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*Revision Materials and Past Papers* 2.1.1 Atomic structure and isotopes 2.1.2 Compounds, formulae and equations 2.1.3 Amount of substance 2.1.4 Acids 2.1.5 Redox 2.2.1 Electron structure 2.2.2 Bonding and structure 3.1.1 Periodicity 3.1.2 Group 2 3.1.3 The halogens 3.1.4 Qualitative analysis 3.2.1 Enthalpy 3.2.2 Reaction Rates 3.2.3 Chemical equilibrium 4.1 Basic concepts and hydrocarbons 4.1.2 Alkanes 4.1.3 Alkenes 4.2.1 Alcohols 4.2.2 Haloalkanes 4.2.3 Organic synthesis 4.2.4 Analytical techniques 5.1.1 How fast? 5.1.2 How far? 5.1.3 Acids, bases and buffers 5.2.1 Lattice enthalpy 5.2.2 Enthalpy and entropy 5.2.3 Redox and electrode potentials 5.3.1 Transition elements 5.3.2 Qualitative analysis 6.1.1 Aromatic compounds 6.1.2 Carbonyl compounds 6.1.3 Carboxylic acids and esters 6.2.1 Amines 6.2.2 Amino acids, amides and chirality 6.2.3 Polyesters and polyamides 6.2.4 Carbon–carbon bond formation 6.2.5 Organic synthesis 6.3.1 Chromatography and qualitative analysis 6.3.2 Spectroscopy Required Practicals

Required Practicals

1 Moles determination 2 Acid–base titration 3 Enthalpy determination 4 Qualitative analysis of ions 5 Synthesis of an organic liquid 6 Synthesis of an organic solid 7 Qualitative analysis of organic functional groups 8 Electrochemical cells 9 Rates of reaction – continuous monitoring method 10 Rates of reaction – initial rates method 11 pH measurement 12 Research skills

Core Practical 4: Qualitative analysis of ions

Aim: To identify unknown anions and cations using simple test-tube reactions and confirmatory tests.

Safety and Disposal

Tests for anions

1) Carbonate ions, CO32−

Reagent: Dilute aqueous acid (e.g. HCl).

Test for carbonate: addition of dilute acid releases CO2 bubbles; confirm with limewater.

Observation: Effervescence as CO2 gas is released.

Ionic equations:
CO32−(aq) + 2H+(aq) → CO2(g) + H2O(l)
HCO3−(aq) + H+(aq) → CO2(g) + H2O(l)

Confirmation of CO2: Bubble gas through limewater; turns cloudy due to CaCO3 formation.
Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l)

2) Sulfate ions, SO42−

Reagent: Barium chloride solution acidified with dilute HCl.

Test for sulfate: acidified barium chloride gives white barium sulfate precipitate.

Observation: White precipitate of BaSO4 forms.

Ionic equation: Ba2+(aq) + SO42−(aq) → BaSO4(s)

Why acidified? To remove CO32− which would also give a white precipitate with Ba2+ and cause a false positive.

3) Halide ions, Cl−, Br−, I−

Test: Add dilute HNO3 (to remove carbonates), then add AgNO3.

Chloride gives white AgCl precipitate.
Cl−: white AgCl
Bromide gives cream AgBr precipitate.
Br−: cream AgBr
Iodide gives yellow AgI precipitate.
I−: yellow AgI

Further test (NH3): AgCl dissolves in dilute NH3; AgBr dissolves in concentrated NH3; AgI insoluble.

Ionic equation: Ag+(aq) + X−(aq) → AgX(s)   (X = Cl−, Br−, I−)

Tests for cations

Ammonium ions, NH4+

Reagent: NaOH(aq), then warm gently.

Test for ammonium: warming with NaOH releases ammonia that turns damp red litmus blue.

Observation: NH3 gas evolved; sharp smell; turns damp red litmus paper blue.

Ionic equation: NH4+(aq) + OH−(aq) → NH3(g) + H2O(l)

Group 2 metal ions (Mg2+, Ca2+, Sr2+, Ba2+)

Reagent: NaOH(aq)

Adding sodium hydroxide to Group 2 solutions: Mg2+ and Ca2+ give white precipitates; Sr2+ and Ba2+ remain in solution but give high pH.

Observations:

Photo of Matt
Matt’s exam tip

Test sequence matters:

  1. Carbonate test first — Ba2+ and Ag+ also form white carbonates.
  2. Sulfate test second — Ag+ can form insoluble Ag2SO4.
  3. Halide test last.

Incorrect order can create confusing precipitates (e.g. BaCO3 or Ag2SO4).