How do you identify an unknown salt in Secondary 3 Chemistry without guessing from one precipitate? A salt contains positive ions (cations) and negative ions (anions). A suitable identification therefore needs evidence for both sides of the formula. In the 2027 Singapore–Cambridge SEC G3 Chemistry K324 qualitative-analysis method, aqueous sodium hydroxide and aqueous ammonia help distinguish selected cations by the colour of a precipitate and whether it dissolves in excess reagent. The anion tests use particular reagents and characteristic observations—for example, acidified silver nitrate for chloride or iodide, and acidified barium nitrate for sulfate. The safe scientific decision is to work from observations toward a justified ion identity, not from an attractive colour directly to an entire salt name.
I would begin a lesson with this question: “A white precipitate formed. Does that prove the unknown contains zinc?” It does not. Aluminium, zinc, calcium and several other ions can give white precipitates in suitable tests. We need the identity of the reagent, whether it was added in excess, and sometimes the result of a different reagent on a fresh portion. Even a reliable cation result does not identify the anion.
For students and parents searching for a Secondary 3 Chemistry tutor in Hougang, unknown salt identification, qualitative analysis decision tree, cation and anion tests, aqueous ammonia versus sodium hydroxide, or G3 Pure Chemistry tuition near Kovan, this article gives an evidence-first method with worked examples, explanations of ambiguity and exam practice.
The official 2027 SEC G3 Chemistry K324 syllabus covers qualitative analysis in Section 6; the prescribed Notes for Qualitative Analysis appear on page 30 of the document and are supplied with the practical examination. They are the reference for named cation, anion and gas tests. G1, G2 and Combined Science pupils may have different test lists and assessment requirements, so always check the student’s actual course and year.
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The one-minute unknown-salt decision method
Before examining a table, identify what each kind of observation can actually establish.
- Step 1: read the sample information. Is it a solution, a solid or a mixture? Has the question stated that the unknown is a single pure salt, or could several substances be present?
- Step 2: inspect the cation tests. Record any precipitate’s initial colour and its behaviour in excess aqueous NaOH and excess aqueous ammonia, where those results are supplied.
- Step 3: identify any gas test. If warming with alkali produces gas that turns damp red litmus blue, consider ammonium ions and the ammonia gas evidence.
- Step 4: inspect the anion tests on an appropriate separate portion. The prescribed acidified silver nitrate or barium nitrate tests are not interchangeable, and carbonate is assessed differently.
- Step 5: cross-check both ions. Construct a neutral compound formula using the ion charges, such as Cu²⁺ with two Cl⁻ giving CuCl₂.
- Step 6: name the boundary of the conclusion. If the test information is insufficient or multiple salts could produce the observations, say what is supported and what remains unknown.
A three-pupil tutorial can work through the same procedure using different unknowns, but each student should have to justify the result independently.
Observations come before inferences
A blue precipitate is an observation. “Copper(II) ions are indicated” is an inference supported by that observation under the specified test. “The unknown is copper(II) chloride” is a further compound-identity claim requiring evidence for chloride as well.
The difference matters in questions that ask state the observation, identify the ion, or suggest an explanation. Answering a different level of the question can lose marks, even if the pupil remembers the correct chemical substance.
I teach students to use three short headings on rough paper: Reagent and condition / Observation / Ion indicated. The reagent and its concentration or excess state are part of the evidence. A precipitate colour without the reagent name is rarely enough.
If a question says “in excess”, take that word seriously. Some precipitates dissolve in a large excess of one reagent but not the other, which is often the clue distinguishing two cations.
Why anions and cations must be identified separately
A typical salt consists of a positive ion paired with a negative ion. CuSO₄, for example, contains copper(II) ions and sulfate ions in a 1:1 ratio. CuCl₂ contains the same copper(II) ion but chloride ions in a 1:2 ratio.
The two compounds can share observations associated with Cu²⁺ yet respond differently to tests for sulfate or chloride. A positive copper(II) cation test alone therefore does not tell you which salt is present.
Conversely, two different salts such as Na₂SO₄ and ZnSO₄ can both contain sulfate, giving the characteristic sulfate anion result, yet have different positive ions.
To name an entire salt, you need a justified combination of its ions and evidence that the simple single-salt model is appropriate. When the unknown may be a mixture, detecting two ions does not necessarily prove they originated from one compound.
This is a general scientific lesson about combinations: observing the parts of a mixture is not the same as determining how those parts were originally paired.
The cation information you should be able to read
The official K324 qualitative-analysis notes include Al³⁺, NH₄⁺, Ca²⁺, Cu²⁺, Fe²⁺, Fe³⁺ and Zn²⁺. Their familiar sodium hydroxide observations are different enough to guide identification when the context is appropriate:
- Al³⁺: white precipitate with aqueous NaOH, soluble in excess, giving a colourless solution.
- Ca²⁺: white precipitate with aqueous NaOH, insoluble in excess.
- Cu²⁺: light blue precipitate, insoluble in excess NaOH.
- Fe²⁺: green precipitate, insoluble in excess NaOH.
- Fe³⁺: red-brown precipitate, insoluble in excess NaOH.
- Zn²⁺: white precipitate, soluble in excess NaOH, giving a colourless solution.
- NH₄⁺: gives ammonia on warming with alkali, rather than a characteristic metal-hydroxide precipitate in the listed test.
These are the expected school observations under the specified conditions, not a guarantee that arbitrary real-world mixtures will always show clean colours without interference.
The cation tests using NaOH and aqueous ammonia guide covers the individual precipitate patterns in more depth.
Why the excess-ammonia test changes the conclusion
Some cations give a white precipitate with aqueous NaOH that later dissolves in excess alkali. That alone can leave more than one candidate.
The aqueous ammonia observations help:
- Al³⁺: white precipitate insoluble in excess aqueous ammonia.
- Zn²⁺: white precipitate that dissolves in excess aqueous ammonia, giving a colourless solution.
- Ca²⁺: no precipitate in the official ammonia-test notes.
- Cu²⁺: light blue precipitate that dissolves in excess aqueous ammonia, giving a dark blue solution.
- Fe²⁺ and Fe³⁺: familiar green and red-brown precipitates that remain insoluble in excess aqueous ammonia.
Why does a white precipitate sometimes dissolve? The larger body of coordination and equilibrium Chemistry may involve soluble complex-ion formation or amphoteric hydroxide behaviour, depending on the reagent and ion. The official G3 syllabus states that formulae of complex ions are not required for this cation-testing outcome.
A student should be able to describe the change correctly without inventing complex formulas beyond the assessed level.
The ion that can give both a white precipitate and a colourless solution
Imagine a fresh portion of unknown solution gives a white precipitate on adding NaOH. It dissolves in excess NaOH. Another portion gives a white precipitate with aqueous ammonia, which also dissolves in excess ammonia to form a colourless solution.
Under the official K324 table and the simplified single-ion question, this pattern strongly indicates Zn²⁺. Aluminium is less consistent because its precipitate is insoluble in excess aqueous ammonia.
The key observation is not just white. It is white → dissolves in excess NaOH and white → dissolves in excess ammonia.
This is a good example of why a student who memorises one colour word will struggle when a Paper 3 table adds a second column.
A different white-precipitate pattern indicates aluminium
Now change the second observation: the unknown forms a white precipitate with NaOH, dissolves in excess NaOH, but remains a white precipitate insoluble in excess aqueous ammonia.
This matches the standard Al³⁺ pattern. The first test resembles Zn²⁺; the second distinguishes them under the simple school comparison.
What if the question provides only the NaOH result? Then Al³⁺ and Zn²⁺ may remain possible. A careful answer lists or qualifies the candidates, or suggests an appropriate further test, rather than inventing certainty.
I ask pupils to say: “Which observation excluded zinc?” They should point directly to the behaviour in excess aqueous ammonia.
This form of evidence-based elimination is powerful across many science disciplines.
The copper(II) clue is not just blue
Aqueous Cu²⁺ can produce a light blue precipitate with NaOH or with an initial addition of aqueous ammonia. In the prescribed excess-ammonia comparison, that light blue precipitate dissolves to give a dark blue solution.
A student who writes simply “blue solution” loses a useful observational distinction between a precipitate and a solution. The process matters: light blue precipitate initially, dark blue solution in excess ammonia.
If the exam asks for an observation, describe both stages when relevant. If it asks for the ion, the pattern indicates copper(II) under the prescribed method.
This is also an example of how observations can change when more reagent is added—something a pupil should not dismiss as irrelevant once the first colour appears.
Iron(II) versus iron(III): green and red-brown
Under the official K324 table, Fe²⁺ typically gives a green precipitate with NaOH or aqueous ammonia, while Fe³⁺ gives a red-brown precipitate. Both are listed as insoluble in excess of the relevant reagent.
A student should distinguish iron(II) from iron(III) in both name and charge. The correct ions are Fe²⁺ and Fe³⁺; their hydroxide formulas differ correspondingly.
In real samples, iron(II) hydroxide can darken or become brownish when exposed to oxygen as Fe(II) species are oxidised. That is why an observed “brown after standing” result should not be treated as identical to an immediate red-brown precipitate under a controlled test.
Read the timing and conditions. The iron(II) hydroxide turning brown guide explains this specific extension.
The ammonium exception: gas rather than metal hydroxide
An aqueous sample containing NH₄⁺ can release ammonia gas, NH₃, when treated with aqueous sodium hydroxide and warmed under the prescribed school test. The net ionic equation is:
NH₄⁺(aq) + OH⁻(aq) → NH₃(g) + H₂O(l).
The gas turns damp red litmus paper blue under the standard observation. The positive indication is for ammonium ions in the original sample, not nitrate by default.
Be cautious because the separate nitrate test also generates ammonia by reduction using aluminium and alkali. The mechanism and possible source of the gas differ. Treating “ammonia appeared” as unambiguous proof of nitrate would be wrong if ammonium is also a possible source.
This is a useful place to separate gas identification from identification of the ion that generated the gas.
Anion tests: why the acidification reagent matters
The 2027 K324 notes include five anion tests: carbonate, chloride, iodide, nitrate and sulfate. The standard results are:
- CO₃²⁻: add dilute acid; effervescence may occur and CO₂ is indicated by the suitable limewater test.
- Cl⁻ in solution: acidify with dilute nitric acid, then add aqueous silver nitrate; white precipitate.
- I⁻ in solution: acidify with dilute nitric acid, then add aqueous silver nitrate; yellow precipitate.
- NO₃⁻ in solution: add aqueous sodium hydroxide, then aluminium foil, warm carefully; ammonia is produced, to be identified with damp red litmus as appropriate.
- SO₄²⁻ in solution: acidify with dilute nitric acid, then add aqueous barium nitrate; white precipitate.
For the sulfate test, barium ions form insoluble BaSO₄. For chloride, silver ions form insoluble AgCl. These are different reactions, despite both potentially producing white solids.
The student should not substitute sulfuric acid when testing for sulfate: that acid can supply the very ion being investigated, creating misleading evidence.
Why two different white precipitates cannot be swapped
In the chloride test:
Ag⁺(aq) + Cl⁻(aq) → AgCl(s).
In the sulfate test:
Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s).
Both reactions can produce white precipitates, but the reagent ions and solid formulas are different. A student who writes “white means chloride” without naming the silver nitrate conditions has not completed the inference.
This is why the procedure matters. Acidification controls some potential interferences, while the selected metal-ion reagent supplies the partner that precipitates the target anion.
The sulfate test with barium nitrate and chloride test with nitric acid and silver nitrate are separate companion guides.
Worked unknown 1: a possible zinc sulfate sample
A hypothetical school question says an unknown single salt has been dissolved appropriately. Two separately described cation tests give white precipitates that dissolve in excess aqueous NaOH and in excess aqueous ammonia, each giving colourless solutions. A fresh portion is acidified with nitric acid and then treated with barium nitrate; a white precipitate forms.
The cation observations support Zn²⁺ under the standard K324 table. The acidified barium nitrate result supports SO₄²⁻. Combining their charges gives ZnSO₄, zinc sulfate, as the consistent single-salt identification.
Why not Zn₂SO₄? Zn²⁺ and SO₄²⁻ balance at a 1:1 ion ratio. The extra zinc would leave excess positive charge.
The words single salt matter. If the unknown could be a mixture, finding zinc and sulfate ions would not necessarily prove that all ions originated as zinc sulfate rather than separate salts.
Worked unknown 2: copper(II) chloride
Now imagine a single unknown salt solution gives a light blue precipitate with aqueous ammonia, which dissolves in excess ammonia to form a dark blue solution. A different portion is acidified with dilute nitric acid, then treated with aqueous silver nitrate, giving a white precipitate.
The cation evidence supports Cu²⁺. The anion evidence supports Cl⁻. Charge balance requires two Cl⁻ for one Cu²⁺, so the formula is CuCl₂.
Students often write CuCl₂ correctly but cannot justify the 2. The reason is the ion charges: +2 plus two contributions of −1 gives zero. A balanced formula is not selected because the solution looked blue; it follows from ionic charge conservation.
The corresponding simplified identity inference assumes the sample really is one suitable salt, rather than a complex mixture.
Worked unknown 3: iron(III) nitrate and the limits of a gas inference
Imagine an unknown’s cation test indicates Fe³⁺ from a red-brown precipitate, and the question separately provides reliable evidence for NO₃⁻. Under a single pure salt assumption, the neutral formula is Fe(NO₃)₃.
Each Fe³⁺ requires three nitrate ions to balance charge. Do not write FeNO₃ simply because one nitrate appears in the name, and do not call the salt iron(II) nitrate when the positive ion was identified as iron(III).
If the anion evidence comes from a procedure generating ammonia, make sure the test conditions and possible ammonium interference have been addressed or excluded by the question. The nitrate conclusion cannot rest solely on the fact that ammonia was observed when ammonium was also a plausible initial ion.
That boundary between observing NH₃ and proving nitrate as its source is part of analytical accuracy.
The unknown might not be a single compound
A laboratory sample may contain multiple ionic substances. Suppose tests independently support Cu²⁺ and Zn²⁺ cations as well as chloride and sulfate anions. You cannot determine precisely how the ions were paired in the original materials from those observations alone.
Possible original mixtures can yield the same set of dissolved ions after mixing. A test of the final solution identifies species present, not necessarily the packaging of every original salt.
In a structured school question, the phrase “a single salt” or a limited option list often makes a specific formula deduction appropriate. In a real environmental sample, further analytical work would be needed.
I teach students to honour the information in the question, neither under-answering a well-specified simple salt problem nor overstating certainty when the sample could be complex.
How to record results without turning observation into conclusion
A reliable result statement might read: “A light blue precipitate formed; it dissolved upon adding excess aqueous ammonia to give a dark blue solution.” That is much more informative than writing “copper detected” in an observation column.
The inference column can then state: “Copper(II) ions are indicated under the specified cation test.”
For a chloride anion test, write “A white precipitate formed after the acidified sample was treated with aqueous silver nitrate”, rather than “the test was positive” alone.
Precise language allows someone else to evaluate whether the conclusion follows from the actual observations. That is the heart of practical Chemistry, not merely a neat notebook style.
A common question: why must tests use separate portions?
When one test adds a strong reagent, the sample’s chemical composition may change. Adding barium nitrate introduces Ba²⁺ and nitrate ions; adding silver nitrate introduces Ag⁺ and nitrate; alkali changes pH. Later observations could then reflect what an earlier reagent added, not what the unknown originally contained.
School analysis therefore often specifies separate portions or an appropriate test sequence. A student should follow that design rather than assume every reagent can be added into one tube successively without affecting interpretation.
For written examination questions, the requirement is to understand contamination of the evidence by the testing process. Any actual reagents should be handled only according to the teacher’s laboratory instructions.
The principle is general: a test is more credible when the way the observation was produced is traceable.
Never substitute an arbitrary acid for a prescribed one
The K324 anion notes specify dilute nitric acid before the silver nitrate tests for chloride and iodide and before the barium nitrate test for sulfate.
Adding hydrochloric acid to a sample before testing for chloride would introduce Cl⁻. A subsequent silver chloride precipitate could then come from the reagent rather than the sample. Adding sulfuric acid before a sulfate test could supply SO₄²⁻ and produce an equally misleading result.
Students should learn the purpose of the specified acidification rather than memorise “add any acid.” A chemical test only works as evidence if the method does not introduce the target species inadvertently.
This is a practical illustration of scientific controls and false-positive avoidance.
What if there is no precipitate?
“No visible precipitate under the prescribed conditions” is an observation. It can help rule out certain ions within the scope and detection limits of the school test.
But no precipitate cannot prove a substance contains absolutely none of a species in every possible trace amount. Concentration, interfering ions and the sensitivity of the observation can matter in real chemical analysis.
For an exam, if the question supplies a set of candidate ions and assumes ordinary working concentrations, the negative result may be decisive among the options. Use that information rather than inventing a different ion not in the task.
A careful learner can reason with both positive and negative evidence, while recognising that certainty depends on the experimental model.
A four-question rapid diagnostic for tutoring
Give the student four statements and ask which are justified.
- “White precipitate formed with NaOH.” Can you name the cation uniquely? No, not from this alone.
- “White with NaOH dissolves in excess, white with ammonia remains.” Which is indicated among Al³⁺ and Zn²⁺? Al³⁺.
- “White precipitate after dilute nitric acid and aqueous silver nitrate.” Which anion is indicated in the standard school test? Cl⁻.
- “Both Zn²⁺ and SO₄²⁻ are detected in a sample known to be a single salt.” What formula is consistent? ZnSO₄.
Then remove the words “single salt” from the fourth question and ask whether the formula is still uniquely established as the original chemical. The improved answer is not necessarily.
Those changed assumptions reveal whether a child understands the limits of evidence, not just the colours.
The role of charge balance after ion identification
An unknown salt’s name follows from the ions; its formula follows from charge balance. Examples:
- Cu²⁺ and Cl⁻ give CuCl₂.
- Zn²⁺ and SO₄²⁻ give ZnSO₄.
- Al³⁺ and NO₃⁻ give Al(NO₃)₃.
- Fe³⁺ and SO₄²⁻ give Fe₂(SO₄)₃.
- NH₄⁺ and Cl⁻ give NH₄Cl.
The subscripts are not selected by precipitate darkness, mass or how many reagents were used. They are fixed by the smallest ratio that gives an electrically neutral compound.
A pupil who identifies both ions yet writes the wrong formula needs a charge-balance repair, not another qualitative-analysis colour table.
Why the practical syllabus gives you notes
For the 2027 SEC G3 K324 Paper 3 practical, the Notes for Qualitative Analysis are printed with the paper. That makes it especially important for students to know how to read and use those notes quickly, rather than rely exclusively on memorising every line.
A provided reference does not perform the interpretation for you. A candidate must still recognise a precipitate as a solid, notice whether it dissolves in excess, identify what gas was actually evolved and integrate separate cation and anion clues.
There is no benefit in pretending an examination reference sheet removes the need for chemical understanding. It changes the task from raw recall to correct retrieval, selection and evidence-based inference.
Home learning without chemical risks
Qualitative-analysis reagents may include hazardous acids, alkalis, silver and barium salts. Families should not prepare unknown-salt tests with improvised household equipment. These reactions belong to properly supervised school laboratory work with approved risk controls and waste disposal.
A parent can still help effectively at home. Present a completed observation table with the ion names covered. Ask the pupil to name the reagent, the observed change and the ion supported. Then reveal an alternative possible ion and ask what further test would distinguish it.
This builds decision-making without exposure to chemicals. It also prepares students for data-based examination questions where results are supplied rather than physically observed.
How a responsive three-student tutorial can teach this
Imagine three Secondary 3 learners. Student A can name a copper(II) ion but writes CuCl instead of CuCl₂. Student B identifies chloride from a white silver nitrate precipitate but forgets nitric acid pretreatment. Student C uses the reagent sequence correctly yet claims any detected cation and anion must have originated from one salt, even when the unknown is a mixture.
The repairs differ. Student A needs charge balance. Student B needs test specificity and contamination control. Student C needs limits of inference.
In a deliberately small three-pupil tutorial, these different conceptual errors can be corrected with targeted practice. Then all three should solve a new unknown, without the answer table beside them, and explain which observation supported each step.
This is the teaching outcome parents should look for: independent reasoning after feedback, not simply a class of three students copying the same colour chart.
Independent exam practice: fifteen questions
- 1. Why is a cation test alone insufficient to identify an entire unknown salt?
- 2. Which K324 test reagents are used to distinguish aluminium and zinc ions by excess-reagent behaviour?
- 3. Which cation can give a light blue precipitate and a dark blue solution in excess aqueous ammonia?
- 4. Which cation commonly gives a red-brown hydroxide precipitate?
- 5. Which cation gives a green precipitate in the standard K324 table?
- 6. A white precipitate dissolves in excess NaOH but not excess ammonia. Which is consistent among Al³⁺ and Zn²⁺?
- 7. A white precipitate dissolves in excess both NaOH and ammonia. Which is consistent among Al³⁺ and Zn²⁺?
- 8. Which anion gives a white precipitate with acidified silver nitrate?
- 9. Which anion gives a yellow precipitate with acidified silver nitrate?
- 10. Which anion is indicated by a white precipitate with acidified barium nitrate?
- 11. Which gas turns damp red litmus blue, and which cation may release it with warmed alkali?
- 12. Why should hydrochloric acid not be used to acidify a sample for a chloride test?
- 13. If tests indicate Cu²⁺ and SO₄²⁻ in a sample known to be one pure salt, what formula fits?
- 14. If tests indicate Fe³⁺ and Cl⁻ in a sample known to be one pure salt, what formula fits?
- 15. If a mixture may contain two cations and two anions, can you always identify the original salt pairings from the final dissolved ion tests alone?
Attempt the questions before reading the answers. In particular, explain why an observation is enough or not enough for a conclusion.
Answer key: reasons matter
- 1. A salt needs a cation and anion, and identifying only one leaves the other unknown.
- 2. Aqueous NaOH and aqueous ammonia, noting precipitate formation and solubility in excess.
- 3. Cu²⁺, copper(II).
- 4. Fe³⁺, iron(III).
- 5. Fe²⁺, iron(II).
- 6. Al³⁺.
- 7. Zn²⁺.
- 8. Cl⁻, chloride.
- 9. I⁻, iodide.
- 10. SO₄²⁻, sulfate.
- 11. Ammonia, NH₃; NH₄⁺ may release it with alkali on warming.
- 12. HCl introduces Cl⁻, risking a misleading silver chloride precipitate.
- 13. CuSO₄.
- 14. FeCl₃, since three Cl⁻ balance Fe³⁺.
- 15. No. The dissolved-ion results may not uniquely establish the original pairings.
A correct identification without the reagent and inference behind it may still be fragile. Retest with a fresh unknown a week later.
Seven days from colour memorisation to identification mastery
Day 1: revise cation names, charges and hydroxide colours. Day 2: focus on the effect of excess NaOH. Day 3: add the effect of excess ammonia, contrasting aluminium, zinc, calcium and copper.
Day 4: learn the five anion test families using the actual 2027 K324 qualitative-analysis notes. Day 5: practise gas-test evidence without confusing nitrate and ammonium. Day 6: combine one cation and one anion and derive the salt formula by charge balance.
Day 7: read two new unknown-sample reports—one known to be a single salt and one possible mixture—and explain how the strength of the conclusion differs. Repeat one case later without hints, because true learning is tested by transfer, not immediate recognition.
Questions Hougang parents often ask
Does any white precipitate mean zinc?
No. Multiple ions can give white precipitates. The reagent and excess behaviour are essential.
Is the dark blue solution the same as a blue precipitate?
No. The K324 copper(II) test describes a light blue precipitate dissolving in excess ammonia to give a dark blue solution.
Does the practical examination provide the qualitative-analysis notes?
The 2027 G3 K324 Paper 3 includes the notes as specified in the official syllabus.
Does a positive chloride test prove that the salt is sodium chloride?
No. It supports chloride ions, not the identity of the cation.
Why do we use nitric acid before silver nitrate?
It is the prescribed acidification reagent and avoids introducing chloride through the acid.
What if a sample is a mixture?
A positive set of ion tests may identify ions present but not unique original salt pairings.
Can children perform these tests at home?
No. The reagents require school laboratory supervision. Interpretation and equation practice can be done at home without chemicals.
Is this part of 2027 G3 Pure Chemistry?
Yes, Section 6 and the Notes for Qualitative Analysis. Check other subject routes separately.
The idea that turns qualitative analysis into real Chemistry
The most useful question is not “What colour did we memorise?” It is what did we add, what changed, which ion could cause that change, and what else could cause it? Only then do we combine ion evidence into a chemical formula where the question supports that inference.
A student who can explain those steps from an unfamiliar observation sheet is beginning to use Chemistry as a method of investigation rather than a list of colours.
Secondary 3 Cation Tests: NaOH vs Aqueous Ammonia · Sulfate Test Using Nitric Acid and Barium Nitrate · Nitrate Ion Test and Ammonium Interference · Solubility Rules and Precipitates · 2027 SEC G3 K324 Official Syllabus and QA Notes · Hougang Chemistry Reference Library · eduKate 3-Pax Teaching Reference · Parent–Student Consultation.