₄OH, AgBr sparingly, AgI insoluble
Cation groups	HCl → H₂S/HCl → NH₄OH → H₂S/NH₄OH → (NH₄)₂CO₃ → soluble Mg/Na/K; NH₄⁺ tested first
Spot colours	Fe³⁺/SCN⁻ red; Cu²⁺/NH₃ deep blue; Ni²⁺/DMG rosy red; flames Na yellow, K lilac, Ca brick, Sr crimson, Ba green
Lassaigne	N Prussian blue; S violet nitroprusside/black PbS; N+S blood-red SCN⁻; X as AgX after HNO₃ boil
Food tests	Molisch violet ring; Benedict brick-red; starch-iodine blue-black; Biuret violet; Sudan/grease spot fats
Separations	Sublime NH₄Cl/I₂; fractional for close b.p.; steam below 100 °C; R_f = spot/front; H₂ atmolyses 4× O₂
Indicators	Methyl orange acidic end point; phenolphthalein basic; weak–weak no sharp visual end point
Redox	KMnO₄ self-indicating in H₂SO₄, 5 e⁻ in acid; iodometry: liberated I₂ vs S₂O₃²⁻, starch late
EDTA	1:1 metal binding; EBT at pH 10 wine-red → blue; hardness as mg L⁻¹ CaCO₃

Rapid-revision tables and high-yield facts

Item	Value or rule
Group II medium	H₂S in acid
Group III	NH₄Cl then NH₄OH
Phenolphthalein range	8.2–10.0
Methyl orange range	3.1–4.4
KMnO₄	self-indicator
Iodometric titrant	Na₂S₂O₃ with starch
R_f	solute distance / solvent distance
Graham's law	rate ∝ 1/√density
Titration balance	N₁V₁ = N₂V₂
Steam distillation	aniline, essential oils

Chapter summary

• Anion analysis begins with acid-evolved gases and ends with precipitate confirmations whose acid/ammonia solubilities separate look-alikes.
• Cation groups are solubility cliffs climbed in order; group reagents work because pH controls S²⁻, OH⁻ and CO₃²⁻ availability.
• Organic tests are the functional-group reactions of Chapters 17–20 read backwards from product to structure, with benzene and phenol as the standing bromine-test traps.
• Lassaigne's fusion converts covalent N, S and halogens into CN⁻, S²⁻ and X⁻ so inorganic colour chemistry can see them.
• Molisch, Benedict, iodine, Biuret, ninhydrin and grease-spot tests identify the three great food classes by their reactive groups.
• Separation choice is property choice: volatility for distillation, solubility for crystallization, partition for chromatography and √M for atmolysis.
• An indicator is chosen from the pH of the salt at equivalence, not from habit; the end point is an observation that must be engineered onto the equivalence point.
• Permanganate counts five electrons in acid and announces itself; iodine methods split into direct iodimetry and indirect iodometry with late starch.
• EDTA's 1:1 chelation plus a metal-sensitive dye turns water hardness into a clean wine-red-to-blue measurement.

Practice questions

Set A — Recall (25 questions)

1. A carbonate with dilute acid evolves a gas that turns lime waterabluebblackcgreendmilky
2. Hydrogen sulfide is confirmed when lead-ethanoate paper turnsablackbredcbluedyellow
3. The sulfate test gives a white precipitate with BaCl₂ that is insoluble in diluteaammoniabwater onlycHCldethanol
4. The brown-ring test is the classic confirmation forachloridebnitrateccarbonatedsulfide
5. Silver chloride precipitate dissolves readily in diluteaNH₄OHbHNO₃cH₂SO₄dNaCl solution
6. Silver iodide is yellow and in ammonium hydroxide it isasolublebredissolved as Ag metalcgaseousdinsoluble
7. Group I cations are precipitated by diluteaH₂SbHClcNH₄OHd(NH₄)₂CO₃
8. Fe³⁺ with potassium thiocyanate gives a colour described asaPrussian bluebcanary yellowcblood-reddapple green
9. Excess NH₄OH added to Cu²⁺ solution gives a deep-blue colour due toaCu(OH)₂b[Cu(NH₃)₄]²⁺cCu₂OdCuS
10. Dimethylglyoxime gives a rosy-red precipitate withaNi²⁺bZn²⁺cCa²⁺dNa⁺
11. A C=C bond is commonly detected by decolorization of bromine inawater onlybFehling solutionclime waterdCCl₄
12. Neutral FeCl₃ gives a violet colour withaethanolbbenzenecphenoldacetone
13. A carboxylic acid is distinguished from phenol by brisk effervescence withaNaHCO₃bNa metalcNaOHdNH₄OH
14. Tollens' reagent is reduced to a silver mirror by anaordinary ketonebethercaldehydedalkane
15. Prussian blue in Lassaigne's test confirms the elementasulfurbnitrogencchlorinedphosphorus
16. Sodium nitroprusside gives a violet colour with the Lassaigne ionaCN⁻bSCN⁻ onlycCl⁻dS²⁻
17. A violet ring in Molisch's test indicates aaproteinbfatccarbohydratednucleic acid only
18. The violet colour in Biuret's test indicatesareducing sugarbpeptide bonds (protein)cstarchdcholesterol
19. Iodine solution gives a blue-black colour withastarchbglucoseccellulose powder onlydalbumin
20. Ammonium chloride can be separated from sand byafiltration drybprecipitationcatmolysisdsublimation
21. Two miscible liquids with close boiling points are best separated byasimple evaporationbfractional distillationcfiltrationdsublimation
22. Steam distillation is chosen for a water-immiscible organic liquid that isaionicbsoluble in waterca gasdheat-sensitive at its normal b.p.
23. In paper chromatography, R_f is the ratio of distance moved by the spot to distance moved by theapaper edgebbaselinecsolvent frontdspray reagent
24. In acidified KMnO₄ titrations, the indicator isaKMnO₄ itselfbphenolphthaleincstarchdmethyl orange
25. In the EDTA hardness titration at pH 10, Eriochrome Black T changes at the end point from wine-red toayellowbcolourlesscbluedgreen

Answers and explanations — Set A

1. (d) — CaCO₃ precipitates in lime water; excess CO₂ clears it as soluble Ca(HCO₃)₂.

2. (a) — black PbS is one of qualitative analysis's least ambiguous colours.

3. (c) — BaSO₄ survives acid that would dissolve BaCO₃/BaSO₃.

4. (b) — the brown nitrosyl–iron complex forms only at the acid junction.

5. (a) — [Ag(NH₃)₂]⁺ complex formation dissolves AgCl.

6. (d) — AgI's lattice/covalency defeats even concentrated ammonia.

7. (b) — AgCl, Hg₂Cl₂ and PbCl₂ are the insoluble-chloride trio.

8. (c) — thiocyanatoiron(III) complexes.

9. (b) — tetraamminecopper(II) (with axial waters in solution).

10. (a) — Ni(dmg)₂, a square-planar chelate precipitate.

11. (d) — the inert solvent keeps substitution side-reactions quieter than water.

12. (c) — an iron–phenoxide charge-transfer colour.

13. (a) — RCO₂H is stronger than carbonic acid; phenol is not (Ch. 18).

14. (c) — Ag(I) is reduced while the aldehyde is oxidized to carboxylate.

15. (b) — cyanide from the fusion becomes the ferrocyanide that precipitates with Fe³⁺.

16. (d) — the thionitroprusside ion is intensely violet even at low S²⁻ levels.

17. (c) — acid-dehydrated sugar fragments condense with α-naphthol.

18. (b) — Cu²⁺ coordinates to peptide nitrogens in alkali.

19. (a) — iodine sits inside the amylose helix.

20. (d) — NH₄Cl dissociates/sublimes and recombines on the cool surface.

21. (b) — the column multiplies simple distillations.

22. (d) — combined vapour pressures let the mixture boil below 100 °C.

23. (c) — both distances are measured from the origin line.

24. (a) — the first permanent pale pink is excess MnO₄⁻ itself.

25. (c) — free Eriochrome Black T is blue at pH 10; its Mg/Ca complex is wine-red.

Set B — Understanding (15 questions)

1. Lime water turned milky by CO₂ becomes clear again in excess CO₂ because insoluble CaCO₃ converts to solubleaCaObCaCl₂cCa(HCO₃)₂dCa(OH)₂
2. The sulfate precipitate is tested with dilute HCl mainly to distinguish BaSO₄ from acid-soluble barium salts ofanitratebphosphate onlycchloridedcarbonate/sulfite
3. Pb²⁺ can be separated from Ag⁺ and Hg₂²⁺ in Group I because PbCl₂ is appreciably soluble inahot waterbcold ammoniacbenzenedconc. HNO₃ only
4. Benzene fails Baeyer's test despite DU = 4 because oxidation/addition would destroy itsaσ bondsbaromatic stabilizationchydrogen atomsdring size
5. Phenol decolorizes bromine water, yet this is not evidence of an alkene because the reaction is electrophilic substitution giving aadibromide oilbblue complexcsilver mirrordwhite tribromophenol precipitate
6. Lassaigne fusion is necessary before ionic tests because elements in organic compounds are mostly boundaas free ionsbcovalentlycas metalsdin crystals only
7. If an organic compound contains both N and S, FeCl₃ on the fusion extract may give blood-red colour due to formation ofaFeSbPrussian bluecthiocyanate complexdFe(OH)₃ only
8. Sucrose does not reduce Benedict's solution until hydrolysed because its anomeric carbons are locked in theaglycosidic bond (no free reducing end)bpeptide bondcester bonddhydrogen bond
9. A single free amino acid gives little/no Biuret colour because Biuret requires at least twoaamino groupsbpeptide bondsccarboxyl groupsdside chains
10. Recrystallization usually purifies better than simple evaporation because soluble impurities tend to remain in theamother liquorbvapourcfilter paperdcrystal lattice always
11. Vacuum distillation protects a heat-sensitive liquid by lowering its boiling point through reducedamolar massbvolumecpolaritydexternal pressure
12. In atmolysis, H₂ diffuses through a porous wall faster than O₂ by a factor of abouta2b8c4d16
13. Phenolphthalein is the right indicator for CH₃CO₂H titrated with NaOH because at equivalence the acetate solution isabasic (pH about 8.7)bstrongly acidiccneutral alwaysdbuffered at pH 4
14. Dilute HCl must not acidify a permanganate titration because MnO₄⁻ can oxidize chloride toaClO⁻ onlybMnCl₂ precipitatecHCl gas onlydCl₂ (extra oxidant consumption)
15. In iodometry, starch is added near the end point rather than at the start to avoid a sluggish end point from a very strong/adsorbedathiosulfate–starch saltbstarch–iodine complexcCuI precipitatedblue litmus change

Answers and explanations — Set B

1. (c) — CO₂ + CaCO₃ + H₂O ⇌ Ca(HCO₃)₂, the temporary-hardness equilibrium.

2. (d) — both would also precipitate with Ba²⁺ from neutral solution.

3. (a) — collect the white Group I ppt, wash with hot water, and Pb²⁺ alone enters the filtrate.

4. (b) — DU counts the ring; reagents meet a stabilized π sextet.

5. (d) — precipitate + HBr formation mark substitution, unlike clean alkene addition in CCl₄.

6. (b) — sodium converts them to NaCN/Na₂S/NaX that ordinary reagents can see.

7. (c) — SCN⁻ + Fe³⁺ gives red thiocyanato complexes, masking the separate N test.

8. (a) — glucose C1 and fructose C2 are both engaged in the sucrose link.

9. (b) — the chromophore is a Cu²⁺ complex of consecutive peptide nitrogens.

10. (a) — evaporation, by contrast, deposits everything that was dissolved.

11. (d) — boiling begins when vapour pressure equals the (now lower) external pressure.

12. (c) — √(32/2) = √16 = 4.

13. (a) — acetate hydrolysis puts the steep part of the curve inside phenolphthalein's range.

14. (d) — permanganate would be spent on the acid medium, inflating the titre.

15. (b) — with much I₂ present the complex releases iodine too slowly for a sharp disappearance.

Set C — Application (10 questions)

1. A salt with dilute H₂SO₄ gives a choking gas that turns acidified dichromate paper green. The anion isacarbonatebsulfitecsulfidednitrite
2. A white Group I precipitate turns black on adding NH₄OH. The cation isaAg⁺bPb²⁺cHg₂²⁺dBa²⁺
3. Which C₅H₁₀O ketone gives a positive iodoform test?apentan-2-onebpentan-3-oneccyclopentanoldpentanal
4. A Lassaigne extract, boiled with dilute HNO₃ and treated with AgNO₃, gives a pale-yellow precipitate only sparingly soluble in conc. NH₄OH. The halogen isafluorinebiodinecchlorinedbromine
5. A food extract gives a violet Biuret test and a yellow xanthoproteic test. It contains mainlyastarchbfatcproteindglucose
6. Iodine crystals mixed with sand are best separated byasublimationbfiltrationcatmolysisdprecipitation
7. Benzene (b.p. 80 °C) and toluene (b.p. 111 °C) in one liquid mixture are best separated byapaper chromatographybfiltrationcsublimationdfractional distillation
8. On a chromatogram a spot moves 4.2 cm while the solvent front moves 7.0 cm. Its R_f isa0.42b0.60c1.67d0.70
9. For titrating aqueous NH₃ (weak base) with standard HCl, the suitable indicator isamethyl orangebphenolphthaleincstarchdEriochrome Black T
10. A 50.0 mL water sample at pH 10 needs 10.0 mL of 0.0100 mol L⁻¹ EDTA. Its hardness as mg L⁻¹ CaCO₃ isa100b200c50d400

Answers and explanations — Set C

1. (b) — SO₂ reduces orange Cr₂O₇²⁻ to green Cr³⁺ (Ch. 13's sulfur-dioxide signature).

2. (c) — Hg₂Cl₂ disproportionates in ammonia to black Hg mixed with white HgNH₂Cl.

3. (a) — only pentan-2-one contains the CH₃CO– unit demanded by the haloform reaction.

4. (d) — AgBr is the middle member of the solubility ladder.

5. (c) — peptide bonds plus aromatic residues: a protein answering both of its name tests.

6. (a) — iodine vaporizes and re-deposits as crystals on the cold funnel; sand stays.

7. (d) — an 31 K gap still rewards a fractionating column for clean cuts at scale or purity.

8. (b) — R_f = 4.2/7.0 = 0.60 (dimensionless, and never above 1 for the spot).

9. (a) — equivalence gives acidic NH₄Cl solution near pH 5, inside methyl orange's range.

10. (b) — moles EDTA = 0.0100×0.0100 = 1.00×10⁻⁴ = moles CaCO₃ equivalent = 10.0 mg in 50.0 mL, i.e. 200 mg L⁻¹.