Testing for Ions: Flame Tests, Hydroxides, Halides and Sulfates

(Separate Chemistry only) The identification of ions by chemical means in section 4.8.3 is assessed only in AQA GCSE Chemistry (8462). Students taking Combined Science do not need it.

Identifying ions means finding which positive ion (cation) and which negative ion (anion) a compound contains. This lesson covers the chemical tests: flame tests and metal hydroxide precipitates for cations, and carbonate, halide and sulfate tests for anions. Together they let you work out the identity of an unknown ionic compound, which is the aim of Required practical 7.

Throughout, the standard warning about mixtures applies: these tests are most reliable on a single unknown compound, because a mixture of ions can hide or confuse individual results.

A flame test identifies some metal ions (cations) by the colour they give to a flame. Dip a clean nichrome wire loop in the sample, hold it in a blue Bunsen flame, and note the colour.

You must know five colours exactly:

Metal ion	Flame colour
Lithium (Li⁺)	Crimson (red)
Sodium (Na⁺)	Yellow
Potassium (K⁺)	Lilac (purple)
Calcium (Ca²⁺)	Orange-red
Copper (Cu²⁺)	Green

A mixture of ions may mask some colours: the strong yellow of sodium in particular can hide the flame colours of other ions present. This is one reason chemical tests can be less reliable than instrumental methods.

The wire must be clean each time (dip it in acid and reheat until no colour shows) so that a leftover ion from a previous test does not give a false result. Flame colours for metal ions other than these five are not required.

Adding sodium hydroxide solution to a solution of a metal salt can produce a coloured, insoluble metal hydroxide precipitate, which identifies the cation.

Metal ion	Precipitate colour with NaOH
Aluminium (Al³⁺)	White (dissolves in excess NaOH)
Calcium (Ca²⁺)	White
Magnesium (Mg²⁺)	White
Copper(II) (Cu²⁺)	Blue
Iron(II) (Fe²⁺)	Green
Iron(III) (Fe³⁺)	Brown

Three ions (Al³⁺, Ca²⁺, Mg²⁺) all give a white precipitate, so you need a second step to tell them apart. Add excess sodium hydroxide: only the aluminium hydroxide precipitate dissolves again. Calcium and magnesium can then be separated by a flame test (calcium gives orange-red; magnesium gives no flame colour).

Only aluminium hydroxide redissolves in excess sodium hydroxide. Calcium and magnesium hydroxides stay as white solids. (Equations for the sodium aluminate formed when it redissolves are not required.)

The spec requires you to write balanced equations for the reactions that produce the insoluble hydroxides. The metal ion combines with hydroxide ions; the number of hydroxide ions equals the charge on the metal ion so that charges balance.

Full ionic equations:

$C{a}^{2+}(aq)+2O{H}^{-}(aq)\to Ca(OH{)}_2(s)$

$M{g}^{2+}(aq)+2O{H}^{-}(aq)\to Mg(OH{)}_2(s)$

$A{l}^{3+}(aq)+3O{H}^{-}(aq)\to Al(OH{)}_3(s)$

$C{u}^{2+}(aq)+2O{H}^{-}(aq)\to Cu(OH{)}_2(s)$

$F{e}^{2+}(aq)+2O{H}^{-}(aq)\to Fe(OH{)}_2(s)$

$F{e}^{3+}(aq)+3O{H}^{-}(aq)\to Fe(OH{)}_3(s)$

Check the charge balance on iron(III): the ion is 3+, so it needs three OH⁻ (each 1−), giving a neutral Fe(OH)₃. Iron(II) is only 2+, so it takes two OH⁻. Getting the number of hydroxide ions to match the metal-ion charge is what balances the equation.

Three anions are identified by producing a gas or a coloured precipitate. Learn the reagent and the result for each.

Carbonates (CO₃²⁻) react with dilute acid to give carbon dioxide, which turns limewater milky. For example:

$CaC{O}_3(s)+2HCl(aq)\to CaC{l}_2(aq)+H_{2}O(l)+C{O}_2(g)$

Halides (Cl⁻, Br⁻, I⁻) in solution give a precipitate with silver nitrate solution in the presence of dilute nitric acid. For chloride:

$A{g}^{+}(aq)+C{l}^{-}(aq)\to AgCl(s)$

Halide ion	Precipitate with silver nitrate	Colour
Chloride (Cl⁻)	Silver chloride, AgCl	White
Bromide (Br⁻)	Silver bromide, AgBr	Cream
Iodide (I⁻)	Silver iodide, AgI	Yellow

Sulfates (SO₄²⁻) in solution give a white precipitate with barium chloride solution in the presence of dilute hydrochloric acid:

$B{a}^{2+}(aq)+S{O}_4^{2-}(aq)\to BaS{O}_4(s)$

The dilute acid in the halide and sulfate tests is not optional. It removes ions that would otherwise give a false positive.

Add the acid before the test reagent. It reacts away carbonate ions, which would themselves form a precipitate with silver nitrate or barium chloride and be mistaken for a halide or sulfate.

• In the sulfate test, dilute hydrochloric acid is used, not sulfuric acid, because sulfuric acid would add sulfate ions and guarantee a white precipitate whatever the sample was.
• In the halide test, dilute nitric acid is used, not hydrochloric acid, because hydrochloric acid would add chloride ions and give a false white precipitate with the silver nitrate.

Choosing the wrong acid ruins the test, so match the acid to the anion you are looking for: nitric acid for halides, hydrochloric acid for sulfates.

Required practical 7 uses these chemical tests to identify the ions in unknown single ionic compounds, working through cations and then anions.

A sensible order to work in:

1. Cation, flame test: clean the wire, hold the sample in the flame, and record the colour against the five known colours.
2. Cation, sodium hydroxide: add NaOH to a fresh sample and note the precipitate colour; add excess to check whether a white precipitate redissolves (aluminium).
3. Anion, carbonate: add dilute acid; fizzing that turns limewater milky shows a carbonate.
4. Anion, sulfate: add dilute hydrochloric acid then barium chloride; a white precipitate shows sulfate.
5. Anion, halide: add dilute nitric acid then silver nitrate; white, cream or yellow shows chloride, bromide or iodide.

Variable	In this practical
Independent variable	The unknown compound being tested
Dependent variable	The observations (flame colour, precipitate colour, gas produced)
Control variables	Same volumes/concentrations of reagents, clean apparatus each time

Cleaning apparatus between tests and doing the anion tests on fresh samples matters, so that a reagent left from an earlier test does not contaminate the next observation.

1. Mixing up the flame colours

Lithium is crimson, sodium yellow, potassium lilac, calcium orange-red, copper green. Lithium (crimson) and calcium (orange-red) are the pair most often confused; learn all five precisely.

2. Not using excess NaOH to separate the white precipitates

Aluminium, calcium and magnesium all give white precipitates with sodium hydroxide. You must add excess to see that only aluminium hydroxide redissolves.

3. Getting the hydroxide equations unbalanced

The number of OH⁻ ions must equal the charge on the metal ion: two for 2+ ions (Ca²⁺, Mg²⁺, Cu²⁺, Fe²⁺) and three for 3+ ions (Al³⁺, Fe³⁺). Check the charges balance to zero.

4. Using the wrong acid before silver nitrate or barium chloride

Use nitric acid before silver nitrate (halides) and hydrochloric acid before barium chloride (sulfates). Using hydrochloric acid before silver nitrate, or sulfuric acid before barium chloride, gives a false positive.

5. Confusing the silver halide colours

Silver chloride is white, silver bromide is cream, silver iodide is yellow. Cream sits between white and yellow; describe it as cream, not "pale yellow" or "off-white".