To identify an unknown salt in O-Level Chemistry, use the prescribed sodium-hydroxide and aqueous-ammonia results to infer its cation, then use a separate appropriate test to infer its anion. Combine charges only after both ions are supported by evidence.
For Secondary 4 Chemistry Hougang, the 2027 G3 reference includes aluminium, ammonium, calcium, copper(II), iron(II), iron(III), zinc, and carbonate, chloride, iodide, nitrate and sulfate tests. This complete decision guide uses original salt examples, observations, ion equations and ambiguity checks rather than guessing from one precipitate.
- Infer the cation from colour and NaOH/ammonia excess behaviour
- Infer the anion from a separate correct reagent system
- Distinguish Al³⁺ from Zn²⁺ using excess ammonia
- Check the test reagent as the source of Ag⁺ or Ba²⁺
- Balance the verified cation and anion charges to get the salt formula
The first precipitate rarely tells the whole formula
A Secondary 4 student sees a white precipitate after sodium hydroxide is added to an unknown salt solution and announces that the salt is aluminium chloride. I ask which of the observations established chloride, and how zinc or calcium were ruled out. Neither was established by that one observation. A salt formula requires compatible positive and negative ions, and a defensible conclusion needs evidence for both.
This guide teaches students in Hougang how to combine the 2027 SEC G3 qualitative analysis tables to identify an unknown salt, distinguish zinc from aluminium, compare iron ions, select the correct anion result and finally write a charge-balanced formula. It does not replace either standalone cation or anion guide; it focuses on joining independent results into one chemically valid identity without guessing.
What the official K324 notes actually provide
The 2027 Singapore-Cambridge SEC G3 Chemistry syllabus K324 lists tests for Al³⁺, NH₄⁺, Ca²⁺, Cu²⁺, Fe²⁺, Fe³⁺ and Zn²⁺, and for carbonate, chloride, iodide, nitrate and sulfate anions. It also lists common gas tests. The notes are provided in the practical examination paper, but students still need to apply them to the given observations.
For 2026 GCE O-Level Pure Chemistry the subject code is 6092. G2 and Combined Science candidates should check their own practical requirements. The full 2027 SEAB K324 syllabus, Notes for Qualitative Analysis is the current reference. Candidates are not expected to carry out sulfur dioxide gas testing in Paper 3; this guide remains a written reasoning exercise rather than a home laboratory procedure.
The four decisions in an unknown-salt question
The method is to identify the sample form, infer a cation from its appropriate test, infer an anion from a separate appropriate test, and combine charges into an electrically neutral formula. Each stage must be supported by the given observations. If a result fits more than one candidate, do not jump ahead to a full formula; look for the discriminating comparison.
A good written answer can say “the results are consistent with Zn²⁺ and SO₄²⁻, so ZnSO₄ is a possible formula within the stated single-salt model.” If the sample might contain several salts, the same observations may establish ions without uniquely identifying how they were originally paired. The premise of one pure salt matters to how strong the conclusion can be.
Begin with the material the question supplies
An unknown may be a solid or an aqueous solution. Some inorganic salts are soluble; others are not. A test that requires the anion to be in aqueous solution must be interpreted with the actual sample preparation supplied by the question. A carbonate solid may release gas with dilute acid, while a soluble sulfate may be tested through a suitable solution.
The word unknown does not authorise students to invent extra reagents or dissolve every material by any means they choose. In a practical examination, follow the stated instructions and reagent restrictions. In a paper question, use the reported evidence and named test conditions rather than assume all unknown powders are identical white ionic solids.
Record the observation before naming an ion
A practical observation includes the colour of any precipitate, whether it dissolves in excess reagent and any gas or solution colour change at the specified stage. “A white precipitate forms and dissolves in excess sodium hydroxide, giving a colourless solution” conveys more than “the mixture went white”.
The chemical inference comes afterward. White initial precipitates occur for more than one cation, and the excess response helps distinguish them. If the question asks for the observation, give the visible result; if it asks to deduce the ion, connect that result with the prescribed table. A one-word ion name is not a substitute for a complete observation when the command is describe.
Why excess reagent is part of the test
For Al³⁺ and Zn²⁺, aqueous sodium hydroxide can initially give a white precipitate that dissolves in excess, producing a colourless solution under the standard notes. The initial and excess stages therefore do not distinguish those ions by NaOH alone. Aqueous ammonia offers a useful further contrast.
Al³⁺ gives a white precipitate with aqueous ammonia that remains insoluble in excess, whereas Zn²⁺ gives a white precipitate that dissolves in excess aqueous ammonia. The word excess changes the diagnostic meaning, not just the intensity of the reagent’s colour. The relevant full sequence is essential to a reliable conclusion.
The aluminium–zinc decision
An unknown cation sample produces white precipitates with both sodium hydroxide and ammonia. In excess NaOH the solid dissolves. With excess ammonia the precipitate remains. Within the official candidate set, that is consistent with aluminium(III), Al³⁺, rather than zinc(II).
Reverse only the final ammonia observation: the white precipitate dissolves in excess ammonia, leaving a colourless solution. That supports zinc(II), Zn²⁺, in the familiar set. This paired changed-case question is one of the most effective checks of understanding because all the earlier white precipitate information was the same.
Net equations for aluminium hydroxide and zinc hydroxide
The familiar initial precipitation equations are Al³⁺(aq) + 3OH⁻(aq) → Al(OH)₃(s) and Zn²⁺(aq) + 2OH⁻(aq) → Zn(OH)₂(s). The different coefficients follow directly from the ion charges. Each describes formation of the initial solid, not the detailed complex-ion processes that can occur in excess reagent.
The official G3 syllabus says formulas of complex ions are not required for this part. A student should not invent a complex-ion formula when a school answer only asks which precipitate dissolves. A correct initial ionic equation plus a precise excess-reagent observation is a useful level of chemical explanation.
Copper(II) produces a distinctive colour pattern
Aqueous copper(II) commonly gives a light blue precipitate with aqueous sodium hydroxide, insoluble in excess under the official test. With aqueous ammonia it gives an initial light blue precipitate that dissolves in excess to yield a dark blue solution. The deep colour is part of the prescribed cation distinction, not proof that the whole salt formula has been identified.
If the anion result is not yet known, Cu²⁺ could be paired with more than one suitable negative ion in different salts. A student who immediately writes CuSO₄ solely because the starting liquid looks blue has guessed the sulfate. The anion must be tested independently or provided through other reliable information.
A worked Cu²⁺ plus sulfate identity
Suppose a single-salt aqueous unknown has the characteristic Cu²⁺ ammonia and NaOH observations. A separate portion, acidified with the prescribed nitric acid and tested with aqueous barium nitrate, produces a white precipitate. That independent anion evidence supports sulfate, SO₄²⁻.
Charge balance now gives CuSO₄: one Cu²⁺ ion with one SO₄²⁻ ion. The familiar initial cation precipitate can be represented by Cu²⁺ + 2OH⁻ → Cu(OH)₂(s), while the sulfate test is Ba²⁺ + SO₄²⁻ → BaSO₄(s). Two different solids arise from two different test reagent systems. Their observations should not be merged into one imaginary beaker event.
Why iron(II) and iron(III) need separate conclusions
In the official sodium hydroxide and aqueous ammonia tables, iron(II) gives a green precipitate insoluble in excess, while iron(III) gives a red-brown precipitate insoluble in excess. Their simple hydroxide formulas are Fe(OH)₂ and Fe(OH)₃ respectively.
The Roman numeral identifies the ion charge, not a colour label assigned without chemistry. If Fe²⁺ is present, two hydroxides balance the charge; for Fe³⁺, three are needed. A student who correctly sees red-brown but writes FeCl₂ as the eventual chloride formula has forgotten that the anion stoichiometry also depends on which iron ion was identified.
An iron(III) chloride unknown
Imagine an aqueous single-salt sample gives a red-brown precipitate with appropriate hydroxide testing and a separate halide test produces a white precipitate after nitric acid and silver nitrate. The cation result supports Fe³⁺; the anion result supports Cl⁻. The corresponding neutral compound formula is FeCl₃.
Do not write FeCl₂ from an everyday familiarity with iron(II) salts. One Fe³⁺ requires three Cl⁻ ions to give total charge zero. The white anion-test precipitate is AgCl(s), while the red-brown cation-test solid is Fe(OH)₃(s). Both are evidence about the original unknown but are not themselves the unknown compound’s full formula.
An iron(II) sulfate unknown
If a suitable unknown gives the familiar green iron(II) hydroxide precipitate and a separate sulfate test gives white barium sulfate, the ion pair is Fe²⁺ and SO₄²⁻. Those charges balance 1:1, giving FeSO₄ as a possible single-salt identity.
An iron(II) solution can undergo oxidation on exposure to air under some conditions, so real observations may be complicated by sample age or preparation. The school problem normally supplies the intended tests and scope. A reliable written conclusion uses the actual green result provided and does not substitute a later brown colour without evidence that it occurred in the question.
The calcium result differs from zinc and aluminium
Ca²⁺ gives a white precipitate with aqueous sodium hydroxide that does not dissolve in excess in the official G3 notes. With aqueous ammonia, calcium gives no precipitate in the prescribed test. That pattern differs from both zinc and aluminium, which give white precipitates with the ammonia reagent.
A student who sees no precipitate with aqueous ammonia and declares no metal ion exists has missed calcium in the standard table. The correct inference uses the full comparison. If the unknown’s anion is chloride, CaCl₂ is the charge-balanced formula; if sulfate, CaSO₄ has limited solubility and the actual sample preparation must be considered.
A worked calcium chloride identification
A hypothetical single-salt solution produces a white precipitate with NaOH, insoluble in excess, and no precipitate with aqueous ammonia. It separately produces a white AgCl precipitate after appropriate nitric acid and aqueous silver nitrate. Within the stated candidate set, the findings support Ca²⁺ and Cl⁻.
The formula is CaCl₂ because the calcium ion has charge +2 and two chloride ions each contribute −1. A strong student can explain why CaCl cannot be electrically neutral and why the NaOH precipitate is Ca(OH)₂ rather than elemental calcium. The chemical formulas arise from charges, not from the number of test tubes used.
Ammonium produces a gas rather than the usual metal precipitate
The NH₄⁺ test in the official G3 cation notes involves aqueous sodium hydroxide and warming, producing ammonia gas. A suitable prescribed damp red litmus test gives a blue colour in response to the basic gas. This differs from the metal cations identified by hydroxide precipitates.
The presence of ammonia should be supported by the specific reagent sequence and gas-test observation. A student should not assume every evolved alkaline gas proves nitrate, because nitrate has a different prescribed reduction test. Ammonium and nitrate can be particularly confusing if the test conditions overlap or both ions might be present.
The nitrate and ammonium trap
Nitrate in solution is tested through a prescribed sequence involving sodium hydroxide, aluminium foil and careful warming, with ammonia produced. Ammonium can also produce ammonia when warmed with sodium hydroxide without the aluminium step. Therefore ammonia gas is not a unique nitrate identifier in an unknown already known to contain ammonium.
A good experimental interpretation must consider possible interference and the actual instructions. The school question may provide other information or a controlled separate sample that makes nitrate a reasonable inference. If all the evidence proves is that ammonia formed, do not announce nitrate without checking which reagent, original ions and conditions were involved.
The anion half of the investigation must be independent
After identifying a cation, anion testing supplies separate evidence about the negative ion. The standard K324 anion set includes carbonate, chloride, iodide, nitrate and sulfate. Several give white precipitates or gas formation, but their reagents differ. An anion name cannot be deduced from a cation hydroxide colour alone.
I ask students to keep the original cation observation and separate anion observation in different columns. The final salt formula is written only when both sides are supported. This method is especially important when the starting solution is coloured: a blue copper(II) solution might tempt the student to guess sulfate because CuSO₄ is familiar, even if the anion is actually chloride or nitrate.
Carbonate acid test and gas confirmation
A carbonate-containing sample can give effervescence with suitable dilute acid as CO₂ is produced. Under the prescribed gas test, carbon dioxide turns limewater milky through formation of calcium carbonate solid. The net ionic carbonate reaction in solution can be written CO₃²⁻ + 2H⁺ → CO₂ + H₂O.
Fizzing alone does not identify the gas uniquely. The gas test and known starting reagent support the inference. If the carbonate sample is insoluble, its cation may need a separately appropriate analysis route supplied by the question. Do not assume the solution for a cation test is prepared simply by performing all possible anion tests in the same contaminated tube.
Why an insoluble carbonate changes sample handling
Some carbonates, such as calcium carbonate, are insoluble under the familiar school classification; sodium carbonate is soluble. An unknown solid producing CO₂ with dilute acid might be a carbonate, but its metal cation is not automatically identified from the bubbles. Different carbonates can release the same gas.
A sample’s physical state matters. If a practical question describes preparing a solution for cation testing, follow those given instructions and interpret the resulting ions. The student should not treat the gas identity as proof of the metal or invent a formula from the colour of the original solid. The complete identity needs another evidence source.
Chloride and silver nitrate
The official chloride test acidifies the aqueous unknown with dilute nitric acid and then adds aqueous silver nitrate. A positive test gives a white AgCl precipitate. The net equation is Ag⁺ + Cl⁻ → AgCl(s). The acidification step helps address interference and does not introduce chloride in the way hydrochloric acid would.
A student who writes “add HCl before silver nitrate” has supplied chloride from the test reagent and undermined the inference about the original unknown. The exact acid matters. Record the reagent sequence before stating the positive result. The white precipitate belongs to silver chloride, not to the original unknown salt suddenly becoming white for no chemical reason.
Iodide and a yellow silver precipitate
The prescribed iodide test uses the same nitric-acid then aqueous silver-nitrate sequence but gives a yellow AgI precipitate. The net equation is Ag⁺ + I⁻ → AgI(s). Within the familiar school options, this colour contrast distinguishes chloride and iodide under the same reagent system.
Do not confuse iodide ion I⁻ with iodine molecule I₂. A yellow silver iodide precipitate is not the same observation as purple iodine vapour or the coloured products of halogen displacement. Reading the original ionic species and precipitate formula prevents the student from transferring an unrelated halogen reaction into a qualitative-analysis test.
Sulfate and barium nitrate
The official sulfate test acidifies the aqueous sample with dilute nitric acid and then adds aqueous barium nitrate. A white precipitate of barium sulfate indicates sulfate under the stated comparison: Ba²⁺ + SO₄²⁻ → BaSO₄(s). The barium reagent is not interchangeable with silver nitrate simply because both can produce white precipitates in their respective tests.
A student who writes “white precipitate, therefore chloride” without noting which reagent was used has insufficient evidence. The same colour can support different ions when different chemicals form the solids. To infer sulfate, name the acidified barium-nitrate test and its BaSO₄ product rather than rely on the word white alone.
Why nitric acid is the prescribed acid in those tests
Acidification helps prevent some interfering species, such as carbonate, from generating misleading precipitates in subsequent halide or sulfate tests. Nitric acid is used because nitrate is not the target precipitating halide or sulfate in those reagent systems. Adding hydrochloric acid to a chloride test could introduce exactly the species whose presence is being investigated.
The useful school lesson is that reagent choice supports the interpretation. If the question supplies a complete procedure, follow it. Do not give a different acid because both have the everyday name acid. A chemical test has meaning only when the reagent system and observed product are correctly matched.
A white precipitate does not equal one unique anion
Silver chloride and barium sulfate both appear white in the ordinary school tests. Aqueous hydroxide can also precipitate white metal hydroxides from some cations. Without a stated reagent, the observation “white solid formed” is inadequate to identify one chemical ion.
I ask students to write each white precipitate with its formula: AgCl, BaSO₄, Al(OH)₃, Zn(OH)₂ or Ca(OH)₂ as appropriate to the given scenario. Then name the reagent that supplied the other participating ion. This turns a colour memory into a specific chemical explanation and prevents confusing cation and anion tests.
The correct order in a complete unknown-salt answer
Begin with the given sample or supplied solution. Identify the cation from sodium-hydroxide and/or ammonia responses, including what happens in excess. Independently identify the anion from the prescribed reagent and observation. Only then write an electrically neutral formula from the charges. If the problem supplies a list of possible salts, check whether that formula fits the list and the stated solubility.
This order prevents premature guesses. If the cation evidence is incomplete, give possible cations rather than invent one. If the anion has been identified but the sample may contain several salts, distinguish ions present from original ion pairing. The final claim should be no stronger than the experiment’s stated design and evidence.
Charge balance gives the final formula
Suppose the tests support Al³⁺ and SO₄²⁻. To make a neutral compound, two Al³⁺ ions give +6 and three sulfate ions give −6, so the formula is Al₂(SO₄)₃. It is not AlSO₄ simply because one cation and one anion have been named.
The formula subscripts follow the charges, not the number of tests performed or colours observed. If the cation changes to Zn²⁺ while sulfate stays SO₄²⁻, the ratio becomes 1:1 and the formula is ZnSO₄. This changed-case comparison is an excellent way to test whether a student can build a salt formula rather than recalling a familiar list.
Worked unknown A: zinc sulfate
Unknown A gives a white precipitate with NaOH that dissolves in excess, and a white precipitate with aqueous ammonia that also dissolves in excess to a colourless solution. That supports Zn²⁺ in the official set. A separate appropriate sulfate test gives white BaSO₄ after nitric acid and barium nitrate.
Combining Zn²⁺ with SO₄²⁻ gives ZnSO₄ for a stated pure single salt. The hydroxide precipitate is Zn(OH)₂, and the sulfate test precipitate is BaSO₄. Neither test product is identical to the original unknown compound. The observations come from separate reactions that reveal its constituent ions.
Worked unknown B: aluminium chloride
Unknown B’s white hydroxide precipitate dissolves in excess NaOH but remains insoluble in excess aqueous ammonia. That supports Al³⁺. A separately acidified silver nitrate test gives white AgCl, supporting Cl⁻. For a single dissolved salt, the resulting formula is AlCl₃.
One Al³⁺ balances three chloride ions. The initial cation precipitate can be represented by Al³⁺ + 3OH⁻ → Al(OH)₃(s). The anion precipitate is Ag⁺ + Cl⁻ → AgCl(s). Correct identification depends on the different ammoniacal behaviour of Al³⁺ and Zn²⁺, not simply on the fact that both yield white hydroxides.
Worked unknown C: copper(II) sulfate
Unknown C gives a light blue hydroxide precipitate and, with excess aqueous ammonia, a deep blue solution. Those observations support Cu²⁺. A separate nitric-acid/barium-nitrate test gives a white BaSO₄ precipitate, supporting sulfate. If the sample is specified as one pure salt, CuSO₄ is the charge-balanced candidate.
It is not enough to recognise the solution colour and instantly write copper sulfate. Copper(II) chloride and nitrate are different possible compounds in a broader chemical context. The sulfate inference comes from the anion test, while copper(II) follows from the cation reagent behaviour. The two conclusions work together to justify the final formula.
Worked unknown D: iron(III) chloride
Unknown D’s cation testing gives a red-brown precipitate insoluble in excess NaOH and ammonia, supporting Fe³⁺. The acidified silver-nitrate anion test gives a white precipitate, supporting Cl⁻. The corresponding single salt is FeCl₃.
The relevant hydroxide formula is Fe(OH)₃ and the halide precipitate AgCl. A student who writes FeCl₂ has used the iron(II) charge despite the red-brown result supporting iron(III). The final formula should be checked by total charge: +3 from iron and three −1 chloride ions give zero.
Worked unknown E: calcium chloride
Unknown E gives a white precipitate with sodium hydroxide that remains in excess, while aqueous ammonia gives no precipitate under the official test. That is consistent with Ca²⁺. Its separately observed white silver-chloride precipitate supports chloride. The single-salt formula is CaCl₂.
The calcium cation also has charge +2, but its ammonia response differs from the white precipitates given by Zn²⁺ and Al³⁺ in that reagent. A learner who chooses zinc solely because both can give white hydroxide precipitates has ignored the discriminating ammonia observation. This is why the full reagent sequence belongs in the evidence column.
Worked unknown F: zinc iodide
Suppose a soluble single-salt unknown gives the zinc hydroxide response with both NaOH and ammonia, and a separate nitric-acid/silver-nitrate test gives a yellow precipitate. The results support Zn²⁺ and I⁻ under the familiar tables. The electrically neutral formula is ZnI₂.
The yellow precipitate is silver iodide, AgI(s), while the cation tests involve zinc hydroxide. The two iodide ions balance the +2 zinc charge. If a student writes ZnI, the net charge would be +1 rather than zero. A little charge arithmetic completes the reasoning after the qualitative observations identify the ions.
When one formula does not fit the sample state
If an unknown is described as a stable clear aqueous solution, predicting a highly insoluble salt as the only dissolved compound may contradict the sample description. Solubility rules can act as an additional plausibility check on the proposed ion pair. For example, calcium carbonate is ordinarily insoluble in water rather than a typical concentrated clear CaCO₃ solution.
This does not mean a trace amount of an insoluble material can never dissolve, or that all real samples contain only one species. It means the school inference should agree with the stated physical preparation. A candidate whose cation and anion tests appear to imply an unlikely dissolved salt should review the observations and sample assumptions before confidently writing the formula.
A single unknown may be a mixture
A practical question can involve a mixture of salts rather than one pure compound. Finding two cations or several anions then identifies constituents but does not necessarily reveal their original pairings. For example, detecting Na⁺, K⁺, Cl⁻ and NO₃⁻ in solution would not on its own prove which initial combination of salts was added.
The wording of the question matters. “Identify the ions present” can be answered with the tested ion set. “Suggest the formula of the single pure salt” allows pairing the confirmed positive and negative ions. A student who ignores the word mixture may force all observations into one chemically impossible formula. Scientific conclusions must match the evidence’s scope.
What to do when the observations conflict
If the supplied test results appear inconsistent—for example, the cation evidence supports Fe³⁺ while the proposed compound formula contains Fe²⁺—the first response is to recheck charge, reagent and observation interpretation. A real laboratory might also have contamination, oxidation or mixed substances. A paper question may include a deliberate trap requiring identification of the inconsistency.
Do not erase a reported observation simply because it does not fit the expected textbook salt. The student’s job is to explain what the evidence shows and what additional information would be needed. In a normal school problem, one compatible interpretation is intended; if none is supported, an honest qualified response is better than inventing extra ions or reagent additions.
The test with silver nitrate is not a copper test
When an aqueous unknown is acidified appropriately and tested with AgNO₃, the reagent supplies silver ions to react with anions such as chloride or iodide. It does not primarily identify whether the original cation is copper. The white or yellow silver halide is a test product whose positive ion came from the reagent.
This distinction is easy to miss because the newly formed precipitate is a salt with a metal cation. A learner should track which ions originated in the unknown and which entered with the reagent. If silver was introduced during the test, a silver compound in the product does not establish that the original sample contained silver.
The barium reagent is not proof of barium in the sample
Similarly, Ba(NO₃)₂ supplies Ba²⁺ for the sulfate test. The resulting white BaSO₄ precipitate contains barium from the test reagent and sulfate from the unknown. A student who identifies the unknown’s cation as barium because barium sulfate formed has mistaken the source of one ion.
The correct inference from that anion test is the presence of sulfate under the prescribed conditions. Separate cation tests are needed to identify the unknown’s positive ions. This simple reagent-origin question makes the whole logic of qualitative analysis much clearer: test reagents contribute their own particles to the observed products.
Why the excess ammonia test is so useful
Zinc and aluminium commonly produce initial white precipitates with aqueous NaOH and also white precipitates with aqueous ammonia. The discriminating observation appears after adding excess ammonia under the official notes: zinc’s precipitate dissolves to a colourless solution, while aluminium’s remains insoluble.
A student who sees the same white starting colour and stops has not completed the diagnostic comparison. In a written response, use the phrase in excess where the table requires it. The condition changes what can be inferred even if the reagent’s chemical name is unchanged. This is the main reason a complete stage-by-stage observation can outperform a vague colour-only answer.
Why calcium needs a separate ammonia observation
Calcium can give a white precipitate with NaOH that remains insoluble in excess, while the prescribed aqueous ammonia test gives no precipitate. This helps distinguish Ca²⁺ from certain other ions producing white hydroxides. The absence of an ammonia precipitate is therefore meaningful data, not proof that the sample was empty.
A strong student uses both positive and negative results. “White with NaOH, no precipitate with ammonia” is a two-test pattern. The word no is a legitimate observation and can rule out candidates when the conditions are appropriate. A cation table should not be read as only a list of colourful reactions.
A gas result should name its confirming observation
If ammonia is evolved under the appropriate prescribed test, it turns damp red litmus blue. If carbon dioxide is produced by carbonate–acid reaction, it forms a white precipitate in limewater. These observations have different chemicals and meanings. The student should not use the limewater test for every gas or call any alkaline gas carbon dioxide.
This is particularly important when a question supplies a gas but not its name. Ask which test actually changed, then identify the gas within the familiar candidates. The anion and gas tests guide
A cation observation can change after air exposure
Iron(II) species can be oxidised by air under suitable conditions, potentially complicating the colour of a precipitate over time. A sample that begins with a green iron(II) hydroxide precipitate may become browner as chemical oxidation proceeds. A student should record when observations occur rather than treating all colours as simultaneous original results.
The school QA reference gives the familiar initial green Fe²⁺ and red-brown Fe³⁺ patterns. If a question provides a time sequence, report that sequence and reason from it. Do not replace the stated initial observation with a later textbook colour simply because it is more familiar. Time and exposure can matter.
The compound’s charge should always sum to zero
A simple neutral ionic salt pairs positive and negative ions so total charge is zero. Cu²⁺ and Cl⁻ yield CuCl₂. Fe³⁺ and SO₄²⁻ require a 2:3 ratio, giving Fe₂(SO₄)₃. NH₄⁺ and NO₃⁻ give NH₄NO₃. The formulas are derived from ionic charges rather than the order in which the tests were performed.
A common error is writing FeSO₄ for a sample whose cation test proved Fe³⁺. That formula corresponds to Fe²⁺ paired with sulfate under the simple salt model. The correct Fe³⁺ sulfate formula requires a different ratio. Formula validation is the last check after ion inference, and it can catch an otherwise hidden mismatch between identification and final answer.
An optional worked ammonium nitrate scenario
A hypothetical single-salt sample is known from a suitable independent analysis to contain NH₄⁺ and NO₃⁻. Its formula is NH₄NO₃ because each ion has one unit of opposite charge. But if the only observations were ammonia after NaOH warming and ammonia after adding aluminium foil under alkaline conditions, that is not automatically two independent proofs.
The original ammonium can interfere with interpreting ammonia as evidence for nitrate. A robust question must provide additional suitable information or a method that resolves the interference. This is an example where chemical honesty matters more than a tidy answer key. Knowing which gas appears is not enough if another starting ion can generate it in the same test system.
No precipitate is an observation too
A solution that gives no precipitate with aqueous ammonia may still contain Ca²⁺ or other ions depending on the candidate set and test conditions. A negative result can help eliminate some candidates, but it does not prove the sample contains no ions. Soluble salts can remain transparent while carrying mobile charged particles.
Students should be able to state “No precipitate observed” and then use the rest of the reference table for inference. For example, a calcium-like sodium-hydroxide result and no ammonia precipitate together are more informative than either alone. Negative tests become useful evidence when interpreted as part of a controlled sequence.
A table with one ambiguous result
Imagine unknown X gives a white precipitate soluble in excess NaOH, and no other cation tests are reported. Under the official set, at least aluminium and zinc remain plausible. The student should not write a definitive Al³⁺ or Zn²⁺ solely from that row. If a separate ammoniacal result is then given, it can help discriminate them.
The task may ask for a further suitable test rather than a final ion identity. A good answer explains why the current evidence is insufficient and which prescribed observation would distinguish the possibilities. This is a scientific reasoning skill: the correct response is sometimes a conditional statement rather than one confident formula.
Original case one: colour and excess reagent
A fictional unknown gives a light blue precipitate with NaOH that is insoluble in excess, and a deep blue solution after excess aqueous ammonia. The question asks for the cation. The best familiar answer is Cu²⁺, with the deep-blue ammonia observation serving as distinctive supporting evidence.
If a sulfate test is subsequently positive under the prescribed reagent system, a single-salt formula can be CuSO₄. If the anion information is not supplied, do not assume sulfate merely because copper sulfate is a common blue laboratory reagent. The discipline is to keep the cation inference separate until anion evidence arrives.
Original case two: white hydroxide and halide
Unknown Y gives a white precipitate that dissolves in excess NaOH and stays insoluble in excess ammonia. In a separate nitric-acid/silver-nitrate test, it gives white AgCl. Under the intended single-salt model, Al³⁺ and Cl⁻ are supported, leading to AlCl₃.
A student who writes ZnCl₂ has missed the excess-ammonia distinction. Another who writes AlCl has not balanced ion charges. The complete answer requires both the reagent comparison and correct final formula. Merely naming aluminium correctly does not finish an unknown salt question that also requests the anion and compound.
Original case three: change only the silver precipitate
Keep the zinc-like NaOH and ammonia observations but change the anion silver-nitrate result from white to yellow under the same prescribed acidification. The cation remains Zn²⁺, while the anion inference changes from Cl⁻ to I⁻. The candidate single-salt formula becomes ZnI₂.
This is a useful transfer question because only one observation changed. A student who keeps writing ZnCl₂ has memorised the whole example as one phrase rather than connected each reagent result to a specific ion. A strong learner can explain precisely which part of the salt formula changed and why the zinc count did not.
Original case four: change the iron colour
Unknown A gives a green precipitate insoluble in excess hydroxide reagents and chloride evidence from the correct silver-nitrate test. That supports Fe²⁺ and Cl⁻, giving FeCl₂ as the familiar single-salt formula. Change the cation observation to red-brown under the corresponding official conditions, and it instead supports Fe³⁺, giving FeCl₃.
The anion remains chloride in both cases, so the chloride-to-iron ratio changes only because iron’s charge changed. This is a compact connection between a visible observation, oxidation state and chemical formula. It helps prevent a student from regarding Roman numerals as optional decorations in a compound name.
An unfamiliar mixture calls for restrained conclusions
A paper may describe a sample with two overlapping reactions or mixed observations inconsistent with one pure salt. Rather than force one compound formula, identify the ions individually where evidence supports them and say if the original pairing cannot be uniquely determined. In a real aqueous mixture, different salts can dissociate into the same ions.
If the task explicitly asks for the unknown single salt and the results cleanly support one cation and one anion, charge balance gives the intended answer. If it instead asks which ions are present, listing the correctly inferred ions may be sufficient. The candidate should read the command word and sample description rather than assume every problem ends with exactly one chemical formula.
The source of precipitate ions is a diagnostic check
In a chloride test, silver came from the silver nitrate reagent. In a sulfate test, barium came from barium nitrate. In a hydroxide test, hydroxide usually came from the sodium hydroxide or ammonia-related reaction environment. The original unknown contributes the ion being investigated, not necessarily all atoms in the newly formed solid.
This source ledger prevents misidentifications. The existence of AgCl does not mean silver was originally present in the sample; the existence of BaSO₄ does not make the unknown a barium salt. Each test is meaningful because a known reagent selectively supplies a reaction partner for an unknown ion. The chemical equation makes that relationship explicit.
A full net-ionic equation check
For Zn²⁺ + 2OH⁻ → Zn(OH)₂(s), zinc and hydroxide counts and charge balance agree. For Ba²⁺ + SO₄²⁻ → BaSO₄(s), the +2 and −2 charges neutralise. For Ag⁺ + I⁻ → AgI(s), the charges also neutralise. These short equations are useful when a student must explain why the precipitation occurs.
A wrong final salt formula can sometimes be diagnosed from an unbalanced charge. If a learner writes ZnI from Zn²⁺ and I⁻, the one iodide only provides −1, leaving an overall positive charge. Write ZnI₂ instead. Charge balance is a separate check after the identity inference, and it is far more reliable than guessing subscripts from the number of test steps.
Avoid adding all reagents to the same sample in an imaginary plan
Separate chemical tests can consume, precipitate or alter ions, and previous reagents may introduce interfering species. A rigorous practical plan uses the provided approved procedure and appropriate separate sample portions when needed; a student should not automatically combine every prescribed test in one tube and still expect clean independent results.
This is a planning principle rather than a request to conduct hazardous reactions. The 2027 practical paper supplies instructions and a qualitative-analysis reference, and students are expected to follow the prescribed methods and safety rules. At home, one can practise interpreting independent result tables instead of combining laboratory acids, silver salts, ammonia or barium compounds.
Why the gas itself may be hazardous
Ammonia can irritate the respiratory system; chlorine and sulfur dioxide can be hazardous gases; acids, hydroxides, silver salts and barium reagents require appropriate safety and disposal controls. These real risks make paper simulations and supervised school laboratories the right contexts for learning qualitative analysis.
The official practical guidance says candidates are not required to carry out sulfur dioxide gas tests even though the qualitative-analysis reference describes them. A student should not interpret the presence of a test in a syllabus as authorisation to perform it alone. Good Chemistry education includes scientific observation and risk awareness, not improvised production of gases at home.
Seven days to combine the tables
Day one, identify cation colours in the official NaOH column. Day two, compare excess sodium hydroxide and excess aqueous ammonia. Day three, practise chloride, iodide and sulfate reagent systems. Day four, revise carbonate and gas confirmation. Day five, build neutral salt formulas from ion charges. Day six, solve two independent cation–anion case studies. Day seven, solve a mixture or ambiguous observation question without being told which candidate to choose.
This revision sequence starts with individual results and ends with integration. A child who cannot distinguish Al³⁺ from Zn²⁺ needs a focused excess-ammonia exercise before attempting six-salt inference. A student who identifies both ions but writes the wrong formula needs charge-balance practice. The first error determines the shortest useful repair.
How a three-student tutorial diagnoses different gaps
One learner sees an aluminium-like precipitate but misreads excess ammonia. Another identifies both ions correctly but writes AlCl instead of AlCl₃. A third uses the correct compound but claims the white AgCl proves the unknown contained silver. These are reagent reading, charge balance and ion-source errors.
A carefully run small-group tutorial can inspect those distinct decisions and then give each student a changed case. The Secondary 4 Chemistry Tutor Hougang overview describes the local teaching context; actual classes, venue and availability need confirmation. A useful outcome is independent reasoning on a new unknown, not a promise of a particular examination grade.
What parents can ask from a three-row result table
Make a fictional table: NaOH gives white precipitate soluble in excess; aqueous ammonia gives white precipitate insoluble in excess; the separate nitric-acid/silver-nitrate test gives white precipitate. Ask your child first for the cation, then the anion, then the complete formula. The intended answers are Al³⁺, Cl⁻ and AlCl₃.
Now change only “ammonia precipitate insoluble” to “ammonia precipitate soluble in excess”. The cation becomes zinc-like under the familiar notes, giving ZnCl₂ when the chloride evidence stays unchanged. The experiment is entirely on paper. If the formula changes correctly, the learner is applying the two independent ion deductions rather than memorising one salt name.
Frequently asked unknown-salt questions
Can one white precipitate identify a salt? Not reliably; reagent sequence and further discrimination matter. How are Al³⁺ and Zn²⁺ distinguished? Their excess-aqueous-ammonia responses differ in the official notes. What distinguishes Fe²⁺ and Fe³⁺? Green versus red-brown hydroxide precipitates in the familiar prescribed tests. What does a deep blue ammonia solution suggest? Cu²⁺ in the stated school context.
Does white AgCl prove silver was originally in the unknown? No; Ag⁺ came from the test reagent. Does white BaSO₄ prove barium was originally present? No; the sulfate ion may come from the unknown. Why use nitric acid before AgNO₃? To follow the prescribed test and reduce interference without adding chloride. Can a mixture be named as one salt? Not necessarily. How is the final formula chosen? Identify both ions and balance their charges.
A whole-salt answer is an evidence chain
Unknown-salt identification is not a game of guessing a familiar compound after seeing one colour. The observations from cation tests identify a positive ion within the available set; independent anion tests identify a negative ion; charges determine a chemically neutral formula. Reagent contributions, solubility and the sample’s single-salt or mixture status all matter to the strength of the conclusion.
For individual-test revision, visit cation tests and anion and gas tests. For formula and reaction validation, use solubility rules. The final question is: which ion did the unknown contribute, which came from the reagent, and what can the combined evidence really prove?
Official 2027 source: SEAB SEC G3 Chemistry K324, Notes for Qualitative Analysis, page 30.