Every other chemistry topic asks what does this substance do? C8 asks the opposite question — which substance is this? — and gives you a toolkit of qualitative tests to answer it. Each test is built on a reaction with an unmistakable signal: a gas with a distinctive property, a sharp colour change, or an insoluble solid that drops out of solution as a precipitate. Learn to read those signals and you can name an unknown from a few drops on a bench.
C8 (4.8) has a clean Combined/Triple split — check what you actually need:
- Everyone — sections 1–4: pure substances and mixtures, formulations, paper chromatography and Rf values (Required practical 6), and the tests for the four common gases.
- Triple (Chemistry only) T — sections 5–8: identifying ions by flame tests, sodium hydroxide precipitates, and the carbonate, halide and sulfate tests (Required practical 7), then instrumental methods and flame emission spectroscopy.
There is no Higher-only content in C8 — the whole topic is accessible at both tiers. The challenge here is precision of language: examiners want the exact reagent, the exact observation, and the exact words (“colourless” not “clear”; “milky” or “cloudy” for limewater). Marks are won and lost on a single word, so these notes flag the wording the mark scheme rewards.
1Pure Substances & Mixtures
The word pure means something very specific in chemistry — and it is not what it means on a carton of orange juice. Getting that distinction right, and knowing the one physical test that reveals purity, is worth easy marks — it comes up early and often in the exam.
- Pure substance (chemistry) — a single element or a single compound, with nothing else mixed in. A beaker of pure water contains only H2O molecules.
- Pure substance (everyday language) — something natural or unadulterated, that has had nothing added to it, e.g. “pure” milk or “pure” orange juice. In chemistry, milk is a mixture.
- Mixture — two or more elements or compounds physically mixed together, not chemically combined. The components keep their own properties and can be separated by physical means — the filtration, crystallisation and distillation methods from C1, plus the chromatography you meet below.
Black and red circles are two different types of atom (the same diagram you met in C1). An element and a compound are both pure — one kind of particle throughout; a mixture contains more than one kind, jumbled together and not chemically bonded.
Telling pure from impure: the melting point test
A pure substance melts and boils at one specific, sharp temperature — pure ice melts at exactly 0 °C and pure water boils at exactly 100 °C. A mixture melts and boils over a range of temperatures, and the impurities usually lower the melting point as well as spreading it out. So melting point and boiling point data can be used to distinguish a pure substance from a mixture: compare the measured value with the data-book value, and check whether the change of state is sharp or smeared over a range.
On a cooling curve, a pure substance freezes at a single sharp temperature (a flat plateau); a mixture freezes over a range, so the line only slopes.
A liquid is thought to be pure water. It is heated and boils steadily at 102 °C. Is it pure?
- Pure water boils at exactly 100 °C (the data-book value).
- The sample boils above 100 °C, so it is not pure — a dissolved solid (e.g. salt) raises the boiling point.
- For a fuller check you would look for a sharp boiling temperature: a pure liquid holds steady at one value, a mixture boils over a range.
- Using the everyday meaning. If a question says “in chemistry”, pure means a single element or compound — never “natural” or “nothing added”.
- Saying a mixture melts at one temperature. Mixtures melt and boil over a range; only pure substances have a sharp melting/boiling point.
- Forgetting impurities lower the melting point. This is why gritting salt melts ice and why a smeared melting range signals an impure drug.
🧪 Exam-style questions
Which two of these are mixtures? Tick (✓) two boxes, then press Check.
A solid is heated. It melts gradually between 68 °C and 75 °C. What does this show? Tick (✓) one box.
In chemistry, what is meant by a pure substance? Tick (✓) one box.
2Formulations
Some mixtures aren’t accidental — they are designed. A formulation is a mixture that has been put together as a useful product, with every component there for a reason and present in a carefully measured quantity so the final product has exactly the properties it needs.
A formulation is a mixture that has been designed as a useful product, in which:
- each component is included for a particular purpose, and
- the components are mixed in carefully controlled (measured) quantities so the product has the required properties.
The AQA examples to remember: fuels, cleaning agents, paints, medicines, alloys, fertilisers and foods.
Take paint — the example examiners use most. It is not one substance but a recipe, each part doing a job:
Each component has a purpose and a measured amount — change the recipe and you change the product. That is what makes paint a formulation, not just a mixture.
You are given information about a product and asked to decide whether it is a formulation. Look for both of these:
- it is a mixture of several components (often shown as a list or a percentage composition on a label), and
- each component is present in an exact, measured amount for a particular job.
A label reading “42% active ingredient, 31% solvent, 18% …” is the classic tell — precise proportions chosen to give required properties.
- A formulation is not a compound. The components are physically mixed, not chemically combined, so they keep their own properties and could in principle be separated.
- A single pure substance is never a formulation — pure copper, distilled water or oxygen are one substance each, not a designed mixture.
- You don’t need to know the named ingredients of any real branded product — only the idea of measured components with purposes.
🧪 Exam-style questions
Which of these is a formulation? Tick (✓) one box.
A weed-killer label lists: 36% active weed-killer, 4% wetting agent, 60% water. Which statement best explains why this product is a formulation? Tick (✓) one box.
Which statement about formulations is correct? Tick (✓) one box.
3Chromatography & Rf Values
Chromatography separates the components of a mixture and gives information that helps identify them — it is how you prove that “black” ink is really a blend of dyes, or that a food colouring contains a banned additive. It works because every chromatography method has two phases, and different substances spend different amounts of time in each.
- Mobile phase — the phase that moves. In paper chromatography this is the solvent (e.g. water or ethanol) that travels up the paper.
- Stationary phase — the phase that stays put. In paper chromatography this is the chromatography paper itself.
Separation depends on how each substance is distributed between the two phases. A component that is more soluble in the solvent (and less attracted to the paper) spends more time in the mobile phase and is carried further up the paper; a less soluble component lags behind. Because the components travel at different speeds, they spread out into separate spots.
How a paper chromatogram is made
A pencil start line (baseline) is drawn near the bottom of the paper and a small spot of each mixture is placed on it. The paper is stood in solvent so that the solvent level is below the start line. The solvent rises up the paper by capillary action, carries the soluble components with it, and is left to run until it is near the top. The result is a chromatogram: the developed paper with its separated spots.
Rf = d1d2 — the distance the substance moved over the distance the solvent moved. Always measure from the start line to the centre of the spot.
A single spot of three-dye “black” ink sits on the pencil start line, just above the solvent.
Pure or impure? What the spots tell you
A chromatogram is also a quick purity test:
- a pure compound produces a single spot — and it stays a single spot in every solvent;
- a mixture separates into two or more spots;
- two samples that are the same substance produce spots that travel the same distance (the same Rf) in the same solvent.
That last point is how chromatography identifies things: run your unknown next to a known reference and compare the spots.
Rf = distance moved by the substancedistance moved by the solvent
- Both distances are measured from the start line — to the centre of the spot, and to the solvent front.
- Rf is a ratio, so it has no units and is always less than 1.
- The same compound gives the same Rf in the same solvent; change the solvent and the Rf changes. Comparing to reference Rf values (measured under the same conditions) identifies an unknown.
Worked example. A spot moves 3.0 cm; the solvent moves 6.0 cm. Rf = 3.06.0 = 0.50.
Required practical 6: paper chromatography
Investigate how paper chromatography can be used to separate and identify a mixture of coloured substances (e.g. food colourings), and calculate Rf values.
The one rule the practical is built on: the start line and spots must sit above the solvent, or the samples wash straight off into the beaker.
- Draw a pencil start line about 2 cm from the bottom of the chromatography paper (pencil is insoluble, so it won’t run).
- Use a fresh capillary tube to put a small spot (2–3 mm) of each known colouring on the line, plus a spot of the unknown mixture. Label each in pencil.
- Add solvent to the beaker to a depth of less than 1 cm, so the level is below the start line. Hang the paper so its bottom edge dips in.
- Leave it undisturbed until the solvent has risen about three-quarters of the way up.
- Remove the paper, immediately mark the solvent front in pencil, and let it dry.
- Measure, in mm, the distance from the start line to the centre of each spot, and from the start line to the solvent front. Calculate Rf for each.
- Independent variable: the substance spotted on the start line (each different colouring or ink).
- Dependent variable: the distance each spot moves — used to work out its Rf value.
- Control variables: the same solvent, the same paper and the same start line for every sample; always measure to the centre of the spot.
- Compare each colouring’s Rf (and spot position) with the unknown to see which dyes the unknown contains.
- A colouring that gives a single spot is a pure substance; one that separates into several spots is a mixture.
- Rf values let you compare results between papers and labs — but only if the same solvent was used, since Rf depends on the solvent.
- Start line in ink. The ink dissolves and travels up with the samples, ruining the chromatogram. Always pencil.
- Solvent above the start line. The spots dissolve straight into the solvent and are lost.
- Measuring to the top of a spot. Always measure to the centre of the spot.
- An Rf greater than 1, or with units. Rf is a ratio of two distances — it has no units and must be less than 1. If yours is bigger than 1, you’ve divided the wrong way round.
🧪 Try the common mistakes
Run the chromatography three ways and watch what goes wrong with each mistake.
Correct setup: a pencil start line, solvent below the spots. Press play.
🧪 Exam-style questions
On a chromatogram, a colour moves 7.2 cm and the solvent moves 9.0 cm. Calculate the Rf value of the colour.
Show answer
- Rf = distance moved by substance ÷ distance moved by solvent = 7.2 ÷ 9.0 1 mark
- = 0.80 (no units, less than 1). 1 mark
A student set up their chromatography incorrectly. Which two of these are mistakes? Tick (✓) two boxes, then press Check.
On a chromatogram, a sample produces a single spot in two different solvents. What does this show? Tick (✓) one box.
A colour has an Rf value of 0.65 and moves 3.2 cm up the paper. Calculate the distance moved by the solvent.
Show answer
- Rearrange Rf = substance ÷ solvent → solvent = substance ÷ Rf = 3.2 ÷ 0.65 1 mark
- = 4.9 cm (4.92 cm). 1 mark
4Tests for the Common Gases
Four gases, four quick bench tests — this is the most reliably examined recall in C8, and the marks go to the exact method and the exact positive result. Learn each as a pair: what you do and what you see.
The four gas tests at a glance — each is a method paired with a single, unmistakable observation.
- Hydrogen (H2) — hold a lighted (burning) splint at the open end of the tube. Positive result: the gas burns rapidly with a squeaky “pop”.
- Oxygen (O2) — place a glowing splint inside the tube. Positive result: the splint relights.
- Carbon dioxide (CO2) — bubble the gas through limewater (calcium hydroxide solution). Positive result: the limewater turns milky / cloudy.
- Chlorine (Cl2) — hold damp litmus paper in the gas. Positive result: the litmus is bleached white.
- A ligHted splint tests for Hydrogen; a glOwing splint tests for Oxygen.
- Chlorine bleaches — think of chlorine bleach. The colour is removed, leaving the paper white.
- CO2 putting out a lit splint is not the test. It is a property, but other gases (e.g. nitrogen) do the same, so it isn’t definitive. Quote the limewater test.
- “Chlorine smells like a swimming pool” is not a test. It’s a characteristic, not an identification — and chlorine is toxic, so use damp litmus in a fume cupboard.
- State the result precisely: oxygen relights the splint (not just “glows brighter”); limewater goes milky/cloudy (not “white”); litmus is bleached/white (not “clear”).
🧪 Exam-style questions
What is the test for chlorine gas? Tick (✓) one box.
A gas relights a glowing splint. Which gas is it? Tick (✓) one box.
A student says “carbon dioxide is the gas that puts out a lighted splint”. Why is this not a good test for carbon dioxide? Tick (✓) one box.
Which test correctly identifies hydrogen? Tick (✓) one box.
5Flame Tests T
Sections 5–8 (identifying ions and instrumental methods) are Chemistry only — they are not on the Combined Science (Trilogy) papers. They are also the home of Required practical 7.
Heat certain metal ions (cations) strongly in a flame and they glow a characteristic colour — a quick way to identify the metal. You need five ions and their colours, and nothing else: flame colours of other metals are not required.
- Clean a nichrome (or platinum) wire loop by dipping it in dilute acid and holding it in a blue Bunsen flame until it adds no colour — this removes any ions left from before, so you test only your sample.
- Dip the clean loop in the sample and hold it in the edge of the blue flame.
- Observe the colour. Don’t let the wire glow red-hot — that orange glow can be mistaken for a flame colour.
- A strong colour masks a weaker one. In a mixture of ions, an intense colour (sodium’s yellow especially) can hide the others, so you may not see them. This is exactly why flame emission spectroscopy (section 8) is better for mixtures — it can pick out every ion at once.
- A flame test identifies one cation. To fully identify a compound you still need a test for the anion (section 7).
- Calcium (orange-red) is the only one of the five that also forms a white hydroxide precipitate; the flame test is a handy way to tell Ca2+ from Mg2+, which look identical with sodium hydroxide.
🧪 Exam-style questions
What colour flame do sodium ions produce in a flame test? Tick (✓) one box.
A metal compound gives a green flame. Which metal ion does it contain? Tick (✓) one box.
A salt contains both sodium and potassium ions. Why is it difficult to identify the potassium from a flame test? Tick (✓) one box.
Two white solids both give a white precipitate with sodium hydroxide. One gives an orange-red flame; the other gives no flame colour. Which ion does the orange-red solid contain? Tick (✓) one box.
6Testing Cations with Sodium Hydroxide T
Add sodium hydroxide solution to a solution of a metal salt and an insoluble metal hydroxide drops out as a precipitate. The colour of that precipitate — and whether it dissolves in excess — identifies the metal ion.
Three coloured precipitates (copper blue, iron(II) green, iron(III) brown) and three that are all white (aluminium, calcium, magnesium) — the whites need a second test to tell apart.
Add a few drops of NaOH first (slowly), then add it in excess:
- Copper(II), Cu2+ → blue precipitate.
- Iron(II), Fe2+ → green precipitate.
- Iron(III), Fe3+ → brown precipitate.
- Aluminium, Al3+ → white precipitate that dissolves in excess NaOH (giving a colourless solution).
- Calcium, Ca2+ → white precipitate, does not dissolve in excess.
- Magnesium, Mg2+ → white precipitate, does not dissolve in excess.
So of the three whites, only aluminium re-dissolves in excess. To separate the remaining two, use a flame test: calcium is orange-red, magnesium gives no colour.
You should be able to write balanced ionic equations for forming the insoluble hydroxides — it uses the same charge-balancing skill you met in C4. The metal ion’s charge fixes the number of OH–:
- Cu2+(aq) + 2OH–(aq) → Cu(OH)2(s)
- Fe2+(aq) + 2OH–(aq) → Fe(OH)2(s)
- Fe3+(aq) + 3OH–(aq) → Fe(OH)3(s)
- Al3+(aq) + 3OH–(aq) → Al(OH)3(s)
- Mg2+(aq) + 2OH–(aq) → Mg(OH)2(s)
- Ca2+(aq) + 2OH–(aq) → Ca(OH)2(s)
The pattern: a 2+ ion takes two OH–, a 3+ ion takes three — so the charges cancel and the formula is neutral.
- “Clear” when you mean “colourless”. A solution that loses its colour becomes colourless; “clear” only means see-through (copper sulfate solution is clear and blue). Examiners penalise “clear” here.
- Adding NaOH too fast. Add a few drops first — if you flood aluminium with excess straight away, the white precipitate forms and re-dissolves before you spot it.
- Over-writing the equation. You are not expected to write the equation for aluminium hydroxide dissolving (forming sodium aluminate) — only for forming the precipitates.
🧪 Exam-style questions
A few drops of sodium hydroxide solution are added to a metal salt solution and a brown precipitate forms. Which ion is present? Tick (✓) one box.
A white precipitate forms with sodium hydroxide solution, then dissolves when excess sodium hydroxide is added. Which ion is present? Tick (✓) one box.
Which is the correct balanced ionic equation for forming the iron(III) precipitate? Tick (✓) one box.
Two solutions both give a white precipitate with sodium hydroxide that does not dissolve in excess. What further test would tell them apart, and what result shows calcium? Tick (✓) one box.
7Testing for Anions T
With the cation found, the other half of a compound is the negative ion (anion). Three to know: carbonates, halides and sulfates. Each has its own reagent and its own unmistakable result.
| Anion | Add this… | Positive result |
|---|---|---|
| Carbonate (CO32–) | dilute acid | fizzes; the gas (CO2) turns limewater milky |
| Halide (Cl–, Br–, I–) | dilute nitric acid, then silver nitrate | silver halide precipitate — white (Cl), cream (Br) or yellow (I) |
| Sulfate (SO42–) | dilute hydrochloric acid, then barium chloride | white precipitate (barium sulfate) |
The silver halides step through white → cream → yellow as you go down the group (Cl → Br → I). Barium sulfate is a separate white precipitate, made with barium chloride.
- Carbonate: CO32–(aq) + 2H+(aq) → CO2(g) + H2O(l) — then CO2(g) + Ca(OH)2(aq) → CaCO3(s) + H2O(l) (the white solid that turns limewater milky).
- Halide: Ag+(aq) + X–(aq) → AgX(s) — e.g. Ag+ + Cl– → AgCl.
- Sulfate: Ba2+(aq) + SO42–(aq) → BaSO4(s).
Carbonate ions would also form a precipitate with silver nitrate or barium chloride, giving a false positive. Adding acid first reacts away any carbonate (it fizzes off as CO2) so it can’t interfere. The catch is choosing an acid that doesn’t add the very ion you’re testing for:
- Halide test → nitric acid (not hydrochloric — HCl would add chloride ions and give a false positive).
- Sulfate test → hydrochloric acid (not sulfuric — that would add sulfate ions).
Required practical 7: identifying the ions in an unknown salt T
Use chemical tests to identify the cation and the anion in unknown single ionic compounds, covering all the tests in sections 5–7. The skill is to combine them: one test names the metal, another names the non-metal, and together they name the salt.
- If the sample is solid, dissolve a little in distilled water. Work tidily — you’ll have several tubes and reagents, so a neat results table is essential.
- Find the cation: a flame test, and/or sodium hydroxide for the precipitate colour.
- Find the anion: carbonate (acid), then halide (nitric acid + silver nitrate) and/or sulfate (hydrochloric acid + barium chloride) as needed.
- You won’t need every test on every sample — choose the ones that narrow it down. Repeat any test that gives an unclear result.
This practical is largely qualitative — you record what you see rather than measure a quantity.
- Independent variable: which chemical test you apply (flame test, sodium hydroxide, carbonate, halide or sulfate test).
- Dependent variable (what you observe): the colour, precipitate or gas produced — the result that names the ion.
- Control variables: test a fresh separate portion of the sample for each test; acidify before the halide and sulfate precipitation tests to remove carbonate ions that would give a false positive.
A white solid gives these results:
| Test | Observation | Conclusion |
|---|---|---|
| Flame test | yellow flame | sodium, Na+ |
| Dilute nitric acid, then silver nitrate | cream precipitate | bromide, Br– |
Cation Na+ (1+) and anion Br– (1–) balance one-to-one, so the salt is sodium bromide, NaBr. (To get a formula, always balance the charges: a 2+ ion with a 2– ion is 1:1; a 2+ ion with 1– ions needs two of them, e.g. CaCl2.)
🧪 Identify the unknown salt
You’re handed an unknown ionic compound. Run bench tests to find its cation (metal ion) and anion (non-metal ion), then name it — in as few tests as you can. For each result, try to recall what it tells you before you tap Reveal. This is exactly the Required practical 7 skill.
Pick a test to start.
| Test | Observation | What does it tell you? |
|---|---|---|
| No tests run yet — click a test above. | ||
- Silver bromide is “cream”, not yellow. Yellow is silver iodide. Cream is the trickiest colour to name — learn the white → cream → yellow sequence.
- Wrong acid. Nitric acid for the halide test (HCl adds chloride); hydrochloric acid for the sulfate test (sulfuric adds sulfate).
- Forgetting to acidify. Skipping the acid lets carbonate ions give a false precipitate.
- Letting the gas escape. In the carbonate test, connect to the limewater quickly so the CO2 doesn’t escape before it can turn it milky.
🧪 Exam-style questions
Dilute nitric acid and silver nitrate solution are added to a salt solution. A cream precipitate forms. Which ion is present? Tick (✓) one box.
Describe how you would test a solution for sulfate ions, and give the result you would see.
Show answer
- Add dilute hydrochloric acid to the solution. 1 mark (Allow: any reference to acidifying first to remove carbonate ions.)
- Then add barium chloride solution. 1 mark
- A white precipitate (of barium sulfate) forms. 1 mark
Do not accept: dilute sulfuric acid (it adds sulfate ions), or silver nitrate (that is the halide test).
Why is dilute nitric acid added before the silver nitrate in the test for halide ions? Tick (✓) one box.
A salt gives no flame colour, a green precipitate with sodium hydroxide, and a white precipitate when dilute hydrochloric acid then barium chloride are added. What is the salt? Tick (✓) one box.
A chemist needs to find the concentration of a solution of barium hydroxide (an alkali). It could be found by one of two methods.
Method 1: an excess of sodium sulfate solution is added to 25 cm³ of the barium hydroxide solution; a precipitate of barium sulfate forms, which is filtered, dried and weighed; the concentration is calculated from the mass of barium sulfate produced.
Method 2: 25 cm³ of the barium hydroxide solution is titrated with hydrochloric acid of known concentration; the concentration is calculated from the titration result.
Compare the advantages and disadvantages of the two methods.
Show answer
Any five from the points below (1 mark each, maximum 5; converse arguments are allowed):
Method 1 — precipitate and weigh:
- weighing is accurate
- not all the barium sulfate may be precipitated (allow: not all the barium hydroxide has reacted)
- some precipitate may be lost (e.g. when filtering)
- the precipitate may not be (fully) dry
- it takes longer
- it requires energy (to dry the precipitate)
Method 2 — titration:
- it is accurate
- it works for low concentrations (allow: reliable / precise)
Source: AQA GCSE Chemistry.
8Instrumental Methods & Flame Emission T
The bench tests in this topic are cheap and quick, but modern labs mostly use instruments — machines that detect and identify elements and compounds automatically. They matter most when the sample is tiny or the answer must be exact, which is why forensic and drug-control scientists rely on them.
Compared with the chemical tests in this topic, instruments are:
- more accurate,
- more sensitive — they work with very small samples,
- faster, and able to run many samples automatically.
The exam phrasing to reach for is the trio accurate, sensitive and rapid.
Flame emission spectroscopy
This is the instrumental method AQA names — an example used to analyse metal ions in solution. The sample is put into a flame; the hot ions emit light; that light is passed through a spectroscope, which spreads it into a line spectrum. Because every metal ion gives its own pattern of lines, the output identifies the ions present — and the intensity of the lines measures their concentration.
You don’t need to know any element’s lines — just how to compare. The unknown’s line sits exactly under sodium’s, so the sample contains sodium ions.
🧪 Flame-emission spectroscope
The flame makes the sample’s metal ions emit light; the spectroscope spreads it into a line spectrum. Match the sample’s lines against the references to identify the ion(s) — and see how a mixture shows every ion’s lines at once, which a flame test cannot separate.
A flame test looks at one ion at a time and a strong colour can mask others. Flame emission spectroscopy gives a full line spectrum, so it can identify several ions in a mixture at once and measure how much of each is present — faster, more sensitive and more accurate than the naked eye.
🧪 Exam-style questions
Give one advantage of using an instrumental method rather than a chemical test. Tick (✓) one box.
What is one advantage of flame emission spectroscopy over a flame test? Tick (✓) one box.
An unknown solution gives a line spectrum whose lines match exactly the reference spectrum for potassium. What does this show? Tick (✓) one box.
Besides identifying which metal ions are present, what else can flame emission spectroscopy measure? Tick (✓) one box.
★Capstone: Identify the Unknown T
The whole of this Triple-only run of tests builds to one skill: read a result and name the ion that made it. Each observation below is the unmistakable signature of exactly one species — drag or tap each one into the box for the ion it identifies.
Drag or tap each observation into the ion it identifies, then press Check.
- Pure substances & mixtures — a pure substance (single element or compound) melts and boils at one sharp temperature; a mixture melts/boils over a range, and impurities lower the melting point.
- Formulations — a mixture designed as a useful product, each component in a measured amount for a purpose (fuels, paints, alloys, medicines, fertilisers, foods).
- Chromatography & Rf — mobile phase (solvent) + stationary phase (paper); the more soluble a substance, the further it travels. Rf = distance moved by substance ÷ distance moved by solvent (no units, < 1). A pure substance gives one spot in every solvent. (Required practical 6)
- Gas tests — hydrogen: lit splint → squeaky pop; oxygen: glowing splint relights; carbon dioxide: limewater turns milky; chlorine: damp litmus bleached white.
- Flame tests T — lithium crimson, sodium yellow, potassium lilac, calcium orange-red, copper green; a strong colour can mask a weaker one.
- Cations with NaOH T — Cu2+ blue, Fe2+ green, Fe3+ brown; Al3+/Ca2+/Mg2+ white (only aluminium redissolves in excess). A 2+ ion takes two OH–, a 3+ ion takes three.
- Anion tests T — carbonate + dilute acid → CO2 (limewater milky); halide + nitric acid + silver nitrate → AgCl white / AgBr cream / AgI yellow; sulfate + hydrochloric acid + barium chloride → white BaSO4. (Required practical 7)
- Instrumental methods T — accurate, sensitive and rapid; flame emission spectroscopy gives a line spectrum that identifies metal ions (even several in a mixture at once) and measures their concentration.
That completes the analyst’s toolkit: melting point for purity, chromatography for the components of a mixture, a splint or a drop of limewater for a gas, and a flame or a precipitate for an ion — with instrumental methods waiting in the modern lab when speed, sensitivity and tiny samples matter. For Paper 2, pair these notes with C7 Organic Chemistry (crude oil is the classic mixture, separated by physical means) and carry the same gas-testing skills into C9 Chemistry of the Atmosphere (where carbon dioxide and the products of combustion take centre stage). The ionic equations and charge-balanced formulae behind every precipitate are set out in C4 Chemical Changes.