Identification of ions and gases Revision Notes

Carbonate ions react with dilute acids to produce carbon dioxide gas, water, and a salt. The test involves adding dilute hydrochloric acid (HCl) or dilute sulfuric acid (H₂SO₄) to the sample. Effervescence (bubbling) indicates gas evolution.

To confirm the gas is carbon dioxide, bubble it through limewater (a saturated solution of calcium hydroxide, Ca(OH)₂). Carbon dioxide reacts with limewater to form calcium carbonate, which causes the limewater to turn milky (cloudy).

This test is specific for carbonate ions because few other common anions produce carbon dioxide gas with dilute acid.

• Add dilute acid to the sample; observe effervescence.
• Collect gas and bubble through limewater.
• Limewater turns milky if CO₂ is present.
• Balanced equation for reaction with acid:
• CO₃²⁻(aq) + 2H⁺(aq) → CO₂(g) + H₂O(l)
• Calcium hydroxide reaction:
• Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) ↓ + H₂O(l)

Halide ions are identified by first acidifying the sample with dilute nitric acid (HNO₃) to remove interfering ions such as carbonate or hydroxide, which could form precipitates.

Then, aqueous silver nitrate (AgNO₃) is added. Silver ions react with halide ions to form silver halide precipitates, each with a characteristic colour.

The precipitates can be further tested by adding dilute ammonia solution to distinguish between them based on solubility.

• Add dilute nitric acid to the sample to acidify.
• Add aqueous silver nitrate solution.
• Observe precipitate colour:
• Chloride (Cl⁻): white precipitate (AgCl)
• Bromide (Br⁻): cream precipitate (AgBr)
• Iodide (I⁻): yellow precipitate (AgI)
• Solubility in ammonia:
• AgCl dissolves in dilute NH₃ (forming [Ag(NH₃)₂]⁺) solution, precipitate disappears; confirms chloride ion presence,

Nitrate ions are not detected by simple precipitation. Instead, they are identified by reduction to ammonia gas using aluminium foil and sodium hydroxide solution.

The sample is warmed with aluminium foil and concentrated sodium hydroxide (NaOH). Aluminium reduces nitrate ions to ammonia gas (NH₃).

Ammonia gas is detected by its characteristic pungent smell and by turning damp red litmus paper blue, indicating alkalinity.

• Mix sample with aluminium foil and concentrated NaOH solution.
• Warm gently to produce ammonia gas.
• Detect ammonia by smell and by turning damp red litmus paper blue.
• Ionic equation for reduction:
• NO₃⁻(aq) + 4Al(s) + 5OH⁻(aq) + 2H₂O(l) → NH₃(g) + 4AlO₂⁻(aq) + 3OH⁻(aq) (simplified)
• Overall reaction is complex but results in ammonia formation.

Sulfate ions are identified by adding dilute nitric acid to acidify the sample, preventing precipitation of other barium salts.

Then, barium nitrate (Ba(NO₃)₂) solution is added. Barium ions react with sulfate ions to form a white precipitate of barium sulfate (BaSO₄), which is insoluble in water.

This test is specific because barium sulfate is very insoluble, and acidification prevents false positives from carbonate or sulfite ions.

• Add dilute nitric acid to the sample to acidify.
• Add aqueous barium nitrate solution.
• Observe formation of white precipitate (BaSO₄).
• Ionic equation:
• Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s) ↓

Sulfite ions can be detected by their ability to reduce acidified potassium manganate(VII) solution.

Potassium manganate(VII) (KMnO₄) is purple and acts as an oxidising agent. When added to an acidified solution containing sulfite ions, the purple colour fades as MnO₄⁻ is reduced to colourless Mn²⁺ ions.

This colour change from purple to colourless indicates the presence of sulfite ions.

• Add dilute sulfuric acid (H₂SO₄) to acidify the sample.
• Add acidified potassium manganate(VII) solution dropwise.
• Observe the purple colour fading to colourless.
• Redox reaction:
• 3SO₃²⁻(aq) + 2MnO₄⁻(aq) + 6H⁺(aq) → 3SO₄²⁻(aq) + 2Mn²⁺(aq) + 3H₂O(l)

Aqueous sodium hydroxide (NaOH) and aqueous ammonia (NH₃) are commonly used reagents in qualitative inorganic analysis to identify metal cations in solution. When these reagents are added to solutions containing metal ions, characteristic precipitates often form. The colour, solubility, and behaviour of these precipitates in excess reagent provide diagnostic clues to the identity of the cation present.

The tests rely on the formation of metal hydroxides, which are often insoluble and precipitate out of solution. Some hydroxides dissolve in excess sodium hydroxide or ammonia due to the formation of complex ions. Additionally, ammonium ions (NH₄⁺) can be identified by the release of ammonia gas when heated with sodium hydroxide.

• Sodium hydroxide reacts with metal ions to form metal hydroxide precipitates.
• Aqueous ammonia can also form precipitates or dissolve them depending on the metal ion.
• Observing precipitate colour and solubility in excess reagent helps identify the cation.
• Heating ammonium salts with sodium hydroxide releases ammonia gas, detected by its smell or damp red litmus paper turning blue.

When aqueous sodium hydroxide is added to a solution containing metal cations, insoluble metal hydroxides often form as precipitates. The colour and solubility of these precipitates in excess sodium hydroxide are key to identification.

Below are the observations and ionic equations for each cation:

• Aluminium (Al³⁺): A white precipitate of aluminium hydroxide, Al(OH)₃, forms initially. This precipitate dissolves in excess NaOH to form a colourless, soluble aluminate ion, [Al(OH)₄]⁻.
• Equation: Al³⁺(aq) + 3OH⁻(aq) → Al(OH)₃(s) (white ppt)
• Al(OH)₃(s) + OH⁻(aq) → [Al(OH)₄]⁻(aq) (colourless solution)
• Ammonium (NH₄⁺): No precipitate forms. On heating with NaOH, ammonia gas (NH₃) is released, which turns damp red litmus paper blue.
• Equation: NH₄⁺(aq) + OH⁻(aq) → NH₃(g) + H₂O(l)
• Calcium (Ca²⁺): A white precipitate of calcium hydroxide, Ca(OH)₂, forms. This precipitate does not dissolve in excess NaOH.
• Equation: Ca²⁺(aq) + 2OH⁻(aq) → Ca(OH)₂(s) (white ppt)
• Chromium(III) (Cr³⁺): A green precipitate of chromium(III) hydroxide, Cr(OH)₃, forms. This precipitate dissolves in excess NaOH to give a dark green solution containing the chromite ion, [Cr(OH)₄]⁻ or [Cr(OH)₆]³⁻ depending on conditions, but commonly represented as [Cr(OH)₄]⁻ for simplicity in IGCSE context. The solution is greenish in colour, indicating complex formation. Equation: Cr³⁺(aq) + 3OH⁻(aq) → Cr(OH)₃(s) (green ppt) Cr(OH)₃(s) + OH⁻(aq) → [Cr(OH)₄]⁻(aq) (green solution)

Continuing with other cations: Cu²⁺, Fe³⁺, Fe²⁺, and Zn²⁺ show distinct precipitates and solubility behaviour with NaOH.

• Copper(II) (Cu²⁺): A blue precipitate of copper(II) hydroxide, Cu(OH)₂, forms. This precipitate dissolves in excess NaOH to give a dark blue solution due to the formation of complex ions.
• Equation: Cu²⁺(aq) + 2OH⁻(aq) → Cu(OH)₂(s) (blue ppt) Cu(OH)₂(s) + 2OH⁻(aq) → [Cu(OH)₄]²⁻(aq) (dark blue solution)
• Iron(III) (Fe³⁺): A reddish-brown precipitate of iron(III) hydroxide, Fe(OH)₃, forms. This precipitate does not dissolve in excess NaOH.
• Equation: Fe³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s) (reddish-brown ppt)
• Iron(II) (Fe²⁺): A green precipitate of iron(II) hydroxide, Fe(OH)₂, forms. This precipitate does not dissolve in excess NaOH and slowly oxidises to iron(III) hydroxide on exposure to air, changing colour from green to reddish-brown.
• Equation: Fe²⁺(aq) + 2OH⁻(aq) → Fe(OH)₂(s) (green ppt)
• Zinc (Zn²⁺): A white precipitate of zinc hydroxide, Zn(OH)₂, forms. This precipitate dissolves in excess NaOH to form a colourless, soluble zincate ion, [Zn(OH)₄]²⁻.
• Equation: Zn²⁺(aq) + 2OH⁻(aq) → Zn(OH)₂(s) (white ppt) Zn(OH)₂(s) + 2OH⁻(aq) → [Zn(OH)₄]²⁻(aq) (colourless solution)

Aqueous ammonia (NH₃ solution) can also be used to test for these cations. It provides hydroxide ions but in lower concentration than sodium hydroxide, so some hydroxides precipitate while others dissolve. The solubility in excess ammonia helps distinguish cations further.

Observations with aqueous ammonia: Al³⁺, Ca²⁺, Fe³⁺, Fe²⁺, and Cr³⁺ form precipitates that do not dissolve in excess ammonia. Cu²⁺ and Zn²⁺ form precipitates that dissolve in excess ammonia due to complex ion formation. Ammonium ions do not form precipitates with ammonia.

Details: - Aluminium (Al³⁺): White precipitate of Al(OH)₃, insoluble in excess ammonia. - Calcium (Ca²⁺): No precipitate or very slight white precipitate, generally no reaction. - Chromium(III) (Cr³⁺): Grey-green precipitate, insoluble in excess ammonia. - Copper(II) (Cu²⁺): Blue precipitate of Cu(OH)₂, dissolves in excess ammonia to form a deep blue complex ion [Cu(NH₃)₄]²⁺. - Iron(III) (Fe³⁺): Reddish-brown precipitate, insoluble in excess ammonia. - Iron(II) (Fe²⁺): Green precipitate, insoluble in excess ammonia. - Zinc (Zn²⁺): White precipitate, dissolves in excess ammonia to form a colourless complex ion.

• Cu²⁺ + 4NH₃ → [Cu(NH₃)₄]²⁺ (deep blue complex)
• Zn²⁺ + 4NH₃ → [Zn(NH₃)₄]²⁺ (colourless complex)

The ammonium ion is identified differently from metal cations because it does not form a precipitate with sodium hydroxide or ammonia. Instead, when aqueous sodium hydroxide is added to a solution containing ammonium ions and the mixture is gently heated, ammonia gas (NH₃) is released.

This ammonia gas can be detected by its characteristic pungent smell and by its ability to turn damp red litmus paper blue, indicating its alkaline nature.

Equation: NH₄⁺(aq) + OH⁻(aq) → NH₃(g) + H₂O(l) (upon heating)

• Add aqueous sodium hydroxide to the suspected ammonium salt solution.
• Gently warm the mixture to release ammonia gas.
• Test the gas with damp red litmus paper; a colour change to blue confirms ammonia.
• This test distinguishes ammonium ions from metal cations which form precipitates.

Ammonia is a colourless gas with a characteristic pungent smell. It is alkaline and soluble in water, forming ammonium hydroxide (NH₄OH). To test for ammonia gas, damp red litmus paper is used as an indicator.

When ammonia gas comes into contact with damp red litmus paper, the paper turns blue due to the alkaline nature of ammonia. This is a simple and reliable test for ammonia.

The reaction responsible for this colour change is the formation of ammonium hydroxide in the moisture of the paper: NH₃(g) + H₂O(l) → NH₄⁺(aq) + OH⁻(aq) (ammonium hydroxide). The hydroxide ions cause the red litmus paper to turn blue.

• Use damp red litmus paper to test for ammonia gas.
• Observation: Red litmus paper turns blue in the presence of ammonia.
• Reason: Ammonia dissolves in water to form alkaline ammonium hydroxide.

Carbon dioxide is a colourless, odourless gas that is slightly acidic when dissolved in water. To test for carbon dioxide, limewater (a saturated solution of calcium hydroxide, Ca(OH)₂) is used.

When carbon dioxide gas is bubbled through limewater, the solution turns milky or cloudy due to the formation of insoluble calcium carbonate (CaCO₃). This is a classic test for carbon dioxide.

The chemical reaction is: CO₂(g) + Ca(OH)₂(aq) → CaCO₃(s) ↓ + H₂O(l) (calcium carbonate precipitate forms). If excess CO₂ is passed, the milky precipitate dissolves to form soluble calcium hydrogen carbonate: CaCO₃(s) + CO₂(g) + H₂O(l) → Ca(HCO₃)₂(aq).

• Bubble suspected gas through limewater.
• Observation: Limewater turns milky (white precipitate forms).
• Reason: Formation of insoluble calcium carbonate.

Chlorine is a greenish-yellow gas with a strong, choking smell. It is a strong oxidising agent and bleaches coloured substances.

To test for chlorine gas, damp blue or red litmus paper is used. When chlorine gas comes into contact with damp litmus paper, it bleaches the colour, turning the paper white.

This bleaching occurs because chlorine reacts with water in the damp paper to form hydrochloric acid (HCl) and hypochlorous acid (HClO), which are both oxidising and decolourising agents: Cl₂(g) + H₂O(l) ⇌ HCl(aq) + HClO(aq).

• Use damp litmus paper to test for chlorine gas.
• Observation: Damp litmus paper is bleached white (loses colour).
• Reason: Chlorine reacts with water to form bleaching agents.

Hydrogen is a colourless, odourless gas that is highly flammable. To test for hydrogen, a lighted splint is brought near the gas sample.

If hydrogen is present, it burns with a distinctive 'pop' sound due to the rapid combustion reaction: 2H₂(g) + O₂(g) → 2H₂O(g).

This test is qualitative and the 'pop' sound is a key identifying feature of hydrogen gas.

• Bring a lighted splint near the gas.
• Observation: A 'pop' sound is heard if hydrogen is present.
• Reason: Hydrogen combusts rapidly with oxygen producing water vapour.

Oxygen is a colourless, odourless gas essential for combustion. To test for oxygen, a glowing splint (a wooden splint that has been lit and then blown out so it is glowing but not flaming) is introduced into the gas.

If oxygen is present, the glowing splint will reignite or burst into flame because oxygen supports combustion.

The reaction is the combustion of the splint's material in oxygen: CₓHᵧ + O₂ → CO₂ + H₂O (general combustion).

• Insert a glowing splint into the gas.
• Observation: The glowing splint reignites or bursts into flame.
• Reason: Oxygen supports combustion.

Sulfur dioxide is a colourless gas with a pungent, choking smell. It is a reducing agent and acidic when dissolved in water.

To test for sulfur dioxide, acidified potassium manganate(VII) solution (acidified KMnO₄) is used. Potassium manganate(VII) is purple in colour.

When sulfur dioxide is bubbled through acidified potassium manganate(VII), the purple colour fades and the solution becomes colourless due to the reduction of Mn⁷⁺ to Mn²⁺ ions: 2KMnO₄ + 5SO₂ + 2H₂O → K₂SO₄ + 2MnSO₄ + 2H₂SO₄. This is a redox reaction where sulfur dioxide is oxidised to sulfate and manganate(VII) is reduced to manganese(II).

• Add acidified potassium manganate(VII) to the gas sample.
• Observation: Purple solution decolourises to colourless.
• Reason: Sulfur dioxide reduces Mn⁷⁺ to Mn²⁺, causing loss of purple colour.

When metal ions are heated in a flame, their electrons absorb energy and become excited to higher energy levels. As the electrons return to their original energy levels, they release energy in the form of visible light. The wavelength (and thus colour) of this light is characteristic of the metal ion present, allowing identification by flame colour.

The flame test is a qualitative method because it indicates the presence of certain metal ions but does not measure their concentration.

The energy transitions correspond to specific colours: for example, sodium ions emit yellow light, while copper(II) ions emit blue-green light.

• Excited electrons emit light of characteristic wavelengths.
• Flame colour depends on the metal ion's electronic structure.
• Useful for detecting alkali and some alkaline earth metals and transition metals.

A clean platinum or nichrome wire loop is used to hold the sample. The loop must be thoroughly cleaned by dipping in dilute hydrochloric acid and then rinsing with distilled water to avoid contamination.

A small amount of the solid metal salt or a solution containing the metal ion is placed on the wire loop.

The loop is then placed in the hottest part of a Bunsen burner flame (usually the blue cone) and the colour of the flame is observed carefully against a white background for clarity.

• Use a clean wire loop to prevent contamination.
• Clean the loop by acid washing and rinsing before each test.
• Place the sample in the hottest part of the flame for best colour intensity.
• Observe flame colour against a white background for accuracy.
• Repeat tests to confirm results.

Each metal ion produces a distinctive flame colour due to its unique electron transitions. The following table summarises the characteristic flame colours for the specified cations:

• Lithium (Li⁺): Crimson red flame.
• Sodium (Na⁺): Bright yellow flame.
• Potassium (K⁺): Lilac (light purple) flame.
• Calcium (Ca²⁺): Orange-red flame.
• Barium (Ba²⁺): Pale green (light green) flame.
• Copper(II) (Cu²⁺): Blue-green flame.

The flame colours arise because electrons in the metal ions absorb thermal energy and jump to higher energy orbitals. When they fall back, they emit photons of specific wavelengths.

For example, sodium ions emit light at about 589 nm, which appears bright yellow.

A worked example: Testing a sample suspected to contain sodium ions. The sample is placed on a clean wire loop and introduced into the flame. Observation of a bright yellow flame confirms the presence of sodium ions.

• Electron excitation and relaxation cause emission of characteristic light.
• Different metals emit different wavelengths, producing distinct flame colours.
• Example: Na⁺ ions produce a yellow flame at 589 nm wavelength.

Always handle acids and metal salts with care, wearing safety goggles and gloves.

Avoid inhaling fumes from heated samples, as some metal salts can be toxic.

Ensure the wire loop is clean before each test to prevent false colours from contamination (e.g., sodium contamination can mask other colours).

• Wear appropriate personal protective equipment (PPE).
• Work in a well-ventilated area or under a fume hood.
• Clean wire loop thoroughly between tests.
• Repeat tests to confirm flame colour consistency.