Group 17 (Halogens): Properties, Reactions, Trends, and Uses
1. General Characteristics
- Group 17 elements are called halogens.
- All exist as diatomic molecules (X₂) in the elemental state.
- Electronic configuration:
ns²np⁵ → 7 valence electrons.
- Typical oxidation states:
- –1 (halides)
- +1, +3, +5, +7 (in oxo‑compounds or when bonded to fluorine)
- All are strong oxidising agents; reactivity decreases down the group (F > Cl > Br > I > At).
2. Physical Properties
| Element |
Atomic No. |
Atomic radius (pm) |
Melting point (°C) |
Boiling point (°C) |
Colour (elemental) |
Density (g cm⁻³) |
| Fluorine (F) | 9 | 42 | -219.6 | -188.1 | pale‑yellow gas | 0.0015 |
| Chlorine (Cl) | 17 | 79 | -101.5 | -34.0 | yellow‑green gas | 0.0032 |
| Bromine (Br) | 35 | 94 | -7.2 | 58.8 | reddish‑brown liquid | 3.12 |
| Iodine (I) | 53 | 115 | 113.7 | 184.3 | violet solid | 4.93 |
| Astatine (At) | 85 | 150 | ≈302 | ≈337 | metallic‑grey (predicted) | ≈7 (estimated) |
3. Chemical Properties
3.1. Reaction with Metals – Halide Formation
General equation:
Metal + ½ X₂ → MetalX
Examples:
- 2 Na + Cl₂ → 2 NaCl
- Mg + Br₂ → MgBr₂
- 2 K + I₂ → 2 KI
3.2. Hydrogen Halides
H₂ + X₂ → 2 HX (ΔH < 0)
All HX are gases at STP except HF (a liquid). In water they give acids.
Acid strength of hydrogen halides
| Acid | pKₐ (approx.) | Relative strength |
|---|
| HF | 3.2 | weak (strong H–F bond) |
| HCl | -7 | strong |
| HBr | -9 | stronger |
| HI | -10 | strongest |
Trend: acidity increases down the group because the H–X bond strength decreases, making proton release easier.
3.3. Reaction with Water – Hydrolysis
- Fluorine: explosive oxidation of water
2 F₂ + 2 H₂O → 4 HF + O₂
- Chlorine, bromine, iodine: partial hydrolysis to give a mixture of halide and hypohalous acid
Cl₂ + H₂O ⇌ HCl + HOCl
Br₂ + H₂O ⇌ HBr + HOBr
I₂ + H₂O ⇌ HI + HOI
The equilibrium lies far to the left for Br₂ and I₂; only a small amount of HOX is formed.
Predicting pH of a chlorine‑water solution (syllabus requirement)
Given 0.10 M Cl₂ dissolved in water:
- Write the equilibrium:
Cl₂ + H₂O ⇌ HCl + HOCl
- HCl is a strong acid → [H⁺] ≈ 0.10 M.
- HOCl is a weak acid (Kₐ ≈ 3.5 × 10⁻⁸). Its contribution to [H⁺] is negligible compared with HCl.
- pH ≈ –log(0.10) = 1.0 (the solution is strongly acidic).
3.4. Disproportionation with Sodium Hydroxide
Cold NaOH (halide + hypohalite):
X₂ + 2 NaOH → NaX + NaOX + H₂O
- Cl₂ + 2 NaOH → NaCl + NaOCl + H₂O
- Br₂ + 2 NaOH → NaBr + NaOBr + H₂O
- I₂ + 2 NaOH → NaI + NaOI + H₂O
Hot NaOH (halide + halate):
3 X₂ + 6 NaOH → 5 NaX + NaXO₃ + 3 H₂O
- 3 Cl₂ + 6 NaOH → 5 NaCl + NaClO₃ + 3 H₂O
- 3 Br₂ + 6 NaOH → 5 NaBr + NaBrO₃ + 3 H₂O
- 3 I₂ + 6 NaOH → 5 NaI + NaIO₃ + 3 H₂O
3.5. Reaction with Aqueous Silver(I) Ions – Halide Identification
Ag⁺ + X⁻ → AgX(s)
- AgF – soluble (no precipitate)
- AgCl – white precipitate
- AgBr – pale‑yellow precipitate
- AgI – yellow‑brown precipitate (least soluble)
3.6. Chlorine‑Water Disinfection Chemistry
Cl₂ + H₂O ⇌ HCl + HOCl
HOCl ↔ H⁺ + OCl⁻ (pKₐ ≈ 7.5)
- HOCl and OCl⁻ are the active germicidal species.
- pH < 7 → HOCl predominates (more effective).
- pH > 8 → OCl⁻ predominates (still active but less potent).
3.7. Formation of Halogen Oxo‑Anions
| Oxidation State of X |
Name of Anion |
General Formula |
Example (Cl) |
| +1 | Hypohalite | XO⁻ | ClO⁻ |
| +3 | Halite | XO₂⁻ | ClO₂⁻ |
| +5 | Halate | XO₃⁻ | ClO₃⁻ |
| +7 | Perhalate | XO₄⁻ | ClO₄⁻ |
Students should be able to write the correct formula when given the oxidation state (e.g., “X in +5 → X O₃⁻”).
Environmental / industrial relevance of oxo‑anions
- Chlorate (ClO₃⁻) – used in bleaching of paper and textiles; can contaminate groundwater.
- Perchlorate (ClO₄⁻) – component of solid rocket propellants and fireworks; persistent environmental pollutant that interferes with thyroid function.
- Bromate (BrO₃⁻) – formed as a by‑product during ozonation of drinking water; regulated because of carcinogenic potential.
3.8. Reactions with Organic Compounds
- Free‑radical halogenation of alkanes (Cl₂ or Br₂, hv):
CH₄ + Cl₂ → CH₃Cl + HCl
- Nucleophilic substitution (formation of alkyl halides):
R–OH + HX → R–X + H₂O
- Electrophilic addition to alkenes:
C₂H₄ + Br₂ → C₂H₄Br₂
- Halogenation of phenols (e.g., bromination):
Phenol + Br₂ → 2,4,6‑Tribromophenol + 2 HBr
4. Periodic Trends in Group 17
| Property |
Trend (F → At) |
Reason |
| Atomic radius | Increase | Additional electron shells are added down the group. |
| Electronegativity (Pauling) | Decrease | Electron‑nuclear attraction weakens as radius increases. |
| X–X bond dissociation energy | Decrease | Bond length increases; orbital overlap diminishes (see Box 1). |
| Standard reduction potential, E° (X₂ + 2e⁻ → 2X⁻) | Become less positive | Halogen becomes a weaker oxidising agent down the group. |
| Acidity of hydrogen halides (HX) | Increase (HF < HCl < HBr < HI) | H–X bond strength decreases, facilitating proton release. |
| Stability of higher oxidation states | Increase down the group (Cl → Br → I) | Larger atoms can accommodate more oxygen ligands; lattice energies of salts increase. |
Box 1 – Representative X–X Bond Energies
| Bond | Bond dissociation energy (kJ mol⁻¹) |
|---|
| F–F | ≈ 158 (kJ mol⁻¹) – unusually low because of repulsion between lone pairs on small 2p orbitals. |
| Cl–Cl | ≈ 242 |
| Br–Br | ≈ 193 |
| I–I | ≈ 151 |
Box 2 – Standard Reduction Potentials (E°) for X₂/2X⁻
| Half‑reaction | E° (V) |
|---|
| F₂ + 2e⁻ → 2F⁻ | +2.87 |
| Cl₂ + 2e⁻ → 2Cl⁻ | +1.36 |
| Br₂ + 2e⁻ → 2Br⁻ | +1.07 |
| I₂ + 2e⁻ → 2I⁻ | +0.54 |
Interpretation: The more positive the potential, the stronger the halogen is as an oxidising agent. This explains why fluorine oxidises water, while iodine does not.
5. Oxidation‑State Range & Practice
- In compounds containing oxygen or fluorine, halogens exhibit oxidation states +1, +3, +5, +7 (in addition to –1).
- When writing formulas, the oxidation state determines the number of oxygen atoms in the oxo‑anion (see Table 3.7).
Practice: Write the formula for a chlorine species in the +5 oxidation state.
Solution: +5 → halate → ClO₃⁻.
6. Important Uses of Halogens and Their Compounds
- Fluorine – PTFE (Teflon), refrigerants (hydrofluoro‑olefins), uranium enrichment (UF₆), fluoride in toothpaste & water fluoridation.
- Chlorine – water & swimming‑pool disinfection, PVC production, chlorinated solvents (chloroform, carbon tetrachloride), bleaching powder (Ca(OCl)₂).
- Bromine – flame retardants (brominated diphenyl ethers), photographic film (AgBr), brominated pharmaceuticals, fire‑suppressant agents.
- Iodine – antiseptic (tincture of iodine), iodised salt, thyroid hormone synthesis, contrast agents for X‑ray imaging, organic synthesis (iodination).
- Astatine – research isotope; potential use in targeted α‑particle radiotherapy (experimental).
- Oxidised halogen species – chlorate in paper bleaching, perchlorate in solid‑rocket propellants, bromate as a water‑treatment by‑product (environmental monitoring required).
7. Summary of Key Points
- Halogens have the configuration
ns²np⁵ and exist as diatomic molecules.
- Physical properties change markedly down the group (state, colour, density, melting/boiling points).
- Reactivity as oxidising agents decreases down the group; fluorine is the strongest.
- Core reactions:
- Halide formation with metals.
- Formation of hydrogen halides and their acid behaviour (strength increases down the group).
- Hydrolysis in water – gives HX + HOX; pH of resulting solutions can be predicted.
- Disproportionation with cold and hot NaOH (halide + hypohalite/halate).
- Precipitation with Ag⁺ for halide identification.
- Chlorine‑water disinfection (HOCl/OCl⁻ equilibrium).
- Formation of oxo‑anions; link oxidation state ↔ anion name.
- Typical organic reactions (free‑radical halogenation, substitution, addition).
- Periodic trends: atomic radius ↑, electronegativity ↓, X–X bond energy ↓, reduction potential ↓, acid strength ↑, higher oxidation‑state stability ↑.
- Bond‑energy and reduction‑potential tables help predict redox behaviour.
- Environmental/industrial relevance of halogen compounds (bleaching, propellants, water treatment) is an essential context for the syllabus.