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Halide ions (F⁻, Cl⁻, Br⁻, I⁻) are the anions formed when halogens gain an electron. Testing for halide ions and understanding their reactions with concentrated sulfuric acid are key topics for Edexcel A-Level chemistry. These reactions reveal important differences in the reducing power of the halide ions.
The standard test for halide ions involves adding dilute nitric acid (to remove any carbonate or sulfate ions that might interfere) followed by silver nitrate solution (AgNO₃).
The silver ions react with halide ions to form insoluble silver halide precipitates:
Exam tip: The distinction between "white" and "cream" can be subtle in the laboratory. That is why the ammonia confirmation test is so important. Always describe both the precipitate colour AND the ammonia result.
If the precipitate colours are difficult to distinguish, ammonia (NH₃) can be added as a further test:
| Precipitate | Colour | Ksp / mol² dm⁻⁶ | Dilute NH₃ | Concentrated NH₃ |
|---|---|---|---|---|
| AgCl | White | 1.8 × 10⁻¹⁰ | Dissolves | Dissolves |
| AgBr | Cream | 5.0 × 10⁻¹³ | Does not dissolve | Dissolves |
| AgI | Yellow | 8.3 × 10⁻¹⁷ | Does not dissolve | Does not dissolve |
AgCl dissolves in dilute ammonia because it forms the soluble complex ion [Ag(NH₃)₂]⁺:
AgCl(s) + 2NH₃(aq) → [Ag(NH₃)₂]⁺(aq) + Cl⁻(aq)
AgBr requires concentrated ammonia to dissolve (it has a much lower Ksp, so a higher ammonia concentration is needed to shift the equilibrium). AgI does not dissolve in either — its Ksp is so low that even concentrated ammonia cannot produce enough [Ag(NH₃)₂]⁺ to dissolve it.
flowchart TD
A["Add dilute HNO₃ then AgNO₃"] --> B{"Precipitate colour?"}
B -->|"White"| C["Likely Cl⁻"]
B -->|"Cream"| D["Likely Br⁻"]
B -->|"Yellow"| E["Likely I⁻"]
C --> F{"Add dilute NH₃"}
D --> G{"Add dilute NH₃"}
E --> H{"Add dilute NH₃"}
F -->|"Dissolves"| I["Confirmed Cl⁻ (AgCl)"]
G -->|"Does not dissolve"| J{"Add conc NH₃"}
H -->|"Does not dissolve"| K{"Add conc NH₃"}
J -->|"Dissolves"| L["Confirmed Br⁻ (AgBr)"]
K -->|"Does not dissolve"| M["Confirmed I⁻ (AgI)"]
When a solid sodium halide is treated with concentrated sulfuric acid (H₂SO₄), the initial reaction in all cases is an acid-base reaction producing the hydrogen halide:
NaX(s) + H₂SO₄(l) → NaHSO₄(s) + HX(g)
However, what happens next depends on the reducing power of the halide ion.
NaF and NaCl produce HF and HCl gas respectively. Steamy/misty fumes are seen. Neither F⁻ nor Cl⁻ is a strong enough reducing agent to reduce the sulfuric acid, so the reaction stops here.
Br⁻ is a stronger reducing agent than Cl⁻. Initially, HBr is formed:
NaBr(s) + H₂SO₄(l) → NaHSO₄(s) + HBr(g)
But HBr then reduces the H₂SO₄:
2HBr + H₂SO₄ → Br₂ + SO₂ + 2H₂O
Observations: steamy fumes (HBr), orange/brown fumes or vapour (Br₂), and a choking gas with a sharp smell (SO₂). The sulfur in H₂SO₄ is reduced from +6 to +4 in SO₂.
I⁻ is the strongest reducing agent among the common halide ions. It can reduce H₂SO₄ further than bromide can.
NaI(s) + H₂SO₄(l) → NaHSO₄(s) + HI(g)
HI then reduces H₂SO₄ in stages:
Observations: steamy fumes, purple/black solid or vapour (I₂), choking gas (SO₂), smell of rotten eggs (H₂S), and possibly yellow solid sulfur deposited.
The reducing power of halide ions increases down the group:
F⁻ < Cl⁻ < Br⁻ < I⁻
| Halide | Ionic Radius / pm | Can Reduce H₂SO₄? | Products with H₂SO₄ |
|---|---|---|---|
| F⁻ | 133 | No | HF only |
| Cl⁻ | 181 | No | HCl only |
| Br⁻ | 196 | Yes — partially | HBr, Br₂, SO₂ |
| I⁻ | 220 | Yes — fully | HI, I₂, SO₂, S, H₂S |
This is because:
| Halide | Products | Key Observations | Sulfur Reduced To |
|---|---|---|---|
| NaCl | HCl only | Steamy fumes | Not reduced |
| NaBr | HBr, Br₂, SO₂ | Steamy fumes, orange vapour, choking gas | +4 (SO₂) |
| NaI | HI, I₂, SO₂, H₂S, S | Steamy fumes, purple/black solid, choking gas, rotten egg smell, yellow solid | −2 (H₂S) and 0 (S) |
The reactions with concentrated H₂SO₄ are a classic way to demonstrate the increasing reducing power of halide ions. They also illustrate why HCl can be prepared using H₂SO₄ (since Cl⁻ does not reduce it) but HBr and HI cannot be prepared this way.
To prepare HBr and HI, an alternative non-oxidising acid such as concentrated phosphoric acid (H₃PO₄) is used:
NaBr(s) + H₃PO₄(l) → NaH₂PO₄(s) + HBr(g)
Phosphoric acid is not a strong enough oxidising agent to be reduced by Br⁻ or I⁻, so the hydrogen halide is the only product.
Question: A student adds concentrated H₂SO₄ to an unknown sodium halide. They observe steamy fumes, a purple-black solid forming, a choking smell, and a rotten egg smell. Identify the halide and explain each observation.
Answer:
Common exam mistake: Students sometimes forget to mention the initial steamy fumes of HI, focusing only on the reduction products. The acid-base reaction producing HI happens first, then the redox reactions follow. Always mention both stages.
Silver halides have historically been of great importance in photography. AgBr was the primary light-sensitive compound in photographic film. When light strikes AgBr crystals, the energy promotes electrons that reduce Ag⁺ to metallic Ag atoms, forming a latent image. Chemical developing then amplifies this to produce a visible image.
All three silver halides decompose in sunlight (photodecomposition):
2AgX(s) → 2Ag(s) + X₂(g)
This can be observed in the laboratory: AgCl turns grey/purple when exposed to light as metallic silver is deposited. This reaction is used to confirm the identity of silver halide precipitates — if a white precipitate of AgCl darkens in sunlight, it confirms the presence of chloride ions.
When working through the reactions with concentrated H₂SO₄, it is important to track oxidation states explicitly:
| Reaction | Species Oxidised | Change | Species Reduced | Change |
|---|---|---|---|---|
| NaCl + H₂SO₄ | None | — | None | — |
| 2HBr + H₂SO₄ → Br₂ + SO₂ + 2H₂O | Br⁻ | −1 → 0 | S in H₂SO₄ | +6 → +4 |
| 2HI + H₂SO₄ → I₂ + SO₂ + 2H₂O | I⁻ | −1 → 0 | S in H₂SO₄ | +6 → +4 |
| 6HI + SO₂ → H₂S + 3I₂ + 2H₂O | I⁻ | −1 → 0 | S in SO₂ | +4 → −2 |
The key insight is that I⁻ is such a strong reducing agent that it can reduce sulfur through multiple stages: +6 → +4 (SO₂) → 0 (S) → −2 (H₂S).
Testing for halide ions uses silver nitrate (precipitate colour) confirmed by ammonia solubility. Reactions with concentrated H₂SO₄ demonstrate the increasing reducing power down Group 7: Cl⁻ gives only HCl, Br⁻ reduces sulfur to +4, and I⁻ reduces sulfur all the way to −2. Use H₃PO₄ (non-oxidising) to prepare pure HBr or HI. Track oxidation states explicitly in exam answers to demonstrate understanding.
Edexcel 9CH0 specification Topic 4 — Inorganic Chemistry and the Periodic Table, sub-topic 4.2 covers the tests for halide ions using silver nitrate (AgNO₃) and confirmation with dilute and concentrated ammonia, the trend in solubility/colour of silver halides, and the reactions of solid halides with concentrated sulfuric acid (H₂SO₄), where the products differ depending on the reducing power of X⁻ (refer to the official specification document for exact wording). Examined in Paper 1 (9CH0/01) and prominently in Paper 3 (9CH0/03) through CP4 and CP16. Synoptic dependencies: Topic 8 (Redox I) for the H₂SO₄ oxidation reactions, Topic 13 (Energetics II) for AgX lattice/hydration and ammonia solubility, Topic 11 (Equilibrium I) for solubility products of AgX, and Topic 15 (later) for [Ag(NH₃)₂]⁺ complex formation.
Question (8 marks):
(a) Describe the colour of the precipitate formed and the result of adding dilute then concentrated NH₃ when AgNO₃(aq) is added to (i) NaCl(aq), (ii) NaBr(aq), and (iii) NaI(aq). (4)
(b) State and explain what is observed when concentrated H₂SO₄ is added to solid NaBr at room temperature. Give the equations for any redox reactions. (4)
Solution with mark scheme:
(a)
| Halide | Initial precipitate (with AgNO₃ + dil. HNO₃) | Dilute NH₃ | Concentrated NH₃ |
|---|---|---|---|
| Cl⁻ | White (AgCl) | Dissolves | Dissolves |
| Br⁻ | Cream (AgBr) | Insoluble | Dissolves |
| I⁻ | Pale yellow (AgI) | Insoluble | Insoluble |
B1 for each row of correct precipitate colour + ammonia behaviour. (3 marks for table)
B1 — explanation: solubility of AgX in NH₃ depends on Ksp; [Ag(NH₃)₂]⁺ formation removes Ag⁺ from solution, but only AgCl has Ksp small enough to be redissolved by dilute NH₃; AgBr requires concentrated; AgI's Ksp is so small (≈ 8 × 10⁻¹⁷) that even concentrated NH₃ cannot dissolve it.
(b) Step 1 — initial reaction.
Concentrated H₂SO₄ acts first as an acid: NaBr(s) + H₂SO₄(l) → NaHSO₄(s) + HBr(g) — steamy white/cream fumes of HBr.
M1 — HBr formation written / observation of fumes.
Step 2 — redox follow-up.
Br⁻ is a moderately strong reducing agent. The HBr produced reduces some H₂SO₄: 2HBr + H₂SO₄ → Br₂ + SO₂ + 2H₂O. Visible orange/brown fumes of Br₂ and a colourless choking gas SO₂.
M1 — redox stage recognised + observations.
Step 3 — equation balance and oxidation states.
Net redox: 2Br⁻ → Br₂ + 2e⁻ (oxidation, Br: −1 → 0); SO₄²⁻ + 4H⁺ + 2e⁻ → SO₂ + 2H₂O (reduction, S: +6 → +4).
A1 — half-equations or overall balanced redox.
A1 — final synthesis: NaBr's reducing power is intermediate; oxidation goes only as far as SO₂ (S +6 → +4), not to S(0) or H₂S as I⁻ would drive.
Total: 8 marks (M2 A2 B4).
Question (6 marks): A student is given an unknown sodium halide solid and is asked to identify the halide ion using two tests.
(a) Describe Test 1 (silver-nitrate sequence) and Test 2 (concentrated H₂SO₄), stating exactly what would be observed for the three possibilities Cl⁻, Br⁻, I⁻. (4)
(b) The student observes a pale yellow precipitate insoluble in concentrated NH₃ in Test 1, and dark grey solid (with rotten-egg smell) plus violet vapour in Test 2. Identify the halide and justify with equations. (2)
Mark scheme decomposition by AO:
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