Qualitative Inorganic Analysis

01 / UNDERSTAND & APPLY

1. Identification is a chain of evidence

Qualitative analysis asks which species are present. Quantitative analysis asks how much is present. A coloured solid, gas evolution or precipitate can narrow the possibilities, but a single observation rarely identifies every component of an unknown mixture. The useful question is which competing explanations survive each observation.

A systematic scheme first separates groups whose reactions overlap, then confirms individual ions within a group. Keep track of both the solid and the solution above it. The supernatant is not chemically empty simply because a precipitate has appeared: it contains ions that did not precipitate, soluble complexes and excess reagent.

A preliminary colour is a clue rather than proof. A blue aqueous sample can suggest copper(II), while iron(II) salts are often pale green and iron(III) solutions may appear yellow or brown depending on hydrolysis and ligands. Concentration, mixtures and oxidation during storage can change the appearance. A colourless sample does not exclude metal ions.

The sample history matters. A reagent containing sodium cannot be introduced and then used as evidence that sodium was present originally. Separate fresh portions for tests that would contaminate later observations. A blank contains the test reagents without the unknown; a positive control shows that a test can respond to a known relevant species.

“No visible precipitate” means no precipitate was detected under those conditions. It does not prove absolute absence of an ion. The concentration may be below the threshold, the pH may be unsuitable, or complex formation may keep the metal dissolved. Conclusions should match the sensitivity and selectivity of the evidence.

The reactions here are conceptual teaching cases for supervised laboratory study. Unknown salts, heavy-metal reagents and potentially toxic gases require institutional procedures and waste controls. Gas identity is inferred from specified analytical observations, never by deliberately inhaling an unknown gas.

02 / UNDERSTAND & APPLY

2. Solubility products explain why a separation works

For a sparingly soluble salt MX, the ion product is Q=[M⁺][X⁻] in the ideal dilute model. Precipitation becomes thermodynamically favourable when Q exceeds Ksp. At the first appearance of solid, the threshold anion concentration is approximately Ksp divided by the free-metal concentration before appreciable precipitation.

For MX₂, the threshold instead follows Ksp=[M²⁺][X⁻]². The square matters. Comparing two numerical Ksp values without their stoichiometric expressions can rank precipitation incorrectly. Activities replace concentrations in a more exact treatment; numerical exercises must state the dilute ideal approximation.

Once one salt is precipitating, its dissolved free-metal concentration is constrained by Ksp and the current free precipitant concentration. At the threshold for a second salt, calculate how much of the first metal remains. “First to precipitate” does not automatically mean “quantitatively separated”.

Adding a common ion can lower solubility, but excess reagent can also produce soluble complexes. Silver chloride, for example, is affected by sufficiently high chloride through chlorido-complex formation. A simple Ksp-only calculation assumes that such complexes are negligible over the stated concentration range.

Solids can carry down other species through adsorption or trapped solution. Washing reduces soluble contamination, but a wash that changes pH or ligand concentration can partly dissolve the desired precipitate. The analytical aim is a chemically justified separation, rather than treating every solid as perfectly pure.

03 / UNDERSTAND & APPLY

3. pH controls the availability of sulfide and hydroxide

Traditional cation-group labels differ among laboratory schemes. Learn the chemical basis and the specified membership rather than assuming a group number is universal. Insoluble chlorides may be removed first; later separations can use sulfide, hydroxide or carbonate under different acid–base conditions.

For a diprotic sulfide system, H₂S ⇌ H⁺+HS⁻ and HS⁻ ⇌ H⁺+S²⁻. Multiplying the two equilibrium expressions gives [S²⁻]=K₁K₂[H₂S]/[H⁺]² in the ideal model. Acidity strongly suppresses free sulfide. At maintained dissolved H₂S, raising pH by one unit increases free S²⁻ by a factor of one hundred.

Consequently, extremely insoluble metal sulfides can precipitate in acidic medium while less insoluble sulfides require less acidic conditions. This is an equilibrium explanation, not an instruction to generate hydrogen sulfide. The fixed-H₂S approximation must be stated; changing pH in a closed total-sulfide system also changes the distribution of sulfide species.

Hydroxide precipitation likewise depends on pH. For M(OH)₃, Q=[M³⁺][OH⁻]³, so a modest change in hydroxide can produce a large change in Q. An ammonium/ammonia buffer can limit hydroxide concentration and help separate hydroxides with different precipitation thresholds.

Ammonium salts suppress ammonia ionization through the common-ion effect, but ammonia also acts as a ligand. Therefore a metal’s behaviour in aqueous ammonia may reflect both hydroxide precipitation and ammine complexation. A buffer formula alone cannot predict every metal’s fate without checking both equilibria.

04 / UNDERSTAND & APPLY

4. Chloride precipitates can be distinguished by subsequent chemistry

AgCl, PbCl₂ and Hg₂Cl₂ are familiar sparingly soluble chlorides in classical schemes. Their shared precipitation response is a group clue, not an individual identification. Lead(II) chloride is appreciably more soluble in hot water than in cold water, which supports a temperature-based separation from the other two in the usual teaching scheme.

Silver chloride dissolves in aqueous ammonia through formation of [Ag(NH₃)₂]⁺. The equilibrium combines dissolution and complex formation: AgCl(s)+2NH₃ ⇌ [Ag(NH₃)₂]⁺+Cl⁻. Acidifying this solution protonates ammonia, lowers the free ligand concentration and can restore AgCl precipitation.

Mercury(I) chloride has the dimeric cation Hg₂²⁺, not two independent Hg⁺ ions in the formula explanation. With ammonia it gives a dark mixture including finely divided mercury and mercury(II) amidochloride. That distinguishes it conceptually from the transparent dissolution of AgCl; mercury-containing tests require specialist waste handling.

Lead(II) can be confirmed by another appropriate insoluble salt after separation. Yellow PbI₂ is a familiar example. A yellow precipitate alone is not a unique lead test in an arbitrary mixture, because other substances may be yellow. The preceding separation establishes which alternatives remain plausible.

For halide identification using silver salts, AgCl is white, AgBr cream or pale yellow, and AgI yellow. Their ammonia responses differ: AgCl dissolves readily in dilute ammonia, AgBr requires more concentrated ammonia in the standard comparison, and AgI does not appreciably dissolve under the usual test conditions. Colours overlap and should be combined with chemistry.

05 / UNDERSTAND & APPLY

5. Amphoterism and ligand exchange separate similar-looking solids

Aluminium(III) and zinc(II) can form white hydroxides that dissolve in excess strong base as hydroxo complexes. Representative balances are Al(OH)₃+OH⁻ ⇌ [Al(OH)₄]⁻ and Zn(OH)₂+2OH⁻ ⇌ [Zn(OH)₄]²⁻. This amphoteric behaviour distinguishes them from hydroxides that remain insoluble in the same comparison.

Zinc hydroxide can also dissolve in excess ammonia through ammine formation, whereas aluminium hydroxide does not behave identically in the usual test. Thus “white hydroxide soluble in base” alone does not separate Al³⁺ from Zn²⁺; the response to a second ligand supplies independent evidence.

Copper(II) with a small amount of aqueous ammonia may first form a pale-blue hydroxide precipitate. Excess ammonia can give a deep-blue ammine-containing solution, commonly represented as [Cu(NH₃)₄(H₂O)₂]²⁺. The observed sequence reflects the competition between basicity and coordination, not a change of copper from +2 to +4.

Iron(III) gives a red-brown hydroxide in the usual aqueous comparison and a red thiocyanato-complex response under appropriate conditions. The simple formula [FeSCN]²⁺ is useful for an equilibrium model, but real solution speciation depends on concentrations and other ligands. Iron(II) hydroxide is initially greenish and can oxidize in air, so delayed brown coloration is not proof that the original sample contained only Fe³⁺.

Nickel(II) can give a characteristic red precipitate with dimethylglyoxime in an appropriate medium. Calcium can form calcium oxalate, and barium gives very insoluble barium sulfate. These are examples of confirmatory chemical selectivity after interference control; a reagent name by itself is not a guarantee of specificity.

Always state oxidation state and formula together. Chromium(III) hydroxide and chromate(VI) are different chemical species, and converting between them involves redox chemistry in addition to acid–base behaviour. Likewise, Fe²⁺ and Fe³⁺ cannot be interchanged in a precipitation equation merely because both contain iron.

06 / UNDERSTAND & APPLY

6. Anion tests require compatible reagents and separate portions

Carbonate reacts with acid according to CO₃²⁻+2H⁺→CO₂+H₂O; hydrogen carbonate requires one proton per ion. Both can therefore give carbon dioxide, so effervescence alone does not distinguish them. In a specified confirmatory test, CO₂ can produce CaCO₃ turbidity in limewater; excess CO₂ can convert carbonate to soluble hydrogen carbonate.

Sulfite also reacts with acid, but its sulfur dioxide product has reducing chemistry. Thus “gas with acid” is not uniquely a carbonate test. Sulfide can form toxic H₂S, and nitrite may generate nitrogen-oxide mixtures under acidic conditions. These possibilities make unsupervised sniffing or acidifying an arbitrary unknown inappropriate.

Sulfate is commonly supported by a white BaSO₄ precipitate that persists in an appropriately acidified medium. Carbonate and sulfite can also form barium precipitates but are removed or distinguished by their acid reactions. The acid must not itself introduce sulfate, so sulfuric acid would invalidate a sulfate test.

For chloride, bromide and iodide, silver-ion precipitation is interpreted in a suitable nitric-acid medium. Hydrochloric acid would introduce chloride and create a false positive. Acidification can reduce interference from carbonate and hydroxide, while halide-specific colour and ammonia response provide additional discrimination.

The classical nitrate brown-ring observation involves reduction of nitrate to NO and formation of a nitrosyl iron complex. The redox core can be balanced NO₃⁻+3Fe²⁺+4H⁺→NO+3Fe³⁺+2H₂O. The ring-forming complex is not simply a brown precipitate of iron hydroxide, and nitrite can interfere. This is a reaction-interpretation example rather than an acid-layering protocol.

Phosphate can form a yellow phosphomolybdate response in an appropriate acidic reagent system. Oxalate gives characteristic precipitation and oxidation chemistry, while acetate participates in acid–base and coordination reactions. Tests must be interpreted with the specified mixture and medium; smell descriptions in old manuals are not a recommended detection method.

Anions and cations should not be inferred as a unique original salt pairing merely because both are detected in solution. A mixture containing NaCl and KNO₃ has the same dissolved-ion inventory as a suitable mixture containing NaNO₃ and KCl. Dissolution usually removes that pairing information.

07 / UNDERSTAND & APPLY

7. Flame structure, heating and emission answer different questions

A burner flame results from fuel oxidation. With insufficient premixed air, a luminous yellow flame contains glowing carbon-rich particles and tends to deposit soot. With suitable air mixing, a nonluminous blue flame provides cleaner heating. The yellow luminosity of such a fuel flame is not automatically a sodium flame-test signal.

In a simplified blue-flame cross-section, the inner cone contains relatively unburnt gas and an adjacent primary combustion region. The outer region has greater access to oxygen and more complete combustion. The commonly used hot region lies just above the tip of the inner blue cone; the exact temperature field depends on burner design, gas composition and airflow.

An oxygen-poor region can favour reduction of suitable substances, while an oxygen-rich outer region can favour oxidation. “Reducing flame” and “oxidizing flame” describe the chemical environment, not two universal fixed temperatures. A flame diagram is a spatial aid rather than a map valid to millimetre precision for every burner.

A water bath gives gentler, more uniform heating for an appropriate sample and is limited by the bath liquid’s boiling point at the prevailing pressure. Direct flame heating can create local hot spots. Neither method should be used with sealed vessels or incompatible flammable materials; actual apparatus selection belongs to the supervised procedure.

Thermal observations include loss of water of crystallization, melting, sublimation and decomposition. These are different processes. A solid disappearing from the hot region and redepositing on a cooler surface may be subliming or undergoing reversible dissociation; mass loss alone does not establish that the only escaping substance is water.

Copper(II) sulfate pentahydrate can lose hydration water and change colour; a carbonate may instead release CO₂ on decomposition. Heating a nitrate may involve oxygen or nitrogen oxides depending on the cation and conditions. Before inferring a hydrate number, explicitly exclude concurrent decomposition and loss of other volatile substances.

08 / UNDERSTAND & APPLY

8. A flame colour is an emission pattern, with limitations

Heating can produce gaseous atoms and other emitting species from a salt. Excited states emit photons as they relax to lower energy states. The characteristic wavelengths depend on the emitting species’ energy-level differences. The visible colour is a combined response to multiple lines or bands and their intensities, not necessarily one isolated transition.

Familiar school-level observations include sodium yellow, potassium lilac, lithium crimson, calcium brick-red, strontium crimson and barium apple-green. Copper compounds often give blue-green or green shades depending on the species and conditions. These verbal labels are approximate and should not be used as precise wavelength measurements.

Sodium’s strong yellow emission can mask weaker potassium colour. Cobalt-blue glass can reduce the yellow background in the traditional potassium observation, but it is a filter rather than a chemical separation. A spectroscope that resolves characteristic lines offers stronger discrimination than an unaided colour name.

Not every cation gives a useful visible flame colour under ordinary conditions. Aluminium and magnesium, for example, do not have a reliable characteristic salt-flame colour in the standard comparison. Bright combustion of magnesium metal is a different phenomenon and should not be confused with a diagnostic magnesium-salt flame test.

A clean blank and an uncontaminated sample are essential. Sodium is common in the environment, so a weak yellow observation must be compared with the background. A colour shared approximately by two candidates needs another test. Colour blindness, ambient lighting and the chemical form of the sample further limit visual identification.

09 / UNDERSTAND & APPLY

9. Double salts crystallize with fixed composition

Mohr’s salt is ferrous ammonium sulfate hexahydrate, (NH₄)₂Fe(SO₄)₂·6H₂O, also written FeSO₄·(NH₄)₂SO₄·6H₂O. It contains Fe²⁺, two ammonium ions, two sulfate ions and six waters per formula unit. The charges sum to zero: +2+2−4=0.

Its preparation principle is crystallization from a compatible solution containing ferrous sulfate and ammonium sulfate in a 1:1 formula-unit ratio. Acidic conditions limit hydrolysis, and avoiding oxidation of Fe²⁺ is important for the intended product. Acidification does not make Fe²⁺ indefinitely immune to oxidation.

When ferrous sulfate heptahydrate is the starting solid, the composition balance is FeSO₄·7H₂O+(NH₄)₂SO₄→(NH₄)₂Fe(SO₄)₂·6H₂O+H₂O. Water is part of the balance and solvent pool; the product does not inherit seven waters merely because the starting iron salt was a heptahydrate.

Potash alum is KAl(SO₄)₂·12H₂O. A doubled formula, K₂SO₄·Al₂(SO₄)₃·24H₂O, represents two of these formula units. Confusing the two conventions doubles the assumed molar mass or changes the apparent water count. State which formula unit is being counted.

Combining one mole K₂SO₄ with one mole Al₂(SO₄)₃ supplies enough ions for two moles KAl(SO₄)₂·12H₂O, provided water is available and crystallization is ideal. The ion-assembly equation is K⁺+Al³⁺+2SO₄²⁻+12H₂O→KAl(SO₄)₂·12H₂O(s).

These are double salts: on ordinary dissolution their constituent ions become available for the usual ion tests, subject to hydration and aqueous equilibria. Contrast a coordination salt whose complex ion can remain an identifiable unit in solution. Calling alum a double salt does not mean that bare, unhydrated Al³⁺ floats in water.

10 / UNDERSTAND & APPLY

10. Preparation succeeds only when identity and recovery both check out

Crystallization separates a solid whose solubility is lower under the final conditions. A conceptual preparation therefore includes forming the correct ion inventory, removing insoluble contamination, allowing the desired hydrate to crystallize and separating crystals from the mother liquor. The mother liquor still contains dissolved product and soluble impurities.

Cooling can improve recovery when solubility falls with temperature, but it cannot guarantee complete recovery. Excessive solvent leaves more product dissolved. Evaporating everything to dryness may also leave impurities and alter hydration, so “maximum recovered mass” is not equivalent to “pure desired crystals”.

Wet crystals carry mother liquor and give an artificially high apparent yield. Overheating can instead remove crystal water or decompose the salt, lowering the mass and changing identity. A yield above 100% is therefore a prompt to investigate contamination, retained solvent, incorrect formula mass or a faulty limiting-reagent calculation.

For an alternative alum synthesis from aluminium metal, the metal first enters an oxidized soluble aluminium species; acid–base changes then permit alum crystallization with potassium and sulfate. The oxidation state changes from Al(0) to Al(III), whereas combining pre-existing aluminium sulfate with potassium sulfate is not itself an aluminium redox step.

Characterization should combine composition and chemical evidence. Sulfate precipitation and potassium emission support two components of potash alum, but they do not by themselves prove the exact hydrate number. A mass-loss experiment can support hydration only if the endpoint and absence of other decomposition have been established.

For a pure hydrate with a verified anhydrous residue, n=moles of lost water divided by moles of anhydrous formula units. Round to an integer only when experimental uncertainty supports it. A noninteger result can indicate incomplete dehydration, mixed hydrates, retained liquid or decomposition rather than a new exotic formula.

11 / UNDERSTAND & APPLY

11. Original solved practice: defend the conclusion

For each problem, identify which information is an observation and which is a model assumption. A correct numerical answer with an unjustified identification is not a complete analytical argument.

12 / UNDERSTAND & APPLY

12. Final audit: separate chemical evidence from assumptions

Before accepting an ion assignment, check that the reagent did not introduce that ion, the pH and ligand conditions were appropriate, interfering species were removed or excluded, and the observation was strong enough to detect. A positive control establishes responsiveness; a blank exposes reagent contamination.

Before accepting a separation, calculate thresholds using the correct stoichiometric exponents and free-ion concentrations. Then test the amount of the first ion remaining when the next precipitate begins. Complex formation, acid–base speciation and volume change may be the reason a memorized Ksp shortcut fails.

Before accepting a prepared solid, state the exact hydrate formula, count formula units consistently and identify the limiting reagent. Compare isolated mass with theoretical yield and check whether the solid’s chemical responses and water content support the intended identity.

Before interpreting heating, distinguish emitted light from chemical decomposition and distinguish hydration water from other volatile products. Flame colours are useful clues within a controlled test; burner zones explain heating conditions, and spectral lines explain atomic energy changes. These are connected ideas with different evidential roles.