Qualitative Analysis

AS · 13 min

One question in every Paper 3 is an observational problem: you are given one or more unknown substances, told which tests to carry out, and asked to record what you see and deduce what the substances are. The Qualitative analysis notes, printed in the exam paper, list the tests for cations, anions, gases and iodine. This note sets out those tests with the reasoning behind them, the ionic equations that explain each observation, how to record observations in the way examiners reward, and a systematic strategy for identifying unknowns. The tables match the syllabus notes so that you can practise using them exactly as you will in the exam.

Good practice

The syllabus lists the methods you are expected to follow:

treat all unknown materials with caution;
use an appropriate (small) quantity of the material: about $1\ \text{cm}^3$ depth of solution in a test-tube, or a small spatula-tip of solid;
add only the specified amount of reagent;
work safely, for example using a test-tube holder when heating a solid in a hard-glass test-tube;
record all observations, even "no change" or "remains a colourless solution";
use excess alkali when a precipitate forms with $\ce{NaOH(aq)}$ or $\ce{NH3(aq)}$ , to see whether it dissolves;
identify any gas given off as effervescence.

Definition

A precipitate (abbreviated ppt.) is an insoluble solid that forms when two solutions are mixed. Effervescence is the bubbling seen when a gas is produced in a liquid.

Reactions of aqueous cations

Most metal hydroxides are insoluble, so adding hydroxide ions to a solution of a metal ion gives a precipitate of the hydroxide:

\ce{M^{n+}(aq) + nOH-(aq) -> M(OH)_n(s)}

Aqueous ammonia also contains $\ce{OH-}$ ions ( $\ce{NH3 + H2O <=> NH4+ + OH-}$ ), but at a lower concentration, and ammonia molecules can form soluble complexes with some metal ions. So comparing $\ce{NaOH(aq)}$ and $\ce{NH3(aq)}$ , each added first dropwise and then in excess, distinguishes the ions.

cation	with $\ce{NaOH(aq)}$	with $\ce{NH3(aq)}$
aluminium, $\ce{Al^3+}$	white ppt., soluble in excess	white ppt., insoluble in excess
ammonium, $\ce{NH4+}$	no ppt.; ammonia produced on warming	–
barium, $\ce{Ba^2+}$	faint white ppt. is observed unless $\ce{[Ba^2+]}$ is very low	no ppt.
calcium, $\ce{Ca^2+}$	white ppt. unless $\ce{[Ca^2+]}$ is very low	no ppt.
chromium(III), $\ce{Cr^3+}$	grey-green ppt., soluble in excess giving dark green solution	grey-green ppt., insoluble in excess
copper(II), $\ce{Cu^2+}$	pale blue ppt., insoluble in excess	pale blue ppt., soluble in excess giving dark blue solution
iron(II), $\ce{Fe^2+}$	green ppt., turning brown on contact with air; insoluble in excess	green ppt., turning brown on contact with air; insoluble in excess
iron(III), $\ce{Fe^3+}$	red-brown ppt., insoluble in excess	red-brown ppt., insoluble in excess
magnesium, $\ce{Mg^2+}$	white ppt., insoluble in excess	white ppt., insoluble in excess
manganese(II), $\ce{Mn^2+}$	off-white ppt., rapidly turning brown on contact with air; insoluble in excess	off-white ppt., rapidly turning brown on contact with air; insoluble in excess
zinc, $\ce{Zn^2+}$	white ppt., soluble in excess	white ppt., soluble in excess

Why the results differ

Precipitates. $\ce{Al^3+ + 3OH- -> Al(OH)3}$ ; $\ce{Cu^2+ + 2OH- -> Cu(OH)2}$ (pale blue); $\ce{Fe^2+ + 2OH- -> Fe(OH)2}$ (green); $\ce{Fe^3+ + 3OH- -> Fe(OH)3}$ (red-brown); $\ce{Mg^2+ + 2OH- -> Mg(OH)2}$ .
Amphoteric hydroxides dissolve in excess $\ce{NaOH}$ . Aluminium, zinc and chromium(III) hydroxides react as acids with excess hydroxide: $\ce{Al(OH)3(s) + OH-(aq) -> [Al(OH)4]-(aq)}$ and $\ce{Zn(OH)2(s) + 2OH-(aq) -> [Zn(OH)4]^2-(aq)}$ .
Ammonia complexes. Copper(II) and zinc hydroxides dissolve in excess ammonia because ammonia molecules form soluble complex ions with the metal ions (dark blue for copper). Aluminium hydroxide does not. This is how $\ce{Al^3+}$ and $\ce{Zn^2+}$ are told apart.
Group 2 solubility. Calcium hydroxide is slightly soluble and barium hydroxide fairly soluble, so with $\ce{NaOH(aq)}$ they give only a white or faint white precipitate, and none at all at low concentration. With $\ce{NH3(aq)}$ the $\ce{[OH-]}$ is too low to precipitate them. Magnesium hydroxide is very insoluble and precipitates with both.
Oxidation in air. Iron(II) hydroxide (green) and manganese(II) hydroxide (off-white) are oxidised by oxygen in the air to brown iron(III) and manganese compounds, so the precipitate darkens at the surface.
Ammonium. $\ce{NH4+ + OH- -> NH3 + H2O}$ on warming: no precipitate, but ammonia gas.

Tip

Two pairs need care. $\ce{Al^3+}$ or $\ce{Zn^2+}$ : both give white precipitates soluble in excess $\ce{NaOH}$ ; only zinc's dissolves in excess $\ce{NH3}$ . $\ce{Mg^2+}$ or $\ce{Ca^2+}$ : both give white precipitates with $\ce{NaOH}$ ; only magnesium gives one with $\ce{NH3}$ .

Reactions of anions

anion	test and result
carbonate, $\ce{CO3^2-}$	$\ce{CO2}$ liberated by dilute acids (effervescence; gas turns limewater milky)
chloride, $\ce{Cl-(aq)}$	white ppt. with $\ce{Ag+(aq)}$ , soluble in $\ce{NH3(aq)}$
bromide, $\ce{Br-(aq)}$	cream / off-white ppt. with $\ce{Ag+(aq)}$ , partially soluble in $\ce{NH3(aq)}$
iodide, $\ce{I-(aq)}$	pale yellow ppt. with $\ce{Ag+(aq)}$ , insoluble in $\ce{NH3(aq)}$
nitrate, $\ce{NO3-(aq)}$	$\ce{NH3}$ liberated on heating with $\ce{OH-(aq)}$ and Al foil
nitrite, $\ce{NO2-(aq)}$	$\ce{NH3}$ liberated on heating with $\ce{OH-(aq)}$ and Al foil; decolourises acidified aqueous $\ce{KMnO4}$
sulfate, $\ce{SO4^2-(aq)}$	white ppt. with $\ce{Ba^2+(aq)}$ , insoluble in excess dilute strong acids; white ppt. with high $\ce{[Ca^2+(aq)]}$
sulfite, $\ce{SO3^2-(aq)}$	white ppt. with $\ce{Ba^2+(aq)}$ , soluble in excess dilute strong acids; decolourises acidified aqueous $\ce{KMnO4}$
thiosulfate, $\ce{S2O3^2-(aq)}$	off-white / pale yellow ppt. slowly with $\ce{H+}$

The chemistry behind each test:

Carbonate: $\ce{CO3^2- + 2H+(aq) -> CO2(g) + H2O(l)}$ (the carbonate may be a solid or in solution).
Halides: $\ce{Ag+(aq) + X-(aq) -> AgX(s)}$ ; see Reactions of the halide ions for why ammonia distinguishes them.
Nitrate and nitrite are reduced to ammonia by aluminium in alkaline solution: $\ce{3NO3- + 8Al + 5OH- + 18H2O -> 3NH3 + 8[Al(OH)4]-}$ . Ammonium ions must be absent (or removed by warming with $\ce{NaOH}$ first), otherwise they give ammonia anyway.
Nitrite and sulfite are reducing agents, so they decolourise purple manganate(VII): for example $\ce{2MnO4- + 5SO3^2- + 6H+ -> 2Mn^2+ + 5SO4^2- + 3H2O}$ . Nitrate and sulfate do not.
Sulfate: $\ce{Ba^2+(aq) + SO4^2-(aq) -> BaSO4(s)}$ ; barium sulfate does not dissolve in acid.
Sulfite: $\ce{Ba^2+(aq) + SO3^2-(aq) -> BaSO3(s)}$ , but this dissolves in acid: $\ce{BaSO3(s) + 2H+(aq) -> Ba^2+(aq) + SO2(g) + H2O(l)}$ .
Thiosulfate: $\ce{S2O3^2-(aq) + 2H+(aq) -> S(s) + SO2(g) + H2O(l)}$ ; the sulfur forms a slowly thickening pale yellow cloudiness (the basis of the rate experiment).

Tests for gases and for iodine

gas	test and result
ammonia, $\ce{NH3}$	turns damp red litmus paper blue
carbon dioxide, $\ce{CO2}$	gives a white ppt. with limewater (limewater turns milky)
hydrogen, $\ce{H2}$	"pops" with a lighted splint
oxygen, $\ce{O2}$	relights a glowing splint

element	test and result
iodine, $\ce{I2}$	gives a blue-black colour on addition of starch solution

Other gases may be produced in an exam (for example sulfur dioxide or nitrogen dioxide). You are not expected to have a named test for them, but you should describe what you see (colour, smell, effect on damp litmus) and use any reagent the question supplies, such as acidified $\ce{KMnO4}$ , which $\ce{SO2}$ decolourises.

Organic tests that may appear

The syllabus also lists four organic tests you may carry out or interpret in Paper 3. They are taught fully in the organic notes; the positive results are:

reagent	positive result	indicates
Fehling's reagent (warm)	orange-red ppt.	aldehyde
Tollens' reagent (warm)	silver mirror / black ppt.	aldehyde
alkaline aqueous iodine	yellow ppt.	$\ce{CH3CO}$ or $\ce{CH3CH(OH)}$ group
acidified potassium manganate(VII)	purple to colourless	a compound that can be oxidised

Recording observations

The syllabus asks for observations in simple, precise language, with "pale" or "dark" for fine distinctions.

Key result

How to write observations that score

Say what you see, not what you conclude: "white ppt." not "aluminium hydroxide formed"; "effervescence" not "carbon dioxide produced".
For every precipitate, state its colour and then what happens in excess reagent: "white ppt., soluble in excess giving a colourless solution".
Record "no change" or "no ppt." when nothing happens. A blank cell scores nothing.
For a gas, state the observation and the test: "effervescence; gas turns limewater milky".
Use the syllabus colour words: white, cream, pale yellow, pale blue, dark blue, green, grey-green, red-brown, off-white, brown, blue-black.

Record results in a table with the test in the first column and observations for each unknown in the next columns, drawn before you start testing.

A strategy for identifying an unknown

Method

Identifying a salt

Note the appearance: colour of solid or solution. Blue/green solutions suggest $\ce{Cu^2+}$ , $\ce{Fe^2+}$ or $\ce{Cr^3+}$ ; yellow-brown suggests $\ce{Fe^3+}$ ; colourless suggests $\ce{Al^3+}$ , $\ce{Zn^2+}$ , Group 1, Group 2 or $\ce{NH4+}$ .
Cation: add $\ce{NaOH(aq)}$ dropwise then in excess; warm if there is no precipitate (ammonium). Then repeat with $\ce{NH3(aq)}$ on a fresh sample. Compare with the cation table.
Anion: add dilute acid (effervescence suggests carbonate, or sulfite with a choking gas). Then test separate samples with acidified $\ce{AgNO3(aq)}$ (halides), with $\ce{Ba^2+(aq)}$ and then dilute acid (sulfate or sulfite), and with acidified $\ce{KMnO4}$ (reducing anions: nitrite, sulfite).
Confirm with a second, independent test where possible, and write the formula of the salt with charges balanced.

Worked examples

Aluminium or zinc

A colourless solution gives a white precipitate with aqueous sodium hydroxide that dissolves in excess. With aqueous ammonia, it gives a white precipitate that does not dissolve in excess. Identify the cation and write ionic equations for the reactions with sodium hydroxide.

Solution

White precipitate soluble in excess $\ce{NaOH}$ : $\ce{Al^3+}$ or $\ce{Zn^2+}$ . Insoluble in excess $\ce{NH3}$ : $\ce{Al^3+}$ (zinc's would dissolve).

$\ce{Al^3+(aq) + 3OH-(aq) -> Al(OH)3(s)}$

$\ce{Al(OH)3(s) + OH-(aq) -> [Al(OH)4]-(aq)}$

Choosing reagents to distinguish two ions

Suggest one reagent that would distinguish between solutions of magnesium chloride and calcium chloride, and give the observations for each.

Solution

Aqueous ammonia.

Magnesium chloride: white precipitate (of $\ce{Mg(OH)2}$ ), insoluble in excess.

Calcium chloride: no precipitate.

(Aqueous NaOH would not work well: both give white precipitates.)

Identifying an iron salt

A pale green solution P is tested.

Test 1: $\ce{NaOH(aq)}$ added: green precipitate, insoluble in excess, turning brown at the surface on standing.

Test 2: dilute HCl then $\ce{BaCl2(aq)}$ added: white precipitate.

Identify P and write ionic equations for both precipitates. Explain the brown colour.

Solution

Test 1: green precipitate turning brown in air: $\ce{Fe^2+}$ . $\ce{Fe^2+(aq) + 2OH-(aq) -> Fe(OH)2(s)}$

Test 2: white precipitate with $\ce{Ba^2+}$ in the presence of acid: $\ce{SO4^2-}$ . $\ce{Ba^2+(aq) + SO4^2-(aq) -> BaSO4(s)}$

P is iron(II) sulfate, $\ce{FeSO4}$ .

The brown colour is iron(III) hydroxide (or hydrated iron(III) oxide), formed as oxygen in the air oxidises iron(II) to iron(III) at the surface of the precipitate.

Exam-style: nitrate or nitrite

Solutions Q and R each contain a sodium salt. When each is heated with aqueous sodium hydroxide and aluminium foil, a gas is given off that turns damp red litmus blue. When acidified aqueous potassium manganate(VII) is added, Q decolourises it but R does not. Identify the anions in Q and R and explain the reasoning.

Solution

The gas is ammonia. Ammonia from heating with $\ce{OH-}$ and Al foil indicates nitrate or nitrite (the salts are sodium salts, so ammonium ions are absent).

Q decolourises acidified $\ce{KMnO4}$ , so it contains a reducing anion: nitrite, $\ce{NO2-}$ .

R does not, so it contains nitrate, $\ce{NO3-}$ (nitrogen in nitrate is already at $+5$ and cannot be oxidised further).

Exam-hard: a full unknown

A white solid S dissolves in water to give a colourless solution. The following observations are made.

test	observation
add $\ce{NaOH(aq)}$ , then warm	no ppt.; on warming, a gas is given off that turns damp red litmus blue
add $\ce{Ba(NO3)2(aq)}$	white ppt.
add dilute $\ce{HNO3}$ to the mixture from the previous test	ppt. dissolves; effervescence; the gas decolourises acidified $\ce{KMnO4}$ on a filter paper
add acidified $\ce{KMnO4(aq)}$ to a fresh solution of S	purple colour disappears

Identify S, explaining each step, and write ionic equations for three of the reactions.

Solution

No precipitate with $\ce{NaOH}$ , ammonia on warming: cation is ammonium, $\ce{NH4+}$ . $\ce{NH4+ + OH- -> NH3 + H2O}$
White precipitate with $\ce{Ba^2+}$ : sulfate or sulfite (or carbonate).
Precipitate dissolves in acid with a gas: not sulfate (barium sulfate is insoluble in acid). The gas decolourises $\ce{KMnO4}$ , so it is a reducing gas, $\ce{SO2}$ , not $\ce{CO2}$ : the anion is sulfite, $\ce{SO3^2-}$ . $\ce{Ba^2+ + SO3^2- -> BaSO3}$ , then $\ce{BaSO3 + 2H+ -> Ba^2+ + SO2 + H2O}$ .
Decolourising acidified $\ce{KMnO4}$ confirms a reducing anion (sulfite).

S is ammonium sulfite, $\ce{(NH4)2SO3}$ .

Watch out

Do not write "solution turns white". A solution cannot be white; a precipitate can be white, and a solution can be colourless or cloudy.
"Clear" does not mean "colourless". A clear solution can be blue. Use "colourless" for no colour.
Always add excess reagent and record the result. "White ppt." for $\ce{Al^3+}$ earns less than "white ppt., soluble in excess".
Acidify with nitric acid before silver nitrate, and with hydrochloric or nitric acid before barium ions, never with sulfuric acid (it adds sulfate).
Ammonia from a solution with $\ce{NaOH}$ without aluminium means ammonium; with Al foil after removing ammonium, it means nitrate or nitrite.

Exam tip

Observation marks are lost for conclusions written in the observation column and for vague words ("reacts", "fizzes a bit", "goes cloudy"). Write "effervescence", "white ppt.", "no change".
Conclusions must follow from your observations. If your observation does not match the notes exactly, deduce what you can and say what the evidence suggests; examiners give credit for logical deductions from your own results.
When the question asks for a test to distinguish two ions, name the reagent and give the observation for each ion.
If asked for a confirmatory test, choose a different reagent from the one already used (for example $\ce{NH3(aq)}$ to confirm $\ce{Cu^2+}$ after $\ce{NaOH(aq)}$ ).

Practical skills

Practical details that make the tests work:

Add reagents dropwise at first, shaking after each addition, so you see the precipitate form before it redissolves; then add excess (several $\text{cm}^3$ , up to about half-filling the tube).
To test a gas, hold damp litmus at the mouth of the tube; bubble a suspected $\ce{CO2}$ through limewater using a delivery tube, or draw it into a dropping pipette and bubble it through limewater.
For a mixture of anions, use separate portions of the solution for each test so that one reagent does not interfere with another.
When heating a solid, use a hard-glass tube held in a test-tube holder, pointing away from people, and heat gently at first.

Summary

Cations: test with $\ce{NaOH(aq)}$ and $\ce{NH3(aq)}$ , dropwise then in excess; learn the colours and solubilities in the table.
Soluble in excess $\ce{NaOH}$ : $\ce{Al^3+}$ , $\ce{Zn^2+}$ , $\ce{Cr^3+}$ (amphoteric hydroxides). Soluble in excess $\ce{NH3}$ : $\ce{Cu^2+}$ (dark blue), $\ce{Zn^2+}$ .
$\ce{Fe^2+}$ green and $\ce{Mn^2+}$ off-white precipitates turn brown in air; $\ce{Fe^3+}$ is red-brown.
Anions: carbonate (acid, $\ce{CO2}$ ); halides ( $\ce{AgNO3}$ then $\ce{NH3}$ ); sulfate and sulfite ( $\ce{Ba^2+}$ , then acid); nitrate and nitrite ( $\ce{OH-}$ + Al foil gives $\ce{NH3}$ ); nitrite and sulfite decolourise $\ce{KMnO4}$ ; thiosulfate gives pale yellow sulfur with acid.
Gases: $\ce{NH3}$ damp red litmus blue; $\ce{CO2}$ limewater milky; $\ce{H2}$ pops; $\ce{O2}$ relights glowing splint. Iodine: blue-black with starch.
Record every observation precisely, including "no change".

Practice

Question

State what you would see when aqueous ammonia is added dropwise, then in excess, to a solution containing copper(II) ions.
Give one reagent that distinguishes between $\ce{Fe^2+(aq)}$ and $\ce{Fe^3+(aq)}$ , with the observations for each.
Describe a test to show that a solution contains sulfate ions, and explain why the acid is added.
Write an ionic equation for the reaction of zinc hydroxide with excess aqueous sodium hydroxide.
A gas relights a glowing splint. Identify it, and name a reaction in this course that produces it together with a brown gas.
Suggest how you could distinguish between solutions of sodium sulfate and sodium sulfite using barium nitrate solution and dilute nitric acid.
A colourless solution gives no precipitate with $\ce{NaOH(aq)}$ , even on adding excess, and no gas on warming. With acidified silver nitrate it gives a white precipitate soluble in aqueous ammonia. Suggest what the salt might be, and explain why you cannot identify the cation fully from these tests.
Explain why calcium ions give a white precipitate with aqueous sodium hydroxide but not with aqueous ammonia.
A solid T is a single salt. Its solution is green. With $\ce{NaOH(aq)}$ it gives a grey-green precipitate which dissolves in excess to give a dark green solution. With acidified silver nitrate it gives a pale yellow precipitate insoluble in concentrated ammonia. Identify T, and write ionic equations for the two precipitation reactions.
Three unlabelled bottles contain aqueous solutions of ammonium chloride, aluminium chloride and zinc chloride. Using only aqueous sodium hydroxide and aqueous ammonia (and a Bunsen burner and red litmus paper), describe how you would identify each solution, giving all expected observations.

Answers

A pale blue precipitate forms; in excess ammonia it dissolves to give a dark blue solution.
Aqueous sodium hydroxide (or ammonia). $\ce{Fe^2+}$ : green precipitate, turning brown on contact with air, insoluble in excess. $\ce{Fe^3+}$ : red-brown precipitate, insoluble in excess.
Add dilute hydrochloric (or nitric) acid, then aqueous barium chloride (or nitrate). A white precipitate (of $\ce{BaSO4}$ ) that does not dissolve in the acid shows sulfate. The acid removes carbonate and sulfite ions, which would also give white precipitates with $\ce{Ba^2+}$ (those precipitates dissolve in acid).
$\ce{Zn(OH)2(s) + 2OH-(aq) -> [Zn(OH)4]^2-(aq)}$ .
Oxygen. Thermal decomposition of a Group 2 nitrate, e.g. $\ce{2Mg(NO3)2 -> 2MgO + 4NO2 + O2}$ (brown $\ce{NO2}$ ).
Add barium nitrate to each: both give a white precipitate. Add excess dilute nitric acid: the barium sulfite precipitate dissolves (with effervescence of $\ce{SO2}$ ), while barium sulfate does not dissolve.
The anion is chloride. No precipitate with NaOH and no ammonia on warming rules out the metal ions in the notes that form hydroxides and ammonium; the cation is probably a Group 1 ion such as $\ce{Na+}$ or $\ce{K+}$ (or $\ce{Ba^2+}$ at low concentration). Group 1 ions give no observable result with any test in the notes, so the cation cannot be identified, only described ("no ppt.; probably a Group 1 cation").
Calcium hydroxide is only slightly soluble. Aqueous sodium hydroxide provides a high concentration of $\ce{OH-}$ , so $\ce{Ca(OH)2}$ precipitates. Aqueous ammonia is a weak base and provides only a low concentration of $\ce{OH-}$ , too low for calcium hydroxide to precipitate.
Grey-green precipitate soluble in excess NaOH giving dark green solution: $\ce{Cr^3+}$ . Pale yellow precipitate insoluble in ammonia: $\ce{I-}$ . T is chromium(III) iodide, $\ce{CrI3}$ . $\ce{Cr^3+(aq) + 3OH-(aq) -> Cr(OH)3(s)}$ ; $\ce{Ag+(aq) + I-(aq) -> AgI(s)}$ .
Add $\ce{NaOH(aq)}$ dropwise then in excess to a sample of each. Ammonium chloride: no precipitate; on warming, a gas turns damp red litmus blue (ammonia). Aluminium chloride and zinc chloride: white precipitate, soluble in excess, no gas. To separate these two, add $\ce{NH3(aq)}$ dropwise then in excess to fresh samples: aluminium chloride gives a white precipitate insoluble in excess; zinc chloride gives a white precipitate that dissolves in excess.

Good practice

Reactions of aqueous cations

Why the results differ

Reactions of anions

Tests for gases and for iodine

Organic tests that may appear

Recording observations

A strategy for identifying an unknown

Worked examples

Practice

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