Group 17 (Halogens)
The halogens — fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At) — form Group 17 of the periodic table. They are the most electronegative non-metals, each requiring just one electron to complete their valence shell (ns²np⁵ configuration). This makes them powerfully reactive oxidising agents, particularly towards metals and hydrogen. Their physical states at room temperature vary strikingly: F₂ and Cl₂ are gases, Br₂ is the only non-metallic liquid element, and I₂ is a lustrous purple-black solid. All form diatomic molecules held together by a single covalent bond, and all exhibit vivid colours ranging from pale yellow (F₂) through yellow-green (Cl₂) and reddish-brown (Br₂) to deep violet vapour (I₂). A key trend across Group 17 is the progressive decrease in oxidising power from F₂ to I₂. Standard reduction potentials (E° values) reflect this: F₂/F⁻ (+2.87 V), Cl₂/Cl⁻ (+1.36 V), Br₂/Br⁻ (+1.07 V), I₂/I⁻ (+0.54 V). Because fluorine has the highest E°, it can oxidise every other halide ion, displacing them from solution. A lighter halogen will always displace a heavier halogen from its salt: for example, chlorine water added to potassium iodide solution turns the solution brown as free iodine is liberated. These displacement reactions are the basis of the 'halogen activity series' and have important consequences in analytical and industrial chemistry. Beyond the pure elements, halogens react with one another to form interhalogen compounds — covalent molecules of the general formula XYₙ, where X is the heavier halogen and Y is the lighter. These compounds display properties intermediate between their parent halogens and are themselves strong oxidising agents. Industrial preparation of chlorine — historically achieved by electrolysis of brine — can also be accomplished by the catalytic Deacon process, which oxidises hydrogen chloride with atmospheric oxygen over a copper(II) chloride catalyst. Understanding Group 17 chemistry integrates redox theory, molecular geometry, and industrial catalysis in one cohesive topic.
- Oxidising Power and the Electronegativity Trend
- Interhalogen Compounds: Structure and Bonding
- The Deacon Process for Chlorine Preparation
- Oxidation States and Chemical Versatility of Halogens
Oxidising Power and the Electronegativity Trend
Oxidising power measures the ability of a species to gain electrons. For halogens, this corresponds to the half-reaction X₂ + 2e⁻ → 2X⁻. The ease of this process decreases down the group for three reasons: (1) Atomic radius increases, so the incoming electron experiences less nuclear attraction. (2) Electron shielding by inner shells increases, further reducing effective nuclear charge. (3) Bond dissociation energy of X₂ (energy needed to break X–X to give atoms) decreases steadily: F₂ 158 kJ mol⁻¹, Cl₂ 243 kJ mol⁻¹, Br₂ 193 kJ mol⁻¹, I₂ 151 kJ mol⁻¹. Note F₂ has a surprisingly low bond energy due to lone-pair repulsion in the very short F–F bond, yet it still has the highest oxidising power because of its enormous electron affinity and hydration energy of F⁻. The displacement reactions Cl₂ + 2KBr → 2KCl + Br₂ and Cl₂ + 2KI → 2KCl + I₂ confirm the order F₂ > Cl₂ > Br₂ > I₂ in oxidising ability.
Interhalogen Compounds: Structure and Bonding
Interhalogen compounds form when two different halogens react. General types include AB (e.g. ClF, BrF, ICl, IBr), AB₃ (e.g. ClF₃, BrF₃, IF₃), AB₅ (e.g. ClF₅, BrF₅, IF₅), and AB₇ (e.g. IF₇). The central atom X is always the heavier, larger halogen because it can accommodate the expanded octet. Bonding in AB molecules: the molecule is polar covalent with the partial negative charge on the lighter, more electronegative halogen. In ClF₃, VSEPR predicts a T-shaped geometry (3 bonding pairs + 2 lone pairs on Cl, sp³d hybridisation). In IF₅ the geometry is square pyramidal (5 bonding pairs + 1 lone pair, sp³d² hybridisation). In IF₇ the geometry is pentagonal bipyramidal (7 bonding pairs, sp³d³). Interhalogens are powerful halogenating and fluorinating agents; BrF₃ is even used as a non-aqueous solvent. Importantly, interhalogens are more reactive than either parent halogen because the X–Y bond (joining unlike atoms) is weaker and more polar than either X–X or Y–Y bond.
The Deacon Process for Chlorine Preparation
The Deacon process (1868) oxidises hydrogen chloride gas using atmospheric oxygen over a copper(II) chloride catalyst supported on pumice at 400–450 °C: 4HCl(g) + O₂(g) ⇌ 2Cl₂(g) + 2H₂O(g), ΔH° = −114 kJ mol⁻¹. The reaction is exothermic and reversible; lower temperatures favour products but reduce reaction rate, so the operating temperature is a compromise. The catalytic cycle involves: (i) CuCl₂ oxidises HCl, forming CuCl and liberating Cl₂ (2CuCl₂ + oxidation step); (ii) CuCl is re-oxidised by O₂ back to CuCl₂, completing the cycle — this is a redox catalytic cycle. The process is industrially significant because it recovers chlorine from HCl by-product streams in chlorination reactions, making it economically and environmentally valuable. Modern variants use RuO₂ catalysts (Sumitomo process) at lower temperatures with higher conversion. Compare this with electrolysis of brine (chlor-alkali process): 2NaCl(aq) + 2H₂O(l) → Cl₂(g) + H₂(g) + 2NaOH(aq), which requires electrical energy but simultaneously produces valuable NaOH and H₂.
Oxidation States and Chemical Versatility of Halogens
Fluorine, being the most electronegative element, has only the −1 oxidation state in compounds. The other halogens can exhibit oxidation states of −1, 0, +1, +3, +5, and +7, expanding their chemistry enormously. Chlorine demonstrates this range in its oxoacids: HClO (hypochlorous acid, Cl = +1), HClO₂ (chlorous acid, Cl = +3), HClO₃ (chloric acid, Cl = +5), HClO₄ (perchloric acid, Cl = +7). Acid strength increases with oxidation state because higher oxidation states increase the polarity of the O–H bond, facilitating proton donation. The halogens also exhibit disproportionation in alkaline solution: Cl₂ + 2NaOH → NaCl + NaOCl + H₂O (cold), and 3Cl₂ + 6NaOH → 5NaCl + NaClO₃ + 3H₂O (hot). These reactions are fundamental to manufacturing bleaching powder, sodium hypochlorite bleach, and chlorates.
The Key Equation
- \ce{HCl} — Hydrogen chloride (gas): The reactant and industrial by-product being valorised; acts as the reducing agent (Cl goes from −1 to 0)
- \ce{O2} — Oxygen (gas): The oxidising agent; oxygen is reduced from 0 to −2 in water
- \ce{CuCl2} — Copper(II) chloride catalyst (solid / surface): Heterogeneous catalyst that participates in a redox cycle: CuCl₂ → CuCl → CuCl₂, regenerating active species
- \ce{Cl2} — Chlorine gas (gas): The desired product; a yellow-green diatomic gas used in water treatment, PVC manufacture, and pharmaceuticals
- \Delta H° — Standard enthalpy change (kJ mol⁻¹): −114 kJ mol⁻¹; the reaction is exothermic, releasing heat and making lower temperatures thermodynamically favourable
Worked Example
Problem. Chlorine water is added to a solution of potassium iodide (KI). (a) Write a balanced ionic equation for the reaction. (b) Calculate the standard cell EMF for this reaction using E°(Cl₂/Cl⁻) = +1.36 V and E°(I₂/I⁻) = +0.54 V. (c) Predict what would happen if bromine water were used instead of chlorine water with KI solution.
Answer. E°cell = +0.82 V for Cl₂/I⁻; reaction is spontaneous. Ionic equation: Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) + I₂(aq). Bromine water also displaces iodine (E°cell = +0.53 V).
Origins · 1 minHenry Deacon and how this was discovered
Henry Deacon was a British industrial chemist who developed the Deacon process in 1868 while working at Widnes, England. Seeking to address the enormous pollution caused by the Leblanc soda process (which produced vast quantities of waste HCl gas), Deacon designed a catalytic method to convert HCl into useful chlorine gas using atmospheric oxygen over a copper chloride catalyst. His process represented one of the earliest examples of industrial heterogeneous catalysis and laid the groundwork for modern green chemistry principles — converting a waste by-product into a valuable product. It also demonstrated the economic power of catalysis: the same reaction without a catalyst proceeds negligibly slowly at practical temperatures. Deacon operated in the era before environmental legislation, when hydrogen chloride gas from the Leblanc process was simply vented to atmosphere and was destroying vegetation across entire regions of industrial England. His catalytic solution was partly inspired by watching the acid fumes bleach and kill the surrounding landscape — he turned pollution into profit.
Related Concepts
- Redox Reactions and Standard Electrode Potentials — Halogen displacement reactions and the Deacon process both rely on redox thermodynamics; E° values directly predict which halogen displaces another.
- VSEPR Theory and Molecular Geometry — The T-shape of ClF₃ and square pyramidal geometry of IF₅ are classic VSEPR examples illustrating how lone pairs dictate interhalogen compound shapes.
- Industrial Catalysis and Reaction Rates — The Deacon process is a primary example of heterogeneous catalysis with a redox catalytic cycle (Cu²⁺/Cu⁺), connecting kinetics and thermodynamics.
- Periodic Trends: Electronegativity and Atomic Radius — The decline in oxidising power down Group 17 directly reflects trends in atomic radius, electron affinity, and effective nuclear charge.
- Oxoacids of Halogens and Variable Oxidation States — Chlorine's oxoacids (HClO through HClO₄) illustrate the expanding range of oxidation states (+1 to +7) available to heavier halogens, in contrast to fluorine.