§-Syllabus dot point

NSWChemistrySyllabus dot point

Inquiry Question 1: How are the ions present in the environment identified and measured?

Conduct qualitative investigations to test for the presence in aqueous solutions of cations and anions using flame tests, precipitation reactions and complexation reactions

A focused answer to the HSC Chemistry Module 8 dot point on qualitative ion identification. Flame tests for group 1 and 2 cations, precipitation tests for transition metals and halides, complexation tests for copper, iron and silver, a structured systematic analysis, and worked HSC past exam questions.

Generated by Claude Opus 4.812 min answerUpdated 2026-06-22

Reviewed by: AI editorial process; not yet individually human-reviewed

Have a quick question? Jump to the Q&A page

Jump to a section

What this dot point is asking
The answer
Examples in context
Try this

In plain English

Think of an unknown solution as a stranger you have to identify using only a handful of yes/no questions, and each reagent is one of those questions. The flame test asks "what colour do you burn?", which screens the group 1 and group 2 metal cations in one quick look. Adding sodium hydroxide asks "do you form a coloured jelly-like precipitate?", which separates the transition metal cations by colour: blue for copper, rust-brown for iron(III), dirty green for iron(II). Silver nitrate asks anions the halide question, and the ammonia that follows is a second, more specific question: a chloride precipitate says yes to weak ammonia, a bromide precipitate only half-answers with strong ammonia, and an iodide precipitate refuses to answer at all. You never ask two questions of the same sample at once, because the answers would tangle together, so every test gets a fresh portion of the solution. String enough correct yes/no answers together and there is only one ion left that fits them all.

What this dot point is asking

NESA wants you to be able to identify a list of named cations ( $Ba^{2+}, Ca^{2+}, Mg^{2+}, Pb^{2+}, Ag^+, Cu^{2+}, Fe^{2+}, Fe^{3+}$ ) and anions (chloride, bromide, iodide, hydroxide, acetate, carbonate, sulfate, phosphate) by their colour, flame test, precipitation behaviour with named reagents, and complexation behaviour. You should also be able to sequence tests so that one ion does not interfere with another.

The answer

Flame tests (group 1 and 2 cations mostly)

Heat a clean platinum or nichrome wire in a Bunsen flame until it glows colourless. Dip it in concentrated HCl to clean residual ions, then in the unknown, and observe the flame colour.

Cation	Flame colour
$Li^+$	Carmine red
$Na^+$	Persistent yellow-orange
$K^+$	Lilac (use cobalt-blue glass to block sodium)
$Ca^{2+}$	Brick red
$Sr^{2+}$	Crimson
$Ba^{2+}$	Apple green
$Cu^{2+}$	Blue-green
$Pb^{2+}$	Pale blue (variable)

Flame tests are qualitative only. They are excellent for group 1 and group 2 cations because the electron transitions are in the visible range and the colours are characteristic.

Checkpoints for every qualitative test

Flame test electrons: the flame's heat excites an outer electron to a higher energy level; as it falls back it emits a photon of a characteristic wavelength, giving each metal ion its signature colour.
Fresh portion rule: every reagent test uses a NEW sample. Adding a second reagent to a tube that already has a precipitate in it produces an unreadable mixed solid.
Acidify before $AgNO_3$ / $BaCl_2$ : dilute $HNO_3$ (for silver nitrate) or dilute $HCl$ (for barium chloride) destroys carbonate first, so a genuine halide or sulfate precipitate is not masked by a false-positive $BaCO_3$ or $Ag_2CO_3$ .
Halide colour and solubility ladder: $AgCl$ white/dissolves in dilute $NH_3$ ; $AgBr$ cream/partly dissolves in concentrated $NH_3$ ; $AgI$ yellow/insoluble in $NH_3$ . Increasing $K_{sp}$ decrease and increasing covalent character down the group explain the trend.

Precipitation tests for cations

Add a named reagent and observe the precipitate (colour, texture, solubility in excess).

Reagent	$Cu^{2+}$	$Fe^{2+}$	$Fe^{3+}$	$Pb^{2+}$	$Ag^+$	$Mg^{2+}$	$Ba^{2+}$ , $Ca^{2+}$
Dilute NaOH	Blue gel	Dirty green	Rust brown	White (redissolves in excess)	Brown $Ag_2O$	White	No ppt (Ca slow)
Dilute $NH_3$	Blue, then deep blue with excess (complex)	Green, darkens	Brown	White	Brown then dissolves in excess	White	No ppt
$Na_2CO_3$	Green-blue	Green-white	Brown	White	Pale yellow	White	White
$K_2CrO_4$	Brown	(Pale)	(Brown)	Yellow $PbCrO_4$	Red $Ag_2CrO_4$	(No)	Yellow $BaCrO_4$ , faint $CaCrO_4$

The $Fe^{2+}$ green hydroxide darkens on standing as it oxidises in air to $Fe^{3+}$ brown. State that change explicitly if you see it in a question.

Precipitation tests for anions

Anion	Reagent	Observation
$Cl^-$	$AgNO_3$ in dilute $HNO_3$	White ppt of $AgCl$ , dissolves in dilute $NH_3$
$Br^-$	$AgNO_3$ in dilute $HNO_3$	Cream ppt of $AgBr$ , partly dissolves in concentrated $NH_3$
$I^-$	$AgNO_3$ in dilute $HNO_3$	Yellow ppt of $AgI$ , insoluble in $NH_3$
$SO_4^{2-}$	$BaCl_2$ in dilute $HCl$	White ppt of $BaSO_4$ , insoluble in acid
$CO_3^{2-}$	Dilute $HCl$	Effervescence of $CO_2$ , turns limewater milky
$PO_4^{3-}$	$AgNO_3$ in neutral solution	Yellow ppt of $Ag_3PO_4$
$OH^-$	Universal indicator or pH	Blue/purple, pH > 10
$CH_3COO^-$	Warm with conc. $H_2SO_4$ and ethanol	Fruity smell of ethyl ethanoate (ester)

The order halide colours (white, cream, yellow) and ammonia solubility (yes, partial, no) is the standard halide differentiation.

The acidification step (dilute $HNO_3$ for $AgNO_3$ , dilute $HCl$ for $BaCl_2$ ) destroys any carbonate, which would otherwise also precipitate and give a false positive.

Complexation tests

Complexation distinguishes ions that give similar precipitates by re-dissolving one in excess reagent through formation of a soluble complex ion.

Silver halides with ammonia is the canonical example. $AgCl$ dissolves in dilute $NH_3$ , $AgBr$ partly dissolves in concentrated $NH_3$ , $AgI$ does not dissolve:

AgCl_{(s)} + 2NH_{3(aq)} \rightarrow [Ag(NH_3)_2]^+_{(aq)} + Cl^-_{(aq)}

Copper with ammonia. Add dilute $NH_3$ to a $Cu^{2+}$ solution; pale blue $Cu(OH)_2$ forms, then with excess ammonia it dissolves to give the deep blue tetraammine complex:

Cu(OH)_{2(s)} + 4NH_{3(aq)} \rightarrow [Cu(NH_3)_4]^{2+}_{(aq)} + 2OH^-_{(aq)}

Iron(III) with thiocyanate. Add $KSCN$ to a $Fe^{3+}$ solution; a deep blood-red complex forms:

Fe^{3+}_{(aq)} + SCN^-_{(aq)} \rightarrow [FeSCN]^{2+}_{(aq)}

This test is so sensitive it picks up traces of $Fe^{3+}$ at sub-ppm levels.

Iron(III) with hydroxide vs iron(II) with hydroxide. $Fe^{3+}$ gives rust-brown $Fe(OH)_3$ ; $Fe^{2+}$ gives dirty green $Fe(OH)_2$ that browns on standing. Adding $KSCN$ confirms which is present, since only $Fe^{3+}$ gives the red colour.

A flame test apparatus produces the emission colours used to screen group 1 and group 2 cations before any wet-chemistry reagent is added:

Sequencing a mixed-ion identification

Worked example: identifying two ions from three observations.

Observation 1: the solution is pale green.
Observation 2: dilute NaOH gives a dirty green precipitate that turns brown on standing.
Observation 3: acidified $BaCl_2$ gives a dense white precipitate that does not dissolve in excess dilute acid.

Reasoning. Pale green colour and a dirty-green hydroxide that browns on air exposure both point to $Fe^{2+}$ , since $Fe^{2+}$ slowly oxidises to $Fe^{3+}$ once exposed to atmospheric oxygen:

Fe^{2+}_{(aq)} + 2OH^-_{(aq)} \rightarrow Fe(OH)_{2(s)}

A white precipitate with acidified $BaCl_2$ that survives excess dilute acid is diagnostic of sulfate, since the acidification step has already ruled out a carbonate false positive:

Ba^{2+}_{(aq)} + {SO_4^{2-}}_{(aq)} \rightarrow BaSO_{4(s)}

Conclusion. The unknown is iron(II) sulfate, $FeSO_4$ .

Marker's note: full marks require citing BOTH the colour AND the "browns on standing" behaviour for $Fe^{2+}$ (a rust-brown hydroxide observed immediately, with no colour change, is more consistent with $Fe^{3+}$ already present), plus the acidified reagent and a balanced ionic equation for each identification.

A systematic procedure

When you do not know what is in the sample:

Look. Coloured solution suggests $Cu^{2+}$ (blue), $Fe^{3+}$ (yellow-brown), $Fe^{2+}$ (pale green), $CrO_4^{2-}$ (yellow), $MnO_4^-$ (purple).
Flame test on a small portion to screen group 1/2 cations.
Add NaOH to a fresh portion to test for transition metal hydroxides.
Targeted tests for suspected ions on fresh portions: $AgNO_3$ for halides, $BaCl_2$ for sulfate, dilute HCl for carbonate, $KSCN$ for $Fe^{3+}$ .
Always use a fresh portion for each test. Acidify with the appropriate acid to suppress interferences (carbonate is the most common false-positive).

The extent of precipitation for a sparingly soluble salt like $BaSO_4$ depends on how far the ion product exceeds $K_{sp}$ ; adding excess reagent drives the reaction to near-completion, which is why quantitative gravimetric analysis (Module 8, Inquiry Question 2) can build directly on this qualitative test:

Exam technique for ion identification questions

State the OBSERVATION and the CONCLUSION separately, then back the conclusion with a balanced ionic equation including states. "White precipitate forms" earns partial credit; "white precipitate of $AgCl_{(s)}$ forms, confirming chloride, because $Ag^+_{(aq)} + Cl^-_{(aq)} \rightarrow AgCl_{(s)}$ " earns full credit. When a question gives MULTIPLE observations, use each one to rule ions IN or OUT explicitly rather than jumping straight to the answer; markers award process marks for the elimination logic, not just the final identity. For calculation-style questions built on a qualitative test (e.g. mass of precipitate recovered), always show molar mass, moles from the mass given, the mole ratio from the balanced equation, and the final answer to the correct number of significant figures.

Common traps

Confusing $Fe^{2+}$ and $Fe^{3+}$ hydroxides. Dirty green = $Fe^{2+}$ , rust brown = $Fe^{3+}$ . The $Fe^{2+}$ precipitate goes brown on standing.

Forgetting to acidify before $AgNO_3$ or $BaCl_2$ tests. Carbonate gives a false positive with both silver and barium. A drop of dilute acid before the reagent removes it.

Stating that copper hydroxide is soluble in ammonia in one step: It precipitates first (pale blue), then redissolves in excess to a deep blue complex. Examiners want both stages described.
Treating flame tests as quantitative: They are not. Two ions in the same flame mask each other (sodium is so bright it hides almost everything; use cobalt-blue glass).
Mixing test reagents into one tube: Each test gets a fresh sample portion. Sequential addition produces unreadable mixtures.

Examples in context

Example 1. Identifying contaminants in groundwater near Botany Bay. Following a 1990s legacy spill of dense non-aqueous phase liquids at the Botany Industrial Park, NSW EPA contractors run a qualitative ion screen on every new groundwater sample before quantitative GC-MS. The test sequence covers: flame test (no yellow $Na^+$ confirms no halite intrusion), acidified $AgNO_3$ (a white $AgCl$ precipitate indicates chloride), and acidified $BaCl_2$ (a white $BaSO_4$ flag for sulfate from acid mine drainage). A positive reaction triggers escalation to quantitative AAS for the cation profile. The HSC qualitative framework students apply at the bench is the same first-pass screen that scopes the contamination plume.

Example 2. NSW HSC depth study unknown solution identification. A common Stage 6 task gives students six unlabelled solutions and asks them to identify the cation and anion in each using only the qualitative test reagents on the bench. A solution that turns blue with NaOH (forming $Cu(OH)_2$ ) and then dissolves to deep blue with excess ammonia ( $[Cu(NH_3)_4]^{2+}$ ) is copper. The same solution, treated with $BaCl_2$ in acid and giving a white precipitate, contains sulfate. The student writes the two ion identification net ionic equations: $Cu^{2+} + 4NH_3 \rightarrow [Cu(NH_3)_4]^{2+}$ and $Ba^{2+} + SO_4^{2-} \rightarrow BaSO_4$ . NESA markers reward both the observation and the equation.

Try this

Q1. State the flame test colour for each of: $Li^+$ , $Na^+$ , $K^+$ , $Ca^{2+}$ , $Cu^{2+}$ . [3 marks]

Cue. $Li^+$ crimson; $Na^+$ yellow; $K^+$ lilac; $Ca^{2+}$ brick red; $Cu^{2+}$ blue-green.

Q2. A solution gives a white precipitate with acidified $AgNO_3$ that dissolves in dilute ammonia. Calculate the mass of chloride ion in 50 mL of the solution if 0.072 g of $AgCl$ is recovered. [3 marks]

Cue. $n(AgCl) = 0.072 / 143.32 = 5.02 \times 10^{-4}$ mol; $n(Cl^-) = 5.02 \times 10^{-4}$ mol; mass = $5.02 \times 10^{-4} \times 35.45 = 0.0178$ g.

Q3. A solution gives a pale-green precipitate with NaOH that darkens on standing. (a) Identify the cation. (b) Write the net ionic equation. (c) Describe one additional test to confirm. [2+1+2 marks]

Cue. (a) Iron(II), $Fe^{2+}$ . (b) $Fe^{2+} + 2OH^- \rightarrow Fe(OH)_2$ . (c) Oxidation to brown $Fe(OH)_3$ on air exposure or red colour with $SCN^-$ after oxidation.

Exam-style practice questions

Practice questions written in the style of NESA exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.

2022 HSC5 marksAn unknown aqueous solution gives a brick-red flame test, a white precipitate with silver nitrate that dissolves in dilute ammonia, and no precipitate with sodium hydroxide. Identify the cation and the anion present. Write balanced ionic equations for the precipitation and the complexation steps.

Show worked answer →

A 5 mark answer needs both ions, both equations, and one sentence of reasoning per step.

Brick-red flame test points to calcium $Ca^{2+}$ . (Strontium is crimson, lithium is bright red but more carmine; brick-red is the standard description for calcium.)

No precipitate with NaOH rules out transition metal cations (which form coloured hydroxides) and rules out $Mg^{2+}$ (which gives white $Mg(OH)_2$ ). $Ca(OH)_2$ has $K_{sp} \approx 5 \times 10^{-6}$ and is moderately soluble, so dilute NaOH does not precipitate calcium. Consistent with $Ca^{2+}$ .

White precipitate with silver nitrate that dissolves in dilute ammonia is the classic test for chloride $Cl^-$ :

Ag^+_{(aq)} + Cl^-_{(aq)} \rightarrow AgCl_{(s)}

AgCl_{(s)} + 2NH_{3(aq)} \rightarrow [Ag(NH_3)_2]^+_{(aq)} + Cl^-_{(aq)}

The first equation is the precipitation; the second is the complexation that distinguishes chloride from bromide (which is cream and only partially dissolves in concentrated ammonia) and iodide (which is yellow and does not dissolve).

Conclusion. The unknown is calcium chloride, $CaCl_2$ .

Markers reward (1) calcium from the flame, (2) ruling out other cations using the hydroxide result, (3) chloride from the silver nitrate result, (4) the complexation equation, (5) using both pieces of evidence consistently.

2019 HSC4 marksDescribe a systematic procedure to identify which of the following ions are present in an unknown aqueous solution:

Pb^{2+}

Cu^{2+}

Fe^{3+}

Cl^-

SO_4^{2-}

Show worked answer →

A 4 mark answer needs a sequence of tests with observations and conclusions for each candidate.

Step 1: Inspect the colour: $Cu^{2+}$ is blue, $Fe^{3+}$ is yellow-brown. If the solution is colourless, both are absent. If blue, $Cu^{2+}$ is likely; if yellow-brown, $Fe^{3+}$ is likely.
Step 2: Add dilute NaOH dropwise: A blue gelatinous precipitate confirms $Cu(OH)_2$ (Cu present). A rust-brown precipitate confirms $Fe(OH)_3$ (Fe present). A white precipitate that dissolves in excess NaOH suggests $Pb(OH)_2$ , which is amphoteric.
Step 3: Confirm $Pb^{2+}$: To a fresh portion, add dilute KI. A bright yellow precipitate of $PbI_2$ confirms lead:

Pb^{2+} + 2I^- \rightarrow PbI_{2(s)}

Step 4: Test for chloride. To a fresh portion (acidified with dilute $HNO_3$ to suppress carbonate interference), add silver nitrate. A white precipitate that dissolves in ammonia confirms $Cl^-$ .

Step 5: Test for sulfate. To a fresh portion (acidified with dilute HCl to remove carbonate), add barium chloride. A white precipitate of $BaSO_4$ that does not dissolve in acid confirms $SO_4^{2-}$ .

Always use fresh portions for each test; never add reagents sequentially to one sample because precipitates and complexes interfere with later tests.

Markers reward (1) a logical sequence (inspection then NaOH then specific tests), (2) at least one named precipitate per ion, (3) the acidification step to remove carbonate interference, (4) the principle of using fresh portions.

Practice questions

Original practice questions graded from foundation to exam level, each with a full worked solution. Try them before revealing the solution.

foundation3 marksState the flame colour, precipitate colour with dilute NaOH, and one complexation test result for Cu2+.

Show worked solution →

A 3-mark identify needs all three pieces of evidence stated correctly.

Flame colour: Blue-green.
Precipitate with dilute NaOH: A blue gelatinous precipitate of $Cu(OH)_2$ forms.
Complexation test: Adding excess ammonia dissolves the blue precipitate to give a deep blue solution of the tetraamminecopper(II) complex:

Cu(OH)_{2(s)} + 4NH_{3(aq)} \rightarrow [Cu(NH_3)_4]^{2+}_{(aq)} + 2OH^-_{(aq)}

Marking criteria: 1 mark for the correct flame colour, 1 mark for the correct precipitate colour and formula, 1 mark for the correct complexation observation and equation.

foundation4 marksOutline a two-step test sequence, using fresh portions, that distinguishes chloride, bromide and iodide ions in three separate unknown solutions.

Show worked solution →

A 4-mark outline needs the reagent, the acidified background, and the distinguishing observation for all three halides.

Step 1 (each fresh portion): add acidified $AgNO_3$ (in dilute $HNO_3$ , to exclude carbonate interference).

Chloride gives a white precipitate of $AgCl$ .
Bromide gives a cream precipitate of $AgBr$ .
Iodide gives a yellow precipitate of $AgI$ .

Step 2: add dilute ammonia to each precipitate.

The white ( $Cl^-$ ) precipitate dissolves completely in dilute ammonia.
The cream ( $Br^-$ ) precipitate only partly dissolves, and only in concentrated ammonia.
The yellow ( $I^-$ ) precipitate does not dissolve at all.

Marking criteria: 1 mark for the correct acidified reagent, 1 mark for the three correct precipitate colours, 1 mark for the correct ammonia-solubility ladder, 1 mark for explicitly using fresh portions of each unknown.

core5 marksA 25.00 mL sample of

0.0500\ \text{mol L}^{-1}

sodium sulfate solution is treated with excess barium chloride solution, precipitating all the sulfate as barium sulfate. Calculate the theoretical mass of

BaSO_4

recovered, to 3 significant figures. (

M(BaSO_4) = 233.40\ \text{g mol}^{-1}

Show worked solution →

Step 1: write the balanced ionic equation.

Ba^{2+}_{(aq)} + {SO_4^{2-}}_{(aq)} \rightarrow BaSO_{4(s)}

One mole of sulfate gives one mole of barium sulfate precipitate.

Step 2: moles of sulfate ion available.

n(SO_4^{2-}) = c \times V = 0.0500\ \text{mol L}^{-1} \times 0.02500\ \text{L} = 1.25 \times 10^{-3}\ \text{mol}

Step 3: moles of $BaSO_4$ (1:1 ratio, barium chloride in excess so sulfate is the limiting reagent).

n(BaSO_4) = 1.25 \times 10^{-3}\ \text{mol}

Step 4: mass of $BaSO_4$ .

m = n \times M = 1.25 \times 10^{-3}\ \text{mol} \times 233.40\ \text{g mol}^{-1} = 0.29175\ \text{g}

Step 5: round to 3 significant figures (matching 0.0500 and 25.00, both 3 s.f.).

m(BaSO_4) = 0.292\ \text{g}

Marking criteria: 1 mark for the correct balanced 1:1 ionic equation, 1 mark for correct moles of sulfate from concentration and volume, 1 mark for correctly identifying barium chloride as being in excess (sulfate limiting), 1 mark for the mass calculation, 1 mark for the correct answer to 3 significant figures with units. This is a THEORETICAL yield; real gravimetric practice loses some mass on filtration and washing.

core5 marksThe graph below is an owned illustrative plot of the mass of BaSO4 precipitate formed as

0.100\ \text{mol L}^{-1}

BaCl2 solution is added dropwise to a fixed volume of unknown sulfate sample, reaching a plateau at 0.292 g once about 12.5 mL has been added. (a) Explain the shape of the curve in terms of limiting and excess reagent. (b) State how this graph could be used to determine the original moles of sulfate present.

Show worked solution →

(a) Explanation of shape. While $BaCl_2$ is the limiting reagent (the rising, roughly linear region), each additional mL added reacts completely and precipitates proportionally more $BaSO_4$ , so the mass rises steadily. Once enough $BaCl_2$ has been added to react with ALL the sulfate present (around 12.5 mL here), sulfate becomes the limiting reagent and any further $BaCl_2$ added is in excess with nothing left to precipitate, so the mass of precipitate plateaus at a constant maximum value (0.292 g).

(b) Using the graph to find moles of sulfate. The plateau mass (0.292 g) is the maximum possible mass of $BaSO_4$ , corresponding to complete precipitation of all the original sulfate. Converting: $n(BaSO_4) = 0.292 / 233.40 = 1.25 \times 10^{-3}$ mol, and since the stoichiometric ratio is 1:1, $n(SO_4^{2-})_{\text{original}} = 1.25 \times 10^{-3}$ mol. The volume at which the plateau BEGINS (the "elbow" of the curve, about 12.5 mL of 0.100 mol/L BaCl2) is an independent cross-check: $n(Ba^{2+})$ added at that point should also equal $1.25 \times 10^{-3}$ mol.

Marking criteria: (a) 1 mark for correctly identifying the rising region as BaCl2-limited, 1 mark for identifying the plateau as sulfate exhausted/BaCl2 in excess. (b) 1 mark for reading the plateau mass correctly, 1 mark for the mass-to-moles conversion using the 1:1 ratio, 1 mark for the independent cross-check using the elbow volume.

exam7 marksA council water-testing laboratory receives an unlabelled aqueous sample and must identify BOTH the cation and the anion present using only qualitative bench tests, given this evidence: a pale-green solution; dilute NaOH gives a dirty-green precipitate that turns rust brown on standing; acidified BaCl2 gives a dense white precipitate insoluble in excess dilute acid. Justify a full identification, including balanced ionic equations, and evaluate ONE limitation of relying only on qualitative tests for a council report on water quality.

Show worked solution →

This is a 7-mark JUSTIFY and EVALUATE: markers reward reasoned identification using ALL the evidence, correct equations, and a genuine evaluative limitation.

Band 6 PLAN.

Use the colour evidence, the NaOH precipitate colour and its standing behaviour, and the BaCl2 result together, ruling other ions in or out at each stage.
Write balanced ionic equations with states for both the cation and anion identification steps.
Evaluate a real limitation of qualitative-only testing for a regulatory water report (e.g. no quantitative concentration, or interference/sensitivity limits), not a vague "tests can go wrong" statement.

Model answer.

Pale-green colour is consistent with $Fe^{2+}$ (not $Fe^{3+}$ , which is yellow-brown, and not $Cu^{2+}$ , which is blue). Adding dilute NaOH gives a dirty-green precipitate:

Fe^{2+}_{(aq)} + 2OH^-_{(aq)} \rightarrow Fe(OH)_{2(s)}

This precipitate darkening to rust brown on standing is diagnostic of $Fe^{2+}$ specifically, because atmospheric oxygen slowly oxidises the pale $Fe(OH)_2$ to $Fe(OH)_3$ ; a precipitate that was rust brown FROM THE START would instead indicate $Fe^{3+}$ was already present, so the standing-colour-change evidence rules out $Fe^{3+}$ as the original cation.

Acidifying a fresh portion with dilute HCl and adding $BaCl_2$ gives a dense white precipitate that survives excess acid. The acidification step first destroys any carbonate, so a precipitate that persists in excess acid is not $BaCO_3$ ; it must be:

Ba^{2+}_{(aq)} + {SO_4^{2-}}_{(aq)} \rightarrow BaSO_{4(s)}

confirming sulfate. Combining both identifications, the sample is iron(II) sulfate, $FeSO_4$ .

Evaluation. A genuine limitation is that qualitative tests only confirm PRESENCE, not concentration: a council report needs to know whether the iron and sulfate levels breach a numerical guideline (e.g. Australian Drinking Water Guidelines), which qualitative colour/precipitate tests cannot supply. This is why a positive qualitative screen is always followed by a quantitative technique (such as AAS for the cation or gravimetric/titrimetric analysis for the anion) before a compliance decision is reported; qualitative tests are a fast, cheap first-pass screen, not a substitute for measurement.

Marker's note: top-band answers (1) use every piece of given evidence, not just the most obvious one, (2) explicitly rule an alternative ion IN or OUT using the standing-colour-change reasoning, (3) give both balanced ionic equations with states, and (4) evaluate a real, specific limitation (lack of quantification) rather than a generic "human error" statement.